GIFT OF 
 Dean Frank H. Probert 
 
 Mining Dept 
 
AMERICAN SCIENCE SERIES ADVANCED COURSE 
 
 INORGANIC CHEMISTRY 
 
 BY 
 
 IEA REMSEN 
 
 Professor of Chemistry in the Johns Hopkins University 
 
 FIFTH EDITION, REVISED 
 
 NEW YORK 
 
 HENRY HOLT AND COMPANY 
 
 1899 
 
MM* 
 
 OEPI. 
 
 QlffI 
 
 DEAN FRANK H f 
 
 DEPT. 
 
 Copyright 1889, 1898, 
 
 LLi V/t:/: 
 
 HENRY HOLT'e v C 
 
 ROBERT DRUMMOND, Printed by 
 
 Electrotyper, D G. F, CLASS, 
 
 NEW YORK. NEW YORK. 
 
PREFACE TO THE FIRST EDITION. 
 
 IN the preparation of this book I have been much 
 encouraged by the cordial reception which has been 
 given my earlier text-books, both in this country and 
 abroad.' While those earlier works are intended to form 
 a series of which the present volume is the most advanced 
 member, it has little in common with the lower mem- 
 bers except the general method of treatment. An occa- 
 sional paragraph from the "Briefer Course" has been 
 incorporated, but the two books are quite distinct. 
 
 In classifying the elements the periodic system has 
 been adopted, and this has been pretty closely adhered 
 to. In order to secure as logical treatment as possible 
 it has been thought best not to give detailed descriptions 
 of apparatus and specific directions for the preparation 
 of substances, in the text proper. By avoiding these the 
 attention can be better directed to the principles involved 
 and a clearer conception of these principles will be formed, 
 than when the attention is distracted by the reading of such 
 details. On the other hand, full descriptions of apparatus 
 and processes will be found in the Appendix ; and these, 
 it is believed, will be of service to the teacher in the 
 lecture-room, as well as to the student in the laboratory. 
 
 The feature of the book which perhaps most distin- 
 guishes it from others is the fulness with which general 
 relations are discussed in it. Attention is constantly 
 called to analogies between properties of substances and 
 between chemical reactions, so that the thoughtful 
 student will, it is hoped, be led to look upon the sub- 
 stances and the reactions not as independent of one 
 another, but as related in many ways, and thus forming 
 
 iii 
 
ir PREFACE. 
 
 parts of a system. All thinking chemists have no doubt at 
 times an indistinct vision of a perfect science of Chemis- 
 try yet to come, in which the relations of the parts will 
 be clearly seen, and in which much that now appears of 
 little or no importance will be recognized as significant. 
 The subject cannot as yet, however, be treated as if that 
 perfection had been reached. Much progress has been 
 made of late years in the classification of the facts, and 
 it is of prime importance to the student that general rela- 
 tions should be pointed out for him as clearly as possible. 
 Of course in the classification of facts the end is not 
 reached. In every case of chemical action there are 
 certain features which call for much deeper study than 
 is usually given to them. For the most part chemists 
 have been content to know what chemical changes take 
 place when two or more substances are brought into 
 action, and have paid much less attention to the accom- 
 panying phenomena ; and yet it is evident that, in 
 order to get a clear conception of the nature of the 
 chemical act, it is necessary that we should learn all 
 we possibly can in regard to that act. Of late years 
 more and more attention has been given to the study of 
 the phenomena accompanying chemical changes ; and 
 a clearer view has been gained regarding chemical 
 action. A great field of study is thus opened, which 
 bears to the science of Chemistry as a whole somewhat 
 the same relation that Physiology bears to Biology, while 
 the study of chemical substances and their changes as 
 usually carried on is in the same way the counterpart 
 of Morphology. Neither of the parts taken separately 
 is Chemistry in the fullest sense. It will never be pos- 
 sible to study Chemistry without taking up and working 
 with chemical substances ; but as knowledge grows, 
 more and more attention will surely be given to chemical 
 action. In this book considerable space is devoted 
 to the discussion of the results obtained in the latter 
 kind of study. Some, no doubt, will hold that even 
 more prominence should have been given to this side of 
 the subject. Indeed I shall be glad if some of those who 
 use the book become interested in the new problems, 
 
PREFACE. V 
 
 and go further into their study. It has not, however, 
 appeared to me advisable, considering the purposes for 
 which this book has been written, to discuss them more 
 fully. 
 
 The subject of the Constitution of Chemical Compounds 
 receives a due share of attention. Constitutional for- 
 mulas are not, however, used recklessly as though they 
 were provided by nature ready-made ; but the effort is 
 made to keep clearly in mind the facts which they ex- 
 press so that they may be used intelligently. In this 
 connection I may call special attention to the way in 
 which the constitution of the so-called double salts of 
 the halogens is treated. To those who have not care- 
 fully looked into the evidence, the formulas used will 
 perhaps appear too speculative. I should be sorry to 
 err in this direction. For some time past the view put 
 forward has seemed to me to be justified, and I find that 
 others whose judgment I respect have held the same 
 view at least in regard to some of the compounds in 
 question. As, generally speaking, these compounds are 
 treated inadequately, and as they are commonly regarded 
 as inexplicable, I propose soon to present, in the proper 
 place, the evidence upon which my present view rests, 
 when it will, I think, be found that the evidence is fully 
 as strong as that upon which our views concerning the 
 constitution of most compounds are founded. 
 
 IRA EEMSEN, 
 BALTIMORE, March, 1889. 
 
 PREFACE TO SECOND EDITION. 
 
 THE call for a new edition of this book has given me 
 an opportunity to make some desirable changes, and to 
 correct those errors to which my attention has been 
 directed by others or which I have myself discovered. 
 The revision is based upon the labors of a very consid- 
 
VI PREFACE. 
 
 erable number of readers who have given me the benefit 
 of their criticisms, and I take this opportunity to express 
 my sincere thanks to all those who have aided me. 
 Should any one using the new edition discover errors in 
 it, I shall be thankful to be informed of the fact. It 
 seems fair to say that I have heard only words of com- 
 mendation in regard to the general plan and spirit of 
 the book. 
 
 IRA EEMSEN. 
 BALTIMORE, Decembw 6, 1889. 
 
 PREFACE TO FIFTH EDITION. 
 
 DURING the eight years that have passed since the 
 second edition of this book was published a number of 
 minor corrections have been made in it from time to 
 time. It has now, however, been subjected to a thorough 
 revision, and it is hoped that this new edition will be 
 found to contain everything that can fairly be looked for 
 in a book of its size. A new Appendix has been added 
 containing much information concerning the properties 
 of a large number of compounds which are necessarily 
 treated briefly or not at all in the text. 
 
 It may not be inappropriate to mention the fact that 
 the book has been well received not only in this country, 
 but in England and, in Germany, a German translation 
 having appeared shortly after the publication of the first 
 American edition. 
 
 IRA EEMSEN. 
 
 BALTIMORE, February 21, 1898. 
 
CONTENTS. 
 
 CHAPTER I. 
 
 CHEMICAL AND PHYSICAL CHANGE EARLIEST CHEMICAL KNOWLEDGE 
 LAW OF THE INDESTRUCTIBILITY OF MATTER LAW OF DEFINITE 
 PROPORTIONS LAW OF MULTIPLE PROPORTIONS THE ELEMENTS. 
 
 PAGE 
 
 Matter and Energy Chemical Change Physical Change Physics 
 and Chemistry Earliest Chemical Knowledge Alchemy 
 Chemistry as a Science Lavoisier's Work Law of the Inde- 
 structibility of Matter Conservation of Energy Early 
 Views regarding the Composition of Matter Elements Chem- 
 ical Action Chemical Affinity Chemical Compounds and 
 Mechanical Mixtures Qualitative and Quantitative Study of 
 Chemical Changes Law of Definite Proportions Law of Mul- 
 tiple Proportions Combining Weights of the Elements The 
 Elements, their Symbols and Atomic Weights Symbols of 
 Compounds Chemical Equations The Scope of Chemistry 
 Chemical Action accompanied by other Kinds of Action, . . 1 
 
 CHAPTER II. 
 
 A STUDY OF THE ELEMENT OXYGEN. 
 
 Historical Occurrence Preparation Physical Properties Chem- 
 ical Properties Burning in the Air and Burning in Oxygen 
 Phlogiston Theory Lavoisier's Explanation of Combustion 
 Kindling Temperature Slow Oxidation Heat of Combustion 
 Heat of Decomposition Chemical Energy and Chemical 
 Work-Oxides, 28 
 
 CHAPTER III. 
 
 A STUDY OF THE ELEMENT HYDROGEN. 
 
 Historical Occurrence Preparation Physical Properties Chem- 
 ical Properties Comparison of Oxygen and Hydrogen, ... 40 
 
 vii 
 
viii CONTENTS. 
 
 CHAPTER IV. 
 
 STUDY OP THE ACTION OP HYDROGEN ON OXYGEN. 
 
 PAGB 
 
 Burning of Hydrogen Method of Dumas Eudiometric Method 
 Calculation of the Results obtained in exploding Mixtures of 
 Hydrogen and Oxygen Determination of tho Volume of Water 
 Vapor formed by Union of Definite Volumes of Hydrogen and 
 Oxygen Heat evolved in the Union of Hydrogen and Oxy- 
 gen Applications of the Heat formed by the Combination of 
 Hydrogen and Oxygen Oxyhydrogen Light Velocity of 
 Combination of a Mixture of Hydrogen and Oxygen Sum- 
 mary, 49 
 
 CHAPTER V. 
 
 WATER. 
 
 Historical Occurrence Formation of Water and Proofs of its 
 Composition Properties of Water Chemical Properties of 
 Water Water as a Solvent Solution as an Aid to Chemical 
 Action Natural Waters What constitutes a Bad Drinking 
 Water Purification of Water, 57 
 
 CHAPTER VI. 
 
 CONSTITUTION OP MATTER ATOMIC THEORY ATOMS AND MOLECULES- 
 CONSTITUTION VALENCE. 
 
 Early Views The Atomic Theory as proposed by Dalton Use and 
 Value of a Theory Atomic Weights and Combining Weights 
 Molecules Avogadro's Law Distinction between Molecules 
 and Atoms Molecular Weights Deduction of Atomic 
 Weights from Molecular Weights Exact Atomic Weights 
 determined by the Aid of Analysis Molecular Formulas 
 Constitution Valence Replacing Power of Elements, ... 68 
 
 CHAPTER VII. 
 
 OZONE ALLOTROPY NASCENT STATE HYDROGEN DIOXIDE. 
 
 Occurrence Preparation Properties Relation between Oxygen 
 and Ozone Ozone in the Air Allotropy Varying Number 
 of Atoms in the Molecules of one and the same Element 
 Nascent State Hydrogen Dioxide or Hydrogen Peroxide 
 Properties Occurrence in the Air Characteristic Reactions 
 Thermochemical Considerations, . ... 85 
 
CONTENTS. ix 
 
 CHAPTER VIII. 
 
 CHLORINE HYDROCHLORIC ACID. 
 
 PAGE 
 
 Historical Occurrence of Chlorine Preparation Weldon's Proc- 
 essElectrolytic Process Properties Different Kinds of 
 Action Chlorine Hydrate and Liquid Chorine Applications 
 of Chlorine Hydrochloric Acid Historical Study of the 
 Action of Hydrogen upon Chlorine Preparation Properties 
 Chemical Action of Hydrochloric Acid, 96 
 
 CHAPTER IX. 
 
 COMPOUNDS OP CHLORINE WITH OXYGEN AND WITH HYDROGEN AND 
 
 OXYGEN. 
 
 General Principal Reactions for Making Compounds of Chlorine 
 with Hydrogen and Oxygen Chloric Acid Properties 
 Hypochlorous Acid Chlorous Acid Perchloric Acid Gen- 
 eral Compounds of Chlorine with Oxygen Constitution of 
 the Compounds of Chlorine with Hydrogen and Oxygen 
 Comparison of Chlorine and Oxygen, 113 
 
 CHAPTER X. 
 
 ACIDS BASES NEUTRALIZATION SALTS. 
 
 General A Study of the Act of Neutralization General Statements 
 Definitions Comparison of the Reaction between Acids and 
 Hydroxides, and between Acids and Chlorides Other Similar 
 Reactions Distinction between Acids and Bases Metals or 
 Base-forming Elements Constitution of Acids and Bases 
 Constitution of Salts Basicity of Acids Acidity of Bases- 
 Salts Acid Properties and Oxygen Nomenclature of Acids 
 Nomenclature of Bases Nomenclature of Salts, 127 
 
 CHAPTER XI. 
 
 NATURAL CLASSIFICATION OF THE ELEMENTS THE PERIODIC LAW. 
 
 Historical Arrangement of the Elements Connection between 
 the Position of the Elements in the Natural System and their 
 Chemical Properties Plan to be followed, 147 
 
 CHAPTER XII. 
 
 THE ELEMENTS OF FAMILY VII, GROUP B: 
 FLUORINE CHLORINE BROMINE IODINE. 
 
 General Bromine Occurrence Preparation Properties Chem- 
 ical Conduct of Bromine Uses of Bromine Hydrobromic 
 
CONTENTS. 
 
 PACK 
 
 Acids Properties Compounds of Bromine with Hydrogen and 
 Oxygen Compounds of Bromine and Chlorine Iodine 
 Occurrence Preparation Properties Hydriodic Acid lodic 
 Acid Iodine Pentoxide or lodic Anhydride Anhydrides, or 
 Acidic Oxides Periodic Acid Periodates Constitution of 
 Periodic Acid Constitution of lodic Acid and the Oxygen 
 Acids of Bromine Compounds of Iodine with Chlorine 
 Compounds of Iodine with Bromine Fluorine Occurrence 
 Properties Hydrofluoric Acid Constitution of Hydrofluoric 
 Acid and the Fluorides Compound of Fluorine with Iodine- 
 Tabular Presentation of the Compounds of the Members of the 
 Chlorine Family with Hydrogen, with Oxygen, with Hydrogen 
 and Oxygen, and with One Another Relative Affinities of the 
 Elements of the Chlorine Group Family VII, Group A 
 
 . 160 
 
 CHAPTER XIII. 
 
 THE ELEMENTS OP FAMILY VI, GROUP B : 
 SULPHUR SELENIUM TELLURIUM. 
 
 Introductory Sulphur Occurrence Extraction of Sulphur from 
 its Ores Properties Uses of Sulphur Compounds of Sulphur 
 with Hydrogen Hydrogen Sulphide, Sulphuretted Hydrogen 
 Properties Action of Hydrogen Sulphide upon Solutions of 
 Salts, Use in Chemical Analysis Hydrosulphides Hydrogen 
 Persulphide Compounds of Sulphur with Members of the 
 Chlorine Group Selenium Occurrence Properties Hydro- 
 gen Selenide Tellurium Occurrence Properties Hydro- 
 gen Telluride, 185 
 
 CHAPTER XIV. 
 
 COMPOUNDS OF SULPHUR, SELENIUM, AND TELLURIUM WITH OXYGEN 
 AND WITH OXYGEN AND HYDROGEN. 
 
 Introductory Sulphuric Acid Pure Sulphuric Acid Tetrahy- 
 droxyl Sulphuric Acid Normal Sulphuric Acid Disulphuric 
 Acid, Pyrosulphuric Acid Sulphurous Acid Hyposulphurous 
 Acid Thiosulphuric Acid Other Acids of Sulphur Per- 
 sulphuric Acid Constitution of the Acids of Sulphur Com- 
 pound of Sulphur with Oxygen Sulphur Dioxide Sulphur 
 Trioxide Acid Chlorides of Sulphur Thionyl Chloride 
 Sulphtiryl Chloride Chlorsulphuric Acid, or Sulphuryl- 
 hydroxyl Chloride Compounds of Selenium and Tellurium 
 with Oxygen and with Oxygen and Hydrogen Selenious Acid 
 Selenic Acid Selenium Dioxide Acid Chlorides of Sele- 
 nium Tellurious Acid Telluric Acid Oxides of Tellurium 
 Sulphotelluric Acid Family VI, Group A . .206 
 
CONTENTS. xi 
 
 CHAPTER XV. 
 
 NITEOGEN THE AIK ARGON. 
 
 PAGE 
 
 Nitrogen General Occurrence of Nitrogen Preparation Prop- 
 erties The Air Analysis of Air Argon, 248 
 
 CHAPTER XVI. 
 
 COMPOUNDS OF NITROGEN WITH HYDROGEN WITH HYDROGEN AND 
 OXYGEN WITH OXYGEN, ETC. 
 
 General Conditions which give Rise to the Formation of the Sim- 
 pler Compounds of Nitrogen RelatioDs between the Principal 
 Compounds of Nitrogen Ammonia Composition of Am- 
 monia Ammonium Amalgam Metallic Derivatives of Am- 
 monium Compounds and of Ammonia Structure of Ammoni- 
 um Compounds Hydraziue Hydroxylamine Triazoic Acid 
 Nitric Acid Red Fuming Nitric Acid Nitrous Acid Hy- 
 ponitrous Acid Nitrous Oxide Nitric Oxide Nitrogen Tri- 
 oxide Nitrogen Peroxide Nitrogen Pentoxide Structure of 
 the Compounds of Nitrogen with Oxygen and Hydrogen 
 Compounds of Nitrogen with the Elements of the Chlorine 
 Group Compounds of Nitrogen with the Members of the 
 Sulphur Group, 260 
 
 CHAPTER XVII. 
 
 ELEMENTS OP FAMILY Y, GROUP B: 
 
 PHOSPHORUS ARSENIC ANTIMONY BISMUTH. THE ELEMENTS AND 
 THEIR COMPOUNDS WITH HYDROGEN. 
 
 General Phosphorus Occurrence Preparation Properties Ap- 
 plications of Phosphorus Compounds of Phosphorus with 
 Hydrogen Phosphine, Gaseous Phosphuretted Hydrogen- 
 Arsenic Occurrence Preparation Properties Arsine, Ar- 
 seniuretted Hydrogen Antimony Occurrence Properties 
 Applications of Antimony Stibine Methods of distinguishing 
 between Arsenic and Antimony Bismuth Occurrence Com- 
 pounds of the Members of the Phosphorus Group with the 
 Members of the Chlorine Group Phosphorus Trichloride 
 Phosphorus Pentachloride Arsenic Trichloride Compounds 
 of Antimony and Chlorine Bismuth and Chlorine Double 
 Salts, 294 
 
 CHAPTER XVIII. 
 
 COMPOUNDS OF THE ELEMENTS OF THE PHOSPHORUS GROUP WITH 
 OXYGEN AND WITH OXYGEN AND HYDROGEN. 
 
 Introduction Phosphoric Acid, Orthophosphoric Acid Proper- 
 ties Pyrophosphoric Acid Metaphosphoric Acid Phosphor- 
 
xii CONTENTS. 
 
 PAGE 
 
 ous Acid Hypophosphoric Acid Hypopliosphorous Acid- 
 Phosphorus Peutoxide, Phosphoric Anhydride Phosphorus 
 Trioxide or Phosphorous Anhydride Phosphorus Suboxide 
 Phosphorus Tetroxide Constitution of the Acids of Phos- 
 phorus Phosphorus Oxychloride Arsenic Acid Arsenious 
 Acid Arsenic Trioxide Arsenic Pentoxide Sulphides 
 Arsenic Disulphide Arsenic Trisulphide Arsenic Pentasul- 
 phide Antimonic Acid Antimony Trioxide Salts of Anti- 
 mony Antimony Tetroxide Antimony Peutoxide Antimony 
 Trisulphide Antimony Pentasulphide Constitution of the 
 Acids of Arsenic and Antimony Oxychlorides of Antimony 
 Oxides of Bismuth Salts of Bismuth Bismuth Dioxide 
 Bismuth Peutoxide Bismuth Trisulphide Bismuth Oxy- 
 chloride Family V, Group A Vanadium Vauadic Acid 
 Tantalum Columbium Didymium Boron General Oc- 
 currence Preparation Properties Boron Trichloride Boron 
 Trifluoride Boric Acid Salts of Boron Nitrogen Boride, . 321 
 
 CHAPTER XIX. 
 
 CAKBON AND ITS SIMPLER COMPOUNDS WITH HYDROGEN AND CHLO- 
 RINE. 
 
 Introductory Occurrence of Carbon Diamond Graphite Amor- 
 phous Carbon Coal Diamond, Graphite, and Charcoal are 
 Different Forms of the Element Carbon Chemical Conduct of 
 Carbon Compounds of Carbon with Hydrogen, or Hydrocar- 
 bons. Conditions under which Hydrocarbons are formed 
 Number of Hydrocarbons Homology, Homologous Series- 
 Cause of the Homology among Compounds of Carbon Other 
 Series of Hydrocarbons Marsh Gas, Methane, Fire-damp 
 Ethylene, Olefiant Gas Acetylene Simpler Compounds of 
 Carbon with the Members of the Chlorine Group, 357 
 
 CHAPTER XX. 
 
 SIMPLER COMPOUNDS OF CARBON WITH OXYGEN, AND WITH OXYGEN 
 AND HYDROGEN. 
 
 General Relations between the Compounds of Carbon with Hy- 
 drogen and Oxygen Carbon Dioxide Preparation Proper- 
 tiesRelations of Carbon Dioxide to Chemical Energy 
 Respiration Carbon Dioxide and Life Energy Stored up in 
 Plants Carbonic Acid and Carbonates Carbon Monoxide 
 Formic Acid Carbonyl Chloride, Phosgene, 376 
 
 CHAPTER XXI. 
 
 ILLUMINATION FLAME BLOW-PIPE. 
 COMPOUNDS OF CARBON WITH NITROGEN AND SULPHUR. 
 
 Introduction Illuminating Gas, Coal Gas Flames Kindling 
 Temperature of Gases Miner's Safety-lamp Structure of 
 
CONTENTS. xiii 
 
 PAGE 
 
 Flames Blow-pipe Causes of the Luminosity of Flames 
 Bunsen Burner Compounds of Carbon with Nitrogen and 
 with Sulphur Cyanogen Hydrocyanic Acid, Prussic Acid- 
 Cyanic Acid Carbon Bisulphide Sulphocarbonic Acid, Thio- 
 carbonic Acid Oxysulphides Sulphocyanic Acid Constitu- 
 tion of Cyanogen and its Simpler Compounds, 394 
 
 CHAPTER XXII. 
 
 ELEMENTS OP FAMILY IV, GROUP A : 
 SILICON TITANIUM ZIRCONIUM CERIUM THORIUM. 
 
 General Silicon Occurrence Preparation Silicon Hydride 
 Titanium Zirconium Thorium Cerium Compounds of the 
 Elements of the Silicon Group with those of the Chlorine 
 Group Silicon Tetrachloride Silicon Hexachloride Silicon 
 Tetrafluoride Constitution of Fluosilicic Acid Titanium Tet- 
 rachloride Titanium Tetrafluoride Zirconium Tetrachloride 
 Thorium Tetrachloride Thorium Tetrafluoride Compari- 
 son of the Chlorides of Family IV with those of Fainity V 
 Compounds of the Members of the Silicon Group with Oxygen, 
 and with Oxygen and Hydrogen Silicon Dioxide Properties 
 Uses Silicic Acid Polysilicic Acids Disilicic Acids 
 Trisilicic Acids Titanium Dioxide Zirconium Dioxide 
 Thorium Dioxide Silicides Family IV, Group B, . . . .409 
 
 CHAPTER XXIII. 
 
 CHEMICAL ACTION. 
 
 Retrospective Classification of Reactions of the Elements and 
 Compounds Studied Kinds of Chemical Reactions Direct 
 Combination Direct Decomposition Metathesis The Cause 
 of Chemical Reactions An Ideal Chemical Reaction Influ- 
 ence of Mass Reactions may be complete if one of the Prod- 
 ucts formed is Insoluble or Volatile Thermochemical Study 
 of Aflinity Value of Thermochemical Measurements Heat 
 of Neutralization Avidity of Acids Other Methods for De- 
 termining the Avidity of Acids Study of Chemical Decom- 
 positions Dissociation Electrolysis Electrolytic Dissociation 
 Relations between Specific Heat and Atomic Weights Ex- 
 ceptions to the Law of Specific Heats Raoult's Method for 
 the Determination of Molecular Weights Determination of 
 the Extent of Dissociation of a Dissolved Substance, .... 426 
 
 CHAPTER XXIV. 
 
 BASE-FORMING ELEMENTS GENERAL CONSIDERATIONS. 
 
 Introductory Metallic Properties Order in which the Base-form- 
 ing Elements will be taken up Occurrence of the Metals 
 Extraction of the Metals from their Ores The Properties of 
 
Xiv CONTENTS. 
 
 PAGE! 
 
 the Metals Compounds of the Metals Chlorides Forma- 
 tion of Salts in General General Properties of the Chlorides 
 The so-called Double Chlorides and similar Compounds of 
 Fluorine, Bromine, and Iodine Different Chlorides of the 
 same Metal Oxides Differen-t Oxides of the same Metal 
 Hydroxides Decomposition of Salts by Bases Sulphides 
 Hydrosulphides Sulpho-salts Nitrates Chlorates Sul- 
 phatesCarbonates Phosphates Silicates, 455 
 
 CHAPTER XXV. 
 
 ELEMENTS OF FAMILY I, GROUP B : 
 
 THE ALKALI METALS : LITHIUM SODIUM POTASSIUM RUBIDIUM 
 CvESIUM AKMONIUM. 
 
 General Potassium Occurrence Preparation Properties Po- 
 tassium Hydride Potassium Fluoride, Chloride, Bromide, 
 Iodide Properties Applications Potassium Hydroxide 
 Potassium Oxide Potassium Hydrosulpbide Potassium Sul- 
 phide Potassium Nitrate Applications Gunpowder Potas- 
 sium Nitrite Potassium Chlorate Potassium Perchlorate 
 Potassium Periodate Potassium Cyanide Potassium Cyanate 
 Potassium Sulphocyanate Potassium Sulphate Primary, 
 or Acid, Potassium Sulphate Sulphites Carbonates Acid 
 Potassium Carbonate Phosphates Potassium Silicate Rubi- 
 dium Caesium Sodium Occurrence Preparation Proper- 
 ties Applications Sodium Hydride Sodium Chloride So- 
 dium Hydroxide Oxides Sodium Peroxide Sodium Sul- 
 phantimonate Sodium Nitrate Sodium Sulphate Sodium 
 Thiosulphate Sodium Carbonate Properties Applications 
 The Le Blanc Process for the Manufacture of Sodium Carbon- 
 ate Ammonia Process for the Manufacture of Soda Manu- 
 facture of Soda from Cryolite Mono-Sodium Carbonate, Pri- 
 mary Sodium Carbonate Sodium-Potassium Carbonate 
 Phosphates Sodium Metaphosphate Di-sodium Pyro-anti- 
 monate Sodium Borate Sodium Silicate Lithium Lithium 
 Phosphate Lithium Carbonate Lithium -Chloride Ammo- 
 nium Salts Ammonium Chloride Ammonium Sulphocyan- 
 ate Ammonium Sulphide Ammonium Nitrate Ammonium 
 Carbonate Sodium-ammonium Phosphate Reactions of the 
 Members of the Sodium Group which are of Value in Chemical 
 Analysis Flame Reactions and the Spectroscope, 482 
 
 CHAPTER XXVI. 
 
 ELEMENTS OF FAMILY II, GROUP A : 
 GLUCINUM MAGNESIUM CALCIUM STRONTIUM BARIUM [ERBIUM]. 
 
 General Calcium Sub-Group Calcium Occurrence Prepara- 
 tion Properties Calcium Chloride Calcium Fluoride Cal- 
 cium Oxide Calcium Hydroxide Bleaching-powder Cal- 
 
CONTENTS. XV 
 
 PAGE 
 
 cium Carbonate Applications Calcium Sulphate Calcium 
 Phosphates Calcium Silicate Glass Mortar Calcium Sul- 
 phide Calcium Nitride Calcium Carbide Strontium 
 Occurrence and Preparation Properties Compounds of 
 Strontium Barium Occurrence and Preparation Properties 
 Barium Chloride Barium Hydroxide Barium Oxide 
 Barium Peroxide or Dioxide Barium Sulphide Barium 
 Nitrate Barium Sulphate Barium Carbonate Phosphates of 
 Barium Reactions which are of Special Value in Analysis 
 Magnesium Snb-Group Gluciuum Occurrence and Prepara- 
 tion Properties Compounds of Glucinum Glucinum 
 Chloride Glucinum Hydroxide Glucinum Sulphate Glu- 
 cinum Carbonate Weak Basic Character of Gluciuum Mag- 
 nesium Occurrence Preparation Properties Applications 
 Compounds of Magnesium Magnesium Chloride Mag- 
 nesium Oxide Magnesium Sulphates Magnesium Carbonate 
 Phosphates Borates Silicates Magnesium Silicide Re- 
 actions of Magnesium Salts which are of Special Value in 
 Chemical Analysis Erbium General, 527 
 
 CHAPTER XXVII. 
 
 ELEMENTS OF FAMILY III, GROUP A : 
 ALUMINIUM SCANDIUM YTTKIUM YTTERBIUM SAMARIUM HELIUM. 
 
 General Aluminium Occurrence Preparation Properties 
 Applications Aluminium Chloride Chloroaluminates, or 
 Double Chlorides of Aluminium and analogous Compounds 
 Aluminium Hydroxide Aluminates Aluminium Oxide 
 Aluminium Sulphate Basic Aluminium Sulphates Alums 
 Potassium Alum, Potassium-Aluminium Sulphate Ammo- 
 nium Alum, Ammonium-Aluminium Sulphate Sodium Alum 
 Aluminium Silicate Kaoline Clay Ultramarine Por- 
 celain Earthenware Reactions of Aluminium Salts which 
 are of Special Value in Chemical Analysis Other Members of 
 Family III, Group A Scandium Yttrium Ytterbium Sa- 
 marium, Terbium, and Gadolinium Helium The Boron- 
 Aluminium Group in General 563 
 
 CHAPTER XXVIII. 
 
 ELEMENTS OF FAMILY I, GROUP B : 
 COPPER SILVER GOLD. 
 
 General Copper General Forms in which Copper occurs in 
 Nature Metallurgy of Copper Properties Applications 
 Alloys Cuprous Hydride Cupric Hydride Cuprous Chloride 
 Cupric Chloride Cuprous Iodide Cuprous Hydroxide Cu- 
 prous Oxide Cupric Hydroxide Cupric Oxide Other Oxides 
 of Copper -Cupric Sulphate Cupric Nitrate Cupric Arseuite 
 Cupric Carbonates Cyanides of Copper Cuprous Sulpho- 
 
xvi CONTENTS. 
 
 PAGE 
 
 cyanate Cupric Sulphocyanate Cuprous Sulphide Cupric 
 Sulphide Copper-plating Reactions which are of Special 
 Value in Chemical Analysis Silver General Forms in which 
 Silver occurs in Nature Metallurgy of Silver Properties 
 Allotropic Forms of Silver Alloys of Silver Argentous 
 Chloride Silver Chloride, Argentic Chloride Silver Bromide 
 and Iodide Application of the Chloride, Bromide, and Iodide 
 of Silver in the Art of Photography Silver Triazoate Silver 
 Oxide Other Oxides of Silver Sulphides of Silver Silver 
 Nitrate, Argentic Nitrate Silver Cyanide Silver Sulpho- 
 cyanate Berates of Silver Reactions which are of Special 
 Value in Chemical Analysis Gold General Forms in which 
 Gold occurs in Nature Metallurgy of Gold Properties- 
 Alloys of Gold Chlorides of Gold Chlorauric Acid Cyan- 
 auric Acid Auric Hydroxide Gold Sulphide, 587 
 
 CHAPTER XXIX. 
 
 ELEMENTS OF FAMILY II, GROUP B : 
 ZINC CADMIUM MERCURY. 
 
 General Zinc General Forms in which it occurs in Nature- 
 Metallurgy Properties Applications Alloys Zinc Chloride 
 Zinc Hydroxide Zinc Oxide Zinc Sulphide Zinc Sulphate 
 Zinc Carbonate Reactions which are of Special Value in 
 Chemical Analysis Cadmium General Preparation and 
 Properties Cadmium Sulphide Cadmium Cyanide Analyti- 
 cal Reactions Mercury General Forms in which Mercury 
 occurs in Nature Metallurgy of Mercury Properties Appli- 
 cationsAmalgams Mercurous Chloride Mercuric Chloride, 
 or Corrosive Sublimate Mercurous Iodide Mercuric Iodide 
 Mercurous Oxide Mercuric Oxide Mercurous Sulphide 
 Mercuric Sulphide Mercuric Cyanide Mercurous Nitrate 
 Mercuric Nitrate Compounds formed by Salts of Mercury 
 with Ammonia Reactions which are of Special Value in 
 Chemical Analysis, 616 
 
 ELEMENTS OF FAMILY III, GROUP B : 
 GALLIUM INDIUM THALLIUM. 
 
 General Gallium Compounds of Gallium Indium Compounds 
 of Indium Thallium Compounds of Thallium, 635 
 
 CHAPTER XXX. 
 
 ELEMENTS OF FAMILY IV, GROUP B: 
 GERMANIUM TIN LEAD. 
 
 General Germanium Tin General Occurrence Metallurgy 
 Properties Applications Alloys Stannous Chloride Stan- 
 nic Chloride Stannous Hydroxide Stannic Hydroxide 
 Metastannic Acid Stannous Oxide Stannic Oxide Stannous 
 
CONTENTS. xvii 
 
 PAGE 
 
 Sulphide Stannic Sulphide Stannous and Stannic Salts 
 Reactions which are of Special Value in Chemical Analysis 
 Lead General Forms in which Lead occurs in Nature Met- 
 allurgy Properties - Applications Lead Chloride Lead 
 Tetrachloride Lead Iodide Lead Hydroxide Oxides of Lead 
 Lead Suboxide Lead Oxide Lead Sesquioxide Lead Per- 
 oxide Red Lead, Minium Lead Sulphide Lead Nitrate 
 Lead CarbonateLead Sulphate Reactions which are of 
 Special Value in Chemical Analysis Lanthanum Cerium 
 Didymium, Praseodymium and Neodymium, 638 
 
 CHAPTER XXXI. 
 
 ELEMENTS OF FAMILY VI, GROUP A. 
 CHROMIUM MOLYBDENUM TUNGSTEN URANIUM/ 
 
 General Chromium General Forms in which Chromium occurs 
 in Nature Preparation Properties Chromous Chloride 
 Chromic Chloride Chromous Hydroxide Chromic Hydroxide 
 Chromic Oxide Chromic Sulphate Chrome-Alums 
 Chromic Acid and the Chromates Potassium Chromate Po- 
 tassium Bichromate Chromium Trioxide Relations between 
 the Chromates and Dichromates Sodium Chromate and So- 
 dium Dichromate Barium Chromate Lead Chromate 
 Chromium Oxychloride, Chromyl Chloride Reactions which 
 are of Special Value in Chemical Analysis Molybdenum 
 General Occurrence and Preparation Properties Chlorides 
 Oxides Molybdic Acid and the Molybdates Lead Molyb- 
 date Phospho-inolybdic Acid Tungsten General Occur- 
 rence and Preparation Properties Chlorides Oxides 
 Tungstic Acid and the Tuugstates Silico-tungstic Acids- 
 Uranium General Occurrence and Preparation Properties 
 Chlorides Oxides Uranous Salts Urauyl Salts Uranates, 657 
 
 CHAPTER XXXII. 
 
 ELEMENTS OF FAMILY VII, GROUP A : 
 MANGANESE. 
 
 General Forms in which Manganese occurs in Nature Prepara- 
 tion and Properties Manganous Chloride General Remarks 
 concerning the Oxides Manganous Oxide Mnnganous Hy- 
 droxide Manganous-manganic Oxide Manganic Oxide 
 Mangenese Dioxide Manganites Weldou's Process for the . 
 Regeneration of Manganese Dioxide in the Preparation of 
 Chlorine Sulphides Mangauous Cyanide Mauganous Car- 
 bonate Mangauous Sulphate Manganic Sulphate Manganic 
 Acid and the Manganates Permanganic Acid and the Per- 
 manganatesPotassium Permanganate Reactions which are 
 of Special Value in Chemical Analysis, 678 
 
xviii CONTENTS. 
 
 CHAPTER XXXIII. 
 
 ELEMENTS OF FAMILY VIII, SUB-GROUP A : 
 IRON COBALT NICKEL. 
 
 PAGK 
 
 Oeneral Iron Introductory Forms in which Iron occurs in 
 Nature Metallurgy Varieties .of Iron Steel Properties of 
 Iron Ferrous Chloride Ferric Chloride Cyanides Potas- 
 sium Ferrocyauide Ferrohydrocyaiiic Acid Ferric Ferro- 
 cyanide, or Prussian Blue Potassium Ferricyauide Ferri- 
 hydrocyauic Acid Ferrous Ferricyanide Nitroprussiates 
 Ferrous Hydroxide Ferrous Oxide Ferric Hydroxide Fer- 
 rous-ferric Oxide Soluble Ferric Hydroxide Ferric Oxide 
 Ferrous Sulphide Ferric Sulphide Ferrous Carbonate - 
 Ferrous Sulphate Ferric Sulphate Ferrous Phosphate Fer- 
 ric Acid Iron Disulphide Iron Carbonyls Reactions which 
 are of Special Value in Chemical Analysis Ferrous Compounds 
 Ferric Compounds Cobalt General Occurrence and Prep- 
 arationProperties Cobaltous Chloride Cobaltous Hydrox- 
 ide Cobaltous Oxide Cobaltic Hydroxide Cobalt Sulphide 
 Cyanides Smalt Compounds of Ammonia with Salts of 
 Cobalt Nickel General Occurrence and Preparation Prop- 
 erties Alloys Other Applications of Nickel Nickelous 
 Chloride Nickelous Hydroxide Nickelic Hydroxide Cyan, 
 ides Reactions of Cobalt and Nickel which are of Special 
 Value in Chemical Analysis, 691 
 
 CHAPTER XXXIV. 
 
 ELEMENTS OF FAMILY VIII, SUB-GROUP B: 
 
 RUTHENIUM RHODIUM PALLADIUM. 
 
 ELEMENTS OF FAMILY VIII, SUB-GROUP C : 
 
 OSMIUM IRIDIUM PLATINUM. 
 
 General The Platinum Metals Metallurgy Ruthenium Prop- 
 erties Chlorides Oxides Osmium Preparation Properties 
 Chlorides Oxides Rhodium Iridium Preparation 
 Properties Chlorides Oxides Palladium Preparation 
 Properties Palladium Hydrogen Chlorides Oxides Plati- 
 num Preparation Properties Applications of Platinum 
 Alleys of Platinum Chlorides Chlorplatinic Acid Cyanides 
 Hydroxides and Oxides Sulphides Compounds with Am- 
 monia The Platinum Bases, 719 
 
 APPENDIX I. 
 
 CONTAINING SPECIAL DIRECTIONS FOR LABORATORY WORK. 
 
 Introduction, 733 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER I. 
 
 Chemical Change caused by Heat Chemical Changes can be 
 effected by an Electric Current Mechanical Mixtures and! 
 Chemical Compounds Other Examples of Chemical Action, . 784 
 
CONTENTS. xix 
 
 EXPEKTMENT8 TO ACCOMPANY CHAPTEB H. 
 
 PAGE 
 
 Preparation of Oxygen Measurement of the Volume of Gases 
 Determination of the Amount of Oxygen liberated when a 
 known Weight of Potassium Chlorate is decomposed Physical 
 Properties of Oxygen Chemical Properties of Oxygen Oxy- 
 gen 5s used up in Combustion The Products of Combustion 
 weigh more than the Body burned, 740 
 
 EXPERIMENTS TO ACCOMPANY CHAPTEB III. 
 
 Preparation of Hydrogen Something besides Hydrogen is formed 
 Determination of the Amount of Hydrogen evolved when a 
 Known Weight of Zinc is dissolved in Sulphuric Acid Hydro- 
 gen is purified by passing through a Solution of Potassium 
 Permanganate Hydrogen passes readily through Porous 
 Vessels Diffusion Chemical Properties of Hydrogen Prod- 
 uct formed when Hydrogen is Burned Reduction, .... 751 
 
 EXPERIMENTS TO ACCOMPANY CHAPTEB IV. 
 
 Composition of Water Eudiometric Experiments Oxyhydrogen 
 Blow-pipe, .../>. 760 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER V. 
 
 Organic Substances contain Water Water of Crystallization- 
 Efflorescent Salts Deliquescent Salts Purification of Water 
 by Distillation 763 
 
 EXPERIMENTS TO ACCOMPANY CHAPTEB VI. 
 
 Method of Dumas Method of Victor Meyer, 765 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER VH. 
 
 Ozone Hydrogen Dioxide, 767 
 
 EXPERIMENTS TO ACCOMPANY CHAPTEB VIII. 
 
 Preparation of Chlorine Chlorine decomposes Water in the Sun- 
 light Chlorine Hydrate Formation of Hydrochloric Acid 
 Preparation of Hydrochloric Acid 768 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER IX. 
 
 Chloric Acid and Potassium Chlorate Perchloric Acid, .... 772 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER X. 
 
 Neutralization of Acids and Bases ; Formation of Salts Study of 
 the Products formed, . . 774 
 
 FOR CHAPTER XI., 
 
 776 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XII. 
 
 Preparation of Bromine Hydrobromic Acid Iodine Iodine can 
 be detected by Means of its Action upon Starch -paste Action 
 of Sulphuric Acid upon Potassium Iodide lodic Acid, 
 Hydrofluoric Acid, 776 
 
XX CONTENTS. 
 
 EXPERIMENTS~TO ACCOMPANY CHAPTER XIII. 
 
 PAGE 
 
 Properties of Sulphur Hydrogen Sulphide, 779> 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XIV. 
 
 Manufacture of Sulphuric Acid Sulphurous Acid and Sulphur 
 Dioxide Sulphurous Acid is a Reducing Agent Sulphur Tri- 
 oxide, 781 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XV. 
 
 Preparation of Nitrogen Analysis of Air, 784 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. 
 
 Preparation and Properties of Ammonia Ammonia burns in Oxy- 
 gen Ammonia forms Ammonium Salts with Acids Compo- 
 sition of Ammonia Preparation and Properties of Nitric Acid 
 Nitric Acid gives up Oxygen readily, and is hence a good 
 Oxidizing Agent Metals dissolve in Nitric Acid, forming 
 Nitrates Nitrates are decomposed by Heat Nitrates are sol- 
 uble in Water Nitric Acid is reduced to Ammonia by Nascent 
 Hydrogen Nitrous Acid Nitrous Oxide Nitric Oxide 
 Nitrogen Trioxide Nitrogen Peroxide, 78$ 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVII. 
 
 Phosphorus Phosphorus abstracts Oxygen from other Substances 
 Phosphine Arsenic Arsine Marsh's Test for Arsenic 
 Antimony Stibine Bismuth Phosphorus Trichloride 
 Phosphorus Pentachloride, . 79ft 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVIII. 
 
 Phosphoric Acid Arsenic Acid Reduction of Arsenic Trioxide 
 Sulphides of Arsenic Sulphides of Antimony Oxychlorides 
 of Antimony Basic Nitrates of Bismuth Boron, 801 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XIX. 
 
 Carbon Bone-black Filters Charcoal absorbs Gases Carbon 
 combines with Oxygen to form Carbon Dioxide Carbon re- 
 duces some Oxides when heated with them Hydrocarbons, . 80 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XX. 
 
 Carbon Dioxide is formed when a Carbonate is treated with an 
 Acid Preparation and Properties of Carbon Dioxide Carbon 
 Dioxide is given off from the Lungs Formation of Carbon- 
 ates Preparation and Properties of Carbon Monoxide Carbon 
 Monoxide is a Good Reducing Agent, 80S 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXI. 
 
 Coal Gas Oxygen burns in an Atmosphere of a Combustible Gas 
 Kindling Temperature of Gases The Blow-pipe and its 
 Uses Cyanogen 807 
 
CONTENTS. xxi 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXH. 
 
 Silicon Silicon Tetrafluoride and Fluosilicic Acid Silicic Acid, . 810 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXIV. 
 
 Chlorides, Bromides, and Iodides Hydroxides Sulphates Re- 
 duction of Sulphates to Sulphides Carbonates, 812 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXV. 
 
 Potassium Salts Sodium Salts 816 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVI. 
 
 Calcium Salts Magnesium and its Salts, 817 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVII. 
 
 Aluminium Chloride, a 818 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVIII. 
 
 Copper and its Salts Silver and its Salts, 818 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXIX. 
 
 Zinc and its Salts Mercury and its Salts, 819 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXX. 
 
 Tin and its Compounds Lead and its Compounds, 819 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXI. 
 
 Chromic Acid and the Chromates, 820 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXII. 
 
 Manganese and its Compounds, 821 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXIII. 
 
 Iron and its Compounds, 821 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXIV. 
 
 Platinum, 821 
 
 Conclusion, 821 
 
 APPENDIX II. 
 
 Note Atomic Weights Melting-points and Boiling-points of the . 
 Elements Melting-points, Boiling-points, and Solubilities of 
 Inorganic Substances Weights of Gases at and 760 mm. 
 Pressure Approximate Composition of a number of Alloys 
 Freezing-mixtures Table of Weights and Measures Com- 
 parison of the Twaddell Scale with the Baume and Gay-Lussac 
 Scales, 823 
 
A TEXT-BOOK OF 
 
 INORGANIC CHEMISTRY. 
 
 CHAPTER I. 
 
 CHEMICAL AND PHYSICAL CHANGE EARLIEST CHEMI- 
 CAL KNOWLEDGE LAW OF THE INDESTRUCTIBILITY 
 OF MATTER LAW OF DEFINITE PROPORTIONS LAW 
 OF MULTIPLE PROPORTIONS THE ELEMENTS. 
 
 Matter and Energy. The sensible universe is made -up 
 of matter and energy. It is difficult to give satisfactory 
 definitions of either of these terms, but, in a general way, 
 it may be said that matter is anything which occupies 
 space, and energy is that which causes change in matter. 
 It requires but little observation to show that there are 
 many kinds of matter, and apparently many kinds of 
 energy. As examples of the different kinds of matter we 
 have the many varieties of rocks and earth, as granite, 
 limestone, quartz, clay, sand, etc. ; the plants and their 
 fruits ; the substances which enter into the composition 
 of animals; and innumerable manufactured products. 
 As examples of the different forms of energy, we have 
 heat, light, motion, etc. Under the influence of the forms 
 of energy the forms of matter are constantly undergoing 
 change. Everywhere these changes are taking place. 
 Changes in position and in temperature appeal most 
 directly to our senses, and are most easily studied. But 
 there are many other kinds of change which are of the 
 highest importance. Thus there are electrical changes, 
 manifestations of- which we see in thunder-storms ; there 
 are magnetic changes which may be studied to some ex- 
 tent by means of the magnetic needle ; and there are, 
 further, what are called chemical changes which affect 
 the composition of substances. 
 
 (1) 
 
3 " , INORGANIC CHEMISTRY. 
 
 Chemical Change. For the purpose of study it is con- 
 venient to distinguish between two classes of changes in 
 matter, the difference between which can best be made 
 clear by means of examples. Consider the changes in- 
 cluded under the head of fire. We see substances de- 
 stroyed by fire, as we say. They disappear as far as we 
 can determine by ordinary observation. When iron is ex- 
 posed to the air a serious change takes place. It becomes 
 covered with a reddish-brown substance which we call 
 rust. If the piece of iron is comparatively thin, and it be 
 allowed to lie in the air long enough, it is completely 
 changed to the reddish-brown substance, and no iron as 
 such is left. If the juices from fruits, as from apples, be 
 allowed to stand in the air, they undergo change, becom- 
 ing sour, and a somewhat similar change takes place in 
 milk. If a spark be brought in contact with gunpowder 
 there is a flash and the powder disappears, a dense cloud 
 appearing in its place. 
 
 In the changes referred to the substances changed dis- 
 appear as such. After the fire, the wood or the coal, 
 or whatever may be burned, is no longer to be found. 
 The rusted iron is no longer iron. The gunpowder after 
 the flash is no longer gunpowder. Changes of this kind 
 in which the substances disappear and something else is 
 formed in their place are known as cJiemical changes. 
 
 Physical Change. There are many changes taking 
 place which do not affect the composition of substances. 
 Iron, for example, may be changed in many ways and 
 still remain iron. It may become hotter or colder. Its 
 position may be changed, or, as we say, it may be moved. 
 The iron may be struck in such a way as to give forth 
 sound. It may be made so hot that it gives light. 
 When, for example, it becomes red-hot, it can be seen in 
 a dark room. A piece of iron may be changed further 
 by connecting it with what is known as a galvanic bat- 
 tery. A current of electricity then passes through it, and 
 we can easily recognize the difference between a piece of 
 iron through which a current of electricit} r is passing and 
 one through which no current is passing. The former 
 when brought into certain liquids will at once change 
 
PHYSICS AND CHEMISTRY. 3 
 
 their composition, while the latter will not cause such 
 change. Finally, when a piece of iron is brought in con- 
 tact with loadstone, it acquires new properties. It now 
 has the power to attract and hold to itself other pieces 
 of iron. In all these cases, the iron is changed, but it re- 
 mains iron. After the moving iron comes to rest, it is 
 exactly the same thing that it was before it was moved. 
 After the iron which is giving forth sound has ceased to 
 give forth sound, it returns to its original condition. Let 
 the heated iron alone, and it cools down, ceasing soon to 
 give light, and presenting no evidence of being warm. 
 Remove the iron from contact with the galvanic battery, 
 and it loses those properties which are due to the current 
 of electricity. In time, the iron which is magnetized by 
 contact with the loadstone loses its magnetic properties. 
 It then no longer has the power to attract other pieces 
 of iron ; and does not difler from ordinary iron. 
 
 While iron has been taken as an example, other sub- 
 stances undergo similar changes. These changes which 
 do not affect the composition of the substances are called 
 physical changes. 
 
 Physics and Chemistry. According to what has been 
 said, we have two classes of changes presented to us for 
 study : 
 
 (1) Those which do not affect the composition of sub- 
 stances, or physical changes. 
 
 (2) Those which do affect the composition of sub- 
 stances, or chemical changes. 
 
 That branch of science which has to deal with physical 
 changes is known as PHYSICS. And that which has to 
 deal with chemical changes is known as CHEMISTRY. 
 
 Everything that has to do with motion, heat, light, 
 sound, electricity, and magnetism, is studied under the 
 head of Physics. Everything that has to do with the 
 composition of substances is studied under the head of 
 Chemistry. It is, however, impossible to study these two 
 subjects entirely independently of each other. When- 
 ever a chemical change takes place, it is accompanied 
 by physical changes ; and in order that the former may 
 be clearly understood, a study of the latter is necessary. 
 
INORGANIC CHEMISTRY. 
 
 Earliest Chemical Knowledge. Those substances which 
 are most abundant and most widely distributed in nature 
 were, of course, the first known and studied ; and the 
 same is true of those chemical changes which occur most 
 commonly and produce the most striking effects. Simply 
 by observing those things which surround us and those 
 changes in composition which take place naturally, a 
 considerable amount of chemical knowledge might be 
 gained, and indeed the earliest knowledge of chemistry 
 was acquired in this way. It was not, however, until 
 men came to experiment upon the substances which they 
 found in nature, that knowledge of chemical changes 
 made rapid progress. Since then an enormous amount 
 of knowledge has been gained, and every year the stock 
 is increased by new discoveries, until the field appears 
 almost boundless. 
 
 Alchemy. One of the first and one of the most power- 
 ful incentives to experiment upon chemical substances 
 was the desire to transform ordinary metals like lead 
 into gold. As will be seen farther on, there was no good 
 reason to suppose that transformations of this kind could 
 not be effected, indeed there were good reasons for sup- 
 posing them possible. For many hundred years men 
 worked with this object in view, and, though they did 
 not succeed in accomplishing that which they undertook, 
 they did add greatly to our knowledge of the action of 
 chemical substances upon one another, and they laid the 
 foundation of what has since become the great science of 
 chemistry. The work done by the alchemists was neces- 
 sary to prove that the transformations of matter which 
 they tried to effect cannot be effected. The problem 
 which they tried to solve was strictly a chemical prob- 
 lem, and the work they did was chemical work. 
 
 Chemistry as a Science. While the alchemists accumu- 
 lated a vast amount of knowledge of chemical facts, they 
 did not, strictly speaking, build up the science of chem- 
 istry. It was only when chemists came to study the facts 
 in their relations to one another, and when they were en- 
 abled to trace connections between them ; when they suc- 
 ceeded in discovering that certain general truths hold 
 
LAVOISIER 8 WORK. 5 
 
 good for all cases of chemical action ; when, in short, 
 the fundamental laws of chemical action were discovered 
 it was then that knowledge became science. It is 
 impossible to say definitely when chemistry became a 
 science. From the unorganized state to the organized 
 there was a gradual transition. But it is certain that the 
 labors of the French chemist Lavoisier contributed 
 largely to making chemistry what it is to-day, and it 
 is common to refer to his work as the beginning of the 
 science. 
 
 Lavoisier's Work. What distinguished Lavoisier's 
 work most markedly from that of his predecessors was 
 the way in which he used the balance for the purpose of 
 studying chemical changes. The balance had been used 
 to a considerable extent by earlier workers and results 
 of value had been reached by them, but Lavoisier suc- 
 ceeded by means of it in explaining some important 
 Chemical phenomena which had long been the subject of 
 ^tudy. His first investigation, the results of which were 
 published in 1770, was on the transformation of water into 
 earth. It was generally believed that water is trans- 
 formed into earth by boiling, because it was a matter of 
 common observation that, whenever water is boiled for 
 a time in a glass vessel, a deposit of earthy matter is 
 formed. In order to decide whether such transformation 
 takes place or not, Lavoisier boiled some water in a 
 closed vessel which he weighed before and after the boil- 
 ing ; and he found that the vessel decreased in weight by 
 a certain amount. He also determined the weight of the 
 deposit formed in the vessel and found that this was 
 exactly equal to the loss in weight of the vessel. He 
 also showed that there was just as much water after the 
 experiment as before. He therefore concluded that what 
 his predecessors had held to be a transformation of water 
 into earth was nothing but a partial disintegration of the 
 glass vessel caused by the action of the boiling water. 
 What had appeared mysterious became clear and sim- 
 ple. Having succeeded so well in this experiment, La- 
 voisier proceeded to study other chemical changes in the 
 same way, and soon he was able to give a perfectly satis- 
 
6 INORGANIC CHEMISTRY. 
 
 factory explanation of the process of combustion which 
 for a long time had been a subject of study. The ex- 
 planation and the experiments which led to it will be 
 taken up later. Suffice it to say here that the essential 
 feature of the work consisted in the fact that the sub- 
 stances which entered into action and those formed 
 during the action were all carefully weighed, and it was 
 found, in every case, that the weight of the substances 
 formed was exactly equal to the weight of the substances 
 which acted upon one another. 
 
 Law of the Indestructibility of Matter. While, as has 
 been stated, chemists before Lavoisier had used the 
 balance, they do not appear to have been very strongly 
 impressed by the importance of the weight of sub- 
 stances. They seem tacitly to have held that matter 
 can be destroyed. Lavoisier's work, however, showed 
 that whenever matter apparently disappears, it continues 
 to exist in some other form. If it were possible to an- 
 nihilate matter or to call it into being, it would be of 
 little or no value to weigh things. Innumerable experi- 
 ments which have been performed since Lavoisier's time 
 have confirmed the view that matter is indestructible. 
 The first fundamental law bearing upon the changes in 
 composition which the different forms of matter undergo 
 is the law of the indestructibility of matter. While, if we 
 think of it, we can scarcely conceive that this great law 
 should not be true, we must not forget that the only 
 way in which its truth could be established was by ex- 
 periment. The law may be stated thus : 
 
 Whenever a change in the composition of substances takes 
 pt,ace the amount of matter after the change is the same as 
 before the change. 
 
 According to this, and assuming that the law has al- 
 ways held good, it follows that the amount of matter in 
 the universe is the same to-day as it has been from the 
 beginning. Transformations are constantly taking place, 
 but these involve no increase nor decrease in the total 
 amount of matter. 
 
 Conservation of Energy. Just as matter is neither cre- 
 ated nor destroyed, so it has been made probable that 
 
COMPOSITION OF MATTER. 7 
 
 the total amount of energy is unchangeable. One of the 
 greatest discoveries in science was the recognition of the 
 fact that one form of energy can be transformed into 
 others, and that in these transformations nothing is lost. 
 We now know that for a certain amount of heat we can get 
 a certain amount of motion, and that for a certain amount 
 of motion we can get a certain amount of heat. We know 
 that a similar definite relation exists between heat and 
 electrical energy, and between these and chemical energy. 
 We know, for example, that a definite amount of heat 
 can be obtained by burning a definite amount of a given 
 substance, and we know also that with a definite amount 
 of heat we can produce a definite amount of chemical 
 change. Modern investigation has shown that all the 
 different forms of energy are convertible one into the 
 other without loss. This great fact is generally spoken 
 of as the law of the conservation of energy. Trans- 
 formations of energy are taking place constantly, as 
 transformations of matter are, but the total amount in 
 each case remains the same. 
 
 Early Views regarding the Composition of Matter. 
 The fact that first impresses one in studying the various 
 forms of matter found in the earth is their great variety. 
 We find an almost infinite number of kinds of matter, 
 and the question at once suggests itself, of what are 
 these things composed? This question has long been 
 asked, and it will be long before an entirely satisfactory 
 answer is reached. Still, much more is now known in 
 regard to the subject than was known in past ages, and 
 some progress is constantly being made towards a solu- 
 tion of the problem. At first, men attempted to answer 
 the question, as they attempted to answer all important 
 questions, by what is known as the speculative method ; 
 that is to say, they took the facts, as far as they knew 
 them, into consideration, and they endeavored by purely 
 mental processes to furnish an explanation. They were 
 on the whole much bolder in the use of the imagination 
 than the scientific men of the present are, or rather they 
 do not appear to have had as great respect for facts as 
 men now have, and, as a consequence, we find that some 
 
8 INORGANIC CHEMISTRY. 
 
 extremely curious speculations were indulged in. One 
 of the most prominent views in regard to the composi- 
 tion of matter was that put forward by Aristotle. Ac- 
 cording to this view, all forms of matter are made up of 
 four elements, earth, air, fire, and water ; and the vari- 
 ous forms differ from one another in the proportions of 
 these elements contained in them. Aristotle evidently 
 had in mind the fundamental properties of the four ele- 
 ments, rather than the elements themselves, and his idea 
 was that these fundamental properties are found in dif- 
 ferent proportions. Instead of meaning that water as 
 such was contained in substances, he meant that the 
 properties of cold and moisture were met with in sub- 
 stances, and so on for the other elements ; fire repre- 
 senting heat and dryness ; earth, cold and dryness ; and 
 air, heat and moisture. Afterwards it was pointed out 
 that besides the four fundamental properties represented 
 by the four elements, each substance has a special prop- 
 erty of its own which distinguishes it from all others. 
 This was called the quinta essentia, or fifth essence, from 
 which our modern word quintessence is derived. The 
 four or five elements of the older philosophers were, as 
 will be seen, imaginary things. They represented ideas 
 rather than tangible substances. 
 
 Elements. As experimenting upon chemical substances 
 advanced further and further, the fact impressed itself 
 more and more strongly upon investigators, that, of the 
 large number of substances known, some can be con- 
 verted into simpler ones by chemical action and some 
 cannot. In other words, some substances like water 
 can be broken down by various methods into two or 
 more others of different properties, and these when 
 brought together again under proper conditions form 
 the original substance. In the case of water, the action 
 of an electric current breaks it down or decomposes it, 
 and two gases, hydrogen and oxygen, are formed from it. 
 Elaborate experiments have shown that the weight of 
 water decomposed is exactly equal to the weight of the 
 hydrogen plus that of the oxygen obtained, and that 
 when the hydrogen and oxygen are brought together 
 
ELEMENTS. 9 
 
 again under proper conditions exactly as much water is 
 formed as was originally decomposed. It appears, there- 
 fore, that water consists of at least two simpler sub- 
 stances. A similar conclusion is reached by a study of 
 by far the largest number of the substances with which 
 we have to deal. On the other hand, no treatment to 
 which hydrogen and oxygen have been subjected has, as 
 yet, effected their decomposition. They can be made to 
 combine with other substances, as, for example, with ^ach 
 other, and thus form more complex substances, but noth- 
 ing simpler than hydrogen has ever been obtained from 
 hydrogen, and nothing simpler than oxygen has ever been 
 obtained from oxygen. Whether the decomposition of 
 these substances will ever be effected is a question which 
 cannot be answered. All that we know is that at present 
 they cannot be decomposed. We therefore speak of 
 them as elements, meaning by the term, that, with the 
 means now at the disposal of chemists, it is impossible to 
 get simpler substances from them. There are at present 
 about seventy substances known which are called ele- 
 ments for the same reasons that hydrogen and oxygen 
 are called elements. It is quite possible that the num- 
 ber may be increased in the future, and it is also quite 
 possible that the number may be decreased. New ele- 
 ments will in all probability be discovered, and prob- 
 ably some of the substances now included in the list of 
 elements may eventually be shown to be capable of 
 decomposition. 
 
 The view at present held in regard to the forms of 
 matter which go to make up that part of the universe 
 which comes under our observation is that they are all 
 composed of the seventy elementary substances. Many 
 of them, like water, are composed of only two elements ; 
 others of three ; and still others of four, five, six, and 
 more ; but most of them are comparatively simple, and 
 rarely does any one contain more than four or five ele- 
 ments. Of the seventy elements known, only about 
 twelve enter into the composition of most things with 
 which we commonly have to deal. The others occur in 
 .relatively small quantity. 
 
10 INORGANIC CHEMISTRY. 
 
 Chemical Action. In the last paragraph it was stated 
 that most substances can be decomposed, and that under 
 proper conditions the elements combine. We must now 
 inquire more carefully into the meaning of these expres- 
 sions. Among the elements are the well-known sub- 
 stances lead, iron, and sulphur. If some finely divided 
 iron is brought in contact with sulphur, apparently no 
 action takes place. If the two are put in a mortar and 
 mixed no matter how thoroughly, there is no evidence 
 of action. The mixture has, to be sure, a different ap- 
 pearance from that of either constituent, but still both 
 substances are present, and can be recognized by various 
 methods. If, for example, a little of the mixture is ex- 
 amined with the aid of the microscope, particles of iron 
 and of sulphur will be recognized lying side by side. If, 
 further, the mixture is treated with the liquid, carbon 
 disulphide, which has the power to dissolve the sulphur 
 but not the iron, the sulphur will be dissolved while the 
 iron will be left unchanged. Finally, if a dry magnet is 
 introduced into the mixture, the iron will adhere to it, 
 and by careful manipulation the two constituents can be 
 separated. These facts furnish evidence that both iron 
 and sulphur are present in the mixture in unchanged 
 condition, just as sugar and sand are present in a mixture 
 of these two substances. If now the mixture of sulphur 
 and iron is heated in a dry test-tube, marked changes 
 will take place, and there will be formed a black sub- 
 stance entirely different from either of the elements em- 
 ployed in the experiment. Carbon disulphide can no 
 longer extract sulphur from it. The magnet can no 
 longer pick out the iron, and under the microscope one 
 homogeneous substance is seen instead of the two ele- 
 mentary substances. If the experiment is performed 
 with proper precautions, the amount of matter after the 
 action will be found to be exactly the same as before the 
 action. A serious change has taken place, but no change 
 in the amount of matter. The act is one of chemical com- 
 bination, and the substance formed is called a chemical 
 compound. A few other examples will aid in making the 
 conception of chemical combination clear. When a bit 
 
CHEMICAL ACTION AND AFFINITY. 11 
 
 of phosphorus is brought in contact with a little iodine 
 action takes place at once ; the two elements combine, 
 losing their own characteristic properties and forming a 
 compound with properties quite different from those of 
 the constituents. When the gases hydrogen and oxygen 
 are brought together and a spark is passed through the 
 mixture an explosion occurs, and, in place of the gases, the 
 liquid, water, is formed. When sulphur burns in the air 
 the product formed is a pungent gas. It has been shown 
 that the act consists in the combination of the sulphur 
 with the gas, oxygen, which is contained in the air. All 
 these cases are examples of chemical combination. But 
 chemical action may be of the opposite kind, that is to 
 say, instead of being combination, it may be decomposition. 
 Thus, water which is formed by the chemical combina- 
 tion of hydrogen and oxygen may, by proper methods, 
 be decomposed into the same elements. We may con- 
 veniently think of that which causes elements to combine 
 as an attractive force exerted between the elements. 
 Now, when some power which can overcome this attrac- 
 tion is brought to bear upon a compound, decomposition 
 takes place, and the elements are, as we say, set free. 
 When, for example, an electric current is passed through 
 water, the power which holds together the hydrogen and 
 oxygen is overcome and bubbles of the one gas rise from 
 one pole of the battery and bubbles of the other gas rise 
 from the other pole. This is a simple example of chemi- 
 cal decomposition. Again, when the substance known as 
 red oxide of mercury or mercuric oxide is heated to a 
 sufficiently high temperature a colorless gas is given off 
 from it, and globules. of mercury are formed at the same 
 time. The gas, as will be shown later, is oxygen, so that 
 from the red oxide of mercury, which is a chemical com- 
 pound of mercury and oxygen, we get, by heating, the 
 two elements in the free state. In this case, heat over- 
 comes the chemical attraction which, in the compound, 
 holds the elements together. 
 
 Chemical Affinity. It is evident from what has already 
 been said that there is some power which can hold sub- 
 stances together in a very intimate way, so intimate that 
 
12 INORGANIC CHEMISTRY. 
 
 we cannot recognize them by ordinary means. We do 
 not know what causes the sulphur and iron to combine, 
 but we do know that they combine. Similarly, we do not 
 know what causes a stone thrown in the air to fall back 
 again, but we know that it falls back. It is true we may 
 say that the cause of the falling of the stone is the attrac- 
 tion of gravitation, but this does not give us any real in- 
 formation, for, if we ask what the attraction of gravitation 
 is, we can only answer that it is that which causes all 
 bodies to attract one another. We may also say, and do 
 say, that the cause of chemical combination is chemical 
 affinity. But in so doing we only give a name to something 
 about which we know nothing except the effects it pro- 
 duces. All the chemical changes which are taking place 
 around us may, then, be referred to the operation of chemi- 
 cal affinity. If this power should cease to operate, what 
 would be the result? Nature would be infinitely less 
 complex than it now is. All complex substances w r ould 
 be resolved into the elements of which they are com- 
 posed, and, as far as we know, there would be only about 
 seventy different kinds of substances. All living things 
 would cease to exist, and in their place there would be 
 three invisible gases, and something very much like char- 
 coal. Mountains would crumble to pieces, and all water 
 would disappear giving two invisible gases. The pro- 
 cesses of life in its many forms would be impossible, as, 
 however subtle that which we cpll life may be, we cannot 
 imagine it to exist without the existence of certain com- 
 plex forms of matter ; and, as regards the life process 
 of animals and plants, most complex chemical changes 
 are constantly taking place within them, and these 
 changes are essential to the continuance of life. 
 
 Chemical, Compounds and Mechanical Mixtures. The 
 substances formed by chemical combination of the ele- 
 ments are called chemical compounds. Most substances 
 found in nature are made up of several others. Wood, 
 for example, is very complex, containing a large number 
 of distinct chemical compounds intimately mixed together. 
 Some of these can be isolated, but it is impossible to 
 isolate them all wHh the* means at present at our com- 
 
CHEMICAL COMPOUNDS AND MECHANICAL MIXTURES. 13 
 
 mand. Most of the rocks met with are also quite com- 
 plex, and it is difficult to isolate the constituents. If we 
 look at a piece of coarse-grained granite, we see plainly 
 enough that it contains different things mixed together, 
 and if it be broken up we can pick out pieces of different 
 substances from the mass. If we now examine a piece 
 of each of the different substances thus picked out of the 
 granite, it appears to be homogeneous, i.e. we cannot 
 recognize the presence of more than one kind of thing in 
 -any one piece. If the piece is carefully selected it may 
 be powdered finely in an agate mortar, and some of the 
 powder examined with a microscope without the presence 
 of more than one substance being recognized. We are 
 &ble to isolate three substances from granite by simply 
 breaking it up and picking out the pieces of different 
 kinds. We might therefore conclude that granite con- 
 sists of three substances. This is true, but it is not the 
 whole truth. For it is possible by proper means to get 
 simpler substances from each of the three already sep- 
 arated. This is, however, a much more difficult process 
 than the separation first accomplished. To effect the sep- 
 aration of each of the three constituents of granite into 
 its elements requires more powerful means. Substances 
 must be brought in contact with them which act upon 
 them, changing their composition, i.e. act chemically 
 upon them, and high heat must be used to aid the action. 
 By skilful work it is, however, possible to separate the 
 three components of granite into their elements. 
 
 From the above it is evident that substances may be 
 united in different ways. They may be so united that it is 
 a simple thing to separate them by mechanical processes. 
 Or they may be so united that it is impossible to separate 
 them by mechanical processes. By a mechanical process 
 is meant any process which does not involve the use of 
 heat, electricity, or chemical change. Thus, the mechan- 
 ical process made use of in the case of granite consisted 
 in picking out the pieces. The separation of the parti- 
 cles of different sizes by means of a sieve is a mechanical 
 process. The separation of two liquids which do not mix 
 with each other is a mechanical process. Complex sub- 
 
14 INORGANIC CHEMISTRY. 
 
 stances which may be separated into their components by 
 purely mechanical processes are called mechanical mix- 
 tures. Thus granite is a mechanical mixture of three 
 chemical compounds. Similarly, most natural substances 
 are more or less complex mixtures of chemical com- 
 pounds, or, much more rarely, of elements. Air, for ex- 
 ample, is a mechanical mixture consisting mainly of the 
 two elements nitrogen and oxygen. It is not always an 
 easy matter to distinguish between mechanical mixtures 
 and chemical compounds, as there are mixtures which it is 
 extremely difficult to subdivide into their components, and 
 there are, on the other hand, chemical compounds which 
 are extremely unstable. Generally, however, the differ- 
 ence is recognized without serious difficulty. 
 
 Qualitative and Quantitative Study of Chemical Changes. 
 In general there are two ways in which chemical 
 changes may be studied. Substances may be brought 
 together under a variety of conditions and, if action takes 
 place, the properties of the product or products may then 
 be studied and compared with those of the substances 
 brought together. In the early periods of the history of 
 chemistry the study was almost wholly of this kind. This 
 is called qualitative study. But we may go farther than 
 this, and take into consideration the weights or masses 
 of the substances we are dealing with. We should then 
 be studying the changes quantitatively. We have already 
 seen that by means of the quantitative method Lavoisier 
 placed the law of the % indestructibility of matter upon a 
 firm basis, and that he also succeeded by the use of this 
 method in explaining a number of important chemical 
 changes, particularly combustion. By further use of this 
 method other laws of the highest importance to the science 
 of chemistry were soon brought to light. 
 
 Law of Definite Proportions. The fact that sulphur and 
 iron combine chemically when a mixture of the two is 
 heated has been referred to. The question whether they 
 combine in all proportions is one which can be answered 
 only by a quantitative study of the process. If the pro- 
 cess were to be studied for the first time the method of 
 procedure would be this : We should mix the elements 
 
LAW OF DEFINITE AND OF MULTIPLE PROPORTIONS. 15 
 
 in different proportions and, after the action, we should 
 determine whether any of either of the elements is left in 
 the uncombined state ; and, further, by decomposing the 
 product, we should determine whether it always contains 
 the elements in the same proportions. The problem, in this 
 case, is by no means a simple one, but it has been repeat- 
 edly worked over with the greatest possible care, and, as 
 the result of the work, the conclusion is justified that the 
 product always contains the elements in exactly the same 
 proportions. Similar work has been done for most other 
 chemical compounds known, and the general conclusion 
 known as the law of definite proportions has been drawn. 
 This law may be stated thus : 
 
 A chemical compound always contains the same constitu- 
 ents in the same proportion by weight. 
 
 The truth of this general statement or law has not al- 
 ways been acknowledged by chemists. At the beginning 
 of this century a celebrated discussion on the subject took 
 place between Proust and Berthollet. The discussion 
 led to a great deal of careful work which tended to con- 
 firm the law, and since that time it has not been seriously 
 doubted. About twenty years ago a Belgian chemist, 
 Stas, by a long series of probably the most painstaking 
 and accurate chemical experiments ever performed, 
 showed that in the compounds which he worked with there 
 was no variation in composition that could be detected 
 by the most refined methods of chemistry. In the pres- 
 ent state of our knowledge it appears that the law of 
 definite proportions is a law in the strictest sense. 
 
 Law of Multiple Proportions. It does not require a very 
 extended study of chemical phenomena to show that from 
 the same elements it is possible in many cases to get 
 more than one product. Thus iron and sulphur form 
 three distinct compounds with each other. Tin combines 
 with oxygen in two proportions. The elements potassium, 
 chlorine, and oxygen combine in four different ways, form- 
 ing four distinct products. Nitrogen and oxygen form 
 five products. In the early part of this century the Eng- 
 lish chemist Dalton by a study of cases like those men- 
 tioned was led to the discovery of another great law of 
 
16 INORGANIC CHEMISTRY. 
 
 chemistry known as the law of multiple proportions* 
 Many substances had been analyzed before his time, and 
 the percentages of the constituents determined with a fair 
 degree of accuracy. He examined, first, two gases, both 
 of which consist of carbon and hydrogen. He determined 
 the percentages of the constituents, and found them to- 
 be as follows : 
 
 Olefiant gas, 85.7 per cent, carbon and 14.3 per cent, 
 hydrogen. 
 
 Marsh gas, 75.0 per cent, carbon and 25.0 per cent, hy- 
 drogen. 
 
 On comparing these numbers, he found that the ratio of 
 carbon to hydrogen in olefiant gas is as 6 to 1 ; whereas 
 in marsh gas it is as 3 to 1 or 6 to 2. The mass of hy- 
 drogen, combined with a given mass of carbon, is exactly 
 twice as great in the one case as in the other. 
 
 There are, further, two compounds of carbon and oxy- 
 gen, and in analyzing these the following figures were 
 obtained : 
 
 Carbon monoxide, 42.86 per cent, carbon and 57.14 per 
 cent, oxygen. 
 
 Carbon dioxide, 27.27 per cent, carbon and 72.73 per 
 cent, oxygen. 
 
 But 42.86 : 57.14 :: 6 : 8 and 27.27 : 72.73 :: 6 : 16. 
 
 The mass of oxygen combined with a given mass of 
 carbon in carbon dioxide is exactly twice as great as the 
 mass of oxygen combined with the same mass of carbon in 
 carbon monoxide. These facts and other similar ones led 
 to the discovery of the law of multiple proportions, which 
 may be stated thus: 
 
 If two dements A and B form several compounds with 
 each other, and we consider any fixed mass of A, then the 
 different masses of B which combine with the fixed mass of 
 A bear a simple ratio to one another. 
 
 By way of further illustration we may take the three 
 compounds which iron forms with sulphur. In one of 
 these, approximately 7 parts of iron are in combination 
 with 4 parts of sulphur ; in a second, 7 parts of iron are in 
 combination with 6 parts of sulphur ; and in the third, 7 
 of iron are in combination with 8 of sulphur. The figures 
 
COMBINING WEIGHTS OF THE ELEMENTS. 17 
 
 '4, 6, and 8 bear a simple ratio to one another which is 
 2:3:4. The five compounds of nitrogen and oxygen 
 contain 7 parts of nitrogen combined with 8, 16, 24, 32, 
 and 40 parts of oxygen respectively. The figures repre- 
 senting the parts by weight of oxygen combined with 7 
 parts by weight of nitrogen are in the ratio 1:2:3:4:5. 
 In the compounds formed by the elements chlorine, 
 potassium, and oxygen the proportions by weight are 
 represented in the following table : 
 
 Chlorine. Potassium. Oxygen. 
 35.18 38.82 15.88 
 
 35.18 38.82 31.76 
 
 35.18 38.82 47.64 
 
 35.18 38.82 63.52 
 
 It will be observed that the ratio between the chlorine 
 and potassium remains constant, but that the mass of 
 oxygen varies regularly from 15.88 to 63.52 ; the masses 
 bearing to one another the simple ratio 1:2:3:4. 
 
 The law of multiple proportions like the law of defi- 
 nite proportions is simply a statement in accordance 
 with what has been found true by experiment. Although 
 discovered by Dalton at the beginning of this century 
 and put forward upon what appears now to be only a 
 slight basis of facts, all work since that time has con- 
 firmed it, and it forms to-day one of the corner-stones of 
 the science of chemistry. 
 
 Combining Weights of the Elements. A careful study 
 of the figures representing the composition of chemical 
 compounds reveals a remarkable fact regarding the rela- 
 tive quantities of one and the same element which enter 
 into combination with different elements. The propor- 
 tions by weight in which some of the elements combine 
 chemically with one another are stated in the following 
 table : 
 
 1 part Hydrogen combines with 35.18 parts Chlorine. 
 
 1 " " " " 79.34 
 
 1 " " " " 125.89 
 
 35.18 parts Chlorine combine " 38.82 
 
 79.34 " Bromine " " 38.82 
 
 125.89 " Iodine " " 38.82 
 
 Bromine. 
 
 Iodine. 
 
 Potassium. 
 
18 
 
 INORGANIC CHEMISTRY. 
 
 15.88 parts Oxygen combine with 64.91 parts Zinc. 
 
 15.88 
 
 " 
 
 15.88 
 
 " 
 
 15.88 
 
 " 
 
 64.91 
 
 Zinc 
 
 34.10 
 39.76 
 
 Magnesium 
 Calcium 
 
 136.39 
 
 Barium 
 
 24.10 
 b9.76 
 
 Magnesium. 
 Calcium. 
 
 136.39 
 
 Barium. 
 
 31.83 
 31.83 
 
 Sulphur. 
 
 31.83 
 
 " 
 
 31.83 
 
 
 
 It will be seen that the figures which express the rela- 
 tive quantities of chlorine, bromine, and iodine that 
 combine with 1 part of hydrogen also express the rela- 
 tive quantities of these elements that combine with 
 38.82 parts of potassium. So also the figures which ex- 
 press the relative quantities of zinc, magnesium, calcium, 
 and barium that combine with 15.88 parts of oxygen 
 also express the relative quantities of these elements 
 that combine with 31.83 parts of sulphur. Now, an 
 examination of all compounds known has shown that 
 hydrogen enters into combination with the other elements 
 in the smallest proportions ; it is therefore taken as unity 
 in stating the relative quantities of the other elements 
 which enter into combination. The weight of another 
 element that combines with 1 part by weight of hydro- 
 gen may be called its combining iveight. Thus, according 
 to the above, the combining weights of chlorine, bromine, 
 and iodine are respectively 35.18, 79.34, and 125.89. 
 Similarly 38.82 is the combining weight of potassium, 
 as it expresses the weight of potassium that combines 
 with the above quantities of chlorine, bromine, and 
 iodine. Thus for every element a number can be se- 
 lected, such that the proportions by weight in which the 
 element enters into combination with others can be con- 
 veniently expressed by the number or by a simple multiple 
 of it. These numbers are the combining weights. 
 
 It is not by any means an easy matter to determine 
 which numbers are most convenient for all circumstances ; 
 and if the selection is to be determined solely by con- 
 venience, there may be differences of opinion as to what 
 is most convenient. We shall see a little later that, 
 while the numbers primarily express the combining 
 
SYMBOLS AND ATOMIC WEIGHTS OF THE ELEMENTS. 19 
 
 weights and nothing else, and are based solely upon a 
 study of the composition of chemical compounds, they 
 have come to have a deeper significance which will be 
 explained in due time. They are now called atomic 
 iveights because for strong reasons they are believed to 
 express the relative weights or masses of the minute in- 
 divisible particles or atoms of which the various kinds of 
 matter are assumed to be made up. The atomic theory, 
 as it is called, will be treated of farther on, and the re- 
 lation which exists between the theory and the figures 
 called the atomic weights will be discussed at some 
 length. For the present it will be best to use the figures 
 as expressing the combining weights, and as being en- 
 tirely independent of any speculations regarding the con- 
 stitution of matter and the existence of atoms. 
 
 The Elements, their Symbols and Atomic Weights. It 
 has already been stated that there are about seventy 
 elementary substances known, but that of these only a 
 small number enter into the composition of common 
 things to any great extent. It has been calculated that 
 the solid crust of the earth is made up approximately as 
 represented in the subjoined table : 
 
 Oxygen 47.13$ 
 
 Silicon 27.89$ 
 
 Aluminium 8.13$ 
 
 Iron... 4.71$ 
 
 Calcium 3.53$ 
 
 Magnesium 2. 64$ 
 
 Sodium 2.68$ 
 
 Potassium . 2.35$ 
 
 While oxygen forms a large proportion of the solid 
 crust of the earth, it forms a still larger proportion (eight- 
 ninths) of water, and about one-fifth of the air. Carbon 
 is the principal element which enters into the structure 
 of living things, while hydrogen, oxygen, and nitrogen 
 also are essential constituents of animals and plants. 
 Nitrogen forms about four-fifths of the air. 
 
 In representing the results of chemical action, it is con- 
 venient to use abbreviations for the names of the elements 
 and compounds. Thus, instead of oxygen we may write 
 simply O ; for hydrogen, H ; for nitrogen, N ; etc. These 
 symbols are used in expressing what takes place when sub- 
 
20 INORGANIC CHEMISTRY. 
 
 stances act upon one another. Very frequently the first 
 letter of the name is used as the symbol. If the names- 
 of two or more elements begin with the same letter, 
 this letter is used, and some other letter of the name is 
 added. Thus, B is the symbol of boron, Ba of barium, 
 Bi of bismuth ; C is the symbol of carbon, Ca of calcium, 
 Cd of cadmium, Ce of cerium, Cl of chlorine, Cr of 
 chromium, Cs of caesium, Cu of copper. In some cases 
 the symbol is derived from the Latin name of the ele- 
 ment. Thus, the symbol of iron is Fe, from the Latin 
 ferrum; for copper Cu, from cuprum; for mercury Hg> 
 from hydrargyrum; etc. 
 
 The names themselves are formed in a variety of ways. 
 Chlorine is derived from the Greek ^/Vc^po?, which means- 
 yellowish-green, as this is the color of chlorine. Bro- 
 mine comes from /?pc3/*o, a stench, a prominent charac- 
 teristic of bromine being its bad odor. Hydrogen comes, 
 from vdoop, water, and ysreiv, to produce, signifying that 
 it is a constituent of water. Similarly nitrogen comes- 
 from virpov, niter, and yeveiv, to produce, nitrogen be- 
 ing one of the constituents of niter. Potassium is an 
 element found in potash, and sodium is found in soda. 
 Some elements have been named after the country in 
 which they were first discovered. Thus we have gallium, 
 discovered in France ; scandium, discovered in Sweden ; 
 germanium, discovered in Germany. Tantalum was so 
 called on account of the long-continued difficulties ex- 
 perienced in isolating it when it was discovered. Colum- 
 bium received its name from the fact that it occurs in 
 the mineral columbite, and this owes its name to the fact 
 that it was first found in the United States of America. 
 
 Below is given a table containing the names of all 
 the elementary substances now known, together with 
 their symbols and atomic weights. The names of those 
 which are most widely distributed, and form by far the 
 largest part of the earth, are printed in SMALL CAPITALS. 
 The names of those which are rare are printed in italics. 
 
SYMBOLS AND ATOMIC WEIGHTS OF THE ELEMENTS. 21 
 
 Element. Symbol. 
 
 ALUMINIUM Al 
 
 Antimony Sb 
 
 Argon A 
 
 Arsenic As 
 
 Barium Ba 
 
 Bismuth Bi 
 
 Boron B 
 
 Bromine Br 
 
 Cadmium Cd 
 
 Ooesium Cs 
 
 CALCIUM Ca 
 
 CARBON C 
 
 Cerium Ce 
 
 CHLORINE Cl 
 
 Chromium Cr 
 
 Cobalt.. Co 
 
 Columbium Cb 
 
 Copper Cu 
 
 Erbium E 
 
 Fluorine F 
 
 Gadolinium Gd 
 
 Gallium Ga 
 
 Germanium Ge 
 
 Glucinum Gl 
 
 Gold An 
 
 Helium He 
 
 HYDROGEN H 
 
 'Indium In 
 
 Iodine I 
 
 Iridium Ir 
 
 IRON Fe 
 
 Lanthanum La 
 
 Lead Pb 
 
 Lithium Li 
 
 MAGNESIUM Mg 
 
 Manganese. . . , Mn 
 
 Mercury Hg 
 
 Atomic 
 Weight. 
 26.91 
 119.52 
 
 (?) 
 74.44 
 
 Element. 
 Molybdenum 
 Neodymium 
 Nickel 
 NITROGEN . . . . 
 
 Symbol. 
 
 .... Mo 
 .... Nd 
 .... Ni 
 
 N 
 
 136.39 
 
 Osmium.. 
 
 Os 
 
 206.54 
 
 OXYGEN . 
 
 o 
 
 10.86 
 
 Palladium.. 
 
 Pd 
 
 79.34 
 111.10 
 
 Phosphorus 
 Platinum 
 
 .... P 
 Pt 
 
 131.89 
 
 POTASSIUM. 
 
 K 
 
 39.76 
 
 Pi~aseodymium. . . 
 
 Pr 
 
 11.92 
 
 JiJiodium 
 
 Rh 
 
 139.10 
 35.18 
 
 Rubidium 
 Ruthenium 
 
 .... Rb 
 Ru 
 
 51.74 
 
 Samarium.. 
 
 Sm 
 
 58.49 
 
 Scandium... 
 
 Sc 
 
 93.02 
 63.12 
 
 Selenium 
 SILICON 
 
 .... Se 
 Si 
 
 165.06 
 
 Silver 
 
 Ag 
 
 18.91 
 
 SODIUM 
 
 . .. Na 
 
 155.57 
 69.38 
 
 Strontium 
 Sulphur . . 
 
 .... Sr 
 .. S 
 
 71.93 
 
 Tantalum 
 
 Ta 
 
 9.01 
 
 Tellurium 
 
 Te 
 
 195.74 
 
 Terbium. . 
 
 Tr 
 
 (?) 
 1.00 
 
 Thallium 
 Thorium 
 
 .... Tl 
 Th 
 
 112.99 
 
 Thulium 
 
 Tm 
 
 125.89 
 
 Tin 
 
 Sn 
 
 191.66 
 
 Titanium 
 
 Ti 
 
 55.60 
 
 Tungsten 
 
 W 
 
 137 59 
 
 Uranium 
 
 u 
 
 205.36 
 
 Vanadium 
 
 .... V 
 
 697 
 
 Ytterbium 
 
 Yb 
 
 24.10 
 
 Yttrium. 
 
 .. Y 
 
 54.57 
 
 Zinc 
 
 . ... Zn 
 
 198.49 
 
 Zirconium.. . 
 
 ,. Zr 
 
 Atomic 
 Weight. 
 
 95.26 
 
 139.70 
 
 58.24 
 
 13.93 
 
 189.55 
 
 15.88 
 
 105.56 
 
 30.79 
 
 193.41 
 
 38.82 
 
 142.50 
 
 , 102.23 
 
 84.78 
 
 100.91 
 
 149.13 
 
 43.78 
 
 78.42 
 
 28.18 
 
 107.11 
 
 22.88 
 
 86.95 
 
 31.83 
 
 181.45 
 
 126.52 
 
 158.80 
 
 202.61 
 
 230.87 
 
 169.40 
 
 118.15 
 
 47.79 
 
 183.43 
 
 237.77 
 
 50.99 
 
 171.88 
 
 88.35 
 
 64.91 
 
 89.72 
 
 The symbols of the elements represent not only the 
 names but relative quantities. Thus O stands for 15.88 
 parts by weight of oxygen ; N for 13.93 parts by weight 
 of nitrogen ; and hydrogen being that element which 
 enters into combination in the smallest relative quantity, 
 it is taken as the basis of the system, or H stands 
 for 1 part by weight of hydrogen. What the symbol O 
 really means then is that the quantity of matter repre- 
 sented by it is 15.88 times as great as the quantity of 
 matter represented by the symbol H ; and the quantity 
 of matter represented by N is 13.93 times as great as that 
 represented by the symbol H ; and so on through the 
 list. The figures do not represent absolute but relative 
 masses. There are very serious difficulties encountered 
 in determining the combining weights of the elements, 
 and in regard to several given in the above table there is 
 
22 INORGANIC CHEMISTRY. 
 
 considerable doubt as to the accuracy. Those of the ele- 
 ments with which we most frequently have to deal have, 
 however, been determined .with great care. Work in this 
 field is being constantly carried on, and every year our 
 knowledge in regard to the combining weights becomes 
 more and more accurate. 
 
 Symbols of Compounds. As the elements enter into 
 combination in the proportion of their respective combin- 
 ing weights or simple multiples of these weights, it is an 
 easy matter to represent the composition of compounds 
 by means of the symbols. Thus hydrogen and chlorine 
 combine in the proportion of their combining weights to 
 form the compound hydrochloric acid. The compound 
 is represented by the symbol HC1, which signifies that 
 the compound contains hydrogen and chlorine in the 
 proportion of 35.18 parts of chlorine to 1 part of hydro- 
 gen. So the symbol ZnO means a chemical compound 
 consisting of 64.91 parts of zinc and 15.88 parts of oxygen ; 
 HCN means a compound made up of 1 part of hydrogen, 
 11.92 parts of carbon, and 13.93 of nitrogen. Whenever 
 the symbols of the elements are placed side by side with 
 no sign between them, as in the above examples, the re- 
 sulting symbol means that the elements are in chemical 
 combination. But, as has been pointed out, elements 
 may combine in more than one proportion. In one of 
 the two compounds of carbon and oxygen the elements 
 are combined in the proportion of their combining 
 weights, and the compound is represented by the symbol 
 CO ; in the other compound the elements are combined 
 in the proportion of twice the combining weight of oxy- 
 gen to the combining weight of carbon, and the com- 
 pound is represented by the symbol CO 2 . The three 
 compounds of iron and sulphur to which reference has 
 already been made are represented by the symbols FeS, 
 Fe 2 S 3 , and FeS 2 . The first represents a compound in 
 which the elements are combined in the proportion of 
 their combining weights, or 55.60 parts iron to 31.83 
 parts sulphur ; the second represents a compound in 
 which the elements are combined in the proportion of 
 twice the combining weight of iron (2 X 55.60 = 111.20 
 
SYMBOLS OF COMPOUNDS- CHEMICAL EQUATIONS. 23 
 
 parts) to three times the combining weight of sulphur 
 (3 X 31.83 = 95.49 parts) ; and the third represents a com- 
 pound in which the elements are combined in the pro- 
 portion of the combining weight (55.60 parts) of iron to 
 twice the combining weight (2 X 31.83 = 63.66 parts) of 
 sulphur. The four compounds of potassium, chlorine, 
 and oxygen above mentioned are represented by the sym- 
 bols KC1O, KC1O 2 , KC1O 3 , and KC1O 4 , the meaning of 
 which will be clear from the explanation just given. By 
 means of such symbols all chemical compounds can be 
 represented, and they represent not only what elements 
 are contained in the compounds, but in what proportions 
 the elements are combined. They represent facts which 
 have been determined by experiment. Knowing the act- 
 ual weight of one constituent of any compound we can 
 calculate by the aid of the symbol the actual weights of 
 the other constituents and of the compound itself. 
 Thus, if we know the actual weight of the chlorine con- 
 tained in a quantity of potassium chlorate, KC1O 3 , we 
 can calculate how much potassium and how much oxygen 
 are contained in that same quantity, and also what the 
 quantity of potassium chlorate is. Suppose, for example, 
 we know that in a certain quantity of potassium chlorate 
 there is contained 25 grams of chlorine, and it is desired 
 to know how much potassium and how much oxygen 
 there are in this quantity, and also what the quantity of 
 potassium chlorate is. We know that the compound 
 KC1O 8 is made up of 38.82 parts of potassium, 35.18 
 parts of chlorine, and 3 times 15.88, or 47.64, parts of 
 oxygen, the whole making 122.28 parts. The solution 
 of the following equations in proportion will give the 
 quantities desired : 
 
 35.18 
 35.18 
 35.18 
 
 25 38.82 : weight of potassium ; 
 
 25 47.64 : " " oxygen ; 
 
 25 121.64 : " " potassium chlorate. 
 
 Chemical Equations. In dealing with cases of chemi- 
 cal action it is desirable to express by means of the 
 symbols which represent the elements and compounds 
 what takes place. In general, a chemical change is called 
 
24 INORGANIC CHEMISTRY. 
 
 a chemical reaction, and these reactions are of three 
 kinds : 
 
 (1) Two or more elements or compounds may unite 
 directly to form one product. This is called combination. 
 The following examples will suffice. When mercury is kept 
 boiling in the air for a time it becomes covered with a layer 
 of a red substance which is a compound of mercury and 
 oxygen represented by the symbol HgO. Magnesium 
 burns in the air and forms the compound MgO. Hydro- 
 chloric acid, HC1, combines directly with ammonia, NH 3 , 
 and forms the compound known as ammonium chloride, 
 NH 4 C1. Water, H 2 O, combines directly with lime or cal- 
 cium oxide, CaO, to form slaked lime or calcium hydrox- 
 ide, CaO 2 H 2 . To represent these facts, the symbols of 
 the elements or compounds which act upon each other 
 are written with a plus sign between them, and the sign 
 of equality is written before the symbol of the product. 
 The chemical equations which represent the above-men- 
 tioned chemical reactions are : 
 
 Hg +0 -HgO; 
 Mg + O = MgO ; 
 
 H 2 O + CaO = CaO 2 H 2 . 
 
 In reading the equations the plus sign is generally ren- 
 dered by and, and the sign of equality by give. The first 
 equation should accordingly be read, " Mercury and oxy- 
 gen give mercuric oxide ;" but it represents besides this 
 fact the exact relations which exist between the quantities 
 of the elements and the compound which take part in the 
 reaction. 
 
 (2) The second kind of chemical reaction is decomposi- 
 tion or the opposite of combination. Examples are fur- 
 nished by the decomposition of mercuric oxide into mer- 
 cury and oxygen by heat ; of potassium chlorate into 
 potassium chloride and oxygen by heat ; of water into 
 hydrogen and oxygen by the electric current ; and of cal- 
 cium carbonate or limestone, CaCO 3 , into lime or calcium 
 oxide, CaO, and carbon dioxide, CO 2 , by heat. These re- 
 actions are represented by the following equations : 
 
CHEMICAL EQUATIONS THE SCOPE OF CHEMISTRT. 25 
 
 HgO =Hg +0; 
 KC1O 3 = KC1 + 3O ; 
 H 2 = 2H +0; 
 CaCO 3 = CaO + CO 3 . 
 
 The expressions 3O and 2H mean respectively three 
 times the combining weight of oxygen and twice the com- 
 bining weight of hydrogen, the figure being generally 
 used in this way when the element is not in combination. 
 It is, however, sometimes written the same as if the ele- 
 ment were in combination, as will be explained later. 
 
 (3) The third kind of chemical reaction is that in 
 which two or more compounds give rise to the formation 
 of two or more others ; or an element and a compound 
 may act in such a way as to give another compound and 
 another element. This is called double decomposition 
 or metathesis. The following cases will serve as exam- 
 ples : Sulphuric acid, H 2 SO 4 , acts upon potassium nitrate, 
 KNO 3 , or saltpeter, forming potassium sulphate, K 2 SO 4 , 
 and nitric acid, HNO 3 ; nitric acid, HNO 3 , acts upon 
 sodium carbonate, Na 2 CO 3 , forming sodium nitrate, 
 NaNO 3 , carbon dioxide, CO 2 , and water, H 2 O ; hydro- 
 chloric acid, HC1, and zinc, Zn, give zinc chloride, ZnCl 2 , 
 and hydrogen. These facts are represented as below : 
 
 H 2 S0 4 + 2KN0 3 = K 2 SO 4 + 2HNO 3 ; 
 2HN0 3 + Na 2 C0 3 = 2NaNO 3 + CO 2 + H 2 O ; 
 2HC1 +Zn =ZnCl 2 +2H. 
 
 In the expressions 2KNO 3 , 2HNO 3 , 2NaNO 3 , and 2HC1, 
 the large figures placed before the symbol of the com- 
 pounds signify that the quantities of the compounds repre- 
 sented by the symbol are to be multiplied by the figure. 
 Thus, HC1 stands for 1 + 35.18 = 36.1.8 parts of hydro- 
 chloric acid ; but 2HC1 stands for 2(1 + 35.18) = 72.36 
 parts ; 3HC1 stands for 3(1 + 35.18) = 108.54 parts ; etc. 
 
 The Scope of Chemistry. A complete study of chem- 
 istry would involve the study of the action of all the 
 elements upon one another under all possible circum- 
 stances, and a study of the action of all compounds 
 upon one another and upon the elements under all cir- 
 
26 INORGANIC CHEMISTRY. 
 
 cumstances. This indicates that the field is almost 
 boundless ; and if the facts were not related among one 
 another, if every time a reaction is studied we are to ex- 
 pect something entirely different from all other reactions, 
 the task would be practically hopeless. Fortunately a 
 great many general facts are known, and reactions which 
 at first seem to have no connection are by careful study 
 shown to be related. Thus the study is very much sim- 
 plified and made interesting. It must be our purpose to 
 study the facts in as systematic a way as possible, and to 
 be constantly on the alert to detect relations. The habit 
 of comparing a new reaction with others already studied 
 should be cultivated. In this way light will come out of 
 the darkness, and the subject will gradually become clear. 
 While the simplest way to begin the study of chemistry 
 is by a consideration of the elements, the subject is com- 
 plicated by the fact that we cannot readily obtain these 
 elements without the aid of substances which have not 
 been studied, and of processes which are incomprehen- 
 sible. There are, however, two elements that occur in 
 nature in enormous quantities, that can be obtained in 
 the iincombined condition quite easily. As the kinds 
 of action which they exhibit are of great importance 
 and well calculated to give an insight into the nature of 
 chemical action in general, we may profitably begin our 
 study of chemical phenomena by a study of these two 
 elements. They are oxygen and hydrogen. In learning 
 the main facts in regard to these two elements we shall 
 learn a great deal that will be of importance in enabling 
 us to understand other chemical phenomena ; we shall 
 begin to learn how to study things chemically ; and we 
 shall thus prepare ourselves for a systematic study of 
 the science of chemistry. 
 
 Chemical Action accompanied by other kinds ol Action. 
 Whenever a chemical change takes place it is accom- 
 panied by other changes and, in order to gain a com- 
 plete knowledge of the phenomenon, these other changes 
 must be studied. Thus when sulphuric acid acts upon 
 zinc the chemical change is represented both qualitatively 
 and quantitatively by the equation 
 
CHEMICAL REACTIONS. 21 
 
 Zn + H 2 SO 4 = ZnSO 4 + 2H. 
 
 In studying the reaction, the first thing to do is to learn 
 the nature of the substances formed, and the relations 
 between the substances which act upon each other and 
 the products. This may be called the purely chemical 
 study of the reaction. But much more can be learned 
 in regard to it by careful observation. In the first place, 
 we must take into account the fact that a solid and liquid 
 here react to form a solid and a gas ; and the ques- 
 tion suggests itself, does this change to the gaseous con- 
 dition exert any influence on the reaction, or is this a fact 
 of no special importance ? Again, it will be observed 
 that accompanying the chemical change there is a marked 
 rise in temperature, and we naturally inquire whether the 
 quantity of heat evolved is* definite for definite quantities 
 of the substances, and, if so, what relation exists be- 
 tween them. There are still other changes which must be 
 taken into account in order to get a complete knowledge 
 of a chemical reaction, but, as yet, the study of the other 
 changes has not been taken up in a general way, and our 
 information in regard to them is comparatively limited. 
 Within late years much progress has been made in the 
 study of the heat changes which accompany chemical 
 changes. It has been found that every chemical change 
 gives rise either to an evolution or to an absorption of 
 heat, and that for definite quantities of the same sub- 
 stances under the same circumstances the same amount 
 of heat is evolved or absorbed. The special study of the 
 heat changes connected with chemical changes is called 
 thermochemistry. A consideration of the facts and laws 
 of thermochemistry is of assistance in dealing with chemi- 
 cal reactions, and some attention will be paid to the sub- 
 ject in this book. 
 
CHAPTER II. 
 
 A STUDY OF THE ELEMENT OXYGEN. 
 
 Historical. The older chemists considered air to be a 
 simple substance, but the experiments of Priestley (1774) 
 and Scheele (1775) showed that the air contains two gases 
 only one of which has the power to support combustion ; 
 and they succeeded independently of each other in show- 
 ing that oxygen is a distinct substance. The discovery 
 of oxygen had a very important bearing on the work of 
 Lavoisier on combustion, and it was he who gave the 
 name oxygen (or oxygene) to the gas, for the reason that 
 he supposed it to be the essential constituent of all those 
 chemical substances which are known as acids, the word 
 being derived from the Greek ogvcr, acid, and yeveiv, to 
 produce. While this is generally true, it has since been 
 found that there are other elements which have the 
 power to give acid properties to the substances into the 
 composition of which they enter, and, therefore, the name 
 is misleading. 
 
 Occurrence. Oxygen is the most widely distributed 
 and most abundant element of the earth. It forms, as 
 has been stated, about 47 per cent of the solid crust of 
 the earth ; eight-ninths of water ; and about one-fifth 
 of the air. It occurs also in combination with carbon 
 and hydrogen, or with carbon, hydrogen, and nitrogen in 
 the substances which go to make up the structure of liv- 
 ing things, whether vegetable or animal. Besides this it 
 forms a part of most manufactured chemical products. 
 
 Preparation. Notwithstanding the abundant supply of 
 oxygen in nature it is not a simple matter to get it in the 
 free or uncombined state from most substances found in 
 nature. As it forms eight-ninths of water, and water 
 consists of only hydrogen and oxygen, the idea suggests 
 itself at once that it may be made by the decomposition 
 
 (28) 
 
OXYGEN PREPARATION. 29 
 
 of water. This can be accomplished without serious 
 difficulty by means of an electric current, and both hy- 
 drogen and oxygen obtained in this way ; but the method 
 is expensive and more complicated than others which are 
 available, and therefore it is not commonly used for the 
 purpose. In the air the two gases nitrogen and oxygen 
 are mixed together in the proportion of 1 volume of 
 oxygen to 4 volumes of nitrogen. Here then, too, as in 
 water, we have an enormous supply, but it is difficult to 
 separate the oxygen from the nitrogen in such a way as 
 to leave it uncombined. This can, however, be accom- 
 plished, and a method is now in practical use on the 
 large scale for the purpose of preparing oxygen from the 
 air. The method is based upon the fact that when 
 barium oxide, BaO, is heated in a current of air it takes 
 up oxygen and is converted into barium dioxide, BaO 2 ; 
 and when the pressure upon the dioxide is sufficiently 
 diminished, it is decomposed into the oxide, BaO, and 
 oxygen, as represented in the equation 
 
 By this means the oxygen can be extracted from the air 
 and obtained in the free condition. 
 
 Among substances which occur in nature and which 
 can be used for the preparation of oxygen, manganese 
 dioxide or pyrolusite, also called the black oxide of man- 
 ganese, MnO 2 , is the most important. It gives off a part 
 of its oxygen when heated to a comparatively high tem- 
 perature. It has been shown that the decomposition is 
 represented by this equation : 
 
 As will be seen, only one-third of the oxygen contained 
 in the dioxide is thus obtained in the free state. A simi- 
 lar method is that used by Priestley when he discovered 
 oxygen. It consists simply in heating mercuric oxide, 
 HgO, when it is decomposed as represented thus : 
 
 HgO = Hg + O. 
 
 The substance potassium chlorate, KC1O 3 , is manufac- 
 tured on the large scale for a variety of purposes and is, 
 
30 INORGANIC CHEMISTRY. 
 
 therefore, easily obtained. It gives up its oxygen when 
 heated. At first the decomposition represented by this 
 equation takes place : 
 
 8KC10 3 = 5KC1O, + 3KC1 + 4O. 
 
 The products are potassium perchlorate, KC1O 4 , potas- 
 sium chloride, KC1, and oxygen, one-sixth of the total 
 oxygen being given off in the first stage. This part of the 
 decomposition takes place readily, and at a comparatively 
 low temperature. If, after it is complete, the tempera- 
 ture is raised considerably higher, more gas is given off 
 and the change represented by the equation 
 
 KC1O 4 = KC1 + 4O 
 
 is accomplished. The final result is, therefore, the setting 
 free of all the oxygen contained in the chlorate. This- 
 fact is represented thus : 
 
 KC10 3 = 
 
 The best method for use in the laboratory for the 
 preparation of oxygen consists in heating a mixture of 
 equal parts of coarsely powdered manganese dioxide and 
 potassium chlorate. This mixture gives up oxygen very 
 readily under the influence of heat. Potassium chlorate 
 alone requires to be heated to a temperature of over 
 350 C. to effect its decomposition, but when mixed with 
 manganese dioxide the decomposition takes place at 
 about 200 C. The manganese dioxide does not lose any 
 of its oxygen under the circumstances. Other substances, 
 such as ferric oxide, copper oxide, etc., may be used with 
 similar effect. No satisfactory explanation of the action 
 of these substances has been given. Recent experi- 
 ments have shown that, when manganese dioxide is 
 used, oxygen, chlorine, and potassium permanganate, 
 KMnO 4 , are first formed. The permanganate is de- 
 composed by heat, yielding the manganate, K 2 MnO 4 , 
 the dioxide, and oxygen ; and, finally, the manganate 
 is decomposed by chlorine, yielding potassium chloride, 
 the dioxide, and oxygen. 
 
 Physical Properties. Oxygen is a colorless, tasteless, 
 inodorous gas. It is only slightly soluble in water, 100 
 
OXYGEN CHEMICAL PROPERTIES. 31 
 
 volumes of water at dissolving 4.1 volumes of oxygen. 
 It is slightly heavier than air, its specific gravity is 
 1.10563, and 1 liter under 760 mm. pressure and at tem- 
 perature weighs 1.429 grams, while a liter of air 
 weighs 1.2932 grams. In dealing with chemical elements 
 and compounds which are gaseous, it is customary to use 
 hydrogen instead of the air as the standard for specific 
 gravity. While the specific gravity of oxygen in terms 
 of air is 1.10563, in terms of hydrogen it is 15.88, or, 
 in other words, a given volume of oxygen weighs 15.88 
 as much as the same volume of hydrogen under the 
 same conditions. Under a pressure of 50 atmospheres 
 and at a temperature 118 it is condeDsed to a liquid 
 of specific gravity 0.978. Liquid oxygen is a pale steel- 
 blue transparent and very mobile liquid which boils 
 at 181.4 at ordinary pressure. When the pressure is 
 reduced or removed, evaporation takes place so rapidly 
 that a part of the oxygen is often frozen to a white solid. 
 Chemical Properties. At ordinary temperatures oxy- 
 gen does not act readily upon most other things, as can 
 be clearly shown by putting a variety of substances in 
 the gas without'heating them. If they are left for a con- 
 siderable time some evidence of change will be observed, 
 but generally the change is extremely slow unless the 
 temperature is raised. At higher temperatures, different 
 for different substances, it combines with all the elements 
 except fluorine, and it acts readily upon a large number 
 of compounds. Its action is generally accompanied by 
 an evolution of heat and light, and the process under 
 these circumstances is called combustion. This action 
 may be illustrated by first heating and then introducing 
 into vessels containing oxygen, sulphur, charcoal, iron in 
 the shape of a steel watch-spring, and a bit of phos- 
 phorus. The phenomena observed show that chemical 
 action takes place, but they do not show what is formed. 
 It is evident that in each case light and heat are 
 evolved, and that the substances introduced into the 
 oxygen are changed to other things. In the case of 
 phosphorus the light given off is very intense, while in 
 that of carbon and that of sulphur it is only slight. In 
 
32 INORGANIC CHEMISTRY. 
 
 the vessel in which the burning of the iron takes 
 place a reddish-brown substance is deposited, while in 
 that in which the phosphorus is burned dense white 
 fumes are formed and at first the product is partly de- 
 posited upon the walls of the vessel in the form of a 
 white powder that looks like snow. After standing for 
 some time over water it disappears and the water evi- 
 dently contains something in solution. A thorough study 
 of the reactions above mentioned has shown that they 
 consist in the chemical combination of oxygen with the 
 substances burned. The light and heat are results of 
 the chemical action. The reactions are represented by 
 the following equations : 
 
 /t/ # 
 With sulphur, 8 + 2O = SO a ; 
 
 " carbon, C + 2O = CO, ; 
 
 " iron, 3Fe + 4O = Fe 3 O 4 ; 
 
 " phosphorus, 2P + 5O = P 2 O 5 . 
 
 The products are sulphur dioxide, SO 2 , a colorless 
 pungent gas ; carbon dioxide, CO 2 , a colorless gas ; mag- 
 netic oxide of iron, Fe 3 O 4 , a reddish-brown substance ; 
 and phosphorus pentoxide, P a O 5 , a white solid which dis- 
 solves in water. 
 
 Burning in the Air and Burning in Oxygen. One can- 
 not well help noticing a strong resemblance between the 
 burning of substances in the air and in oxygen ; and the 
 question naturally suggests itself, Are these two acts 
 the same in character, or is there a difference between 
 them ? To answer this question, we must burn the same 
 substances in pure oxygen and in air, and determine 
 whether the same products are formed in the two cases, 
 and at the same time whether anything else is formed. 
 If we should make this comparison in any case, we 
 should find that, whether a substance burns in the air or 
 in pure oxygen, the same product is formed, and nothing 
 else. It is therefore certain that the act of burning in 
 the air is due to the presence of oxygen. But substances 
 do not burn as readily in the air as in oxygen, and some 
 which burn in oxygen do not burn in the air. This is 
 
PHLOGISTON THEORY. 33 
 
 due to the fact that only about one-fifth of the volume of 
 the air is oxygen, while most of the remaining four-fifths 
 consists of an extremely inactive element, nitrogen, which 
 takes no part in the process of burning. 
 
 Phlogiston Theory. Fire in its various forms is one of 
 the longest known chemical phenomena. From the earli- 
 est times it has attracted the attention of men and has 
 been the subject of speculative and experimental study. 
 It was one of the elementary substances of Aristotle, as 
 has already been stated. The first comprehensive theory 
 covering all cases of combustion was that put forward by 
 Stahl and known as the pMogiston theory. According to 
 this, every combustible substance contains something, 
 called phlogiston, which escapes in the process of burn- 
 ing. It was repeatedly pointed out that some substances 
 grow heavier by burning, or rather that the products of 
 combustion weigh more than the substance burned. 
 
 This can be shown to be true in the case of zinc by 
 placing some turnings of the metal on one pan of a bal- 
 ance and determining the weight. If now the metal is 
 burned it will change almost completely to a white pow- 
 der, and this will weigh more than the zinc. 
 
 If an ordinary candle is placed on one pan of a bal- 
 ance and a glass vessel open at both ends and filled with 
 large pieces of sodium hydroxide or caustic soda is sus- 
 pended directly over the candle in such way that the 
 smoke from the candle must pass upwards through the 
 tube, and a similar tube is suspended over the other pan 
 of the balance, and equilibrium is then established, it 
 will be found that the pan on which the candle is placed 
 will gradually grow heavier and sink as the candle burns 
 away. The explanation of this is simply that the gases 
 which are formed by the burning of the candle are re- 
 tained by the caustic soda, and they weigh more than the 
 candle from which they were formed by combustion. 
 
 Facts like those mentioned were known when the phlo- 
 giston theory was held, but they do not appear to have 
 made a very strong impression upon chemists. Chemists 
 were at all events not able to give a satisfactory explana- 
 tion. This remained for Lavoisier. 
 
34 INORGANIC CHEMISTRY. 
 
 Lavoisier's Explanation of Combustion. When Lavoi- 
 sier began his work oxygen was unknown, but it was 
 discovered soon afterward ; and this discovery was of the 
 highest importance for the explanation of the phenome- 
 non of combustion. Lavoisier showed that when a sub- 
 stance is burned in oxygen or in the air, some of the gas 
 is used up. He then weighed the substance burned, the 
 oxygen used up, and the product formed, and found that 
 this relation holds good : 
 
 Weight of Sub- , Weight of Oxygen Weight of 
 
 stance burned used up Product formed. 
 
 Having established this relation in a number of cases, 
 it followed that the process of combustion consists in 
 the chemical combination of oxygen with the substance 
 burned. There was no longer room for the hypotheti- 
 cal phlogiston, and since that time it has not occupied 
 a place in the thoughts of most chemists. 
 
 Combustion. By the term combustion in its broadest 
 sense is meant any chemical act which is accompanied 
 by an evolution of light and heat. Ordinarily, however, 
 it is restricted to the union of substances with oxygen, as 
 this union takes place in the air, with evolution of light 
 and heat. Substances which have the power to unite 
 with oxygen are said to be combustible, and substances 
 which have not this power are said to be incombustible. 
 Most elements combine with oxygen under proper con- 
 ditions, and are therefore combustible. Most compounds 
 formed by the union of oxygen with combustible sub- 
 stances are incombustible. For example, the sulphur 
 dioxide, carbon dioxide, magnetic oxide of iron, and phos- 
 phorus pentoxide, formed when sulphur, carbon, iron, and 
 phosphorus are burned in oxygen, are incombustible. 
 They contain oxygen and they cannot combine directly 
 with more. 
 
 Kindling Temperature. We have seen that substances 
 do not usually combine with oxygen at ordinary tempera- 
 tures, but that in order to effect the union the tempera- 
 ture must be raised. If this were not the case, it is plain 
 that every combustible substance in nature would burn up, 
 
KINDLING TEMPERATURE SLOW OXIDATION. 35 
 
 for the air probably supplies a sufficient quantity of oxygen 
 for this purpose. Some substances need to be heated to a 
 high temperature before they will combine with oxygen ; 
 others require to be heated but little. If we were to sub- 
 ject pieces of phosphorus, of sulphur, and of carbon to 
 the same gradual rise in temperature, we should find 
 that the phosphorus takes fire very easily, only a slight 
 elevation of temperature being necessary ; next in order 
 would come the sulphur ; and last the carbon. If we 
 were to repeat these experiments a number of times, we 
 should find that the phosphorus always takes fire at the 
 same temperature, and a similar result would be reached 
 in the case of sulphur and carbon. Every combustible 
 substance has its kindling temperature ; that is, the tem- 
 perature at which it will combine with oxygen. Below 
 this temperature it will not combine with oxygen. If a 
 piece of wood should be heated to its kindling tempera- 
 ture all at once, it would burn up as rapidly as it could 
 secure the necessary oxygen ; but the burning does not 
 usually take place rapidly, for the reason that only a 
 small part of it is at any one time heated to the kind- 
 ling temperature. Watch a stick of wood burning, and 
 see how, as we say, " the fire creeps " slowly along it. 
 The reason of the slow advance is simply this : Only 
 those parts of the stick which are nearest the burning 
 part become heated to the kindling temperature. They 
 take fire and heat the parts nearest them, and so on 
 gradually throughout the length of the stick. 
 
 Slow Oxidation. Substances may combine slowly with 
 oxygen without evolution of light. Thus, if a piece of 
 iron is allowed to lie in moist air, it becomes covered 
 with rust. The rust is similar to the substance formed 
 when iron is burned in oxygen. Both are formed by the 
 union of iron and oxygen. Magnesium burns in the air 
 and forms a white compound containing oxygen. It 
 burns with increased brilliancy in oxygen, forming the 
 same compound. If left in moist air for some days or 
 weeks, it becomes covered with a layer of the same white 
 substance. If this is scraped off, and the magnesium 
 again allowed to lie, it will again become covered 
 
36 INORGANIC CHEMISTRY. 
 
 with a layer of the compound with oxygen, and this may 
 be continued until the magnesium has been completely 
 converted into the same substance that is formed when 
 it burns in oxygen or in the air. Many other cases of 
 slow oxidation might be described, some of which, such 
 as the decay of wood, are constantly taking place. The 
 most important illustration of slow oxidation is that 
 which takes place in our bodies, for, as we shall see, the 
 food of which we partake undergoes a great many 
 changes, some of the substances uniting with oxygen, 
 and thus keeping up the temperature of our bodies. 
 This, however, is done without evolution of light. We 
 take large quantities of oxygen into our lungs in the act 
 of breathing. This acts upon various substances pre- 
 sented to it, oxidizing them to other forms which can 
 easily be got rid of. More will be said in regard to the 
 breathing process of animals and plants when the sub- 
 ject of carbon and its compounds is taken up. 
 
 Heat of Combustion. What is the chief difference be- 
 tween combustion, as we ordinarily understand it, and 
 slow oxidation? As far as can be judged by a cursory 
 examination, it is that in the former there is an evolution 
 of light and much heat, while in the latter there is ap- 
 parently but little heat evolved and no light. Remem- 
 bering that the reason why a body gives light is that it 
 is heated to a sufficiently high temperature, the problem 
 resolves itself into a question of heat. What difference, 
 if any, is there between the quantity of heat given off 
 when a substance burns, and when it undergoes slow 
 oxidation without evolution of light ? Experiment has 
 shown that there is no difference. In one case all the heat 
 is given off in a short space of time, and therefore the 
 temperature of the substance becomes high and it emits 
 light. In the other the heat is given off slowly and con- 
 tinues for a much longer time, and therefore the tem- 
 perature of the substance does not get high, as surround- 
 ing substances conduct off the heat nearly as rapidly as 
 it is evolved. If, however, we were to measure the 
 quantity of the heat, we should find it to be the same in 
 both cases. 
 
HEAT OF COMBUSTION. 37 
 
 We can measure the heat given off or absorbed in a 
 chemical reaction by allowing the reaction to take place 
 in a vessel called a calorimeter, so constructed as to 
 prevent loss of heat, and containing a known weight of 
 water. The temperature of the water is noted at the 
 beginning of the operation and at the end. A quantity 
 of heat is generally stated by giving the number of grams 
 of water which it will raise one degree (Centigrade) in tem- 
 perature. The quantity of heat necessary to raise a gram 
 of water one degree in temperature is the unit used in heat 
 measurements. It is called the calorie. If we say that 
 the quantity of heat evolved in any reaction is 250 cal- 
 ories (written generally 250 cal.), this means simply a 
 quantity of heat capable of raising the temperature of 
 250 grams of water one degree or of one gram of water 
 250 degrees in temperature. Sometimes it is convenient 
 to use a larger unit. The quantity of heat required to 
 raise the temperature of one kilogram of water one de- 
 gree serves the purpose. This is the large calorie. To 
 distinguish it from the smaller one it is written with a 
 capital. Thus, 250 Cal. means 250 large calories. The 
 large calorie is obviously 1000 times greater than the 
 small calorie. 
 
 To repeat, then : by the heat of combustion of a sub- 
 stance is meant simply the quantity of heat given off 
 when a certain weight of the substance combines with 
 oxygen. In order to avoid confusion it is necessary to 
 have an agreement in regard to the weight of substance 
 which shall be used as the standard. This may be a 
 gram or any other weight, but for the purposes of chem- 
 istry it is most convenient to take weights in proportion 
 to the combining or atomic weights. Thus, by the heat 
 of combustion of carbon is meant the quantity of heat 
 evolved by the combination of 11.92 grams of carbon 
 with 2 X 15.88 = 31.76 grams of oxygen. By the heat 
 of combustion of sulphur is meant the quantity of heat 
 evolved in the combination of 31.83 grams sulphur with 
 2 X 15.88 = 31.76 grams oxygen, etc. 
 
 Not only is the heat of combustion the same whether 
 
38 INORGANIC CHEMISTRY. 
 
 the union with oxygen takes place slowly or rapidly, 
 but the heat evolved in any given chemical reaction is 
 always the same, and chemical action is always accom- 
 panied by an evolution or absorption of heat. 
 
 Heat of Decomposition. Just as it is true that a definite 
 quantity of heat is evolved when two or more elements 
 combine chemically, so also it is true that in order to 
 overcome the force which holds these elements together 
 the same quantity of heat is absorbed. Thus, the heat 
 of formation of mercuric oxide, HgO, is 30,660 cal.; or, 
 in other words, when 198.49 grams of metallic mercury 
 and 15.88 grams of oxygen combine, 30,660 calories of 
 heat are evolved. Now, we have seen that when heat is 
 applied to the compound it is decomposed into its 
 elements. To effect this decomposition, as much heat 
 is absorbed as was evolved in the formation of the com- 
 pound. 
 
 Chemical Energy and Chemical Work. Any substance 
 which has the power to unite with others can do chemical 
 work : it possesses chemical energy. Thus, all combustible 
 substances can do work. In uniting with oxygen heat 
 is evolved, and this can be transformed into motion. In 
 the case of the steam-engine, the cause of the motion is 
 the burning of the fuel, which is a chemical act. We 
 thus see that the source of the power of the steam-engine 
 is chemical energy. Substances, on the other hand, 
 which have no power to combine with others have no 
 power to do chemical work, or they have no chemical 
 energy. So far as power to combine with oxygen is con- 
 cerned, water is a substance of this kind, as is also car- 
 bon dioxide, the gas formed when carbon is burned in 
 oxygen. In order that they may do work by combining 
 with oxygen, they must first be decomposed, and their 
 constituents put together in some form in which they 
 have the power of combination. This decomposition 
 of carbon dioxide and water is taking place constantly 
 on the earth. All plant-life is dependent on it. The 
 products of the action, i.e., the different kinds of wood, 
 have chemical energy, they can do chemical work. 
 This power to do work has been acquired from the heat 
 
OXIDES. 39 
 
 of the sun, which is the main force used in decomposing 
 the carbon dioxide and water. We have thus a trans- 
 formation of the sun's heat into chemical energy, which 
 is stored up in the combustible woods. The quantity 
 of heat which is given off in burning wood is believed to 
 be exactly equal to the quantity of heat used up in its 
 formation* 
 
 Oxides. The compounds of oxygen with other elements 
 are called oxides. To distinguish between different ox- 
 ides, the name of the element with which the oxygen is 
 in combination is prefixed. Thus, the compound of zinc 
 and oxygen is called zinc oxide; that of calcium and 
 oxygen, calcium oxide; that of silver and oxygen, silver 
 oxide; etc. When an element forms more than one com- 
 pound with oxygen, suffixes are used to distinguish 
 between them. Thus in the case of copper there are 
 two oxides which have the composition represented by 
 the symbols Cu 2 O and CuO. The former is known as 
 cuprous oxide and the latter as cupric oxide. That oxide 
 which contains the smaller quantity of oxygen in com- 
 bination with a given quantity of the other element is 
 designated by the suffix ous ; that which contains the 
 larger proportion of oxygen is designated by the suffix 
 ic. In other cases the number of combining weights of 
 oxygen contained in the compound is indicated by the 
 name. Thus, manganese dioxide is MnO 2 ; sulphur tri- 
 oxide is SO 3 ; etc. 
 
CHAPTER III. 
 
 A STUDY OF THE ELEMENT HYDROGEN. 
 
 Historical. Hydrogen was discovered as a distinct 
 substance by Cavendish in 1766, although it had been 
 observed as an inflammable gas before that time. 
 
 Occurrence. It occurs to some extent in the free con- 
 dition, and issues from the earth in small quantity in 
 some localities. It is, for example, a constituent of 
 the gases which escape from the petroleum wells in 
 Pennsylvania. It has also been shown to occur in enor- 
 mous quantities in the atmosphere of the sun. On the 
 earth it occurs chiefly, however, in combination in water, 
 of which it forms 11.11 per cent. It occurs also in most 
 substances of animal and vegetable origin, such as the 
 various kinds of wood and fruits, and the tissues of all 
 animals. In these products of life it is contained in 
 combination with carbon and oxygen or with carbon, 
 oxygen, and nitrogen. 
 
 Preparation. The simplest way, theoretically, to pre- 
 pare hydrogen is by the decomposition of water by the 
 electric current. It has already been stated that when 
 an electric current is passed through water the two gases 
 oxygen and hydrogen are liberated. But this method 
 is less convenient and more expensive than other 
 methods which are available, and it is therefore used 
 only under special circumstances. It is particularly 
 well adapted to the preparation of small quantities of 
 pure hydrogen. 
 
 Some elements when brought in contact with water at 
 the ordinary temperature decompose it and set hydrogen 
 free. The two most easily obtained elements which act 
 in this way are sodium and potassium. If a small piece 
 of potassium is thrown upon water, a flame is observed 
 at once. If sodium is used, it is seen to form a small 
 
 (40) 
 
HYDROGEN- PREPARATION. 41 
 
 ball which moves about on the surface of the water with 
 a hissing sound, but under ordinary circumstances no 
 flame is observed. By applying a flame to the ball 
 something takes fire and burns. By filling a good-sized 
 test-tube with water and inverting it in a larger vessel 
 and bringing a small piece of sodium wrapped in a piece 
 of filter-paper below the mouth of the tube, the sodium 
 will rise to the top of the tube when released, and it 
 will then be seen that a gas is evolved which gradually 
 depresses the water in the tube. A similar experi- 
 ment with potassium gives a similar result. The gas 
 given off in each case is hydrogen. By evaporating off 
 the water left in the vessel there will be found in each 
 case a white substance of marked chemical properties. 
 That formed with the potassium is known as potassium 
 hydroxide, or caustic potash, and has the composition 
 represented by the symbol KOH ; that formed with the 
 sodium is known as sodium hydroxide, or caustic soda, 
 and is represented by the symbol NaOH. The reactions 
 which take place between potassium and sodium and 
 water are represented by the equations 
 
 E + H 2 O = KOH + H ; and 
 Na + H 2 = NaOH + H. 
 
 Half the hydrogen of the water which is decomposed is 
 replaced by the potassium or the sodium, as the case 
 may be. These reactions are partly described by say- 
 ing that the potassium or sodium is substituted for half 
 the hydrogen in the water, and the act is called sub- 
 stitution. This is a very common kind of chemical ac- 
 tion, and we shall constantly meet with it in the course of 
 our study. The reaction is one of double decomposition 
 or metathesis, two substances acting upon each other to 
 form two others. The cause of such a reaction is to be 
 sought for in the different degrees of attraction exerted 
 by the elements upon one another. In general terms, if 
 two compounds AB and CD are brought together, and 
 the element A has for C a stronger attraction than A for 
 B, and B has for D a stronger attraction than it has for 
 
42 INORGANIC CHEMISTRY. 
 
 A, then reaction will take place, to some extent at least, 
 according to the equation 
 
 AB + CD = AC+BD. 
 
 The action may be modified by a number of circum- 
 stances which will be treated of in due time. It is evi- 
 dent, however, now that a very important problem for 
 the chemist to solve is the determination of the attrac- 
 tion which the elements exert upon one another. It is 
 extremely difficult to make these determinations, but 
 something can be learned in regard to them by a study 
 of the changes in temperature which accompany them. 
 The attraction is not proportional to the heat evolved, for 
 reasons which will be pointed out later, but there is some 
 relation between them. 
 
 Some substances which decompose water slowly at the 
 ordinary temperature do so readily at a higher tempera- 
 ture. This is true, for example, of iron. At ordinary 
 temperatures it decomposes water, as is seen in the for- 
 mation of a coating upon it when left in contact with 
 water. At higher temperatures when the iron is red- 
 Lot it decomposes water very readily, and hydrogen may 
 be made in quantity by this means. In the laboratory 
 the iron may be heated in a gun-barrel or in a porcelain 
 tube. When steam is passed over it the decomposition 
 represented in this equation takes place : 
 
 3Fe + 4H 2 O = Fe 3 O 4 + 8H. 
 
 The iron combines with the oxygen and liberates the 
 hydrogen. 
 
 Carbon, in the form of charcoal or coal, may be used 
 in a similar way to effect the decomposition of water 
 and the liberation of hydrogen. At a high heat the re- 
 action takes place mainly as represented thus : 
 
 C + H 2 = CO + 2H. 
 
 A mixture of two gases, carbon monoxide and hydrogen, 
 is thus formed. This mixture is the essential part of the 
 gas which has of late years come into such extensive 
 
HYDROGEN PREPARATION. 43 
 
 use under the name " water-gas." This is formed by 
 passing steam over highly heated anthracite coal. 
 
 By far the most convenient method for making hydro- 
 gen consists in treating a metal with an acid. Among 
 the metals best adapted to the purpose are zinc and 
 iron, and indeed zinc is almost exclusively used. As 
 will be seen later, acids are substances that contain 
 hydrogen, and are characterized by the property that 
 they give up this hydrogen very easily and take up 
 other elements in the place of it. Among the common 
 acids found in every laboratory are hydrochloric acid, 
 sulphuric acid, and nitric acid. The chemistry of these 
 compounds will be treated of in due time ; but, as we 
 shall be obliged to use them before they are taken up 
 systematically, a few words in regard to them are desir- 
 able in this place. 
 
 Hydrochloric acid is a compound of hydrogen and 
 chlorine. It is a gas which dissolves easilv in water. 
 
 t/ 
 
 It is this solution which is used in the laboratory, and 
 which is manufactured in enormous quantities in connec- 
 tion with the manufacture of soda or sodium carbonate. 
 Its chemical symbol is HC1. In commerce it is not 
 uncommonly called " muriatic acid." 
 
 Sidphuric acid is a compound of sulphur, oxygen, and 
 hydrogen in the proportions represented by the formula 
 H 2 SO 4 . It is an oily liquid and is frequently called 
 " oil of vitriol." It is manufactured in very large quan- 
 tities, as it plays an important part in many of the most 
 important chemical industries. 
 
 Nitric acid is a compound containing nitrogen, oxygen, 
 and hydrogen in the proportions represented by the 
 formula HNO 3 . It is a colorless liquid, though, as we 
 get it, it is commonly colored straw-yellow. 
 
 When a metal, such as zinc, is brought in contact with 
 hydrochloric or sulphuric acid, an evolution of hydro- 
 gen takes place at once. The reactions are as repre- 
 sented in these equations : 
 
 Zn + 2HCl = ZnCl 2 +2H; 
 Zn + H 2 SO 4 = ZnSO 4 + 2H. 
 
44 INORGANIC CHEMISTRY. 
 
 Each combining weight of zinc liberates and replaces 
 two combining weights of hydrogen. 
 
 The action between iron and these two acids is of the 
 same character : 
 
 Fe + 2HCl = FeCl 2 +2H; 
 Fe + H 2 SO 4 = FeSO 4 + 2H. 
 
 The hydrogen obtained from acids by the action of 
 metals is not pure, but it can be purified by treatment 
 with appropriate substances. That obtained by the de- 
 composition of water by the electric current is pure. 
 
 Physical Properties. Hydrogen is a colorless, inodor- 
 ous, tasteless gas. That made by the action of acids on 
 zinc or iron has a somewhat disagreeable odor which is 
 due to the presence of other gases in small quantity. It 
 is not poisonous, and may therefore be inhaled with im- 
 punity. We could not, however, live in an atmosphere 
 of hydrogen, as we need oxygen. It is the lightest 
 known substance. Its specific gravity in terms of the 
 air standard is 0.06926. A litre under. 760 mm. pressure 
 and at weighs 0.089873 gram. Under Oxygen it was 
 stated that in chemistry hydrogen is commonly taken as 
 the standard of specific gravity, and that, hydrogen being 
 unity, the specific gravity of oxygen is 15'. 88. The gas is 
 only slightly soluble in water. 100 volumes of water 
 take up 1.93 volumes of hydrogen. The fact that hy- 
 drogen is lighter than the air is shown by opening a 
 vessel which contains it and turning the mouth of the 
 vessel upward. The gas escapes at once, and in a very 
 short time no evidence of its presence can be obtained. 
 Light vessels as, for example, soap-bubbles or collodion- 
 balloons filled with the gas rise in the air, and it is used 
 for the purpose of filling large balloons. 
 
 Hydrogen has been converted into the liquid form 
 under a pressure of 20 atmospheres at a temperature of 
 234. 5. Its boiling-point under ordinary atmospheric 
 pressure is 243.5. 
 
 Hydrogen passes readily through porous substances, 
 or it diffuses rapidly. This can easily be demonstrated 
 in the case of porous earthenware and paper. It 
 
HYDROGEN-CHEMICAL PROPERTIES. 45 
 
 also passes readily through, some metals, as iron and 
 platinum, when heated to redness. There is a direct re- 
 lation between the specific gravity of gases and the rate 
 at which they diffuse. The lower the specific gravity 
 the more rapid the diffusion. The law governing these 
 phenomena is : 
 
 The rate of diffusion of gases is approximately inversely 
 proportional to the square roots of their specific gravities. 
 
 The specific gravity of hydrogen being 1 and that of 
 oxygen nearly 16 (15.88), the rate of diffusion of oxy- 
 gen is approximately ^ that of hydrogen. If hydrogen 
 is on one side of a porous wall, and oxygen on the 
 other, the hydrogen will pass through the wall so 
 much more rapidly than the oxygen that there will be 
 an accumulation of hydrogen on one side of the wall, 
 and if the vessel were closed there would be in- 
 creased pressure on that side. The ready passage of 
 gases through porous walls is a matter of great impor- 
 tance in connection with the ventilation of dwellings. 
 Most of the materials used in building are porous and 
 permit the passage of gases through them in both direc- 
 tions, and change of air is secured in this way to some 
 extent. 
 
 Chemical Properties. Under ordinary circumstances, 
 hydrogen is not a particularly active element. It does 
 not unite with oxygen gas at ordinary temperatures, 
 but, like other combustible substances, it must be 
 heated up to the kindling temperature before it will 
 l)urn. If a lighted match is applied to it, it takes 
 fire at once. The flame is colorless or slightly blue. 
 Generally the flame is somewhat colored in consequence 
 of the presence of foreign substances ; but that it is 
 colorless when the gas is burned alone can be shown by 
 burning it as it issues from a platinum tube which is 
 itself not chemically acted upon by the heat. Although 
 the flame is not luminous it is intensely hot, as can be 
 .seen by inserting into it a coil of platinum wire, which 
 will at once become red-hot and emit light accordingly. 
 
 The burning of hydrogen in the air, like the burning 
 of other combustible substances in the air, consists in a 
 
46 INORGANIC CHEMISTRY. 
 
 union of the gas with oxygen. This has been shown to 
 be true by most elaborate experiments on the combus- 
 tion of hydrogen in oxygen and in the air. On the other 
 hand, substances which burn in the air are extinguished 
 when put in a vessel containing hydrogen. This is- 
 equivalent to saying that a substance which is uniting 
 with oxygen does not continue to unite with oxygen 
 when put in an atmosphere of hydrogen, and does not 
 combine with hydrogen. The fact is expressed by say- 
 ing that hydrogen does not support combustion. Thi& 
 can be shown by holding a vessel filled with hydrogen, 
 with the mouth downward, and inserting into it a lighted 
 taper supported on a wire. The gas takes fire at the 
 mouth of the vessel, but the taper is extinguished. 
 
 Ordinarily we say that hydrogen burns in oxygen,, 
 but, as the act consists in the union of the two gases, it 
 would seem probable that oxygen will burn in an at- 
 mosphere of hydrogen. This can be shown to be true 
 by a proper arrangement of apparatus. If we were 
 surrounded by an atmosphere of hydrogen we should 
 probably speak of oxygen as a combustible gas in the 
 same way that we now speak of hydrogen as a combus- 
 tible gas. 
 
 It can easily be shown that, when hydrogen is burned 
 either in oxygen or air, water is formed. The simplest 
 way to show this is by holding a glass plate or some 
 other incombustible object a short distance above a 
 flame of hydrogen. It will be seen that drops of water 
 are condensed upon it. 
 
 Hydrogen combines with many other elements besides 
 oxygen, and forms some of the most important and in- 
 teresting compounds, such as hydrochloric acid, HC1 ; 
 sulphuretted hydrogen or hydrogen sulphide, H 2 S ; 
 ammonia, NH 3 ; marsh gas, CH 4 ; and all the acids. On 
 account of its affinity for oxygen it is used very exten- 
 sively in the laboratory for the purpose of extracting 
 oxygen from compounds containing it. Thus, when 
 hydrogen is passed over heated copper oxide, CuO, it 
 combines- with the oxygen to form water, and the coppel 
 
COMPARISON OF OXYGEN AND HYDROGEN. 47 
 
 is left in the free or uncombined state. The reaction is 
 represented thus : 
 
 CuO + 2H=H 3 + Cu. 
 
 A similar reaction takes place when hydrogen is passed 
 over highly heated oxide of iron, Fe a O 3 : 
 
 Fe 2 O 3 + 6H = 2Fe + 3H 3 O. 
 
 The removal of oxygen from a compound is called 
 reduction. Reduction is therefore plainly the opposite 
 of oxidation. Any substance which has the power to 
 abstract oxygen is spoken of as a reducing agent, just as 
 any substance which has the power to add oxygen to 
 a substance, or to decompose it by the action of oxygen, 
 is called an oxidizing agent. 
 
 A number of metals have the power to absorb a 
 large quantity of hydrogen when they are heated to 
 red heat in the gas. This phenomenon is shown most 
 strikingly by palladium, which under the most favor- 
 able conditions takes up something more than 935 times 
 its own volume of hydrogen. The gas is given up at 
 elevated temperature in a vacuum. When it absorbs 
 hydrogen, palladium undergoes marked changes in 
 properties. Its volume is increased, and its magnetic 
 and electric properties are also changed. It was sug- 
 gested by Graham, who first observed this phenomenon, 
 that the hydrogen held in combination by the palladium 
 is something quite different from ordinary hydrogen, 
 and that it must have some properties like those of the 
 so-called metals. He therefore called the combined 
 hydrogen hydrogenium. 
 
 Comparison of Oxygen and Hydrogen. Hydrogen and 
 oxygen are different kinds of matter, just as heat and 
 electricity are different kinds of energy. Heat can be 
 converted into electrical energy, and electrical energy 
 into heat, but one element cannot by any means known 
 to us be converted into another. They are apparently 
 entirely independent of each other. The question will 
 therefore suggest itself, whether, in spite of their ap- 
 
48 INORGANIC CHEMISTRY. 
 
 parent independence, there is not some relation be- 
 tween the different elements which reveals itself by 
 similarity in properties ? It will be found that the ele- 
 ments can be separated into groups or families accord- 
 ing to their properties. There are some elements, for 
 example, which in their chemical conduct resemble oxy- 
 gen markedly. These elements constitute the oxygen 
 family. So far as hydrogen is concerned, however, it 
 stands by itself. There is no other element which con- 
 ducts itself like it. If we compare it with oxygen, we 
 find very few facts which indicate any analogy between 
 the two elements. In their physical properties they 
 are, to be sure, similar. Both are colorless, inodorous, 
 tasteless gases. On the other hand, oxygen combines 
 readily with a large number of substances with which 
 hydrogen does not combine. Oxygen, as we have seen, 
 combines easily with carbon, sulphur, phosphorus, and 
 iron. It is a difficult matter to get any of these elements 
 to combine directly with hydrogen. Further than this, 
 substances which combine readily with hydrogen do not 
 combine readily with oxygen. The two elements ex- 
 hibit opposite chemical properties. What one can do 
 the other cannot do. This oppositeness of properties is 
 favorable to combination ; for not only do hydrogen and 
 oxygen combine with great ease under proper condi- 
 tions, but, as we shall see later, it is a general rule that 
 elements of like properties do not readily combine with 
 one another, while elements of unlike properties do 
 readily combine. 
 
CHAPTER IV. 
 
 STUDY OF THE ACTION OF HYDROGEN ON OXYGEN. 
 
 Burning of Hydrogen. Attention has already been 
 called to the fact that, when hydrogen burns, water is 
 formed. It is now necessary that this reaction should 
 be studied more thoroughly with the view of discover- 
 ing, as far as possible, exactly what takes place. One 
 of the first questions to be answered is what relation 
 exists between the weights of the hydrogen burned, the 
 oxygen used up in the burning process, and the water 
 formed ? But to weigh gases accurately and to collect 
 small quantities of water and weigh it are by no means 
 simple operations, and a great deal of work has been 
 done upon the problem under consideration. No good 
 method has been devised for the quantitative study of 
 the combination of hydrogen and oxygen by ordinary 
 combustion. On the other hand, very accurate experi- 
 ments on the subject have been made in three other 
 ways, a brief account of which will now be given. 
 
 Method of Dumas. The first accurate experiments 
 on the combustion of hydrogen are those of Dumas. 
 The method employed by this chemist was as follows : 
 He passed carefully purified hydrogen over heated 
 copper oxide and collected the water formed. The 
 reaction involved is that represented by the equation 
 
 CuO + 2H = H 2 O + Cu. 
 
 The weight of the oxygen that entered into combi- 
 nation with hydrogen was obtained by weighing the 
 vessel containing the copper oxide before and after the 
 experiment. The loss in weight represented the weight 
 of the oxygen which had been abstracted. The water 
 was collected by passing the gases formed through an 
 empty glass vessel in which most of the water was con- 
 densed, and then through tubes containing molten caus- 
 tic potash and phosphorus pentoxide, substances which 
 
 (49) 
 
50 INORGANIC CHEMISTRY. 
 
 have a marked power to absorb water and hold it in 
 combination. In nineteen experiments he obtained a 
 total amount of water weighing 945.439 grams, and the 
 total amount of oxygen used in the formation of this 
 amount of water was 840.161 grams. According to these 
 results the ratio between hydrogen and oxygen in water 
 expressed in percentages is : 
 
 Oxygen 88.864 
 
 Hydrogen 11.136 
 
 100.000 
 
 Morley with great pains proceeded as follows: He 
 weighed comparatively large volumes of oxygen in the 
 form of gas ; he then weighed hydrogen after it had been 
 absorbed by palladium. The hydrogen being expelled 
 from the palladium, it was brought together with the 
 oxygen, with which it was caused to combine by means 
 of electric sparks. The water thus formed was carefully 
 collected and weighed. Twelve experiments gave results 
 that agreed closely with one another. The mean of these 
 shows that hydrogen and oxygen combine in the ratio of 
 1 part of hydrogen to 15.8792 parts of oxygen. Dumas' 
 result expressed in the same way is 1 : 15.961. 
 
 Eudiometric Method. When hydrogen and oxygen are 
 mixed at ordinary temperatures no chemical change 
 takes place, and the two gases may be left in contact 
 with each other indefinitely without chemical action. 
 If, however, a spark is brought in contact with the 
 mixture violent action takes place, accompanied by a 
 flame and explosion. The action consists in the sudden 
 combination of the two gases to form water. It may be 
 illustrated in a number of ways : most simply by filling 
 soap-bubbles with the mixture of the two gases and ap- 
 plying a flame to them. The explosion which ensues 
 is harmless. Plainly, to study the combination of 
 hydrogen and oxygen by exploding a mixture of the 
 gases will require special precautions. It can be carried 
 out by the aid of the eudiometer (from evdia, good air, 
 and jteTpor, a measure, an instrument for determin- 
 
HYDROGEN AND OXYGEN EUDIOMETRIC METHOD. 51 
 
 ing the purity of air). The eudiometer is simply a 
 tube graduated in millimeters and having two small 
 platinum wires passed through it at the closed end, 
 nearly meeting inside and ending in loops outside, as 
 shown in Fig. 1. The eudiometer is filled with mercury, 
 inverted in a mercury trough, and held in an upright 
 position by means of proper clamps. For the purpose 
 of the experiment a quantity of hydrogen is passed up 
 into the tube, and its volume accurately measured. 
 About half this volume of oxygen is then introduced 
 and the volume again accurately determined, and 
 after the mixture has been allowed to stand for a 
 few minutes a spark is passed between the wires in 
 the eudiometer by connecting the loops with the poles 
 
 
 FIG. 1. 
 
 of a small induction coil or with a Leyden jar. Under 
 these circumstances the explosion takes place noise- 
 lessly and with little or no danger. If the interior of 
 the tube was dry before the explosion, it will be seen to 
 be moist afterwards, and a marked decrease in the vol- 
 ume of the gases is also observed. That water is the 
 product of the action has been proved beyond any possi- 
 bility of a doubt, over and over again. As the liquid 
 water which is formed occupies an almost inappreciable 
 volume as compared with the volume of the gases which 
 combine, the decrease in volume represents the total 
 volume of hydrogen and oxygen which have combined. 
 Now, if the experiment is performed with the two gases 
 in different proportions, it will be found that only when 
 they are mixed in the proportion of 2- volumes of hy- 
 drogen to 1 volume of oxygen do they completely dis- 
 appear in the explosion. If there is a larger proportion 
 of hydrogen present, the excess is left over; and the 
 same is true of the oxygen. It will thus be seen that 
 when hydrogen and oxygen combine to form water, they 
 do so in the proportion of 2 volumes of hydrogen to 
 1 volume of oxygen or more accurately 2.0008 to 1. 
 
52 INORGANIC CHEMISTRY. 
 
 Calculation of the Results Obtained in Exploding Mix- 
 tures of Hydrogen and Oxygen. Having determined that 
 whenever hydrogen and oxygen combine, they do so in 
 the proportion 1 volume oxygen to 2 volumes hydrogen, 
 and that when they combine the volume of liquid water 
 formed measures so little as to amount to nothing in the 
 measurements, we know that whenever a mixture of hy- 
 drogen and oxygen is exploded, no matter in what propor- 
 tions they may be present, the volume of gas which disap- 
 pears as such consisted of 2 volumes of hydrogen and 1 
 volume of oxygen, or, in other words, one-third of the 
 volume which disappears was oxygen and two-thirds 
 hydrogen. Take this example : A quantity of hydrogen 
 corresponding to 60 cc. under standard conditions is in- 
 troduced into a eudiometer ; 40 cc. of oxygen are added* 
 "What contraction will there be on exploding the mixture ? 
 Plainly the 60 cc. of hydrogen will combine with 30 cc. 
 of oxygen. The 90 cc. of gas will disappear, and 10 cc 
 of oxygen will remain uncombined. From a total vol- 
 ume of 100 cc., therefore, we get a contraction to 10 cc. 
 One-third of the contraction represents the oxygen and 
 two-thirds the hydrogen. 
 
 Determination of the Volume of Water Vapor formed 
 by Union of Definite Volumes of Hydrogen and Oxygen. 
 The experiments which have been described enable 
 us to draw the conclusion that hydrogen and oxygen 
 combine in certain proportions by volume and by weight, 
 and that a definite weight of water is formed ; further, 
 that the volume of liquid water formed when the two 
 gases combine is inappreciable as compared with that of 
 the gases. The question remains to be answered, what 
 relation exists between the volumes of the combining 
 gases and that of the water in the form of vapor ? This 
 can be determined by causing the gases to combine in a 
 eudiometer at a temperature sufficiently high to keep 
 the water in the form of vapor. The simplest arrange- 
 ment for accomplishing this is that shown in Fig. 2. A 
 long eudiometer, the upper half of which is divided into 
 three equal divisions marked on the outside, is filled 
 with mercury and inverted in a bath of mercury. It is 
 
HYDROGEN AND OXYGEN WATER VAPOR. 
 
 53 
 
 then surrounded by a large tube or jacket arranged as 
 shown in Fig. 2. This is connected through the cork 
 at the lower end with a vessel from which a current of 
 steam can be obtained. The steam is passed through 
 the jacket until the temperature of the mercury has 
 reached that of the steam. A mixture of hydrogen and 
 oxygen in the proportions in which they combine, viz., 
 2 of hydrogen to 1 of oxygen, is then introduced into 
 the eudiometer so that it is filled to the third mark. 
 This must be somewhat above the level of the mercury 
 in the bath, so that the gases in the eudiometer shall be 
 under diminished pressure. On now passing the spark 
 the gases unite and the water which 
 is formed remains in the form of 
 vapor, as the temperature inside 
 the eudiometer is nearly that of 
 boiling water, and the vapor is un- 
 der diminished pressure. It is 
 found that the volume of the water 
 vapor is less than that of the gases 
 introduced into the eudiometer. 
 In order to secure the same press- 
 ure as that under which the gases 
 were measured, the eudiometer 
 must be lowered until the height 
 of the mercury column in it is the 
 same as it was before the explosion. 
 On now measuring the volume of 
 water vapor, it will be found to be 
 two-thirds that occupied by the 
 uncombined gases. Therefore, 2 
 volumes of hydrogen combine with 
 1 volume of oxygen and form 2 
 volumes of water vapor. It is an 
 
 interesting fact that these simple relations exist between 
 the volumes of the combining gases and the volume of 
 the product. We shall see that similar relations hold 
 good in the case of other gases ; and the following gen- 
 eral statement is based upon a great deal of careful 
 study : 
 
 FIG. 2. 
 
54 INORGANIC CHEMISTRY. 
 
 When two or more gaseous substances combine to form a 
 gaseous compound, the volumes of the individual constituents 
 as well as their sum bear a simple relation to the volume of 
 the compound. 
 
 This is known as the law of combination by volume. 
 As will be seen farther on, it has a most important bear- 
 ing upon some of the fundamental ideas held in regard 
 to the constitution of matter. 
 
 Heat Evolved in the Union of Hydrogen and Oxygen. To 
 get as complete a knowledge as possible of the reaction 
 which takes place between hydrogen and oxygen we have 
 still to determine the amount of heat evolved. The heat 
 evolved in burning a gram of hydrogen can be deter- 
 mined, and from this we can calculate the heat of forma- 
 tion of water which, according to what was said on page 
 37, is the amount of heat evolved by the combination 
 of 2 grams of hydrogen with 15.88 grams of oxygen. We 
 are to determine the value of x in the equation 
 
 [H 2 , O] = x cal, 
 
 which expresses in thermochemical language the fact that 
 when 2 grams of hydrogen combine with 15.88 grams of 
 oxygen x calories of heat are evolved. 
 
 The determination is made by burning a known weight 
 of hydrogen in a vessel surrounded by water and arranged 
 in such a way that all the heat is absorbed by the water. 
 Experiment has shown that when 1 gram of hydrogen is 
 burned 34,180 calories are evolved. Or, 
 
 [H 2 , O] = 68,360 cal. 
 
 No other substance gives as much heat as this in propor- 
 tion to the weight used. 
 
 Applications of the Heat formed by the Combination of 
 Hydrogen and Oxygen. To burn hydrogen in the air is, 
 as we have seen, a simple matter, but to burn it in oxy- 
 gen requires a special apparatus to prevent the mixing 
 of the gases before they reach the end of the tube where 
 the combustion takes place. The oxyhydrogen bloiv-pipe 
 answers this purpose. This may be constructed in sev- 
 
OXYHYDROGEN BLO W-PIPE OXYHYDROGEN LIGHT. 55 
 
 eral ways, but the simplest is that represented in Fig. 3. 
 It consists of a tube through which a smaller tube passes. 
 The hydrogen is admitted through a and the oxygen 
 through b. It will be seen that they come together only 
 at the end of the tube. The hydrogen is first passed 
 through and lighted ; then the oxygen is passed through 
 slowly, the pressure being increased until the flame ap- 
 pears thin and straight. It gives very little light but is 
 intensely hot. Iron wire, steel, copper, zinc, and other 
 
 a FIG. 3. 
 
 metals burn in the flame with ease. Platinum vessels 
 are made by melting the platinum by means of the oxy- 
 hydrogen flame. 
 
 Oxyhydrogen Light. When the oxyhydrogen flame is 
 allowed to play upon some substance which it cannot 
 melt or burn, the substance becomes heated so high that 
 it gives off an intense light. The substance commonly 
 used for this purpose is quicklime. Hence the light is 
 often called the lime-light. It is also known as the 
 Drummond light. 
 
 The hydrogen is first allowed to pass through the stop- 
 cock, and lighted, when the oxygen is admitted. The 
 flame plays against the piece of lime, and from this the 
 light is given off when it has acquired a high tempera- 
 ture. Coal-gas may be used instead of hydrogen, and it 
 is generally used. Cylinders of compressed coal-gas and 
 of oxygen can be bought, and the gases so prepared are 
 used for the purpose of projecting pictures upon screens 
 in illustrating lectures and for other similar purposes. 
 
 Velocity of Combination of a Mixture of Hydrogen and 
 Oxygen. When a mixture of hydrogen and oxygen ex- 
 plodes, the action appears to take place instantaneously 
 throughout the mass. Whether this is really so or not 
 can be determined only by experiment. The action cer- 
 
56 INORGANIC CHEMISTRY. 
 
 tainly takes place with great rapidity, and a special 
 apparatus is necessary in order to study the rate of 
 transmission. This subject has been studied by ex- 
 ploding the mixture in long tubes arranged with little 
 movable pistons at various distances. When the explo- 
 sion reached these points the fact was indicated by 
 motion of the pistons. The result showed that the 
 action is not instantaneous though extremely rapid. The 
 rate of transmission is about 2500 metres per second. 
 
 Summary. In our study of the action of hydrogen on 
 oxygen, we have learned : (1) the relations between the 
 weights of the two gases which act upon each other ; (2) 
 the relations between the volumes of the combining 
 gases ; (3) the relations between the volumes of the com- 
 bining gases and that of the water vapor formed ; (4) the 
 amount of heat evolved when a given weight of hydrogen 
 combines with oxygen ; and (5) that the act of combina- 
 tion of the two gases does not take place instantaneously 
 though with great rapidity. It remains for us to study 
 more carefully the product formed. This is water. 
 
CHAPTER V. 
 
 WATER. 
 
 Historical. Water was long considered an elementary 
 substance until, towards the end of the last century, the 
 discovery of hydrogen and oxygen, and of the nature of 
 combustion, led to the discovery of its composition. 
 
 Occurrence. Besides the form in which water occurs 
 in such enormous quantities in the earth, it also occurs 
 in forms and conditions which prevent its immediate 
 recognition. Thus all living things contain a large pro- 
 portion of water, which can be driven off by heat. If a 
 piece of wood or of meat is heated, liquids pass off, and 
 by purification these can be shown to consist mainly of 
 water. The proportion of water in animal and vegetable 
 substances is very great. If the body of a man weighing 
 150 pounds were to be put in an oven and thoroughly 
 dried, there would be only about 40 pounds of solid mat- 
 ter left, most of the rest being water. 
 
 Water also occurs in another form in which it does not 
 easily reveal its presence. This is as water of crystalliza- 
 tion. Many chemical compounds found in nature and 
 manufactured are found to give off water when heated. 
 If, for example, the zinc sulphate formed in the prep- 
 aration of hydrogen from zinc and sulphuric acid is 
 dried by exposure to the air or by pressing between lay- 
 ers of filter-paper, it will be found that, when heated in 
 a dry tube, it gives off water, and at the same time 
 changes its appearance. The same is true of gypsum 
 which is found in nature, and of copper sulphate or blue 
 vitriol. In this last case the loss of water is accompanied 
 by a loss of color. After all the water is driven off, the 
 powder left behind is white. 
 
 Many compounds when deposited from solutions in 
 water in the form of crystals combine with definite quan- 
 
 (57) 
 
58 INORGANIC CHEMISTRY. 
 
 titles of water. This water is not present as such, but is 
 held in chemical combination. Hence the substance does 
 not appear moist, though it may contain more than half 
 its weight of water. This water of crystallization is, in 
 some way which is not understood, essential to the form 
 of the crystals. If it is driven off, the crystals generally 
 crumble to pieces. Some compounds combine with dif- 
 ferent quantities of water under different circumstances, 
 the form of the crystals varying with the quantity of 
 water held in combination. 
 
 Compounds differ greatly as regards the ease with 
 which they give up water of crystallization. In general, 
 it is given off when the compound containing it is heated 
 up to the temperature of boiling water. But some com- 
 pounds give it up by simple contact with the air. This 
 is true of sodium sulphate or Glauber's salt, which con- 
 tains a quantity of water of crystallization represented by 
 the formula Na 2 SO 4 . 10H 2 O. If some of the crystals are 
 allowed to lie exposed to the air they undergo a marked 
 change in the course of an hour or two. They lose 
 their lustre and gradually crumble to pieces. Sub- 
 stances which lose their water of crystallization by sim- 
 ple contact with the air are said to be efflorescent. 
 
 Some compounds if deprived of their water of crystal- 
 lization will take it up again when allowed to lie in an 
 atmosphere containing moisture. As the air always con- 
 tains moisture, it is only necessary to expose such com- 
 pounds to the air in order to notice the phenomenon. 
 It is well shown by the compound calcium chloride, 
 CaCl 2 . This substance has a remarkable power of at- 
 tracting water and holding it in combination. If a few 
 pieces are exposed to the air it will be noticed that they 
 soon have a moist appearance, and if they are allowed to 
 lie long enough they will dissolve in the water which is 
 absorbed from the air. Substances which absorb water 
 from the air are said to be deliquescent. 
 
 Formation of Water and Proofs of its Composition. If 
 we had not already learned, in studying the action of hy- 
 drogen upon oxygen, that water is composed of these 
 two gases, we should first subject it to analysis. For 
 
FORMATION OF WATER. 59 
 
 this purpose we should have to bring it under the influ- 
 ence of a number of reagents and study its conduct. If 
 we should pass an electric current through it in the proper 
 way, we should observe that a gas rises from each pole. 
 By placing each pole under the mouth of an inverted 
 tube filled with water the gases are easily collected. 
 When one of the tubes has become full of gas the other 
 one will be only half full. An examination will show 
 that the gas in the tube which is filled is hydrogen, 
 whereas that in the one which is half filled is oxygen. 
 By decomposition of water by means of the electric cur- 
 rent, then, there are obtained two volumes of hydrogen 
 for each volume of oxygen. We already know the rela- 
 tive weights of equal volumes of the two gases, so that 
 we can easily calculate the relative weights of the gases 
 obtained in the experiment. The ratio of the weights 
 of equal volumes of hydrogen and oxygen is 1 : 15.88. 
 Therefore, if we have 2 volumes of hydrogen combined 
 with 1 volume of oxygen, the ratio between the weights 
 is 2 : 15.88 or 1 : 7.94 Although we know from the ex- 
 periment referred to that hydrogen and oxygen are ob- 
 tained from water in certain proportions, it does not fol- 
 low that this is the composition of water. For it may be 
 that other elements besides hydrogen and oxygen are 
 contained in it, and it may be also that all the hydrogen 
 and oxygen are not set free by the action of the electric 
 current. We might determine whether either of these 
 possibilities is true or not by decomposing a weighed 
 quantity of water, and weighing the hydrogen and oxygen 
 obtained from it. If we should find that the sum of the 
 weights of hydrogen and oxygen is equal to the weight 
 of the water decomposed, this fact would be evidence 
 that only hydrogen and oxygen are contained in water, 
 and that they are present in the proportions stated. 
 The same thing can be satisfactorily proved by causing 
 hydrogen and oxygen to combine, or by effecting the syn- 
 thesis of water. How this may be done has already been 
 pointed out. It was shown, in the first place, that by 
 burning hydrogen in oxygen water is formed. This 
 proves that water consists of hydrogen and oxygen, but 
 
60 INORGANIC CHEMISTRY. 
 
 it does not furnish any proof as to the relation between 
 the quantities of the gases which combine. It is a quali- 
 tative synthesis. Other methods were described, the ob- 
 ject of which was to show in what proportion by weight 
 and by volume hydrogen and oxygen combine to form 
 water. These methods are examples of quantitative syn- 
 theses. The results proved that to form water hydrogen 
 and oxygen combine in the proportion of 1 volume of 
 oxygen to 2 of hydrogen ; and it therefore follows that 
 the decomposition of water which is effected by the elec- 
 tric current is complete. 
 
 Properties of Water. Pure water is tasteless and in- 
 odorous, and in small quantities colorless. Thick 
 layers are, however, blue. This is seen by filling a 
 long tube with carefully purified water, and examining it 
 by transmitted light, when it appears blue. Some 
 mountain lakes also have a marked blue color. When 
 cooled, water contracts until it reaches the temperature 
 of 4. At this point it has its maximum density. If 
 cooled below this it expands, and the specific gravity of 
 ice is somewhat less than that of water. Hence ice 
 floats on water. If this were not so there would be 
 great danger in cold climates that the water in the 
 streams would freeze solid. As it is, the lower layers 
 of water are protected by the ice and the cold water 
 just below it, which are poor conductors of heat. Water 
 can be cooled down below its freezing temperature, or 
 0, if it is kept perfectly quiet, protected from the air, or 
 cooled in capillary tubes. Water thus cooled down will 
 suddenly solidify when disturbed, and then its tempera- 
 ture rises to 0. Water boils at 100 under 760 mm. pres- 
 sure. Increased pressure raises the boiling point, and 
 decreased pressure lowers it. 
 
 Chemical Properties of Water. Water is a very stable 
 chemical compound. An indication of the attraction ex- 
 erted by the hydrogen for the oxygen is given in the 
 great evolution of heat when the two combine. In order 
 to decompose it by heat, as much heat must be added as 
 is evolved when it is formed. At high temperatures it 
 is decomposed into its elements. The decomposition 
 
CHEMICAL PROPERTIES OF WATER. 61 
 
 begins at 1000 and is half complete at 2500. This 
 kind of gradual decomposition of a compound by heat 
 is called dissociation. It is a common phenomenon in 
 chemistry, and farther on we shall have occasion to study 
 it more fully. As has been seen, water is decomposed 
 completely by an electric current, and partly by contact 
 with sodium and potassium at ordinary temperatures, 
 and by iron and carbon at higher temperatures. It 
 combines directly with a large number of substances in 
 the form of water of crystallization, and with others to 
 form definite chemical compounds called hydrates or 
 hydroxides. Thus the oxides of potassium, K 2 O, and of 
 sodium, Na 2 O, combine with water with evolution of 
 much heat to form the compounds potassium hydroxide, 
 KOH, and sodium hydroxide, NaOH, which, it will be re- 
 membered, are also formed by the action of the elements 
 potassium and sodium on water with liberation of hydro- 
 gen. The reactions between the oxides and water are 
 represented by the equations 
 
 K 2 O + H 2 0=2KOH; 
 Na 2 O + H 2 = 2NaOH. 
 
 Similarly, lime or calcium oxide, CaO, acts upon water 
 with evolution of heat, as is observed in the process of 
 slaking. The change is like that which takes place with 
 potassium and sodium oxides ; and is represented thus : 
 
 CaO + H 2 = CaO 2 H 2 [or Ca(OH)J. 
 
 The product represented by the symbol Ca(OH) 2 is 
 known as calcium hydroxide or slaked lime. In the 
 same way barium oxide, BaO, forms barium hydroxide : 
 
 BaO + H 2 O = BaOJI 2 [or Ba(OH)J. 
 
 We shall meet with many examples of this kind of 
 action in our study of chemical reactions, and we shall 
 see that the hydroxides form two of the most important 
 classes of compounds, known as acids and bases. The 
 hydroxides of potassium, sodium, calcium, and barium 
 are, for example, bases ; while certain hydroxides con- 
 
62 INORGANIC CHEMISTRY. 
 
 taining sulphur, nitrogen, and carbon are acids, such as 
 sulphuric acid SO 2 (OH) 2 , nitric acid NO 2 (OH), and car- 
 bonic acid, CO(OH) 2 . It is not believed that water as 
 such is contained in these hydroxides. Nevertheless, 
 when heated many of them give off water. Thus, when 
 heated to a red heat, calcium hydroxide is decomposed 
 into the oxide and water according to the equation 
 
 Many substances which contain hydrogen and oxygen 
 act in the same way. This is due to the great sta- 
 bility of water even at elevated temperatures. As the 
 temperature becomes higher and higher the attraction 
 between the constituents of the compound becomes 
 weaker and weaker. When a point is reached at which 
 the attraction of the hydrogen for the oxygen is greater 
 than that required to hold the constituents together, a 
 rearrangement takes place, and compounds which are 
 stable at the higher temperature are formed. 
 
 Water as a Solvent. With a great many substances 
 water forms unstable compounds, the nature of which 
 cannot at present be explained. These unstable com- 
 pounds are called solutions. It is known that many solids, 
 liquids, and gases when brought into water disappear 
 and form colorless liquids which look like water. Some 
 give colored liquids of the same color as the substance 
 dissolved, and others give liquids which have a color 
 quite different from the substance dissolved. On the 
 other hand, there are many substances which do not 
 form such compounds with water or which, as we say, 
 are insoluble in water. In a solution the particles of 
 the substance dissolved are in some way attracted and 
 held in combination by the particles of the liquid. If a 
 very small quantity of substance be dissolved in a large 
 quantity of water and the solution thoroughly stirred, 
 the dissolved substance is uniformly distributed through- 
 out the liquid, as can be shown by refined chemical meth- 
 ods. That the dissolved substance is everywhere present 
 in the solution can be shown further by the aid of certain 
 dye-stuffs, as, for example, fuchsine. A drop of a concen- 
 
WATER AS A SOLVENT. 63 
 
 trated solution of the substance brought into many gallons 
 of water imparts a distinct color to all parts of the liquid. 
 An experiment of this kind gives some idea of the 
 extent to which the division of matter can be carried. 
 For it is evident that in each drop of the dilute solution 
 there must be contained some of the coloring matter, 
 though the quantity must be what we should ordinarily 
 speak of as infinitesimal. While there seems to be no 
 limit to the extent to which a solution can be diluted and 
 still retain the dissolved substance uniformly distributed 
 through its mass, there is a limit to the amount of every 
 substance that can be brought into solution, and this 
 varies with the temperature, and, in the case of gases, 
 with the pressure. Some substances are easily soluble, 
 others are difficultly soluble. When the solutions are 
 boiled the water simply passes off and leaves the dis- 
 solved substance behind, if it is a non-volatile solid. 
 If, however, the substance in solution is a liquid, a par- 
 tial separation will take place, the extent of the separa- 
 tion depending largely upon the difference between the 
 boiling-points of the water and the other liquid. A com- 
 plete separation of two liquids by boiling is difficult and 
 in most cases impossible. If, finally, the substance in 
 solution is a gas, it generally passes off when the solution 
 is heated, though in some cases water is given off leaving 
 the gas in solution, which of course then becomes more 
 concentrated. When a certain concentration is reached a 
 solution of the gas passes over. It is probable that in these 
 cases the gas is in a condition of chemical combination 
 with the water. Solutions, in general, seem to differ from 
 true chemical compounds in some important particulars, 
 and also from mere mechanical mixtures. Definiteness 
 of composition appears to be characteristic of chemical 
 compounds, or, at least, it is characteristic of a large num- 
 ber of compounds which we call chemical compounds. 
 But solutions have no definite composition. We can dis- 
 solve any quantity of a substance from the minutest par- 
 ticle to a certain fixed quantity, and the solutions formed 
 are uniform and appear to be just as truly solutions as 
 that which contains the largest quantity which can be 
 
64 INORGANIC CHEMISTRY. 
 
 held in combination. On the other hand, in a mere me- 
 chanical mixture the constituents may be present in all 
 proportions, while this is not true of solutions. The sub- 
 ject of solution is under investigation, and it has been 
 made highly probable that some substances, when dis- 
 solved, are partly or wholly broken down into smaller 
 parts charged with positive and negative electricity re- 
 spectively. These smaller parts are called ions. 
 
 Solution as an Aid to Chemical Action. When it is 
 desired to secure the chemical action of one solid sub- 
 stance upon another, it is generally necessary to bring 
 them together in solution. One reason why they do 
 not act readily when mixed in the solid condition is to 
 be found in the fact that, under these circumstances, 
 their particles remain separated by sensible distances, 
 no matter how finely the mixture may be powdered. 
 If, however, the substances are dissolved, and the 
 solutions poured together, the particles of the liquid 
 move so freely among one another that they come 
 in intimate contact, thus facilitating chemical action. 
 Many substances which do not act upon one another at 
 all when brought together in dry condition act readily 
 when brought together in solution. It is believed that 
 this is due principally to the splitting of the compounds 
 into ions by the action of the water. These ions being 
 free are capable of acting upon other ions which may 
 be brought into the same solution. This idea will be 
 developed farther on in other connections. Although 
 it is highly probable, hen, that when a reaction takes 
 place in a water solution the water itself plays a very 
 important part, the reaction is generally represented 
 by an equation in which the water does not appear. Of 
 course, such an equation is imperfect, but it answers cer- 
 tain purposes quite satisfactorily, and may be used with- 
 out danger of confusion. Thus, when hydrochloric acid 
 acts upon zinc, hydrogen is liberated and zinc chloride 
 is formed. What we call hydrochloric acid in the labo- 
 ratory is the liquid which is formed by the absorption of 
 hydrochloric acid gas, HC1, by water. When this liquid 
 is used, however, the chemical act which makes itself 
 
NATURAL WATEES. 65 
 
 known to us is that which gives hydrogen gas and zinc 
 chloride in solution. Apparently this reaction is inde- 
 pendent of the water, and it may be represented thus : 
 
 Zn + 2HC1 = ZnCl a + 2H. 
 
 Sometimes such reactions are written as follows in order 
 to express the fact that water is present, though, as will 
 be observed, no attempt is made to tell what part the 
 water plays : 
 
 Zn + 2HC1 + Aq = ZnCl 2 + 2H + Aq ; or 
 Zn + 2HC1 + H 2 O = ZnCl 3 + 2H + H 2 O. 
 
 There is no objection to this, certainly, but it is ques- 
 tionable whether, considering the purposes for which 
 chemical equations are used, this increases their value. 
 
 Natural Waters. All water found in nature is more or 
 less impure or, in other words, contains something in so- 
 lution. In the first place, waters which are exposed to 
 the air dissolve some of the gases of which the air is 
 composed, as oxygen, nitrogen, and carbon dioxide. 
 Again, natural waters necessarily are in contact with 
 the earth, they always dissolve some of the earthy 
 substances ; and, finally, many waters come in contact 
 with animal and vegetable substances and dissolve some- 
 thing from these. The water which is carried up as 
 vapor from the surfaces of natural bodies of water is 
 approximately pure. When this is precipitated as rain it 
 dissolves certain substances from the air, and the first rain 
 that falls during a storm is always more or less contami- 
 nated. In a short time, however, the air becomes washed 
 and the rain which falls thereafter is approximately pure 
 water. If it remains in contact with insoluble rocks, as, 
 for example, quartzite or sandstone, it remains pure, and 
 mountain-streams which flow over sandstone beds are, in 
 general, the purest. Water which flows over limestone 
 dissolves some of this and becomes " hard." A similar 
 change is brought about in water by contact with gyp- 
 sum and magnesium sulphate. The condition of hard- 
 ness will be taken up more fully under calcium and 
 magnesium compounds. The many varieties of mineral 
 
66 INORGANIC CHEMISTRY. 
 
 springs have their origin in the presence in the earth of 
 certain substances which are soluble in water. Among 
 those most frequently met with in solution in natural 
 waters are carbonic acid, sodium carbonate, sodium sul- 
 phate or Glauber's salt, sodium chloride or common salt, 
 magnesium sulphate, iron carbonate, and sulphuretted 
 hydrogen. Effervescent waters are those which contain 
 a large quantity of carbonic acid in solution and give 
 off carbon dioxide gas when exposed to the air. Chalyb- 
 eate waters are those which contain some compound of 
 iron in solution ; sulphur waters contain the gas, sulphu- 
 retted hydrogen. Common salt occurs in large quantities 
 in different parts of the earth. As it is easily soluble in 
 water, many streams contain it ; and as most streams find 
 their way to the ocean, we see one reason why the water 
 of the ocean should be salt. 
 
 As streams approach the habitation of man they are 
 subjected to a serious cause of contamination. The 
 drainage from the neighborhood of human dwellings is 
 very apt to find its way into a near stream. The sub- 
 stances thus carried into the stream undergo decompo- 
 sition and give rise to the formation of larger or smaller 
 quantities of new products some of which have the power 
 to produce disturbances when taken into the system, 
 and others to produce disease. This condition of things 
 is most strikingly illustrated by the case of a large town 
 situated on the banks of a river. It frequently happens 
 that the water of the river is used for drinking purposes, 
 and it also frequently happens that the water is contami- 
 nated by drainage. Biver water when once contaminated 
 by drainage tends to become pure again by contact with 
 the air, the change consisting largely in the slow oxida- 
 tion of the substances which are of animal or vegetable 
 origin, and their conversion into harmless products. If 
 water is to be used for drinking purposes, however, it is 
 not well to rely too much upon this process of purifica- 
 tion. So much has of late years been said about drinking- 
 water that excessive alarm has been created, and water is 
 no doubt frequently held responsible for sickness with 
 which it has nothing to do. In some places the war against 
 the water supply has been carried so far that those who 
 
PURIFICATION OF WATER. 67 
 
 can afford it drink only artificially purified and distilled 
 water. It is undoubtedly well to be cautious, but it 
 is possible to be too cautious. 
 
 What Constitutes a Bad Drinking Water. A good drink- 
 ing water should be free from odor and taste and should 
 not contain anything which can act injuriously upon the 
 system. It is, however, difficult to decide by chemical 
 means whether the water contains anything injurious or 
 not, as there may be a very minute quantity of an ex- 
 tremely injurious substance, for example a disease germ, 
 present, and chemical analysis would be powerless to de- 
 tect it. On the other hand, water which is very consid- 
 erably contaminated by sewage may be harmless, and 
 yet the latter might be pronounced " bad " and the 
 former " good." The rule generally adopted by chemists 
 in dealing with water is to pronounce any water danger- 
 ous which is contaminated by sewage. Such contamina- 
 tion can generally be detected by analysis or by analysis 
 and inspection of the sources. 
 
 Purification of Water. Impure water may become purer 
 by natural methods as has been stated, and it may be 
 rendered fit for drinking purposes by filtering through 
 such substances as charcoal, sand, spongy iron, etc. A 
 filter, no matter of what it may be made, will not, how- 
 ever, remain efficient for any length of time, as the sub- 
 stances contained in the impure water are retained by it 
 and, after a time, it becomes a source of pollution in- 
 stead of a purifier. For refined work in chemistry pure 
 water is prepared by distilling natural waters. The 
 process of distillation consists in boiling the water and 
 then passing the steam through a tube or system of tubes 
 surrounded by cold water. Thus the steam is con- 
 densed, and the distilled water is approximately pure. Of 
 course, it is necessary that the tubes in which the con- 
 densation takes place should be of such material that 
 water does not act upon it to any extent. The materials 
 used are glass, block tin, and platinum. Chemically pure 
 water is a very rare substance even in the best chemical 
 laboratories. The slight impurities present in ordinary 
 distilled water are not, however, of special importance 
 under ordinary circumstances. 
 
CHAPTER VI. 
 
 CONSTITUTION OF MATTER-ATOMIC THEORY-ATOMS 
 AND MOLECULES CONSTITUTION VALENCE. 
 
 Early Views. In early times two views were held re- 
 garding the ultimate constitution of matter. The first 
 was, that matter is infinitely divisible that there is no 
 limit to the process of subdivision ; the other was, that 
 there is a limit to the divisibility that when certain in- 
 conceivably small particles are reached the process must 
 stop. These small particles were called atoms, meaning 
 indivisible. As long as the constitution of matter was 
 merely a subject of speculation, the atoms remained with- 
 out a physical basis and were only metaphysical con- 
 ceptions. The facts which come under our ordinary ob- 
 servation do not furnish any evidence for or against the 
 existence of atoms ; and, though we discuss the subject 
 indefinitely, little or no progress can be made without re- 
 fined observations on the properties of matter. So it was, 
 that until the beginning of the present century the atomic 
 theory remained practically what it was when first pro- 
 posed, and as such it was of no value to chemistry. 
 
 The Atomic Theory as proposed by Dalton. We have 
 seen how Dalton, at the beginning of this century, dis- 
 covered the law of multiple proportions. This law, as 
 well as that of definite proportions, required explanation. 
 The questions to be answered are : (1) Why do the ele- 
 ments combine in definite proportions ? (2) Why, when 
 elements combine with each other in more than one way, 
 do the relative quantities which enter into combination 
 in the different cases bear simple relations to one another ? 
 and (3) What is the significance of the figures represent- 
 ing the combining weights ? Dalton saw that the facts re- 
 ferred to could be explained by the atomic theory, while, 
 on the theory that matter is infinitely divisible, they 
 
 (68) 
 
THE ATOMIC THEORY. 69 
 
 appear to be inexplicable. It is only necessary to assume 
 that each element is made up of particles which are not 
 divisible in chemical processes, and that these particles, 
 or atoms, have definite weights. The atoms of any one 
 element must be supposed to have the same weight, 
 while the atoms of different elements have different 
 weights. Now, when chemical combination takes place, 
 Dalton supposed that the action was between the atoms. 
 The simplest case is that in which combination takes place 
 in such way that each atom of one element combines 
 with one atom of another ; but, besides this kind of com- 
 bination, we may have that in which one atom of one ele- 
 ment combines with two atoms of another, or two of one 
 may combine with three of another, etc. Suppose two 
 elements A and B, the weights of whose atoms are to each 
 other as 1 : 10, are brought together, and they combine in 
 the simplest way, i.e., one atom of one with one atom of 
 the other, then it is plain that in the compound AB the 
 elements will be contained in the proportion of 1 part of 
 A to 10 parts of B, whether a small or a large quantity 
 of the compound is formed, and no matter in what pro- 
 portions the elements are brought together. If they 
 should be brought together in the proportion of their 
 atomic weights (1 : 10), then no part of either element will 
 be left uncombined after the act of combination has taken 
 place. If, however, a larger proportion of either element 
 is taken than that stated, then the quantity of the one 
 which is in excess of this proportion will be left uncom- 
 bined. This is in accordance with what we know takes 
 place, and it is a conclusion drawn from the theory. No 
 matter how many atoms of A we may take, the same 
 number of atoms of B will be required to combine with 
 all of them. But each atom of B weighs 10 times as 
 much as each atom of A, therefore the total mass of B 
 which enters into combination must be 10 times that of 
 A with which it combines. It may be, however, that 
 these same elements can form other compounds with each 
 other. If A and B represent the atoms of the elements 
 and we assume these atoms to be chemically indivisible, 
 then the other compounds must be represented by such 
 
70 INORGANIC CHEMISTRY. 
 
 symbols as AB A^B, AB^ A^B, A^B^ etc., which repre- 
 sent compounds in which 1 atom of A is combined with 
 
 2 atoms of B ; 2 of A with 1 of B ; 1 of A with 3 of B ; 
 
 3 of A with 1 of B ; 2 of A with 3 of B ; etc. : or they 
 also represent compounds in w r hich 1 part by weight of 
 A is combined with 20 parts by weight of B ; 2 parts of 
 A with 10 of B ; 1 of A with 30 of B ; 3 of A with 10 of 
 B ; 2 of A with 30 of B ; etc. It is therefore clear that, 
 if the atomic theory as put forward by Dalton is true, the 
 elements must combine according to the laws of definite 
 and multiple proportions ; and it appears that the figures 
 which represent the combining weights of the elements 
 must either bear to one another the same relation as the 
 weights of the atoms, or, at all events, the atomic weights 
 or the relative weights of the atoms must be closely re- 
 lated to the combining weights, as will be pointed out 
 more clearly presently. 
 
 Use and Value of a Theory. The relation of a theory 
 to facts is very simple, but is frequently misunderstood. 
 The relation may be conveniently illustrated by the case 
 under consideration. By a careful investigation of a 
 number of chemical compounds it was shown that in each 
 of them the same elements always occurred in the same 
 proportion. This led to the belief that this is true of 
 every chemical compound, and after further investigation 
 which, as far as it went, showed the surmise to be correct, 
 the law of definite proportions was proposed. This law 
 is simply a statement of what has been found true in all 
 cases examined. It involves no speculation. It is a 
 statement of fact. It may be said that the statement or 
 law must be open to some doubt for the reason that all 
 possible cases have not been examined, and it may not 
 hold true for some of these unexamined cases. The reply 
 to this is that it has been found true in a very large num- 
 ber of cases and in all cases which have been investigated. 
 It is true for the present state of our knowledge, and that 
 is all we can demand of any law. Again, further investi- 
 gation led to the discovery of the law of multiple propor- 
 tions, which is also a statement of what has been found 
 true in all cases investigated. It, like the law of definite 
 
ATOMIC WEIGHTS AND COMBINING WEIGHTS. 71 
 
 proportions and in the same sense, is a statement of fact. 
 But having gone thus far, we now ask, what is the ex- 
 planation of these laws? We simply know the facts 
 what is the explanation ? By experiment we cannot go 
 beyond these facts, but it is possible to imagine a cause 
 and then proceed to see whether the imagined cause is 
 sufficient to account for the facts. This is what Dalton 
 did. He imagined that matter is made up of atoms of 
 definite weights, and that chemical combination takes 
 place in simple ways between these atoms. This imag- 
 ined cause is the atomic theory. It is not a statement of 
 anything found by investigation. It is not a statement 
 of an established fact. It may or may not be literally 
 true, but at all events it is the best guess that has ever 
 been made as to the cause of the fundamental laws of 
 chemical action, and it furnishes a very convenient means 
 of interpreting the facts of chemistry. Since the atomic 
 theory was first proposed it has been accepted by nearly 
 all chemists. It has been of great value in suggesting 
 methods of work, and has contributed largely to the ad- 
 vance of chemistry. Any theory which is in accord- 
 ance with the facts and leads to the discovery of new 
 facts is of value, whether it should eventually prove 
 to be true or false. At the same time a false theory 
 may do much harm, as it may lead men to misinterpret 
 the facts which they observe, and thus retard progress. 
 
 Atomic Weights and Combining "Weights. If the atomic 
 theory is true the atoms of each element must have defi- 
 nite weights, and the determination of these atomic 
 weights must evidently be of great importance. By 
 analysis of compounds we can only determine the pro- 
 portions by weight in which the elements combine with 
 one another. Can we in this way determine the atomic 
 weights ? In the first place, it is clear that it is out of 
 the question to think of determining the absolute weights 
 of the atoms, and all that we can possibly do is to deter- 
 mine their relative weights. As of all the elements hy- 
 drogen enters into combination in smallest relative quan- 
 tity, its atomic weight is taken as the unit of the system, 
 and the problem before us is to determine how many 
 
72 INORGANIC CHEMISTRY. 
 
 times heavier the atoms of the other elements are than 
 that of hydrogen. If every element combined with hydro- 
 gen in only one proportion the problem would be a com- 
 paratively simple one. Thus the three elements chlorine, 
 bromine, and iodine combine with hydrogen, forming only 
 one compound each. On analysis these are found to 
 contain respectively 
 
 1 part of hydrogen to 35.18 parts of chlorine ; 
 
 1 " " " 79,34 " " bromine; and 
 
 1 " " " 125.89 " " iodine. 
 
 There is no reason for believing that in these compounds 
 the elements are combined in any but the simplest way, 
 i.e., that each atom of hydrogen is combined with one 
 atom of chlorine to form hydrochloric acid, etc. If this 
 is true, then the atom of chlorine must weigh 35.37 times ; 
 that of bromine 79.76 times; and that of iodine 126.54 
 times as much as that of hydrogen, or, in other words, 
 the atomic weights of chlorine, bromine, and iodine are 
 respectively 35.37, 79.76, and 126.54. It will, however, 
 be observed that there is no evidence as to whether the 
 elements in these compounds are combined in the simplest 
 way or not. It is possible, as far as we know, that one 
 atom of hydrogen may combine with two or three of 
 chlorine, or that one of chlorine may combine with two 
 or three of hydrogen. As there is, however, no evidence 
 upon this point the simplest assumption is made. 
 
 If we take the case of oxygen the problem is more com- 
 plex. In water the elements are combined in the propor- 
 tion of 7.94 parts of oxygen to 1 part of hydrogen, and 
 from this we should naturally conclude that the atomic 
 weight of oxygen is 7.94 ; but further study shows that 
 this conclusion is not justified. Hydrogen and oxygen 
 form a second compound known as hydrogen dioxide or 
 hydrogen peroxide in which there are 15.88 parts of oxy- 
 gen to 1 of hydrogen. This may be explained in the 
 terms of the atomic theory by assuming that water is 
 represented by the formula HO, and hydrogen dioxide 
 by HO 3 . But it may be that the atomic weight of oxygen 
 is 15.88, and then water must be represented by the 
 
MOLECULES- AVOGADRO'S LAW. 73 
 
 formula H 2 O, and hydrogen dioxide by HO. Simple 
 analysis of the compounds is not sufficient to enable us to 
 decide between these possibilities. It is, therefore, evi- 
 dent that in order to determine the atomic weights some- 
 thing besides the determination of the composition of 
 compounds is necessary. The figures representing the 
 combining weights found in this way will, however, either 
 be identical with the atomic weights or will bear a simple 
 numerical relation to them. 
 
 Molecules. Investigation of certain phenomena of 
 light, of electricity, of liquid films and the conduct of 
 gases has led physicists to the conclusion that matter is 
 not continuous, but made up of small particles, which are 
 called molecules. A gaseous molecule is defined as " that 
 minute portion of a substance which moves about as a whole, 
 so that its parts, if it has any, do not part company during 
 the motion of agitation of the gas." It would be out of 
 place here to present the physical facts upon which the 
 molecular theory rests. Suffice it to say that it is the 
 only theory which has been found adequate to account 
 for the behavior of gases. 
 
 Avogadro's Law. The fact that gases conduct them- 
 selves in the same way under the influence of changes in 
 temperature and pressure can only be explained by as- 
 suming that equal volumes of all gases and vapors contain 
 the same number of ultimate particles or molecules at the 
 mme temperature and pressure. 
 
 This is a deduction from the well-tested molecular 
 theory of gases. It was, however, originally put forward 
 by the Italian chemist Avogadro from a study of chemical 
 as well as of physical facts, and a little later it was sug- 
 gested as probable by the French physicist Ampere. It 
 is therefore generally spoken of as Avogadro's law, and 
 sometimes as Ampere's law. Absolute proof of its truth 
 cannot be given, but it is in thorough accordance with a 
 la*rge number of w T ell-known facts, and it is undoubtedly 
 true, if the molecular theory of matter is true. It may 
 therefore be considered as furnishing a solid foundation 
 for further conclusions bearing upon the problem of the 
 determination of the atomic weights. 
 
74 INORGANIC CHEMISTRY. 
 
 Distinction between Molecules and Atoms. If we con- 
 sider any chemical compound, as water or hydrochloric 
 acid, it is evident that the smallest particle or the mole- 
 cule of the compound must be made up of still smaller 
 particles. Thus, the smallest particle of water must con- 
 tain smaller particles of hydrogen and oxygen, and the 
 smallest particle of hydrochloric acid must contain smaller 
 particles of hydrogen and chlorine. These smallest par- 
 ticles of the molecules are the atoms. The molecules of 
 the compounds are, according to this view, made up of 
 the atoms of the elements. Similarly the elements them- 
 selves are, for good reasons which will be presented, be- 
 lieved to consist of molecules which are in turn made up 
 of -atoms of the same kind, though in a few cases the 
 molecule of the element is identical with the atom. The 
 difference between a compound and an element then is, 
 in general, that the molecule of the compound consists of 
 atoms of different kinds, while the molecule of an ele- 
 ment consists of atoms of the same kind or, in a few 
 oases, of one atom. Generally the atoms do not exist in 
 the free or uncombined state, but, if they are set free by 
 chemical action, they unite to form molecules. The fol- 
 lowing may serve as a definition of the conception of 
 atoms at present held by chemists : 
 
 Atoms are the indivisible constituents of molecules. They 
 are the smallest particles of the elements that take part in 
 chemical reactions, and are, for the greater part, incapable 
 of existence in the free state, being generally found in combi- 
 nation with other atoms, either of the same kind or of differ- 
 ent kinds. 
 
 It cannot be too strongly emphasized that the views 
 held in regard to the relations between molecules and 
 atoms are based upon an enormous amount of painstak- 
 ing study of facts, and in order fully to comprehend their 
 value a study of most of these facts would be necessary. 
 These views have gradually become firmly established as 
 knowledge of the facts has grown more and more pro- 
 found. Accepting them, we are now to see how they aid 
 us in the problem with which we are dealing, viz., the 
 determination of the atomic weights. 
 
MOLECULAR WEIGHTS. 75 
 
 Molecular Weights. If equal volumes of gases contain 
 the same number of molecules at the same temperature 
 and pressure, it is only necessary to determine the 
 weights of equal volumes of gases to learn the relative 
 weights of their molecules. Thus, if we weigh a liter of 
 each of three gases, and find that the weights are to one 
 another as 1 to 2 to 3, then it follows that the relation 
 between the weights of the molecules of these gases is 
 expressed by these figures, or, in other words, the mole- 
 cule of the second gas is twice as heavy ; and that of the 
 third gas is three times as heavy as that of the lightest. 
 The determination of the relative weights of the mole- 
 cules of substances which either are gaseous or can be 
 converted into gases resolves itself simply into a deter- 
 mination of the weights of equal volumes. In represent- 
 ing the molecular weights we may use any figures which 
 are most convenient, provided only that they bear to one 
 another the relations determined by experiment. If, 
 however, we call the atomic weight of hydrogen 1, then 
 our system of molecular weights must be based upon 
 this, and the molecular weight of a compound should 
 state how much heavier the molecule is than an atom of 
 hydrogen. Thus, if we say that the molecular weights 
 of water and hydrochloric acid are respectively 17.88 and 
 36.18, we mean that, if the hydrogen atom weighs 1, then 
 the weights of the molecules of water and hydrochloric 
 acid are represented by the figures given. Now, if the 
 molecule of hydrogen were identical with the atom or, in 
 other words, if the molecular weight were equal to 1, then 
 the adjustment of the system of molecular weights to the 
 atomic weight of hydrogen would be perfectly simple. 
 It would only be necessary to determine the weight of a 
 given volume of hydrogen and compare the weights of 
 equal volumes of other gases with it. If the weight of a 
 certain volume of any compound should be found to be 
 10 times that of an equal volume of hydrogen, it would 
 follow that the molecular weight of the compound is 10. 
 But the molecule and atom of hydrogen are not identical, 
 as can be shown without difficulty. When a given vol- 
 ume of hydrogen combines with chlorine it combines with 
 
76 
 
 INORGANIC CHEMISTRY. 
 
 an equal volume of this element, and the two volumes 
 which combine form an equal volume of the compound 
 hydrochloric acid. These facts may be graphically repre- 
 sented as follows : 
 
 combine and form 
 
 1vol. 
 
 
 
 HC1 
 
 
 
 
 
 2 volumes of 
 
 
 
 - hydrochloric 
 acid gas. 
 
 
 1vol. 
 
 
 HC1 
 
 
 
 1 
 
 j 
 
 
 Now, bearing in mind Avogadro's law that equal vol- 
 umes of all gases contain the same number of molecules, 
 it follows that if, in the volume of hydrogen taken, there 
 is any finite number of molecules, say 100, then in the 
 same volume of chlorine there must be 100 molecules, 
 and in the two volumes of hydrochloric acid gas obtained 
 there must be 200 molecules of the compound. There- 
 fore, from 100 molecules of hydrogen and 100 molecules 
 of chlorine there are formed 200 molecules of hydro- 
 chloric acid. But in each molecule of hydrochloric acid 
 there must be at least one atom of hydrogen and one 
 atom of chlorine, and in the 200 molecules there must be 
 at least 200 atoms of hydrogen and 200 atoms of chlo- 
 rine. Now, these 200 atoms of hydrogen have come from 
 the 100 molecules, and the same is true of the chlorine. 
 It, therefore, follows that each molecule of hydrogen and 
 each molecule of chlorine must consist of at least two atoms. 
 Or, we may say that, if there is one atom of hydro- 
 gen in the molecule of hydrochloric acid, then there 
 are two atoms of hydrogen in the molecule of hydrogen. 
 The assumption that the molecule of hydrogen is twice 
 as heavy as the atom is found to satisfy every require- 
 ment of the present state of our knowledge. From the 
 above, the following rule for the determination of molec- 
 ular weights is deduced : Determine the specific gravity 
 of the substance in terms of hydrogen and multiply the 
 
DEDUCTION OF ATOMIC WEIGHTS. 
 
 result by ,2. Tims the specific gravity of water vapor in 
 terms of hydrogen is 8.94, or, in other words, a given vol- 
 ume of water vapor weighs 8.94 times as much as the 
 same volume of hydrogen. But the molecular weight of 
 hydrogen being 2, that of water must be 17.88. As the 
 specific gravity of gases is frequently stated in terms of 
 the air standard, it is desirable to know the relation be- 
 tween these figures and those based upon hydrogen. The 
 specific gravity of hydrogen as compared with air is 
 .06926 ; when taken as the standard it is represented by 
 1. But 1 -f- 0.06926 is 14.44 ; therefore, to convert the 
 specific gravities on the air standard into those on the 
 hydrogen standard it is only necessary to multiply by 
 14.44. Thus the specific gravity of water vapor, air be- 
 ing the standard, is 0.623. To find its specific gravity, hy- 
 drogen being the standard, multiply by 14.44. Now 0.623 
 X 14.44 is very nearly 9. This being the specific gravity, 
 the molecular weight is obtained by multiplying by 2. 
 Of course, we should reach the same result by multiply- 
 ing the specific gravities in terms of air directly by 28.88. 
 Below are given the molecular weights of a few elements 
 and compounds which have been determined by the 
 method just described : 
 
 Name. 
 
 Sp. Grav. 
 H = 1. 
 
 Molec. 
 Weight. 
 
 Molecular 
 Formula. 
 
 Hydrogen 
 
 1 
 
 2 
 
 Ho 
 
 Nitrogen 
 
 13 93 
 
 27 86 
 
 N 2 
 
 Water 
 
 8 94 
 
 17 88 
 
 H 2 O 
 
 Hydrochloric acid 
 
 18 09 
 
 36 18 
 
 Hen 
 
 Ammonia , 
 
 8 46 
 
 16 93 
 
 NH 3 
 
 
 7 96 
 
 15 92 
 
 CHd 
 
 Carbon monoxide , 
 
 13 90 
 
 27 80 
 
 CO 
 
 Carbon dioxide 
 
 21 84 
 
 43 68 
 
 CO 3 
 
 
 
 
 
 Deduction of Atomic Weights from Molecular Weights. 
 The determination of molecular weights does not neces- 
 sarily carry with it the determination of the atomic 
 weights. It is plain from what has already been said 
 that a knowledge of the molecular weight of an element 
 does not convey a knowledge of its atomic weight. If, 
 
78 INORGANIC CHEMISTRY. 
 
 for example, we learn that the molecular weight of nitro- 
 gen is approximately 28, we have no means of judging 
 from this what the atomic weight is. It is plainly nec- 
 essary to know of how many atoms each molecule of 
 nitrogen is made up, and to learn this is not a simple 
 matter. It is easier to determine the atomic weight of 
 an element through a study of its compounds. Suppose 
 it is desired to determine the atomic weight of oxygen. 
 We first determine the molecular weights of a number of 
 compounds which contain oxygen, and then analyze these 
 compounds. We then see what the smallest figure is 
 that is required to express the weight of the oxygen 
 that enters into the composition of the molecules, and 
 this figure is selected as the atomic weight. The mo- 
 lecular weights and the composition of several oxygen 
 compounds are given in the following table : 
 
 Compound. Mol. Wt. Approx. Composition. 
 
 Water 17.88 2 parts hydrogen, 
 
 15.88 
 
 Carboii monoxide 27.80 11.92 carbon, 
 
 15.88 
 
 Carbon dioxide 43.68 11.92 carbon, 
 
 31.76 
 
 Nitric oxide 29.81 13.93 nitrogen 
 
 15.88 
 
 Nitrous oxide 43.74 27.86 nitrogen, 
 
 15.88 
 
 Sulphur dioxide 63.59 31.83 sulphur, 
 
 31.76 
 
 Sulphur trioxide 79.47 31.83 sulphur, 
 
 47.64 
 
 oxygen. 
 
 oxygen. 
 
 oxygen. 
 
 oxygen. 
 
 oxygen. 
 
 oxygen. 
 
 oxygen. 
 
 The figures in the third column are of course determined 
 by analysis, an example of the methods used having been 
 given in the chapter on water. Stated in ordinary lan- 
 guage, the figures in the case of carbon monoxide mean that 
 the molecule of this compound weighs 27.80 times as much 
 as the atom of hydrogen, and the 27.80 parts of matter 
 are made up of 11.92 parts of carbon and 15.88 parts of 
 oxygen. Considering now the composition of the com- 
 pounds in the table, it will be seen that the smallest mass 
 of oxygen which enters into the composition of any of the 
 
'MOLECULAR FORMULAS. 79 
 
 molecules weighs 15.88 times as much as the atom of 
 hydrogen. We find twice this mass as in carbon dioxide 
 and sulphur dioxide ; and three times as in sulphur tri- 
 oxide, but no smaller mass. Now, if we should examine 
 all compounds of oxygen which can exist in the form of 
 gas or vapor we should find the same thing true ; that is 
 to say, the smallest mass of oxygen which enters into the 
 composition of molecules is 15.88 as great as that of the 
 atom of hydrogen. The conclusion is therefore drawn 
 that 15.88 is the atomic weight of oxygen. The possi- 
 bility that the atomic weight of oxygen is less than this 
 figure is not excluded. It may be that in the simplest 
 oxygen compounds now known there are two or more 
 atoms of this element in the molecules. But in the total 
 absence of evidence on this point all we can do is to 
 accept the figure 15.88 as in perfect accordance with all 
 our knowledge of oxygen compounds. 
 
 In this way the atomic weights of all elements which 
 form gaseous compounds or compounds that can be con- 
 verted into vapor have been determined ; and the deter- 
 minations made in this way are regarded as the most 
 reliable. 
 
 Exact Atomic Weights determined by the Aid of Analy- 
 sis. By determining molecular weights it is possible 
 to decide approximately what figure represents the atomic 
 weight of an element, but the methods employed in making 
 determinations of molecular weights are liable to slight 
 errors, and therefore the atomic weights obtained directly 
 from the molecular weights deviate slightly from the true 
 figures. In order to determine the atomic weights with 
 the greatest possible accuracy, the most refined methods 
 of chemical analysis are brought into play, and the figures 
 in the table on page 21 have been determined in this way 
 by a combination of a study of the specific gravity of 
 gases and by the most careful analyses, together with 
 some other methods which will be taken up later. 
 
 Molecular Formulas. The symbols of chemical com- 
 pounds first used were intended to express simply the 
 composition of the compounds, and this can be done as 
 was explained in Chapter I. by adopting a system of 
 
80 INORGANIC CHEMISTRY. 
 
 combining weights of the elements. According to the 
 theory explained in the last chapter the smallest particle 
 of every compound is a molecule, and each molecule is 
 made up of atoms. It appears, therefore, desirable for 
 the sake of uniformity that the symbols used to repre- 
 sent chemical compounds should represent molecules. 
 Where the molecular weight of a compound, the atomic 
 weights of the elements of which it is composed, and its 
 composition are known, there is no difficulty in represent- 
 ing it by a molecular formula. Thus, the molecular 
 weight of ammonia is found by experiment to be approxi- 
 mately 17, and the 17 parts are made up of 14 parts of 
 nitrogen and 3 parts of hydrogen. The atomic weight 
 of nitrogen is found by the method which has just been 
 described to be very nearly 14. Therefore the mole- 
 cule of ammonia weighing 17 parts is composed of 1 
 atom of nitrogen weighing 14 parts and 3 atoms of hy- 
 drogen weighing 3 parts. The composition of the mole- 
 cule is therefore represented by the* formula NH 3 . Simi- 
 larly the composition of the molecule of water is repre- 
 sented by the formula H 2 O ; that of hydrochloric acid 
 by HOI ; that of marsh gas by CH 4 ; etc., etc. Every 
 formula now in use is intended to represent a molecule of 
 the compound for which it stands. In regard to the 
 molecular weights of compounds that are not gaseous 
 nor convertible into vapor, Avogadro's method is plainly 
 of no avail. Methods have, however, been devised which 
 are applicable to a number of these (see Chapter XXIII). 
 Constitution. When hydrochloric acid is formed, we 
 conceive that each atom of hydrogen combines with one 
 atom of chlorine, and that the molecules of the resulting 
 compound are made up each of an atom of hydrogen and 
 an atom of chlorine. What the act of combination con- 
 sists in we do not know. We simply know that something 
 very remarkable takes place, and that as a consequence 
 the hydrogen and chlorine cease to exist in their original 
 forms. It is idle at present even to speculate in regard 
 to the character of the change. The fact of union is ex- 
 pressed by writing the symbols of the elements side by 
 side without any sign between them, as HC1, or; some- 
 
VALENCE. 81 
 
 times, it is convenient to use a line to indicate chemical 
 union, thus : H-C1. According to the molecular theory 
 the molecule of water consists of two atoms of hydrogen 
 and one of oxygen, as represented by the formula H 2 O, 
 and the question now suggests itself whether all three 
 atoms are in combination with one another or whether 
 each of the hydrogen atoms is in combination with the 
 oxygen atom, but not with each other, as represented by 
 the formula H-O-H. So too in the case of ammonia, 
 the molecular formula of which is NH 3 , the question sug- 
 gests itself : Are the three atoms of hydrogen in combi- 
 nation with the atom of nitrogen, but not with one an- 
 
 / H 
 
 other, as represented in the formula N^-H ? It is ex- 
 Mi 
 
 tremely difficult to answer such questions, but, at the 
 same time, certain facts are known which enable us to 
 draw probable conclusions. Formulas which express the 
 composition of molecules and at the same time express 
 the relations or the connections which exist between the 
 atoms are called constitutional formulas. These constitu- 
 tional formulas are very frequently used at present, but 
 sometimes without a sufficient basis of facts to justify 
 them. Whenever they are used in this book, the rea- 
 sons for them will be stated as fully as may appear nec- 
 essary. 
 
 Valence. The formulas of the hydrogen compounds 
 of chlorine, oxygen, nitrogen, and carbon, all determined 
 by the same method, are 
 
 C1H OH, NH 3 CH 4 . 
 
 A consideration of these formulas and of many similar 
 ones has led to the belief that the atoms of different ele- 
 ments differ in their power of holding other atoms in 
 combination. The simplest explanation of the composi- 
 tion of the compounds above represented is that the 
 atoms of chlorine, oxygen, nitrogen, and carbon differ in 
 their power of holding hydrogen atoms in combination. 
 Hydrogen and chlorine combine in only one way, 1 atom 
 of chlorine combining with 1 of hydrogen ; 1 of oxygen 
 
82 INORGANIC CHEMISTRY. 
 
 combines with 2 of hydrogen; 1 of nitrogen with 3 of 
 hydrogen ; and 1 of carbon with 4 of hydrogen. The 
 limit of the combining power of the atom of chlorine is 
 reached when it has combined with one atom of hydro- 
 gen. And as one chlorine atom can hold but one atom 
 of hydrogen in combination, so one atom of hydrogen 
 can hold but one atom of chlorine. Either the hydrogen 
 atom or the chlorine atom may be taken as an example 
 of the simplest kind of atom. Any element like hydro- 
 gen or chlorine is called a univalent element ; an element 
 like oxygen whose atom can hold two unit atoms in 
 combination is called a bivalent element ; an element like 
 nitrogen whose atom can hold three unit atoms in com- 
 bination is called a trivalent element ; and an element like 
 carbon whose atom can hold four unit atoms in combina- 
 tion is called a quadrivalent element. Most elements be- 
 long to one or the other of these four classes, though 
 there are some which can hold five, six, and even seven 
 unit atoms in combination. These are called quinqui- 
 valent, sexivalent, and septivalent respectively. 
 
 Valence is defined as that property of an element by 
 virtue of which its atom can hold a definite number of 
 other atoms in combination. In the formation of com- 
 pounds the valence of the elements determines how many 
 atoms of any element can enter into combination with 
 any other. The atoms are sometimes spoken of as hav- 
 ing bonds which are graphically represented by lines. 
 Thus, a univalent element is said to have one bond, as 
 represented by H-, C1-, etc. ; a bivalent element is said 
 to have two bonds, -O-, -S-, etc. ; .a trivalent element 
 
 three, -N- ; and a quadrivalent element four, -C-. Of 
 
 course, this is merely a symbolical representation of the 
 idea that each atom has a definite power of combining 
 with others. It is further said that when the atoms unite 
 these bonds become satisfied. Thus when one atom of 
 hydrogen unites with one of chlorine, the bond of each 
 is regarded as uniting with the bond of the other, and this 
 is represented by the symbol H-C1. So too, when two 
 atoms of hydrogen unite with one of oxygen, the com- 
 
REPLACING POWER OF ELEMENTS. 83 
 
 TT 
 
 pound is represented in this way : H-O-H or O<TJ- In 
 
 the union of atoms, further, to use the figurative lan- 
 guage, two bonds may be satisfied by two univalent atoms 
 or by one bivalent atom. Thus, in marsh gas, CH 4 , the 
 four affinities or bonds of the quadrivalent carbon atom 
 are regarded as being satisfied by the four univalent 
 
 H 
 
 hydrogen atoms. In the compound H-C=O, however, 
 two of the bonds are regarded as satisfied by two univa- 
 lent hydrogen atoms, and the other two by one bivalent 
 oxygen atom ; and, again, in the compound O=C O, the 
 four bonds of the carbon atom are regarded as satisfied 
 by two bivalent oxygen atoms. 
 
 Replacing Power of Elements. As has been seen, 
 when potassium acts upon water the action is represented 
 by the equation 
 
 Expressing this reaction by means of formulas which 
 take the valence of the elements into consideration we 
 have the following : 
 
 K + H-O-H = K-O-H + H. 
 
 According to this, one atom of potassium is substituted 
 for one atom of hydrogen in water, and, in the compound 
 formed, the atom of potassium is regarded as occupying 
 the place of the hydrogen atom. The elements calcium 
 and barium are bivalent like oxygen, as shown in the 
 compounds CaCl 2 and BaCl 2 , in which the atom of cal- 
 cium as well as that of barium holds two univalent atoms 
 of chlorine in combination. When these elements act 
 upon water, hydrogen is liberated as with potassium, but 
 each atom replaces two atoms of hydrogen, as represent- 
 ed in the equations 
 
 Ca + 2H 2 O = CaH a O 3 + 2H ; 
 Ba + 2H 2 O = BaH 2 O a + 2H. 
 
84 INORGANIC CHEMISTRY. 
 
 Or expressing the changes by valence formulas, we have 
 
 -o , H-O-H -o 0-H 
 Ba + H-O-H = Ba< 0-H + H '* 
 
 A trivalent element acting in the same way would give a 
 
 /0-H 
 compound of the general formula M( O-H, in which M 
 
 \O-H 
 represents any trivalent element. 
 
 So, too, in the action of various elements upon hydro- 
 chloric acid, HC1, a univalent element like sodium or 
 potassium replaces one hydrogen atom in one molecule 
 of the acid, and forms a compound of the general formula 
 MCI, as, for example, KC1 and NaCl. A bivalent ele- 
 ment, like zinc, replaces two atoms of hydrogen in two 
 molecules of the acid, forming the compound ZnCl 2 : 
 
 A trivalent element forms a compound of the general 
 formula MC1 3 , as, for example, aluminium chloride : 
 
 /Cl 
 
 AlCl. 
 
 The above explanation of the hypothesis of valence will 
 suffice by way of introduction to the subject. The re- 
 lation which it bears to the facts is simply this : On study- 
 ing the composition of chemical compounds we find that, 
 in general, the elements combine with one another in 
 comparatively few proportions. Thus, hydrogen and 
 chlorine combine with each other in only one proportion ; 
 hydrogen and oxygen in two ; nitrogen and hydrogen 
 in four ; etc. There is something limiting the com- 
 plexity of compounds. We might study the laws govern- 
 ing the complexity of compounds without any hypothesis 
 as to the cause, but the hypothesis of valence is a con- 
 venient explanation of these laws, and it has been of much 
 service in furnishing chemists with a simple language for 
 representing chemical changes. 
 
CHAPTER VII. 
 
 OZONE ALLOTROPY NASCENT STATE HYDROGEN 
 DIOXIDE. 
 
 Occurrence. Ozone lias long been thought to be pres- 
 ent in the air in small quantity, but careful research has 
 made this occurrence appear doubtful. 
 
 Preparation. It was observed in the last century that, 
 when a powerful electric machine is worked in a room, 
 something is formed which has a strong odor, and the 
 same odor was noticed during thunder-showers. After- 
 wards the same thing was noticed when water is decom- 
 posed by an electric current, and when phosphorus is ex- 
 posed to moist air. The substance was at first supposed 
 to be a compound of water and oxygen, but long-contin- 
 ued investigation showed that it could be made from pure 
 dry oxygen, and that by heat and other means it is de- 
 composed into nothing but ordinary oxygen. 
 
 It is prepared mixed with oxygen by passing electric 
 sparks through ordinary oxygen in an apparatus con- 
 structed on the principle of that represented in Fig. 4. 
 
 FIG. 4. 
 
 AA is a glass tube about an inch in diameter closed at 
 the ends with brass caps or with corks, covered with shel- 
 lac on the inner side. A metallic cylinder BB covered 
 with tin-foil is placed inside this tube. The cylinder is 
 connected with the tubes (70, and through these and the 
 
 85 
 
86 INORGANIC CHEMISTRY. 
 
 cylinder a current of cold water is kept flowing. The 
 oxygen passes through the glass tube by means of the 
 tubes DD, and necessarily passes through the narrow 
 space between the glass tube and the metallic cylinder. 
 Around the outside of the glass tube is wound a strip of 
 tin-foil. By means of the wires F and E connection is 
 established with the poles of an induction-coil, or a Holtz 
 electrical machine. 
 
 Ozone is also made by placing a few pieces of phos- 
 phorus in the bottom of a good-sized bottle and partly 
 covering them with water ; and, finally, it is made by 
 treating barium dioxide, BaO 2 , and some other compounds 
 rich in oxygen with sulphuric acid. 
 
 Properties. Ozone is a gas which can be condensed to 
 the liquid form, the liquid having a blue color. It has a 
 strong odor, and acts in an irritating way upon the mem- 
 branes lining the throat. Its chemical conduct is entirely 
 different from that of oxygen. While the latter at ordi- 
 nary temperatures is not an active element, ozone is. 
 It acts upon most substances of animal or vegetable 
 origin, and oxidizes nearly all the metals, besides pro- 
 ducing a variety of other changes which are not produced 
 by oxygen at the ordinary temperature. Among the 
 characteristic changes which may be made use of for the 
 purpose of detecting ozone are the following : (1) It 
 liberates iodine when brought in contact with potassium 
 iodide ; the action being represented thus : 
 
 2KI + H 2 O + O = 2KOH + 21. 
 
 Now, iodine has the power to turn starch blue. There- 
 fore, if ozone is brought into a solution containing starch 
 and potassium iodide, a blue color is produced. (2) Ozone 
 combines with metallic silver to form silver peroxide, 
 which is brown. In order to detect ozone, therefore, a 
 strip of polished silver may be exposed to the gas, and if 
 it turns brown ozone is present. 
 
 When heated to 300 ozone loses its characteristic odor 
 and is converted into ordinary oxygen. It thus appears 
 that we can start with ordinary oxygen, and by means of 
 an electric current convert it into a substance with much 
 
RELATION BETWEEN OXYGEN AND OZONE. 
 
 more active chemical properties and differing from it so 
 markedly that one would hardly suspect the close relation 
 between the two ; and then, further, by simply passing 
 the active substance, or ozone, through a tube heated to 
 300, it is converted back again into oxygen without loss 
 of weight. 
 
 Relation between Oxygen and Ozone. When experi- 
 ment had shown that oxygen and ozone are convertible 
 one into the other without change of weight, the suggestion 
 was made that the difference between them might be due 
 to a difference in the number of atoms of oxygen contained 
 in the molecule of each. It might be, for example, that 
 in the molecule of oxygen there are two atoms and in that 
 of ozone three or more. If this view is correct there 
 should be a difference between the specific gravities of 
 the two gases, and by a study of this difference it should 
 be possible to draw a conclusion as to the constitution of 
 the molecule of ozone. That there is a change of volume 
 when oxygen is changed to ozone was shown by enclosing 
 the former in a tube constructed as shown in Fig. 5. The 
 large part of the tube is furnished with two platinum 
 wires which pass through the glass. By means of these 
 wires a silent discharge of electricity is 
 kept up through the oxygen, and it is thus 
 partly converted into ozone. In the smaller 
 bent part of the tube there is a small col- 
 umn of concentrated sulphuric acid which 
 serves as a stopper. Any change in the 
 volume of the gas in the tube will cause 
 this sulphuric acid to change its position. 
 Now, during the conversion of the oxygen 
 into ozone it was noticed that the volume 
 of the gas decreased, and that, on heating 
 the tube and thus converting the ozone 
 into oxygen again, the original volume of 
 the latter was restored. Unfortunately 
 for this purpose it is not possible by any 
 means to convert more than a comparatively 
 small proportion of the oxygen into ozone, 
 and it is not possible in the experiment described to de- 
 
 FIG. 5. 
 
88 INORGANIC CHEMISTRY. 
 
 termine the relation between the decrease in volume and 
 the extent of the conversion. 
 
 Certain ethereal oils, as oil of turpentine, have the 
 power to absorb ozone without decomposing it. This 
 fact has been taken advantage of for the purpose of de- 
 termining the constitution of ozone. The method con- 
 sisted in ozonizing oxygen, heating some of the mixture 
 and measuring the increase in volume caused by conver- 
 sion into oxygen ; and then treating some of the same 
 mixture with oil of turpentine, and observing the de- 
 crease in volume. By this means it is possible to cal- 
 culate the change in volume which takes place when 
 ozone is converted into oxygen. The results showed 
 that two volumes of ozone give three volumes of oxygen, 
 or when a given volume of oxygen is converted into 
 ozone there is contraction to two-thirds the original vol- 
 ume. According to this, the molecular weight of ozone 
 by Avogadro's law must be about 48, for from 3 volumes 
 of oxygen there are formed 2 volumes of ozone, which, 
 interpreted according to the molecular theory, means 
 that three molecules of oxygen must be converted into 
 two of ozone ; but 3 molecules of oxygen weigh 96 parts ; 
 therefore, 2 molecules weighing the same, the molecular 
 weight of ozone is 48. The further conclusion is justified 
 that, while the molecule of ordinary oxygen consists oi 
 two atoms, that of ozone consists of three atoms. The 
 difference may be represented by the formulas O 2 or 
 
 O 
 
 A 
 O=O for oxygen, and O 3 or O-O for ozone. 
 
 Ozone in the Air. No satisfactory evidence of the oc- 
 currence of ozone in the air has ever been furnished. Yery 
 little if anything is known in regard to the effect of ozone 
 on the health. Larger quantities would undoubtedly act 
 injuriously. It is commonly believed that small quanti- 
 ties are advantageous, as it tends to destroy substances 
 which are unwholesome. This subject has not, however, 
 been studied with sufficient care to justify a positive 
 opinion in regard to it. The difficulty in drawing a con- 
 clusion is increased by the fact that there are other sub- 
 
ALLOTROPT. 89 
 
 stances present in the air which resemble ozone in some 
 respects, as, for example, hydrogen dioxide, H 2 O a (which 
 see). 
 
 Allotropy. The occurrence of an element in two or 
 more different modifications is called cdlotropy. Thus, 
 ozone is called an allotropic form of oxygen. There are, 
 however, a number of other cases of a similar kind. 
 Phosphorus, for example, presents itself in three or four 
 different varieties which, in some respects, differ mark- 
 edly from one another. Carbon also appears in three 
 different forms. Whether in all cases the difference be- 
 tween the allotropic forms of an element is due to a dif- 
 ference in the molecular constitution, as in the case of 
 oxygen, is impossible to decide at present, for the 
 reason that the molecular weights of the substances can- 
 not always be determined. The facts learned in regard to 
 the relations between oxygen and ozone make it appear 
 quite probable that the explanation which holds good for 
 this case will probably hold good for others. 
 
 Varying Number of Atoms in the Molecules of one and 
 the same Element. It has been pointed out that the mol- 
 ecules of hydrogen, oxygen, and some other elements 
 consist of two atoms each. We have just seen that the 
 molecule of ozone consists of three atoms. There are 
 some elements w r hich contain a different number of atoms 
 in their molecules, according to the temperature. Thus, 
 sulphur is a solid substance which boils at 440 C. The 
 specific gravity of its vapor between 450 and 500 leads 
 to the molecular weight about 192. But the atomic 
 weight of sulphur is very nearly 32, therefore the mole- 
 cule of sulphur just above its boiling-point would appear 
 to be made up of six atoms. As the temperature is 
 raised, the specific gravity becomes less, and above 800 
 it corresponds to the molecular weight 64, showing that 
 at this high temperature the molecule apparently con- 
 sists of two atoms. Above that point the specific gravity 
 does not change materially. According to the latest 
 investigations on this subject, it appears probable that 
 the higher specific gravity of sulphur vapor at the 
 
90 INORGANIC CHEMISTRY. 
 
 lower temperature is to be ascribed to imperfect vaporiza- 
 tion of the substance, and that the evidence that the 
 molecule is more complicated at lower than at higher 
 temperature is not conclusive. Facts somewhat similar to 
 those just mentioned have been brought to light in the case 
 of iodine, only here the vapor at the lower temperature has 
 a specific gravity showing that the molecule consists of two 
 atoms, but at very high temperatures the specific gravity 
 shows that the molecule contains only one atom. The 
 atomic weight of mercury is the same as its molecular 
 weight, or, in other words, the molecule and atom are in 
 this case identical. 
 
 Nascent State. One of the most curious phenomena 
 connected with chemical action is that of the nascent 
 state. Some elements, which under ordinary circum- 
 stances are inactive, show themselves possessed of marked 
 activity if they are allowed to act the instant they are set 
 free from their compounds. Thus, if carbon monoxide, 
 CO, and ordinary oxygen are brought together at the or- 
 dinary temperature, no action takes place. If, however, 
 carbon monoxide is brought in contact with something 
 which easily gives off oxygen, it is converted into carbon 
 dioxide. At ordinary temperatures, for example, chro- 
 mium trioxide gives up oxygen to carbon monoxide as 
 represented thus : 
 
 SCO + 2CrO 3 = 3CO 2 + Cr 2 O 3 . 
 
 So, too, hydrogen at ordinary temperatures is an inactive 
 element, but, if brought in contact with substances the 
 instant it is set free from a compound, it produces many 
 marked changes. Many substances which would be left 
 unchanged, if hydrogen were passed over them at ordinary 
 temperatures, are changed if put in the liquid from which 
 the hydrogen is being evolved. The most plausible ex- 
 planation of facts like these is to be found in the molec- 
 ular theory. Free oxygen and free hydrogen consist of 
 molecules made up of atoms which are in combination 
 with each other. Before these molecules can act chemi- 
 cally they must be broken up into atoms. When carbon 
 
HYDROGEN DIOXIDE OR HYDROGEN PEROXIDE. 91 
 
 monoxide is brought together with oxygen in the mole- 
 cular state, the condition is represented thus : 
 
 co + o, 
 
 Before combination can take place the molecule, O 2 , 
 must be decomposed as represented thus : 
 
 0, = + O ; 
 
 and the atoms of oxygen then combine with the carbon 
 monoxide. Now there are some substances which yield 
 oxygen atoms more readily than ordinary oxygen does. 
 This is true of chromium trioxide, CrO 3 , which in con- 
 tact with substances that have the power to combine 
 with oxygen breaks up thus : 
 
 2CrO 3 = Cr,O 3 + O + O + O. 
 
 When hydrogen is liberated from a compound by 
 chemical action the uncombined atoms are believed to be 
 given off first ; if there is nothing present with which 
 the hydrogen can combine, these atoms combine with 
 each other in pairs to form the molecules of ordinary 
 hydrogen. 
 
 The examples furnished by allotropism and the phe- 
 nomena of the nascent state show how chemical facts 
 which otherwise appear entirely incomprehensible are 
 satisfactorily explained by the aid of the molecular 
 theory. 
 
 Hydrogen Dioxide or Hydrogen Peroxide, H 2 O 2 . When 
 barium dioxide is treated with hydrochloric acid or di- 
 lute sulphuric acid the reactions represented by the fol- 
 lowing equations take place : 
 
 BaO 2 + 2HC1 = BaCl 2 + H 2 O 2 ; and 
 Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 O 2 . 
 
 In the latter case the compound BaSO 4 , or barium 
 sulphate, is insoluble, and the liquid can easily be sep- 
 arated from it by filtration. If the liquid is boiled, 
 decomposition of the hydrogen dioxide takes place, oxy- 
 
92 INORGANIC CHEMISTRY. 
 
 gen being liberated and water left behind. The decom- 
 position is represented thus : 
 
 2H a 2 = 2H 2 + O a ; or H 3 O 9 = H,O + O. 
 
 The dioxide can be obtained in the form of a color- 
 less liquid by allowing the solution to stand in a vacuum 
 over sulphuric acid, or by distilling its solution in 
 a vacuum. It boils at 84 85 under a pressure of 
 68 mm. 
 
 Properties. Hydrogen dioxide is a clear, syrupy liquid 
 which does not solidify at 30. It is characterized by 
 marked instability. It breaks down slowly even at or- 
 dinary temperatures if simply allowed to stand. The de- 
 composition takes place easily under the influence of 
 heat, and if heated rapidly to 100 explosion is apt to 
 take place. The products are water and oxygen, as indi- 
 cated above. In consequence of the ease with which hy- 
 drogen dioxide gives off oxygen it is a good oxidizing 
 agent, and it is now manufactured on the large scale for 
 use in bleaching and in medicine, and comes into the 
 market in solutions of different strengths. The solution 
 has a bitter taste ; if concentrated, it affects the skin, 
 causing a pricking sensation, and making white spots, 
 which, however, disappear in a few hours. As hydrogen 
 dioxide cannot be converted into vapor, its molecular 
 weight cannot be determined by the method of Avo- 
 gadro. 
 
 Determinations made by other methods, depending 
 upon observations on the freezing-point of its solutions 
 (see Chapter XXIII), show that its molecule is repre- 
 sented by the formula H 2 O 2 , and not by HO. Assuming 
 that the formula is H 2 O 2 , what is the constitution ? "We 
 have very little clear evidence on this point, but the most 
 prevalent view is that the atoms are connected as repre- 
 sented in the formula H-O-O-H, in which the oxygen 
 atoms are in combination with each other, and also with 
 hydrogen. The principal ground upon which this con- 
 ception rests is that this is the only way in which two 
 hydrogen atoms and two oxygen atoms can be joined to- 
 
HYDROGEN DIOXIDE CHARACTERISTIC REACTIONS. 93 
 
 gether, if the oxygen atoms are bivalent and the hydro- 
 gen atoms univalent. This, clearly, is not good evidence, 
 as it is quite possible that oxygen may not be bivalent in 
 this compound. Indeed, there are some facts which 
 seem to show pretty conclusively that the oxygen atoms 
 are quadrivalent, and that the constitution of hydrogen 
 dioxide should be represented by the formula 
 
 H-0 = 0-H. 
 
 It is impossible here to discuss the significance of this 
 formula. Suffice it to say that it is largely based upon 
 the results of observations on the refractive power of the 
 compound. 
 
 Occurrence in the Air. That hydrogen dioxide occurs 
 in the air has already been stated. It is also found in 
 rain and snow. The quantity in the air is extremely 
 small, and it varies at different times of the day, the 
 .action of sunlight being evidently favorable to its forma- 
 tion. 
 
 Characteristic Reactions. Like ozone, hydrogen diox- 
 ide decomposes potassium iodide, setting iodine free : 
 
 H 2 O 2 + 2KI = 2KOH + I 2 . 
 
 This fact may be utilized for the purpose of detecting 
 the compound. The separation of the iodine does not 
 take place readily as in the case of ozone, but the action 
 is hastened by the addition of a very small quantity of 
 .a dilute solution of ferrous sulphate, FeSO 4 . An acid 
 solution of potassium permanganate is decolorized by 
 Jiydrogen dioxide. If in a glass cylinder a layer of ether 
 is poured upon a solution of hydrogen dioxide and 
 a drop of a solution of potassium dichromate is then 
 added, and the cylinder thoroughly shaken, the ether 
 will take up a blue compound, and will itself become 
 blue. 
 
 When hydrogen dioxide is brought together with sub- 
 stances which give up oxygen readily, action generally 
 takes place involving decomposition of the hydrogen 
 dioxide as well as of the other substance. Thus, when 
 
94 INORGANIC CHEMISTRY. 
 
 it is brought together with silver oxide, Ag a O, this reac- 
 tion takes place : 
 
 So, also, it undergoes decomposition with ozone as rep- 
 resented thus : 
 
 O, + HA = H,0 + 20,. 
 
 The explanation of these facts is to be found partly in 
 the attraction of the atoms of oxygen for each other. 
 In the molecule of silver oxide and of hydrogen dioxide 
 there is an atom of oxygen which is held loosely. When 
 the substances are brought together these loosely com- 
 bined atoms attract and combine with each other, as 
 may be represented thus : 
 
 (0,)0 + 0(H,0) = O, + O, + H,0. 
 
 Thermochemical Considerations. By methods which 
 need not be described here, it has been determined that 
 when ozone is converted into oxygen heat is evolved, 
 and the thermochemical equation expressing the facts is 
 this : 
 
 2O 3 = 3O 2 = + 2 X 36,200 cal. 
 
 In accordance with the explanation given on page 37, 
 this means that when two molecules of ozone are con- 
 verted into three molecules of oxygen 72,400 c. are 
 evolved. So, also, when oxygen .is converted into ozone 
 heat is absorbed, the equation being 
 
 (O 2 , O) = O 3 = - 36,200 cal. 
 
 To convert oxygen into ozone therefore requires an 
 addition of energy. A reaction which requires an addi- 
 tion of heat is called an endothermic reaction, and one 
 which takes place with an evolution of heat is called an 
 exothermic reaction. In general, that exothermic reaction 
 which is accompanied by the greatest evolution of heat 
 takes place most readily, and endothermic reactions do 
 not take place without the addition of energy from with- 
 
THERMOCHEMICAL CONSIDERATIONS. 95 
 
 out. In the language of physics, we say that ozone con- 
 tains more energy than oxygen, and therefore it acts 
 more readily. 
 
 Somewhat similar relations are observed between hy- 
 drogen dioxide and water. The decomposition of the 
 dioxide into water and oxygen is accompanied by an evo- 
 lution of heat : 
 
 H a O 3 = H a O + ; = + 23,100 cal. ; 
 
 and the formation of the dioxide requires the addition of 
 the same amount of heat : 
 
 (H 2 O, O) = - 23,100 cal. 
 
 In order to get the dioxide, therefore, the reaction must 
 be of such a character as to furnish this amount of heat. 
 In the action of hydrochloric acid upon barium dioxide 
 there is more heat evolved than is required in the forma- 
 tion of hydrogen dioxide, and therefore the formation in 
 this way is possible. 
 
CHAPTER VIII. 
 
 CHLORINE HYDROCHLORIC ACID. 
 
 Historical. Sodium chloride or common salt, which is 
 the principal chlorine compound found in nature, has- 
 been known for a very long time. In 1774 Scheele first 
 called attention to chlorine in his treatise on the black 
 oxide of manganese or manganese dioxide. In accord- 
 ance with the ideas then prevailing, he called it dephlo- 
 gisticated muriatic acid. Berthollet suggested in 1785 
 that it was oxidized hydrochloric acid, and it was then 
 regarded as consisting of the hypothetical element, mu- 
 rium, in combination with oxygen. In 1810 Davy pointed 
 out that the idea, previously expressed by Gay-Lussac 
 and Thenard, that the substance is an element, is in the 
 highest degree probable, and he gave it the name chlorine 
 (from XXoopos, greenish-yellow). Since that time every- 
 thing learned in regard to chlorine has gone to show that 
 it is an element. 
 
 Occurrence of Chlorine. Though widely distributed in 
 nature, chlorine never occurs in the uncombined state, for 
 the reason that it combines with other substances with 
 great ease, and, if it were set free, it would at once enter 
 into combination. It does not occur in very large quantity 
 as compared with oxygen and hydrogen. It is found 
 chiefly in combination with the element sodium as common 
 salt, or sodium chloride, a compound of the composition 
 represented by the formula NaCl. It is also found in 
 combination with other elements, as potassium, magne- 
 sium, etc., as in the celebrated mines at Stassfurt, Ger- 
 many. In comparatively small quantity it occurs in 
 combination with silver, forming one of the most valua- 
 ble silver ores. All the chlorine which we have to deal 
 with is made from common salt. 
 
 (96) 
 
PREPARATION OF CHLORINE. 97 
 
 Preparation. The problem to be solved in the prepa- 
 ration of chlorine from common salt is the separation 
 of the two elements sodium and chlorine. This cannot 
 be accomplished directly as the separation of mercury 
 and oxygen in the decomposition of mercuric oxide, HgO, 
 and there is no easily obtained compound which gives off 
 chlorine when heated. The method adopted consists in 
 making hydrochloric acid, HC1, from sodium chloride, and 
 then treating this acid with some substance which readily 
 gives off oxygen. The change of sodium chloride to hydro- 
 chloric acid is readily accomplished by treating salt with 
 ordinary sulphuric acid a reaction carried on on the 
 large scale in the manufacture of sodium carbonate or 
 " soda." When the two are brought together a change 
 takes place which will be studied more in detail farther 
 on. The reaction is represented by the equation 
 
 (1) 2NaCl + H 2 SO 4 = Na 2 SO 4 + 2HC1. 
 
 Sodium Sulphuric Sodium Hydrochloric 
 
 chloride acid sulphate acid 
 
 As will be seen, the sodium of the sodium chloride and 
 the hydrogen of the sulphuric acid exchange places a 
 kind of action which is quite common. 
 
 The decomposition of the hydrochloric acid and libera- 
 tion of chlorine under the influence of oxygen takes place 
 as represented in this equation : 
 
 (2) 2HCl + O = H a O + Cl 2 . 
 
 As there is an unlimited supply of oxygen in the air, it 
 would be advantageous if the decomposition of the hydro- 
 chloric acid could be effected by means of the element in 
 the free state. But free oxygen alone will not accom- 
 plish the change. A process has been invented, how- 
 ever, for the manufacture of chlorine on the large scale 
 which depends upon the decomposition of hydrochloric 
 acid by the oxygen of the air. This is Deacon's process. 
 It consists in passing hydrochloric acid and air together 
 through a heated tube containing clay balls saturated 
 with a solution of copper sulphate, and then dried. If 
 
98 INORGANIC CHEMISTRY. 
 
 the temperature of the tube is not raised too high the 
 copper sulphate remains unchanged. Exactly why the 
 oxidation of the hydrochloric acid takes place under 
 these circumstances is not positively known, but it prob- 
 ably depends upon the formation and decomposition of 
 intermediate products. Deacon's process is used quits 
 extensively on the continent of Europe, while in England 
 and Scotland another method, apparently more compli- 
 cated, known as Weldon's process, is chiefly used. In the 
 laboratory chlorine is generally made by treating hydro- 
 chloric acid with manganese dioxide. The reaction is 
 represented thus : 
 
 (1) MnO 2 + 4HCl = MnCl 2 + 2H 2 O + Cl 2 . 
 
 This is explained by the tendency of hydrogen to com- 
 bine with oxygen to form water. When the compound 
 MnO is treated with hydrochloric acid, this reaction 
 takes place : 
 
 MnO + 2HC1 = MnCl 2 + H 2 O. 
 
 In this case there is a simple exchange of places by the 
 manganese and hydrogen and the oxygen and chlorine, 
 the great affinity of hydrogen for oxygen being a promi- 
 nent cause of the change. 
 
 So, also, when manganese dioxide is treated with hy- 
 drochloric acid, the oxygen may be first replaced by 
 chlorine, as represented in the equation 
 
 MnO 2 + 4HC1 = MnCl 4 + 2H 2 O. 
 
 But the compound MnCl 4 gives up half its chlorine when 
 heated : 
 
 MnCl 4 = MnCl 2 + C1 2 ; 
 
 so that the action of hydrochloric acid on manganese 
 dioxide is represented by the equation (1) above. Some 
 recent investigations make it appear probable that the 
 reaction is more complicated than it is here repre- 
 sented to be ; that the first product of the action of 
 hydrochloric acid on manganese dioxide is a compound 
 
CHLORINE WELDON'S PROCESS. 99 
 
 of the composition H 2 MnCl 6 (or H 2 MnCl B ) ; and that this 
 breaks down tinder the influence of heat thus : 
 
 H 2 MnCl 6 = MnCl, + 2HC1 + C1 2 . 
 
 These reactions will be taken up under the head of Man- 
 ganese (which see). 
 
 Instead of making hydrochloric acid from salt, and 
 then treating it with manganese dioxide, it is better to 
 mix the manganese dioxide and common salt together, 
 and pour upon the mixture the necessary quantity of 
 sulphuric acid. In this case the manganese dioxide and 
 sulphuric acid give off oxygen, and the common salt and 
 sulphuric acid give off hydrochloric acid. The oxygen 
 then oxidizes the hydrochloric acid, and chlorine is liber- 
 ated. At least this is a probable explanation of the reac- 
 tion, for it is known that when manganese dioxide is 
 heated with sulphuric acid oxygen is liberated, as repre- 
 sented in the equation 
 
 Mn0 2 + H 3 S0 4 = MnS0 4 + H 2 O + O. 
 
 Weldon's Process. As there is a large demand for 
 chlorine, much attention has been given to the improve- 
 ment of the methods for its preparation. One of the 
 objections to the ordinary method is the comparatively 
 high price of the mineral, manganese dioxide. As this is 
 converted into the chloride, MnCl 2 , in the preparation of 
 chlorine, and the chloride is of no value, the expense of 
 preparation is quite high. A process has been invented 
 for the regeneration of the manganous chloride, MnCl,, 
 or for the conversion of this compound into an oxygen 
 compound which with hydrochloric acid will give chlo- 
 rine. This is Weldon's process. It will be taken up 
 under the head of Manganese (which see). 
 
 Electrolytic Process. Various processes have been 
 devised for the preparation of chlorine by the action of 
 an electric current on a solution of sodium chloride or of 
 hydrochloric acid. Such electrolytic processes are now 
 in use, and by means of them the price of chlorine has 
 been much lowered. 1 
 
 Properties. Chlorine is a greenish-yellow gas. It has 
 
100 INORGANIC CHEMISTRY. 
 
 a disagreeable odor, and acts upon the membranes lining 
 the throat and nose, causing irritation and inflammation, 
 suggesting a " cold in the head." Inhaled in concen- 
 trated form it causes death. Its specific gravity at 20 
 is 2.48 (air = 1), and as compared with hydrogen it is 
 35.18. A liter of chlorine gas, under standard conditions, 
 weighs 3.162 grams. It is soluble in water and acts 
 upon mercury, and therefore cannot be collected by dis- 
 placement of either of these liquids. The most conven- 
 ient way to collect it is by displacement of air. It can 
 also be collected over warm water in which it is less 
 soluble than in cold water, or over a saturated solution 
 of sodium chloride in which it is but slightly soluble. 
 It is easily compressed to a liquid, and is now sold in 
 this form in cast-iron cylinders. 
 
 It is a remarkably active substance, combining with 
 or acting in some way upon most other substances even 
 at ordinary temperature. This activity may be illus- 
 trated by introducing into vessels containing chlorine a 
 little finely powdered antimony, a few pieces of thin 
 copper-foil, a piece of paper with ink-marks on it, some 
 flowers, and pieces of cotton-prints. The antimony will 
 take fire and a white substance will be formed. The 
 reaction is represented by the following equation : 
 
 Sb + 301 = Sb01 3 . 
 
 The copper-foil also takes fire, and is converted into a 
 chloride as represented thus : 
 
 Cu + 201 = CuCl a . 
 
 Many other substances unite directly with chlorine 
 with evolution of heat and light, and form compounds 
 which are called chlorides. This kind of action is of 
 the same character as that which takes place in 
 oxygen. There is, however, this difference between re- 
 actions in oxygen and in chlorine : the latter frequently 
 take place at ordinary temperatures, whereas those in 
 oxygen require an elevation of temperature to start 
 them. In both cases the gases combine with the other 
 
PROPERTIES OF CHLORINK^ /''>-" 10 
 
 substances directly, and disappear as such, and the 
 light and heat are caused by the act of combination. 
 
 Dry liquid chlorine, at its boiling temperature (33. 6) 
 does not act on potassium, sodium, or aluminium. 
 
 The action of chlorine upon ink, flowers, and cotton- 
 print illustrates its power to bleach. It is important to 
 notice that if the colored objects be introduced dry into 
 dry chlorine the action does not take place. Moisture 
 is generally essential to the bleaching by chlorine. 
 Chlorine acts directly upon some dye-stuffs, converting 
 them into colorless substances. In other cases it has 
 been shown that the destruction of the color is due to 
 the action of oxygen, which is set free from water by 
 chlorine. In direct sunlight chlorine decomposes water 
 according to the equation 
 
 This decomposition can be illustrated by filling a 
 long tube with a solution of chlorine in water, and in- 
 verting it in a shallow vessel containing some of the 
 same solution. If this is placed in the direct sunlight, 
 bubbles of gas will be seen to rise in the tube and these 
 will collect at the top, while the color of the solution, 
 which was at first greenish-yellow like that of chlorine, 
 will disappear. The gas which collects in the upper 
 part of the tube is oxygen. 
 
 The disintegrating action of chlorine upon substances 
 of animal and vegetable origin may be illustrated by 
 moistening a piece of filter-paper with some oil of turpen- 
 tine, and introducing it into a vessel of chlorine. A flash 
 of light is seen, and a dense black cloud is formed. 
 The black substance is mainly carbon. Oil of turpen- 
 tine is a compound of carbon and hydrogen. Chlorine 
 abstracts the hydrogen from the carbon, leaving the 
 latter mainly in the uncombined state. If the chlorine 
 is allowed to act slowly upon the oil of turpentine and 
 similar organic substances, the chlorine is substituted 
 atom for atom for the hydrogen, and a series of so-called 
 substitution-products is obtained. 
 
"102 INORGANIC CHEMISTRY. 
 
 Chlorine dissolves readily in water and forms a solu- 
 tion known as chlorine water. It has the odor and color 
 of the gas, and it is frequently used in the laboratory 
 instead of the gas. From what has been said it is evi- 
 dent that it must be kept protected from the sunlight, 
 or decomposition will take place, resulting in the forma- 
 tion of hydrochloric acid and oxygen. 
 
 Different Kinds of Action. A careful study of the dif- 
 ferent kinds of action exhibited by chlorine shows that 
 they may be classified under three heads : 
 
 (1) First it acts by direct combination with elements 
 as in the experiments with antimony and copper, and, as 
 will be shown, with hydrogen and many other elements. 
 Just as the compounds of oxygen with other elements 
 are called oxides, so the compounds of chlorine with 
 other elements are called chlorides. Thus the compound 
 of antimony and chlorine, SbCl 3 , is called antimony chlo- 
 ride ; that of zinc and chlorine, ZnCl 2 , is called zinc 
 chloride ; etc. In case an element forms more than one 
 compound with chlorine, the names used to distinguish 
 between them are similar to those used for oxides. 
 Mercury forms two chlorides which have the composi- 
 tion represented by the formulas HgCl and HgCl 2 . The 
 one with the smaller proportion of chlorine is called 
 mercurous chloride , and the one with the larger propor- 
 tion of chlorine is called mercuric chloride. So, too, 
 there are two chlorides of iron which correspond to the 
 formulas FeCl 2 and FeCl 3 . The former is called ferrous 
 chloride, and the latter ferric chloride. 
 
 (2) The second kind of action of chlorine is that which 
 is called substitution. This was referred to in connec- 
 tion with the action of chlorine on the oil of turpentine. 
 The general character of this kind of action may be ex- 
 plained by the aid of the following example. There is 
 an important compound of carbon and hydrogen called 
 benzene, which has the formula C 6 H 6 . When chlorine is 
 passed through this compound, which is a liquid, a gas 
 is given off which can easily be shown to be hydro- 
 chloric, acid, HC1. This action continues until there is 
 no hydrogen left in combination with the carbon, but in 
 
DIFFERENT KINDS OF ACTION OF CHLORINE. 103 
 
 place of the benzene there is now a compound of the 
 formula C 6 C1 6 . This has been shown to be the final 
 product of a series of reactions represented by the fol- 
 lowing equations : 
 
 C 6 H 6 + Cl, = C.H t Cl +HC1; 
 C,H 6 C1 + Cl, = C,H,Cl a + HCl; 
 C.H.C1, + Cl, = O.H.01. + HC1 ; 
 C.H 3 C1 S + Cl, = C.H,C1 4 + HC1 ; 
 C e H,Cl. + Cl, = C 6 HC1, + HC1 ; 
 C 6 HC1 6 + Cl, = C.C1, + HC1. 
 
 In each stage one atom of chlorine is substituted for an 
 atom of hydrogen, but the hydrogen does not escape as 
 such. It combines with chlorine and passes off in the 
 form of hydrochloric acid. 
 
 (3) The third kind of action is that noticed in bleach- 
 ing, which depends upon the decomposition of water 
 and the escape of oxygen as already explained. This 
 action does not take place in the dark, but does take 
 place readily in the direct sunlight. We have, however, 
 seen that when oxygen acts upon hydrochloric acid 
 under proper conditions water is formed and chlorine 
 set free. It appears, therefore, that, under some cir- 
 cumstances, this reaction is possible : 
 
 2HC1 + O = H 2 O + Cl, ; 
 and, under other circumstances, this one : 
 H 2 O + Cl, = 2HC1 + O. 
 
 These facts appear to be contradictory* What part the 
 sunlight plays is not known, though it is well known that 
 it is capable of producing a great variety of chemical 
 changes. We shall soon see that it is only necessary to 
 allow it to act for an instant upon a mixture of hydrogen 
 and chlorine to cause them to combine with violence. 
 Then, too, the various processes known under the genera] 
 name of photography depend upon chemical changes 
 brought about by the sunlight. Leaving out of consid- 
 eration this kind of action, the decomposition of hydro- 
 chloric acid by oxygen and that of water by chlorine can 
 
104 INORGANIC CHEMISTRY. 
 
 be explained by a consideration of the heat relations, 
 The heat evoIVed in the formation of 1 molecule of water 
 in the gaseous form is 58,069 cal., while that absorbed in 
 the decomposition of 2 molecules of hydrochloric acid is 
 44,002 cal. Therefore, the reaction, 
 
 2H<21 + O = H 2 O + 01,, 
 
 is accompanied by Vn evolution of heat. It is exo- 
 thermic, and can take place without the addition of 
 energy from without. If water is formed as a liquid, 
 the heat evolved for 1 molecule is 68,357 cal., while that 
 evolved by the formation of 2 molecules of hydrochloric 
 acid in solution is 78,630 cal. Therefore, the heat 
 evolved in the formation of hydrochloric acid in solu- 
 tion is greater than that required to decompose water, 
 and this reaction takes place. This does not, however, 
 explain what part the sunlight plays in the process. 
 
 Chlorine Hydrate and Liquid Chlorine. When chlo- 
 rine gas is passed into water cooled down almost to the 
 freezing-point, crystals appear in the vessel. These con- 
 sist of chlorine and water as represented by the formula 
 Cl -f- 5H 2 O ; or, assuming that it is formed by the com- 
 bination of the molecules of chlorine with water, the 
 formula should be written Cl a + 10H 2 O. It gives off 
 chlorine at the ordinary temperature and, if allowed to 
 stand, undergoes complete decomposition into chlorine 
 and water. If gently heated the chlorine is given off 
 rapidly. This fact was taken advantage of by Faraday 
 for the purpose of subjecting the gas to high pressure 
 and low temperature. For this purpose he placed some 
 of the hydrate in a strong glass tube of the form repre- 
 sented in Fig. 6. The com- 
 pound was put in the part of 
 the tube marked db, and the 
 other end, c, then sealed. The 
 arm ab was warmed by dip- 
 ping it in warm water, while 
 the other arm was placed in 
 Fl0 - 6 - a freezing mixture. Under 
 
 these circumstances the chlorine is given off from the 
 hydrate, but being unable to escape from the tube the 
 
HYDROCHLORIC ACID. 105 
 
 pressure is increased to such an extent that at the low 
 temperature the gas assumes the liquid form. 
 
 Applications of Chlorine. Chlorine is used very exten- 
 sively in the arts, particularly for the purpose of bleach- 
 ing. It is also used for the manufacture of a large num- 
 ber of compounds which contain chlorine, the principal 
 ones being bleaching powder or calcium hypochlorite, 
 and potassium chlorate. If used in sufficient quantity 
 chlorine is an excellent disinfectant and deodorizer. 
 By far the largest quantity of the chlorine manufactured 
 is converted into bleaching powder or calcium hypochlo- 
 rite, as this can be conveniently transported, and the 
 chlorine can be obtained from it when desired. It is 
 only necessary to expose it to the air to effect a partial 
 decomposition accompanied by a liberation of chlorine ; 
 and the addition of hydrochloric or sulphuric acid causes 
 it to give it up completely, as will be shown farther on. 
 This bleaching powder is now used almost exclusively 
 instead of chlorine gas for bleaching. 
 
 HYDEOCHLOKIC ACID. 
 
 Historical. Hydrochloric acid was first prepared in 
 large quantity by Glauber in the seventeenth century, 
 and his description is not unlike those which one fre- 
 quently reads nowadays referring to some patent medi- 
 cine. The method of preparation used by him was the 
 same as that used at present, viz., the action of sulphuric 
 acid upon common salt. 
 
 Study of the Action of Hydrogen upon Chlorine. If hy- 
 drogen and chlorine are brought together in the dark no 
 action takes place, no matter how long they are allowed 
 to stand together. If, however, the mixture is put in 
 diffused sunlight, gradual combination takes place ; 
 and if the direct light of the sun is allowed to shine 
 for an instant on the mixture, explosion occurs, and this 
 is the sign of the combination of the two gases. The 
 same sudden combination is effected by applying a 
 flame or spark to the mixture, or by illuminating it in- 
 
106 INORQANIC CHEMISTRY. 
 
 stantaneously with the light from burning magnesium 
 or an electric light. On comparing these facts with 
 those learned in studying the action of hydrogen on 
 oxygen a marked difference is evident. Hydrogen and 
 oxygen do not combine either in the dark or the direct 
 sunlight, but only when a spark is brought in contact 
 with the mixture. 
 
 Another way in which hydrogen can be made to 
 combine with chlorine is by introducing a jet of burning 
 hydrogen into a vessel containing chlorine. The hydro- 
 gen will continue to burn, but the character of the flame 
 will change completely, and above the vessel white 
 fumes will be observed. This burning of hydrogen in 
 chlorine is entirely analogous to the burning of hydro- 
 gen in oxygen. It is simply an act of combination of 
 the two gases, in each case, accompanied by an evolution 
 of light and heat. And just as oxygen can be burned in 
 hydrogen by a proper arrangement of apparatus, so 
 chlorine can also be burned in hydrogen. 
 
 To determine the relation between the volumes of hy- 
 drogen and chlorine which combine with each other and 
 the volume of the product formed is more difficult than 
 in the case of hydrogen and oxygen, mainly for the rea- 
 son that chlorine acts upon mercury and is dissolved by 
 water. It is necessary to proceed indirectly. 
 
 Instead of causing hydrogen and chlorine to combine, 
 hydrochloric acid is decomposed and the volumes of the 
 hydrogen and chlorine obtained are determined. One 
 method of effecting this consists in decomposing hydro- 
 chloric acid by an electric current in an apparatus 
 like that referred to in connection with the decom- 
 position of water. As chlorine is, however, soluble 
 in water, the apparatus is filled with a saturated solution 
 of common salt to which a strong solution of hydro- 
 chloric acid in water is added. On passing a fairly 
 strong current, the hydrochloric acid is decomposed, 
 hydrogen being given off at one pole and chlorine at 
 the other. For a given volume of hydrogen the same 
 volume of chlorine is liberated, which makes it ap- 
 pear probable that hydrogen and chlorine are combined 
 
PREPARATION OF HYDROCHLORIC ACID. 107 
 
 in hydrochloric acid in the proportion of volume to 
 volume. 
 
 For the purpose of further studying the volume re- 
 lations, the following experiment is of value. A tube is 
 filled with hydrochloric acid gas. A small piece of po- 
 tassium is then introduced, when decomposition takes 
 place as represented in the equation 
 
 The gas left in the vessel is hydrogen, as can easily be 
 shown. On measuring its volume it is found to be just 
 half that of the hydrochloric acid gas decomposed. 
 Taking this fact into consideration with the fact that 
 whenever hydrochloric acid is decomposed by an electric 
 current equal volumes of hydrogen and chlorine are ob- 
 tained, it appears that in the formation of hydrochloric 
 acid gas 1 volume of hydrogen combines with 1 volume 
 of chlorine to form 2 volumes of hydrochloric acid, a 
 fact which was referred to in the chapter on the Atomic 
 Theory. The weight of the hydrogen is found to bear 
 to the weight of the hydrochloric acid the proportion 
 1 : 36.18. In other words, in 36.18 parts of hydro- 
 chloric acid there are 35.18 parts of chlorine and 1 part 
 of hydrogen. 
 
 Preparation. For the preparation of hydrochloric acid 
 in the laboratory as well as on the large scale, ordinary 
 sulphuric acid is poured upon common salt. Two reac- 
 tions may take place between these substances, depend- 
 ing largely upon the amount of sulphuric acid used. If 
 the two substances are brought together in the propor- 
 tion of the weights of their molecules or their molecular 
 weights, the principal reaction is the one represented in 
 the following equation : 
 
 NaCl + H 2 S0 4 = NaHS0 4 + HC1. 
 
 In this case the atom of sodium of the molecule of 
 sodium chloride is substituted for one atom of hydrogen 
 in the molecule of sulphuric acid, while the hydrogen 
 and chlorine unite to form hydrochloric acid. If, on the 
 other hand, the substances are brought together in the 
 
108 INORGANIC CHEMISTRY. 
 
 proportion of 2 molecules of sodium chloride and 1 
 molecule of sulphuric acid the principal reaction is the 
 following : 
 
 2NaCl + H 2 S0 4 = Na 3 SO 4 + 2HC1. 
 
 Properties. Hydrochloric acid is a colorless trans- 
 parent gas, and has a sharp penetrating taste and smell. 
 Inhaled it produces suffocation. It is extremely easily 
 soluble in water, 1 volume of water at ordinary temper- 
 atures dissolving 450 times its own volume of the gas, and 
 at 0, 500 times. The solution of the gas in water is what 
 is generally called hydrochloric acid. So great is the at- 
 traction of the gas for water that it condenses moisture 
 from the air ; hence, although the gas itself is quite color- 
 less and transparent, when it comes in contact with the air 
 dense white clouds are formed, which are not formed if it 
 is kept from contact with the air, as can easily be shown 
 by filling glass vessels with the gas. Hydrochloric acid 
 does not burn and does not support combustion. This 
 is equivalent to saying that it does not combine with 
 oxygen under ordinary circumstances, and that sub- 
 stances which combine with the oxygen of the air do not 
 combine with hydrochloric acid. On the other hand, we 
 have seen that under some circumstances oxygen does 
 act upon hydrochloric acid and cause an evolution of 
 chlorine. The gas is comparatively easily condensed to 
 the liquid form. 
 
 When a concentrated solution of hydrochloric acid in 
 water is heated, gas is given off, but if a dilute solution 
 is heated water is given off. In either case, when the 
 composition of the liquid is that represented by the 
 formula HC1 -f- 8H 2 O, it boils under the ordinary pres- 
 sure of the atmosphere unchanged. If the pressure is 
 lowered the composition of the liquid which passes over 
 in the process of distillation changes, so that it contains 
 a larger percentage of hydrochloric acid the lower the 
 pressure becomes. This fact seems to show that the 
 liquid of the composition HC1 + 8H 2 O, which boils 
 unchanged at the temperature 110 under the ordinary 
 pressure of the atmosphere is not a chemical compound. 
 
HYDROCHLORIC ACID. 109 
 
 On the other hand, it certainly does not conduct itself 
 like most ordinary solutions of gases. 
 
 There is a definite compound of hydrochloric acid 
 with water called hydrochloric acid hydrate, which has 
 the composition HC1 -|- 2H 2 O. This is formed by pass- 
 ing hydrochloric acid gas into the concentrated aqueous 
 solution cooled down to 22. Under these circum- 
 stances the hydrate separates in the form of crystals. 
 
 Commercial hydrochloric acid is a yellowish liquid, the 
 color being due to the presence of impurities, such as 
 iron and organic substances. The solution is obtained 
 in the factories in which " soda " or sodium carbonate 
 is made. This is an extremely important substance in 
 the arts. It does not occur in nature, but is manu- 
 factured from common salt. In the process most com- 
 monly used salt is first converted into sodium sulphate, 
 Na 2 SO 4 , by treating it with sulphuric acid. Hydro- 
 chloric acid is necessarily given off. When the factories 
 were first established in England, the gas was allowed to 
 escape as a waste product, but the effects produced by 
 it upon the vegetation of the surrounding country were 
 so injurious that a law was passed prohibiting the 
 manufacturers from allowing the gas to escape. It is 
 now collected by causing it to pass through towers so 
 constructed as to expose a large surface of bricks or 
 sandstone plates over which a current of cold water is 
 constantly kept flowing. This water dissolves the hy- 
 drochloric acid, and the solution collected below is 
 commercial hydrochloric acid. In this way enormous 
 quantities of the acid are produced, but its uses are nu- 
 merous and it always commands a price. 
 
 Pure hydrochloric acid is a solution of the pure gas in 
 pure water. It is colorless, and when concentrated it 
 gives off fumes when exposed to the air. The solution 
 when heated gives off a large part of the gas contained in 
 it, and by boiling it can all be evaporated. 
 
 Chemical Action of Hydrochloric Acid. If the action of 
 hydrochloric acid towards the elements should be studied 
 systematically it would be found that many of them 
 act by simply taking the place of the hydrogen, as has 
 
110 INORGANIC CHEMISTRY. 
 
 already been illustrated in the preparation of hydrogen 
 by treating hydrochloric acid with zinc when this reac- 
 tion takes place : 
 
 Zn + 2HC1 = ZnCl 2 + H 2 . 
 With iron the reaction is : 
 
 Fe + 2HC1 = FeCl 2 + H 3 ; 
 with tin : 
 
 Sn + 2HC1 = SnCl a + H 2 ; 
 with potassium : 
 
 K + HC1 ' = KOI + H, or 
 K a + 2HC1 = 2KC1 + H 2 ; 
 
 with sodium : 
 
 Na a + 2HC1 = 2NaCl + H 2 ; 
 
 with calcium : 
 
 Ca + 2HC1 = Ca01 2 + H 3 ; etc., etc. 
 
 On the other hand, there are many elements which do 
 not act in this way towards hydrochloric acid. Sul- 
 phur, nitrogen, phosphorus, carbon, and boron may be 
 taken as examples. These elements do not act upon hy- 
 drochloric acid at all. We might, therefore, divide the 
 elements into two classes : (1) those which act upon hy- 
 drochloric acid setting hydrogen free and forming chlo- 
 rides ; and (2) those which do not act upon hydrochloric 
 acid. 
 
 Again, when hydrochloric acid acts upon the oxides of 
 those elements which have the power to liberate hydro- 
 gen from it, it forms the same chlorides as are formed 
 when the element alone acts, but instead of hydrogen 
 being liberated water is formed. Thus when the acid 
 acts upon zinc oxide, ZnO, the reaction takes place thus : 
 
 ZnO + 2HC1 = ZnCl 2 + H 2 O ; 
 
CHEMICAL ACTION OF HYDROCHLORIC ACID. .111 
 with lime or calcium oxide, CaO, it is : 
 
 CaO + 2HC1 = CaCl a + H a O ; 
 with potassium oxide : 
 
 K 2 O + 2HC1 = 2KC1 + H 2 O ; etc., etc. 
 
 In these reactions there are two forces at work tending 
 to effect the change. There is, first, the affinity of the 
 element, which is combined with oxygen, for chlorine, 
 and> second, the affinity of the hydrogen for oxygen. 
 "We should, therefore, naturally expect hydrochloric 
 acid to act more readily upon the oxides than upon ele- 
 ments, and it has been found that this is the case. 
 
 If the so-called hydroxides of those elements which 
 act upon hydrochloric acid be brought in contact with 
 the acid, action takes place even more readily than with 
 the oxides, and the products are the same. Thus potas- 
 sium hydroxide and hydrochloric acid give potassium 
 chloride and water : 
 
 KOH + HC1 = KC1 + H 2 O ; 
 
 calcium hydroxide and hydrochloric acid give calcium 
 chloride and water, thus : 
 
 Ca(OH), + 2HC1 = CaCl a + 2H 2 O ; 
 
 aluminium hydroxide and hydrochloric acid give alumin- 
 ium chloride and water, thus : 
 
 Al(OH), + 3HC1 = A1C1, + 3H 2 O ; etc., etc. 
 
 There are then elements which act upon hydrochloric 
 acid liberating hydrogen and forming chlorides ; and 
 the oxides and hydroxides of these elements act upon 
 hydrochloric acid forming chlorides and water. The 
 elements which act in this way are commonly called 
 metals or, for reasons which will be discussed farther on, 
 base-forming elements. 
 
 If those elements which do not set hydrogen free from 
 
112 INORGANIC CHEMISTRY. 
 
 hydrochloric acid are treated directly with chlorine, they 
 generally combine with it to form chlorides. But these 
 chlorides differ markedly from the chlorides of the met- 
 als, especially in their conduct towards water. Two ex- 
 amples will suffice for the present. Phosphorus forms 
 a chloride of the formula PC1 3 , known as phosphorus- 
 trichloride. In contact with water it undergoes decom- 
 position according to this equation : 
 
 PC1 3 + 3H 3 = PO 3 H 3 + 3HC1 ; 
 
 so, too, the chloride of boron, BC1 3 , undergoes the same 
 kind of change : 
 
 BC1 3 + 3H 2 O = BO 3 H 3 + 3HC1. 
 
 Similarly, the other chlorides of the elements of this 
 class tend to pass into the oxides or hydroxides when 
 brought in contact with water. Those elements which 
 do not act upon hydrochloric acid setting hydrogen free 
 and forming chlorides are generally called non-metals or, 
 for reasons which will appear later, acid-forming elements. 
 The chlorides of the acid-forming elements are generally 
 decomposed by water and the corresponding oxides or 
 hydroxides are formed. In general terms, the oxide 
 of a base-forming element or of a metal is acted upon by 
 hydrochoric acid, a chloride and water being formed ; 
 and the chloride of an acid-forming element or of a non- 
 metal is acted upon by water, a hydroxide, or oxide, and 
 hydrochloric acid being formed. We shall have many 
 illustrations of the opposite chemical character of these 
 two classes of elements, and we shall see that many of 
 the most important and characteristic chemical reac- 
 tions are associated with these differences. 
 
CHAPTER IX. 
 
 COMPOUNDS OF CHLORINE WITH OXYGEN AND WITH 
 HYDROGEN AND OXYGEN. 
 
 General. As has been seen, chlorine combines very 
 readily with hydrogen, and hydrogen with oxygen, and 
 the products are stable compounds. On the other 
 hand, chlorine cannot be made to combine directly with 
 oxygen. By indirect processes they can be combined, 
 but the compounds undergo decomposition easily, yield- 
 ing back the chlorine and oxygen contained in them. 
 Before taking up the compounds of chlorine and oxygen, 
 however, it will be best to discuss, as far as may be 
 necessary, the compounds of chlorine, hydrogen, and 
 oxygen which are more easily made, and from which the 
 oxides are made. 
 
 Principal Reactions for making Compounds of Chlorine 
 with Hydrogen and Oxygen. One of the principal reac- 
 tions made use of for the preparation of compounds of 
 chlorine, oxygen, and hydrogen consists in treating po- 
 tassium hydroxide with chlorine. The strong affinity of 
 chlorine for potassium shown by the decomposition of 
 hydrochloric acid by potassium would lead us to expect 
 that when chlorine acts upon potassium hydroxide, po- 
 tassium chloride, KC1, would be formed : 
 
 KOH + 01 = KC1 + O + H. 
 
 But it also has a strong affinity for hydrogen, so that 
 hydrochloric acid would be formed as well as potassium 
 chloride : 
 
 KOH + 2C1 = KC1 + HC1 + O. 
 
 The oxygen can, however, combine with potassium 
 chloride and form compounds, KC1O, KC1O 2 , KC1O 3 , and 
 
 (113) 
 
114 INORGANIC CHEMISTRY. 
 
 KC1O 4 ; and hydrochloric acid, if formed, would com> 
 bine with potassium hydroxide, thus : 
 
 KOH + HC1 = KC1 + H 2 O. 
 
 When potassium hydroxide is treated with chlorine 
 we may therefore expect to obtain potassium chloride, 
 KC1 ; some compound containing potassium, chlorine, 
 and oxygen ; and water. Experiment has shown that 
 the action takes place as we should expect, and that the 
 compound of potassium, chlorine, and oxygen is differ- 
 ent according to the conditions of the experiment. If 
 the solution of caustic potash is warm and concentrated 
 the product is richer in oxygen than when the solution 
 is dilute and cold. With the concentrated solution the 
 reaction takes place thus : 
 
 6KOH + 3C1 2 = 5KC1 + KC1O 3 + 3H 2 O. 
 
 While five atoms of potassium appear in the form of the 
 chloride, one appears in the form of an oxygen com- 
 pound, KC1O 3 , potassium chlorate, which we have already 
 had to deal with in connection with the preparation of 
 oxygen. With the dilute solution of caustic potash the 
 reaction takes place thus : 
 
 2KOH + Cl a = KC1 + KC1O + H 2 O. 
 
 The oxygen product in this case is potassium hypocMo- 
 rite, KC1O. 
 
 Potassium chlorate, KC1O 3 , and potassium hypocMorite, 
 KC1O, bear the same relation to two compounds, HC1O 3 
 and HC1O, that potassium chloride, KC1, and sodium 
 chloride, NaCl, bear to hydrochloric acid. But we have 
 seen that hydrochloric acid can easily be obtained from 
 sodium chloride by treating it with sulphuric acid. 
 Potassium chloride undergoes the same change when 
 treated with sulphuric acid. Indeed, we shall see that 
 nearly all compounds containing sodium or potassium 
 give up these metals when treated with sulphuric acid, 
 and take up hydrogen in the place of them. 
 
CHLORIC ACID. 115 
 
 Treating potassium chloride with sulphuric acid this 
 reaction takes place : 
 
 2KC1 + H 2 S0 4 = K 2 S0 4 + 2HCL 
 
 Similarly, treating potassium chlorate with sulphuric 
 acid, this reaction takes place : 
 
 2KC1O 3 + H a SO 4 = K 8 SO 4 + 2HC1O 3 . 
 
 The compound HC1O 3 is called cUoric acid. Further, 
 when potassium hypochlorite is decomposed by sul- 
 phuric acid under proper circumstances, it undergoes 
 the same kind of decomposition : 
 
 2KC10 + H 2 SO 4 = K 2 SO 4 + 2HC1O. 
 
 The compound HC1O is called hypochlorous acid. 
 
 Chloric Acid, HC1O 3 . The preparation of chloric acid 
 from potassium chlorate is accomplished by treating 
 a water solution of the chlorate with fluosilicic acid, 
 H 2 SiF 6 , with which potassium forms an insoluble com- 
 pound. The reaction which takes place is represented 
 thus: 
 
 2KC1O 3 + H a SiF 6 = K 2 SiF 6 + 2HC1O 3 . 
 
 The compound K 2 SiF 6 is known as potassium fluosili- 
 cate. It will be observed that this reaction is of the 
 same general character .as that represented above as tak- 
 ing place between potassium chlorate and sulphuric acid. 
 The difference between the two reactions is that potas- 
 sium fluosilicate is insoluble in water, while potassium 
 sulphate is soluble. By using fluosilicic acid, therefore, 
 a solution of chloric acid is obtained free from other sub- 
 stances, provided just enough of the fluosilicic acid is 
 added to form potassium fluosilicate with all the potas- 
 sium. This solution is unstable, and if heated above 40 
 the chloric acid undergoes decomposition according to 
 the equation 
 
 4HC1O, = 01, + 3O + 2HC1O 4 + " 
 
116 INORGANIC CHEMISTRY. 
 
 Properties. Chloric acid acts upon metals in the same 
 general way that hydrochloric acid does. It gives up its 
 hydrogen and takes up metal in its place forming com- 
 pounds like potassium chlorate, KC1O 3 , sodium chlorate, 
 NaClO 3 , etc. In consequence of the ease with which it 
 gives up oxygen, it is used extensively for the prepara- 
 tion of oxygen, and for the purpose of adding oxygen to 
 other substances, or as an oxidizing agent. Potassium 
 chlorate and other compounds of similar character de- 
 rived from chloric acid are used in the manufacture of 
 fire- works.* 
 
 Hypochlorous Acid, HC1O. The formation of potas- 
 sium hypochlorite, KC1O, by treating caustic potash with 
 chlorine has been mentioned. A similar reaction is em- 
 ployed on the large scale in the manufacture of bleach- 
 ing powder or " chloride of lime." This consists in 
 treating slaked lime or calcium hydroxide with chlo- 
 rine. The action is represented thus : 
 
 2Ca (OH) 2 + 2C1 2 = Ca(ClO) 2 + CaCl 2 + 2H 2 O. 
 
 The compound Ca(ClO) 2 known as calcium hypochlorite 
 is derived from hypochlorous acid by replacing two 
 atoms of hydrogen in two molecules of the acid by one 
 atom of the bivalent metal calcium 
 
 ; ggjg gives Ca(gg) or Ca(ClO),. ., J 
 
 Just as a mixture of potassium chloride and potassium 
 hypochlorite is formed when potassium hydroxide is 
 used, so apparently a mixture of calcium chloride and 
 calcium hypochlorite is formed when calcium hydroxide 
 is used. This point will be discussed to some extent 
 under the head of Calcium Hypochlorite (which see), 
 when it will be shown that there are good reasons for 
 
 * Great care is necessary when working with potassium chlorate, as 
 with many substances it forms explosive mixtures. Treated with con- 
 centrated acids it undergoes rapid and violent decomposition. 
 
HYPOCHLOBOUS ACID. 11? 
 
 believing that the product called bleaching powder is a 
 distinct chemical compound and not a mixture of the 
 chloride and hypochlorite. For our present purpose, 
 however, it may be considered as such a mixture, for 
 under most circumstances it acts as if it were. When 
 treated with sulphuric acid or hydrochloric acid, bleach- 
 ing powder gives up hypochlorous acid first, and then 
 chlorine. The character of the action will be clear by 
 considering first the conduct of the corresponding potas- 
 sium compounds. When a mixture of potassium hypo- 
 chlorite and potassium chloride is treated with dilute 
 sulphuric acid, the hypochlorite is decomposed with 
 liberation of hypochlorous acid and formation of potas- 
 sium sulphate : 
 
 2KC1O + H 2 SO 4 = K 2 SO 4 + 2HC1O. 
 
 At the same time the sulphuric acid acts upon the chlo- 
 ride liberating hydrochloric acid : 
 
 2KC1 + H 2 SO 4 = K 2 SO 4 + 2HC1. 
 
 But when hydrochloric acid and hypochlorous acid are 
 brought together they react as represented in this 
 equation : 
 
 HC1O + HC1 = H 2 O + C1 2 . 
 
 So that the result of treating a mixture of hypochlorite 
 and chloride with sulphuric acid is the liberation of 
 chlorine : 
 
 KC1O + KC1 + H 2 SO 4 = K 2 S0 4 + H 2 O + C1 2 . 
 With bleaching powder the reaction is : 
 Ca(01O) 2 + CaCl 2 + 2H 2 SO 4 = 2CaSO 4 + 2H 2 O + 201,. 
 
 Hypochlorous acid can also be made by passing chlo- 
 rine gas into water in which mercury oxide is suspended. 
 The reaction is : 
 
 HgO + H 2 O + 2C1 2 = HgCl 2 -f 2HOCL 
 
118 INORGANIC CHEMISTRY. 
 
 The concentrated solution of hypochlorous acid has a 
 peculiar odor suggesting that of chlorine. It is the odor 
 which is familiar as that of bleaching powder or chloride 
 of lime. The acid undergoes decomposition very readily, 
 forming chlorine and a compound of chlorine and oxygen. 
 A solution of the acid bleaches about as well as chlo- 
 rine, and when bleaching powder is used for bleaching it 
 is largely the hypochlorous acid set free from the hypo- 
 chlorite which effects the desired changes. 
 
 Like chloric acid, hypochlorous acid is an excellent 
 oxidizing agent, and is used in the laboratory in this 
 capacity. 
 
 Chlorous Acid, HClOa. Although a substance of this 
 composition is not known, a number of compounds have 
 been made which are closely related to it. Such, for 
 example, are the compounds potassium chlorite, KC1O 2 , 
 silver chlorite, AgClO 2 , etc. Potassium chlorite is formed 
 when a solution of chlorine dioxide, C1O 2 , in water is 
 treated with a solution of potassium hydroxide. ' From 
 the solution of the potassium compound, the silver com- 
 pound can be made by adding a solution of silver nitrate. 
 
 Perchloric Acid, HC1O 4 . When the preparation of 
 oxygen by heating potassium chlorate was considered, 
 it was pointed out that in the first stage of the decompo- 
 sition a reaction of this kind takes place : 
 
 2KC10 3 = KC1 + KC10 4 + 2 . 
 
 The compound KC1O 4 , or potassium perchlorate, can be 
 separated from the chloride by treating the mixture with 
 cold water in which the chloride is easily soluble, while 
 the perchlorate is practically insoluble. From the per- 
 chlorate, perchloric acid can be made in the same way 
 that chloric acid is made from potassium chlorate, by 
 treating with fluosilicic acid (see preparation of chloric 
 acid). Perchloric acid is, however, much more stable in 
 concentrated solution than the other oxygen compounds 
 of chlorine, and if the perchlorate is treated with sul- 
 phuric acid, perchloric acid can be obtained from the 
 mixture by distillation. . ; i > 
 
COMPOUNDS 0V CHLORINE GENERAL. 119 
 
 Pure perchloric acid, HC1O 4 , can be obtained in the 
 form of a colorless fuming liquid. It is a dangerous sub- 
 stance to deal with, as it produces bad wounds when 
 brought in contact with the flesh, and is very unstable and 
 explosive. In contact with combustible substances in gen- 
 eral it causes explosion in consequence of the ease with 
 which it gives up oxygen and converts the combustible 
 substances into gaseous products. 
 
 A hydrate of the formula HC1O 4 + H 2 O or H 3 C1O 6 
 is known. Further, there are some facts known that 
 point to the existence of a second hydrate of the formula 
 HC10 4 + 2H,0. 
 
 General. From the above it will be seen that the com- 
 pounds of chlorine with hydrogen and oxygen form a 
 series, the members of which bear a simple relation to 
 one another. Beginning with hydrochloric acid the series 
 is as follows : 
 
 Hydrochloric acid, .... HC1 
 
 Hypochlorous acid, .... HC1O 
 
 Chlorous acid, HC1O 3 
 
 Chloric aid, HC1O 3 
 
 Perchloric acid, . . . . . HC1O 4 
 
 The successive members differ from each other by one 
 atom of oxgen, the ratio between the hydrogen and the 
 chlorine remaining the same throughout the series. 
 While the compounds differ markedly from one another 
 in many ways, they have some common features. Upon 
 metals, and their oxides and hydroxides, all the members 
 of the series act in general in the same way that hydro- 
 chloric acid does, the result being the formation of 
 products which do not contain hydrogen, but do contain 
 a metal in the place of the hydrogen. We have examples 
 of these compounds in potassium chlorate, KC1O 3 , cal- 
 cium hypochlorite, Ca(ClO) 2 , potassium chlorite, KC1O 2 , 
 and potassium perchlorate, KC1O 4 . All these compounds 
 belong to the class called salts, which will presently be 
 taken up. On the other hand, while there is a class of 
 elements upon which hydrochloric acid does not act, the 
 
120 INORGANIC CHEMISTRY. 
 
 oxygen compounds of the above series will in many 
 cases act upon these elements and convert them into 
 oxides. Thus sulphur and phosphorus, which are not 
 acted on by hydrochloric acid, are converted into oxides 
 by the oxygen compounds of chlorine. 
 
 Finally, the addition of oxygen to hydrogen and chlo- 
 rine decreases the stability of the compound. Hydro- 
 chloric acid, for example, is characterized by great 
 stability, while hypochlorous acid, HC1O, as well as all 
 the other members of the series, is characterized by 
 instability. The larger the proportion of oxygen, how- 
 ever, the greater the stability of the compound. The 
 most stable member of the series of oxygen compounds 
 is perchloric acid. Another fact that is worthy of special 
 notice is that the metal derivatives or salts of these acids 
 are more stable than the acids themselves. Many of them 
 can be heated to a comparatively high temperature without 
 undergoing decomposition. This is most marked in the 
 case of the perchlorates. It will be remembered that in 
 decomposing potassium chlorate for the purpose of mak- 
 ing oxygen the change takes place in two stages. In the 
 first, potassium perchlorate is formed. In order to de- 
 compose this, however, the temperature must be raised 
 considerably higher than that which was required to effect 
 the breaking down of the chlorate. 
 
 Compounds of Chlorine with Oxygen. The compounds 
 of chlorine with oxygen are : 
 
 Chlorine monoxide, C1 2 O, and chlorine dioxide, C1Q 2 . 
 
 The first or chlorine monoxide, C1 2 O, is formed by the 
 action of chlorine on dry mercury oxide : 
 
 HgO + 2C1 2 = HgCl a + Cl t O. 
 
 It is a gas which can easily be condensed to the liquid 
 form. The specific gravity of its vapor gives the molec- 
 ular weight corresponding to the formula C1 2 CX It is 
 extremely unstable, breaking down under the influence 
 of heat into chlorine and oxygen. With water it forms 
 hypochlorous acid, thus : 
 
COMPOUNDS OF CHLORINE WITH OXYGEN. 121 
 
 Cl a O+H a O = 2HOCl. 
 
 A substance formed by treating a mixture of arsenic 
 trioxide, As 2 O 3 , and potassium chlorate with nitric acid 
 and by other methods has been described as a greenish- 
 yellow gas which can be condensed to an extremely un- 
 stable liquid ; and it is generally referred to under the 
 name chlorine trioxide, the formula C1 2 O 3 being ascribed 
 to it. The most careful investigation of this substance 
 has, however, shown that it is not chlorine trioxide, but 
 a mixture of chlorine dioxide with varying quantities of 
 chlorine or free oxygen. 
 
 Chlorine dioxide, C1O 2 , is a greenish-yellow gas of great 
 instability. It can be condensed to a liquid which boils 
 at + 9. It is always one of the products of the action 
 of concentrated sulphuric acid upon potassium chlorate, 
 and is formed in consequence of the decomposition of 
 the chloric acid which is first set free : 
 
 2KC1O 3 + H 2 S0 4 = K 2 S0 4 + 2HC1O, ; 
 6HC10 3 = 2HC10 4 + 2H 3 O + 4ClO a 
 
 a . 
 
 When moderately dilute hydrochloric acid acts upon 
 potassium chlorate a greenish-yellow gas is formed 
 which has been called euchlorine. It is a mixture of 
 chlorine and chlorine dioxide, formed thus : 
 
 2KC10 3 + 4HC1 = 2KC1 + 2H 2 O + 201O a + C1 2 . 
 
 Combustible substances burn in chlorine dioxide with 
 violence. This action can be shown by putting a few 
 small pieces of phosphorus under water in a glass vessel, 
 and upon this a little potassium chlorate. If now a few 
 drops of concentrated sulphuric acid are added through a 
 long narrow tube or pipette, the phosphorus will be seen 
 to burn under water. This is due to the liberation of chlo- 
 rine dioxide, and the action of this compound upon the 
 phosphorus. 
 
 When a solution of chlorine dioxide in water is treated 
 with potassium hydroxide, potassium chlorite, KC1O 2 , 
 is formed (see p. 118). 
 
122 INORGANIC CHEMISTRY. 
 
 Constitution of the Compounds of Chlorine with Hydro* 
 gen and Oxygen. To determine the constitution of un- 
 stable compounds which, when they break down at all, 
 are almost completely disintegrated is difficult, and in 
 many cases impossible. Our knowledge of the constitu- 
 tion of the oxygen acids of chlorine is for this reason 
 extremely limited. Regarding the series of these com- 
 pounds and comparing their composition with that of 
 hydrochloric acid, it would appear that they are com- 
 pounds of hydrochloric acid with different amounts of 
 oxygen. But we have seen that chlorine and hydrogen 
 are univalent elements, as is shown in hydrochloric acid, 
 H-C1. If each of the atoms in the molecule of hydro- 
 chloric acid is doing all it can in holding the other atom 
 in combination, then, plainly, it is impossible for the 
 molecule to take up oxygen directly and form a com- 
 pound of the formula H-C1-O or O-H-C1. On the 
 other hand, it is possible to conceive of the atoms chlo- 
 rine, hydrogen, and oxygen as being unite'd in such a 
 way that hydrogen and chlorine shall be univalent and 
 oxygen bivalent. This is represented in the formula 
 Cl O H. Further, extensive study of compounds con- 
 taining hydrogen and oxygen has made it appear ex- 
 tremely probable that in them the hydrogen is generally 
 in combination with oxygen as represented in the above 
 formula, and as is represented also in the formulas 
 of such compounds as potassium hydroxide, K O-H, 
 
 O H 
 
 calcium hydroxide, Ca < QTT > aluminium hydroxide, 
 
 /O-H 
 
 Al^-O-H , etc. All the reasons for this cannot possibly 
 X 0-H 
 
 be made clear without a knowledge of a great many facts 
 which must be acquired gradually. As regards such 
 hydroxides as those just referred to, the views expressed 
 in the formulas given are certainly simpler than any 
 others which have been proposed, and they are not con- 
 tradicted by any known facts. Assuming then that in 
 the acids of chlorine the hydrogen is in combination 
 with oxygen, or that the compounds are hydroxides, we 
 
CONSTITUTION OF COMPOUNDS OF CHLORINE. 123 
 
 have the formulas H-O-C1, H-O-C1O, H-O-ClO a , and 
 H-O-C1O 3 for the four compounds. If, however, chlo- 
 rine is univalent the additional oxygen cannot be in 
 direct combination with the chlorine, and the only 
 way in which the constitution of these compounds can 
 be represented on the assumption that hydrogen and 
 chlorine are univalent and oxygen bivalent is this: 
 H-O-C1, H-O-O-C1, H-O-O-O-C1, and H-O-O-O-O-C1. 
 These formulas have been used for some time, but 
 strong opposition has been raised to them and they are 
 now rapidly losing ground. They represent compounds 
 in which oxygen atoms are in combination with oxygen 
 atoms. But judging by the conduct of hydrogen dioxide 
 and ozone, this kind of combination is a very unstable 
 one. The acids of chlorine are unstable enough, but the 
 one which contains the most oxygen is the most stable, 
 and this we should hardly expect if the oxygen atoms 
 are arranged as shown in the above formulas. This is 
 not a fatal objection to these formulas, but it makes 
 them, at least, appear improbable. 
 
 The view which now finds most support is based upon 
 the conception that the valence of chlorine towards oxy- 
 gen and towards oxygen and hydrogen, or towards hy- 
 droxyl as the group O-H is called, varies from univa- 
 lence to septivalence ; that it is univalent in hydrochloric 
 acid and in hypochlorous acid, trivalent in chlorous 
 acid, quinquivalent in chloric acid, and septivalent in 
 perchloric acid. This is shown in the formulas 
 
 O O 
 
 H-O-C1, H-O-C1=O, H-O-C1 and H-O-C1=O. 
 
 From a study of other similar compounds these com- 
 pounds are regarded as derived from hydroxides, thus : 
 Chlorous acid,O=Cl-O-H, is supposed to be formed from 
 
 the hydroxide Cl^ OH by loss of water ; chloric acid 
 \OH 
 
124 INORGANIC CHEMISTRY. 
 
 from the hydroxide C1(OH) 5 by loss of two molecules of 
 water : 
 
 C1(OH) 5 = C1O 2 (OH) + 2H 2 O ; 
 
 perchloric acid from the hydroxide C1(OH) 7 by the loss 
 of three molecules of water : 
 
 C1(OH) 7 = ClO 3 (pH) + 3H 3 O. 
 
 When the acids are dissolved in water it is quite prob- 
 able that in many cases the hydroxides are formed, but 
 being unstable they cannot generally be isolated. The 
 hydrate of perchloric acid, HC1O 4 + H,O(H 3 C1O 6 ), ap- 
 pears to be such a compound. It probably has the con- 
 stitution represented by the formula O 2 C1(OH) 3 . From 
 the compound C1(OH) 7 a series of compounds can be 
 derived by successive losses of one molecule of water, as 
 here shown : 
 
 C1(OH) 7 - H 9 = OC1(OH) 5 ; 
 OC1(OH) 6 - H,0 = 2 C1(OH) 3 ; and 
 O a Cl(OH) 3 - H a O = 3 C1(OH). 
 
 The second and last products are the hydrate and the 
 compound known as perchloric acid. While the evidence 
 in favor of this view presented by these compounds 
 themselves is very slight, the view is strongly supported 
 by the conduct of certain analogous derivatives of the 
 element iodine, which in many respects conducts itself 
 like chlorine. 
 
 Comparison of Chlorine and Oxygen. The power of 
 chlorine to combine with other elements is nearly as 
 great as that of oxygen. It combines with all other ele- 
 ments except fluorine, but does not form quite as great a 
 variety of compounds as oxygen. Oxygen combines with 
 several elements in three or four different proportions, 
 but chlorine rarely combines in more than two propor- 
 tions with one element. While in general chlorine and 
 oxygen conduct themselves in the same way, there is a 
 
COMPARISON OF CHLORINE AND OXYGEN. 125 
 
 very important difference between them, to which atten- 
 tion has already been called indirectly. There are some 
 elements towards which oxygen has a stronger affinity 
 than chlorine ; and there are others towards which chlo- 
 rine has a stronger affinity than oxygen. The former are 
 the non-metals or acid-forming elements ; the latter are 
 the metals or base-forming elements. The difference is 
 shown most readily by the fact that the chlorine com- 
 pounds of the acid-forming elements are converted by 
 water into oxygen compounds or hydroxides, while the 
 oxides or hydroxides of the base-forming elements are 
 converted into chlorides by the action of hydrochloric 
 acid. This distinction is not a sharp one which can 
 easily be made, for the conduct of an element is to a con- 
 siderable extent dependent upon conditions, particularly 
 of temperature, and a distinction which holds good under 
 one set of conditions may possibly not hold good under 
 another set. Still the above statements in regard to the 
 conduct of the metals and non-metals towards chlorine 
 and oxygen are in accordance with well-marked tenden- 
 cies of the two classes of elements, as will appear more 
 clearly. 
 
 Chlorine and oxygen being in general similar elements 
 we should not expect them to combine readily with each 
 other. Although they do combine indirectly in a number 
 of proportions, none of the compounds are stable. There is 
 a marked difference between these compounds and hydro- 
 chloric acid as regards the ease with which they are 
 formed, and also as regards their stability. A marked 
 difference between chlorine and oxygen is also to be 
 found in their relations to hydrogen. While one volume 
 -of chlorine combines with one of hydrogen to form two 
 volumes of hydrochloric acid gas, one volume of oxygen 
 combines with two volumes of hydrogen, the three vol- 
 umes of gas condensing to two volumes of water vapor. 
 So^also while, as we say, one atom of chlorine combines 
 with one atom of hydrogen, one atom of oxygen com- 
 bines with two atoms of hydrogen. What the difference 
 between a bivalent and a univalent element consists in is 
 
126 INORGANIC CHEMISTRY. 
 
 not known, but that the difference is something deep 
 seated appears from the marked difference in conduct 
 between chlorine and oxygen in combining with hy- 
 drogen. 
 
CHAPTER X. 
 
 ACIDS BASES NEUTRALIZATION-SALTS. 
 
 General One cannot deal with chemical phenomena 
 without constant reference to acids, and in the course of 
 our study thus far a number of substances belonging to 
 this class have been met with. It is now time to inquire 
 what features these substances have in common which 
 lead chemists to call them all acids. What is there in 
 common between the heavy, oily liquid, sulphuric acid, 
 the colorless gas, hydrochloric acid, and the unstable 
 compounds chloric and hydrochlorous acids? To un- 
 derstand the common features requires some knowl- 
 edge of a class of substances to which attention has 
 already been given. These are substances like caustic 
 potash and caustic soda, or potassium and sodium hy- 
 droxides which are called alkalies, which are the most 
 marked representatives of the class of substances known 
 as bases. These two classes, acids and bases, have 
 the power to destroy the characteristic properties of 
 each other. When an acid is brought in contact with 
 a base in proper proportions, the characteristic prop- 
 erties of both the acid and the base are destroyed. 
 They are said to neutralize each other. They form new 
 products which are said to be neutral, which means that 
 they have not the properties of an acid nor those of a 
 base. This act of neutralization is an extremely im- 
 portant one, with which we have constantly to deal in 
 chemical operations. 
 
 A Study of the Act of Neutralization. The fact hav- 
 ing been learned that acids and bases neutralize one 
 another, the next thing to do is to study the act of neu- 
 tralization as carefully as possible, and learn what chemi- 
 cal changes are involved in it. For this purpose we 
 
 (127) 
 
128 INORGANIC CHEMISTRY. 
 
 should select a number of acids and a number of basest 
 and study their action upon one another. We may take 
 sulphuric, hydrochloric, and nitric acids ; and potassium,, 
 sodium, and calcium hydroxides. We know from many 
 analyses that have been made that the composition of 
 these substances is as follows : 
 
 - + 
 
 Hydrochloric acid, HC1 
 
 Nitric acid, . . . . - HNO 3 
 
 Sulphuric acid, H 2 SO 4 
 
 Potassium hydroxide, .... KOH 
 
 Sodium hydroxide, . . . . . NaOH 
 
 Calcium hydroxide, Ca(OH) 2 
 
 The first question to be answered is whether, in order 
 to effect neutralization, definite quantities of the sub- 
 stances are necessary. To decide this, solutions of the 
 acids and of the bases should be prepared and allowed 
 to act upon one another in different proportions. But 
 how shall we determine whether the solutions we are- 
 working with are acid, basic, or neutral ? It has been 
 found that all acids have the power to change the color 
 of certain substances. For example, the dye litmus is 
 blue. If a solution which is colored blue with litmus is 
 treated with a drop or two of an acid, the color is 
 changed to red. If now the red solution is treated with 
 a few drops of a solution of a strong base, the blue color 
 is restored. There are many other substances which 
 change markedly in color by the addition of acids or 
 bases. These facts furnish a means of recognizing 
 whether a solution is acid or basic. Now, suppose that to 
 a carefully measured quantity of one of the acid solutions 
 a few drops of blue litmus is added. It will at once 
 turn red. On adding slowly a solution of one of the 
 bases the color will remain red as long as the solution is 
 acid, but the instant it is basic it will turn blue. By 
 noticing when the change in color takes place, it is pos- 
 sible to determine exactly how much of a certain basic 
 solution is required to neutralize the quantity of the acid 
 solution taken. If it is found in the case studied that 
 
NEUTRALIZATION. 129 
 
 to neutralize 20 cc. of the acid solution 30 cc. of the basic 
 solution are required, then, using the same solutions, it 
 will be found in every experiment that the same quanti- 
 ties are required to effect neutralization, or that the 
 change of color takes place whenever these proportions 
 are reached. And no matter how the quantity of one of 
 the liquids is varied, the quantity of the other required 
 for neutralization varies in the same proportion. A great 
 many experiments of this kind have been performed with 
 many different acids, and what is true in one case has 
 been found true in all. It appears, therefore, that tfie 
 act of neutralization is a definite one, loJiich takes place be- 
 tween definite quantities of acid and base; that for a certain 
 quantity of base a certain quantity of acid is required to 
 effect neutralization, and vice versa. 
 
 The next question to be answered is, What is formed 
 when the acid and base are neutralized ? To determine 
 this, larger quantities of acids should be neutralized 
 with bases, and the substance or substances- formed 
 should then be studied. If hydrochloric acid is neu- 
 tralized with sodium hydroxide a solid product, sodium 
 chloride, is formed. The action takes place according 
 to the following equation : 
 
 HC1 + NaOH = NaCl + H 2 O. 
 
 Hydrochloric acid and calcium hydroxide act thus : 
 2HC1 + Ca(OH) 2 = CaCl 2 + 2H 2 O. 
 
 Nitric acid acts upon the three bases mentioned above as 
 represented in these equations : 
 
 HNO 3 +KOH =KN0 3 + H 2 O ; 
 HNO 3 +NaOH = NaNO 3 + H 2 O ; 
 2HN0 3 + Ca(OH) 2 = Ca(NO 3 ) 2 + H 2 O. 
 
 Sulphuric acid acts upon these same bases thus : 
 
 H 2 S0 4 +2KOH = K 2 S0 4 + 2H 2 O ; 
 H 2 SO 4 +2NaOH = Na SO 4 + 2H 2 O ; 
 H,S0 4 + Ca(OH) 2 = CaS0 4 + 2H 2 O. 
 
130 INORGANIC CHEMISTRY. 
 
 The reactions which take place in these cases are typi- 
 cal of all reactions between acids and bases. One of the 
 products formed is always water, the other is a com- 
 pound which is without acid and basic properties, or 
 which is neutral and differs from the acid in that it 
 contains some other element in place of the hydrogen. 
 This other element is the one which in the base is in 
 combination with hydrogen and oxygen as a hydroxide. 
 The simplest case is that of hydrochloric acid and either 
 potassium or sodium hydroxide : 
 
 HC1 + KOH = KC1 +H 2 0. 
 
 As has already been stated (see p. Ill), we have here two 
 forces operating to bring about the change : (1) the ten- 
 dency of hydrogen to combine with hydroxyl (OH) to 
 form water; and (2) the tendency of chlorine to unite 
 with potassium. A similar statement may be made 
 in regard to every reaction between an acid and a 
 base. 
 
 General Statements. Considering the facts treated of 
 in the last paragraph, it appears : 
 
 (1) That an acid contains hydrogen ; 
 
 (2) That a base contains a metal ; 
 
 (3) That when an acid acts upon a base the hydrogen 
 and metal exchange places ; 
 
 (4) That the substance formed by substituting hydro- 
 gen for the metal of the base is water ; 
 
 (5) That the substance obtained from the acid by sub- 
 stituting a metal for the hydrogen is neither an acid nor 
 a base, but is generally neutral. 
 
 The last statement is subject to some modification, for 
 reasons which in some cases are clear but in others are 
 not apparent. It is true that in some cases after substi- 
 tuting a metal for the hydrogen the substance has an 
 alkaline reaction, and in other cases an acid reaction. 
 
 Definitions. We have already seen that hydrochloric 
 acid and sulphuric acid act upon certain metals, as iron 
 and zinc, and that the action consists in giving up hy- 
 drogen and taking up metal in its place. The products 
 
REACTION BETWEEN ACIDS AND BASES. 131 
 
 of this action are the same in character as those formed 
 by the action of acids on bases. 
 
 An acid is a substance containing hydrogen, which it 
 easily exchanges for a metal, when treated with a metal 
 itself, or with a compound of a metal, called a base. 
 
 A base is a substance containing a metal combined 
 with hydrogen and oxygen. It easily exchanges its metal 
 for hydrogen when treated with an acid. 
 
 The products of the action of an acid on a base are, 
 first, water, and, second, a neutral substance called a salt. 
 
 In the examples above cited the products KNO 3 , po- 
 tassium nitrate ; NaNO 3 , sodium nitrate ; Ca(NO 3 ) 2 , cal- 
 cium nitrate ; K Q SO 4 , potassium sulphate ; Na 2 SO 4 , sodium 
 sulphate ; CaSO 4 , calcium sulphate, are salts. The rela- 
 tions between them and the acids from which they are 
 derived will be easily recognized on comparing their 
 formulas with those of the acids. 
 
 Comparison of the Reaction between Acids and Hy- 
 droxides, and between Acids and Chlorides. The reac- 
 tion between acids and hydroxides, or, as it is generally 
 spoken of, between acids and bases, is quite similar in 
 character to that which takes place between some acids 
 and chlorides. This is illustrated by the reaction be- 
 tween sulphuric acid and sodium chloride, represented 
 by the equation 
 
 2NaCl + H 2 SO 4 = Na u SO 4 + 2HC1. 
 
 Here, as when the hydroxide is used, the acid is neutral- 
 ized and the salt, sodium sulphate, Na 2 SO 4 , is formed. 
 The other product, however, is hydrochloric acid instead 
 of water. For the sake of closer comparison the two reac- 
 tions may be written thus : 
 
 NaOH , H ) Qn _ Na ) Qn , HOH . 
 NaOH + H f SU < - Na f M<J < + HOH > 
 
 NaCl , H ) Na ) so , HC1 
 NaCl ' - H f ^ = Na f ^'< + HC1. 
 
 The two reactions are thus seen to be of the same general 
 character. That with the chloride does not take place 
 
132 INORGANIC CHEMISTRY. 
 
 as readily as that with the hydroxide, and therefore is 
 not as general. There are many acids which have not 
 the power to decompose chlorides as sulphuric acid does ; 
 whereas, in general, any acid is neutralized by any me- 
 tallic hydroxide. In some cases this reaction is an ener- 
 getic one accompanied by a great evolution of heat ; in 
 others the reaction is not at all energetic. Both acids 
 and bases differ very markedly from one another in some 
 property which is spoken of in a vague sort .of way as 
 the strength. For the present it is sufficient to recog- 
 nize that this difference is similar to the difference no- 
 ticed between elements. Hydrogen and chlorine, for 
 example, differ markedly in their power to act upon 
 other substances, and chlorine is spoken of as the more 
 energetic or active element. 
 
 Other Similar Reactions. There are many other reac- 
 tions like those which take place between acids and 
 chlorides, and between acids and hydroxides. Another 
 example is furnished by the sulphides and hydrosuphides, 
 which are compounds that in some respects resemble 
 oxides and hydroxides. The reactions which take place 
 between the sulphur compounds and acids, and between 
 the oxygen compounds and acids, are entirely analogous, 
 as shown in the following equations : 
 
 K 2 S + 2HC1 = 2KC1 + H 2 S ; 
 K 2 O +2HC1 =:2KC1 +H 2 O; 
 CaS + H 2 S0 4 = CaS0 4 + H 2 S ; 
 CaO + H 2 S0 4 = CaS0 4 + H 2 O ; 
 KSH +HC1 =KC1 + H 2 S; 
 KOH +HC1 =KC1 +H 2 O; 
 Ca(SH) 2 + H 2 S0 4 = CaS0 4 + 2H 2 S ; 
 Ca(OH) 2 + H 2 S0 4 = CaS0 4 + 2H 2 O. 
 
 The product formed in place of water is the correspond- 
 ing compound of sulphur, H 2 S. It will be observed that 
 the hydrosulphides, or compounds which have the general 
 composition MSH, neutralize the acids in the same sense 
 that the hydroxides do. If hydroxides were not known, 
 our conceptions of acids might easily be based upon the 
 
DISTINCTION BETWEEN ACIDS AND BASES. 133 
 
 relations of compounds to the hydrosulphides, and the 
 substances now classed with the acids would be classed 
 with them upon this basis. As we go on we shall see that 
 there are other reactions of the same general character. 
 
 Distinction between Acids and Bases. Although there 
 is no difficulty in distinguishing between most acids and 
 most bases, there are some compounds which act some- 
 times in one way and sometimes in the other. Sulphuric 
 acid, nitric acid, and hydrochloric acid always act as 
 acids, and sodium and potassium hydroxides always act 
 as bases, but some substances which are generally basic 
 will under some circumstances act as acids, and some 
 which act as acids will occasionally act as bases. What 
 is the standard? How shall we tell whether a sub- 
 stance is an acid or a base ? We may take a pronounced 
 acid, such as hydrochloric acid, and say that any hy- 
 droxide which has the power to neutralize this acid and 
 form with it a salt shall be called a base ; and in the 
 same way we may take a pronounced base, like potas- 
 sium hydroxide, and say that any hydroxide which has 
 the power to neutralize this shall be called an acid. 
 Having made the division in this way, it would be found 
 that a few substances would be included in both lists, or, 
 in other words, some substances which are basic toward 
 hydrochloric acid are acid toward potassium hydroxide. 
 As an example, we may take aluminium hydroxide, 
 A1(OH) 3 . This neutralizes hydrochoric acid and forms 
 aluminium chloride according to the equation 
 
 A1(OH) 3 + 3HC1 = A1G1, + 3H 2 O. 
 
 But it also neutralizes potassium hydroxide according to 
 the equation 
 
 A1(OH) 3 + 3KOH = A1(OK) 3 + 3H 2 O. 
 
 It may be said in regard to this case, as in regard to 
 most other cases of the kind, that the hydroxide in ques- 
 tion is basic toward nearly all substances toward which 
 potassium hydroxide is basic ; whereas it is acid toward 
 only three or four of the most energetic bases. Bearing 
 
134 INORGANIC CHEMISTRY. 
 
 in mind, then, the fact that there are some exceptional 
 cases, it may be said that the distinction between acids 
 and bases is easily recognized. 
 
 Metals or Base-forming Elements. The question, What 
 is a metal ? may fairly be asked. But unfortunately it is 
 by no means an easy matter to give a satisfactory answer 
 to the question. We can give examples of metals, such 
 as iron, zinc, silver, calcium, magnesium, etc. ; but when 
 we attempt to find the distinguishing features of these 
 substances we are somewhat at a loss to state them. In 
 general, it may be said that to the chemist any element 
 is a metal which with hydrogen and oxygen forms a base, 
 or a product which has the power to neutralize acids. 
 In general, any element which has the power to enter 
 into an acid in the place of the hydrogen is called a 
 metal, or is said to have metallic properties. This is the 
 sense in which the word metal is used in this book. A 
 better, though a longer, name for the metals is base-form- 
 ing elements. 
 
 Constitution of Acids and Bases. As has been pointed 
 out, the bases are hydroxides, and these hydroxides are 
 regarded as derived from water by the replacement of 
 the hydrogen by metals. Examples of the hydroxides 
 of univalent, bivalent, and trivalent metals were given 
 in a previous chapter (see pp. 83-84). Similarly, the 
 acids which contain oxygen are regarded as hydroxides, 
 or as derived from water, as was stated when the sub- 
 ject of the constitution of the acids of chlorine was under 
 consideration. This view is illustrated by the following 
 formulas of some of the more common acids : 
 
 Nitric acid, (HO)NO, 
 
 Sulphuric acid, (HO) 2 SO 2 
 
 Phosphoric acid, . . * , .- . (HO) 3 PO 
 
 Carbonic acid, .... . . (HO) 2 CO 
 
 Metaphosphoric acid, .... (HO)PO 2 
 
 Nitrous acid, . (HO)NO 
 
 Arsenious acid, (HO) 3 As 
 
 Hypochlorous acid, .... (HO)C1 
 
 Perchloric acid, (HO)CIO, 
 
CONSTITUTION OF ACIDS AND BASES. 135 
 
 There are three classes of acids represented in this 
 list: (1) those with one, (2) those with two, and (3) 
 those with three atoms of hydrogen in the molecule. Or, 
 considering the compounds as hydroxides, these classes 
 are : (1) those derived from one molecule, (2) those de- 
 rived from two molecules, and (3) those derived from three 
 molecules of water by replacement of half the hydrogen 
 by something else. It is interesting to observe, also, that 
 this something which replaces the hydrogen is in most 
 cases an element in combination with oxygen or, if it is 
 not in combination with oxygen, it has the power to take 
 up more oxygen. Thus hypochlorous and arsenious 
 acids are regarded as derived from water by the replace- 
 ment of hydrogen in water by chlorine and arsenic as 
 
 shown thus : H-O-C1 and H-O-)As. But, in each case, 
 
 H-0/ 
 
 the element which is in combination with hydroxyl has 
 the power to combine with oxygen. Hypochlorous acid 
 forms the products (HO)CIO, (HO)C1O 2 , and (HO)C1O 3 , 
 while arsenious acid forms arsenic acid (HO) 3 AsO. 
 
 We may consider water as forming the connecting link 
 between the oxygen acids and bases. If A stands for any 
 acid-forming element, and B for any base-forming ele- 
 ment, then the general formula of a base is B(OH), and 
 that of an oxygen acid A(OH) or O X A(OH), in which O x 
 stands for some number of oxygen atoms from one to 
 three or four. We should then have these relations : 
 
 Water. Acids. 
 
 I. B'(OH) HOH (O X A)'(OH) 
 
 II.B"(OH), ggg (O x Ar(OH) 2 
 
 HOH 
 
 III. B /// (OH) 3 HOH (O X A) /// (OH) 3 
 HOH 
 
 In these general formulas B" means any bivalent metal, 
 and B //7 .any trivalent metal; and (O X A)" means any 
 
136 INORGANIC CHEMISTRY. 
 
 group of atoms which has the power to hold two hydroxyl 
 groups in combination, and is therefore bivalent like (O 9 S), 
 and (O X A)'" means a trivalent group like (OAs). 
 
 Constitution of Salts. The view held in regard to the 
 constitution of salts is based directly upon those held 
 in regard to the constitution of acids and bases. It 
 is believed that when an oxygen acid acts upon a base 
 the action takes place as represented in the following 
 equation : 
 
 O 2 N-O-|H + H-0|-K = O 3 N-O-K + H-O-H ; 
 or in this : 
 
 2 N-|0-H + H|-O-K = O 2 N-O-K + H-O-H. 
 
 In either case the salt formed appears as the acid, the 
 hydrogen of which has been replaced by the metal. 
 "Whether the hydroxyl of the base unites with the hydro- 
 gen of the acid, or the hydroxyl of the acid unites with 
 the hydrogen of the base, cannot be determined ; and, 
 as far as the constitution of the salt is concerned, it 
 evidently makes no difference. The case stands thus : 
 For reasons partly pointed out above, the bases are 
 regarded as hydroxides ; for similar reasons the acids are 
 also regarded as hydroxides. Now, when an acid acts 
 upon a base water and a product which differs from the 
 acid in having the metal of the base in place of its hydro- 
 gen are formed. The simplest interpretation of this 
 reaction is that given above. A case in which there 
 appears to be no room for doubt as to what takes place 
 is that of hydrochloric acid and a simple base like sodium 
 hydroxide : 
 
 Cl-H + H-O-Na = CINa + H-O-H. 
 
 It is highly probable that the reaction between acids and 
 bases is always of this character. 
 
 Basicity of Acids. In working with acids and bases it 
 is noticed that some acids have the power to form but 
 
BASICITY OF ACIDS. 137 
 
 one salt with a base like potassium hydroxide, while 
 others have the power to form two or more salts with 
 such a base. Thus, for example, hydrochloric acid, HC1, 
 and nitric acid, HNO 3 , can form but one salt with potas- 
 sium hydroxide, and the reactions are represented in the 
 following equations : 
 
 KOH + HC1 = KC1 + H 2 O ; 
 KOH + HNO 3 = KN0 3 + H 2 O. 
 
 If only half the quantity of base which is required to 
 neutralize the acid is added, half the acid remains un- 
 changed, and on evaporating the solution the excess of 
 acid will pass off. So also, if only half the quantity of 
 acid which is required to neutralize the base is added, 
 half the base will remain unchanged. On the other hand, 
 if an acid like sulphuric acid is taken, it is found that this 
 has the power to form two distinct salts with potassium 
 hydroxide, in one of which there is twice as much of the 
 metal as in the other. The reactions are represented 
 thus : 
 
 KOH + H 2 SO 4 = KHSO 4 + H 2 O ; 
 2KOH + H 2 SO 4 = K 2 S0 4 + H 2 O. 
 
 If to a given quantity of sulphuric acid only half the 
 quantity of potassium hydroxide which is required to 
 neutralize it is added, the first reaction takes place ; but 
 if the act of neutralization is complete the second reac- 
 tion takes place. An acid of this kind can, further, form 
 one salt with two bases, in which one of the hydrogen 
 atoms of the acid is replaced by one metal and the other 
 by a second metal. 
 
 The different properties of the two kinds of acids re- 
 ferred to are ascribed to differences in constitution. .In 
 the molecule of hydrochloric acid, as in that of nitric 
 acid, there is but one atom of hydrogen according to 
 the views at present held. If, therefore, the act of neu- 
 tralization takes place in each molecule it is complete, 
 and the salt is said to be a neutral or normal salt. In 
 sulphuric acid, however, there are two atoms of hydrogen 
 
138 INORGANIC CHEMISTRY. 
 
 in each molecule, and either one or both of these may be 
 replaced. If only one is replaced, a salt of the general 
 formula MHSO 4 is obtained. This is still an acid, while 
 also partly a salt. It is in fact an acid salt or a salt acid. 
 
 Acids like hydrochloric and nitric acids have not the 
 power to form acid salts. They are called monobasic 
 acids. While acids like sulphuric acid, which can form 
 two salts with one base, one of which is acid, are called 
 dibasic acids. 
 
 Monobasic acids are those which contain but one re- 
 placeable hydrogen atom in the molecule. Dibasic acids 
 are those which contain two replaceable hydrogen atoms 
 in the molecule. 
 
 Similarly, there are tribasic acids, like phosphoric acid, 
 H 3 PO 4 , arsenic acid, H 3 AsO 4 , etc. ; tetrdbasic acids, like 
 pyrophosphoric acid, H 4 P 2 O 7 ; pentdbasic acids, like peri- 
 odic acid, H 5 IO 6 ; etc., etc. The higher the basicity of 
 the acid the greater the variety of salts it can yield. 
 
 Acidity of Bases. Just as we speak of monobasic, 
 dibasic, tribasic acids, etc., so we distinguish between 
 bases of different acidity. Thus there are the monacid 
 bases, like potassium and sodium hydroxides, KOH and 
 NaOH ; diacid bases, like calcium and barium hydoxides, 
 Ca(OH) 2 and Ba(OH) 2 ; triacid bases, like aluminium and 
 ferric hydoxides, Al(OH) 3 -and Fe(OH) 3 ; etc., etc. 
 
 If a monobasic acid acts upon a monacid base, one 
 molecule of one forms a salt with one molecule of the 
 other, and, in general, no other reaction between the two 
 is possicl-j. If a monobasic acid acts upon a diacid base 
 two reactions are possible, just as when a monacid base 
 acts upon a dibasic acid. Thus, when, for example, hy- 
 drochloric acid acts upon zinc hydroxide, Zn(OH) 2 , two 
 reactions are possible : 
 
 Zn(OH) 2 + HC1 = Zn < g^ + H 2 O; 
 Zn(OH) 2 + 2HC1 = ZnCl 2 + 2H 2 O. 
 
 The compound ZnCl(OH) is still basic, just as the salt 
 KHSO 4 is still acid, and it is called a basic salt. Simi- 
 
SALTS. 139 
 
 larly, a triacid base can form three salts with a monobasic 
 acid as, for example, in the case of bismuth hydroxide 
 and nitric acid, in which three reactions are possible : 
 
 ( OH ( N0 3 
 
 Bi J OH + HNO 3 = Bi^ OH + H,O; 
 OH OH 
 
 Bi j 
 
 OH ( N0 3 
 
 OH + 2HN0 3 = -Bi\ N0 3 + 2H 2 O; 
 
 OH (OH 
 
 OH ( N0 3 
 
 , OH + 3HN0 3 = Bi-| N0 3 + 3H 2 O. 
 ( OH ( N0 3 
 
 The salts Bi j fojj\ and Bi j Q^ 2 are basic salts or 
 
 basic nitrates of bismuth, while the salt Bi(NO 3 ) 3 is the 
 neutral or normal salt. 
 
 Salts. From the above it appears that there are three 
 classes of salts : (1) Normal salts, which are derived from 
 the acids by replacement of all the acid hydrogen atoms 
 by metal atoms ; (2) Add salts, which are derived from 
 the acids by replacement of part of the hydrogen by 
 metal atoms ; and (3) Basic salts, which are derived from 
 the bases by neutralization of part of the basic hydroxyl 
 by acids. Normal salts are generally neutral ; or, if by 
 a neutral substance is meant one which has not the power 
 to form salts with acids nor with bases, then the expres- 
 sion normal salt is synonymous with neutral salt. But, 
 strange to say, some normal salts have what is called an 
 acid reaction, and others have an alkaline or basic reac- 
 tion. Thus a normal salt of a weak acid with a strong 
 base as sodium carbonate, Na 2 CO 3 , has an alkaline reac- 
 tion. So also a normal salt of a strong acid with a weak 
 base may have an acid reaction, as in the case of copper 
 sulphate, CuSO 4 . As generally used, the expression neu- 
 tral salt means a salt which exhibits neither an acid nor 
 an alkaline reaction. 
 
 In naming acid salts various methods are adopted. In 
 the case of a dibasic acid, the only distinction necessary 
 is between the acid and the normal salts. The expres- 
 
140 INORGANIC CHEMISTRY. 
 
 sions acid potassium sulphate and normal potassium sul- 
 phate mean, of course, the salts which have the formulas 
 KHSO 4 and K 2 SO 4 , and there is no danger of confusion. 
 We may, however, use the names mono-potassium sul- 
 phate and di-potassium sulphate, or primary and secondary 
 potassium sulphates. The last names are convenient and 
 readily convey to the mind the nature of the salt spoken 
 of. Just as dibasic acids yield primary and second- 
 ary salts, so tribasic acids yield primary, secondary, and 
 tertiary salts. For example, phosphoric acid yields three 
 classes of salts : primary phosphates, of the general formula 
 MH 2 PO 4 ; secondary phosphates, of the general formula 
 M 2 HPO 4 ; and tertiary phosphates, of the general formula 
 M 3 PO 4 . The phosphates of the first two classes are called, 
 in general, acid phosphates. The tertiary phosphate is 
 identical with the normal phosphate. In naming basic 
 salts there is no difficulty in the simplest cases. Thus, tak- 
 ing the three bismuth nitrates the formulas of which are 
 
 given above, the one of the formula Bi j /QTT\ is called 
 
 the mono-nitrate ; that of the formula Bi < XTT *' a , the di- 
 
 nitrate; and that of the formula Bi(NO 3 ) 3 , the tri-nitrate 
 or normal nitrate. 
 
 There are many cases which are much more complicated 
 than any of those referred to above. Thus, there are 
 basic salts formed by dibasic acids and diacid bases, by 
 dibasic acids and triacid bases, etc. There is, for ex- 
 ample, a basic copper carbonate formed by the partial 
 neutralization of two molecules of copper hydroxide, 
 Cu(OH) 2 , by one molecule of carbonic acid, CO(OH) 2 . 
 The relations will be seen by the aid of the following 
 equation, in which the structural formulas of copper hy- 
 droxide and of carbonic acid are used : 
 
 OH , HO, ^s\ ^^-004 
 
 OH Cu< OH 
 
 The salt is basic. 
 
ACID PROPERTIES AND OXYGEN. 141 
 
 Acid Properties and Oxygen. Almost all those sub- 
 stances which are called acids contain oxygen, as, for 
 example, nitric acid, HNO 3 ; sulphuric acid, H 2 SO 4 ; phos- 
 phoric acid, H 3 PO 4 ; silicic acid, H 2 SiO 3 ; carbonic acid, 
 H 2 CO 3 ; boric acid, H 3 BO 3 ; etc. The presence of oxygen 
 in acids was recognized by Lavoisier. As he showed 
 its presence in acids to be general, and as he found 
 that several elements and some compounds are con- 
 verted into acids by combination with oxygen, he con- 
 cluded that this element is an essential constituent of all 
 acids, and therefore called it oxygen, a name which, as 
 already stated (see p. 28), means the acid-former. Ac- 
 cording to Lavoisier, hydrochloric acid, like other acids, 
 contained oxygen, and this view prevailed for many years. 
 As has been pointed out under the head of Chlorine, 
 many investigations were undertaken with the object of 
 determining whether this element does or does not con- 
 tain oxygen, the result being to show that in chlorine, 
 and consequently in hydrochloric acid, there is no oxygen. 
 Several acids are now known which are like hydrochloric 
 acid in this respect, but the latter is the best known ex- 
 ample. Similar compounds are hydrobromic acid, HBr; 
 hydriodic acid, HI ; and hydrocyanic acid, HON. The 
 number of these acids is, however, quite small, and it is 
 undoubtedly true that, of the compounds which we com- 
 monly call acids, by far the larger number contain 
 oxygen as an essential constituent. Further, some com- 
 pounds which are basic can be converted into acids by 
 introducing oxygen into them. 
 
 On the other hand, there are many compounds which 
 do not contain oxygen which exhibit reactions entirely 
 analogous to those of the acids. There are for example 
 compounds containing sulphur which combine with sul- 
 phides, and others containing chlorine which combine 
 with chlorides in much the same way that the oxygen 
 acids combine with oxides, and the compounds formed 
 are analogous to ordinary salts, only they contain 
 sulphur or chlorine in place of oxygen. Thus, there is 
 a compound of arsenic and sulphur of the composition 
 H 3 AsS 4 , known as sulpharsenic acid, which is analo- 
 
14:2 INORGANIC CHEMISTRY. 
 
 gous to the oxygen compound arsenic acid, H 3 AsO 4 . When 
 arsenic acid is treated with potassium hydroxide, KOH, 
 this reaction takes place : 
 
 H 3 As0 4 + 3KOH = K 3 As0 4 + 3H 2 O. 
 
 So, too, when sulpharsenic acid is treated with potassium 
 hydrosulphide this reaction takes place : 
 
 H 3 AsS 4 + 3KSH = K 3 AsS 4 + 3H 2 S. 
 
 As many such sulphur compounds are decomposed by 
 water yielding the corresponding oxygen compounds, and 
 as most such reactions must be studied in solution in water, 
 a good reason for the fact that they are not as numerous 
 as the oxygen acids will be seen. 
 
 Just as sulphur acids act upon sulphur bases to form 
 sulphur salts, so there are what may be called chlorine 
 acids which act upon chlorine bases to form chlorine 
 salts. For example, there is a compound, H 2 PtCl 6 , known 
 as chlorplatinic acid, which with chlorides forms well- 
 marked salts : 
 
 H 2 PtCl 6 + 2KC1 = K 2 PtCl 6 + 2HC1. 
 
 The product formed in the reaction represented by this 
 equation is known as potassium chlorplatinate. The 
 reaction is analogous to the following, in which oxygen 
 compounds take part : 
 
 H 2 Pt0 3 + 2KOH = K 2 Pt0 3 + H 2 0. 
 
 In the chlorine compounds two atoms of the univalent 
 element chlorine take the place of each atom of the biva- 
 lent oxygen. Many such compounds are known ; but in 
 working with them the same difficulty arises that was 
 referred to above in speaking of the sulphur compounds ; 
 many of the chlorides which are capable of forming 
 chlorine salts are decomposed by water and converted 
 into oxygen acids. Therefore, if we start with a chlorine 
 acid and work in water solution the probability is that 
 
NOMENCLATURE OF ACIDS. 143 
 
 the product obtained will be an oxygen compound. The 
 fact that the oxygen acids are the most prominent is 
 partly to be ascribed to the fact that water is in such 
 general use as a solvent. The analogous solvent for 
 the sulphur compounds would be liquid hydrogen sul- 
 phide, H 2 S, but at ordinary temperatures this is a gas, 
 and it is, therefore, impossible to work with the sulphur 
 compounds under conditions analogous to those under 
 which we work with the oxygen compounds. The same 
 statement applies to the chlorine compounds for which 
 the analogous solvent would be liquid hydrochloric acid, 
 HC1, not the solution of the gas in water. 
 
 Nomenclature of Acids. The names of the acids of 
 chlorine illustrate some of the principles of nomenclature 
 in use in chemistry. The acid of the series which is best 
 known is called chloric acid. In naming acids the suffix 
 ic is always used in naming the principal member of a 
 group of acids containing the same elements. This is 
 seen in the names hydrochloric, sulphuric, nitric, phos- 
 phoric, silicic, carbonic, acetic, etc. If there are two 
 acids containing the same elements, that one of the two 
 which contains the smaller proportion of oxygen is given 
 a name ending in ous. Thus we have the two series : 
 
 Chloric acid, . . HC1O 3 Chlorous acid, . . HC1O 2 
 
 Sulphuric acid, . H 2 SO 4 Sulphurous acid, . H 2 SO 3 
 
 Nitric acid, . . . HNO 3 Nitrous acid, . . HNO 2 
 
 Phosphoric acid, . H 3 PO 4 Phosphorous acid, H 3 PO 3 
 
 For most cases which present themselves this method 
 of naming will suffice, but in others the number of acids 
 known is larger than two, as, for example, in the series 
 of chlorine acids. In such cases recourse is had to 
 prefixes. If there is an acid known containing a smaller 
 proportion of oxygen than the one whose name ends in 
 ous, it is generally designated by means of the prefix 
 hypo, which is derived from the Greek vno, signifying 
 tinder. Thus there are the following examples' : Hy- 
 pochlorous acid, HC1O ; hyposulphurous acid, H 2 SO Q ; 
 hyponitrous acid, H 2 N 2 O 2 ; and hypophosphorous acid, 
 
144 INORGANIC CHEMISTRY. 
 
 H 3 PO 2 . It will be seen on comparing the formulas of 
 these acids with those above given that they differ from 
 them in a very simple way. 
 
 In the series of chlorine acids there is one which con- 
 tains a larger proportion of oxygen than chloric acid. It 
 is called perchloric acid, the Latin prefix per signifying 
 here very or fully. Similarly there is a perbromic acid 
 and a permanganic -acid. Other cases arise, but they are 
 of a more or less special character, and the compounds 
 are given special names according to circumstances. 
 
 Nomenclature of Bases. As pointed out above, a base 
 is a compound of a metal with hydrogen and oxygen. 
 The bases are commonly known as hydroxides ; and in 
 order to distinguish between the hydroxides of the differ- 
 ent metals, the names of the metals are put before the 
 name hydroxide, as in naming the oxides and chlorides. 
 Thus, as has been seen, caustic soda, NaOH, is called 
 sodium hydroxide, etc. It is necessary in some cases to- 
 distinguish between two hydroxides of the same metal. 
 This is done by using the suffixes ous and ic in the same 
 sense as they are used in naming oxides and chlorides. 
 Thus ferric hydroxide has the composition Fe(OH) 3 , and 
 ferrous hydroxide the composition Fe(OH) 2 ; cuprous hy- 
 droxide is Cu(OH), and cupric hydroxide Cu(OH) 2 , etc. 
 These compounds are sometimes called hydrates, and 
 there are some good reasons for using this name, as 
 will be more fully shown in a later paragraph. On the 
 other hand, compounds in which water as such is re*- 
 garded as present are called hydrates, and there is- 
 danger of confusion if the same name is used to desig- 
 nate what are believed to be two entirely different classes 
 of compounds. As examples of hydrates we have salts 
 with their water of crystallization, chlorine hydrate, 
 C1 9 + 8H 2 O ; hydrochloric acid hydrate, HC1 + 2H 2 O ; 
 etc. While some of the compounds which are commonly 
 regarded as hydrates should probably be classed with the 
 hydroxides, there seem to be two classes, and it is there- 
 fore desirable to have two names. 
 
 Nomenclature of Salts. Theoretically every metal caji 
 yield a salt with every acid. The salts derived from a 
 
NOMENCLATURE OF SALTS. 145 
 
 given acid receive a general name, and this general 
 name is qualified in each case by the name of the metal 
 contained in the salt. Thus, all the salts derived from 
 nitric acid are called nitrates ; all the salts derived from 
 chloric acid are called chlorates ; the salts of sulphuric 
 acid are called sulphates ; * the salts of phosphoric acid 
 are called phosphates; * etc. So too, further, the salts of 
 chlorous acid are called chlorites; those of nitrous acid, 
 nitrites ; those of sulphurous acid, sulphites ; etc., etc. 
 It will be noticed that the final syllable of the name 
 of the salt differs according to the name of the acid. If 
 the name of the acid ends in ic, the name of the salt de- 
 rived from it ends in ate. If the name of the acid ends 
 in ous, the name of the salt ends in ite. To dis- 
 tinguish between the different salts of the same acid, 
 the name of the metal contained in it is prefixed. 
 Thus, the potassium salt of nitric acid is called po- 
 tassium nitrate, the sodium salt is called sodium ni- 
 trate ; the calcium salt of sulphuric acid is called calcium 
 sulphate ; the magnesium salt of nitrous acid is magne- 
 sium nitrite ; the calcium salt of hypochlorous acid is 
 calcium hypochlorite ; etc., etc. If a metal forms two 
 salts with the same acid in one of which the valence of 
 the metal is lower than in the other, the one in which 
 the valence of the metal is lower is designated by 
 means of the suffix ous, while the one in which the 
 valence of the metal is higher is designated by means of 
 the suffix ic. Thus there are two series of salts of iron 
 which correspond to the two chlorides FeCl 2 and Fed,. 
 In one series the iron appears to be bivalent, in the 
 other trivalent. Examples are, Fe(NO 3 ) 2 and Fe(NO 3 ) 3 ; 
 FeSO 4 and Fe 2 (SO 4 ) 8 ; etc. Those salts in which the 
 iron is bivalent are called ferrous salts, as ferrous 
 nitrate, ferrous sulphate, etc. ; and those in which it is 
 trivalent are called ferric salts, as ferric nitrate, ferric 
 sulphate, etc. Similarly there are two series of copper 
 
 * Strictly speaking, the salts of sulphuric acid should be called sul- 
 phurates, and those of phosphoric acid phosphorates, but for the sake of 
 euphony and convenience these names are shortened to the above forms. 
 
14:6 INORGANIC CHEMISTRY. 
 
 salts known as cuprous and cupric salts ; and two series 
 of mercury salts known as mercurous and mercuric salts. 
 If the salts of hydrochloric acid were named in ac- 
 cordance with the principle just explained, they would 
 be called hydrocMorates, and this name is sometimes used 
 for complex salts, but in the case of the salts of the 
 metals it will be observed that these are identical with 
 the products formed by direct combination of the metals 
 with chlorine. Thus, hydrochloric acid and zinc act as 
 represented in the equation 
 
 Zn + 2HC1 = ZnCl, + H 2 ; 
 while zinc and chlorine act thus : 
 Zn + 01, = ZnCl,. 
 
 In each case the same product, ZnCl 2 , is formed. But 
 these compounds of metals with chlorine are called chlo- 
 rides, as has already been explained. Hence for these 
 cases the name hydrocJdorate is unnecessary. 
 
 The name hydrate to which reference was made in a 
 paragraph above suggests a salt of hydric acid. Potas- 
 sium hydrate signifies the potassium salt of this acid or 
 of water. In one sense this is a proper name for the 
 compound. It is water in which a part of the hydrogen 
 is replaced by a metal, and it is in this respect like a 
 salt. While, however, there is an unmistakable analogy 
 between the formation of a metallic hydroxide from 
 water and that of a salt from an acid, it appears, on the 
 whole, wise not to class water with the acids nor with 
 the bases, but rather to regard it as the connecting link 
 between the two classes. We shall see later that the 
 similar compounds hydrogen sulphide, H 2 S, and hy- 
 drogen selenide, H 2 Se, have much more marked acid 
 properties than water. When treated with metallic hy- 
 droxides they form salts of the general formulas M,S 
 and M a Se. 
 
CHAPTER XI. 
 
 NATURAL CLASSIFICATION OF THE ELEMENTS THE 
 PERIODIC LAW. 
 
 Historical. It has long been known that simple rela- 
 tions exist between the atomic weights of some elements 
 which resemble one another closely. Thus chlorine, 
 bromine, and iodine are very similar elements. Their 
 atomic weights are 35.18, 79.34, and 125.89 respectively. 
 It will be seen that the atomic weight of bromine, 79.34, 
 is approximately the mean of those of chlorine and 
 iodine. We have 
 
 35.18 + 125.89 = Q() 53 
 
 2 
 
 A similar group is that of sulphur, selenium, and tellu- 
 rium, which resemble one another as closely as chlorine, 
 bromine, and iodine do. The atomic weights are S = 
 31.83, Se = 78.42, and Te = 126.52. W T e have here 
 
 I *u* + i*.n , 
 
 Other groups are those of phosphorus, 30.79, vanadi- 
 um, 50.99, and arsenic, 74.44 : 
 
 30.79 + 7444 = g261; 
 
 lithium, 6.97, sodium, 22.82, and potassium, 38.82 : 
 ..." 6.97 + 38.82 = 22m 
 
 a 
 
 (147) 
 
148 INORGANIC CHEMISTRY. 
 
 In 1863-64 J. A. E. Newlands called attention to the 
 fact that if all the elements are arranged in a table in the 
 order of their atomic weights, beginning with that one 
 which has the lowest atomic weight and ending with that 
 one which has the highest atomic weight, provided they 
 are arranged horizontally in groups of seven, placing the 
 eighth under the first, the ninth under the second, etc., 
 then similar elements would fall in the same perpen- 
 dicular line. Newlands' arrangement was quite imper- 
 fect, and it required considerable modification in order to 
 make it appear at all satisfactory. In 1869 and 1870 two 
 papers appeared, one by D. Mendeleeff and the other by 
 Lothar Meyer, in which these relations are treated in a 
 masterly manner, and it was then seen that one of the 
 most important laws of chemistry had been discovered. 
 Everything learned since then has only made it appear 
 more and more certain that the law which is known as 
 the periodic laiv is a fundamental law of chemistry. 
 
 Arrangement of the Elements. Mendeleeff and Lothar 
 Meyer have proposed several arrangements for the pur- 
 pose of making clear the connection between the proper- 
 ties and atomic weights of the elements. Those which 
 have proved most useful will first be given, and then the 
 connection between the atomic weights and properties 
 will be discussed briefly. The different arrangements 
 are to be regarded only as different ways of expressing 
 the same law, and no one of them is perfect. The inves- 
 tigation of the relations between the atomic weights and 
 the properties of the elements has not yet been pushed 
 far enough to justify a final opinion as to the character 
 of the relations, but it has nevertheless reached a stage 
 in which we are justified in stating that these relations 
 are general and deep-seated. 
 
MENDELEEFF'S TABLES OF THE ELEMENTS. 149 
 
 
 
 TO 
 
 oo'oi 
 
 
 QO't^ 
 
 
 
 
 WTO 
 
 O O 
 
 
 og 
 
 . 
 
 tJ 
 
 
 "I 
 
 
 
 
 1 
 
 fc 1 
 
 
 S* 
 
 SJf 
 
 1 
 
 
 1 
 
 
 
 
 s's 
 
 g 
 
 
 OS (JJ 
 
 
 o 
 
 
 II II 
 
 II II 
 
 1 
 
 
 1 
 
 
 
 5 
 
 S2 
 
 
 Ao 
 
 
 . 
 
 
 i "^ 
 
 SB ^~ 
 
 OS 
 
 1 
 
 I 
 
 ** ^ 
 
 P, W o 
 
 1 * * 
 
 O 
 
 II 
 fe 
 
 " : 
 
 S ii 
 
 
 
 ii 
 
 ii 
 i 
 
 II 
 
 H 
 
 1 
 
 1 
 
 1 
 
 M 
 
 
 8 
 
 g 
 
 5 
 
 1 
 
 1 
 
 
 
 II TO 
 
 
 
 
 
 MM 
 
 ila 
 
 o 
 II 
 
 03 g 
 
 II 
 
 OJ 
 02 
 
 II 
 
 n 
 
 1 
 
 II 
 
 
 o 
 
 
 
 
 
 i 
 
 i 
 
 
 
 
 
 TO 
 
 B 
 
 TO 
 
 
 
 
 I rf 4 
 
 
 O 
 
 II 
 
 II 
 P-l ** 
 
 2 
 
 ll 
 
 II c, 
 
 3 a6 
 
 II 
 
 g 
 
 m 
 II ^ 
 
 02 II 
 
 5 
 
 oo 
 
 S 
 
 II 
 
 II 
 
 S 
 
 1 
 
 
 
 oo 
 
 
 
 00 
 
 
 Oi 
 
 
 
 
 
 
 
 1 
 
 
 HH 
 
 
 II 
 
 II ^ 
 
 II 10 - 
 
 
 
 S w o 
 
 JSS 
 
 II 
 
 
 II 
 H 
 
 O OT 
 
 ii 
 
 3 
 
 1 
 
 n 1 
 
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 fe 
 
 Oi 
 
 z> 
 
 
 ,_, 
 
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 o 
 
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 II 
 
 
 
 
 
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 J s 
 ii 
 
 TO 
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 i 
 
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 .0 
 
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 1 
 
 j 
 
 
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 3 
 
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 "* 
 
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 a " 
 
 S ; 
 
 
 II 
 
 o 
 
 1 
 
 6 
 
 
 M 
 
 1 
 
 1 
 
 
 ^ 
 
 TO 
 
 TO 
 
 OS 
 
 
 ^ 
 
 1-4 
 
 It 
 
 
 i 
 
 1 0* 
 
 x 
 
 
 
 & i : 
 
 W 
 
 5z os 
 
 n ^ 
 
 II TO' 
 
 
 II 
 
 ' 
 
 *" 
 
 
 o II 
 
 'So 1 
 
 
 3^ 
 
 
 II 
 
 3 
 
 II 
 M 
 
 " 3 
 
 Sg 
 
 1 
 
 ^ 1 
 
 1 
 
 _ ,,, 
 
 TO ** 
 
 CO 
 
 l> 00 
 
 
 
 S S 
 
 02 
 
 
 
 
 
 
 
150 
 
 INOEGANIG CHEMISTRY. 
 
 MENDELEEFF'S TABLE II. 
 
 
 
 
 I 
 
 
 
 
 
 
 
 
 I. 
 
 II. 
 
 III. 
 
 IV. 
 
 V. 
 
 VI. 
 
 R a O 
 
 I. 
 
 
 Li = 7 
 
 K 39 
 
 Rb 85 
 
 Cs 133 
 
 - - 
 
 - - 
 
 RO 
 
 II. 
 
 
 Be = 9 
 
 Ca 40 
 
 Sr 87 
 
 Ba 137 
 
 
 
 
 
 R 3 3 
 
 III. 
 
 
 B = ll 
 
 Sc 44 
 
 Y 89 
 
 La 138 
 
 Yb 173 
 
 
 
 RO, 
 
 IV. 
 
 (H 4 C) 
 
 C = 12 
 
 Ti 48 
 
 Zr 90 
 
 Ce 142 
 
 
 
 Th 231 
 
 R a 6 
 
 V. 
 
 (H 3 N; 
 
 N = 14 
 
 V 51 
 
 Cb 94 
 
 Di 146 
 
 Ta 182 
 
 _ _ 
 
 RO, 
 
 VI. 
 
 (H a O) 
 
 O = 16 
 
 Cr 52 
 
 Mo 96 
 
 
 
 W 184 
 
 U 240 
 
 R 3 7 
 
 VII. 
 
 (HF) 
 
 F = 19 
 
 Mn 55 
 
 _ _ 
 
 
 
 
 
 
 
 RO 4 
 
 
 
 
 Fe 56 
 
 Ru 103 
 
 
 
 Os 192? 
 
 
 
 
 VIII. 
 
 
 
 Co 58 
 
 Rh 104 
 
 _ _ 
 
 Ir 193 
 
 - 
 
 
 
 
 
 Ni 59 
 
 Pd 106 
 
 
 
 Pt 195 
 
 
 
 R a O 
 
 I. 
 
 H = l 
 
 Na=23 
 
 Cu 63 
 
 Ag 108 
 
 
 
 Au 196 
 
 
 
 RO 
 
 II. 
 
 
 Mg 24 
 
 Zn 65 
 
 Cd 112 
 
 
 
 Hg200 
 
 _ 
 
 R a O, 
 
 III. 
 
 
 Al 2? 
 
 Ga 69 
 
 In 113 
 
 
 
 Tl 204 
 
 
 
 RO 2 
 
 IV. 
 
 (H 4 R) 
 
 Si 28 
 
 Ge 72 
 
 Sn 118 
 
 
 
 Pb 206 
 
 
 
 R.,0 6 
 
 V. 
 
 (H 3 R) 
 
 P 31 
 
 As 75 
 
 Sb 120 
 
 
 
 Bi 209 
 
 
 
 RO 8 
 
 VI. 
 
 (H a R) 
 
 S 32 
 
 Se 79 
 
 Te 125? 
 
 
 
 
 
 _ _ 
 
 R,0 T 
 
 VII. 
 
 (HR) 
 
 01 35.5 
 
 Br 80 
 
 I 127 
 
 - - 
 
 - - 
 
 - - 
 
 In the above tables the approximate atomic weights 
 are used instead of those which have been determined 
 and calculated with the greatest care. For most pur- 
 poses in the laboratory the approximate figures answer 
 well enough, and they are most commonly used. In 
 the following table of Lothar Meyer the refined atomic 
 weights (as calculated by Meyer and Seubert) are used. 
 The difference between the two sets of figures is in most 
 cases very slight. The atomic weights adopted in this 
 book are those calculated by F. W. Clarke (see " The 
 Constants of Nature," Part V, 1897). 
 
MEYERS TABLE OF THE ELEMENTS. 
 
 151 
 
 S 85 
 
 " sf 
 
 5 
 
 S s 
 
 s a 
 
152 INORGANIC CHEMISTRY. 
 
 In Mendeleeff's Table I the elements are arranged in 
 horizontal lines, beginning with lithium. When the 
 eighth element in the order of the increasing atomic 
 weights is reached it is found that it is very much like 
 lithium. It is sodium. If this is placed below lithium, 
 and the next six elements in the same horizontal line, 
 when the fifteenth element is reached, it is found like 
 the eighth to be similar to lithium. Up to and includ- 
 ing manganese there are twenty-one elements excluding 
 hydrogen. These fall then naturally into three series of 
 seven members each, and placing these horizontally, 
 those elements which fall in the same perpendicular 
 lines have the same general character. This is seen 
 most strikingly in Group I, in which lithium, sodium, 
 and potassium fall, and in Group Y, in which nitrogen, 
 phosphorus, and vanadium fall ; but there is no difficulty 
 in recognizing the similarity in the other groups. The 
 three elements following manganese, viz., iron, nickel, 
 and cobalt, are very much alike, and they certainly do 
 not belong in Groups I, II, and III, while the next 
 element, copper, has some properties which ally it to the 
 members of Group I. The next six elements fall in 
 Groups II to VII, and are evidently in place, and the 
 six following fall in Groups I to YI, and are also in their 
 proper places, as far as their properties are concerned. 
 After molybdenum in the sixth series comes a blank 
 which means that there is no element to fill that place, 
 but that probably there is one undiscovered which has 
 the atomic weight approximately 100, and has properties 
 which are similar to those of manganese. Then follow 
 three elements which resemble one another as closely as 
 iron, nickel, and cobalt do. These do not belong in 
 Groups I, II, and III, but form a small independent 
 group. These two groups of three elements occur at the 
 end of the fourth and sixth series respectively. We 
 should therefore expect to find a similar group at the end 
 of the eighth series. No such group is known, however, 
 though at the end of the tenth series, where we should 
 look for the next similar small group, there are the three 
 elements iridium, platinum, and osmium. The elements 
 
ARRANGEMENT IN MENDELEEFF'S SECOND TABLE. 153 
 
 of series 2 beginning with lithium and ending with fluo- 
 rine differ in some respects quite markedly from all the 
 other elements, as will be seen when they are taken 
 up. Beginning with sodium, it will be seen that there 
 are two series of seven elements and a short series of 
 three ; then again two series of seven and a series of 
 three ; and, although the following series are imperfect, 
 it is not difficult to recognize that the same general ar- 
 rangement of the elements holds good to the end. A 
 series of seven elements is called a short period ; while 
 two short periods with the accompanying three similar 
 elements constitute what is called a long period. 
 
 In MendeleefFs Table II the long periods are ar- 
 ranged in perpendicular lines, each long period begin- 
 ning and ending with a short period and having a side 
 group of three elements in the middle. Thus in the 
 column beginning with potassium there is, first, the 
 short period potassium to manganese, then the side 
 group iron, cobalt, nickel, and then the short period 
 copper to bromine. In this table similar elements occur 
 in the same horizontal lines. Thus in one line there are 
 lithium, potassium, rubidium, and caesium ; in another 
 sulphur, selenium, and tellurium ; and in another chlo- 
 rine, bromine, and iodine. 
 
 The symbols at the top of each column in Table I 
 have reference to the general formulas of the compounds 
 which the elements in each group form with oxygen and 
 with hydrogen. Beginning with Group I, the general 
 formula of the oxygen compounds of the members of this 
 group is B 2 O, in which B represents any element of that 
 group ; the general formula of the oxygen compounds of 
 the members of Group II is BO ; and so on. It will be 
 observed that the oxygen compounds grow more and 
 more complex from Group I to Group VII. Writing 
 BO, BO 2 , and BO 3 with doubled formulas, thus : B 2 O 2 , 
 B 2 O and B 2 O 6 , the series of oxygen compounds is 
 represented as below : 
 
 B 9 0, B 2 2 , B 2 3 , B 2 4 , B 2 6 , B 2 O 6 , B 2 O 7 . 
 
154 INORGANIC CHEMISTRY. 
 
 As regards the general formula of the oxygen com- 
 pounds of the members of Group VIII, it must be said 
 that it does not in general correspond to the composition 
 EO 4 . Osmium and ruthenium do, however, form the 
 oxides OsO 4 and EuO 4 . 
 
 There is also regularity in the composition of the hy- 
 drogen compounds. Beginning with Group VII, those 
 members which combine with hydrogen form compounds 
 of the general formula EH, as, for example, C1H, hydro- 
 chloric acid ; FH, hydrofluoric acid ; etc. Those mem- 
 bers of Group VI which combine with hydrogen form 
 compounds of the general formula EH 2 , as, for example, 
 water, H 2 O, and hydrogen sulphide, H 2 S. The maxi- 
 mum power of combining with hydrogen is met with in 
 Group IV, in which occur the elements carbon and sili- 
 con. These form the hydrogen compounds CH 4 and 
 SiH 4 . The members of Groups I, II, and III do not 
 readily form compounds with hydrogen. A few are 
 known, but they are quite unstable. 
 
 The hydroxides vary in composition from the simple 
 form E(OH) to E(OH) 7 . While in the first four groups 
 well-marked examples of the hydroxides E(OH), E(OH) 2 , 
 E(OH) 3 , and E(OH) 4 are found, in the fifth group the 
 hydroxides have not the general formula E(OH) 5 , though 
 several of them have the formula OE(OH) 3 , as phos- 
 phoric, arsenic, and antimonic acids, which are respec- 
 tively OP(OH) 3 , OAs(OH) 3 , and OSb(OH) 3 . These may 
 be regarded as derived from the hydroxides of the gen- 
 eral formula E(OH) 5 by loss of water : 
 
 K(OH) S = OE(OH) 3 + H,0. 
 
 Hydroxyl derivatives of the members of Group VI 
 corresponding to the general formula E(OH) 6 are known, 
 as, for example, the so-called hydrate of sulphuric acid, 
 S(OH) 6 . The maximum hydroxides of Group VII should 
 have the general formula E(OH) 7 , but those known do 
 not correspond to this. The nearest approach to it is 
 found in crystallized periodic acid, H 5 IO 6 , which may be 
 
LOTHAR MEYER'S ARRANGEMENT. 155 
 
 regarded as derived from the hydroxide I(OH) 7 by loss 
 of one molecule of water, thus : 
 
 I(OH), = OI(OH). + H,0. 
 
 The arrangement of Lothar Meyer is a continuous 
 one. The elements are arranged on a spiral beginning 
 with lithium and ending with uranium. The divisions 
 are such that when the two ends of the table are brought 
 together on a cylinder, the line ending with fluorine will 
 join that beginning with sodium ; that ending with nickel 
 will join that beginning with copper ; and so on. In other 
 respects the arrangement is much like that in Mende- 
 leefF s Table I. "What Mendeleeff calls a Group, Lothar 
 Meyer calls a Natural Family, while those elements which 
 fall in the same horizontal line are said to form a series. 
 Now, each natural family falls into two groups indicated 
 by the letters A and B placed above. Thus the first 
 natural family falls into Group A, consisting of lithium, 
 sodium, potassium, rubidium, and caesium, and Group B, 
 consisting of copper, silver, and gold. Family II falls 
 into Group A, consisting of glucinum, magnesium, cal- 
 cium, strontium, barium, and perhaps erbium, and Group 
 B, consisting of zinc, cadmium, and mercury ; etc. The 
 members of each group in a family resemble one another 
 much more closely than they resemble the members of 
 the other group. * 
 
 The figures at the bottom of the table of Lothar 
 Meyer refer to the valence of the elements in each group. 
 Judging by the composition of the oxides and hydrox- 
 ides, the valence increases from 1 to 8 from Families I to 
 VIII. But the valence of the elements in each family 
 varies according to conditions. Thus the valence of the 
 elements of Family IY is generally 4, as shown in the. 
 compounds CH 4 , CC1 4 , CO 2 , SiH 4 , SiCl 4 , SiO 2 , etc. ; but 
 they may also appear as bivalent elements, as seen in the 
 compound CO. So too, while the elements of Family V 
 are quinquivalent, as in PC1 5 , NH 4 C1, etc., they may also 
 be trivalent, as in PC1 3 , NH 3 , etc. 
 
 What we call valence does not then appear to be an 
 
156 INORGANIC CHEMISTRY. 
 
 unchangeable property of the elements, but a property 
 Avhich may change according to conditions. It appears 
 further that a given element may have one valence tow- 
 ards one element and another valence towards another 
 element. This is most strikingly seen on comparing 
 the formulas of the hydrogen compounds of the ele- 
 ments of Families V, VI, and VII with those of their 
 oxygen compounds. The members of Family VII com- 
 bine with hydrogen in only one proportion, and that is 
 the simplest possible. Towards hydrogen these elements 
 are univalent, and their valence towards hydrogen is con- 
 stant. On the other hand, they combine with oxygen 
 and with hydroxyl in several proportions, and judging by 
 the composition of these compounds, the valence tow- 
 ards oxygen varies from 1 to 7. The members of 
 Family VI are bivalent towards hydrogen, and their hy- 
 drogen valence is constant ; but they combine with oxy- 
 gen and hydroxyl in several proportions, and the compo- 
 sition of the compounds indicates that their valence 
 towards oxygen varies from 2 to 6. The hydrogen val- 
 ence of the members of Family V is 3, while the oxygen 
 valence varies from 1 to 5. Finally, the hydrogen val- 
 ence of the members of Family IV is 4, while the oxygen 
 valence varies from 2 to 4. As regards the hydrogen 
 valence of the members of Families I to III, but little is 
 known. These elements do not generally combine with 
 hydrogen, though some of them do. 'Towards oxygen 
 their valence is fairly constant, though some variations 
 are observed as in the case of copper and mercury. 
 
 Judging then by the composition of the compounds, 
 we are justified in making a distinction between the 
 hydrogen-valence of some elements and their oxygen- 
 valence. While the former is constant, the latter is sub- 
 ject to variations. In those cases in which there is a 
 marked difference between the hydrogen-valence of an ele- 
 ment and its maximum oxygen-valence, the maximum 
 valence towards chlorine is greater than the hydrogen- 
 valence and less than the maximum oxygen-valence. 
 This is shown in the case of sulphur ; the formulas of its 
 
THE ELEMENTS IN THE NATURAL SYSTEM. 157 
 
 hydrogen compound and of its highest compounds with 
 chlorine and oxygen being respectively 
 
 SH 2 , SC1 4 , SO,. 
 
 Prom this it appears that the maximum valence of sul- 
 phur towards hydrogen is 2, towards chlorine 4, and 
 towards oxygen 6. 
 
 Connection between the Position of the Elements in the 
 Natural System and their Chemical Properties. The 
 changes in composition of the oxygen and hydrogen 
 compounds and of the hydroxides from Family I to VII 
 have been referred to. Another fact of great impor- 
 tance is that the elements of Group I are the most 
 strongly marked base-forming elements, while those of 
 Group YII are the most strongly marked acid-forming 
 elements. Passing in either direction the character of 
 the elements becomes less pronounced, until in the mid- 
 dle (Group IV), elements which form neither strongly 
 marked acids nor strongly marked bases are found. 
 Thus, beginning with sodium, this element forms a 
 strong base, magnesium forms a weaker base, the hy- 
 droxide of aluminium is a still weaker base. Beginning, 
 on the other hand, with chlorine at the other end of the 
 same series, its hydrogen compound is a strongly marked 
 acid ; that of sulphur is an acid, but less marked in char- 
 acter than hydrochloric acid ; that of phosphorus has no 
 acid properties, nor has that of silicon. The hydroxides 
 of these four elements have acid properties. Each one, 
 however, forms several acids, and it is difficult to com- 
 pare them, as some of those of chlorine are strongly 
 marked and others not, as we have seen. 
 
 Some very interesting variations in properties are also 
 noticed in passing from one end of a group of a natural 
 family to the other. Thus in Group B, Family VII, the 
 activity of the elements grows less from fluorine to iodine, 
 or, as we commonly say, fluorine is the strongest ele- 
 ment in the group, and then follow, in order, chlorine, 
 bromine, and iodine. 
 
158 INORGANIC CHEMISTRY. 
 
 The remarkable relations above referred to are summed 
 up in the periodic law : 
 
 The properties of an element are periodic functions of the 
 atomic weight. 
 
 It appears that if an element has a certain atomic 
 weight it must have certain properties, and that if the 
 atomic weight is known the properties can be stated, 
 just as, if the properties are known, the atomic weight 
 can be approximately stated. When the law was first 
 stated, Mendeleeff predicted the discovery of certain ele- 
 ments to fill some of the vacant places in the table. At 
 that time the elements gallium, Ga, scandium, Sc, and 
 germanium, Ge, were not known. Not only was their 
 discovery predicted, but their properties were clearly 
 stated years before they were brought to light. Within 
 the last few years these three elements have been dis- 
 covered, and a remarkable agreement is observed be- 
 tween their properties as determined by observation and 
 as foretold by Mendeleeff by the aid of the periodic 
 law. 
 
 The relations between the atomic weights and proper- 
 ties will appear more and more clearly as our study of 
 the elements proceeds. The natural arrangement of the 
 elements suggested by the periodic law is adopted in 
 this book. The elements hydrogen, oxygen, and chlo- 
 rine were studied at the outset in order to illustrate the 
 methods of studying chemical problems, and as exam- 
 ples of chemical elements in general. It is, however, 
 now time to take up the elements systematically, and to 
 learn what may be necessary in regard to them in order 
 to get as clear a notion as possible of the facts and prin- 
 ciples of the science of chemistry. 
 
 Plan to be followed. The most systematic method of 
 procedure in studying the elements would be to begin 
 with Family I, Group A (see Lothar Meyer's Table, 
 p. 151), then to take up Group B of the same family ; 
 and so on in order, ending with Family VIII. It seems 
 better, however, to begin with Family VII ; to follow 
 with Families VI, V, and IV ; and then to take up in 
 order Families I, II, III, and VIII. The main reason 
 
THE ELEMENTS IN THE NATURAL SYSTEM. 159 
 
 for this is that it is impossible to study most of the mem- 
 bers of Families I, II, III, and VIII without a knowl- 
 edge of several of the elements of Families VII, VI, V, 
 and IV, while these last families can be studied with 
 only slight reference to the others. It is proposed then 
 to begin with Group B, Family VII, the members of 
 which are very much like chlorine. The only member 
 of Group A of this family is manganese. While man- 
 ganese resembles the members of the chlorine group in 
 some respects, it has other properties which ally it to 
 the so-called base-forming elements. So also the mem- 
 bers of Group A, Family VI, are like the members of the 
 oxygen or sulphur group, but they are also allied to the 
 base-forming elements. A similar difference is observed 
 between the members of Groups A and B, Family V. 
 
 While the plan above sketched takes into considera- 
 tion the greater number of the analogies of the elements, 
 there are other analogies which are not brought out. 
 Thus, as will be seen in due time, the elements alumin- 
 ium, chromium, manganese, and iron are analogous in 
 some respects, but by following the plan sketched they 
 will be taken up in different groups. This appears to be 
 justified, however, when we consider the entire conduct 
 of these elements, and do not confine ourselves to a study 
 of only a few reactions which, being useful for some 
 purposes, have been studied more carefully than others 
 which from a scientific point of view are perhaps just as 
 important. 
 
CHAPTER XII. 
 
 THE ELEMENTS OF FAMILY VII, GROUP B: 
 FLUORINE-CHLORINEBROMINEIODINE. 
 
 General. The elements of this group are commonly 
 called the halogens. The best known member of the 
 group is chlorine, which has already been treated. Al- 
 though fluorine is in general like the other members of 
 the group, it differs from them in some respects, and it 
 certainly is not as much like them as they are like one 
 another. While chlorine, bromine, and iodine accom- 
 pany one another in nature, fluorine compounds are not 
 generally found in company with compounds of the other 
 elements of the family. In those cases in which chlorine, 
 bromine, and iodine are found together, chlorine is gen- 
 erally present in largest quantity, and iodine in smallest 
 quantity. Fluorine and chlorine are gases under ordi- 
 nary conditions, while bromine is a liquid and iodine is 
 a solid. Fluorine, bromine, and iodine form with hy- 
 drogen the compounds hydrofluoric acid, HF, hydro- 
 bromic acid, HBr, and hydriodic acid, HI, which 
 are analogous to hydrochloric acid. All these com- 
 pounds are gases which have marked acid properties. 
 With oxygen, fluorine does not combine, whereas chlo- 
 rine, bromine, and iodine combine with it in a number of 
 proportions, as has already been seen in the case of chlo- 
 rine. Among themselves these elements also form some 
 compounds : thus bromine and chlorine form the com- 
 pound BrCl; iodine forms the compounds IC1, IC1 3 , 
 IBr, and IF 5 . It appears from this that the valence of 
 iodine towards bromine is 1, towards chlorine 3, and 
 towards fluorine 5. 
 
 Towards base-forming members the elements of this 
 group are univalent, as shown in such compounds as 
 NaCl, KBr, CaCl a , KI, etc. They, however, appear to 
 
 (160) 
 
BROMINE: OCCURRENCE-PREPARATION. 161 
 
 have a valence greater than 1 in some compounds known 
 as double salts. These can be explained satisfactorily 
 only by assuming that in them the element is in combi- 
 nation with itself and has a valence greater than 1. 
 
 BROMINE, Br (At. Wt. 79.34). 
 
 Occurrence. This element occurs in nature in com- 
 pany with chlorine. Chlorine, as has been stated, occurs 
 mostly in combination with sodium, as sodium chloride, 
 or common salt. In several of the great salt-beds 
 bromine occurs in the form of sodium bromide, NaBr, and 
 in some places it occurs as magnesium bromide, MgBr 2 . 
 The chief source of bromine is the mother-liquors from 
 the salt works. When a solution containing a large 
 quantity of sodium chloride and a small quantity of bro- 
 mide is evaporated, the chloride is first deposited, and 
 from the mother-liquors the bromide mixed with chlo- 
 ride is deposited. The great beds at Stassfurt are par- 
 ticularly rich in bromides, and a great deal of bromine 
 is made from the salts which occur in this locality. 
 
 Preparation. Bromine can be prepared from the 
 bromides in the same way that chlorine is made from the 
 chlorides : by first treating with sulphuric acid, thus lib- 
 erating hydrobromic acid, and then treating with man- 
 ganese dioxide, or, better, by mixing the bromide with 
 manganese dioxide and treating the mixture with sul- 
 phuric acid. The reaction is represented by the equation 
 
 2NaBr + MnO 2 + 2H 2 SO 4 = 
 
 Na 2 SO 4 + MnSO 4 + 2H 2 O + Br 2 . 
 
 Or it may be represented as taking place in different 
 stages. First the sulphuric acid would liberate hydro- 
 bromic acid from the bromide, and this would act upon 
 the manganese dioxide thus : 
 
 MnO 2 + 4HBr -= MnBr 2 -f 2H 2 + Br 2 . 
 
 But sulphuric acid would act upon manganous bromide, 
 MnBr 2 , thus : 
 
 MnBr 3 -f H 2 S0 4 = MnS0 4 + 2HBr ; 
 
162 INORGANIC CHEMISTRY. 
 
 and the hydrobromic acid would then again react with 
 manganese dioxide, etc. 
 
 Another method for the preparation of bromine de- 
 pends upon the fact that chlorine has the power to set 
 bromine free from its compounds. If, therefore, a solu- 
 tion containing a bromide is treated with manganese 
 dioxide and hydrochloric acid, the chlorine which is 
 formed from the hydrochloric acid will act upon the 
 bromide and bromine will be given off. This method is 
 used at Stassfurt. 
 
 Properties. Bromine is a heavy, dark-red liquid at 
 ordinary temperatures. If exposed to the air it is con- 
 verted into a vapor of a brownish-red color. It boils at 
 58-58.6, and at 7.3 it is solid. It has an extremely 
 disagreeable odor, to which fact it owes its name (from 
 flpcdfiios, a stench). From carbon disulphide at 90 it 
 crystallizes in fine dark-red needles. 
 
 Its properties are similar to those of chlorine. It acts 
 violently upon organic substances ; attacking the skin, 
 and the membranes lining the passages of the throat and 
 lungs. Wounds caused by the liquid coming in contact 
 with the skin are painful and serious, and it must there- 
 fore be handled with great care. 
 
 Like chlorine, bromine is dissolved by water, one part 
 dissolving in 33.3 parts at 15. The solution, which has 
 a reddish color and the odor of bromine, is called bro- 
 mine tvater. At a low temperature bromine forms with 
 water a compound in every way analogous to chlorine 
 hydrate, wz., bromine hydrate, Br 2 -|- 10H 2 O. This de- 
 composes when left in contact with the air at ordinary 
 temperatures. 
 
 Chemical Conduct of Bromine. Bromine acts chemi- 
 cally like chlorine. It was pointed out that chlorine 
 acts in three different ways : (1) By direct addition ; (2) 
 by substitution; and (3) by liberating oxygen from 
 water, as in bleaching and other oxidizing processes. 
 Bromine is capable of acting in all three ways. It com- 
 bines directly with base-forming elements or metals, as 
 iron, aluminium, potassium, etc. ; also with the acid- 
 forming elements, as sulphur, phosphorus, etc. It com- 
 
HYDROBROMIC ACID. 163 
 
 bines with hydrogen almost as readily as chlorine does. 
 With oxygen it does not combine directly, and in this 
 respect also it is like chlorine. 
 
 It acts upon compounds containing hydrogen almost 
 as readily as chlorine does, replacing the hydrogen and 
 forming bromine substitution-products. Thus benzene, 
 C 6 H B , yields the products C 6 H 5 Br, C 6 H 4 Br 2 , C 6 H 3 Br 3 , 
 C 6 H 2 Br 4 , etc., and the hydrogen which leaves the com- 
 pound passes off in combination with bromine in the 
 form of hydrobromic acid. 
 
 It bleaches like chlorine, partly by direct action and 
 disintegration of the organic dye-stuffs, partly by action 
 upon water, liberating oxygen. 
 
 A solution of bromine in water left exposed to the 
 direct sunlight loses its color and becomes acid in conse- 
 quence of the decomposition of the water, as in the case 
 of chlorine : 
 
 Br 2 +H,O = 
 or 2Br a + 2H 2 O = 4HBr + O 2 . 
 
 Uses of Bromine. Bromine and its compounds are 
 used in photography, medicine, and to some extent in 
 the manufacture of coal-tar colors. It is manufactured 
 in large quantity, and a good proportion of it is manu- 
 factured in the United States. According to the official 
 report the production of bromine in the United States 
 in the year 1896 amounted to over 500,000 pounds. 
 
 Hydrobromic Acid, HBr. The only compound which 
 bromine forms with hydrogen alone is hydrobromic acid. 
 This is in all respects very much like hydrochloric acid. 
 It is set free from bromides by the action of sulphuric 
 acid, but owing to its instability it acts upon the sul- 
 phuric acid, causing decomposition. The elements hy- 
 drogen and bromine are not held together as firmly in 
 hydrobromic acid as hydrogen and chlorine are in hy- 
 drochloric acid. Consequently, if hydrobromic acid is 
 brought together with certain substances which contain 
 oxygen it gives up its hydrogen to the oxygen. This is 
 
164 INORGANIC CHEMISTRY. 
 
 seen in the conduct towards manganese dioxide. But 
 towards this substance both hydrochloric and hydro- 
 bromic acids act in essentially the same way. Sulphuric 
 acid does not, however, give up its oxygen as readily as 
 manganese dioxide, and the difference in the stability of 
 the hydrogen compounds of chlorine and bromine is 
 seen very clearly in their conduct towards sulphuric 
 acid. Hydrochloric acid does not act upon sulphuric 
 acid at all. Hydrobromic acid acts according to the 
 following equation : 
 
 2HBr + H 2 SO 4 = 2H 2 O + SO 2 + Br 2 . 
 
 The action consists in the decomposition of the hydro- 
 bromic acid into bromine and hydrogen, and the subse- 
 quent action of the nascent hydrogen upon the sulphuric 
 acid thus : 
 
 2HBr = 2H + Br 2 ; and 
 H 2 S0 4 + 2H = 2H 2 + S0 2 . 
 
 The hydrobromic acid acts here, then, as a reducing agent, 
 and the sulphuric acid as an oxidizing agent. It is plain 
 that hydrobromic acid cannot be made in pure condition 
 by the action of sulphuric acid upon a bromide. Some 
 of the hydrobromic acid, to be sure, escapes the action 
 of the sulphuric acid, but at best it is always mixed with 
 the compound SO 2 , or sulphur dioxide, which is a gas, 
 and with bromine. 
 
 It can be made by passing a mixture of hydrogen and 
 bromine over heated finely divided platinum. An ap- 
 paratus has been devised for making hydrobromic acid 
 in this way in quantity. 
 
 It can also be made by allowing bromine to act upon 
 an organic compound containing hydrogen. Substitut- 
 ing action takes place and hydrobromic acid is given off. 
 Thus, if a compound of the formula C 10 H 22 were used, the 
 reaction would be represented in this way : 
 
HTDROBROMIC ACID. 165 
 
 The product C 10 H 21 Br, or the bromine substitution- 
 product, would not be volatile at ordinary temperatures, 
 and therefore only the hydrobromic would be given off. 
 
 The method most commonly adopted in the labora- 
 tory consists in treating phosphorus with bromine and 
 water. In all probability the bromine acts first upon the 
 phosphorus, forming the product PBr 3 or PBr 5 accord- 
 ing to the proportions of the substances used. Both 
 these substances are decomposed by water, the first 
 forming phosphorous acid and hydrobromic acid, accord- 
 ing to this equation : 
 
 ( Br HHO 
 
 P^ Br + HHO = PO 3 H 3 + 3HBr, 
 ( Br HHO 
 
 or PBr 3 + 3H,0 = PO 3 H 3 + 3HBr ; 
 
 the second forming phosphoric acid and hydrobromic 
 acid : 
 
 PBr 5 + 4H 2 O = P0 4 H 3 + 5HBr. 
 
 The gas thus formed can be freed from bromine by 
 passing it through a tube containing phosphorus. 
 
 Properties. Hydrobromic acid is a colorless gas which 
 forms fumes in contact with the air in consequence of its 
 attraction for moisture. It dissolves in water in large 
 proportion. The solution conducts itself much like hy- 
 drochloric acid. When boiled a compound of definite 
 composition passes over under ordinary conditions. It 
 corresponds to the formula HBr -f- 5H 2 O, but here, as 
 with the hydrate of hydrochloric acid, the composition 
 changes with the pressure. With metallic hydroxides 
 or bases, hydrobromic acid forms bromides, as hydro- 
 chloric acid forms chlorides : 
 
 KOH + HBr = KBr + H 2 O. 
 
 Compounds of Bromine with Hydrogen and Oxygen. 
 With hydrogen and oxygen bromine forms compounds 
 which closely resemble those which chlorine forms with 
 the same elements. They are : Hypobromous acid, 
 
166 INORGANIC CHEMISTRY. 
 
 HBrO ; bromic acid, HBrO 3 ; and perhaps perbromic 
 acid, HBrO 4 . 
 
 Hypobromom acid, HBrO, is made by reactions which 
 are entirely analogous to those used in making hypo- 
 chlorous acid. When bromine acts upon a dilute solu- 
 tion of sodium or potassium hydroxide, reaction takes 
 place thus : 
 
 2KOH + Br a = KBr + KBrO + H 9 O. 
 
 So also bromine vapor acting upon slaked lime or cal- 
 cium hydroxide forms a compound similar to bleaching 
 powder. Hypobromous acid has not been prepared in 
 pure condition owing to its instability. 
 
 Bromic acid, HBrO 3 , is not known in pure condition. 
 Its salts are made in the same way as the chlorates 
 are ; the principal reaction made use of for the purpose 
 being that between bromine and concentrated potassium 
 hydroxide : 
 
 3Br a + 6KOH = 5KBr + KBrO 3 + 3H 2 O. 
 
 The decompositions of the bromates are much like 
 those of the chlorates. 
 
 As regards the existence of perbromic acid there is 
 some doubt. It is stated by one observer that he ob- 
 tained it by treating perchloric acid with bromine : 
 
 HC1O 4 + Br = HBr0 4 + 01. 
 
 Others have not succeeded in getting it in this or in any 
 other way. 
 
 Compound of Bromine and Chlorine. When chlorine 
 is passed into liquid bromine it is absorbed in large 
 quantity. If the process is carried on at a low tempera- 
 ture the product BrCl is formed. Above 10 it under- 
 goes decomposition. Although it is unstable, there is 
 no good reason for regarding this substance as anything 
 but a chemical compound. There are many chemical 
 compounds known which are less stable than this. 
 
IODINE: OCCURRENCE-PREPARATION". 167 
 
 IODINE, I (At. Wt. 125.89). 
 
 Occurrence. Iodine, as has already been stated, occurs 
 in company with chlorine and bromine in nature, but in 
 smaller quantity than these. The relative quantity in 
 sea water is extremely small. The sea plants, however, 
 assimilate it, and the ashes of these plants contain a 
 considerable quantity of compounds of iodine. It also 
 occurs in small quantity in the great beds of soda salt- 
 peter, or sodium nitrate, which are found in Chili, South 
 America. It occurs in small quantity in combination 
 with silver, and also in combination with lead and with 
 mercury. 
 
 Preparation. The method of obtaining iodine from 
 its salts is like that used in making chlorine and bromine 
 from the chlorides and bromides. It consists in treat- 
 ing the iodides with sulphuric acid and manganese 
 dioxide. 
 
 2KI + MnO 8 + 2H 2 SO 4 = K 2 SO 4 + MnSO 4 + 2H 2 + I 2 . 
 
 The iodine, although solid at the ordinary tempera- 
 ture, is easily volatilized, and if the mixture mentioned 
 is heated, iodine vapor passes over and may be con- 
 densed in appropriately arranged vessels. 
 
 On the large scale iodine is obtained mostly from sea- 
 weed. On the coasts of Scotland, Ireland, and France 
 the sea- weed which is thrown up by storms is gathered, 
 dried, and burned. The organic portions are thus de- 
 stroyed, and the mineral or earthy portions are left be- 
 hind as ashes. This incombustible residue is called 
 kelp. It contains a small percentage of potassium 
 iodide, from .5 to 2 per cent according to the sea- weed 
 used. The dried weed was formerly burned in cavities 
 dug in the earth, but of late years the process has in 
 some places been much improved, and the yield in kelp 
 increased. 
 
 In Scotland the iodine is liberated by means of sul- 
 phuric acid and manganese dioxide. In France, how- 
 
168 INORGANIC CHEMISTRY. 
 
 ever, this is effected by passing chlorine into the solution 
 containing the iodide. If too little chlorine is used all 
 the iodine is not separated ; if too much, a compound of 
 iodine and chlorine is formed, or an iodate, in conse- 
 quence of the oxidizing action of the chlorine on the 
 iodine. 
 
 The iodine which occurs in Chili saltpeter, NaNO 3 , is 
 in the form of sodium iodate, NaIO 3 , and iodide, Nal, 
 and to some extent as magnesium iodide, MgI 2 . Most 
 of the iodine now in the market is made from this ma- 
 terial, and the competition created in this way has led 
 to a careful study of the process for obtaining iodine 
 from kelp. Sea- weed is now collected from certain parts 
 of the ocean where it grows in large quantity, vessels 
 being sent out for the purpose. 
 
 Properties. Iodine is a grayish-black crystallized 
 solid. At ordinary temperatures it is volatile. Accord- 
 ing to the most reliable determinations it melts at 
 113-115, and boils at 250. The vapor has a violet 
 color when mixed with air. When in pure condition 
 it is intensely blue. At temperatures considerably 
 above the boiling point the specific gravity of iodine 
 vapor is such as to show that its molecular weight is 
 approximately 254, or twice the atomic weight. As the 
 temperature is raised, however, the specific gravity is 
 lowered, until, finally, at a very high temperature, it 
 becomes about half what it is at lower temperatures. 
 This is accounted for by supposing that at the lower 
 temperatures the molecules of iodine consist of two 
 atoms each, while as the temperature is raised these 
 molecules are gradually broken down, so that at the 
 temperature at which the lowest specific gravity is 
 reached the iodine vapor consists of free atoms, or the 
 atoms and molecules are then identical, and the specific 
 gravity is therefore only half what it is when the mole- 
 cules consist of two atoms. 
 
 Iodine has a characteristic strong taste. It acts upon 
 the mucous membranes, but much less energetically 
 than chlorine or bromine. It colors the skin yellowish- 
 
HYDRIODIC ACID. 169 
 
 brown, and acts as an absorbent, causing the reduction 
 of some kinds of swellings. 
 
 It dissolves slightly in water, easily in alcohol, and 
 easily in a water solution of potassium iodide. The 
 solution in alcohol is known as tincture of iodine. It 
 -dissolves also in carbon disulphide, CS 2 , and in chloro- 
 form forming solutions which have a beautiful deep 
 violet color. 
 
 In general, iodine conducts itself chemically like bro- 
 mine and chlorine, only it acts in almost all reactions 
 less energetically than the other two elements. It com- 
 bines directly with a number of elements, as with hydro- 
 gen, sulphur, phosphorus, iron, mercury, etc. In pres- 
 ence of water it acts as an oxidizer just as chlorine and 
 bromine do, but less energetically. Thus it oxidizes 
 sulphurous acid, H 2 SO 3 , to sulphuric acid, H 2 SO 4 : 
 
 H 2 S0 3 + I 2 + H 2 O = H 2 S0 4 + 2HI. 
 
 As a substituting agent it does not act as readily as 
 chlorine and bromine, though iodine substitution-prod- 
 ucts are made in large quantities, particularly in con- 
 nection with the manufacture of dye-stuffs. 
 
 Iodine is used extensively in the dye-stuff industry, in 
 photography, and in medicine. One factory in Scotland 
 makes on an average 60 tons of iodine a year. 
 
 Hydriodic Acid, HI. The affinity of hydrogen for 
 iodine is less than for bromine, and therefore hydriodic 
 acid cannot be made pure by treating an iodide with 
 sulphuric acid. The hydrogen of the hydriodic acid 
 acts upon the sulphuric acid very readily, and according 
 to the conditions the following reactions may take 
 place : 
 
 H 2 S0 4 + SHI = 4H 2 + SH 3 + 41, ; 
 H a S0 4 + 6HI = 4H 2 + S + 31, ; 
 H 2 S0 4 + 2HI = 2H 2 + S0 2 + I 2 . 
 
 On treating potassium iodide with sulphuric acid, 
 therefore, there may be formed, in addition to hydriodic 
 
170 INORGANIC CHEMISTRY. 
 
 acid and free iodine, sulphur dioxide, sulphur and hydro- 
 gen sulphide. 
 
 The method adopted for the preparation of hydriodic 
 acid is like that used for the preparation of hydrobromic 
 acid. It consists in treating phosphorus with iodine and 
 water. The reactions involved are of the same kind as 
 those which were discussed under hydrobromic acid. 
 The iodine probably acts at first on the phosphorus, 
 forming a compound of phosphorus and iodine, which 
 then in turn is decomposed by the water. The reactions 
 which generally take place are those represented by the 
 following equations : 
 
 P +31 = PI 8 ; 
 
 PI 3 + 3H a O = P0 3 H 3 + SHI. 
 
 Hydriodic acid is a colorless transparent gas like hy- 
 drochloric and hydrobromic acids. It also like these dis- 
 solves in water in large quantity, and when brought in 
 contact with the air it forms dense white fumes. When 
 boiled the water solution conducts itself like those of 
 hydrochloric and hydrobromic acids. The liquid, which 
 boils at 127 under the ordinary atmospheric pressure, 
 contains 57 per cent hydriodic acid. If the solution of 
 the gas in water is allowed to stand, decomposition be- 
 gins in consequence of the action of the oxygen of the 
 air. The hydrogen is oxidized to water and the iodine 
 is set free, coloring the solution brown. 
 
 When heated, the gas begins to decompose at 180, 
 and at higher temperatures the decomposition takes 
 place rapidly. The products are simply hydrogen and 
 iodine. In consequence of the ease with which hydrio- 
 dic acid breaks down, yielding free hydrogen, it is an ex- 
 cellent reducing agent, and it is frequently used in the 
 laboratory for the purpose of extracting oxygen from 
 substances. Its action upon sulphuric acid has already 
 been spoken of. The reason why it acts so well is that 
 the hydrogen is separated from the iodine with little ex- 
 penditure of energy, and the hydrogen thus separated 
 is in the nascent state, or, as is believed, in the atomic 
 state. 
 
IODIC ACID. 171 
 
 lodic Acid, HIO 3 . This compound is strictly analo- 
 gous to chloric and bromic acids, but differs from them 
 in being much more stable. It can be made by treat- 
 ing iodine with strong oxidizing agents, as, for example, 
 concentrated nitric acid. It is also formed very easily 
 by passing chlorine through water in which iodine is 
 suspended, when hydrochloric acid and iodic acid are 
 formed, as represented in this equation : 
 
 I 2 + 5C1 2 + 6H 2 O = 2HI0 3 + 10HC1. 
 
 The reaction is probably somewhat more complicated 
 than it appears from this equation, for when chlorine 
 acts upon iodine a compound of the two elements is 
 first formed. Iodine trichloride is decomposed by water 
 thus : 
 
 2IC1 3 + 3H 2 O = 5HC1 + HIO 3 + 101. 
 
 Iodine monochloride is also decomposed by water, giv- 
 ing iodic acid, hydrochloric acid, and free iodine : 
 
 10IC1 + 6H 2 O = 10HC1 + 2HIO 3 + 4I 2 . 
 
 Whether these chlorides of iodine are formed or not, 
 the prime causes of the formation of iodic acid when 
 chlorine acts upon iodine in water are the oxidizing 
 power of the chlorine and the affinity of iodine for 
 oxygen. 
 
 When iodine is dissolved in an alkali the reaction 
 which takes place is the same as that which takes place 
 with chlorine and bromine under like circumstances. A 
 mixture of the iodide and iodate is formed : 
 
 6KOH + 3I 2 = 5KI + KIO 3 + 3H 2 O. 
 
 Iodic acid is a crystallized solid, which when heated 
 to 170 loses water and is converted into iodine pent- 
 oxide, I 2 O 5 : 
 
 2HI0 3 = 1,0. + H 2 O. 
 
172 INORGANIC CHEMISTRY. 
 
 Its salts have the general formula MIO 3 , though it alsa 
 forms salts MH(IO 3 ) 2 and MH 2 (IO 3 ) 3 . It gives up its 
 oxygen readily and is therefore a good oxidizing agent, 
 just as hydriodic acid is a good reducing agent. 
 
 Iodine Pentoxide or lodic Anhydride, I 2 O a . This com- 
 pound is formed, as was stated in the last paragraph, 
 by heating iodic acid to 170. It is a white solid which 
 is easily soluble in water, forming iodic acid. It is de- 
 composed when heated to 300. It will be observed, 
 therefore, that this compound of iodine and oxygen is 
 very much more stable than any of the compounds of 
 chlorine or bromine and oxygen ; and it is interesting to 
 note that as the affinity for oxygen increases, that for 
 hydrogen decreases. In the group chlorine, bromine, 
 and iodine, chlorine has the strongest affinity for hydro- 
 gen and the weakest for oxygen, while iodine has the 
 strongest affinity for oxygen and the weakest for hydro- 
 gen. We shall presently see that fluorine, which does 
 not unite with oxygen, has a stronger affinity for hydro- 
 gen than chlorine has. 
 
 Anhydrides, or Acidic Oxides. An oxide which, like 
 iodine pentoxide, forms an acid when dissolved in water, 
 or which forms salts by treatment with basic hydroxides, 
 is called an anhydride or acidic oxide. The oxides of the 
 base-forming elements form bases when dissolved in 
 water, and they are, therefore, called basic oxides. As 
 examples of acidic oxides or anhydrides, there may be 
 mentioned besides iodic anhydride, sulphuric anhydride, 
 SO 3 ; sulphurous anhydride, SO 2 ; phosphoric anhy- 
 dride, P 2 O 6 ; carbonic anhydride, CO 2 . When dissolved 
 in water these oxides are converted into acids as repre- 
 sented in these equations : 
 
 S0 3 + H 2 = H 2 S0 4 ; 
 S0 2 +H 2 = H 2 S0 8 ; 
 P 2 5 + H 2 = 2HP0 3 ; 
 C0 2 + H 2 = H 2 C0 3 . 
 
 Silicic anhydride, SiO 2 , is an example of an acidic oxide 
 which does not dissolve in water, but which does form 
 salts when treated with basic hydroxides : 
 
PERIODIC ACID. 173 
 
 SiO 2 + 2KOH = K 2 Si0 3 + H a O. 
 
 As examples of basic oxides or oxides which when 
 treated with water yield bases, the following may be 
 taken : calcium oxide, CaO ; potassium oxide, K 2 O ; ba- 
 rium oxide, BaO. As has already been shown, when 
 treated with water these are respectively converted into 
 calcium hydroxide, Ca(OH) 2 ; potassium hydroxide, 
 KOH; and barium hydroxide, Ba(OH) 2 . There are, 
 however, many basic oxides which do not dissolve in 
 water, but which, nevertheless, have the power to neu- 
 tralize acids and form salts. This is true, for example, 
 of aluminium oxide, A1 2 O 3 , lead oxide, PbO, manganous 
 oxide, MnO, cupric oxide, CuO, etc. The action of such 
 oxides upon acids takes place as represented below : 
 
 A1 2 3 + 3H 2 S0 4 = A1 2 (S0 4 ) 3 + 3H 2 ; 
 PbO + 2HN0 3 = Pb(N0 3 ) 2 + H 2 ; 
 MnO + 2HC1 = MnCl 2 + H 2 O ; 
 CuO + H 2 S0 4 = CuS0 4 + H 2 O, 
 
 Periodic Acid, H 5 IO 6 . This acid is analogous to per- 
 chloric acid. Its salts are formed by oxidation of iodates 
 or by heating iodates, just as perchlorates are formed by 
 heating chlorates. The simplest way to make a peri- 
 odate is to pass chlorine into a solution containing so- 
 dium hydroxide and sodium iodate, when a reaction 
 takes place which is at least partly represented by the 
 following equation : 
 
 NaIO 3 + 3NaOH + C1 2 = Na 2 H 3 IO. + 2NaCl. 
 
 The salt Na 2 H 3 IO 6 is difficultly soluble in water, and 
 therefore separates from the solution. From the sodium 
 salt the corresponding silver salt, Ag 2 H 3 IO 6 , can be ob- 
 tained, and when this silver salt is treated with nitric 
 acid it is converted into the simpler salt, AgIO 4 , which 
 is evidently derived from the simpler acid, HIO 4 : 
 
 2A&H.IO. + 2HNO S At 2AgNO 3 + 4H S O + 2AgIO.. 
 
174 INORGANIC CHEMISTRY. 
 
 The acid when separated from its solutions is a crys- 
 tallized solid which has the composition H 6 IO 6 . When 
 heated it undergoes decomposition, losing water and 
 oxygen, and yielding iodic acid. It cannot, however, bo 
 converted into a compound of the composition HIO 4 , 
 for the loss of water is always accompanied by a loss of 
 oxygen. Like iodic acid, periodic acid is a good oxidiz- 
 ing agent in consequence of the ease with which it gives 
 up its oxygen. 
 
 Periodates. Periodic acid yields a large number of 
 salts the connection between which and the acid does 
 not appear clear at first sight. A few examples will 
 suffice for the present purpose : KIO 4 , Na 5 IO 6 , Ag 3 IO 5 , 
 AgJA, Zn,I,O n . 
 
 Constitution of Periodic Acid. The complicated salts 
 of periodic acid are apparently inexplicable on any other 
 theory than that they are derived from acids which are 
 closely related to the hypothetical acid I(OH) 7 . This i 
 now commonly regarded as normal periodic acid. It, 
 however, breaks down into the ordinary form of the acid 
 by loss of water. The relation is expressed thus : 
 
 H 2 0. 
 
 The salts Na 2 H 3 IO 6 , Na 5 IO 6 , and others of the same kind 
 are derived from this acid by replacement of one or more 
 of the hydrogen atoms by metallic elements. The acid 
 of the formula H 5 IO 6 can also be imagined to break 
 down into H 3 IO 5 and water thus : 
 
 ,. 
 
 roH 
 
 OH 
 OH 
 OH = I 
 OH 
 OH 
 OH 
 
 OH 
 OH 
 OH 
 OH 
 OH 
 
CONSTITUTION OF PERIODIC ACID. 
 
 175 
 
 The salt Ag 3 IO 5 and similar known salts are plainly 
 derived from this hypothetical acid H 3 IO 5 . Finally, the 
 acid H 3 IO 6 can also be imagined to break down into 
 HIO 4 and water thus : 
 
 = I 
 
 O 
 
 OH 
 
 and the salts like KIO 4 are derived from this hypo- 
 thetical acid. It appears, therefore, that the assump- 
 tion of the fundamental normal acid, I(OH),, is com- 
 petent to explain the existence of the salts which are 
 derived from the acids H 5 IO 6 , H 3 IO 5 , and HIO 4 . More 
 complicated acids can be formed by the loss of water 
 from two or more molecules of any one of these simpler 
 acids. Thus, if from two molecules of the acid H 3 IO 6 
 one molecule of water is taken, an acid of the formula 
 H 4 I 2 O 9 would be formed ; or if two molecules of the acid 
 H 6 IO 6 lose one molecule of water, the acid H 8 I 2 O n would 
 be formed. These relations are made clear by the 
 equations here given : 
 
 21 < 
 
 O 
 
 O 
 
 OH = (HO) 2 2 I-0-I0 2 (OH) 2 + H 2 ; 
 
 OH 
 
 OH 
 
 'O 
 OH 
 
 21 H 
 
 - (HO) 4 OI-0-IO(OH) 4 
 
 OH 
 OH 
 
 A salt of the formula Ag 4 I 2 O 9 , and another of the for- 
 mula Zn 4 I 2 O n , are known. The former is derived from 
 the acid H 4 I 2 O 8 , the latter from the acid H 6 I,O n , by the 
 substitution of four bivalent atoms of zinc for the eight 
 atoms of hydrogen. There are many more complicated 
 
176 INORGANIC CHEMISTRY. 
 
 salts than those mentioned, but they can all be satisfac- 
 torily explained by the assumption that they are related 
 to the normal acid I (OH) 7 , in which iodine is septivalent. 
 The existence of the periodates, the ease with which they 
 can be explained by the above method, and the apparent 
 impossibility of explaining them on the assumption that 
 iodine is univalent, form an exceedingly strong argument 
 in favor of the view that iodine is septivalent in these 
 compounds. 
 
 Constitution of Iodic Acid and the Oxygen Acids of 
 Bromine. The conclusion reached in regard to the con- 
 stitution of periodic acid makes it appear highly probable 
 that perchloric acid has a similar constitution, and this 
 view is now commonly accepted, as was stated when the 
 acid was discussed. Applying a similar method to iodic 
 acid, it appears probable that this is derived from the acid 
 I(OH) B by loss of water : 
 
 = I0 2 (OH) + 2H 2 0; or 
 
 r OH 
 OH (O 
 
 !<! OH = I-J O + 2H a O. 
 
 OH (OH 
 
 OH 
 
 The iodine is regarded as quinquivalent in both forms of 
 the acid. This is represented in the case of the acid 
 O 2 I(OH) by the structural formula 
 
 in ; 
 
 The corresponding compound of bromine is regarded as- 
 having the same constitution as the iodine compound. 
 
 Compounds of Iodine with Chlorine. When chlorine 
 is passed over dry crystallized iodine it is absorbed, and 
 a compound of the formula IC1 is formed. This is a thick 
 reddish-brown, very volatile liquid. Under proper con- 
 ditions it solidifies in crystals. Iodine chloride is decom- 
 posed by water, the products being iodic acid, hydro- 
 
FL UORINE: OCCURRENCE- PROPERTIES. 17? 
 
 chloric acid, and free iodine, as stated under lodic Acid 
 (p. 171). 
 
 If the passage of chlorine over iodine is continued be- 
 yond the point required for the formation of the simple 
 compound IC1, the trichloride IC1 3 is formed. This is 
 a crystallized compound of a yellow color. When heated 
 it breaks down into chlorine and iodine monochloride. 
 When treated with water it is partly dissolved without 
 decomposition, but it is partly decomposed, yielding iodic 
 acid, iodine monochloride, and hydrochloric acid. 
 
 Compound of Iodine with Bromine. There is only one 
 compound of iodine and bromine known, and that is the 
 one having the formula IBr. It is a crystallized com- 
 pound which is formed by direct combination of the two 
 elements. It is decomposed by heat and by water. 
 
 FLUORINE, F (At. Wt. 18.91). 
 
 Occurrence. This element occurs in large quantity in 
 nature, and is widely distributed, but it is always in com- 
 bination with other elements. It is found chiefly in com- 
 bination with calcium, as fluor-spar or calcium fluoride, 
 CaF 2 , and in combination with sodium and aluminium, as 
 cryolite, a mineral which occurs abundantly in Greenland 
 and has the composition represented by the formula 
 Na 3 AlF 6 or AlF 8 .3NaF. It is called fluorine from the fact 
 that it occurs in fluor-spar, which in turn receives its 
 name for the reason that it melts when heated and is 
 therefore used as a flux in heating chemical substances 
 together (from fluo, I flow). On account of the remark- 
 able affinity of fluorine for other elements, all attempts 
 to prepare it in the free condition failed until a few years 
 ago, when its isolation was effected by passing an electric 
 current through liquid hydrofluoric acid contained in a 
 platinum vessel. 
 
 Properties. Fluorine is a light greenish-yellow gas, 
 that has recently been converted into a liquid at a very low 
 temperature. It acts upon almost all substances. Thus, 
 it decomposes water, yielding ozone and hydrofluoric 
 acid ; it combines directly with hydrogen at the ordinary 
 
178 INORGANIC CHEMISTRY. 
 
 temperature; and with sulphur, phosphorus, iron, etc., 
 with evolution of light and heat. It does not, however, 
 act upon platinum. Owing to its active properties it is 
 of course a difficult matter to isolate and preserve it. 
 
 Hydrofluoric Acid, HF. Hydrofluoric acid is made 
 by treating a fluoride with sulphuric acid. Thus, when 
 calcium fluoride or fluor-spar is used, this reaction takes 
 place : 
 
 CaF 2 + H 2 S0 4 = CaSO 4 + 2HF. 
 
 The reaction must be performed in vessels of platinum 
 or lead, as glass is disintegrated by the acid. In perfectly 
 pure anhydrous condition it can be obtained by heating 
 the pure dry salt KHF 2 , known as acid potassium fluor- 
 ide. It is a liquid which boils at 19.4 and does not so- 
 lidify even at a very low temperature. The pure dry acid 
 in the liquid form does not act upon glass. It does not 
 clissolve the acid-forming elements, but does dissolve 
 most of the base-forming elements with evolution of hy- 
 drogen and formation of fluorides. The gas acts upon 
 the skin, causing swellings and violent pains. Inhaled it 
 is poisonous. To preserve it, vessels of platinum or caout- 
 chouc must be used. In the moist condition it attacks 
 glass, converting the . silicon into the fluoride, SiF 4 , and 
 the metals into their fluorides. A silicate of the formula 
 CaSiO 3 would undergo the changes represented in the 
 following equation : 
 
 CaSiO 3 + 6HF = CaF 2 + SiF 4 + 3H 2 O. 
 
 Silicon fluoride is a gas, and calcium fluoride is soluble 
 in acids. Thus calcium silicate, which is insoluble in 
 water, is so changed by hydrofluoric acid as to be ren- 
 dered soluble. In a similar way glass, which is a com- 
 pound resembling calcium silicate, is rendered soluble, 
 or is, as we commonly say, dissolved, by hydrofluoric 
 acid. 
 
 When an aqueous solution of hydrofluoric acid is boiled 
 it passes over at 120, and the distillate contains 36 to 38 
 per cent of the acid. 
 
HYDROFLUORIC ACID THE FLUORIDES. 179 
 
 Hydrofluoric acid is used for the purpose of etching 
 glass, particularly for marking scales on thermometers 
 and other graduated glass instruments. The glass is 
 covered with a thin layer of wax or paraffin and, at the 
 places where the etching is wanted, marks are made 
 through the paraffin, so that the glass is exposed. Those 
 parts of the glass which are covered are not acted upon 
 by the hydrofluoric acid, while those parts which are not 
 covered are corroded and, when the paraffin is removed, 
 permanent marks are found corresponding to those made 
 through the paraffin. A solution of hydrofluoric acid in 
 water is manufactured and sold in rubber bottles. 
 
 The specific gravity of hydrofluoric acid gas at about 
 100 leads to the molecular weight corresponding to the 
 formula HF, fluorine having the atomic weight 19. At 
 about 30 the specific gravity corresponds to the for- 
 mula H 2 F 2 . At lower temperatures the molecular weight 
 appears to be still greater. In solutions of ordinary 
 concentration in water the substance appears to have 
 the formula H 2 F 2 , while when the solution is much 
 diluted the formula becomes HF. 
 
 Constitution of Hydrofluoric Acid and the Fluorides. 
 Hydrofluoric acid forms two series of salts correspond- 
 ing to the two general formulas MHF 2 and M 2 F 2 or MF. 
 The former, of which the salt KHF 2 is an example, are 
 called acid fluorides, the latter simply fluorides. The 
 fluorides are commonly represented by the simpler 
 general formula MF, though it appears probable that 
 the doubled formula is more correct. It will be seen 
 later that fluorine forms a large number of so-called 
 double salts or double fluorides, which it is difficult to 
 explain in any other way than that they are derived from 
 the acid H 2 F 2 . Thus cryolite, to which reference has 
 been made, is called a double fluoride of aluminium and 
 sodium, and is generally expressed by the formula 
 A1F 3 . 3NaF, which means simply that in some way alu- 
 minium fluoride is combined with three molecules of 
 sodium fluoride ; but it is difficult to see how this union 
 can be effected without assuming that fluorine has a 
 greater valence than one. If hydrofluoric acid has the 
 
180 INORGANIC CHEMISTRY. 
 
 formula H 2 F 2 , its constitution is probably this: H-F-F-H; 
 or possibly H-F=F-H. In the one case the fluorine 
 is represented as bivalent, in the other as trivalent, but 
 we have no evidence in favor of either view as opposed 
 to the other. Still it is generally observed that when 
 the valence of an element varies, it changes from odd to 
 odd or from even to even. Thus in the case of the 
 oxygen acids of chlorine, it appears that the valence of 
 chlorine varies from 1 to 3 to 5 to 7. Similarly the 
 valence of sulphur varies from 2 to 4 to 6, etc. For this 
 reason the view that fluorine is trivalent in hydrofluoric 
 acid is perhaps to be preferred to the simpler view that 
 it is bivalent. If then the constitution of hydrofluoric 
 acid be expressed thus, H-F=F-H, the formation of the 
 so-called double fluorides is not difficult to understand. 
 The double fluoride above referred to, viz., cryolite, has 
 probably the constitution represented bv the formula 
 
 /F-F-Na 
 Al^ F=F-Na, and the other double fluorides are to be 
 
 \F=F-Na 
 regarded as having a similar constitution. 
 
 Compound of Fluorine with Iodine. The only com- 
 pound of fluorine with the members of the chlorine group 
 is iodine pentafluoride, IF 6 .* This is a liquid which is 
 formed by treating silver fluoride, AgF, with iodine. 
 Water decomposes it, forming iodic acid : 
 
 IF 5 + 3H 2 O = 5HF + HIO 3 . 
 
 Considering the compounds which the halogens form 
 with one another, it appears that iodine combines with 
 bromine to form the compound IBr, with chlorine it 
 forms IC1 3 , and with fluorine IF 5 ; or its valence towards 
 bromine is 1, towards chlorine 3, and towards fluorine 
 5. The farther removed in the series the element is 
 from iodine the greater is the valence of iodine for it. 
 
 * There seems to be some doubt in regard to the existence of this 
 compound. 
 
RELATIVE AFFINITIES OF THE CHLORINE GROUP. 181 
 
 Tabular Presentation of the Compounds of the Members 
 of the Chlorine Family with Hydrogen, with Oxygen, with 
 Hydrogen and Oxgen, and with One Another. 
 
 Compounds with Hydrogen. 
 HF(H 2 F 2 ) HC1 HBr HI 
 
 Compounds with Oxygen. 
 
 C1 2 O 
 
 CIA 
 
 C1O 2 
 
 Compounds ivith Hydrogen and Oxygen. 
 
 HC10 HBrO 
 
 HC10 2 
 
 HC10 3 HBrO 3 HIO 3 
 
 HC10 4 HI0 4 (H 5 I0 6 ) 
 
 Compounds with One Another. 
 
 ClBr 101, IBr 
 IC1 3 ,IF 6 
 
 Relative Affinities of the Elements of the Chlorine 
 Group. The difference between the affinities of these 
 elements, which has already been commented upon, is 
 illustrated in a number of ways. From iodides, chlorine 
 and bromine set iodine free ; and from bromides, chlorine 
 sets bromine free. When chlorine is added to a solution 
 containing a bromide and an iodide, it first sets the 
 iodine free, and forms the corresponding chloride. Thus, 
 in the case of potassium iodide : 
 
 2KI + 01, = 2KC1 + I 2 . 
 
 After this reaction is complete, the chlorine acts upon 
 the water, decomposing it, oxidizing the iodine to iodic 
 acid (which see). The solution is then colorless. After 
 all the iodine is converted into iodic acid the bromine 
 is liberated and colors the solution yellowish red. 
 
1.82 INORGANIC CHEMISTRY. 
 
 Again, as we shall see, there are some oxidizing agents 
 which decompose iodides but which do not decompose 
 bromides and chlorides, and others which decompose 
 chlorides but do not decompose bromides and iodides. 
 
 FAMILY VII, GROUP A MANGANESE. 
 
 There is one element which belongs in the same family 
 as those which have just been treated, and resembles 
 them in some respects ; but at the same time it differs 
 from them quite markedly in other respects. This is 
 manganese. It acts in fact in two different ways, and is 
 one of those elements, already referred to, which are 
 both acid-forming and base-forming. Some of its com- 
 pounds with hydrogen and oxygen are distinctly acid, 
 others are distinctly basic. So far as it acts like the 
 members of the chlorine family a brief reference to it 
 here is desirable. On the other hand, it will be dealt 
 with chiefly in connection with those base-forming 
 elements which it most resembles, as, for example, 
 iron. 
 
 Manganese occurs in nature principally in the form of 
 pyrolusite or manganese dioxide, MnO 2 , also known as 
 the black oxide of manganese. It forms with oxygen 
 compounds of the following formulas : MnO, Mn 2 O 3 , 
 Mn 3 O 4 , MnO 2 , and Mn 2 O 7 . When a compound of man- 
 ganese is subjected to the influence of powerful oxidizing 
 agents in the presence of an alkali it is converted into a 
 salt of manganic acid, H 2 MnO 4 , which in its composition 
 resembles sulphuric acid. If the salt of manganic acid 
 thus obtained is dissolved in water it undergoes partial 
 decomposition, which is complete if the solution is boiled, 
 or if carbon dioxide is passed through it. The change 
 consists in the transformation of manganic acid into per- 
 manganic acid, HMnO 4 : 
 
 3H 2 MnO 4 = 2HMnO 4 + MnO 2 + 2H 2 O ; or 
 3K 2 MnO 4 + 2H 2 = 2KMnO 4 + MnO, + 4KOH. 
 
 Permanganic acid, HMnO 4 , is a compound which in 
 many respects resembles perchloric acid. It can be ob- 
 
MANGANESE. 183 
 
 tained in water solution by decomposition of certain 
 of its salts, but like perchloric acid it is easily decom- 
 posed. In consequence of the ease with which it gives 
 up oxygen it is a good oxidizing agent, and is extensively 
 used in the laboratory in this capacity. It is employed in 
 the form of the potassium salt, potassium permanganate, 
 KMnO 4 , which will receive special attention under the 
 head of Manganese Compounds. In order, however, to 
 make clear the difference in conduct between perchloric 
 and permanganic acids a few characteristic facts will be 
 mentioned here. The conduct of permanganic acid and 
 of potassium permanganate will be understood, if it is 
 borne in mind that in the presence of substances of 
 strongly acid character manganese tends to act as a base- 
 forming element, and in this capacity to form salts with 
 the acids. Thus in presence of sulphuric acid potassium 
 permanganate forms potassium sulphate, manganous 
 sulphate, and oxygen, if there is anything present which 
 has the power to take up oxygen. In the salts in which 
 it plays the part of a metal manganese is generally biva- 
 lent. With hydrochloric acid, as we have already seen 
 in studying the action of hydrochloric acid upon manga- 
 nese dioxide, it forms the chloride MnCl 2 . When now 
 potassium permanganate is treated with hydrochloric acid 
 it is decomposed according to the following equation : 
 
 2KMnO 4 + 6HC1 = 2KC1 + 2MnCl 2 + 3H 2 O + 5O. 
 
 Similarly, with sulphuric acid manganous sulphate is 
 formed, thus : 
 
 2KMnO 4 + 3H 2 S0 4 = K 2 SO 4 + 2MnSO 4 + 3H 2 O + 5O. 
 
 Such reactions do not take place with perchloric acid, as 
 chlorine is entirely lacking in the power to enter into 
 acids in the place of the hydrogen and form salts. 
 
 Manganese forms some other acids besides perman- 
 ganic acid, but they exhibit little or no analogy with 
 compounds of chlorine, and their study will therefore be 
 postponed until manganese is taken up. The point of 
 
184 INORGANIC CHEMISTRY. 
 
 chief interest to be noted here is that this element is un* 
 mistakably like chlorine in its highest oxygen com- 
 pounds, but entirely different from it in most of its com- 
 pounds. The compound manganese heptoxide, Mn 2 O 7 , 
 stands in the relation of an anhydride to permanganic 
 acid. In water solution it passes over into the acid : 
 
 Mn,O 7 + H 2 O = 2HMnO 4 . 
 
 It is formed by treating potassium permanganate with 
 the most concentrated sulphuric acid : 
 
 2KMn0 4 + H 2 S0 4 = K 2 SO 4 + Mn 2 O 7 + H 2 O. 
 
 It is extremely unstable, giving up oxygen readily. In 
 contact with organic substances or other substances 
 which have the power to take up oxygen it decomposes 
 so rapidly as frequently to lead to explosions. It is of 
 interest to note that this is the only oxide of Family YII 
 in which the maximum valence of 7 is shown. 
 
 Judging by analogy, it seems probable that the consti- 
 tution of permanganic acid is like that of periodic and 
 perchloric acids, and is represented by the formula 
 
 O 
 
 ll 
 
 O=Mn-0-H, in which the manganese is septivalent. 
 I) 
 O 
 
CHAPTER XIII. 
 
 THE ELEMENTS OF FAMILY VI, GROUP B : 
 SULPHUR SELENIUM TELLURIUM. 
 
 Introductory. The elements of this group bear to oxy- 
 gen a relation somewhat similar to that which the ele- 
 ments of Group B, Family VII, bear to fluorine. The 
 three members sulphur, selenium, and tellurium resem- 
 ble one another fully as strikingly as chlorine, bromine, 
 and iodine do. Their compounds bear a general resem- 
 blance, to those of oxygen, and yet they form very char- 
 acteristic compounds with oxygen, while oxygen forms no 
 analogous compounds with any of them. Just as iodine 
 forms a compound with fluorine of the formula IF 5 , but 
 fluorine does not form with iodine a compound FI 5 , so 
 sulphur, selenium, and tellurium form with oxygen the 
 -compounds SO 3 , SeO 3 , and TeO 3 , while oxygen does not 
 form analogous compounds with these other elements. 
 The valence of the elements of this group towards hydro- 
 gen is 2, as shown in the compounds H 2 O, H 2 S, H 2 Se, and 
 H 3 Te. Of oxygen and sulphur there are other hydrogen 
 compounds, as hydrogen dioxide, H 2 O 2 , and an analogous 
 compound of sulphur, but there is great uncertainty in 
 regard to the constitution of the latter compound, and 
 practically nothing is known in regard to the valence of 
 sulphur in it. In general the hydrogen valence is con- 
 stant. Towards the members of the chlorine group the 
 valence varies from 2 to 6. Oxygen never exhibits a 
 higher valence than 2 towards chlorine and its analogues. 
 The compound OC1 2 illustrates the bivalence of oxygen 
 towards chlorine. Sulphur forms with chlorine the com- 
 pounds S 2 C1 2 and SC1 2 , which are analogous to the hy- 
 drogen compounds H 2 S 2 and H 2 S, and in both of them 
 the sulphur is probably bivalent. It also forms the com- 
 pound SC1 4 , in which it is quadrivalent. With iodine it 
 
 (185) 
 
186 INOEGANIC CHEMISTRY. 
 
 is said to form the unstable compounds S 2 I 2 and SI 6 . Se- 
 lenium and tellurium form similar compounds, and, in 
 general, the stability of the compounds of the members 
 of the sulphur group with the members of the chlorine 
 group increases in the order sulphur, selenium, tellurium. 
 Sexivalence of these elements towards members of the 
 chlorine group is rare, being shown only in the com- 
 pound SI 8 , and even this is doubtful. 
 
 Towards oxygen the three elements of the sulphur 
 group are quadrivalent and sexivalent, as seen in the 
 compounds SO 2 , SeO 2 , TeO a , and SO 3 , SeO 3 , and TeO 3 . Of 
 course, it is possible that in these oxygen compound o 
 the elements are bivalent. Thus, sulphur dioxide, SO 2 , 
 
 may be represented by the formula S^ I , in which 
 
 both the oxygen and the sulphur appear as bivalent ; and, 
 in a similar way, the trioxide may be represented by the 
 
 O 
 formula S O ; but the only reason for doing this is 
 
 O 
 
 the desire to represent sulphur as always bivalent. The 
 existence of the compounds SC1 4 , SeCl 4 , and SI 6 cannot 
 be explained, however, on the assumption that sulphur 
 is bivalent, and the simplest view which can be taken of 
 the matter is that the members of the sulphur group are 
 in general bivalent towards hydrogen ; bivalent, quadri- 
 valent, and, exceptionally, sexivalent towards the mem- 
 bers of the chlorine group ; and quadrivalent and sexi- 
 valent towards oxygen. Towards hydroxyl the valence 
 of the members of the sulphur group appears to vary 
 from 4 to 6. The quadrivalence is shown in the com- 
 pound hydrosulphurous acid, H 2 SO 2 , which probably has 
 H 
 
 the constitution O=S-O-H ; the sexi valence is seen in 
 
 O 
 
 II 
 sulphurous acid, H 2 SO 3 , or H-S-O-H, and in sulphuric 
 
 II 
 O 
 
SULPHUR. 187 
 
 acid, S(OH) 6 , or in the ordinary form H 2 S0 4 , or 
 O 
 
 H-O-S-O-H. 
 
 II 
 O 
 
 Of the three elements of this group sulphur occurs in 
 greatest abundance in nature, selenium next in order, 
 and finally tellurium. Just as bromine frequently 
 accompanies chlorine, so selenium frequently accompa- 
 nies sulphur, but it is always present in much smaller 
 quantity than sulphur. Tellurium occurs in very small 
 quantity relatively, and not uncommonly in combina- 
 tion with valuable metals like gold and silver. Large 
 quantities of sulphur are found in the native or uncom- 
 bined condition. Only extremely small quantities of 
 selenium and tellurium are found native. 
 
 SULPHTJK, S (At. Wt. 31.83). 
 
 Occurrence. The principal deposits of native sulphur 
 are found in Sicily, Italy, and Spain. In California also 
 there is a considerable deposit. In general, sulphur is 
 likely to be found near dying or extinct volcanoes. 
 Sulphur occurs in nature, further, in the form of the 
 hydrogen compound, hydrogen sulphide, H 2 S, issuing 
 from the earth in volcanic regions, and in solution in 
 some natural waters, known as "sulphur waters." The 
 oxide SO 2 is likewise found issuing from the earth in 
 volcanic regions. Compounds of sulphur with metallic 
 elements, as with iron, copper, lead, zinc, are very abun- 
 dant. Such compounds are iron pyrites, FeS 2 ; copper 
 pyrites, CuFeS 2 ; galenite, PbS ; and zinc blende, ZnS. 
 Some sulphates are widely distributed and occur in 
 large quantities ; for example, gypsum or calcium sul- 
 phate, CaSO 4 + 2H 2 O ; barium sulphate, or heavy spar, 
 BaSO 4 ; lead sulphate, PbSO 4 . Finally, sulphur occurs 
 in a few animal and vegetable products in combination 
 with carbon, hydrogen, and, generally, with nitrogen. 
 
 Extraction of Sulphur from its Ores. By far the largest 
 quantity of sulphur found in the market is taken from 
 
188 INORGANIC CHEMISTRY. 
 
 the mines in Sicily. Of these mines there are between 
 250 and 300. When taken from the mines it is mixed 
 with many earthy substances, from which it must be 
 separated. The separation is generally effected in Sicily 
 by piling the ore so as to leave passages for air, covering 
 the piles with earthy matter to prevent free access of air 
 and then setting fire to them. A part of the sulphur 
 burns, and this causes the rest of the sulphur to melt. 
 The molten sulphur runs down to the bottom of the 
 pile, and by a proper arrangement it is drawn off from 
 time to time. If the pile of ore were not covered up, 
 and the air were allowed to enter freely, the sulphur 
 would burn up, and be converted into the gas sulphur 
 dioxide. The " crude brimstone " obtained in the manner 
 described is afterwards refined by distillation, and it is 
 this refined or distilled sulphur which is met with in the 
 market under the names " roll brimstone " and " flowers 
 of sulphur." 
 
 The distillation is carried on in retorts made of earth- 
 enware, and these are connected with large chambers of 
 brick-work. When the vapor of sulphur first comes into 
 the condensing- chamber it is suddenly cooled, and hence 
 deposited in the form of a fine powder. This is what is 
 called " flowers of sulphur." After the distillation has 
 continued for some time the vapor condenses in the form 
 of a liquid, which collects at the bottom of the chamber. 
 This is drawn off into slightly conical wooden moulds, 
 and takes the form of "roll brimstone" or "stick sul- 
 phur." 
 
 Some sulphur is obtained from iron pyrites by heating 
 in closed vessels. The change which takes place on 
 heating iron pyrites is perfectly analogous to that which 
 takes place on heating manganese dioxide in preparing 
 oxygen. A sulphur compound of the formula Fe 3 S 4 and 
 free sulphur are formed in the former case, as the com- 
 pound of manganese and oxygen, Mn 3 O 4 , and free oxygen 
 are formed in the latter : 
 
 3FeS 2 = Fe 3 S 4 +S 2 ; 
 3MnO 2 = Mn 3 O 4 + O 2 . 
 
PROPERTIES OF SULPHUR. 189 
 
 Properties. Sulphur is a yellow, brittle substance 
 which at 50 is almost colorless. It melts at 114.5, 
 forming a thin, straw-colored liquid. When heated to 
 a higher temperature it becomes darker and darker in 
 color, and at 200 to 250 it is so viscid that the vessel 
 in which it is contained may be turned upside down with- 
 out danger of its running out. Finally, at 448.4 it boils, 
 and is then converted into a brownish-yellow vapor. 
 When molten sulphur solidifies, or when it is deposited 
 from a solution, it takes the form of crystals. But, 
 strange to say, the crystals formed from molten sulphur 
 are entirely different from those deposited from cold 
 solutions of sulphur. The latter belong to the rhombic 
 system. They are octahedrons with a rhombic base. 
 And this is the form of the sulphur found in nature. The 
 former are honey-yellow needles. An examination of 
 these needles shows that the angles which their faces 
 form with one another are not the same as the angles 
 formed by the faces of the octahedrons, and that they 
 belong to an entirely different system the monoclinic. 
 
 The rhombic crystals of sulphur can be made by dis- 
 solving " roll brimstone" in carbon disulphide and allow- 
 ing the solution to stand. When the liquid has suffi- 
 ciently evaporated, the sulphur will appear in larger or 
 smaller rhombic crystals, according to the conditions. 
 A comparison of the crystals thus obtained with natural 
 crystals will show that the two have identical or very 
 similar forms. The formation of the needles or mono- 
 clinic crystals may be shown by melting a considerable 
 quantity, say a pound or two, of roll brimstone in a sand 
 or Hessian crucible, and allowing the liquid mass to cool 
 slowly. When a thin crust has formed on the surface 
 this should be perforated, and the remainder of the 
 liquid sulphur poured off. The crucible will then be 
 found lined with long, dark yellow, lustrous needles 
 which do not look at all like those obtained from the 
 solution in carbon disulphide. If the monoclinic needles 
 a-re allowed to lie unmolested they gradually undergo 
 change spontaneously. They lose their lustre and be- 
 come lighter in color ; and now, if examined carefully, they 
 
190 INORGANIC CHEMISTRY. 
 
 are found to consist of minute crystals like those found 
 in nature. They have changed to the rhombic form. It 
 is evident, therefore, that the arrangement of the parti- 
 cles in the monoclinic crystals of sulphur is not a stable 
 one. The change is accompanied by a considerable evo- 
 lution of heat. 
 
 Substances which crystallize in two distinct forms are 
 called dimorphous. We shall see that carbon also crys- 
 tallizes in two different forms, and that this kind of phe- 
 nomenon is met with not unfrequently among chemical 
 compounds. The difference between the two varieties 
 of sulphur suggests that observed between the two 
 forms of oxygen. Whether the explanation is the same 
 in the two cases is doubtful. It appears more probable 
 that the difference in the former case is due to different 
 arrangements of the molecules in the crystals, rather 
 than to different arrangements of the atoms in the mole- 
 cules. The chemical properties of the two varieties are 
 practically identical. This could hardly be the case if 
 the number of atoms in the molecule were different in 
 the two cases. 
 
 Besides the two forms mentioned sulphur can also be 
 obtained in the amorphous, or uncrystallized, condition. 
 If molten sulphur is quickly cooled under water it re- 
 mains for some time soft and dough-like, and while in 
 this condition it is amorphous. If allowed to stand it 
 gradually becomes hard and brittle. 
 
 When separated from certain compounds which are 
 in solution in water, the sulphur is in a very finely 
 divided condition, and gives the liquid the appearance of 
 milk. This is seen on adding hydrochloric acid to a 
 solution of sodium thiosulphate, or hyposulphite, as it 
 is generally called. This substance has the formula 
 Na 2 S 2 O 3 , and the reactions which take place between it 
 and hydrochloric acid are these : 
 
 Na 2 S 2 3 + 2HC1 = H 2 S 2 3 + 2NaCl ; 
 H,S 2 3 = SO, + H 2 + S. 
 
 On treating certain varieties of sulphur with carbon 
 disulphide they are found to dissolve completely. This 
 
PROPERTIES OF SULPHUR. 191 
 
 is true, for example, of the natural crystals and of those 
 made artificially by deposition from a solution in carbon 
 disulphide. On the other hand, sulphur in the form of 
 " flowers of sulphur " is only partly soluble in the liquid. 
 There are therefore two forms of sulphur to be dis- 
 tinguished between, the soluble and the insoluble. The 
 cause of the difference between these modifications is 
 not known. " Stick sulphur " is mostly soluble, while 
 in the " flowers of sulphur " there is at times a consid- 
 erable percentage of the insoluble variety. Sulphur is 
 insoluble in water, and slightly soluble in alcohol and 
 ether. 
 
 Sulphur is a much less active element chemically than 
 the members of the chlorine group, and also less active 
 than oxygen. Generally speaking, however, it conducts 
 itself like oxygen. It combines directly though not 
 easily with hydrogen, and it combines readily with most 
 metals, forming compounds called sulphides which, so far 
 as their composition is concerned, are analogous to the 
 oxides. Thus when heated together with iron, copper, 
 or lead, combination takes place readily with evolution 
 of heat and light. In its chemical conduct it differs 
 markedly from the members of the chlorine group. 
 When heated to a sufficiently high temperature in the air 
 or in oxygen, sulphur forms the compound sulphur 
 dioxide, SO 2 , and, under certain conditions which will be 
 described farther on, this combines with more oxygen 
 to form the trioxide SO 3 . Further, it combines with 
 nearly all the acid-forming elements if heated with them 
 to a sufficiently high temperature. It combines with 
 most other elements less readily than oxygen, and it 
 forms less stable compounds. Thus, its compound with 
 hydrogen is decomposed very much more readily into its 
 elements than water is ; and the sulphides or its com- 
 pounds with metals are decomposed when they are 
 heated in oxygen, the oxygen displacing the sulphur, 
 much as chlorine displaces bromine ; though there is a 
 difference between the two cases to be found in the fact 
 that sulphur itself unites readily with oxygen, and this 
 facilitates the decomposition of the sulphides by the 
 action of oxygen. 
 
192 INORGANIC CHEMISTRY. 
 
 When treated with powerful oxidizing agents sulphur 
 is converted into sulphuric acid. Thus the action of 
 concentrated nitric acid takes place in the main accord- 
 ing to the equation 
 
 2HNO 3 + S = H 2 SO 4 + 2NO. 
 
 The action of sulphur upon the so-called caustic al- 
 kalies, sodium and potassium hydroxides, is somewhat 
 like that of chlorine, bromine, and iodine. It will be 
 remembered that with potassium hydroxide chlorine 
 forms potassium chloride and potassium chlorate or 
 hypochlorite, according to the concentration and tem- 
 perature of the solution. When sulphur acts upon 
 sodium hydroxide the sulphide is formed, but oxygen is 
 thus rendered available and some of it combines with the 
 sulphide, and sulphur also combines with a part of the 
 sulphide. The principal reactions involved are : 
 
 (1) 2NaOH + S 
 
 (2) Na 2 S + 4S 
 
 (3) Na 2 S + S -f 30 = Na 2 S 2 O 3 . 
 
 Expressing these reactions in one equation we have 
 GNaOH + 12S = 2Na 2 S & + Na 2 S 2 O 3 + 3H 2 O. 
 
 The action is of the same general character as that 
 which takes place in the case of chlorine, but differs from 
 it in the fact that sodium sulphide, Na 2 S, has the power 
 to take up sulphur, and also to take up sulphur and 
 oxygen. 
 
 The specific gravity of the vapor of sulphur varies with 
 the temperature, and is such as to indicate that at tem- 
 peratures not far above the boiling point the molecule 
 consists of eight atoms, while at the temperature 800 
 and higher the molecule consists of two atoms. This 
 suggests the question whether the molecule of sulphur 
 in the solid form may not be more complex than the 
 condition represented by the symbol S b . We have no 
 means of answering this question with any certainty at 
 present. 
 
COMPOUNDS OF SULPHUR WITH HYDROGEN. 193 
 
 (Jses of Sulphur. Enormous quantities of sulphur are 
 used in the manufacture of sulphuric acid, and of gun- 
 powder. It is also used in the manufacture of fire-works 
 of various kinds. Burning sulphur gives sulphur dioxide, 
 which is extensively employed for bleaching wool, silk, 
 straw, etc. When caoutchouc is thoroughly mixed with 
 sulphur or some sulphur compound it becomes vulcan- 
 ized. 
 
 COMPOUNDS -OF SULPHUE WITH HYDEOGEN. 
 
 The principal compound of sulphur and hydrogen is 
 analogous in composition to water. It is known as hy- 
 drogen sulphide or sulphuretted hydrogen, and has the 
 formula H 2 S. Besides this there is at least one other 
 compound which contains a larger proportion of sulphur, 
 and probably has the composition H 2 S 2 . There are 
 reasons for supposing, further, that still more complex 
 compounds can exist, but owing to their instability it is 
 impossible to isolate them in pure condition and study 
 them. 
 
 Hydrogen Sulphide, Sulphuretted Hydrogen, H 2 S. 
 When hydrogen is passed over highly heated sulphur 
 the two elements combine to form hydrogen sulphide. 
 The action is, however, quite incomplete and is not to be 
 compared with that which takes place when hydrogen and 
 oxygen are heated together. This compound of sulphur 
 and hydrogen occurs in nature in solution in the so-called 
 " sulphur waters," which are met with in many parts of 
 the world. It is formed by heating organic substances 
 which contain sulphur, just as water is formed by heat- 
 ing organic substances which contain oxygen. It is 
 formed, further, by decomposition of organic substances 
 which contain sulphur, as, for example, the albumen of 
 eggs. The odor of rotten eggs is partly due to the for- 
 mation of hydrogen sulphide. It is formed by the action 
 of acids upon sulphides or hydrosulphides, just as water 
 is formed by the action of acids upon oxides or hydrox- 
 ides (see p. 132). Thus hydrochloric acid and ferrous 
 sulphide, FeS, give ferrous chloride, FeCl 2 , and hydro- 
 gen sulphide : 
 
194 INORGANIC CHEMISTRY. 
 
 FeS + 2HC1 = FeCl a + H 2 S ; 
 
 just as ferrous oxide, FeO, and hydrochloric acid give 
 ferrous chloride and water : 
 
 FeO + 2HC1 = FeCl 2 + H 2 O. 
 
 So also potassium hydroxide and potassium hydrosul- 
 phide act in the same way, as has been pointed out : 
 
 KSH + HC1 = KC1 + H 2 S ; 
 KOH + HC1 = KC1 +'H 2 0. 
 
 It is generally formed by the action of nascent hydro- 
 gen upon sulphur compounds. Thus, it has been shown 
 that the hydrogen from hydriodic acid has the power to 
 reduce sulphuric acid to hydrogen sulphide : 
 
 H 2 S0 4 + 8HI = H 2 S + 4H 2 + 81. 
 
 In the laboratory, where the gas is extensively used, 
 it is generally prepared from ferrous sulphide, FeS, and 
 dilute sulphuric acid, which are simply brought together 
 at the ordinary temperature in a flask such as is used in 
 making hydrogen. The reaction is like that which takes 
 place between ferrous sulphide and hydrochloric acid. 
 It is represented by this equation : 
 
 FeS + H 2 S0 4 = FeS0 4 + H 2 S. 
 
 Properties. Hydrogen sulphide is a colorless, trans- 
 parent gas of the specific gravity 1.178. It has an ex- 
 tremely disagreeable odor, somewhat suggestive of that 
 of rotten eggs. It is poisonous, even small quantities 
 causing headache, vertigo, nausea, and other bad symp- 
 toms. It burns with a blue flame, forming water and 
 sulphur dioxide : 
 
 H 2 S + 3O = H 2 + S0 2 . 
 
 If, however, the air has not free access, as when the gas 
 is burned in a cylinder open at one end, only a part of 
 the sulphur is burned, the rest being deposited upon 
 the walls of the vessel, while the hydrogen burns. The 
 gas is soluble in water, about three volumes being taken 
 
PROPERTIES OF HYDROGEN SULPHIDE. 195 
 
 up at ordinary temperatures. This solution is used to 
 some extent in the laboratory instead of the gas, but, 
 owing to the fact that it readily undergoes change in 
 consequence of the action of the oxygen of the air, it is 
 not as valuable as it would be if it were more stable. 
 The change consists simply in the oxidation of the hy- 
 drogen and the separation of the sulphur. If a bottle 
 containing a solution of hydrogen sulphide is allowed 
 to stand for a few days, particularly if it is opened from 
 time to time, the odor of the gas will disappear and a 
 deposit of sulphur will be noticed in the bottle. The 
 liquid is then nothing but water. When the solution is 
 boiled it loses all the gas contained in it. 
 
 Hydrogen sulphide is easily decomposed into its ele- 
 ments. It requires a temperature of only a little above 
 400 to effect direct decomposition. In consequence of 
 this instability it causes a number of changes which the 
 analogous compound water cannot effect. The relations 
 here are similar to those which exist between hydro- 
 chloric and hydriodic acids. Hydrochloric acid is very 
 stable, while hydriodic acid breaks down readily into 
 hydrogen and iodine. Therefore hydriodic acid, as we 
 have seen, acts as a reducing agent, while hydrochloric 
 acid does not. So, also, hydrogen sulphide acts as a 
 reducing agent. Thus, if it be passed into concentrated 
 sulphuric acid this reaction takes place : 
 
 H 2 S0 4 + H 2 S = 2H 2 + S + S0 2 . 
 
 The action is to be traced to the decomposition of the 
 hydrogen sulphide into hydrogen and free sulphur, the 
 hydrogen then acting upon the sulphuric acid thus : 
 
 H 2 S0 4 + 2H = 2H 2 + S0 2 . 
 
 With hydriodic acid the reduction may go farther, as 
 has been seen ; with hydrobromic acid, however, the 
 action takes place practically in the same way as with 
 hydrogen sulphide. 
 
 Chlorine, bromine, and iodine act upon hydrogen 
 
196 INORGANIC CHEMISTRY. 
 
 sulphide, setting the sulphur free and uniting with the 
 hydrogen : 
 
 H,S + 01, = 2HC1 + S. 
 
 This reaction suggests the decomposition of water by 
 chlorine, but what a difference there is between the two 
 cases ! Chlorine decomposes water very slowly and 
 only under the influence of the direct sunlight, while it 
 decomposes hydrogen sulphide completely and instantly. 
 Similarly, hydrogen sulphide has the power to abstract 
 chlorine from some of its compounds, as, for example, 
 from ferric chloride, FeCl 3 . When this is treated with 
 nascent hydrogen from any source, it is reduced to fer- 
 rous chloride, FeCl 2 , thus : 
 
 FeCl 3 + H = FeCl 2 + HC1. 
 
 So, also, when it is treated with hydrogen sulphide it 
 is reduced in the same way in consequence of the action 
 of the hydrogen : 
 
 2FeCl 3 + H 2 S = 2FeCl 2 + 2HC1 + S. 
 
 The instability of hydrogen sulphide is further shown 
 by the ease with which it is decomposed by metals with 
 liberation of hydrogen and formation of sulphides. It 
 will be remembered that at high temperatures several 
 metals decompose water, but that at ordinary tempera- 
 tures only a few decompose it easily. Hydrogen sul- 
 phide acts much more readily ; a number of metals which 
 do not act upon water even at high temperatures, as sil- 
 ver, gold, and mercury, decompose this gas at ordinary 
 temperatures. 
 
 Hydrogen sulphide acts upon metallic oxides, convert- 
 ing them into sulphides, as for example : 
 
 CuO + H 2 S = CuS + H 2 O. 
 
 Action of Hydrogen Sulphide upon Solutions of Salts- 
 Use in Chemical Analysis. Hydrogen sulphide is exten- 
 sively used in every chemical laboratory as a reagent in 
 chemical analysis. In order that its action may be 
 understood a few words of explanation are necessary. 
 Sulphur, as we have seen, has a strong affinity for the 
 
HYDROGEN SULPHIDE IN CHEMICAL ANALYSIS. 197 
 
 metallic or base-forming elements, forming with them 
 the sulphides. Further, hydrogen sulphide is easily 
 decomposed, and the substitution of metals for the 
 hydrogen is facilitated by this fact. If now a salt is in 
 solution in water and hydrogen sulphide is passed into 
 the solution there will, of course, be the tendency to the 
 formation of the sulphide of the metal contained in the 
 salt. Thus, suppose the salt in solution is silver nitrate, 
 AgNO 3 . On passing hydrogen sulphide into this solu- 
 tion the silver will tend to combine with the sulphur to 
 form the sulphide, Ag 2 S. If this is formed, hydrogen 
 must be freed from the hydrogen sulphide, and this 
 would probably take the place of the silver in the nitrate, 
 forming nitric acid, according to this equation : 
 
 2AgNO 3 + H 2 S = Ag 2 S + 2HNO 3 . 
 
 If the dilute acid thus formed has the power to decom- 
 pose silver sulphide the sulphide will not be formed ; 
 but if it has not this power the sulphide will be formed, 
 and it will be thrown down or precipitated if it is an 
 insoluble compound. The sulphides of some metals 
 are not decomposed by dilute acids, and are insoluble 
 in water. If hydrogen sulphide is passed through solu- 
 tions of the salts of these metals the sulphides are 
 thrown down. 
 
 Secondly, the sulphides of some metals are decom- 
 posed by dilute acids. Plainly, these cannot be thrown 
 down by simply passing hydrogen sulphide through the 
 solutions of their salts, whether they are soluble or in- 
 soluble in water. Thus, for example, the sulphide of 
 iron, FeS, is insoluble in water, but it is easily decom- 
 posed by dilute acids, and therefore when hydrogen 
 sulphide is passed through a solution of an iron salt the 
 sulphide is not precipitated. In the case of the sulr 
 phate the reaction would be 
 
 FeS0 4 + H 2 S = FeS + H 2 SO 4 . 
 
 But the dilute sulphuric acid would decompose the sul- 
 phide, and the reaction does not take place. If, however, 
 a soluble sulphide is added to a solution of such a metal, 
 
198 INORGANIC CHEMISTRY. 
 
 decomposition takes place. Thus, if, instead of passing 
 hydrogen sulphide, a solution of potassium sulphide is 
 added, reaction takes place thus : 
 
 FeS0 4 -f K 2 S = K 2 S0 4 + FeS. 
 
 Here no sulphuric acid is formed, but, instead of it, 
 neutral potassium sulphate, which does not act upon the 
 insoluble sulphide. Advantage is taken of this fact in 
 chemical analysis, but, in place of potassium sulphide, 
 the analogous compound, ammonium sulphide, (NH 4 ) 2 S, 
 is used. This acts in the same way. Thus, in the case 
 above cited the action with ammonium sulphide is repre- 
 sented by this equation : 
 
 FeSO 4 + (NH 4 ) a S = (NH 4 ) 2 SO 4 + FeS. 
 
 There are several metals which act towards hydrogen 
 sulphide in the same way that iron does. 
 
 Thirdly, there are some metals whose sulphides are 
 soluble in water, and, therefore, if solutions of their salts 
 are treated with hydrogen sulphide or with ammonium 
 sulphide no apparent action takes place. 
 
 Facts like those just referred to form a good basis 
 for the division of the metallic elements into groups for 
 purposes of analysis. According to the above these ele- 
 ments can be divided into three great groups, as follows : 
 
 I. Metals whose sulphides are insoluble in water and 
 not decomposed by dilute acids. This is called the hy- 
 drogen sulphide group. It includes lead, bismuth, silver, 
 mercury, copper, cadmium, gold, platinum, tin, anti- 
 mony, and arsenic. 
 
 II. Metals whose sulphides are insoluble in water but 
 are decomposed by dilute acids. They are therefore 
 not precipitated by hydrogen sulphide, but are precipi- 
 tated by soluble sulphides. As ammonium sulphide is 
 used for the purpose of effecting the precipitation the 
 group is known as the ammonium sulphide group. It in- 
 cludes iron, nickel, cobalt, manganese, thallium, zinc, 
 and uranium. Further, the two elements aluminium and 
 chromium are thrown down with the above, not as sul- 
 
HYDROGEN SULPHIDE IN CHEMICAL ANALYSIS. 199 
 
 phides but as hydroxides, and they are therefore in- 
 cluded in the group. 
 
 III. Metals whose sulphides are soluble in water. 
 This group includes all the metals not included in the 
 above two. 
 
 By taking advantage, then, of the properties of the 
 sulphides of the metals they can be divided into these 
 three groups, and the detection of any particular element 
 is thus facilitated. If hydrogen sulphide is passed 
 through a solution, and a precipitate formed, we know 
 that this can contain only those metals which belong to 
 the hydrogen sulphide group ; and so on. Now, if the 
 precipitate formed with hydrogen sulphide is treated 
 with certain other reagents other changes take place, 
 and by further study it is quite possible, and indeed 
 comparatively simple, to determine which of the mem- 
 bers of the group are present. 
 
 One reaction made use of in further examination of 
 the hydrogen sulphide precipitate is particularly inter- 
 esting in this connection. Under the head of Acids and 
 Bases compounds were referred to which were called sul- 
 phur acids and sulphur bases. Corresponding to the 
 oxygen acid known as arsenic acid, which has the formula 
 H 3 AsO 4 , there are salts which are plainly derived from 
 the corresponding sulphur acid H 3 AsS 4 . Such salts are 
 formed by treating arsenic sulphide with soluble sul- 
 phides, as for example ammonium sulphide : 
 
 As 2 S 5 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 4 . 
 
 So, too, tin forms an oxygen acid, H 2 SnO 3 ; and salts of 
 the corresponding sulphur acid, H 2 SnS 3 , are formed by 
 treating the sulphide of tin with soluble sulphides : 
 
 SnS 2 + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3 . 
 
 Now some of these sulphur salts are soluble in water. 
 This is true of the ammonium salts. So that when 
 the sulphides of metals which have the power to form 
 salts of this kind are treated with ammonium sul- 
 phide they pass into solution. Of the metals of the 
 
200 INORGANIC CHEMISTRY. 
 
 hydrogen sulphide group only arsenic, antimony, and 
 tin have this power ; so that, if the hydrogen sul- 
 phide precipitate is treated with ammonium sulphide, 
 arsenic, antimony, and tin sulphides are dissolved if 
 present, whereas the other sulphides are not changed 
 by this treatment. Thus a means'is afforded of subdivid- 
 ing the hydrogen sulphide group into two sub-groups: 
 (a) Metals whose sulphides are insoluble in ammonium 
 sulphide ; and (6) metals whose sulphides are soluble 
 in ammonium sulphide. 
 
 Hydrosulphid.es. The action of hydrogen sulphide 
 shows that it belongs to the class of acids. When it 
 acts upon an oxide the corresponding sulphide and 
 water are formed. But just as there are sulphides 
 which are derived from hydrogen sulphide by the re- 
 placement of both hydrogen atoms by metallic elements, 
 so there are hydrosulphides which are derived from it 
 by the replacement of only one of the two atoms of hy- 
 drogen in the molecule. The sulphides correspond to 
 the oxides, and the hydrosulphides to the hydroxides. 
 Thus the analogous oxygen and sulphur compounds of 
 potassium are : 
 
 K 2 
 
 KOH KSH. 
 
 We speak of sulphides and hydrosulphides as salts of hy- 
 drogen sulphide. In consequence of this ability to form 
 salts in the same way in general that acids do, hydrogen 
 sulphide is sometimes called sulphydric acid, and the salts 
 of the formula MSH, sulphydrates. The name sulphydrate 
 is analogous to hydrate, which, as has been pointed out, 
 is used by some to designate the compounds of the for- 
 mula MOH or the hydroxides. Between the acid, hydro- 
 gen sulphide, and the neutral compound, water, there is no 
 fundamental difference. The difference is simply one of 
 degree. In the present system of chemistry, which is 
 largely an oxygen system, water is regarded as the con- 
 necting link between acids and bases, as was shown on 
 p. 135. But we might with equal right base our defi- 
 
HYDROGEN PERSULPHIDE. 201 
 
 nitions and conceptions of acids and bases upon the 
 conduct of sulphur compounds, and thus build up a 
 sulphur system. In such a system hydrogen sulphide 
 would be the connecting link between acids and bases. 
 
 Hydrogen Persulphide, H a S a (P). The sulphides of cer- 
 tain metals, particularly the so-called alkali metals, so- 
 dium and potassium, combine with sulphur to form the 
 polysulphides, examples of which are K 2 S 2 , K 2 S 3 , K 2 S 4 , 
 and K 2 S 5 . When these are decomposed with dilute acids 
 compounds of hydrogen and sulphur are formed. It 
 has thus far, however, been impossible to determine 
 whether more than one such compound is formed, for 
 the reason that there is no means of deciding whether 
 the substances formed are chemical compounds or mere 
 mixtures of sulphur and some one compound of sulphur 
 and hydrogen. Hydrogen persulphide is a liquid with a 
 very disagreeable odor. Just as hydrogen dioxide de- 
 composes readily into water and oxygen, so hydrogen 
 persulphide decomposes readily into hydrogen sulphide 
 and sulphur. 
 
 Concerning the constitution of hydrogen persulphide 
 nothing definite is known. If the constitution of hydro- 
 gen dioxide is H-O-O-H, and hydrogen persulphide is 
 a disulphide, it seems probable that it has the constitu- 
 tion H-S-S-H, but there is no evidence bearing upon 
 this point. The fact that the sulphides can take up four 
 and only four atoms of sulphur, just as they can take up 
 four and only four atoms of oxygen, taken together with 
 the general similarity between the conduct of oxygen and 
 that of sulphur, suggests that these two reactions may 
 be of the same kind : 
 
 K 2 S + 4S = K 2 SS 4 ; 
 K 2 S + 40 = K 2 SO 4 . 
 
 But for reasons which will be more fully presented 
 under Sulphuric Acid (which see), it is generally believed 
 that in this acid two of the oxygen atoms are in direct 
 combination with sulphur alone, while two are in com- 
 bination with sulphur and hydrogen. If the polysul- 
 
202 INORGANIC CHEMISTRY. 
 
 phide has a similar constitution it must be represented 
 
 S 
 
 II 
 by the formula K-S-S-S-K, and, further, if the structure 
 
 II 
 8 
 
 of the polysulphide be as here represented, it is possible 
 that persulphide of hydrogen has a similar constitution. 
 Compounds of Sulphur with Members of the Chlorine 
 Group. Sulphur combines directly with chlorine and 
 forms the compounds S 2 C1 2 , SC1 2 , and SC1 4 . Of these 
 the first is the most stable. This can be boiled without 
 undergoing decomposition. The second, sulphur dichlo- 
 ride, SC1 2 , undergoes decomposition into chlorine and 
 sulphur monochloride at the boiling point ; while sulphur 
 tetrachloride exists only at low temperatures. All these 
 compounds are decomposed by water, yielding oxygen 
 compounds. In referring to the differences between the 
 acid-forming and the base-forming elements, attention 
 was called (see p. 125) to the fact that, in general, the 
 oxides of the base-forming elements are decomposed by 
 hydrochloric acid, yielding metallic chlorides and water ; 
 whereas with the acid-forming elements the reverse is 
 true, that is to say, the chlorides of the acid-forming 
 elements are, in general, decomposed by water yielding 
 oxides or hydroxides and hydrochloric acid. The truth 
 of this general statement is illustrated by the compounds 
 of sulphur and chlorine. But the products formed, ex- 
 cept in the case of the tetrachloride, are not strictly 
 analogous to the chlorides which are decomposed. Thus, 
 if in the monochloride S 2 C1 2 oxygen were simply substi- 
 tuted for chlorine, the product would be S 2 O ; but there 
 is no compound of oxygen of this composition, the sim- 
 plest one being sulphur dioxide, SO 2 , and this is formed. 
 The excess of sulphur set free, above that required for 
 the formation of the dioxide, is precipitated. The main 
 part of the reaction is represented thus : 
 
 2S 2 C1 2 + 2H 2 = 4HC1 + SO 2 + 38. 
 The decomposition of the other chlorides takes place in 
 
SELENIUM. 203 
 
 a similar way. That of the dichloride is represented 
 by the equation 
 
 2SC1 2 + 2H 2 O = SO 2 + 4HC1 + S. 
 
 That of the tetrachloride consists simply in the direct 
 replacement of the chlorine by oxygen. 
 
 Of the other compounds of sulphur with members of 
 the chlorine group, the hexiodide, SI 6 , is perhaps the 
 most interesting, as it shows that sulphur can be sexiva- 
 lent towards other elements than oxygen. This hexio- 
 dide is an extremely unstable compound. There is 
 doubt as to its existence. 
 
 SELENIUM, Se (At. Wt. 78.42). 
 
 Occurrence. Selenium occurs only in small quantity 
 yn nature. It was first found in the deposit formed in a 
 sulphuric acid chamber (see p. 215), and owed its origin 
 to the presence of small quantities of selenides in the sul- 
 phides used in the operation. It was found to resemble 
 tellurium, and for that reason was called selenium, from 
 <re\rjvrj, the moon, tellurium receiving its name from 
 tdlus, the earth. Selenium occurs in a number of the 
 natural sulphides, such as iron pyrites, copper pyrites, . 
 zinc blende, etc. ; and when these are treated in a current 
 of air to decompose them and convert the sulphur into 
 sulphur dioxide, the selenium is also oxidized, and the 
 .selenium dioxide thus formed is carried into the flues 
 and other parts of the apparatus. As it is a solid it is 
 easily condensed, and gradually a considerable quantity 
 collects in the flues. This flue-dust is the best material 
 from which to make selenium. For the purpose of ex- 
 traction the dust is treated with oxidizing agents, and 
 the selenium converted either into selenious acid, H 2 SeO 3 , 
 or a salt of selenic acid, H 2 SeO 4 . This is then reduced 
 by proper means. Thus selenious acid is reduced to 
 selenium by passing sulphur dioxide through its solu- 
 tion : 
 
 HJ3eO a + 2SO. + H,O = Se 4- 2HJ3O.. 
 
204 INORGANIC CHEMISTRY. 
 
 Hydrogen sulphide also reduces selenious acid, forming, 
 however, not free selenium, but selenium sulphide, SeS 2 , 
 analogous to sulphur dioxide and selenium dioxide : 
 
 H 2 Se0 3 + 2H 2 S = 3H 2 O + SeS 2 . 
 
 Properties. There are two modifications of selenium, 
 corresponding to those of sulphur. One is soluble in 
 carbon disulphide, the other is not. The soluble form 
 is obtained by reducing selenious acid by means of sul- 
 phurous acid, or by means of other reducing agents. The 
 insoluble variety is obtained by melting selenium, then 
 cooling it down suddenly to 210, and keeping it at 
 this temperature for some time. When it solidifies it is 
 found to be no longer soluble in carbon disulphide. Se- 
 lenium burns as sulphur does, forming the oxide SeO 2 , 
 which has the odor of rotten horse-radishes. 
 
 Hydrogen Selenide, H 2 Se. This compound is made in 
 the same way as the corresponding compound of sulphur, 
 by treating a selenide, as, for example, ferrous selenide, 
 FeSe, with an acid. It is a gas, which is soluble in 
 water, and has an odor something like that of hydrogen 
 sulphide, but it produces much more powerful effects 
 upon the olfactory nerves, temporarily destroying the 
 .sense of smell and causing painful sensations. 
 
 The compounds of selenium with the members of the 
 chlorine group present no features of special interest or 
 importance. In general they resemble the correspond- 
 ing compounds of sulphur, but are more stable. No 
 iodide of selenium corresponding to the hexiodide of 
 sulphur has been discovered. 
 
 TELLUKIUM, (Te At. Wt. 126.52). 
 
 Occurrence. Tellurium occurs in some gold ores in 
 the native or uncombined condition, and also in com- 
 bination with gold, silver, antimony, lead, and other 
 metals. Tellurides occur, among other places in the 
 United States, in California and Virginia. The general 
 method of preparing tellurium from its ores is the same 
 as that by which selenium is made. The tellurium is 
 
HYDROGEN TELLURIDE. 205 
 
 oxidized to tellurious acid, H 2 TeO 3 , which is isolated, 
 and then reduced by means of sulphurous acid. 
 
 Properties. Tellurium is silver-white, and crystallizes 
 easily. Heated in the air it burns, forming a thick 
 white cloud of tellurium dioxide, TeO 2 . Treated with 
 sulphuric acid it is oxidized to tellurious acid, H 2 TeO 3 , 
 the sulphuric acid being reduced. Nitric acid oxidizes 
 It likewise to tellurious acid. Melted together with 
 potassium nitrate it is converted into potassium tellurate, 
 K,TeO, 
 
 Hydrogen Telluride, H 2 Te. This compound is made 
 by treating zinc telluride, ZnTe, with an acid. It is a 
 gas, which resembles hydrogen sulphide in most of its 
 properties. 
 
 Like sulphur and selenium, tellurium combines with 
 the members of the chlorine group. The compounds 
 are more stable than those of the other two elements of 
 the group. Thus, taking the three tetrachlorides, that 
 of sulphur is extremely unstable, being capable of exist- 
 ence only at low temperatures ; that of selenium is more 
 stable, but still it is decomposed when heated to the 
 boiling temperature ; while that of tellurium can be 
 heated far above the temperature of boiling without de- 
 composition. 
 
 When tellurium tetrabromide is treated with potas- 
 sium chloride a compound of the formula K 2 TeBr 6 is 
 formed, which in composition, as will be observed, is 
 analogous to potassium tellurite, K 2 Te0 3 , two atoms of 
 bromine taking the place of each atom of oxygen. This 
 appears to be a derivative of the acid H 2 TeBr 6 , which is 
 similar in composition to the chlorplatinic acid, H 2 PtCl 6 , 
 referred to on p. 142, and to fluosilicic acid, H 2 SiF 6 . As 
 has already been remarked, not many such compounds 
 of the acid-forming elements are commonly met with, 
 largely for the reason that they are for the most part 
 easily decomposed by water and converted into the cor- 
 responding oxygen compounds. 
 
CHAPTER. XIV. 
 
 COMPOUNDS OF SULPHUR, SELENIUM, AND TELLURIUM 
 WITH OXYGEN AND WITH OXYGEN AND HYDROGEN. 
 
 Introductory. It has already been stated that, when 
 the three elements of the sulphur group are burned in 
 the air, they are converted into the corresponding diox- 
 ides. Under certain conditions it is possible to convert 
 sulphur dioxide and tellurium dioxide into the trioxides 
 SO 3 and TeO 3 , while the corresponding compound of 
 selenium has not been made. A lower oxide of sul- 
 phur, S 2 O 3 , and a higher one of the formula S 2 O 7 have 
 also been made. The dioxides dissolve in water, and 
 from these solutions salts of the general formula M 2 SO 3r 
 M 2 SeO 3 , and M 2 TeO 3 are obtained. The trioxide of sul- 
 phur dissolves in water with great evolution of heat, 
 forming compounds S(OH) 6 , OS(OH) 4 , and O 2 S(OH) 2 . 
 
 Most of the salts obtained from this solution are derived 
 from an acid of the formula H 2 SO 4 , and therefore this 
 is generally called sulphuric acid. By treating sulphuric 
 acid with reducing agents of various kinds it can be con- 
 verted successively into other acids containing a smaller 
 proportion of oxygen ; and if the reduction is pushed 
 far enough sulphur and hydrogen sulphide are obtained. 
 The limit of reduction is reached in hydrogen sulphide ; 
 and, on the other hand, if hydrogen sulphide is oxidized 
 the limit of oxidation is reached in sulphuric acid. The 
 intermediate products which have been isolated are 
 represented in the following table, by the side of which 
 for the sake of comparison is placed the similar series 
 of chlorine acids : 
 
 H 2 S HC1 
 
 HC10 
 
 H 2 S 2 O 4 HC10, 
 
 H 2 S0 3 HC10 8 
 
 H,SO 4 HC10 4 
 
 (206) 
 
COMPOUNDS OF SULPHUR. 207 
 
 In the chlorine series, however, one member more is 
 known than in the sulphur series. A remarkable fact 
 which has not yet been explained is that in each of the 
 ordinary forms of the acids, both in the chlorine series 
 and in the sulphur series, the number of hydrogen atoms 
 is the same as in the hydrogen compound. Throughout 
 the chlorine series there is but one hydrogen atom in 
 the molecule ; throughout the sulphur series there are 
 two. This refers only to those forms of the acid from 
 which most of the salts are derived, and plainly not to 
 such compounds as OS(OH) 4 and S(OH) 6 , which, as we 
 shall see, are closely related to sulphuric acid, just as 
 OC1(OH) 5 or H 5 C1O 6 is closely related to perchloric acid. 
 The number of compounds of sulphur is increased by 
 the power the element possesses of uniting with itself, 
 much as it does with oxygen. Thus the sulphides take 
 up sulphur and are converted into polysulphides, and 
 the limit of action of this kind is a compound of the gen- 
 eral formula M 2 SS 4 . This series is complete, the mem- 
 bers having the composition represented in the follow- 
 ing table : 
 
 M 2 S M 2 S 
 
 M 2 SS 
 
 M 2 SS 2 
 
 M 2 SS 3 M 2 S0 3 
 
 M 2 SS 4 M 2 SO 4 . 
 
 Comparing this with the series of oxygen compounds 
 an analogy is apparent, though the second and third 
 members of the oxygen series are lacking. The limit is 
 reached in a compound of the same order in both series. 
 Again, sulphur can take the place of part of the oxy- 
 gen in the oxygen acids. Thus compounds of the com- 
 position represented by the following formulas are con- 
 ceivable : .' - ; 
 
 H 2 SOS or H 2 S 2 ; 
 
 H 2 S0 2 S or H 2 S 2 2 ; 
 
 H 2 SOS 2 or H 2 S 3 O ; 
 
 H 2 S0 3 S or H 2 S 2 3 ; 
 
 H 2 S0 2 S 2 or H 2 S 3 2 ; 
 
 H 2 SOS 3 or H 2 S 4 0. 
 
208 INORGANIC CHEMISTRY. 
 
 Of these compounds only that one which has the compo- 
 sition H 2 S 2 O 3 is known. This is thiosulphuric acid or, as 
 it is sometimes called, hyposulphurous acid. 
 
 Other complications are occasioned by the combina- 
 tion of sulphur with sulphur in some of the oxygen acids 
 above mentioned. In this way probably is formed the 
 series represented in the following table : 
 
 Dithionic acid, H 2 S 2 O 6 
 
 Trithionic " . . ; . .V ';; I' H 2 S 3 O 6 
 
 Tetrathionic " H 2 S 4 O 6 
 
 Pentathionic " . ... . . .' '. H 2 S 5 O 6 . 
 
 While the number of compounds which sulphur forms 
 with hydrogen and oxygen is comparatively large, the 
 relations between them can be traced without serious 
 difficulty, and, studied in the right way, they are seen 
 to follow certain laws which govern their composition. 
 These relations will be discussed after the main facts 
 concerning the principal compounds have been studied. 
 
 The number of compounds which selenium and tellu- 
 rium form with hydrogen and oxygen is much smaller 
 than in the case of sulphur. The only ones known are 
 those which are analogous in composition to sulphurous 
 acid, H 2 SO 3 , and sulphuric acid, H 2 SO 4 ; besides these, 
 however, there are a few compounds known which con- 
 tain hydrogen, sulphur, selenium, and oxygen. These 
 are closely related to some of the compounds of sulphur 
 with hydrogen and oxygen, being derived from them by 
 the introduction of selenium for a part of the oxygen. 
 
 Sulphur combines with oxygen and members of the 
 chlorine group, and also with these elements and hydro- 
 gen. The simplest compound of this kind is that which 
 has the composition SOC1 2 . This is plainly analogous 
 to sulphur dioxide, from which it can be made by sub- 
 stituting two atoms of chlorine for one atom of oxygen in 
 the molecule. While this compound cannot unite directly 
 with more chlorine, a compound of the composition 
 SO 2 Cl a can be made. This is in accordance with what 
 we should expect from a consideration of the conduct 
 of sulphur towards oxygen and towards chlorine respec- 
 
SULPHURIC ACID. 209 
 
 tively. With oxygen it forms a stable compound, SO S , 
 while the limit of combination with chlorine is reached 
 in the compound SC1 4 , and even this is very unstable. 
 As we commonly say, sulphur is sexivalent towards 
 oxygen and quadrivalent towards chlorine. Towards 
 both together it is also sexivalent, as shown in the com- 
 pound SO 2 C1 2 , though the compound SOC1 4 does not 
 exist. Just as the simple compounds of sulphur and 
 chlorine are decomposed readily by water, forming oxy- 
 gen compounds, so these mixed compounds containing 
 oxygen and chlorine are also readily decomposed by 
 water. The decomposition appears to take place as 
 represented in the following equations : 
 
 
 In the first case, the product formed breaks down into 
 sulphur dioxide and water, so that the main action is 
 represented thus : 
 
 + H 2 = S0 2 + 2HC1, 
 
 and the change therefore consists in substituting one 
 atom of oxygen for the two atoms of chlorine. It is 
 analogous to the transformation of sulphur tetrachloride 
 into sulphur dioxide : 
 
 SC1 4 + 2H 2 = S0 2 + 4HC1. 
 
 As these mixed chlorides and oxides are readily con- 
 verted into the corresponding acids by treatment with 
 water, they are frequently spoken of as the chlorides of 
 the acids, or the acid chlorides. Between them and the 
 anhydrides, or acidic oxides, there is plainly an analogy. 
 
 Sulphuric Acid, H 2 SO 4 . The many salts of sulphuric 
 acid can be best explained on the assumption that the 
 forms of sulphuric acid are derived from a compound 
 S(OH) 6 , just as the salts of periodic acid can be best ex- 
 plained on the assumption of a fundamental compound, 
 I(OH) 7 . As the latter is called normal periodic acid, so 
 
210 INORGANIC CHEMISTRY. 
 
 the former is called normal sulphuric acid. From normal 
 sulphuric acid by successive losses of water are formed 
 the compounds H 4 SO 5 , H 2 SO 4 , and SO 3 : 
 
 S(OH) 6 = OS(OH) 4 + H 2 O ; 
 OS(OH) 4 = 2 S(OH) 2 + H 2 0; 
 2 S(OH) 2 = 3 S + H 2 0. 
 
 While these compounds are all known, the salts of sul- 
 phuric acid are for the most part derived from the acid 
 containing two hydrogen atoms, viz., O 2 S(OH) 2 . In some 
 salts two of the hydrogen atoms in normal sulphuric 
 acid are replaced by metals, the others remaining. Salts 
 of the general formula S(OH) 4 (OM) 2 are thus formed. 
 These are generally represented as containing two 
 molecules of water of crystallization, thus: M 2 SO 4 +2H 2 O. 
 To decide between the two views is at present impos- 
 sible. In the first place, the question as to the nature 
 of water of crystallization must be answered before it 
 can be said whether there is any conflict between the 
 two views. As far as the assumption of normal sul- 
 phuric acid is concerned, it can only be said that similar 
 assumptions in the case of iodine and in that of phos- 
 phorus seem to be well founded, and, although it would 
 be difficult to furnish proof positive of the reality of the 
 relations, the above suggestion is by far the simplest 
 which has been made, and it is of great assistance in 
 dealing with the acid derivatives of the elements of 
 Families VII, VI, V, and IV. 
 
 For the science as well as for the art of chemistry 
 sulphuric acid is of fundamental importance. It is 
 used daily in every chemical laboratory and in every 
 chemical factory, and in some of the most important 
 branches of chemical industry enormous quantities of 
 it are used. In consequence of the large demand for 
 the acid the process used in preparing it has been 
 studied with great care, and it has reached a high state 
 of perfection. As it furnishes an excellent example of 
 the applications of the facts of science to the building 
 up of an industry, it will be studied with some degree of 
 fulness. 
 
SULPHURIC ACID. 211 
 
 Sulphuric acid has been known for a long time. It 
 was made in the eighteenth century by heating calcined 
 iron vitriol (ferrous sulphate, FeSO 4 ) with sand, and 
 was hence called oil of vitriol, a name which still sur- 
 vives. It was also prepared by treating sulphur with 
 saltpeter (potassium nitrate, KNO 3 ). The acid was first 
 prepared on the large scale in England with the use of 
 comparatively large chambers lined with lead. It was 
 known as English sulphuric acid, and is still called by 
 this name. 
 
 Sulphuric acid occurs in nature in the form of salts or 
 sulphates, such as calcium sulphate or gypsum, barium 
 sulphate or heavy spar, and others. Although these 
 salts occur in large quantity, the acid is not obtained 
 from them, as there is no economical way of substituting 
 hydrogen for the metal. The preparation of the acid by 
 the substitution of hydrogen for the calcium in calcium 
 sulphate, CaSO 4 , suggests itself, but this cannot be easily 
 effected by any substance available in quantity. Hydro- 
 chloric and nitric acids are made from their salts which 
 occur in nature, the former, as we have seen, from 
 sodium chloride, NaCl, the latter from saltpeter or po- 
 tassium nitrate, KNO 3 , by treating with sulphuric acid. 
 There. is, however, no acid which acts upon the sulphates 
 as sulphuric acid acts upon chlorides, nitrates, and many 
 other salts. 
 
 The manufacture of sulphuric acid is based upon the 
 two fundamental facts that (1) when sulphur is burned 
 it is converted into sulphur dioxide, SO 2 ; and (2) when 
 sulphur dioxide is treated with an oxidizing agent in the 
 presence of water it is converted into sulphuric acid : 
 
 SO 2 + H 2 O + O = H 2 SO 4 ; or 
 SO 2 + H 2 O = H 2 SO 3 ; and 
 
 H 2 S0 3 + = H 2 S0 4 . 
 
 The chief difficulty is, of course, experienced in effect- 
 ing the oxidation of the sulphurous acid. It is accom- 
 plished by introducing the gas, sulphur dioxide, into 
 large chambers together with compounds of nitrogen 
 
212 INORGANIC CHEMISTRY. 
 
 and oxygen, and steam. That which plays the chief 
 part in effecting the transformation is a mixture of the 
 two oxides NO and NO 2 . This acts like the trioxide,. 
 N 2 O 3 , and it may be represented by this formula. 
 Instead of starting with the trioxide, nitric acid is used 
 in the manufacture of sulphuric acid. This at first 
 reacts with sulphur dioxide and steam, as represented in. 
 the equation 
 
 2HNO 3 + 2S0 2 + H 2 = 2H 2 SO 4 + N,O 3 - 
 
 After this the main reactions are (1) the formation of a, 
 
 NO 
 compound of the formula SO 2 <Q-rr 2 , called nitrosyl-am- 
 
 phuric acid; and (2) the decomposition of the nitrosyl- 
 sulphuric acid by water. These reactions are repre- 
 sented in the two equations following : 
 
 280, + N.O, + O, + H,0 = 2SO,(OH)(NO,) ; 
 
 The nitrogen trioxide formed in the second reaction 
 then again enters into combination with sulphur dioxide, 
 oxygen, and water to form nitrosyl- sulphuric aci$, which 
 again undergoes decomposition with water. It will be 
 seen, therefore, that the trioxide is not lost, but that 
 it simply serves the purpose of effecting the oxidation of 
 the sulphur dioxide, and, theoretically, a small quantity 
 of the gas should be capable of transforming an infinite 
 quantity of sulphur dioxide into sulphuric acid. 
 
 Other reactions besides those mentione/i above are 
 undoubtedly involved in the manufacture of sulphuric 
 acid, but, according to the most elaborate researches 
 made on the subject in a factory in operation, those 
 mentioned are the principal ones. Whatever the details. 
 may be, the oxidation is effected without difficulty, and 
 the waste in nitrogen compounds is now but slight. 
 
 The reactions, as has been said, are carried on in large 
 chambers lined with lead, and known as the leaden 
 chambers. The sulphur is burned in a special furnace 
 
SULPHURIC ACID. 
 
 213 
 
 so arranged that air has free access to it. The dioxide 
 thus formed is then conducted through a tower so con- 
 structed that it presents a large surface to the action of 
 the gas. Through this tower dilute sulphuric acid, taken 
 from another tower at the end of the system, flows from 
 
 above, and the hot gases coming in contact with this 
 serve the purpose of concentrating it, while the gas itself 
 is cooled before entering the leaden chamber. The 
 general arrangement of the essential parts of a sulphuric 
 acid factory are shown in Fig. 7. 
 
214 INORGANIC CHEMISTRY. 
 
 The sulphur dioxide passes through the large tube x 
 into the tower G, called the Glover tower. This is filled 
 from d to e with pieces of fire-brick over which from the 
 cistern b a continual stream of dilute sulphuric acid 
 flows. The gases are thus cooled down and the acid 
 concentrated. From a second cistern there flows at the 
 same time concentrated sulphuric acid from the tower 
 G', or the Gay Lussac tower. This stronger acid con- 
 tains oxides of nitrogen in combination, and by contact 
 with the dilute acid the oxides are set free and are thus 
 mixed with the sulphur dioxide. The nitric acid is in- 
 troduced into the first chamber, No. 1, in the form of 
 vapor, together with the oxides of nitrogen and sulphur 
 dioxide, and here also the gases meet with water vapor. 
 The reactions above referred to now take place. From 
 the first chamber the gases pass through the pipe v into 
 the second, from this into the third at w, and finally, 
 in order to prevent the escape of any unused oxides 
 of nitrogen, the gases are passed through the Gay 
 Lussac tower G'. This contains coke over which is 
 kept flowing concentrated sulphuric acid which takes up 
 the oxides of nitrogen and will give them up again when 
 diluted. This liberation of the oxides of nitrogen is 
 accomplished, as has been said, in the Glover tower. 
 The concentrated acid collected at the bottom of the 
 Glover tower is well adapted for use in the Gay Lussac 
 tower. The leaden chambers in some factories are 
 nearly 100 feet long, 18 to 30 feet broad, and 15 feet 
 high. 
 
 The acid taken from the chambers contains about 64 
 per cent of the compound H 2 SO 4 , and has the specific 
 gravity 1.5. This is evaporated first in lead pans until 
 it reaches the specific gravity 1.75. As stronger acid 
 acts upon lead, the evaporation beyond this point is 
 carried on in platinum, glass, or iron. The strong acid 
 thus obtained, which has a specific gravity of about 1.830, 
 is the concentrated sulphuric acid of commerce. 
 
 Instead of sulphur, iron pyrites is now extensively 
 used for the preparation of sulphur dioxide in the manu- 
 facture of sulphuric acid. This is a compound of iron 
 
SULPHURIC ACID. 215 
 
 and sulphur of the composition FeS 2 . When it is heated 
 in contact with the air it is converted into the oxide, and 
 the sulphur passes off in the form of sulphur dioxide. 
 If the sulphide used contains selenides the selenium 
 dioxide formed in the roasting process is carried into 
 the flues, and is there deposited with the flue-dust, as 
 was stated in speaking of the source of selenium. 
 
 Commercial sulphuric acid is an oily liquid, usually 
 somewhat colored by impurities. The nature of the 
 impurities is dependent upon the substances used in the 
 manufacture and upon the conditions. It often con- 
 tains some lead sulphate in solution, and when it is 
 diluted with water this separates, giving the liquid a 
 more or less cloudy appearance. By standing, however, 
 the liquid becomes clear, as the lead sulphate collects 
 at the bottom. For obvious reasons, some oxides of 
 nitrogen are also generally present. Among other im- 
 purities frequently met with in the commercial acid are 
 arsenic from the pyrites, and a little selenium. 
 
 Pure Sulphuric Acid is made from the commercial acid 
 by treating it with such substances as will remove the 
 oxides of nitrogen and arsenic, and by distilling. It is 
 a colorless liquid at the ordinary temperature. The 
 pure acid generally made has about the same concentra- 
 tion as the commercial crude acid. By taking special pre- 
 cautions in the distillation a product having very nearly 
 the composition H 2 SO 4 can be obtained. This is a thick, 
 clear liquid of specific gravity 1.854 at 0. When cooled 
 down to a low temperature it forms large crystals which 
 melt at 10. 5. When heated it gives off some sulphur 
 trioxide in consequence of partial decomposition into 
 this substance and water : 
 
 H 2 S0 4 = H 2 + S0 3 . 
 
 It finally boils, however, at the temperature 338, and 
 the distillate has the composition represented by the 
 formula H 2 SO 4 + -^-H a O. Heated somewhat above its 
 boiling point it is completely decomposed into sulphur 
 trioxide and water, according to the above equation. If 
 the heating is carried to a higher temperature the sul- 
 
216 INORGANIC CHEMISTRY. 
 
 phur trioxide breaks down into sulphur dioxide and 
 oxygen. 
 
 Sulphuric acid has a strong tendency to absorb water, 
 and to form compounds with it. In consequence of the 
 formation of these compounds a great deal of heat is 
 evolved when sulphuric acid is mixed with water. One 
 molecule of sulphuric acid when mixed with about 1600 
 molecules of water gives 17,850 cal. Among the com- 
 pounds thus formed are the so-called hydrates of the 
 composition H 2 SO 4 + H 2 O, and H 2 SO 4 + 2H 2 O, which 
 should probably be regarded as having the constitution 
 OS(OH) 4 and S(OH) 6 . 
 
 So strong is the tendency of the acid to combine with 
 the elements of water that it abstracts them in the pro- 
 portions to form water from many organic substances. 
 Thus a piece of wood placed in sulphuric acid soon 
 turns black and is completely disintegrated. The cause 
 of this is to be found in the fact that wood consists 
 essentially of carbon, hydrogen, and oxygen ; and the 
 sulphuric acid abstracts the hydrogen and oxygen, 
 leaving the carbon mainly in the uncombined state. 
 Similarly it abstracts moisture from gases, and it is used 
 extensively in the laboratory for this purpose. In con- 
 tact with the skin it acts as upon wood, causing wounds 
 which are painful and difficult to heal. 
 
 Sulphuric acid is what is called a strong acid, a term 
 which needs further explanation, and the subject of the 
 relative strengths of acids will be discussed in due time. 
 As used here, it means simply that the acid has the 
 power to decompose the salts of most other acids, ap- 
 propriating the metals and setting the other acids free. 
 This is illustrated by the formation of hydrochloric and 
 nitric acids by treating sodium chloride and potassium 
 nitrate respectively with sulphuric acid. We shall see, 
 however, that the fact that sulphuric acid decomposes 
 the salts of hydrochloric and nitric acids does not prove 
 that it is a stronger acid than they are. There are other 
 facts besides the strengths of the acids which determine 
 whether such decompositions take place or not. 
 
 Sulphuric acid gives up its oxygen to other substances 
 
SULPHURIC ACID. 217 
 
 comparatively easily, and is generally reduced to sul- 
 phurous acid which is decomposed into sulphur dioxide 
 and water. Thus hydrogen sulphide and hydrobromic 
 acid both act upon it, as we have seen ; the products be- 
 ing, in the former case, sulphur dioxide, water, and sul- 
 phur ; in the latter, sulphur dioxide, water, and bromine : 
 
 H 2 S0 4 + H 2 S = 2H 2 + SO 2 + S ; 
 H 2 S0 4 + 2HBr = 2E 2 O + SO 2 + Br 2 . 
 
 Generally, in acting upon metals it forms the correspond- 
 ing salt and hydrogen is given off, but if the temperature 
 is high and the acid concentrated the hydrogen acts 
 upon the acid, reducing it. Thus, in making hydrogen 
 by treating zinc with sulphuric acid, if the acid is dilute 
 and is kept cool the reaction takes place in the simplest 
 way ; but if the acid is concentrated and is allowed to 
 get hot some hydrogen sulphide is always formed. 
 Copper does not act upon sulphuric acid at ordinary 
 temperatures. If it is treated with concentrated sul- 
 phuric acid, sulphur dioxide is the chief reduction- 
 product. Even free hydrogen if passed into sulphuric 
 acid heated to 160 reduces it, forming sulphur dioxide. 
 The complete reduction of sulphuric acid to sulphur and 
 hydrogen sulphide is beautifully shown by the action of 
 hydriodic acid. As was stated in speaking of the action 
 of sulphuric acid upon potassium iodide, four reactions 
 may take place when these substances act upon each 
 other. They are represented by the equations 
 
 2KI + H 2 S0 4 = K 2 SO 4 + 2HI ; 
 H 2 S0 4 + 2HI = 2H 2 + S0 2 + 21 ; 
 H 2 S0 4 + 6HI =4H 2 + S +61; 
 H 2 SO 4 + 8HI =4H 2 +H.S + 8I. 
 
 Carbon and sulphur also act upon sulphuric acid and 
 reduce it to sulphur dioxide. With sulphur the action 
 is represented thus : 
 
 2H 2 S0 4 + S = 3S0 2 + 2H S O. 
 
218 INORGANIC CHEMISTRY. 
 
 Sulphuric acid is a dibasic acid (see p. 138) and forms 
 two series of sajts, normal salts of the general formula 
 M^O^ and acid salts of the formula MHSO 4 . The sul- 
 phates are very stable salts. Those of the strongest 
 bases are not decomposed by the highest heat. Those of 
 weaker bases break down, giving up sulphur trioxide ; or 
 if the decomposition takes place at a temperature above 
 that at which sulphur trioxide breaks down, this de- 
 composes into sulphur dioxide and oxygen. 
 
 Tetrah.ydroxyl-Sulph.uric Acid, OS(OH) 4 . This com- 
 pound has already been referred to as being formed by 
 the addition of water to ordinary sulphuric acid. It is 
 a crystallized compound which melts at 7.5. No salts 
 of this acid are known, or, rather, no salts derived from 
 it by the replacement of all the hydrogen by metal are 
 known. Only two of the hydrogen atoms appear to be 
 replaceable. This compound is generally represented 
 by the formula H 2 SO 4 -f- H 2 O, and called a hydrate. 
 
 Normal Sulphuric Acid, S(OH) 8 . This is commonly 
 represented by the formula H 2 SO 4 -j- 2H 2 O and, like the 
 preceding compound, regarded as a hydrate. It appears 
 to be formed by the action of water on sulphuric acid. 
 On mixing sulphuric acid and water the maximum con- 
 traction takes place when the quantities necessary to 
 form this compound are brought together, and there are 
 other good reasons for believing that the compound 
 exists in the solution. It is not a solid like the pre- 
 ceding compound. The acid forms no salts bearing sim- 
 ple relations to it. There are some salts known, how- 
 ever, which appear to be related to it, as, for example, 
 the salt K 3 HS 2 O 8 and some others similar to it. The 
 acid from which these salts are derived is believed to be 
 formed from the normal acid by elimination of water, as 
 represented below : 
 
 fOH HOI 
 
 18> S 1 OH + HO* S< OH = (HO) 2 OS<g>SO(OH) 2 + 4H 2 
 OH HO I 
 
 An acid formed in this way would have the formula 
 
DISULPHURIC ACID. 219 
 
 H 4 S 2 O g , and the salt K 3 HS 2 O 8 is the tertiary potassium 
 salt of this acid. 
 
 Disulphuric Acid, Pyrosulphuric Acid, H 2 S 2 O 7 . This 
 compound, which is also known by the names fuming 
 sulphuric acid and Nordhausen sulphuric acid, is closely 
 related to ordinary sulphuric acid, and is made from it 
 by treating it with sulphur trioxide, the two combining 
 directly, as represented thus : 
 
 H 2 SO 4 + SO 3 = H 2 S a O 7 . 
 
 It is made by distilling ferrous sulphate which is not 
 perfectly dry : 
 
 4FeS0 4 + H 2 = 2Fe a O 3 + 2SO 2 + H a S a 7 . 
 
 A so-called solid sulphuric acid is now manufactured by 
 a process which will be referred to under Sulphur Triox- 
 ide. It is essentially disulphuric acid. 
 
 Disulphuric acid, as it is found in the market, is gen- 
 erally a thick liquid which gives off dense fumes when 
 exposed to the air, and breaks down completely into sul- 
 phur trioxide and sulphuric acid when heated. When 
 pure it crystallizes in large crystals which melt at 35. 
 
 It is believed that the relation between disulphuric 
 acid and ordinary sulphuric acid should be expressed by 
 these formulas : 
 
 *ro >so = s - -, 
 
 H0 \OH HO/ 
 
 Or the formula of the acid may be written thus : 
 
 2 S-OH 
 
 This relation is similar to that believed to exist between 
 the normal acid S(OH) 6 and the acid H 4 S 2 O 8 (see p. 218). 
 
220 INORGANIC CHEMISTRY. 
 
 Disulphuric acid forms normal salts of the general for- 
 mula M 2 S 2 O 7 and acid salts of the general formula 
 MHS 2 O 7 . When heated the normal salts break down, 
 yielding ordinary sulphates and sulphur trioxide : 
 
 M 2 S 2 O 7 = M 2 SO 4 + S0 3 . 
 
 When an acid sulphate like KHSO 4 is heated to a 
 sufficiently high temperature it breaks down into water 
 and a disulphate : 
 
 2KHS0 4 = K 2 S 2 7 + H 2 0. 
 
 Sulphurous Acid, H 2 SO 3 . While no acid of the formula 
 H 2 SO 3 is known in the free condition, a large number of 
 salts which are derived from this acid are known. They 
 are made by treating a water solution of sulphur dioxide 
 with bases, and therefore it is believed that the solution 
 contains the acid which is formed by the action of sulphur 
 dioxide on water, thus : 
 
 S0 2 + H 2 = H 2 S0 3 . 
 
 It is, however, so unstable that it breaks down into the 
 dioxide and water at every attempt to isolate it. The 
 dioxide, as has been stated and as will be shown more 
 fully, is formed by the burning of sulphur and by the 
 reduction of sulphuric acid. The acid forms a number 
 of unstable hydrates, apparently of complicated com- 
 position. Owing to their great instability, however, the 
 Investigation of these substances is extremely difficult 
 and unsatisfactory. 
 
 Sulphurous acid takes up oxygen readily and is thus 
 transformed into sulphuric acid. It is only necessary to 
 allow a solution to stand for a time to find that the odor 
 of the gas disappears and that sulphuric acid is then 
 present in the solution. Sulphurous acid is frequently 
 used in the laboratory as a reducing agent. Its action in 
 this way has been illustrated in the method for the ex- 
 traction of selenium from selenious acid (see p. 203). The 
 reaction in this case is represented thus : 
 
SULPHUROUS ACID. 221 
 
 H 2 Se0 3 + 2S0 2 + H 2 O = Se + 2H 2 SO 4 ; or 
 H 2 SeO 3 + 2H 2 S0 3 = Se + 2H 2 SO 4 + H 2 O. 
 
 Another illustration of its power as a reducing agent is 
 -shown in its action upon iodic acid. When it is added 
 to a solution containing iodic acid, HIO 3 , iodine separates, 
 the reaction taking place in accordance with the follow- 
 ing equation : 
 
 2HI0 3 + 5H 2 SO 8 = H 2 + 5H 2 SO 4 + I 2 . 
 
 If sulphurous acid is added to the liquid in which the 
 iodine is suspended further action takes place, the iodine 
 being reduced to hydriodic acid, thus : 
 
 H 2 S0 3 + H 2 + I 2 = H 2 S0 4 + 2HI. 
 
 This reaction takes place only in dilute solution. Con- 
 centrated sulphuric acid acts upon hydriodic acid and is 
 reduced by it, as we have seen. Towards some sub- 
 stances sulphurous acid acts as an oxidizing agent, and 
 is itself reduced to lower forms, as hyposulphurous acid, 
 H 2 S,,O 4 , and hydrogen sulphide. This has been illustrated 
 in the action of hydriodic acid upon sulphuric acid. It 
 is also illustrated in the action of zinc upon sulphurous 
 acid in the presence of hydrochloric acid, when this re- 
 action takes place : 
 
 3Zn + 6HC1 + H 2 SO 3 = 3ZnCl 2 + 3H 2 O + H 2 S. 
 
 Treated with zinc alone the acid sodium salt, NaHSO 3 , 
 is reduced as shown in the following equation, a salt of 
 hyposulphurous acid, H 2 S 2 O 4 , being formed : 
 
 Zn + 4NaHSO 3 = ZnSO 3 + Na 2 SO 3 + Na 2 S 2 O 4 + 2H 2 O. 
 
 Sulphurous acid when heated in a sealed tube breaks 
 down into sulphuric acid and sulphur : 
 
 3H 2 S0 3 = 2H 2 S0 4 + H 2 + S. 
 This kind of decomposition is also characteristic of the 
 
"222 INORGANIC CHEMISTRY. 
 
 salts of the acid with the alkali metals, as, for example,, 
 sodium sulphite : 
 
 4Na a SO 3 = 3Na 2 S0 4 + Na 2 S. 
 
 In fact, whenever an alkali salt of any of the oxygen 
 acids of sulphur is heated the tendency is towards the 
 formation of the sulphate, all the oxygen in the salt 
 going to form sulphate ; and the other elements in excess 
 of what may be needed for the sulphate arranging them- 
 selves in other forms. 
 
 Just as the sulphites take up oxygen to form sul- 
 phates, they also take up sulphur to form thiosulphates. 
 The two reactions appear to be perfectly analogous : 
 
 Na 2 SO 3 + 0=Na 2 S0 4 ; 
 
 Na 2 SO 3 + S = Na 2 S 2 O 3 (or Na 2 SO 3 S). 
 
 t 
 
 Sulphurous acid forms two series of salts, the normal 
 sulphites of the general formula M 2 SO 3 , and the acid 
 sulphites of the general formula MHSO 3 . These are un- 
 stable, though not as markedly so as the acid itself. 
 When treated with most acids they are decomposed,, 
 yielding sulphur dioxide instead of sulphurous acid. 
 The decomposition of sodium sulphite with hydrochloric 
 acid is represented by the equation 
 
 Na 2 SO 3 + 2HC1 = 2NaCl + H 2 O + SO, ; 
 with sulphuric acid thus : 
 
 Na 2 S0 3 + H 2 S0 4 = Na 2 SO 4 + H 2 O + SO 2 . 
 
 The sulphites, like sulphurous acid, combine readily 
 with oxygen, tending to pass over into the sulphates, and, 
 as has been remarked, they also tend to unite with sul- 
 phur to form the thiosulphates. 
 
 Hyposulphurtfus Acid, H 2 S 2 O 4 . This acid is also called 
 hydrosulphurous acid, but the name hyposulphurous 
 acid is more in accordance with the nomenclature adopted 
 for the other acids, and is now preferred. But little is 
 
THIOSULPHURIC ACID. 223 
 
 known of the compound. It is formed by reduction of 
 .a salt of sulphurous acid by treating with zinc, when this 
 reaction takes place : 
 
 Zn + 4NaHS0 8 = ZnSO 8 + Na a SO. + Na a S a O, + 2H a O. 
 
 The free acid cannot be obtained from this sodium salt, 
 .and notwithstanding the fact that it has been the subject 
 of a number of elaborate investigations there is still 
 some doubt as to the composition of the salt, the dis- 
 coverer still claiming that it has the composition repre- 
 sented by the formula Na a SO a , according to which it is to 
 be referred to an acid of the formula H 3 SO a . 
 
 Hyposulphurous acid, like sulphurous acid, combines 
 readily with oxygen, and passes first into sulphurous and 
 then into sulphuric acid. Its reducing action is stronger 
 than that of sulphurous acid. It is decomposed by 
 standing in the air, yielding first thiosulphuric acid, 
 H 2 S 2 O 3 , and then sulphur dioxide, sulphur, and water. 
 
 Hyposulphurous acid, so far as composition is con- 
 cerned, occupies a position between thiosulphuric acid, 
 H 2 S 2 O 3 , and pyrosulphurous acid, H 2 S a O 6 , which latter is 
 related to sulphurous acid in the same way that disul- 
 phuric or pyrosulphuric acid, H 2 S 2 O 7 , is related to sul- 
 phuric acid. The relations between hyposulphurous, 
 thiosulphuric, and pyrosulphurous acids are shown in 
 the table below : 
 
 H,S 2 3 ; 
 IW) 4 ; 
 H a S 2 0, 
 
 Thiosulphuric Acid, H 2 S 2 O 3 . This acid was formerly, 
 and is still by many, called hyposulphurous acid. Its 
 formation, or the formation of its salts by the addition 
 of sulphur to the sulphites, has been mentioned, and the 
 analogy between this reaction and that of the formation 
 of sulphates by the addition of oxygen to sulphites has 
 been commented upon. It may be regarded as sulphuric 
 acid in which sulphur has been substituted for one atom 
 
224 INORGANIC CHEMISTRY. 
 
 of oxygen, and hence the name thiosulphuric acid is 
 appropriate, whereas the name hyposulphurous acid 
 suggests at once a compound similar to sulphurous acid, 
 but containing less oxygen, and is therefore inappro- 
 priate. Sodium thiosulphate is formed together with the 
 pentasulphide by the action of sulphur upon sodium 
 hydroxide : 
 
 GNaOH + 12S = 2Na 3 S 6 + Na,S a O 3 + 3H a O. 
 
 The sulphides of the alkali metals pass over into the 
 corresponding thiosulphates by the action of oxygen. 
 Thus the pentasulphide is changed when exposed to the 
 air in aqueous solution. The action consists in a sub- 
 stitution of three atoms of oxygen for three of sulphur : 
 
 Na 2 S 5 + 30 = Na 2 S 2 3 + 38. 
 
 Sodium thiosulphate is formed, further, by the action of 
 iodine upon a mixture of sodium sulphide and sodium 
 sulphite : 
 
 21 = > 2NaI. 
 
 The acid itself is very unstable, breaking down into sul- 
 phur dioxide, sulphur, and water (see p. 190). By acids its 
 salts are decomposed in a similar way with evolution of 
 sulphur dioxide and separation of sulphur, which appears 
 suspended in the liquid in a very fine state of division. 
 With hydrochloric acid the decomposition takes place 
 thus : 
 
 Na 2 S 2 O 3 + 2HC1 = 2NaCl + SO 2 + S + H 2 O. 
 
 When heated the thiosulphate of sodium breaks down 
 according to the rule stated in speaking of the decompo- 
 sition of the sulphite by heat. All the oxygen goes to 
 the formation of the sulphate, and the elements left over 
 
OTHER ACIDS OF SULPHUR. 225 
 
 in excess of what is required for the sulphate unite to 
 form another compound, thus : 
 
 4Na 2 S 2 O 3 = 3Na 2 SO 4 + Na 2 S 5 . 
 
 Other Acids of Sulphur. Of the other acids of sulphur 
 but little need be said here. As was stated on page 208, 
 these acids form a series the members of which are closely 
 related to one another. The series is as follows : 
 
 Dithionic acid, ....... H 2 S 2 O 6 
 
 Trithionic acid, ....... H 2 S 3 O 6 
 
 Tetrathionic acid, ...... H 2 S 4 O 6 
 
 Pentathionic acid, . . . . . . H 2 S 5 O 6 
 
 Dithionic Acid, or a salt of the acid, is made by passing 
 sulphur dioxide into water having finely powdered man- 
 ganese dioxide in suspension. This reaction takes place : 
 
 MnO 2 + 2SO 2 = MnS 2 O 6 . 
 
 The product is the manganese salt of dithionic acid, 
 and from this other salts can be prepared. The free 
 acid breaks down readily into sulphuric acid and sulphur 
 dioxide : 
 
 So, too, when a salt of the acid is heated it breaks down, 
 forming a sulphate and sulphur dioxide : 
 
 K 2 S 2 6 = K 2 S0 4 + S0 2 . 
 
 Trithionic Acid, H 2 S 3 O fl , or its potassium salt, is formed 
 by treating a solution of acid potassium sulphite, KHSO 3 , 
 with "flowers of sulphur," when reaction takes place 
 thus: 
 
 6KHSO 3 + 28 = 2K 2 S 3 6 + K 2 S 2 O 3 + 3H 2 O. 
 "When the dry potassium salt is heated it is de- 
 
226 INORGANIC CHEMISTRY. 
 
 composed, yielding the sulphate, sulphur dioxide, and 
 sulphur : 
 
 Similarly, when treated with acids, decomposition takes 
 place with evolution of sulphur dioxide, separation of 
 sulphur, and formation of sulphuric acid : 
 
 H 2 S 3 6 = H 2 S0 4 + SO, + S. 
 
 Tetrathionic Acid, H a S 4 O 6 , is made from salts of thio- 
 sulphuric acid by treating them with iodine. Thus with 
 sodium thiosulphate the reaction is 
 
 2Na 2 S 2 O 3 + 21 = Na 2 S 4 O 6 + 2NaI. 
 
 The acid is moderately stable, so that a dilute solution 
 can be boiled without undergoing decomposition. When 
 the concentrated acid is heated, however, it breaks down 
 into sulphuric acid, sulphur dioxide, and sulphur : 
 
 H 2 S 4 6 = H 2 S0 4 + S0 2 + 28. 
 
 Pentathionic Acid, H 2 S B O 6 , is formed by the action of 
 hydrogen sulphide upon a solution of sulphur dioxide in 
 water. 
 
 Persulphuric Acid, H 2 S 3 O e , is formed by dissolving the 
 oxide S 2 O T in water, the oxide itself being formed by 
 subjecting a mixture of sulphur dioxide and oxygen to 
 the silent discharge in an ozone tube (see p. 85). The 
 potassium salt of persulphuric acid is easily obtained by 
 subjecting to electrolysis a saturated solution of acid 
 potassium sulphate. 
 
 Constitution of the Acids of Sulphur. The existence 
 of the oxide SO, and of the iodide SI, seems to show that 
 sulphur is sexivalent towards oxygen and towards iodine. 
 Considering, further, the facts presented under the head 
 of Periodic Acid (which see), which can only be explained 
 satisfactorily by the aid of the assumption that the differ- 
 ent varieties of periodic acid are derived from the normal 
 acid I(OH),, and the analogous facts presented under Sul- 
 
CONSTITUTION OF THE ACIDS OF SULPHUR. 
 
 phuric Acid, which lead to the belief that this acid is de- 
 rived from the normal acid S(OH) 6 , the arguments in favor 
 of the sexivalence of sulphur in sulphuric acid are seen 
 to be strong, though not conclusive. On the other hand, if 
 sulphur is sexivalent in sulphur trioxide and in sulphuric 
 acid, it is quadrivalent in sulphur dioxide and sulphur 
 tetrachloride, and bivalent in hydrogen sulphide and sul- 
 phur dichloride. But if sulphur is sexivalent in sul- 
 phuric acid the constitution of the acid must be repre- 
 
 OH 
 
 rented by the formula H-O-S-O-H or S4 H . Of 
 
 O 
 
 course such a formula as this involves the hypothesis 
 that when an oxygen atom is combined with only one 
 other atom two valences or affinities are brought into 
 play, and in regard to this we have very little if any evi- 
 dence. It may be said, however, that if oxygen which 
 is thus combined is replaced by univalent atoms its 
 place is always taken by two of these, indicating that 
 whatever the power may be which holds the oxygen atom 
 in combination, that same power can hold two chlorine 
 atoms, etc., and it is convenient to use the double line to 
 indicate the existence of this power. 
 
 The view expressed by the above formula in regard to 
 the structure of sulphuric acid has been tested experi- 
 mentally by methods which appear somewhat compli- 
 cated, but the principle involved can be easily explained. 
 If the formula is correct, then both hydroxyl groups bear 
 the same relation to the sulphur, and so also do the two 
 hydrogen atoms. Whether one or the other of these hy- 
 drogen atoms be replaced by another atom or group of 
 atoms, the same product should result. Or, if one of the 
 hydrogen atoms is replaced by one group and the other 
 by another group, it should make no difference in which 
 order the two groups are introduced. The product should 
 be the same in the two cases. Thus, suppose one hydro- 
 gen atom is replaced by a group X, and the other by Y, 
 the product should be represented by the formula 
 
228 INORGANIC CHEMISTRY. 
 
 o o 
 
 II II 
 
 Y-O-S-O-X in one case, and by X-O-S-O-1 in the 
 
 II II 
 
 O O 
 
 other case. But if the formula given for sulphuric acid 
 is correct the two compounds are identical. By methods 
 which involve the use of apparently complex organic 
 compounds the two hydrogen atoms have been thus re- 
 placed by two different groups, first in one way and then 
 in the reverse order ; and the two products have been 
 found to be identical. Further, it has been shown that 
 when the two hydroxyl groups of sulphuric acid are re- 
 
 :C1 
 
 placed by chlorine, forming the compound SK Cl, and 
 
 the chlorine atoms then replaced by certain groups of 
 atoms, these groups are in direct combination with sul- 
 phur and not with oxygen. All the evidence points to 
 the conclusion that the view represented above is correct, 
 In attempting to determine the constitution of sulphur- 
 ous acid a new difficulty arises. Just as hydrogen which 
 is in combination with oxygen is replaceable by metals, 
 or is acidic, so, also, is hydrogen which is in combination 
 with sulphur. It is possible, therefore, to conceive of 
 two arrangements of the atoms composing sulphurous 
 acid, both representing dibasic acids. One of these ar^ 
 
 H 
 
 i 
 
 rangements is this, O=S-O-H, in which the sulphur is 
 
 II 
 O 
 
 O 
 
 n 
 
 represented as sexivalent ; the other is this, H-O-S-O-H, 
 in which the sulphur is represented as quadrivalent. 
 The facts that sulphurous acid is formed so readily by 
 simple contact of sulphur dioxide with water ; that it 
 breaks down as readily into sulphur dioxide and water ; 
 and that it takes up oxygen and sulphur so readily to 
 form sulphuric and thiosulphuric acids, seem to speak in 
 favor of the second of the above formulas. But, on the 
 
CONSTITUTION OF THE ACIDS OF SULPHUR. 229 
 
 other hand, certain facts established in the study of the 
 organic derivatives of sulphurous acid form a strong ar- 
 gument in favor of the first formula. It is possible by 
 means of reactions with certain compounds of carbon to 
 replace one of the metal atoms in a sulphite in such a way 
 
 ( E 
 
 as to produce a compound of the formula S< ONa, the 
 
 conduct of which is such as to show that in it the group 
 B is in direct combination with sulphur. As the com- 
 pound is formed apparently by direct replacement of a 
 sodium atom in the sulphite it appears that this sodium 
 atom was in combination with sulphur. Further, when 
 the second sodium is replaced by the group B a com- 
 
 (K 
 
 pound of the formula S-< OB is obtained, in which it 
 
 (o, 
 
 appears that one of the groups is in combination with 
 sulphur and the other with oxygen. Now, it is possible 
 
 (01 
 by starting with the compound S -< 01, which is made, as 
 
 (o 
 
 we shall see, by replacing one oxygen atom of sulphui 
 dioxide by two chlorine atoms, to introduce in the place 
 of the two chlorine atoms two groups, OB, and thus ob- 
 
 (OB 
 
 tain the compound S-< OB, which plainly has the same 
 
 (o 
 
 composition as the one represented by the formula 
 
 (E 
 
 S-< OB, but it has a different constitution. It was found 
 
 (o, 
 
 by experiment that the two compounds have different 
 properties, and it seems probable, therefore, that the two 
 formulas given represent the structure of the two com- 
 pounds. As the one which has one group B in combina- 
 tion with sulphur is obtained from sodium sulphite by 
 replacement of sodium, the conclusion seems to be justi- 
 fied that sulphurous acid has the constitution represented 
 
230 INORGANIC CHEMISTRY. 
 
 H f TT 
 
 i j -n- 
 
 by the formula O=S=O or SK OH, which may also be 
 
 6 i* 
 i 
 
 TT 
 
 written O 2 S <QTT' At the same time it appears quite pos- 
 
 sible that a sulphurous acid of the other constitution, 
 
 OTT 
 
 jj, maybe found, as there are innumerable ex- 
 
 amples furnished by chemistry of the existence of two 
 or more compounds of the same percentage composition 
 but different constitution. Two or more substances hav- 
 ing the same composition but different constitution are 
 said to be isomeric. Although examples of isomerism 
 are rare among the compounds of most elements, yet 
 among the compounds of carbon they are met with in 
 large numbers, and in studying these compounds a great 
 deal of attention has been paid to the phenomena of 
 isomerism. It is possible that the second form of sul- 
 phurous acid cannot exist on account of the tendency of 
 sulphur to act as a sexivalent element. This is, how- 
 ever, mere speculation, and, unless the suggestion can be 
 tested experimentally, it is of very little value. If sul- 
 phurous acid has the constitution above assigned to it 
 then the transformation of sulphurous acid into sulphuric 
 acid is not simply a direct combination of oxygen with 
 sulphur, but the act must involve a partial breaking down 
 of the sulphurous acid and a recombination of the con- 
 stituents thus : 
 
 O O 
 
 H-O-S-H + O = H-O-S-O-H. 
 
 ll I'l 
 
 O O 
 
 If this view is correct the oxygen displaces the hydrogen 
 and then combines with it and the sulphur. 
 
 If the action takes place in the same way with sulphur 
 it must be represented thus : 
 
CONSTITUTION OF THE ACIDS OF SULPHUR. 231 
 
 o o 
 
 H-O-S-H + S = H-O-S-S-H ; 
 
 II II 
 
 O O 
 
 o 
 
 II 
 and, according to this, the formula H-O-S-S-H repre- 
 
 II 
 O 
 
 sents the structure of thiosulphuric acid. The same con- 
 clusion regarding the structure of thiosulphuric acid is 
 reached by a consideration of one of the methods by 
 which its sodium salt is made. This is the method which 
 consists in treating a mixture of sodium sulphide and 
 sodium sulphite with iodine. The iodine simply extracts 
 two atoms of sodium from the two molecules, and it 
 seems probable that the residues of the two molecules 
 unite. If this is correct the following equation represents 
 what takes place : 
 
 Q .Na 
 
 5< Na S-Na 
 
 Na +2I = 
 
 O=S-0-Na 
 
 II O 
 
 O 
 
 and the constitution of thiosulphuric acid is represented 
 S-H 
 
 by the formula O=S-O-H, which is identical with that 
 
 II 
 O 
 
 given above. This latter method, however, it should be 
 remarked, might also be used as an argument in favor of 
 H 
 
 the constitution O=S-O-S-H for thiosulphuric acid ; for 
 
 II 
 O 
 
 if, when the iodine acts upon the sulphite, it extracts the 
 sodium atom which is in combination with oxygen, then 
 the union of the two residues would, in all probability, 
 
INORGANIC CHEMISTRY. 
 
 take place at this point, and the constitution of the acid 
 would be that represented by the last formula given. It 
 will be seen that our knowledge in regard to the structure 
 of thiosulphuric acid is very unsatisfactory. 
 
 The same remark may be made in regard to our knowl- 
 edge of the structure of the other acids of sulphur. The 
 method used in making the salts of tetrathionic acid is 
 suggestive, and if we knew the structure of thiosulphuric 
 acid, we might draw a conclusion in regard to that of 
 tetrathionic acid. The method consists in treating the 
 sodium salt of thiosulphuric acid with iodine. It appears 
 probable that of the two formulas above given for thio- 
 
 O 
 
 II 
 
 sulphuric acid this one, H-S-S-O-H, is to be preferred, 
 
 II 
 O 
 
 for the reason that it is more probable that when iodine 
 acts upon sodium thiosulphate it removes the sodium 
 atom which is in combination with sulphur, than that it 
 removes the one which is in combination with oxygen. 
 If this is true the above formula for thiosulphuric acid 
 follows. Now, further, reasoning in the same way, it 
 appears probable that when iodine acts upon sodium 
 thiosulphate it removes the sodium which is in combina- 
 tion with sulphur, and in this case the formation of tetra- 
 thionic acid must be represented in the following way : 
 
 O O 00 
 
 (O-S- 
 
 6 
 
 Na + Na 
 I I 
 
 -S-&-O-Na 
 
 6 
 
 = NaO-^-S-S-S-O-Na + 
 
 According to this, in tetrathionic acid there are four 
 sulphur atoms in combination by one affinity each. 
 This compound is, however, comparatively stable, while 
 thiosulphuric acid, in which a similar combination of 
 only two sulphur atoms is assumed, is extremely un- 
 stable. What value to attach to such considerations as 
 these we do not know. 
 
 Compounds of Sulphur with Oxygen. Sulphur, as has 
 been repeatedly stated, combines with oxygen in two pro- 
 
SULPHUR DIOXIDE. 233 
 
 portions, forming the oxides SO 2 and SO S , or sulphur di- 
 oxide and sulphur trioxide. Besides these two it also 
 forms a sesquioxide, S 2 O 3 , and a heptoxide, S a O 7 , but com- 
 paratively little is known in regard to the last two. The 
 one best known is sulphur dioxide. 
 
 Sulphur Dioxide, SO 2 . This, as has been seen, is 
 formed when sulphur is burned in the air or in oxygen ; 
 and It is also easily formed by reduction of the higher 
 oxides and acids of sulphur. Owing to the fact that 
 with water it forms sulphurous acid, it is frequently 
 called sulphurous anhydride. The methods for making 
 it were relerred to under sulphuric acid, and the reac- 
 tions involved were discussed then with a sufficient degree 
 of fulness. It need only be said here that in the labora- 
 tory the methods most commonly employed are : (1) Heat- 
 ing sulphuric acid with copper ; (2) heating the acid with 
 carbon (charcoal) ; and (3) heating the acid with sulphur. 
 When carbon is used two gases are formed, viz., carbon 
 dioxide and sulphur dioxide : 
 
 2H 2 SO 4 + C = CO 3 + 2SO 2 + 2H,O. 
 
 The gas can also be made by heating a mixture of a 
 metallic oxide and sulphur. Thus, when cupric oxide and 
 sulphur are heated together this reaction takes place : 
 
 2CuO + 28 = Cu 2 S + SO 2 . 
 
 Sulphur dioxide is a colorless, transparent gas, which 
 has a pungent, suffocating odor, familiar as the odor of 
 burning sulphur matches. It is poisonous, causing death 
 when inhaled in any quantity, and giving rise to bad 
 symptoms even in comparatively small quantities. It 
 does not readily give up its oxygen, so that burning 
 bodies are extinguished when introduced into it. It acts 
 something like water in this respect. It is more than 
 twice as heavy as air, its specific gravity being 2.26. 
 When sulphur is burned in oxygen gas the sulphur di- 
 oxide formed occupies the same volume as the oxygen 
 used up, so that there is no change in the volume. This 
 
234 INORGANIC CHEMISTRY. 
 
 can be shown by the experiment here described. In a 
 bent glass tube, of the form shown in Fig. 8, there is 
 placed a piece of sulphur, and the 
 tube is then half filled with pure 
 oxygen over mercury. On now 
 heating the tube at the part where 
 the sulphur is, this burns and is 
 converted into sulphur dioxide. 
 After the tube has cooled down to 
 FIG. s. the ordinary temperature the gas is 
 
 found to occupy the same volume as before. This will 
 be readily understood by the aid of the following con- 
 siderations : In the reaction 
 
 one molecule of sulphur dioxide is formed for every 
 molecule of oxygen used up. But a molecule of sulphur 
 dioxide in the form of gas occupies the same space as a 
 molecule of oxygen, so that, as the space occupied by the 
 sulphur in the experiment is insignificant, there is no 
 change in volume occasioned by the above reaction. 
 
 Sulphur dioxide dissolves in water, as we have seen, 
 and forms a liquid in which, judging by its conduct, sul- 
 phurous acid is present. 
 
 The gas is easily liquefied by cold alone. It is only 
 necessary for this purpose to pass the dry gas through a 
 tube surrounded by a freezing mixture of ice and salt. 
 The liquid changes rapidly into gas under ordinary pres- 
 sure at the ordinary temperature: In this change so 
 much heat is absorbed that a temperature of about 60 
 can be produced by means of it ; and a portion of the 
 liquid can be solidified. 
 
 Sulphur dioxide is very stable. If heated to 1200 
 under pressure, however, it breaks down into sulphur 
 trioxide and sulphur : 
 
 '* 3SO 2 = 2SO 3 + S. 
 
 When conducted into solutions of bases or of carbonates 
 the corresponding sulphites are formed : 
 
SULPHUR TRIOXIDE. 235 
 
 2KOH + SO 2 = K 2 SO 3 + H 2 O ; 
 K 2 C0 3 + SO i 
 
 Under certain conditions, as when the gas is passed 
 into a hot solution of an alkali carbonate, a salt of the 
 general formula M 2 S 2 O 5 is formed. This bears to the 
 sulphite the same relation that the pyrosulphate bears 
 to the sulphate : 
 
 and 2KHS0 4 = K 2 S 2 O 7 + H 2 O. 
 
 Sulphur dioxide is used extensively for the purpose of 
 bleaching silk, wool, straw, and basket-ware. In order 
 that it may bleach, however, water must be present, so 
 that it appears that the true bleaching agent in this case 
 is sulphurous acid and not the dioxide. When we con- 
 sider that sulphur dioxide does not readily take up nor 
 give up oxygen, while sulphurous acid does readily take 
 it up, the necessity of having water present in the 
 bleaching process at once becomes apparent. The 
 bleaching in some cases certainly consists in abstracting 
 oxygen from the colored substances, and thus converting 
 them into colorless products. In other cases it is due to 
 the formation of compounds of sulphurous acid with the 
 dye-stuffs. 
 
 Sulphur dioxide is not only a bleaching agent like 
 chlorine, but like chlorine it is also a disinfectant. It 
 has to some extent the power to destroy the organisms 
 which cause changes in organic substances. It prevents 
 fermentation and is therefore used as a preservative. Its 
 power to destroy the germs of disease, that is, to disinfect, 
 is not as great as is frequently supposed. Much larger 
 quantities are necessary for this purpose than are com- 
 monly used. 
 
 Sulphur Trioxide, SO 3 . This compound is made by 
 passing sulphur dioxide and oxygen together over heated 
 platinum in a finely divided state. It is obtained most 
 readily by heating disulphuric acid, which breaks up 
 
236 INORGANIC CHEMISTRY. 
 
 easily into sulphur trioxide and ordinary sulphuric acid 
 according to the equation 
 
 H a S a 7 = H a S0 4 + S0 3 . 
 
 Similarly, the acid sulphates of the alkali metals yield 
 the corresponding normal sulphates and sulphur trioxide : 
 
 2NaHS0 4 = Na 2 SO 4 + SO 3 + H 2 O. 
 
 It is now manufactured on the large scale by passing sul- 
 phur dioxide and oxygen together over asbestos covered 
 with finely divided platinum, and the product thus ob- 
 tained is passed into ordinary sulphuric acid for the pur- 
 pose of making " solid sulphuric acid " which, as has 
 been stated, is almost pure disulphuric acid, H 2 S 2 O 7 . 
 
 Sulphur trioxide is a white crystallized solid which 
 appears to exist in two modifications. The one is a liquid 
 at ordinary temperatures, but it solidifies at about 16. 
 According to the latest investigations there is but one 
 modification of the oxide. It is a solid which melts at 
 14.8, forming a liquid which boils at 46. In contact 
 with the air the oxide gives off thick fumes which are 
 partly due to the great power of the compound to com- 
 bine with water. Water acts with violence upon it, the 
 heat evolved in the act being 39,170 cal. It also acts 
 upon substances containing hydrogen and oxygen in much 
 the same way that concentrated sulphuric acid does, 
 charring them by abstracting the hydrogen and oxygen. 
 It acts, however, more violently in this way than sulphuric 
 acid does. With water it forms sulphuric acid, and it is, 
 therefore, called sulphuric anhydride. The reaction in- 
 volved in passing from sulphur trioxide to sulphuric acid 
 is of a kind which, as we have seen, is frequently met with 
 both with acidic oxides or anhydrides, and with basic or 
 metallic oxides ; and it is desirable that it should here 
 be studied a little more carefully than it has yet been. 
 What we know is that when sulphur trioxide acts upon 
 water there is a great deal of heat evolved, and com- 
 pounds of different composition are obtained. The com- 
 position of these compounds is represented by the formu- 
 
ACIDIC OXIDES AND WATER. 237 
 
 las SO 3 + H 2 O, SO 3 + 2H 2 O, and SO 3 + 3H 2 O. So, too, 
 when calcium oxide or lime, CaO, acts upon water, there 
 is great evolution of heat, and a compound is formed 
 the composition of which is represented by the formula 
 CaO 4- H 2 O. But these formulas do not attempt to give 
 any account of what takes place in the chemical acts re- 
 ferred to. That the water is not present in the com- 
 pounds as water seems evident, in the first place from 
 the conduct of the compounds, and in the second place 
 from the amount of heat evolved in the act of combina- 
 tion. Now, taking the chemical conduct of the substances 
 into consideration, they appear to contain hydrogen in 
 combination with oxygen, and their conduct becomes 
 comprehensible on the supposition that the group known 
 as hydroxyl, (-O-H), is present. This view has been 
 found to be in accordance with a large number of facts, 
 and it is of great assistance in dealing with these facts. 
 The view is distinctly this : When an acidic oxide acts 
 upon water it is converted into a hydroxyl compound 
 which has acid properties, as shown in this equation : 
 
 roH 
 TT o OH 
 
 II H,U 
 
 O^S^O + H 2 = S 
 
 TT r\ 
 
 H 2 
 
 LOH 
 
 According to this, each molecule of water is decomposed 
 and each hydrogen atom in the resulting compound 
 is in combination with oxygen. The same kind of action 
 takes place in the case of some basic oxides, as shown, 
 for example, in the case of calcium oxide : 
 
 By means of certain reagents which will be taken up 
 later it is possible to replace oxygen in compounds by 
 chlorine, two chlorine atoms taking the place of one 
 oxygen atom. So, also, when the oxygen of the hydrox- 
 ides is replaced by chlorine, the result is that each 
 
238 INORGANIC CHEMISTRY. 
 
 hydroxyl group is replaced by one atom of chlorine. This 
 is easily understood. For, if in calcium hydroxide each 
 oxygen atom should be replaced by two chlorine atoms, 
 the result would be a combination of atoms represented 
 
 C1C1 -H 
 thus, C a <cici-H an ^ fr m ^ s > hydrochloric acid 
 
 would be given off, leaving calcium chloride : 
 
 Cl 
 
 O TT 
 
 The result, therefore, of treating the hydroxide C/a<Q_Tj- 
 
 in such a way as to replace each of the two oxygen atoms 
 by two chlorine atoms is the formation of a compound 
 
 Cl 
 Ca<, which is derived from the hydroxide by replacing 
 
 each hydroxyl group by one atom of chlorine. These 
 facts lend support to the hydroxyl hypothesis. 
 
 Just as sulphur trioxide acts upon water to form sul- 
 phuric acid, so it acts upon metallic oxides, forming 
 sulphates. Thus with calcium oxide it forms calcium 
 sulphate : 
 
 CaO + SO 3 = CaSO 4 . 
 
 It combines also with hydrochloric acid, forming the 
 compound SO 3 HC1. The action in this case is analogous 
 to that which takes place with water, and is represented 
 thus : 
 
 O O 
 
 O=S=0 + H-C1 = 0-S-O-H. 
 
 I 
 Cl 
 
 The analogy between this and the action with water will 
 be apparent by a consideration of the following equation : 
 
 O O 
 
 O=S=O + H-O-H = O=S-O-H. 
 
 I 
 
 O 
 
 I 
 
 H 
 
ACID CHLORIDES OF SULPHUR. 239 
 
 The chlorine compound appears to be ordinary sul- 
 phuric acid in which one hydroxyl group has been replaced 
 by a chlorine atom. When the compound is treated with 
 water hydrochloric acid is given off and sulphuric acid 
 is formed, thus : 
 
 ( 01 ( OH 
 
 S^ OH + HOH = S-J OH + HC1. 
 (O, (O, 
 
 Acid Chlorides of Sulphur. In the introduction to 
 this chapter reference was made to certain compounds 
 which sulphur forms with oxygen and chlorine, known 
 as acid chlorides, because when brought in contact with 
 water they form acids. Strictly speaking, all the com- 
 pounds of sulphur and chlorine are of this order ; the 
 name is, however, generally applied to those compounds 
 which are derived from the acids by replacement of the 
 hydroxyl by chlorine. Thus the acid chloride of sul- 
 
 01 
 phuric acid is SO 2 < C j, which is plainly sulphuric acid in 
 
 which the two hydroxyl groups have been replaced by 
 chlorine. So, also, the compound SO< is related in 
 
 the same way to a sulphurous acid of the formula 
 
 OH 
 
 , which, however, does not appear to be ordi- 
 
 nary sulphurous acid. But the existence of the com- 
 
 Pl PI 
 
 pounds SO 2 <^ and SO<QJ, which are to be regarded 
 
 respectively as sulphuric and sulphurous acids in which 
 both hydroxyls are replaced by chlorine, suggests the 
 
 Cl PI 
 
 existence of the two compounds SO 2 < and 
 
 derived from the acids by the replacement of one hy- 
 droxyl in each by chlorine. The former compound is 
 known, the latter is not. 
 
 Thionyl Chloride, SOC1 2 , can be made from sulphur 
 dioxide by replacing one oxygen atom by two chlorine 
 atoms. This is effected by treating sulphur dioxide with 
 phosphorus pentachloride, PC1 5 , a compound which 
 readily gives up a part or all of its chlorine and takes up 
 
240 INORGANIC CHEMISTRY. 
 
 oxygen in its place, and is therefore extensively used in 
 chemistry for the purpose of replacing oxygen by chlo- 
 rine. With sulphur dioxide the reaction takes place 
 thus: 
 
 SO 3 + PC1 6 = SOC1 2 + POC1 3 . 
 
 The compound is also formed by treating sodium sul- 
 phite with phosphorus pentachloride, when this reaction 
 takes place : 
 
 SO(ONa) 2 + 2PC1 5 = SOC1 2 + 2NaCl + 2POC1 3 . 
 
 In order to understand this reaction it appears neces- 
 sary that sodium sulphite should be represented by the 
 formula given, which, as will be seen, is not in accordance 
 with the conclusion reached regarding the structure of 
 sulphurous acid. It may be that the reaction is more 
 complicated than here represented, and that sulphur 
 dioxide is first given off, and that this then acts upon the 
 pentachloride. 
 
 Thionyl chloride is a liquid with a strong characteristic 
 odor. It acts readily upon water, forming sulphur dioxide 
 and hydrochloric acid according to the equation 
 
 SOC1 2 + H 2 O = SO 2 + 2HC1. 
 
 Sulphuryl Chloride, SO 2 C1 2 is formed by the direct 
 action of chlorine upon sulphur dioxide in the sunlight, 
 the action being similar to that which takes place when 
 sulphur dioxide and oxygen unite to form the trioxide : 
 
 SO 2 + = SO 3 . 
 
 01 
 It is best prepared by heating the compound SO a 
 
 to 170-180, when the following reaction takes place : 
 
 2SO - < OH = H * SO < + SO * C1 *' 
 
 It is a liquid which is easily decomposed by water, as 
 represented in the equation 
 
 S0 3 C1 2 + 2H 2 = S0 2 (OH) 2 + 2HC1. 
 
OXYGEN COMPO UNDS OF SELENIUM AND TELL URIUM. 241 
 With half the quantity of water required to effect 
 the above decomposition cJdorsulphuric acid, S0 a </yrr, 
 is formed : 
 
 so *<ci + H = 8 <oa + HCL 
 
 Chlorsulphuric Acid, or Sulphuryl-hydroxyl Chloride, 
 SO a <Q H , is also formed, as has been shown, by the direct 
 action of hydrochloric acid upon sulphur trioxide : 
 
 and, further, by the action of phosphorus pentachloride 
 upon sulphuric acid : 
 
 SO <OH + PC1 = 8 .<OH + POC1 + HCL 
 
 Like sulphuryl chloride it is decomposed by water, yield- 
 ing hydrochloric acid and sulphuric acid. 
 
 Compounds of Selenium and Tellurium with Oxygen and 
 with Oxygen and Hydrogen. For the purposes of this 
 book it is not necessary to go into details in regard to 
 the compounds of selenium and tellurium. In the intro- 
 duction to this chapter it was stated that selenium and 
 tellurium form compounds with oxygen corresponding 
 to sulphur dioxide and sulphur trioxide. Both of these 
 oxides of tellurium are known, together with a third of 
 the composition represented by the formula, TeO, while 
 only one oxide of selenium, the dioxide, SeO 2 , is known. 
 On the other hand, the acids of selenium and tellurium, 
 corresponding to sulphurous and sulphuric acids, are 
 known. 
 
 Selenious Acid, H 2 SeO 3 . This compound is formed by 
 oxidation of selenium in presence of water or by dissolv- 
 ing selenium dioxide in water. It is a crystallized solid, 
 which attracts moisture from the air, and when heated 
 breaks down into selenium dioxide and water. It forms 
 two series of salts, the acid selenites, MHSeO 3 , and the nor- 
 
242 INORGANIC CHEMISTRY. 
 
 mal selenites, M a SeO 3 . While in composition the acid 
 and its salts are analogous to sulphurous acid and the 
 sulphites, in conduct there is a marked difference. Sul- 
 phurous acid tends, as we have seen, to take up more 
 oxygen and form sulphuric acid, while selenious acid 
 gives up its oxygen with great ease and yields selenium. 
 It is evident that the affinity of selenium for oxygen is 
 much less than that of sulphur for oxygen. When a 
 solution of selenious acid is treated with sulphur dioxide 
 the acid is reduced to selenium : 
 
 H a Se0 3 + 2SO 2 + H 2 = 2H 2 SO 4 + Se. 
 
 The same method of investigation which leads to the 
 conclusion that sulphurous acid has the formula 
 
 TT 
 
 SO a <Q-rr, when applied to selenious acid, leads to the 
 
 OH 
 
 conclusion that this has the constitution SeO<Q-rr. 
 
 Selenic Acid, H 2 SeO 4 , is formed by the action of power- 
 ful oxidizing agents like saltpeter, or chlorine or bromine 
 in water solution on selenium. The action with chlorine 
 is represented thus : 
 
 Se + 601 + 4H 2 O = H 2 SeO 4 + 6HC1. 
 
 It is a liquid resembling sulphuric acid. It cannot, how- 
 ever, be obtained in pure condition on account of its ten- 
 dency to give up oxygen and pass over into selenious 
 acid, in which respect it plainly differs markedly from 
 sulphuric acid. It gives up its oxygen to other substances 
 much more readily than sulphuric acid does. This is 
 seen in its action upon hydrochloric acid, which takes 
 place as represented in the equation 
 
 H 2 Se0 4 + 2HC1 = H 2 SeO 3 + 01, + H 2 O. 
 
 Sulphuric acid does not act upon hydrochloric acid, 
 but it does act upon hydrobromic and hydriodic acids, 
 and its action upon hydrobromic acid is very similar to 
 
ACID CHLORIDES OF SELENIUM. 243 
 
 that which takes place between selenic acid and hydro- 
 chloric acid : 
 
 H 3 S0 4 + 2HBr = SO 2 + Br 2 + 2H 2 O ; 
 
 the only essential difference is that the sulphurous acid 
 breaks down into sulphur dioxide and water, while the 
 selenious acid remains in solution as such. 
 
 Selenium Dioxide, SeO 2 . This analogue of sulphur 
 dioxide is made by burning selenium, or by heating 
 selenious acid. It is a solid crystallized substance 
 which can be sublimed without undergoing decompo- 
 sition and without melting. As we have seen, it dissolves 
 readily in water, forming selenious acid. It loses its oxy- 
 gen readily in contact with substances which have the 
 power to unite with oxygen, and it is thus reduced to 
 selenium. Hence when it is sublimed particles of dust 
 which may be present in the vessel effect partial reduction, 
 and the product, instead of being white, as the oxide is 
 when pure, is colored by the particles of selenium. 
 
 Acid Chlorides of Selenium. Selenium combines with 
 oxygen and chlorine, forming compounds analogous to the 
 acid chlorides of sulphur. The principal one of these is 
 sdenyl chloride, SeOCl 2 , which is formed by bringing to- 
 gether selenium tetrachloride and selenium dioxide : 
 
 SeCl 4 + SeO 2 = 2SeOCl 2 . 
 
 By water it is decomposed, forming selenious and hydro- 
 chloric acids : 
 
 Just as sulphur trioxide combines with hydrochloric 
 acid, so also does selenium dioxide. It forms products 
 which are represented by the formulas SeO 2 .2HCl and 
 SeO 2 .4HCl. It is quite possible that the former is a 
 
 Cl >Se< OH 
 chloride of selenious acid of the constitution Cl OH ' 
 
244 INORGANIC CHEMISTRY. 
 
 and that the second should be represented by the for. 
 mula -rr_/ni 2 \>Se<QTT, a few compounds of this kind 
 
 being known, as will be pointed out. 
 
 Tellurious Acid, H 2 TeO 3 . Tellurious acid is formed by 
 treating tellurium tetrachloride with water. It is possible 
 that the first action causes the formation of normal tel- 
 lurious acid, Te(OH) 4 , and that this then breaks down 
 into tellurious acid, H 2 TeO 3 , and water : 
 
 Cl HOH f OH 
 
 T , Cl , HOH_ I 
 
 le ^ Cl +HOH - 1 
 
 Cl HOH L OH 
 
 (OH 
 Te^g| = Te]oH + H 2 0. 
 
 OH ^ 
 
 The potassium salt of the acid is formed by melting 
 together tellurium dioxide and potassium carbonate : 
 
 K 2 CO 3 + TeO 2 = K 2 TeO 3 + CO,. 
 
 If the salt thus formed is dissolved in water and nitric 
 acid added to the solution, tellurious acid is thrown down: 
 
 K 2 TeO 3 + 2HNO 3 = 2KNO 3 + H 2 TeO 3 . 
 
 It is a solid which easily loses water and is thus trans- 
 formed into tellurium dioxide. 
 
 Telluric Acid, H 2 TeO 4 . This acid is formed by melting 
 tellurious acid with saltpeter and other oxidizing agents. 
 When the solution of the acid is evaporated to crystal- 
 lization the solid compound deposited has the compo- 
 sition H 6 TeO 6 , and, according to what was learned in 
 studying sulphuric acid, it appears probable that this is 
 normal telluric acid, Te(OH) 6 . When normal telluric 
 acid is heated to a little above 100 it loses water and is 
 
OXIDES OF TELLURIUM. 245 
 
 transformed into the acid H 2 TeO 4 , corresponding to or- 
 dinary sulphuric acid, from which most of the tellurates 
 are derived : 
 
 Te(OH) 6 = Te0 2 (OH) 2 + 2H 2 O. 
 
 Heated higher, to about 160, the acid is decomposed 
 into tellurium trioxide and water : 
 
 Te0 2 (OH) a = TeO 3 + H 2 O. 
 
 Although most of the tellurates are simple salts of the 
 acid H 2 TeO 4 , others are derived from more complex forms 
 of the acid. One of these is analogous to disulphuric acid. 
 Another is derived from four molecules of telluric acid, 
 Te(OH) 6 , by loss of eleven molecules of water. 
 
 Oxides of Tellurium. Tellurium monoxide is formed by 
 heating sulphur trioxide and tellurium together in a 
 vacuum. Tellurium dioxide is formed by burning tellu- 
 rium or by oxidizing it with nitric acid ; and, further, by 
 the decomposition of tellurious acid by heat. It crystal- 
 lizes and is but slightly soluble in water. 
 
 The trioxide, TeO 3 , is formed by heating telluric acid 
 to a high temperature. Its conduct is entirely different 
 from that of sulphur trioxide. While the latter acts with 
 violence upon water, and readily upon metallic oxides 
 and hydrochloric acid, the former does not act readily 
 upon any of these substances. It is insoluble in hot as 
 well as cold water. 
 
 Sulphotelluric Acid is an example of the sulphur acids 
 referred to on p. 141. While the acid itself is not known, 
 a potassium salt of the formula K 2 TeS 4 is known. This 
 is plainly analogous to the salt of the oxygen acid, 
 K 2 TeO 4 , differing from it only in containing sulphur in 
 place of the oxygen. 
 
 FAMILY VI, GBOUP A. 
 
 Group A, Family VI, includes chromium, molyb- 
 denum, tungsten, and uranium. All of these show 
 some resemblance to the elements of the sulphur group, 
 
246 INORGANIC CHEMISTRY. 
 
 but they also appear in entirely different characters, 
 forming compounds of a kind unknown among the 
 derivatives of sulphur and its analogues. The relation 
 which these elements bear to sulphur is much like the 
 relation which manganese bears to chlorine. The re- 
 semblance to sulphur is seen mainly in the formation of 
 acids of the formulas H 2 CrO 4 , H 2 MoO 4 , H 2 WO 4 , and 
 H 2 UO 4 ; and the oxides CrO 3 , MoO 3 , WO 3 , and UO 3 . Most 
 of these acids yield complicated derivatives, all of which 
 can, however, be explained by the same method as that 
 used in the case of periodic acid. The common salts of 
 chromic acid are derived from dichromic acid, which is 
 analogous to disulphuric acid. They have the general 
 formula M 2 Cr 2 O 7 . So, too, salts of molybdic acid are 
 known which are derived from the simple form of the 
 acid, H 2 MoO 4 , and others which are derived from a dimo- 
 lybdic acid, H 2 Mo 2 O 7 , and from more complicated forms. 
 Tungsten has a wonderful power of forming complex 
 acids. All of them, however, can be referred to the 
 simple form H 2 WO 4 . And, finally, uranic acid forms salts 
 which for the most part are derived from diuranic acid, 
 H 2 U 2 O 7 . All the most important of these compounds 
 will be taken up later. 
 
 When the acids of chromium, molybdenum, tungsten, 
 and uranium lose oxygen they form compounds which 
 have little or no acid character. The lower oxides of 
 chromium form salts with acids, and these bear a general 
 resemblance to the salts of aluminium, iron, and manga- 
 nese. The chromates lose their oxygen quite readily 
 when acids are present with which the chromium can 
 enter into combination in its capacity as a base-forming 
 element. Thus, when potassium chromate K 2 CrO 4 is 
 treated with hydrochloric acid in the presence of some- 
 thing which can take up oxygen, decomposition takes 
 place thus : 
 
 2K 2 Cr0 4 + 10HC1 = 4KC1 + 2CrCl 3 + 5H 2 O + 30. 
 
 With sulphuric acid the action takes place as repre- 
 sented in this equation : 
 
 2K 2 Cr0 4 + 5H 2 S0 4 = 2K 2 SO 4 + Cr a (S0 4 ) 3 + 5H 2 + 3O. 
 
COMPOUNDS OF CHROMIUM, MOLYBDENUM, ETC 247 
 
 In both these cases the chromium enters into combina- 
 tion as a trivalent base-forming element, taking the place 
 of three atoms of hydrogen in hydrochloric acid in the 
 first case, and of three atoms of hydrogen in the sulphuric 
 acid in the second. Molybdenum and tungsten do not 
 form salts of this character ; indeed they seem to be 
 practically devoid of basic properties. Uranium, on the 
 other hand, forms some curious salts which differ from 
 the simple metallic salts which we commonly have to 
 deal with. These are the so-called uranyl salts, which 
 are regarded as acids in which the hydrogen is either 
 wholly or partly replaced by the group UO 2 , which is 
 bivalent. Thus, the nitrate has the formula (UO 2 )(NO 3 ) 2 , 
 the sulphate is (UO 2 )SO 4 , etc. These salts are derived 
 from the compound UO 2 (OH) 2 acting as a base, whereas 
 this compound has also distinctly acid properties. 
 
CHAPTER XV. 
 
 NITROGEN THE AIR ARGON. 
 
 NITROGEN, N (At. Wt. 13.93). 
 
 General. Nitrogen bears to a group of elements rela- 
 tions very similar to those which oxygen bears to the sul- 
 phur group, and fluorine to the chlorine group. There 
 are easily recognized resemblances between it and the 
 members of the group, and yet there are some marked 
 differences. As has been stated, and as is seen from its 
 position in the periodic system, nitrogen is trivalent 
 towards hydrogen, as shown in the compound NH 3 , while 
 it is both trivalent and quinquivalent towards oxygen, as 
 appears to be shown in N 2 O 3 and N 2 6 . The principal 
 hydrogen compound, ammonia, is entirely different in 
 character from those of chlorine and sulphur, for, while 
 these are acid, ammonia has in a marked way the char- 
 acter of a base, acting, however, in a peculiar way upon 
 acids to form salts. The two oxides above referred to 
 are acidic, forming the acids HNO 2 and HNO 3 , which are 
 known as nitrous and nitric acids respectively. 
 
 Occurrence of Nitrogen. It was discovered by Lavoi- 
 sier and Scheele towards the end of the last century that 
 the air consists of two gases, one of which is oxygen, and 
 they showed that when the oxygen is removed the gas 
 which is left has not the power to support combustion 
 nor to support respiration. This gas was first called 
 azote (from a, privitive, and goon/co?, life), and this name 
 is still retained in France, the symbol in use in that 
 country being Az, whereas in all others the symbol is N. 
 This is the only case in which there is a difference of 
 usage in respect to the symbols of the chemical elements 
 in different countries. The name nitrogene was given to 
 it later, from the fact that it is a constituent of niter or 
 saltpeter, KNO 3 (nitrum, saltpeter, and yeveiv, to pro- 
 
 (248) 
 
PREPARATION OF NITROGEN. 249 
 
 duce), and this is the origin of the English name nitro- 
 gen. Not only is nitrogen found free in the air, but it is 
 found in combination in a large number of substances in 
 nature. It is found in the nitrates, or salts of nitric acid, 
 particularly as the potassium salt KNO 3 , and the sodium 
 salt NaNO 3 , which occurs in enormous quantities in 
 Chili, and is therefore known as Chili saltpeter. It is 
 also found in the form of ammonia, which is a compound 
 of nitrogen and hydrogen of the formula NH 3 . Ammonia 
 occurs in small quantity in the air, and is formed under 
 a variety of conditions, to which reference will be made 
 when the substance is treated. Nitrogen occurs, further, 
 in combination in many animal substances. 
 
 Preparation. The most convenient way to prepare 
 nitrogen is by burning in a closed vessel something which 
 does not give a gaseous product of combustion ; or by 
 passing air over something which has the power to unite 
 with oxygen. The best substance to use for the first 
 purpose is phosphorus, which burns readily and yields a 
 solid product, soluble in water. It is only necessary, 
 therefore, to place a piece of phosphorus in a floating 
 vessel on the surface of water, set fire to it, and immedi- 
 ately place over it a closed bell-jar. As soon as the oxygen 
 is used up the combustion stops, and the vessel then 
 contains the residual nitrogen, and the walls are covered 
 with a thin layer of phosphorus pentoxide, P 2 O 5 . This 
 is soon converted by the water into phosphoric acid, 
 which dissolves. Another convenient method for pre- 
 paring nitrogen consists in passing air over copper 
 heated in a tube. The copper takes up the oxygen 
 readily, and the nitrogen passes on. Another good 
 method consists in exposing to the air copper turnings 
 partly covered with a solution of ammonia in a vessel so 
 arranged as to allow free access of air while the escape 
 of the gas in the vessel is prevented. This mixture ab- 
 sorbs oxygen slowly at the ordinary temperatures. 
 Nitrogen can also be made from other substances than 
 the air. Thus, when chlorine is passed into a water 
 solution of ammonia this reaction takes place : 
 NH S + 3C1 = N + 3HC1 ; 
 
250 INORGANIC CHEMISTRY. 
 
 but the hydrochloric acid combines at once with ammonia 
 to form ammonium chloride, NH 4 C1 : 
 
 NH 3 + HC1 = NH 4 C1 ; 
 
 so that the only gaseous product is nitrogen. This ex- 
 periment is more or less dangerous, for if all the ammonia 
 should be used up, and the passage of chlorine continued, 
 a compound of nitrogen and chlorine which is extremely 
 explosive is formed. Finally, nitrogen can be made by 
 heating ammonium nitrite, NH 4 NO 2 , either dry or in solu- 
 tion. The hydrogen and oxygen of the compound unite 
 to form water and the nitrogen is set free : 
 
 The nitrogen prepared from the air is never pure, as 
 there are always present in the air other substances be- 
 sides nitrogen and oxygen ; and while some of these can 
 be removed without serious difficulty, others cannot be. 
 Properties. Nitrogen is a colorless, tasteless, inodor- 
 ous gas. It has been converted into a liquid by subject- 
 ing it to a very low temperature and high pressure. The 
 liquid solidifies at 203. A liter of nitrogen under 
 standard conditions weighs 1.250 grams. Its specific 
 gravity (air 1) is 0.971. It does not support combus- 
 tion, nor does it burn. This latter fact is obvious, for, if 
 nitrogen had the power to combine with oxygen when 
 the temperature of the mixture is elevated, it is plain 
 that this process of combustion would long ago have taken 
 place, leaving one or the other of the two gases and the 
 product of combustion as the constituents of the air. Ni- 
 trogen not only does not combine with oxygen readily, but 
 it does not combine readily with any other element ex- 
 cept at very high temperature, and then with only a few. 
 Just as it does not support combustion, so also it does 
 not support respiration. Animals would die in it, not on 
 account of any active poisonous properties possessed by 
 it, but for lack of oxygen. In the air it serves the useful 
 purpose of diluting the oxygen. If the air consisted only 
 of oxygen, all processes of combustion would certainly be 
 much more active than they now are. What effect the 
 
THE AIR. 251 
 
 continued breathing of oxygen would have upon animals 
 it is impossible to say. 
 
 The Air. The atmosphere of the earth, commonly 
 called the air, consists essentially of the two elements 
 nitrogen and oxygen in the proportion of 79 volumes of 
 nitrogen to 21 volumes of oxygen, or, by weight, of 77 
 per cent of nitrogen and 23 per cent of oxygen. Wher- 
 ever air has been collected and analyzed it has been 
 found to have practically the same composition. Never- 
 theless very accurate analyses have shown that the com- 
 position of the air is subject to slight variations. To de- 
 cide whether the air is a chemical compound or a me- 
 chanical mixture requires a careful examination of a 
 number of facts. The evidence may be summed up as 
 follows : 
 
 (1) If nitrogen and oxygen are mixed together the mix- 
 ture conducts itself in exactly the same way as air. The 
 mixing is not attended by any phenomena indicating 
 chemical action. Generally the chemical combination of 
 two elements is accompanied by an evolution of heat, and 
 whenever a chemical act takes place there is some 
 change in the temperature of the substances. When 
 nitrogen and oxygen are brought together there is no 
 change in the temperature of the gases. 
 
 (2) The composition of a chemical compound is con- 
 stant. The law of definite proportions is founded upon 
 a very large number of observations, and in all cases in 
 which we have independent evidence that chemical action 
 takes place it is found that the substances combine in 
 exactly the same proportions to form the same product. 
 Variation in the composition of a chemical compound is 
 not known. The composition of the air varies slightly, 
 according to circumstances, and this fact may be regarded 
 as evidence that the air is not a chemical compound. 
 
 (3) Air dissolves somewhat in water. If air which *s 
 in solution in water is pumped out and analyzed, it is 
 found to have a different composition from that of or- 
 dinary air. Instead of containing nearly 4 volumes of 
 nitrogen to 1 of oxygen, it contains only 1.87 volumes of 
 nitrogen to 1 of oxygen. The proportion of oxygen is 
 
252 INORGANIC CHEMISTRY. 
 
 much larger in the air which has been dissolved in water 
 than it is in ordinary air. This is due to the fact that 
 oxygen is more soluble in water than nitrogen. There- 
 fore, when air is shaken with water, relatively more oxy- 
 gen than nitrogen is dissolved. If the gases were in 
 chemical combination we should expect the compound to 
 dissolve as such and without change of composition. 
 
 The above evidence shows that nitrogen and oxygen 
 are not combined chemically in the air, but that they are 
 simply mixed together. 
 
 As enormous quantities of oxygen are constantly em- 
 ployed in the processes of respiration of animals, com- 
 bustion, and various kinds of decay, the question will 
 suggest itself : Is the quantity of oxygen in the air decreas- 
 ing ? In regard to this point some ingenious calculations 
 have been made the results of which are reassuring. An 
 approximate estimate of the extent of the atmosphere, 
 and therefore of the supply of oxygen, can easily be made. 
 Assuming that the population of the earth is 1000 million 
 human beings, the quantity of oxygen used by them in 
 respiration in a year would amount only to about 
 swornnr P ar * f ^ ne supply. Suppose, further, that for 
 all other purposes nine times as much oxygen is required 
 as for the respiration of human beings, then the total 
 amount used up in a year would be only g-g-gVoT ^ * ne 
 whole supply. In 3800 years the decrease in the amount 
 of oxygen in the air would be only 1 per cent. Whether 
 there has been such a decrease or not it is impossible to 
 say, as it is only within a comparatively few years that 
 accurate analyses have been made. It appears probable, 
 however, from other considerations that the quantity of 
 oxygen in the air is not decreasing. It is known that 
 the process of plant life involves a giving off of oxygen 
 which comes from other compounds. The plants have 
 the power to decompose the carbon dioxide found in the 
 air, and they utilize the carbon and a part of the oxygen, 
 but another part they give back to the air, so that in the 
 process of vegetable growth we have a constant source 
 of supply of oxygen. 
 
ANALYSIS OF AIR. 253 
 
 Analysis of Air. The earliest examinations of Priest- 
 ley, Lavoisier, and Scheele were made by burning sub- 
 stances in air contained in closed vessels. They con- 
 cluded that the air is made up of I oxygen and f nitrogen 
 by bulk. In order to determine the composition of the 
 air to-day, we should proceed as follows : A qualitative 
 examination would easily show the presence of nitrogen 
 and oxygen. If a solution of calcium hydroxide, Ca(OH) 2 , 
 which is known as lime-water, or a solution of barium 
 hydroxide, Ba(OH) 2 , is exposed to the air it becomes 
 turbid, and a precipitate is formed. Neither nitrogen nor 
 oxygen nor an artificially prepared mixture of the two 
 gases can produce this change. It has been shown that 
 the change is due to the presence in the air of a small 
 quantity of the gaseous compound, carbon dioxide, CO 2 . 
 If calcium chloride or phosphorus pentoxide is exposed 
 to the air it soon becomes moist and after a time turns 
 liquid. This effect has been found to be due to the pres- 
 ence of water vapor in the air. Nitrogen obtained from 
 the air by passing the latter over copper, which abstracts 
 the oxygen, has been shown to contain a small amount of 
 an extremely inert gas, argon (which see). By other 
 methods which need Hot be considered here it can be 
 shown that there are many other substances in the air 
 besides those mentioned. Among them are ammonia, 
 hydrogen dioxide, and organic matters of various kinds, 
 including a large variety of germs the presence of 
 which can be detected by the changes which exposure to 
 the air produces in certain liquids, as milk and fruit 
 juices. 
 
 Having thus learned what the chief constituents of the 
 air are, the next thing is to determine in what quantities 
 they are present, or to make a quantitative analysis of 
 the air. For this purpose advantage may be taken of 
 the fact that phosphorus when exposed to the air at 
 ordinary temperatures combines slowly with the oxygen, 
 leaving the nitrogen. If, therefore, a piece of ordinary 
 phosphorus is inserted into a measured volume of air 
 contained in a graduated glass tube over water or mer- 
 cury, a diminution in volume will take place slowly. If, 
 
254 INORGANIC CHEMISTRY. 
 
 in the course of a few hours, the volume is again meas- 
 ured, the difference will give the volume of oxygen 
 absorbed, while the gas remaining is nitrogen. Of 
 course, in this case as in all others in which gas volumes 
 are measured, corrections for temperature, pressure, and 
 tension of aqueous vapor must be made. 
 
 Another method by which the ratio between the nitro- 
 gen and oxygen in air can be determined is that which 
 was first employed by Dumas and Boussingault. It con- 
 sists in passing air over heated copper, collecting and 
 measuring the nitrogen, and weighing the copper oxide. 
 The apparatus is arranged as shown in Fig. 9. 
 
 FIG. 9. 
 
 The copper is contained in the glass tube ab on the 
 combustion furnace. At the ends of this tube are the 
 stop-cocks rr. V is a glass globe provided with a stop- 
 cock u. Before the experiment the air is exhausted from 
 the globe and the tube ab, and the tube then carefully 
 weighed. The tubes B and C and the apparatus A con- 
 tain substances which have the power to absorb the car- 
 bon dioxide of the air. The tube ab is now heated and 
 air admitted after passing through (7, B, and A. The 
 copper takes up the oxygen, and the nitrogen enters the 
 globe V. After the globe is full it is weighed, then ex- 
 hausted and weighed again, and the difference gives the 
 weight of the nitrogen. The tube is also exhausted and 
 weighed, and the difference between this weight and that 
 
ANALYSIS OF AIR. 255 
 
 of the exhausted tube before the experiment gives the 
 weight of the oxygen. 
 
 The most refined method for the analysis of the air is 
 the eudiometric method of Bunsen. This consists in 
 adding some pure hydrogen to a measured volume of air 
 contained in a eudiometer over mercury, and then ex- 
 ploding the mixture by means of an electric spark. If 
 the conditions are right all the oxygen present will com- 
 bine with hydrogen, and in consequence of this there 
 will be a corresponding contraction in the volume of the 
 gases. The amount of contraction will be equal to the vol- 
 ume of hydrogen and that of oxygen which have combined 
 to form water. But we know from previous experiments 
 on these two gases that they combine in the ratio of one 
 volume of oxygen to two of hydrogen. Consequently 
 the volume of oxygen which was present is equal to one 
 third of the total contraction. Of course it is necessary 
 that there should be enough hydrogen present to com- 
 bine with all the oxygen. This method is capable of 
 great exactness. The most accurate analyses made by 
 this method by Bunsen and others have shown that in 
 100 volumes of air there are 20.9 to 21 volumes of oxygen. 
 
 The estimation of the quantity of water vapor present 
 in the air is an important problem. The quantity pres- 
 ent depends upon a variety of causes, the temperature 
 and the direction of the wind being the chief ones. A 
 good chemical method for estimating the water consists 
 in drawing a known volume of air over calcium chloride 
 in a weighed tube. This substance has the power to 
 take up water, as we have repeatedly seen. If the tube 
 is weighed after a certain volume of air has been drawn 
 through it, the increase in weight will show the weight 
 of water contained in that volume of air. 
 
 The quantity of water vapor present in the air varies be- 
 tween comparatively wide limits. At any given temper- 
 ature the air cannot hold more than a certain quantity. 
 When it contains this quantity it is said to be saturated. 
 If cooled down below this temperature the vapor partly 
 condenses, and appears now as water. When a vessel 
 containing ice is placed in the air, that which immedi- 
 
256 INORGANIC CHEMISTRY. 
 
 ately surrounds the vessel is cooled down below the point 
 at which the quantity of water vapor present would satu- 
 rate the air, and water condenses on the outside of the 
 vessel. Every one has noticed that on a warm cloudy day 
 more water condenses on such a vessel than on a clear 
 cool day. The water vapor present in the air has an 
 important effect on man. The inhabitants of countries 
 with moist climates apparently have characteristics which 
 are not generally met with in those who inhabit countries 
 with dry climates. The difference in the effects of moist 
 and of dry air on an individual is well known. 
 
 When air which is charged with water vapor comes in 
 contact with cooler air, the vapor condenses and falls as 
 rain. 
 
 The method employed for the purpose of estimating 
 the quantity of carbon dioxide in the air consists in draw- 
 ing a known volume of air over something which has the 
 power to absorb the carbon dioxide, and then determin- 
 ing the increase in weight of the absorbing substance. 
 Potassium or sodium hydroxide is well adapted to this. 
 An apparatus has been constructed in which barium 
 hydroxide, Ba(OH) 3 , is used as the absorbent. When 
 carbon dioxide is passed through a solution of this sub- 
 stance insoluble barium carbonate, BaCO 3 , is thrown 
 down according to the equation 
 
 Ba(OH) 2 + C0 2 = BaC0 3 + H 2 O. 
 
 This may be filtered off and weighed, and the quantity 
 of carbon dioxide estimated from the results ; or, if a 
 known quantity of the hydroxide is taken, the quantity 
 left unacted upon after the experiment can be determined 
 by neutralizing with an acid, the neutralizing power of 
 which has previously been determined with care. The 
 quantity of carbon dioxide present in the air is relatively 
 very small, being about 3 parts in 10,000. It varies 
 slightly according to the locality and season, being 
 greater in cities and in summer than in the country and 
 in winter ; and greater in warm countries than in cold. 
 It is as essential to the life of plants as oxygen is to the 
 life of animals. 
 
AIR AND LIFE. 257 
 
 It is not an easy matter to determine the quantities of 
 the other constituents of the air, as the ammonia, organic 
 substances, etc., though there is no difficulty in deter- 
 mining that they are present in very small quantities. 
 
 The relations of the air to the most important chemical 
 changes which are taking place upon the earth form one 
 of the most interesting subjects for all men. We have 
 had a slight glimpse of the action of oxygen, and of that 
 of carbon dioxide ; both are essential to the life of plants 
 and animals. So, too, the water vapor acts chemically 
 upon plants, and probably to some extent in the respira- 
 tion of animals. As regards the nitrogen, this element is 
 frequently referred to as inert, and as serving the purpose 
 of diluting the oxygen. Inert it undoubtedly is, and 
 there is also no doubt that it dilutes the oxygen, but 
 these statements give a very inadequate conception of 
 the important part played by it in the processes of nature. 
 Nitrogen in some form of cpmbination is an essential 
 constituent of plants and animals. The animals get their 
 nitrogenous compounds from the plants, and the plants 
 get theirs partly at least from the soil. By the growth 
 of plants, therefore, nitrogenous compounds are con- 
 stantly being withdVawn from the soil. When plants 
 and animals undergo decomposition in the soil, the nitro- 
 gen contained in them is gradually converted into salts 
 of nitric acid or nitrates, and if the decomposition takes 
 place in the air the nitrogen is converted principally into 
 ammonia. Both in the form of nitrates and of ammonia 
 the nitrogen can be utilized by plants, so that if the 
 plants and animals which have received their nourish- 
 ment from a certain tract of land should be allowed to 
 decay upon this land after death, and the products thus 
 formed should be uniformly distributed in the soil, the 
 latter would not become exhausted. But the products 
 of the soil are removed, and, therefore, the nitrogen re- 
 quired for the growth of other plants is removed, and 
 the soil becomes unproductive. In order that it may be 
 rendered fertile again, the lost nitrogen must be sup- 
 plied. It appears from recent very elaborate experi- 
 ments that the plants have the power to take up from 
 
258 INORGANIC CHEMISTRY. 
 
 the air a part of the nitrogen which they need. Whether 
 they take it up directly, or it is first taken up by the 
 soil and converted into some compound which the plants 
 can utilize, are questions which apparently have not 
 been answered satisfactorily as yet. It appears possible, 
 further, that in this absorption of nitrogen from the 
 air certain minute organisms which exist in the soil play 
 a part. 
 
 Pure air may be defined as air which consists of nitro- 
 gen, oxygen, and carbon dioxide in the proportions stated 
 above, together with some water vapor, ammonia, hy- 
 drogen dioxide, and argon, and nothing of an injurious 
 nature. It is evident from what has been said that there 
 is constant danger of contamination from natural causes. 
 The most common cause of contamination is the breathing 
 of human beings in rooms which are inadequately sup- 
 plied with air. The breathing process involves the using 
 up of oxygen and the giving off of carbon dioxide and 
 small quantities of organic matter which is undergoing de- 
 composition. If the quantity of oxygen is reduced below 
 a certain limit the air becomes unfit for breathing pur- 
 poses, and evil effects follow. An ordinary inhalation 
 will not then be sufficient to supply the blood with the 
 oxygen necessary to purify it, and the system will begin 
 to suffer. Headache, drowsiness, and a general sense of 
 discomfort follow. The ill effects of breathing the air of 
 a badly ventilated room occupied by a number of human 
 beings appear to be due for the most part to the pres- 
 ence of the small quantities of decomposing organic mat- 
 ters which are given off from the lungs with the carbon 
 dioxide and other gases. These act as poisons. They 
 have been thrown off from the lungs because they are 
 unfit for use, and when they are taken back again the 
 normal processes of the body are interfered with. The 
 subject of ventilation has been so thoroughly discussed 
 of late years that great improvement has been made 
 in the arrangements for supplying pure air to dwelling 
 apartments and audience halls, but there is still room for 
 improvement. Fortunately, in most buildings there is 
 one source of supply of pure air which is independent of 
 
IMPURITIES IN AIR. 259 
 
 architects' plans. This is the diffusion of gases through 
 the porous materials of which the buildings are con- 
 structed. This diffusion was referred to under the head 
 of Hydrogen (see p. 45). There is also a good deal of 
 ventilation through the cracks and other apertures which 
 are always to be found in our buildings. 
 
 Under some conditions not thoroughly understood the 
 air becomes contaminated by the decomposition of ani- 
 mal and vegetable matter. Air thus contaminated may 
 cause specific diseases, and some of these are spoken of 
 as being caused by malaria, a word which signifies sim- 
 ply bad air. It has been shown that there are in the air 
 microscopic germs which have the power to develop in 
 the body, and then to cause the symptoms which are 
 referred to malaria. Disease germs and germs of other 
 kinds are present in the air in great variety, and they 
 play an important part in connection with the life and 
 health of mankind. 
 
 Argon. Lord Eayleigh has shown that a liter of nitro- 
 gen prepared from the air by abstracting the oxygen and 
 the small quantities of other substances known to be 
 present weighs 1.2572 grams, while a litre of nitrogen 
 made from some chemical compound, such as ammonia, 
 weighs 1.2505 grams. Chemically prepared nitrogen is 
 lighter than that obtained from the air. This observa- 
 tion led Lord Rayleigh and "W. Eamsay to a more 
 thorough chemical examination of the air, the result of 
 which was the discovery of a constituent previously un- 
 known. This is the gas argon. It can be obtained by 
 either of two methods : (1) By passing air over heated 
 copper until all the oxygen is abstracted, and then over 
 heated magnesium which unites with the nitrogen but 
 leaves the argon ; (2) By mixing air with oxygen and 
 passing electric sparks through the mixture. The oxygen 
 and nitrogen combine and the product can easily be re- 
 moved. After all the nitrogen has thus been removed 
 argon remains behind. Argon is an element with the 
 atomic weight about 40. It is present in the air to the 
 extent of about 1 per cent of the nitrogen. It cannot be 
 made to combine with any other element. 
 
CHAPTER XVI. 
 
 COMPOUNDS OF NITROGEN WITH HYDROGEN WITH 
 HYDROGEN AND OXYGEN WITH OXYGEN, ETC. 
 
 General Conditions which give Rise to the Formation of 
 the Simpler Compounds of Nitrogen. We have seen that 
 nitrogen is an inactive element, showing little tendency 
 to combine with other elements. It is nevertheless an 
 easy matter to get compounds of nitrogen with many 
 other elements, and among these compounds, some of 
 those which it forms with hydrogen and oxygen are of 
 much importance. 
 
 When a compound which contains carbon, hydrogen, and 
 nitrogen, and is not volatile, is heated in a closed vessel, 
 so that the air does not have access to it, the nitrogen 
 passes out of the compound, not as nitrogen, but partly 
 in combination with hydrogen, in the form of the com- 
 pound ammonia. Nearly all animal substances contain 
 carbon, hydrogen, oxygen, and nitrogen in many forms 
 of combination, some of which are quite complicated. 
 Many of these give off ammonia when heated. Similarly, 
 compounds containing carbon, oxygen, and hydrogen, 
 even though they are thoroughly dry, when heated give 
 off oxygen in combination with hydrogen in the form of 
 water. Both these kinds of decomposition, that which 
 gives ammonia and that which gives water, are to be 
 ascribed to the fact that the compounds of carbon which 
 are heated are unstable at higher temperatures, and 
 when they are broken down the elements contained 
 in them arrange themselves in combination in stable 
 forms, such as the comparatively simple compounds, 
 water and ammonia. An illustration of this kind of 
 action was referred to in speaking of the preparation of 
 nitrogen by heating ammonium nitrite. This compound 
 
 (260) 
 
COMPOUNDS OF NITROGEN". 261 
 
 breaks down very easily under the influence of heat, the 
 hydrogen and oxygen combining to form the stable com- 
 pound water, while the nitrogen remains uncombined. 
 Some animal substances, as, for example, urine, give off 
 ammonia when they undergo spontaneous decomposition 
 in the air. This decomposition is generally, if not always, 
 due to the action of minute organisms, the germs of 
 which are in the air, which develop when they come in 
 contact with certain substances. The coal which is used 
 for making illuminating gas contains some hydrogen and 
 nitrogen in chemical combination, and when the coal is 
 heated ammonia is given off with the other products. 
 
 When animal substances undergo decomposition in 
 the presence of basic compounds where the temperature 
 is comparatively high, the nitrogen combines with oxygen 
 and with the metal of the base. Either a salt of nitric 
 acid, HNO 3 , or of nitrous acid, HNO 2 , is formed. In 
 some countries where the conditions are favorable to the 
 process, immense quantities of nitrates are found, chiefly 
 potassium nitrate, KNO 3 , and sodium nitrate, NaNO 3 . 
 Nitrates are, however, found everywhere in the soil. The 
 change of animal and vegetable nitrogenous substances 
 to the form of nitrates is probably caused by the action 
 of minute living organisms, which are found everywhere, 
 and serve an important purpose in converting the waste 
 animal and vegetable matter into simple compounds 
 which can be utilized by plants. How they effect the 
 change is not known. From the salts of nitric acid 
 which are found in nature, nitric acid itself can easily be 
 made. 
 
 Nearly all the compounds of nitrogen with which we 
 have to deal are made either from ammonia or from 
 nitric acid. 
 
 Relations between the Principal Compounds of Ni- 
 trogen. In studying the compounds of sulphur, we saw 
 that whenever a compound of sulphur is oxidized with a 
 strong oxidizing agent the final product of the action is 
 sulphuric acid ; and that, on the other hand, under the 
 influence of strong reducing agents these compounds 
 yield hydrogen sulphide as the final product. So also 
 
262 INORGANIC CHEMISTRY. 
 
 the limit of oxidizing action in the case of chlorine is 
 perchloric acid, and of reducing action, hydrochloric 
 acid. The other compounds of sulphur with hydrogen 
 and oxygen are products intermediate between the two 
 limits, sulphuric acid and hydrogen sulphide, as the 
 other compounds of chlorine with hydrogen and oxygen 
 are intermediate between hydrochloric acid and per- 
 chloric acid. The limit of reduction of nitrogen com- 
 pounds is ammonia, NH 3 , and of oxidation, nitric acid, 
 HNO 3 . As has been noticed, the valence of nitrogen 
 towards oxygen is greater than it is towards hydrogen, 
 but the difference is not as marked as in the case of the 
 members of the chlorine group, and in that of the sulphur 
 group. Its hydrogen valence is 3, its maximum oxygen 
 valence is 5. For reasons similar to those which have 
 already been discussed under the head of Hydroxides, 
 the oxidation products of ammonia are believed to have 
 their oxygen in combination with hydrogen in the form 
 of hydroxyl, so far as these two elements are present 
 in the right proportions to form this group. If this 
 is true the first product of oxidation of ammonia 
 would have the structure represented by the formula 
 
 (OH 
 NK H . - A compound of the composition represented by 
 
 ( H 
 
 NH 3 O is known, and it is believed that it has the above 
 structure. The next product of oxidation formed by the 
 
 (OH 
 same process would have the formula NK OH. No such 
 
 (H 
 
 compound is known, however. If a compound of this 
 structure should lose water the product would have the 
 composition NOH, and probably the structure O=N-H. 
 A compound of the composition represented by this 
 formula is known, but it seems probable from its con- 
 duct that the formula should be doubled, N 2 O 2 H 2 , and 
 that in the compound there are two hydroxyl groups, as 
 
 N-OH 
 represented in the formula II . Continuing the pro- 
 
 N-OH 
 
 cess of oxidation of ammonia, the next product which 
 
RELATIONS BETWEEN COMPOUNDS OF NITROGEN. 263 
 
 we should expect would be the trihydroxyl derivative 
 
 (OH 
 NS OH. Salts derived from an acid of this composition, 
 
 (OH 
 
 which should be called normal nitrous acid, are known. 
 But this normal acid breaks down readily by loss of water 
 
 into an acid of the formula N \ QTT or O=N-O-H, 
 
 which is the acid from which nearly all the nitrites are 
 derived. By continued oxidation the quinquivalence of 
 the nitrogen makes it possible to add another oxygen 
 atom to the molecule of nitrous acid, and thus to form 
 the final product of oxidation, nitric acid of the structure 
 O 
 
 N-O-H ; or, if the oxidation takes place in solution, it 
 I! 
 
 O 
 is probable that the final product is a hydroxyl deriva- 
 
 roH 
 
 I OH 
 tive of the structure N-{ OH, which should be called 
 
 I OH 
 
 LOH 
 
 normal nitric acid. There are some salts known which 
 are derived from an acid of this composition, but most 
 of the nitrates are derived from an acid which is formed 
 from this by loss of two molecules of water : 
 
 N(OH) 5 = N0 2 (OH) + 2H 2 O. 
 
 It is impossible at present to furnish satisfactory evi- 
 dence in favor of the views just pointed out. All that 
 can be said is that the views are in accordance with a 
 number of facts, and that they simplify the study of the 
 relations of the compounds. 
 
 When nitric acid is reduced, of course, the above pro- 
 cesses must be regarded as reversed. Thus, the first 
 result of the action of nascent hydrogen upon nitric acid, 
 for example, would be the elimination of one atom of 
 oxygen and the formation of nitrous acid : 
 
 NO 2 (OH) + 2H = NO(OH) + H a O. 
 
264 INORGANIC CHEMISTRY. 
 
 The next result would be the formation of the compound 
 N(OH) or N 2 O 2 H 2 ; the next the formation of the com- 
 pound NH 2 (OH) ; and finally this would be converted 
 into ammonia. But the matter is complicated by the 
 fact that some of the compounds referred to above break 
 down into water and oxides of nitrogen. Thus, nitrous 
 acid, NO(OH), breaks down into water and nitrogen tri- 
 oxide or nitrous anhydride, N 2 O 3 ; and the compound 
 N 2 O 2 H 2 , or hyponitrous acid, breaks down into water and 
 nitrous oxide, N 2 O : 
 
 N(OH) 
 
 II = N 2 O + H 2 O. 
 
 N(OH) 
 
 Further, there are some compounds of nitrogen and 
 oxygen for which there are no analogous hydroxyl com- 
 pounds known, as, for example, NO and NO 2 or N 2 O 4 . In 
 these compounds nitrogen appears to be bivalent and 
 quadrivalent, while in nitrous oxide, N 2 O, it appears to 
 be univalent. Only those oxygen compounds of nitro- 
 gen in which the nitrogen is univalent, trivalent, or 
 quinquivalent appear to form corresponding hydroxyl 
 compounds. Taking the simplest view of the matter, 
 nitrogen appears to be univalent, bivalent, trivalent, 
 quadrivalent, and quinquivalent towards oxygen, as 
 shown in the compounds 
 
 N 2 O, NO, N 2 O 3 , NO 2 , and N 2 5 . 
 
 Corresponding to nitrous oxide, N 2 O, nitrogen trioxide, 
 N 2 O 3 , and nitrogen pentoxide, N 2 O 5 , there are hydroxyl 
 derivatives which have acid properties : 
 
 N 2 + H 2 = N 2 (OH) 2 ; 
 N 2 O 3 + H 2 = 2NO(OH) ; 
 N 2 6 + H 2 = 2N0 2 (OH). 
 
 But there are no hydroxyl derivatives corresponding 
 to those oxides in which the valence appears to be 2 and 
 4. Considering the ease with which those hydrox- 
 
RELATIONS BETWEEN COMPO UNDS OF NITROGEN. 265 
 
 ides of nitrogen break down which contain two hydroxyl 
 groups in the molecule, the fact that hydroxides of the 
 formulas N(OH) 2 and N(OH) 4 do not exist is not surpris- 
 ing. We should expect them to break down spontane- 
 ously into NO and NO a . 
 
 The methods by which it is possible to pass from one 
 of the oxides or hydroxides of nitrogen to the other will 
 be pointed out below. A tabular list of the compounds 
 is first given : 
 
 Oxides. Acids. 
 
 Nitrous oxide, N 3 O Hyponitrous acid, N 2 (OH), 
 
 Nitric oxide, NO(N 2 O 2 ) Nitrous acid, NO(OH) 
 
 Nitrogen trioxide, ) XT r> Nitric acid, NO 2 (OH) 
 (Nitrous anhydride), f 1 ^ 3 
 Nitrogen peroxide, NO 2 (N 2 O 4 ) 
 
 Nitrogen pentoxide, ) v r| 
 (Nitric anhydride), j 2 * 
 
 Besides the above there are the basic compounds, 
 hydroxylamine, NH 2 (OH), ammonia, NH 3 , and hydrazine, 
 N 2 H 4 ; and with water ammonia probably forms the 
 hydroxide NH 4 (OH), which is a base. In addition to 
 these there is a compound of nitrogen with hydrogen of 
 the formula N S H. This is a strong acid. (See Triazoic 
 Acid.) And of the same composition as hyponitrous 
 acid but differing from it in properties is nitramide, 
 N 2 H 2 O 2 , the relation of which to hyponitrous acid is not 
 yet quite clear. 
 
 With the members of the chlorine group nitrogen 
 forms a few compounds which are characterized by 
 marked instability. So unstable are they that they ex- 
 plode violently when simply touched. The chloride of 
 nitrogen explodes with terrific violence when the direct 
 rays of the sun are allowed to shine upon it. With sul- 
 phur, nitrogen forms two compounds. With sulphur, 
 hydrogen, and oxygen, however, nitrogen forms a number 
 of compounds, one of which has already been referred to 
 in connection with the manufacture of sulphuric acid. 
 This is the so-called nitrosyl-sulphuric acid which is 
 formed by the action of a mixture of the peroxide, NO,, 
 
266 INORGANIC CHEMISTRY. 
 
 and the monoxide, NO, on sulphurous acid, oxygen, and 
 water. 
 
 In studying the compounds of nitrogen, it will be best 
 to begin with the end products, ammonia and nitric acid. 
 
 Ammonia, NH 3 . The conditions under which ammonia 
 is formed have been mentioned. The chief source at 
 present is the " ammoniacal liquor" of the gas-works. 
 This is the water through which the gas has been passed 
 for the purpose of removing the ammonia, and it contains 
 ammonia, or ammonium hydroxide, NH 4 (OH), in solu- 
 tion. By adding hydrochloric acid to this solution the 
 salt ammonium cJdoride, NH 4 C1, is formed. This is the 
 well-known substance sal ammoniac. It appears that this 
 name has its origin in the fact that common salt or 
 sodium chloride, NaCl, was formerly called sal armenia- 
 cum, and that afterward, through a misunderstanding, 
 ammonium chloride came to be known by the same name 
 which underwent change to the form sal ammoniacum, or 
 sal ammoniac. When the ammoniacal liquor is treated 
 with sulphuric acid, ammonium sulphate is formed. 
 From one or the other of these salts it is a simple matter 
 to obtain ammonia. For this purpose it is only necessary 
 to treat the salt with some strongly basic compound, as, 
 for example, potassium or sodium hydroxide, or calcium 
 hydroxide. Thus, when a solution of potassium hydrox- 
 ide is poured on ammonium chloride or sulphate the 
 strong penetrating odor of ammonia is at once noticed. 
 The first reaction probably results in the formation of 
 ammonium hydroxide, thus : 
 
 NH 4 C1 +KOH =NH 4 OH + KC1 ; 
 (NH 4 ) 2 SO 4 + 2KOH = 2NH 4 OH + K 2 SO 4 ; 
 2NH 4 C1 + Ca(OH) 2 = 2NH 4 OH + CaCl 2 ; 
 (NH 4 ) 2 SO 4 + Ca(OH) 2 = 2NH 4 OH.+ CaSO 4 . 
 
 But the ammonium hydroxide breaks down very readily 
 into water and ammonia, which escapes as a gas : 
 
 Further, ammonia can also be made by bringing 
 
AMMONIA. 26? 
 
 nascent nitrogen and nascent hydrogen together, as, for 
 example, by heating a mixture of iron filings, potassium 
 nitrate, and potassium hydroxide. Under these circum- 
 stances the iron sets hydrogen free from the hydroxide, 
 and nitrogen from the nitrate, and they unite to 
 form ammonia. The formation of ammonia by reduction 
 of nitric acid can be shown by treating some granulated 
 zinc with dilute sulphuric acid, and while the action is 
 in progress adding nitric acid drop by drop. The nitric 
 acid is thus reduced, and the ammonia which is formed 
 remains in combination with the sulphuric acid as am- 
 monium sulphate. Other interesting modes of formation 
 of ammonia are by the action of electric sparks on nitro- 
 gen in the presence of water, and, in general, by the 
 evaporation of water : 
 
 N 2 + 2H 2 O = NH 4 N0 2 . 
 
 The product ammonium nitrite is always found in the 
 air in small quantities. 
 
 In the laboratory ammonia is prepared by treating 
 ammonium chloride with slaked lime or calcium hydrox- 
 ide. The two are mixed in the proportion of two parts 
 of slaked lime to one of ammonium chloride, placed in a 
 flask and gently heated, when the ammonia is given off 
 at once. 
 
 It is frequently more convenient to heat a strong 
 aqueous solution of ammonia, such as is found in every 
 chemical laboratory. Such a solution when gently 
 heated readily gives off ammonia. 
 
 Ammonia is a colorless, transparent gas with a very 
 penetrating, characteristic odor. In concentrated form 
 it causes suffocation. Its specific gravity is 0.589; that 
 is to say, it is but little more than half as heavy as air. . 
 A liter of the gas under standard conditions weighs 
 0.7635 gram. It can easily be compressed to the liquid 
 form by pressure and cold. When the pressure is re- 
 moved from the liquefied ammonia it passes back to the 
 form of gas, and in so doing it absorbs a great deal of 
 heat. These facts are taken advantage of for the arti- 
 
268 INORGANIC CHEMISTRY. 
 
 ficial preparation of ice. This application will be clear 
 from the following explanation and Fig. 10. 
 
 An aqueous solution of ammonia saturated at is 
 brought into the strong iron cylinder A, and then gently 
 warmed, while the vessel B is cooled by cold water. 
 The gas given off from A passes though the bent tubes 
 into B, where it is condensed 
 to a liquid. The cylinder 
 A is now placed in a ves- 
 sel of cold water, and the 
 water which is to be frozen is 
 placed in a cylinder D, and 
 this into the hollow space E 
 in the vessel B. The liquid 
 ammonia passes rapidly into 
 the form of gas which is ab- 
 sorbed in the water in A, 
 while at the same time so FIG. 10. 
 
 much heat is absorbed that the water in D is frozen. 
 
 Ammonia does not burn in the air, but does burn in 
 oxygen with a pale yellowish flame. It is absorbed by 
 water in very large quantity. One volume of water at 
 the ordinary temperature dissolves about 600 volumes of 
 ammonia gas, and at about 1000 volumes. The sub- 
 stance with which we commonly have to deal under the 
 name of ammonia is a solution of ammonia in water. It 
 is called " spirits of hartshorn" in common language. The 
 solution has the odor of the gas. It loses all its gas when 
 heated to the boiling temperature. The solution shows 
 a strong alkaline reaction, and has the power to neutral- 
 ize acids and form salts. The conduct of the solution is, 
 in fact, strikingly like that of sodium and potassium 
 hydroxides, and it is believed that in the solution there 
 is contained a compound of the formula NH 4 (OH), known 
 as ammonium hydroxide, and formed by the direct action 
 of ammonia upon water. If this is true, then the action 
 of ammonia upon acids is to be explained as follows : 
 Ammonium hydroxide is analogous to potassium hydrox- 
 ide, but differs from it in that it contains the group of 
 atoms NH 4 in place of the atom K. In some way this 
 
AMMONIUM SALTS. 269 
 
 group plays in the salts formed by ammonia the same 
 part that the elementary atom potassium plays in the 
 salts of potassium, and just as the latter are called potas- 
 sium salts, so the former are called ammonium salts. 
 According to this the ammonium salts are salts which 
 contain the group NH 4 , known as the ammonium group, 
 in place of the hydrogen of the acids. They are formed 
 by direct combination of ammonia with the acids, or by 
 the action of ammonium hydroxide upon acids. The 
 analogy between the action of ammonium hydroxide and 
 that of potassium hydroxide upon acids is clearly shown 
 by the aid of the following equations : 
 
 K(OH) +HC1 =KC1 +H 2 O; 
 
 NH 4 (OH) +HC1 = NH 4 C1 +H 2 O; 
 K(OH) + HN0 3 =KN0 3 +H 2 O; 
 NH 4 (OH) + HN0 3 = NH 4 N0 3 + H 2 O; 
 2K(OH) + H 2 S0 4 = K 2 S0 4 + 2H 2 O ; 
 2NH 4 (OH) + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2H 2 O. 
 
 The formation of the ammonium salts by direct action 
 of ammonia, NH 3 , upon acids is represented in the fol- 
 lowing equations : 
 
 NH 3 +HC1 = 
 
 + HNO 3 = NH 4 NO 3 ; 
 
 The strong tendency of ammonia to combine directly 
 with acids is shown by bringing two uncovered vessels, 
 one containing a solution of ammonia, and the other a 
 solution of hydrochloric acid, near each other. A dense 
 cloud will at once be noticed, if the solutions are concen- 
 trated. This is due to the direct combination of the gases 
 which escape from the solutions. 
 
 While the assumption of the existence of the group 
 ammonium, NH 4 , in the ammonium salts is of great ser- 
 vice in dealing with these salts, and while this assump- 
 tion appears to be entirely justified by the facts, no 
 compound of this composition has as yet been isolated. 
 
270 INORGANIG CHEMISTRY. 
 
 The name ammonium is given to the hypothetical com- 
 pound on account of the fact that it evidently plays the 
 part of a metallic element, and it is customary to give 
 such elements names ending in ium. While, further, it 
 is generally believed that the compound ammonium 
 hydroxide, NH 4 (OH), is formed when ammonia dissolves 
 in water, the compound itself has not been isolated, 
 owing to its instability and tendency to break down into 
 ammonia and water. On the other hand, some very in- 
 teresting derivatives of this hydroxide have been isolated. 
 There is one of these which is derived from the hydrox- 
 ide by the replacement of the four hydrogen atoms of 
 the ammonium by groups of carbon and hydrogen 
 atoms. This compound is stable and can be isolated as 
 a hydroxide of the general formula NE 4 (OH), in which 
 E 4 represents the groups of carbon and hydrogen atoms. 
 It acts almost exactly like potassium hydroxide. 
 
 Composition of Ammonia. By oxidation under the 
 proper conditions it is possible to convert the hydrogen 
 of ammonia into water and leave the nitrogen in the free 
 state. As water and nitrogen are the only products 
 formed, and the quantity of oxygen used up in the oxida- 
 tion is equal to the quantity of oxygen found in the- 
 water formed, it follows that nitrogen and hydrogen are 
 the only elements contained in ammonia. 
 
 When electric sparks are passed for some time through 
 a mixture of nitrogen and hydrogen, some ammonia is 
 formed. Conversely, when electric sparks are passed for 
 a time through ammonia, nitrogen and hydrogen are 
 obtained. 
 
 If, in the oxidation of a known weight of ammonia, the 
 water formed and the nitrogen left uncombined are accu- 
 rately determined, it will be found that in ammonia the 
 elements are combined very nearly in the proportion of 
 fourteen parts by weight of nitrogen to three parts by weight 
 of hydrogen. Further, the molecular weight determined 
 by the method of Avogadro is approximately 17. There- 
 fore, the molecular formula of ammonia is NH 3 , the 
 atomic weight of nitrogen being 14. 
 
 The proportion by volume in which the two elements 
 
COMPOSITION OF AMMONIA. 271 
 
 combine can be determined by the following method : 
 A glass tube closed at one end and provided with a glass 
 stop- cock at the other is filled with pure chlorine gas. 
 By means of a small funnel attached to the open end a 
 little of a strong aqueous solution of ammonia is slowly 
 introduced into the tube. Reaction takes place at once 
 between the chlorine and the ammonia, according to the 
 equation 
 
 and the hydrochloric acid unites with ammonia to form 
 ammonium chloride : 
 
 The entire change is therefore represented by the equa- 
 tion 
 
 4NH, + 301 = N + 3NH 4 C1. 
 
 Hydrogen and chlorine combine in equal volumes, as we 
 have already learned. Now, if we start with a measured 
 volume of chlorine, and add ammonia to it until it is all 
 used up, we know that the volume of hydrogen which 
 has been extracted from ammonia is equal to the volume 
 of chlorine with which we started. If we measure the 
 volume of nitrogen left over, we know the volume of the 
 nitrogen which was in combination with a volume of 
 hydrogen equal to that of the chlorine originally taken. 
 This experiment has been tried repeatedly, and it has 
 been found that the ratio of the volume of nitrogen to 
 that of the hydrogen with which it was combined is as 1 
 to 3. The tube being full of chlorine at the beginning 
 of the experiment, it will be found to be one-third full 
 of nitrogen at the end. Therefore, in ammonia 1 volume 
 of nitrogen is combined ivith 3 volumes of hydrogen. 
 
 The experiment just referred to will perhaps be better 
 
272 
 
 INORGANIC CHEMISTRY. 
 
 understood by the aid of the accompanying diagram. 
 The chlorine in the tube may be represented as made 
 up of three equal parts or volumes. Each volume of 
 chlorine combines with an equal volume of hydrogen, 
 leaving the nitrogen uncombined. The volume of nitro- 
 gen left is only one third of that of the chlorine, or for 
 three volumes of chlorine there is one volume of mitrogen : 
 
 combine 
 with 
 
 leaving 
 
 Therefore in ammonia the gases nitrogen and hydrogen 
 are combined in the proportion of 1 volume of nitrogen 
 to 3 volumes of hydrogen : 
 
 Volume relations in ammonia. 
 
 Another question in regard to the volume relations 
 remains to be answered, and that is : When nitrogen and 
 hydrogen unite in the proportions above stated, how 
 many volumes of ammonia gas do the four volumes of 
 
AMMONIUM AMALGAM. 273 
 
 the constituents form ? It is not possible to determine 
 this by direct combination of the two gases, but am- 
 monia can be decomposed into its constituents by con- 
 tinued passage of electric sparks through it. When this 
 is done it is found that after the decomposition the 
 gases occupy twice the volume which was occupied by 
 the ammonia. It appears, therefore, that when hydrogen 
 and nitrogen combine to form ammonia the volume is 
 reduced to one half, or, what is the same thing, when 
 three volumes of hydrogen combine with one volume of 
 nitrogen the four volumes form two volumes of ammonia 
 gas. 
 
 The above facts have already been commented upon 
 in speaking of the combination of gases in general ; and 
 it has been shown that chlorine, oxygen, and nitrogen 
 combine with hydrogen in entirely different ways : 
 
 (1) 1 volume of chlorine combines with 1 volume of 
 hydrogen to form 2 volumes of hydrochloric acid gas. 
 
 (2) 1 volume of oxygen combines with 2 volumes of 
 hydrogen to form 2 volumes of water vapor. 
 
 (3) 1 volume of nitrogen combines with 3 volumes of 
 hydrogen to form 2 volumes of ammonia gas. 
 
 What the cause of these differences is we do not know. 
 In some way these facts are directly connected with the 
 law of Avogadro that equal volumes of all gases contain 
 the same number of molecules, and with the power of 
 the atoms of chlorine, oxygen, and nitrogen to combine 
 with one, two, and three atoms of hydrogen respectively. 
 When one atom of chlorine unites with one atom of 
 hydrogen the result is a molecule. So also when one 
 atom of oxygen unites with two atoms of hydrogen, and 
 when one atom of nitrogen unites with three atoms of 
 hydrogen, the result in each case is a molecule, and, ac- 
 cording to the law of Avogadro, a gaseous molecule 
 whether it consists of one atom or a hundred atoms oc- 
 cupies the same space. 
 
 Ammonium Amalgam. A very curious substance 
 which appears to consist of mercury and ammonium 
 is formed when a solution of ammonium chloride 
 is treated with a compound of sodium and mercury 
 
274 INORGANIC CHEMISTRY. 
 
 known as sodium amalgam. The action is thought to 
 take place thus : 
 
 2NH 4 C1 + Na 2 Hg = 2NaCl + (NH 4 ) 2 Hg. 
 
 The product, ammonium amalgam, is unstable, break- 
 ing down very soon into ammonia, hydrogen, and mer- 
 cury. It will be referred to again and somewhat more 
 fully under the head of Mercury. 
 
 Metallic Derivatives of Ammonium Compounds and of 
 Ammonia. Ammonia acts upon a number of metallic 
 salts, forming with them complex derivatives which in 
 some cases appear to be ammonia, and in others am- 
 monium, in which one or more hydrogen atoms are re- 
 placed by metal. Thus mercuric chloride, HgCl 2 , acts 
 upon ammonia, forming a compound of the formula 
 HgClNH 2 . This appears to be ammonium chloride in 
 which two hydrogen atoms are replaced by a mercury 
 atom, or as mercuric chloride in which one chlorine 
 atom is replaced by the group NH 2 , known as the amide 
 group. According to the latter view the structure of 
 
 Cl 
 this compound is represented by the formula Hg<^-rr . 
 
 Similarly, copper chloride, CuCl 2 , forms a compound the 
 composition of which is represented by the formula 
 CuCl 2 .2NH 3 . It seems probable that this is ammonium 
 chloride in which two hydrogen atoms are replaced by 
 
 an atom of copper, C u <>jT 3 ni- There are many such 
 
 compounds, particularly of the metals mercury, copper, 
 cobalt, and platinum. The structural formulas above 
 given are to be regarded merely as suggestions. Experi- 
 mental evidence in favor of them is lacking. 
 
 Structure of Ammonium Compounds. In ammonia ni- 
 trogen is unquestionably trivalent. But what takes place 
 when ammonia acts upon water, upon hydrochloric acid, 
 and upon acids in general? What structure is to be 
 ascribed to the ammonium compounds? The view 
 generally held is that in the ammonium compounds ni- 
 trogen is quinquivalent. According to this, ammonia is 
 an unsaturated compound, and when brought in contact 
 
STRUCTURE OF AMMONIUM COMPOUNDS. 
 
 275 
 
 with water, hydrochloric acid, etc., it saturates itself. 
 The conception is represented thus : 
 
 H 
 
 = N^ 
 
 HN0 3 = 
 
 H 
 H 
 H ; 
 H 
 
 01 
 
 H 
 H 
 
 H ; 
 H 
 OH 
 
 H 
 H 
 H 
 H 
 
 [NO 3 
 
 In all the resulting compounds the nitrogen is believed 
 to be quinquivalent. It must be confessed that, while 
 this is a convenient hypothesis, further evidence for or 
 against it is desirable. One objection which may be 
 raised to it is this. It is assumed that when the stable 
 compound hydrochloric acid acts upon ammonia the 
 hydrogen separates from the chlorine and both combine 
 with nitrogen ; but for nitrogen alone chlorine has very 
 little attraction. This objection may not be a real one. 
 It may be shown that, while chlorine has for nitrogen 
 alone very little attraction, it has great attraction for 
 nitrogen which is in combination with hydrogen. 
 
 Hydrazine, N 2 H 4 . A compound closely related to am- 
 monia and ammonium has recently been prepared by a 
 complicated method which cannot be explained here. 
 This is known as hydrazine. Its composition and mo- 
 lecular weight are represented by the formula N 2 H 4 . A 
 large number of derivatives of hydrazine are known, and 
 have been studied exhaustively. From what has been 
 learned in regard to them it appears that the constitution 
 
 of hydrazine is represented by the formula J,** Accord- 
 
276 INOEGANIC CHEMISTRY. 
 
 ing to this it bears to ammonia a relation similar to that 
 which hydrogen peroxide is believed to bear to water : 
 
 OH 3 NH 3 
 
 OH NH 2 
 
 Hydrazine is a liquid that boils at 113.5 in a current 
 of hydrogen. It acts upon acids much as ammonia does, 
 forming the hydrazine salts. 
 
 Hydroxylamine, NH 2 (OH). This compound is pre- 
 pared by reducing nitric acid : 
 
 NO 2 (OH) + 6H = NH 2 (OH) + 2H 2 O. 
 
 It can also be prepared by reduction of nitric oxide : 
 
 2NO + 6H = 2NH 2 (OH). 
 
 Hydroxylamine is a solid consisting of leaflets or hard 
 needles. It melts at 33.05, and boils at 58 under a 
 pressure of 22 mm. When the water solution is evap- 
 orated, both ammonia and hydroxylamine pass over with 
 the water. Its salts are easily obtained by treating the 
 solution with acids : 
 
 NH 2 (OH) + HC1 = NH 3 (OH)C1 ; 
 
 NH 2 (OH) + HNO 3 = NH.(QH)NQ,, 
 
 From the method of formation and the composition it 
 appears that these salts are ammonium salts in which 
 one hydrogen of the ammonium is replaced by hydroxyl. 
 They should therefore be called Jiydroxyl-ammonium salts. 
 One of the most characteristic properties of hydroxyl- 
 amine is the ease with which it breaks down into ammonia, 
 nitrogen, and water : 
 
 3NH 2 OH = N 3 + NH 3 + 3H 2 O. 
 
 If brought in contact with compounds capable of reduc- 
 tion, it reduces them, the nitrogen in these cases gener- 
 ally combining with oxygen to form nitrous oxide, N 2 O. 
 The reduction of cupric oxide, CuO, takes place accord- 
 ing to the equation 
 
 2NH 2 (OH) + 4CuO = N 2 O + 2Cu 2 + 3H 2 0. 
 
NITRIC ACID. 277 
 
 By nascent hydrogen hydroxylamine is reduced to 
 ammonia. 
 
 Triazoic Acid, N 3 H. This compound, which is also 
 called hydrazoic and hydronitric acid, can be made by a 
 number of reactions involving the use of complex organic 
 substances. A simpler method consists in passing am- 
 monia gas over heated metallic sodium, and then passing 
 nitrous oxide, N a O, over the resulting product, which is 
 sodium amide. The following reaction takes place: 
 
 NH 2 Na + N 2 = N 3 Na + H 3 O. 
 
 By dissolving in water the sodium salt thus formed, 
 treating with dilute sulphuric acid, and distilling, a solu- 
 tion of the acid is obtained. The compound is a color- 
 less liquid that boils at 37. It has a fearfully penetrat- 
 ing odor, and produces bad effects upon one who inhales 
 it. It is a strong acid resembling hydrochloric acid. It, 
 as weljl as some of its salts, is extremely explosive. Its 
 ammonium salt formed by direct union with ammonia is 
 interesting as it has the composition N,NH 4 or N 4 H 4 . 
 
 Nitric Acid, HNO 3 . This important chemical compound 
 was first made, though not in pure condition, about the 
 ninth century by distilling saltpeter, copper sulphate, 
 and alum. The name nitric acid has its origin in the 
 fact that the compound is formed from niter. In German 
 it is called Salpetersdure, which literally translated is 
 saltpeter acid. It has already been stated that the salts 
 of this acid, particularly the potassium and sodium salts, 
 occur very widely distributed in the earth, and that there 
 is a great accumulation of the sodium salt in South 
 America, whence the name Chili saltpeter. Wherever 
 organic matter, particularly that of animal origin, under- 
 goes spontaneous decomposition in the presence of basic 
 substances, nitrates are formed, probably in consequence 
 of the action of an organism known as the nitrifying fer- 
 ment. This process of nitrification has already been re- 
 ferred to in a general way. It is one of great importance 
 for the welfare of the human race, and indeed of most 
 
278 INORGANIC CHEMISTRY. 
 
 living beings, as by its aid the useless nitrogenous sub- 
 stances of dead plants and animals are converted back 
 into the useful nitrates which in the soil aid the pro- 
 cesses of plant growth. 
 
 Nitric acid can be formed by the action of electric 
 sparks on nitrogen and oxygen in the presence of water. 
 It is also easily formed by the action of oxidizing agents 
 on ammonium compounds. 
 
 Nitric acid is always prepared by treating potassium 
 or sodium nitrate with concentrated sulphuric acid. When 
 sodium nitrate is treated with sulphuric acid action takes 
 place thus : 
 
 NaNO 3 + H a SO 4 = NaHSO 4 + HNO 3 . 
 
 The salt formed in this way is primary sodium sulphate, or 
 acid sodium sulphate. If sufficient of the saltpeter is 
 present and the temperature is raised, a second reaction 
 takes place, resulting in the formation of normal sodium 
 sulphate : 
 
 NaN0 3 + NaHSO 4 = Na 2 SO 4 + HNO 3 . 
 
 But the temperature required for this reaction is so high 
 that a considerable part of the nitric acid is decomposed. 
 In the preparation of nitric acid, therefore, the first re- 
 action is the one used, and for this purpose the sub- 
 stances are brought together in a retort in the proportion 
 of their molecular weights (about equal weights), and 
 the retort gently heated. The nitric acid distils over 
 slowly, and is condensed by cooling the receiver. 
 
 On the large scale the acid is made by bringing Chili 
 saltpeter and concentrated sulphuric acid together in cast- 
 iron cylinders or retorts. 
 
 Nitric acid is a colorless volatile liquid. It begins to 
 boil at 86, but at this temperature it undergoes partial 
 decomposition into nitrogen peroxide, water, and oxygen : 
 
 2HNO 3 = 2N0 2 + H 2 + O. 
 
 It undergoes the same change slowly when exposed to 
 the direct rays of the sun. In consequence of this de- 
 
NITRIC ACID. 279 
 
 composition the distillate collected in the manufacture of 
 nitric acid, and, in general, whenever the acid is distilled, 
 always contains a considerable percentage of water, and 
 is colored more or less yellow by the nitrogen peroxide 
 present. In order to abstract the water from ordinary 
 nitric acid it is mixed with concentrated sulphuric acid and 
 slowly distilled ; but even under these circumstances the 
 product is colored in consequence of some decomposition, 
 and it also contains some water. By conducting carbon 
 dioxide gas through the gently warmed acid the nitrogen 
 peroxide can be removed, and in this way an acid con- 
 taining about 99.5 per cent of the compound HNO 3 has 
 been obtained. 
 
 Pure nitric acid is a very active substance chemically. 
 It gives up its oxygen readily and is itself thus reduced 
 to other compounds of nitrogen and oxygen, or of nitro- 
 gen, oxygen, and hydrogen, as has already been pointed 
 out. When it acts upon the metals it forms nitrates, 
 metal atoms being substituted for the hydrogen. Accord- 
 ing to the conditions, nitrogen peroxide, NO 2 , nitrous 
 acid, HNO 2 , nitric oxide, NO, nitrous oxide, N 2 O, nitro- 
 gen, hydroxylamine, NH 2 (OH), and, finally, ammonia are 
 formed by reduction of the acid. The formation of am- 
 monia and of hydroxylamine by reduction of nitric acid 
 has already been specially referred to. Of these reac- 
 tions, that which gives nitric oxide, NO, is the one which 
 commonly takes place on treating metals with nitric 
 acid. The oxides NO 2 and N 2 O 3 are themselves readily 
 reduced to nitric oxide. 
 
 If the element upon which the acid acts has not the 
 power to replace the hydrogen, the action consists in 
 oxidation. This is shown in the action of strong nitric 
 acid upon tin, phosphorus, carbon, sulphur, etc. In each 
 case the highest oxidation-product is formed. Tin is 
 converted into normal stannic acid, Sn(OH) 4 ; phosphorus 
 into phosphoric acid, PO(OH) 3 ; carbon into carbon di- 
 oxide, CO a ; and sulphur into sulphuric acid, SO 2 (OH) 2 . 
 It disintegrates carbon compounds very readily, convert- 
 ing them into their final products of oxidation. In con- 
 tact with the skin it causes bad and dangerous wounds. 
 
280 INORGANIC CHEMISTRY. 
 
 Upon some stable compounds of carbon it acts forming 
 so-called nitro-compounds, very much as chlorine acts 
 upon them forming chlorine substitution-products, as 
 was explained under Chlorine. The formation of a nitro- 
 compound takes place as represented in the equation 
 
 C.H. + HO.NO, = C 9 H S .NO, + H.O. 
 
 The compound C 6 H 6 is benzene, and the compound 
 C 6 H B .NO 2 is nitro-benzene. 
 
 The acid mostly used in the laboratory has the specific 
 gravity 1.2 and contains 32 per cent nitric acid, HNO 3 . 
 The commercial acid contains about 68 per cent of the 
 acid. 
 
 When a mixture of nitric acid and water is boiled 
 under the ordinary atmospheric pressure it loses either 
 water or nitric acid until it contains 68 per cent of the 
 acid which passes over. This does not correspond to 
 any definite hydrate of nitric acid, though it approxi- 
 mates the composition required by normal nitric acid, 
 N(OH) B , or HNO 3 + 2H 2 O, and it is probable that this 
 hydrate is the chief constituent of the mixture. 
 
 Nitric acid is a strong monobasic acid, forming salts 
 of the general formula MNO 3 , all of which are soluble 
 in water. Because nitric acid is a strong acid, and all 
 normal nitrates are soluble in water, nitric acid is one of 
 the best solvents. On the other hand, the acid forms 
 basic salts, some of which are difficultly soluble or in- 
 soluble in water. An example of such insoluble basic 
 
 (N0 3 
 U 
 
 nitrates is the nitrate of bismuth of the formula Bi < OH , 
 
 (OH 
 
 which is to be regarded as bismuth hydroxide one-third 
 neutralized by nitric acid. There are some apparently 
 complex salts of nitric acid which are derived from the 
 normal acid, N(OH) 5 , as, for example, the salt HPb 3 O 6 N, 
 
 which should probably be expressed thus, N -I O -p-, , 
 
 
 
 [O.Pb.OH 
 
 being a basic lead salt of the normal acid. There are, 
 
-RED FUMING NITRIC ACID NITROUS ACID. 281 
 
 further, some salts which are derived from the acid 
 NO(OH) 3 , which is formed by abstracting one molecule of 
 water from normal nitric acid : 
 
 N(OH) 6 = NO(OH) 8 + H,0. 
 
 By far the largest number of nitrates, however, are ' de- 
 rived from the acid of the formula HNO 3 . 
 
 Ked Fuming Nitric Acid is formed in the manufacture 
 of nitric acid from saltpeter and sulphuric acid if the 
 temperature is raised to a sufficient extent to cause the 
 acid sulphate to act upon the nitrate : 
 
 NaN0 3 + HNaSO 4 = Na 2 SO 4 + HNO 3 . 
 
 At this temperature the nitric acid undergoes consider- 
 able decomposition. The nitrogen peroxide formed is 
 absorbed by the nitric acid, and the product thus ob- 
 tained is the red fuming acid. It acts more energetically 
 than nitric acid, and finds some applications in the lab- 
 oratory and in the arts. When heated it gives off 
 nitrogen peroxide, and if diluted with water it is changed 
 to ordinary nitric acid, as nitrogen peroxide is decom- 
 posed by water, forming nitric acid and nitric oxide or 
 nitrous acid, according to the temperature of the water. 
 
 Nitro-Tiydrochloric Acid or Aqua Regia is a liquid 
 formed by mixing concentrated nitric and hydrochloric 
 acids. It was called aqua regia because it can dissolve 
 gold, the king of the metals. The active power of this 
 liquid as a solvent of metallic substances is due to the 
 fact that it gives off chlorine, and a compound of nitrogen, 
 oxygen, and chlorine which readily gives up its chlorine. 
 This compound has the composition represented by 
 the formula NOC1, and it is best designated by the 
 name nitrosyl chloride. The product of the action, of 
 mtro-hydrochloric acid upon a metal is the correspond- 
 ing chloride. 
 
 Nitrous Acid, HNO 2 . When certain salts of nitric acid 
 are reduced they yield the corresponding nitrites. Thus, 
 when potassium nitrate is heated with metallic lead 
 this reaction takes place : 
 
 KNO 3 + Pb = KNO a + PbO. 
 
282 INORGANIC CHEMISTRY. 
 
 Indeed, if potassium nitrate is heated alone it loses 
 oxygen and is converted into the nitrite : 
 
 2KN0 3 = 2KN0 2 + O 2 ; 
 
 but the reaction is not complete, and the salt thus ob- 
 tained always contains more or less nitrate. 
 
 If an attempt is made to isolate nitrous acid from a 
 nitrite the product is the anhydride, nitrogen trioxide, 
 N 2 O 3 . Thus, if sulphuric acid is added to potassium 
 nitrite the following reaction takes place : 
 
 2KN0 2 + H 2 S0 4 = N 2 3 + H 2 O + K 2 SO 4 . 
 
 It may be that the first action is the liberation of 
 nitrous acid, and that this then breaks down by loss of 
 water. The two reactions are represented thus : 
 
 2KN0 2 + H 2 S0 4 = 2HNO 2 + K a SO 4 ; 
 2HNO 3 = N a 3 + H 2 0. 
 
 A certain analogy will be observed between this action 
 and that which takes place in the action of potassium 
 hydroxide upon an ammonium salt, when ammonium hy- 
 droxide is probably first given off, and then breaks down 
 into ammonia and water. 
 
 Salts are known which are derived from normal nitrous 
 acid, N(OH) 8 , but most of the nitrites are derived from 
 the acid of the formula NO(OH), which is to be regarded 
 as formed from the normal acid by loss of one molecule 
 of water : 
 
 N(OH) 3 = NO(OH) + H 2 O. 
 
 Hyponitrous Acid, H 2 N 2 O 2 . The sodium salt of this 
 acid is made by reducing sodium nitrite in solution by 
 means of sodium amalgam : 
 
 2NaN0 3 + 8H = Na 2 N,O a + 4H a O. 
 
 The acid can also be made by oxidation of hydroxyl- 
 amine. 
 
 It is a solid consisting of white crystalline plates. It 
 is very explosive when freed from water. In water solu- 
 
NITROUS OXIDE. 283 
 
 tion it is much more stable, but at the ordinary tempera- 
 ture it breaks down gradually, the principal products 
 being nitrous oxide and water : 
 
 It would appear from this that nitrous oxide, N 2 O, bears 
 to hyponitrous acid the relation of an anhydride, but 
 salts of the acid cannot be obtained from nitrous oxide 
 directly. The molecular weight of hyponitrous acid in 
 solution has been determined, and the result shows 
 that the acid has the double formula H 2 N a O 2 . Its 
 structure is probably to be represented by the for- 
 
 N(OH) 
 
 mula II 
 
 N(OH) 
 
 Nitrous Oxide, N 2 O. This compound can be obtained 
 by reduction of nitric acid, and is sometimes formed in 
 considerable quantity when copper is treated with the 
 concentrated acid, though when made in this way it 
 is always mixed with a large proportion of nitric oxide. 
 The best way to make it is to heat ammonium nitrate, 
 NH 4 NO 3 , which breaks down into nitrous oxide and water : 
 
 NH 4 N0 3 = N 2 -f 2H 2 O. 
 
 In the same way we have seen that ammonium nitrite 
 breaks down into free nitrogen and water when heated : 
 
 NH 4 N0 2 = N 2 + 2H 2 0. 
 
 In these reactions we see exhibited the tendency of 
 hydrogen and oxygen to combine at elevated tempera- 
 tures. At ordinary temperature this tendency is not 
 strong enough to cause a disturbance of the equilibrium 
 of the parts of the compound. As the temperature is 
 raised and the equilibrium thus disturbed/the affinity 
 of the hydrogen for the oxygen asserts itself. The two 
 elements combine to form water, and the decomposition 
 above represented takes place. 
 
 Nitrous oxide is a colorless, transparent gas which has 
 
284 INORGANIC CHEMISTRY. 
 
 a sweetish taste and odor. Its specific gravity is 1.527. 
 It is somewhat soluble in water ; one volume of water at 
 dissolving somewhat more than its own volume of the 
 gas. It supports combustion almost as well as pure 
 oxygen. Some substances which burn in oxygen do not, 
 however, burn in nitrous oxide. Sulphur which burns 
 in oxygen is extinguished in nitrous oxide, unless it is 
 previously heated to a high temperature. To under- 
 stand the action of this compound in supporting com- 
 bustion it must be borne in mind that, when anything 
 burns in oxygen, the oxygen molecules must first be 
 broken down into atoms before the combination can take 
 place. Thus, when carbon and oxygen are brought to- 
 gether we have at first a condition represented by these 
 symbols : 
 
 the question as to the condition of the carbon being left 
 open. When the temperature is raised to a sufficiently 
 high point the condition is represented thus : 
 
 O + O + C; 
 and now the reaction takes place : 
 
 o + o + c = co 2 . 
 
 In the act of burning in free oxygen, therefore, there is 
 always a certain resistance to be overcome. Now, when a 
 combustible substance is brought into a gas which gives 
 up its oxygen easily the condition is much like that in 
 free oxygen. If the temperature is raised the gas is de- 
 composed, and the oxidation then follows. In the case 
 of nitrous oxide this decomposition takes place : 
 
 and this oxygen in the atomic condition effects the oxi- 
 dation. 
 
 When inhaled, nitrous oxide causes a kind oi intoxica- 
 tion, which is apt to show itself in the form of hysterical 
 laughing. Hence the gas is called laughing gas. In- 
 
NITRIC OXIDE. 285 
 
 haled in larger quantity it causes unconsciousness 
 and insensibility to pain. It is therefore used ex- 
 tensively to prevent pain in some surgical operations, 
 particularly in extracting teeth. 
 
 When subjected to a low temperature and high pres- 
 sure the gas is easily liquefied, and enclosed in properly 
 constructed metallic cylinders the liquid is now sent into 
 the market. In order to get the gas it is only necessary 
 to open the stop- cock of the cylinder. When the liquid 
 comes in contact with the air it rapidly turns to gas, and 
 the temperature is very much lowered in consequence. 
 This causes a part of the liquid to solidify. 
 
 Nitric Oxide, NO. This is the most stable compound 
 of nitrogen and oxygen, and is the common product 
 of the reduction of nitric acid. Thus, when nitric acid 
 acts upon copper and other metallic elements the chief 
 product is generally nitric oxide, though, as we have 
 seen, the reduction may be carried farther. The prin- 
 cipal action in the case of copper is represented thus : 
 
 8HNO, + 3Cu = 3Cu(N0 3 ) 3 + 2NO + 4H,O. 
 
 Considering the ease with which nitric acid gives up 
 its oxygen, and the ease with which copper takes up 
 oxygen, it is probable that the copper abstracts oxygen 
 directly from the acid as represented thus : 
 
 2HNO 3 + 3Cu = CuO + H,O + 2NO. 
 
 In this case the copper oxide would at once form copper 
 nitrate with the excess of nitric acid : 
 
 6HNO 3 + 3CuO = 3Cu(NO 3 ) 2 + 3H 2 O. 
 
 Or, combining the two equations, the total action is rep- 
 resented in the same way as it is above. The nitric 
 acid must not have a specific gravity higher than 1.2, and 
 the temperature must be kept down, otherwise the reduc- 
 tion of the nitric acid is carried farther and considerable 
 nitrous oxide is formed. 
 
286 INORGANIC CHEMISTRY. 
 
 It is possible that to some extent the hydrogen liber- 
 ated from the acid may act as a reducing agent, thus- 
 causing the formation of the lower oxides of nitrogen, 
 as, for example, 
 
 2HN0 3 + 6H = 2NO + 4H a O. 
 
 A good method for making pure nitric oxide consists in 
 treating ferrous chloride, FeCl 2 , with saltpeter. The re- 
 action is represented thus : 
 
 3Fe01 2 + KNO 3 + 4HC1 = 3FeCl 3 + KC1 + 2H 3 O + NO. 
 
 The reducing action is here effected by the ferrous 
 chloride, FeCl a , which tends to pass over into ferric 
 chloride, FeCl 3 , in the presence of hydrochloric acid, and 
 anything which has the power to take up hydrogen. 
 With hydrochloric acid alone it does not form ferric 
 chloride, but if any reducible compound is present ac- 
 tion takes place thus : 
 
 2FeCl 2 + 2HC1 + O = 2FeCl 3 + H 2 O. 
 
 In the above reaction saltpeter furnishes the oxygen, 
 and it is consequently reduced and breaks down, yielding 
 nitric oxide, while the potassium forms potassium chlo- 
 ride with chlorine. 
 
 Nitric oxide is a colorless, transparent gas. Its most 
 remarkable property is its power to combine directly 
 with oxygen when the two are brought together. The 
 act of combination is not accompanied by the appear- 
 ance of light, though heat is evolved. In the reaction 
 which takes place at ordinary temperatures nitrogen 
 peroxide, NO, , is formed : 
 
 NO + O = NO 2 . 
 
 The product is a colored gas, and the change of the color- 
 less nitric oxide to this colored product can therefore 
 easily be recognized. This reaction is, further, the chief 
 cause of the reddish-brown fumes seen when nitric acid 
 acts upon metals and other elements. At a low temper- 
 ature some nitrogen trioxide is formed when oxygen acts 
 upon nitric oxide. 
 
NITROGEN TRIOXIDE. 287 
 
 From what has already been said, it will appear that 
 in nitric oxide the oxygen and nitrogen are more firmly 
 united than in the other oxides. Most burning sub- 
 stances are extinguished when introduced into it, though 
 a few when heated in it to a high temperature extract all 
 or a part of the oxygen. Zinc and iron extract half the 
 oxygen and convert nitric oxide into nitrous oxide. 
 Potassium and sodium decompose it, leaving the nitrogen 
 free. 
 
 A curious reaction by means of which it is possible to 
 separate nitric oxide from other gases takes place, when 
 the oxide is passed into a solution of ferrous sulphate, 
 FeSO 4 . Under these circumstances an unstable dark- 
 colored compound is formed, which appears to have the 
 composition FeSO 4 -|- 2NO. When the solution contain- 
 ing it is heated the pure gas is given off. 
 
 By nascent hydrogen nitric oxide is reduced to am- 
 monia and hydroxylamine. 
 
 Nitrogen Trioxide, N 2 O 3 . This oxide is formed by ad- 
 dition of oxygen to nitric oxide at low temperatures ; by 
 decomposition of the nitrites by means of acids ; and by 
 the combination of nitric oxide with the peroxide at a 
 temperature below 21. The gas given off when nitric 
 acid is reduced with starch or arsenious oxide, As a O 3 , 
 appears to be a mixture of nitric oxide and the peroxide. 
 Pure nitrogen trioxide is a liquid of an indigo-blue color. 
 At a temperature below it undergoes partial decom- 
 position into" nitrogen peroxide and nitric oxide : 
 
 With cold water nitrogen trioxide undergoes decomposi- 
 tion accompanied by an evolution of nitric oxide. Pos- 
 sibly this reaction takes place : 
 
 3N,O 3 +H a O = 2HN0 3 + 4NO. 
 
 By treating the oxide with a solution of sodium hydrox- 
 ide or potassium hydroxide the corresponding nitrite is 
 formed : 
 
 2KOH + N,O 3 = 2KNO 2 + H,O. 
 
288 INORGANIC CHEMISTRY. 
 
 Nitrogen Peroxide, NO 2 . When nitric oxide and oxygen 
 are brought together in the proportion of 2 volumes of 
 the former to 1 volume of the latter they combine com- 
 pletely to form nitrogen peroxide. These relations will 
 be readily understood when it is borne in mind that 2 
 molecules of nitric oxide require 1 molecule of oxygen 
 to effect the change, as is shown in the equation 
 
 The compound is most easily obtained by heating lead 
 nitrate, when nitrogen peroxide and oxygen are given off, 
 and lead oxide remains behind in the vessel : 
 
 Pb(NO 3 ) 2 = PbO + 2NO 2 + O. 
 
 If the gases are passed through a tube surrounded by a 
 freezing mixture the peroxide is condensed to the form 
 of liquid, while the oxygen passes on. When perfectly 
 dry the peroxide is easily solidified. It acts energetic- 
 ally upon compounds which have the power to take up 
 oxygen. When treated with water it undergoes decom- 
 position. If the temperature is low, nitrous and nitric 
 acids are formed : 
 
 2N0 2 + H 2 O = HNO 2 + HNO 3 . 
 
 If the water is hot, however, the products are nitric acid 
 and nitric oxide : 
 
 3NO 2 + H 2 O = 2HNO 3 + NO. 
 
 The nitric oxide thus formed will take up oxygen from 
 the air and yield nitrogen peroxide again, and this, in 
 contact with hot water, will be decomposed, forming 
 nitric acid and nitric oxide, until all the peroxide is con- 
 verted into nitric acid. 
 
 The determinations of the specific gravity of the gas 
 from the peroxide show that at low temperatures the 
 molecular formula is N a O 4 , but that when the tempera- 
 ture 150 is reached the molecule is represented by the 
 formula NO 2 . The compound appears therefore to 
 undergo gradual decomposition or dissociation by heat, 
 so that until the temperature 150 is reached the gas is 
 a mixture of the compounds N a O 4 and 
 
STRUCTURE OF COMPOUNDS OF NITROGEN. 289 
 
 Nitrogen Pentoxide, N 2 O 5 . This compound, which 
 bears to nitric acid the relation of an anhydride, is 
 formed by passing chlorine over silver nitrate and con- 
 densing the product. The reaction takes place thus : 
 
 2AgN0 3 + 01, = N,0 5 + 2AgCl + O. 
 
 It is also formed by treating nitric acid with phosphorus 
 pentoxide, P 2 O 6 , a compound which has a very marked 
 power to unite with water. The action is represented 
 thus : 
 
 2HNO 3 = NA + H 2 O. 
 
 The pentoxide is a crystallized substance, which readily 
 decomposes into nitrogen peroxide and oxygen. In con- 
 sequence of the ease with which it gives up its oxygen it 
 acts violently upon many oxidizable substances. With 
 water it forms nitric acid : 
 
 NA + H 2 = 2HNO,. 
 
 Structure of the Compounds of Nitrogen with Oxygen 
 and Hydrogen. Our knowledge of the structure of the 
 compounds with which we have just been dealing is un- 
 satisfactory. There is at present no way of deciding 
 whether in a compound like nitrous oxide, for example, 
 the oxygen is in combination with both nitrogen atoms : 
 there are no reactions of the compound which throw any 
 light upon this question. Similar difficulties are met 
 with in connection with the other compounds of nitrogen 
 and oxygen. As was remarked on page 264, the simplest 
 view which can be held in regard to these oxides is that 
 in them the nitrogen is univalent, bivalent, trivalent, 
 quadrivalent, and quinquivalent, a view which is ex- 
 pressed by the following formulas : 
 
 ^ N=O 0=N=0 
 
 >0, N=0, 0, 0=N=0, 
 
 These formulas are, however, purely speculative and 
 represent nothing known to us. But if the valence of 
 
290 INORGANIC CHEMISTRY. 
 
 nitrogen can vary in this way, we may also conceive that 
 the oxygen is univalent in all the compounds except 
 nitrous oxide. Thus nitric oxide may be represented by 
 
 the formula N-O, nitrogen peroxide by N < Q, etc. On 
 
 the other hand, there is an unmistakable tendency on 
 the part of the elements to act either with even valences, 
 as 2, 4, 6, etc., or with odd, as 1, 3, 5, etc. This is beauti- 
 fully shown by the members of the chlorine group and 
 those of the sulphur group. It has been pointed out 
 that the relations between the compounds of chlorine, 
 bromine, and iodine can be explained, by assuming that 
 these elements are univalent, trivalent, quinquivalent, 
 and septivalent ; and that the relations between the 
 compounds of sulphur, selenium, and tellurium can be 
 equally easily explained by assuming that these elements 
 are bivalent, quadrivalent, and sexivalent. In the case 
 of nitrogen and the elements belonging to the same group 
 we should naturally expect to find a similar law of com- 
 position holding good. As far as the hydroxyl deriva- 
 tives, represented by nitrous acid and nitric acid, are 
 concerned, the same regularity is observed as in the case 
 of sulphur. In nitric acid the nitrogen is probably quin- 
 quivalent, and in nitrous acid trivalent. Further, in 
 ammonia nitrogen is trivalent, while it is probably quin- 
 quivalent in the ammonium compounds, as has been 
 pointed out (see p. 275). It is clear that nitrogen tends 
 to act either as a trivalent or quinquivalent element. 
 "Whether it ever acts as a univalent element it is impos- 
 sible to say, for, while the existence of the compound 
 N 2 O seems to show that it does, this same compound 
 may be explained on the assumption that 111 it the ni- 
 
 N 
 trogen is trivalent, as shown in the formula || >O; and 
 
 N 
 
 indeed there is no difficulty in assuming any desired 
 valence for the nitrogen. Taking the compound nitric 
 oxide, there seems to be no escape here from the con- 
 clusion that the nitrogen is bivalent if the oxygen is bi- 
 valent ; and the compound forms a striking exception to 
 
STRUCTURE OF COMPOUNDS OF NITROGEN. 291 
 
 the rule above referred to that the valence of an element 
 generally changes from odd to odd or from even to even. 
 It may be said that this compound is unsaturated, and 
 that one of its bonds is unemployed, a condition which 
 may be symbolized by this expression, -N= O, but this 
 does not help us out of the difficulty, and, further, this 
 conception is not in accordance with the fact that nitric 
 oxide takes up one atom of oxygen to form nitrogen per- 
 oxide, NO 2 . And then the question arises, What is the 
 structure of this last-mentioned compound ? Should it 
 be represented thus : O=N=O? If so the nitrogen is 
 quadrivalent. But it passes readily into the form N 2 O 4 . 
 It may be that this act consists simply in the union of 
 the two molecules by means of the fifth bond of quin- 
 quivalent nitrogen, the structure of the resulting mole- 
 cule being represented thus : I . All this is, how- 
 ever, almost pure speculation, and, at the present stage 
 of our knowledge of the subject of structure, the above 
 formulas have very little value. Still it must not be 
 forgotten that the structure of all chemical compounds 
 is a legitimate subject of investigation. 
 
 When we come to the acids of nitrogen it is seen, as 
 has already been pointed out, that these can be explained 
 very satisfactorily by the aid of the same hypothesis 
 that served so well in dealing with the acids of iodine 
 and of sulphur. Nitric acid is to be regarded as derived 
 from the maximum hydroxyl compound of quinquivalent 
 nitrogen, known as normal nitric acid, by loss of water ; 
 and in a similar way nitrous acid is to be regarded as 
 derived from the maximum hydroxyl compound of tri- 
 valent nitrogen, or normal nitrous acid, by loss of water. 
 A few salts are known which appear to be derived from 
 the normal acids, but for the most part all the hydrogen 
 atoms of these normal acids are not replaceable by 
 metals, and the formation of salts generally involves a 
 breaking down of the compound into water and the com- 
 mon form of the acid. 
 
 Compounds of Nitrogen with the Elements of the Chlo- 
 rine Group. Notwithstanding the ease with which chlo- 
 
292 INORGANIC CHEMISTRY. 
 
 rine combines with most elements, and the stability of 
 the compounds which it forms with them, its compound 
 with nitrogen is extremely unstable. It can be made by 
 the action of chlorine on ammonia, and by decomposing 
 a solution of ammonium chloride by means of an electric 
 current. In the latter case chlorine is liberated at one 
 of the poles and then acts upon the ammonium chloride : 
 
 NH 4 C1 + 601 = 4HC1 + NCI,. 
 
 It appears that when chlorine acts upon ammonia differ- 
 ent products are formed by successive replacement of 
 the hydrogen atoms of the ammonia by chlorine, thus : 
 
 NH 3 + C1 2 = NH,C1 + HC1 ; 
 NH 2 C1 + Cl, = NHC1 2 + HC1 ; 
 
 NHCl a + C1 2 = NC1 3 + HC1. 
 
 According to this, the trichloride of nitrogen is the final 
 product of the substituting action of chlorine upon am- 
 monia. The compound is an oil, which undergoes de- 
 composition very readily. It is, indeed, one of the most 
 explosive substances known. It is decomposed by heat, 
 and especially by contact with certain substances, among 
 which are oil of turpentine and caoutchouc. It is slowly 
 decomposed by water, though, probably owing to the 
 slight affinity of nitrogen for oxygen, the decomposition 
 does not take place as readily as that of the compounds 
 of sulphur and chlorine. Direct sunlight causes explo- 
 sion of the chloride. 
 
 When ammonia is treated with iodine reactions take 
 place similar to those which take place with chlorine. 
 The products are the iodides of nitrogen, the final 
 product of the action being the tri-iodide, NI 3 . These 
 compounds, like the corresponding chlorine compounds, 
 are extremely explosive. The simplest way to prepare 
 them is to place a little powdered iodine on a filter and 
 pour concentrated ammonia over it. The substance 
 should be made in only very small quantities at a time. 
 "When dried it decomposes with violent explosion by 
 contact even with soft substances ; and it will also ex- 
 
COMPOUNDS OF NITROGEN WITH SULPHUR, ETC. 293 
 
 plode if left entirely undisturbed. The different com- 
 pounds called nitrogen iodide are slowly decomposed by 
 water. 
 
 Compounds of Nitrogen with the Members of the Sul- 
 phur Group. Nitrogen combines with sulphur forming 
 two compounds, N 4 S 4 and N 6 S 2 , both of which are well 
 characterized. Among the most interesting compounds 
 containing sulphur and nitrogen is that which has been 
 referred to as nitrosyl-sulphuric acid in connection with 
 the account of the manufacture of sulphuric acid. It is 
 formed by the action of sulphur dioxide on fuming nitric 
 acid : 
 
 also in the manufacture of sulphuric acid by the action 
 of sulphur dioxide, water, and oxygen upon a mixture 
 of nitrogen peroxide and nitric oxide : 
 
 280, + H.O + N,0, + 20 = 280,<gg>; 
 
 and by the action of the nitrogen peroxide upon sul- 
 phuric acid : 
 
 2N0 2 + S0 g < = 80,< +HNO a . 
 
 When treated with water it breaks down into sulphuric 
 acid and nitrogen trioxide. It will be remembered that, 
 in order to prevent loss of oxides of nitrogen in the sul- 
 phuric acid factories, the gases are brought in contact 
 with concentrated sulphuric acid in the Gay Lussac tower 
 before being allowed to escape. The oxides form with 
 sulphuric acid compounds similar to nitrosyl-sulphuric 
 acid, and when these are diluted with " chamber acid " 
 and heated by the hot gases from the sulphur furnace 
 the oxides of nitrogen are given up. 
 
CHAPTER XVII. 
 
 ELEMENTS OF FAMILY V, GROUP B : 
 
 PHOSPHORUS ARSENIC ANTIMONY BISMUTH. 
 
 THE ELEMENTS AND THEIR COMPOUNDS WITH 
 
 HYDROGEN. 
 
 General. The elements of this group bear to nitrogen 
 very much the same relations that the members of the 
 sulphur group bear to oxygen, and those of the chlorine 
 group bear to fluorine. In general they form compounds 
 of the same character and of similar composition. At 
 the same time gradations in properties are noticed in 
 passing from one end of the group to the other. Like 
 nitrogen, the elements of the group are strongly marked 
 acid-formers, though this character grows less marked 
 from nitrogen to bismuth. Antimony is both an acid- 
 forming and a base-forming element, while bismuth is 
 more basic than acid. The stability of the hydrogen 
 compounds decreases from nitrogen to antimony ; while 
 bismuth does not form a compound with hydrogen. 
 Ammonia, as we have s"een, is strongly basic ; the cor- 
 responding compound of phosphorus and hydrogen has 
 weak basic properties, while those of arsenic and anti- 
 mony have no basic properties. These hydrogen com- 
 pounds correspond in composition to ammonia. They 
 are : 
 
 NH 3 : PH 3 AsH 3 SbH 3 
 
 With chlorine they all form compounds corresponding 
 to nitrogen trichloride, and phosphorus and antimony 
 form compounds in which they are quinquivalent, while 
 
 (294) 
 
ELEMENTS OF FAMILY V, GROUP B. 295 
 
 bismuth forms a chloride, Bi 2 Cl 4 , in addition to the tri- 
 chloride. The compounds referred to are : 
 
 Bi 2 Cl 4 
 
 NC1 3 : PC1 3 AsCl 3 SbCl s BiCl 3 
 PC1 5 SbCl 6 
 
 They all form two oxides corresponding to nitrogen tri- 
 oxide and pentoxide : 
 
 N 2 3 : P 2 3 As 2 3 Sb 2 3 Bi 2 O 3 
 N 2 6 : P 2 5 As A Sb 2 5 BiA 
 
 No one of the elements of the group forms as great a 
 variety of compounds with oxygen as nitrogen does. 
 Antimony, however, forms the oxide Sb 2 O 4 , correspond- 
 ing to nitrogen peroxide, N 2 O 4 ; and bismuth forms the 
 oxide Bi 2 O 2 or BiO, corresponding to nitric oxide, NO. 
 
 The hydroxyl compounds or acids, like those of nitro- 
 gen, are related to the maximum hydroxyl compounds 
 of the elements with the valence 5, and to the maximum 
 hydroxyl compounds of the elements with the valence 3. 
 That is to say, they may be regarded as derived from a 
 hydroxide of the general formula M(OH) B , and another 
 of the formula M(OH) 3 . Where M is nitrogen these acids 
 break down to the forms NO 2 (OH) and NO(OH) by loss 
 of one or two molecules of water. In the case of the 
 elements of the phosphorus group, however, the breaking 
 down is not generally carried as far as with nitrogen. 
 The general rule is the same as in the sulphur and chlo- 
 rine groups : the normal acid breaks down to form com- 
 pounds containing the same number of hydrogen atoms 
 as the hydrogen compounds of the elements. Thus the 
 hydroxyl derivatives of chlorine generally break down to 
 form compounds containing one atom of hydrogen, or the 
 same number that is contained in the hydrogen compound, 
 hydrochloric acid, thus: Cl(OH), yields C1O 3 (OH) ; 
 C1(OH) 6 yields GLO a (OH), etc. So, also, in the sulphur 
 group, S(OH) 6 yields SO 9 (OH),, etc., the number of hy- 
 drogen atoms in the common form of the acid being the 
 same as that in the hydrogen compound of sulphur, SH a . 
 
296 INORGANIC CHEMISTRY. 
 
 This rule does not hold good for nitrogen, for the 
 tendency here is to break down to compounds contain- 
 ing one atom of hydrogen. On the other hand, phos- 
 phorus, arsenic, and antimony follow the rule, the 
 principal acids of these elements containing three atoms 
 of hydrogen in the molecule. As already stated, there 
 are two series of these represented by the following for- 
 mulas : 
 
 H 3 P0 3 H 3 As0 3 H 3 Sb0 3 
 H 3 PO 4 H 3 AsO 4 H 3 SbO 4 
 
 Besides the above, however, phosphorus forms several 
 other acids. The principal ones bear simple relations to 
 the acid H 3 PO 4 , which is called orthophosphoric acid. 
 The simplest view in regard to the acid of phosphorus 
 of the formula H 3 PO 4 , and the corresponding compounds 
 of arsenic and antimony, is that it is derived from the 
 corresponding normal acid by loss of one molecule of 
 water. Thus, normal phosphoric acid is P(OH) 5 . By 
 loss of one molecule of water this yields ordinary or 
 orthophosphoric acid : 
 
 P(OH) 5 = PO(OH) 3 + H 2 0. 
 
 Normal phosphorous acid is P(OH) 3 . Whether or- 
 dinary phosphorous acid has this structure is a ques- 
 tion very difficult to answer at present. By loss of 
 another molecule of water orthophosphoric acid is con- 
 verted into metaphosphoric acid, which in composition 
 corresponds to nitric acid. Its formation is represented 
 thus : 
 
 PO(OH) 3 = P0 2 (OH) + H 2 0. 
 
 The series of phosphorus acids, H 3 PO 2 , H 3 PO 3 , and 
 H 3 PO 4 is strongly suggestive of the series of sulphur 
 acids, H 2 SO 2 , H 2 SO 3 , and H 2 SO 4 , and of the series of 
 chlorine acids, HC1O, HC1O 2 , HC1O 3 , and HC1O 4 . 
 
 Arsenic and antimony also form acids corresponding 
 to metaphosphoric acid, known respectively as metarsenic 
 and metantimonic acids. 
 
ELEMENTS OF FAMILY V, GROUP B. 297 
 
 By elimination of one molecule of water from two 
 molecules of ordinary phosphoric acid there is formed 
 an acid H 4 P 2 O,, known as pyrophosphoric acid, which 
 bears to ordinary phosphoric acid much the same rela- 
 tion that pyrosulphuric or disulphuric acid bears to 
 ordinary sulphuric acid : 
 
 /OH 
 SO/ 
 
 >0 + H 3 0. 
 
 /OH 
 PO-OH 
 
 = \ H 
 
 /OH PO/OH 
 ^-(Jl \rm" 
 
 \OH 
 
 Ordinary arsenic and antimonic acids yield corre- 
 sponding derivatives known as pyroarsenic and pyroanti- 
 monic acids. The elements of the phosphorus group 
 form compounds with oxygen and chlorine known as the 
 oxy chlorides, which in general resemble the oxychlorides 
 of the members of the sulphur group. Examples of these 
 compounds are phosphorus oxychloride, POC1 3 , anti- 
 mony oxychloride, SbOCl, and bismuth oxychloride, 
 BiOCl. Phosphorus oxychloride is readily decom- 
 posed by water, forming phosphoric and hydrochlo- 
 ric acids : 
 
 ( Cl HOH ( OH 
 
 PO^ Cl + HOH = PO^ OH + 3HC1. 
 ( Cl HOH ( OH 
 
 The oxychlorides of antimony and bismuth are not 
 completely decomposed by water. This is in accordance 
 with the fact to which attention has been called that the 
 chlorides of the acid-forming elements are in general 
 easily decomposed by water and converted into hydroxyl 
 compounds, while the chlorides of the base-forming ele- 
 ments are not readily decomposed in this way, but, on 
 the contrary, their oxides and hydroxides are, as a rule, 
 
298 INORGANIC CHEMISTRY. 
 
 readily converted into chlorides by hydrocliloric acid 
 Elements which, like antimony and bismuth, play the 
 part of base-formers and acid-formers form stable 
 oxy chlorides. 
 
 Of the elements of this group phosphorus occurs 
 most abundantly in nature, arsenic and antimony next, 
 and bismuth least abundantly. Arsenic and bismuth 
 occur to some extent in the uncombined condition. 
 Phosphorus and antimony occur in combination. All 
 the elements of the group find applications in the arts, 
 either as the elements or in the form of compounds. 
 
 PHOSPHORUS, P (At. Wt. 30.79). 
 
 
 
 Occurrence. The name phosphorus is derived from 
 the Greek 0c3?, light, and (pspeiv, to carry, on account 
 of the fact that it always gives light and takes fire very 
 easily. The element occurs in nature in the form of 
 phosphates derived from orthophosphoric acid, H 3 PO 4 . 
 The chief of these is calcium phosphate, Ca 3 (PO 4 ) 2 , which 
 is the principal constituent of the minerals phosphorite 
 and apatite, and of the ashes of bones. The phosphates, 
 like the nitrates, are widely distributed in the soil and 
 are of fundamental importance in the process of plant 
 life. The phosphates found in the bones are taken into 
 the animal body in the food. All plants used as food 
 contain small quantities of the phosphates which they 
 get from the soil. The phosphates taken into the body 
 are partly given off in the excrement and urine, and it 
 was in an examination of urine made in the hope of 
 finding the philosopher's stone that phosphorus was 
 first discovered in 1669. At present phosphorus is 
 made almost entirely from bones. 
 
 Preparation. Besides the phosphates, considerable 
 quantities of organic materials are contained in bones. 
 When the bones are burned the organic materials pass 
 off for the most part in the form of carbon dioxide, water, 
 and volatile compounds containing nitrogen, and the so- 
 called mineral or earthy portions, the chief constituent 
 of which is tertiary calcium phosphate, or phosphoric 
 
PHOSPHORUS: OCCURRENCE- PREPARATION. 299 
 
 acid in which all the hydrogen is replaced by calcium, 
 remain behind. As calcium is bivalent and there are 
 three atoms of hydrogen in the molecule of phosphoric 
 acid, H 3 PO 4 , the simplest way in which all the hydrogen 
 atoms of the acid can be replaced by calcium is that 
 represented by the formula Ca 3 (PO 4 ) 2 , the six atoms of 
 hydrogen in two molecules of the acid being replaced by 
 three bivalent atoms of calcium. The problem now is 
 to isolate the phosphorus from this calcium phosphate. 
 The salt is insoluble in water, and there is no simple 
 way by which the phosphorus can be set free from it. 
 When it is treated with sulphuric acid calcium sulphate 
 which is difficultly soluble is deposited and phosphoric 
 acid is set free and remains in solution. The reaction is 
 represented as follows : 
 
 Ca 3 (P0 4 ) 2 + 3H 2 SO 4 = 2H,PO t + 3CaSO 4 . 
 
 The calcium sulphate, or gypsum, is allowed to settle 
 and is then filtered off and washed. The solution of 
 phosphoric acid is evaporated until it has the specific 
 gravity 1.325 to 1.5. Then it is mixed with coarsely- 
 ground wood charcoal, coke, or sawdust, and carefully 
 dried in a cast-iron pot or muffle furnace. In this pro- 
 cess the orthophosphoric acid, H 8 PO 4 , is converted into 
 metaphosphoric acid, HPO 3 , 
 
 H 3 P0 4 = HP0 3 + H 2 ; 
 
 and in case sawdust is used this is changed, so that 
 the resulting mixture consists of charcoal and metaphos- 
 phoric acid. This mixture is next subjected to distilla- 
 tion in clay retorts, when the metaphosphoric acid is re- 
 duced according to the following equation : 
 
 2HP0 3 + 60 = H 2 + 6CO + 2P. 
 
 The phosphorus passes over in the form of vapor, and 
 is collected under water. The crude phosphorus thus 
 obtained must be subjected to a cleansing process before 
 it can be used. For this purpose it is pressed, while in 
 the molten condition under water, through chamois 
 leather, or it is distilled again from iron retorts ; or, still 
 better, it is treated with chromic acid as follows : It is 
 fused under water, then a little potassium or sodium 
 
300 INORGANIC CHEMISTRY. 
 
 bichromate in solution is added, and afterwards an equiv- 
 alent proportion of sulphuric acid, and the whole al- 
 lowed to stand for two hours or more. The phosphorus 
 is then washed with hot water, and after being si- 
 phoned off it is filtered through canvas bags. The 
 phosphorus is then cast into sticks in tin tubes. In this 
 form it generally comes into the market. 
 
 At the time of the last report available there were 
 manufactured in one year about 1200 tons of phospho- 
 rus in two factories, one of which is in England and the 
 other in France. Quite recently phosphorus has been 
 manufactured to some extent in Sweden. 
 
 Properties. Ordinary phosphorus is colorless or 
 slightly yellowish, translucent, and at ordinary tempera- 
 tures it can be cut like wax, but it becomes hard and 
 brittle at low temperatures. It melts at 44, and boils 
 at 290. It is insoluble in water. When kept under 
 water for any length of time in dispersed light it be- 
 comes opaque, crystalline on the surface, and yellow. 
 It is soluble in carbon disulphide, and crystallizes when 
 deposited from this solution. It gives off fumes in con- 
 tact with the air, and emits a pale light which is known 
 as a phosphorescent light. It is very poisonous, the in- 
 halation of the vapor in small quantities causing very 
 serious disturbance of the system. The workmen in the 
 factories where phosphorus is made or used are fre- 
 quently affected by phosphorus-poisoning. Among the 
 prominent symptoms is gradual decomposition of the 
 bones. When taken into the stomach phosphorus also 
 acts as a poison and causes death. When heated in the 
 air it takes fire at 50. It also takes fire by rubbing, 
 and it must be handled with the greatest care, as wounds 
 caused by it are dangerous and difficult to heal. When 
 it burns in the air it is converted into the pentoxide, 
 P 2 O 5 , which is also the product of its combustion in 
 oxygen, as we have seen. It combines also with other 
 elements directly, frequently with evolution of light. 
 Thus, when it is brought together with chlorine, bromine, 
 and iodine, it forms the compounds PC1 3 , PBr 3 , and PI 3 . 
 It also combines with sulphur. When a piece is put in 
 water and the water boiled, a part of the phosphorus 
 
PROPERTIES OF PHOSPHORUS. 301 
 
 passes over, and if the water vapor is condensed in a 
 glass tube in a dark room, it is seen to be phosphores- 
 cent. This furnishes a convenient method for its detec- 
 tion, as, for example, in a case of suspected poisoning by 
 phosphorus. 
 
 Owing to its strong tendency to combine with oxygen, 
 it abstracts the element from some of its compounds. 
 Thus, if a solution of phosphorus in carbon disulphide 
 is added to a solution of copper sulphate, metallic cop- 
 per is thrown down, while at the same time copper 
 phosphate and a compound of copper and phosphorus 
 are formed. 
 
 When phosphorus is left for a long time under water 
 and subjected to the action of light, it becomes at first 
 yellow, then reddish, and finally red. The same change 
 takes place when phosphorus is heated for a time in an 
 atmosphere which is free from oxygen; and rapidly 
 when it is heated to 300 in an hermetically sealed tube. 
 The red substance thus obtained has properties entirely 
 different from those of ordinary phosphorus. It is a 
 red powder, which frequently has a crystalline struc- 
 ture. It does not emit light. It does not melt at a low 
 temperature. It is not poisonous, and cannot be easily 
 ignited. Further, it is perfectly insoluble in carbon 
 disulphide. In every respect this red modification of 
 phosphorus conducts itself as a much less active sub- 
 stance chemically than ordinary phosphorus. In an 
 atmosphere of carbon dioxide it is converted into ordi- 
 ary phosphorus when heated to 261, and if heated to 
 this temperature in the air it takes fire, and then forms 
 the same product that ordinary phosphorus does in 
 burning. 
 
 When phosphorus is heated with lead for eight to ten 
 hours to a very high temperature in sealed tubes from 
 which the air has been exhausted, and the whole then 
 allowed to cool, the surface of the lead is found covered 
 with black, laminated crystals, which undergo no change 
 in the air. Crystals are also found in the interior of the 
 lead. This variety of phosphorus is called crystallized, 
 metallic phosphorus on account of the metallic lustre. It 
 is not as volatile as the ordinary variety. 
 
302 INORGANIC CHEMISTRY. 
 
 When the vapor of phosphorus is suddenly cooled by 
 ice water in an atmosphere of hydrogen, it is deposited 
 in the form of a snow-white powder on the surface of 
 the water. Under water this variety undergoes very 
 little change even when exposed for a long time to the 
 action of the sunlight. When exposed to the air on 
 filter-paper it gives off dense fumes, and then melts, 
 forming ordinary phosphorus, but it does not generally 
 take fire under these circumstances. 
 
 Treated with oxidizing agents, as, for example, nitric 
 acid, phosphorus is slowly converted into phosphoric 
 acid, just as sulphur is converted into sulphuric acid 
 under the same conditions. 
 
 Applications of Phosphorus.- Phosphorus is used prin- 
 cipally in the manufacture of matches and as a poison 
 for vermin. Various mixtures are used for making 
 matches. Nearly all of them contain phosphorus to- 
 gether with some oxidizing compound, and some neutral 
 substance to act as a medium for holding the constitu- 
 ents together. An example is a mixture consisting of 2 
 parts phosphorus, 1 part manganese dioxide, 3 parts 
 chalk, -J part lamp-black, and 5 parts glue. The mix- 
 ture used in the manufacture of the so-called " safety 
 matches" consists of potassium chlorate, potassium 
 dichromate, minium, and antimony trisulphide. This 
 will not ignite by simple friction, but will ignite when 
 drawn across a paper upon which is a mixture of red 
 phosphorus and antimony pentasulphide. 
 
 Compounds of Phosphorus with Hydrogen. There are 
 three compounds of phosphorus with hydrogen, a gaseous 
 compound of the formula PH 3 , corresponding to am- 
 monia ; a liquid of the formula PH 2 , or P Q H 4 , correspond- 
 ing to hydrazine ; and a solid of the formula P,H 3 , or P 4 H a . 
 
 Phosphine, Gaseous Phosphuretted Hydrogen, PH 3 . 
 This compound is formed when phosphorous or hypo- 
 phosphorous acid is heated. The decompositions take 
 place as represented in these equations : 
 
 4H 3 P0 3 :=3H 3 P0 4 
 
 2 = H 3 P0 4 -fPH 8 . 
 
PHOSPHINE. 303 
 
 "We see here an example of the same kind of action that 
 was referred to in connection with the sulphur com- 
 pounds. It will be remembered that, in general, when 
 a salt of any oxygen acid of sulphur except sulphuric 
 acid is heated it is converted into the sulphate, and that 
 the other elements arrange themselves in simpler forms 
 of combination. Thus, when sodium thiosulphate is 
 heated it is converted into sodium sulphate and sodium 
 pentasulphide, as represented in the following equation : 
 
 4Na 2 S 2 O 3 = 3Na 2 SO 4 + Na 2 S 5 . 
 
 So also sodium sulphite yields sodium sulphate and 
 sodium sulphide : 
 
 4Na 2 SO 8 = 3Na 2 SO 4 + Na a S. 
 
 Other ways of making phosphine are : (1) By treating 
 a strong solution of potassium hydroxide with phos- 
 phorus, when reaction takes place as follows : 
 
 3KOH + 4P + 3H 2 = 3KH 2 PO 2 + PH 3 . 
 
 The compound KH 2 PO 2 is known as potassium hypo- 
 phosphite, being derived from hypophosphorous acid, 
 H 3 PO. 2 . (2) By treating zinc phosphide with dilute 
 hydrochloric acid. Assuming that zinc phosphide has 
 the composition represented by the formula Zn 3 P 2 , the 
 reaction with hydrochloric acid takes place according to 
 the equation 
 
 Zn 3 P 2 + 6HC1 = 3ZnCl 2 + 2PH 3 . 
 
 (3) By treating phosphonium iodide, PH 4 I, with water or 
 a dilute solution of potassium hydroxide : 
 
 PH 4 I + H 2 =PH 3 + HI + H 2 O; 
 
 PHJ + KOH = PH 3 + KI + H 2 0. 
 
 When made from phosphorus and potassium hydroxide 
 it always contains a considerable proportion of hydrogen, 
 for the reason that potassium hypophosphite gives off 
 hydrogen when heated with a solution of potassium 
 hydroxide. From calcium phosphide and from phos- 
 phonium iodide it can be obtained in pure condition. 
 
304 INORGANIC CHEMISTRY. 
 
 Phosphine is a colorless gas with an unpleasant, gar- 
 lic-like odor. It is insoluble in water, and is poisonous. 
 It burns, but does not take fire spontaneously when 
 pure. When burned with free access of air the products 
 of combustion are phosphorus pentoxide and water : 
 
 whereas when it is burned in a cylinder so that the air 
 has not free access to it, the products are water and 
 phosphorus, which is deposited in a reddish layer upon 
 the glass. 
 
 Although pure phosphine does not take fire spontane- 
 ously when brought in contact with the air, the gas 
 made by any one of the methods above referred to is 
 pretty sure to contain some of the liquid compound of 
 phosphorus and hydrogen, P 2 H 4 , which is spontaneously 
 inflammable, and therefore the gas takes fire. If it is 
 collected in a glass vessel over water, and allowed to 
 stand so that the light acts upon it, the liquid phosphine 
 is decomposed into the gaseous and solid varieties, and 
 the gas which is left no longer has the property of tak- 
 ing fire spontaneously. Phosphine is much less stable 
 than ammonia. When heated or when treated with elec- 
 tric sparks it is easily decomposed into phosphorus and 
 hydrogen. While ammonia dissolves in water, probably 
 forming the hydroxide NH 4 (OH), phosphine is only very 
 slightly soluble in water. Ammonia combines with acids 
 very energetically, forming the ammonium salts, and we 
 should expect to find that similar salts are formed by 
 the action of phosphine on acids ; but only a few such 
 compounds are known, and these are unstable. Thus, 
 when phosphine is brought together with hydrochloric, 
 hydrobromic, and hydriodic acids, the reactions repre- 
 sented by the following equations take place : 
 
 PH 3 + HC1 = PH 4 C1 ; 
 PH 3 + HBr = PH 4 Br ; 
 PH 3 + HI =PH 4 I. 
 
 The products are called respectively phosphonium Mo- 
 ride, bromide, and iodide. The reactions are, as will be 
 
ARSENIC : OCCURRENCE PREPARATION. 305 
 
 seen, perfectly analogous to those which take place be- 
 tween the same acids and ammonia. But the products 
 are much less stable than the ammonium salts. The 
 bromide when exposed to the air attracts water and 
 decomposes rapidly, forming hydrobromic acid and 
 phosphine. Phosphonium iodide undergoes a similar 
 decomposition. 
 
 AESENIC, As (At. Wt. 74.44). 
 
 Occurrence. Arsenic occurs in nature to some extent 
 in the uncombined condition or native. Compounds of 
 the metals with arsenic, or the arsenides, occur very 
 widely distributed, and they frequently accompany, and 
 are similar to, the sulphides. The most common com- 
 pound of this kind is the so-called arsenical pyrites, 
 which has the composition FeAsS, and may therefore 
 be regarded as iron pyrites, FeS 2 , in which one atom of 
 sulphur has been replaced by one atom of arsenic. 
 Among other arsenic compounds deserving special men- 
 tion are the two arsenides of iron of the formulas FeAs 2 
 and Fe 2 As 3 , which are apparently analogous to the sul- 
 phides FeS 2 and Fe 2 S 3 ; and, further, the sulphides of 
 arsenic, orpiment, As 2 S 3 , and realgar, As 2 S 2 . The oxide 
 As 2 O 3 occurs in considerable quantity, and also salts of 
 arsenic acid, or the arsenates, which in composition are 
 analogous to the phosphates. 
 
 Preparation. The arsenic which conies into the market 
 is either that which occurs native or it is made from 
 arsenical pyrites by heating : 
 
 FeAsS = FeS + As. 
 
 The arsenic thus obtained is not pure. By bringing a 
 little iodine in the bottom of a porcelain crucible, put- 
 ting the arsenic upon it, and heating, the arsenic ac- 
 quires a high metallic lustre, and once in this condition 
 it will remain so for some time even when exposed to 
 the air. 
 
 Properties. Arsenic has a metallic lustre and steel 
 color. It is very brittle. When heated it volatilizes 
 
306 INORGANIC CHEMISTRY. 
 
 without melting. At red heat it burns with a bluish 
 flame, and the vapor given off has the odor of garlic. 
 This odor produced under such circumstances is very 
 characteristic of arsenic, and furnishes one of the means 
 for detecting it. Arsenic combines with most elements 
 directly, the action being accompanied in some cases, as 
 in that of chlorine, by an evolution of light. As an ele- 
 ment it is not poisonous, but when oxidized to the form 
 of the oxide As 2 O 3 it is extremely poisonous. As it is 
 easily oxidized, the element itself may act as a poison. 
 When boiled with nitric acid arsenic is converted into 
 arsenic acid, H 3 AsO 4 , just as phosphorus is converted by 
 nitric acid into phosphoric acid, H 3 PO 4 . 
 
 One peculiarity in the conduct of arsenic is suggestive, 
 and that is its power to form compounds which are an- 
 alogous to the compounds of sulphur. There are a 
 number of compounds similar to arsenical pyrites which 
 appear to be perfectly analogous to the sulphur com- 
 pounds, and in them it seems necessary to assume that 
 the arsenic plays the same part as the sulphur. On the 
 other hand, arsenic conducts itself in nearly all its com- 
 pounds like phosphorus. This power to play double 
 parts is not uncommon among the elements, and we shall 
 hereafter meet with a number of examples. The case of 
 manganese is one in point. While it conducts itself in 
 some of its compounds like the members of the chlo- 
 rine group, to which on account of its position in the 
 periodic system we should expect to find it related, yet 
 it is perhaps more closely related to iron and chromium, 
 which belong to different groups ; and so, also, chromium, 
 which in many respects resembles sulphur very striking- 
 ly, is like iron and aluminium in other respects. 
 
 Arsine, Arseniuretted Hydrogen, AsH 3 . This com- 
 pound is analogous to ammonia and to gaseous phos- 
 phine. It is made by reduction of compounds of arsenic 
 containing oxygen, as arsenic trioxide or arsenic acid ; 
 and also by treating a compound of zinc and arsenic with 
 dilute sulphuric acid. The reactions involved in the 
 first method are 
 
ARSINE. 307 
 
 As 2 O 3 + 6H 2 = 2AsH 3 + 3H 2 O ; 
 H 3 AsO 4 + 4H 2 = AsH 3 + 4H 2 O. 
 
 - --_ 
 
 That involved in the second method mentioned is : 
 As 3 Zn 3 + 3H 2 SO 4 = 2AsH 3 + 3ZnSO 4 . 
 
 Arsine is a colorless gas with an odor suggestive 
 of garlic. It is extremely poisonous, even very small 
 quantities being capable of producing bad effects, and it 
 requires but little to cause death. When ignited in the 
 air it takes fire and burns with a pale blue flame, the 
 products of the combustion being arsenic trioxide, As 2 O 3 , 
 and water. If the air is prevented from gaining free 
 access to it the hydrogen burns, but the arsenic is 
 deposited as a brownish layer. The gas is so unstable 
 that, when it is passed through a glass tube heated to 
 redness, it is decomposed into arsenic and hydrogen, the 
 former being deposited just in front of the heated por- 
 tion of the tube as a thin, almost black, layer with a high 
 metallic lustre. 
 
 Arsine is easily decomposed by most active chemical 
 substances. Water and concentrated acids decompose 
 it ; as do chlorine, bromine, and iodine, which form with 
 it the corresponding acids, and compounds of chlorine, 
 bromine, and iodine with arsenic. 
 
 Passed into a solution of a metallic salt, arsine either 
 reduces the salt and throws down the metal as in the 
 case of silver ; or it forms an arsenide of the metal, acting 
 in this case very much as hydrogen sulphide does when 
 passed into similar solutions. Considering the instability 
 of arsine, it is not surprising that it acts as a reducing 
 agent. It will be remembered that hydriodic acid and 
 hydrogen sulphide act in the same way towards some 
 oxygen compounds, and the action is due to their break- 
 ing down into hydrogen and the other element. Thus, 
 when hydriodic acid acts as a reducing agent the iodine 
 is left uncombined, and when hydrogen sulphide acts in 
 this way the sulphur is left. But when arsine acts as a 
 reducing agent both the hydrogen and the arsenic com- 
 
308 INORGANIC CHEMISTRY. 
 
 bine with oxygen. Thus, when arsine is passed into a 
 solution of silver nitrate this reaction take? place : 
 
 AsH 3 + 6AgN0 3 + 3H 2 = As(OH) 3 + 6HNO 3 + 6Ag. 
 
 When, on the other hand, arsine is passed through a 
 solution of a salt of a difficultly reducible metal, the ar- 
 senide of the metal is thrown down : 
 
 2AsH 3 + 3CuS0 4 = As 2 Cu 3 + 3H 2 SO 4 . 
 
 Arsine does not combine with acids to form arsonium 
 compounds such as AsHJ, analogous to ammonium and 
 phosphonium compounds. 
 
 There is a second compound of arsenic and hydrogen 
 which is solid and appears to have the composition rep- 
 resented by the formula As 2 H 2 . 
 
 ANTIMONY, Sb (At. Wt. 119.52). 
 
 Occurrence. Antimony occurs in nature chiefly in the 
 form of stibnite, which is the trisulphide Sb 2 S 3 . This also 
 occurs very widely distributed in nature in combination 
 with sulphides of various metals, as copper, lead, and 
 silver. The element is made from the sulphide either by 
 heating it with iron, with which the sulphur combines, 
 leaving the antimony free ; or by roasting it, that is, heat- 
 ing it in combination with the air, thus converting the anti- 
 mony into the tetroxide Sb 2 O 4 , and the sulphur into the 
 dioxide SO 2 , and then treating the oxide of antimony with 
 reducing agents, as, for example, carbon : 
 
 Sb 2 4 + 4C = 2Sb + 4CO. 
 
 Properties. Antimony is hard and brittle ; has a silver- 
 white color ; and a high metallic lustre. It can be dis- 
 tilled at white heat. At ordinary temperature it is not 
 changed by contact with the air. When heated to a suffi- 
 ciently high temperature in the air it takes fire and burns, 
 forming the white oxide Sb 2 O 3 . It combines directly with 
 chlorine, forming the chloride SbCl B . Nitric acid oxidizes 
 it either to antimony oxide, Sb 2 O 3 , or antimonic acid, 
 
APPLICATIONS OF ANTIMONY STIBINE. 309 
 
 H 3 SbO 4 . Aqua regia dissolves it. Hot concentrated sul- 
 phuric acid dissolves it, forming antimony sulphate, and 
 sulphur dioxide escapes. This action is similar to that 
 which takes place when sulphuric acid acts upon copper, 
 It is probable that the formation of the sulphur dioxide 
 is due to the action of the hydrogen liberated from the 
 sulphuric acid by the antimony in forming antimony sul- 
 phate : 
 
 2Sb + 3H 2 S0 4 = Sb 2 (S0 4 ) 3 + 3H 2 ; 
 3H 2 SO 4 + 3H 2 = 3SO 2 + 6H 2 O. 
 
 This power to replace the hydrogen of some acids dis- 
 tinguishes antimony from arsenic and phosphorus, while 
 its power to form acids corresponding to those of phos- 
 phorus and arsenic shows its analogy to these elements. 
 
 Applications of Antimony. Antimony is used as a 
 constituent of several alloys which are somewhat in- 
 definite compounds which metallic elements form with 
 one another. Among the alloys of antimony are type- 
 metal, from which type is made, and britannia metal. 
 The former consists of lead and antimony, and the latter 
 of tin and antimony. There are a number of alloys which 
 contain antimony which will be referred to under the 
 other constituents. 
 
 Stibine, SbH 3 . This analogue of ammonia, phosphine, 
 and arsine is more like arsine than it is like the others. 
 It is made by the same methods as those used in making 
 arsine, i.e., by treating an alloy of zinc and antimony 
 with sulphuric acid, or by reducing oxides of antimony 
 by means of nascent hydrogen. The latter method gives 
 a gas which contains a large percentage of hydrogen, but 
 for most purposes this is not objectionable. It is only 
 necessary to introduce into a flask containing zinc and 
 dilute sulphuric acid a little of a solution of some oxy- 
 gen compound of antimony, when the reduction is at once 
 effected, and the escaping hydrogen contains stibine. 
 
 Stibine is a colorless, inodorous gas, which burns with 
 a greenish-white flame. In general, it conducts itself 
 much like arsine. It is unstable and breaks down when 
 the tube through which it is passing is heated to red- 
 
310 INORGANIC CHEMISTRY. 
 
 ness. It then leaves a deposit which looks like that 
 formed in the case of arsine. When a cold object, as a 
 piece of porcelain, is held for a moment in a flame of 
 stibine a dark deposit is formed which resembles that 
 formed with arsine. 
 
 Methods of Distinguishing between Arsenic and Anti- 
 mony. As arsenic is frequently used in cases of poison- 
 ing the question of deciding whether it is present in a 
 given liquid or mixture is of great importance. One of 
 the chief difficulties encountered is the similarity of 
 the two elements arsenic and antimony. The method 
 commonly employed in examining a substance for arsenic 
 is known as Marsh's test. This consists in getting the 
 substance in solution, and then pouring some of the 
 liquid into a vessel containing pure zinc and pure dilute 
 sulphuric acid. If arsenic is present in the solution it 
 will, under these circumstances, be converted into arsine, 
 the presence of which can be recognized by heating the 
 tube through which the gas is passing, and by holding a 
 piece of porcelain in the flame. If deposits are not 
 formed in the tube or on the porcelain, arsenic is not 
 present ; but if deposits are formed, the only conclusion 
 that can be drawn is that either arsenic or antimony is 
 present, or possibly both may be present. For the pur- 
 pose of distinguishing between the two elements, advan- 
 tage is taken of the following differences between the 
 spots : The antimony spots are darker than those formed 
 by arsenic, and they have a smoky appearance, while 
 those of arsenic have not ; further, the arsenic deposits 
 are quite volatile, and can therefore be driven before the 
 flame in the tube or upon the porcelain, while those of 
 antimony are not volatile ; again, the deposits of arsenic 
 are easily soluble in a solution of sodium hypochlorite 
 or hypobromite, while the antimony deposits are in- 
 soluble in these solutions. There are other differences, 
 but those mentioned will suffice to enable a careful 
 worker and observer to distinguish between the two with- 
 out any possibility of doubt. Another difficulty always 
 encountered in examining for arsenic is the fact that the 
 sulphuric acid, the zinc, and the glass of which the ves- 
 
BISMUTH. 311 
 
 sels are made may contain arsenic. It is quite possible 
 to overcome all the difficulties and to decide positively 
 whether arsenic is present or not. If it is found that on 
 heating the tube through which the hydrogen is passing 
 no deposit is formed, even after continued heating, and 
 that the hydrogen flame gives no deposit upon a piece 
 of porcelain introduced into it, then it is safe to proceed 
 with the examination of the suspected liquid. If the 
 substance which is to be examined for arsenic has to be 
 treated with chemical compounds in order to prepare it 
 for analysis, every compound used in this part of the 
 process must be separately examined for arsenic. 
 
 BISMUTH, Bi (At. Wt. 206.54). 
 
 Occurrence, etc. Bismuth is not abundant nor widely 
 distributed in nature. It occurs for the most part native 
 in veins of granite and clay slate. Among the compounds 
 of bismuth found in nature are the oxide Bi 2 O 3 and the 
 corresponding sulphide Bi 2 S 3 . 
 
 The ores are roasted and then treated with appropriate 
 reducing agents. In different places different methods 
 of extraction are employed. As the chief applications of 
 bismuth are for pharmaceutical purposes, it is necessary 
 that the element should be specially pure ; above all, 
 that it should not be contaminated with arsenic. In 
 order to remove the last traces of this element the pow- 
 dered bismuth is generally melted with saltpeter. 
 
 Bismuth is a hard, brittle, reddish- white substance 
 with a metallic lustre. It looks very much like antimony, 
 but is distinguished from it by its reddish tint. At or- 
 dinary temperatures it remains unchanged in the air. 
 When heated to red heat it burns with a bluish flame, 
 forming the yellow oxide Bi 2 O 3 . 
 
 Hydrochloric acid scarcely acts upon it ; concentrated 
 sulphuric acid forms bismuth sulphate, Bi 2 (SO 4 ) 3 , in which 
 the bismuth evidently plays the part of a base-forming 
 element ; nitric acid gives bismuth nitrate, Bi(NO 3 ) 3 , which 
 is partly decomposed by water, forming so-called basic 
 nitrates which are difficultly soluble in water. These 
 salts will be taken up in the next chapter. 
 
312 INORGANIC CHEMISTRY. 
 
 Some bismuth is used in the preparation of alloys 
 which are easily fusible, as, for example, Newton's metal, 
 which contains bismuth, lead, and tin ; Hose's metal, 
 which consists of the same constituents in slightly dif- 
 ferent proportions ; and Wood's metal, which consists 
 of bismuth, lead, tin, and cadmium. 
 
 Bismuth does not combine with hydrogen. 
 
 Compounds of the Members of the Phosphorus Group 
 with the Members of the Chlorine Group. In the intro- 
 duction to this chapter it was stated that the elements of 
 the phosphorus group combine with chlorine in two 
 proportions, forming compounds of the general formulas 
 MC1 3 and MC1 5 . Arsenic, however, forms only one com- 
 pound with chlorine, AsCl 3 , while bismuth forms one of 
 the formula BiCl 3 , and another, Bi 2 Cl 4 . The compounds 
 of phosphorus and chlorine are the best known, and a 
 brief study of these will give a fair idea of the methods 
 of preparation and the conduct of the analogous com- 
 pounds of the other members of the group. 
 
 Phosphorus Trichloride, PC1 3 , is made by conducting 
 dry chlorine gas upon phosphorus in a retort connected 
 with a receiver. Action takes place at once with evo- 
 lution of heat, and the trichloride distils over and is 
 condensed as a liquid into the receiver. It is purified 
 by distillation on a water-bath. It is a clear, color- 
 less liquid, which boils at 74. In contact with air it 
 fumes in consequence of the action of the water vapor 
 which decomposes it. It has a disagreeable odor of its 
 own mixed with that of hydrochloric acid. Its most 
 characteristic decomposition is that which it undergoes 
 with water, which is of the same kind as that which the 
 chlorides of sulphur, selenium, and tellurium undergo 
 with water. The general tendency of the chlorides of 
 the acid-forming elements is to undergo decomposition 
 with water in such a way that the corresponding hydroxyl 
 compound is formed, together with hydrochloric acid. 
 This is shown in the case of tellurium tetrachloride, which 
 with water forms normal tellurious acid, Te(OH) 4 , and 
 hydrochloric acid : 
 
PHOSPHORUS TRICHLORIDE. 313 
 
 HOH f OH 
 
 HOH [ OH 
 
 In the case of sulphur tetrachloride the hydroxyl de- 
 rivative, if formed, breaks down into water and sulphur 
 dioxide. When phosphorus trichloride is treated with 
 water the decomposition is probably represented by the 
 equation 
 
 HOH ( OH 
 
 PCI 4- HOH = P^ OH + 3HC1. 
 HOH ( OH 
 
 From some experiments it appears possible that this 
 form of compound is unstable, and that, owing to the 
 marked tendency of phosphorus to act as a quinquivalent 
 element, the constituents arrange themselves differently, 
 
 ( H 
 
 as represented in the formula O=P-< OH. Thisques- 
 
 (OH 
 
 tion will be referred to under the head of Phosphorous 
 Acid. 
 
 The trichloride shows a strong tendency to take up 
 chlorine, bromine, iodine, oxygen, and sulphur, and thus 
 to become saturated as a quinquivalent element. With 
 chlorine it forms the pentachloride, PC1 5 , with oxygen 
 the oxychloride, POC1 3 , and with sulphur the sulphochlo- 
 ride, PSC1 3 . It does not, however, readily take up free 
 oxygen or free sulphur directly, but will take up these 
 elements from compounds in which they are not firmly 
 held. Thus, when the trichloride is brought together 
 with sulphur trioxide this reaction takes place : 
 
 S0 3 + PC1 3 = POC1 3 + SO 2 ; 
 
 and when it is brought together with a polysulphide, 
 as Na 2 S & , it takes up a part of the sulphur and forms 
 the sulphochloride, PSC1 3 . So, further, it is converted 
 into the oxychloride when treated with ozone. These 
 reactions show the marked tendency which the trichlo- 
 
314 INORGANIC CHEMISTRY. 
 
 ride lias to pass over into compounds of quinquivalent 
 phosphorus a tendency which is characteristic of phos- 
 phorus compounds in general. 
 
 Phosphorus Pentachloride, PC1 5 , is formed by treating 
 phosphorus or the trichloride with dry chlorine. It 
 is best prepared by passing chlorine through a wide 
 tube upon the surface of the trichloride, contained 
 in a vessel, which is kept cool. Gradually the 
 liquid becomes thicker and thicker, and finally, if well 
 stirred, it becomes solid. It is a white solid, but it 
 generally has a slightly yellowish or greenish color in 
 consequence of a slight decomposition into the tri- 
 chloride and free chlorine. It sublimes below 100 
 without melting. When heated to boiling it under- 
 goes partial decomposition into chlorine and the trichlo- 
 ride, and this decomposition is complete at about 300. 
 As the temperature is raised from the apparent boiling 
 point to the point at which the decomposition is com- 
 plete, the color of the vapor is seen to grow darker 
 in consequence of the increased quantity of free chlo- 
 rine present. The decomposition is gradual, and, for 
 any given temperature, the amount of decomposition 
 is constant. This kind of decomposition, which is known 
 as dissociation, has been studied very carefully, and is 
 found to be capable of explanation by the aid of the 
 kinetic theory of gases. In a later chapter this subject 
 will be treated, and a number of other examples will be 
 given. Owing to this decomposition under the influence 
 of heat the specific gravity of the vapor of phosphorus 
 pentachloride is not what it should be, if the formula is 
 PC1 5 . On the other hand, the specific gravity of the 
 vapor of the trichloride leads to the formula PC1 3 , and 
 that of the oxychloride to the formula POC1 3 . The ap- 
 parent anomaly presented by the pentachloride is easily 
 understood. When a molecule of the compound is con- 
 verted into vapor, or is heated to a sufficiently high 
 temperature, it is broken down in accordance with this 
 equation : 
 
 PC1 6 = PC1 3 + C1 2 . 
 
PHOSPHORUS PENTACHLOEIDE. 315 
 
 From the one molecule, therefore, two gaseous mole- 
 cules are 'obtained. Consequently the vapor formed oc- 
 cupies twice as much space as it would if there were no 
 decomposition. It follows that the specific gravity of 
 the vapor must be only half what it would be if there 
 were no decomposition. When the compound is con- 
 verted into vapor in an atmosphere of phosphorus tri- 
 chloride, the decomposition referred to does not take 
 place, and, under these circumstances, the specific gravity 
 is found to be in accordance with Avogadro's law, and 
 with the formula PC1 5 . This case is a particularly in- 
 teresting one, as it has played an important part in 
 the discussions in regard to the validity of Avogadro's 
 law. 
 
 The conduct of phosphorus pentachloride towards 
 water is in general like that of the other chlorides of 
 acid-forming elements. But, owing probably to a second- 
 ary action, the product is not the corresponding hydroxyl 
 compound. It is probable that the first action of the 
 water is represented by the equation 
 
 TTTT/-V ( OH 
 
 PC1 *~ = P OH 2H01. 
 
 But this product, if formed, breaks down at once into 
 phosphorus oxychloride and water, and the water thus 
 given off acts upon a further quantity of the penta- 
 chloride : 
 
 The formation of the oxychloride from the penta- 
 chloride by the action of water takes place very easily. 
 The oxychloride is then further acted upon by the water, 
 and each chlorine atom is replaced by hydroxyl : 
 
 HOH ( OH 
 
 OPC1 3 + HOH = OP^ OH + 3HC1. 
 HOH OH 
 
316 INORGANIC CHEMISTRY. 
 
 The final product is the acid H 3 PO 4 , or phosphoric 
 acid. 
 
 It will be seen that the effect of phosphorus penta- 
 chloride upon water is to replace the hydroxyl of the 
 water by chlorine. Thus, one molecule of the penta- 
 chloride and five molecules of water give one molecule 
 of phosphoric acid and five of hydrochloric acid : 
 
 HOH HC1 
 
 HOH HC1 
 
 PC1 5 + HOH = OP(OH) 3 + HC1 + H 2 O. 
 HOH HC1 
 
 HOH HC1 
 
 In the reaction, the hydroxyl of the water and the 
 chlorine of the chloride exchange places. Similarly, 
 when any compound which contains hydroxyl is treated 
 with phosphorus pentachloride the same reaction takes 
 place, the hydroxyl being replaced by chlorine. There- 
 fore phosphorus pentachloride may be used as a reagent 
 for testing for the hydroxyl condition in compounds. If 
 a compound which contains hydrogen and oxygen is 
 treated with the pentachloride, and an atom of hydrogen 
 and one of oxygen is replaced by an atom of chlorine, 
 the conclusion is drawn that the compound contains 
 hydroxyl. This, of course, amounts to saying that the 
 compound resembles water in its reaction with the penta- 
 chloride, and this is most easily explained by the as- 
 sumption that the same condition exists in both. It 
 should be borne in mind, further, that, in general, any 
 compound of chlorine with an acid-forming element 
 which undergoes decomposition with water might be used 
 for the same purpose. The action of the pentachloride 
 upon a hydroxyl compound is well illustrated in the case 
 of sulphuric acid : 
 
 SO '<OH + PC1 * = POC1 * + SO <OH + HCL 
 
 The action of the oxychloride would take place as rep- 
 resented thus : 
 
ARSENIC TRICHLORIDE. 31 
 
 POC1 = 
 
 Phosphorus forms a tribromide, PBr 3 , and a penta- 
 bromide, PBr B , similar to the chlorides. With iodine it 
 forms the compounds P,I 4 and PI 3 , and with fluorine a 
 trifluoride, PF 3 , and a pentafluoride, PF B . 
 
 Arsenic Trichloride, AsCl 3 , is the only compound which 
 arsenic forms with chlorine. It is easily made by passing 
 chlorine into a retort containing powdered arsenic. It 
 is also formed by the action of phosphorus pentachloride 
 on arsenic trioxide, in which case the oxygen of the tri- 
 oxide is simply replaced by chlorine : 
 
 As 2 3 + 3PC1 5 = 3POC1 3 + 2AsCl 3 . 
 
 Further, it is obtained when dry hydrochloric acid gas 
 is passed over arsenic trioxide, and it is present in solu- 
 tion when the trioxide is dissolved in strong hydrochloric 
 acid. 
 
 Arsenic trichloride is a liquid which boils at about 
 134. It is extremely poisonous. If a hydrochloric acid 
 solution of the trichloride is boiled, it passes over when 
 the temperature rises above 100. When substances 
 containing arsenic are heated in a retort with sulphuric 
 acid and sodium chloride, arsenic chloride is formed and 
 passes over into the receiver. 
 
 When arsenic trichloride is treated with a little water 
 two of the chlorine atoms are replaced by hydroxyl, thus : 
 
 2HC1. 
 
 If treated with much water it yields arsenious acid ;_ . . 
 AsCl 3 + 3H 2 O = As(OH) 3 + 3HCL 
 
 We have here a curious illustration of the effect of the 
 relative quantities of the substances which take part in 
 a reaction upon the nature of the reaction. When 
 
318 INORGANIC CHEMISTRY. 
 
 arsenic trioxide is treated with an excess of hydrochloric 
 acid it is converted into the chloride : 
 
 As 2 O 3 + 6HC1 = 2AsCl 3 + 3H 2 O. 
 
 If, on the other hand, the chloride is treated with a 
 large excess of water it is completely broken down, yield- 
 ing the trihydroxyl derivative, or arsenious acid. There 
 are many other illustrations of this kind met with among 
 chemical reactions, and the subject of the influence of 
 mass in determining the character of a reaction has been 
 studied with much care. This subject will receive at- 
 tention in a later chapter under the head of Mass Action. 
 Compounds of Antimony and Chlorine. Antimony forms 
 two compounds with chlorine, analogous in compo- 
 sition to the two chlorides of phosphorus. These 
 are antimony trichloride, SbCl 8 , and antimony penta- 
 chloride, SbCl 6 . The trichloride is formed by direct 
 treatment of the element with chlorine. It is also 
 formed by dissolving antimony in hydrochloric acid 
 with addition of nitric acid, and when this solution 
 is distilled the chloride of antimony passes over. At the 
 ordinary temperatures it is a solid, colorless, crystal- 
 line, soft substance, which, on account of its consistency, 
 has received the common name " butter of antimony " 
 (Butyrum Antimonii). When treated with water it forms 
 an oxychloride, the composition of which varies according 
 to circumstances, but it generally approximates to that 
 represented by the formula Sb 4 O 6 Cl 2 . If treated with 
 cold water the decomposition is simple, the product 
 being antimony oxychloride, SbOCl : 
 
 SbCl 3 + H 2 = SbOCl + 2HC1. 
 
 The oxychlorides formed by treating antimony tri- 
 chloride with hot water, which, as already stated, have 
 a composition approximating that represented by the 
 formula Sb 4 O 6 Cl 2 , are known as " Powder of Algaroth" 
 The relation of this compound to the simple oxychloride, 
 SbOCl, is indicated by the equation 
 
DOUBLE SALTS. 319 
 
 SSbOCl = Sb 4 6 Cl a + SbCl 3 . 
 
 Liquor Stibii Muriatici is a solution of antimony tri- 
 chloride prepared by dissolving antimony sulphide in 
 concentrated hydrochloric acid : 
 
 Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. 
 
 Antimony chloride has a caustic action and is used in 
 medicine. It is also used for the purpose of burnishing 
 iron ware, as gun-barrels. When treated with the chloride 
 they acquire a brownish, bronze-like color. 
 
 PentachLoride of Antimony is in general like the pen- 
 tachloride of phosphorus. It gives up its chlorine, 
 however, somewhat more readily. When treated with 
 water it yields first an oxychloride, SbO 2 Cl, and this is 
 further acted upon and converted into antimonic acid, 
 H 3 Sb0 4 . 
 
 Bismuth and Chlorine. Bismuth differs from the other 
 members of the phosphorus group in its conduct towards 
 chlorine. It forms the chloride, Bi 2 Cl 4 , of which 
 there is no analogue among the compounds of the other 
 elements of the group with chlorine. The compound 
 may, however, be regarded as analogous to the hydrogen 
 compounds hydrazine, N 2 H 4 , and liquid phosphine, P 2 H 4 . 
 Bismuth dichloride, Bi 2 Cl 4 , is formed by reduction of the 
 trichloride, BiCl 3 , when the latter is treated with hydro- 
 gen at a temperature of about 300. Bismuth trichloride, 
 BiCl 3 , is formed by treating bismuth with chlorine, and 
 by dissolving bismuth oxide, Bi Q O 3 , in hydrochloric acid. 
 From this solution a crystallized compound of the 
 formula BiCl 3 + H 2 O is deposited, and it is impossible 
 to drive all the water off from this compound without 
 causing decomposition. Treated with water the chloride 
 is decomposed, forming the oxychloride, BiOCl : 
 
 BiCl 3 + H 2 O = BiOCl + 2HC1. 
 
 Double Salts. When the chlorides, bromides, iodides, 
 and fluorides of the members of the phosphorus group are 
 treated with the corresponding salts of potassium, sodium, 
 
320 INORGANIC CHEMISTRY. 
 
 and some other strongly marked base-forming elements, 
 compounds known as double salts are formed. Thus, 
 when antimony trichloride is treated with a strong water 
 solution of potassium chloride, a salt of the composition 
 SbCl 3 .3KCl is formed. This compound appears to be 
 analogous to the oxygen compound K 3 SbO 3 , or potassium 
 antimonite, differing from it in containing chlorine in 
 place of the oxygen. While the latter is derived from 
 an acid of the composition H 3 SbO 3 , the former appears 
 to be derived from the corresponding chlorine acid, 
 H 3 SbCl 6 . Similar compounds of sulphur and tellurium 
 are known, and one was referred to on page 205. This 
 is the potassium compound K 2 TeBr 6 or TeBr 4 .2KBr, 
 which, as explained there, is analogous to potassium 
 tellurite, K 3 Te0 3 . 
 
CHAPTER XVIII. 
 
 COMPOUNDS OF THE ELEMENTS OF THE PHOSPHORUS 
 GROUP WITH OXYGEN AND WITH OXYGEN AND HY- 
 DROGEN. 
 
 Introduction. The product of the direct action of 
 oxygen upon phosphorus is the pentoxide PA- Ar- 
 senic, antimony, and bismuth, however, form the tri- 
 oxides As 2 O 3 , Sb 2 O 3 , and Bi 2 O 3 . It is possible to obtain 
 a compound of arsenic and oxygen of the formula As 2 O 6 , 
 one of antimony, Sb 2 O 4 , and another, Sb A, and, finally, 
 two oxides of bismuth, Bi 2 O 2 and Bi 2 O 5 . Phosphorus, 
 further, forms the oxides P 4 O, P,O 8 , and P 3 O 4 . The 
 table below contains the formulas of the above-men- 
 tioned compounds systematically arranged : 
 
 P 4 
 
 PA AsA Sb,0, 
 
 PA Sb,o 4 
 
 PA AsA SbA 
 
 The final products of the oxidation of the elements of 
 this group, if water is present, are phosphoric, arsenic, 
 antimonic, and bismuthic acids. All of these are well- 
 marked acids except the last. They can all be regarded 
 as derived from the normal acids of the general formula 
 M(OH) 5 by loss of one or two molecules of water. The 
 common forms of phosphoric, arsenic, and antimonic 
 acids are those which are formed from the normal acids 
 by loss of one molecule of water : 
 
 P(OH) 6 P {(OH) S + H 0; 
 
 Normal phosphoric acid Orthophosphoric acid 
 
 As(OH) 5 As {( 
 
 (OH) s > 
 
 Normal arsenic acid Orthoarsenic acid 
 
 (321) 
 
322 INORQANIG CHEMISTRY. 
 
 Sb(OH) s Sb {(OH) s 
 
 Normal antimonic acid Orthoantimonic acid 
 
 Bismuthic acid appears, however, to be formed from 
 the normal acid by loss of two molecules of water, just 
 as the so-called metaphosphoric, metarsenic, and metanti- 
 monic acids are : 
 
 Bi(OH) 5 = BiJQH + 2H 2 0; 
 
 Normal bismuthic acid Bismuthic acid 
 
 P(OH) 5 P|gk + 2H,0; 
 
 Metaphosphoric acid 
 
 As(OH) 5 = As -j Q|J -f- 2H 2 O ; 
 
 Metarsenic acid 
 
 Sb(OH) 5 = s M8b + 2H '- 
 
 Metantimonic acid 
 
 From the ordinary or ortho acids, and from the meta 
 acids, more complex forms can be derived by loss of 
 different quantities of water. The most common form 
 besides those mentioned is that seen in the so-called 
 pyro acids, of which pyrophosphoric acid is the best 
 known example. It is formed from the ortho acid by 
 loss of one molecule of water from two molecules of 
 the acid, just as pyrosulphuric or disulphuric acid is 
 formed from two molecules of ordinary sulphuric acid 
 by loss of one molecule of water. The formation of 
 pyrophosphoric acid from orthophosphoric acid takes 
 place according to the equation 
 
 OH 
 OH 
 
 OH 
 pi OH p " 
 
 18 H A ; 
 
COMPOUNDS OF THE PHOSPHORUS GROUP. 323 
 
 or 
 
 2H.PO, = HJP.O, + H,0. 
 
 Orthophosphoric acid Pyrophosphoric acid 
 
 Pyroarsenic and pyroantimonic acids bear the same 
 relations to the ortho acids that pyrophosphoric acid 
 bears to orthophosphoric acid. 
 
 By partial oxidation of phosphorus in presence of 
 water, phosphorous acid, H 3 PO 3 , is formed. The same 
 acid is formed by the action of phosphorus trichloride 
 on water. According to the latter method of formation 
 we should expect to find that this acid is normal phos- 
 phorous acid, P(OH) 3 . As already stated, however, it 
 appears probable that the acid has the constitution 
 
 /H 
 O=P -OH. The acids of arsenic and antimony of 
 
 \OH 
 
 similar composition seem to be the normal acids As(OH) 8 
 and ' Sb(OH) 3 . The hydroxyl derivative of bismuth 
 corresponding to these acids has no acid properties, 
 but on the contrary is basic. Hypophosphorous acid 
 has the composition H 3 PO 2 . It is monobasic, and it 
 appears therefore that it contains but one hydroxyl, as 
 represented in the formula H 2 OP(OH). It is possible 
 that the relations between phosphoric, phosphorous, 
 and hypophosphorous acids should be represented by 
 the formulas 
 
 (OH ( H ( H 
 
 OP^OH, OP^OH, OP^H . 
 (OH (OH (OH 
 
 Phosphoric acid Phosphorous acid Hypophosphorous acid 
 
 The fundamental compound, then, from which these may 
 be regarded as derived is the unknown oxyphosphine 
 OPH 3 . By oxidation we should expect phosphine 
 to yield in successive stages the three products above 
 named : 
 
 (H (H (H (H (OH 
 
 P^H,OP^H, OP^H , OP^OH,OP^OH. 
 
 (H (H (OH (OH (OH 
 
 Unknown Hypophosphorous Phosphorous Phosphoric 
 acid acid acid 
 
3M INORGANIC CHEMISTRY. 
 
 The oxidation of hydrogen sulphide takes place sim* 
 ilarly, as has been shown : 
 
 , , 
 
 Unknown Sulphurous acid Sulphuric acid 
 
 With oxygen and chlorine the elements of the phos- 
 phorus group form a number of compounds known as 
 oxychlorides. Towards chlorine as well as towards 
 oxygen all these elements except bismuth are quin- 
 quivalent. A part or all of the oxygen of the oxygen 
 compounds can be replaced by chlorine. Starting with 
 the chlorine compound on the one hand, oxychlorides. 
 can be obtained from it, until all the chlorine is replaced 
 by oxygen, and the limit is reached in the oxide. So 
 also the chlorine can be replaced by hydroxyl and the 
 acids thus obtained. 
 
 (1) PC1 5 gives POC1 3 and P 2 O 6 as final product ; 
 
 (2) PC1 5 gives POC1 3 and with water PO(OH) 3 . 
 
 (3) PC1 3 gives as final product P 2 O 3 ; 
 
 (4) PC1 3 gives with water P(OH) 3 . 
 
 Intermediate products are supposable, but not known, 
 as, for example : 
 
 Cl (Cl (Cl (OH 
 
 Cl P^ Cl , PK OH, Pi OH. 
 Cl OH OH OH 
 
 ( Cl 
 A compound of arsenic of the formula As-j /QTT\ is 
 
 known, however, and this plainly corresponds to one of 
 these intermediate products. 
 
 With sulphur phosphorus apparently forms a large 
 number of compounds. Among them are two which 
 have the formulas P 2 S S and P 2 S 5 . These plainly are 
 analogous to the two oxides of phosphorus, P 2 O 3 and 
 P 2 O B . When treated with water these sulphur com- 
 pounds like the corresponding chlorine compounds 
 yield the oxygen acids. Thus the trisulphide undergoes 
 
COMPOUNDS OF THE PHOSPHORUS GROUP. 325 
 
 decomposition with water according to the following 
 equation : 
 
 P 2 S 3 + 6H 2 O = 2H 3 PO 3 + 3H a S ; 
 
 and the pentasulphide is converted by water into phos* 
 phoric acid : 
 
 P a S 6 + 8H a O = 2H 3 P0 4 + 5H a S. 
 
 Arsenic forms with sulphur several compounds, the 
 principal of which are the disulphide, As 2 S 2 , the trisul- 
 phide, As 2 S 3 , and the pentasulphide, As 2 S 6 . The principal 
 sulphides of antimony are those of the formulas Sb 2 S 9 
 and Sb 2 S 5 , and of bismuth those of the formulas Bi 2 S a 
 and Bi 2 S 3 . In general, therefore, the sulphur compounds 
 are analogous in composition to the oxygen compounds, 
 while the number of sulphur compounds of these ele- 
 ments is larger than that of the oxygen compounds. 
 The formulas of the principal sulphur compounds of this 
 group are given systematically arranged in the table 
 below : 
 
 As 2 S a Bi 2 S a 
 
 P 2 S 3 As 2 S 3 Sb a S 8 Bi a S 3 
 P a S 5 As 2 S, Sb a S 6 
 
 Further, there are sulphur acids which are to be re- 
 garded as the oxygen acids, a part or all of whose 
 oxygen is replaced by sulphur. Thus, in the case of 
 arsenic the following possibilities suggest themselves, 
 starting with arsenious acid : 
 
 (SH 
 
 -{ SH ; 
 
 (SH 
 
 and starting with arsenic acid, the following possibilities 
 suggest themselves : 
 
 OH (OH (OH (OH ( 
 
 OH , SAs^ OH , SAs^ OH , SAs^ SH, SAsJ 
 
 OH (OH (SH (SH (SH 
 
326 INORGANIC CHEMISTRY. 
 
 While none of these compounds is known, many com- 
 pounds are known which are to be regarded as salts of 
 one or another of these acids. Thus salts of the general 
 formulas M 3 AsS 3 and M 3 AsS 4 are well known, as are also 
 salts of the general formula MAsS 2 , which are derived 
 
 ( ^ 
 from the acid As -j QTT , corresponding to the oxygen 
 
 compound As j QTT , which in turn is derived from arsen- 
 ious acid by loss of one molecule of water : 
 
 So, too, we have : 
 
 SH , 
 
 +H 2 S. 
 
 Similar compounds of antimony are also well known. 
 The possibility of making analogous compounds contain- 
 ing selenium and tellurium will suggest itself. 
 
 Phosphoric Acid, Orthophosphoric Acid, H 3 PO 4 . The 
 compound to which the name phosphoric acid is gener- 
 ally applied, and from which the best known phosphates 
 are derived, is that which has the formula H 3 PO 4 . To 
 distinguish it from the other varieties it is called ortho- 
 phosphoric acid. As has been stated, this is the final 
 product of oxidation of phosphorus in the presence of 
 water. Thus, when phosphorus is boiled with nitric acid 
 it is converted into orthophosphoric acid ; and also when 
 phosphorus is burned in the air, and the product dis- 
 solved in water, phosphoric acid is formed. In this case 
 the first product of the oxidation is the pentoxide P 2 O 5 , 
 also known as phosphoric anhydride, and when this is 
 treated with water it is converted into phosphoric acid : 
 
 P a 6 + 3H 3 O = 2H 3 PO 4 . 
 The occurrence of phosphoric acid in nature has already 
 
PHOSPHORIC ACID. 327 
 
 been referred to in connection with the occurrence of 
 phosphorus, which is found in nature almost exclusively 
 in the form of phosphates, principally as calcium phos- 
 phate, Ca 3 (PO 4 ) 2 , in phosphorite, apatite, and the ashes 
 of bones. It is formed when either phosphorus penta- 
 chloride or the oxychloride is decomposed by water : 
 
 PC1 5 + 4H 2 = PO(OH) 3 + 5HC1 ; 
 POC1 3 + 3H 2 O = PO(OH) 3 + 3HC1 ; 
 
 and from the analogous bromine and iodine compounds 
 in the same way. In order to prepare the acid two ways 
 suggest themselves : (1) by oxidizing phosphorus with 
 nitric acid ; and (2) by extracting it from one of the natu- 
 ral phosphates, as phosphorite or bone-ash. The first 
 of these methods is better adapted to the preparation of 
 pure phosphoric acid, such as is needed for medicinal 
 purposes ; the latter is used where absolute purity of the 
 product is not required. It should be said, however, 
 that the acid obtained by oxidation of phosphorus is not 
 pure, as commercial phosphorus almost always contains 
 arsenic and small quantities of other impurities. The 
 arsenic can easily be removed by passing hydrogen sul- 
 phide through the solution after the nitric acid has been 
 evaporated. If the solution is then filtered and evapo- 
 rated to dryness, the orthophosphoric acid is transformed 
 into pyrophosphoric or metaphosphoric acid according 
 to the temperature : 
 
 H 3 P0 4 = HP0 3 +H 2 0. 
 
 The preparation of phosphoric acid from a phosphate 
 is not a simple matter. If the acid were volatile or in 
 soluble there would be no difficulty in separating it. In 
 the former case it would only be necessary to proceed as 
 in preparing hydrochloric and nitric acids. By adding 
 an acid which is not volatile except at a high temperature, 
 such, for example, as sulphuric acid, and heating, the 
 non-volatile acid replaces the volatile. On the other 
 
328 INORGANIC CHEMISTRY. 
 
 hand, if phosphoric acid were insoluble in water, it could 
 be separated by adding a soluble acid to one of its soluble 
 salts. When, for example, nitric acid is added to a solu- 
 tion of potassium tellurite, K a TeO 3 , tellurious acid, being 
 insoluble, is thrown down : 
 
 K 2 Te0 3 + 2HN0 3 = 2KNO 3 + H a TeO 3 . 
 
 But phosphoric acid is not volatile and is soluble, so that 
 plainly neither of these methods can be used. By treat- 
 ing the calcium salt with sulphuric acid the calcium can 
 be completely separated in the form of calcium sul- 
 phate, which is difficultly soluble in water and insoluble 
 in alcohol. The ideal reaction to be accomplished is that 
 represented in the following equation : 
 
 Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 P0 4 . 
 
 But when sulphuric acid is added to calcium phosphate, 
 only a part of the calcium is thrown down as sulphate, 
 the rest remaining in the form of primary calcium phos- 
 phate : 
 
 Ca/PO,), + 2H,SO, = 2CaSO. + CaH,(PO 4 ) s . 
 
 The phosphate thus formed is soluble in water, and the 
 calcium is not easily precipitated from it. By evapora- 
 tion and addition of sufficient sulphuric acid and alcohol 
 the precipitation can be effected, and a solution of phos- 
 phoric acid thus obtained. This acid is not pure, as 
 there are substances in bone-ash which are not removed 
 by the method described. 
 
 Phosphoric acid for medicinal purposes is often made 
 by dissolving the pentoxide in water. 
 
 Properties. When evaporated to the proper consis- 
 tency the acid forms a thick syrup which slowly solidifies 
 in the form of large crystals. The crystals are deliques- 
 cent. When heated to a sufficiently high temperature 
 the acid loses water, as already explained, and yields, 
 first, pyrophosphoric, and then metaphosphoric acid. 
 It is a tribasic acid, capable of yielding three classes of 
 
PHOSPHATES. 329 
 
 (OH (OH 
 
 salts of the general formulas OP-< OH, OP-< OM,and 
 
 ( OM ( OM 
 
 (OM 
 
 OP -< OM , which are known respectively as the primary, 
 
 (OM 
 
 secondary, and tertiary phosphates. The primary and 
 secondary phosphates are also known as acid phosphates, 
 and the tertiary salts as neutral or normal phosphates. In 
 these salts it is not necessary that all the hydrogen should 
 be replaced by the same metal. There are salts in which 
 two or three metals take the place of the hydrogen atoms. 
 A phosphate much used in the laboratory, for example, 
 is one in which one hydrogen atom of phosphoric acid is 
 replaced by a sodium atom, and another by the ammoni- 
 
 (OH 
 um group, NH,. This salt has the formula OP-< ONa , 
 
 (ONH. 
 
 and is called ammonium sodium phosphate. Another 
 phosphate commonly met with is ammonium magnesium 
 
 (ONH 4 
 phosphate, OP-< O jyr , which is derived from the acid 
 
 by replacement of two hydrogen atoms in the molecule 
 by one bivalent magnesium atom, and one by the am- 
 monium group. The changes which these three classes 
 of phosphates undergo when heated are of special inter- 
 est. The tertiary phosphates are stable. The primary 
 and secondary phosphates give up all their hydrogen, 
 which passes off in the form of water. Thus, primary 
 sodium phosphate, H 2 NaPO 4 , loses one molecule of water 
 from each molecule of the salt, and is converted into the 
 metaphosphate, NaPO 3 : 
 
 (OH 
 
 OP^ OH = O 2 P(ONa) + H,0. 
 (ONa 
 
 In general, the primary phosphates are converted into meta- 
 phosphates by heat. 
 
 When a secondary phosphate is heated the product is 
 
330 INORGANIC CHEMISTRY. 
 
 a pyrophosphate, as when secondary sodium phosphate 
 is heated to a sufficiently high temperature it is converted 
 into sodium pyrophosphate : 
 
 w >- 
 
 j ONa OP ONa 
 
 -< ui>a J Q^_ 
 
 (ONa 
 
 In general, a secondary phosphate is converted into a py- 
 rophosphate by heat. 
 
 The above rules do not hold good for ammonium salts, 
 for these always undergo another kind of decomposition 
 when heated. When sodium ammonium phosphate is 
 heated, ammonia is first given off, thus : 
 
 HNa(NH 4 )P0 4 = H 2 NaPO 4 + NH 3 ; 
 
 and the primary salt formed breaks down according to 
 the above rule, forming the metaphosphate. So, also, 
 when ammonium magnesium phosphate is heated, the 
 first change consists in the giving off of ammonia, thus . 
 
 (NH,)MgPO, = HMgPO, + NH S ; 
 
 and the secondary magnesium phosphate thus formed 
 then breaks down, forming the pyrophosphate, Mg 2 P 2 O 7 : 
 
 The presence of phosphoric acid can be detected by 
 means of the following characteristic reactions : With 
 silver nitrate it gives a yellow precipitate of tertiary sil- 
 ver phosphate, Ag 3 PO 4 ; with a soluble magnesium salt 
 and ammonia it gives ammonium magnesium phosphate, 
 (NH 4 )MgPO 4 , which is insoluble in water ; with a solu- 
 tion of ammonium molybdate, (NH 4 ) 2 MoO 4 , which con- 
 tains nitric acid, it gives a complicated insoluble salt, 
 ammonium phospho-molybdate (which see). 
 
 Pyrophosphoric Acid, H 4 P 2 O 7 . When phosphoric acid 
 is heated to 200-300 until a specimen neutralized with 
 ammonia gives a pure white precipitate with silver nitrate. 
 it is completely transformed into pyrophosphoric acid by 
 
METAPHOSPHORIC ACID. 331 
 
 loss of water. The white precipitate referred to is the 
 silver salt of pyrophosphoric acid. The silver salt of 
 orthophosphoric acid is yellow. This difference in color 
 led, many years ago, to a careful investigation of the 
 change in composition which phosphoric acid undergoes 
 when heated, and to the recognition of the existence of 
 pyrophosphoric acid as distinct from orthophosphoric 
 acid ; and the study of the relations existing between 
 these acids and metaphosphoric acid has had a strong 
 influence in shaping the views of chemists in regard to 
 the relations between other similar acids. The views at 
 present held in regard to the relations between the com- 
 mon forms of oxygen acids and the so-called normal acids 
 or maximum hydroxides are simply an extension of the 
 ideas first introduced into chemistry in connection with 
 the three varieties of phosphoric acid. The different 
 varieties of periodic acid, and the modifications of sul- 
 phuric acid seen in the normal acid, S(OH) 6 , the ordinary 
 acid, SO 2 (OH) 2 , and the pyro-acid, H 2 S 2 O 7 , are examples 
 of the same kind of relations. 
 
 Pyrophosphates are formed, as we have seen, when 
 the secondary phosphates, like disodium phosphate, 
 HNa 2 PO 4 , are heated. 
 
 Metaphosphoric Acid, HPO 3 . This acid is formed by 
 dissolving phosphorus pentoxide, P 2 O 5 , in cold water : 
 
 P 2 5 + H 2 = 2HP0 3 . 
 
 It is also formed by heating phosphoric acid to 400 : 
 H 3 P0 4 = HP0 3 + H 2 O. 
 
 Further, the metaphosphates are formed by heating 
 the primary phosphates like primary sodium phosphate, 
 H. 2 NaPO 4 . The acid is a vitreous translucent mass, known 
 in the market as glacial phosphoric acid (Acidum phos- 
 pJioricum glaciale). It is the more common commercial 
 form of phosphoric acid. It is a monobasic acid, and in 
 composition is analogous to nitric and chloric acids : 
 
 HPO 3 , . . . c . Metaphosphoric acid. 
 
 HNO 3 , Nitric acid. 
 
 HC1O 3 , . . . . . Chloric acid. 
 
332 INORGANIC CHEMISTRY. 
 
 When boiled with water in which there is a little nitric 
 acid metaphosphoric acid is readily converted into ortho- 
 phosphoric acid : 
 
 HPO 3 + H 2 = H 3 P0 4 . 
 
 This transformation is effected also by simply allowing 
 the solution of the meta-acid in water to stand for a time, 
 and by boiling the solution. 
 
 When a metaphosphate, as, for example, sodium meta- 
 phosphate, NaPO 3 , is heated in contact with a metallic 
 oxide, it takes up the oxide as the free acid takes up 
 water, and phosphates are thus formed in which two or 
 more metals take the place of the hydrogen of the acid. 
 With a metallic oxide of the formula M 2 O it combines 
 to form a phosphate, M 2 NaPO 4 , thus : 
 
 NaPO 3 + M 2 = M 2 NaPO 4 
 
 a kind of action which is plainly analogous to the con- 
 version of metaphosphoric into orthophosphoric acid. 
 So, also, when an oxide of the formula MO is heated 
 with sodium metaphosphate a phosphate of the formula 
 MNaPO 4 , in which M represents a bivalent metal, is 
 formed : 
 
 NaP0 3 + MO = MNaPO 4 . 
 
 Upon facts of this kind depends the power of sodium 
 metaphosphate to dissolve metallic oxides, as when beads 
 formed by heating sodium ammonium phosphate are 
 used in analysis. The first effect of heating the phos- 
 phate is, as explained above, the formation of sodium 
 metaphosphate which melts, forming a clear liquid known 
 as the " bead of microcosmic salt." 
 
 Phosphorous Acid, H 3 PO 3 . This acid is formed when 
 pliosphorus trichloride is treated with water. It is also 
 formed together with phosphoric and hypophosphoric 
 .acids when phosphorus is allowed to lie in contact with 
 moist air. The acid can be obtained from its solutions 
 by evaporation, when it is deposited in transparent crys- 
 tals. When heated it is converted into phosphoric acid, 
 phosphine being given off : 
 
 4H 3 P0 3 = 3H 3 PO 4 + PH 3 . 
 
HYPOPHOSPHORIC AND HYPOPHOSPHOROUS ACID. 333 
 
 This reaction has been discussed under Phosphine (which 
 see). The tendency of phosphorous acid to take up 
 oxygen and form phosphoric acid makes it a good reduc- 
 ing agent. Its action is well illustrated in the case of 
 mercuric chloride, HgCl 2 , which it transforms into mer- 
 curous chloride, HgCl. Water being present, the phos- 
 phorous acid appropriates the oxygen of a part of it, 
 leaving the hydrogen to act upon the mercuric chloride : 
 
 2HgCl 2 + H 3 P0 3 + H 2 = H 3 P0 4 + 2HgCl + 2HC1. 
 
 Phosphorous acid is only dibasic, its salts having the 
 general formula HM 2 PO 3 . This fact has led to the be- 
 lief that in the acid two of the hydrogen atoms are in 
 combination with oxygen in the form of hydroxyl, while 
 the third is in combination with phosphorus as repre- 
 
 (H 
 sented in the formula OP < OH . This conclusion finds 
 
 (OH 
 
 further support in the conduct of some derivatives of 
 phosphorous acid. 
 
 Hypophosphoric Acid, H 4 P 2 O 6 , is formed together with 
 phosphoric and phosphorous acids when sticks of ordi- 
 nary phosphorus placed in glass tubes drawn out to a small 
 opening at one end are exposed to the action of moist air. 
 By arranging a number of such tubes on a funnel the 
 lower end of which is in a bottle, a solution is gradually 
 collected which contains hypophosphoric acid together 
 with the other two acids mentioned. The salts o the 
 acid show that it is tetrabasic. It has been suggested 
 that this acid has the constitution represented by the 
 
 OP(OH), 
 formula i 
 
 OP(OH), 
 
 Hypophosphorous Acid, H 3 PO 2 , has already been re- 
 ferred to, as its potassium salt is formed in the prepara- 
 tion of phosphine by the action of phosphorus upon a 
 solution of potassium hydroxide : 
 
 3KOH + 4P + 3H 2 = 3KH 2 PO 2 + PH 3 . 
 
 The acid is a solid which crystallizes well. The most 
 characteristic fact in its conduct is its marked tendency 
 
334 INORGANIC CHEMISTRY. 
 
 to pass over into phosphoric acid by taking up oxygen. 
 It is therefore a good reducing agent. It reduces sul- 
 phuric acid to sulphurous acid and even to sulphur, as 
 represented in the two equations below : 
 
 2H 2 SO 4 + H 3 P0 2 = 2H 2 S0 3 + H 3 PO 4 ; 
 H,S0 3 + H 3 P0 2 = S + H 2 + H 3 PO, 
 
 When heated, also, it forms phosphoric acid and phos- 
 phine just as phosphorous acid does : 
 
 mfO, =, H a PO, + PH 3 . 
 
 The acid is monobasic, and this has led to the belief that 
 only one of the hydrogen atoms in the molecule of the 
 acid is in combination with oxygen as hydroxyl, and that 
 the two others are in combination with phosphorus as 
 
 ( H 
 
 represented in the formula OP-< H . The relation be- 
 
 (OH 
 
 tween this acid and phosphorous and phosphoric acids 
 has already been commented upon (see page 323). 
 
 Phosphorus Pentoxide, Phosphoric Anhydride, P 2 O 5 . 
 This highest oxidation-product of phosphorus is formed 
 by burning the element in air or in oxygen. It is a 
 white powder which attracts moisture from the air and 
 becomes liquid. This power to combine with water is 
 its most characteristic property. It forms first, as we 
 have seen, metaphosphoric acid and, by further action, 
 orthophosphoric acid. Its action towards water is 
 strongly suggestive of the action of sulphur trioxide 
 or sulphuric anhydride towards water. Owing to this 
 power to combine with water, phosphorus pentoxide is 
 used for the purpose of drying gases, and as- a dehy- 
 drating agent. 
 
 Phosphorus Trioxide, or Phosphorous Anhydride, P 2 O 3 
 (or P 4 O 6 ), is formed by burning phosphorus so that the 
 air does not have free access to it, as by putting a piece 
 of phosphorus in a glass tube drawn out to a fine open- 
 ing, drawing air over the phosphorus, and warming it 
 gently. In this way not enough air can get access to the 
 
CONSTITUTION OF TEE ACIDS OF PHOSPHORUS. 335 
 
 phosphorus to convert it into the pentoxide. The tri- 
 oxide has such a strong tendency to pass over into the 
 pentoxide that when brought into the air it takes fire and 
 burns, forming the higher oxide. It is readily converted 
 into phosphorous acid by water. 
 
 Phosphorus Suboxide, P^O, is one of the products 
 formed by the burning of phosphorus in a limited sup- 
 ply of air. It is also formed in several other ways, and 
 two varieties of it have been described. 
 
 Phosphorus Tetroxide, P 2 O 4 , is also formed when phos- 
 phorus is burned in the air. It can be obtained pure in 
 the form of colorless crystals. 
 
 Constitution of the Acids of Phosphorus. Considerable 
 has already been said on this subject in dealing with the 
 relations between the acids. The view that phosphoric 
 acid contains three hydroxyl groups is based upon the 
 fact that the acid is tribasic, which, taken together with 
 what is known in regard to the conduct of other acids, 
 suggests that all three hydrogen atoms in the molecule 
 are in combination with oxygen. This view is the sim- 
 plest, and all facts known in regard to the conduct of 
 phosphoric acid are in accordance with it. The constitu- 
 
 /0-H 
 tion is represented by the formula O=P O-H , which 
 
 \O-H 
 
 may also be written in this way : OP(OH) 3 . Two views 
 suggest themselves in considering the constitution of 
 phosphorous acid. It may be, like phosphoric acid, a tri- 
 hydroxyl derivative of the formula P(OH) 3 , or it may have 
 
 /H 
 
 the structure represented by the formula O=P^-OH or 
 
 \OH 
 
 ( H 
 
 OP -j /QTT\ The easy formation of the acid from phos- 
 phorus trichloride and water is in accordance with the 
 former view. On the other hand, as has already been 
 remarked, the fact that the acid is dibasic speaks against 
 this view, and in favor of the latter. A somewhat com- 
 plex reaction of an organic derivative of phosphorous acid 
 also furnishes evidence in favor of the view 'that there are 
 only two hydroxyl groups contained in the molecule of 
 
336 INORGANIC CHEMISTRY. 
 
 phosphorous acid, and that its structure is represented by 
 
 / H 
 the formula O-P-O-H. 
 
 Phosphorus Oxy chloride, POC1 3 . This compound has 
 been referred to in connection with the chlorides of phos- 
 phorus. It is formed by the action of ozone on phos- 
 phorus trichloride and by the action of water upon the 
 pentachloride : 
 
 PC1 3 + = POC1 3 ; 
 
 PC1 B + H 3 O = POC1 3 + 2HC1. 
 
 It may be regarded as phosphoric acid in which all three 
 of the hydroxyl groups are replaced by chlorine, just as 
 sulphuryl chloride, SO 2 C1 2 , is to be regarded as sulphuric 
 acid in which both hydroxyls are replaced by chlorine". 
 The fact that when treated with water and other com- 
 pounds containing hydroxyl it yields phosphoric acid has 
 been mentioned, and the value of this reaction and the 
 similar reaction of phosphorus pentachloride as a means 
 of detecting the hydroxyl condition in compounds has 
 been pointed out (see p. 316). 
 
 Arsenic Acid, H 3 AsO 4 . The compound of arsenic and 
 oxygen which is most readily obtained is the trioxide, 
 As 2 O 3 , and this is formed by direct combination of the two 
 elements. When this is oxidized either with aqua regia 
 or by passing chlorine into water in which the trioxide is 
 suspended it is converted into arsenic acid : 
 
 As 2 O 3 + 3H 2 + 2O = 2H 3 AsO 4 . 
 
 From its solutions it is obtained in crystallized form. 
 According to the temperature to which it is heated 
 the deposit has the composition of the ortho-acid, H 3 AsO 4 , 
 of the pyro-acid, H 4 As 2 O 7 , or of the meta-acid, HAsO 3 . 
 Perfect analogy with the phosphorus compounds is here 
 observed. When the pyro- and meta-acids are dis- 
 solved in water they pass at once into the form of the 
 ortho-acid. Arsenic acid, like phosphoric acid, is a strong 
 tribasic acid, forming three series of salts which under 
 
ARSENIOUS ACID. 337 
 
 the influence of heat conduct themselves like the cor- 
 responding phosphates, the primary salts yielding pyro- 
 arsenates, and the secondary salts yielding meta-arsen- 
 ates. When these are dissolved in water they pass at 
 once into the corresponding salts of ortho-arsenic acid. 
 Arsenic acid is easily reduced to the form of arsenic. 
 When hydrogen sulphide is passed through a hydro- 
 chloric acid solution of arsenic acid different reactions 
 take place according to the conditions. The three pos- 
 sibilities are : (1) The formation of the pentasulphide ; 
 
 (2) the formation of sulphoxyarsenic acid, H 3 AsO 3 S ; and 
 
 (3) the formation of arsenic pentasulphide, arsenic trisul- 
 phide, and sulphur. These reactions are represented by 
 the following equations : 
 
 (1) H 3 As0 4 + H 2 S = H 3 AsO 3 S + H 2 O ; 
 
 (2) 2H 3 As0 3 S + 3H 2 S = As 2 S 5 + 6H 2 O ; 
 
 , Q x j 2H 3 As0 3 S + 6HC1 = 2AsCl 3 + 6H 2 O + 28 ; 
 
 I 
 
 33 3 2 
 
 2AsCl 3 + 3H 2 S = As 2 S 3 + 6HC1. 
 
 The first action is that represented by equation (1). 
 The acid thus formed, known as sulphoxyarsenic acid, 
 differs from arsenic acid only in the fact that it contains 
 a sulphur atom in the place of one oxygen atom. It is 
 soluble in water, and, therefore, when hydrogen sulphide 
 is passed into a solution of arsenic acid there is at 
 first no precipitate formed ; but gradually, where the hy- 
 drogen sulphide is in excess, some of the sulphoxyar- 
 senic acid is changed to arsenic pentasulphide, while an- 
 other part of the acid is decomposed by hydrochloric 
 acid, forming arsenic chloride and sulphur, and the tri- 
 sulphide is then precipitated. Therefore, the precipitate 
 formed by passing hydrogen sulphide into a solution 
 of arsenic acid is likely to consist of a mixture of arsenic 
 pentasulphide, trisulphide, and sulphur. 
 
 Arsenious Acid, H 3 AsO 3 , is not known in the free state, 
 but salts related to it are formed by treating arsenic tri- 
 oxide with bases. Thus, when it is treated with potas- 
 sium hydroxide the salt KAsO 2 , or potassium meta- 
 arsenite, is formed : 
 
 As a O, + 2KOH = 2KAs0 2 + H 9 O. 
 
338 INORGANIC CHEMISTRY. 
 
 Salts of meta-arsenious acid, AsO.OH, are more com- 
 monly obtained than those of the normal acid, As(OH) 3 . 
 In alkaline solution arsenious acid tends to pass into 
 the form of arsenic acid, and it is therefore a useful 
 reducing agent. Its action in this way is, however, not 
 as strong as that of phosphorous acid. 
 
 Arsenic Trioxide, As 2 O 3 . This compound is commonly 
 called arsenic or white arsenic. It is the most important 
 of all the compounds of the element arsenic. It finds 
 applications for many purposes, and is manufactured in 
 large quantities. It occurs in small quantity in nature, 
 but that which comes into the market is manufactured 
 by roa-sting natural arsenides, particularly arsenical py- 
 rites, FeAsS. The products of roasting this compound 
 are ferric oxide, Fe Q O 3 , sulphur dioxide, SO 2 , and arsenic 
 trioxide, As 2 O 3 . Of these, the first is a non-volatile solid, 
 the second a gas, and the third a volatile solid. By pass- 
 ing the volatile products through properly constructed 
 canals the arsenic trioxide is condensed on the walls. 
 Some of the powder thus obtained must be subjected to 
 a second process of distillation to make it pure enough 
 for the market. In a recent year over 6000 tons of this 
 substance were produced in England and Saxony. 
 
 Arsenic trioxide is a colorless, amorphous, vitreous 
 mass. Gradually it becomes opaque and crystalline, with 
 an appearance like that of porcelain. It crystallizes in 
 two forms, the common one being that of regular octa- 
 hedrons. Under exceptional conditions it crystallizes in 
 the form of rhombic prisms. When heated it sublimes, 
 and is deposited on a cold surface in the form of octa- 
 hedrons. Arsenic trioxide is difficultly soluble in water, 
 but more easily in hydrochloric acid. The solution in 
 hydrochloric acid contains arsenic trichloride (see p. 317), 
 and when the solution is boiled the chloride is carried 
 over. When the solution of the amorphous oxide in 
 hydrochloric acid is concentrated enough it deposits the 
 oxide in crystalline form, and the formation of the 
 crystals is accompanied by an evolution of light which 
 can be seen in a dark room. When the crystalline 
 variety is dissolved it is deposited in crystals without 
 
ARSENIC TRIOXIDE. 339 
 
 evolution of light. The formation of arsenic trichloride 
 by the action of hydrochloric acid on the oxide is perfectly 
 analogous to the formation of the chloride of any base- 
 forming element by the action of hydrochloric acid upon 
 the oxide, as, for example, ferric oxide. The reactions 
 are represented thus : 
 
 As 2 O 3 + 6HC1 = 2AsCl 3 + 3H 2 O ; 
 and Fe 2 3 + 6HC1 = 2FeCl 3 + 3H 2 O. 
 
 While in this reaction arsenic appears as a base-forming 
 element, its character as an acid-forming element shows 
 itself when the chloride is treated with a large excess of 
 water, under which circumstances it is completely con- 
 verted into the oxide. Towards some acids also arsenic 
 trioxide acts as a weak base. A somewhat complex sul- 
 phate is known in which the arsenic replaces a part of 
 the hydrogen of the acid. It is formed by treating the 
 trioxide with fuming sulphuric acid. 
 
 The trioxide is easily reduced. When heated with 
 potassium cyanide, KCN, or with charcoal in a dry glass 
 tube arsenic is deposited above the flame in the form of 
 a dark lustrous layer. When brought into a vessel from 
 which hydrogen is being evolved it is reduced to arsine. 
 
 The specific gravity of the vapor of the oxide shows 
 that it has the formula As 4 O 6 , and not As 2 O 3 ; as, however, 
 most of its reactions can be more conveniently expressed 
 by the aid of the simpler formula, the latter is commonly 
 used. 
 
 Arsenic trioxide has a weak, disagreeable, sweet taste, 
 and is an active poison. A dose of from two to three 
 grains is sufficient to cause death unless it is ejected by 
 vomiting, or rendered harmless by being converted into 
 an insoluble compound. It is possible, by beginning with 
 small doses, and gradually increasing them, to accus- 
 tom the human body to considerably larger doses 
 than that mentioned. It strengthens the power of the 
 respiratory organs, and consequently facilitates mountain- 
 climbing. The peasants in some mountain regions are 
 said to use it habitually. It is much used in medicine, 
 especially in skin diseases. It is also used extensively 
 
340 INORGANIC CHEMISTRY. 
 
 as a rat-poison. The most efficient antidote is a mixture 
 of ferric hydroxide, Fe(OH) 3 , and magnesia, which forms 
 with arsenic trioxide an insoluble compound. 
 
 Arsenic Pentoxide, As 2 O 5 , is formed by igniting arsenic 
 acid. If heated too high the pentoxide breaks down into 
 arsenic trioxide and oxygen. A marked difference will be 
 observed between the conduct of the oxides of phosphorus 
 and that of the corresponding oxides of arsenic. While 
 phosphorus trioxide takes up oxygen spontaneously when 
 exposed to the air, and the pentoxide is not decomposed 
 by heat, the trioxide of arsenic does not under any cir- 
 cumstances take up oxygen directly, and the pentoxide 
 easily breaks down into the trioxide and oxygen when 
 heated. 
 
 Sulphides. There are three compounds of arsenic with 
 sulphur the disulphide, As 2 S 2 , the trisulphide, As 2 S 3 , 
 and the pentasulphide, As 2 S 5 . 
 
 Arsenic Disulphide, As 2 S 2 , occurs in nature and is 
 known as realgar. It can also be obtained by melting 
 arsenic and sulphur together in the right proportions. 
 It forms an orange-red powder which was formerly used 
 as a pigment. 
 
 Arsenic Trisulphide, As 2 S 3 , is found in nature and is 
 called orpiment or king's yellow. It can be prepared by 
 melting together arsenic and sulphur in the proper pro- 
 portions, and by precipitating a solution of arsenic tri- 
 oxide in hydrochloric acid with hydrogen sulphide. It 
 melts, forming a red liquid. The natural substance, as 
 well as that which is precipitated by means of hydrogen 
 sulphide, is yellow. It dissolves in soluble sulphides, 
 forming salts of sulpharsenious acid, H 3 AsS 3 , or HAsS 2 . 
 The salts are, for the most part, derived from the acid of 
 the latter formula. There is, therefore, perfect analogy 
 between the oxygen and sulphur compounds, for, as we 
 have seen, when arsenic trioxide is dissolved in potassium 
 hydroxide a salt of the formula KAsO 2 is formed. The 
 analogy is clearly shown by means of the equations 
 
 As 2 O 8 + 2KOH = 2KAsO 2 + H 2 O ; 
 
 As a S 3 + 2KSH = 2KAsS 2 + H 2 S. 
 
ARSENIC TRISULPHIDE. 341 
 
 The acid HAsS 2 is derived from the corresponding nor- 
 
 (SH 
 mal acid As -< SH , by loss of one molecule of hydrogen 
 
 ( SH 
 sulphide : 
 
 (SH 
 
 As^ SH = 
 
 (SH 
 
 just as the acid HAsO 2 is derived from the normal oxy- 
 
 (OH 
 gen acid As -< OH , by loss of one molecule of water : 
 
 (OH 
 
 (OH 
 
 As^OH = 
 (OH 
 
 When a solution of a sulpharsenite is treated with one 
 of the stronger acids, as, for example, hydrochloric acid, 
 arsenic trisulphide is precipitated. We should naturally 
 look for the separation of the free acid according to the 
 equation 
 
 KAsS 2 + HC1 = HAsS 2 + KC1 ; 
 
 but, if this is formed, it breaks down at once into hydro- 
 gen sulphide and arsenic trisulphide : 
 
 2HAsS 2 = As 2 S 3 + H 2 S. 
 
 There is a striking analogy between this action and that 
 which takes place when a stronger acid is added to a 
 solution of a carbonate, when carbon dioxide is set free : 
 
 K,C0 3 + 2HC1 = H,C0 3 + 2KC1 ; 
 H 2 C0 3 = C0 2 + H 2 0. 
 
 A marked difference between the two cases is to be 
 found in the fact that the trisulphide of arsenic is insol- 
 uble in water and therefore appears as a precipitate, 
 while carbon dioxide escapes as a gas. 
 
 Besides salts of the acids H 3 AsS s and HAsS 2 , there are 
 others derived from the more complex acid H 4 As 3 S 6 . 
 
342 INORGANIC CHEMISTRY. 
 
 This bears to normal sulpharsenious acid, As(SH) 3 , a re- 
 lation similar to that which pyrophosphoric acid bears 
 to orthophosphoric. If two molecules of the normal acid 
 lose one molecule of hydrogen sulphide, this pyrosulph- 
 arsenious acid is the product : 
 
 (SH 
 
 2As-{ SH = As 2 S(SH) 4 + H 2 S. 
 (SH 
 
 It is a salt of this acid which is formed when arsenic 
 trisulphide is dissolved in ammonium sulphide : 
 
 As 2 S 3 + 2(NH 4 ) 2 S = As 2 S(SNH 4 ) 4 . 
 
 Arsenic Pentasulphide, As 2 S 3 , is formed by melting sul- 
 phur and arsenic together in the proper proportions, 
 and by precipitating a solution of sodium sulpharsenate 
 with hydrochloric acid : 
 
 2Na 3 AsS 4 + 6HC1 = 6NaCl + As 2 S 5 + 3H 2 S. 
 
 Sulpharsenic acids corresponding to the oxygen acids 
 suggest themselves. We might, for example, expect to 
 find salts derived from the acids H 3 AsS 4 , HAsS 3 , and 
 H 4 As 2 S 7 , corresponding to ortho-, meta-, and pyro-arsenic 
 acids. When arsenic pentasulphide is dissolved in solu- 
 tions of metallic sulphides the products are generally 
 salts of pyrosulpharsenic acid, H 4 As 2 S 7 , and these under- 
 go decomposition into salts of the ortho- and meta-acids. 
 When, for example, arsenic pentasulphide is dissolved in 
 ammonium sulphide reaction takes place thus : 
 
 AsA + 2(NH 4 ) Z S = (NH 4 ) 4 As,S, 
 
 The ammonium salt formed in this way is, however, de- 
 composed thus : 
 
 (NH 4 ) 4 As,S, = (NH 4 ),AsS 4 + (NH 4 )AsS s . 
 
 Only one compound intermediate between arsenic and 
 sulpharsenic acids is known. This is the sulphoxyarsenic 
 acid formed as the first product of the action of hydro- 
 
ANTIMONIC ACID ANTIMONY TRIOXIDE. 343 
 
 gen sulphide upon a solution of arsenic acid, which was 
 referred to under Arsenic Acid (p. 337). The possibility 
 of other products of the formulas H 3 AsO 2 S 2 and 
 H 3 AsOS 3 will occur to every one. 
 
 Antimonic Acid, H 3 SbO 4 . This acid is the final product 
 of the oxidation of antimony when treated with aqua 
 regia. It need only be said that it is very similar to 
 phosphoric and arsenic acids ; and that, like these, it 
 yields a meta- and a pyro-acid of the formulas HSbO 3 
 and H 4 Sb 2 O 7 . The acid of the formula OSb(OH) 3 , or 
 orthoantimonic acid, is known in the free state, and is 
 formed by treating a soluble salt of antimonic acid with 
 sulphuric or nitric acid : 
 
 OSb(OK) 3 + 3HNO 3 = 3KNO 3 + OSb(OH) 3 . 
 
 An acid Sb 2 O(OH) 8 is also known in the free state, be- 
 ing formed by the action of antimony pentachloride upon 
 water. The lower oxides of antimony, the trioxide, Sb 2 O 3 , 
 and the tetroxide, Sb 2 O 4 , are not strongly acidic; that is 
 to say, they do not readily form salts when treated with 
 bases. In this respect the trioxide of antimony differs 
 markedly from the corresponding oxides of phosphorus 
 and arsenic. 
 
 Antimony Trioxide, Sb 2 O 3 . This compound is found 
 in nature as white ore of antimony, and is easily formed 
 by burning antimony in the air and by oxidizing it with 
 nitric acid or saltpeter. That formed by burning anti- 
 mony in the air always contains some of the tetroxide, and 
 by heating it long enough in the air and to a temperature 
 high enough it is completely transformed into the tetrox- 
 ide. When the trioxide is dissolved in caustic soda a 
 salt of the formula NaSbO 2 is formed. This is plainly 
 derived from an acid of the formula HSbO 2 , which bears 
 a simple relation to normal antimonious acid. Towards 
 most bases, however, antimony trioxide does not conduct 
 itself as an acid. On the other hand, towards the stronger 
 acids it acts as a base. 
 
 Salts of Antimony. The salts of antimony are derived 
 either from the hydroxide Sb(OH) 3 , or from the hydrox- 
 ide SbO.OH. The salts of the first class are called anti- 
 
344 INORGANIC CHEMISTRY. 
 
 mony salts / those of the second class are called antimonyl 
 salts. In the salts formed when the trihydroxide of an- 
 timony is completely neutralized by acids, the antimony 
 takes the place of three atoms of hydrogen. Thus, the 
 nitrate has the formula Sb(NO 3 ) 3 ; the sulphate has the 
 formula Sb 3 (SO 4 ) 3 ; etc. Besides these normal salts there 
 are, however, basic salts. Thus there are two basic 
 
 (OH (OH 
 
 nitrates possible of the formulas Sb -< OH and Sb -< NO 3 . 
 
 I NO, ( NO, 
 
 The formation of antimonyl salts may be illustrated by 
 the sulphate. This may be regarded as formed by the 
 action of sulphuric acid upon the hydroxide SbO.OH, 
 which is analogous in composition to the acid of arsenic 
 of the formula AsO.OH : 
 
 2SbO.OH 
 
 The product is antimonyl sulphate. The weak basic 
 character of the hydroxides of antimony is shown by the 
 fact that many of its salts are decomposed by water. 
 The salt of antimony which is most commonly met with 
 is the so-called tartar emetic, which appears to be an anti- 
 monyl potassium salt of tartaric acid. Tartaric acid is a 
 
 ( OTT 
 dibasic acid of the formula C 4 H 4 O 4 -j QTT . When one of 
 
 its acid hydrogen atoms is replaced by potassium, and 
 the other by the antimonyl group SbO, the salt thus 
 
 formed is tartar emetic, C 4 H 4 O 4 \ QTT- . It is also pos- 
 sible that this salt may be derived from the trihydroxide 
 Sb(OH) 3 by replacement of one hydrogen atom by potas- 
 sium, and neutralization of the rest of the compound by 
 the dibasic tartaric acid. It seems more probable, how- 
 ever, that when tartaric acid acts upon the compound 
 
 Sb j x-r ' 2 it first appropriates the potassium atom, form- 
 ing acid potassium tartrate, and that the antimony triox- 
 ide being basic is neutralized by the acid tartrate. To 
 decide between the two views is at present impos- 
 sible. 
 
OXIDES AND SULPHIDES OF ANTIMONY. 345 
 
 Antimony trioxide dissolves in hydrochloric acid, 
 forming the trichloride, and this, as has been stated, is 
 decomposed by water yielding oxychloricies. 
 
 Antimony Tetroxide, Sb 2 O 4 . This compound is most 
 easily obtained by igniting antimonic acid, H 3 SbO 4 . Two 
 reactions are of course involved : 
 
 2H 3 Sb0 4 = Sb 2 5 + 3H 2 ; 
 Sb 2 6 = Sb 2 4 + O. 
 
 It is also formed by igniting the trioxide in the air. At 
 ordinary temperatures the tetroxide is white, but it be- 
 comes yellow when heated. Towards strong acids this 
 oxide acts like a weak base. A potassium salt of the 
 formula K 2 Sb 2 O 5 is known, which is derived from the 
 acid H 2 Sb 2 O 5 , and this in turn from the simpler acid 
 SbO(OH) 2 by loss of water. The oxide itself is regarded 
 by some as an antimonyl salt of metantimonic acid, 
 SbO 2 .OH, of the formula SbO 2 .O.SbO. 
 
 Antimony Pentoxide, Sb 2 O 5 . The tetroxide of anti- 
 mony does not combine with oxygen to form the pentox- 
 ide. The latter can be obtained only by gentle ignition 
 of antimonic acid, care being taken not to raise the tem- 
 perature high enough to decompose the pentoxide into 
 the tetroxide and oxygen. The fact that the pentoxide 
 readily yields salts of antimonic acid when treated with 
 basic solutions was mentioned under Antimonic Acid. 
 
 Antimony Trisulphide, Sb 2 S 3 . This compound occurs 
 in nature in considerable quantity and is the chief source 
 of antimony. It is known as stibnite and antimony blende. 
 In some localities, especially in Japan, it occurs in large 
 rystals of great beauty. When heated in the air, or 
 roasted, it is converted into the trioxide, and finally into 
 the tetroxide, while the sulphur escapes as the dioxide. 
 Hydrochloric acid dissolves the trisulphide in the form 
 of the chloride with evolution of hydrogen sulphide : 
 
 Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. 
 
 Nitric acid converts it into the oxide with separation of 
 sulphur. When a solution of antimony chloride is 
 treated with hydrogen sulphide, the trisulphide is thrown 
 
346 INORGANIC CHEMISTRY. 
 
 down. This artificially prepared trisulphide has an 
 orange-red color, while that which occurs in nature is 
 black or gray. The sulphide dissolves in solutions of 
 metallic sulphides, forming salts of sulphantimonious 
 acid, either SbS.SH or Sb(SH) 3 . 
 
 Antimony Pentasulphide, Sb 2 S 5 , is formed by passing 
 hydrogen sulphide into a solution of antiinonic acid or 
 by decomposing a salt of sulphantimonic acid by means 
 of an acid. The action takes place thus : 
 
 2H 3 Sb0 4 + 5H 2 S = Sb 2 S B + 8H 2 ; 
 2Na,SbS 4 + 6HC1 = GNaCl + Sb 2 S & + 3H 2 S. 
 
 It is, when dry, a golden-yellow powder known as sul- 
 phur auratum. It dissolves easily in solutions of metallic 
 sulphides, forming the sulphantimonates, of which the 
 sodium salt, Na 3 SbS 4 , known as Schlippe's salt, is a good 
 example. The action is represented by this equation : 
 
 Sb 2 S & + GNaSH = 2Na 3 SbS 4 + 3H 2 S. 
 
 When heated in the air the pentasulphide gives 
 off enough sulphur to form the trisulphide ; while 
 when the pentoxide is heated it is converted into the 
 tetroxide. The sulphantimonates are decomposed when 
 treated with acids and the pentasulphide is thrown down. 
 
 Constitution of the Acids of Arsenic and Antimony. 
 There is, in general, marked analogy between the com- 
 pounds of phosphorus and those of arsenic and anti- 
 mony. In one particular, however, there is a difference 
 which is worthy of special mention. It appears that, 
 while phosphorous acid is dibasic and probably has the 
 
 (H 
 structure OP < OH , arsenious and antimonious acids are 
 
 (OH 
 
 the normal compounds represented by the formulas 
 As(OH) 3 and Sb(OH) 3 . Arsenic and antimonic acids ap- 
 pear to have the same structure as phosphoric acid 
 
 represented by the formulas As ] /^TTX and Sb -! /ri -rj x . 
 
 ( (Ui) 3 ( (Oi) 3 
 
 The difference between phosphorous and arsenious acids 
 suggests the difference between sulphurous and selenious 
 
OXIDES OF BISMUTH. 347 
 
 acids. While, according to the evidence, the constitution 
 of sulphurous acid is that represented by the formula 
 
 {TT 
 ^TT , that of selenious acid is represented by the 
 
 formulaOSe 
 
 Oxychlorides of Antimony. Under the head of Anti- 
 mony Trichloride the fact was mentioned that this com- 
 pound is decomposed by cold water as represented in 
 the equation 
 
 SbCl 3 + H 2 O = SbOCl + 2HC1. 
 
 If, however, hot water is used, the composition of the 
 product approximates to that represented by the formula 
 Sb 4 O 5 Cl 2 . This complex mixture of oxychlorides is 
 known as the "Poivder of Algaroth" It may be regarded 
 as derived from the simple oxychloride by loss of anti- 
 mony trichloride, thus : 
 
 SSbOCl = Sb 4 O B Cl 2 + SbCl 3 . 
 
 Many other oxychlorides besides the two mentioned 
 have been obtained, but they are all more or less closely 
 related to the simple compound SbOCl. 
 
 Oxides of Bismuth. The principal compound of bis- 
 muth and oxygen is the trioxide, Bi 2 O 3 , which is formed 
 when bismuth is burned in the air. It is a yellow pow- 
 der. Besides the method just mentioned, it is formed 
 by decomposing bismuth nitrate by high heat. If a so- 
 lution of bismuth nitrate, Bi(NO 3 ) 3 , is treated with a 
 cold solution of potassium hydroxide, bismuth hydrox- 
 ide, Bi(OH) 3 , is thrown down. When this is dried at 
 100 it loses water and is converted into the hydroxide, 
 BiO(OH) ; and if the hydroxide first precipitated is boiled 
 with the solution it is converted into the yellow oxide, 
 Bi 2 O 3 . The reactions involved are , . 
 
 Bi(NO 3 ) 3 + 3KOH = Bi(OH) 3 + 3KN0 3 ; 
 Bi(OH) 3 = BiO.OH + H 2 O ; 
 2Bi(OH) 3 = Bi 2 O 3 + 3H 2 0. 
 
 The trioxide of bismuth is basic and forms salts which 
 in composition correspond to the salts of antimony. 
 
348 INORGANIC CHEMISTRY. 
 
 Like the latter, they are of two classes the bismuth salts 
 and the bismuthyl salts. The former are derived from 
 the triacid base, Bi(OH) 3 , the latter from the monacid 
 base, BiO(OH). 
 
 Salts of Bismuth. The best known salts of bismuth 
 are those which it forms with sulphuric and with nitric 
 acids. There is a sulphate of the formula BiH(SO 4 ) 3 
 formed by dissolving bismuth oxide in dilute sulphuric 
 acid. The sulphate which is most stable in the presence 
 of water is the bismuthyl salt, (BiO) 2 SO 4 . When 
 bismuth is dissolved in nitric acid and the solution 
 evaporated to dryness the salt Bi(NO 3 ) 3 + 10H 2 O is ob- 
 tained. This salt is decomposed when heated, and by 
 water, forming basic nitrates of bismuth. The composition 
 of the basic nitrate obtained by decomposing the neutral 
 nitrate with water differs according to the conditions. 
 Hot and cold water produce different results. A solu- 
 tion containing much nitric acid does not give the same 
 result as one which contains little, etc. As basic bismuth 
 nitrate is used in medicine it is necessary that specific 
 directions should be given for its preparation, in order 
 that a substance of the same composition should always 
 be obtained. Among the basic nitrates which have been 
 
 isolated are the following : Bi j ^\ BiO.NO 3 and 
 
 ( O.BiO 
 Bi-< O.NO 2 . Besides these many of much more complex 
 
 (OH 
 
 composition are known, but all of them can be referred 
 to the simple forms. Some of them are of special inter- 
 est, as they appear to be derived from complex forms of 
 nitric acid, as, for example, an acid of the formula 
 N 2 O 3 (OH) 4 or H 4 N 2 O 7 , which is analogous to pyrophos- 
 phoric, pyroarsenic, and pyroantimonic acids. The basic 
 nitrate of bismuth, or the subnitrate, as it is frequently 
 called in pharmacy, is much used in medicine as a rem- 
 edy in dysentery and cholera. It is also used as a 
 cosmetic. 
 
 Bismuth Dioxide, Bi 2 O 2 , is formed as a brown precipi- 
 tate when potassium hydroxide is added to a solution of 
 
COMPOUNDS OF BISMUTH. 349 
 
 bismuth chloride and stannous chloride, SnCl 2 . Stan- 
 nous chloride combines very readily with chlorine to 
 form stannic chloride, SnCl 4 . When, therefore, stannous 
 chloride and bismuth chloride are brought together, it 
 is probable that the former extracts a part of the chlo- 
 rine from the latter, forming a chloride of the formula 
 BiCl 2 , and this with the potassium hydroxide breaks 
 down, yielding the dioxide : 
 
 2BiCl 3 + Sn01 2 = 2BiCl 2 + SnCl 4 ; 
 2BiCl a + 4KOH = Bi 2 2 + 4KC1 + 2H 2 O. 
 
 Bismuth. Pentoxide, Bi 2 O 5 , is formed by oxidizing the 
 trioxide, by means of chlorine, in alkaline solution. Al- 
 though some experimenters appear to have obtained salts 
 of bismuthic acid, as, for example, KBiO 3 , others have 
 failed to obtain them. In any case it is evident that the 
 acid properties of the oxide are very weak. 
 
 Bismuth. Trisulphide, Bi 2 S 3 , occurs in nature, and is 
 formed by precipitating bismuth from solutions of its 
 salts with hydrogen sulphide. It dissolves in hot con- 
 centrated hydrochloric acid and in nitric acid. It does 
 not dissolve in solutions of the sulphides as the sulphides 
 of arsenic and antimony do. 
 
 Bismuth Oxychloride, BiOCl, which in composition is 
 analogous to the simplest form of antimony oxy chloride, 
 is thrown down as a white powder when a solution con- 
 taining bismuth chloride is treated with water : 
 
 BiCl 3 + H 3 O = BiOCl + 2HC1. 
 
 FAMILY Y, GROUP A. 
 
 As the members of Group A, Family VII, are related 
 to Group B of the same family ; and as the members of 
 Group A, Family VI, are related to the members of Group 
 B of the same family, so the members of Group A,. Family 
 V, are related to the members of Group B, which have just 
 been studied. The members of Group A are vanadium, 
 columbium, tantalum, and didymium, all of which are 
 rare. Of these vanadium has been most thoroughly in- 
 vestigated, and columbium next, 
 
350 INORGANIC CHEMISTRY. 
 
 Vanadium, V (At. Wt. 50.99). This element occurs in 
 nature in the form of vanadates or salts of vanadic acid, 
 H 3 VO 4 , which is analogous to phosphoric acid. The 
 methods employed in separating the element from its 
 compounds depend upon the composition of the com- 
 pound. In the separation advantage is frequently taken 
 of the fact that the ammonium salt of vanadic acid is 
 difficultly soluble in a solution of ammonium chloride. 
 When this ammonium salt is ignited it is converted into 
 the pentoxide V 2 O 6 . With chlorine, vanadium forms the 
 compounds YC1 2 , YC1 3 , and YC1 4 ; with oxygen, the 
 compounds V 2 O, Y 2 O 2 , Y 2 O 3 , Y 2 O 4 , and Y 2 O 5 . In its re- 
 lations to oxygen it suggests nitrogen. The oxide, Y 2 O 4 , 
 conducts itself something like the tetroxide of antimony. 
 Towards strong bases it acts like an acid, forming salts 
 of the general formula Y 4 O 7 (OM) 2 . (See Antimony 
 Tetroxide.) 
 
 Vanadic Acid, H 3 VO 4 , is the most important and best 
 known of the compounds of vanadium. It is the final 
 product of the oxidation of vanadium, and bears to this 
 element the same relation that phosphoric, arsenic, and 
 antimonic acids bear to phosphorus, arsenic, and anti- 
 mony. The vanadates are derived from ortho-, meta-, 
 and pyro-vanadic acids, though the most stable ones are 
 the metavanadates, MYO 3 . The free metavanadic acid is 
 known. It is a beautiful golden-yellow compound, which 
 may be used as a substitute for gold bronze. An oxy- 
 chloride of the formula YOC1 3 , corresponding to phos- 
 phorus oxychloride, is made by direct addition of chlorine 
 to vanadium dioxide. 
 
 Tantalum, Ta (At. Wt. 181.45). Tantalum occurs in the 
 minerals columbite and tantalite, accompanied by nio- 
 bium. With the members of the chlorine group it forms 
 the compounds TaF 6 , TaCl 6 , TaBr 5 , and TaI 6 . Tantalum 
 fluoride combines easily with the fluorides of other metals 
 forming the fluotantalates. These may be regarded as 
 salts of fluotantalic acid, which are derived from the oxy- 
 gen acids by replacement of a part or all of the oxygen by 
 fluorine. Thus, the salt K 2 TaF 7 is easily obtained by 
 treating tantalum fluoride with a solution of potassium 
 
BORON. 351 
 
 fluoride. This is a salt of the acid H 2 TaF 7 or H 4 Ta 2 F 14 , 
 which is analogous to the oxygen acid H 4 Ta 2 O 7 . With 
 oxygen it forms Ta 2 O 4 and Ta 2 O 5 . The latter forms the 
 tantalates with bases. When tantalum pentachloride is 
 decomposed with water it forms the acid H 4 Ta a O, or 
 pyrotantalic acid : 
 
 2TaCl 5 + 7H a O = Ta 2 3 (OH) 4 + 10HC1. 
 
 The tantalates are derived from the meta-acid HTa0 3 , 
 and from the hexa-acid H 8 Ta 6 O 19 , which is derived from 
 the ortho-acid as represented in this equation : 
 
 6H 3 TaO 4 = H 8 Ta 6 O 19 + 5H 3 O. 
 
 Columbium, Cb (At. Wt. 93.7). This element, which is 
 sometimes called niobium, occurs in the mineral colum- 
 bite. It forms two chlorides, CbCl 3 and CbCl 6 , and a 
 bromide and fluoride corresponding to the latter chlo- 
 ride. The fluoride readily forms fluocolumbates, similar 
 to the fluotantalates. The niobates are- derived from a 
 number of forms of the acid which are, however, closely 
 related to the ortho-acid H 3 CbO 4 . 
 
 Didymium consists of two very similar elements, neo- 
 dymium and praseodymium. In some of their compounds 
 they show a resemblance to the members of this group. 
 They form, for example, an oxide of the formula Di 2 O 6 . 
 On the other hand, they seem to be more closely related 
 to cerium and lanthanum, which are also very rare ele- 
 ments, occurring associated with didymium. These will 
 be further treated of in connection with lanthanum and 
 cerium. 
 
 BOEON, B (At. Wt. 10.86). 
 
 General. Although the element boron is not a mem- 
 ber of the family to which nitrogen and phosphorus be- 
 long, it nevertheless resembles the members of this 
 family in some respects. It belongs to the same family 
 as aluminium, and in the composition of its compounds 
 it is undoubtedly similar to aluminium ; but, on the other 
 hand, its oxide is distinctly acidic, while that of aluminium 
 is basic. 
 
352 INORGANIC CHEMISTRY. 
 
 Occurrence. Boron occurs in nature chiefly in the form 
 of boric acid, or as salts of this acid, particularly a 
 sodium salt known as borax. 
 
 Preparation. From borax and the other borates the 
 acid can easily be obtained. When heated, water is given 
 off, and boron trioxide, B 2 O 8 , is left : 
 
 2B(OH) 3 = B 3 3 + 3H,0. 
 
 By heating the oxide with potassium amorphous boron 
 is obtained. By melting the oxide with aluminium, 
 boron is formed and is dissolved in the molten alumin- 
 ium, from which, on cooling, it is deposited in crystals. 
 Amorphous boron in almost pure form is obtained by 
 heating borax with magnesium powder. One of the 
 chief difficulties encountered in preparing boron is to 
 prevent the element from combining with the nitrogen 
 of the air. At the high temperature at which the 
 reduction takes place the two elements combine very 
 readily to form, the compound boron nitride, BN. The 
 crystals obtained in the process described are not pure 
 boron, but contain aluminium, or carbon and aluminium, 
 apparently in combination with the boron. The crystals 
 are very hard, and some of them have a high lustre. 
 
 Properties. Amorphous boron is a greenish-brown 
 powder. It burns when heated in the air or in oxygen, 
 the product being the trioxide B 2 O 3 . Strong oxidizing 
 agents, like nitric acid and saltpeter, readily oxidize it> 
 forming boric acid. It combines readily also with many 
 other elements, as with chlorine, nitrogen, and sulphur. 
 When it is brought into the melting hydroxides or car- 
 bonates of potassium or sodium, it forms borates of the 
 corresponding metals. 
 
 Boron Trichloride, BC1 9 . This compound is formed by 
 heating boron in a current of dry chlorine, and by heat- 
 ing a mixture of boron trioxide and charcoal in chlorine : 
 
 2B,0 8 + 30 + 601, = 4BC1 3 + 3CO,. 
 
 This reaction is especially interesting on account of its 
 double character. Carbon alone could not reduce the 
 
BORON TRIFLUORIDE. 353 
 
 boron trioxide at the temperature employed ; nor could 
 the chlorine alone displace the oxygen and form the 
 chloride, but when both chlorine and carbon act together 
 these changes take place, one aiding the other. 
 
 The chloride is a liquid which boils at 17. Like 
 phosphorus trichloride, it is easily decomposed by water, 
 forming boric acid, which, as will be seen, is analogous 
 in composition to phosphorous acid and arsenious acid : 
 
 BC1 3 + 3H 2 = B(OH) 3 + 3HC1. 
 
 This decomposition is analogous to that of arsenic tri- 
 chloride rather than to that of phosphorus trichloride, 
 for in the latter case a secondary change takes place, re- 
 sulting in the formation of an acid of the constitution 
 H 
 
 OP - 1 (OH); 
 
 Boron Trifluoride, BF 3 , is obtained by treating a mix- 
 ture of fluor-spar and boron trioxide with concentrated 
 sulphuric acid. The reaction is a double one, consisting, 
 first, in the setting free of hydrofluoric acid from the 
 fluor-spar : 
 
 CaF 2 + H 2 SO 4 = CaS0 4 + 2HF ; 
 
 and, second, in the action of the hydrofluoric acid upon 
 the oxide of boron : 
 
 B 2 O 3 + 6HF = 2BF 3 + 3H 2 O. 
 
 It is a colorless gas, which acts upon water, and therefore 
 forms a thick white cloud in the air. The action upon 
 water is represented by the equation 
 
 4BF 3 + 3H 2 O = B(OH) 3 + 3HBF 4 . 
 
 The first action which we should expect is the formation 
 of normal boric acid, thus : 
 
 BF S + 3H 2 O = B(OH) 3 + 3HF. 
 
 But the hydrofluoric acid combines with some of the 
 trifluoride of boron, forming the compound HBF 4 , which 
 is known as fluoboric acid. Several elements act in this 
 
354 INORGANIC CHEMISTRY. 
 
 way, particularly the members of the silicon group. 
 Silicon itself forms the well-known compound fluosilicic 
 acid. Fluoboric acid is to be regarded as metaboric acid, 
 HBO 2 , in which the two oxygen atoms have been re- 
 placed by fluorine. The acid has been obtained in the 
 free state, and is a liquid boiling at 120. It forms salts 
 of the general formula MBF 4 , of which the potassium 
 salt, KBF 4 , is the best example. 
 
 Boric Acid, B(OH) 3 . Boric acid occurs free in nature 
 and in the form of salts, of which the principal one is 
 borax. Besides borax, which is a sodium salt derived 
 from tetraboric acid, H 2 B 4 O 7 , there are other natural 
 borates, as boracite, which is a magnesium salt combined 
 with magnesium chloride ; and datholite, which is made 
 up of silicic acid, boric acid, and the element calcium. 
 One of the most interesting natural forms of boric acid 
 is that which is given off from the earth with steam. 
 Such jets of steam are met with in many volcanic regions, 
 and are called fumaroles. In Tuscany many of the 
 fumarbles are charged with small quantities of boric 
 acid, which is somewhat volatile with steam. Those at 
 Monte Cerboli and Monte Kotundo in Tuscany are util- 
 ized for the purpose of obtaining the boric acid. For 
 this purpose basins are built over the fumaroles and filled 
 with water, so that the steam is condensed and the boric 
 acid dissolved in the water. The solutions formed at the 
 higher levels flow into basins at lower levels, and finally 
 become charged with a considerable quantity of the acid, 
 when it is evaporated to crystallization by the aid of the 
 heat furnished by the fumaroles. The acid obtained in 
 this way is not pure, but by recrystallization it is easily 
 purified. 
 
 Boric acid can also be made from borax by heating the 
 salt in solution with dilute sulphuric acid : 
 
 Na 3 B 4 T + H 2 S0 4 + 5H 3 O = Na 2 SO 4 + 4B(OH) 3 . 
 
 If the solution is sufficiently concentrated the boric acid 
 crystallizes out on cooling. 
 
 Boric acid is easily soluble in water, and crystallizes 
 from the solution. It is also soluble in alcohol, and this 
 
BORIC ACID. 355 
 
 solution burns with a characteristic green flame. The 
 acid is quite volatile with water vapor. When heated at 
 100 orthoboric acid loses one molecule of water, and is 
 converted into metaboric acid, HBO 2 ; at 160 it yields 
 tetraboric acid, H 2 B 4 O 7 ; and at a higher temperature it is 
 converted into boron trioxide or boric anhydride, B a O 8 . 
 These changes are represented in the equations follow- 
 ing: 
 
 OH 
 
 /OH 
 
 B(O 
 
 M>H 
 
 OH 
 
 H 
 OH 
 
 \OH 
 
 The most stable salts are the tetraborates and meta- 
 borates. Borax is the sodium salt of tetraboric acid, 
 Na 2 B 4 O 7 . The salts of orthoboric acid are unstable. 
 They break down when treated with water, forming free 
 boric acid and either metaborates or tetraborates. 
 
 When heated together with oxides, boric oxide forms 
 salts just as boric oxide and water form boric acid. 
 Borax also, when treated with metallic oxides, forms 
 double borates, which are derived from normal boric 
 acid. Thus with copper oxide action takes place which 
 should probably be represented thus : 
 
 Na 2 B 4 O 7 + 5CuO = Na 2 Cu 6 (BO 3 ) 4 ; 
 and B a O 3 + 3CuO = Cu 3 (BO 3 ) 2 . 
 
 Many of these salts are colored, and the action of 
 metallic compounds upon boron trioxide and upon borax 
 ig utilized for the purpose of determining their nature 
 
356 INORGANIG CHEMISTRY. 
 
 by the color of the mass formed. It will be remembered 
 that sodium metaphosphate is used in the same way* 
 With it the oxides form salts of phosphoric acid. 
 
 Most of the boric acid obtained from Tuscany is used 
 in the manufacture of borax, a salt which finds extensive 
 application. 
 
 Salts of Boron. Although the most characteristic com- 
 pounds of boron are those in which it acts as an acid- 
 forming element, it forms some compounds in which its 
 power as a base-former is shown. Thus, with concen- 
 trated sulphuric acid the trioxide forms a compound 
 which appears to be pyrosulphuric acid, H 2 S 2 O 7 , in which 
 one hydrogen is replaced by the group BO, which 
 is analogous to antimonyl, SbO, and bismuthyl, BiO. 
 It has the composition (BO)HS 2 O 7 . Further, when con- 
 centrated phosphoric acid acts upon crystallized boric 
 acid, boron phosphate, BPO 4 , is formed. This compound 
 is characterized by great stability. Concentrated acids, 
 for example, do not decompose itt It also forms a salt 
 which appears to be analogous to tartar emetic, which, 
 as has been pointed out, is probably antimonyl potas- 
 
 sium tartrate, C 4 H 4 O 6 -j ^ . This is the salt represented 
 
 ( "RO 
 
 by the formula C 4 H 4 O 6 \ -g- , which may be called boryl 
 
 potassium tartrate. 
 
 Nitrogen Boride, BN. This Compound has been re- 
 ferred to in connection with the preparation of boron. 
 It is easily obtained by igniting a mixture of dehydrated 
 borax and ammonium chloride. It forms a white pow- 
 der, which is insoluble in water, and is characterized by 
 great stability. At red heat it is decomposed by water 
 vapor into ammonia and boric acid : 
 
 2BN + 6H 3 = 2B(OH) 3 + 2NH,. 
 
 IJJ.S 
 
CHAPTER XIX. 
 
 CARBON (C, At. Wt. 11.92) AND ITS SIMPLER COMPOUNDS 
 WITH HYDROGEN AND CHLORINE. 
 
 Introductory. Carbon bears to Family IV relations 
 similar to those which nitrogen, oxygen, and fluorine 
 bear to Families V, VI, and VII. Towards hydrogen, 
 as well as towards chlorine and oxygen, carbon is quadri- 
 valent, and towards oxygen it is also bivalent. In this 
 family the maximum oxygen-valence coincides with the 
 hydrogen- valence, while, as has been seen, in Families V, 
 VI, and VII, the oxygen- valence is higher than the hy- 
 drogen-valence, the difference becoming greater from 
 Family V to VII. While the higher oxygen compounds 
 of Family IV are acidic, forming acids which are derived 
 from the normal acid, E(OH) 4 , the lower oxides are not 
 generally acid. The hydrogen compounds of the general 
 formula MH 4 , of which there are but two, those of car- 
 bon and silicon, have neither acid nor basic properties. 
 Carbon is distinguished by the large number of the 
 compounds into which it enters, all of which are more or 
 less closely related to a comparatively small number of 
 fundamental forms. Silicon also forms a large number 
 of compounds, as we shall see ; but these are of a differ- 
 ent kind from those obtained from carbon. 
 
 Occurrence of Carbon. In general, substances which 
 are obtained from the vegetable or animal kingdom blacken 
 when heated to a sufficiently high temperature, and after- 
 wards, if they are heated in the air, they burn up, as we 
 say. When we consider the great variety of substances 
 found in living things, it certainly appears remarkable 
 that nearly all have this property in common. It is due 
 to the fact that nearly all animal and vegetable substances 
 contain the element carbon. When they are heated the 
 
 (357) 
 
358 INORGANIC CHEMISTRY. 
 
 other elements present are first driven off in various 
 forms of combination, while the carbon is the last to go. 
 Hydrogen and oxygen pass off as water ; hydrogen and 
 nitrogen as ammonia ; and much of the carbon also passes 
 off in combination with hydrogen, with hydrogen and 
 oxygen, and with nitrogen and hydrogen. If the heat- 
 ing is carried on in the air, the carbon finally combines 
 with oxygen to form a colorless gas it burns up. Car- 
 bon is the central element of organic nature. There is 
 not a living thing, from the minutest microscopic animal 
 to the mammoth, from the moss to the giant tree, which 
 does not contain this element as an essential constituent. 
 The number of the compounds which it forms is almost 
 infinite, and they present such peculiarities that they are 
 commonly treated of under a separate head, " Organic 
 Chemistry." There is no good reason for this, except the 
 large number of the compounds. For our present pur- 
 pose it will suffice to consider the chemistry of the ele- 
 ment itself, and of a few of its more important simple 
 compounds. 
 
 From what has already been said, it will be seen that 
 the principal form in which carbon occurs in nature is 
 in combination with other elements. It occurs not only 
 in living things, but in their fossil remains, as in coal. 
 Coal-oil, or petroleum, the formation of which is perhaps 
 due to the action of water on metallic carbides, consists 
 of a large number of compounds which contain only car- 
 bon and hydrogen. Most products of plant-life contain 
 the elements carbon, hydrogen, and oxygen. Among the 
 more common of these products may be mentioned sugar, 
 starch, and cellulose. Most products of animal life con- 
 tain carbon, hydrogen, oxygen, and nitrogen. Among 
 them may be mentioned albumen, fibrin, casein, etc. 
 Carbon occurs in the air in the form of carbon dioxide. 
 It also occurs in the form of salts of carbonic acid ; the 
 carbonates, which are very widely distributed, forming 
 whole mountain ranges. Limestone, marble, and chalk 
 are varieties of calcium carbonate. 
 
 Uncombined, the element occurs pure in two very dif- 
 
DIAMONDGRAPHITE. 359 
 
 ferent forms in nature : (1) As diamond ; and (2) as 
 graphite, or plumbago. 
 
 Diamond. The diamond occurs in but few places 
 on the earth, and practically nothing is known as to the 
 conditions which gave rise to its formation. The cele- 
 brated diamond beds are in India, Borneo, Brazil, and 
 South Africa. When found, diamonds are generally 
 covered with an opaque layer, which must be removed 
 before its beautiful properties are apparent. The crys- 
 tals are sometimes regular octahedrons, though usually 
 they are somewhat more complicated, and the faces are 
 frequently curved. It is the hardest substance known. 
 For use as a gem it must be cut and polished. 
 The object in view is to bring out as strikingly as possi- 
 ble its brilliancy by exposing the faces favorably to 
 the action of the' light. If heated to a very high tem- 
 perature without access of air, it swells up and is 
 converted into a black mass resembling coke. The 
 change takes place without loss in weight. Heated to a 
 high temperature in oxygen, it burns up, yielding only 
 carbon dioxide. It is insoluble in all known liquids at 
 ordinary temperatures. It dissolves, however, in molten 
 cast iron and in some other molten metals. Small 
 diamonds have recently (1897) been made by Moissan by 
 dissolving carbon in cast iron with the aid of an electric 
 furnace, and suddenly cooling the mass. Under these 
 conditions the carbon is under great pressure. 
 
 Graphite. Graphite, or plumbago, is found in nature 
 in large quantities. Sometimes it is crystallized, but in 
 forms entirely different from those assumed by the dia- 
 mond. It can be prepared artificially by dissolving char- 
 coal in molten iron, from which solution, on cooling, it is 
 partly deposited as graphite. It has a grayish-black 
 color and a metallic lustre. It is quite soft, leaving a 
 leaden-gray mark on paper when drawn across it, and it 
 is hence used in the manufacture of so-called lead pencils. 
 It is sometimes called black-lead. When heated without 
 access of air it remains unchanged. Heated to a very high 
 temperature in the air, or in oxygen, it burns up, forming 
 
360 INORGANIC CHEMISTRY. 
 
 only carbon dioxide. Like the diamond, it is insoluble 
 in all known liquids at ordinary temperatures. 
 
 Amorphous Carbon. All forms of carbon which are 
 not diamond, nor graphite, are included under the name 
 amorphous carbon. The name signifies simply that it is 
 not crystallized. The most common form of amorphous 
 carbon is ordinary charcoal. 
 
 Charcoal is that form of carbon made by the charring 
 process, which consists simply in heating wood without 
 a sufficient supply of air to effect complete combus- 
 tion. The substance almost exclusively used in the 
 manufacture of charcoal is wood. As has already been 
 stated, wood is made up of a large number of substances, 
 nearly all of which, however, consist of the three elements 
 carbon, hydrogen, and oxygen. One of the chief con- 
 stituents of all kinds of wood is cellulose. Now, when we 
 set fire to a piece of wood, that is to say, when we heat 
 it up to the temperature at which oxygen begins to act 
 on it, it burns, if air is present. Under ordinary cir- 
 cumstances the chemical changes which take place are 
 complex;. but if care is taken, the combustion can be 
 made complete, when all the carbon is converted into car- 
 bon dioxide, and all the hydrogen into water. If, on the 
 other hand, the air is prevented from coming in contact 
 with the wood, as by heating it in a closed vessel, or if it 
 is prevented from coming in contact with it sufficiently 
 to effect complete combustion, the hydrogen is given 
 off partly as water and partly in the form of volatile 
 products containing carbon and oxygen, as wood spirits 
 or methyl alcohol, pyroligneous or acetic acid, acetone, etc. 
 The carbon, however, is for the most part left behind 
 as charcoal, as there is not enough oxygen to convert 
 it into carbon dioxide. Such a process as that just de- 
 scribed, when carried on in closed vessels, is known as 
 destructive distillation or dry distillation. It is also known 
 as the charring process. It is a complex example of a 
 kind of change which we have already had to deal with. 
 Whenever chemical compounds are heated the constitu- 
 ents tend to arrange themselves in forms which are stable 
 at the higher temperature. Sulphites become sulphates ; 
 
CHARCOAL. 361 
 
 phosphites become phosphates; chlorates become per- 
 chlorates ; ammonium salts break down into the acids and 
 ammonia ; ammonium nitrite is decomposed into nitro- 
 gen and water ; ammonium nitrate yields nitrous oxide 
 and water ; primary phosphates yield metaphosphates ; 
 secondary phosphates yield pyrophosphates, etc., etc. 
 Carbon compounds are, in general, more sensitive to 
 the influence of heat than the compounds of other 
 elements, and all are decomposed even at comparatively 
 low temperatures. 
 
 The above statements will make it possible to under- 
 stand the working of a charcoal-kiln. This consists es- 
 sentially of a pile of wood arranged to leave spaces be- 
 tween the pieces, and covered with some rough material 
 through which the air will not pass easily, as, for exam- 
 ple, a mixture of powdered charcoal, turf, and earth. 
 Small openings are left in this covering, so that after the 
 wood is kindled it will continue to burn slowly. The 
 process is sometimes carried on in structures of brick- 
 work with the necessary number of small openings in the 
 walls. The changes above mentioned take place, the 
 gases or volatile substances passing out of the top of the 
 kiln, and appearing as a dense cloud. In due time 
 the holes through which the air gains access to the 
 wood, thus making the burning possible, are stopped up, 
 and the burning ceases. Charcoal, which is impure 
 amorphous carbon, is left behind. As wood always 
 contains some incombustible substances in small quan- 
 tity, these are, of course, found in the charcoal. When 
 the wood or charcoal is burned, these substances remain 
 behind as the ash. 
 
 Ordinary charcoal is a black, comparatively soft sub- 
 stance. It burns in the air, though not easily, unless 
 the gases which are formed are constantly removed and 
 fresh air is supplied, conditions which are met by a 
 good draught, or by blowing upon the fire with a bellows. 
 It burns readily in oxygen. The product of the com- 
 bustion in the air and in oxygen, when the conditions are 
 favorable, is carbon dioxide, CO 2 . In the air, when the 
 draught is bad, another compound of carbon and oxygen, 
 
362 INORGANIC CHEMISTRY. 
 
 carbon monoxide, CO, is formed. Heated without access 
 of air, charcoal remains unchanged. Charcoal is insolu- 
 ble in liquids generally, though it is soluble in molten 
 iron, and it crystallizes from the solution, as we have 
 seen, in the form of graphite, and sometimes as diamond. 
 
 Coke. Besides wood charcoal, there are other forms 
 of amorphous carbon, which are manufactured for special 
 purposes, or are formed in processes carried on for the 
 sake of other products. Coke is a form of amorphous 
 carbon which is made by heating ordinary gas-coal with- 
 out access of air, as is done on the large scale in the 
 manufacture of illuminating gas. Coke bears to coal 
 much the same relation that charcoal bears to wood. 
 
 Lamp-black is a very finely divided form of charcoal 
 which is deposited on cold objects placed in the flames 
 of burning oils. The oils consist almost exclusively of 
 carbon and hydrogen. When burned in the air they 
 yield carbon dioxide and water. If the flame is cooled 
 down by any means, or if the supply of air is partly cut 
 off, the carbon is not completely burned, the flame 
 "smokes," as we say, and deposits soot. This process 
 is chemically analogous to the deposit of metallic arsenic 
 from a flame of arsine, to the deposit of sulphur from 
 a flame of hydrogen sulphide, and that of phosphorus 
 from a flame of phosphine, when these gases are burned 
 in a supply of air insufficient to effect complete combus- 
 tion of both constituents. The soot obtained from the 
 flames of burning oils is made up largely of fine particles 
 of carbon, though some of the unchanged oils are con- 
 tained in it. It is used in the manufacture of printing- 
 ink. As carbon is acted upon directly by very few sub- 
 stances, and is not soluble, it is almost impossible to 
 destroy the color of printing-ink without destroying the 
 material upon which it is impressed. 
 
 Bone-Hack, or Animal Charcoal, is a form of amorphous 
 carbon which is made by charring bones. Bones consist 
 of about one third organic matter and two thirds incom- 
 bustible matter, mostly calcium phosphate. When 
 charred, the organic matter undergoes the changes 
 briefly described under the head of Charcoal, while the 
 
CHARCOAL. 363 
 
 incombustible constituents remain unchanged. As the 
 organic matter is distributed through the substance of 
 the bones the charcoal obtained in this way is in a very 
 fine state of division, but it is mixed with several times 
 its own weight of mineral matter. In order to remove 
 the latter the bone-black must be treated with hydro- 
 chloric acid, and afterwards thoroughly washed with 
 water. An efficient variety of animal charcoal is made, 
 further, by mixing blood with sodium carbonate, char- 
 ring, and afterwards dissolving out the sodium carbon- 
 ate with water. 
 
 Bone-black and wood-charcoal are very porous, and 
 have the power to absorb gases. When placed in air 
 containing bad-smelling gases these are absorbed, and 
 the air is thus to some extent purified. When water con- 
 taining disagreeable substances is treated with charcoal, 
 these are wholly or partly absorbed, and the water is im- 
 proved. Charcoal-filters are therefore extensively used. 
 A charcoal-filter to be efficient should be of good size, 
 and from time to time the charcoal should be taken out 
 and renewed. The small filters which are screwed into 
 faucets are of little value, as the charcoal soon becomes 
 charged with the objectionable material which is pres- 
 ent in the water, and is then a source of contamination 
 rather than a means of purification. The power of 
 charcoal to absorb gases depends upon its porosity. 
 That from some varieties of wood is more porous than 
 that from other varieties. Box-wood charcoal has 
 been shown to absorb 90 times its own volume of am- 
 monia gas, 35 times its volume of carbon dioxide, and 
 nearly twice its volume of hydrogen. Charcoal from 
 cocoa-nut wood absorbs 172 times its volume of am- 
 monia, and 68 times its volume of carbon dioxide. 
 
 Some coloring matters can be removed from liquids 
 by passing the liquids through bone-black filters. On the 
 large scale, this fact is taken advantage of in the refining 
 of sugar. The solution of sugar first obtained from the 
 cane or beet is highly colored ; and, if it were evapo- 
 rated, the sugar deposited from it would be dark-colored. 
 If, however, the solution is first passed through bone- 
 
364 INORGANIC CHEMISTRY. 
 
 black filters, the color is removed, and now, on evaporat- 
 ing, white sugar is deposited. In the laboratory con- 
 stant use is made of this method for decolorizing liquids. 
 The action can easily be shown by adding a little bone- 
 black to a solution containing some litmus or indigo. 
 If the solution is digested for a short time with the 
 bone-black, and then passed through a filter, it will be 
 found that the coloring matter is removed. 
 
 Charcoal does not undergo decay in the air or under 
 water nearly as readily as wood. That is another way 
 of stating the chemical fact that the substances of which 
 wood is made up are more susceptible to the action of 
 other chemical substances than charcoal is. We have 
 one good illustration of this, indeed, in the relative ease 
 with which charcoal and wood burn in the air. Piles 
 which are driven below the surface of water are some- 
 times charred to protect them from the action of those 
 substances which cause decay. 
 
 Coal. Under this head are included a great many 
 kinds of impure amorphous carbon which occur in na- 
 ture. Although we might distinguish between an almost 
 infinite number of kinds of coal, for ordinary purposes 
 they are divided into hard and soft coals, or anthracite 
 and bituminous coals. Then there are substances more 
 nearly allied to wood called lignite,, and those which rep- 
 resent a very early stage in the process of coal-forma- 
 tion, viz., peat. A close examination 01 all these varie- 
 ties has shown that they have been formed by the 
 gradual decomposition of vegetable matter in an insuf- 
 ficient supply of air. The process has been going on 
 for ages. Sometimes the substances have, at the same 
 time, been subjected to great pressure; ac can be seen 
 from the position in which they occur in the earth. The 
 products in the earlier stages of the coal-forming pro- 
 cess are more closely allied to wood than those in the 
 later stages. All forms of coal contain other substances 
 in addition to the carbon. The soft coals are particu-= 
 larly rich in other substances. When heated they give 
 off a mixture of gases and the vapors of volatile liquids. 
 The gases are, for the most part, useful for illuminating 
 
DIAMOND, GRAPHITE, AND CHARCOAL. 365 
 
 purposes. The liquids form a black, tarry mass known 
 as coal-tar, from which many valuable compounds of car- 
 bon are obtained. The gases are passed through water 
 for the purpose of removing certain impurities. This 
 water absorbs ammonia, and forms the ammoniacal liquor 
 of the gas-works, which, as has been stated, is the prin- 
 cipal source of ammonia. 
 
 Diamond, Graphite, and Charcoal are Different Forms of 
 the Element Carbon. We have seen that oxygen presents 
 itself in two forms ordinary oxygen and ozone. Ozone 
 is made from oxygen, and oxygen from ozone, without any 
 increase or decrease in weight ; and the compounds ob- 
 tained by the combination of other elements with oxygen 
 are identical with those obtained by the combination of 
 the same elements with ozone. So also there are several 
 varieties of sulphur, two of which crystallize in' different 
 forms. There are, further, three or four different modi- 
 fications of the element phosphorus, and these differ 
 from one another in a very marked way. The explana- 
 tion of the difference between oxygen and ozone is that 
 the molecule of the former is made up of two atoms, 
 while that of ozone is made up of three, which are in 
 a state of unstable equilibrium. This explanation is 
 reached through a study of the specific gravity of the 
 two gases. At present no satisfactory explanation can 
 be given of the difference between the varieties of phos- 
 phorus and between the varieties of sulphur. It will 
 probably be shown to be due to the way in which the 
 atoms are grouped together in the molecules, and also 
 to the way in which the molecules are grouped together 
 to form the masses. Carbon, as we have seen, occurs in 
 three distinct forms It is difficult to conceive that the 
 black, porous charcoal, and the dull, gray, soft graphite 
 are chemically identical with the hard, transparent, 
 brilliant diamond. Yet this is undoubtedly the case, as 
 can be shown by a very simple experiment. Each of 
 the substances when burned in oxygen yields carbon 
 dioxide. Now ? the composition of carbon dioxide is 
 known, so that, if the weight of the carbon dioxide 
 formed in a given experiment is known, the weight of the 
 
366 INORG'ANIC CHEMISTRY. 
 
 carbon in it is also known. When a gram of pure char- 
 coal is burned it yields 3f grams carbon dioxide, and in 
 this quantity of carbon dioxide there is contained ex- 
 actly one gram of carbon. Further, when a gram of 
 graphite is burned the same weight (3f grams) of carbon 
 dioxide is obtained as in the case of charcoal ; and the 
 same thing is true of diamond. It follows from these 
 facts that the three forms of matter known as char- 
 coal, graphite, and diamond consist only of the element 
 carbon. The explanation of the diffei ^nce is not known, 
 but, as in the cases of phosphorus an \ sulphur, it will 
 probably be found to be in the differeii ways in which 
 the atoms are arranged in the molecule, and the mole- 
 cules in the masses. 
 
 Notwithstanding the marked differences in their ap- 
 pearance and in many of their physical properties, the 
 three forms of carbon have, as we have seen, some prop- 
 erties in common. They are insoluble in all known 
 liquids at ordinary temperatures. They are tasteless, 
 inodorous, and infusible. When heated without access 
 of air they remain unchanged, unless the temperature is 
 very high, when the diamond swells up and is converted 
 into a mass resembling coke a change which is con- 
 nected with a rearrangement of the particles in an irreg- 
 ular way, so that the substances cease to be crystalline, 
 or become amorphous. 
 
 Chemical Conduct of Carbon. At ordinary tempera- 
 tures carbon is an inactive element. If it is left in contact 
 with any one of the elements, no chemical change takes 
 place. It will not combine with any of them unless the tem- 
 perature is raised. At higher temperatures, however, it 
 combines with several of them with great ease, especially 
 with oxygen. Under proper conditions it combines also 
 with nitrogen, with hydrogen, with sulphur, and with 
 many other elements. It combines with oxygen either 
 directly, as when it burns in the air or in oxygen ; or it 
 abstracts oxygen from some of the oxides. The direct 
 combination of oxygen and carbon has already been seen 
 in the burning of charcoal in oxygen, and is familiar to 
 every one in the fire in a charcoal furnace. That car- 
 
CHEMICAL CONDUCT OF CARBON. 367 
 
 ton dioxide is the product formed can be shown by 
 passing the gas through lime-water or baryta-water, 
 when insoluble calcium or barium carbonate will be 
 thrown down. The reason why lime-water or baryta- 
 water is used is simply that an insoluble compound is 
 formed, and this can be seen, and it can be separated 
 from the liquid and examined. The reaction which 
 takes place is represented thus : 
 
 Ca(OH) 3 + CO 2 = CaCO 3 + H 2 O; 
 
 Calcium Carbon Calcium 
 
 hydroxide dioxide carbonate 
 
 Ba(OH) 2 + C0 2 = BaC0 3 +H 2 O. 
 
 Barium Carbon Barium 
 
 hydroxide dioxide carbonate 
 
 No other common gas acts in this way on these solu- 
 tions. Hence, when, under ordinary circumstances, a 
 gas is passed into lime-water and an insoluble compound 
 is formed, we may conclude that the gas is carbon diox- 
 ide, though this conclusion may require further proof. 
 
 The abstraction of oxygen from compounds by means 
 of carbon may be illustrated in a number of ways. Thus, 
 when powdered copper oxide, CuO, is mixed with pow- 
 dered charcoal, and the mixture heated in a tube, car- 
 bon dioxide is given off, and can be detected as in the 
 last experiment mentioned. Copper is left behind, and, 
 if the proportions are properly selected, all the carbon 
 will pass off as carbon dioxide, and only the copper be 
 left behind : 
 
 2CuO + C = 2Cu + CO,. 
 
 In a similar way, arsenious oxide, As^O 3 , gives up its 
 oxygen to carbon. This fact furnishes indeed a delicate 
 method for the detection of the substance. If a little 
 is placed in th'3 bottom of a small tube, and above it a 
 small piece of charcoal, then when heat is applied the 
 arsenious oxide sublimes, and as its vapor passes the 
 heated charcoal the oxygen is abstracted, and the ele- 
 ment arsenic, being also somewhat volatile, is deposited 
 
368 INORGANIC CHEMISTRY. 
 
 just above the charcoal in the form of a lustrous mirror 
 on the walls of the tube. The reaction is 
 
 As has already been explained, the abstraction of oxy- 
 gen from a compound is known as reduction. Hence, 
 carbon is called a reducing agent. It is indeed the re- 
 ducing agent which is most extensively used in the arts. 
 Its chief use is in extracting metals from their ores, 
 which are the forms in which they occur in nature. 
 Thus, iron does not occur in nature as iron, but in com- 
 bination with other elements, especially with oxygen. 
 In order to get the metal the ore must be reduced, or, 
 in other words, the oxygen must be extracted. This is 
 invariably accomplished by heating it with some form 
 of carbon, either coke or charcoal. 
 
 The elements chlorine, oxygen, nitrogen, and hydro- 
 gen being gases, and the products formed when the first 
 three combine with hydrogen being also gaseous or con- 
 vertible into vapor, it is a comparatively easy matter to 
 study the relations between the volumes of the combin- 
 ing gases and the volumes of the products formed. It is, 
 however, impossible to determine the ratio between the 
 volume of carbon gas and that of other gases with which 
 it combines. 
 
 Compounds of Carbon with Hydrogen, or Hydrocarbons. 
 Conditions under which Hydrocarbons are Formed. When 
 the carbon pencils connected with a powerful battery, 
 as in the production of the electric arc-light, are sur- 
 rounded by an atmosphere of hydrogen the two elements 
 combine to some extent to form the compound acetylene, 
 C,H a . When organic matter undergoes decomposition 
 without free access of air, as for example under water, 
 the carbon compounds are reduced to the final product 
 known as marsh-gas or methane, CH 4 , just as compounds 
 containing nitrogen yield ammonia. Compounds of 
 various metals with carbon, known as carbides, yield 
 hydrocarbons, such as acetylene and marsh-gas, when 
 treated with water. The compounds which make up 
 petroleum are hydrocarbons which have probably been 
 
NUMBER OF HYDROCARBONS. 369 
 
 formed, either by the action of water on metallic carbides 
 in the interior of the earth, or by decomposition of organic 
 matter without free access of air. Finally, when wood 
 or coal is heated, hydrocarbons are given off, and com- 
 pounds of this kind are therefore contained in coal-gas. 
 
 Number of Hydrocarbons. The number of hydrocar- 
 bons is very great, and new ones are constantly being 
 made. The fact that carbon is distinguished for the 
 large number of its compounds has already been men- 
 tioned. The simplest of these are the hydrocarbons. 
 It is safe to say that there are as many as two hun- 
 dred hydrocarbons known. Fortunately, however, most 
 of these have been found to bear comparatively simple 
 relations to one another, and therefore, though the 
 number is large and the variety great, their study is not 
 as difficult as one would be inclined to think. 
 
 Petroleum is an oily liquid found in many places in the 
 earth in large quantity, and escapes when a cavity in 
 which it is contained is punctured. In the earth it 
 contains both gases and liquids. When it is brought 
 into the air the gases and the liquids which are vola- 
 tile at the ordinary temperature are given off. There 
 are several gases given off, and a large number of 
 liquids left behind. The simplest gas corresponds to 
 the formula CH 4 , the next to C 2 H 6 , the next to C 3 H 8 , the 
 next to C 4 H 10 . An examination of the liquid has shown 
 that it contains other hydrocarbons of the formulas 
 C 5 H 12 , C 6 H 14 , C 7 H 16 , C 8 H 18 , etc. It will be seen that these 
 compounds bear a simple relation to one another, as far 
 as composition is concerned. Arranging them in a series, 
 tne first eight members are 
 
 CH 4 , Methane, or Marsh-gas ; 
 
 C 2 H 6 , Ethane; 
 
 C 3 H 8 , Propane ; 
 
 C 4 H 10 , Butane ; 
 
 C 6 H 12 , Pentane ; 
 
 C 6 H 14 , Hexane ; 
 
 C 7 H 16 , Heptane ; 
 
 C 8 H 18 , Octane. 
 
370 INORGANIC CHEMISTRY. 
 
 Homology, Homologous Series. In the above series the 
 first member differs from the second by CH 2 ; and there 
 is also this same difference between the second and 
 third, the third and fourth, and in general between any 
 two consecutive members in the series. This relation is 
 known as homology, and such a series is known as an 
 homologous series. Carbon is distinguished from all 
 other elements by its power to form homologous series. 
 Other elements form similar series, but the homology is 
 not of the same kind as that which is met with among 
 carbon compounds. Thus the series of chlorine acids, 
 and the similar series of acids of nitrogen and sulphur, 
 are homologous series, in which the constant difference 
 between any two consecutive members is represented by 
 an atom of oxygen : 
 
 HC1O H 2 SO 2 HNO 
 
 HC10 2 H 2 S0 3 HNO 2 
 
 HC10 3 H 2 SO 4 HN0 3 . 
 HC10 4 
 
 These series, are, however, much less extensive than the 
 homologous series of compounds of carbon. 
 
 Cause of the Homology among Compounds of Carbon. 
 The explanation of the homology observed between the 
 compounds of carbon is founded on the view that carbon 
 is quadrivalent, and that it has the power to unite with 
 itself in chains. The quadrivalence is shown in the com- 
 pounds CH 4 , CC1 4 , CHC1 3 , CO 2 , etc. When marsh-gas, 
 CH 4 , is treated with chlorine the first product of the ac- 
 tion is chlor-methane, CH 3 C1, which according to the 
 prevailing views is marsh-gas in whose molecule one 
 atom of hydrogen has been replaced by chlorine. The 
 
 H 
 
 structure of this compound is represented thus, H-C-C1, 
 
 * H 
 H 
 if that of marsh-gas is represented in this way, H-C-H. 
 
OTHER SERIES OF HYDROCARBONS. 371 
 
 Now, when chlor-methane is treated with sodium the 
 chlorine is extracted, and a compound of the formula 
 C a H 6 is formed : 
 
 2CH 3 C1 + 2Na = C a H fl + 2NaCl. 
 
 It appears that, the chlorine being extracted from the 
 compound, the residues of the composition CH 3 unite 
 in pairs to form the compound C 2 H 6 , which is ethane, or 
 the second member of the series of hydrocarbons above 
 given. The simplest explanation of the facts stated is 
 that the carbon atoms unite by means of the bonds or 
 valences left free when the chlorine is extracted. The 
 residues after the extraction of the chlorine may be rep- 
 
 H 
 
 resented by the formula H-C-; when two of these 
 
 i 
 
 unite, the resulting compound will have the structure 
 
 H H 
 
 represented by the formula H-C-C-H. In a similar 
 
 ii ii 
 
 way the relation between the other members of the 
 series can be explained, and the explanation is in perfect 
 accordance with a large number of facts. 
 
 Other Series of Hydrocarbons. Besides the series above 
 mentioned, which, as its simplest member is marsh-gas, 
 is known as the marsh-gas series, there are other homolo- 
 gous series of hydrocarbons. There is one beginning 
 with ethylene, or olefiant gas, C 2 H 4 , examples of which 
 are 
 
 Ethylene, C 2 H 4 ; 
 
 Propylene, C 3 H 6 ; 
 
 Butylene, C 4 H 8 . 
 
 There is another beginning with acetylene, C a H a , ex- 
 amples of which are 
 
 Acetylene, C 2 H a ; 
 Allylene, C 3 H 4 . 
 
372 INORGANIC CHEMISTRY. 
 
 Another series begins with benzene, C 6 H a . Some of 
 the members of this series are 
 
 Benzene, C 6 H 
 Toluene, C,H 
 Xylene, C 8 H 
 
 These are the hydrocarbons which are obtained from 
 coal-tar. 
 
 The relations between these hydrocarbons and those 
 of the marsh-gas series have been extensively studied, 
 and a great deal of light has been thrown upon the sub- 
 ject. It would, however, lead too far to take up this 
 subject here. A word may be said, however, in regard 
 to the relations believed to exist between the hydrocar- 
 bons of the ethylene and the acetylene series, and those 
 of the marsh-gas series. It is believed that in ethylene 
 the two carbon atoms are united in a different way from 
 that in ethane. The condition, whatever it may be, is 
 thought to be similar to that which exists in the mole- 
 cule of a compound consisting of two bivalent atoms, 
 as, for example, calcium oxide. The condition is called 
 double union, and it is represented by a double line as 
 in the formulas Ca=O, H 2 C=CH 2 , etc. Whatever this 
 condition may be, it carries with it the power to com- 
 bine with other atoms. Thus, when ethylene is treated 
 with chlorine it takes up two atoms, and is converted 
 into dichlorethane, C 2 H 4 C1 2 , the double union being de- 
 stroyed and single union existing in the resulting com- 
 pound as in ethane. So also, when calcium oxide, 
 Ca=O, is treated with water, it is converted into the hy- 
 droxide, in which the condition of double union does not 
 exist : 
 
 Ca OH Ca-OH 
 
 ii + i =i ; 
 O H OH 
 
 CH 2 Cl CH 2 C1 
 
 II +1=1 
 CH 2 Cl CH a Cl 
 
 Similar reasons have led to the conclusion that in 
 acetylene, C a H 2 , the carbon atoms are held together in 
 
MABSH-GAS, METHANE, FIRE-DAMP. 373 
 
 a -different way from that in ethane, and that in ethy- 
 lene. This is believed to be similar to the kind of union 
 which exists in a molecule consisting of two trivalent 
 atoms, as in the compound boron nitride, B=N, a con- 
 dition which is called triple union or triple linkage. This 
 view is expressed by the formula HC=CH. Wherever 
 this condition exists we find the power to take up four 
 univalent atoms, the compound thus becoming saturated, 
 as we say. Thus, acetylene can take up four atoms of 
 chlorine, or two- of hydrogen and two of chlorine, form- 
 ing in the former case tetra-chlor-ethane, and in the latter 
 di-chl or- ethane : 
 
 HteCH + 4C1 = C1 2 HC-CHC1 2 (C 2 H 2 C1 4 ) ; 
 HC=CH + 2HC1 = C1H 2 C-CH 2 C1(C 2 H 4 C1 2 ). 
 
 Marsh-gas, Methane, Fire-damp, CH 4 . Marsh-gas is 
 found in nature in petroleum, and is given off when the 
 oil is taken out of the earth and the pressure removed. 
 It is formed, as the name implies, in marshes, as the 
 product of a reducing process. Vegetable matter is com- 
 posed essentially of carbon, hydrogen, and oxygen. 
 When it undergoes decomposition in the air in a free 
 supply of oxygen, the final products formed are carbon 
 dioxide and water. When the decomposition takes 
 place without access of oxygen, as under water, marsh- 
 gas, which is a reduction-product, is formed. The gas 
 can be made in the laboratory by passing a mixture of 
 hydrogen sulphide and the vapor of carbon disulphide 
 over heated copper, and also by the action of water on 
 several metallic carbides, more especially aluminium 
 carbide, C 3 A1 4 : 
 
 C 3 A1 4 + 12H 2 = 3CH 4 + 4A1(OH),. 
 
 The gas is met with in coal-mines, and is known to 
 the miners as Jire-damp, damp being the general name 
 applied to a gas, and the name fire-damp meaning a gas 
 that burns. To prepare it in the laboratory it is most 
 convenient to heat a mixture of sodium acetate and 
 
374 INORGANIC CHEMISTRY. 
 
 quick -lime. The change which takes place will be most 
 readily understood by regarding it as a simple decom- 
 position of acetic acid. Acetic acid has the formula 
 C 2 H 4 O 2 . When heated alone it boils, and does not suffer 
 decomposition. If it is converted into a salt, and heated 
 in the presence of a base, it breaks down into marsh-gas 
 and carbon dioxide : 
 
 The carbon dioxide, which with bases forms salts, does 
 not pass off, but remains behind in the form of a salt of 
 carbonic acid. 
 
 Marsh-gas is a colorless, transparent, tasteless, inodor- 
 ous gas. It is slightly soluble in water, and burns, 
 forming carbon dioxide and water. When mixed with 
 air the mixture explodes if a flame or spark comes in 
 contact with it. This is one of the causes of the explo- 
 sions which so frequently occur in coal-mines. To pre- 
 vent these explosions a special lamp was invented by Sir 
 Humphry Davy, which is known as the safety -lamp. The 
 simple principles involved in its construction will be ex- 
 plained when the subject of flame is taken up. 
 
 Ethylene, defiant Gas, C 2 H 4 . This liydrocarbon is 
 formed by heating a mixture of ordinary alcohol and 
 concentrated sulphuric acid. The reaction is represented 
 thus : 
 
 C.H.O - H.O + C Q H, 
 
 Alcohol Ethylene 
 
 Ethylene is a colorless gas, which can be condensed to 
 a liquid. It burns with a luminous flame, and forms an 
 explosive mixture with oxygen. 
 
 Acetylene, C 2 H 2 . Acetylene is formed when a current 
 of hydrogen is passed between carbon poles, which are 
 incandescent in consequence of the passage of a power- 
 ful electric current. In this case carbon and hydrogen 
 combine directly. It is formed also when the flame of 
 an ordinary laboratory gas-burner (Bunsen burner) 
 " strikes back," or burns at the base without a free sup- 
 
SIMPLER COMPOUNDS OF CARBON. 375 
 
 ply of air. It is most easily obtained by treating certain 
 metallic carbides, especially calcium carbide, with water : 
 
 C 2 Ca + H 2 = C 2 H 8 + CaO. 
 
 Its odor is unpleasant. It burns with a luminous, 
 smoky flame. 
 
 Simpler Compounds of Carbon with, the Members of the 
 Chlorine Group. When chlorine acts upon a hydrocar- 
 bon it generally takes the place of the hydrogen, atom 
 for atom. Thus, when it acts upon marsh-gas, the 
 following reactions take place ; 
 
 CH 4 + C1 2 = CH 3 C1 + HC1 ; 
 CH 3 C1 + C1 2 = CH 2 C1 2 + HC1 ; 
 CH 2 C1 2 + 01, = CHC1 3 + HC1 ; 
 CHC1 3 + 01, = CC1 4 + HC1. 
 
 The four products are known respectively as mono-cUor- 
 methane, di-cJdor-methane, tri-chlor-metharie, and tetra-chlor- 
 methane, or carbon tetrachloride. By treating these com- 
 pounds with nascent hydrogen they can all be converted 
 back again into marsh-gas. The fact that the hydrogen 
 in marsh-gas can be replaced in four steps, one fourth 
 of the hydrogen being replaced at each step, furnishes a 
 strong confirmation of the correctness of the view ex- 
 pressed by the formula CH 4 , which signifies that in the 
 molecule of marsh-gas there are four atoms of hydrogen. 
 The most important of the four compounds is the third, 
 tri-chlor-methane or chloroform. While chloroform can 
 be made by treating marsh-gas with chlorine, it is much 
 more easily obtained in other ways, as, for example, by 
 treating alcohol or acetone with bleaching-powder. 
 Without a study of the relations which exist between 
 several classes of compounds of carbon these reactions 
 cannot well be explained, and their study, as well as that 
 of chloroform, had better be postponed until the subject 
 of Organic Chemistry is taken up systematically. Cor- 
 responding to chloroform there are bromine and iodine 
 derivatives, known as bromoform, CHBr 3 , and iodoform, 
 CHI 3 . 
 
CHAPTER XX. 
 
 SIMPLER COMPOUNDS OF CARBON WITH OXYGEN, AND 
 WITH OXYGEN AND HYDROGEN. 
 
 General. The final product of oxidation of carbon is 
 carbon dioxide, and the final product of reduction is 
 marsh-gas, but between these two limits there are a 
 number of interesting derivatives, just as there are a 
 number of compounds of sulphur between hydrogen sul- 
 phide and sulphuric acid ; a number of compounds of 
 nitrogen between ammonia and nitric acid ; and a number 
 of compounds of phosphorus between phosphine and 
 phosphoric acid. We have seen that in the cases men- 
 tioned two kinds of change are brought about by oxida- 
 tion : (1) Owing to the fact that the valence of chlo- 
 rine, sulphur, nitrogen, and phosphorus towards oxygen 
 is greater than towards hydrogen, the act of oxidation 
 involves the addition of oxygen to the element ; (2) hy- 
 droxyl appears to take the place of the hydrogen atoms 
 one by one. In the case of carbon, the valence towards 
 hydrogen and oxygen being the same, only the latter 
 kind of change takes place. 
 
 Relations between the Compounds of Carbon with Hy- 
 drogen and Oxygen. In order that the relations between 
 the simpler compounds of carbon with hydrogen and 
 oxygen may be made clear, it will be of assistance to 
 compare the oxidation of hydrogen sulphide, ammonia, 
 phosphine, and methane. In the cases of hydrogen sul- 
 phide and ammonia the oxidation appears to take place 
 as represented below : 
 
 Hydrogen Unknown Sulphurous Sulphuric 
 
 Sulphide acid acid 
 
 (376) 
 
RELATIONS BETWEEN COMPOUNDS OF CARBON. 377 
 
 (OH (OH 
 
 N^OH, ON -{OH. 
 (OH (OH 
 
 Yields hypo- Yields nitrous Yields nitric 
 nitrous acid acid acid 
 
 The last three products, if formed, break down, losing 
 water, and forming respectively hyponitrous, nitrous, and 
 nitric acids. The change to hyponitrous acid does not 
 appear to be of an altogether simple kind. The changes 
 to nitrous and nitric acids, however, are apparently of a 
 kind which we are constantly meeting with, as has al- 
 ready been pointed out (see pp. 261-265). 
 
 It is possible that the changes involved in the gradual 
 oxidation of ammonia take place primarily just as in the 
 oxidation of sulphur, the nitrogen first becoming satu- 
 rated. According to this view the changes should be 
 represented as follows : 
 
 H (H 
 
 H, ON-{ H, 
 
 H H 
 
 If nitrous acid were formed by the breaking down of 
 
 (OH 
 
 the compound ON < OH , its structure would probably 
 
 be that represented by the formula N < O , or O 2 NH, 
 
 ( H 
 
 the hydrogen being in combination with nitrogen and not 
 with oxygen. Some facts seem to show that this view is 
 probable. While these changes cannot be followed very 
 well in the case of the compounds of nitrogen, and there 
 is, therefore, considerable speculation in what has just 
 been said regarding them, the case of phosphorus is much 
 clearer, as has already been shown. Here, starting with 
 phosphine, the changes are apparently correctly repre- 
 sented by the following formulas : 
 
378 
 
 INORGANIC CHEMISTRY. 
 
 (H 
 
 P^H, 
 (H 
 
 Phosphine 
 
 (H (OH 
 
 OP^H, OP^H , 
 (H (H 
 
 (OH (OH 
 
 OP^OH, OP^OH. 
 (H (OH 
 
 Unknown 
 
 Hypophosphor- Phosphorous 
 ous acid acid 
 
 Phosphoric 
 acid 
 
 With methane the changes effected by oxidation are 
 apparently perfectly analogous to those considered. We 
 should expect to find the following : 
 
 OH 
 OH 
 OH 
 OH 
 
 But the tendency of the compounds containing two 
 hydroxyl groups to break down, yielding water as one of 
 the products, is as marked as in the compounds of nitro- 
 gen. Consequently the products 3, 4, and 5 do not exist 
 in the free state. They break down as represented in 
 the following equations : 
 
 (H 
 CJH + H,0; 
 
 = C 
 
 H,0; 
 
 = CJg + 2H,0. 
 
 The products actually obtained, therefore, are as 
 follow : 
 
RELATIONS BETWEEN COMPOUNDS OF CARBON. 379 
 
 (O 
 
 |o 
 
 Methane Methyl Formic Formic Carbon 
 
 alcohol aldehyde acfd dioxide 
 
 By oxidation of formic acid we should expect the for- 
 
 (OH 
 mation of a product, C < O . While this is not known, 
 
 (OH 
 
 salts of an acid of this composition are known. It is 
 ordinary carbonic acid, which when set free breaks down 
 into carbon dioxide and water. 
 
 It will be seen from the above considerations that 
 there is a general analogy between the changes which 
 take place in passing by oxidation from the lowest re- 
 duction-products of the elements to their highest oxida- 
 tion-products. 
 
 The intermediate stages have been studied with special 
 care in the case of carbon ; the intermediate products rep- 
 resent classes of compounds vrhich are not met with among 
 any derivatives of the other elements. The intermediate 
 products in the case of sulphur and phosphorus are all 
 acids. One of the intermediate products in the case of 
 nitrogen, hydroxylamine, is basic, while the rest are acid. 
 The first oxidation-product of marsh-gas, methyl alcohol, 
 CH 3 .OH, is somewhat basic, but in some respects differs 
 from ordinary bases. It is the simplest representative 
 of a great class of compounds of carbon known as alco- 
 hols, of which our ordinary alcohol, or spirits of wine, is 
 the best known example. The next product, or formic 
 aldehyde, which has the structure represented by the 
 
 ( H 
 
 formula C 1 H or H 2 C=O, is the simplest representative 
 
 (O 
 
 of another great class of compounds of carbon known as 
 aldehydes. The aldehydes are neither acid nor basic, but 
 are easily converted into acids by oxidation, and into bases 
 
380 .INORGANIC CHEMISTRY. 
 
 by reduction. The third oxidation-product of the formula 
 
 , H H 
 
 I " i 
 
 C < OH or O=C is the simplest representative of a 
 
 ( 6n 
 
 great class of carbon compounds known as the organic 
 acids. It is called formic acid. It would lead too far to 
 pursue this subject now. The relations referred to are 
 studied under the head of Organic Chemistry, or the 
 Chemistry of the Compounds of Carbon. Only the sim- 
 plest oxygen compounds will be taken up here. 
 
 Carbon Dioxide, CO 2 . The principal compound of car- 
 bon and oxygen is carbon dioxide, CO 2 , commonly known 
 as carbonic acid gas. Under the head of The Air atten- 
 tion was called to the fact that this gas is a constant con- 
 stituent of the air, though its relative quantity is small 
 about 3 parts in 10,000. It issues from the earth in many 
 places, particularly in the neighborhood of volcanoes. 
 Many mineral waters contain it in large quantity, promi- 
 nent among which are the waters of Pyrmont, Selters, and 
 the Geyser Spring of Saratoga. In small quantity it is 
 present in all natural waters. In combination with bases 
 it occurs in enormous quantities, particularly in the form 
 of calcium carbonate, CaCO 3 , varieties of which are 
 ordinary limestone, chalk, marble, and calc-spar. Dolo- 
 mite, which forms mountain-ranges, being particularly 
 abundant in the Swiss Alps, is a compound containing 
 calcium carbonate and magnesium carbonate, MgCO 3 . 
 
 Carbon dioxide is constantly formed in many natural 
 processes. Thus, all animals. that breathe in the air give 
 off carbon dioxide from the lungs. That the gases from 
 the lungs contain carbon dioxide can easily be shown by 
 passing them through lime-water, when a precipitate of 
 calcium carbonate is formed. 
 
 That carbon dioxide is formed in the combustion of 
 charcoal and wood has already been shown. In a similar 
 way it can be shown that the gas is formed whenever any 
 of our ordinary combustible substances are burned. 
 From our fires, as from our lungs, and from the lungs of 
 all animals, then, carbon dioxide is constantly given off. 
 
CARBON DIOXIDE. 381 
 
 Further, the natural processes of decay of both vegetable 
 and animal matter tend to convert the carbon of this 
 matter into carbon dioxide, which then finds its way 
 principally into the air. The process of alcoholic fer- 
 mentation, and some other similar processes, also give 
 rise to the formation of carbon dioxide. In all fruit- 
 juices there is contained sugar. When the fruits ripen, 
 fall to the earth, and undergo spontaneous change, the 
 sugar is converted into alcohol and carbon dioxide. We 
 see, thus, that there are many important sources of 
 supply of carbon dioxide, and it will be readily under- 
 stood why the gas should be found everywhere in the air. 
 Preparation. The easiest way to get carbon dioxide 
 not mixed with other substances is by adding an acid to 
 a salt of carbonic acid or a carbonate. In the decompo- 
 sition of the carbonates by other acids we see exemplified 
 the same principle as that which is involved in setting 
 nitric acid free from a nitrate, or hydrochloric acid from 
 sodium chloride by sulphuric acid, and more particularly 
 in the liberation of sulphur dioxide from a sulphite. 
 In all these cases the products are volatile, and there- 
 fore, when a non-volatile acid is added to the salts, 
 decomposition takes place. Sulphites do not yield the 
 corresponding acid, but this breaks down into water^and 
 the anhydride : 
 
 j 2NaN0 3 + H 2 SO 4 = Na 2 SO 4 + 2HNO 8 ; 
 |2NaCl + H a SO 4 = Na a S0 4 + 2HC1. 
 
 ] Na 2 S0 3 + H 2 SO 4 = Na 2 SO 4 + H 2 SO 3 ; 
 |H 2 S0 3 = S0 2 +H 2 0. 
 
 ( Na 2 C0 3 + H S0 4 = Na 2 S0 4 + H 2 CO 3 ; 
 |H 2 C0 3 =C0 2 +H a O. 
 
 Any acid which is not volatile at the ordinary tempera- 
 ture will decompose a carbonate and cause an evolution 
 of carbon dioxide. The action between sodium carbon- 
 ate and hydrochloric acid is represented in this way : 
 
 Na 2 C0 8 + 2HC1 = 2NaCl + CO 3 + H a O ; 
 
382 INORGANIC CHEMISTRY. 
 
 that between nitric acid and sodium carbonate in this 
 way : 
 
 Na a CO 3 + 2HNO 3 = 2NaNO, + CO, + H 2 O. 
 
 For the purpose of preparing carbon dioxide in the 
 laboratory, calcium carbonate, in the form of marble or 
 limestone, and hydrochloric acid are commonly used. 
 The reaction involved is represented thus : 
 
 CaCO 3 + 2HC1 = CaCl 2 + CO 3 + H 2 O. 
 
 The apparatus used is the same as that used in making 
 hydrogen from zinc and sulphuric acid. As the gas is 
 somewhat soluble in water, it is best for ordinary pur- 
 poses to collect it by displacement of air, the vessel being 
 placed with the mouth upward, as the gas is considerably 
 heavier than air. 
 
 Properties. Carbon dioxide is a colorless gas at or- 
 dinary temperatures. When subjected to a low tempera- 
 ture and high pressure it is converted into a liquid. 
 Liquid carbon dioxide is now manufactured on the large 
 scale for use as a fire-extinguisher, and for the purpose 
 of charging liquids with the gas. When some of the 
 liquid is exposed to the air evaporation takes place so 
 rapidly that a great deal of heat is absorbed, and some 
 of the liquid becomes solid. The gas has a slightly acid 
 taste and smell. It is not combustible, nor does it sup- 
 port combustion. It is not combustible for the same 
 reason that water is not : because it already holds in 
 combination all the oxygen it has the power to combine 
 with. Before it can burn again it must first be decom- 
 posed. As regards the statement that it does not support 
 combustion, it should be remarked that this is only rela- 
 tively true. The compound does not easily give up 
 oxygen, but to some substances it does give it up, and 
 some such substances burn in it. For example, the ele- 
 ment potassium, which, as we have seen, has the power 
 to decompose water, has also the power to decompose 
 carbon dioxide if heated in it to a sufficiently high tem- 
 perature, and when the decomposition once begins, it 
 
RESPIRATION. 383 
 
 proceeds with brilliancy, the act being accompanied by 
 a marked evolution of heat and light. Carbon dioxide 
 is much heavier than air, its specific gravity being 1.529. 
 A liter of the gas under standard conditions of tempera- 
 ture and pressure weighs 1.977 grams. It dissolves in 
 water, one volume of water dissolving about one volume 
 of the gas at the ordinary temperature. As is the case 
 with all gases, when the pressure is increased the water 
 dissolves more gas, and when the pressure is removed 
 the gas again escapes. The so-called " soda-water" is 
 simply water charged with carbon dioxide under pressure. 
 The escape of the gas when the water is drawn is famil- 
 iar to every one. The name soda-water has its origin in 
 the fact that the carbon dioxide used in charging the 
 water is frequently made from primary or acid sodium 
 carbonate, NaHCO 3 , which is also called soda or bicar- 
 bonate of soda. 
 
 Relations of Carbon Dioxide to Chemical Energy. 
 Carbon has the power to combine with oxygen, and in so 
 doing a definite quantity of heat is evolved. A kilogram 
 of carbon represents a certain quantity of chemical 
 energy, which we can get from it first in the form of heat, 
 and by transformation, in other forms of energy, as mo- 
 tion, electrical energy, etc. After the kilogram of carbon 
 has been burned it no longer represents the energy it 
 did in the form of carbon. A body of water elevated 
 ten or fifteen feet represents a certain quantity of energy 
 which can be obtained by allowing the water to fall upon 
 the paddles of a water-wheel connected with the machin- 
 ery of a mill. After the water has fallen, however, it no 
 longer has the power to do work, or it has none of the 
 energy which it possessed by virtue of its position. In 
 order that it may again do work it must again be lifted. 
 So, too, in order that the carbon in carbon dioxide may 
 again do work the compound must be decomposed. 
 
 Respiration. It was stated above that carbon dioxide 
 is given off from the lungs just as it is from a fire. It is 
 a waste-product of the processes taking place in the ani- 
 mal body. Just as it cannot support combustion, so also 
 it cannot support respiration. It is not poisonous any 
 
384 TNOEGANIC CHEMISTRY. 
 
 more than water is ; but it is not able to supply the oxy* 
 gen which is needed for breathing purposes, and hence 
 animals die when placed in it. They die by suffocation, 
 very much as they do in drowning. Any considerable 
 increase in the quantity of carbon dioxide in the air above 
 that which is normally present is objectionable, for the 
 reason that it decreases the proportion of oxygen in the 
 air which is breathed. If, however, pure carbon dioxide 
 is introduced into the air, it has been found that as much 
 as 5 per cent may be present without serious results to 
 those who breathe it. In a badly ventilated room in 
 which a number of people are collected, and lights are 
 burning, it is well known that in a short time the air be- 
 comes foul, and bad effects, such as headache, drowsi- 
 ness, etc., are produced on the occupants of the room. 
 These effects appear to be due, not to the carbon 
 dioxide, but largely to other waste-products which are 
 given off from the lungs in the process of breathing. 
 The gases given off from the lungs consist of nitrogen, 
 oxygen, carbon dioxide, and water vapor. Besides these,, 
 however, there are many substances in small quantity, 
 in a finely divided condition, which contain carbon, and 
 are in a state of decomposition. These act as poisons, 
 and they are the chief cause of the bad effects experi- 
 enced in breathing air which is contaminated by the exha- 
 lations from the lungs. As carbon dioxide is given off 
 from the lungs at the same time, the quantity of this 
 gas present is proportional to the quantity of the or- 
 ganic impurities. Hence, by determining the quantity 
 of carbon dioxide it is possible to form an opinion a& 
 to whether the air of a room occupied by human beings 
 is fit for use or not. 
 
 As carbon dioxide is formed in the earth wherever an 
 acid solution comes in contact with a carbonate, the gas 
 is frequently given off from fissures in the earth. It is 
 hence not unfrequently found in old wells which have not 
 been in use for some time, and deaths have been caused 
 by descending these wells for the purpose of repairing 
 them. The gas is also frequently met with in mines, and 
 is called choke-damp by the miners. The miners are 
 
CARBON DIOXIDE AND LIFE. 385 
 
 ;iware that after an explosion caused by fire-damp there 
 is danger of death from choke-damp. The reason of the 
 presence of this gas after an explosion is simple. When 
 fire-damp, or marsh-gas, explodes with air the carbon is 
 oxidized to choke-damp, or carbon dioxide, and the hy- 
 drogen to water. Air in which a candle will not burn 
 is not fit for breathing purposes. 
 
 Carbon Dioxide and Life. The role played by carbon 
 dioxide in nature is extremely important and interesting. 
 The carbon contained in living things is obtained from 
 carbon dioxide, and generally returns to this form when 
 life ceases. We have seen that all living things contain 
 carbon as an essential constituent. Whence comes this 
 carbon? Animals eat either the products of plant-life 
 or other animals which derive their sustenance from the 
 vegetable kingdom. The food of animals comes, then, 
 either directly or indirectly from plants. But plants 
 derive their sustenance largely from the carbon dioxide 
 of the air. The plants have the power to decompose the 
 gas with the aid of the direct light of the sun, and they 
 then build up the complex compounds of carbon which 
 form their tissues, using for this purpose the carbon of 
 the carbon dioxide which they decompose. Many of 
 these compounds are fit for food for animals ; that is to 
 say, they are of such composition that the forces at work 
 in the animal body are capable of transforming them 
 into animal tissues, or of oxidizing them, and thus keep- 
 ing the temperature of the body up to the necessary 
 point. That part of the food which undergoes oxidation 
 in the body plays the same part as fuel in a stove. It is 
 burned up with an evolution of heat, the carbon being 
 converted into carbon dioxide, which is given off from 
 the lungs. From fires and from living animals carbon 
 dioxide is returned to the air, where it again serves as 
 food for plants. When the life process stops in the ani- 
 mal or the plant, decomposition begins ; and the final 
 result of this, under ordinary circumstances, is the con- 
 version of the carbon into the dioxide. 
 
 Energy Stored up in Plants. It will thus be seen that 
 under the influence of life and sunlight carbon dioxide is 
 
386 INORGANIC CHEMISTRY. 
 
 constantly being converted into compounds containing 
 carbon which are stored up in the plant. These com- 
 pounds are capable of burning, and thus giving heat ; or 
 some of them may be used as food by animals, when they 
 assume other forms under the influence of the life-process 
 of the animals. As long as life continues, plants and 
 animals are storehouses of energy. When death occurs, 
 the carbon compounds begin to pass back to the form of 
 carbon dioxide, and the chemical energy is transformed 
 partly into heat, and is thus, as we say, dissipated. The 
 power to do work, which the carbon compounds of plants 
 and animals possess, comes from the heat of the sun. 
 It takes a certain quantity of this heat, operating under 
 proper conditions, to decompose a certain quantity of 
 carbon dioxide, and elaborate the compounds contained 
 in the plants. When these compounds are burned they 
 give out the heat which was absorbed in their formation 
 during the growth of the plants. These compounds are 
 said to possess chemical energy. This has its origin in 
 heat, and is capable of reconversion into heat. The 
 transformation of the energy of the sun's heat into chemi- 
 cal energy lies at the foundation of all life. As the heat 
 of the sun acting upon the great bodies of water and on 
 the air gives rise to the movements of water which are 
 so essential to the existence of the world as it is, so the 
 action of the sun's rays on carbon dioxide, under the 
 influence of the delicate and inexplicable mechanism of 
 the leaf of the plant, gives rise to those changes in the 
 forms of combination of the element carbon which ac- 
 company and are fundamental to the wonderful process 
 of life. 
 
 Carbonic Acid and Carbonates. When carbon dioxide 
 is passed into water the solution has a slightly acid re- 
 action. The solution will act upon bases and form salts. 
 The formula of the sodium salt formed in this way has 
 been shown to be Na 2 CO 3 ; that of the potassium sail, 
 K 2 CO 8 ; etc. These salts are plainly derived from an 
 acid, H 2 CO 3 , which is called carbonic acid. It is prob- 
 able that this acid is contained in the solution of carbon 
 
CARBONIC ACID AND CARBONATES. 387 
 
 dioxide in water. It is, however, so unstable that it 
 
 breaks up into carbon dioxide and water : 
 
 - -~~-.. 
 
 H 2 C0 3 = C0 2 + H 2 0. 
 
 The formation of a salt by the action of carbon di- 
 oxide on a base takes place as shown in the following 
 equations : 
 
 2KOH + CO 2 = K 2 CO 3 + H 2 O ; 
 Ca(OH) 2 + CO a = CaGO, + H 3 O. 
 
 "With the acid the action would take place as represented 
 thus : 
 
 2KOH + H 2 C0 3 = K 2 CO 3 + 2H 2 O ; 
 
 Ca(OH) 2 + H 2 C0 3 = CaC0 3 + 2H 2 O. 
 
 There is perfect analogy between the action of carbon 
 dioxide and that of sulphur dioxide on basic solutions. 
 With potassium hydroxide and calcium hydroxide, sul- 
 phur dioxide acts as represented in the following equa- 
 tions : 
 
 2KOH + SO 2 = K 2 SO 3 + H 2 O ; 
 Ca(OH) 2 + SO, = CaS0 3 + H 2 O. 
 
 The products formed are sulphites or salts of sulphur- 
 ous acid. 
 
 Like sulphurous acid, carbonic acid is dibasic, and 
 forms two series of salts, the primary and secondary, 
 or the acid and normal salts. The primary or acid salts 
 have the general formula HMCO 3 , and the secondary or 
 normal salts have the general formula M 2 CO 3 . Exam- 
 ples of the former are HKCO 3 , HNaCO 3 , CaH 2 (CO 3 ) 2 , 
 etc. ; and of the latter K.CO,, Na 2 CO 3 , CaCO 3 , BaCO 3 , 
 etc. It also readily forms basic salts, as, for example, 
 basic copper carbonate. Neutral copper carbonate is to 
 be regarded as formed by the action of one molecule of 
 the dibasic carbonic acid upon one molecule of the di- 
 acid copper hydroxide, Cu(OH) 2 : 
 
388 INORGANIC CHEMISTRY. 
 
 OC< OH + HO >Cu = oc <o >Cu 
 
 One of the simplest basic carbonates of copper is that 
 formed by the action of two molecules of copper hydrox- 
 ide upon one molecule of carbonic acid : 
 
 or OH , (HO)Cu(OH) op OCuOH ,o H o 
 00< OH + (HO)Cu(OH) = 00< OOuOH + 2H *- 
 
 Another basic salt of more complicated composition is 
 that of magnesium. It is to be regarded as derived from 
 carbonic acid as represented in this formula : 
 
 There are some salts which are derived from a pyro- 
 carbonic acid, that is, a form of the acid derived from 
 two molecules of the acid by loss of one molecule of 
 water : 
 
 200 < OH = H 2 C 2 5 + H 2 0. 
 
 Such a salt is formed by loss of water from the primary 
 sodium salt : 
 
 2HNaCO s = Na 2 C 2 5 + H 2 0. 
 
 There are no salts known derived from normal carbonic 
 acid, C(OH) 4 , though there are some compounds analo- 
 gous to salts which are derivatives of this normal acid. 
 The secondary or normal salts which carbonic acid forms 
 with the most strongly marked metallic elements, viz., 
 potassium and sodium, are not decomposed by heat, 
 but all other carbonates are decomposed by heat more 
 or less easily, according to the strength of the base. 
 Calcium carbonate when ignited loses carbon dioxide, 
 and lime, or calcium oxide, remains behind : 
 
 CaCO 3 = CaO + CO 2 . 
 
CARBON MONOXIDE. 389 
 
 Carbon Monoxide, CO. When a substance containing 
 carbon burns in an insufficient supply of air, as, for 
 example, when the draught in a furnace is not strong 
 enough to remove the products of combustion and sup- 
 ply fresh air, the oxidation of the carbon is not com- 
 plete, and the product, instead of being carbon dioxide, 
 is carbon monoxide, CO. This compound can also be 
 made by extracting oxygen from carbon dioxide. It is 
 only necessary to pass the dioxide over heated carbon, 
 when reaction takes place as represented thus : 
 
 CO 2 + C = 2CO. 
 
 This method of formation is illustrated in coal fires, and 
 can be well observed in an open grate. The air has free 
 access to the coal, and at the surface complete oxidation 
 takes place. But that part of the carbon dioxide which 
 is formed at the lower part of the grate is drawn up 
 through the heated coal, and is partly reduced to carbon 
 monoxide. When the monoxide escapes from the upper 
 part of the grate it again combines with oxygen, or burns, 
 giving rise to the characteristic blue flame always noticed 
 above a mass of burning anthracite coal. Should any- 
 thing occur to prevent free access of air, carbon monox- 
 ide may readily escape complete oxidation. 
 
 The monoxide is also formed by passing steam over 
 highly heated carbon, when this reaction takes place : 
 
 C + H 2 = CO + H 3 . 
 
 This is the reaction made use of in the manufacture of 
 " water-gas." The gas thus obtained is largely a mixture 
 of hydrogen and carbon monoxide. The gas is enriched 
 by passing it through a furnace in which it is mixed with 
 highly heated vapors of hydrocarbons from petroleum. 
 The main reaction, the decomposition of water by heated 
 carbon, is effected in large furnaces filled with anthracite 
 coal. The coal is first heated to a high temperature 
 by setting fire to it, the products of combustion being 
 allowed to escape. When it is hot enough, the air is 
 
390 INORGANIC CHEMISTRY. 
 
 shut off and steam passed rapidly in, when the decom- 
 position of the water by the carbon takes place. Soon 
 the mass becomes so much cooled that the reaction 
 stops. The steam is then cut off and air turned on again, 
 and so on. 
 
 The easiest way to make carbon monoxide is by heat- 
 ing oxalic acid, which is a compound of carbon, hydro- 
 gen, and oxygen, of the formula C 2 H 2 O 4 , with five to six 
 times its weight of concentrated sulphuric acid. The 
 change which takes place is represented thus : 
 
 C 2 H 3 4 = C0 2 + CO + H 2 0. 
 
 Both the dioxide and monoxide of carbon are formed. 
 Both are gases. In order to separate them the mixture 
 is passed through a solution of sodium hydroxide, which 
 takes up the carbon dioxide, forming sodium carbonate, 
 and allows the monoxide to pass. 
 
 Carbon monoxide is a colorless, tasteless, inodorous 
 gas, insoluble in water. It burns with a pale-blue flame, 
 forming carbon dioxide. It is exceedingly poisonous 
 when inhaled. Hence it is very important that it should 
 not be allowed to escape into rooms occupied by human 
 beings. We not unfrequently hear of deaths caused by 
 the gases from coal stoves. The most dangerous of the 
 gases given off from these stoves is probably carbon 
 monoxide. A pan of smouldering charcoal gives off this 
 gas, and the fact that it is poisonous is well known. It 
 has been used to a considerable extent for the purpose 
 of suicide. The poisonous character of carbon monoxide 
 has led to a great deal of discussion and to some legisla- 
 tion on the subject of " water-gas." The question has 
 been repeatedly raised whether government should allow 
 the manufacture of the gas. There is no doubt of the 
 fact that it is a dangerous substance, and that it should 
 not be allowed to escape into the air is obvious. Where- 
 ever it is used special precautions should be taken to 
 guard against leaking. There is no doubt that it is 
 somewhat more poisonous than coal-gas. 
 
 At high temperatures carbon monoxide has a very 
 
FORMIC ACID. 391 
 
 strong tendency to combine with oxygen, and is hence a 
 good reducing agent. In the reduction of iron from its 
 ores, the carbon monoxide formed in the blast-furnace 
 plays an important part in the reducing process. At 
 ordinary temperatures the gas does not combine readily 
 with oxygen. Thus, it does not act upon ozone, even 
 when heated with it somewhat above the temperature at 
 which ozone is converted into ordinary oxygen. When 
 passed over some substances which are rich in oxygen, 
 as, for example, chromic anhydride, CrO 3 , and potassium 
 permanganate, KMnO 4 , in acid solution, it takes up oxy- 
 gen even at the ordinary temperature. It unites with 
 chlorine in the direct sunlight, and forms the com- 
 pound known as carbonyl chloride, COC1 2 . 
 
 Formic Acid, H 2 CO 2 . Just as carbon dioxide may be 
 regarded as the anhydride of carbonic acid, so carbon 
 monoxide may be regarded as the anhydride of an acid 
 of the formula H 2 CO 2 . While, however, an acid of this 
 formula is known, it is not formed by action of carbon 
 monoxide upon water, nor are its salts easily formed by 
 the action of carbon monoxide upon bases. By passing 
 it over certain basic substances, however, as, for example, 
 potassium hydroxide and calcium hydroxide, at a com- 
 paratively high temperature action takes place, and salts 
 of the acid are formed. Thus, in the case of potas- 
 sium hydroxide, the action takes place as represented in 
 the equation 
 
 CO + KOH = HCO 2 K. 
 
 Although it contains two atoms of hydrogen in the 
 molecule, formic acid is monobasic. This fact finds its 
 explanation in the structure of the acid. All its reactions 
 show that only one of the hydrogen atoms of formic 
 acid is in combination with oxygen, while the other is in 
 combination with carbon, as represented in the formula 
 
 ? (H 
 
 HC-OH or C K O . The relations here are similar to 
 
 (OH 
 those met with in phosphorous and sulphurous acids, 
 
392 i INORGANIC CHEMISTRY. 
 
 which have been so frequently referred to. Formic 
 acid bears to carbonic acid the same relation that 
 sulphurous bears to sulphuric acid, and phosphorous 
 to phosphoric acid, as shown in the formulas : 
 
 Formic acid Carbonic acid 
 
 o,sj 
 
 H q ( OH 
 
 OH u ' b \ OH 
 
 Sulphurous acid Sulphuric acid 
 
 (H (OH 
 
 OP \ OH OP \ OH 
 
 (OH (OH 
 
 Phosphorous acid Phosphoric acid 
 
 Formic acid occurs in nature in red ants, in stinging 
 nettles, and elsewhere. It is a colorless liquid, which 
 solidifies at 8. 6. When treated with concentrated sul- 
 phuric acid it breaks down into carbon monoxide and 
 water : 
 
 H 3 CO, = CO + H a O. 
 
 By oxidation it is converted into carbon dioxide and 
 water. 
 
 Carbonyl Chloride, Phosgene, COC1 2 . This compound 
 was referred to above as being formed when chlorine acts 
 upon carbon monoxide under the influence of the sun's 
 rays. It is also formed when the two gases are passed 
 together through a tube filled with pieces of animal char- 
 coal. It is a colorless gas, which is easily condensed to 
 a liquid boiling at 8. 2. It is now manufactured on the 
 large scale for use in the preparation of certain classes 
 of dye-stuffs. 
 
 Like the oxychlorides of sulphur and of phosphorus, 
 this compound, which is an oxychloride of carbon, is de- 
 composed by water, forming carbonic acid or its products 
 of decomposition : 
 
CARBONTL CHLORIDE OR PHOSGENE. 393 
 
 + 181 
 
 Cl HOH ( OH 
 
 Cl + HOH = PO^ OH + 3HC1. 
 
 Cl HOH ( OH 
 
 It is interesting to note that, while the chlorides of 
 sulphur and phosphorus, SC1 2 and PC1 3 , as well as 
 SC1 4 and PC1 6 , are easily decomposed by water, the 
 tetrachloride of carbon, CC1 4 , is not. On the other 
 hand, the tetrachloride is not formed when the oxides 
 of carbon are treated with hydrochloric acid. It 
 will be remembered that, in discussing the relations be- 
 tween the acid-forming and the base-forming elements, 
 attention was called to the fact that, in general, the chlo- 
 rides of the acid-forming elements are easily decom- 
 posed by water, forming the corresponding hydroxides or 
 oxides, while the oxides and hydroxides of the base-form- 
 ing elements are converted into chlorides by treatment 
 with hydrochloric acid. In carbon we have an example 
 of an element which occupies what may be called almost 
 a neutral ground between the two classes of elements. 
 It forms both acids and bases, to be sure, but these are, 
 generally speaking, not as strongly marked as the acids 
 and bases formed by other elements. This neutral char- 
 acter of the element is also well shown in the conduct of 
 its chloride towards water, and of its oxides towards 
 hydrochloric acid. 
 
CHAPTER XXI. 
 
 ILLUMINATiON-FLAME BLOW-PIPE. 
 COMPOUNDS OF CARBON WITH NITROGEN AND SULPHUR 
 
 Introduction. As the substances used for illumina- 
 tion contain carbon, and the chemical processes involved 
 consist largely in the oxidation of the carbon of these 
 compounds, this is an appropriate place to consider 
 briefly the subject of illumination from a chemical point 
 of view, as well as that of flame, and the blow -pipe, 
 which gives an extremely useful form of flame constantly 
 used in the laboratory. 
 
 In all ordinary kinds of illumination we are dependent 
 upon flames for the light. Whether we use illuminating 
 gas, a lamp, or a candle, the light comes from a flame. 
 In the first case, the gas is burned directly ; in the case 
 of the lamp, the oil is first drawn up the wick, then con- 
 verted into a gas, and this burns ; while, finally, in the 
 case of the candle, the solid material of the candle is 
 first melted, then drawn up the wick, converted into gas, 
 and the gas burns, forming the flame. In each case we 
 have, then, to deal with a burning gas, and this burning 
 gas is called a flame. 
 
 Illuminating Gas, Coal-gas. Illuminating gas is some- 
 times made from coal by heating in closed retorts. As 
 has already been explained, coal, particularly the softer 
 kinds, contains compounds of carbon and hydrogen, 
 together with some nitrogen and other elements. When 
 it is subjected to destructive distillation, as in the manu- 
 facture of coal-gas, the hydrogen passes off partly in 
 combination with carbon, as hydrocarbons, and partly in 
 the free state. The nitrogen passes off as ammonia, and 
 a large percentage of the carbon remains behind in the 
 retort in the uncombined state as coke. The gases given 
 
 (394) 
 
FLAMES. 395 
 
 .' " f 
 
 off are purified, and form illuminating gas. One ton of 
 coal yields on an average 10,000 cubic feet of gas. The 
 value of a gas depends upon the quantity of light given 
 by the burning of a definite quantity. It is measured 
 by comparing it with the light given by a candle burn- 
 ing at a certain rate. The standard candle is one made 
 of spermaceti, which burns at the rate of 120 grains per 
 hour ; that is to say, a candle which, burning under 
 ordinary conditions, loses 120 grains in one hour. The 
 standard burner used for the gas is one through which 
 five cubic feet of gas pass per hour. Now, to determine 
 the illuminating power of a gas, it is passed through the 
 standard burner at the rate mentioned, and the light 
 which it gives is compared with the light given by the 
 standard candle. This comparison is easily made by 
 means of an instrument called the photometer. The 
 illuminating power of the gas is then stated in terms of 
 the standard candle. The statement that the illumi- 
 nating power of a gas is fourteen candles, signifies that, 
 when burning at the rate of five cubic feet per hour, its 
 flame gives fourteen times as much light as that of the 
 standard candle. 
 
 Flames. Ordinarily when we speak of a flame wo 
 mean a gas which is combining with oxygen. The hy- 
 drogen flame is simply the phenomenon accompanying 
 the act of combination of the two gases hydrogen and 
 oxygen. Owing to the fact that we are surrounded by 
 oxygen, we speak of hydrogen as the burning gas. How 
 would it be if we were surrounded by an atmosphere of 
 hydrogen? Plainly, oxygen would then be a burning 
 gas. If we allow a jet of oxygen to escape into a vessel 
 containing hydrogen, a flame will appear where the oxy- 
 gen escapes from the jet, if a light is applied. This is 
 an experiment which requires special precautions, and, 
 as the principle can be illustrated as well by means of 
 illuminating gas, this may be used instead. Just as 
 illuminating gas burns in an atmosphere of oxygen, so 
 oxygen burns in an atmosphere of illuminating gas. 
 
 Kindling Temperature of Gases. In studying the action 
 of oxygen upon other substances, we learned that it is 
 
396 INORGANIC CHEMISTRY, 
 
 necessary that each of these substances should be raised 
 to a certain temperature before it will combine with 
 oxygen. This statement is as true of gases as of other 
 substances. When a current of hydrogen is allowed to 
 escape into the air, or into oxygen, no action takes place 
 unless it is heated up to its burning temperature, when 
 it takes fire and continues to burn, as the burning of 
 one part of the gas heats up the part which follows it, 
 and hence it is heated up to the burning tempera- 
 ture as fast as it escapes into the air. If the gas is 
 cooled down even very slightly below this temperature, 
 it is extinguished. This can easily be shown by bring- 
 ing down upon the flame of a Bunsen burner a piece of 
 wire gauze. There will be no flame above the gauze, 
 but gas will pass through unburned, and this will burn 
 if it is lighted above the gauze. In this case, by simply 
 passing through the thin wire gauze, the gases are cooled 
 down below their burning temperatures, and the flame 
 does not pass through. So, also, if the gas is turned 
 on and not lighted, and the gauze held an inch or two 
 above the outlet, the gas will burn above the gauze if 
 lighted above, and will not pass downward through the 
 gauze, unless this becomes very hot. 
 
 Miner's Safety -lamp. The principle illustrated in the 
 experiments referred to in the last para- 
 graph is utilized in the miner's safety- 
 lamp, to which reference has already been 
 made. One of the dangers which the coal- 
 miner has to encounter is the occurrence 
 in the mines of fire-damp, or methane, 
 CH 4 , which with air forms an explosive 
 mixture. The explosion can only be 
 brought about by contact of flame with 
 the mixture. In order to avoid the con- 
 tact, the flame of the safety-lamp is sur- 
 rounded by wire gauze, as shown in Fig. 
 11. When a lamp of this kind is brought 
 into an explosive mixture of marsh-gas 
 FIG. 11. and air, the mixture passes through the 
 wire gauze and comes in contact with the flame, and a 
 
STRUCTURE OF FLAMES. 39? 
 
 small explosion or a series of small explosions inside 
 the gauze occurs, but the flame of the burning gas 
 inside the wire gauze cannot pass through and raise the 
 temperature of the gas outside to the burning tempera- 
 ture. Hence no serious , explosion can take place. The 
 flickering of the flame of the lamp, and the occurrence 
 of small explosions inside, furnish the miner with the 
 information that he is in a dangerous atmosphere. While 
 the safety-lamp does undoubtedly afford much protection, 
 still explosions occur. These have been shown to be caused 
 by the presence of coal-dust in the mines, and by the com- 
 motion of the air produced in blasting. By the aid of the 
 coal-dust, and by sudden and violent movements of the air, 
 it is possible for a flame surrounded by wire gauze to 
 explode a mixture of marsh-gas and air on the other side 
 of the gauze. 
 
 Structure of Flames. The hydrogen flame consists of 
 .a thin envelope of burning hydrogen enclosing unburned 
 .gas, and surrounded by water vapor, which is the prod- 
 uct of the combustion. The structure of other flames 
 depends upon the complexity of the gases burned, and 
 the conditions under which the burning takes place. In 
 .general, a flame consists of an outer envelope of gas 
 combining with oxygen, and hence hot, and an inner 
 part which contains unburned gas, which is compara- 
 tively cool. A part of the unburned gas is, however, 
 quite hot, and it would combine with oxygen were it not 
 for the fact that it is surrounded by an envelope which 
 prevents access of air. The outer hot part of the flame 
 is called the oxidizing flame, because it presents condi- 
 tions favorable to the oxidation of substances introduced 
 into it. The inner hot part is called the reducing flame, 
 because it consists of highly heated substances which 
 have the power to combine with oxygen ; and hence many 
 compounds containing oxygen lose it, or are reduced, 
 when introduced into this part of the flame. The hot- 
 test part of the flame is about half-way between the 
 bottom and the top. Here oxidation is taking place 
 most energetically. The hottest part of the unburned 
 gases is at the tip of the dark central part of the flame. 
 
398 INORGANIC CHEMISTRY. 
 
 In the flame of a Bunsen burner the two parts can be 
 easily distinguished. The dark central part of the flame 
 extends for some distance above the outlet of the burner. 
 If the holes at the base of the burner are partly closed, 
 the tip of the central part of the flame becomes lumi- 
 nous. This luminous tip is most efficient for the pur- 
 pose of reduction. The principal parts of 
 the flame are those marked in Fig. 12. The 
 part marked b is the central cone of un- 
 burned gases ; that marked c is the lumi- 
 nous tip, the best part of the flame for re- 
 duction. A is the envelope of burning 
 gas. The hottest part of the flame is at a ; 
 that which is most efficient in causing oxi- 
 dation is at d. This is further surrounded 
 by a non-luminous envelope consisting of 
 the products of combustion, carbon diox- 
 ide and water vapor. Certain metals placed 
 in the upper end of the flame take up 
 oxygen, because they are highly heated in the presence 
 of oxygen. Certain oxides lose their oxygen when placed 
 in the tip of the central cone, because the gases are here 
 heated to the temperature at which they have the 
 power to combine with oxygen. 
 
 Blow-pipe. The oxidizing and reducing flames are 
 frequently utilized in the laboratory. For the purpose 
 of increasing their efficiency a blow-pipe is used. This is 
 simply a tube with a convenient mouth-piece and a nozzle 
 with a small opening through which air is blown into a 
 flame by means of the mouth. The blow-pipe may be 
 used with the flame of a candle, an alcohol-lamp, or a 
 gas-lamp. It is commonly used with a gas-lamp. By 
 regulating the current of air and slightly changing the 
 position of the tip of the blow-pipe a good oxidizing 
 flame or a good reducing flame can be produced. Some 
 oxides are very easily reduced when heated in the re- 
 ducing blow-pipe flame. Others are not. We can fre- 
 quently judge of the composition of a substance by heat- 
 ing in the blow-pipe flame, and noticing its conduct. 
 Some metals are easily oxidized in the oxidizing flame. 
 
LUMINOSITY OF FLAMES. 399 
 
 Some form characteristic films, or thin layers of oxides, 
 on the substance upon which they are heated, which is 
 usually charcoal ; and, in some cases, it is possible to 
 detect the presence of certain substances by the color of 
 the film of oxide. The blow-pipe is therefore of much 
 value as affording a method for the detection of the 
 presence of certain elements in mixtures or compounds 
 of unknown composition. The chemical principles in- 
 volved in its use will be clear from what has already 
 been said. 
 
 Causes of the Luminosity of Flames. It is evident from 
 what we have seen that flames differ greatly in their 
 light-giving power. The hydrogen flame, for example, 
 though extremely hot, gives practically no light. This is 
 also the case with the flame of the Bunsen burner ; 
 while, on the other hand, the flame of coal-gas, burning 
 under ordinary circumstances, and that of a candle, etc., 
 give light. To what is the difference due ? This subject 
 has been studied very thoroughly, and it has been found 
 that there are several causes which operate to make a 
 flame give light, and vice versa. In the first place, if a 
 solid substance which does not burn is introduced into 
 a non-luminous flame, a part of the heat appears as 
 light. This is seen when a spiral of platinum wire is 
 introduced into a hydrogen flame. It is also seen when 
 a piece of lime is introduced into the hot non-luminous 
 flame of the oxyhydrogen blow-pipe. A similar cause 
 operates in ordinary gas-flames to make them luminous. 
 There are always present particles of unburned carbon, 
 as can be shown by putting a piece of porcelain or any 
 solid substance into the flame, when there will be de- 
 posited on it a layer of soot, which consists mainly of 
 finely divided carbon. In the flame such particles are 
 heated up to incandescence, or to the temperature at 
 which they give light. Again, it has been found that a 
 candle gives more light at the level of the sea than it does 
 when at the top of a high mountain, as Mount Blanc, on 
 which the experiment was actually performed. This is 
 partly due to a difference in the density of the gases. 
 Naturally, the denser the gas the more active the com- 
 
400 INORGANIC CHEMISTRY. 
 
 bustion, the greater the heat, and the brighter the light. 
 This last statement ceases to be true when the oxidation 
 becomes sufficient to burn up all the solid particles in 
 the flame. If gases, which in burning give light, are 
 cooled down before they are burned, the luminosity is 
 diminished, and, conversely, non-luminous flames may 
 be rendered luminous by heating the gases before burn- 
 ing them. Gases which otherwise give luminous flames 
 give non- luminous flames when diluted to a sufficient 
 extent with neutral gases, such as nitrogen and carbon 
 dioxide, which neither burn nor support combustion. 
 
 Bunsen Burner. All the statements made in regard to 
 the causes of the luminosity of flames are based upon 
 carefully performed experiments. These experiments, 
 however, cannot, for the most part, be readily repeated 
 by the student in the laboratory in a satisfactory way. 
 One constant reminder of the possibility of rendering a 
 luminous flame non-luminous, and vice versa, is fur- 
 nished by the burner universally used in chemical labora- 
 tories, and called, after the name of the inventor, the 
 Bunsen burner. The construction of this burner is easily 
 understood. It consists of a base and an upper tube. 
 The base is connected by means of a rubber tube with 
 the gas supply. The gas escapes from a small opening 
 in the base, and passes upward through the tube. At 
 the lower part of the tube there are two holes, which 
 may be opened or closed by turning a ring with two cor- 
 responding holes in it. When the gas is turned on, it is 
 lighted at the top of the tube. Air is at the same time 
 drawn through the holes at the base. The result is that 
 the flame is practically non-luminous. If the ring at the 
 base is turned so that the air-holes are closed, the 
 flame becomes luminous. The advantage of the non- 
 luminous flame for laboratory use consists in the fact 
 that it does not deposit soot, and, at the same time, it 
 is hot. 
 
 The non-luminosity of the flame of the Bunsen burner 
 appears to be due to several causes : (1) Dilution of the 
 gases by means of the nitrogen of the air ; (2) Cooling 
 of the gases by the entrance of the air ; (3) Burning of 
 
CYANOGEN. 401 
 
 the solid particles by tlie aid of the oxygen of the air 
 admitted to the interior of the flame. 
 
 COMPOUNDS or CARBON WITH NITROGEN AND WITH 
 SULPHUR. 
 
 Cyanogen, C 2 N 2 . Carbon does not combine with ni- 
 trogen under ordinary circumstances. If, however, they 
 are brought together at very high temperatures in the 
 presence of metals, they combine to form compounds 
 known as cyanides. Thus, when nitrogen is passed over 
 a highly heated mixture of carbon and potassium car- 
 bonate, potassium cyanide, KCN, is formed. Carbon 
 containing nitrogen, as animal charcoal, when ignited 
 with potassium carbonate, reduces the carbonate, form- 
 ing potassium, in presence of which carbon and nitro- 
 gen combine, forming potassium cyanide. When refuse 
 animal substances, such as blood, horns, claws, hair, 
 etc., are heated together with potassium carbonate and 
 iron, a substance known as potassium ferro-cyanide, or 
 yellow prussiate of potash, K 4 Fe(CN) 6 -J- 3H 2 O, is formed. 
 When this is simply heated it is decomposed, yielding 
 potassium cyanide : 
 
 K 4 Fe(CN) 6 = 4KCN + FeC 2 + N 2 . 
 
 It is an easy matter to make the mercury salt, Hg(CN) a , 
 from the potassium salt. By heating mercuric cyanide 
 it breaks up, yielding metallic mercury and cyanogen 
 gas: 
 
 just as mercuric oxide yields mercury and oxygen when 
 heated : 
 
 HgO = Hg + 0. 
 
 Cyanogen (from KIMXVOS, Hue) owes its name to the fact 
 that several of its compounds have a blue color. It is a 
 colorless gas, which is easily soluble in water and alco- 
 
402 INORGANIC CHEMISTRY. 
 
 hoi, and is extremely poisonous. It burns with a purple- 
 colored flame. In aqueous solution cyanogen soon un- 
 dergoes change, and a brown amorphous substance is 
 deposited. In the solution are found hydrocyanic acid, 
 oxalic acid, ammonia, and carbon dioxide. The princi- 
 pal cause of this decomposition is apparently the ten- 
 dency of the nitrogen to combine with hydrogen to form 
 the stable compound ammonia, and of carbon to com- 
 bine with hydrogen and oxygen to form stable com- 
 pounds like oxalic acid and carbon dioxide. One of the 
 chief decompositions which cyanogen undergoes with 
 water is that represented in the equation 
 
 H 
 
 The compound I 2 or H 2 C 2 O 4 is oxalic acid. This 
 
 kind of decomposition with water is characteristic of cy- 
 anogen compounds. It consists, as will be seen, in the 
 union of the nitrogen with hydrogen to form ammonia, 
 and the union of the carbon with oxygen and hydroxyl. 
 
 Hydrocyanic Acid, Prussia Acid, t HCN. This acid, 
 which is commonly known by the name prussic acid, oc- 
 curs in nature in amygdalin, in combination with other 
 substances, in bitter almonds, the leaves of the cherry, 
 laurel, etc. It is prepared by decomposing metallic cy- 
 anides with hydrochloric acid. It is volatile and passes 
 over. The action is represented thus : 
 
 KCN + HC1 = KC1 + HON. 
 
 It can also be made by treating chloroform with 
 ammonia : 
 
 CHC1 8 + NH 3 = HCN + 3HC1. 
 
 Of course, the hydrochloric acid and the hydrocyanic 
 acid formed combine with ammonia, so that the complete 
 action is represented by this equation : 
 
 CHC1 3 + 5NH 3 = NH 4 CN + 3NH 4 C1. 
 The product NH 4 CN is ammonium cyanide. 
 
HYDROCYANIC ACID. 403 
 
 Hydrocyanic acid is a volatile liquid, boiling at 26. 1, 
 and solidifying at 14. It has a very characteristic 
 odor suggestive of bitter almonds. It is extremely poi- 
 sonous. It dissolves in water in all proportions, and it is 
 such a solution which is known as prussic acid. Pure 
 hydrocyanic acid is very unstable. By standing, a brown 
 substance is deposited from its solution. By boiling 
 with alkalies or acids it is converted into formic acid and 
 ammonia. This is another example of the tendency of 
 cyanogen compounds to decompose in the presence of 
 water, yielding ammonia and oxygen compounds of car- 
 bon. The decomposition of hydrocyanic acid takes place 
 as represented in the equation 
 
 HHO 
 
 The relations between chloroform, formic acid, and 
 hydrocyanic acid are instructive. By replacing all the 
 chlorine atoms of chloroform by hydroxyl a compound of 
 
 roH 
 
 OTT 
 
 the formula C -j QJJ would be formed ; but this would 
 
 break down by loss of water, yielding formic acid, C 
 
 By replacing the three chlorine atoms by one nitrogen 
 atom hydrocyanic acid is formed ; and this in turn when 
 decomposed in presence of water yields formic acid. 
 These relations will be made clear by the aid of the fol- 
 lowing formulas and equations : 
 
 oH 
 
 fHOH=Cj^ + 3H01; 
 OH 
 
 TTr\TT 
 
 H H 
 
404 INORGANIC CHEMISTRY. 
 
 Cyanic Acid, HCNO. By gentle oxidation of a cyan- 
 ide it is converted into a cyanate. Thus, by melting 
 together potassium cyanide and lead oxide, potassium 
 cyanate is formed : 
 
 KCN + PbO = KCNO + Pb. 
 
 Cyanic acid is a volatile, acrid, unstable liquid. It 
 breaks down at once into carbon dioxide and ammonia 
 in presence of water : 
 
 COKE + H 3 O = NH 3 + C0 a . 
 
 The potassium salt is easily soluble in water, but is 
 decomposed by it, yielding ammonia and acid potassium 
 carbonate : 
 
 CONK + 2H 2 O = KHC0 3 + NH 3 . 
 
 These decompositions of cyanic acid and the cyanates 
 further exemplify the tendency of cyanogen compounds 
 to undergo decomposition in presence of water. 
 
 Carbon Bisulphide, CS 2 . Just as carbon combines di- 
 rectly with oxygen to form the dioxide, so it combines 
 directly with sulphur to form the disulphide ; but there 
 is a great difference in the ease with which carbon com- 
 bines with the two elements. In order to effect combina- 
 tion with sulphur a very high temperature is necessary. 
 The compound is prepared on the large scale by heating 
 charcoal to a high temperature in an upright cast-iron 
 cylinder, and adding sulphur in such a way that it enters 
 the bottom of the cylinder. The product is passed 
 
CARBON BISULPHIDE. 405 
 
 through a series of tubes arranged so as to secure 
 condensation. 
 
 Carbon disulphide is a clear liquid which has a high 
 refractive power. It boils at 46. 2. When pure it has 
 a pleasant odor, but if kept for a time, particularly if 
 water is present in the vessel, it undergoes slight decom- 
 position, and products of extremely disagreeable odor 
 are formed. It can generally be freed from these by 
 shaking the liquid with a little mercury and then redis- 
 tilling. It burns readily, forming carbon dioxide and 
 sulphur dixoide : 
 
 CS, + 30 2 = C0 2 + 2S0 2 . 
 
 In nitric oxide it burns with an intensely brilliant flame, 
 as can be shown by filling a cylinder with the gas, adding 
 a few drops of the disulphide, shaking and then apply- 
 ing a flame. A column of brilliant flame rises from the 
 mouth of the cylinder for an instant. A lamp has been 
 constructed in which this flame is utilized. It is of 
 special value in photographic work. 
 
 Carbon disulphide is only very slightly soluble in 
 water, and is decomposed by it only very slowly. The 
 disulphide is an excellent solvent for many substances 
 which are not soluble in water, as, for example, fats, 
 resins, iodine, and one of the modifications of sulphur 
 and of phosphorus. The solution of iodine in it has a 
 beautiful violet color ; and when a water solution con- 
 taining a little free iodine is shaken with carbon disul- 
 phide the latter acquires a violet color and separates 
 below the water. 
 
 When the vapors of carbon disulphide and hydrogen 
 sulphide are passed together over heated copper, methane 
 and cuprous sulphide are formed, as has been stated. 
 Methane is also formed when the vapors of carbon disul- 
 phide and of water are passed over ignited iron. While 
 the disulphide is not easily decomposed by water at the 
 ordinary temperature, the two compounds react when 
 
406 INORGANIC CHEMISTRY. 
 
 heated in a sealed tube to 150, the products being car* 
 bon dioxide and hydrogen sulphide : 
 
 CS 2 + 2H 2 O = CO 2 + 2H 2 S. 
 
 Carbon disulphide finds extensive application as a 
 solvent, and it is also used for the purpose of destroying 
 phylloxera, the insect which is so destructive to grape- 
 vines, particularly in the wine districts of France. 
 
 Sulphocarbonic Acid, Thio-carbonic Acid, H 2 CS 3 . Salts 
 of this acid are formed by dissolving carbon disulphide 
 in concentrated solutions of the hydrosulphides. Thus, 
 when it is dissolved in a solution of sodium hydrosulphide 
 this reaction takes place : 
 
 CS 2 + 2NaSH = Na 2 CS 3 + H 2 S. 
 
 The reaction, as will be seen, is perfectly analogous to 
 that of carbon dioxide upon a solution of sodium 
 hydroxide : 
 
 CO 2 + 2NaOH = Na 2 CO 3 + H 2 O. 
 
 The salts of sulphocarbonic acid are easily decom- 
 posed by water if the temperature is slightly elevated, 
 the products being the corresponding carbonates and 
 hydrogen sulphide : 
 
 Na.CS, + 3H 2 O = Na.CC), + 3H 2 S. 
 
 When a sulphocarbonate is treated with cold dilute 
 hydrochloric acid, the free acid separates as a dark yel- 
 low oil of a very disagreeable odor. This readily 
 undergoes decomposition into carbon disulphide and 
 hydrogen sulphide : 
 
 H 2 GS 3 = CS 2 -4- H 2 S. 
 
 This reaction is again perfectly analogous to the decom- 
 position of ordinary carbonic acid into carbon dioxide 
 and water. 
 
 Oxysulphides. Products intermediate between car- 
 bonic acid and sulphocarbonic acid are possible. Such, 
 
CONSTITUTION OF CYANOGEN. 407 
 
 for example, are the compounds represented by the for- 
 mulas CO < Q , CS -j , etc. 
 
 Sulphocyanic Acid, HCNS. Just as the cyanides take 
 up oxygen and are converted into cyanates, so also they 
 take up sulphur and are converted into sulphocyanates : 
 
 KCN + S = KCNS. 
 
 By suspending in water a salt of the acid, the metal of 
 which forms an insoluble sulphide with hydrogen sul- 
 phide, and passing this gas through the liquid, a 
 solution of the acid is obtained. When the solution is 
 boiled the acid passes over partly unchanged, though a 
 part is decomposed by the water into carbon dioxide, 
 carbon disulphide, and ammonia : 
 
 2HCNS + 2H 2 O = CO 2 + CS 3 + 2NH 3 . 
 
 The ammonium salt of sulphocyanic acid is formed by 
 dissolving carbon disulphide in an alcoholic solution of 
 ammonia : 
 
 CS 2 + 4NH 3 = (NH 4 )CNS + (NH 4 ) 2 S. 
 
 Constitution of Cyanogen and its Simpler Compounds. 
 The compounds of cyanogen show, in general, a remark- 
 able similarity to the compounds of the chlorine group. 
 The hydrogen compound is a monobasic acid and forms 
 a series of salts, the cyanides, which in general are ana- 
 logous to the chlorides. Comparing the cyanides with 
 the chlorides it is clear that in the former the group (CN), 
 or the cyanogen group, plays the same part that the atom 
 chlorine plays in the chlorides : 
 
 H(ON) HC1 
 K(CN) KC1 
 Hg(CN) a HgCl, 
 
 So also cyanic acid and hypochlorous acid are analo- 
 gous : 
 
 HO(CN) HOC1. 
 
408 INORGANIC CHEMISTRY. 
 
 This relation suggests that which is observed between 
 the ammonium compounds and those of potassium and 
 sodium. The cyanogen group is evidently univalent, as 
 it combines with one atom of hydrogen, one of potassium, 
 etc., and there are two ways in which we can conceive 
 the atoms carbon and nitrogen combined to form a uni- 
 valent group. If the nitrogen is trivalent and the carbon 
 quadrivalent the structure is that represented by the 
 formula -C=N. If, on the other hand, the nitrogen is 
 quinquivalent and the carbon quadrivalent the structure 
 is C=N-. By combination of the first group with hydro- 
 gen a compound of the structure H-C=N would result 
 while with the second group the structure of the hydro- 
 gen compound would be C=N-H, In the one case the 
 hydrogen is in combination with carbon, in the other with 
 nitrogen. It appears probable that in ordinary hydro- 
 cyanic acid the hydrogen is in combination with carbon, 
 the structure being H-C=N. This is in accordance with 
 the formation of the acid by the action of ammonia upon 
 chloroform, which is most readily understood on the as- 
 sumption that the three atoms of chlorine are replaced 
 by an atom of nitrogen. It has not been positively de- 
 termined which of the two possible structures above 
 given the cyanogen group has in cyanic acid. In one 
 case the acid would have the structure N=C OH ; in the 
 other it would be C=N-OH. It may also be O=C=NH. 
 There are some reasons for believing that the ordinary 
 cyanates are derived from an acid of the structure rep- 
 resented by the last formula. 
 
CHAPTER XXII. 
 
 ELEMENTS OF FAMILY IV, GROUP A: 
 SILICON TITANIUM ZIRCONIUM CERIUM THORIUM. 
 
 General. While the elements of this group in some 
 respects exhibit resemblances to carbon and bear to it 
 relations similar to those which the members of the 
 chlorine group bear to fluorine, the members of the sul- 
 phur group to oxygen, and the members of the phos- 
 phorus group to nitrogen, yet between them and carbon 
 there are some remarkable differences. All the members 
 of the group except titanium combine with hydrogen. 
 The compounds formed have the formulas SiH 4 , ZrH 2 , 
 CeH 2 , and ThH 2 . The power to form homologous series 
 which is so characteristic of carbon is entirely wanting 
 in the other members of the group. With the members 
 of the chlorine group they all form compounds analogous 
 to carbon tetrachloride, examples of which are : 
 
 SiCl 4 , TiCl 4 , ZrCl 4 , CeF 4 , ThCl 4 . 
 
 Compounds analogous to hexa-chlor-ethane, C 2 Cl g , to 
 tetra-chlor-ethylene, C 2 C1 4 , and to octo-chlor-propane, 
 C 8 Cl e ,are: 
 
 Si 2 Cl 8 Si 2 Cl 4 Si 3 Cl 8 
 Ti 2 Cl 6 Ti,Cl 4 
 
 All the elements of the group form oxygen compounds 
 analogous to carbon dioxide. They are : 
 
 Si0 2 , Ti0 2 , Zr0 2 , Ce0 2 , ThO a . 
 
 The first three are acidic, and form salts which in com- 
 position are analogous to the carbonates. These are the 
 
 (409) 
 
410 INORGANIC CHEMISTRY. 
 
 silicates, titanates, and zirconates of the general formulas 
 
 M 3 SiO 3 , M a TiO 3 , M 3 ZrO 8 . 
 
 Cerium and thorium oxides are basic. These facts 
 suggest the relations between the members of the phos- 
 phorus group. The oxides of the last two members, 
 antimony and bismuth, are basic, although the oxide 
 of antimony is also acidic in its conduct towards the 
 stronger bases. 
 
 The compounds of silicon are very abundant in nature ; 
 those of the other members of the group are rare. 
 
 SILICON, Si (At. Wt. 28.18). 
 
 Occurrence. We have already seen what an exceed- 
 ingly important part carbon plays in animate nature. It 
 is interesting to note that silicon, which in some respects 
 from a chemical point of view resembles carbon, is one 
 of the most important constituents of the mineral or in- 
 organic parts of the earth. It occurs chiefly in the form 
 of the dioxide, SiO 2 , commonly called silica, or silicon 
 dioxide ; and in combination with oxygen and several of 
 the common metallic elements, particularly with sodium, 
 potassium, aluminium, and calcium, in the form of the 
 silicates. Next to oxygen, silicon is the most abundant ele- 
 ment in the earth. There are extensive mountain-ranges 
 consisting almost entirely of the dioxide, SiO 2 , in the form 
 known as quartz or qiiartzite. Other ranges are made up 
 of silicates, which are compounds formed by the com- 
 bination of silicon dioxide and bases. The clay of 
 the vallevs and river-beds also contains silicon in 
 large quantity, while the sand found so abundantly 
 on the deserts and at the sea-shore is largely silicon 
 dioxide. 
 
 Preparation. Silicon does not occur in nature in the 
 free state. The oxide, SiO a , which is most abundant in 
 the form of sand, is decomposed by heating it with potas- 
 sium or magnesium, and silicon is thus set free. When 
 magnesium is used the action is violent, and besides the 
 silicon a compound of silicon and magnesium is formed. 
 
SILICON. 411 
 
 Silicon has also been made by heating the oxide and 
 carbon in the electric furnace, and by decomposing the 
 chloride with potassium : 
 
 SiCl 4 + 4K = Si + 4KC1. 
 
 The best way to make it is by heating together potassium 
 fluosilicate, K 2 SiF t , sodium, and zinc: 
 
 K 2 SiF 6 + 4Na = 4NaF + 2KF -f Si. 
 
 At the same time the zinc melts and the silicon which 
 separates dissolves in the molten zinc. On cooling, it is 
 deposited from the solution in beautiful needle-shaped 
 crystals, around which the zinc solidifies at a lower tem- 
 perature. By treating the mass with hydrochloric acid 
 the zinc is dissolved and the crystals of silicon are left 
 behind. When obtained by reduction of the oxide or the 
 chloride by means of potassium, it is a brown amorphous 
 powder. If made by decomposition of potassium fluo- 
 silicate by aluminium, it is deposited from the molten 
 aluminium in crystals somewhat resembling graphite. 
 Just as there are three forms of carbon, the amorphous, 
 graphite, and diamond, so there are three corresponding 
 forms of silicon, the amorphous brown powder, the 
 graphitoidal, and the needles. The amorphous variety 
 is converted into crystallized silicon by continued heat- 
 ing at a high temperature. 
 
 Amorphous silicon acts upon hydrofluoric acid, form- 
 ing silicon tetrafluoride, SiF 4 , and setting hydrogen free : 
 
 In this reaction it exhibits one of the properties of a 
 base-forming element. Towards other acids, however, it 
 is indifferent. It is not acted upon by sulphuric acid, 
 nor by nitric acid, nor aqua regia. It dissolves, however, 
 in potassium hydroxide, forming potassium silicate, in 
 this case acting like an acid-forming element : 
 
 Si + 2KOH + H,O = K 2 SiO 3 + 2H q . 
 
412 INORGANIC CHEMISTRY. 
 
 This form of silicon also burns in the air, forming the 
 dioxide. 
 
 Crystallized silicon, on the other hand, does not burn 
 in oxygen at the highest temperatures. It, however, re- 
 duces carbon dioxide and decomposes carbonates at a 
 high temperature. It is also oxidized by a melting mix- 
 ture of potassium nitrate and the hydroxide or carbonate. 
 It combines with nitrogen at a high temperature. 
 
 Both the graphitoidal and needle-formed crystals of 
 silicon consist of regular octahedrons. Both forms have 
 a blackish-gray color and a metallic lustre. 
 
 Silicon Hydride, SiH 4 . This gas is obtained mixed 
 with hydrogen when a compound of magnesium and sili- 
 con is treated with hydrochloric acid : 
 
 Mg 2 Si + 4HC1 = SiH 4 + 2MgCl 2 . 
 
 Thus made, it takes fire when it comes in contact with 
 the air, and the act is accompanied by explosion. The 
 products of its combustion are silicon dioxide and water. 
 When pure it forms a colorless gas which does not take 
 fire spontaneously in the air at the ordinary temperature. 
 If it is diluted with hydrogen, or if it is heated, it does 
 take fire. When burned in a cylinder or narrow tube, 
 so that free access of air is not possible, amorphous sili- 
 con is deposited upon the walls of the vessel. 
 
 Titanium, Ti (At. Wt. 47.79). Titanium occurs in nature 
 as titanium dioxide, TiO 2 , in three distinct forms, known 
 as rutile, brookite, and anatase ; in combination with iron, 
 as titaniferous iron which contains ferrous titanate, 
 JFeTiO 3 ; and in a number of iron ores and rare minerals. 
 The element is obtained in the free state by decomposing 
 potassium fluotitanate, K 2 TiF 6 , with potassium, just as 
 silicon is obtained by decomposing potassium fluosilicate, 
 . K 2 SiF 6 , with potassium or sodium. It burns when heated 
 in the air. It acts upon water at 100, causing the evolu- 
 tion of hydrogen. It is dissolved by hydrochloric acid, 
 forming the chloride, Ti 2 Cl 6 . At a high temperature it 
 unites directly with nitrogen as silicon does. Titanium 
 does not form a compound with hydrogen. 
 
SILICON TETRACHLOR1DE. 413 
 
 Zirconium, Zr (At. Wt. 89.72). The principal form in 
 which zirconium occurs in nature is as zircon, which is a 
 silicate of the formula ZrSiO 4 , derived from normal 
 silicic acid, Si(OH) 4 , by the replacement of the four hy- 
 drogen atoms by a quadrivalent atom of zirconium. The 
 element is obtained in the free condition by decomposing 
 potassium fluozirconate by heating it with aluminium to 
 a high temperature. In this way it is obtained in crystal- 
 lized form, somewhat resembling antimony. It does not 
 burn in the air. It is dissolved by hot concentrated hy- 
 drochloric acid, and when heated in a current of hydro- 
 chloric acid gas. The product is the tetrachloride, ZrCl 4 ; 
 and the same compound is formed when chlorine acts 
 directly upon zirconium. 
 
 Thorium, Th. (At. Wt. 230.87). This element occurs 
 principally in the mineral thorite, which is essentially a 
 silicate of thorium, ThSiO 4 , analogous to zircon. It is 
 obtained free by treating the chloride with silicon or 
 potassium. At high temperatures it burns in the air, 
 forming thorium dioxide, ThO 2 . 
 
 Cerium so much resembles the two elements lanthanum 
 and didymium, that although it falls in the same group as 
 silicon, and resembles the elements of this group in some 
 respects, it seems advisable to postpone its study until 
 lanthanum and didymium are taken up. 
 
 COMPOUNDS OF THE ELEMENTS OF THE SILICON GKOUP 
 
 WITH THOSE OF THE CHLOKINE GROUP. 
 
 Silicon Tetrachloride, SiCl 4 . This compound is formed 
 when silicon is heated in a current of chlorine, and by 
 passing a current of dry chlorine over a heated mixture 
 of silicon dioxide and carbon. Under these latter cir- 
 cumstances the following reaction takes place : 
 
 SiO, + 20 + 2C1, = SiCl 4 + 2CO. 
 
 Carbon acting alone upon silicon dioxide cannot reduce 
 it, nor has chlorine acting alone the power to convert it 
 into the chloride. When, however, carbon and chlorine 
 
414 INORGANIC CHEMISTRY. 
 
 act together both reactions take place. The tetrachlo- 
 ride is a colorless liquid. It is decomposed by water, 
 forming silicic acid and hydrochloric acid. The reaction 
 probably takes place as represented in the following 
 equation : 
 
 SiCl 4 + 4H 2 = Si(OH) 4 + 4HC1. 
 
 The normal acid thus formed breaks down very readily, 
 however, forming the ordinary acid of the formula 
 SiO(OH) 2 or H 2 SiO 3 , corresponding to carbonic acid, 
 H 2 CO, 
 
 Silicon Hexachloride, Si 2 Cl 6 , is formed when silicon 
 tetrachloride is heated with silicon : 
 
 3SiCl 4 + Si = 2Si 2 Cl e . 
 
 When heated to a sufficiently high temperature it is de- 
 composed, yielding silicon and the tetrachloride : 
 
 2Si 2 Cl 6 = 3SiCl 4 + Si. 
 
 "Water decomposes it, forming the corresponding hy- 
 droxyl derivative, which loses water and forms the acid 
 Si 2 2 (OH) 2 : 
 
 Si 2 Cl 6 + 6H 2 O = Si 2 (OH) 6 + 6HC1. 
 Si 2 (OH) 6 = Si 2 2 (OH) 2 + 2H 2 0. 
 
 The product is a disilicic acid, in some respects analo- 
 gous to disulphuric acid. 
 
 Similar compounds of silicon with bromine and iodine 
 are known. 
 
 Silicon Tetrafluoride, SiF 4 . This is one of the most 
 interesting of the compounds which silicon forms with 
 the members of Family VII. It is made by treating 
 silicon dioxide with hydrofluoric acid. This action is 
 secured by treating a mixture of silicon dioxide (sand) 
 and calcium fluoride (fluor-spar) with concentrated sul- 
 phuric acid, when two reactions take place : 
 
 CaF 2 + H 2 S0 4 = CaSO 4 + 2HF ; 
 Si0 2 + 4HF = 2H 2 + SiF 4 . 
 
FLUOSILICIC ACID. 415 
 
 The tetrafluoride escapes as a colorless gas, which forms 
 thick clouds in moist air on account of the action of 
 water upon it. 
 
 Water decomposes the tetrafluoride, as it does the 
 tetrachloride. The first action probably consists in the 
 formation of normal silicic acid and hydrofluoric acid, the 
 normal acid then breaking down by loss of water and 
 yielding the ordinary form of silicic acid : 
 
 SiF 4 + 4H 2 O = Si(OH) 4 + 4HF ; 
 Si(OH) 4 = SiO(OH) 2 + H 2 O. 
 
 The silicic acid thus formed separates as a gelatinous 
 mass. At the same time the hydrofluoric acid acts upon 
 some of the silicon tetrafluoride, forming the compound 
 fluosilicic acid, which has the formula H 2 SiF 6 : 
 
 SiF 4 + 2HF = H 2 SiF 6 . 
 
 The complete action may be represented in one equa- 
 tion, as follows : 
 
 3SiF 4 + 3H 2 O = H 2 Si0 3 + 2H 2 SiF 6 . 
 
 The fluosilicic acid remains in solution in the water, and 
 by treating this solution with carbonates or hydroxides 
 of the metallic elements the salts known as the fluosili- 
 cates are obtained. The solution of the acid can be con- 
 centrated to a certain extent in a platinum vessel, but it 
 breaks down into silicon tetrafluoride and hydrofluoric 
 acid when it becomes concentrated. If more potassium 
 hydroxide than is required to neutralize the acid is added 
 to the solution, decomposition ensues, with formation of 
 silicic acid : 
 
 H 2 SiF 6 + 6KOH = 6KF + H 2 SiO 3 + 3H 2 O. 
 
 By water alone, however, the acid is not decomposed, 
 and the salts are fairly stable. When heated, the salts 
 give off silicon tetrafluoride, and fluorides are left behind 
 
 K 2 SiF 6 = 2KF + SiF 4 . 
 
416 INORGANIC CHEMISTRY. 
 
 Constitution of Pluosilicic Acid. Attention has already 
 been called to the fact that fluosilicic acid and silicic acid 
 seem to be analogous substances, and that the former 
 may be regarded as derived from the latter by the sub- 
 stitution of six fluorine atoms for the three oxygen 
 atoms. According to this, fluorine has a valence higher 
 than one, and this accords with the fact that at the or- 
 dinary temperature the density of hydrofluoric acid is 
 greater than that required by the formula, HF. Assum- 
 ing, then, that fluorine may act as a bivalent or a tri- 
 valent element, and for the present purpose it is imma- 
 terial which view is taken, the relation between silicic 
 acid and fluosilicic acid is shown by the following for- 
 mulas : 
 
 Silicic acid Fluosilicic acid 
 
 It is commonly held that the acid is a " double com- 
 pound " made by the union of one molecule of silicon tetra- 
 fluoride with two molecules of hydrofluoric acid, and rep- 
 resented by the formula SiF 4 .2HF. This is not even an 
 attempt at an explanation of the fact that the composition 
 of the acid is so similar to that of silicic acid. The above 
 explanation is, however, in accordance with the composi- 
 tion of a large number of similar " double compounds," 
 in which not only fluorine, but chlorine, bromine, and 
 iodine enter. 
 
 Titanium Tetrachloride, TiCl 4 , is formed by the direct 
 action of chlorine on titanium, and also by passing dry 
 chlorine over a mixture of carbon and titanium dioxide. 
 It is, like silicon tetrachloride, a liquid. It forms crys- 
 tallized compounds with water. When heated with 
 water it is decomposed, yielding titanium dioxide and 
 hydrochloric acid : 
 
 TiCl 4 + 2H,0 = TiO, + 4HC1. 
 
 Titanium also forms with chlorine the compounds Ti 2 Cl 9 
 and Ti 2 Cl 4 . 
 
THORIUM TETRAFLUORIDE. 41? 
 
 Titanium Tetraihioride, TiF 4 , is formed in the same 
 way as silicon tetrafluoride, by treating a mixture of 
 titanium dioxide and fluor-spar with concentrated sul- 
 phuric acid, and by dissolving titanium dioxide in hydro- 
 fluoric acid. When treated with water it forms a com- 
 pound analogous to fluosilicic acid, called fluotitanic 
 acid, H 2 TiF 6 , which yields well characterized salts, the 
 fluotitanates. 
 
 Zirconium Tetrachloride, ZrCl 4 , is not completely decom- 
 posed by water, only half the chlorine being replaced by 
 oxygen, forming a product, zirconium oxychloride, ZrOCl 2 : 
 
 ZrCl 4 + H 2 = ZrOCl 2 + 2HC1. 
 
 This is in accordance with the fact that zirconium acts 
 as a base-forming as well as an acid-forming element. 
 The chlorides of silicon and titanium are completely de- 
 composed by water, as they are acid-forming. 
 
 The tetrafluoride of zirconium is obtained from zircon 
 or zirconium silicate by mixing the finely powdered 
 mineral with fluor-spar and passing hydrochloric acid 
 gas over it at a high temperature : 
 
 ZrSi0 4 + 2CaF 2 + 2HC1 = CaCl 2 + CaSiO 3 +ZrF 4 + H 2 O. 
 
 With metallic fluorides the tetrafluoride forms salts of 
 fluozirconic acid, H 2 ZrF 6 , analogous to fluosilicic and fluo- 
 titanic acids. 
 
 Thorium Tetrachloride, ThCl 4 , is not decomposed by 
 water at the ordinary temperature, but if its solution 
 is evaporated to dryness hydrochloric acid is given off 
 and thorium dioxide is left : 
 
 ThCl 4 + 2H 2 = Th0 2 + 4HC1. 
 
 Thorium Tetrafluoride, ThF 4 , is easily made by treat- 
 ing the tetrachloride with hydrofluoric acid. With po- 
 tassium fluoride it forms a salt of the formula K 2 ThF 6 , 
 or potassium fluothorate. The chloride also forms a simi- 
 lar salt with potassium chloride, potassium chlorthorate, 
 K 2 ThCl 6 . 
 
418 INORGANIC CHEMISTRY. 
 
 Comparison of the Chlorides of Family IV with those 
 of Family V. In studying the chlorides formed by the 
 members of the phosphorus group it was found that the 
 chlorides of phosphorus are readily decomposed by water, 
 forming the corresponding acids, and that the same is 
 true of the chloride of arsenic ; but that the trichlorides 
 of antimony and bismuth are only partly decomposed 
 by water, yielding oxychlorides. In the silicon group 
 we find now similar differences between the members 
 with low atomic weights and those with high atomic 
 weights. The chlorides of silicon and titanium are com- 
 pletely decomposed by water at the ordinary tempera- 
 ture, while that of zirconium is only half decomposed, 
 and that of thorium is not decomposed except at high 
 temperature. 
 
 COMPOUNDS OF THE MEMBERS or THE SILICON GROUP 
 WITH OXYGEN, AND WITH OXYGEN AND HYDROGEN. 
 
 Silicon Dioxide, SiO 2 . This compound occurs very 
 abundantly in nature in many different forms, both crys- 
 tallized and amorphous. Quartz is a form of crystallized 
 silicon dioxide which is found very widely distributed. 
 It crystallizes in the hexagonal system in prisms and 
 pyramids, the crystals sometimes attaining great size 
 and beauty. Another form of the crystallized compound 
 is that known as tridymite. Like quartz it crystallizes in 
 the hexagonal system, but the characteristic forms are 
 not the same as those of quartz. Further, it nearly 
 always occurs in triplet crystals. The finer crystals of 
 quartz are generally called rock-crystal ; the crystalline 
 variety in which the crystals are not well developed is 
 called quartzite. The amorphous varieties of silicon di- 
 oxide frequently contain water in combination, or, rather, 
 they are hydroxides of silicon. Examples of these forms 
 are opal, agate, amethyst, carnelian, flint, sand, chalced- 
 ony. Some of these are colored by small quantities of 
 other substances contained in them. Carnelian owes 
 its color to a compound of iron, probably ferric oxide ; 
 flint contains small quantities of organic matter. The 
 
SILICON DIOXIDE. 419 
 
 specific gravity of the crystallized varieties is higher 
 than that of the amorphous varieties, and there are also 
 some chemical differences between them which will be 
 referred to below. 
 
 Pure silicon dioxide can be made by melting sand or 
 a finely powdered silicate with sodium carbonate, when 
 sodium silicate is formed. This is soluble in water, and 
 when hydrochloric acid is added to the solution silicic 
 acid separates in the form of a gelatinous mass. By 
 evaporating the mass to complete dryness, moistening 
 with concentrated hydrochloric acid, and after a time 
 treating with water, everything dissolves except silicon 
 dioxide, which is perfectly pure and in a very finely di- 
 vided state. It can also be obtained pure by passing sili- 
 con tetrafluoride into water. As we have seen, a form 
 of silicic acid separates under these circumstances. 
 This, when filtered, dried, and ignited, yields perfectly 
 pure silicon dioxide. 
 
 Properties. Silicon dioxide is insoluble in water and 
 in most acids. It dissolves, however, in hydrofluoric 
 acid, forming the tetrafluoride. It requires the tempera- 
 ture produced by the oxyhydrogen blow-pipe to melt it. 
 The amorphous varieties are more easily acted upon by 
 other substances than the crystallized. Thus, hydro- 
 fluoric acid acts much more readily upon them. When 
 the amorphous compound is boiled with solutions of 
 potassium or sodium hydroxide, or of the carbonates of 
 these metals, it dissolves, forming the corresponding sili- 
 cate : 
 
 2KOH + SiO 2 = K 2 SiO 3 + H 2 O. 
 
 The crystallized varieties are not dissolved in this way. 
 All forms of the dioxide act upon melting hydroxides or 
 carbonates of potassium or sodium, and form the corre- 
 sponding silicates. 
 
 Uses. Plants take up silicon dioxide from the soil, 
 and this being deposited in various part of their tissues, 
 gives them the necessary firmness. - Straw, for example, 
 
420 INORGANIC CHEMISTRY. 
 
 is rich in silicon dioxide. Horse-tail, a plant of the 
 genus Equisetum, is so rich in finely divided silicon di- 
 oxide that it is used for polishing. There are great 
 natural deposits of finely divided silicon dioxide known 
 as infusorial earth. This consists of the remains of dia- 
 toms. And finally silicon dioxide is found in the hair, 
 in feathers, and in egg albumen. Silicon dioxide finds 
 extensive application in the manufacture of mortar, glass, 
 and porcelain. Ordinary glass, as we shall see, is a 
 silicate of calcium and potassium or sodium, which is 
 made by melting together sand and the carbonates of 
 the metals mentioned. 
 
 Silicic Acid. There are many varieties of silicic acid, 
 all of which can, however, be referred to the normal acid, 
 Si(OH) 4 . This normal acid is not known in the free 
 state in pure condition, but it is probably contained in 
 the gelatinous precipitate which is formed when silicon 
 tetrachloride or tetrafluoride is decomposed by water : 
 
 SiCl 4 + 4H 2 O = Si(OH) 4 + 4HC1. 
 
 This cannot, however, be isolated, as, even by standing, 
 it loses a molecule of water, and passes into the form 
 H 2 Si0 3 : 
 
 Si(OH) 4 = OSi(OH) 2 + H 2 O. 
 
 This is the form from which most of the ordinary sili- 
 cates are derived. It cannot be isolated in pure con- 
 dition, for when filtered off and exposed to the air it 
 loses more water, and when heated to a sufficiently high 
 temperature it is converted into silicon dioxide. 
 
 OSi(OH), = Si0. 2 + H 2 O. 
 
 When potassium or sodium silicate in solution is treated 
 with hydrochloric acid, most of the silicic acid separates 
 in the form of a gelatinous mass if the solution is con- 
 centrated. If, however, the solution is dilute, a consid- 
 erable part of the acid remains in solution. Further, if 
 a concentrated solution of the silicate of potassium or 
 
SILICIC ACID. 421 
 
 sodium is poured quickly into hydrochloric acid, or if 
 the acid is poured quickly into the solution of the sili- 
 cate, the silicic acid remains in solution. If, however, the 
 solutions are brought together drop by drop the silicic 
 acid separates. From these solutions of silicic acid am- 
 monia or ammonium carbonate throws down the acid. 
 
 A solution of pure silicic acid can be obtained by 
 means of dialysis. It has been found that solutions of 
 different substances pass with different degrees of ease 
 through porous membranes, just as gases differ as re- 
 gards the ease with which they pass through porous dia- 
 phragms. This fact concerning gases was referred to in 
 connection with hydrogen. Now, while some solutions 
 pass readily through parchment paper, others pass 
 through with difficulty, and some do not pass through at 
 all. A dialyser, or an apparatus used in dialysis, may be 
 made by tying a piece of parchment paper over the mouth 
 of a ring-formed glass or rubber vessel, and placing this in 
 another shallow vessel. Pure water is put in the outer 
 vessel, and the solution for dialysis in the inner one. 
 The arrangement is illustrated in Fig. 13. 
 
 FIG. 13. 
 
 In the figure aa is the hoop of gutta-percha, and 5 is the 
 parchment paper. When now the solution containing hy- 
 drochloric acid, sodium chloride, and silicic acid is put in 
 the dialyser, the hydrochloric acid and sodium chloride 
 pass readily through the membrane, while the silicic acid 
 is left behind, and in the course of a few days, if the water 
 in the outer vessel is renewed, the solution of silicic 
 acid in the inner vessel will be found to be free from the 
 other substances. This solution can be evaporated to 
 
422 INORGANIC CHEMISTRY, 
 
 some extent by boiling, but when a certain concentration 
 is reached the acid separates. In a vacuum such a solu- 
 tion can be evaporated further without the formation of 
 a deposit. Finally, there is left a transparent mass which 
 has approximately the composition represented by the 
 formula H 2 SiO 3 . The dialysed solution of silicic acid is 
 coagulated by a very dilute solution of sodium or potas- 
 sium carbonate, and by carbon dioxide itself. 
 
 When the solutions containing silicic acid are evapo- 
 rated to complete dryness the acid is converted into sili- 
 con dioxide and other insoluble hydrates. This residue 
 is called insoluble silicic acid. When this is treated with 
 hydrochloric acid and water it remains undissolved, and 
 if filtered off and ignited it leaves a residue of silicon di- 
 oxide. To sum up, then : Whenever silicic acid is formed 
 in a solution it is a more or less complex derivative of 
 normal silicic acid, and is somewhat soluble in water, but 
 by the processes just described the soluble acid is con- 
 verted into insoluble silicic acid, as explained. 
 
 Poly silicic Acids. Silicic acid is remarkable for the 
 great number of derivatives which it yields. Most of 
 these bear to the normal acid relations similar to those 
 which the various forms of phosphoric acid bear to nor- 
 mal phosphoric acid, and the various forms of periodic 
 acid to normal periodic acid. It has already been stated 
 that salts of the acid H 2 SiO 3 are more common than 
 those of the normal acid. Among the salts of the normal 
 acid are zircon, ZrSiO 4 , and thorite, ThSiO 4 . The or- 
 dinary silicates of potassium and sodium are derived 
 from the acid H 2 SiO 3 ; so also are wollastonite, CaSiO 3 , 
 and enstatite, MgSiO 3 . 
 
 Disilicic Acid is derived from "ordinary silicic acid by 
 loss of one molecule of water from two molecules of the 
 acid : 
 
 Its composition is, therefore, H^Si^O,, which may be 
 W 
 
TRISIL1C1C ACIDS. 423 
 
 written O 3 Si,(OH) 2 . Another form of disilicic acid is de- 
 rived from two molecules of the normal acid by loss of 
 one molecule of water : 
 
 2Si(OH) 4 = OSi,(OH) 6 + H 2 0. 
 
 The well-known mineral serpentine is apparently the 
 magnesium salt of this acid. It is represented by the 
 formula Mg,Si 9 O,. 
 
 Trisilicic Acids are derived from three molecules of 
 the normal acid or the ordinary acid by loss of different 
 numbers of molecules of water. Thus, by loss of two 
 molecules the normal acid would yield a product H fl Si 3 O 10 . 
 By loss of two molecules of water this trisilicic acid 
 would yield an acid of the formula H 4 Si 3 O 8 . The struc- 
 ture of the first acid is expressed by formula I, and of 
 the second by formula II, below given : 
 
 Si 
 
 Sij(OH), 
 
 i (OH), Si 
 
 OH (OK 
 
 O Si J 
 
 O O 
 
 O Si]o 
 O O 
 
 (OH), Si|o 8 Al 
 n. in. 
 
 Orthoclase or ordinary feldspar is the aluminium- 
 potassium salt of the second form of trisilicic acid, in 
 which one atom of hydrogen is replaced by potassium, 
 and three by an aluminium atom, as shown in formula 
 III above, 
 
 Titanium Dioxide, TiO a As has been stated, this is 
 one of the principal forms in which titanium is found in 
 nature. There are three natural crystallized varieties 
 rutile, brookite, and anatase. In order to prepare the 
 pure dioxide from one of the natural forms, it is melted 
 in finely powdered condition with potassium carbonate, 
 when it is converted into potassium titanate, K 2 TiO 3 , the 
 reaction being entirely analogous to that which takes 
 place when silicon dioxide is treated in the same way : 
 
 K 2 C0 3 + Si0 2 = K 2 Si0 3 + C0 2 ; 
 K 2 C0 3 + TiO 3 = K 2 Ti0 3 + C0 a . 
 
424 INORGANIC CHEMISTRY. 
 
 When titanic acid is precipitated from a solution of a 
 titanate it appears as a hydroxide, the composition of 
 which varies from Ti(OH) 4 , or normal titanic acid, to 
 H 2 Ti 2 O 5 , a dititanic acid. When these substances are 
 ignited they yield titanium dioxide. The hydroxides of 
 titanium conduct themselves somewhat like those of sili- 
 con. They are to some extent soluble in water, and when 
 these solutions containing sulphuric acid are much 
 diluted and boiled, the titanium is all precipitated as a 
 hydroxide. Titanium dioxide forms some salts with 
 acids, among which the following are examples Ti(SO 4 ) a 
 and TiO(SO 4 ). The former is normal titanium sulphate, 
 the latter titanyl sulphate, in which the bivalent group, 
 TiO, or titanyl, takes the place of two hydrogen atoms. 
 
 Zirconium Dioxide, ZrO 2 , is obtained by a rather com- 
 plicated series of reactions from zircon. It dissolves in 
 molten potassium or sodium carbonate, forming the cor- 
 responding zirconate, K 2 ZrO 3 . The sodium salt of nor- 
 mal zir conic acid, Na 4 ZrO 4 , has also been obtained. 
 
 The dioxide forms salts with acidrf, among which two 
 of the sulphates are of special interest. One has the 
 composition ZrSO 6 , and the other Zr 3 (SO 6 ) 2 . The former 
 is to be regarded as the salt of the acid OS(OH) 4 , formed 
 by substituting one atom of zirconium for the four atoms 
 of hydrogen ; the other is the salt of normal sulphuric 
 acid, S(OH) e , formed by substituting zirconium for all 
 the hydrogen. 
 
 Thorium Dioxide, ThO 2 , does not form thorates as the 
 dioxides of the other members of the group. It does, 
 however, form salts with acids. In these, thorium acts 
 as a quadrivalent element. 
 
 Silicides are compounds of silicon with other elements, 
 as, for example, with carbon. These two elements com- 
 bine forming an extremely interesting compound carbon 
 silicide, CSi, which is manufactured on the large scale and 
 known in the market as carborundum. This is made by 
 heating a mixture of quartz sand, coke, and common 
 salt, or sodium chloride, in the electric furnace to 3500, 
 when the reaction represented below takes place : 
 
 SiO, + 20 = CSi + 2CO. 
 
FAMILY IV, GROUP B. 425 
 
 The product is in the main crystallized, the crystals 
 being bluish or yellowish-green. They have the specific 
 gravity 3.22 to 3.12. The silicide is said to be colorless 
 when perfectly pure. It scratches ruby and chrome- 
 steel, and on account of its hardness it is much prized 
 as a polishing agent, being used to a considerable ex- 
 tent in place of emery. Pure carbon silicide is insoluble 
 in nearly all ordinary solvents, including hydrochloric, 
 nitric, sulphuric, and hydrofluoric acids. It is, however, 
 decomposed by fusing caustic alkalies or their car- 
 bonates. 
 
 Many other silicides have been made, several of which 
 are well crystallized compounds. 
 
 FAMILY IV, GROUP B. 
 
 Allied to the members of the silicon group, yet differ- 
 ing from them in some important particulars, are the 
 three elements germanium, tin, and lead. Of these the 
 first two are more acidic in character than the last. The*y 
 combine with chlorine in two proportions, forming the 
 chlorides GeCl a , SnCl a , PbCl,', GeCl 4 , SnCl 4 , PbCl 4 . With 
 oxygen they unite, forming the compounds GeO 2 , SnO 2 , 
 and PbO 2 . Stannic oxide, SnO 2 , and lead peroxide, PbO 2 , 
 form salts with bases, and these have the composition 
 represented by the general formulas M 2 SnO 3 and M 2 PbO 3 , 
 and are therefore analogous to the silicates and titanates. 
 On the other hand, further, salts are known which are 
 derived from the oxide PbO. These have the general 
 formula M 2 PbO 2 , and are to be regarded as salts of an 
 acid, Pb(OH) 2 . These salts are not stable, and are not 
 easily obtained. Most of the derivatives of lead are 
 those in which it plays the part of a base-forming ele- 
 ment. It will therefore be better to postpone its study 
 until it is taken up under the general head of the base- 
 forming elements. Notwithstanding, further, the marked 
 analogy between some of the compounds of tin and those 
 of the members of the silicon group, it appears on the 
 whole advisable to treat of this element in company with 
 lead, which it also resembles in many respects. 
 
CHAPTER XXIII. 
 
 CHEMICAL ACTION. 
 
 Retrospective. "We have been studying the principal 
 elements of four families and the compounds which they 
 form with one another. No matter how simple or how 
 complex the chemical changes studied were, certain fun- 
 damental laws governing all cases of chemical action were 
 found to hold good. These laws have been discussed, 
 but it will be well to recall them here before taking 
 up other laws which are intimately connected with them. 
 The first great law of chemical change is 
 
 I. The laiv of conservation of mass. 
 
 * According to this, the amount of matter is not changed 
 by a chemical act. 
 The second law is 
 
 II. The law of definite proportions. 
 
 According to this, the composition of every compound 
 is always the same. 
 The third law is 
 
 III. The law of multiple proportions. 
 
 According to this, the different masses of any element 
 which combine with a fixed mass of another or others 
 bear simple relations to one another. 
 
 To account for the laws of definite and multiple pro- 
 portions the Atomic Theory was proposed. 
 
 According to this, each element is made up of particles 
 of definite weight, which are chemically indivisible, and 
 chemical action consists in union or separation of these 
 particles. These hypothetical particles are called atoms. 
 The elements must combine in the proportion of their 
 atomic weights or of simple multiples of these, if the 
 atomic theory is true. 
 
 Further study showed that it is necessary to assume 
 
 (426) 
 
CHEMICAL ACTION. 427 
 
 the existence of larger particles than the atoms, viz., the 
 molecules. According to the theory of molecules, every 
 chemical compound and element is made up of mole- 
 cules, which are the smallest particles having the same 
 general properties as the mass. These molecules are 
 made up of atoms which, in the case of compounds, are 
 of different kinds, and in the case of elements, of the 
 same kind. In the case of a few elements the atom ap- 
 pears to be identical with the molecule. 
 
 From the study of gases the conclusion is reached that 
 in equal volumes of all gases under standard conditions 
 there is always the same number of molecules (Avoga- 
 dro's law). This gives us a means of determining the 
 relative weights of molecules of gaseous substances ; and 
 from these molecular weights it is possible to draw con- 
 clusions in regard to the atomic weights of those elements 
 which enter into the composition of the compounds thus 
 studied. 
 
 The formulas of chemical compounds are intended to 
 be molecular formulas. They are intended to tell of what 
 atoms and of how many atoms the molecules represented 
 are made up. 
 
 The method of determining molecular weights based 
 upon Avogadro's law is applicable only to gaseous sub- 
 stances, or to such as can be converted into gas without 
 undergoing decomposition. While many of the com- 
 pounds with which we have had to deal are of this char- 
 acter, many of them are not, and in regard to the mole- 
 cular weights of these, we must be in doubt unless some 
 other method applicable to liquids and solids is avail- 
 able. So, too, the atomic weights of those elements which 
 enter into the composition of gaseous compounds can be 
 deduced from the molecular weights, but plainly those 
 which do not enter into the composition of such com- 
 pounds demand some other method. For determining 
 the atomic weights of such elements an excellent method 
 is based upon the study of specific heats ; while for the 
 determination of the molecular weights of solid substances 
 which can be dissolved without decomposition a method 
 has quite recently come into play which is based upon 
 
428 INORGANIC CHEMISTRY. 
 
 the extent to which the compound raises the boiling, 
 point or lowers the freezing-point of its solution. Both 
 these methods will be briefly described in this chapter. 
 
 Next, it is found that there is a limit to the law of 
 multiple proportions. "While, according to this law, the 
 masses of any element which unite with a given mass of 
 another element bear simple relations to one another, 
 the law is silent as to frow many kinds of compounds are 
 possible between any two elements. A careful examina- 
 tion of the composition of the compounds of the ele- 
 ments shows, however, that there is a limit to the num- 
 ber of atoms of one element which can combine with 
 one atom of another element. This limit is determined 
 by what is called the valence of the elements. Observa- 
 tions on the composition of compounds led to the hy- 
 pothesis of the linking of atoms the linking taking place 
 according to the laws of valence. The arrangement of 
 the atoms in a molecule is, according to this, the consti- 
 tution of a compound. 
 
 Yalence, as we have seen, is not a constant property 
 of the atoms. Towards oxygen the elements which we 
 have thus far studied have the highest valence ; towards 
 hydrogen the lowest ; and, in general, towards tjie mem- 
 bers of the chlorine group they exhibit an intermediate 
 valence. The valence towards hydrogen is in most 
 cases constant, while the valence towards oxygen and 
 towards the members of the chlorine group varies, in 
 some cases between comparatively wide limits, as between 
 1 and 7 in the chlorine group, and between 2 and 6 in the 
 sulphur group. Further, the variations in the valence 
 of an element generally take place from odd to odd or 
 from even to even. In the case of chlorine it appears to 
 vary from 1 to 3 to 5 to 7 ; in that of sulphur, from 2 to 
 4 to 6 ; in that of phosphorus, from 3 to 5. A knowledge 
 of the valence of the elements is of great assistance in 
 dealing with their compounds, as, knowing their valence, 
 we know in general the composition of their principal 
 compounds. 
 
 A comparison of the atomic weights finally led to the 
 discovery that the properties of the elements are a pen- 
 
CHEMICAL ACTION. 429 
 
 odic function of these weights. This is the great periodic 
 laiv of chemistry. This makes a systematic classification 
 of the elements according to their atomic weights and 
 their properties possible, and is full of suggestion as to 
 the relations which the forms of matter we call elements 
 bear to one another. 
 
 Classification of Reactions of the Elements and Com- 
 pounds Studied. While there is undoubtedly something 
 confusing in the number of the compounds and their reac- 
 tions which we have been studying, still, when these are 
 interpreted in the light of the atomic theory, of the law of 
 valence, and of the periodic law, the study is much sim- 
 plified, and those things which seem to have little or no 
 connection are found to form parts of a general system. 
 In studying chemistry, one of the first things to be done 
 is to learn how elements and compounds act upon one 
 another, and what products are formed. The question 
 of composition is one of the first which presents itself, 
 and this must be studied before other questions can be 
 intelligently discussed. What, then, are the most promi- 
 nent facts which we have learned in studying the ele- 
 ments and compounds which have thus far been taken 
 up? 
 
 In the first place, it will have been noticed that, gener- 
 ally speaking, the compounds which any element forms 
 with oxygen and hydrogen are the most prominent ; that, 
 taking the maximum oxygen compound of an element as 
 one end of a series, the other end is formed by the hy- 
 drogen compound. These end-products in the case of 
 chlorine, sulphur, phosphorus, and silicon are : 
 
 Hydrogen compound. Maximum oxygen compound. 
 
 HC1 C1 2 O 7 
 
 H 2 S S0 3 (S 2 6 ) 
 
 H 3 P PA 
 
 H 4 Si Si0 2 (Si 2 O 4 ) 
 
 The valence towards hydrogen increases while that 
 towards oxygen decreases regularly in the order given. 
 With water these oxides form the acids HC1O 4 , H 2 SO 4 , 
 H.PO 4 , and H 4 SiO 4 . Here the remarkable fact is ob- 
 
430 INORGANIC CHEMISTRY. 
 
 served, that the number of hydrogen atoms in each oi 
 these acids is the same as that in the hydrogen com- 
 pounds, and the limit of the addition of oxygen is reached 
 in each case with four atoms of oxygen. Further, each 
 of the first three acids appears to be related to so-called 
 normal acids which are formed by union of the chlorine, 
 sulphur, and phosphorus with a number of hydroxyl 
 groups corresponding to the oxygen valence. These 
 normal acids are 
 
 Cl(OH),, S(OH) 6) P(OH),, Si(OH),. 
 
 Now, whenever a chlorine compound is subjected to 
 oxidation under proper circumstances the final product 
 is perchloric acid, which when isolated has probably the 
 composition represented by the formula HC1O 4 . So 
 when a sulphur compound is oxidized the final product 
 is sulphuric acid, H 2 SO 4 ; when a phosphorus compound 
 is fully oxidized the final product is phosphoric acid,, 
 H 3 PO 4 ; and the final product of oxidation of a silicon 
 compound is silicic acid, H 4 SiO 4 . 
 
 By reduction of the above compounds the final prod- 
 ucts are the hydrogen compounds ; but before the limit 
 of reduction is reached intermediate products are formed. 
 All these intermediate products are comparatively un- 
 stable, and tend to take up oxygen under ordinary cir- 
 cumstances and to form the stable derivative of the 
 highest oxygen compound. Thus phosphites pass over 
 into phosphates, sulphites into sulphates, and chlorates 
 into perchlorates when heated. These changes are repre- 
 sented by the following equations : 
 
 2KC1O 3 = KC1 + KC10 4 + 2 ; 
 4K 2 S0 3 = K 2 S + 3K 2 S0 4 ; 
 4H 3 PO 3 = PH 3 + 3H 3 PO 4 . 
 
 The highest forms are therefore evidently most stable. 
 Turning to the compounds which the elements of Families 
 IY, Y, YI, and YII form with the members of the chlorine 
 group, attention has repeatedly been called to the fact that 
 
DIRECT COMBINATION. 431 
 
 these are for the most part decomposed by water with 
 the formation of the corresponding hydroxyl compounds. 
 
 The elements of Families IY, Y, VI, and VII do not 
 form compounds with the members of the sulphur group, 
 nor with those of the nitrogen group, as readily as they do 
 with hydrogen, with oxygen, and with the members of the 
 chlorine group. Those elements which have basic char- 
 acter, however, like antimony and bismuth, form very 
 characteristic compounds with sulphur. The sulphur 
 compounds, in general, have a composition similar to 
 that of the oxygen compounds of the same elements. 
 
 Kinds of Chemical Reactions. As was pointed out in 
 the early part of this book, all chemical reactions may be 
 classified under three heads : 
 
 (1) Those which consist in direct combination ; 
 
 (2) Those which consist in direct decomposition ; and, 
 
 (3) Those which involve the interaction of two or more 
 elements or compounds and the formation of two or more 
 compounds. This is known as double decomposition or 
 metathesis. 
 
 Direct Combination. We have had to deal with a 
 number of examples of each of these kinds of reactions. 
 As examples of the first kind already studied the fol- 
 lowing may be mentioned : 
 
 The combination of hydrogen and chlorine to form 
 hydrochloric acid ; the formation of ammonium chloride 
 from ammonia and hydrochloric acid ; the formation of 
 calcium hydroxide from calcium oxide and water ; the 
 formation of nitrogen peroxide from nitric oxide and 
 oxygen ; and the formation of carbon disulphide from 
 carbon and sulphur. 
 
 As regards the combination of hydrogen and chlorine, 
 it should be remarked that this act is the. same in princi- 
 ple as that of metathesis. Strictly speaking, it is not a case 
 of direct combination, as we understand it. For, as we 
 have seen, according to the molecular theory, free chlo- 
 rine and free hydrogen consist of molecules which are 
 made up of two atoms each. Therefore, when these ele- 
 ments are brought together the molecules are first de- 
 composed into atoms before the act of union can take 
 
432 INORGANIC CHEMISTRY. 
 
 place. The two acts are represented by the two equa* 
 tions following : 
 
 C1 2 + H 2 = Cl + Cl + H + H ; 
 d -f Cl + H + H = 2HC1. 
 
 In the case also of the union of hydrochloric acid 
 and ammonia it appears probable that a serious disar- 
 rangement of the constituent atoms is necessary in order 
 that the act of combination may take place. According 
 to the ammonium theory, ammonium chloride is repre- 
 
 sented by the formula 
 
 H 
 H 
 
 H, which means that the 
 
 H 
 
 Cl 
 
 atom of chlorine and four atoms of hydrogen are in com- 
 bination with the atom of nitrogen. But in order that 
 a compound of this constitution may be formed from 
 ammonia and hydrochloric acid, it is necessary that the 
 molecule of hydrochloric acid should be broken down 
 into its constituent atoms. So that this case of apparent 
 direct combination is, as far as we can judge, in reality 
 more complicated than it appears, and should be repre- 
 sented by the two equations : 
 
 NH 3 + HC1 = NH 3 + H + 01; 
 NH. + H + Cl = NH 4 C1. 
 
 All other cases of apparent direct combination are 
 probably of the same character, so "that it is doubtful 
 whether a single case of simple direct combination is 
 known. 
 
 Direct Decomposition. As examples of direct decom- 
 position the following cases may be cited : 
 
 The decomposition of mercuric oxide by means of 
 heat into mercury and oxygen ; that of ammonium chlo- 
 ride into ammonia and hydrochloric acid by heat ; that 
 of potassium nitrate into potassium nitrite and oxygen 
 
METATHESIS. 433 
 
 by 'heat ; that of phosphorus pentachloride into the 
 trichloride and chlorine by heat ; that of ammonia into 
 hydrogen and nitrogen by continued action of electric 
 sparks ; and that of nitrogen iodide by contact with a 
 solid substance. 
 
 On close examination of each of the above cases, which 
 are fairly typical and as simple as any that could be 
 chosen, it will be seen that no one of them is merely a 
 case of decomposition ; for even though we must assume 
 that the first result in each case is the setting free of the 
 atoms of one or two elements, we must also assume that 
 these atoms unite again to form other molecules either 
 of elements or compounds. Thus, when mercuric oxide 
 is decomposed we get mercury and oxygen. As far as 
 can be determined, the mercury atoms do not unite with 
 each other, but the oxygen atoms do, so that the total ac- 
 tion involves decomposition and afterwards combination 
 as represented in the equations 
 
 In the case of the pentachloride of phosphorus, it is 
 probable that the two atoms of chlorine are first given 
 off from each molecule of the chloride, leaving a molecule 
 of the trichloride, but the atoms of chlorine afterwards 
 unite to form molecules as represented thus : 
 
 PC1 5 = PC1 3 + Cl + Cl ; 
 PC1 3 + Cl + Cl = PC1 3 + C1 2 . 
 
 Similar statements hold good for all other cases of 
 direct decomposition. 
 
 Metathesis. This is the most common kind of chemi- 
 cal action, and indeed from what has been said in regard, 
 to direct combination and direct decomposition it will be 
 seen that there is no essential difference between them 
 and metathesis. Most of the reactions with which we 
 have had to deal are examples of double decomposition 
 
434 INORGANIC CHEMISTRY. 
 
 or metathesis, as : The formation of salts by the action of 
 bases upon acids ; the formation of the sulphides of 
 arsenic, antimony, and bismuth by the action of hydro- 
 gen sulphide upon solutions of compounds of these ele- 
 ments ; the setting free of hydrochloric and nitric acids 
 by the action of sulphuric acid upon chlorides and 
 nitrates ; of carbon dioxide and oxides of nitrogen by the 
 action of acids upon carbonates and nitrites ; and of am- 
 monia by treating ammonium salts with lime. Among 
 the more complicated examples which have been dealt 
 with are : The action of sulphuric acid upon potassium 
 iodide, giving rise to the formation of potassium sul- 
 phate, hydriodic acid, free iodine, sulphur dioxide, sul- 
 phur, and hydrogen sulphide ; the action of chlorine 
 upon a mixture of silicon dioxide and charcoal ; the action 
 of silicon fluoride upon water, giving rise to the forma- 
 tion of silicic acid and fluosilicic acid ; and the action 
 of phosphorus pentachloride upon water, forming phos- 
 phoric and hydrochloric acids. As simple an example 
 of this kind of action as can be cited is that of the for- 
 mation of hydrogen and potassium chloride from potas- 
 sium and hydrochloric acid gas. The molecular weight 
 of potassium is not positively known, but, assuming its 
 molecule to be made up of two atoms, the action must be 
 represented in this way : 
 
 K 2 + 2HC1 = 2KC1 + H 2 . 
 
 The next stage of complication is exhibited in the re- 
 action following : 
 
 KI + HC1 = KC1 + HI. 
 
 Examples similar to the latter, but somewhat more com- 
 plicated, are these : 
 
 2KOH + H 2 S0 4 = K 2 S0 4 + HjO ; 
 CaCl 2 + H 2 S0 4 = CaSO 4 + 2HC1. 
 
 The Cause of Chemical Reactions. The prime cause of 
 chemical reactions is something which we think of as an 
 
AN IDEAL CHEMICAL REACTION. 435 
 
 attractive force exerted in different degrees between the 
 different elements. When any elements or compounds 
 are brought together under certain conditions the ten- 
 dency is always towards the formation of the most stable 
 compounds of those elements which can be formed un- 
 der the given conditions. Thus, potassium sulphate and 
 water are more stable forms of combination of the ele- 
 ments hydrogen and oxygen, and potassium, sulphur and 
 oxygen, than sulphuric acid and potassium hydroxide 
 are under the conditions under which the action takes 
 place. So also the system composed of potassium chlo- 
 ride and hydriodic acid is more stable than that com- 
 posed of potassium iodide and hydrochloric acid under 
 the conditions of the action. Why the one system is 
 more stable than the other we do not know, for we do 
 not know what relations exist between the atoms in the 
 molecules. It is convenient to think of that which 
 causes the atoms to unite to form compounds as chemical 
 affinity. It is evident that this affinity is more strongly 
 exerted between some elements than between others. 
 The affinity of chlorine for hydrogen is, for example, 
 much stronger than that of chlorine for nitrogen or for 
 oxygen. Owing, however, to the complicated character 
 of most chemical reactions, it is extremely difficult to 
 make measurements of the affinities of the elements, and 
 but little progress has been made in this direction. Still, 
 one of the great objects in view in the study of chemical 
 phenomena is to learn as much about chemical affinity 
 as possible. 
 
 An Ideal Chemical Reaction. In every case in which 
 two compounds act upon each other to form two new 
 ones, several forces must be at work, as we have seen. 
 Suppose, for example, AB and CD act upon each other 
 in the gaseous condition to form two compounds BC and 
 AD, also both gaseous. The normal course of such a 
 reaction would lead to the formation of not only the two 
 compounds BC and AD, but AB and CD would also be 
 present in the resulting system. For A has an affinity 
 for B as well as for D, and C has an affinity for D as 
 well as for B. In the system we should have operating 
 
436 INORGANIC CHEMISTRY. 
 
 the affinity of A for B, and A for D ; of C for D, and oi 
 C for J?. As these operate simultaneously, equilibrium is 
 established when certain quantities of the four possible 
 compounds are formed, the quantities depending in the 
 first instance upon the relative strengths of the various 
 affinities. The same remarks apply to the case in which 
 two substances react in solution and form two products 
 which are soluble. Here the action is not complete in 
 any one direction, but an equilibrium is established be- 
 tween the four possible compounds. 
 
 Influence of Mass. The proportions between the pro- 
 ducts formed in any given case is markedly influenced 
 by the relative masses of the reacting substances. Thus, 
 sulphuric acid acts upon potassium nitrate when the 
 acid is in excess, forming primary potassium sulphate, 
 KHSO 4r and nitric acid. On the other hand, if a large 
 excess of nitric acid is allowed to act upon primary 
 potassium sulphate, sulphuric acid and potassium nitrate 
 are produced. Much attention has been given to the 
 study of mass action, and the result is to show that in 
 reactions generally this kind of action comes prom- 
 inently into play. The law has been established that 
 chemical action is proportional to the product of the active 
 masses of the substances taking part in the change. It would 
 appear from this that the decomposition of two com- 
 pounds to form two new ones would not be complete, if 
 the conditions are such that the two new compounds can 
 act upon each other. If a large excess of one of the re- 
 acting compounds is taken, however, the reaction may be 
 made approximately complete by reason of the mass 
 action. 
 
 Reactions May be Complete if one of the Products 
 Formed is Insoluble or Volatile. When two substances 
 which by interaction can form an insoluble product are 
 brought together, the reaction generally takes place and 
 is complete. When the substances are brought together 
 we may imagine that, owing to interaction, a small quan- 
 tity of the insoluble compound is formed at once. If 
 this product were soluble, the action would stop before 
 it is complete, because this new product would itself 
 
WHEN REACTIONS MAT BE COMPLETE. 437 
 
 exert its action upon the system. Being insoluble, how- 
 ever, it is removed from the sphere of action, and the 
 same reaction which caused the formation of the first 
 particles of it can now be repeated, and so on, until the 
 reaction is complete. This is illustrated in the action of 
 sulphuric acid upon barium chloride in solution. The 
 two substances react as represented in this equation : 
 
 BaCl 2 + H 3 S0 4 + Aq = BaSO 4 + HC1 + Aq. 
 
 The symbol Aq is simply intended to indicate that the 
 reaction takes place in solution. If barium sulphate were 
 soluble, all four substances barium chloride, sulphuric 
 acid, barium sulphate, and hydrochloric acid would be 
 present in the solution after the establishment of equi- 
 librium. But, being insoluble, it is removed, and new 
 quantities are formed as long as the substances necessary 
 for its formation are present in the solution ; that is, until 
 either all the barium chloride is decomposed or all the 
 sulphuric acid is removed. Reactions involving the for- 
 mation of insoluble compounds or precipitates are among 
 the most common with which we have to deal, particu- 
 larly in the various operations of analytical chemistry. 
 
 Again, when two substances which can form a volatile- 
 product are brought together the reaction generally takes 
 place and is complete. The reason why a reaction of 
 this kind is complete is the same as that given in the 
 case of the formation of an insoluble compound. Each 
 successive portion of the volatile product formed is re- 
 moved, and the reaction which gave rise to it proceeds as 
 long as the necessary substances are present. This kind 
 of action has been repeatedly illustrated. It is that, for 
 example, which is seen in the liberation of hydrochloric 
 acid from a chloride by the action of sulphuric acid ; of 
 carbon dioxide by the action of an acid upon a carbonate ; 
 and of ammonia by the action of lime upon ammonium 
 chloride. 
 
 An interesting example of the combined influence of 
 mass and the volatility of the product is seen in the action 
 of heated iron upon an excess of steam, and of the oxide 
 
438 INORGANIC CHEMISTRY. 
 
 of iron upon an excess of hydrogen. When steam is 
 passed over heated iron, action takes place thus : 
 
 4H a O + 3Fe = Fe 3 O 4 + 4H 2 . 
 
 Hydrogen is liberated and the oxide of iron formed. 
 When, however, hydrogen is passed over heated oxide 
 of iron the reverse reaction takes place : 
 
 Fe 3 O 4 + 4H 2 = 3Fe + 4H 2 O. 
 
 Owing to the excess of steam always present in the first 
 reaction, hydrogen is constantly formed and constantly 
 being removed. Undoubtedly the hydrogen formed acts 
 to some extent upon the oxide, but the other reaction 
 always takes place to a greater extent. The opposite is 
 true when the oxide is heated in an excess of hydrogen. 
 The principal reaction which takes place in this case is 
 that of the hydrogen upon the oxide of iron, and the 
 steam is carried out of the field almost as soon as formed, 
 so that the reduction of the oxide of iron continues. 
 
 Thermochemical Study of Affinity . If a mass of hydro- 
 gen and a mass of chlorine consisted of isolated atoms at 
 rest, and, after combination, the molecules as well as 
 their constituent atoms were at rest, then the heat evolved 
 in the act of combination would be the result of the trans- 
 formation of the potential energy of the atoms into kinetic 
 energy, and it would be a measure of the affinity exerted 
 between the atoms. But none of these conditions can 
 be assumed with any confidence, and most of them un- 
 doubtedly do not exist. We have abundant evidence to 
 show that the mass of hydrogen and that of chlorine do 
 not consist of isolated atoms. Taking, then, the reaction 
 between hydrogen and chlorine, it is clear, as lias already 
 been explained, that it is not simply a combination of 
 atoms, but that the act of combination between the atoms 
 must be preceded by the decomposition of the molecules 
 of hydrogen and those of chlorine. The heat which is 
 evolved in the reaction is therefore not simply the result 
 of the combination of hydrogen and chlorine, but it is 
 
VALUE OF THERMOCHEMICAL MEASUREMENTS. 439 
 
 this heat less that which is required to decompose the 
 molecules of hydrogen and those of chlorine into atoms. 
 The heat measured is the difference between two quan- 
 tities ; and we have no means of estimating the value of 
 these quantities. This is true of every chemical reaction. 
 The heat evolved or absorbed in the reaction is the dif- 
 ference between two or more quantities, and it is not 
 therefore a measure of affinity. 
 
 Nevertheless, some knowledge regarding the relations 
 which the affinities of elements bear to one another can 
 be gained by a study of the heat evolved in their re- 
 actions. Thus, the following results have been obtained 
 in the study of chlorine, bromine, and iodine in their re- 
 lations to hydrogen : 
 
 [H 2 , C1J = 2[H, Cl] - [H, H] - [01, Cl] = 44,000 c. 
 [H 2 , BrJ = 2[H, Br] - [H, H] - [Br, Br] = 16,880 c. 
 [H 2 ,I 3 ] =2[H,I] - [H, H] - [I, I] = 12,072 c. 
 
 The meaning of these three equations will appear from 
 an interpretation of the first. This means that when a 
 molecule of hydrogen acts upon a molecule of chlorine 
 to form two molecules of hydrochloric acid gas 44,000 c. 
 of heat are evolved ; and this quantity is the difference 
 between that which is evolved in the combination of two 
 atoms of hydrogen with two atoms of chlorine, and that 
 which is absorbed in the decomposition of one molecule 
 of hydrogen into two atoms, and in the decomposition of 
 one molecule of chlorine into two atoms. The figures 
 thus obtained are not proportional to the affinities of 
 chlorine, bromine, and iodine for hydrogen, but never- 
 theless the affinities in all probability vary in the same 
 order. 
 
 The difficulties are much increased in more complicated 
 cases, and it will therefore be seen that it is impossible 
 to measure the affinity between the atoms by means of 
 the heat evolved in reactions. 
 
 Value of Thermochemical Measurements. Although 
 the affinities of the elements for one another cannot be 
 directly estimated by means of thermochemical measure- 
 
440 INORGANIC CHEMISTR Y. 
 
 ments, nevertheless these measurements are valuable, as 
 they show a direct relation between the quantity of heat 
 evolved and the character of the reaction which takes 
 place in any given case. In the case above cited, for 
 example, it is seen that the heat of formation of hydro- 
 chloric acid is greater than that of hydrobromic acid, 
 and that of hydrobromic acid is in turn greater than 
 that of hydriodic acid. Now, on page 94, it was stated 
 that in general that exothermic reaction takes place 
 which is accompanied by the greatest evolution of heat.* 
 Accordingly, in a case in which both hydrochloric and 
 hydrobromic acid could be produced the former would 
 certainly be produced in larger quantity. 
 
 Heat of Neutralization Avidity of Acids. Among the 
 measurements which have proved of value in connection 
 with the study of the general problem of affinity, are 
 those furnished by the heat of neutralization of acids 
 and bases. The general method of work consisted in 
 determining the heat evolved when equivalent quantities 
 of different acids are neutralized by the same base and 
 equivalent quantities of different bases are neutralized by 
 the same acid. Knowing the heat evolved in the reactions 
 between the various acids and bases, it .is possible to de- 
 termine what takes place when acids act upon salts in 
 which decomposition is not evident, either from the for- 
 mation of a precipitate or the evolution of a gas. Thus, 
 when two molecules of nitric acid act upon one of sodium 
 sulphate in solution, several changes are possible, as 
 represented in the equations 
 
 (1) Na 2 SO 4 + HNO 3 = NaHS0 4 + NaNO 3 ; 
 
 (2) Na 2 S0 4 + 2HNO 3 = H 2 SO 4 + 2NaNO 3 ; 
 
 (3) 2Na 2 SO 4 + 4HNO 3 = Na 2 SO 4 + 2NaNO 3 + H 2 SO 4 
 
 + 2HNO 3 . 
 
 As all the substances involved in these reactions are 
 soluble in water, and the reactions are studied in water 
 solution, it is clear that by ordinary methods it would 
 be impossible to tell which of them takes place. By 
 measuring the heat evolved, however, it has been shown 
 
 * This statement is known as Berthelot's Law of Maximum Work. 
 
AVIDITY OF ACIDS. 441 
 
 that in this and in all similar cases the base is divided 
 between the two acids, and generally more goes to one 
 acid than to the other. Further, it is possible to meas- 
 ure the division of the base between the acids, and in 
 this way measurements of the relative strengths of acids 
 are obtained. The figures representing the strengths of 
 the acids measured in this way are called the avidities of 
 the acids. In the case taken above as an illustration, it 
 was found that in dilute aqueous solution two-thirds of 
 the sodium goes to the nitric acid and one-third to the 
 sulphuric acid. Therefore, it appears that the avidity 
 of nitric acid is twice as great as that of sulphuric 
 acid. Of all acids investigated, nitric and hydrochloric 
 acids were found to have the greatest avidity. Calling 
 this 100, the avidities of some other acids determined by 
 this method are as given in the following table : 
 
 Acids. Avidity. 
 
 Nitric acid, 100 
 
 Hydrochloric acid, 100 
 
 Hydrobromic acid, ...... 89 
 
 Hydriodic acid, V 79 
 
 Sulphuric acid, , 49 
 
 Selenic acid, 45 
 
 Hydrofluoric acid, 5 
 
 Boric acid, . . . . , . . 1 
 
 Silicic acid, .......... 
 
 Hydrocyanic acid, 
 
 The figures given refer to equivalent quantities of the 
 acids, i.e., quantities which can be neutralized by equal 
 quantities of a base. Thus, 1 molecule of nitric acid, 
 HNO 3 , is neutralized by 1 molecule of sodium hydroxide, 
 NaOH ; but only molecule of sulphuric acid is neutral- 
 ized by 1 molecule of sodium hydroxide, and only % 
 molecule of orthophosphoric acid would be neutralized 
 by the same quantity of base. Therefore, we say that 
 1 molecule of nitric acid is equivalent to molecule of 
 sulphuric acid, and to -J- molecule of orthophosphoric 
 acid. 
 
442 INORGANIC CHEMISTRY. 
 
 It is impossible at present to give an exact interpretation 
 of the results above recorded, but it appears that the 
 figures given represent the numerical relations between 
 some common property possessed by acids, a property 
 which we have vaguely in mind when we speak of the 
 strength of acids. This appears more clearly when acids 
 and bases are studied in other ways. 
 
 Other Methods for Determining the Avidity of Acids. 
 Besides the thermochemical method of studying the 
 action of acids on bases, several other methods have been 
 devised. Among these are the volume-chemical method, 
 the optical method, the action of acids on insoluble salts, 
 and the electrical method. The object in view is in all 
 cases practically the same to compare the influence 
 exerted by different acids under the same circumstances, 
 and thus to measure their avidity or, as this has also 
 been called, their specific coefficient of affinity. 
 
 (1) The volume-chemical method depends upon the 
 fact that chemical processes which take place in homo- 
 geneous liquids generally cause changes in volume. 
 " Thus, the specific gravity of a normal caustic soda 
 solution was found to be 1.04051, that of an equivalent 
 solution of sulphuric acid 1.0297, that of an equivalent 
 of nitric acid 1.03089. When equal volumes of soda 
 solution were mixed with each of the acids, the specific 
 gravity of the sodium sulphate solution was 1.02959, and 
 that of the nitrate solution 1.02633. Finally, when to 
 the solution of sodium sulphate (2 vols.) one equivalent 
 (1 vol.) of nitric acid was added, the specific gravity be- 
 came 1.02781." By means of these figures it is possible 
 to determine to what extent the nitric acid acts upon the 
 sulphate, and thus to draw conclusions regarding the 
 distribution of the base between the acids. The results 
 reached by this method agree in general with those 
 reached by the thermochemical method. 
 
 (2) In the optical method the coefficient of refraction 
 of various solutions is determined, and also the changes 
 in the coefficient of refraction produced by mixing these 
 solutions in certain ways, and thus it is possible to draw 
 
DISSOCIATION. 443 
 
 conclusions in regard to the character of reactions which 
 take place in solutions. 
 
 (3) An illustration of the method involving the action 
 of acids on insoluble salts will make the method clear. 
 A weighed quantity of calcium oxalate is treated with 
 equivalent quantities of different acids in dilute solutions, 
 and the quantity of the salt dissolved in a given time 
 then determined. From the result it is possible to 
 calculate the specific coefficients of afiinity of the acids. 
 
 (4) The simplest method of all is the electrical. This 
 consists in determining the conducting power of different 
 substances at the same dilutions. In this way figures 
 are* obtained which bear to one another the same rela- 
 tions as those expressing the coefficients of affinity. 
 
 It is impossible to go into details in regard to these 
 methods here, and it need only be said that when acids 
 and bases are compared by the above methods, they are 
 found to differ markedly from one another, and the order 
 in which they are arranged by the results of the different 
 methods is always essentially the same. 
 
 Study of Chemical Decompositions. As we have seen, 
 practically every case of chemical combination with which 
 we have to deal is associated with the decomposition of 
 molecules, so that it is impossible perfectly to sepa- 
 rate the two acts of combination and decomposition. 
 Nevertheless there are some comparatively simple cases 
 of decomposition which have been studied with special 
 care, and results of much importance have been obtained. 
 The most interesting are those cases of decomposition 
 which are included under the heads of dissociation and 
 electrolysis. While many chemical decompositions are 
 brought about by concussion that is, by mechanical dis- 
 turbance of the mass the very instability of the com- 
 pounds which makes these decompositions possible pre- 
 vents any very profitable study of the phenomena. 
 
 Dissociation. Attention has been called to the fact that 
 many compounds, when heated to sufficiently high tem- 
 peratures, are decomposed. Thus, water is partly de- 
 composed into hydrogen and oxygen when heated to 1000; 
 
444 INORGANIC CHEMISTRY. 
 
 ammonium chloride is decomposed into ammonia and hy- 
 drochloric acid ; phosphorus pentachloride, into the tri- 
 chloride and chlorine ; nitrogen peroxide of the formula 
 N 2 O 4 , into the simpler compound of the formula NO 2 , etc. 
 Careful study of any one of these cases shows the follow- 
 ing facts : (1) That the decomposition takes place gradu- 
 ally ; (2) that the extent of the decomposition depends 
 upon the temperature and pressure, and for the same 
 compound is always the same for the same temperature 
 and pressure ; (3) that if the full amount of decomposition 
 possible at a certain temperature is effected, and the tem- 
 perature then lowered, the constituents will recombine to 
 some extent until equilibrium at the lower temperature is 
 established. 
 
 In a case of dissociation by heat, then, the decomposi- 
 tion is carried farther and farther as the temperature is 
 raised higher and higher, and it is finally complete. On 
 lowering the temperature again, more and more of the 
 compound is formed by the recombination of the constit- 
 ents until, when the lower temperature is again reached, 
 there is no decomposition. 
 
 The explanation of the phenomenon of dissociation is 
 found in the kinetic theory of gases. According to this 
 theory, the molecules of a gas at a given temperature are 
 moving with different velocities, though the average 
 velocity of all the molecules is the same at the same 
 temperature. Now, it is highly probable that the motion 
 of the atoms within the molecules partakes of that of 
 the molecules themselves, so that the motion of the 
 atoms in the molecules with the greatest velocity is 
 probably the greatest, and, in these, decomposition 
 will take place first. When a compound gas is heated, 
 we can easily conceive that even at a comparatively low 
 temperature the motion of some of the molecules will 
 be sufficient to cause their decomposition, and, as the 
 average motion of all the molecules is constant for a 
 given temperature, the amount of decomposition will 
 be constant for that temperature. As the molecules are, 
 however, moving in every direction and constantly col- 
 liding, a molecule which is decomposed at one instant 
 
ELECTROLYSIS. 445 
 
 may be re-formed at the next, and one that is not decom- 
 posed may acquire motion enough to cause its decompo- 
 sition. Though, as is believed, these changes are con- 
 stantly taking place at every temperature, still, as has 
 been said, the number of molecules which will be decom- 
 posed in a given mass at a given temperature and pres- 
 sure will always be the same. The higher the tem- 
 perature, then, the greater the number of molecules in 
 the conditions which cause decomposition, and the 
 smaller the number of those in the conditions favorable 
 to formation. At each temperature and pressure an 
 equilibrium is established, the number of molecules de- 
 composed being equal to the number formed. It is 
 obvious that, if one of the products of decomposition is 
 removed, the conditions are entirely changed. Then the 
 possibility of recombination will not exist, and total 
 decomposition can be effected at a lower temperature 
 than that required for total decomposition in the process 
 of dissociation proper. 
 
 Electrolysis. Some chemical compounds in solution 
 in water conduct electricity, and at the same time they 
 undergo decomposition. Thus, hydrochloric acid in solu- 
 tion in water conducts electricity and the compound is 
 decomposed into its constituents hydrogen and chlorine, 
 the hydrogen appearing at the negative and the chlorine 
 at the positive pole. Compounds that act in this way 
 are called electrolytes. When a current of electricity 
 acts upon solutions of different salts, equivalent quanti- 
 ties of the metals are deposited by the same current in 
 the same time. This is Faraday's Law. Thus if the same 
 current were passed simultaneously through solutions 
 of silver nitrate, AgNO 3 , mercuric nitrate, Hg(NO 3 ) a , 
 cupric sulphate, CuSO 4 , and ferric chloride FeCl s , it 
 would be found that for every 107.11 parts by weight 
 of silver deposited there would be 99.2 parts by weight 
 of mercury deposited, 31.56 of copper, and 18.53 of 
 iron. These are equivalent quantities of these metals 
 quantities that take the place of one part by 
 weight of hydrogen and are to be distinguished from 
 atomic quantities. Those elements which appear at 
 
446 INORGANIC CHEMISTRY. 
 
 the negative pole are called electro-positive, and those which 
 appear at the positive pole are called electro-negative. 
 Those elements which we call acid-forming are electro- 
 negative, while hydrogen and the base-forming elements 
 are electro-positive. The electrolysis of chemical com- 
 pounds is not generally a simple decomposition into two 
 constituents. Thus, when copper sulphate, CuSO 4 , is 
 decomposed, the copper is deposited at the negative 
 pole ; but no such compound as SO 4 appears at the posi- 
 tive pole. This, if formed, perhaps breaks down into 
 oxygen and sulphur trioxide, and the latter with water 
 would then naturally form sulphuric acid. Both oxygen 
 and sulphuric acid as a matter of fact appear at the posi- 
 tive pole. The changes involved may be represented thus,: 
 
 CuSO 4 =Cu+S0 4 ; S0 4 =S0 3 +0; SO 3 +H a O=H 2 SO 4 . 
 
 Electrolytic Dissociation. It has been known for a 
 long time that a very weak electric current acting upon a 
 solution of an electrolyte is sufficient to cause the ions 
 to appear at the poles. This fact is inexplicable if 
 it is assumed that the current is the cause of the de- 
 composition of the electrolyte. This and some other 
 facts which will be referred to farther on make it prob- 
 able that electrolytes are at least to some extent decom- 
 posed into their constituent ions when they are dissolved ; 
 that these ions charged with electricity transfer their 
 charges in the solution and thus conduct the current ; and 
 that when an ion charged with negative electricity reaches 
 the positive pole its electricity is discharged, and the 
 ion then ceases to be an ion and becomes an element in 
 the free state or some compound which appears either as 
 such or in the form of other products. According to 
 this conception, the act of solution of an electrolyte, in 
 water at least, involves partial breaking down or dis- 
 sociation of the compound into its ions. The extent of 
 this breaking down is determined primarily by the con- 
 centration of the solution the greater the dilution the 
 greater the dissociation. At infinite dilution there is 
 complete dissociation. A water solution of hydrochloric 
 
ELECTROLYTIC DISSOCIATION. 447 
 
 acid containing 36.18 grams of the acid in 1000 liters 
 has been shown to be completely dissociated, or it is to 
 be regarded as containing ions of hydrogen and of 
 chlorine. These and all other ions are carefully to be dis- 
 tinguished from the atoms or definite compounds. An 
 ion always carries with it a certain charge of electricity. 
 When this is discharged the ion becomes either an ele- 
 ment or a compound in the free state. When a solution 
 of one electrolyte acts upon a solution of another the 
 reaction observed is probably due to the interaction of 
 the ions, and it is further probable that, so far as the 
 compounds are present in the undissociated condition, 
 they do not act upon each other. If this view is correct 
 the reactions most familiar to us are reactions of ions, 
 and not of elements or compounds. When, for ex- 
 ample, an acid acts upon a base in solution it appears 
 that, so far as they react, they are in dissociated condi- 
 tion. Thus hydrochloric acid and sodium hydroxide are 
 to be regarded as acting as represented in the following 
 equation : 
 
 H + 01 + Na + OH = Na + Cl + H 2 O. 
 
 The act consists in the union of the hydroxyl ion of 
 the base with the hydrogen ion of the acid to form water, 
 the sodium and chlorine ions remaining as ions as in 
 a dilute solution of sodium chloride. In the case of 
 nitric acid and potassium hydroxide the following equa- 
 tion represents the reaction at infinite dilution : 
 
 H + N0 3 + K + OH = K + NO, + H,O. 
 
 And so also whenever the act of neutralization takes 
 place there is simply a union of hydrogen ions with 
 hydroxyl ions to form water. 
 
 This conception finds strong confirmation in the fact 
 that the heat evolved in neutralizing equivalent quantities 
 of all acids at infinite dilution is always the same a fact 
 that it is difficult to explain if it is assumed that in the 
 act of neutralization a salt is formed in the solution. 
 
 In the following table the heats of neutralization of 
 a few acids and bases are given : 
 
448 INORGANIC 1 CHEMISTRY. 
 
 Acid and Base. Heat of 
 
 Neutral. 
 
 Hydrochloric acid and sodium hydrox., . . 13,700 
 
 " " lithium " . . 13,700 
 
 " " " potassium " . . 13,700 
 
 " " " barium " . . 13,800 
 
 " " calcium " . . 13,900 
 
 Hydrobromic " " sodium " ' . . 13,700 
 
 Nitric " " " " . . 13,700 
 
 lodic " " " " . . 13,700 
 
 As will be seen, the heat of neutralization is the same 
 no matter what the base or what the acid may be, and 
 as has been pointed out this fact is easily understood, if 
 the act of neutralization consists in the union of a hy- 
 droxyl ion with a hydrogen ion to form water. 
 
 If the strength of an acid is determined by the extent 
 to which it is dissociated, then of course those acids that 
 are most readily dissociated are the strongest. By every 
 method available hydrochloric and nitric acids are found 
 to be the most readily dissociated and they are the strong- 
 est acids. The electrical method for the determination 
 of the strength of acids is based upon this theory. The 
 applicability of this theory to the explanation of the most 
 common reactions that take place in solution is at present 
 attracting the attention of chemists, and promises to be 
 of great service to chemistry. Those reactions which are 
 made use of for the purpose of detecting the presence of 
 the various elements appear in fact, as has already been 
 stated, to be reactions of the ions, and when these ions 
 are not present the reactions are not observed. For ex- 
 ample, when silver nitrate in solution is added to sodium 
 chloride in solution a precipitate of silver chloride is 
 formed, the reaction taking place as represented in this 
 equation : 
 
 A g NO 3 + ^aCl = AgCl + NaNO,. 
 
 This reaction seems, however, to be due to the fact that 
 ions of silver and of chlorine are present. These com- 
 ing together form the insoluble molecules silver chloride 
 which is then precipitated. There are many compounds 
 that contain chlorine and yet do not give a precipitate 
 
SPECIFIC HEAT AND ATOMIC WEIGHTS. 449 
 
 of silver chloride when treated with silver nitrate. It is 
 believed that in these cases the compound is not ionised^ 
 by the solvent in such a way as to yield ions of chlorine, 
 and that therefore there are no chlorine ions present. 
 An example of a compound that contains chlorine and 
 yet does not ordinarily give the reactions of this element 
 is potassium chlorate, KC1O 3 . A solution of this salt 
 does not give a precipitate with silver nitrate. It is pro- 
 bable that the reason of this is that the salt gives the 
 ions K and C1O 3 , the latter acting quite differently from 
 the ion 01, as we should naturally expect. 
 
 Relations between Specific Heat and Atomic Weights. 
 The fact that there is a method for the determination of 
 atomic weights founded upon the relations existing be- 
 tween these weights, and the specific heat of the ele- 
 ments, has been mentioned. It has been found that, when 
 equal weights of different elements are exposed to exactly 
 the same source of heat, they require different lengths of 
 time to become heated to the same temperature. Given 
 exactly the same heating power, it requires 32 times as 
 long to raise the temperature of a pound of water 10, 20, 
 or 30 degrees as it does to raise the temperature of a 
 pound of mercury the same number of degrees ; or it takes 
 32 times as much heat to raise a pound of water 10, 20, 
 or 30 degrees as it does to raise a pound of mercury 
 the same number of degrees. Starting at the same tem- 
 perature the quantity of heat required to raise the tem- 
 perature of a certain weight of a substance one degree, 
 as compared with the quantity of heat required to raise 
 the temperature of the same weight of water one degree, 
 is called the specific heat of the substance. Thus, from 
 what was said above, the specific heat of mercury is -%, 
 or, in decimals, 0.03125. In a similar way it can be shown 
 that the specific heat of gold is 0.03244 ; of zinc, 0.0955 ; 
 of silver, 0.057 ; of copper, 0.0952. 
 
 Now, when solid elements are examined with reference 
 to their specific heats, a very simple relation is found to 
 exist between the numbers expressing the specific heats 
 and the atomic weights. This relation will be made 
 clear by a consideration of a few cases : 
 
450 INORGANIC CHEMISTRY. 
 
 Element. Specific Heat. Atomic Weight. 
 
 Silver, 0.0570 107.11 
 
 Zinc,. :.>-' ..;- ; Y.-y.y 0.0955 64.91 
 
 Cadimum, .... 0.0567 111.10 
 
 Copper, 0.0952 63.12 
 
 Tin, 0.0562 118.15 
 
 An examination of this table will show that the atomic 
 weights are inversely proportional to the specific heats. 
 We have 
 
 107.11 : 64.91 
 111.10: 63.12 
 107.11 : 118.15 
 
 0.0955 : 0.0570 ; 
 0.0952 : 0.0567 ; 
 0.0562 : 0.0570; etc. 
 
 These proportions are only approximately correct ; but 
 it must be remembered that the means for the determi- 
 nation of atomic weights and specific heats are not per- 
 fect, and in both sets of figures there are undoubtedly 
 small errors. Hence such slight variations from abso- 
 lute agreement in these proportions should occasion no 
 surprise. The agreement is sufficiently close to indicate 
 a close connection between the two sets of figures. This 
 connection may be stated in another way : The product 
 of the atomic weight by the specific heat is a constant. Thus, 
 in the above cases : 
 
 107.11 X 0.057 =6.11; 
 
 64.91 X 0.0955 = 6.20 ; 
 111.10 X 0.0567 = 6.30; 
 
 63.12 X 0.0952 = 6.01 ; 
 118.15 X 0.0562 = 6.64. 
 
 From the above it appears that the quantity of heat 
 necessary to raise masses of the elements proportional 
 to their atomic weights the same number of degrees is 
 the same in all cases. Suppose two elements to have 
 the atomic weights 2 and 4. Their specific heats would 
 be to each other as 2 to 1. That is to say, it would 
 require twice as much heat to raise the temperature of a 
 given mass of the element with the atomic weight 2 a 
 
SPECIFIC HEAT AND ATOMIC WEIGHTS. 451 
 
 certain number of degrees, as it would require to raise 
 the temperature of the same mass of the element with 
 the atomic weight 4 the same number of degrees. But 
 to raise the temperature of masses of these two elements 
 proportional to their atomic weights would require the 
 same quantity of heat. This fact may be stated thus : 
 
 The atoms of all elements have the same capacity for heat. 
 This is only another way of stating that, to raise the 
 temperature of an atom one degree, the same quantity of 
 heat is always necessary. 
 
 Now, if we assume that the constant obtained by 
 multiplying the specific heats by the atomic weights is 
 6.4, which is about the average of the different values 
 found, then it is plain that, if we divide this number by 
 the specific heat of an element, we shall obtain a number 
 which is very near the atomic weight. If we call the 
 atomic weight A, and the specific heat H, the following 
 equation expresses the relation : 
 
 If this law is without exceptions, it is plain that, in order 
 to determine the atomic weight of an element, it is only 
 necessary to determine its specific heat, and divide this 
 into 6.4. The result will be very nearly the atomic 
 weight. Knowing thus very nearly what the atomic 
 weight is, it is a comparatively simple matter to deter- 
 mine it with great accuracy by means of chemical analy- 
 sis. Unfortunately there are some marked exceptions 
 to the law. 
 
 Exceptions to the Law of Specific Heats. The elements 
 glucinum, carbon, boron, and silicon form exceptions to 
 the law of specific heats as this law has been stated above. 
 At ordinary temperatures they do not follow the law. As 
 the temperature is raised, however, the specific heat of 
 these elements changes markedly, until finally, in the cases 
 of carbon and silicon, a point is reached beyond which 
 there is no marked change. Thus, at 600 the specific 
 heat of diamond is 0.441, and at 985 it is 0.449. That 
 of silicon is 0.201 at 185, and 0.203 at 332. At these 
 
452 INORGANIC CHEMISTRY. 
 
 temperatures the elements obey the law. From elabo- 
 rate studies which have been made on this subject, it ap- 
 pears that the law should be modified to read as follows : 
 
 The specific heats of the elements vary with the tem- 
 perature ; but for every element there is a temperature, T t 
 above which variations are very slight. The product of 
 the atomic weight by the constant value of the specific 
 heat is nearly a constant, lying between 5.5 and 6.5. 
 
 Notwithstanding the irregularities referred to, the law 
 of specific heats, commonly called, from the discoverers, 
 the law of Dulong and Petit, is of great value in the de- 
 termination of atomic weights. 
 
 Raoult's Methods for the Determination of Molecular 
 Weights. One great difficulty encountered in the study 
 of chemical compounds is the determination of the mole- 
 cular weights of those which are not gases or cannot be 
 converted into vapor by heat. From some studies on 
 the freezing-points of solutions, it appears that quantities 
 of compounds proportional to their molecular weights 
 cause the same lowering of the freezing-points, provided 
 the solvent does not act chemically upon the compound. 
 This fact makes it possible to determine the molecular 
 weights of substances which cannot be converted into 
 vapor, but which can be dissolved. The application of 
 the method is simple. Suppose water to be the solvent 
 used. We know that this liquid solidifies or freezes at 
 0. Now, it is found that by dissolving a certain quan- 
 tity of some substance in a certain quantity of water the 
 freezing-point is lowered say .5. Further, the quantities 
 of other substances which are necessary to lower the 
 freezing-point of the same quantity of water to the same 
 extent can be determined. These quantities are propor- 
 tional to the molecular weights according to the law of 
 Eaoult. If, therefore, among the substances studied 
 there is one the molecular weight of which can be deter- 
 mined by the method of Avogadro, it is possible to de- 
 termine the molecular weights of all of them by the 
 method of Baoult, as will readily be seen. 
 
 So also it has been shown that quantities of com- 
 pounds proportional to their molecular weights cause 
 
DISSOCIATION OF A DISSOLVED SUBSTANCE. 453 
 
 the same raising of the boiling-points, provided the 
 solvent does not act chemically upon the compounds or 
 cause them to break down into their ions. Convenient 
 methods have been devised for the determination of 
 molecular weights of dissolved substances, the methods 
 being based upon observations on the boiling-points 
 and freezing-points. It should be noted that, when the 
 molecular weight of a substance in solution has been 
 determined, it does not follow that the substance has 
 the same molecular weight when in the solid condi- 
 tion. This is a matter in regard to which we have prac- 
 tically no knowledge. It is quite possible that the mole- 
 cules of solid substances may be made up of large 
 aggregates of the simple molecules, such as probably 
 exist in solutions or in vapors. There is, however, no 
 method at present known that makes a determination 
 of the complexity of these molecules, or molecular aggre- 
 gates, possible. 
 
 Determination of the Extent of Dissociation of a Dis- 
 solved Substance. The effect upon the boiling-point or 
 freezing-point of a solution caused by the presence of a 
 dissolved substance is proportional to the number of 
 molecules in the solution or the number of individual 
 particles, whether these are undecomposed molecules or 
 the ions formed as a result of dissociation. Any sub- 
 stance that is dissociated in solution will give abnormal 
 results if the attempt is made to determine its molecu- 
 lar weight by observations on the boiling-point or the 
 freezing-point of its solutions. This method is there- 
 fore not applicable to solutions of electrolytes. On the 
 other hand, the study of such solutions has shown that 
 there is an increased lowering of the freezing-point for 
 the same weight of solvent as the dilution becomes 
 greater, a fact that points clearly to the conclusion that 
 as the solution is diluted there is greater and greater 
 dissociation, and advantage can be taken of this fact 
 for the purpose of determining the extent to which dis- 
 sociation has taken place in a solution of an electrolyte 
 in water. 
 
454 INORGANIC CHEMISTRY. 
 
 In general only organic compounds come within range 
 of the methods of Raoult. These methods are now ex- 
 tensively used in the study of the compounds of carbon, 
 simple forms of apparatus having been devised for this 
 purpose. The recognition of the fact that electrolytes 
 do not obey the simple law that holds good in the case 
 of non-electrolytes led Arrhenius to the idea that the 
 former are dissociated in solution an idea which has 
 proved of great service to the science, and is likely to 
 revolutionize the views of chemists in regard to the 
 action of chemical substances upon each other in solu- 
 tion. 
 
CHAPTER XXIV. 
 
 BASE-FORMING ELEMENTS GENERAL CONSIDERATIONS 
 
 Introductory. The elements thus far considered be- 
 long for the most part to the class of acid-forming 
 elements, or those whose compounds with oxygen and 
 hydrogen have acid properties. All the members of 
 Family VII, Group B, are acid-forming, while the single 
 member of Group A of the same family is both acid- 
 forming and base-forming. All the members of Family 
 VI, Group B, are acid-forming, while the members of 
 Group A of this family are .both acid-forming and base- 
 forming. In Family V, Group B, there is observed a 
 gradation of properties, the group beginning with strong- 
 ly marked acid-forming elements and ending with an ele- 
 ment, bismuth, which is more basic than acid in char- 
 acter. The elements of Group A, Family V, are both 
 acid-forming and base-forming, but they have not as 
 sharply marked characteristics as the elements of Fam- 
 ilies VI and VII. Passing now to Family IV, we found 
 that the two most important members, carbon and sili- 
 con, belong to Group A. These two elements always act 
 as acid-formers. A gradation of properties is observed 
 in passing from silicon to thorium. The members of 
 Group B of this family have the properties of the base- 
 forming elements much more strongly marked than those 
 of the acid-formers. There are still four families to be 
 studied. These are families I, II, III, and VIII, the 
 members of which are almost exclusively base-forming 
 elements. The compounds of these elements with hy- 
 drogen and oxygen are bases, or, in other words, have the 
 power to neutralize acids. Their oxides are for the most 
 part basic. An exception to this is found in the case of 
 boron, already considered, which forms a weak acid 
 boric acid. Its oxide is only slightly basic. The most 
 
 (455) 
 
456 INORGANIC CHEMISTRY. 
 
 strongly marked examples of base-forming elements are 
 those which occur in Family I, Group A ; then follow in 
 order those of Group A, Family II, and Group A, Fam- 
 ily III. The resemblance between the members of 
 Group B, Family I v and those of Group A of the same 
 family is less striking than the resemblance between the 
 two groups of any other family. Between the members 
 of Group B, Family II, and those of Group A of the same 
 family there is a general resemblance, while there are 
 also differences. A similar remark applies to the rela- 
 tions between Groups A and B, Family III. The mem- 
 bers of Family VIII occupy a somewhat exceptional 
 position, as has already been pointed out. Each group 
 of which this family consists is made up of three very 
 similar elements with atomic weights which differ but 
 little from one another. 
 
 Metallic Properties. It has long been customary to 
 divide the elements into two classes the metals and the 
 non-metals. This classification was originally based upon 
 differences in the physical properties of the elements, the 
 name metal being applied to those elements which have 
 what is known as a metallic lustre, are opaque, and are 
 good conductors of heat and electricity. All those ele- 
 ments which do not have these properties, are called non- 
 metals. Gradually the name metal came to signify an 
 element which has the power to replace the hydrogen of 
 acids and form salts, and the name non-metal to signify 
 an element which has not this power. This classifica- 
 tion, as will be seen, is practically the same as that which 
 divides the elements into acid-forming and base-forming. 
 The latter are the metals, the former are the non-metals. 
 The imperfection of this classification has already been 
 commented upon, the imperfection arising from the fact 
 that some elements belong to both classes. 
 
 Order in which the Base-forming Elements will be Taken 
 up. In studying the base-forming elements, it appears 
 best to begin with those which have the most strongly 
 marked character. These are the members of Family I, 
 Group A. It further appears best to adhere as closely 
 as possible to the arrangement in the periodic system. 
 
OCCURRENCE OF THE METALS. 457 
 
 Accordingly, the following order will be observed in the 
 presentation of the elements yet to be studied : 
 
 1. Elements of Family I, Group A, or the Potassium 
 Group, consisting of lithium, sodium, potassium, rubid- 
 ium, and caesium. 
 
 2. Elements of Family II, Group A, or the Calcium 
 Group, consisting of glucinum, magnesium, calcium, 
 strontium, barium, and erbium. 
 
 3. Elements of Family III, Group A, or the Aluminium 
 Group, consisting of aluminium, scandium, yttrium, lan- 
 thanum, and ytterbium. 
 
 4. Elements of Family I, Group B, or the Copper 
 Group, consisting of copper, silver, and gold. 
 
 5. Elements of Family II, Group B, or the Zinc Group, 
 consisting of zinc, cadmium, and mercury. 
 
 6. Elements of Family III, Group B, or the Gallium 
 Group, consisting of gallium, indium, and thallium. 
 
 7. Elements of Family IY, Group B, or the Tin Group, 
 consisting of germanium, tin, and lead. 
 
 8. Elements of Family Y, Group A, or the Vanadium 
 Group, consisting of vanadium, columbium, didymium, 
 and tantalum. 
 
 9. Elements of Family VI, Group A, or the Chromium 
 Group, consisting of chromium, molybdenum, tungsten, 
 and uranium. 
 
 10. Elements of Family VII, Group A, or the Manganese 
 Group, of which manganese is the only representative. 
 
 11. Elements of Family VIII, of which there are 
 three groups : 
 
 (A) The Iron Group, consisting of iron, nickel, and 
 <?obalt ; 
 
 (B) The Palladium Group, consisting of ruthenium, 
 rhodium, and palladium ; and 
 
 (C) The Platinum Group, consisting of osmium, irid- 
 ium, and platinum. 
 
 Occurrence of the Metals. One of the first questions 
 that suggests itself in connection with each element is, 
 In what forms of combination does it occur in nature ? 
 The chemical compounds which occur ready-formed in 
 nature are called minerals; and the minerals, and mixtures 
 
458 INORGANIC CHEMISTRY. 
 
 of minerals, from which the metals are extracted for prac- 
 tical purposes are called ores. The most common ores 
 are oxides and sulphides. Examples of these are the 
 ores of iron, tin, copper, lead, and zinc. The carbonates 
 also occur in large quantity in nature, and are used for 
 the purpose of preparing some of the metals. The car- 
 bonate of zinc, for example, is a valuable ore of zinc. 
 
 Extraction of the Metals from their Ores. The detailed 
 study of the methods used in the extraction of the metals 
 from their ores is the object of metallurgy. Besides the 
 methods used on the large scale, there are others which are 
 only used in the laboratory. The most common method 
 of extracting metals from their ores is that used in 'the 
 case of iron, which consists in heating the oxides with char- 
 coal. If the ores used are not oxides, they must first be 
 converted into oxides before this method is applicable. 
 This can generally be accomplished by heating the ores 
 in contact with the air. Under these circumstances the 
 natural carbonates, sulphides, and hydroxides, are con- 
 verted into oxides. These changes are illustrated by the 
 following equations : 
 
 FeC0 3 = FeO + CO 2 ; 
 2FeO + O = Fe 2 O 3 ; 
 2FeS 2 -f HO = Fe 2 O 3 + 4SO 2 ; 
 2Fe(OH) 3 = Fe 2 O 3 + 3H 2 O. 
 
 A second method consists in reducing the oxide by heat- 
 ing it in a current of hydrogen. This has been illustrated 
 in the action of hydrogen upon copper oxide, when the 
 following reaction takes place : 
 
 CuO + H 2 = H 2 + Cu. 
 
 The method is efficient for many oxides, but is expen- 
 sive and is not used on the large scale. 
 
 Another method of extraction consists in treating the 
 chloride of a metal with sodium. This is illustrated in 
 the preparation of magnesium, which is made by heating 
 together magnesium chloride and sodium : 
 
 MgCl 2 + 2Na = 2NaCl + Mg. 
 
COMPOUNDS OF THE METALS. 459 
 
 Such a method is employed only in case it is impossible 
 or extremely difficult to reduce the oxide. 
 
 Besides the above methods, there are others which will 
 be described under the individual metals. 
 
 The Properties of the Metals. As we shall find, the 
 metals differ very markedly from one another. Some 
 are light, floating on water, as lithium, sodium, etc. ; 
 some are extremely heavy, as lead, platinum, etc. Some 
 combine with oxygen with great energy ; others form very 
 unstable compounds with oxygen. Some form strong 
 bases ; others form weak bases. In general, those ele- 
 ments which are lightest, or which have the lowest 
 specific gravity, are the most active chemically, while 
 those which have the highest specific gravity are the 
 least active. Among the former are lithium, sodium, and 
 potassium ; among the latter are lead, gold, and platinum. 
 
 Compounds of the Metals. The principal compounds 
 of the metals may be conveniently classified as : 
 
 a. Compounds with fluorine, chlorine, bromine, and 
 iodine ; or the fluorides, chlorides, bromides, and iodides. 
 
 b. Compounds with oxygen, and with oxygen and hy- 
 drogen ; or the oxides and hydroxides. 
 
 c. Compounds with sulphur, and with sulphur and hy- 
 drogen ; or the sulphides and hydrosulphides. 
 
 d. Compounds with nitrogen ; or the nitrides. 
 
 e. Compounds with carbon and with silicon ; or the 
 carbides and silicides. 
 
 f. Compounds with the acids of nitrogen ; or the ni- 
 trates and nitrites. 
 
 g. Compounds with the acids of chlorine, bromine, and 
 iodine ; or the chlorates, bromates, iodates, hypochlorites, etc. 
 
 h. Compounds with the acids of sulphur, selenium, 
 and tellurium ; or the sulphates, sulphites, etc. 
 
 i. Compounds with carbonic acid ; or the carbonates. 
 
 j. Compounds with the acids of phosphorus, arsenic, 
 and antimony ; or the phosphates, arsenates, etc. 
 
 k. Compounds with silicic acid ; or the silicates. 
 
 1. Compounds with boric acid ; or the borates. 
 
 It is more important to become acquainted with the 
 general methods of preparation and the general properties 
 
460 INORGANIC CHEMISTRY. 
 
 of the more important compounds than to learn details 
 concerning many individual members of each class. Only 
 those compounds which illustrate general principles, 
 or which, owing to some application, happen to be of 
 special interest, need be fully treated in this book. 
 
 The acids of which the salts are derivatives are already 
 known to us, and in dealing with the acids frequent 
 reference has been made to the methods of making the 
 salts, and to some of their more important properties. 
 It will be well, before taking up the metals systemati- 
 cally, briefly to treat of the general methods of prepara- 
 tion, and the general properties of the different classes 
 of metallic compounds. It must be borne in mind, how- 
 ever, that the only way to become familiar with these 
 substances and their relations is by working with them in 
 the laboratory. 
 
 Chlorides. The chlorides, as well as the fluorides, 
 bromides, and iodides, may be regarded as the salts of 
 hydrochloric, hydrofluoric, hydrobromic, and hydriodic 
 acids, or simply as compounds of the metals with the 
 members of the chlorine family. The most important 
 of these compounds are the chlorides, and these well 
 illustrate the conduct of the others. 
 
 The chlorides are made by treating a metal with chlo- 
 rine, or with hydrochloric acid ; by treating an oxide 
 or a hydroxide with hydrochloric acid ; by treating an 
 oxide with chlorine and a reducing agent, like carbon ; 
 by treating a salt of a volatile acid with hydrochloric 
 acid ; by treating a salt of an insoluble acid with hydro- 
 chloric acid ; by adding hydrochloric acid or a soluble 
 chloride to a solution containing a metal with which 
 chlorine forms an insoluble compound ; and by adding to 
 a solution of a chloride a salt, the acid of which forms 
 with the metal of the chloride an insoluble salt, while 
 the metal contained in it forms with chlorine a soluble 
 chloride. 
 
 Only two of the above methods are peculiar to chlo- 
 rides. These are the treatment of the metals with chlo- 
 rine, and the treatment of oxides with chlorine and a 
 reducing agent. The others involve principles which 
 
CHLORIDES. 461 
 
 are also involved in the preparation of all salts, and they 
 may therefore be treated of in a general way. 
 
 The formation of chlorides by direct treatment of the 
 metals with chlorine is the simplest method of all. It 
 has been illustrated in studying chlorine. It was found 
 that chlorine combines with other elements with great 
 ease. Thus, iron, copper, and tin combine with it, as 
 represented in the following equations : 
 
 Cu +01, = CuCl 2 ; 
 2Fe + 3C1 2 = 2FeCl 3 ; 
 Sn 
 
 The preparation of chlorides by treating oxides with 
 chlorine and a reducing agent has been illustrated in 
 the making of boron trichloride and of silicon tetra- 
 chloride. It is used in making aluminium trichloride. 
 For this purpose, chlorine is passed over a heated mix- 
 ture of aluminium oxide and charcoal, when reaction 
 takes place according to the following equation : 
 
 A1 2 3 + 3C + 3C1 2 = 2A1C1 3 + 3CO. 
 
 The interesting character of this reaction was referred to 
 in connection with the similar preparation of the chlo- 
 rides of boron and silicon. In this case, as in those, 
 there are two reactions involved. The carbon alone can- 
 not reduce the oxide ; nor can the chlorine alone decom- 
 pose it to form the chloride. But when the carbon and 
 chlorine act together, they assist each other, and as a 
 consequence the oxide is transformed into the chloride. 
 
 The other methods for preparing chlorides are, as has 
 been said, general in character and are applicable to 
 most salts. 
 
 Formation of Salts in General. 
 
 1. By treating a metal with an acid. This is the sim- 
 plest method. It has been illustrated in the preparation 
 of zinc sulphate by the action of zinc on sulphuric acid : 
 
 Zn + H 2 SO 4 = ZnS0 4 + H 2 . 
 
462 INORGANIC CHEMISTRY. 
 
 Other common examples are those represented in the 
 follow equations : 
 
 Fe + H 2 S0 4 = FeS0 4 + H a ; 
 Zn+2HCl = ZnCl a + H 3 . 
 
 2. By treating an oxide or a hydroxide with an acid. 
 This is of more general application than the preceding 
 method. As it has been studied in some detail in con- 
 nection with the subject of salts (see pp. 129-133), it need 
 not be further considered here. 
 
 3. By treating the salt of a volatile acid with another acid. 
 This method has been repeatedly illustrated in the de- 
 composition of carbonates and nitrites by acids in gen- 
 eral. While carbonic acid and nitrous acid themselves- 
 are perhaps not formed in these reactions, and we can- 
 not say that the carbonates and nitrites are salts of vol- 
 atile acids, yet the decomposition-products of these acids 
 are volatile at ordinary temperatures. The decompo- 
 sition of carbonates by acids has been pretty fully 
 studied, though attention was not directed to the fact 
 that this kind of action may be utilized for the purpose 
 of making salts. As some carbonates occur in large 
 quantity in nature or in the market, salts are frequently 
 made by treating them with acids. Thus, magnesium 
 sulphate is made by treating magnesium carbonate with 
 sulphuric acid : 
 
 MgC0 3 + H,SO, = MgSO, + H,0 + CO, ; 
 
 and calcium chloride is made by dissolving calcium car- 
 bonate in hydrochloric acid : 
 
 CaCO, + 2HC1 = CaCl 2 + H 2 O + CO 2 ; etc. 
 
 4. By treating a salt of an insoluble acid with another 
 acid. This case does not occur practically, as there are 
 no common, insoluble acids. The principle involved is 
 illustrated to some extent by the decomposition of a sol- 
 uble silicate. Sodium silicate, for example, is soluble. 
 When its solution in water is treated with an acid 
 
GENERAL PROPERTIES OF THE CHLORIDES. 463 
 the silicic acid is partly precipitated, as we have seen 
 
 Na 2 SiO 3 + 2HC1 + H a O = Si(OH) 4 + 2NaCl. 
 
 The silicic acid formed is, however, not perfectly in- 
 soluble in water, so that the reaction is not complete. 
 In any case the reaction is not one that is used for the 
 preparation of salts. 
 
 5. By the action of two salts upon each other. This 
 method can be best described by means of an example. 
 Suppose it is desired to prepare copper chloride by the 
 action of two salts upon each other. Copper chloride 
 is soluble. If copper sulphate and barium chloride are 
 brought together in solution, the products are insoluble 
 barium sulphate and soluble copper chloride : 
 
 CuSO 4 + BaCl 2 = BaS0 4 + CuCl 2 . 
 
 By simply filtering off the barium sulphate, a solution 
 of copper chloride is obtained. 
 
 6. By precipitation. This method is illustrated in the 
 formation of barium sulphate, referred to in the last par- 
 agraph. Obviously, it is applicable only to difficultly 
 soluble or insoluble salts. Many carbonates and phos- 
 phates can be made in this way. 
 
 General Properties of the Chlorides. Most of the chlo- 
 rides of the metals are soluble in water without decom- 
 position, though many of them are decomposed when 
 heated to a sufficiently high temperature with water. It 
 will be remembered that the chlorides of the non-metal- 
 lic or acid-forming elements are decomposed by water, 
 yielding the corresponding oxides or hydroxides. The 
 chlorides of some elements which are partly basic and 
 partly acid are only partly decomposed. This is illus- 
 trated by the chloride of antimony, which with water 
 forms an oxy chloride : 
 
 SbCl 3 + H 2 = SbOCl + 2HC1. 
 
464 INORGANIC CHEMISTRY. 
 
 The chlorides of the most strongly marked metals, 
 like potassium, sodium, etc., are not decomposed by 
 water. Calcium chloride dissolves with great ease, and, 
 if the solution is evaporated, the chloride is again 
 obtained. If, however, the attempt is made to drive off 
 all the water by heat, some of the chloride is converted 
 into the oxide as represented in the equation 
 
 CaCl, + H 2 = CaO + 2HC1. 
 
 Magnesium chloride is completely decomposed, if its 
 solution in water is evaporated to dryness, the action 
 being the same in character as that which takes place in 
 the case of calcium chloride. The chlorides of iron and 
 aluminium and of many other metals act in the same way. 
 Silver chloride and mercurous chloride, HgCl, are insol- 
 uble in water. Lead chloride is difficultly soluble in 
 water. If, therefore, on adding hydrochloric acid or a 
 soluble chloride to a solution, a precipitate is formed, 
 the conclusion is generally justified that one or more "of 
 the three metals silver, lead, or mercury is present. 
 By taking into account the differences between these 
 chlorides, it is not difficult to decide of which of them a 
 precipitate consists. 
 
 The chlorides are for the most part stable when heated, 
 though a few lose some of their chlorine just as phos- 
 phorus pentachloride does. An example of this is pre- 
 sented by platinic chloride, PtCl 4 , which when heated 
 breaks down into platinous chloride, PtCl 2 , and chlorine : 
 
 PtCl 4 = PtCl, + C1 2 . 
 
 The chlorides are for the most part decomposed when 
 treated with sulphuric acid, as has been shown in the 
 action of sulphuric acid upon sodium chloride. Under 
 these circumstances hydrochloric acid is given off, and 
 the sulphate of the metal with which the chlorine was 
 in combination is formed. In general, the reaction is 
 represented by such equations as the following : 
 
 2MC1 + H 2 S0 4 = M 2 S0 4 + 2HC1 ; 
 MC1 2 + H 2 S0 4 = MS0 4 + 2HC1 ; etc. 
 
THE SO-CALLED DOUBLE CHLORIDES. 465 
 
 Under ordinary circumstances, chlorides are not decom- 
 posed by any acid except sulphuric acid. 
 
 The So-called Double Chlorides and Similar Compounds 
 of Fluorine, Bromine, and Iodine. These compounds and 
 their relations to the oxygen salts have been repeatedly 
 referred to. Many chlorides combine with the chlorides 
 of the stronger metals, like sodium and potassium, forming 
 well-characterized compounds. Generally, these double 
 chlorides are analogous to the oxygen salts in com- 
 position, differing from them only by containing two 
 atoms of chlorine in the place of each of the oxygen 
 atoms. As examples of these salts of the Moro-acids 
 those which are formed by the chlorides of platinum, 
 antimony, chromium, and gold may be mentioned. 
 Platinic chloride, PtCl 4 , combines with other chlorides, 
 forming salts of the general composition expressed by the 
 formula PtCl 4 + 2MC1, or M 2 PtCl 6 . Antimony chlo- 
 ride, combines with three molecules of potassium chlo- 
 ride forming the compound Sb01 3 + 3KC1, or K 3 SbCl 6 . 
 Chromium chloride forms similar compounds, K 3 CrCl 6 , 
 Na 3 CrCl 6 ; and gold chloride forms compounds of the 
 general formula MAuCl 4 , which may be regarded as 
 made up of one molecule of auric chloride, AuCl 3 , and 
 one molecule of a chloride like potassium chloride. 
 A careful study of the double chlorides and the ' similar 
 compounds of fluorine, bromine, and iodine shows 
 that the chlorides of sodium and potassium, and of the 
 other elements of the group to which these metals be- 
 long, combine with most other chlorides to form so- 
 called double salts, and that the number of molecules of 
 potassium or sodium chloride which combine with another 
 chloride is limited by the number of chlorine atoms contained 
 in the other chloride.* Thus, a chloride of the formula 
 MC1 2 may form the double chlorides MC1 2 .KC1 and 
 MC1 2 .2KC1, but not MC1 2 .3KC1. So, further, a chloride of 
 the formula MC1 3 may form three different double chlo- 
 rides with the same metallic chloride. Those with potas- 
 sium will have the formulas MC1 3 .KC1, MC1 3 .2KC1, and 
 
 * There are a few exceptions to this rule, but it undoubtedly holds 
 good in by far the largest number of cases. 
 
46G INORGANIC CHEMISTRY. 
 
 MC1 3 .3KC1, but a double chloride of the formula 
 MC1 3 .4KC1 and more complicated cases seem to be im- 
 possible. Double fluorides are known in large numbers. 
 Among the best-known are the fluosilicates. Aluminium 
 forms double fluorides, one of which, having the formula 
 Na 3 AlF 6 or AlF 3 .3NaF, is the well-known mineral 
 cryolite. All these so-called " double salts " are easily 
 explained by the aid of the hypothesis that the halogen 
 contained in them has a valence greater than one, and 
 that a double atom, like C1 2 , F 2 , etc., or -C1-C1-, -F-F-, 
 plays the same part that oxygen does in the oxygen 
 salts. The following table contains the general for- 
 mulas of the possible double chlorides with potassium 
 chloride, according to the above view concerning them : 
 
 MC1.KC1 MC1 2 .KC1 MC1 3 .KC1 MC1 4 .KC1 
 MC1 2 .2KC1 MC1 3 .2KC1 MC1 4 .2KC1 
 MC1 3 .3KC1 MC1 4 .3KC1 
 MC1 4 .4KC1 
 
 These may also be written thus : 
 
 KMC1, KMC1 3 KMC1 4 KMC1 5 
 
 K 2 MC1 4 K 2 MC1 6 K 2 MC1 6 
 
 K 3 MC1 6 K 3 MC1 7 
 
 K 4 MC1 8 
 
 Different Chlorides of the Same Metal. Just as sul- 
 phur, selenium, phosphorus, and the other acid-forming 
 elements combine with chlorine and the other members 
 of the chlorine group in more than one proportion, so 
 many of the metals combine with the members of the 
 chlorine group in more than one proportion. Thus, mer- 
 cury forms the two chlorides, HgCl 2 and HgCl, known 
 respectively as mercuric and mercurous chlorides ; iron 
 forms the two chlorides FeCl 3 and FeCl 2 , known as ferric 
 and ferrous chlorides ; and tin forms stannic chloride, 
 SnCl 4 , and stannous chloride, SnCl 2 . The conversion of 
 a higher chloride into a lower one is called an act of 
 
OXIDES. 467 
 
 reduction. The change can generally be effected by 
 means of nascent hydrogen : 
 
 SnCl 4 + 2H = SnCl 2 + 2HC1. 
 FeCl 3 + H = FeCl a + HC1. 
 
 The conversion of a lower chloride into a higher one is 
 generally spoken of as an act of oxidation, for the reason 
 that it is most commonly effected by the action of oxygen. 
 Thus the most convenient way to transform ferrous chlo- 
 ride into ferric chloride is to treat it in solution in hydro- 
 chloric acid with an oxidizing agent, when a double action 
 takes place, as represented in the following equation : 
 
 2FeCl 2 + 2HC1 + O = 2FeCl 3 + H 2 O. 
 
 The same change can be effected by the direct action of 
 chlorine : 
 
 Fed, + 01 = FeCl 3 . 
 
 In this case it would obviously be incorrect to speak of 
 the process as one of oxidation. 
 
 Another method of reduction, besides that referred to 
 above, involving the action of nascent hydrogen, is that 
 illustrated in the equation 
 
 2HgCl 2 + SnCl 2 = 2HgCl + SnCl 4 . 
 
 In this case mercuric chloride is changed to mercurous 
 chloride by the action of stannous chloride. The latter 
 unites with chlorine so readily that it extracts it from 
 some other chlorides, and is itself transformed into 
 stannic chloride. While, therefore, we say that the 
 stannous chloride reduces the mercuric chloride, it is 
 equally true to say that the mercuric chloride chlorinates 
 the stannous chloride. 
 
 Oxides. The oxides occur very extensively in nature, 
 and are among the most common ores of some of the 
 important metals. The oxides of iron, tin, and man- 
 ganese, for example, occur in nature. They can be made 
 by oxidizing the metals, by heating nitrates, carbonates, 
 and hydroxides, and by heating some sulphides in con- 
 tact with the air. 
 
468 INORGANIC CHEMISTRY. 
 
 When magnesium is burned it is converted into mag- 
 nesium oxide : 
 
 Mg + O = MgO. 
 
 When lead nitrate is heated it gives off oxygen and an 
 oxide of nitrogen, and lead oxide is left behind : 
 
 Pb(NO s ) 3 = PbO + 2N0 2 + O. 
 
 When calcium carbonate is heated it yields calcium oxide 
 and carbon dioxide : 
 
 CaCO 3 = CaO + C0 2 . 
 
 When aluminium hydroxide, Al(Ofi),, is heated it loses 
 water, and aluminium oxide is left behind : 
 
 2A1(OH) 3 = A1 2 O 3 + 3H 2 O. 
 
 The sulphide of iron, when heated in contact with the 
 air, or " roasted," is converted into ferric oxide and sul- 
 phur dioxide. 
 
 Most of the oxides of the metals are insoluble in water. 
 Those of Group A, Family I, are soluble, but are con- 
 verted by water into the corresponding hydroxides. 
 
 The oxides are acted upon generally by acids forming 
 the corresponding salts. If the salt with a certain acid 
 is insoluble, the salt is not formed by the action of that 
 acid on the oxide unless the acid or its anhydride is 
 fusible and not volatile, when by fusing them together the 
 salt is formed. 
 
 Different Oxides of the Same Metal. Just as there are 
 different chlorides of the same metal, so there are differ- 
 ent oxides, and indeed there is greater variety among 
 these than among the chlorides. Iron forms three oxides, 
 ferric oxide, Fe 2 O 3 , ferroso-ferric oxide, Fe 3 O 4 , and ferrous 
 oxide, FeO ; mercury forms the two oxides HgO and Hg 2 O ; 
 etc. The lower oxides are converted into the higher 
 by oxidation, and the higher into the lower by reduc- 
 tion. The higher oxides of several of the metals are 
 acidic. This is markedly so in the case of chromium and 
 manganese. 
 
HYDROXIDES. 469 
 
 Hydroxides. The hydroxides are formed by treating 
 oxides with water and by decomposing salts by adding 
 soluble hydroxides to their solutions. In general, when- 
 ever a salt is decomposed by a strong base, the base of 
 the salt separates in the form of the hydroxide. The 
 formation of a hydroxide by the action of water on an 
 oxide is well illustrated by the action of water on lime 
 or calcium oxide, a process which is familiarly known as 
 slaking : 
 
 CaO + H 2 O =t Ca(OH),. 
 
 Most of the hydroxides of the metals are insoluble in 
 water. If a soluble hydroxide is added to a solution 
 containing a metal whose hydroxide is insoluble, the 
 latter is precipitated. Thus, if a solution of sodium hy- 
 droxide is added to a solution of a magnesium salt, 
 magnesium hydroxide is precipitated : 
 
 MgSO 4 + 2NaOH = Na 2 SO 4 + Mg(OH) 2 . 
 
 So, also, when a solution of a ferric salt is treated with 
 sodium hydroxide, a precipitate of ferric hydroxide is 
 formed : 
 
 FeCl 3 + 3NaOH = 3NaCl + Fe(OH) 3 . 
 
 Only the hydroxides of the members of the potassium 
 family, and some of the members of the calcium family, 
 are soluble in water. The hydroxides of sodium and 
 potassium are called alkalies. The solution of ammonia 
 in water acts like a soluble hydroxide, and probably con- 
 tains ammonium hydroxide, NH 4 (OH), formed by the 
 action of water on ammonia : 
 
 NH 3 + H 2 = NH 4 (OH). 
 
 Now, when any one of the soluble hydroxides is added to 
 a salt containing any metal which does not belong to the 
 potassium or calcium family, an insoluble compound is 
 formed. 
 
 Decomposition of Salts by Bases. The decomposition 
 of salts by bases is analogous to the decomposition by 
 
470 INORGANIC CHEMISTRY. 
 
 acids. When a soluble base acts upon a salt, there are 
 four possible kinds of action : 
 
 1. The base from which the salt is derived may be 
 volatile, or may break up, yielding a volatile product. 
 
 In this case, decomposition takes place and the volatile 
 base is given off. This is not a common case except 
 among the compounds of carbon. The one illustration 
 which we have had is the decomposition of ammonium 
 salts by calcium hydroxide and sodium hydroxide, when 
 the volatile compound ammonia, NH 3 , is given off. 
 
 2. The hydroxide, or base from which the salt is de- 
 rived, may be insoluble or difficultly soluble in water, 
 and not volatile. 
 
 In this case, if both the salt and the base are in solu- 
 tion, decomposition takes place, and the insoluble or 
 difficultly soluble hydroxide, or base, is precipitated. 
 This has already been illustrated. 
 
 3. The base from which the salt is derived may be 
 soluble and not volatile. 
 
 This is the case, for example, when sodium hydroxide 
 is added to a solution of potassium nitrate. Here sodium 
 nitrate, potassium nitrate, sodium hydroxide, and potas- 
 sium hydroxide may all be present in the solution, and 
 investigation has shown that all are present and that the 
 quantity of each depends upon the masses of the sub- 
 stances brought together, and upon their affinities. 
 
 4. The fourth case is that in which a soluble hydroxide 
 forms an insoluble salt with the acid of a soluble salt, 
 leaving a soluble hydroxide in solution. 
 
 This is illustrated by the action of calcium hydroxide 
 on a solution of sodium carbonate, when insoluble cal- 
 cium carbonate is thrown down, and sodium hydroxide 
 remains in solution, as represented in the equation 
 
 Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCO 3 . 
 
 Some basic hydroxides, which are precipitated by solu- 
 ble hydroxides, have a weak acid character, and, after they 
 are precipitated, they redissolve in an excess of the solu- 
 ble hydroxide. This is true, for example, of aluminium, 
 
DECOMPOSITION OF SALTS BY BASES. 471 
 
 chromium, and lead. The salt-like compounds thus 
 formed are generally quite unstable. The precipitation 
 and subsequent solution of the hydroxides of the three 
 metals named take place thus : 
 
 A1C1, + 3NaOH = Al(OH), + 3NaCl; 
 A1(OH) 3 + 3NaOH = Al(ONa) 3 + 3H a O ; 
 CrCl 3 + 3NaOH = Cr(OH), + 3NaCl ; 
 Cr(OH) 3 + 3NaOH = Cr(ONa) 3 + 3H O ; 
 Pb(NO 3 ) 2 + 2NaOH = Pb(OH) 2 + 2NaNO 3 ; 
 Pb(OH) 2 + 2NaOH = Pb(ONa) 2 + 2H 2 O. 
 
 In some cases where a soluble hydroxide is added to a 
 salt, an oxide is precipitated instead of the hydroxide. 
 This is analogous to the formation of an anhydride of 
 an acid instead of the acid itself, as when carbonates, 
 sulphites, and nitrites are decomposed. 
 
 When a silver salt is treated with a soluble hydroxide, 
 silver oxide is at once precipitated. The same is true of 
 mercury salts : 
 
 2AgN0 3 + 2KOH = Ag 2 + H 2 O + 2KNO 3 ; 
 HgCl 2 + 2NaOH = HgO + H 2 O + 2NaCl. 
 
 It is probable that the first product is the hydroxide, 
 and that this breaks down into the oxide and water : 
 
 2AgNO 3 + 2KOH = 2AgOH + 2KNO 3 ; 
 
 2AgOH = Ag 2 + H 2 0; 
 
 HgCl 2 + 2NaOH = Hg(OH) 2 + 2NaCl ; 
 
 Some hydroxides are converted into the oxides by 
 simply boiling the liquids in which they are suspended. 
 Thus, when a salt of copper is treated with a soluble hy~ 
 droxide, copper hydroxide is first precipitated ; but if the 
 solution in which it is suspended is boiled, it is soon 
 changed to the oxide : 
 
 CuSO 4 + 2NaOH = Na 2 SO 4 + Cu(OH) 2 ; 
 
472 INORGANIC CHEMISTRY. 
 
 The hydroxides corresponding to some of the highei 
 oxides of the metals, as those of chromium and mangan- 
 ese, are acids. 
 
 The hydroxides of most of the metals are decomposed 
 by heat into water and the corresponding oxides. Those 
 of the alkali metals, as potassium and sodium, are not, 
 however, decomposed by heat. 
 
 Sulphides. Many sulphides are found in nature, as, 
 for example, iron pyrites, FeS 2 ; lead sulphide, or galen- 
 ite, PbS ; copper pyrites, FeCuS 2 ; etc. They are made 
 in the laboratory by heating metals with sulphur ; by 
 treating solutions of salts with hydrogen sulphide ; by 
 treating solutions of salts with soluble sulphides ; and 
 by reducing sulphates. Attention has been called to the 
 fact that the sulphides are analogous in composition to 
 the oxides, and that they are to be regarded as salts of 
 hydrogen sulphide formed by replacing the hydrogen of 
 the acid by metals. 
 
 The formation of sulphides by the direct combination 
 of sulphur with the metals is shown in the formation of 
 lead sulphide and copper sulphide : 
 
 2Cu + S = Cu 2 S. 
 
 The formation of sulphides by the action of hydrogen 
 sulphide upon solutions of salts was discussed at some 
 length under Hydrogen Sulphide (which see). The ex- 
 tensive use made of this reaction in chemical analysis 
 was also referred to. 
 
 The action of soluble sulphides or solutions of salts is 
 in general the same as that of Irydrogen sulphide, but in 
 some cases, in which the former will not act, the latter 
 will. Thus, hydrogen sulphide will not precipitate iron 
 sulphide from a solution of an iron salt, because iron 
 sulphide is easily acted upon by dilute acids. Thus, 
 when hydrogen sulphide is passed into a solution of 
 ferrous chloride, it naturally tends to form the sul- 
 phide FeS: 
 
 FeCl 2 + H 2 S = FeS + 2HC1. 
 
SULPHIDES. 
 
 But ferrous sulphide, FeS, is acted upon by dilute hy 
 drochloric acid, and is converted by it into ferrous chlo- 
 ride and hydrogen sulphide : 
 
 FeS + 2HC1 = FeCl 2 + H 2 S. 
 
 It is therefore obvious that the first reaction cannot take 
 place. 
 
 If, however, a soluble sulphide, as sodium or ammonr 
 ium sulphide, is added to a solution of an iron salt, iron 
 sulphide is precipitated, as in this case no free acid is 
 formed. Thus, when ferrous chloride and ammonium 
 sulphide are brought together the reaction takes place 
 as represented in the equation 
 
 FeCl 2 + (NH 4 ) 2 S = FeS + 2NH 4 C1. 
 
 In ammonium chloride the ferrous sulphide is not 
 soluble. 
 
 The formation of a sulphide by reduction of a sul- 
 phate is illustrated by the formation of barium sulphide 
 by heating a mixture of barium sulphate and charcoal : 
 
 BaSO 4 + 4C = BaS + 4CO ; 
 
 and by the formation of copper sulphide by heating 
 copper sulphate in a current of hydrogen : 
 
 CuS0 4 + 4H 2 = CuS + 4H 2 0. 
 
 The sulphides of the alkali metals are soluble in water. 
 Those of the other metals are insoluble. It should be 
 remarked, however, that aluminium and chromium do 
 not form sulphides, or, at least, if they do, the compounds 
 are decomposed by water into hydroxides and hydrogen 
 sulphide. Barium sulphide is decomposed by water, 
 and probably magnesium sulphide also. 
 
 The sulphides are stable when heated without access 
 of air ; but if heated in the air they are converted into 
 oxides of the metals and sulplmr dioxide, or, in some 
 cases, they take up oxygen and are converted into sul- 
 phates. The conversion of sulphides into oxides and 
 sulphur dioxide by heating in contact with the air has 
 
474 INORGANIC CHEMISTRY. 
 
 been repeatedly referred to. The process is carried on 
 on the large scale in the preparation of iron ores for re- 
 duction, and is called roasting. The conversion of a 
 sulphide into a sulphate by heating is a simple process 
 of oxidation. Copper sulphide is converted into the 
 sulphate when heated for some time : 
 
 CuS + 4O = CuS0 4 . 
 
 This is the reverse of the reaction mentioned by which a 
 sulphate is converted into a sulphide by reduction. 
 
 Some sulphides, as those of sodium, potassium, and 
 ammonium, take up sulphur in much the same way that 
 they take up oxygen, and form the polysulphides. The 
 two reactions appear to be entirely analogous : 
 
 K 2 S + 40 = K 2 S0 4 ; 
 
 K 2 S + 48 = K 2 SS 4 , or K 2 S 6 . 
 
 Hydrosulphid.es. The hydrosulphides bear the same 
 relation to the sulphides that the hydroxides bear to the 
 oxides. They are not, however, as numerous nor as 
 easily obtained as the hydroxides. When a hydrosul- 
 phide, as, for example, potassium hydrosulphide, KSH, 
 is added to a salt containing a metal whose sulphide 
 is insoluble, the sulphide, and not the hydrosulphide, 
 is precipitated. Thus, copper sulphate and potassium 
 hydrosulphide give copper sulphide : 
 
 CuSO 4 + 2KSH = CuS + H 2 S + K 2 SO 4 . 
 
 If the reaction took place in the same way that it does with 
 the hydroxide, the product would be copper hydrosul- 
 phide : 
 
 CuSO 4 + 2KSH = Cu(SH) 2 + K 2 SO 4 . 
 
 If this is formed it certainly breaks down into copper 
 sulphide and hydrogen sulphide, in the same way that 
 copper hydroxide breaks down into copper oxide and 
 water, only more easily : 
 
 Cu(SH) 2 = CuS + H 2 S ; 
 Cu(OH) 2 = CuO + H 2 0. 
 
SULPHO-SALTS. 475 
 
 The only hydrosulphides known are derived from the 
 members of the potassium and calcium groups, and these 
 are soluble. They are formed by saturating solutions of 
 the corresponding hydroxides with hydrogen sulphide. 
 Potassium hydrosulphide is formed thus : 
 
 KOH + H 2 S = KSH + H 2 O. 
 Ammonium hydrosulphide is formed thus : 
 
 NH 4 OH + H 2 S = NH 4 SH + H 2 O. 
 
 It also appears probable that whenever a sulphide is 
 dissolved in water it is converted into a hydrosulphide 
 and a hydroxide. Thus it seems to be true that potas- 
 sium sulphide is converted into the hydrosulphide and 
 hydroxide : 
 
 K 2 S + H 2 = KSH + KOH. 
 
 Sulpho-salts. The relation of the sulpho-salts to the 
 sulphides has already been explained. It is like that of 
 the ordinary oxygen salts to the oxides, and that of the 
 chloro-salts, or double chlorides, to the chlorides. They 
 are formed by dissolving the sulphides of certain metals, 
 particularly tin, arsenic, and antimony, in the sulphides 
 of the members of the potassium group : 
 
 As 2 S 3 + 3K 2 S = 2K 3 AsS 3 ; 
 As 2 S 5 + 3K 2 S = 2K 3 AsS 4 ; 
 SnS 2 + K 2 S = K 2 SnS 3 , etc. 
 
 These sulpho-salts are decomposed by the ordinary acids, 
 the insoluble sulphides being precipitated thus : 
 
 2K 3 AsS 3 + 6HC1 = As 2 S 3 + 6KC1 + 3H 2 S ; 
 2K 3 AsS 4 + 6HC1 = As 2 S 5 + 6KC1 + 3H 2 S. 
 
 Nitrates. The nitrates are formed by dissolving the 
 metals in nitric acid, and by treating oxides, hydroxides, 
 carbonates, and some other easily decomposed salts with 
 nitric acid. The action of nitric acid upon metals was 
 discussed under the head of Nitric Acid (which see). It 
 was pointed out that the acid gives up a part of its 
 
476 INORGANIC CHEMISTRY. 
 
 oxygen to the metal and forms different oxides, accord- 
 ing to the conditions. Thus, when the acid acts upon 
 copper the main product of the reduction is nitric 
 oxide, but by changing the concentration of the acid a 
 considerable quantity of nitrous oxide is formed. When 
 zinc is dissolved in nitric acid a part of the acid is 
 reduced to ammonia. 
 
 The nitrates are soluble in water, and all are decom- 
 posed by heat. Some of them when heated lose only a 
 third of their oxygen and are reduced to nitrites. This 
 is true of potassium nitrate, the decomposition of which 
 is represented by the equation 
 
 KNO 3 = KNO 2 + O. 
 
 Most of the nitrates, however, are decomposed further, 
 forming oxides. This has been shown in the case of lead 
 nitrate, which when heated is converted into lead oxide, 
 while nitrogen peroxide and oxygen are given off: 
 
 Pb(N0 3 ) 2 = PbO + 2NO 2 + O. 
 
 If the oxide of the metal is decomposed by heat, as in 
 the case of mercury, of course the product will be the 
 metal. 
 
 Chlorates. These salts, except potassium chlorate, are 
 not commonly met with. Potassium chlorate is manu- 
 factured in large quantity, and the other chlorates are 
 generally made from it. The chlorates are soluble in 
 water, and are decomposed by heat more easily than the 
 nitrates are. They are first converted into perchlorates, 
 and these are further decomposed by higher heat into 
 chlorides and oxygen. 
 
 The hypocJdorites are formed by treating some of the 
 metallic hydroxides in dilute solution with chlorine. 
 This has been illustrated in the formation of " bleaching 
 powder," which contains calcium hypochlorite or a com- 
 pound closely related to it. The hypochlorites, like the 
 chlorates, are decomposed by heat. 
 
 Sulphates. The general relations of the sulphates to 
 sulphuric acid were treated of under Sulphuric Acid 
 (which see). Some of these salts occur in nature in large 
 
SULPHATES. 477 
 
 quantity, as those of calcium and barium. The former 
 is known as gypsum, the latter as heavy spar. Sulphates 
 are made by treating metals, metallic hydroxides or ox- 
 ides, carbonates, etc., with sulphuric acid ; and by treat- 
 ing a solution containing a metal whose sulphate is 
 insoluble, with sulphuric acid or a soluble sulphate. 
 
 Zinc and iron give hydrogen and a sulphate when 
 treated with sulphuric acid : 
 
 Zn + H 2 S0 4 = ZnSO 4 + H 2 ; 
 Fe + H 2 S0 4 = FeS0 4 + H 2 . 
 
 This kind of action takes place whenever a metal is dis- 
 solved in sulphuric acid at the ordinary temperature. If, 
 however, the temperature is raised the displaced hydro- 
 gen acts upon some of the sulphuric acid, or the metal 
 extracts some of the oxygen of the acid, reducing it 
 partly to sulphurous acid, when sulphur dioxide is given 
 off. This happens in the case of copper, as has been 
 pointed out. It may be represented either by these 
 equations : 
 
 Cu + H 2 SO 4 = CuS0 4 + H 2 ; 
 H 3 S0 4 + 2H = SO a + 2H 2 O ; 
 
 or by these : 
 
 Cu + H 2 SO 4 = CuO + SO 2 + H a O ; 
 CuO + H 2 SO 4 = CuSO 4 + H 3 O. 
 
 The action of sulphuric acid on metallic hydroxides 
 has been fully described. 
 
 Most sulphates are soluble in water. The sulphates 
 of barium, strontium-, and lead are insoluble in water, 
 and the sulphate of calcium is difficultly soluble. There- 
 fore, if sulphuric acid or a soluble sulphate is added to 
 a solution containing either of the metals, barium, stron- 
 tium, or lead, a precipitate is formed. A precipitate 
 is also formed when a concentrated solution of a calcium 
 salt is treated in the same way. 
 
478 INORGANIC CHEMISTRY. 
 
 When heated with charcoal in the reducing flame of 
 the blow-pipe, sulphates are reduced to sulphides : 
 
 K 2 SO 4 + 40 = K 2 S + 4CO, or 
 K 2 SO 4 + 2C = K 3 S + 2C0 2 . 
 
 Sulphites are made from sodium or potassium sul- 
 phite, which are made by treating sodium or potassium 
 hydroxide in solution with sulphur dioxide : 
 
 2NaOH + SO 2 = Na 2 SO 3 + H 2 O. 
 
 All sulphites are decomposed by the common acids, 
 sulphur dioxide being given off : 
 
 Na 2 SO 3 + H 2 SO 4 = Na 2 SO 4 + H 2 O + SO,. 
 
 The sulphites are changed to sulphates by oxidation. 
 Thus, sodium sulphite is changed to the sulphate when 
 its solution is allowed to stand in contact with the air : 
 
 Na 2 S0 3 + O = Na 2 SO 4 . 
 
 The sulphites, like the sulphates, are reduced to sul- 
 phides. 
 
 Carbonates. Many carbonates are found in nature, 
 some of them in great abundance and widely distrib- 
 uted. The principal one is calcium carbonate. They 
 are made by passing carbon dioxide into solutions of hy- 
 droxides, and by adding soluble carbonates to solutions 
 of salts containing metals whose carbonates are insolu- 
 ble. 
 
 The formation of carbonates by the action of carbon 
 dioxide on a solution of hydroxide is illustrated in the 
 case of potassium hydroxide : 
 
 2KOH + CO 2 = K 2 C0 3 + H 2 O. 
 
 The formation of calcium carbonate takes place in the 
 same way, but the carbonate formed is insoluble : 
 
 Ca(OH) 2 + C0 2 = CaCO 3 + H,0. 
 
PHOSPHATES. 479 
 
 If in either case the action is continued, the normal 
 carbonate first formed is converted into acid carbonate : 
 
 K 2 CO 3 + CO, + H 2 O = 2KHCO 3 ; 
 CaCO 3 + CO 2 + H 2 O = Ca ) , pn x 
 
 H 2 f v^Vt/* 1 
 
 All normal carbonates except those of the members of 
 the potassium family are insoluble, and are decomposed 
 by heat into carbon dioxide and the oxide of the metal. 
 The decomposition of calcium carbonate into lime and 
 carbon dioxide is an example : 
 
 CaCO 3 = CaO + CO 2 . 
 
 When a soluble carbonate is added to a solution of a 
 calcium, barium, or strontium salt, the corresponding 
 insoluble carbonates are precipitated. When a magne- 
 sium salt is treated with a soluble carbonate, however, a 
 basic carbonate is precipitated : 
 
 4MgSO 4 + 3Na 2 CO 3 + 2H 2 O = Mg 4 (OH) 2 (CO 3 ) 3 + 3Na 2 SO 4 + H 2 SO 4 . 
 
 This salt, which at first sight appears to be quite com- 
 plicated, is in all probability derived from three mole- 
 cules of carbonic acid and four of magnesium hydroxide, 
 as represented in the formula on page 388. Many other 
 metals give basic carbonates under the same conditions. 
 Further, some of the metals, like aluminium, chromium, 
 and tin do not form salts with carbonic acid. If, there- 
 fore, salts of these metals are treated with soluble car- 
 bonates, the oxides or hydroxides are thrown down, and 
 not the carbonates. 
 
 Phosphates. Calcium phosphate is very abundant in 
 nature, and a few other phosphates are also found. The 
 methods used for making phosphates are the same as 
 those used in making salts in general. 
 
 The normal phosphates of all the metals except the 
 members of the potassium family are insoluble in water. 
 The normal phosphates, as a rule, are not changed by 
 heat. The secondary phosphates, such as secondary 
 
480 INORGANIC CHEMISTRY. 
 
 sodium phosphate, HNa a PO 4 , lose water when heated,, 
 and yield pyrophosphates : 
 
 2HNa a PO 4 = Na 4 P 2 O 7 + H 2 O. 
 
 Sodium 
 pyrophosphate. 
 
 Those phosphates in which only one third of the hy- 
 drogen is replaced by metal as, for example, primary 
 sodium phosphate, H a NaPO 4 lose water when heated, 
 and yield metaphosphates : 
 
 H a NaPO 4 = NaPO 3 + H 2 O. 
 
 Sodium 
 metaphosphate 
 
 Neither the pyrophosphates nor the metaphosphates are 
 changed by heat. 
 
 Silicates. The silicates, as has been stated, are very 
 widely distributed in nature. Those which are most 
 abundant are the feldspars and their decomposition-pro- 
 ducts. The principal feldspar is a complex silicate of 
 aluminium and potassium, of the formula KAlSi 3 O 8 , de- 
 rived from the polysilicic acid H 4 Si 3 O 8 , which is formed 
 from three molecules of normal silicic acid by the loss 
 of four molecules of water : 
 
 3Si(OH) 4 =H 4 Si 3 O 8 + 4H S O. 
 
 Silicates can be made by heating together, at a high 
 temperature, silicon dioxide, in the form of sand, and 
 basic oxides or carbonates : 
 
 CaO + SiO a = CaSi0 3 ; 
 Na a CO 3 + SiO 2 = Na a Si0 3 + C0 2 . 
 
 Only the silicates of the members of the potassium 
 group are soluble in water. When these are treated in 
 solution with dilute acids, they are decomposed, as has 
 been explained under Silicic Acid (which see). 
 
 Some silicates, which are insoluble in water, are decom- 
 posed by the ordinary acids, such as sulphuric and hy- 
 drochloric acids, the silicic acid separating as a difficultly 
 soluble substance, which, if dried on the water-bath, be- 
 comes insoluble. 
 
SILICATES. 481 
 
 Many silicates, which are not acted upon by strong 
 acids, are decomposed by fusion with sodium or potassi- 
 um carbonate, when the silicate of potassium or sodium 
 and the oxide of the metal contained in the silicate are 
 formed. Silicates which are not decomposed in either 
 of the ways mentioned, yield to hydrofluoric acid. The 
 action consists in the formation of the gas, silicon tetra- 
 fluoride, SiF 4 , and the fluorides of the metals present. 
 Thus, the reaction in the case of feldspar takes place in 
 accordance with the equation, 
 
 KAlSi 3 O 8 + 16HF = KF + A1F 3 + 3SiF 4 + 8H 2 O. 
 
 The silicon fluoride is given off as a gas, and the flu- 
 orides formed are soluble in water. Hence, hydrofluoric 
 acid is said to dissolve the silicates. 
 
CHAPTER XXV 
 
 ELEMENTS OF FAMILY I, GROUP A: 
 
 THE ALKALI METALS : LITHIUM SODIUM POTASSIUM- 
 RUBIDIUM CJESIUM AMMONIUM. 
 
 General. The elements of this group which are most 
 abundant in nature are sodium and potassium. While 
 lithium occurs in considerable quantity, the two remain- 
 ing elements, rubidium and caesium, have been found in 
 only very small quantities. , They are all strongly basic, 
 their hydroxides being the strongest bases known. They 
 form well- characterized salts with all acids, and as a 
 rule their salts are very stable. In all their compounds 
 they act as univalent elements, except in those which 
 they form with hydrogen, and in their peroxides ; in the 
 latter they appear to be bivalent. Leaving these com- 
 pounds out of consideration the general formulas of 
 some of the other principal compounds are as follows : 
 
 MCI, M 2 O, M 2 S, M(OH), M(SH), MNO 3 , M 2 SO 4 , etc. 
 
 The valence of the members of the group towards other 
 elements is, in general, constant. 
 
 The relations between the atomic weights are interest- 
 ing. That of sodium, 22.82, is very nearly half the sum of 
 those of lithium, 6.97, and potassium, 38.82. We have 
 
 - ("7 88.* 
 
 i 
 
 A 
 
 So, also, that of rubidium, 84.78, is approximately half 
 the sum of those of potassium, 38.82, and caesium, 131.89. 
 
 38.82 + 131.89 
 
 ~ = oD.OO. 
 
 a 
 
 (482) 
 
POTASSIUM. 483 
 
 The specific gravity of these elements increases with 
 the atomic weight; and their melting-points become 
 lower as the atomic weights become higher. 
 
 At. Wt. Sp. Grav. M. P. 
 
 Lithium, . . . 6.97 0.594 180. 
 
 Sodium, . . . 22.82 0.972 95.6 
 
 Potassium, . . 38.82 0.865 62.5 
 
 Kubidium, . . 84.78 1.52 38.5 
 
 Cesium, . . . 131.89 ? ? 
 
 The regularity is complete in the case of the melting- 
 points, but as regards the specific gravities sodium is an 
 exception to the rule. As sodium and potassium and 
 their compounds are much more commonly met with 
 than the other members of the group, these will form 
 the chief subject of consideration in this chapter. 
 
 POTASSIUM, K (At. Wt. 38.82). 
 
 Occurrence. Potassium is a constituent of many min- 
 erals, particularly of feldspar, the common variety of 
 which, as has already been explained, is a complex sili- 
 cate of aluminium and potassium. It is found also in 
 combination with chlorine as carnallite and sylvite ; with 
 sulphuric acid and aluminium, as alum ; with nitric acid, 
 as saltpeter or potassium nitrate ; and in other forms. 
 The natural decomposition of minerals containing potas- 
 sium gives rise to the presence of this metal in various 
 forms of combination everywhere in the soil. It is taken 
 up by the plants ; and, when vegetable material is burned, 
 the potassium remains behind, chiefly as potassium car- 
 bonate. When wood-ash is treated with water the 
 potassium carbonate is dissolved, and it can be obtained 
 in an impure state by evaporating the solution. The 
 substance thus obtained is called potash. In the juice of 
 the grape there is contained a salt of potassium, mono- 
 potassium tartrate, which is deposited in large quantity 
 from wine. This is commonly called "crude tartar." 
 
 Preparation. Potassium was first prepared by Davy 
 in the year 1807, by the action of a powerful electric cur- 
 rent on potassium hydroxide. It is now prepared by 
 
484 INORGANIC CHEMISTRY. 
 
 heating to a high temperature a mixture of potassium 
 carbonate and carbon : 
 
 K a CO 3 + 20 = 2K + 3CO. 
 
 Such a mixture is best obtained by heating in a closed 
 vessel ordinary mono-potassium tartrate obtained from 
 wine. This contains some calcium tartrate. When the 
 whole is heated decomposition takes place, and there is 
 left behind an intimate mixture of potassium carbonate, 
 calcium carbonate, and charcoal. This mixture is placed 
 in a wrought-iron retort which is connected with a closed 
 flat receiver of sheet-iron. The retort is then heated to 
 a high temperature. The metal distils over into the 
 closed receiver, and at the end of the operation the re- 
 ceiver is placed under petroleum to protect it from the 
 action of the air. The metal obtained in this way is 
 not pure. It can be partly purified by melting it under 
 petroleum and pressing it through a linen bag. It can 
 also be purified, and more completely, by distilling it 
 from a wrought-iron retort. 
 
 Properties. Potassium is a light substance, which 
 floats on water. Its freshly cut surface has a bright 
 metallic lustre, almost white ; it acts energetically upon 
 water, causing the evolution of hydrogen, which, together 
 with some of the potassium, burns, while potassium hy- 
 droxide is formed at the same time. This reaction has 
 been studied in connection with hydrogen. In conse- 
 quence of its action upon water, potassium cannot be 
 kept in the air. It is kept under some oil, as petroleum, 
 upon which it does not act. In an atmosphere upon 
 which it does not act, as, for example, hydrogen, it can 
 be distilled. Its vapor is green. Its specific gravity is 
 0.865; its melting-point 62.5. It combines with chlo- 
 rine and bromine with great energy, and has the power 
 to extract chlorine from its compounds. It can, there- 
 fore, be used for the purpose of isolating some elements, 
 as, for example, magnesium and aluminium, whose oxy- 
 gen compounds cannot be reduced by the ordinary 
 methods. As, however, sodium is generally used for 
 this purpose instead of potassium -, on account of its lower 
 
POTASSIUM SALTS. 485 
 
 price, the action will be referred to more at length under 
 Sodium. Although the metal is converted into vapor, 
 no reliable determination of the specific gravity of the 
 vapor has been made, for the reason that the vessels 
 which have been used for the purpose have always 
 been acted upon, and the results thus vitiated. 
 
 Potassium Hydride, K 2 H. This compound is formed 
 by heating potassium in an atmosphere of hydrogen at 
 about 300. It is a silver-white mass with a metallic 
 lustre. It takes fire in the air. When heated, it begins 
 to dissociate at 200. 
 
 [ Fluoride, KF 
 
 Potassium] gjEdd?; f& Of these salts the only 
 
 L Iodide, KI 
 
 one which occurs in nature in quantity is the chloride. 
 This is found in the great salt deposits at Stassfurt, 
 Germany, and in some other localities in the form of the 
 mineral sylvite, which is more or less impure potassium 
 chloride. It is also found in the form of a compound 
 containing magnesium, potassium, and chlorine, of the 
 formula MgCl 2 .KCl + 6H 2 O, or KMgCl 3 + 6H 2 O, known 
 as carnallite. 
 
 The other salts of the group are made by the general 
 methods for making salts, that is, by neutralizing the 
 acids with the hydroxide or carbonate of potassium. It 
 is, however, easier to make the iodide by other methods, 
 and as there is a large demand for this salt for use in 
 medicine and in the art of photography, several methods 
 have been devised for its preparation. Of these, two 
 may serve as examples : (1) The first consists in treating 
 a solution of potassium hydroxide with iodine until it 
 begins to show a permanent yellow color, which is an 
 indication that no more iodine will be taken up. The 
 action is the same as that which takes place when chlo- 
 rine acts upon warm concentrated caustic potash. Both 
 the iodide and iodate are formed : 
 
 6KOH + 61 = SKI + KIO 3 + 3H 2 O. 
 
 By evaporating the water and heating the residue with 
 very finely powdered charcoal, the iodate is decomposed 
 
486 INORGANIC CHEMISTRY. 
 
 into iodide and oxygen. The reduction of the iodate 
 takes place in accordance with the equation : 
 
 2KIO 3 + 30 = 2KI + 3CO 2 . 
 
 (2) Another method employed in the preparation of 
 potassium iodide consists in treating iron filings under 
 water with iodine. Both the iron and iodine dissolve, 
 forming ferrous iodide, FeI 2 . If to the solution of this 
 compound half as much iodine is added as has already 
 been used in its preparation, ferroso-ferric iodide, Fe 3 I 8 , 
 is formed and remains in solution. By adding a solution 
 of potassium carbonate to this, reaction takes place as 
 represented in the equation : 
 
 Fe 3 I. + 4K 2 C0 3 + 4H 2 = SKI + Fe 3 (OH) 8 + 4CO 2 t 
 
 The hydroxide of iron is insoluble, and can be removed 
 by filtration. 
 
 The fact that the specific gravity of hydrofluoric acid 
 at a low temperature corresponds to the formula H 2 F 2 
 makes it not improbable that potassium fluoride has 
 the formula K 2 F 2 . This appears still more probable 
 from the fact that there is an acid potassium fluoride of 
 the formula KHF 2 , or KF + HF. Similar acid salts 
 have not been obtained from the other acids of the 
 group. 
 
 Properties. All these salts are soluble, and crystallize 
 well in cubes. The fluoride is the most easily soluble 
 in water. If deposited from a water solution at the or- 
 dinary temperature the crystals contain two molecules 
 of water of crystallization, and are deliquescent. The 
 iodide is soluble in 0.7 parts of water at the ordinary 
 temperature, and is also soluble in alcohol (40 parts). 
 The bromide requires about 1^ parts of water for solu- 
 tion at the ordinary temperature, and is but slightly 
 soluble in alcohol. The chloride is soluble in 3 parts of 
 water at the ordinary temperature, and is insoluble in 
 alcohol. All are decomposed by sulphuric acid. The 
 fluoride gives hydrofluoric acid ; the chloride gives hy- 
 drochloric acid. The bromide gives hydrobromic acid, 
 which acts upon the sulphuric acid, giving sulphur di- 
 
DOUBLE HALOGEN SALTS. 
 
 48: 
 
 oxide and free bromine (see Hydrobromic Acid). The 
 action in the case of the iodide is more complicated for 
 the reason that hydriodic acid is less stable than hydro- 
 bromic acid, and, as it gives up hydrogen very easily, it 
 causes deeper-seated decomposition of the sulphuric 
 acid (see Hydriodic Acid). Potassium iodide in solution 
 takes up iodine readily, and a compound of the formula 
 KI 3 can be isolated from a very concentrated solution. 
 No similar compounds of chlorine, bromine, and fluorine 
 are known. 
 
 All the salts of this group combine readily with the 
 fluorides, chlorides, bromides, and iodides of the metallic 
 elements in general, forming salts, of which the double 
 fluorides and double chlorides are examples. The re- 
 lations between these salts and the ordinary oxygen 
 salts have already been discussed to a sufficient extent 
 (see pp. 465 and 466). Those containing fluorine have 
 been studied most fully. Good examples are the fol- 
 lowing : 
 
 F 
 F 
 
 As^jF ,orAsF B .KF; 
 F 
 F 2 K 
 
 rr 
 
 F 
 
 F , or AsF 6 .2KF ; 
 F 2 K 
 [F 2 K 
 
 F 
 
 F (F 
 
 F ,orSbF 5 .2KF; B^F ,orBF 3 .KF; 
 
 F 2 K F 2 K 
 
 F 2 K 
 
 Applications. Potassium chloride is extensively used 
 for the purpose of making other potassium salts, as, for 
 example, the nitrate and carbonate ; the bromide is used 
 in medicine ; the iodide, as stated above, is used in medi- 
 cine and in photography. 
 
488 INORGANIC CHEMISTRY. 
 
 Potassium Hydroxide, KOH. This well-known sub- 
 stance, commonly called caustic potash, is prepared by 
 treating potassium carbonate in solution with calcium 
 hydroxide in a silver or iron vessel. The reaction is 
 based upon the fact that calcium carbonate is insoluble, 
 and that potassium carbonate and calcium hydroxide 
 are soluble : 
 
 K 2 CO 3 + Ca(OH) 2 = 2KOH + CaCO 3 . 
 
 After enough lime has been added, it is found that a 
 little of the liquid taken out of the vessel gives no carbon 
 dioxide when treated with acids. When this point is 
 reached the liquid is drawn off from the deposit of cal- 
 cium carbonate by means of a siphon. In the prepara- 
 tion on the large scale this is then evaporated down in a 
 bright wrought-iron vessel until it has the specific gravity 
 1.16. If the evaporation is carried farther the liquid 
 acts upon the iron. Concentration beyond this point 
 must be carried on in silver vessels, upon which potas- 
 sium hydroxide does not act. Finally, a liquid is ob- 
 tained which on cooling completely solidifies. While in 
 the molten condition it is generally poured into moulds of 
 cast-iron or of brass, plated with silver, in which it solidi- 
 fies in the form of the thin sticks found in the market. 
 This substance is generally not pure. It always con- 
 tains some carbonate formed by the action of the carbon 
 dioxide of the air, and other substances are also present 
 in small quantity. It can be purified by dissolving it 
 in alcohol, in which the impurities are insoluble. The 
 alcoholic solution of the hydroxide is poured off after a 
 time and evaporated to dryness in a silver vessel. The 
 liquid becomes colored in consequence of a partial de- 
 composition of the alcohol, but on melting the residue 
 the color disappears, as the substances formed from the 
 alcohol are thus destroyed. This product is known as 
 " caustic potash by alcohol." Pure potassium hydroxide 
 in solution is easily obtained by the action of potassium 
 upon distilled water. 
 
 Potassium hydroxide is a white brittle substance. In 
 contact with the air it deliquesces, and absorbs carbon 
 
POTASSIUM OXIDE. 489 
 
 dioxide, being completely transformed into potassium 
 carbonate. It is the strongest of the bases. It decom- 
 poses the salts of all other bases, even of those which, 
 like sodium and lithium hydroxides, are soluble in water. 
 
 Animal substances like the skin are disintegrated by 
 the hydroxide. It has a caustic action. It is interest- 
 ing to observe that the strongest bases, like the strongest 
 acids, exert this kind of influence on the complex organic 
 compounds which go to make up the tissues of animals. 
 The action is not by any means always of the same kind, 
 and all that can be said in regard to it, of a general 
 character, is that it tends to break down the complex 
 substances to simpler ones. In the molten condition 
 the hydroxide acts as an oxidizing agent. Hydrogen is 
 given up from it, and substances of acid character are 
 formed with which the potassium combines, forming 
 salts. 
 
 Instead of potassium hydroxide, the corresponding 
 sodium compound is used wherever this is possible, as 
 the latter is cheaper. The chief application of the 
 potassium compound outside of the laboratory is for 
 making soft-soap. For this purpose fats are boiled with 
 a solution of potassium hydroxide or carbonate. 
 
 Potassium Oxide, K 2 O. This compound can be made 
 by burning potassium in the air, and heating the residue 
 to a high temperature. It is also formed by melting 
 potassium hydroxide and metallic potassium together : 
 
 2K + 2KOH = 2K 2 O + H a . 
 
 With water it forms the hydroxide, with a marked evo- 
 lution of heat : 
 
 K 2 + H 2 = 2KOH. 
 
 Potassium also forms other oxides of which the peroxidt 
 of the formula K 2 O 4 is the best studied. This peroxide 
 is the final product of the combustion of potassium in 
 the air or in oxygen. At a high temperature it breaks 
 down into potassium oxide, K 2 O, and oxygen. It also 
 gives up its oxygen very readily to substances which are 
 
490 INORGANIC CHEMISTRY. 
 
 capable of oxidation, acting so energetically upon some 
 as to cause evolution of light. 
 
 Potassium Hydrosulphide, KSH, is analogous to potas- 
 sium hydroxide. Just as the latter is made by the action 
 of potassium on water, so the former can be made by the 
 action of potassium on hydrogen sulphide : 
 
 K 2 + 2H 2 S = 2KSH + H 2 . 
 
 It is, however, obtained most readily by the action of 
 hydrogen sulphide on a solution of potassium hydroxide : 
 
 KOH + H 2 S = KSH + H 2 0. 
 
 When exposed to the action of the air it is oxidized, and 
 becomes colored in consequence of the formation of the 
 disulphide. The action takes place as represented thus : 
 
 2KSH + O = K 2 S 2 + H a O. 
 
 Potassium Sulphide, K 2 S, is made by the reduction of 
 potassium sulphate either by means of hydrogen or car- 
 bon. It is thought by some to be present in a solution 
 prepared by saturating a given quantity of potassium 
 hydroxide with hydrogen sulphide, and then adding the 
 same quantity of potassium hydroxide to the product. 
 The formation is supposed to take place as represented 
 in the equations 
 
 KSH + KOH = K 2 S + H 2 0. 
 
 This is the action which we should expect, as hydrogen 
 sulphide acts like an acid, and with a strong base we 
 should expect it to form two salts, the acid salt KSH 
 and the neutral salt K 2 S. From thermo-chemical investi- 
 gations of this subject, however, the conclusion appears 
 to be justified that the salt K 2 S does not exist in solution, 
 but that it breaks down with water, forming the hydro- 
 sulphide and hydroxide : 
 
 K 2 S + H 2 O = KSH + KOH. 
 
 This reaction is analogous to that of water on the oxide . 
 K,0 + H,O = 2KOH. 
 
POLYSULPHIDES OF POTASSIUM. 491 
 
 The poly sulphides of potassium are compounds having 
 the composition expressed by the formulas K 2 S 2 , K 2 S 3 , 
 K 2 S 4 , and K 2 S 5 . They are formed in general by the 
 action of sulphur on a solution of the hydrosulphide or 
 of the simple sulphide. The disulphide is also formed 
 as explained above by oxidation, when a solution of the 
 hydrosulphide is allowed to stand exposed to the air. 
 They are all colored substances, which readily give up 
 sulphur. If treated with dilute acids each one gives up 
 sufficient sulphur to reduce it to the simple form K 2 S. 
 If the air is allowed to act upon them for a sufficient 
 length of time they all yield the thiosulphate, K 2 S 2 O 3 , 
 the action taking place as represented in the following 
 equations : 
 
 K,S, + 30 = K,S A + S ; 
 K,S 4 + 3O = K,SA + 2S; 
 K A + 30 = K.S A + 38. 
 
 The fact that no higher sulphide of potassium than 
 the pentasulphide exists, suggests that the action of 
 sulphur upon the monosulphide is analogous to that of 
 oxygen, and that the pentasulphide is analogous to the 
 sulphate : 
 
 K 2 S + 40 = K 2 S0 4 ; 
 
 K,S + 4S = K 2 SS 4 , or K 2 S 6 . 
 
 According to this view, the pentasulphide is the salt of 
 
 (mr\ 
 S 2 S<Sjj), or hydrogen 
 
 pentasulphide, H 2 S 5 . 
 
 The substance used in medicine under the name of 
 liver of sulphur or Hepar sulfuris is a brown mass 
 formed by melting together potassium carbonate and 
 sulphur, and consisting of polysulphides of potassium 
 and potassium thiosulphate and sulphate. The chief re- 
 action involved is the one represented in the equation 
 
 3K 2 C0 3 + 8S = 2K 2 S 3 + K 2 S 2 O 3 + 3CO 
 
 23 a . 
 
492 INORGANIC CHEMISTRY. 
 
 If the mass is ignited the thiosulphate is decomposed, 
 forming the sulphate and pentasulphide : 
 
 4K,S,0. = 3K0 4 + K 2 S 8 . 
 
 Potassium sulphide combines with the sulphides of 
 arsenic, antimony, and tin, forming salts of sulpho-acids. 
 Among the best known of these are the following : 
 
 Potassium Sulpharsenate, K 3 AsS 4 -|- H 2 O. This is 
 formed by treating arsenic pentasulphide or the trisul- 
 phide and sulphur with potassium hydrosulphide : 
 
 6KSH + As 2 S 5 = 2K 3 AsS 4 + 3H 2 S ; 
 6KSH + As 2 S 3 + 28 = 2K 3 AsS 4 + 3H 2 S. 
 
 It is also formed by saturating a solution of potassium 
 arsenate with hydrogen sulphide. 
 
 Potassium sulphantimonate, K 3 SbS 4 + 4^H 2 O, and potas- 
 sium sulpharsenite, KAsS 2 -\- 2^H 2 O, are also easily made. 
 The latter is plainly the analogue of the metarsenite, 
 KAsO a metarsenious acid being derived from normal 
 arsenious acid by elimination of one molecule of water : 
 
 H 3 AsO 3 - HAs0 2 + H 2 O. 
 
 Potassium Nitrate, KNO 3 . This salt is commonly called 
 saltpeter. Its occurrence in nature has already been 
 spoken of under Nitric Acid (which see). When refuse 
 animal matter is left to undergo decomposition in the 
 presence of bases nitrates are always the end-products. 
 They are consequently found very widely distributed in 
 the soil. In the East Indies the potassium nitrate 
 formed in the neighborhood of dwellings and stables is 
 collected, and sent into the market. The process of nit- 
 rification is carried on artificially on the large scale in 
 the so-called " saltpeter plantations." In these, refuse 
 animal matter is mixed with earthy material, wood 
 ashes, etc., and piled up. These piles are moistened 
 with the liquid products from stables. After the action 
 has continued for two or three years the outer crust is 
 taken off, and extracted with water. The solution thus 
 obtained contains, besides potassium nitrate, calcium and 
 magnesium nitrates. It is treated with a water-extract 
 
GUNPOWDER. 493 
 
 of wood ashes or with potassium carbonate, by which the 
 calcium and magnesium are thrown down as carbonates. 
 Much of the saltpeter which is now in the market is 
 made from Chili saltpeter, or sodium nitrate, by treating 
 it with potassium chloride, advantage being taken of the 
 fact that sodium chloride is less soluble in water than 
 potassium nitrate. Molecular weights of sodium nitrate 
 and potassium chloride are dissolved in water and the 
 solution evaporated, when sodium chloride is deposited : 
 
 NaNO 3 + KC1 = KNO 3 + NaOl. 
 
 Potassium nitrate crystallizes in long rhombic prisms, 
 of a salty taste. Under some circumstances it crystal- 
 lizes in rhombohedrons. When dissolved in water it 
 causes a lowering of the temperature. At ordinary tem- 
 peratures 100 parts of water dissolve from 20 to 30 parts 
 of the salt ; at 100, 100 parts dissolve 247 parts. 
 
 Applications. Potassium nitrate is used as an oxidiz- 
 ing agent in the laboratory, and in the manufacture of 
 fireworks. Its chief use, however, is in the manufac- 
 ture of gunpowder. 
 
 Gunpowder. The value of gunpowder is due to the 
 iact that it explodes readily, the explosion being a chemi- 
 cal change accompanied by a sudden evolution of gases. 
 It is a mixture of saltpeter, charcoal, and sulphur. 
 When heated, the saltpeter gives off oxygen and nitro- 
 gen ; the oxygen combines with the charcoal, forming 
 carbon dioxide and carbon monoxide ; and the sulphur 
 combines with the potassium, forming potassium sul- 
 phide. When a mixture of saltpeter and charcoal is 
 burned, the reaction which takes place is this : 
 
 2KNO 3 + 30 = C0 2 + CO + N 2 + K 2 CO 3 . 
 
 By adding the necessary quantity of sulphur the car- 
 bon dioxide, which would otherwise remain in combina- 
 tion with the potassium as potassium carbonate, is giveR 
 off, and potassium sulphide formed : 
 
 2KNO 3 + 30 + S = 3C0 2 + N a + K 2 S. 
 
494 INORGANIC CHEMISTRY. 
 
 For this reaction the constituents should be mixed in 
 the proportions : 
 
 Saltpeter, 74.83 
 
 Charcoal, . . . . 13.31 
 
 Sulphur, 11.86 
 
 100.00 
 
 This is approximately the composition of all powder. 
 When gunpowder explodes, the gases formed occupy 
 about 280 times the volume occupied by the powder 
 itself. 
 
 Potassium Nitrite, KNO 2 , is formed simply by heating 
 the nitrate to a sufficiently high temperature. The re- 
 duction is, however, much facilitated by adding to the 
 nitrate some easily oxidized metal, as lead or iron. 
 When the gases formed by the action of arsenic trioxide 
 on nitric acid are passed into potassium hydroxide, both 
 the nitrite and nitrate are formed, and they can be sep- 
 arated by crystallization. 
 
 Potassium Chlorate, KC1O 3 . The character of the reac- 
 tion by which potassium chlorate is formed when chlo- 
 rine acts upon a solution of potassium hydroxide has 
 already been discussed (see p. 114). In the manufac- 
 ture of the chlorate it is found advantageous to make 
 calcium chlorate, and then to treat this with potassium 
 chloride, when, at the proper concentration, potassium 
 chlorate crystallizes out, on account of the fact that it is 
 less soluble than the salts which are brought together. 
 The process in brief consists in passing chlorine into a 
 solution of calcium hydroxide in which an excess of hy- 
 droxide is held in suspension. The first action consists 
 in the formation of calcium hypochlorite. When the 
 solution of this salt is boiled it is decomposed, yielding 
 the chlorate and chloride : 
 
 3Ca(OCl) 2 = Ca(O 3 Cl) 2 + 2CaCl 2 . 
 
 On now treating the solution with potassium chloride 
 the following reaction takes place : 
 
 Ca(O 3 Cl) 3 + 2KC1 = 2KC10 3 + CaCl 
 
POTASSIUM PERCHLOBATE. 495 
 
 Potassium chlorate crystallizes in lustrous crystals of 
 tlie monoclinic system. Its taste is somewhat like that 
 of saltpeter. It melts at a comparatively low tempera- 
 ture (334), and at 352 begins to decompose, with evolu- 
 tion of oxygen. At ordinary temperatures 100 parts of 
 water dissolve 6 parts of the salt, and at the boiling tem- 
 perature 60 parts. In consequence of the ease with 
 which it gives up its oxygen, the chlorate is an excellent 
 oxidizing agent, and it is constantly used in this capacity 
 in the laboratory. Its oxidizing action is well illus- 
 trated by grinding a very little of it in a mortar with a 
 little sulphur,* when an explosion takes place. With 
 phosphorus the action is exceedingly violent. 
 
 The chief uses of potassium chlorate are for the prep- 
 aration of oxygen, and in the manufacture of matches 
 and fireworks. The tips of Swedish safety matches are 
 made of potassium chlorate and antimony sulphide. 
 The surface upon which they are rubbed to ignite them 
 contains red phosphorus. The chlorate is frequently 
 used in medicine, particularly as a gargle in cases of 
 sore throat. 
 
 Potassium Perchlorate, KC1O 4 , is formed in the first 
 stage of the decomposition of the chlorate by heat, as was 
 explained under Oxygen (which see). It is prepared 
 best by heating the chlorate in an open vessel until, af- 
 ter having been liquid, it begins to get solid again. As 
 the salt is difficultly soluble in water, the residue is pow- 
 dered and washed with water to remove the chloride, 
 and then crystallized from water. Owing to the difficult 
 solubility of this salt it is utilized in chemical analysis 
 for detecting the presence of potassium. For this pur- 
 pose a solution of perchloric acid is added to the solu- 
 tion under examination, and, if a precipitate is formed, 
 the presence of potassium may be inferred. When 
 heated to about 400 the perchlorate gives up its oxygen, 
 and is reduced to the chloride. It is used to some ex- 
 tent in the manufacture of fireworks, instead of the chlo- 
 
 * Great care should be taken with all experiments with potassium 
 chlorate. See description of experiments. 
 
496 INORGANIC CHEMISTRY. 
 
 rate, which, owing to its greater instability, is more 
 dangerous. 
 
 Potassium Periodate, KIO 4 , is formed by the action of 
 chlorine on a mixture of potassium hydroxide and po- 
 tassium iodate. As has been stated in discussing 
 the acids of iodine, this salt is only one form of a 
 group of potassium salts called periodates, all of 
 which are closely related to the normal acid, I(OH) 7 . 
 Among these salts, for example, are the mesoperiodate, 
 K 3 IO 6 + 4H 2 O, and the diperiodate, K 4 I 2 O 9 + 9H 2 O. The 
 former is a salt of the acid H 3 IO 6 , which is derived from 
 the normal acid by loss of water, thus : 
 
 Diperiodic acid is derived from the normal acid as 
 represented in this equation : 
 
 2I(OH), = H.I.O. + 6H.O. 
 
 Potassium Cyanide, KCN. Under Cyanogen it was 
 stated that when nitrogen is passed over a highly heated 
 mixture of carbon and potassium carbonate, potassium 
 cyanide is formed ; and that carbon containing nitrogen 
 compounds, as animal charcoal, when ignited with potas- 
 'sium carbonate, reduces the carbonate, forming potas- 
 sium, in presence of which the carbon and nitrogen 
 combine, forming the cyanide. The simplest way to 
 make the cyanide is by heating potassium ferrocyanide, 
 K 4 Fe(CN) 6 , which is the starting-point in the prepara- 
 tion of all cyanogen compounds. It breaks down first 
 into potassium cyanide and ferrous cyanide, thus : 
 
 K 4 Fe(CN) 6 = 4KCN + Fe(CN) 2 . 
 
 The ferrous cyanide is, however, decomposed by heat 
 into free nitrogen and carbide of iron, 
 
 so that the complete decomposition of the ferrocyanide 
 is represented by this equation : 
 
 K 4 Fe(CN) 6 = 4KCN + FeC 2 + N 3 . 
 
POTASSIUM CYANIDE. 497 
 
 As part of the cyanogen is lost in this operation, potas- 
 sium carbonate is commonly added to the ferrocyanide. 
 This acts upon the ferrous cyanide, forming potassium 
 cyanide and ferrous carbonate : 
 
 Ee(CN) 2 + K 2 C0 3 = FeCO 3 + 2KCN. 
 
 But the ferrous carbonate breaks down under the influ- 
 ence of heat into ferrous oxide and carbon dioxide 
 
 FeCO 3 = FeO + CO 2 ; 
 
 and the ferrous oxide then gives up its oxygen to a part 
 of the potassium cyanide, converting it into the cyanate. 
 The complete reaction between the ferrocyanide and car- 
 bonate is therefore represented as follows : 
 
 K 4 Fe(CN) 6 + K 2 C0 3 = 5KCN + KCNO + CO 2 + Fe. 
 
 The product obtained in this way necessarily contains 
 some of the cyanate, but for ordinary purposes this does 
 no harm. Potassium cyanide is extremely easily soluble 
 in water, and is deliquescent in moist air. When boiled 
 with water it is decomposed, forming potassium formate 
 and ammonia : 
 
 KCN + 2H 2 O = KC0 2 H + NH 3 . 
 
 It combines readily with oxygen when in the molten con- 
 dition, as shown in its action upon ferrous oxide and 
 upon lead oxide (see p. 404). In consequence of this 
 power to combine with oxygen to form the cyanate, it is 
 a valuable reducing agent, and is not unfrequently used 
 in the laboratory in this capacity. 
 
 Just as the fluoride, chloride, bromide, and iodide of 
 potassium combine with the fluorides, chlorides, bro- 
 mides, and iodides of the metallic elements in general, so 
 potassium cyanide combines with the cyanides of the 
 metallic elements, forming the double cyanides. These 
 have generally a composition analogous to that of the 
 double chlorides and of similar compounds. Thus, silver 
 cyanide forms the compound AgCy.KCy or AgKCy 2 , in 
 which the cyanogen group CN is represented by the 
 symbol Cy, as is customary ; ferrous cyanide forms the 
 
498 INORGANIC CHEMISTRY. 
 
 compounds K 4 FeCy 6 , or 4KCy.FeCy 2 , and K 3 FeCy 6 , or 
 3KCy.FeCy 3 . These double cyanides are for the most 
 part soluble in water ; hence potassium cyanide dissolves 
 many deposits of metallic salts. It is frequently used 
 in the laboratory in analytical operations. 
 
 Potassium Cyanate, KCNO. The cyanate is formed by 
 oxidation of the cyanide, when this is melted and lead 
 oxide or minium added to the molten mass. It is most 
 easily prepared by heating together potassium ferrocya- 
 nide and manganese dioxide. The action consists first 
 in the decomposition of the ferrocyanide, and the subse- 
 quent oxidation of the potassium cyanide thus formed. 
 The mass is extracted with alcohol, as the cyanate is de- 
 composed by water even at the ordinary temperature, 
 the products being potassium carbonate and ammonium 
 carbonate : 
 
 2KCNO + 4H,O = K 2 C0 3 + (NH 4 ) 2 CO 3 . 
 
 Acids set cyanic acid free from the cyanate, but the 
 acid is at once decomposed by water, thus : 
 
 CNOH + H 2 O = CO 2 -f NH 3 . 
 
 Potassium Sulphocyanate, KCNS. Just as potassium 
 cyanide takes up oxygen to form the cyanate, it also 
 takes up sulphur to form the sulphocyanate : 
 
 KCN + S = KCNS. 
 
 It is easily prepared by adding sulphur to molten 
 potassium cyanide, or by heating a mixture of dehy- 
 drated potassium ferrocyanide, potassium carbonate, 
 and sulphur. It crystallizes particularly well out of its 
 solution in alcohol. It is deliquescent, and when dis- 
 solved in water it causes a very considerable lowering 
 of the temperature. Thus, when 500 grams of the salt 
 are mixed with 400 grams of water at the ordinary tem- 
 perature, the temperature sinks to about 20. Unlike 
 the cyanate, it is not decomposed by water. 
 
 Potassium Sulphate, K 2 SO 4 . This salt occurs in com- 
 bination with others in nature, particularly in the mineral 
 
PRIMARY, OR ACID, POTASSIUM SULPHATE. 499 
 
 kainite, which contains the constituents of potassium 
 sulphate, magnesium sulphate, and magnesium chloride, 
 as indicated in the formula K 2 SO 4 .MgSO 4 .MgCl 2 + 6H 2 O. 
 This occurs in Stassfurt and in Kalusz. Potassium sul- 
 phate is used in medicine, and in the preparation of 
 ordinary alum and of potassium carbonate. 
 
 Primary, or Acid, Potassium Sulphate, KHSO 4 . This 
 salt is obtained as a secondary product in the prepara- 
 tion of nitric acid by the action of sulphuric acid upon 
 saltpeter. It occurs in nature in the Grotto del Solfo, 
 near Naples ; and is made by treating the neutral salt 
 with concentrated sulphuric acid. When heated above 
 its melting point it gives off water, and is transformed 
 into the disulphate, thus : 
 
 2KHS0 4 = K 2 S 2 7 + H 2 0. 
 
 When the disulphate is heated in contact with basic 
 oxides it breaks down, forming sulphates. The decom- 
 position is that represented in the equation 
 
 K 2 S 2 7 = K 2 S0 4 + SO,. 
 
 The nascent sulphur trioxide thus set free acts with great 
 energy upon the oxides which are present. Hence acid 
 potassium sulphate is a valuable agent for the purpose 
 of decomposing some mineral substances which do not 
 readily yield to the ordinary reagents. Its action consists 
 in breaking down into the disulphate and water, the disul- 
 phate then further breaking down into normal sulphate 
 and sulphur trioxide. Besides the salt just mentioned, 
 which is known as the disulphate or pyrosulphate, there 
 are some other salts known, which are derived from an- 
 other form of disulphuric acid. Two good examples 
 are the salts represented by the formulas K 3 H(SO 4 ) 2 and 
 KH 3 (SO 4 ) 2 . The acid from which these are derived has 
 the formula H 4 S 2 O 8 . It is to be regarded as derived 
 from two molecules of normal sulphuric acid by elimina- 
 tion of four molecules of water : 
 
500 INORGANIC CHEMISTRY. 
 
 The acid probably has the constitution represented thus : 
 
 (HO) s OS<g>SO(OH), = H&0 r 
 
 Sulphites. When sulphur dioxide is passed into a 
 solution of potassium carbonate until carbon dioxide 
 ceases to escape, potassium sulphite, K 2 SO 3 , is formed. If 
 the gas is passed to saturation the product is the pri- 
 mary or acid sulphite, KHSO 3 . If the solution of the 
 carbonate is hot and concentrated, the product is the 
 disulphite, K 2 S 2 O 6 , which bears to the sulphite the same 
 relation that potassium disulphate bears to the sulphate. 
 It is the salt of an acid of the formula H 2 S 2 O 5 , which is 
 disulphurous acid : 
 
 This bears to sulphurous acid the same relation that di- 
 sulphuric acid bears to sulphuric acid. 
 
 Carbonates. The normal salt, K 2 CO 3 , is the chief con- 
 stituent of wood-ashes. When these are extracted with 
 water the carbonate passes into solution and the salt 
 thus obtained can be purified in a number of ways. The 
 impure salt is known as potash. Formerly all the potas- 
 sium carbonate made was obtained from wood-ashes, but 
 at present not more than half of the supply comes from 
 this source. The other sources are the residues from 
 the manufacture of beet-sugar, potassium sulphate and 
 chloride, and wool-fat. The preparation of the carbon- 
 ate from the sulphate and chloride is accomplished by 
 the same method as that used in the preparation of 
 sodium carbonate from the chloride. The methods used 
 for this purpose will be treated of under the head of 
 Sodium Carbonate (which see). The salt crystallizes 
 from very concentrated solutions in water. It is deli- 
 quescent, and dissolves in water with an evolution of 
 heat, and the solution has a strong alkaline reaction. 
 
 Acid Potassium Carbonate, HKCO 3 , is formed by pass- 
 ing carbon dioxide over the normal salt, or into the con- 
 centrated aqueous solution of the latter. It is much less 
 easily soluble in water than the normal salt. The dry 
 
RUBIDIUM CESIUM. 501 
 
 salt gives off carbon dioxide and water easily when 
 heatedj and is converted into the normal salt : 
 
 2KHC0 3 = K 2 CO 3 + CO 2 + H 2 O. 
 
 The same decomposition takes place when the water 
 solution is heated, and even on evaporation at the ordi- 
 nary temperature. 
 
 Phosphates. Three phosphates of potassium are 
 known : (1) Tertiary, or normal potassium phosphate, 
 K 3 PO 4 ; (2) secondary, or di-potassium phosphate, K 2 HPO 4 ; 
 and (3) primary, or mono-potassium phosphate, KH 2 PO 4 . 
 There is nothing particularly characteristic about these 
 salts, except the decompositions which the primary and 
 secondary salts undergo when heated. These decompo- 
 sitions have already been referred to (see p. 329 and p. 
 480). 
 
 Potassium Silicate, K 2 SiO 3 . A compound of the definite 
 composition represented by the formula here given has 
 not been prepared. A solution of potassium silicate in 
 water is prepared by dissolving sand or amorphous sili- 
 con dioxide in potassium carbonate or hydroxide. It k 
 prepared on the large scale by melting together quartz 
 powder and purified potash. It is known as water glass, 
 for the reason that its solution dries in the air, forming 
 a glass-like looking mass. To distinguish it from the 
 water glass made with sodium carbonate or hydroxide 
 it is called potash water glass. 
 
 BUBIDIUM, Eb (At. Wt. 84.78). 
 OESIUM, Cs (At. Wt. 131.89). 
 
 Both these elements are widely distributed, but only 
 in small quantities. They generally occur in company 
 with potassium, which they resemble closely. They 
 were discovered by means of the spectroscope by Bun- 
 sen and Kirchhoff. The characteristic spectrum of 
 rubidium consists of two dark red lines, and this is the 
 origin of the name rubidium (from rubidus, dark red). 
 Csesium was found in the Durkheim mineral water, and 
 was recognized by two characteristic blue lines, and the 
 name caesium was given to it on this account (from ccesius, 
 
502 INORGANIC CHEMISTRY. 
 
 sky-blue). Rubidium is found in different varieties of 
 mica, known as lepidolite. The mineral pollux, which is 
 essentially a silicate of caesium and aluminium, contains 
 caesium as one of the chief constituents. 
 
 It is a remarkable fact that the elements rubidium 
 and caesium which are so similar to potassium accom- 
 pany it so generally in nature. Similar facts were noted 
 in the group consisting of chlorine, bromine, and iodine, 
 and that of sulphur, selenium, and tellurium. It will 
 be remembered that chlorine is frequently accompanied 
 by bromine and iodine ; and sulphur by selenium and 
 tellurium ; but that chlorine and sulphur are present in 
 much larger quantities than the elements which accom- 
 pany them. Further, the relations between the atomic 
 weights of the members of each group are approximately 
 the same. 
 
 Rubidium is prepared by the same method as that 
 used in the preparation of potassium. 
 
 'It is silver-white with a yellowish tint. It can be con- 
 verted into vapor which has a blue color. It takes fire 
 in the air at the ordinary temperature. Its action upon 
 water is the same as that of potassium, and its salts are 
 very similar to those of potassium. 
 
 Caesium has not yet been isolated. By subjecting the 
 chloride to the action of a powerful electric current 
 globules of metal are given off at one of the poles, but 
 these take fire in contact with the air at the ordinary 
 temperature. The salts of caesium are much like those 
 of rubidium and potassium. 
 
 SODIUM, Na (At. Wt. 22.82). 
 
 Occurrence. Sodium occurs very widely distributed 
 and in large quantities in nature, principally as sodium 
 chloride. It is found in a number of silicates, and is a 
 constituent of plants, especially of those which grow in 
 the neighborhood of the sea-shore and in the sea. Just 
 as the ashes of inland plants are rich in potassium car- 
 bonate, so the ashes of sea plants and those which grow 
 near the sea are rich in sodium carbonate. It is found 
 evervwhere in the soil, but generally in small quantities. 
 
PREPARATION OF SODIUM. 503 
 
 Its presence in the soil is due to the decomposition of 
 minerals containing it, such as soda feldspar, or albite. It 
 occurs also as sodium nitrate or Chili saltpeter, and in 
 large quantity in Greenland in the form of cryolite, which, 
 as has been explained, is a so-called double fluoride of 
 aluminium and sodium, of the formula Na 3 AlF 6 , or 
 AlF 3 .3NaF. 
 
 Preparation. It is prepared from sodium carbonate by 
 the same method as that used in the preparation of potas- 
 sium, the reaction involved being represented thus ; 
 
 Na 2 C0 3 + 20 = 2Na + SCO. 
 
 The reduction takes place more readily than in the case 
 of potassium, and it is not necessary to prepare the mix- 
 ture of carbonate and charcoal by heating the salt of an 
 organic acid, as is done in the preparation of potassium. 
 The carbonate is mixed with charcoal, or powdered an- 
 thracite coal, and calcium carbonate, and sometimes this 
 mass is mixed with an oil and then ignited in a crucible. 
 A successful method for the preparation of sodium 
 on the large scale has been devised by Castner. This 
 consists essentially in the reduction of sodium hydroxide 
 by heating it with an intimate mixture of finely divided 
 iron and carbon. The mass is prepared by mixing the 
 iron with molten pitch, allowing it to cool, breaking it 
 into pieces, and heating to a comparatively high temper- 
 ature without access of air. The reduction is said to 
 take place at a temperature of 825, instead of 1400 as 
 in the older method. The main reaction is represented 
 by this equation : 
 
 GNaOH + FeC 2 = 2Na 2 CO 3 + 6H + 2Na + Fe. 
 
 The preparation of sodium was formerly of more 
 importance than it is at present, for the separation of 
 aluminium from the compounds found in nature depended 
 upon the preparation of sodium. Methods, depending 
 upon the use of an electric furnace, have been devised 
 for the preparation of aluminium, and the old method 
 involving the use of sodium is no longer employed. 
 
504 INORGANIC CHEMISTRY. 
 
 Properties. The properties of sodium are very similar 
 to those of potassium. It is light, floating on water ; it 
 has a bright metallic lustre ; and at the ordinary tem- 
 perature it is soft like wax. It decomposes water, but 
 not as readily as potassium does. Its specific gravity is 
 0.9735; its melting point 95.6. Its vapor is colorless 
 when seen in thin layers, while thick layers appear pur- 
 ple. When melted and allowed to cool it takes the 
 crystallized form. When exposed to the air it acts upon 
 the moisture, and is converted into the hydroxide. 
 
 Applications. It is used for the purpose of isolating 
 some elements whose oxides cannot easily be reduced, 
 as, for example, aluminium, magnesium, and silicon, 
 which are prepared by treating their chlorides with 
 sodium. Silicon, however, as we have seen, is prepared 
 better by treating potassium fluosilicate, K 2 SiF 6 , with 
 sodium. The element is also used, in combination with 
 mercury as sodium amalgam, a substance which affords 
 a ready means of making nascent hydrogen. It also 
 finds constant application in the laboratory for a variety 
 of purposes. 
 
 Sodium Hydride, Na 2 H, is formed in the same way as 
 the corresponding compound of potassium, and is in 
 every way similar to it. 
 
 Sodium Chloride, NaCl. This is the substance which 
 is generally known simply as salt, or common salt. It 
 occurs very widely distributed, and in immense quantities, 
 in the earth. The most important deposits are those at 
 Wieliczka in Galicia, at Stassfurt and Keichenhall in Ger- 
 many, and at Cheshire in England. Besides these there 
 are, however, many other deposits in the United States 
 of America, in Africa, and in Asia. As it is easily soluble 
 in water, many springs and streams, as well as lakes and 
 the ocean, contain it. Sea- water contains 2.7 per cent. 
 In some places sodium chloride is taken out of mines in 
 solid form. Frequently, however, water is allowed to 
 flow into cavities in the earth, and to remain for some 
 time in contact with the salt. The solution thus formed 
 is afterward drawn or pumped out of the mine and 
 evaporated by appropriate methods. It is generally 
 
SODIUM CHLORIDE. 505 
 
 allowed slowly to run down walls made of twigs, so that 
 a large surface of the liquid is exposed to the air. The 
 concentrated solution thus obtained is then evaporated 
 to crystallization by the aid of heat. 
 
 In hot countries salt is obtained by the evaporation of 
 sea- water, the heat of the sun being used for the purpose. 
 Large shallow cavities are made in the earth, and into 
 these the water flows at high-tide, or it is pumped up into 
 them if they are too high. The process is continued for 
 some months, and then the mother-liquor is drawn off, 
 and the accumulated salt collected and subjected to 
 proper methods of purification. 
 
 The salt obtained by the above methods is not pure. 
 It always contains sodium sulphate, together with mag- 
 nesium and calcium chlorides. The chlorides of magnes- 
 ium and calcium cause it to become moist in the air. 
 Pure salt does not attract moisture. 
 
 Sodium chloride crystallizes in colorless and trans- 
 parent cubes. Some of that which occurs in nature has 
 a blue color. When deposited from an evaporating 
 solution it takes the form of small cubes arranged in 
 groups of the shape of hollow pyramids, known as the 
 hopper-shaped deposits. If urea or boric acid is present 
 in the solution the crystals of sodium chloride are octa- 
 hedrons or combinations of these with cubes. When 
 deposited, the crystals enclose water, not as water of 
 crystallization, and this is given off when the crystals 
 are heated, the action being accompanied by a crackling 
 sound. This is known as decrepitation. 
 
 Sodium chloride melts at 776, and is volatile at a red 
 heat. In hot water it is but little more soluble than in 
 cold. At 100 100 parts of water dissolve 39 parts, and 
 at ordinary temperatures 36 parts. 
 
 Sodium chloride is the starting-point in the preparation 
 of all sodium compounds, as well as of chlorine and hy- 
 drochloric acid. Salt is necessary to the life of man 
 and many other animals. The role played by it in the 
 animal economy is not understood, but it is found 
 generally distributed throughout the body in small 
 quantity. 
 
506 INORGANIC CHEMISTRY. 
 
 The fluoride, bromide, and iodide of sodium are like the 
 corresponding potassium salts and need not be described. 
 
 Sodium Hydroxide, NaOH. This compound resembles 
 potassium hydroxide in all respects. Being cheaper it is 
 used much more extensively. It is prepared in the same 
 way, by treating sodium carbonate in solution with cal- 
 cium hydroxide, when insoluble calcium carbonate and 
 soluble sodium hydroxide are formed : 
 
 Na 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2NaOH. 
 
 The substance is commonly called caustic soda. It is 
 extensively used for the purpose of making soap from 
 fats. 
 
 Oxides. Sodium forms two oxides, the monoxide, 
 Na 2 O, and the peroxide, Na 2 O 2 . In this respect a differ- 
 ence is noticed between sodium and potassium ; the latter 
 forming the compounds K 2 O and K 2 O 4 . 
 
 Sodium Peroxide, Na 2 O 2 , has acquired importance in 
 the arts as a bleaching-agent. It is prepared by heating 
 sodium in a current of dry air at a temperature of 300. 
 When heated to a high temperature it gives off oxygen. 
 Water decomposes^ it, forming sodium hydroxide, and 
 setting oxygen free. 
 
 The hydrosulphide and the sulphides of sodium are made 
 just as the potassium compounds are, and resemble them 
 very closely. 
 
 Sodium Sulphantimonate, Na s SbS 4 , also known as 
 Schlippe's salt, is a particularly beautiful example of the 
 salts of sulpho-acids. It is made, as its composition indi- 
 cates, by dissolving antimony pentasulphide in a solution 
 of sodium sulphide : 
 
 Sb 2 S 6 + 3Na 2 S = 2Na 3 SbS 4 . 
 
 Sodium Nitrate, NaNO 3 . This compound occurs in 
 large quantity in southern Peru on the border of Chili, 
 and is known as Chili saltpeter. The natural salt con- 
 tains, besides the nitrate, sodium chloride, sulphate, and 
 iodate. Sodium nitrate is very similar to potassium 
 nitrate, but it cannot be used in place of the more ex- 
 pensive potassium salt in the manufacture of the finer 
 
SODIUM SULPHATE. 507 
 
 grades of gunpowder, as it becomes moist in the air, and 
 does not decompose quickly enough. It is used ex- 
 tensively in the manufacture of nitric acid, and also for 
 the purpose of preparing ordinary saltpeter. The iodine 
 contained in Chili saltpeter is now extracted on the large 
 scale, and this forms an important source of iodine. 
 
 Sodium Sulphate, Na 2 SO 4 . This salt was first made by 
 Glauber, as it is now made, by the action of sulphuric 
 acid on sodium chloride. It is commonly called Glauber's 
 salt. It occurs in a number of natural waters, as in that 
 of Friedrichshall and Carlsbad. It occurs, further, in 
 solid form in small quantities in some localities. It is 
 made in very large quantities in connection with the 
 manufacture of soda, the first reaction in this process con- 
 sisting in treating sodium chloride with sulphuric acid. 
 It is also formed in the manufacture of nitric acid by 
 the action of sulphuric acid on Chili saltpeter. 
 
 Large quantities of sodium sulphate are now made by 
 the action of magnesium sulphate on sodium chloride. 
 This process is employed at Stassfurt, where both mag- 
 nesium sulphate and sodium chloride occur in immense 
 quantities. The action takes place between concentrated 
 solutions at low temperatures. It is represented by the 
 equation 
 
 2NaCl + MgS0 4 = Na 2 SO 4 + MgCl 2 . 
 
 It crystallizes in large, colorless, monoclinic prisms, 
 which contain ten molecules of water. These crystals 
 are formed, however, only in case the temperature of the 
 solution is below 33 at the time they are deposited. If 
 a saturated solution is cooled down to a point some- 
 where between 33 and 40, the salt is deposited without 
 water of crystallization. When the crystallized salt is 
 heated to 33 it loses a part of its water. The salt is most 
 easily soluble in water at 33 ; above this point the solu- 
 bility decreases. Taking these facts into consideration, 
 it appears probable that in solutions below 33 the com- 
 pound Na 2 SO 4 -f- 10H 2 O is present ; while if the solu- 
 tion is heated above this point the compound breaks 
 
508 INORGANIC CHEMISTRY. 
 
 down, and the anhydrous salt, as well as the salts with 
 less than ten molecules of water, are less easily soluble. 
 One of the ten molecules of water is held in the com- 
 pound more firmly than the rest. It seems probable 
 that this is not present as water but as hydroxyl, the salt 
 
 having the formula OS j ^\ (=Na 2 SO 4 + H 2 O). 
 
 Sodium sulphate easily forms supersaturated solutions 
 which crystallize rapidly if disturbed, if a small crystal 
 of the salt is thrown into them, and if cooled down 
 to 8. This phenomenon is frequently presented by 
 salts, but it is shown in a particularly striking way by 
 this one. 
 
 When expose^d to the air the salt loses its water of 
 crystallization and crumbles to a white powder. This is 
 the process already described as efflorescence (see p. 58). 
 
 Sodium sulphate is used as a purgative in medicine, 
 and in the laboratory for the production of cold arti- 
 ficially. A good freezing mixture is made by bringing it 
 together with concentrated hydrochloric acid. Sodium 
 chloride is formed, and the water of crystallization of the 
 sulphate takes the liquid form. This change from the 
 solid to the liquid form is accompanied by a marked ab- 
 sorption of heat. Ice can be made in this way without 
 difficulty. The chief uses of the sulphate are in 'the 
 manufacture of sodium carbonate and of glass, as will 
 be explained farther on. 
 
 Sodium Thiosulphate, Na 2 S 2 O 3 + 5H 2 O. This is the salt 
 which is commonly called hyposulphite of soda. It is 
 made on the large scale by treating caustic soda with 
 sulphur, and conducting sulphur dioxide into the solu- 
 tion. As has been pointed out, when sulphur acts upon 
 potassium carbonate polysulphides of potassium and the 
 thiosulphate are formed. A similar action takes place 
 when sulphur acts upon caustic soda. The polysul- 
 phides in the solution give up sulphur to the sulphite 
 and convert it thus into the thiosulphate : 
 
 Na,S, + Na 2 SO 3 = Na 2 S + Na 2 S 8 O 3 . 
 
SODIUM CARBONATE. 509 
 
 It is also made by boiling a solution of sodium sulphite 
 and adding sulphur : 
 
 Na a SO 3 + S = Na 2 S 2 O 3 . 
 
 Its chief application is in photography, in which art it is 
 used for the purpose of dissolving the excess of silver 
 salt on the plate which has been exposed to the light, 
 and on which a picture has been developed. The action 
 consists in the formation of salts in which both sodium 
 and silver are contained. These are soluble in water. 
 The thiosulphate will be taken up more in detail under 
 Silver (which see). 
 
 Sodium Carbonate, Na 2 CO 3 . This salt, commonly called 
 soda, is one of the most important of manufactured chemi- 
 cal compounds. The mere mention of the fact that it is 
 essential to the manufacture of glass and soap will serve 
 to give some conception of its importance. It is found 
 in the ashes of sea plants, just as potassium carbonate is 
 found in the ashes of inland plants. Formerly, it was 
 made entirely from plant ashes, but we are no longer de- 
 pendent upon this source for our supply of the salt, as 
 two methods have been devised for preparing it from 
 sodium chloride, with which nature provides us in such 
 abundance. As these methods are of great importance, 
 and are, further, very interesting applications of chemi- 
 cal principles, they will be described below. 
 
 Properties. Anhydrous sodium carbonate is a powder 
 which is formed by heating the crystallized salt. It melts 
 to a clear liquid when heated to a sufficiently high tem- 
 perature. It dissolves in water very readily with evolution 
 of heat. The action is, however, not as marked as in 
 the case of potassium carbonate. When the salt is 
 deposited from a water solution it has the composition 
 Na 2 CO 3 + 10H 2 O. This salt, it will be observed, con- 
 tains the same number of molecules of water of crystal- 
 lization as sodium sulphate. Like this, too, it effloresces 
 when exposed to the air. When heated it melts in its 
 water of crystallization, and the salt Na 2 CO 3 -f- H 2 O, or 
 (HO) 2 C(ONa) 2 , separates. This, however, loses water 
 
510 INORGANIC CHEMISTRY. 
 
 when heated higher, and is converted into the anhydrous 
 salt. The conduct of the carbonate towards water at dif- 
 ferent temperatures is suggestive of that of the sulphate. 
 Its maximum solubility is at temperatures between 38 
 and 70. Above the latter point the solubility decreases. 
 The cause of this phenomenon is, in all probability, 
 the same as that referred to in describing the analogous 
 phenomenon presented by the sulphate ; that is, the ex- 
 istence of the hydrated compound Na 2 CO 3 -|- 10H 2 O in 
 solution at temperatures below 70, and the dissociation 
 of this compound into water and salts containing a smaller 
 number of molecules of water of crystallization, which 
 are less soluble, when the temperature is raised above 
 this point. The crystals of sodium carbonate containing 
 ten molecules of water of crystallization belong to the 
 monoclinic system. 
 
 Applications. Sodium carbonate is used in immense 
 quantities in the manufacture of glass, and in the prepa- 
 ration of caustic soda, which is used in the manufacture 
 of soap. 
 
 The lie Blanc Process for the Manufacture of Sodium 
 Carbonate. In the manufacture of soda the problem to 
 be solved is to convert sodium chloride into sodium car- 
 bonate. The first method devised for this purpose is 
 that of Le Blanc. During the French revolution the 
 supply of potash was cut off from France. This led the 
 government to offer a prize for a practical method for 
 manufacturing soda from common salt. The method 
 proposed by Le Blanc at that time, and which, until re- 
 cently, has been used almost exclusively involves three 
 reactions : 
 
 (1) The sodium chloride is converted into sodium 
 sulphate by treating it with sulphuric acid : 
 
 2NaCl + H 2 S0 4 = Na 2 SO 4 + 2HC1. 
 
 (2) The sodium sulphate thus obtained is heated with 
 charcoal, which reduces it to sodium sulphide : 
 
 Na 2 SO 4 + 20 = Na 3 S +-2CO,. 
 
SODIUM CARBONATE- LE BLANC PROCESS. 511 
 
 (3) The sodium sulphide is heated with calcium car- 
 bonate, when sodium carbonate and calcium sulphide 
 are formed : 
 
 Na 2 S + CaCO 3 = Na a CO 3 + CaS. 
 
 The conversion of the sulphate into the carbonate is, 
 therefore, expressed by the equation 
 
 Na 2 SO 4 + 20 + CaCO 3 = Na 2 CO 3 + CaS + 2CO 2 . 
 
 Calcium sulphide is insoluble in water, so that by 
 treating the resulting mass with water the sodium car- 
 bonate is separated from the sulphide. 
 
 In practice the sodium sulphate is mixed with coal 
 and calcium carbonate, and the mixture heated in ap- 
 propriately constructed furnaces. The coal reduces the 
 sulphate to sulphide, which then reacts upon the cal- 
 cium carbonate as above represented. The product of 
 the action is known as crude soda or black ash. It con- 
 tains, as its chief constituents, sodium carbonate and 
 calcium sulphide, together with some calcium oxide, and 
 a number of other substances in small quantities. In 
 order to purify this product, it is broken to pieces, and 
 treated with water ; and the solution thus obtained 
 evaporated, when the salt of the composition Na. 2 CO 3 -(- 
 2H 2 O is deposited. This is dipped out, and dried by 
 heat, when it loses all its water. The product is the 
 calcined purified soda of commerce. This always contains 
 some sulphate and chloride together with a small quan- 
 tity of sulphite. 
 
 When dissolved in water and allowed to crystallize, 
 the salt is deposited in large crystals which contain 
 water in the proportion represented by the formula 
 Na 2 CO 3 -f- 10H 2 O. This is the so-called crystallized soda. 
 
 Most of the soda which comes into the market is the 
 calcined variety. The mother-liquors from the crystal- 
 lized soda contain some sodium hydroxide in consequence 
 of the action of calcium hydroxide on sodium carbonate. 
 This can be converted into soda by passing carbon di- 
 
512 INORGANIC CHEMISTRY. 
 
 oxide into it ; and it can also be partly separated from 
 the carbonate and brought into the market as such. 
 
 A method has recently been devised for the purpose of 
 avoiding the manufacture and use of sulphuric acid in 
 the soda factories. This consists in passing a hot mixture 
 of sulphur dioxide, air, and steam over sodium chloride. 
 The action which takes place is represented by this 
 equation : 
 
 2NaCl + SO a + H 2 O + O = Na a S0 4 + 2HC1. 
 
 As, in the manufacture of soda, by the Le Blanc pro- 
 cess, the sulphur remains in combination as calcium 
 sulphide, a process, known as the Chance process, has 
 been devised for its recovery. This consists in passing 
 carbon dioxide into the waste, thus liberating hydrogen 
 sulphide ; passing this into another portion of the waste, 
 thus converting the calcium sulphide into the hydro- 
 sulphide ; and then treating this with carbon dioxide, 
 when a gas rich in hydrogen sulphide is given off : 
 
 CO a + CaH a S a + H a O = CaCO 3 + 2H a S. 
 
 By regulating the supply of air the gas is burned either 
 to sulphur dioxide or to sulphur. 
 
 Ammonia Process for the Manufacture of Soda. An- 
 other process now in extensive use for the manufacture 
 of soda is the so-called ammonia process, or the Solvay 
 process. This depends upon the fact that mono-sodium 
 carbonate, HNaCO 3 , is comparatively difficultly soluble 
 in water. If, therefore, mono-ammonium carbonate, or 
 acid ammonium carbonate, HNH 4 CO 3 , is added to a solu- 
 tion of common salt, acid sodium carbonate, HNaCO 3 , 
 crystallizes out, and ammonium chloride remains in the 
 solution : 
 
 NaCl + HNH 4 CO 3 = HNaC0 3 + NH 4 C1. 
 When the acid carbonate thus obtained is heated, it gives 
 
SODA FROM CRYOLITE. 513 
 
 off carbon dioxide, and is converted into the normal salt 
 thus : 
 
 2HNaCO 3 = Na 2 CO 3 + CO 2 + H 2 O. 
 
 The carbon dioxide given off is passed into ammonia, 
 and thus again obtained in the form of acid ammonium 
 carbonate : 
 
 NH 3 + H 2 O + CO 2 = HNH 4 CO 3 . 
 
 The ammonium chloride obtained in the first reaction 
 is treated with lime or magnesia, MgO, and the ammonia 
 set free. This ammonia is used again in the preparation 
 of acid ammonium carbonate. The object of using mag- 
 nesia is to get magnesium chloride, which, when evap- 
 orated to dryness and heated, yields magnesia and hy- 
 drochloric acid : 
 
 MgCl 2 + H 2 = MgO + 2HC1. 
 
 More than half the soda supply of the world is now fur- 
 nished by the Solvay process. 
 
 Manufacture of Soda from Cryolite. As cryolite occurs 
 in nature in large quantities, and can be obtained cheaply, 
 it is used in some places for the manufacture of soda. 
 The reactions involved are : 
 
 (1) The action of calcium carbonate upon cryolite at 
 a high temperature, when sodium aluminate, calcium 
 fluoride, and carbon dioxide are formed as represented 
 in the equation 
 
 Na 3 AlF 6 + 3CaCO 3 = 3CaF 2 + Na 3 AlO 3 + 3CO 2 . 
 
 (2) The action of carbon dioxide upon the solution of 
 the aluminate, when aluminium hydroxide is precipitated, 
 and sodium carbonate formed which remains in solution : 
 
 2Na 3 A10 3 + 3C0 2 + 3H 2 O = 3Na 2 CO 3 + 2A1(OH) 3 . 
 
 After the mixture of cryolite and calcium carbonate, or 
 chalk, has been heated, the mass is treated with water, 
 when the sodium aluminate dissolves, while the calcium 
 fluoride does not. After separating the solution from 
 the insoluble residue, carbon dioxide is passed through it. 
 
514 INORGANIC CHEMISTRY. 
 
 Mono-Sodium Carbonate, Primary Sodium Carbonate, 
 HNaCOg. This salt is commonly called " bi-carbonate of 
 soda." It is easily prepared by passing carbon dioxide 
 over the ordinary carbonate dissolved in its water of 
 crystallization : 
 
 Na a C0 3 + CO a + H 2 O = 2HNaCO 3 . 
 
 When heated it gives up carbon dioxide and water, and 
 is converted into the normal salt. As was stated in con- 
 nection with the ammonia-soda process, primary sodium 
 carbonate is much more difficultly soluble in water than 
 the normal salt. At ordinary temperatures 100 parts of 
 water dissolve about 10 parts of the salt. 
 
 It is used in medicine, and extensively in the prepara- 
 tion of soda-water and other effervescing drinks. 
 
 Sodium-Potassium Carbonate, KNaCO 3 + 12H 2 O, is an 
 interesting example of a salt of a dibasic acid containing 
 two different metals. It is easily made by mixing solu- 
 tions of potassium and sodium carbonates, and is ob- 
 tained in the form of large crystals. 
 
 Phosphates. There are three phosphates of sodium 
 just as there are three phosphates of potassium. The 
 point of chief interest presented by them is that the 
 secondary salt, HNa 2 PO 4 , is the one most easily obtained, 
 and is the substance commonly known as sodium phos- 
 phate. When a solution of this salt is treated with an 
 excess of sodium hydroxide, and the solution evaporated, 
 normal or tertiary sodium phosphate crystallizes out. This 
 has the composition Na 3 PO 4 + 12H 2 O. The solution of 
 the latter salt has an alkaline reaction, and when ex- 
 posed to the air it absorbs carbon dioxide, and is con- 
 verted into the secondary salt : 
 
 2Na 3 P0 4 + CO, + H 2 O = 2HNa 2 PO 4 + Na 2 C0 3 . 
 
 Secondary sodium phosphate, HNa 2 PO 4 -f- 12H 2 O, is 
 easily made by adding sodium carbonate to a solution of 
 phosphoric acid until an alkaline reaction is shown. It 
 is also prepared on the large scale from bone-ash. It 
 forms monoclinic prisms which effloresce in the air. 
 
SODIUM BORATE. 515 
 
 Sodium Metaphosphate, NaPO 3 , is formed when the pri- 
 mary phosphate is ignited. There are several modifica- 
 tions of the salt which appear to differ from one another 
 as represented in the formulas NaPO 3 , Na 2 P 2 O 6 , Na 3 P 3 O 9 , 
 etc. This relation is called polymerism; or substances 
 which have the same composition but different molecu- 
 lar weights are said to be polymeric. Relations of this 
 kind are very common among the compounds of carbon. 
 Among the hydrocarbons mentioned in Chapter XIX, 
 for example, are acetylene, C 2 H 2 , and benzene, C 6 H,. 
 There are, further, two other hydrocarbons of the for- 
 mulas C 4 H 4 and C 8 H 8 . Plainly these hydrocarbons all 
 have the same percentage composition. They are poly- 
 meric in the sense in which that term has been defined. 
 
 Di-sodium Pyro-antimonate, H 2 Na 2 Sb 2 O7 + 6H 2 O, is of 
 special interest because it is insoluble in cold water, and 
 may therefore be used for the purpose of detecting so- 
 dium in analysis. It is formed when a solution of the 
 corresponding potassium salt is added to a solution of 
 a sodium salt. 
 
 Sodium Borate. Normal boric acid, as we have seen, 
 has the composition B(OH) 3 , and there are a number of 
 borates derived from this acid by direct replacement of 
 the hydrogen by metals. The salt which boric acid 
 most readily forms with sodium hydroxide or sodium 
 carbonate, however, is that derived from tetraboric acid, 
 H 2 B 4 O 7 , which is derived from normal boric acid by 
 elimination of water. (See p. 355). This salt is borax, 
 which in crystallized form has the composition repre- 
 sented by the formula Na 2 B 4 O 7 + 10H 2 O. By adding 
 the required quantity of sodium hydroxide to a solution 
 of borax, and evaporating to crystallization, sodium 
 metaborate, NaBO 2 + 4H 2 O, is obtained : 
 
 Na 2 B 4 7 + 2NaOH = 4NaBO 2 + H a O. 
 
 The metaborate is decomposed when its solution is ex- 
 posed to the action of the air. It is thus converted by 
 carbon dioxide into sodium carbonate and borax, or 
 sodium tetraborate. 
 
516 INORGANIC CHEMISTRY. 
 
 Borax occurs in nature in several lakes in Asia and 
 in Clear Lake, Nevada, in the United States. It is man- 
 ufactured by neutralizing, with sodium carbonate, the 
 boric acid found in Tuscany. When heated, borax puffs 
 up, and at red heat it melts, forming a transparent, color- 
 less liquid. The dehydrated salt is known, as anhydrous 
 or calcined borax. In the molten condition, borax has 
 the power to combine with metallic oxides, and, as many 
 of the double borates thus formed are colored, the salt 
 is used in blow-pipe work for the purpose of detecting 
 certain metals. As it dissolves metallic oxides, it is 
 used in the process of soldering, as it is necessary 
 to have bright, untarnished metallic surfaces in order 
 that the solder shall adhere firmly. The action of mol- 
 ten borax upon metallic oxides is similar to that which 
 takes place when sodium hydroxide acts upon a solution 
 of borax. Borates of the metals are formed together 
 with sodium borate, or double borates in which part of 
 the hydrogen is replaced by sodium and part by other 
 metals. 
 
 Borax is extensively used in the manufacture of por- 
 celain and in glass-painting. It is an antiseptic, pre- 
 venting the decomposition of some organic substances. 
 
 Sodium Silicate, Na 2 SiO 3 . Sodium silicate is formed 
 by dissolving silicon dioxide in sodium hydroxide, and 
 can be obtained in crystallized form. It is prepared on 
 the large scale by melting together quartz sand and so- 
 dium carbonate in the proper proportions, and f by melt- 
 ing together sodium sulphate, quartz sand, and charcoal 
 powder. This substance is commonly known as water- 
 glass. It is soluble in water, and, when its solution 
 dries, it leaves a transparent coating on the surface on 
 which it is placed. It is extensively used in the manu- 
 facture of artificial stone. 
 
 LITHIUM, Li (At. Wt. .6.97). 
 
 Lithium occurs in nature in relatively small quantity, 
 chiefly in the minerals lepidolite, petalite, and spodu- 
 mene, and in many mineral waters. It is also found in 
 
AMMONIUM SALTS. 517 
 
 the ashes of a number of plants. It is prepared by the 
 electrolysis of the chloride in the molten condition. 
 The metal is silver- white, and is characterized by its low 
 specific gravity. It acts vigorously upon water, but, if 
 the water is at the ordinary temperature, the hydrogen 
 given off does not take fire. In the air it conducts itself 
 in much the same way that sodium does. 
 
 The most characteristic salts of lithium are the phos- 
 phate, carbonate, and chloride. 
 
 Lithium Phosphate, Li 3 PO 4 + pI 2 O, is precipitated when 
 secondary sodium phosphate is added to a solution of a 
 lithium salt. It is very difficultly soluble in water at the 
 ordinary temperature. 
 
 Lithium Carbonate, Li 2 CO 3 , is also rather difficultly sol- 
 uble in water, and is deposited when a solution of sodi- 
 um carbonate is added to a fairly concentrated solution 
 of lithium chloride. It dissolves uric acid, which is in- 
 soluble in water, and is therefore used in medicine for 
 the purpose of removing pathological deposits of this 
 acid in the body. For this purpose it is generally ad- 
 ministered in the form of a solution in water containing 
 carbon dioxide. 
 
 Lithium Chloride, LiCl, is peculiar on account of the 
 fact that it is soluble in alcohol and in a mixture of al- 
 cohol and ether. In this respect it differs from the 
 chlorides of potassium and sodium, which are insoluble 
 in alcohol. If, therefore, a mixture of the chlorides of 
 the three metals is treated with alcohol, only lithium 
 chloride dissolves ; and in this way lithium can be sep- 
 arated from the other metals. 
 
 AMMONIUM SALTS. 
 
 Attention has already been called to the marked simi- 
 larity of the salts of potassium and sodium to those 
 formed by the action of ammonia on the acids, and 
 known as ammonium salts. The most important of 
 these salts will be briefly considered in this connection. 
 A characteristic property of ammonium salts which dis- 
 tinguishes them from the salts of all the metals is their 
 
518 INORGANIC CHEMISTRY. 
 
 volatility. "When sublimed, they all undergo decomposi- 
 tion, which is either partial or complete. 
 
 The simplest kind of decomposition which they un- 
 dergo is dissociation into ammonia and the acid. This 
 is illustrated in the case of ammonium chloride, which, 
 when heated to a sufficiently high temperature, is dis- 
 sociated into ammonia and hydrochloric acid. This is 
 an example of true dissociation. The amount of decom- 
 position is constant for any given temperature and pres- 
 sure. 
 
 An ammonium salt of a polybasic acid containing some 
 metal gives off ammonia and leaves an acid salt, which 
 generally undergoes further decomposition. Thus, so- 
 dium-ammonium sulphate, NaNH 4 SO 4 , first gives off 
 ammonia and forms mono-sodium sulphate : 
 
 The acid salt thus formed then undergoes further 
 change and the pyrosulphate is formed : 
 
 Another example of this kind of decomposition of am- 
 monium salts is that afforded by sodium - ammonium 
 phosphate, HNaNH 4 PO 4 . When heated, this gives off 
 ammonia and then water, the final product being sodium 
 metaphosphate : 
 
 (ONa 
 PO-^ ONH 4 
 
 (OH 
 
 (ONa 
 = POl OH +NH 3 ; 
 
 (OH 
 
 (ONa 
 PO^ OH 
 OH 
 
 = PO 2 ONa + H 2 O. 
 
 Some ammonium salts undergo deeper-seated decom- 
 positions, and do not give ammonia as one of the prod- 
 ucts. This is true especially of such salts as readily 
 give off oxygen. In such cases the ammonia is oxidized, 
 
AMMONIUM SALTS. 519 
 
 so that the hydrogen forms water. This is illustrated 
 in the decomposition of ammonium nitrate and nitrite L 
 
 NH 4 N0 3 = N 2 + 2H 2 ; and 
 NH 4 N0 2 :=N 2 +2H 2 0. 
 
 Further, all ammonium salts are decomposed with evo- 
 lution of ammonia when treated with basic hydroxides. 
 This has been illustrated in the preparation of ammonia 
 from ammonium chloride by treatment with calcium 
 hydroxide : 
 
 2NH 4 C1 + Ca(OH) 2 = CaCl 2 + 2NH 3 + 2H 2 O. 
 
 The ammonium salts are made by neutralizing acids 
 with ammonia. 
 
 Ammonium Chloride, NH 4 C1. This salt is commonly 
 called sal ammoniac. At present its principal source is 
 the so-called ammoniacal liquor of the gas-works. This 
 liquid contains a considerable quantity of ammonium 
 carbonate, and, when it is treated with lime, ammonia is 
 given off. This is passed into hydrochloric acid, and the 
 solution of ammonium chloride thus formed evaporated 
 to crystallization. The salt has a sharp, salty taste, and 
 is easily soluble in water. When heated, it is converted 
 into vapor without melting and with very slight decom- 
 position ; and when the vapor comes in contact with a 
 cold surface, it condenses in crystalline form. This pro- 
 cess of vaporizing and condensing a solid is called sub- 
 limation. Some of the ammonium chloride met with in 
 the market has been sublimed. The salt is used in the 
 preparation of ammonia, in medicine, and for other pur- 
 poses. When it is dissolved in water, a considerable 
 lowering of temperature is caused. 
 
 Ammonium Sulphocyanate, NH 4 CNS. This salt is pre- 
 pared by bringing together aqueous ammonia, carbon 
 disulphide, and alcohol. The first product is ammonium 
 thiocarbamate, the formation of which is perfectly analo- 
 gous to the formation of the ordinary carbamate by the 
 action of carbon dioxide on ammonia : 
 
520 INORGANIC CHEMISTRY. 
 
 The thiocarbamate afterwards breaks down when 
 heated, forming the sulphocyanate and hydrogen sul- 
 phide : 
 
 CS = CNSNH 4 + H 3 S. 
 
 The salt, like so many other ammonium salts, causes 
 a marked lowering of temperature when dissolved in 
 water. When 100 grams are dissolved in the same 
 weight of water at 17, the temperature falls to 12. It 
 is now much used in analytical processes for the esti- 
 mation of silver and copper. 
 
 Ammonium Sulphide, (NH 4 )aS. This compound is ex- 
 tensively used in chemical analysis for the purpose of 
 precipitating those sulphides which are soluble in dilute 
 hydrochloric acid (see p. 198 and p. 473). As will be re- 
 membered, in the usual method of analyzing a mixture 
 of substances, the first step consists in adding hydro- 
 chloric acid to the solution. This precipitates silver, 
 lead, and, under certain conditions, mercury. This pre- 
 cipitate having been filtered off, hydrogen sulphide is 
 passed through the filtrate, when those metals are pre- 
 cipitated whose sulphides are insoluble in dilute hydro- 
 chloric acid. The precipitate is filtered off, and ammo- 
 nium sulphide added to the filtrate, when the metals 
 whose sulphides are soluble in dilute hydrochloric acid 
 are thrown down. Among these are iron, cobalt, nickel, 
 manganese, etc. Any other soluble sulphide might be 
 used ; but the advantage of ammonium sulphide is that it 
 is volatile, and hence, by evaporating the solution and 
 heating, it can be got rid of after it has served its pur- 
 pose. Another use to which it is put in analysis is for 
 the purpose of dissolving the sulphides of tin, arsenic, 
 and antimony, which are precipitated by hydrogen sul- 
 phide, and thus separating these from the other sul- 
 phides of the group. This solution depends upon the 
 
AMMONIUM SULPHIDE. 521 
 
 poVer of the sulphides to form salts of sulpho- acids, as 
 has been repeatedly explained. 
 
 Ammonium sulphide is made by passing hydrogen 
 sulphide into an aqueous solution of ammonia. If the 
 gas is passed until the solution is saturated, the product 
 is the hydrosulphide : 
 
 NH 3 + H 2 S = NH 4 HS. 
 
 If only half this quantity of the gas is passed, the pro- 
 duct is the sulphide : 
 
 2NH 3 + H 2 S = (NH 4 ) 2 S. 
 
 The simplest way to make it, however, is to divide a 
 quantity of a solution of ammonia into two equal parts ; 
 saturate one half, thus forming the hydrosulphide, and 
 add the other half, when this reaction takes place : 
 HNH 4 S + NH 3 = (NH 4 ) 2 S. 
 
 The product is a colorless liquid of a disagreeable 
 odor. It soon changes color, becoming yellow, and after 
 a time a yellow deposit is formed in the vessel in which 
 it is contained. This change of color is due to the action 
 of the oxygen of the air. Some of the sulphide is de- 
 composed into ammonia, water, and sulphur, thus : 
 
 (NH 4 ) 2 S + O = 2NH 3 + H 2 + S. 
 
 The sulphur set free in this way combines with the 
 undecomposed ammonium sulphide, forming the com- 
 pounds (NH 4 ) 2 S 2 , (NH 4 ) 2 S 3 , (NH 4 ) 2 S 4 , and (NH 4 ) 2 S 5 . When 
 as much sulphur has been set free as is required to form 
 the pentasulphide, further decomposition by the oxygen 
 of the air causes a deposit of sulphur. Therefore,, in 
 bottles containing ammonium sulphide which are al- 
 lowed to stand for a long time a deposit of sulphur is 
 always found. A solution containing the polysulphides 
 is called yellow ammonium sulphide. It is this which is 
 used for the purpose of dissolving the sulphides of 
 arsenic, antimony, and tin in analytical operations. 
 
 As stated above, a solution of ammonium hydrosul- 
 phide, HNH 4 S, is made by passing hydrogen sulphide 
 into a solution of ammonia until no more is taken up. 
 
522 INORGANIC CHEMISTRY. 
 
 Ammonium Nitrate, NH 4 NO 3 , is obtained in crystals, 
 which are easily soluble in water. It is of use chiefly in 
 the preparation of nitrous oxide. When heated sud- 
 denly to a high temperature it is decomposed rapidly 
 into nitrogen, water, and nitric oxide : 
 
 2NH 4 N0 3 = N 2 + 2NO + 4H 2 O. 
 
 This decomposition may take place in the preparation of 
 nitrous oxide if in the last stages of the operation the 
 heat is raised too high, and explosions may be caused 
 in this way. When dissolved in water a marked lower- 
 ing of temperature takes place. 
 
 Ammonium Carbonate, (NH 4 ) 2 CO 3 When dry ammonia 
 gas and dry carbon dioxide are brought together, they 
 unite and form the salt known as ammonium carbamate, 
 
 which has the composition CO 
 
 This is the salt of an acid, CO j QTT S J known as car- 
 
 bamic acid. When the carbamate is dissolved in water, 
 it is converted into the carbonate : 
 
 ONH 
 
 When heated to 58, the normal carbonate is decomposed, 
 forming carbon dioxide, water, and ammonia. The sub- 
 stance found in the market under the name of ammoni- 
 um carbonate is made by heating together ammonium 
 chloride or sulphate and chalk. It consists of normal 
 ammonium carbonate, (NH 4 ) 2 CO 3 , primary ammonium 
 carbonate, HNH 4 (CO 3 ), and ammonium carbamate. 
 
 Primary Ammonium Carbonate, HNH 4 CO 3 , is formed by 
 treating the normal carbonate with carbon dioxide, and 
 by allowing the commercial carbonate to lie exposed to 
 the air, when the carbamate is converted into the car- 
 bonate by the moisture, and the carbonate loses am- 
 monia : 
 
REACTIONS OF THE MEMBERS OF THE SODIUM GRO UP. 523 
 
 It is easily decomposed into ammonia, water, and car- 
 bon dioxide. 
 
 Sodium-ammonium Phosphate, HNaNH 4 PO 4 . This salt 
 is known as microcosmic salt, and is much used in the 
 laboratory in blow-pipe work. It is contained in guano 
 and in decomposed urine. It is easily made by mixing 
 solutions of di-sodium phosphate and ammonium chlo- 
 ride, and allowing to crystallize. In crystallized form it 
 contains four molecules of water, HNaNH 4 PO 4 -f- 4H 2 O. 
 The changes which the anhydrous salt undergoes when 
 heated were described on page 518. When the crystal- 
 lized salt is heated, the water of crystallization is first 
 given off. The value of the salt in blow-pipe work de- 
 pends upon the fact that at high temperatures the meta- 
 phosphate combines with metallic oxides, forming mixed 
 phosphates, the reactions being like those which meta- 
 phosphoric acid undergoes with water : 
 
 2HP0 3 + H 2 = H 4 P 2 7 ; 
 HP0 3 + H 2 = H 3 P0 4 ; 
 2NaP0 3 + M 2 = Na 2 M 2 P 2 7 ; 
 NaP0 3 + M 2 O = NaM 2 P0 4 . 
 
 Many of these double phosphates and pyrophosphates 
 are colored, and, like the double borates (see p. 516) 
 they furnish a means of detecting some of the metals. 
 
 Reactions of the Members of the Sodium Group which 
 are of Value in Chemical Analysis. The chief difficulty 
 experienced in chemical analysis is in distinguishing 
 between similar elements. Sodium and potassium, for 
 example, conduct themselves so much alike in so 
 many respects that we might subject them to the in- 
 fluence of a number of reagents without being able 
 to tell which one we are working with. For pur- 
 poses of analysis, therefore, it is necessary to take 
 advantage of differences between the elements, and the 
 more striking the differences the better. Those reac- 
 tions which give rise to the formation of insoluble com- 
 pounds or precipitates are most frequently used in 
 analysis. Very few salts of the members of the sodium 
 
524 INORGANIC CHEMISTRY. 
 
 group are insoluble, and the difficulty of distinguishing 
 between these elements is increased by this fact. In 
 ordinary analyses the elements of this group which are 
 of most importance are potassium and sodium, the 
 other elements of the group being but rarely met 
 with. Ammonium compounds are easily distinguished 
 from those of potassium and sodium by the fact that, 
 when treated with caustic soda or potash, they give off 
 ammonia, which is recognized by its characteristic odor. 
 The chief reactions which are of value in distinguishing 
 between potassium and sodium are the following : 
 
 Platinum Chloride, PtCl 4 , forms difficultly soluble salts 
 with potassium and ammonium chlorides. These are 
 the cUoroplatinates, K 2 PtCl 6 and (NH 4 ) 9 PtCl e . The cor- 
 responding salt of sodium is easily soluble. 
 
 Perchloric Acid, HC1O 4 , forms difficultly soluble potas- 
 sium perchlorate, KC1O 4 , when added to solutions of po- 
 tassium salts. 
 
 Fluosilicic Acid, H 2 SiF 6 , forms difficultly soluble salts 
 with potassium and sodium, K 2 SiF 6 and Na 2 SiF 6 , but not 
 with ammonium. 
 
 Tartaric Acid, H 2 (C 4 H 4 O 6 ), forms a difficultly soluble 
 potassium salt of the formula KH(C 4 H 4 O 6 ). The corre- 
 sponding salt of sodium is easily soluble. The forma- 
 tion of mono-potassium tartrate takes place as repre- 
 sented in the equation : 
 
 KOI + H a (C,H,0.) = KH(C,H 4 0.) + HC1. 
 
 Normal or neutral potassium tartrate is soluble in water, 
 so that, if the difficultly soluble acid tartrate is filtered 
 off, and potassium carbonate added to it, it dissolves in 
 consequence of the formation of the neutral salt, which 
 takes place as represented in the equation 
 
 2KH(C.H 4 0.) + K,CO, = 2K,(C.H,O.) + CO, + H 3 O. 
 
 If, to the solution of the neutral salt, hydrochloric acid 
 is added, the acid salt is again formed and precipitated : 
 
 K 2 (C 4 H 4 O 6 ) + HC1 = KH(C 4 H 4 O 6 ) + KC1. 
 
 Di-sodium Pyro-antimonate, Na 2 H 2 Sb 2 O 7 , is insoluble in 
 cold water, and is formed when a solution of the corre- 
 
FLAME REACTIONS AND THE SPECTROSCOPE. 525 
 
 spending potassium salt is added to a solution of a so- 
 dium salt. 
 
 Flame Reactions and the Spectroscope. When a clean 
 piece of platinum wire is held for some time in the flame 
 of the Bunsen burner, it then imparts no color to the 
 flame. If now a small piece of sodium carbonate or any 
 other salt of sodium is put on it, the flame is colored 
 intensely yellow. All sodium compounds have this 
 power, and hence the chemist makes use of this fact for 
 the purpose of detecting the presence of sodium. Simi- 
 larly, potassium compounds color the flame violet ; lith- 
 ium compounds color the flame red ; rubidium and 
 caesium produce colors similar to that of the potassium 
 flame. While it is an easy matter to recognize potas- 
 sium alone, or any one of the other metals alone, it is 
 difficult to do so when they are together in the same 
 compound. For example, when sodium and potassium 
 .are together, the intense yellow caused by the sodium 
 completely masks the more delicate violet caused by 
 the potassium, so that the latter cannot be seen by the 
 unaided eye. In this particular case the difficulty can 
 be got over by letting the light from the flame pass 
 through a blue glass, or through a thin vessel of glass 
 containing a solution of indigo. The yellow light is 
 thus cut off, while the violet light passes through and 
 can be recognized. A more general method for de- 
 tecting the constituents of light is by means of a prism 
 of glass. Lights of different colors, which are pro- 
 duced by ether waves of different lengths, are turned 
 out of their course to different extents when passed 
 through a prism, as is seen when white sunlight is 
 passed through a prism. A narrow beam of white light 
 passing in emerges as a band of various colors, called its 
 spectrum. We thus see that white light is made up of 
 lights of different colors ; or, to speak in the language of 
 physics, that motion of the light-ether which produces 
 upon the eye the sensation of white light is made up of a 
 number of motions, each of which alone produces upon the 
 eye the sensation of a color. Similarly, we can determine 
 what any light is composed of. Every light has its char- 
 
526 INORGANIC CHEMISTRY. 
 
 acteristic spectrum. The light given off from any solid 
 heated to a white heat gives a continuous spectrum, like 
 that of the sunlight. An incandescent gaseous substance, 
 on the other hand, gives a spectrum made up of separate 
 bands of color, or a banded spectrum. The light produced 
 by burning sodium, or by introducing a sodium com- 
 pound in a colorless flame, gives a spectrum consisting of 
 a narrow yellow band. The spectrum of the potassium 
 flame consists essentially of two bands, one red and one 
 violet. Further, these bands always occupy definite posi- 
 tions relatively to one another, so that, in looking through 
 a prism at the light caused by potassium and sodium, the 
 yellow band of sodium is seen in its position, and the 
 two potassium bands in their proper positions. There 
 is therefore no difficulty in detecting these elements 
 when present in the same substance or in the presence 
 of other elements which give characteristic spectra. 
 
 The instrument used for the purpose of observing the 
 spectra of different lights is called the spectroscope* It 
 consists essentially of a prism and two telescopes. 
 Through one of the telescopes the light to be examined 
 is allowed to pass so as to strike the prism properly. 
 The light emerges from the other side of the prism, and 
 is observed through the other telescope, which is pro- 
 vided with lenses for the purpose of magnifying the 
 spectrum. By means of a third telescope, an image of a 
 scale is thrown upon the face of the prism from which 
 the spectrum emerges, and is reflected thence into the 
 observing-tube, together with the spectrum, so that the 
 position of the bands can be accurately determined. By 
 means of the spectroscope, it is possible to detect the 
 minutest quantities of some elements, and, since it was 
 devised, several new elements have been discovered 
 through its aid ; as, for example, csesium, rubidium, thal- 
 lium, indium, gallium, and others. 
 
 * For an account of the spectroscope and its uses, the student should 
 consult some work on physics. The principles involved in its construc- 
 tion and application are physical principles, and cannot properly be 
 taken up in detail in a text-book of chemistry. 
 
CHAPTER XXVI. 
 
 ELEMENTS OF FAMILY II, GROUP A: 
 GLUCINUM MAGNESIUM CALCIUM STRONTIUM- 
 BARIUM [ERBIUM]. 
 
 General. The elements of this group fall into two sub- 
 groups. Calcium, strontium, and barium are strikingly 
 alike. They also have some points in common with the 
 members of the potassium family, and at the same time 
 are related in some degree to the metals of Family III, 
 Group A, which are known as the earth metals. There- 
 fore, calcium, barium, and strontium are generally called 
 the metals of the alkaline earths. Glucinum and mag- 
 nesium resemble the metals of the alkaline earths in some 
 ways, but they also resemble the members of Group B, 
 of the same family, which includes zinc and cadmium. 
 On comparing the group with the elements presented in 
 the last chapter, some analogous facts are noticed. Ar- 
 ranging the five elements of the potassium group in the 
 order of their atomic weights, and the elements of Family 
 II, Group A, in the same way, we have this table : 
 
 Li 
 
 Na 
 
 K 
 
 Eb 
 
 Cs 
 
 6.97 
 
 22.82 
 
 38.82 
 
 84.78 
 
 131.89 
 
 Gl 
 9.01 
 
 Mg 
 24.10 
 
 Ca 
 39.76 
 
 Sr 
 86.95 
 
 Ba 
 136.39 
 
 As regards the analogies between the elements in each 
 group, the general statement can be made that the last 
 three members of each group resemble one another more 
 closely than they resemble the first two members of the 
 group, while the first two members in each group also 
 resemble each other closely. The natural grouping 
 according to the properties is into the sub-groups : 
 
 (527) 
 
528 INORGANIC CHEMISTRY. 
 
 a b 
 
 Lithium, Potassium, 
 
 Sodium, and Rubidium, 
 
 Caesium. 
 
 GluciBum, Calcium, 
 
 Magnesium, and Strontium, 
 
 Barium. 
 
 The relations between the atomic weights of the ele- 
 ments of Family II, Group A, are similar to those of the 
 elements of Family I, Group A. That of magnesium, 
 24.10, is nearly half the sum of those of glucinum, 9.01, 
 and calcium, 39.76. We have 
 
 9.01 + 39. T6 
 
 2 
 
 = 24.38. 
 
 So, also, that of strontium, 86.95, is approximately half 
 the sum of those of calcium, 39.76, and barium, 136.39 
 
 39.76 + 136.39 
 
 In the calcium group the specific gravities increase in 
 the order of the atomic weights : 
 
 At. Wt. Sp. Gr. 
 Calcium,. . . . 39.76 1.57 
 
 Strontium, . . . 86.95 2.5 
 
 Barium, .... 136.39 3.75 
 
 All the elements of the group are bivalent. The general 
 formulas of the principal compounds are as follows : 
 
 MC1 2 , M(OH) 2 , M(N0 3 ) 2 , MS0 4 , M 3 (PO 4 ) 2 , MSiO 3 , etc. 
 
 The chlorides, hydroxides, and nitrates are soluble in 
 water. The sulphates decrease in solubility as the 
 atomic weights increase. Glucinum sulphate, G1SO 4 , is 
 soluble in its own weight of water ;' magnesium sulphate, 
 MgSO 4 , is soluble in about three times its weight of 
 water ; calcium sulphate, CaSO 4 , dissolves in 400 parts ; 
 strontium sulphate, SrSO 4 , in about 8000 parts; and 
 
CALCIUM: OCCURRENCE PREPARATION. 529 
 
 barium sulphate, BaSO 4 , in about 400,000 parts of water. 
 Barium sulphate, as will be seen, is practically insoluble 
 in water. The normal carbonates of all except glucinum 
 are insoluble in water. The solubility of the hydroxides 
 increases as the atomic weight increases. Glucinum hy- 
 droxide is insoluble ; magnesium hydroxide is but 
 slightly soluble. One hundred parts of water at the 
 ordinary temperature dissolve 0.1368 parts of calcium hy- 
 droxide, 2 parts of strontium hydroxide, and 3.5 parts 
 of barium hydroxide. The solubility of strontium and 
 barium hydroxides is, however, much increased at higher 
 temperatures. 
 
 CALCIUM SUB-GROUP. 
 
 This sub-group, as has been stated, consists of the 
 three very similar elements, calcium, strontium, and 
 barium. Of these calcium occurs most abundantly in 
 nature. Barium and strontium frequently accompany 
 each other, and both are found in some localities in com- 
 pany with calcium. They are much less abundant in 
 nature than calcium. 
 
 CALCIUM, Ca (At. Wt. 39.76X 
 
 Occurrence. Calcium is found in nature in enormous 
 quantities, chiefly in the form of the carbonate, CaCO 3 , 
 as limestone, marble, and chalk. It also occurs in the form 
 of the sulphate, CaSO 4 , as gypsum ; of the phosphate, 
 Ca 3 (PO 4 ) 2 , as phosphorite and apatite ; of the fluoride, 
 CaF 2 , as fluor-spar. It is found in solution in most 
 natural waters either as the carbonate or sulphate ; and 
 in the organs of plants and animals. Bones contain a 
 large proportion of calcium phosphate ; egg-shells and 
 coral contain calcium carbonate. 
 
 Preparation. The element is made by decomposing 
 molten calcium chloride by means of the electric current ; 
 and by first making zinc-calcium and distilling off the 
 zinc by heating to a high temperature in a crucible made 
 of carbon from a gas-retort. The zinc-calcium is made 
 by melting together a mixture of calcium chloride, zinc, 
 
530 INORGANIC CHEMISTRY. 
 
 and sodium. The sodium decomposes the chloride, and 
 the reduced metal dissolves in or combines with the zinc 
 as soon as it is formed. 
 
 Properties. It is a brass-yellow, lustrous metal, which 
 in moist air becomes covered with a layer of hydroxide 
 and carbonate. At ordinary temperatures it decomposes 
 water just as potassium and sodium do, but heat is not 
 evolved rapidly enough to set fire to the hydrogen. 
 Heated to a high temperature, it burns in the air, forming 
 the oxide. It is not made in quantity, and has found no 
 practical application. 
 
 Calcium Chloride, CaCL. This salt is found in nature 
 in combination with other chlorides, particularly in the 
 mineral tachydrite, which occurs in the salt deposits at 
 Stassfurt, and has the composition represented by the 
 formula CaCl 2 .MgCl 2 -(- 12H 2 O. It is also found in solu- 
 tion in sea- water. It is obtained as a by-product in the 
 preparation of ammonia from ammonium chloride and 
 lime ; in the preparation of potassium chlorate from cal- 
 cium chlorate and potassium chloride (see p. 494); and in 
 the ammonia-soda process. It is made by dissolving 
 calcium carbonate in hydrochloric acid, as in the prepa- 
 ration of carbon dioxide. From very concentrated 
 solutions it crystallizes with six molecules of water, 
 CaCl 2 -J- 6H 2 O. ,When these crystals are exposed to the 
 air they soon deliquesce. When a solution of calcium 
 chloride is evaporated, and care is taken to keep the 
 temperature below 200, it solidifies, forming a porous 
 mass which has the composition represented by the for- 
 mula CaCl 2 -f- 2H 2 O. This is much used in laboratories 
 as a drying agent, as it absorbs water with great ease. If 
 this salt is heated above 200 it loses all its water, and 
 the dehydrated chloride melts, forming fused calcium 
 chloride. This is also much used on account of its dry- 
 ing power. Gases are passed through tubes filled with 
 granulated calcium chloride for the purpose of drying 
 them, and the salt is also placed in vessels in which it is 
 necessary that the air should be dry, as in balance-cases, 
 desiccators, etc. The fused salt generally has a slight 
 alkaline reaction, which is caused by the presence of a 
 
COMPOUNDS OF CALCIUM. 531 
 
 small quantity of lime. This is formed by the action of 
 steam at high temperature on the chloride, the reaction 
 being represented by this equation : 
 
 CaCl 2 + H 2 O = CaO + 2HC1. 
 
 This decomposition takes place only to a slight extent. 
 The porous chloride, which contains two molecules of 
 water, does not contain any hydroxide, and it is therefore 
 better adapted for use' in cases in which it is necessary 
 that it should not absorb carbon dioxide, as in the analysis 
 of organic compounds. 
 
 Calcium chloride forms crystallized compounds with 
 ammonia and with alcohol, as well as with water. It is 
 obvious from this that calcium chloride cannot be used 
 for the purpose of drying ammonia gas. When the com- 
 pounds with ammonia and with alcohol are heated they 
 break down, yielding ammonia and aksohol respectively, 
 as the compound with water gives up the latter. 
 
 Calcium Fluoride, CaF 2 . This compound occurs in 
 large quantities in nature as the mineral fluor-spar. It 
 occurs beautifully crystallized in cubes, and is insoluble 
 in water. It is the source of fluorine compounds in gen- 
 eral, and is used in metallurgical operations for the 
 reason that it melts readily and does not act upon other 
 substances easily. It therefore simply serves as a liq- 
 uid medium in which reactions take place at high tem- 
 peratures. A substance which acts in this way and is 
 used for this purpose is called a flux. The name fluor- 
 spar has its origin in this use of the substance. A flux 
 plays to some extent the same part at elevated tempera- 
 ture in facilitating reactions that water plays at ordinary 
 temperatures. 
 
 Calcium Oxide, CaO. This important compound is 
 commonly called lime, or, to distinguish it from the hy- 
 droxide or slaked lime, it is called quick-lime. It is made 
 in large quantity by heating calcium carbonate in ap- 
 propriately constructed furnaces, known as lime-kilns. 
 Pure lime is made by decomposing some pure form of 
 calcium carbonate, as marble or calc-spar. The decom- 
 position of calcium carbonate is not complete in an at- 
 
532 INORGANIC CHEMISTRY. 
 
 mosphere of carbon dioxide, hence precautions must be 
 taken to remove the gas formed by the decomposition. 
 Further, when lime is heated to a temperature consider- 
 ably higher than that necessary to effect the first decom- 
 position it again absorbs carbon dioxide. 
 
 Lime is a white, amorphous, infusible substance. When 
 heated in the flame of the compound blow-pipe it gives 
 an intense light, as any other infusible substance would 
 do under the same circumstances. When exposed to the 
 air it attracts moisture and carbon dioxide, and is con- 
 verted into, the carbonate. It must therefore be protected 
 from the air. Lime which has been converted into the 
 carbonate by exposure to the air is said to be air-slaked. 
 
 Calcium Hydroxide, Ca(OH) 2 . When calcium oxide or 
 quick-lime is treated with water it becomes hot and crum- 
 bles to a fine powder. The substance which is formed 
 in this operation is somewhat soluble in water, the solu- 
 tion being known as lime-water. The chemical change 
 which takes place when lime is treated with water has 
 been explained. It consists in the formation of a com- 
 pound of the formula Ca(OH) 2 , known as slaked lime ; 
 and the operation is known as slaking. The action is of 
 the same kind as that with which we have so frequently 
 had to deal in the transformation of oxides into the cor- 
 responding hydroxides. Thus when potassium oxide is 
 treated with water it is changed to the hydroxide, with 
 a marked evolution of heat, the reaction being repre- 
 sented in this way : 
 
 So, too, when sulphur trioxide is brought in contact with 
 water it appears to form the hydroxide, normal sulphuric 
 acid : 
 
 = S- 
 
 OH 
 OH 
 OH 
 OH 
 OH 
 OH 
 
CALCIUM HYDROXIDE. 533 
 
 The action in the case of calcium oxide is represented in 
 a similar way : 
 
 The hydroxide is a fine white powder. At red heat it 
 loses water and is reconverted into the oxide : 
 
 
 When lime-water is exposed to the air it becomes cov- 
 ered with a crust of calcium carbonate, and finally all the 
 calcium is precipitated as calcium carbonate. A solution 
 of calcium hydroxide affords a convenient means of de- 
 tecting the presence of carbon dioxide, as has been shown 
 in dealing with this gas. The solution has an alkaline 
 reaction, and acts in many respects like the hydroxides 
 of potassium and sodium. Attention has been called to 
 the fact that the hydroxides of most of the metals are 
 insoluble in water, and that when a soluble hydroxide is 
 added to the salt of such a metal the insoluble hydrox- 
 ide is precipitated. The same kind of decomposition of 
 salts is effected by a solution of calcium hydroxide. 
 Thus, when it is added to ferric chloride, ferric hydrox- 
 ide is thrown down : 
 
 Cl p Q /OH (OH 
 
 Fe Cl < OH Fe^ OH 
 
 Fe Cl n ^OH Fe-| OH 
 
 Cl Ua ^OH OH 
 
 This reaction is entirely analogous to that which takes 
 place between ferric chloride and potassium hydroxide : 
 
 (Cl KOH (OH KC1 
 
 Fe^ Cl + KOH = Fe \ OH + KC1 . 
 Cl KOH OH KC1 
 
534 INORGANIC CHEMISTRY. 
 
 Lime is extensively used in the arts, generally in the 
 form of the hydroxide. As we have seen, it is used in 
 the preparation of ammonia and the caustic alkalies, 
 potassium and sodium hydroxides ; and of bleaching- 
 powder and potassium chlorate. It is further used in 
 large quantity in the process of tanning for the pur- 
 pose of removing the hair from hides ; in decomposing 
 fats for the purpose of making stearin for candles ; 
 for purifying gas ; and especially in the preparation of 
 mortar. 
 
 Bleaching-powder. The preparation of bleaching-pow- 
 der was referred to under Chlorine (which see). The 
 main reaction involved is that represented in the equa- 
 tion 
 
 2Ca(OH) a + 401 = Ca(ClO) 2 + CaCl 2 + 2H 2 O. 
 
 Bleaching-powder 
 
 The compound is commonly called " chloride of lime." 
 Assuming that the reaction takes place in the same way 
 as that of chlorine on caustic potash, the product is a 
 mixture of calcium hypochlorite, Ca(ClO) 2 , and calcium 
 chloride, for it is held that the reaction with potassium 
 hydroxide takes place as represented in this equation : 
 
 2KOH + 2C1 = KC10 + KC1 + H 2 O. 
 
 An objection to the view that calcium chloride is pres- 
 ent as such in bleaching-powder is found in the fact that 
 the substance is not deliquescent, as it should be if cal- 
 cium chloride were present. This has led to the sug- 
 gestion that bleaching-powder in the dry form is not a 
 mixture of two compounds, as represented above, but 
 
 {Cl 
 Od 
 
 or CaOCl 2 . A compound of this formula would plainly 
 have the same composition as a mixture of calcium hy- 
 pochlorite and calcium chloride in the proportion of 
 their molecular weights. For we have 
 
BLEACHING-POWDER. 535 
 
 Ca(ClO) 2 + CaCl 2 = 2CaOCl 2 . 
 
 The point is a difficult one to decide, but at present the 
 evidence appears to be rather in favor of the view that 
 bleaching-powder in the dry form is a single compound 
 of the constitution represented by the last formula 
 given. When treated with water, however, it appears to 
 be resolved into a mixture of the hypochlorite and chlo- 
 ride. 
 
 Bleaching-powder is a white powder which has the 
 odor of hypochlorous acid. It is soluble in about 
 twenty parts of water, though the commercial product 
 always leaves a slight residue, which consists mainly of 
 calcium hydroxide. When treated with an acid, as sul- 
 phuric or hydrochloric acid, it gives up all its chlorine. 
 Thus, with hydrochloric acid the reaction takes place 
 as represented in these equations : 
 
 Ca(ClO) 2 + 2HC1 = CaCl 2 + 2HC1O ; 
 2HC1 + 2HC10 = 2H 2 O + 201,. 
 
 With sulphuric acid the action also probably takes place 
 in two stages. The acid acts upon the hypochlorite, 
 setting hypochlorous acid free ; and upon the chloride, 
 setting hydrochloric acid free. The hydrochloric and 
 hypochlorous acids then react with each other as repre- 
 sented above : 
 
 Ca(C10) 2 + H 2 S0 4 = CaS0 4 + 2HC10 ; 
 CaCl 2 +H 2 S0 4 =CaSO 4 + 2HCl; 
 2HC1 + 2HC1O = 2H 2 O + 2C1 2 . 
 
 When exposed to the action of carbon dioxide hypo- 
 chlorous acid is liberated. Hence, when it is allowed to 
 lie in the air this decomposition takes place slowly. The 
 hypochlorous acid acts further upon the calcium chlo- 
 ride, liberating chlorine : 
 
 CaCl 2 + 2HOC1 + C0 2 = CaCO 3 + H 2 O + 201,. 
 
 It may be, however, that the action takes place between 
 carbon dioxide and the compound CaOCl 2 , thus : 
 
 CaOCl 3 + C0 2 = CaCO 3 + Cl a . 
 
536 INORGANIC CHEMISTRY. 
 
 In any case, the fact remains that carbon dioxide sets 
 the chlorine free from bleaching-powder. 
 
 A solution of bleaching-powder alone is not capable 
 of bleaching except very slowly. If, however, something 
 is added which has the power to decompose it, bleach- 
 ing takes place, the action being due to the presence of 
 hypochlorous acid and chlorine. As is clear from what 
 was said above, the passage of carbon dioxide through 
 the solution or the addition of an acid would cause it to 
 bleach. So, too, certain salts produce a similar effect. 
 The explanation of this is the instability of the hypo- 
 chlorites formed by the salts added. When a concen- 
 trated solution of bleaching-powder is heated it gives off 
 oxygen, and the salt is converted into the chloride. In 
 dilute solution, however, the hypochlorite is converted 
 into chlorate and chloride : 
 
 3Ca(ClO) 2 = Ca(ClO 3 ) 2 + 2CaCl 2 . 
 
 This fact is taken advantage of, as has been shown, for 
 the purpose of making calcium chlorate, and from this 
 potassium chlorate (see p. 494). In contact with certain 
 oxides, as copper oxide, ferric oxide, and with hydroxides, 
 as cobalt and nickel hydroxides, a solution of bleaching- 
 powder readily gives up oxygen when heated. 
 
 The chief application of bleaching-powder is, as its 
 name implies, for bleaching. It is also used as a disin- 
 fectant, and as an antiseptic, that is, for the purpose of 
 destroying disease germs, and of preventing decomposi- 
 tion of organic substances. 
 
 Calcium Carbonate, CaCO 3 . This salt occurs in im- 
 mense quantities in nature in the well-known forms lime- 
 stone, calc-spar, marble, and chalk. The variety of 
 calc-spar found in Iceland, and known as Iceland spar, 
 is particularly pure calcium carbonate. It crystallizes 
 in a number of different forms, the most common being 
 in rhombohedrons, as seen in ordinary calc-spar. A 
 second variety of crystallized calcium carbonate is ara- 
 gonite. This is found in nature crystallized in rhombic 
 prisms, and in forms derived from this. When heated 
 
CALCIUM CARBONATE. 537 
 
 aragonite falls to pieces, the particles being small crys- 
 tals of the form characteristic of calc-spar. This is a 
 case of dimorphism similar to that presented by sul- 
 phur, which, it will be remembered, crystallizes in two 
 forms, the rhombic and monoclinic, the latter of which 
 passes into the former spontaneously. These forms are 
 produced artificially very readily. When calcium car- 
 bonate is precipitated from a solution of a calcium salt 
 by adding a soluble carbonate at ordinary temperatures 
 the precipitate is made up of microscopic crystals which 
 have the same form as calc-spar. If, however, the solu- 
 tion from which the carbonate is precipitated is hot, the 
 salt consists of microscopic crystals of the form of ara- 
 gonite. 
 
 The most abundant form of calcium carbonate is lime- 
 stone, of which many great mountain-ranges are largely 
 made up. This is a compact form of the compound, 
 which has a gray color, and frequently consists of mi- 
 nute crystals. It is always more or less impure, contain- 
 ing clay and other substances. Limestone which is 
 mixed with a considerable proportion of clay is called 
 marl. Many natural waters contain calcium carbonate 
 in solution probably in the form of the acid carbonate. 
 When such a water evaporates the carbonate is again 
 deposited. It happens in some places that a water 
 charged with the carbonate works its way slowly through 
 the earth and drops from the top of a cave. Under these 
 circumstances there is a gradual deposit of the salt 
 which remains suspended. Such hanging formations of 
 the carbonate are known as stalactites. At the same 
 time that part of the liquid which falls to the bottom of 
 the cave forms a projecting mass below the stalactite. 
 Such projecting masses are called stalagmites. The for- 
 mation of stalactites takes place in much the same way as 
 that of icicles. 
 
 Much of the calcium carbonate found in nature has its 
 origin in the remains of animals, and fossils are very 
 abundant in it. Chalk consists almost exclusively of the 
 shells of microscopic animals. 
 
 When carbon dioxide is passed into a solution of cal- 
 
538 INORGANIC CHEMISTRY. 
 
 cium hydroxide, the carbonate is precipitated ; and, it the 
 current of gas is continued long enough, the carbonate 
 is redissolved. It appears, therefore, that calcium car- 
 bonate is soluble in water that contains carbonic acid. 
 It is probable that the cause of this is to be found in the 
 formation of an acid carbonate, possibly the one of the 
 formula HO-OC-O-Ca-O-CO-OH. No positive evi- 
 dence, of the formation of this substance has, however, 
 been furnished. If it is formed, it is certainly very un- 
 stable ; for, on heating the solution to boiling, the normal 
 carbonate is precipitated and carbon dioxide is given off. 
 Natural waters which come in contact with limestone 
 gradually tske up more or less of the carbonate, with the 
 aid of the carbon dioxide of the air, and when such a 
 water is boiled, the carbonate is thrown down. A water 
 containing calcium carbonate in solution is called a hard 
 water; and, as this kind of hardness is easily removed by 
 boiling, it is called temporary hardness in order to dis- 
 tinguish it from a kind which is not removed by boiling, 
 and is therefore called permanent hardness. Temporary 
 hardness is further removed by adding lime to the water, 
 when normal carbonate is formed, which is at once pre- 
 cipitated. 
 
 The decomposition of calcium carbonate by heat, lead- 
 ing to the formation of lime, or calcium oxide, and carbon 
 dioxide, was referred to on p. 468. 
 
 Applications. Calcium carbonate is used, in the arts, 
 for a great many purposes, as in the manufacture of glass ; 
 as a flux (see p. 531) in many important metallurgical 
 operations, as in the reduction of iron from its ores ; in 
 the preparation of lime for mortar ; etc. As is well known, 
 further, marble and some of the varieties of limestone 
 are extensively used in building ; and large quantities of 
 chalk are also used. 
 
 Calcium Sulphate, CaSO 4 - This compound is very 
 abundant in nature. The principal natural variety is 
 gypsum, which occurs in crystals containing two mole- 
 cules of water, CaSO 4 + 2H 2 O. This is perhaps derived 
 directly from the normal acid S(OH) 6 , having the con- 
 
CALCIUM SULPHATE. 539 
 
 stitution represented by the formula (HO) 4 S<Q>Oa. 
 
 The salt of the formula CaS0 4 also occurs in nature, and 
 is called anhydrite. A granular form of gypsum is called 
 alabaster. Calcium sulphate is difficultly soluble in hot 
 and cold water, but its solubility is markedly increased 
 by the presence of certain other salts ; as, for example, 
 sodium chloride. It is comparatively easily soluble in 
 hydrochloric acid and in nitric acid. When heated to 
 100, or a little above, it loses nearly all its water and 
 forms a powder known as plaster of Paris, which has the 
 power of taking up water and forming a solid substance. 
 This process of solidification is known as " setting.' 1 
 Plaster of Paris is very largely used in making casts, on 
 account of its power to harden after having been made 
 into a paste with water. The hardening is a chemical 
 process, and is caused by the combination of water with 
 the salt to form the crystallized variety : 
 
 CaSO 4 + 2H 2 O = (HO) 4 S<^>Ca. 
 
 "When heated to 200, and above, all the water is given off 
 from gypsum, and the product now combines with water 
 only very slowly, and is of no value for making casts. In 
 general, the higher the temperature to which the gypsum 
 is heated, the greater the difficulty with which the pro- 
 duct combines with water. 
 
 Many natural waters contain gypsum in solution. Such 
 waters act in some respects like those which contain cal- 
 cium carbonate. With soap, for example, they form in- 
 soluble compounds. They are called hard waters. 
 This kind of hardness is not removed by boiling, and 
 it is therefore called permanent hardness. Magnesium 
 sulphate acts in the same way, producing permanent 
 hardness. 
 
 When calcium sulphate is treated with a solution of a 
 soluble carbonate, it is decomposed, forming the carbon- 
 ate as represented in the equation 
 
 CaS0 4 + Na 3 C0 3 = Na 2 S0 4 + CaCO,. 
 
540 INORGANIC CHEMISTRY. 
 
 This change is effected simply by allowing the two to 
 stand in contact at the ordinary temperature. 
 
 Besides being used for making casts, calcined gypsum 
 is used also in surgery for making plaster-of-Paris band- 
 ages, and as a fertilizer. Its action as a fertilizer is be- 
 lieved by some to be due to the fact that it has the power 
 to hold ammonia and ammonium carbonate in combina- 
 tion, and thus to make them available for the plants. It 
 has recently been shown that it in some way facilitates 
 the process of nitrification, and perhaps it is in conse- 
 quence of this that it facilitates plant-growth. 
 
 Calcium Phosphates. There are three phosphates of 
 calcium : (1) The normal phosphate, Ca 3 (PO 4 ) 2 ; (2) the 
 secondary phosphate, CaHPO 4 ; and (3) the primary phos- 
 phate, CaH 4 (PO 4 ),. 
 
 (1) Normal calcium phosphate, Ca 3 (PO 4 ) 2 , is derived from 
 phosphoric acid by the replacement of all the hydrogen 
 by calcium. It is found in nature in large quantity as 
 phosphorite, and in combination with calcium fluoride or 
 chloride as apatite. It is, further, the chief inorganic 
 constituent of bones, forming 85 per cent of bone-ash, 
 and is contained in the excrement of animals, as in guano, 
 etc. It is found everywhere in the soil, and is taken up 
 by the plants for whose development it is essential. That 
 it is also essential to the life of animals is obvious from 
 the fact that the bones consist so largely of it. The 
 phosphate needed for the building up of bones is taken 
 into the system with the food. From these statements, 
 it is clear that calcium phosphate is of fundamental 
 importance, and that a fertile soil must either contain 
 this salt or something from which it can be formed. 
 Now, when a crop is raised on a given area, a certain 
 amount of the phosphate contained in it is withdrawn. 
 If the plants were allowed to decay where they grow, the 
 phosphate would be returned and the soil would continue 
 fertile ; but in cultivated lands this is not the case. The 
 crops are removed, and with them the calcium phos- 
 phates contained in them, and the soil therefore becomes 
 exhausted. If the substances removed are used as food, 
 some of the phosphate is found in the excrement of the 
 
CALCIUM PHOSPHATES. 541 
 
 animals ; and, if this excrement is put on the soil, it is 
 again rendered fertile. There are, however, other sources 
 of calcium phosphate, and some of these are utilized ex- 
 tensively in the preparation of artificial fertilizers. The 
 natural form of the phosphate, as that in bone-ash, in 
 phosphorite, and in guano, is mainly the normal or neu- 
 tral phosphate. This is insoluble in water, and is there- 
 fore taken up by the plants with difficulty. To mate it 
 quickly available, it must be converted into a soluble 
 phosphate. This is done by treating it with sulphuric 
 acid in order to effect the reaction represented in this 
 equation : 
 
 Ca s (PO 4 ) 2 + 2H 2 S0 4 = CaH 4 (P0 4 ) 2 + 2CaSO 4 . 
 
 The primary phosphate thus formed is soluble in water, 
 and is of great value as a fertilizer. The mixture of the 
 soluble phosphate and of calcium sulphate is known as 
 " superphosphate of lime." The sulphate, as we have 
 seen, is also of value as a fertilizer. The value of super- 
 phosphates depends mostly upon the amount of soluble 
 phosphate contained in them ; and in dealing with them 
 it is customary to state how much " soluble " and how 
 much "insoluble phosphoric acid" they contain. When 
 a superphosphate is allowed to stand for a time, some of 
 the soluble primary phosphate is converted into insol- 
 uble phosphates by contact with basic hydroxides and 
 water. This is known as the process of "reversion," 
 and that part of the phosphoric acid which is contained 
 in the insoluble phosphate is spoken of as " reverted 
 phosphoric acid." 
 
 Normal calcium phosphate, as has been stated, is in- 
 soluble in water, and is formed when a soluble normal 
 phosphate is added to a solution of a calcium salt. It is 
 also formed when di-sodium phosphate and ammonia are 
 added to a solution of a calcium salt, thus : 
 
 + 3CaCl, -f 2NH 3 = Ca 3 (PO 4 ) 2 + 4NaCl 
 
 Di-sodium phosphate alone at first produces a precipitate 
 of the normal phosphate, while the primary phosphate 
 which is formed at the same time remains in solution. 
 
542 INORGANIC CHEMISTRY. 
 
 The reaction takes place thus : 
 
 4HNa 2 P0 4 + 4CaCl a .= CaH 4 (PO 4 ) 3 + Ca 3 (PO 4 ) 3 + SNaCL 
 
 On standing, the primary acts upon the tertiary salt, 
 forming the secondary phosphate thus : 
 
 CaH 4 (P0 4 ) 2 + Ca 3 (P0 4 ) 2 = 4HCaPO 4 . 
 
 But even on long standing this reaction is not complete. 
 Normal or tertiary calcium phosphate is soluble in hy- 
 drochloric acid and in nitric acid, in consequence of the 
 formation of calcium chloride, or nitrate, and the primary 
 phosphate. If ammonia is added to this solution, the 
 tertiary phosphate is again precipitated, as represented 
 below : 
 
 Ca 8 (PO 4 ) 2 + 4HC1 = 2CaCl 2 + H 4 Ca(PO 4 ) 2 ; 
 2CaCl 2 + H 4 Ca(P0 4 ) 2 + 4NH 8 = Ca 3 (PO 4 ) 2 + 4NH 4 C1. 
 
 (2) Secondary calcium phosphate, CaHPO 4 , is formed, 
 as above described, when a solution of a calcium salt 
 is treated with secondary sodium phosphate. 
 
 (3) Primary calcium phosphate, H 4 Ca(PO 4 ) 2 , is com- 
 monly called the acid phosphate of calcium. It is formed 
 when ordinary insoluble calcium phosphate is treated 
 with concentrated sulphuric acid, and is contained in the 
 so-called superphosphates. It is also formed by treat- 
 ing the neutral phosphate with phosphoric acid and with 
 hydrochloric acid. When treated with but little water, it 
 is converted into the secondary salt and free acid : 
 
 H 4 Ca(PO 4 ) 2 =HCaP0 4 + H 3 PO 4 . 
 
 Calcium Silicate, CaSiO 3 , occurs in nature as the mineral 
 wollastonite, and, in combination with other silicates, in 
 a large number of minerals, as garnet, mica, the zeolites, 
 etc. It is formed when a solution of sodium silicate is 
 added to a solution of calcium chloride, and when a 
 mixture of calcium carbonate and quartz is heated to a 
 high temperature. 
 
 Glass. Ordinary glass is a silicate of calcium and 
 sodium made by melting together sand (silicon dioxide, 
 SiO 2 ) with lime and sodium carbonate or soda. In- 
 
GLASS. 543 
 
 stead of calcium carbonate, lead oxide may be used ; and 
 instead of sodium carbonate, potassium carbonate. The 
 properties of the glass are dependent upon the materials 
 used in its manufacture. 
 
 Ordinary window glass is a sodium-calcium glass. 
 The purer the calcium carbonate and silica, the better 
 the quality of the glass. This glass is comparatively 
 easily acted upon by chemical substances, and is there- 
 fore not adapted to the preparation of vessels which are 
 to be used to hold acids and other chemically active 
 substances. It answers, however, very well for windows. 
 The difference between ordinary window glass and plate 
 glass is essentially that the former is blown and then 
 cut into pieces, while the latter, when in the molten con- 
 dition, is run into flat moulds and there allowed to solidify. 
 
 Bohemian glass is made with potassium carbonate. If 
 pure carbonate is used, as well as pure calcium carbonate 
 and silica, a very beautiful glass is the result. It is 
 characterized by great hardness, by its difficult fusibility, 
 and by its resistance to the action of chemical substances. 
 It is particularly well adapted to the manufacture of 
 vessels and tubes for use in chemical laboratories. 
 
 Flint-glass is made by melting together lead oxide, 
 potassium carbonate, and silicon dioxide. It is charac- 
 terized by its power to refract light, its high specific 
 gravity, its low melting-point, and the ease with which 
 it is acted upon by reagents. Owing to its high refrac- 
 tive power, it is largely used in the manufacture of lenses 
 for optical instruments. 
 
 Strass is a variety of lead-glass which is particularly 
 rich in lead. Its refracting power is so great that it is 
 used in the manufacture of artificial gems. 
 
 Colors are given to glass by putting in the fused mass 
 small quantities of various substances. Thus, a cobalt 
 compound makes glass blue ; copper and chromium make 
 it green ; one of the oxides of copper makes it red ; ura- 
 nium gives it a yellow color ; etc. The most common 
 variety of glass is that used in the manufacture of ordi- 
 nary bottles. It is generally green to black, and some- 
 times brown. In its manufacture, impure materials are 
 
544 INORGANIC CHEMISTRY. 
 
 used, chiefly ordinary sand, limestone, sodium sulphate, 
 common salt, clay, etc. 
 
 When glass is suddenly cooled, it is very brittle and 
 breaks into small pieces when scratched or slightly 
 broken in any way. This is shown by the so-called 
 Prince Rupert's drops, which are made by dropping 
 glass, in the molten condition, into water. When the end 
 of such a drop is broken off, the entire mass is completely 
 shattered into minute pieces. It is clear from this that, 
 in the manufacture of glass objects, care must be taken 
 not to cool them suddenly. In fact they are cooled very 
 slowly, the process being known as annealing. For this- 
 purpose they are placed in furnaces the temperature of 
 which is but little below that of fusion, and they are 
 kept there for some time, the heat being gradually 
 lowered. If red-hot glass is introduced into heated oil 
 or paraffin, and allowed to cool very slowly, it is found 
 to be extremely hard and elastic. The glass of De 
 la Bastie is made in this way. Vessels made of it can 
 be thrown about upon hard objects without breaking, 
 but sometimes a slight scratch will cause the glass to fly 
 in pieces as the Rupert's drops do. 
 
 Mortar. Mortar is made of slaked lime and sand. 
 When this mixture is exposed to the air, calcium carbo- 
 nate is slowly formed and the mass becomes extremely 
 hard. The water contained in the mortar soon passes 
 off, but nevertheless freshly plastered rooms remain 
 moist for a considerable time. This is due to the fact 
 that a reaction is constantly taking place between the 
 carbon dioxide and calcium hydroxide by which calcium 
 carbonate and water are formed, 
 
 Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O, 
 
 and it is the water thus liberated which keeps the air 
 moist. The complete conversion of the lime into car- 
 bonate requires a very long time, because the carbonate 
 which is formed on the surface protects, to some extent, 
 the lime in the interior. 
 
 It is generally regarded as unhealthy to live in rooms 
 with freshly plastered walls, because the air is constantly 
 
CALCIUM CARBIDE. 545 
 
 kept moist in consequence of the reaction above men- 
 tioned. It is, however, difficult to see why the presence 
 of a little extra moisture 1 in the air should be unhealthy ; 
 and, if there is any danger from freshly plastered walls, 
 it seems probable that the cause must be sought for else- 
 where. It is possible that the constant presence of 
 moisture in the pores of the wall interferes with the im- 
 portant process of diffusion, and that therefore when the 
 room is closed this natural method of ventilation cannot 
 come into play. 
 
 When lime-stones which contain magnesium carbonate 
 and aluminium silicate in considerable quantities are 
 heated for the preparation of lime, the product does 
 not act with water as calcium oxide does,' and this lime 
 is not adapted to the preparation of ordinary mortar. 
 On the other hand, it gradually becomes solid, in con- 
 tact with water, for reasons which are not known. Such 
 substances are known as cements, or hydraulic cements. 
 Other cements besides those made in the manner men- 
 tioned are known. 
 
 Calcium Sulphide, CaS, is formed by heating calcium 
 sulphate with charcoal. It is remarkable on account 
 of the fact that it is phosphorescent. After having 
 been exposed to sun-light, it continues to give light for 
 some time afterward. This and the similar compound, 
 barium sulphide, are now used quite extensively in the 
 preparation of luminous objects, such as match-boxes, 
 clock-faces, plates for house-numbers, etc. 
 
 Calcium Nitride, Ca 3 N 2 , a brown mass, is formed by 
 heating a 15-20 per cent amalgam of calcium in dry 
 atmosphere to dull redness. It is somewhat volatile, 
 and is decomposed by water, with the formation of cal- 
 cium hydroxide and ammonia. 
 
 Calcium Carbide, C 2 Ca, is formed by heating lime and 
 carbon together in an electric furnace, when the reac- 
 tion represented by the following equation takes place : 
 
 CaO + 30 = C,Ca + CO. 
 
 It is manufactured on the large scale, the form of 
 carbon used being coke. The carbide forms a crystal- 
 
546 INORGANIC CHEMISTRY. 
 
 line mass. That of average quality has a reddish color. 
 That of bad quality has a grayish or black color. 
 When pure it is colorless. When treated with water it 
 yields acetylene, C a H 2 , a gas that is coming into practi- 
 cal use on account of its value for purposes of illumina- 
 tion (see Acetylene, p. 374). 
 
 STKONTIUM, Sr (At. Wt. 86.95). 
 
 Occurrence and Preparation. Strontium occurs in 
 nature in the form of the sulphate, SrSO 4 , as celestite, 
 and in the form of the carbonate, SrCO 3 , as strontianite. 
 The latter is found in large quantities in Westphalia. 
 The element is isolated by the action of an electric cur- 
 rent on the molten chloride. 
 
 Properties. It is very similar to calcium, having a 
 metallic lustre and a brass-yellow color. It is oxidized 
 by contact with the air, and decomposes water rapidly 
 with evolution of hydrogen, which does not, however, 
 take fire spontaneously. 
 
 Compounds of Strontium. The compounds of stron- 
 tium are very similar to those of calcium. Its chloride 
 has not the same attraction for water that calcium chlo- 
 ride has, though it deliquesces when left in contact with 
 the air. The oxide is not easily made by decomposition 
 of the carbonate by heat, as the carbonate is much more 
 stable than that of calcium. It is, however, prepared 
 without difficulty by heating the nitrate. When brought 
 in contact with water, the oxide forms the hydroxide, 
 which is analogous to calcium hydroxide. It is more 
 easily soluble in water than the latter. 
 
 Strontium nitrate, Sr(NO 3 ) 2 , is made in considerable 
 quantity for the purpose of preparing a mixture which, 
 when burned, gives a red light (red-fire, Bengal-fire). It 
 is easily made by dissolving strontianite or strontium 
 carbonate in nitric acid. 
 
 Strontium sulphate, SrSO 4 , occurs in nature in beauti- 
 ful crystals as the mineral celestite. It is formed when a 
 soluble sulphate is added to a solution of a strontium 
 salt. In solubility it lies between calcium sulphate and 
 barium sulphate. 
 
BARIUM. 
 
 547 
 
 BARIUM, Ba (At. Wt. 136.39). 
 
 Occurrence and Preparation. Barium occurs in nature 
 in the same forms of combination as strontium, viz., as 
 the carbonate, BaCO 3 , in witherite ; and as the sulphate, 
 BaSO 4 , in barite or heavy spar. It is prepared by elec- 
 trolysis of the molten chloride. 
 
 Properties. It closely resembles calcium and stron- 
 tium, being a yellow metal, which is oxidized by contact 
 with the air and readily decomposes water at the ordi- 
 nary temperature. 
 
 Barium Chloride, BaCl 2 + 2H 2 O, is prepared by dissolv- 
 ing barium carbonate in hydrochloric acid. It dissolves 
 easily in water, but not as easily as the chlorides of 
 strontium and calcium. The order of solubility, begin- 
 ning with the most soluble, is, calcium, strontium, bar- 
 ium, the same as in the case of the sulphates. 
 
 Barium Hydroxide, Ba(OH) 2 , is formed by dissolving 
 barium oxide in water, just as calcium hydroxide is 
 formed by treating calcium oxide with water. In hot 
 water it is much more easily soluble than calcium hy- 
 droxide, and it is also more easily soluble in cold water. 
 As such a solution acts in the same general way as lime- 
 water, it is frequently used in the laboratory for the 
 purpose of detecting carbon dioxide, barium carbonate 
 being insoluble. Like lime-water, it has an alkaline 
 reaction. 
 
 Barium Oxide, BaO, is made by heating the nitrate, 
 as the carbonate is not easily decomposed by heat. The 
 most interesting property of the oxide is its power to 
 take up oxygen when heated to a dark red heat in the 
 air or in oxygen, when it forms 
 
 Barium Peroxide or Dioxide, BaO 2 . This peroxide is a 
 white powder which looks like the simple oxide. When 
 heated to a temperature a little higher than that re- 
 quired for its formation, it breaks down into barium 
 oxide and oxygen. The formation of the peroxide by 
 heating the oxide in the air, and the decomposition of 
 the peroxide at a higher heat, make it possible to extract 
 oxygen from the air and to obtain it in the free state. 
 
548 INORGANIC CHEMISTRY. 
 
 This method of preparing oxygen on the large scale 
 from the air was referred to under Oxygen. It is stated 
 that the oxide improves with use. Specimens which 
 have been in use for two years are said to be as 
 efficient as at first. When a solution of hydrogen di- 
 oxide, H 2 O 2 , is added to a solution of barium hydroxide, 
 a precipitate is formed which has the composition BaO 9 
 -f- 8H 2 O. When filtered and put in a vacuum over sul- 
 phuric acid, it loses all its water and leaves behind pure 
 dioxide. The dioxide is a convenient starting-point in 
 the preparation of hydrogen dioxide. It is only necessary 
 to treat it with hydrochloric acid in order to make a solu- 
 tion of hydrogen dioxide. The solution made in this 
 way, however, contains barium chloride. To make a 
 solution containing nothing but the dioxide, pure barium 
 peroxide is treated with dilute sulphuric acid, when in- 
 soluble barium sulphate is formed and the hydrogen 
 dioxide remains in solution : 
 
 BaO 2 + H 2 SO 4 = BaSO 4 + H 2 O 2 . 
 
 It is interesting to compare the action of hydrochloric 
 acid on barium peroxide and on the corresponding com- 
 pound of manganese. With the latter, as we have seen, 
 the reaction takes place as represented in this equation : 
 
 MnO a + 4HC1 = MnCl a + 2H a O + 01, ; 
 
 while with barium peroxide the reaction takes place 
 thus: 
 
 BaO a + 2HC1 = BaCl 3 + H 3 O a - 
 
 It is probable that in the case of manganese dioxide 
 some intermediate reactions take place which are im- 
 possible in the other case. (See Manganese Dioxide.) 
 
 Barium Sulphide, BaS, is made as calcium sulphide is, 
 by reducing the sulphate by heating with charcoal. It 
 is phosphorescent, like the calcium compound. When 
 dissolved in water, it is decomposed, forming the hydro- 
 sulphide and hydroxide, thus : 
 
COMPOUNDS OF BARIUM. 549 
 
 2BaS + 2H 2 = Ba(SH) 2 + Ba(OH) 2 . 
 
 It will be remembered that thermochemical investiga- 
 tions have made it appear probable that similar reac- 
 tions take place when potassium and sodium sulphides 
 are dissolved in water. In the case of barium sulphide 
 the evidence is more tangible, for, on evaporating a so- 
 lution of this compound, both the hydrosulphide and 
 hydroxide crystallize out. 
 
 Barium Nitrate, Ba(NO 3 ) 2 , is easily soluble in water, 
 but difficultly soluble in acids, and is precipitated from 
 its solution in water by the addition of nitric acid. 
 When heated to a sufficiently high temperature, it is de- 
 composed, and barium oxide is left behind. 
 
 Barium Sulphate, BaSO 4 . This occurs in nature as 
 barite, or heavy spar, and is precipitated when a soluble 
 sulphate or sulphuric acid is added to a solution of a 
 barium salt. It is insoluble in water ; when freshly pre- 
 cipitated, it is easily soluble in concentrated sulphuric 
 acid. It is artificially prepared for use as a pigment 
 and is known as permanent ivhite. On account of its in- 
 solubility it is much used in chemical analysis for the 
 purpose of detecting and estimating sulphuric acid. It 
 differs markedly from calcium and strontium sulphate, in 
 the fact that, when treated with a solution of ammonium 
 carbonate, it is not converted into the carbonate, whereas 
 calcium and strontium sulphates are by this means com- 
 pletely converted into the carbonates. This fact is taken 
 advantage of in analysis. There are other differences, 
 which will be stated at the end of this chapter. 
 
 Barium Carbonate, BaCO 3 , occurs in nature as witherite, 
 and is made pure by adding ammonium carbonate and a 
 little ammonia to a solution of barium chloride. The 
 carbonate usually found in the market is made by pre- 
 cipitating a solution of the crude sulphide with sodium 
 carbonate, or by heating together sodium carbonate and 
 natural barium sulphate, or heavy spar. Made in either 
 of these ways it contains alkaline carbonate, from which 
 it is impossible to separate it by washing. The carbonate, 
 like the other salts of barium, is poisonous. It has the 
 power to unite, and form insoluble compounds, with me- 
 
550 INORGANIC CHEMISTRY. 
 
 tallic oxides of the formula M 2 O 3 , as, for example, ferric 
 oxide, Fe 2 O 3 , and is used in analytical operations for the 
 purpose of separating iron from other metals, like man- 
 ganese, which are not precipitated by it. 
 
 Phosphates of Barium. The phosphates of barium cor- 
 respond in general to those of calcium. When ordinary 
 sodium phosphate and ammonia are added to a solution 
 of a barium salt, normal or tertiary phosphate is pre- 
 cipitated : 
 
 3BaCl a + 2HNa 3 PO 4 4- 2NH 3 = Ba 3 (PO 4 ) 2 + 4NaCl + 2NH 4 C1. 
 
 When sodium phosphate alone is added, the first reac- 
 tion which takes place is that represented in the equa- 
 tion 
 
 4BaCl, + 4HNa a P0 4 = BaH 4 (PO 4 ) 2 + Ba 3 (P0 4 ), + SNaCl. 
 
 The precipitate is the tertiary phosphate, while the 
 primary phosphate is in solution. On standing, the solu- 
 ble salt acts upon the insoluble one, forming the second- 
 ary phosphate thus : 
 
 BaH 4 (PO,), + Ba 3 (P0 4 ), = 4HBaPO, 
 
 Reactions which are of Special Value in Analysis. The 
 sulphates of calcium and strontium are completely con- 
 verted into the carbonates by contact with a solution 
 of ammonium carbonate in ammonia. The sulphate of 
 barium is not changed in this way. Consequently, if a 
 mixture of the three sulphates is treated with ammoni- 
 um carbonate, those of calcium and strontium will be 
 converted into carbonates, while that of barium will re- 
 main unchanged. By filtering, washing with water, and 
 treating with dilute nitric or hydrochloric acid, the car- 
 bonates will be dissolved, while the sulphate will not. 
 If nitric acid is used, the solution may be evaporated to 
 dryness and treated with a mixture of alcohol and ether. 
 Calcium nitrate will dissolve ; strontium nitrate will not. 
 
 Fluosilicic acid produces a precipitate of barium fluo- 
 silicate, BaSiF a , in solutions of barium salts. This is 
 
GLUCINUM. 551 
 
 insoluble in a mixture of alcohol and water, and difficultly 
 soluble in water. The corresponding salts of calcium 
 and strontium are soluble. 
 
 Calcium sulphate solution produces a precipitate in a 
 solution of a strontium salt or of a barium salt, but not 
 in one of a calcium salt. 
 
 Strontium sulphate solution precipitates barium sul- 
 phate from a solution of a barium salt, but forms no pre- 
 cipitate in a solution of a strontium salt. 
 
 When boiled with a solution of one part of sodium 
 carbonate and three parts of sodium sulphate, the sul- 
 phates of strontium and calcium are completely con- 
 verted into carbonates, while the sulphate of barium 
 remains unchanged. 
 
 Barium chloride is insoluble in absolute alcohol ; cal- 
 cium chloride is easily soluble ; while strontium chloride 
 dissolves in warm absolute alcohol. 
 
 Ammonium oxalate, (NH 4 ) 2 C 2 4 , produces precipitates 
 of the oxalates in solutions of calcium, barium, and 
 strontium. Only the calcium salt is insoluble in dilute 
 acetic acid. 
 
 Potassium dichromate, K 2 Cr 2 O 7 , precipitates barium 
 chromate, BaCrO 4 . The corresponding salts of calcium 
 and strontium are soluble in water. Barium chromate 
 is easily soluble in hydrochloric or nitric acid. 
 
 All three elements of the group give colored flames 
 which have characteristic spectra. Calcium compounds 
 color the flame reddish yellow ; strontium compounds give 
 an intense red ; and barium compounds a yellowish green 
 color. The spectra are more complicated than those of the 
 elements of the potassium group, but each one contains 
 highly characteristic lines which are easily recognized. 
 
 MAGNESIUM SUB-GROUP. 
 GLUCINUM, Gl (At. Wt. 9.01). 
 
 Occurrence and Preparation The principal form in 
 which the element glucinum occurs in nature is in the 
 mineral beryl, which is a silicate of aluminium and glu- 
 cinum of the formula Al 2 Gl 3 (SiO 3 ) 6 . Emerald has the 
 
552 INORGANIC CHEMISTRY. 
 
 same composition, but is colored green by the presence 
 of a little chromic oxide. The element can be isolated 
 by decomposing the chloride by heating it with potas- 
 sium or sodium. 
 
 Properties. The statements concerning the properties 
 of glucinum, made by those who have prepared it in 
 different ways, differ somewhat from one another, evi- 
 dently in consequence of the fact that it has not gener- 
 ally been pure. It has a metallic lustre. When heated 
 in the flame of the blow-pipe it becomes covered with a 
 thin layer of oxide, which prevents further action ; it 
 dissolves readily in hydrochloric and sulphuric acids, 
 but only with difficulty in nitric acid. It is dissolved 
 by potassium hydroxide, forming in all probability a 
 glucinate of the composition Gl(OK) a : 
 
 Gl + 2KOH = Gl(OK) a + H 2 . 
 
 The specific heat of glucinum at ordinary temperature 
 is 0.425. This multiplied by the atomic weight 9.01 gives 
 3.83 instead of 6.24. But the analysis and the determina- 
 tion of the specific gravity of the vapor of the chloride 
 show that it has the formula G1C1,, the atomic weight of 
 glucinum being 9.01. At 257 C., however, the specific 
 heat of glucinum is 0.582, and this multiplied by 9.01 
 gives 5.24. At ordinary temperature, therefore, glu- 
 cinum, like carbon, boron, and silicon, is an exception to 
 the law of Dulong and Petit, while at a higher tempera- 
 ture, like the elements named, it conforms to the law. 
 
 Compounds of Glucinum. The compounds of glu- 
 cinum differ in many respects from those of the group 
 calcium, barium, strontium. The hydroxide is entirely 
 insoluble in water ; the sulphate is easily soluble in 
 water ; the chloride is completely decomposed when its 
 water solution is evaporated to dryness, the products 
 being hydrochloric acid and glucinum oxide. It shows 
 a marked tendency to form basic salts. 
 
 Glucinum Chloride, G1C1 2 , is formed by the action of 
 chlorine on glucinum, and more easily by treating a mix- 
 ture of glucinum oxide and carbon with chlorine, the 
 reaction being similar to that employed in making sili- 
 
COMPOUNDS OF OLUCINUM. 553 
 
 con chloride (see p. 413) and boron chloride (see p. 352). 
 It is volatile, and it is therefore possible to determine 
 the specific gravity of its vapor. This has been done, 
 with the result of showing its molecular weight to be 
 79.9. Taking this fact into consideration, together with 
 the percentage composition of the compound, the con- 
 clusion is justified that the atomic weight of glucinum is 
 9.01. For a long time it was thought to be 13.65, with 
 which figure the specific heat, 0.425, is in accordance ; 
 for 13.65 X 0.425 = 5.79, but the evidence furnished by 
 the specific gravity of the vapor of the chloride is re- 
 garded as conclusive in favor of the atomic weight 9.01. 
 
 Glucinum Hydroxide, G1(OH) 2 , is thrown down as a 
 precipitate when a soluble hydroxide is added to a solu- 
 tion of a glucinum salt : 
 
 G1SO 4 + 2NaOH = Gl(OH), + Na 2 SO 4 . 
 
 It is a white, gelatinous mass, which is soluble in potas- 
 sium and sodium hydroxides and in ammonia, so that, 
 after precipitation from glucinum salts by these reagents, 
 it redissolves. This solution is due to the formation of 
 glucinates of the formula G1(OM) 2 : 
 
 Gl(OH), + 2NaOH = Gl(ONa) 2 + 2H 2 O. 
 
 When sufficiently diluted with water, the potassium and 
 sodium salts are completely decomposed, and the hy- 
 droxide reprecipitated. This is an illustration of mass 
 action, a large quantity of water effecting a decomposition 
 which a small quantity does not effect. The power of 
 the hydroxide to form salts with the strong bases shows 
 that it has slight acid properties. The hydroxides of cal- 
 cium, barium, and strontium do not possess this power. 
 Glucinum Sulphate, G1SO 4 , is formed by dissolving 
 glucinum hydroxide in dilute sulphuric acid, and has 
 the composition G1SO 4 -)- 4H 2 O when crystallized from 
 water. When a solution of this salt is heated with glu- 
 cinum hydroxide, basic salts are formed, of which the 
 following are examples : G1 2 SO 5 and G1 3 SO 6 . The first 
 of these is to be regarded as derived from a hydroxide 
 of the formula HO-G1-O-G1-OH, by neutralization 
 
554 INORGANIC CHEMISTRY. 
 
 with sulphuric acid, as represented in the formula 
 
 r^i o 
 
 O<Qi_Q>SO 2 ; the second from a hydroxide of the 
 
 formula HO-G1-O-G1-O-G1-OH, by neutralization with 
 sulphuric acid, as represented in the formula 
 
 Gl-O ) 
 "<G1 I SO, 
 
 <G1-OJ 
 
 With potassium sulphate, glucinum sulphate forms a 
 double salt of the formula K 5 SO,. G1SO, + 2H,O, 
 
 , (HO),SO< K 
 
 or X>CH, or possibly X>G1. 
 
 Glucinum Carbonate, G-1CO 3 . When a slight excess of 
 sodium carbonate is added to a solution of glucinum 
 sulphate, a basic carbonate of the formula G1 3 CO 5 is 
 formed. This is similar to the second of the above-men- 
 tioned basic sulphates. It is to be regarded as derived 
 from the hydroxide HO-G1-O-G1-O-G1-OH by neu- 
 tralization with carbonic acid, as represented in the 
 
 formula XG1 CO. 
 
 Weak Basic Character of Glucinum. The power of 
 glucinum hydroxide to form salts with strong bases, 
 such as potassium and sodium hydroxides, which was re- 
 ferred to above, shows that the hydroxide has slight acid 
 properties. At the same time, as we should expect, its 
 basic properties are weaker than those of the other base- 
 forming elements thus far considered. This is shown in 
 the ready formation of basic salts, such as the basic sul- 
 phates and basic carbonates mentioned. The strongest 
 bases do not readily form basic salts, but are, on the other 
 hand, more competent to form stable acid salts. Thus, 
 potassium and sodium form acid carbonates ; calcium 
 appears to form an extremely unstable acid carbonate, 
 but preferably all the members of the calcium group 
 form normal carbonates of the general formula MCO 3 ; 
 
MAGNESIUM. 555 
 
 glucinum, however, and, as we shall see, magnesium, 
 preferably form basic carbonates. We shall see, further, 
 that the members of the next family, of which aluminium 
 is the principal one, form only extremely unstable com- 
 pounds with carbonic acid, their basic properties not 
 being sufficiently strong to hold them in combination 
 with the weak acid, except apparently at a very low 
 temperature. 
 
 This resemblance to the acid-forming elements is 
 shown by glucinum also by the ease with which its 
 chloride is decomposed into the oxide and hydrochloric 
 acid when its water solution is evaporated to dryness. 
 This reaction does not take place in the case of sodium 
 and potassium at all, nor with barium and strontium. 
 With calcium it takes place to a slight extent, but with 
 glucinum it is complete, as it is with the similar metal 
 magnesium. In general, the more acidic the element 
 the more easily is its chloride decomposed in this way. 
 
 MAGNESIUM, Mg (At. Wt. 24.10). 
 
 Occurrence. Magnesium occurs very abundantly in 
 nature, though by no means as abundantly as calcium. 
 Among the widely distributed minerals which contain 
 the element are magnesite, which is the carbonate, 
 MgCO 3 ; dolomite, a double carbonate of magnesium 
 and calcium ; serpentine, talc, soapstone, meerschaum, 
 hornblende, all of which contain magnesium silicates. 
 Further, the metal is found in solution in many spring- 
 waters in the form of the sulphate, or, as it is called, 
 Epsom salt. Kainite is a sulphate and chloride of the 
 composition expressed by the formula 
 
 K 2 S0 4 .MgS0 4 .MgCl 2 + 6H,0 ; 
 
 kieserite is magnesium sulphate, MgSO 4 -\- H 2 O ; car- 
 nallite is a double chloride, KMgCl 3 + 6H 2 O. 
 
 Magnesium compounds are contained in the soil in 
 consequence of the decomposition of minerals contain- 
 ing it. It is to some extent taken up by the plants, and 
 
556 INORGANIC CHEMISTRY. 
 
 subsequently into the animal body. It is found in the 
 bones and in the blood in small quantities. 
 
 Preparation. The metal can be made by the electrol- 
 ysis of its chloride, but is most conveniently made by 
 decomposing the chloride by means of sodium. It is 
 now manufactured in considerable quantity by this 
 method. The operation consists in bringing together 
 dry magnesium chloride, fluor-spar, and sodium in cer- 
 tain proportions, and heating to a high temperature in a 
 crucible. The metal is purified by distillation. Instead 
 of using the chloride, which it is difficult to prepare dry 
 in large quantity, the double chloride of magnesium and 
 potassium, KMgCl 3 or MgCl 2 .KCl, is frequently used. 
 
 Properties. It is a silver-white metal with a high 
 lustre. In the air it changes only slowly, but it gradu- 
 ally becomes covered with a layer of the hydroxide. At 
 ordinary temperatures magnesium does not decompose 
 water ; at 100 it decomposes it slowly. When heated 
 above its melting-point in oxygen or in the air, it takes 
 fire and burns with a bright flame, forming the white 
 oxide. The light of the flame is very efficient in produc- 
 ing certain chemical changes, such as those involved in 
 photography, when a permanent impression is made by 
 the light upon a sensitive plate. It has also the power 
 to cause hydrogen and chlorine to combine just as the 
 sunlight and the electric light do. 
 
 Applications. The principal use to which magnesium 
 is put is for the purpose of producing a bright light, as for 
 photographing in spaces to which the sunlight does not 
 have access, and for signaling. It is also used to some 
 extent as an ingredient of materials employed in making 
 fireworks. 
 
 Compounds of Magnesium. The compounds of mag- 
 nesium present a general resemblance to those of gluci- 
 iium. As the element is much more abundant in nature, 
 its compounds have been studied more extensively. Its 
 acid properties are somewhat weaker, and its basic 
 properties stronger, than those of glucinum. Its hy- 
 droxide does not form salts with the hydroxides of 
 potassium and sodium. On the other hand, its chlo- 
 
MAGNESIUM CHLORIDE. 557 
 
 ride is decomposed when its water solution is evap- 
 orated to dryness. The hydroxide is very slightly sol- 
 uble in water, and this solution has a slightly alkaline 
 reaction. With carbonic acid it forms basic carbon- 
 ates similar to those formed by glucinum. On the 
 other hand, it does not readily form basic salts with 
 sulphuric acid. In character, it is plainly more closely 
 allied to the members of the calcium group than gluci- 
 num is. 
 
 Magnesium Chloride, MgCl 2 . This salt, as has been 
 stated, occurs in nature. It is easily formed by dissolv- 
 ing magnesium oxide or carbonate in hydrochloric acid. 
 On evaporating at as low a temperature as possible, 
 there finally crystallizes out of the very concentrated 
 solution, a salt of the composition MgCl 2 -f- 6H 2 O, anal- 
 ogous to crystallized calcium chloride, CaCl 2 + 6H 2 O, 
 and strontium chloride, SrCl 2 -f- 6H 2 O. When this 
 crystallized salt is heated for the purpose of driving 
 off the water, it is completely decomposed in accordance 
 with the following equation : 
 
 MgCl 2 + H 2 = MgO + 2HC1. 
 
 It is most conveniently prepared in the dry form by 
 first making ammonium-magnesium chloride, and de- 
 composing this by heat. For this purpose, a solution of 
 ammonium chloride is added to a solution of magnesium 
 chloride and the whole evaporated to dryness. There is 
 formed in the solution the double salt of the composi- 
 tion NH 4 MgCl 3 (MgCl 2 .NH 4 Cl), which can be evaporated 
 to complete dryness. When perfectly dry, this double 
 salt breaks down into magnesium chloride and ammo- 
 nium chloride, if heated to a sufficiently high tempera- 
 ture. The ammonium chloride under these circum- 
 stances is volatilized, and the magnesium chloride re- 
 mains behind in the molten condition. 
 
 The chloride is a white, crystalline mass which de- 
 liquesces in the air. At a bright red heat, it is volatile 
 and can be distilled in an atmosphere of hydrogen. It 
 dissolves in water with marked evolution of heat. It 
 -combines readily with the chlorides of potassium, so- 
 
558 INORGANIC CHEMISTRY. 
 
 dium, and ammonium, forming crystallizing compounds 
 of the formulas KMgCl 3 , NaMgCl,, and NH.MgCl,, which 
 may be regarded as formed by the combination of one 
 molecule of magnesium chloride with one molecule of each 
 of the other chlorides. A second compound with potas- 
 sium chloride, of the formula K 2 MgCl 4 , is also known. 
 It seems probable that the latter is analogous to the po- 
 
 OK 
 
 tassium compound of glucinum of the formula GI 
 
 It corresponds to an oxygen compound of the formula, 
 
 OK 
 
 , which, however, does not seem to be formed. 
 
 If in this compound we imagine each of the two oxygen 
 atoms to be replaced by two chlorine atoms, the com- 
 
 //^l \_T7" 
 
 pound would have the formula Mg < >QI 2 <_g- . The exist- 
 
 ence of two double chlorides of magnesium and potas- 
 sium is suggested by what has been said regarding 
 compounds of this kind (see p. 465). One of these 
 
 Cl 
 would be represented by the formula Mg < / ^ -, yrr, the 
 
 other by the above formula. Both are known. 
 Further, magnesium bromide forms the salt K 2 MgBr t 
 
 or 
 
 Magnesium Oxide, MgO. This compound is commonly 
 called magnesia. A fine white variety which is known 
 as magnesia usta, is made by heating precipitated basic 
 magnesium carbonate. It is a white, loose powder, 
 which is very difficultly soluble in water, forming with it 
 the hydroxide, Mg(OH) 2 , which is also very difficultly 
 soluble. Magnesia is used, in medicine, as an applica- 
 tion to wounds, and, mixed with a solution of ferric sul- 
 phate, as an antidote in cases of poisoning by arsenic. 
 As magnesia is infusible, it is used to protect vessels 
 which are subjected to a high temperature. When 
 mixed with water and allowed to lie in the air, it be- 
 comes very hard. Mixtures of magnesia with sand also 
 have this property, and are used as hydraulic cements. 
 It is used, further, in the manufacture of fire-bricks. 
 
MAGNESIUM CARBONATE. 559 
 
 Magnesium Sulphate, MgSO 4 . The mineral kieserite, 
 which occurs in Stassfurt, has the composition expressed 
 by the formula MgSO 4 + H 2 O ; or, more probably, this 
 
 oH 
 
 should be written (HO) 2 MgSO 3 , or OS - , in which 
 
 it appears as a derivative of the acid SO(OH) 4 . The 
 salt MgSO 4 + 7H 2 O (or H 2 MgSO B + 6H 2 O), also occurs 
 in nature. It is this variety which is generally obtained 
 when a solution of magnesium sulphate is evaporated to 
 crystallization. It crystallizes in large rhombic prisms, 
 or, if rapidly deposited from very concentrated solutions, 
 in small, needle-shaped crystals. At ordinary tempera- 
 tures, 100 parts of water dissolve 125 parts of the salt. 
 The water solution has a bitter, salty taste. When 
 heated, it readily loses 6 molecules of water, but it re- 
 quires a temperature of over 200 to drive off the last 
 molecule. This has led to the belief that the salt with 
 one molecule has the constitution above given, being a 
 derivative of the acid SO(OH) 4 . 
 
 Magnesium sulphate finds extensive application. It is 
 used in medicine as a purgative, and is known as Ep- 
 som salt, for the reason that it is contained in the water 
 of Epsom springs ; it is used further in the manufac- 
 ture of sodium sulphate and potassium sulphate, and as 
 a fertilizer in place of gypsum, it having been shown to 
 be advantageous in some cases. Its chief use is for 
 loading cotton fabrics. 
 
 Magnesium sulphate forms double salts with other 
 sulphates ; as, for example, one with potassium sulphate, 
 similar to that formed by glucinum sulphate (see p. 554). 
 The constitution of the double sulphate of magnesium 
 and potassium is probably that expressed by the for- 
 
 o,s<g K 
 
 mula n > Mg. 
 
 Magnesium Carbonate, MgCO 3 . Like glucinum, mag- 
 nesium shows a marked tendency to form basic salts 
 
560 INORGANIC CHEMISTRY. 
 
 with carbonic acid. When a neutral magnesium salt is 
 treated with a soluble carbonate, a basic carbonate is 
 precipitated, the composition of which varies according 
 to the conditions under which it is prepared. The salt 
 obtained by adding an excess of sodium carbonate to a 
 solution of magnesium sulphate has the composition 
 Mg 3 (OH) 2 (CO 3 ) 2 It is derived from three molecules 
 of magnesium hydroxide and two of carbonic acid, as is 
 
 00< 0-Mg-OH 
 
 more clearly shown in the formula n >Mg . The 
 
 c*c\ *>** 
 
 < O-Mg-OH 
 
 salt which is manufactured on the large scale is more 
 complicated than this, being derived from four molecules 
 of magnesium hydroxide and three of carbonic acid. It 
 is known as magnesia alba. It is this form of the car- 
 bonate which is used in the preparation of magnesia usta. 
 Normal magnesium carbonate, MgCO 3 , occurs in nature 
 as magnesite. It crystallizes in the same form as cal- 
 cium carbonate, or is isomorphous with it. It is insolu- 
 ble in water, but like calcium carbonate it dissolves in 
 water containing carbon dioxide in solution. From this 
 solution crystals having the composition MgCO 3 -f~~ 
 3H 2 O and MgCO 3 -f- 5H 2 O are deposited under the 
 proper conditions. 
 
 Phosphates. The conduct of the phosphates of mag- 
 nesium is very similar to that of the phosphates of cal- 
 cium. All three are known ; and of these only the primary 
 salt is soluble in water. A salt much utilized in analysis 
 is ammonium-magnesium phosphate, Mg (NH 4 )PO 4 . This 
 is difficultly soluble in water, and may therefore be used 
 either for the purpose of detecting magnesium or phos- 
 phoric acid. In order to produce this salt, ammonia 
 and some ammonium salt, together with a soluble mag- 
 nesium salt, must be added to a soluble phosphate. 
 If ammonia alone were added to a solution containing a 
 magnesium salt, magnesium hydroxide would be precipi- 
 tated : 
 
 MgS0 4 + 2NH 4 OH = Mg(OH) 2 + (NH 4 ) 2 SO 4 . 
 
BORATES SILICATES MAGNESIUM SILICIDE. 561 
 
 ' ' 
 
 With ammonium salts, however, magnesium salts form 
 compounds, which are not decomposed on the addition 
 of ammonia. When a soluble phosphate is added, the 
 difficultly soluble ammonium-magnesium salt is thrown 
 down. When heated, this salt loses ammonia, then 
 water, and is converted into magnesium pyrophosphate : 
 
 Mg(NH 4 )P0 4 = MgHP0 4 + NH 3 ; 
 SMgHPO, = Mg,PA + H,0. 
 
 The corresponding salt of arsenic acid, Mg(NH 4 )AsO 4 , is 
 very similar to the phosphate, and on account of its in- 
 solubility it is also used in chemical analysis. 
 
 Borates. A borate of magnesium together with mag- 
 nesium chloride occurs in nature, and is known as bora- 
 cite. It has the composition expressed by the formula 
 2Mg 3 B 8 O 15 + MgCl 2 . The borate, Mg 3 B 8 O 15 , is derived 
 from the acid, H 6 B 8 O 15 , which is related to normal boric 
 acid, as is shown by the equation 
 
 Silicates. The simplest silicate of magnesium found 
 in nature is olivine, which is represented by the formula 
 Mg 2 SiO 4 . It is the neutral salt of normal silicic acid. 
 
 Serpentine is derived from the acid, O<i(Qjj\ 3 > and has 
 
 the composition Mg 3 Si 2 O 7 + 2H 2 O. 
 
 Magnesium Silicide, Mg 2 Si, is made by heating together 
 magnesium chloride, sodium fluosilicate, sodium chlo- 
 ride, and sodium. Under these circumstances the so- 
 dium sets magnesium free from the chloride, and silicon 
 from the fluosilicate. Both unite to form magnesium 
 silicide. When treated with hydrochloric acid it gives 
 silicon hydride, SiH 4 , and hydrogen : 
 
 M ga Si + 4HC1 = 2MgCl 2 + SiH 4 . 
 
 The liberation of hydrogen is due to the presence of 
 an excess of magnesium. 
 
 Reactions of Magnesium Salts which are of Special 
 
562 INORGANIC CHEMISTRY. 
 
 
 
 Value in Chemical Analysis. Soluble hydroxides (KOH, 
 NaOH, NH 4 OH) precipitate magnesium hydroxide. If 
 ammonium chloride is present ammonia does not pre- 
 cipitate the hydroxide. 
 
 Di-sodium phosphate with ammonia and ammonium 
 chloride precipitates ammonium-magnesium phosphate 
 from the solution of a magnesium salt. 
 
 Sodium and potassium carbonates precipitate basic mag- 
 nesium carbonate. 
 
 ERBIUM, E (At. Wt. 165.06). 
 
 General. As regards the position of erbium in the pe- 
 riodic system, a final statement cannot as yet be made. 
 According to its atomic weight, assuming it to be 165.06, it 
 falls in the second family. On the other hand, the com- 
 position of its compounds seems to indicate rather that it 
 belongs in the third family, as it resembles aluminium 
 in some respects. It occurs in some rare minerals, as 
 cerite, gadolinite, euxenite, and orthite, which are found 
 in Sweden and Greenland. It is always accompanied 
 by other rare metals, a few of which have been studied 
 with care. Among these may be mentioned lanthanum, 
 cerium, didymium, and scandium. These metals will 
 be treated of in later chapters. It need only be said 
 further in regard to erbium, that our knowledge con- 
 cerning it is as yet quite imperfect, and the cause of 
 this is to be found in the fact that the minerals in which 
 it occurs are exceedingly complex, and it is therefore 
 very difficult to separate the various metals present. It 
 appears that the formula of the oxide of erbium is E 2 O 3 . 
 If this is so, it is in this respect like aluminium oxide, 
 
 AlA. 
 
 It appears probable that the oxide of erbium that is 
 generally obtained from the complex minerals named is 
 a mixture of oxides of rare elements. There is good evi- 
 dence of the presence of thulium and of holmium, though 
 it is still an open question whether these as well as erbium 
 may not be capable of further decomposition. 
 
CHAPTER XXVII. 
 
 ELEMENTS OF FAMILY III. GROUP A : 
 ALUMINIUM SCANDIUM YTTRIUM YTTERBIUM- 
 SAMARIUM HELIUM. 
 
 General. There is in some respects a resemblance 
 between boron and the principal member of this group ; 
 but as boron acts almost exclusively as an acid-forming 
 element, it was taken up in connection with the elements 
 of Family V, Group B, or the nitrogen group. Atten- 
 tion was, however, called to the fact that the analogy 
 between these elements and boron is but slight. The 
 points of resemblance between boron and the members 
 of Family III, Group A, will be pointed out below. 
 The principal member of this group is aluminium. The 
 others are all rare, and some have been but imperfectly 
 studied, owing to serious difficulties in the way of ob- 
 taining their compounds in pure condition. They are 
 trivalent in their compounds, the general formulas being 
 such as the following : 
 
 MCI,, M(OH) a , M(N0 3 ) 3 , M,(S0 4 ),, M/CO,)* MPO 4) etc. 
 
 Aluminium oxide is weakly basic, and somewhat acidic, 
 though less so than boron. Aluminium hydroxide has 
 the power to neutralize most acids, and also to form salts 
 with strong bases. Boron oxide, on the other hand, has 
 scarcely any basic properties, though it does form a few 
 extremely stable compounds, in which the boron replaces 
 the hydrogen of acids. (See Boron Phosphate, p. 356.) 
 
 ALUMINIUM, Al (At. Wt. 26.91). 
 
 Occurrence. Aluminium is an extremely important 
 element in nature and in the arts. It occurs very 
 
 (563) 
 
564 INORGANIC CHEMISTRY. 
 
 widely distributed, and very abundantly in many different 
 forms of combination. Among them are feldspar, mica, 
 cryolite, bauxite. Feldspar is a silicate of aluminium 
 and potassium of the formula AlKSi 3 O 8 . Mica is a gen- 
 eral name applied to a large number of minerals which 
 are silicates of aluminium and some other metal, as po- 
 tassium, lithium, magnesium, etc. The simplest form 
 of mica is that represented by the general formula 
 KAlSiO 4 , according to which the mineral is a salt of 
 orthosilicic acid, Si(OH) 4 . Cryolite is a double fluoride 
 of aluminium and sodium, or the sodium salt of fluo- 
 aluminic acid, Na 3 AlF 6 . Bauxite is a hydroxide of 
 aluminium in combination with a hydroxide of iron. 
 Besides in the above forms, aluminium occurs in the 
 products of decomposition of minerals. One of the most 
 important of these is clay, which is found in all condi- 
 tions of purity from the white kaoline to ordinary 
 dark-colored clay. Kaoline is the aluminium salt of 
 orthosilicic acid of the formula Al 4 (SiO 4 ) 3 -f- 4H 2 O. Alu- 
 minium silicate is found in all soils, but is not taken up 
 by plants, and does not find entrance into the animal 
 body. The name aluminium has its origin in the fact 
 that the salt alum was known at an early date, and the 
 metal was afterwards isolated from it. 
 
 Preparation. The preparation of aluminium on the 
 large scale presents a problem of the highest importance 
 to the human race. The element has properties which 
 adapt it to many uses to which iron is put, and for many 
 purposes it has many advantages over iron. Further, we 
 are supplied by nature with unlimited quantities of the 
 compounds of aluminium, which are distributed every- 
 where over the earth. While, however, iron, lead, tin,, 
 copper, and other metals can be isolated from their 
 natural compounds without serious difficulty, aluminium, 
 which is more abundant than any of them, and in many 
 respects more valuable than any of them, is locked in 
 its compounds so firmly, that it is only by comparatively 
 complicated and expensive methods that it can be iso- 
 lated ; and up to the present it cannot be made at a 
 price sufficiently low to bring it into common use. At 
 
ALUMINIUM. 565 
 
 the same time work is constantly in progress with refer- 
 ence to this important practical problem, and it seems 
 probable that through a thorough study of the Jaws of 
 chemistry some method for the cheap preparation of alu- 
 minium on the large scale will eventually be discovered. 
 The first method devised for the preparation of alu- 
 minium on the large scale consisted in heating aluminium 
 chloride with sodium. The chloride was heated to boil- 
 ing in a retort ; the vapor passed through a vessel contain- 
 ing pieces of iron heated to redness, and then into a long 
 tube containing sodium. Instead of aluminium chloride, 
 the double chloride of aluminium and sodium, which is 
 more easily prepared in the dry condition, is now used. 
 The double chloride and cryolite are heated together 
 with sodium in a properly constructed furnace. It is, 
 further, possible to prepare aluminium by electrolysis 
 of the chloride or of the double chloride above men- 
 tioned ; and the oxide can be reduced by mixing it with 
 charcoal and passing the current from a powerful dy- 
 namo-machine through it. By the latter method an 
 alloy of aluminium and copper is now prepared, but the 
 preparation of aluminium alone by this method does not 
 appear to be entirely successful. New methods for the 
 preparation of the metal are constantly being devised, 
 and the price is constantly being lowered. The latest 
 method of promise consists in the electrolysis of alu- 
 minium oxide, in the form of corundum, in a bath of 
 molten cryolite contained in a carbon crucible. A large 
 number of patents have been issued, covering methods 
 for the preparation of aluminium; but these are fre- 
 quently so imperfectly described, and the evidence of 
 their value so unsatisfactory, that it is difficult to pass 
 judgment upon them. Until recently the commercial 
 preparation of aluminium has appeared to be intimately 
 connected with that of the commercial preparation of 
 sodium ; but, in view of the success of the electrolytic 
 method, this is no longer the case. 
 
 Properties. The color of aluminium is like that of 
 tin, and it has a high lustre. It is very strong, .and yet 
 malleable. It is lighter than most metals in common use, 
 
566 INORGANIC CHEMISTRY. 
 
 its specific gravity being 2.5 to 2.7 according to the con- 
 dition, while that of iron is 7.8, that of silver 10.57, that 
 of tin 7,3, and that of lead 11.37. It does not change in 
 dry or in moist air ; and in the compact form it does not 
 act upon water even at elevated temperatures. It melts 
 at about 700, which is higher than the melting-point of 
 zinc, and lower than that of silver. Hydrochloric acid 
 dissolves it with ease, forming aluminium chloride. At 
 the ordinary temperatures nitric and sulphuric acids do 
 not act upon it ; at higher temperatures, however, action 
 takes place, and the corresponding salts are formed. It 
 dissolves in solutions of the caustic alkalies, forming the 
 so-called aluminates. It reduces many oxides when 
 heated with them to a sufficiently high temperature ; 
 and is used in the preparation of boron and silicon. 
 
 Applications. The metal is used to a considerable 
 extent in the preparation of ornaments, and of useful 
 articles in which lightness is a matter of importance, as 
 in telescopes and opera-glasses. An alloy with a small 
 percentage of silver is used for the beams of chemical 
 balances. Aluminium bronze, which is an alloy with 
 copper, is also used quite extensively. It will be again 
 referred to under Copper. 
 
 Aluminium Chloride, A1C1 3 . When aluminium hydrox- 
 ide is dissolved in hydrochloric acid a solution of alu- 
 minium chloride is formed, and from this solution a 
 compound of the formula A1C1 3 + 6H 2 O can be obtained 
 in crystallized form. Like calcium and magnesium chlo- 
 rides, this salt is deliquescent. When heated to drive off 
 the water the salt conducts itself like magnesium chlo- 
 ride, but the decomposition into the oxide and hydro- 
 chloric acid takes place more easily than that of mag- 
 nesium chloride. The reaction is represented by the 
 equation 
 
 2A1C1 3 + 3H 2 = A1 3 3 + 6HC1. 
 
 The dry chloride is prepared by the same method as 
 that used in the preparation of silicon chloride and boron 
 chloride, viz., by passing chlorine over a heated mixture 
 of the oxide and carbon. The chloride, being volatile, 
 
ALUMINIUM CHLORIDE. 567 
 
 sublimes, and is deposited in the cool part of the vessel, 
 when pure, as a white laminated crystalline mass. Gen- 
 erally, however, it is more or less colored in consequence 
 of the presence of impurities. When exposed to the air 
 it attracts moisture and gives off hydrochloric acid. It 
 dissolves in water very easily, with a marked evolution of 
 heat, but, from what was said above, it is evident that it 
 cannot be obtained from this solution again by evapora- 
 tion. It is volatile without change. The specific gravity 
 of its vapor has been determined by different observers, 
 and, unfortunately, with different results. According to 
 Deville and Troost, it is such as to lead to the formula 
 A1 2 C1 6 . Quite recently, however, Nilson and Pettersson 
 have found it to correspond to that required by the for- 
 mula A1C1 3 , their determinations having been made at a 
 higher temperature than those of Deville and Troost. 
 Still later determinations by Crafts again lead to the 
 formula A1 2 C1 6 . Upon the basis of the determinations by 
 Deville and Troost, chemists have for some time past used 
 the formula Al a Cl 6 to represent the compound. Accord- 
 ing to this, aluminium would appear to be quadrivalent, 
 as represented in the following formula for the chloride : 
 Ck /Cl 
 
 C1-)A1-A1(C1 . On the other hand, in a compound 
 CK \C1 
 
 made by replacing the chlorine of this chloride by certain 
 organic groups the aluminium appears to be trivalent, as 
 represented in the formula A1(CH 3 ) 3 , in which the group 
 CH 3 , known as methyl, is univalent. Further, the posi- 
 tion of aluminium in the periodic system makes it appear 
 extremely probable that it is trivalent, and not quadriv- 
 alent. What, then, is the explanation of the discrepancy 
 above noted in the evidence regarding the constitution 
 of the chloride ? When we come to examine the conduct 
 of aluminium chloride towards the chlorides of other 
 metals, and find with what ease it forms double chlorides, 
 it seems not improbable that aluminium chloride itsolf, 
 at ordinary temperatures, and even in the form of vapor 
 at lower temperatures, may be a compound of the same 
 order as the double chlorides. It has been suggested 
 
568 INORGANIC CHEMISTRY. 
 
 that in these compounds chlorine is probably in com- 
 bination with chlorine, as fluorine is with fluorine in hy- 
 drofluoric acid, in such a way that two chlorine atoms 
 can exert a linking function between two other atoms. 
 
 /Cl 
 Just as there is a compound of the formula A1-C1 
 
 so it is possible that aluminium chloride may have the 
 constitution represented by the formula A1^(C1 2 )^)A1, 
 
 in which the aluminium is trivalent. By replacing the 
 chlorine in a compound of this constitution by groups 
 like methyl, which cannot exert the linking function, 
 the product would not be a double compound. Further, 
 by heating a compound of this constitution it would 
 probably dissociate into two molecules of the simple 
 compound A1C1 3 , and it would be this which comes into 
 play in chemical reactions. In view of the conflicting 
 state of the evidence and the plausibility of the above 
 explanation, the formula for aluminium chloride used 
 here is the simpler one. By means of it and similar for- 
 mulas for the other compounds of aluminium, the reac- 
 tions of the element can be expressed somewhat more 
 easily and probably just as truthfully as by means of the 
 more complicated formula. 
 
 Chloroaluminates or Double Chlorides of Aluminium 
 and Analogous Compounds. These compounds have 
 been repeatedly referred to, and but very little need be 
 added to what has already been said concerning them. 
 In general, the chloride, bromide, and iodide of aluminium 
 combine with the chlorides, bromides, and iodides of the 
 most strongly marked metals, such as potassium and 
 sodium. Those with potassium and sodium have the for- 
 
 XSL 
 mulas A1C1 3 .KC1 and AlCl 3 .NaCl, or probably Al(-Cl 
 
 \C1,)K 
 /Cl 
 and A1^-C1 . The fluoride forms two compounds 
 
 \Cl a )Na 
 with potassium fluoride and two with sodium fluoride. 
 
ALUMINIUM HYDROXIDE. 569 
 
 These have the composition represented by the for- 
 mulas A1F 3 .2KF, AlF 3 .2NaF, and A1F 3 .3KF, AlF 3 .3NaF, 
 
 X F 
 and the constitution expressed thus, A1(-(F 2 )K and 
 
 A1(~(F,)K . The tri-sodium fluoaluminate is the min- 
 
 eral cryolite, which occurs in such large quantity as to 
 be exported, and form the starting-point in the prepa- 
 ration of aluminium and even sodium compounds. A 
 method for making sodium carbonate from cryolite has 
 already been described. Its use in the preparation of 
 aluminium compounds will be taken up as far as may be 
 necessary in this chapter. 
 
 Besides the compounds with metallic chlorides, alu- 
 minium chloride also forms compounds with the chlorides 
 of the acid-forming elements. Such, for example, are 
 the compounds with sulphur tetrachloride and with 
 phosphorus pentachloride. These have the composition 
 represented by the formulas (A1C1 3 ) 2 SC1 4 and A1C1 3 .PC1 5 . 
 The latter may be the chlorine analogue of aluminium 
 phosphate, A1PO 4 . If the oxygen in the phosphate 
 should be replaced by an equivalent quantity of chlorine 
 the result would be a compound of the formula A1PC1 8 , 
 which is that of the above compound. These double 
 chlorides, like the chlorides of the acid-forming elements 
 in general, are easily decomposed by water, yielding the 
 corresponding oxygen compounds. A compound inter- 
 mediate between the oxygen and the chlorine compounds 
 is that formed by the combination of aluminium chloride 
 with phosphorus oxychloride, which is represented by 
 the formula A1POC1 6 , or A1C1 3 .POC1 3 . This may be re- 
 garded as aluminium phosphate, in which three of the 
 oxygen atoms have been replaced by six chlorine atoms. 
 
 Aluminium Hydroxide, A1(OH) 3 . Normal aluminium 
 hydroxide, A1(OH) 3 , occurs in nature as the mineral hy- 
 drargillite. It is precipitated from a solution of alu- 
 minium chloride by ammonia : 
 
 A1C1, + 3NH 4 OH = A1(OH) 3 + 3NH 4 C1. 
 
570 INORGANIC CHEMISTRY. 
 
 Obtained by precipitation it forms a gelatinous mass, 
 which is suggestive of starch-paste, and it is on this 
 account extremely difficult to wash it completely free 
 from the substances in the solution. It dries in the air, 
 forming a gummy substance which has the composition 
 A1(OH) 8 . When heated under proper conditions it loses 
 water, and forms the compound A1O 2 H : 
 
 A1(OH) 3 = A1O.OH + H 2 O. 
 
 This compound is found in nature as the mineral dias- 
 pore. If heated to a higher temperature it is converted 
 into the oxide, A1 2 O 3 : 
 
 2A1(OH) 3 = A1 2 3 + 3H 2 0. 
 
 In the conduct of the chloride and of the hydroxide 
 aluminium exhibits a certain resemblance to boron. The 
 acidic character of the latter is, however, more strongly 
 marked than that of the former. Boron chloride is more 
 easily decomposed by water than aluminium chloride, 
 and, as the decomposition takes place at the ordinary 
 temperature, the product is the hydroxide instead of the 
 oxide, as in the case of aluminium. The hydroxide, 
 B(OH) 3 , readily loses water and forms metaboric acid, 
 which in composition is analogous to diaspore ; and at a 
 higher temperature the oxide, B 2 O 3 , is formed. 
 
 Besides the normal hydroxide, A1(OH) 3 , and that of 
 the formula AIO(OH), there is a third one known. This 
 has the composition A1 2 O(OH) 4 , and, as is plain, is derived 
 from two molecules of the normal hydroxide by loss of 
 one molecule of water : 
 
 /OH TT/^ r\TT 
 
 H6 >A1 -- A1< OH + H >- 
 
 This has been obtained in solution ; or, rather, it has 
 been obtained by evaporation of a solution of hydroxide 
 made by continued boiling of a solution of basic acetate 
 of aluminium which decomposes into hydroxide and 
 acetic acid, the latter then evaporating. From this 
 solution, by evaporation in a water-bath, the above hy- 
 
ALUMINATES. 571 
 
 droxide is obtained. As already stated, bauxite is, in all 
 probability, a compound of this constitution combined 
 with a similar hydroxide of iron. A hydroxide of the 
 same composition is obtained when a solution of the 
 normal hydroxide in caustic soda is boiled with am- 
 monium chloride. The precipitate formed in this way is 
 not gelatinous, and, when dried, it has the composition 
 A1 2 0(OH), 
 
 The preparation of aluminium hydroxide from natural 
 compounds of the element is based upon the fact that 
 aluminium oxide forms with sodium a soluble compound, 
 and that this is decomposed by carbon dioxide with pre- 
 cipitation of the hydroxide. The sodium compound 
 formed has probably the composition Al(ONa) 8 , being a 
 salt of the normal hydroxide. When this is treated in 
 solution with carbon dioxide, the decomposition takes 
 place as represented in this equation : 
 
 2Al(ONa) 3 + 3CO 2 + 3H 2 O = 3Na 2 CO 3 + 2A1(OH) 3 . 
 
 When cryolite is ignited with lime, the products are 
 probably calcium fluoride, sodium oxide, and another 
 variety of sodium aluminate : 
 
 Na 8 AlF 6 + 3CaO = NaA10 2 + Na 2 O + 3CaF,. 
 
 When the mass is treated with water, the calcium fluor- 
 ide remains undissolved, while the sodium and aluminium 
 form the compound Al(ONa) 3 . This undergoes decompo- 
 sition, as above represented, when treated with carbon 
 dioxide. Two valuable products aluminium hydroxide 
 and sodium carbonate are thus obtained. 
 
 In order to prepare the hydroxide from bauxite, this 
 is heated to a high temperature with sodium carbonate. 
 Water extracts sodium aluminate, from which the hy- 
 droxide is precipitated by means of carbon dioxide. 
 
 Aluminium hydroxide forms the material. for the prep- 
 aration of aluminium salts ; as, the chloride, sulphate, 
 alum, etc. 
 
 Aluminates. When sodium or potassium hydroxide is 
 added to a solution of an aluminium salt, aluminium hy- 
 droxide is at first precipitated, but an excess of the re- 
 
572 INORGANIC CHEMISTRY. 
 
 agent used dissolves the precipitate. This action is the 
 same in character as that which takes place in the case 
 of glucinum, and is due to the acidic character of alu- 
 minium hydroxide. It is probable that in solution the 
 action with potassium and sodium hydroxides is of the 
 same kind as represented in the equations 
 
 A1(OH) 3 + 3KOH =A1(OK) 3 + 3H,O, and 
 A1(OH), + 3NaOH = Al(ONa) 3 + 3H 2 O. 
 
 On evaporating the solution of the potassium salt, how- 
 ever, the product obtained has the formula A1O.OK, 
 and is plainly the salt of the hydroxide A1O.OH, which 
 may be called meta-aluminic acid, to suggest its analogy 
 to metaboric acid, BO. OH. When aluminium hydroxide 
 and sodium carbonate are melted together, the salt 
 AlO.ONais formed, as has been shown by determining 
 the amount of carbon dioxide given off when a known 
 weight of the hydroxide is employed. When, however, 
 the solution of the hydroxide in caustic soda is evap- 
 orated, the salt Al(ONa) 3 is deposited. 
 
 These salts are very unstable, though their solutions 
 can be boiled without undergoing decomposition. Car- 
 bon dioxide decomposes them at once with precipitation 
 of aluminium hydroxide, as was stated in describing the 
 method for the preparation of the hydroxide from cryolite 
 and from bauxite. Similar salts are formed with calcium 
 and barium. Among them may be mentioned those of the 
 following formulas : Ca 3 (AlO 3 ) 2 , Ca(AlO 2 ) 2 , Ba 3 (AlO 3 ) 2 , and 
 Ba(AlO 2 ) 2 . The calcium salts are insoluble in water, and 
 some of them become hard in contact with water. They 
 are therefore of importance in the manufacture of hy- 
 draulic cements. The barium salts are soluble in water. 
 
 Many aluminates occur in nature, forming the import- 
 ant group of minerals known as the spinels. Of these, spi- 
 nel itself is the magnesium salt of the hydroxide A1O.OH, 
 
 and is represented by the formula ^Q *Q>Mg, or 
 
 Mg(AlO,) 2 . Chrysoberyl is the corresponding glucinum 
 salt G1(A1O 2 ) 2 ; and gahnite is the zinc salt Zn(AlO 2 ) 2 . 
 These salts are extremely stable, differing markedly in this 
 
ALUMINATES. 573 
 
 respect from those above referred to, which are made in 
 the laboratory. They are decomposed by heating them, 
 in finely powdered condition, with primary or acid po- 
 tassium sulphate, the action of which was described on 
 p. 499. As will be seen farther on, there are other salts 
 similar to the aluminates in structure which occur in 
 nature. Among these there may be mentioned here 
 chromic iron, or chromite, which is an iron salt of a hy- 
 droxide of chromium of the formula CrO.OH. The salt 
 is to be regarded as made up according to the formula 
 
 CrO O >Fe> or Fe ( Cr 3 - Further, magnetic oxide of 
 iron or magnetite, Fe 3 O 4 , is regarded as belonging to 
 the same group, and its constitution represented thus : 
 
 , or Fe(FeO 2 ) 2 ; and there is also a compound 
 
 of magnesium, ^ r\ 'o>Mg. For the sake of empha- 
 
 sizing these analogies, the formulas of the compounds 
 above mentioned are here presented in tabular form : 
 
 Potassium aluminate, . . A1O.OK 
 Sodium aluminate, . . . AlO.ONa 
 
 Calcium aluminate, . . . ~\]r\ *Q>Ca 
 Barium aluminate, . . . ' 
 
 Spinel, . . ..... . . . 
 
 Chrysoberyll, ..... AlO.'o >G1 
 
 Gahnite, ....... 
 
 Chromite, ...... ' 
 
 Magnetite, ...... FeO.'o >Fe 
 
 Magnesio-ferrite, .... 
 
 There is a highly instructive analogy between the 
 aluminates and the double chlorides and other similar 
 
574 INORGANIC CHEMISTRY. 
 
 compounds. In general, aluminium hydroxide acts upon 
 the hydroxides of the strongest base-forming elements 
 to form aluminates. So, also, aluminium chloride acts 
 upon the chlorides of the strongest base-forming ele- 
 ments to form double chlorides. By melting together 
 aluminium hydroxide and potassium or sodium hydrox- 
 ide, compounds of the formulas A1O.OK and AlO.ONa 
 are formed. So, also, by melting together aluminium 
 chloride and potassium or sodium chloride, compounds 
 of the formulas A1C1 4 K and AlCl 4 Na are formed. Com- 
 paring these oxygen and chlorine compounds, it is clear 
 that they are analogous. If the oxygen of the former 
 is replaced by an equivalent quantity of chlorine, the 
 chlorine compounds result : 
 
 EA1O 2 KA1C1 4 
 
 NaAlO, NaAlCl 4 
 
 Or, if their constitutional formulas are written in accord- 
 ance with the views already expressed regarding the 
 double chlorides, the analogy is also seen, thus : 
 
 O / Cl 
 
 AlCl 
 
 \C1 2 )K 
 
 A1 \ONa A1 ^ C1 
 
 \Cl a )Na. 
 
 The compounds of the same order as cryolite have their 
 analogues in such oxygen compounds as Al(ONa) 8 , etc., 
 as is shown by the following formulas : 
 
 Na 3 AlO 8 ; Na 3 AlF 6 ; 
 
 ONa 
 
 f -(F 2 )Na . 
 
 It is not improbable that by fusion with other chlorides 
 besides those of potassium and sodium, aluminium chlo- 
 ride will be found to yield other double chlorides analo- 
 gous to the spinels. According to what was said in 
 
ALUMINIUM OXIDE. 5~5 
 
 discussing the subject of double chlorides in general, 
 three series of such salts may be looked for, correspond- 
 ing to the formulas 
 
 /Cl /Cl /(C1 2 )M - 
 
 AlfCl , A1(--(C1 2 )M, and Alf (C1 2 )M ; 
 
 \C1 2 )M \C1 2 )M \C1 2 )M 
 
 and representatives of all these classes are known. Oxy- 
 gen compounds corresponding to the first and last of 
 these have been mentioned. As an example of an oxygen 
 compound corresponding to the second one, barium 
 aluminate, of the formula Ba 2 Al 2 O 5 , may be cited. 
 
 Aluminium Oxide, A1 2 O 3 . As has been stated, the 
 oxide is formed by heating the hydroxide. It is found 
 in nature in the form of ruby, sapphire, and corundum. 
 The natural variety is extremely hard ; and granular 
 corundum, which is known as emery, is used for polish- 
 ing. The red color of the ruby is caused by the presence 
 of a trace of a chromium compound ; while the blue color 
 of the sapphire is probably due to the presence of 
 a trace of a cobalt compound. Aluminium oxide is 
 infusible in the hottest furnace fire, but it melts in the 
 flame of the oxyhydrogen blow-pipe, and on cooling it 
 becomes crystalline. By mixing it with various easily 
 fusible substances and heating, it is obtained in the form 
 of crystals, and by adding certain metallic oxides these 
 crystals can be colored. In this way artificial rubies and 
 sapphires have been prepared, which have all the prop- 
 erties of the natural ones. When the oxide is moistened 
 with a few drops of a solution of cobaltous nitrate and 
 then ignited, it turns blue. This fact is taken advantage 
 of in chemical analysis for the purpose of detecting alu- 
 minium. When the oxide is made by gently igniting 
 the hydroxide, it dissolves in strong acids. If, however, 
 it is heated to a high temperature, acids will not dissolve 
 it. The natural varieties of the oxide, further, are not 
 soluble in acids. By fusion with acid potassium sulphate 
 insoluble aluminium oxide is converted into a soluble 
 compound. 
 
576 INORGANIC CHEMISTRY. 
 
 Aluminium Sulphate, A1 2 (SO 4 ) 3 . This salt is made by 
 dissolving the hydroxide of aluminium in dilute sulphuric 
 acid, and evaporating to crystallization, when a salt of 
 the composition A1 2 (SO 4 ) 3 + 18H 2 O is deposited. When 
 heated the salt loses its water of crystallization, and, if 
 the temperature is raised to that of red heat, the anhy- 
 drous salt is decomposed with loss of sulphur trioxide 
 and formation of aluminium oxide. This decomposition 
 is, however, not complete. The sulphate is manufactured 
 on the large scale for various purposes, as, for example, 
 for a mordant, for sizing paper, etc. 
 
 Basic Aluminium Sulphates. A solution of ordinary 
 aluminium sulphate has an acid reaction, and has the 
 power to dissolve metals, such as zinc, and hydroxides, 
 such as aluminium hydroxide. When a solution of the 
 sulphate is treated with the hydroxide, a basic salt of 
 the formula A1 2 O(SO 4 ) 2 + H 2 O is formed. This should 
 
 (o so 
 
 probably be represented by the formula Al < O 2 or 
 
 ( OH 
 
 A1(OH)SO 4 . Another basic salt has the formula 
 (A1O) 2 SO 4 , the salt being derived from the hydroxide, 
 
 A1O.OH, as represented thus : ^JO Q > SO 2 . The former 
 
 salt is soluble in water. When, therefore, a solution of 
 sodium or ammonium carbonate is added to a solution 
 of ordinary aluminium sulphate, the first portions of hy- 
 droxide which are precipitated redissolve in the excess 
 of the ordinary salt. There are other basic salts, some 
 of which occur in nature. 
 
 Alums. When a solution of aluminium sulphate is 
 brought together with a solution of potassium sulphate 
 in the proportion of their molecular weights, a salt 
 crystallizes out which has the composition represented 
 by the formula 
 
 KA1(S0 4 ) 2 + 12H 2 O or K 2 S0 4 + A1 3 (SO 4 ) 3 + 24H,O. 
 
 The most rational view which has been expressed re- 
 garding this compound is that it has the constitution 
 
ALUMS. 577 
 
 ,OK 
 2 < 
 
 O--A1, with perhaps some of the so-called water 
 
 of crystallization present in the form of hydroxyl. This 
 salt, which has long been known under the name of alum, 
 is the type of a class of similar compounds, all of which 
 are called alums. These alums may be regarded as 
 derived from the ordinary form by replacing the potas- 
 sium by sodium, ammonium, or any other member of 
 the sodium group, besides some other metals. Thus a 
 series of alums is obtained, of which the following are 
 examples : 
 
 NaAl(SO 4 ) s H 
 LiAl(S0 4 ) s 
 (NH 4 )A1(S0 4 )H 
 CsAl(SO.X H 
 TIA1(SO.X 
 
 - 12H 2 O ; 
 H12H 2 0; 
 
 i- 12H[o ; 
 1- 12H,0. 
 
 Again, alums are derived from the ordinary form by re- 
 placing the aluminium by some other elements which 
 have the power to form compounds resembling those of 
 aluminium, as, for example, iron, chromium, and man- 
 ganese. Such alums are those represented by the follow- 
 ing formulas : 
 
 KFe(SO 4 ) 2 + 12H 2 O ; 
 KCr(S0 4 ) 2 +12H 2 0; 
 KMn(SO 4 ) 2 + 12H 2 O. 
 
 In each of these, again, the potassium can be replaced as 
 in the case of ordinary alum ; so that the class includes 
 a comparatively large number of salts. All have certain 
 properties in common. They are all soluble in water, 
 and all crystallize in the same forms, which are regular 
 octahedrons combined with cubes. If a crystal of one 
 alum be suspended in the solution of any other one it 
 will continue to grow. They are all strictly isomorphous. 
 The principal alums containing aluminium are those of 
 
578 INORGANIC CHEMISTRY. 
 
 potassium and ammonium, both of which are manufac- 
 tured on the large scale. 
 
 Potassium Alum, Potassium - Aluminium Sulphate, 
 KA1(SO 4 ) 2 + 12H a O. Ordinary alum is found in nature 
 in some volcanic regions. The mineral alunite, which is 
 a basic salt of the formula .K(A1O) 3 (SO 4 ) 2 + 3H 2 O, or 
 perhaps K[A1(OH) 2 ] 3 (SO 4 ) 2 , occurs in larger quantities. 
 When this salt is heated and treated with water, ordinary 
 alum dissolves, and is easily obtained from the solution. 
 Another source of alum is alum shale. This occurs in 
 large quantities in nature, and consists of coal, clay, and 
 iron pyrites. When it is heated in contact with the air 
 the coal burns, as do also the sulphur and pyrites, and 
 sulphuric acid is formed. When allowed to lie for a time 
 in contact with the air the iron pyrites is converted into 
 sulphate and sulphuric acid. The latter acts upon the 
 clay or aluminium silicate, forming aluminium sulphate, 
 from which alum can easily be made. It is easier to 
 treat the shale and similar substances with sulphuric 
 acid directly, and this method is now generally employed. 
 
 Alum dissolves readily in hot water, 357.5 parts of the 
 crystallized salt dissolving in 100 parts of water at 100, 
 At only 3.9 parts dissolve, and at the ordinary tem- 
 perature about 12 parts. It crystallizes beautifully in 
 regular octahedrons, occasionally with cube faces devel- 
 oped on them. Under some circumstances it crystallizes 
 in cubes. When heated, alum melts in its water of 
 crystallization, and if heated to a sufficiently high tem- 
 perature the water passes off, leaving burnt alum. Heated 
 higher the salt decomposes, forming aluminium oxide 
 and potassium sulphate, and finally potassium aluminate 
 is formed. When potassium hydroxide, ammonia, or 
 the carbonate of potassium, sodium, or ammonium, is 
 added in small quantity to a solution of alum the pre- 
 cipitate first formed redissolves. If this is continued 
 until the reaction is neutral, or until a point is reached 
 beyond which the addition of the reagent produces a 
 precipitate which does not redissolve, there is then con- 
 tained in the solution a basic compound, known as basic 
 alum, which probably has the composition K 2 (A1 2 O)(SO 4 ) 3 . 
 
ALUMINIUM SILICATE. 579 
 
 When the solution is boiled the salt contained in it is 
 decomposed, forming ordinary alum and another basic 
 alum which is insoluble : 
 
 3[K,(Al a O)(SO.) 3 ] = K(A10) s (SO,) a + 3KA1(SO.), + K,SO, 
 
 The insoluble compound is known as insoluble alum, 
 
 Alum crystallized in cubes is obtained by evaporating a 
 solution to which some sodium or potassium carbonate 
 has been added. Alum is used very extensively in the 
 preparation of pigments, as a mordant, in the sizing of 
 paper, for clarifying water, etc. 
 
 Ammonium Alum, Ammonium - Aluminium Sulphate, 
 <NH4)A1(SO 4 ) 2 + 12H 2 O, is in every way much like the 
 potassium compound, and can be used in place of it for 
 almost all purposes for which alum is used. It is some- 
 what more easily soluble than ordinary alum. As it is 
 cheaper than the latter it is largely manufactured in 
 place of it. 
 
 Sodium Alum is much more easily soluble in water 
 than either potassium or ammonium alum, and this 
 makes it difficult to prepare it in pure condition. It is 
 therefore not manufactured, although sodium compounds 
 are cheaper than those of either potassium or ammonium. 
 
 Aluminium Silicate. It has been stated that alumin- 
 ium silicate enters into the composition of a number of 
 important minerals. It occurs in enormous quantities 
 in nature. The most important of the minerals contain- 
 ing it are the feldspars, of which ordinary feldspar, 
 KAlSi 3 O 8 , and albite, or sodium feldspar, NaAlSi 3 O 8 , are 
 the most abundant. These again enter into the compo- 
 sition of granite together with quartz and mica, and mica 
 is itself a double silicate of aluminium. As remarked 
 under Silicic Acid (which see), the natural silicates are for 
 the most part salts of polysilicic acids which are derived 
 from orthosilicic acid by loss of water from two' or more 
 molecules. Up to the present but little more has been 
 done with the many natural silicates than to determine 
 their percentage composition. It appears probable from 
 
580 INORGANIC CHEMISTRY. 
 
 what lias already been learned regarding their constitu- 
 tion that investigations in this direction will before long 
 yield interesting results. As yet, however, the methods 
 for such investigations are quite unsatisfactory, owing 
 largely to the fact that the compounds are so extremely 
 stable that but few reagents decompose them, and if 
 they are decomposed at all, the products are such that 
 no conclusion can be drawn regarding the constitution. 
 A careful study of the relations in which minerals occur 
 in nature will undoubtedly be of assistance, as this will 
 throw some light upon the conditions under which they 
 were formed. 
 
 One of the most common decompositions of minerals, 
 constantly taking place, and that has taken place to an 
 enormous extent, is that of feldspar. Under the in- 
 fluence of moisture and the carbon dioxide of the air, 
 this substance slowly decomposes, the products being 
 mainly potassium or sodium silicate and aluminium sili- 
 cate. The salt of the alkali metal, principally potas- 
 sium, being soluble, is carried away, and finds its way 
 into the soil. The silicate of aluminium is not soluble, 
 but it easily forms emulsions with water, and is there- 
 fore carried down the sides of the hills and mountains 
 upon which it is formed into the valleys, and much of it 
 finds its way into streams. Sometimes this carrying 
 away is prevented, and then large beds of comparatively 
 pure clay, known as kaoline, are formed. The clay 
 found in the valleys is always more or less impure and 
 colored. 
 
 Kaoline. This is the purest form of aluminium sili- 
 cate found in nature. It always contains water. Its 
 composition varies, some specimens on analysis giving 
 results which lead to the formula Al 4 (SiO 4 ) 3 + 4H 2 O, 
 according to which the substance is the salt of normal 
 silicic acid, Si(OH) 4 . Other specimens have the compo- 
 
 roH 
 
 sition HAlSiO 4 + H 3 O, or Si J Q\ M , H Q . When 
 
 [o/ 
 
CLA T ULTRAMARINE. , 581 
 
 heated alone kaoline does not melt ; but if feldspar is 
 added to it, the whole melts, and forms a translucent 
 mass known as porcelain. Other substances besides 
 feldspar may be used for this purpose. 
 
 Clay. Ordinary clay, as has been stated, is a name 
 given to the impure varieties of aluminium silicate which 
 have been carried down from the place of formation. 
 Among the substances besides aluminium silicate found 
 in clays are calcium carbonate, magnesium carbonate, 
 sand, and hydroxides of iron. The color is largely deter- 
 mined by the amount of the hydroxides of iron present. 
 The better varieties are used in the manufacture of the 
 so-called " stone-ware," gas-retorts, and fire-bricks. The 
 colored varieties are used for making ordinary earthen- 
 ware and bricks. Marl is clay mixed with considerable 
 quantities of calcium carbonate. 
 
 Ultramarine. The substance occurring in nature and 
 known as lapis lazuli consists of a silicate of sodium and 
 aluminium together with a sulphur compound, probably 
 a polysulphide of sodium. The coloring matter, known 
 as ultramarine, obtained by powdering it was formerly 
 very expensive, but it is now made artificially by the ton, 
 and the color of the artificially prepared substance is 
 even more beautiful than that of the natural. A great 
 deal of work has been done in the way of investigating 
 the chemical constitution of ultramarine, but the problem 
 has not yet been fully solved. The artificial preparation 
 is effected by melting together kaoline, anhydrous so- 
 dium carbonate, and sulphur ; or clay, calcined sodium 
 sulphate, and charcoal. By varying the conditions of 
 the preparation products of different colors are obtained. 
 Besides the deep-blue ultramarine, there are now manu- 
 factured ultramarines of different shades of blue, and a 
 green variety. A white, a red, a yellow, and a violet, 
 variety are also known. The substance which gives to 
 ultramarine its color is destroyed by acids, but not by 
 alkalies. It can be heated in a closed vessel without 
 change, but if heated to a high temperature in the air or 
 in oxygen the color is destroyed. 
 
582 INORGANIC CHEMISTRY. 
 
 Ultramarine is now manufactured in very large quan- 
 tity according to a recent report, to the extent of nearly 
 9000 tons a year. It is the most extensively used blue 
 coloring matter. 
 
 Porcelain. It was stated above that when kaoline is 
 heated alone it does not melt, but that if feldspar is 
 added to it, or if that found in nature contains feldspar, 
 as is frequently the case, it either fuses together forming 
 a compact mass, or it melts and forms a translucent 
 mass. Further, when kaoline or any other variety of 
 clay is mixed with water, a plastic substance results, 
 which can be kneaded and worked into any desired 
 form. These facts form the basis of the manufacture of 
 earthenware, porcelain, etc. The ease with which the 
 mass melts depends upon the quantity of feldspar or 
 other flux added to it. If but little is added it melts 
 with difficulty ; if much is added it melts easily. 
 
 In the manufacture of the finest kinds of porcelain 
 kaoline is used. This is generally mixed with a little 
 feldspar or chalk, gypsum or some other flux, and sand 
 is also added. All these substances must be very finely 
 ground. The mixture is then worked into the desired 
 forms, and carefully dried. After the objects are dried 
 they are next burned, first at a red heat at which the 
 mass becomes solid, afterwards at a white heat for the 
 purpose of forming a glaze upon the surface. The prod- 
 uct after the first burning is that which is familiar as 
 porous earthenware ; that formed in the second burning 
 is the porcelain with glaze as it is commonly used. 
 
 In order to form the glaze upon the porcelain the 
 porous earthenware first formed is drawn through a 
 vessel containing proper materials in finely powdered 
 condition and suspended in water. The materials used 
 are generally the same as those used for the porcelain 
 itself, but they are mixed in different proportions, witL 
 less kaolin, and more sand and feldspar, so as to be 
 more easily fusible. After this treatment the objects 
 are again heated to a high temperature. 
 
 Earthenware. The ordinary varieties of earthenware 
 are made from varieties of clay which are much less pure 
 
REACTIONS OF ALUMINIUM SALTS. 583 
 
 than kaoline. Ordinary colored clay is used. The ob- 
 jects are formed, and then subjected in general to the 
 same kind of treatment as porcelain. They are glazed 
 in different ways. One method consists in bringing the 
 glazing material on the earthenware before it is burned ; 
 another method consists in putting the objects in the 
 furnace without a glaze, and towards the end of the 
 firing process sodium chloride is thrown into the fur- 
 nace, and is thus brought in contact with the ware in 
 the form of vapor. A chemical change takes place, re- 
 sulting in the formation of a silicate of aluminium and 
 sodium upon the surface. This melts, and forms a glaze. 
 
 Bricks are the most common variety of unglazed 
 earthenware. Owing to the presence of other sub- 
 stances besides aluminium silicate, as, for example, cal- 
 cium carbonate, the material is comparatively easily 
 fusible. The color of bricks is largely due to the pres- 
 ence of oxides of iron. 
 
 Reactions of Aluminium Salts which are of Special 
 Value in Chemical Analysis. Potassium and sodium hy- 
 droxides precipitate aluminium hydroxide, which is solu- 
 ble in an excess of the reagents. 
 
 Ammonia precipitates the hydroxide, which is only 
 slightly soluble in an excess of the reagent. 
 
 Hydrogen sulphide and carbon dioxide precipitate alu- 
 minium hydroxide from a solution of an aluminate ; that 
 is, from a solution of aluminium hydroxide in a caustic 
 alkali. 
 
 Ammonium sulphide and other soluble sulphides pre- 
 cipitate the hydroxide. This is due to the instability of 
 the sulphide of aluminium, or, going farther back, to 
 the weak basic character of the hydroxide. The reaction 
 of ammonium sulphide with aluminium sulphate takes 
 place as represented in the following equation : 
 
 A1 2 (SO 4 ) 3 + 3(NH 4 ) 2 S + 6H 2 O = 3(NH 4 ) 2 S0 4 + 3H 2 S + 2A1(OH) 8 . 
 
 Soluble carbonates precipitate aluminium hydroxide for 
 the same reason that the soluble sulphides do. The re- 
 
584 INORGANIC CHEMISTRY. 
 
 action between aluminium sulphate and sodium carbon- 
 ate takes place thus : 
 
 Al 3 (S0 4 ) 3 +3Na 2 C0 8 +3H 2 0=3Na a SO 4 +3CO 2 +2Al(OH),. 
 
 OTHEK MEMBERS OF FAMILY III, GROUP A. 
 
 Scandium, Sc (At. Wt. 43.78). This element was dis- 
 covered in 1880 in the minerals euxenite and gadolinite. 
 Its compounds are similar to those of aluminium. It 
 forms an oxide of the formula Sc 2 O 3 ; a sulphate, Se 2 (SO 4 ) 3 ; 
 a double sulphate, KSc(SO 4 ) 2 ; etc. It is of special in- 
 terest for the reason that its properties were foretold by 
 Mendeleeff several years before it was discovered. The 
 prophecy was based upon the position of the element in 
 the periodic system. When the relations between the 
 atomic weights and properties of the elements were first 
 described in a comprehensive way by Lothar Meyer and 
 Mendeleeff, the latter described the properties of an ele- 
 ment then unknown, which he called ekaboron. This 
 should have the atomic weight about 44, should form an 
 oxide of the formula M 2 O 3 , etc. It has been shown that 
 the properties of scandium agree very closely with those 
 foretold. 
 
 Yttrium, Y (At. Wt. 88.35), like scandium, is found in 
 gadolinite, euxenite, and some other rare minerals. 
 The element itself has not been isolated. Its chloride, 
 YC1 3 -\- 6H 2 O, is easily made. With potassium and so- 
 dium chlorides it forms double chlorides analogous to 
 those formed by aluminium chloride. Its oxide has the 
 formula, T 2 O 3 , and is formed by heating the hydroxide, 
 Y(OH) 3 , or nitrate, Y(NO 3 ) 3 . The hydroxide, Y(OH) 3 , is 
 precipitated by adding potassium hydroxide to a solu- 
 tion of an yttrium salt, and is not dissolved by an excess 
 of the alkali. The hydroxide, while being less acidic 
 than aluminium hydroxide, is also more strongly basic, 
 as is shown by its power to unite with weak acids. 
 When exposed to the air, it attracts and combines with 
 carbonic acid. 
 
THE BORON-ALUMINIUM GROUP IN GENERAL. 585 
 
 Ytterbium, Yb (At. Wt. 171.88). This rare element, like 
 scandium and yttrium, is found in gadolinite and euxen- 
 ite most abundantly in the latter. Its compounds in 
 general resemble those of yttrium. Its hydroxide is not 
 soluble in alkalies, but it absorbs and combines with car- 
 bon dioxide. Its oxide has the formula Yb 2 O 3 , its sul- 
 phate, Yb 2 (SO 4 ) 3 , etc. 
 
 Samarium, Sm, Terbium, Tr, and Gadolinium, Gd, are 
 elements that have been obtained from samarskite, a 
 North Carolina mineral. They have also been found in 
 small quantities in some other minerals, as cerite, gado- 
 linite, and orthite. In general these elements resemble 
 aluminium in their chemical conduct. 
 
 Helium. When cleveite and some other rare minerals 
 are treated with acids a gas is given off. This was at 
 first held to be nitrogen, but a careful spectroscopic 
 examination by W. Ramsay has shown that it contains 
 a hitherto unknown gas. The spectrum contains lines 
 which are identical with lines that have been observed 
 in the solar spectrum. These have been ascribed to an 
 unknown element, helium. This has now been isolated 
 and studied. It proves to be a very light gas, as incapa- 
 ble of forming chemical compounds as argon. Practically 
 all the evidence goes to show that it is an element, but 
 as yet little is known in regard to its relations to other 
 elements. 
 
 The chemistry of lanthanum is so intimately connected 
 with that of cerium and didymium, that, although these 
 three elements appear to belong to different families, 
 they will be briefly considered together. 
 
 The Boron- Aluminium Group in General. Comparing 
 the group of which boron and aluminium are the prin- 
 cipal members with the potassium, calcium, and mag- 
 nesium groups, it will be seen that the members of this 
 group do not resemble one another as closely as the 
 members of the other groups do. There is, however, the 
 same strengthening of the basic properties and weaken- 
 ing of the acid properties as the atomic weight increases. 
 Boron is the most strongly acid and the least basic ; 
 -aluminium is more basic, but has still some acid prop- 
 
586 
 
 INORGANIC CHEMISTRY. 
 
 erties ; while the other members are more strongly 
 basic, and do not exhibit any acid properties. Compar- 
 ing the first members of Group A, of Families I, II, and 
 III, it is clear that with increasing atomic weight the 
 acid properties and the valence increase. The elements 
 referred to are lithium, glucinum, and boron. 
 
 Members of Group B, Families I, II, III, and IV. The 
 base-forming elements thus far considered form the prin- 
 cipal groups of the first three families. In these principal 
 groups the most characteristic elements of these families 
 occur. But besides the principal group of each family 
 there is a secondary group, the members of which differ 
 in some respects from those of the principal group, 
 though they resemble one another. Between the second- 
 ary group of Family I and that of Family II, further, 
 there are some points of resemblance. The secondary 
 group in Family IV bears to the principal group much 
 the same relation that the secondary groups of the first 
 three families bear to the principal groups. In the table 
 p. 151 the principal groups are those which fall under 
 the letter A, and the secondary groups are those which 
 fall under the letter B in each of the first four families. 
 These secondary groups are : 
 
 Family 
 
 Group 
 
 
 
 
 I 
 
 B 
 
 Copper 
 
 Silver 
 
 Gold 
 
 II 
 
 B 
 
 Zinc 
 
 Cadmium 
 
 Mercury 
 
 III 
 
 B 
 
 Gallium 
 
 Indium 
 
 Thallium 
 
 IV 
 
 B 
 
 Germanium 
 
 Tin 
 
 Lead 
 
 In the fifth, sixth, and seventh families the most char- 
 acteristic elements are those which occur in Group B. 
 These have already been studied. 
 
 Haying thus considered the members of the principal 
 groups of the first four families, let us next turn to the 
 study of the members of the secondary groups. 
 
CHAPTER XXVIII. 
 
 ELEMENTS OF FAMILY I, GROUP B : 
 COPPER SILVER GOLD. 
 
 General. The facts which strike one most forcibly on 
 comparing the elements of this group with those of 
 Group A of the same family are, that they are much less 
 active chemically, and that they furnish a greater variety 
 of compounds. Sodium and potassium and the other 
 members of Group A display the greatest activity, as we 
 have seen. The basic character is most strongly devel- 
 oped in them. Further, in nearly all their compounds 
 they act with the same valence. They are univalent in 
 all their salts. Copper, silver, and gold, however, are 
 not chemically active elements, and the activity grows 
 less with increasing atomic weight. Copper and gold 
 form two series of compounds each, and silver also forms 
 a few compounds, in which it appears with a valence 
 greater than one. In the two series of salts formed by 
 copper the element appears to be univalent and bivalent, 
 as in the chlorides CuCl and CuCl 2 . Gold, however, is 
 univalent and trivaleiit, while silver is almost exclusively 
 univalent. It must be said that the resemblance between 
 gold and the other members of Group B is apparently 
 not as marked as that between mercury and copper and 
 silver. It is, however, possible that as investigation 
 proceeds the resemblance will appear more striking than 
 it does at present. 
 
 COPPER, Cu (At. Wt. 63.12). 
 
 General. The compounds of .copper which are most 
 commonly met with are those in which it acts as a biva- 
 lent element. Its principal compounds are copper oxide, 
 
 (587) 
 
588 INORGANIC CHEMISTRY. 
 
 CuO ; the sulphate, CuSO 4 ; and the sulphide, CuS. In 
 all these the copper is bivalent. But besides these there 
 are such compounds as CuCl and Cu 2 O, in which the 
 element appears to be univalent. There are, then, two 
 series of salts, of which the following will serve as 
 examples : 
 
 CuCl CuCl a 
 
 CuBr CuBr, 
 
 Cu 2 CuO 
 
 Those compounds which are of the first order, corre- 
 sponding to the chloride CuCl, are called cuprous com- 
 pounds. Thus, CuCl is cuprous chloride; Cu 2 O, cuprous 
 oxide, etc. On the other hand, compounds of the second 
 order are called cupric compounds. Thus, CuCl 2 is cupric 
 chloride; CuO, cupric oxide; CuSO 4 , cupric sulphate, etc. 
 It has been suggested that perhaps the formula of the 
 simpler cuprous compounds, like CuCl, etc., should be 
 doubled, and written Cu 2 Cl 2 , Cu 2 I 2 , etc. This suggestion 
 has its origin in the valence hypothesis. In cupric 
 chloride, CuCl 2 , and cupric oxide, CuO, copper is evi- 
 dently bivalent ; whereas, if the formulas of the cuprous 
 compounds are the simpler ones, CuCl, Cul, etc., copper 
 is univalent in them. If, however, cuprous chloride is 
 Cu 2 Cl 2 , it may be that in it the copper is bivalent. It is 
 only necessary to assume that in the molecule of cu- 
 prous chloride two atoms of copper are combined as 
 represented thus : 
 
 Cu- 
 
 If, then, each of the copper atoms should combine with 
 a chlorine atom, the compound would have the formula 
 Cu 2 Cl 2 . The question here presented is similar to that 
 concerning the molecular formula of aluminium chloride. 
 A determination of the specific gravity of the vapor of 
 cuprous chloride has been made, and it has been found 
 to correspond to that required by the formula Cu 2 Cl 2 . 
 It is possible, however, that at a higher temperature a 
 different result may be obtained, as in the case of alu- 
 
COPPER. 589 
 
 minium chloride, and it is possible that the compound 
 may be a double chloride, formed by union through the 
 chlorine atoms as represented in the formula CuCl-ClOu 
 or Cu-(Cl 2 )-Cu. Then there would be complete analogy 
 between cuprous chloride and cuprous oxide, Cu-O-Cu. 
 Our knowledge in regard to this matter is extremely 
 limited at present, and it seems perfectly justifiable to 
 use the simpler formulas for the cuprous compounds 
 until further evidence has been produced. Whatever 
 the explanation may be, it is undoubtedly a fact that 
 there are two series of salts of copper, in one of which 
 there is relatively half as much copper as in the other, 
 and it is also a fact that by comparatively simple methods 
 the salts of one series can be converted into those of the 
 other, as will be pointed out below. 
 
 Forms in which Copper occurs in Nature. Copper is 
 a widely distributed element, and it occurs also in large 
 quantities. It occurs in the uncombined condition, or 
 as native copper, in large quantity in the United States 
 in the neighborhood of Lake Superior, in China, Japan, 
 Siberia, and Sweden. The most valuable ores of cop- 
 per are the oxides, ruby copper or cuprous oxide, Cu 2 O, 
 and cupric oxide, CuO ; the carbonates, as malachite, 
 Cu 2 (OH) 2 CO 3 ; the sulphides, as chalcocite, Cu 2 S, copper 
 pyrites, Cu 2 S.Fe 2 S 3 ; and others. 
 
 Metallurgy of Copper. The metallurgy of copper is 
 comparatively complicated, owing to the difficulty of 
 converting the ores of copper into the oxide. In most 
 of the ores used sulphur and iron are contained, as well 
 as smaller quantities of other elements, as arsenic, anti- 
 mony, lead, etc. The ores are first roasted with the 
 object of converting the sulphides partly into oxides. 
 Under these circumstances the sulphides of iron are 
 more easily converted into the oxides than the sulphides 
 of copper. By adding a material rich in silicic acid, and 
 melting the roasted ore in a blast furnace with charcoal, 
 the oxide of iron is partly reduced, and converted into 
 silicate, which runs off with the slag. In this way a 
 product is obtained which is richer in copper than the 
 roasted ore. This, which is called the matte, contains 
 
590 INORGANIC CHEMISTRY. 
 
 copper sulphide and iron sulphide. The matte is again 
 roasted and melted in the same way as the ore, and a 
 further quantity of iron is removed, while* some of the 
 copper is reduced. A reaction which plays an impor- 
 tant part in these processes is that which takes place 
 between cuprous oxide and cuprous sulphide, forming 
 metallic copper and sulphur dioxide : 
 
 2Cu 2 O + Cu 2 S = 6Cu + S0 2 . 
 
 Sometimes it is necessary to repeat the roasting and 
 melting with charcoal and sand a number of times, the 
 matte becoming richer in copper at each successive stage. 
 Properties. Copper is a hard metal, of a reddish color 
 and metallic lustre. It does not change in dry air, but 
 in moist air it gradually becomes covered with a green 
 layer of a basic carbonate. It melts at a somewhat 
 lower temperature than gold, and at a somewhat higher 
 temperature than silver. It is very malleable and tena- 
 cious. It decomposes water only at bright-red heat. 
 When heated in the air to a comparatively high tempera- 
 ture it becomes covered with a layer of cupric oxide ; at 
 a lower temperature cuprous oxide is formed. Nitric 
 acid dissolves it, copper nitrate, Cu(NO 3 ) 2 , being formed, 
 and the oxides of nitrogen being evolved (see p. 285) ; 
 dilute sulphuric acid does not act upon it unless the air 
 has access to it ; concentrated sulphuric acid when 
 heated with it forms cupric sulphate, CuSO 4 , and sul- 
 phur dioxide (see p. 217). Dilute acids in general do 
 not act upon it unless the air has access to it. This fact 
 is of importance in connection with the use of copper 
 vessels in culinary operations- Substances containing 
 vegetable acids can be boiled in bright copper vessels 
 with impunity, for the water vapor prevents the access of 
 air, but, on cooling, the air is admitted, and then action 
 takes place, causing solution of some of the copper, which 
 is objectionable. Ammonia in contact with copper ab- 
 sorbs oxygen, and the copper dissolves in consequence 
 of the formation of a compound of cupric oxide and am- 
 monia. This fact is sometimes taken advantage of for 
 
ALLOTS OF COPPER, 591 
 
 the preparation of nitrogen, as was stated in speaking of 
 this gas (see p. 249). 
 
 Applications. As is well known, copper is used very 
 extensively for a variety of purposes, among which the 
 following may be mentioned : for electrical apparatus, 
 coins, copper vessels, roofs, for covering the bottoms of 
 ships, etc. It is also used for copper-plating ; and in the 
 preparation of a number of valuable alloys, such as 
 brass, bronze, gun-metal, bell-metal, etc. 
 
 Alloys. Brass is a mixture or compound of about one 
 part of zinc and two parts of copper ; these proportions 
 may, however, be varied between quite wide limits. 
 There is a variety of brass containing equal parts of 
 zinc and copper, and another containing one part of zinc 
 and five parts of copper. Pinchbeck is made by combin- 
 ing two parts of copper and one of brass. 
 
 Bronze consists of copper, zinc, and tin. The propor- 
 tion of copper varies from 65 to 84 per cent ; that of 
 zinc from 31.5 to 11 per cent ; and that of tin from 2.5 
 to 4 per cent. When exposed to the air bronze becomes 
 covered with a green coating of basic copper carbonate, 
 which protects it from further action. This coating is 
 now generally produced artificially by a variety of meth- 
 ods, as by washing the surface with a solution of salts 
 and acids. 
 
 Gun-metal consists generally of copper and tin in the 
 proportion of 11 parts of tin and 100 parts of copper. 
 
 Bell-metal contains a larger proportion (from 20 to 25 
 per cent) of tin than gun-metal. 
 
 Alloys with Aluminium containing aluminium and cop- 
 per in widely different proportions are made. That with 
 3 per cent of copper has a whiter color than aluminium, 
 the color being more like that of silver. On the other 
 hand, an alloy of copper with 5 to 10 per cent of alu-. 
 minium has a color resembling that of gold. This, which 
 is known as aluminium bronze, is very hard and elastic, 
 and is not easily acted upon by chemical reagents. It is 
 now used to a considerable extent in the manufacture of 
 ornamental and useful articles. 
 
 German silver is an alloy consisting of copper, zinc, and 
 
592 INORGANIC CHEMISTRY. 
 
 nickel. The proportion of copper varies from 40 to 60 
 per cent ; that of zinc from 19 to 44 per cent ; and that 
 of nickel from 6 to 22 per cent. 
 
 Cuprous Hydride, CuH. This compound is made by 
 treating a solution of barium hypophosphite with ,a solu- 
 tion of copper sulphate. It is thrown down as a yellow 
 precipitate which gradually becomes darker. At 60 it 
 decomposes into copper and hydrogen. With hydro- 
 chloric acid it yields cuprous chloride and hydrogen. 
 
 Cupric Hydride, CuH 2 , is formed by treating a solution 
 of copper sulphate with hypophosphorous acid. When 
 freshly prepared it is a reddish-brown sponge-like mass, 
 which, however, changes to a chocolate-colored powder 
 on being freed from acid and boiled for some time. It 
 is not readily changed when heated in the air. It dis- 
 solves in hydrochloric acid with evolution of hydrogen. 
 
 Cuprous Chloride, CuCl, is formed by heating cupric 
 chloride, CuCl 2 ; by passing hydrochloric acid over highly 
 heated copper ; by boiling an excess of copper filings with 
 concentrated hydrochloric acid with the addition of a 
 little nitric acid, filtering through asbestos, and pouring 
 into water ; and by treating cupric chloride with reducing 
 agents, as, for example, sulphurous acid. The action with 
 sulphur dioxide takes place as represented in the equation 
 
 2CuCl a + SO 3 + 2H 2 = 2CuCl + H 2 SO 4 +.2HC1. 
 
 It is a white crystalline compound, and is difficultly solu- 
 ble in water. When exposed to the air it rapidly turns 
 green in consequence of the formation of a basic chloride, 
 as, for example, 2CuO.CuCl 2 , or Cl-Cu-O-Cu-O-Cu-Cl. 
 It is volatile at a high temperature, and a determination 
 of the specific gravity of the vapor gave a result corre- 
 sponding to the formula Cu 2 Cl 2 , as was stated above. It 
 has markedly the power to absorb chlorine, and there- 
 fore acts as a reducing agent. Ammonia dissolves it r 
 forming a compound of the composition represented by 
 the formula CuCl.NH 3 , which may be regarded as de- 
 rived from ammonium chloride by replacing an atom of 
 hydrogen by an atom of copper. 
 
 Cupric Chloride, CuCl 2 . This compound is formed by 
 treating copper or cuprous chloride with chlorine. It is 
 
COMPOUNDS OF COPPER. 593 
 
 also easily made by dissolving cupric hydroxide, or car- 
 bonate, in hydrochloric acid. From its solution in water 
 the chloride crystallizes with two molecules of water, 
 CuCl a -|- 2H 2 O. The crystals when heated lose their 
 water without suffering further decomposition, except at 
 high heat, when a part of the chlorine is given off, and 
 cuprous chloride is formed. Cupric chloride combines 
 with ammonia gas, forming a compound of the formula 
 CuCl 2 .6NH 3 , which is soluble in water, with a dark blue 
 color. When heated it loses four molecules of ammonia, 
 and the compound CuCl 2 .2NH 3 is left behind. This may 
 be regarded as ammonium chloride, in which two hydro- 
 gen atoms have been replaced by an atom of bivalent 
 
 copper, as represented in the formula ^ u <]jTi 3 ni- 
 
 There is, however, no direct experimental evidence in 
 favor of this view. 
 
 With other chlorides cupric chloride forms double 
 chlorides similar to those formed by magnesium and 
 aluminium ; such, for example, as CuCl 2 .2NH 4 Cl or 
 (NH 4 ) 2 CuCl 4 , CuCl 2 .2KCl, or K 2 CuCl 4 , etc. 
 
 Cuprous Iodide, Cul. When a solution of a cupric salt 
 is treated with potassium iodide, cuprous iodide is pre- 
 cipitated and iodine is set free, owing to the instability 
 of cupric iodide : 
 
 2CuS0 4 + 4KI = 2K 2 SO 4 + 2CuI + I 2 . 
 
 If a reducing agent is added at the same time, iodine is 
 not set free. Thus, for example, when sulphur dioxide 
 is used, the reaction takes place as represented in the 
 equation 
 
 2CuSO 4 + SO 2 + 2KI + 2H 2 O = K 2 SO 4 + 2H 2 SO 4 + 2CuL 
 
 It forms a white precipitate. Cupric iodide is not 
 known. A similar conduct is shown by the cyanides. 
 
 Cuprous Hydroxide, Cu(OH). The simple compound 
 of the formula here given is not known, but a derivative 
 of this, of the formula Cu 8 O 3 (OH) 2 or 4Cu 2 O.H 2 O, is 
 easily made by adding sodium hydroxide to a solution of 
 
594 INORGANIC CHEMISTRY. 
 
 a cuprous salt. It passes readily over into cuprous oxide 
 by gently heating it. 
 
 Cuprous Oxide, Cu 2 O, occurs in nature, and is known as 
 r'uhy copper or cuprite. It is easily prepared by treating 
 a solution of glucose, or starch sugar, with copper sul- 
 phate and potassium hydroxide. By boiling, the copper 
 is thrown down in the form of cuprous oxide. At first 
 this is yellow, and it is supposed by some that the yel- 
 low compound is the hydroxide, but satisfactory evi- 
 dence of the correctness of this view has not been fur- 
 nished. The yellow precipitate is soon converted into the 
 red oxide. Cuprous oxide is not changed when allowed 
 to lie in contact with the air. It dissolves in nitric and 
 sulphuric acids, forming cupric salts ; and if the acids 
 are dilute, copper is deposited. This will be clear from 
 a consideration of the following equation : 
 
 Cu 2 + H 2 S0 4 = CuSO 4 + H 2 O + Cu. 
 
 Cupric Hydroxide, Cu(OH) 2 , like the hydroxides of most 
 base-forming elements, is thrown down by the addition of 
 a soluble hydroxide to a cupric salt. It is a voluminous, 
 blue precipitate. When allowed to stand in a solution, or 
 when the solution is boiled, the hydroxide loses a part of 
 its hydroxyl, and is converted into a black compound of 
 the formula 
 
 Cu(OH) a + 2CuO, or HO-Cu-O-Cu-O-Cu-OH, 
 
 and this when dried and heated is converted into the 
 oxide CuO. 
 
 Cupric Oxide, CuO. Cupric oxide is found in nature in 
 tjie neighborhood of Lake Superior in the United States, 
 and is formed by heating copper to redness in contact 
 with the air, or by heating the nitrate. It loses its oxy- 
 gen very readily when treated with reducing agents, such 
 as hydrogen and carbon. It is used extensively in 
 quantitative analysis for the purpose of estimating the 
 composition of organic compounds, or such as contain 
 carbon and hydrogen. Its use is based upon the fact 
 that when organic compounds are heated with the oxide 
 
CUPRIC SULPHATE, 595 
 
 they are oxidized, the carbon being converted into car- 
 bon dioxide and the hydrogen into water. By passing 
 the products of the oxidation through calcium chloride, 
 and a solution of potassium hydroxide, the water is re- 
 tained in the first, and the carbon dioxide in the second, 
 and the weight of each formed can easily be determined. 
 Cupric oxide is dissolved by ammonia in the presence of 
 air and a little of some ammonium salt. The composition 
 of the compound in the solution is not known. 
 
 Other Oxides of Copper. Besides cuprous and cupric 
 oxides, copper forms two other -compounds with oxygen. 
 These are copper siiboxide, Cu 4 O, and the peroxide, CuO 2 . 
 The former is prepared by treating a solution of copper 
 sulphate with stannous chloride. It takes up oxygen 
 from the air, and is converted into higher oxides. The 
 peroxide is said to be formed by treating cupric hydrox- 
 ide with hydrogen peroxide. 
 
 Cupric Sulphate, CuSO 4 . This salt is manufactured on 
 the large scale, and in the crystallized form, containing 
 five molecules of water, CuSO 4 -f- 5H 2 O, is commonly 
 called " blue vitriol." It is found in nature to some ex- 
 tent, being formed by the action of the oxygen of the air 
 on the sulphide. It is most conveniently made by dis- 
 solving metallic copper in concentrated sulphuric acid, 
 or by treating cupric sulphide with sulphuric acid. The 
 action of sulphuric acid on the metal has already been 
 referred to. It consists essentially in the formation of 
 cupric sulphate, sulphur dioxide, and water, as expressed 
 in the equation 
 
 Cu + 2H 2 S0 4 = CuSO 4 + SO 2 + 2H 2 O. 
 
 The question whether the copper reduces the sulphuric 
 acid directly, or the hydrogen given off from the acid 
 effects the reduction, is an open one. But there are 
 other products formed besides those mentioned. At 
 first a brown substance of the composition Cu 2 S is 
 deposited. As the action proceeds oxysulphides are 
 formed, the final product of a series of changes being 
 Cu 2 OS, or CuO.CuS, which is black, and insoluble in 
 water. Under some conditions a considerable Droportion 
 
596 INORGANIC CHEMISTRY. 
 
 of the copper is transformed into the oxy sulphides by 
 sulphuric acid, Cupric sulphate is obtained in large 
 blue crystals of the triclinic system, which have the com- 
 position CuSO 4 + 5H 2 O. When heated to 100, four 
 molecules of water are given off, and the last is not 
 given off until the temperature 200 is reached. This 
 makes it appear probable that the salt has the con- 
 stitution represented by the formula CuSO 3 (OHj 2 or 
 
 [g>Cu 
 OS \ ^TT , corresponding in this respect to magnesi- 
 
 [OH 
 
 um sulphate (which see). When heated higher, it loses 
 all its hydroxyl, and the salt, CuSO 4 , is left in the form 
 of a white powder, which has the power to take up water 
 from the air, becoming blue again. It dissolves in three 
 parts of cold water and one-half part boiling water. Cop- 
 per sulphate, containing seven molecules of water, CuSO 4 
 -f- 7H 2 O, is obtained when mixed with solutions of the 
 sulphates of iron, zinc, or magnesium, all of which crys- 
 tallize with seven molecules of water. In this form cu- 
 pric sulphate is isomorphous with the other sulphates. 
 These salts have in general received the name of vitriols, 
 and the old names " green vitriol," " white vitriol," and 
 "blue vitriol" are still used to some extent, though 
 rarely by chemists. Among the similar salts included 
 under the same general head are the following : 
 
 Zinc sulphate (white vitriol), .... ZnSO 4 + 7H 2 O 
 
 Magnesium sulphate, MgSO 4 -f- 7H 2 O 
 
 Glucinum sulphate, G1SO 4 + 7H 2 O 
 
 Ferrous sulphate (green vitriol), . . . FeSO 4 -j- 7H 2 O 
 
 Nickel sulphate, MSO 4 + 7H 2 O 
 
 Cobalt sulphate, CoSO 4 + 7H 2 O 
 
 Copper sulphate (blue vitriol), 
 
 CuSO 4 + 7H 2 O,(CuSO 4 + 5H 2 O) 
 
 Cupric sulphate is used extensively in the preparation 
 of blue and green pigments, in copper-plating by elec- 
 trolysis, in galvanic batteries, for the purpose of pre- 
 serving wood, and as a remedy against phylloxera (see 
 p. 406), etc. 
 
CUPRIC SULPHATE. 597 
 
 Cupric sulphate combines with other sulphates, form- 
 ing double salts similar to those formed by aluminium 
 and magnesium. The potassium and ammonium com- 
 pounds have the formulas K 2 SO 4 .CuSO 4 -f- 6H 2 O and 
 (NH 4 ) 2 SO 4 .CuSO 4 + 6H 2 O, and probably have the consti- 
 tution represented by the general formula 
 
 " so,< M 
 
 )<>Cu. 
 
 When a solution of cupric sulphate is treated with 
 ammonia a precipitate of a basic salt is at first formed, 
 but this dissolves when more ammonia is added, forming 
 a deep-blue solution. The precipitate first formed is 
 a basic sulphate of copper ; while the solution contains 
 a compound of cupric sulphate, ammonia, and water, of 
 the composition represented by the formula CuSO 4 . 
 4NH 3 -f- H 2 O. When heated the salt loses water and 
 ammonia, until it has the composition CuSO 4 .2NH 3 . 
 This is probably analogous to the ammonia compound 
 of cupric chloride, CuCl 2 .2NH 3 , and may be regarded as 
 having a similar constitution, that is, as ammonium 
 sulphate, in which two hydrogen atoms have been re- 
 placed by an atom of bivalent copper, as expressed in 
 
 the formula SO 4 <^jj 3 > Cu. It is a curious and inter- 
 esting, though at present inexplicable, fact, that anhy- 
 drous copper sulphate combines with five molecules of 
 ammonia just as it does with five molecules of water, 
 and that by lying in moist air the molecules of ammonia 
 in the compound are successively replaced by water, so 
 that the following series of compounds is formed : 
 
 CuSO 4 .5NH 3 ; 
 CuSO 4 .4NH 3 .H 2 O ; 
 CuS0 4 .3NH 3 .2H 2 O; 
 CuSO 4 .2NH 3 .3H 2 O; 
 CuS0 4 .NH,.4H a O ; 
 CuS0 4 .5H 2 O, 
 
598 INORGANIC CHEMISTRY. 
 
 From this it would appear that the ammonia in these 
 compounds plays a part analogous to that played by 
 the " water of crystallization." This does not speak 
 in favor of the view above expressed concerning the 
 constitution of cuprammonium sulphate, in which the 
 copper is held to be in combination with nitrogen. 
 There is in fact no satisfactory theory for most of the 
 salts containing water of crystallization, nor for most 
 of those containing ammonia. The power to combine 
 with ammonia is very commonly met with among me- 
 tallic salts probably fully as much so as the power 
 to combine with water. Some metals indeed, as cobalt 
 and platinum, form a very large number of complex 
 compounds with ammonia, and with ammonium salts. 
 
 Cupric Nitrate, Cu(NO 3 ) 2 , is easily formed by dissolv- 
 ing copper in dilute nitric acid. It is easily soluble in 
 water, and is deposited in crystallized form, the crystals 
 containing three or six molecules of water according to 
 the temperature, the salt with six molecules being 
 formed at the lower temperature. Like other copper 
 salts, it has a blue color. It combines with ammonia 
 and with ammonium nitrate. 
 
 Cupric Arsenite, CuHAsO 3 , is formed as a greenish- 
 yellow precipitate when an ammoniacal solution of arse- 
 nious acid is added to a solution of cupric sulphate. It 
 is known as Scheele's green. A compound of cupric arse- 
 nite and cupric acetate, which is made by treating a 
 basic acetate of copper with arsenious acid, is known as 
 Schweinfurt green. On account of their poisonous char- 
 acter these compounds are not now used as extensively 
 as formerly. 
 
 Cupric Carbonates. When a soluble carbonate is 
 added to a solution of cupric sulphate a voluminous 
 greenish precipitate is formed, which has the composition 
 
 Cu<g H 
 
 Cu(OH) 2 .CuC0 3 , or Q>CO. This is plainly a 
 
 Cu< OH 
 
 basic carbonate. The mineral malachite, which has a 
 beautiful green color, has the same composition as the 
 precipitate just mentioned. 
 
CYANIDES OF COPPER. 599 
 
 Cyanides of Copper. Both cuprous and cupric cya- 
 nides are known, but while generally the cupric com- 
 pounds are the more stable, cupric cyanide, like cupric 
 iodide, is extremely unstable. It is readily changed to 
 a compound intermediate between the cupric and the 
 cuprous salt. This has the composition CuCy 2 .2CuCy. 
 By heating in suspension in water this intermediate com- 
 pound is converted into the cuprous salt. The cuprous 
 compound is quite stable. Cupric cyanide is formed as a 
 yellow precipitate, when potassium cyanide is added to 
 a solution of a copper salt. It soon changes sponta- 
 neously into the compound above mentioned, which has 
 a green color. When this is heated it yields cuprous 
 cyanide, CuCN, which is white. Cuprous cyanide is in- 
 soluble in water. If an excess of potassium cyanide is 
 added to a solution of a copper salt, the precipitate dis- 
 solves in consequence of the formation of double cya- 
 nides similar to the double chlorides. One of these has 
 the composition KCN.CuCN or KCu(CN) 2 . The one 
 formed under ordinary circumstances is SKCN.CuCN or 
 K 3 Cu(CN) 4 . The double cyanides are in general more 
 complicated in composition than the double chlorides, as 
 we shall see in studying those which contain iron such, 
 for example, as the salt already referred to under the 
 name potassium ferrocyanide or yellow prussiate of pot- 
 ash, of the composition K 4 Fe(CN) 6 . 
 
 Cuprous Sulphocyanate, CuSCN", and Cupric Sulpho- 
 cyanate, Cu(SCN) 2 , bear to each other relations similar to 
 those which exist between the cyanides. When potassium 
 sulphocyanate is added to a concentrated solution of a cu- 
 pric salt cupric sulphocyanate is precipitated as a black 
 powder. If the solution is diluted, decomposition into the 
 cuprous salt takes place. When a reducing agent such as 
 sulphur dioxide is added at the same time, only the 
 cuprous salt is formed. ' This is a white, granular pow- 
 der, insoluble in water. 
 
 Cuprous Sulphide, Cu 2 S. This compound occurs in 
 nature, and is known as chalcocite. It is, further, a con- 
 stituent of copper pyrites, which is a compound of cu- 
 prous and ferric sulphides, Cu a S.Fe 9 S, or CuFeS 2 . It 
 
600 INORGANIC CHEMISTRY. 
 
 can be made by heating copper and sulphur together in 
 the right proportions. It has a grayish-black color; 
 does not give up its sulphur, even when heated in hy- 
 drogen ; and is the more stable of the two sulphides of 
 copper. 
 
 Cupric Sulphide, CuS. This is formed as a black pre- 
 cipitate when hydrogen sulphide is passed into a 
 solution of a cupric salt. In water alone cupric sul- 
 phide is somewhat soluble. Hence in washing out a 
 precipitate of copper sulphide with water a little of it 
 will pass through in solution. It also easily undergoes 
 oxid.ation, and, as it forms the sulphate, some is dissolved 
 in this way unless proper precautions are taken. It is 
 slightly soluble in ammonium sulphide, but insoluble in 
 sodium sulphide. The above facts are of importance in 
 the analysis of compounds containing copper, as will 
 readily be seen. When heated, cupric sulphide loses 
 half its sulphur, and is converted into cuprous sulphide. 
 
 Copper-plating. The process of copper-plating con- 
 sists in brief in depositing upon an object a layer of cop- 
 per by putting it in a bath containing some copper salt, 
 and connecting it with one pole of an electric battery. 
 Decomposition of the copper salt takes place, and cop- 
 per is deposited upon the object. Alkaline solutions of 
 the double cyanides are best adapted to the purpose. 
 The process is extensively used in the preparation of 
 electrotype plates. These are plates which are prepared 
 either from wood-cuts or from type by making a mould 
 of gutta-percha, covering this with graphite, and immers- 
 ing the plate thus prepared in the copper-plating bath. 
 The plate thus made is an exact reproduction of the 
 wood-cut or type of which the impression in gutta- 
 percha was taken. 
 
 Reactions which are of Special Value in Chemical 
 Analysis. Potassium or sodium hydroxide forms a blue 
 precipitate which becomes black on standing or when 
 heated. (See Cupric Hydroxide.) 
 
 Ammonia first forms a greenish precipitate, which is a 
 basic salt. With cupric sulphate the reaction takes 
 place thus : 
 
SILVER. 601 
 
 S0 2 <>Cu 
 
 + 2NH 3 + 2H 2 = > S0 2 + (NH 4 ) a S0 4 . 
 
 S0 2 <o>Cu Cu< OH 
 
 If the action is carried farther, the basic salt dis- 
 solves, forming the compound referred to under Cupric 
 Sulphate (which see), the solution being dark blue. 
 
 Potassium or sodium carbonate precipitates the basic 
 carbonate referred to under Cupric Carbonate (which 
 see). The change in color from blue to green which 
 takes place in this precipitate is probably due to a loss 
 of water. 
 
 Potassium ferrocyanide, K 4 Fe(CN) 6 , forms a reddish- 
 brown precipitate, which is the corresponding copper 
 salt, Cu 2 Fe(CN) 6 . This compound is decomposed by 
 caustic alkalies, forming cupric oxide and the corre- 
 sponding alkali salt, Na 4 Fe(CN) 6 or K 4 Fe(CN) 6 . The 
 reactions with potassium iodide, cyanide, and sulpho- 
 cyanide have been explained above. 
 
 In the oxidizing flame the bead of borax or microcosmic 
 salt is greenish blue, while when heated in the reducing 
 flame it appears opaque and red. The red color is due 
 to the reduction of the oxide to copper or cuprous oxide. 
 
 SILVER, Ag (At. Wt. 107.11). 
 
 General. In nearly all the compounds of silver the 
 element is univalent. It, however, forms three oxides of 
 the formulas Ag 4 O, Ag 2 O, and AgO. The compounds 
 correspond closely in many respects to the cuprous com- 
 pounds. There is the same question here as in the case 
 of copper as to whether the molecular weights corre- 
 spond to the simple formulas AgCl, AgBr, AgNO 3 , etc., 
 or to the doubled formulas Ag 2 Cl 2 , Ag 2 Br 2 , Ag 2 (NO 3 ) 2 , 
 etc. There is no evidence at present known by which a 
 decision between the two possibilities can be reached. 
 The simpler formulas will therefore be used here. 
 
 Forms in which Silver Occurs in Nature. Silver occurs 
 to some extent native, but for the most part in combina- 
 tion, particularly with sulphur, and in company with 
 
602 INORGANIC CHEMISTRY. 
 
 lead. The principal ore of silver is the sulphide, Ag 2 S, 
 which occurs in combination with other sulphides, as of 
 lead, copper, arsenic, antimony, etc. The compounds 
 with chlorine, bromine, and iodine are also found, but 
 in smaller quantity than the sulphide. Small quantities 
 of the sulphide are found in almost all varieties of ga- 
 lenite or lead sulphide. 
 
 Metallurgy of Silver. Much of the silver in use is ob- 
 tained from galenite, PbS. This mineral is treated in 
 such a way as to cause the separation of the lead (which 
 see), and the silver is separated from sulphur at the 
 same time. But it is dissolved in a large quantity of 
 lead, and the problem which presents itself to the me- 
 tallurgist is how to separate the small quantity of 
 silver from the large quantity of lead. This is accom- 
 plished by melting the mixture and allowing it to cool 
 until crystals appear. These are almost pure lead. 
 They are dipped out by means of a sieve-like ladle, and 
 the liquid left is again allowed to stand, when another 
 crop of crystals is formed, and can be removed in the 
 same way as before. By this means, and by again melt- 
 ing the crystals removed, allowing the liquid to crystal- 
 lize, and removing the crystals formed, there is finally 
 obtained a product which is rich in silver, but which still 
 contains lead. This is heated in appropriate vessels in 
 contact with the air, when the lead is oxidized, while the 
 silver remains in the metallic state. This method of 
 concentrating by crystallization of lead is known as Pat- 
 tinsoris method. 
 
 Another method of separating lead and silver now ex- 
 tensively used consists in treating the molten alloy with 
 a small quantity of zinc. This takes up all the silver, 
 and the alloy of zinc and silver thus formed is removed, 
 and afterwards treated with superheated steam, by 
 which the zinc is oxidized and the silver left unchanged. 
 
 Some ores of silver are treated in another way, known 
 as the amalgamation process. The ores are mixed with 
 common salt and roasted, when the silver is obtained in 
 the form of the chloride. This is then reduced to silver 
 
METALLURGY OF SILVER. 603 
 
 by means of iron and water, the reaction taking place as 
 represented in the following equation : 
 
 2AgCl + Fe = FeCl 2 + 2Ag. 
 
 The mixture is next treated with mercury, which forms 
 an amalgam with silver, while the other metals present 
 do not combine with the mercury. The amalgam can 
 be separated from the rest of the mass without much 
 difficulty, and when heated to a sufficiently high temper- 
 ature the mercury distils over, leaving the silver. 
 
 A modification of the amalgamation process, known as 
 the American process, consists in grinding the ores very 
 fine, mixing them with sodium chloride, adding roasted 
 copper pyrites, which consists largely of copper sul- 
 phate, and then gradually adding mercury. The silver 
 is slowly converted into silver amalgam. The ex- 
 planation of the process is this : The copper sulphate 
 reacts with the sodium chloride to form cupric chloride 
 and sodium sulphate. Cupric chloride reacts upon the 
 silver sulphide as represented in the equation 
 
 2CuCl a + Ag 2 S = Cu 2 Cl 2 + 2AgCl + S. 
 
 The cuprous chloride thus formed acts upon the rest of 
 the silver sulphide, forming silver chloride and cuprous 
 sulphide : 
 
 Cu 2 Cl 2 + Ag 2 S = Cu 2 S + 2AgCL 
 
 The silver chloride dissolves in sodium chloride, and is 
 then reduced and converted into the amalgam by the 
 mercury. 
 
 The silver in the market is not pure. For chemical 
 purposes it can be purified by dissolving it in nitric 
 acid, precipitating by means of hydrochloric acid, filter- 
 ing and thoroughly washing the chloride, and reducing 
 this either by melting it with sodium carbonate, or by 
 pouring a little dilute hydrochloric acid upon it, and 
 bringing a piece of zinc in contact with it. In the 
 former case the reaction is 
 
 2AgCl + Na 2 CO 3 = 2Ag + C0 2 + O + 2NaCl ; 
 
604 INORGANIC CHEMISTRY. 
 
 in the latter it is 
 
 Zn + 2AgCl = ZnCl 2 + 2Ag. 
 
 Properties. Silver is a white metal with a high lustre, 
 of specific gravity 10.5. It is not acted upon by the air, 
 oxygen, or water. It melts at a lower temperature than 
 copper or gold, the melting-point being about 1000. At 
 the temperature of the oxyhydrogen blowpipe it distils, 
 and in the experiments of Stas on the atomic weights of 
 chlorine and silver the metal used was purified in this 
 way. It is harder than gold and softer than copper, 
 and its "hardness is much increased by the addition of a 
 little copper. It combines very readily with sulphur, 
 forming black silver sulphide, and with chlorine, bro- 
 mine, and iodine. The blackening of silver coins, and 
 other objects carried about the person is caused by the 
 presence of minute quantities of sulphur compounds in 
 the perspiration ; and the blackening of spoons by con- 
 tact with eggs is due to the presence of sulphur in the 
 albumen of the eggs. When pure silver is melted in the 
 air it absorbs about twenty times its volume of oxygen, 
 and this is given off when the metal solidifies, causing in 
 some cases a sputtering of the silver. This phenome- 
 non is observed in the separation of silver from its ores 
 in those processes in which it is necessary to melt the 
 metal. It is known as " spitting." 
 
 At the ordinary temperatures silver is converted into 
 the peroxide, AgO, by ozone. When treated with hy- 
 drochloric acid, the metal becomes covered with a thin 
 layer of the chloride, and no further action takes place, 
 but it is dissolved easily by concentrated sulphuric acid 
 and dilute nitric acid. With the concentrated acids re- 
 duction-products are formed as with copper. Silver is 
 readily dissolved by a solution of potassium cyanide ; 
 hence, such a solution is used in removing stains caused 
 by silver salts. It is not acted upon by the alkaline 
 hydroxides nor by potassium nitrate in the molten con- 
 dition, while platinum is. Therefore, silver vessels are 
 used when it is desired to melt these substances in the 
 
SILVER CHLORIDE. 605 
 
 laboratory, or to evaporate their solution, as in the 
 preparation of the caustic alkalies. 
 
 Allotropic Forms of Silver. M. Carey Lea has dis- 
 covered 'several curious allotropic forms of silver, the 
 principal of which are briefly described by him as fol- 
 lows : "A. Soluble, deep red in solution, mat lilac, blue, 
 or green while moist, brilliant bluish-green metallic 
 when dry. E. Insoluble, derived from A, dark red- 
 dish-brown while moist, when dry somewhat resembling 
 A. 0. Gold silver, dark brown while wet, when dry 
 exactly resembling metallic gold in burnished lumps. 
 Of this form there is a variety which is copper-colored. 
 Insoluble in water; appears to have no corresponding 
 soluble form." 
 
 The form A is soluble in water, and the solution thus 
 formed has a deep red color. The different varieties 
 are formed by the action of reducing agents on solutions 
 of silver salts. For example, the red soluble form is ob- 
 tained by mixing dilute solutions of ferrous citrate and a 
 silver salt. All the allotropic forms of silver are readily 
 changed to the ordinary form. 
 
 Mr. Lea further says : " All the forms of allotropic 
 silver are sensitive to light. A when exposed to the 
 sunlight soon becomes brown. The bright blue-green 
 variety of B is changed into the pure gold-colored variety 
 of C. Other forms of B turn brown on exposure to 
 light." 
 
 The red-yellow variety of C changes to bright gold 
 color. " Continued exposure seems to produce little 
 further change so long as the substance is dry. But if 
 the paper on which the silver is placed is kept moist by 
 a wet pad, with three or four days of good sunshine, the 
 change goes on until the silver becomes perfectly white 
 and is apparently changed to normal silver." 
 
 Alloys of Silver. For practical use, as in making coins 
 and silver-ware, an alloy with copper is used, the pure 
 metal being too soft. The alloy usually contains from 
 7^ to 10 per cent of copper. This alloy is harder than 
 pure silver, and is capable of a higher polish. Silver 
 
606 INORGANIC CHEMISTRY. 
 
 amalgam is an alloy of silver and mercury, which is 
 readily formed by bringing the two metals together. 
 
 Argentous Chloride, Ag 2 Cl or Ag 4 Cl 3 , is formed by treat- 
 ing argentous salts with hydrochloric acid, and, possibly, 
 to some extent when silver chloride, AgCl, is exposed to 
 the light, though this is doubtful. 
 
 Silver Chloride, Argentic Chloride, AgCl, is of special 
 importance on account of its use in photography and in 
 chemical analysis. It occurs in nature to some extent 
 in Mexico and in the United States. It is easily formed 
 as a white precipitate by adding hydrochloric acid to a 
 solution of a silver salt, as, for example, the nitrate. In 
 consequence of its insolubility in water it affords a con- 
 venient means of detecting silver .and chlorine. If 
 allowed to stand in the light it changes color, becoming 
 first violet and finally black. This change in color 
 appears to be due entirely to the reduction of the 
 chloride to the form of metallic silver. Concentrated 
 hydrochloric acid dissolves it somewhat, and from this 
 solution it crystallizes in octahedrons. An aqueous solu- 
 tion of ammonia dissolves it very easily, in consequence 
 of the formation of a compound of the chloride with am- 
 monia analogous to those formed by copper salts. The 
 composition of the compound in the solution is, how- 
 ever, not known. Concentrated solutions of potassium, 
 sodium, and ammonium chlorides dissolve silver chlo- 
 ride, forming double chlorides ; and potassium cyanide 
 also forms an easily soluble double salt with it. The 
 dry compound absorbs ammonia gas, forming a com- 
 pound of the formula 2AgC1.3NH 3 , which readily gives 
 up the ammonia when gently heated. 
 
 Silver Bromide, AgBr, and Silver Iodide, Agl, are very 
 similar to the chloride. Both occur in nature, and both 
 are precipitated from solutions of silver salts by adding 
 the corresponding hydrogen acids. The bromide is less 
 easily soluble in ammonia than the chloride, and the 
 iodide is almost insoluble in it. The bromide is formed 
 by treating the chloride at the ordinary temperature with 
 hydrobromic acid; and the iodide is formed from the 
 chloride and from the bromide by treating these with 
 
8ALT8 OF SILVER IN PHOTOGRAPHY. 607 
 
 hydriodic acid at ordinary temperatures. At higher 
 temperatures, however, both the bromide and iodide are 
 converted into the chloride by hydrochloric acid. Silver 
 dissolves in concentrated hydriodic acid, and from the 
 solution a salt of the formula Agl -\- HI or HAgI a is 
 formed. It seems probable that this is a derivative of 
 the acid H 2 I 2 , from which the double salt KI.AgI is also 
 derived, as indicated in the formula KAgI 2 . Silver 
 bromide at low temperatures is white, but easily changes 
 to yellow, and by exposure it becomes darker, but not as 
 readily as the chloride. The iodide is yellow, and under- 
 goes change in the light only very slowly. The chloride 
 and iodide exist in several modifications, which differ 
 from one another in their conduct towards light, and in 
 their solubility. Probably the differences are due to 
 different complexity of the molecules. Modifications 
 corresponding to the formulas AgCl, Agl, Ag 2 Cl 2 , Ag 2 I 2 , 
 Ag 3 Cl 3 , Ag 3 I 3 , etc., are quite conceivable. The careful 
 study of the effects of light upon the different modifica- 
 tions seems to promise interesting results, which may 
 make it possible to judge as to the relative complexity 
 of the molecules. 
 
 Application of the Chloride, Bromide, and Iodide of 
 Silver in the Art of Photography. The art of photography 
 is based upon the changes which certain compounds, 
 especially salts of silver, undergo when exposed to the 
 light. Silver iodide is best adapted to most purposes. 
 The salt is so changed by the light that when treated 
 with certain compounds, such as ferrous sulphate, pyro- 
 gallic acid, etc., called " developers," a deposit of finely 
 divided silver is formed upon the plate in those places 
 affected by the light. A plate of glass or a sheet of 
 properly prepared paper is covered in the dark with a 
 thin layer of a salt of silver. The plate is then exposed 
 in the camera to the action of the light which is reflected 
 from the object to be photographed. According to the 
 intensity of the light given off from the various parts of 
 the object, the change of the silver salt takes place to a 
 greater or less extent, and thus a perfect image of the 
 object is impressed upon the plate. But after the action 
 
608 INORGANIC CHEMISTRY. 
 
 of the " developer" is complete there is still upon the 
 plate unchanged silver salt, and if this were now exposed 
 to the light it would undergo change and the image 
 would be obliterated. To remove this salt the plate is 
 washed with a solution of sodium thiosulphate, Na 2 S 2 O 3 
 (hyposulphite), which dissolves the salt in consequence 
 of the formation of a double salt of the formula 
 2Na 2 S 2 O 3 .Ag 2 S 2 O 3 , which is readily soluble in water. 
 
 Silver Triazoate, AgN 3 . This is derived from triazoic 
 acid (which see). It is formed by adding a solution of the 
 acid to a solution of a silver salt. It is extremely explo- 
 sive and should be dealt with very cautiously. Serious 
 accidents have been caused by it. In appearance it 
 resembles silver chloride, but it does not darken when 
 exposed to the light. 
 
 Silver Oxide, Ag 2 O. The principal compound of silver 
 and oxygen is that which has the composition Ag 2 O, and 
 in which the silver is univalent, as it is in its compounds 
 with chlorine, bromine, and iodine. It is formed when 
 a soluble hydroxide is added to a solution of a silver 
 salt, and also by the action of concentrated solutions of 
 the caustic alkalies on silver chloride. It is easily de- 
 composed by heat and by reducing agents. 
 
 Other Oxides of Silver. Besides the ordinary oxide, 
 silver forms a sub-oxide, Ag 4 O, corresponding to the sub- 
 oxide of copper, Cu 4 O, and a peroxide of the formula 
 AgO (or Ag 4 O 3 ), which is perhaps analogous to cupric 
 oxide. 
 
 Sulphides of Silver. As has been stated, silver occurs 
 in nature mostly in combination with sulphur as silver 
 glance, Ag 2 S, which is in many minerals in combination 
 with other sulphides. Examples of such double sul- 
 phides are the minerals stromeyerite, Cu 2 S.Ag 2 S, and 
 pyrargyrite, 3Ag 2 S.Sb 2 S 8 . 
 
 Silver Nitrate, Argentic Nitrate, AgNO 3 . This salt is 
 formed by dissolving silver, or silver oxide, in nitric acid, 
 evaporating to dryness, and heating until the salt is 
 melted. It crystallizes in colorless rhombic plates. It 
 is not changed in the light unless it comes in contact 
 
COMPOUNDS OF SILVER. 609 
 
 with organic substances, when it is reduced and metallic 
 silver deposited. Hence the solution produces black 
 spots on the fingers and clothing. As it melts easily, it 
 is generally cast in small cylindrical moulds, and is found 
 in the market in the form of thin sticks, and is known as 
 lunar caustic. It disintegrates flesh, and is used in sur- 
 gery as a caustic to remove superfluous growths. Owing 
 to the formation of a dark deposit when the salt is ex- 
 posed to the light, it is used as a constituent of indelible 
 inks. The dry nitrate absorbs ammonia and forms the 
 compound AgNO 3 + 3NH 3 ; in concentrated solution the 
 compound AgNO s + 2NH 3 is formed. 
 
 Silver Cyanide, AgCN, is formed as a caseous pre- 
 cipitate when a solution of hydrocyanic acid is added to 
 a solution of silver nitrate. It does not change color in 
 the light, is soluble in ammonia, but hot in nitric acid. 
 It readily forms double cyanides with the cyanides of 
 other metals. Of these, the salt with potassium cyanide, 
 KAg(CN) 2 or KCN.AgCN, may be mentioned. 
 
 Silver Sulphocyanate, AgSCN, is very similar to the 
 cyanide, and is formed when solutions of silver nitrate 
 and potassium or ammonium sulphocyanate are brought 
 together. It is soluble in an excess of the soluble cy- 
 anides, double salts similar to the double cyanides being 
 formed. 
 
 Borates of Silver. When a cold concentrated solution 
 of sodium metaborate, NaBO 2 , is mixed with a similar 
 solution of silver nitrate a precipitate of silver meta- 
 borate, AgBO 2 , containing some silver oxide is formed. 
 When dilute solutions of the two compounds are mixed 
 a precipitate of silver oxide is formed ; so, also, silver 
 metaborate is decomposed by water into boric acid and 
 silver oxide, and when the solution in which the pre- 
 cipitate is suspended is boiled the same change takes 
 place. Further, when cold concentrated solutions of 
 silver nitrate and borax are mixed, silver octoborate, 
 Ag 6 B 8 O 15 , is precipitated, and this is mixed with some 
 silver oxide. When the solution is boiled, the silver salt 
 is decomposed into boric acid and silver oxide. When 
 
610 INORGANIC CHEMISTRY. 
 
 the solutions of borax and silver nitrate are mixed hot, 
 the precipitate is the metaborate of silver. 
 
 Reactions which are of Special Value in Chemical 
 Analysis. Hydrochloric acid precipitates insoluble silver 
 chloride from solutions of silver salts, as silver nitrate. 
 
 Soluble hydroxides precipitate silver oxide, not the hy- 
 droxide. Ammonia redissolves the precipitate in conse- 
 quence of the formation of a compound of the oxide with 
 ammonia of the composition Ag 2 O.2NH 3 . In dry con- 
 dition this salt is very explosive, and is known as ful- 
 minating silver. 
 
 Soluble carbonates precipitate the carbonate, Ag 2 CO 3 , 
 which has a yellowish-white color. 
 
 Ammonium carbonate redissolves the precipitate formed 
 by it. 
 
 Sodium phosphate, HNaJPO 4 , gives a precipitate of the 
 normal salt Ag 3 PO 4 , which is yellow. 
 
 Potassium f err ocyanide precipitates white silver ferro- 
 cyanide, Ag 4 Fe(CN) 6 . 
 
 Potassium ferricyanide, K 3 Fe(CN) 6 , gives the corre- 
 sponding silver salt, which is reddish brown. 
 
 Potassium chromate or potassium dichromate (which see) 
 gives a brownish-red precipitate of silver chromate. 
 
 GOLD, Au (At. Wt. 195.74). 
 
 General. Gold forms two series of compounds, in one 
 of which it is univalent and in the other trivalent. In 
 this respect it differs from the other members of the 
 group. Examples of the compounds belonging to the 
 two series are represented by the following formulas : 
 
 AuCl AuCl 3 
 
 AuBr AuBr 3 
 
 Au 2 O Au 2 O 3 
 
 Those of the first series are called aurous compounds t 
 those of the second series auric compounds. The basic 
 character of gold is very weak, so that salts of the ordi- 
 nary acids, as sulphuric, nitric, carbonic, etc., are not 
 
GOLD. 611 
 
 known. On the other hand, its higher oxide and hy- 
 droxide, Au(OH) 3 , have acid properties, and form salts 
 similar in composition to the meta-aluminates MAlO a , 
 and the metaborates MBO 2 . These are the aurates, of 
 which potassium aurate, KAuO 2 , is an example. So, also, 
 the chloride combines readily with the chlorides of potas- 
 sium and sodium, forming the chlor-aurates, KAuCl 4 , and 
 NaAuCl 4 , which are perfectly analogous to the aurates. 
 Further, the chloride and bromide combine respectively 
 with hydrochloric and hydrobromic acids, forming the 
 crystallized compounds HAuCl 4 + 4H 2 O and HAuBr 4 + 
 5H 2 O, which are plainly the acids from which the chlor- 
 aurates and the brom-aurates are derived. 
 
 Besides the compounds of gold in which the element 
 is univalent and those in which it is tnvalent, the chlo- 
 ride AuCl 3 , and the bromide, AuBr 2 , have been described, 
 but their existence is doubtful. 
 
 Forms in which Gold occurs in Nature. Gold is gen- 
 erally found in nature in the native condition a fact 
 which is undoubtedly due to the chemical inactivity of 
 the element. That which is found in nature is never 
 pure, but contains silver, and also, in different localities, 
 iron, copper, and other metals. It is also found to some 
 extent in combination with tellurium in the compounds 
 AuTe 2 and (AuAg) 2 Te 3 . Native gold is frequently found 
 enclosed in quartz, or more commonly in quartz sand. 
 The principal localities in which it is found are California 
 and some of the other Western United States, and 
 Australia, Hungary, Siberia, and Africa. 
 
 Metallurgy of Gold. From the chemical point of view 
 the metallurgy of gold is in general very simple. There 
 are two kinds of gold mining called placer mining and 
 vein mining. In the former the earth and sand which 
 contain gold are washed with water, which carries away 
 the lighter particles, and leaves the gold mixed with 
 other heavy materials. This mixture is then treated 
 with mercury, which forms an amalgam with the gold, as 
 it does with silver, and when this is placed in a prop- 
 erly constructed retort and heated, the mercury passes 
 over and leaves the gold behind. If silver is present, as 
 
612 INORGANIC CHEMISTRY. 
 
 is frequently the case, this is separated with the gold, 
 In vein mining the gold ores are taken out of veins in the 
 earth, and the gold separated by grinding the ores and 
 treating them with mercury, as in the last stage of placer 
 mining. Hydraulic mining is a modification of ordinary 
 placer mining. It consists in forcing water under pres- 
 sure against the sides of hills and mountains in which 
 gold occurs loosely mixed with the earth. The earth is 
 thus carried away and the heavier gold is deposited in 
 sluices. 
 
 Some ores, those especially which contain tellurium, 
 cannot be satisfactorily treated by the amalgamation 
 process, and a method involving the use of potassium 
 cyanide has been devised for them. In a solution of this 
 salt gold dissolves, and from this solution it can be sep- 
 arated in various ways. This method has come into ex- 
 tensive use of late years. 
 
 Another process that is extensively used in the treat- 
 ment of ores that do not give their gold to mercury is 
 known as the chlorination process. This consists in treat- 
 ing the finely ground ore with chlorine made from 
 bleaching powder and sulphuric acid, and then pre- 
 cipitating the gold from the solution of the chloride by 
 means of hydrogen sulphide. From the sulphide the 
 metallic gold can be easily obtained. 
 
 The gold obtained by any of the above methods is not 
 pure. It can be separated from silver by dissolving it 
 in aqua regia, evaporating so as to drive off the nitric 
 acid, then diluting, and treating with a reducing agent, 
 when metallic gold is precipitated. Thus when ferrous 
 sulphate is used the following reaction takes place : 
 
 3FeSO 4 + AuCl 3 = Fe 2 (SO 4 ) 3 + FeCl 3 + Au. 
 
 Another method of separating silver from an alloy 
 with gold consists in treating the metal with nitric acid 
 or with boiling concentrated sulphuric acid, which dis- 
 solves the silver and leaves the gold. This process is 
 not satisfactory, however, unless the amount of gold in 
 the alloy is less than 25 per cent. If the proportion of 
 gold is greater than this, the alloy is melted with silver 
 
PROPERTIES OF GOLD ALLOYS CHLORIDES. 613 
 
 enough to bring the percentage of gold down to that 
 mentioned. This is known as " quartation" 
 
 Properties. Gold is a yellow metal with a high lustre. 
 It is quite soft, and extremely malleable, so that it is 
 possible to make from it sheets the thickness of which is 
 not more than 0.000002 millimeter. Thin sheets are 
 translucent, and the transmitted light appears green. 
 Its specific gravity is 19.3 ; its melting-point higher than 
 that of copper, being about 1200. It crystallizes in the 
 regular system. Gold combines directly with chlorine, 
 but not with oxygen. The three acids, hydrochloric, 
 nitric, and sulphuric, do not act upon it ; but aqua 
 regia dissolves it, forming auric chloride, AuCl 3 , in con- 
 sequence of the evolution of nascent chlorine. Molten 
 caustic alkalies and their nitrates act upon it, probably 
 in consequence of the tendency to form aurates. 
 
 Alloys of Gold. The principal alloy of gold is that 
 which contains copper. The standard gold coin of the 
 "United States contains nine parts of gold to one of cop- 
 per. The composition of gold used for jewelry is usually 
 stated in terms of carats. Pure gold is 24-carat gold ; 
 20-carat gold contains 20 parts of gold and 4 parts of 
 copper ; 18-carat gold contains 18 parts of gold and 6 
 parts of copper, etc. Copper gives gold a reddish color, 
 and makes it harder and more easily fusible. Gold is 
 also alloyed with silver ; and the alloy with mercury, 
 known as gold-amalgam, is extensively used in the pro- 
 cesses for extracting gold from its ores. 
 
 Chlorides of G-old. When gold is dissolved in aqua 
 regia it is converted into auric chloride, AuGl 3 ; and if this 
 solution is evaporated a part of the chloride is decom- 
 posed into aurous chloride, AuCl, and chlorine. When gold 
 is treated with dry chlorine it yields a mixture of auric 
 chloride and metallic gold. This was formerly held to 
 be a chloride of the formula AuCl 2 , but the most careful 
 investigations on the subject have shown that this does 
 not exist. Auric chloride can be obtained in crystal- 
 lized form, the crystals having the composition AuCl 8 -[- 
 2H 2 O. When anhydrous auric chloride is heated to 
 185, it loses chlorine and is converted into aurous chlo- 
 
614 INORGANIC CHEMISTRY. 
 
 ride, AuCl. This yields auric chloride and gold when 
 treated with water. When treated with a solution of 
 stannous chloride a solution of auric chloride gives a 
 purple-colored precipitate, known as the purple of Cas- 
 sius, which appears to consist of finely-divided gold. 
 
 Chlor-auric Acid and its Salts. When a solution of 
 gold in aqua regia containing a large excess of hydro- 
 chloric acid is evaporated a crystallized product of the 
 formula HAuCl 4 + 4H 2 O, or AuCl 3 .HCl + 4H 3 O, is ob- 
 tained. This is chlor-auric acid. It must be regarded 
 as belonging to the same class as fluosilicic acid and 
 the chloro-acids, from which the double chlorides of 
 magnesium, aluminium, copper, etc., are derived. Ac- 
 cordingly its constitution is expressed by the formula 
 
 /ci 
 
 Au^-Cl , being similar to that of the acid from which 
 
 /Cl 
 potassium chlor-aluminate is derived, A1C1 . The 
 
 \(C1,)H 
 
 potassium salt, KAuCl 4 , is obtained by mixing together 
 solutions of auric and potassium chlorides. 
 
 Cyan-auric Acid, HAu(CN) 4 , is perfectly analogous to 
 chlor-auric acid. It is formed by treating the potassium 
 salt, KAu(CN) 4 , with silver nitrate, which gives the silver 
 salt, and then decomposing this with hydrochloric acid. 
 The potassium salt is obtained by mixing solutions of 
 auric chloride and potassium cyanide. The salts of a 
 cyan-aurous acid, HAu(CN) 2 , are also known. 
 
 Auric Hydroxide, Au(OH) 3 . This compound is formed 
 by treating a solution of auric chloride with an excess of 
 magnesia or with sodium hydroxide, and afterwards with 
 sodium sulphate. It is a yellow or brown powder. 
 When exposed to the light it is decomposed with evolu- 
 tion of oxygen. When heated to 100 it yields auric 
 oxide, Au 2 O 3 , and when this is heated to a higher tempera- 
 ture it loses all its oxygen. Aurous oxide, Au 2 O, is formed 
 by treating aurous chloride with caustic potash. It is 
 easily decomposed by heat into gold and oxygen. 
 
 Auric hydroxide dissolves in the soluble hydroxides 
 just as aluminium hydroxide does, and from the solutions 
 
AURIC HYDROXIDE GOLD SULPHIDE. 615 
 
 salts known as the aurates are obtained. In composition 
 these are analogous to the meta-aluminates. Potassium 
 aurate, for example, has the composition KAuO 2 . The 
 analogy between some of the compounds of aluminium 
 and those of gold is shown in the following table : 
 
 A1A Au 2 3 
 
 A1(OH) 3 Au(OH) 3 
 
 A1C1, AuCl 3 
 
 /Cl /Cl 
 
 Al(-Cl Au^-Cl 
 \(C1 2 )K \(C1 3 )K 
 
 Gold Sulphide, Au 2 S 2 . This compound is precipitated 
 together with sulphur from cold solutions of gold salts 
 by means of hydrogen sulphide, and forms a brownish 
 black mass. It forms soluble compounds with the sul- 
 phides of the alkali metals. 
 
 When hydrogen sulphide is passed into hot solutions 
 of gold salts aurous sulphide, Au 2 S, is thrown down as 
 a steel-gray substance. This is soluble in pure water 
 and is reprecipitated by hydrochloric acid, 
 
CHAPTER XXIX. 
 
 ELEMENTS OF FAMILY II, GROUP B: 
 ZINC CADMIUM MERCURY. 
 
 General. There is a very strong resemblance between 
 the first two elements of this group and magnesium, 
 while mercury, in a general way, resembles the first two 
 members of the copper group. Just as gold in the cop- 
 per group furnishes a greater variety of compounds than 
 the first two members of that group, so mercury fur- 
 nishes a greater variety of compounds than the other 
 members of the group to which it belongs. Zinc and 
 cadmium, like magnesium, give only one class of com- 
 pounds and in these they are bivalent. The general for- 
 mulas of some of the principal ones are : 
 
 MC1 2 , M(OH) 2 , MO, MSO 4 , MCO 8 , MS. 
 
 Mercury, on the other hand, furnishes two series of com- 
 pounds, known as the mercurous and mercuric compounds, 
 which correspond closely to the two series of copper 
 salts. The power to form compounds belonging to both 
 series is more strongly developed in mercury than in. 
 copper. Examples of the two classes are represented in 
 the following formulas : 
 
 Mercurous Compounds. Mercuric Compounds. 
 
 HgCl HgCl 2 
 
 Hgl HgI 2 
 
 Hg 2 HgO 
 
 HgN0 3 Hg(N0 3 ) 2 , etc. 
 
 Just as the first member of Group A, Family II, gluci- 
 num, shows a somewhat acidic character in its hydrox- 
 ide, while the other members of that group do not; so 
 also the first member of Group Bj Family II, zinc, is 
 
 (616) 
 
ZING. 617 
 
 acidic, while the other members of the group are not. 
 Glucinum hydroxide and zinc hydroxide dissolve in 
 caustic alkalies, forming glucinates and zincates ; while 
 the hydroxides of all the other members of the two 
 groups of this family are insoluble in caustic alkalies. 
 
 ZINC, Zn (At. Wt. 64.91). 
 
 General. Zinc, in almost all its compounds, exhibits 
 a close resemblance to magnesium. It always acts as a 
 bivalent element. 
 
 Forms in which, it occurs in Nature. Zinc occurs in 
 nature in combination principally as the carbonate, or 
 smithsonite, ZnCO 3 ; as the sulphide, or sphalerite, ZnS ; 
 and as the silicate, Zn 2 SiO 4 . Among other compounds 
 of zinc found in nature are gahnite, Zn(AlO 2 ) 2 , and frank- 
 linite, which contains the compound Zn(FeO 2 ) 2 with 
 Fe(Fe0 2 ) 2 . 
 
 Metallurgy. The metallurgy of zinc is much simpler 
 than that of magnesium, for the reason that the ores are 
 easily converted into the oxide by roasting, and the oxide 
 is easily reduced by heating it with charcoaL Owing to 
 the volatility of the metal the vessels in which the reduc- 
 tion is effected must be so constructed as to facilitate the 
 condensation of the vapors. The vessels used are either 
 earthenware muffles or tubes, open at one end and con- 
 nected with iron receivers. At first the zinc vapor is 
 condensed in the form of a fine dust, as in the case of 
 sulphur. This forms the commercial product called 
 zinc dust. It always contains zinc oxide. Afterwards 
 the zinc condenses to the form of a liquid, and this is 
 cast in plates. The zinc thus obtained is not pure, but 
 contains lead and iron, and sometimes arsenic and cad- 
 mium. It is called spelter. By repeated distillation it 
 can be obtained pure. When distilled under diminished 
 pressure, it is deposited in beautiful lustrous crystals, 
 the forms of which are extremely complicated. 
 
 Properties. Zinc has a bluish-white color and a high 
 lustre. The crystals above referred to, which are per- 
 fectly pure zinc, have a brilliant lustre, and do not 
 
618 INORGANIC CHEMISTRY. 
 
 change in the air. At different temperatures zinc has 
 markedly different properties. At ordinary temperatures 
 it is quite brittle ; at 100-150 it can be rolled out in 
 sheets, but above 200 it becomes brittle again. It melts 
 at 433, and boils at 1040. When heated in the air it 
 takes fire, and burns with a bluish flame, forming zinc 
 oxide. This can be shown by means of the oxyhydro- 
 gen blowpipe. In dry air it does not change. Ordinary 
 zinc dissolves in all the common acids, usually with evo- 
 lution of hydrogen. In the case of nitric acid, however, 
 the hydrogen acts upon the acid, reducing it to ammonia. 
 The purer the zinc the less readily is it acted upon by 
 sulphuric acid, and the pure crystals above referred 
 to are scarcely acted upon at all by this acid. Zinc 
 also dissolves in the caustic alkalies, forming zincates. 
 Pure zinc can be made to act upon sulphuric acid by 
 adding a few drops of platinum chloride. 
 
 Applications. Zinc is extensively used as sheet-zinc, 
 in making galvanic batteries, for galvanizing iron, etc. 
 Zinc dust is a very efficient reducing agent, either in al- 
 kaline or in acid solution. With caustic alkalies as, for 
 example, with potassium hydroxide it gives hydrogen ' 
 and a zincate : 
 
 Zn + 2KOH = Zn(OK) 2 + H 8 . 
 
 With sulphuric acid also it gives hydrogen readily. 
 Zinc is used in the preparation of important alloys. 
 
 Alloys. Iron covered with a layer of zinc is known 
 as galvanized iron. As has been mentioned, zinc is a 
 constituent of brass. It combines readily with mercury 
 to form zinc amalgam, and this fact is taken advantage 
 of for the purpose of preserving the zinc plates in gal- 
 vanic batteries. Zinc plates covered with a layer of 
 the amalgam are acted upon much more slowly than 
 zinc itself. The amalgamation is effected by cleaning 
 the zinc, dipping it in dilute sulphuric acid, and rubbing 
 mercury over the surface with a brush or a piece of 
 cloth. 
 
 Zinc Chloride, ZnCl 2 . This is prepared by treating 
 zinc with chlorine, or by dissolving zinc in hydrochloric 
 
ZINC CHLORIDE ZINC HYDROXIDE. 619 
 
 v 
 
 acid, evaporating to dryness, and distilling the residue. It 
 is a white deliquescent mass. From a very concentrated 
 solution in hydrochloric acid it is obtained in crystals 
 of the composition ZnCl 2 -f- H 2 O. When the solution is 
 evaporated there is always some decomposition into basic 
 chlorides, the hydroxide, and oxide. The basic chloride 
 is formed thus : 
 
 HHO = Zn< + HC1 ; 
 
 the hydroxide thus : 
 
 HHO = Zn< HC1 
 
 and by higher heating the hydroxide yields the oxide 
 and water : 
 
 Zn< OH = Zn + H A 
 
 The chloride has a marked affinity for water, and is used 
 in the laboratory, as sulphuric acid and phosphorus 
 pentoxide are, for the purpose of extracting the elements 
 of water from compounds. It has a caustic action, and 
 is used in surgery on this account. Further, it acts as a 
 disinfectant, and its solution is used for the purpose of 
 preserving wood, particularly railroad sleepers, from de- 
 cay. 
 
 The chloride readily forms double chlorides like those 
 formed by magnesium chloride. Examples of these are 
 the compounds of the formulas ZnCl 2 .2KCl or K 2 ZnCl 4 , 
 ZnCl 2 .2NaCl or Na 2 ZnCl 4 , etc. The double chloride 
 with ammonium chloride, (NH 4 ) 2 ZnCl 4 , is formed by 
 mixing a solution of zinc in hydrochloric acid with a 
 solution of ammonium chloride. This is used in solder- 
 ing, as it cleans the surface of the metal, in consequence 
 of the action of the zinc chloride on the oxides. It also 
 absorbs ammonia, forming compounds analogous to those 
 formed by cupric and cuprous chlorides. 
 
 Zinc Hydroxide, Zn(OH) 2 , is precipitated as a white 
 amorphous powder when a soluble hydroxide is added 
 to a solution of a zinc salt. It is redissolved in an ex- 
 
620 INORGANIC CHEMISTRY. 
 
 cess of the reagent, and the zincate thus formed is de- 
 composed on boiling, the hydroxide being reprecipitated. 
 
 Zinc Oxide, ZnO, is formed in very finely divided con- 
 dition by burning zinc in the air. The product is known 
 as Flores zinci, and is sometimes called philosopher's 
 wool. It is also formed by heating the carbonate or ni- 
 trate, and is found in nature mixed with or in combina- 
 tion with oxide of manganese, Mn 3 O 4 . It is prepared on 
 the large scale both by burning zinc and by heating the 
 basic carbonate, which is formed by adding sodium car- 
 bonate to a solution of zinc sulphate. It is a white pow- 
 der, which turns yellow when heated. Its chief use is as 
 a constituent of paint under the name of zinc ivhite. 
 
 Zinc Sulphide, ZnS. This compound occurs in nature, 
 and is known as zinc blende. The mineral always con- 
 tains a sulphide of iron, and also a small quantity of 
 cadmium sulphide. When hydrogen sulphide is passed 
 into a solution of a zinc salt only a part of the zinc is 
 thrown down as the sulphide, if the salt used is one of a 
 strong acid, like sulphuric, nitric, or hydrochloric acid. 
 The reason of this is that the sulphide is soluble in these 
 acids, even when they are very dilute. In the reaction 
 the acid is set free, and although some sulphide is thrown 
 down, the action soon stops : 
 
 ZnSO 4 + H 2 S = ZnS + H 2 SO 4 . 
 
 If the acetate of zinc is used the precipitation is com- 
 plete, because dilute acetic acid does not dissolve zinc 
 sulphide. If sodium or potassium acetate is added to a 
 solution of a neutral salt of zinc, hydrogen sulphide pre- 
 cipitates all the zinc, for the reason that the acid which is 
 first set free acts upon the acetate and is itself neu- 
 tralized, while acetic acid is then set free. Thus, when 
 hydrogen sulphide acts upon a solution of zinc sulphate 
 containing sodium acetate the action involves two steps, 
 as represented in the two equations : 
 
 ZnSO 4 + H 2 S = ZnS + H 2 SO 4 ; 
 2NaC,H 3 2 + H 2 S0 4 = Na 2 SO 4 + 2C 2 H 4 O a . 
 
COMPOUNDS OF ZINC. 621 
 
 As fast as the sulphuric acid is formed it acts upon the 
 aeetate, and is thus prevented from dissolving the sul- 
 phide. 
 
 The sulphide is, further, completely precipitated by 
 soluble sulphides, as potassium and ammonium sul- 
 phides. Obtained by precipitation, zinc sulphide is a 
 white amorphous substance. 
 
 Zinc Sulphate, ZnSO 4 . This salt is readily formed by 
 oxidation of the sulphide, and is hence found in nature 
 accompanying the sulphide. It is manufactured by care- 
 fully roasting zinc blende, and extracting with water. It 
 crystallizes from the solution in water in large rhombic 
 prisms of the composition ZnSO 4 -|- 7H Q O. Like mag- 
 nesium sulphate, it easily loses six molecules of water, 
 but the last one is removed with difficulty. It appears, 
 therefore, that the constitution of the salt should be ex- 
 
 roH 
 
 OTT 
 
 pressed by the formula SO \ n . Zinc sulphate, as 
 
 [o >Zn 
 
 has been stated (see p. 596), is commonly called white 
 vitriol. It is easily reduced when heated with charcoal. 
 The salt is used extensively in the preparation of cotton- 
 prints and in medicine. 
 
 Zinc Carbonate, ZnCO 3 , occurs in nature as smithson- 
 ite. The precipitate formed by adding a solution of a 
 soluble carbonate to a solution of a zinc salt is generally 
 a basic carbonate, but the composition varies according to 
 the conditions. Dilute solutions of sodium carbonate and 
 
 Zn<g H 
 
 >co 
 
 zinc sulphate give mainly the compound Zn<^ 
 
 . Zn <OH 
 
 or Zn(OH) 2 .ZnCO 3 .ZnO. Much more complicated salts 
 are, however, usually obtained. With ammonia, zinc 
 carbonate forms a soluble compound. 
 
 Reactions which are of Special Value in Chemical Anal- 
 ysis. The principal reactions which are made use of 
 for the purpose of separating zinc from other elements 
 have been mentioned above. These are the reactions 
 
622 INORGANIC CHEMISTRY. 
 
 with hydrogen sulphide and ammonium sulphide ; with 
 potassium and sodium hydroxide ; with potassium and 
 sodium carbonates ; and with ammonium carbonate. An- 
 other reaction which is used in analysis is that which 
 takes place when zinc salts are heated on charcoal before 
 the blowpipe, moistened with cobalt nitrate, and ignited. 
 Under these circumstances a green-colored mass known 
 as Rinmann's green is formed, which is probably a zin- 
 
 cate of cobalt, Zn<Q>Co. 
 
 CADMIUM, Cd (At. Wt. 111.10). 
 
 General. The compounds of cadmium are very similar 
 to those of zinc and magnesium. The element occurs in 
 nature in much smaller quantity than either of these,, 
 frequently in company with zinc, and its compounds ara 
 not as frequently met with. It is always bivalent. A 
 mineral known as greenockite is cadmium sulphide,. 
 CdS. 
 
 Preparation and Properties. Cadmium is obtained 
 principally from different varieties of zinc blende, and 
 separates with the zinc. Being more volatile than zinc, 
 it passes over first when the mixture is distilled. From 
 this first distillate, which contains the oxides of zinc and 
 cadmium, the metals are reduced by heating with char- 
 coal. It has a color like that of tin, and is harder than 
 tin. According to the specific gravity of its vapor, its 
 molecule is identical with its atom, for the molecular 
 weight is approximately 111. 
 
 Cadmium chloride, CdCl 2 , like zinc chloride, is volatile ; 
 the sulphate crystallizes well, but is not analogous in 
 composition to the sulphates of magnesium and zinc, as 
 the composition of the crystallized salt is represented 
 by the formula 3CdSO 4 -f- 8H 2 O ; the normal carbonate, 
 CdCO 3 , is precipitated by soluble carbonates. 
 
 Cadmium Sulphide, CdS, is one of the most character- 
 istic compounds of the element. It is a beautiful yellow 
 substance, which is thrown down from a solution of a 
 cadmium salt by hydrogen sulphide. While it dissolves 
 
CADMIUM MERCURY. G23 
 
 in concentrated acids it does not dissolve in dilute acids, 
 and it is therefore completely precipitated by hydrogen 
 sulphide. It is used as a constituent of yellow paints. 
 
 Cadmium Cyanide, Cd(CN) 2 , is formed as a white pre- 
 cipitate when potassium cyanide is added to a fairly 
 concentrated solution of a cadmium salt. It dissolves 
 in an excess of potassium cyanide in consequence of the 
 formation of the compound K 2 Cd(CN) 4 . 
 
 Analytical Reactions. Cadmium, as has just been 
 stated, is precipitated by hydrogen sulphide. It is 
 thrown down together with the other elements of the 
 hydrogen sulphide group (see p. 198). As the sulphide 
 is not soluble in ammonium sulphide, it is easily sepa- 
 rated from those of arsenic, antimony, and tin by treating 
 with this reagent, when it is left undissolved in company 
 with the sulphides of mercury, lead, bismuth, and copper. 
 The double salt of cuprous cyanide and potassium cya- 
 nide is not decomposed by hydrogen sulphide, whereas 
 the corresponding salt of cadmium is decomposed by it, 
 and the yellow sulphide is precipitated. 
 
 The hydroxide of cadmium differs from that of zinc in 
 not having acid properties. It does not dissolve in the 
 caustic alkalies. 
 
 MEKCUKY, Hg (At. "Wt. 198.49). 
 
 General. As already stated, mercury yields two series 
 of compounds, known as mercurous and mercuric com- 
 pounds, which are analogous to the two series of copper 
 compounds. While, however, copper forms with the 
 oxygen acids only such salts as belong to the cupric 
 series, as CuSO 4 , Cu(NO 3 ) 2 , etc., mercury forms salts be- 
 longing to both series. There is, for example, a mer- 
 curous nitrate, HgNO 3 , and a mercuric nitrate, Hg(NO 3 ) 2 ; 
 a mercurous sulphate, Hg 2 SO 4 , and a mercuric sulphate, 
 HgSO 4 , etc. The mercurous compounds are readily con- 
 verted into the mercuric compounds by the action of 
 oxidizing agents, and the mercuric are converted into 
 mercurous compounds by the action of reducing agents. 
 The action will be treated of under the individual com- 
 pounds. The question as to the correct formula of the 
 
624 INORGANIC CHEMISTHY. 
 
 rnercurous salts is in the same condition as that in regard 
 to the formula of cuprous salts, with this difference : the 
 molecular weight of mercurous chloride leads to the 
 formula HgCl, but there is evidence that when the chlo- 
 ride is heated some mercury is set free, and this has led 
 to the suggestion that the molecule corresponds to the 
 formula Hg 2 Cl 2 , and that the compound breaks down 
 into mercury and mercuric chloride when heated. It is, 
 however, quite possible that the compound has the sim- 
 pler formula, and that this under the influence of heat 
 is partly decomposed, as represented in the equation 
 
 2HgCl := HgCl, + Hg. 
 
 The fact that mercury is set free is, therefore, by no 
 means satisfactory evidence that the formula of mer- 
 curous chloride is Hg 2 Cl 2 , and in the present state of 
 the inquiry it is perfectly justifiable to write the formula 
 HgCl. 
 
 Forms in which Mercury occurs in Nature. Mercury 
 occurs native to some extent, but principally in the 
 form of the sulphide, HgS, which is known as cinnabar. 
 This is sometimes found crystallized, but generally amor- 
 phous. The chief localities are Idria, Almaden in Spain, 
 and New Almaden in California. 
 
 Metallurgy of Mercury. In order to obtain mercury 
 from the sulphide this is roasted in vessels so constructed 
 as to condense and collect the vapor of mercury given 
 off. In the roasting process the sulphur is oxidized to 
 sulphur dioxide, which of course escapes. In some 
 places the ore is mixed with limestone and distilled from 
 clay or iron retorts, when the mercury passes over. 
 Crude mercury is redistilled in order to purify it. It is 
 also purified by treating it with dilute nitric acid or with 
 a solution of ferric chloride. 
 
 Properties. Mercury is a silver-white metal of a high 
 lustre. At ordinary temperatures it is liquid, though at 
 39.5 it becomes solid. Its specific gravity is 13.5959. 
 It does not change in the air at ordinary temperatures. 
 It boils at 357.25, and is converted into a colorless vapor, 
 the specific gravity of which leads to the conclusion that, 
 
AMALGAMS. 625 
 
 as in the case of cadmium, the molecule and atom are 
 identical, or that the molecule consists of only one atom. 
 It is insoluble in hydrochloric acid and in cold sulphuric 
 acid ; but dissolves in hot concentrated sulphuric acid, 
 and is easily soluble in nitric acid. The vapor of mer- 
 cury is very poisonous. 
 
 Applications. Mercury is extensively used in the 
 manufacture of thermometers, barometers, etc. ; as tin- 
 amalgam for mirrors ; and in the processes by which 
 gold and silver are obtained from their ores. 
 
 Amalgams. The alloys which mercury forms with other 
 metals are called amalgams. These compounds are gen- 
 erally obtained without difficulty simply by bringing 
 mercury in contact with other metals. Among the 
 amalgams which are of chief interest are those of sodium, 
 ammonium, silver, and gold. Sodium amalgam is made by 
 bringing mercury and sodium together. A crystallized 
 amalgam containing the constituents in the proportions 
 represented in the formula Hg 6 Na has been obtained. 
 Generally, sodium amalgam is easily decomposed by 
 water, the mercury separating in the free state and the 
 sodium acting upon the water, forming hydrogen and 
 sodium hydroxide. It is much used in the laboratory as 
 a convenient means of producing hydrogen in alkaline 
 solutions. It serves as an excellent reducing agent in 
 some cases. Ammonium amalgam has already been 
 spoken of under the head of Ammonia (which see). It 
 is a curious substance, which is formed when an electric 
 current acts upon a solution of ammonia containing some 
 mercury which is connected with the negative pole, and 
 also very easily by pouring a solution of ammonium 
 chloride upon sodium amalgam. In the latter case 
 sodium chloride and ammonium amalgam are formed. 
 Apparently the reaction takes place in accordance with 
 the following equation : 
 
 NH 4 C1 + NaHg = NaCl + NH 4 Hg. 
 
 The product is extremely voluminous, and swells up 
 during the reaction, so that it occupies under favorable 
 
626 INORGANIC CHEMISTRY. 
 
 circumstances about twenty times the volume occupied 
 by the sodium amalgam. It has a metallic lustre, resem- 
 bling in general the other amalgams. It is very unstable 
 at the ordinary temperature, breaking down into mercury, 
 hydrogen, and ammonia. At a low temperature, how- 
 ever, it has been obtained in crystallized form. The 
 metallic lustre and general outward appearance of the 
 compound suggests that whatever is in combination with 
 mercury in it has probably metallic properties, and this 
 affords some confirmation of the ammonium theory, ac- 
 cording to which the presence of the complex, NH 4 , in 
 the salts formed by ammonia is assumed. Silver amal- 
 gam and gold amalgam vary in composition according to 
 the method of preparation, and when heated are com- 
 paratively easily decomposed. 
 
 Mercurous Chloride, HgCl, is commonly called calomel. 
 Like cuprous chloride, CuCl, and argentic chloride, AgCl, 
 it is insoluble in water. It is formed most readily by re- 
 ducing mercuric chloride. The reduction can be accom- 
 plished by means of sulphurous acid, when the following 
 reaction takes place : 
 
 2HgCl 2 + 2H 2 + S0 2 - 2HgCl + H 2 SO 4 + 2HC1. 
 
 It is also formed by heating together mercuric chloride 
 and mercury, and by subliming a mixture of mercuric 
 sulphate, sodium chloride, and mercury. This method 
 is the one mostly used in the manufacture of calomel. 
 The product obtained by sublimation is crystalline ; the 
 precipitated substance forms a loose powder. As was 
 stated above, the specific gravity of the vapor corre- 
 sponds to that required for the formula HgCl. When 
 acted upon for some time by light it undergoes partial 
 decomposition into mercury and mercuric chloride. This 
 is a fact of great importance, inasmuch as calomel is 
 much used in medicine, and mercuric chloride is an active 
 poison. Bottles in which calomel is kept should be care- 
 fully protected from the action of the light. 
 
 Just as mercuric chloride is converted into mercurous 
 chloride by reducing agents, so the latter is converted 
 into the former by oxidizing agents. When, for example, 
 
MERCURIC CHLORIDE. 627 
 
 mercurous chloride is treated with nitric acid it is con- 
 verted into mercuric chloride and mercuric nitrate, as 
 represented in the equation 
 
 6HgCl + 8HNO 3 = 3Hg(NO 3 ) 2 + 3HgCl 3 + 2NO + 4H 2 O. 
 
 If hydrochloric acid is present in sufficient quantity the 
 action takes place thus : 
 
 3HgCl + 3HC1 + HNO 3 = 3HgCl 2 + 2H 3 O + NO. 
 
 Further, the conversion of mercurous nitrate into mer- 
 curic nitrate is represented by the equation 
 
 3HgNO 3 + 4HN0 3 = 3Hg(NO 3 ) 2 + 2H 2 O + NO. 
 
 Finally, the action of oxidizing agents in general upon 
 mercurous chloride in the presence of hydrochloric acid 
 takes place thus : 
 
 2HgCl + 2HC1 + O = 2HgCl 2 + H 2 O. 
 
 Similar transformations take place by treating ferrous, 
 stannous, and manganous compounds with oxidizing 
 agents ; and they will be taken up farther on. 
 
 Mercuric Chloride, or Corrosive Sublimate, HgCl 2 , which 
 is made by subliming a mixture of sodium chloride and 
 mercuric sulphate, 
 
 HgS0 4 + 2NaCl = HgCl 2 + Na.SO 4 , 
 
 and by dissolving mercury in aqua regia, evaporating to 
 dryness, and subliming the residue, is a white, trans- 
 parent, crystalline mass, which is soluble in water, and 
 can be obtained in crystalline form from the solution. 
 It is more easily soluble in alcohol and ether than in 
 water, and is extracted from a water solution by shaking 
 with ether. It is quite volatile, and the specific gravity 
 of its vapor corresponds to that required for the formula 
 HgCl 2 . It is easily reduced to mercurous chloride by 
 
628 INORGANIC CHEMISTRY. 
 
 contact with organic substances, and by reducing agents 
 in general. The action of sulphur dioxide has already 
 been treated of as furnishing a method for the prepara- 
 tion of mercurous chloride. Stannous chloride abstracts 
 chlorine from it and forms mercurous chloride and me- 
 tallic mercury, while the stannous chloride is converted 
 into stannic chloride : 
 
 2HgCl 2 + SnCl 2 = 2HgCl + SnCl 4 ; 
 HgCl 2 + SnCl 2 = Hg + SnCl 4 . 
 
 Mercuric chloride is an active poison, and has been 
 used extensively in this capacity. It has a very marked 
 influence upon the lower organisms, which play such an 
 important part in producing disease and the decay of 
 organic substances, and is used as a disinfectant. Wood 
 impregnated with a solution of it is partly protected from 
 decay. In surgery it is used for the purpose of pre- 
 venting contamination of wounds by the hands and in- 
 struments of the surgeon, it being customary now for the 
 surgeon to wash his hands and instruments in a dilute 
 solution of the chloride before performing an operation. 
 
 Mercuric chloride unites with other chlorides, forming 
 well-characterized double chlorides, or chlor-mercurates, 
 which are analogous to the double chlorides of magne- 
 sium, zinc, etc. Three potassium salts are known, KHgCl 3 , 
 
 K 2 HgCl 4 , and KHg 2 Cl 5 ; or Hg<g^ Hg<g|, and 
 
 H <C1 
 
 C1 2 . Similar salts are formed with the chlorides 
 
 of sodium, ammonium, calcium, barium, strontium, and 
 other metals. Further, hydrochloric acid forms, with 
 mercuric chloride, a crystallized compound of the for- 
 mula HHg 2 Cl 6 or 2HgCl 2 .HCl, which is plainly the acid 
 from which the potassium salt KHg 2 Cl B is derived. 
 
 Mercurous Iodide, Hgl, can be made by treating mer- 
 cury with iodine ; or by treating mercuric iodide with 
 mercury ; and more easily by adding potassium iodide 
 to a solution of a mercurous salt, when it is thrown down 
 
MERCURIC IODIDE. 629 
 
 as' a greenish-yellow powder. It is unstable, and breaks 
 down, yielding mercury and mercuric iodide. When 
 treated with potassium iodide it suffers the same de- 
 composition, the iodide which is formed combining with 
 the potassium iodide, while mercury is deposited. Mer- 
 curous iodide is used in medicine. As it is more easily 
 decomposed than mercurous chloride, and mercuric 
 iodide is poisonous, care should be taken in its use. 
 
 Mercuric Iodide, HgI 2 , is made by direct combination 
 of mercury and iodine, and by the action of potassium 
 iodide upon a solution of a mercuric salt, when it is pre- 
 cipitated as a beautiful scarlet-red powder. Though 
 insoluble in water, it is soluble in alcohol and ether. It 
 dissolves, further, in a solution of potassium iodide in 
 consequence of the formation of a double iodide, or iodo- 
 mercurate. When heated to about 150 the color 
 changes from red to yellow, and after cooling the yel- 
 low substance changes to red. Sometimes it can be kept 
 in the condition in which it has a yellow color for some 
 time after it has cooled down, and then if touched at one 
 point the change to the red substance takes place rap- 
 idly through the entire mass. Both the red and the yel- 
 low modifications are crystallized, but in different forms. 
 The red crystals are tetragonal ; the yellow ones are 
 rhombic or monoclinic. The monoclinic structure is 
 evidently, from what has been said, an unstable one for 
 this compound. It seems probable that the difference 
 between the two forms is due to a difference in the com- 
 plexity of the molecules ; perhaps the larger molecules of 
 the red iodide are decomposed by heat into the smaller 
 molecules of the yellow iodide. When the iodide is first 
 precipitated from a solution of mercuric chloride by 
 potassium iodide it is yellow, but it rapidly turns red. 
 
 Just as mercuric chloride forms a compound with hy- 
 drochloric acid, so mercuric iodide forms similar com- 
 pounds. These have the composition represented by 
 the formulas HgI 2 .HI or HHgI 3 , and 2HgI 2 .3HI or 
 H 3 Hg 2 I 7 . It also combines with iodine, forming the per- 
 iodide HgI 9 , which is decomposed by water, forming mer- 
 curic iodide. With potassium chloride it forms the salts 
 
G30 INORGANIC CHEMISTRY. 
 
 K 2 HgI 4 or HgI 2 .2KI, and KHgI 3 or HgI a .KI. These salts 
 form colorless solutions. They are also formed by the 
 action of potassium iodide upon mercurous iodide, when 
 mercury separates : 
 
 2HgI + 2KI = K^Hgl, + Hg. 
 
 Similar double iodides are formed with the chlorides of 
 sodium, ammonium, barium, calcium, strontium, magne- 
 sium, and other metals. 
 
 Mercurous Oxide, Hg 2 O, is formed by treating a mer- 
 curous salt with potassium hydroxide. It is a black 
 powder, and decomposes easily into mercuric oxide and 
 mercury when the light shines upon it. 
 
 Mercuric Oxide, HgO, like the iodide, presents itself in 
 two colors the red and the yellow. When mercury is 
 heated for a long time at a temperature near the boiling- 
 point so that the air has access to it, it is converted into 
 the red oxide. When a solution of a mercuric salt is 
 treated with caustic soda or caustic potash the yellow 
 oxide is precipitated. When heated, the red oxide be- 
 comes darker, and finally nearly black, and on cooling 
 the red color reappears. The yellow oxide also becomes 
 darker on heating. Exposed to the light, the red oxide 
 loses some of its oxygen, and mercury is deposited. 
 When heated to a sufficiently high temperature both 
 varieties give up all their oxygen, as was seen in the 
 case of the red oxide in the preparation of oxygen. It 
 will be remembered that oxygen was discovered by heat- 
 ing this compound. Of what overwhelming importance 
 this discovery was the student will now better appreci- 
 ate than when the discovery was first mentioned. It will 
 now be seen that most of the chemical phenomena with 
 which we have to deal involve the action of oxygen. 
 
 Hydroxides of mercury are not known. 
 
 Mercurous Sulphide, Hg 2 S. There seems to be some 
 question whether a compound of the formula Hg 2 S 
 exists or not. According to the latest investigations 
 on the subject, the precipitate which is formed when 
 hydrogen sulphide is passed into a solution of a nier- 
 
MERCURIC SULPHIDE. 631. 
 
 curous salt is a mixture of mercuric sulphide, HgS, 
 and mercury. At all events, it is certain that, even 
 if mercurous sulphide exists, it breaks down with great 
 ease into the mercuric compound and mercury. 
 
 Mercuric Sulphide, HgS, is the principal compound of 
 mercury found in nature. This natural variety, which 
 is known as cinnabar, is a red crystallized compound. 
 When prepared by treating a solution of a mercuric salt 
 with hydrogen sulphide it is an amorphous black pow- 
 der. This black powder can, however, be converted into 
 the red crystallized variety by sublimation, and by allow- 
 ing it to stand in contact with a solution of the sulphide 
 of an alkali metal. The sulphide, whether amorphous 
 or crystallized, is acted upon with difficulty by acids. 
 Dilute nitric acid, which easily dissolves the sulphides 
 of the other metals which resemble mercury, leaves it 
 unacted upon, and advantage is taken of this -fact for the 
 purpose of separating it from the sulphides of lead, bis- 
 muth, copper, and cadmium in qualitative analysis. It is 
 dissolved by concentrated nitric acid and by aqua regia. 
 In the former case a white insoluble compound of mer- 
 curic nitrate and mercuric sulphide, Hg(NO 3 ) 2 .2HgS, 
 is sometimes formed. Potassium hydrosulphide dis- 
 solves mercuric sulphide, forming a compound, K 2 HgS 2 , 
 or K 2 S.HgS, which is easily decomposed by water, mer- 
 curic sulphide being thrown down. The substance used 
 in medicine under the name EtMops martialis is an inti- 
 mate mixture of amorphous sulphide and sulphur. Like 
 the oxide, cinnabar turns dark when heated, and it slowly 
 undergoes the same change when exposed to the light, in 
 consequence of a slight decomposition into mercury and 
 sulphur. When that which has not been heated to too 
 high a temperature is cooled down again, it acquires its 
 original red color. If it has been heated to the tempera- 
 ture of sublimation, the black color of that which does 
 not sublime is permanent. Cinnabar is used as a pig- 
 ment in the manufacture of red paints. 
 
 Mercuric Cyanide, Hg(CN) 2 , is formed by dissolving 
 mercuric oxide in an aqueous solution of hydrocyanic 
 
632 INORGANIC CHEMISTRY. 
 
 acid. It is soluble in water, and crystallizes well in 
 quadratic prisms. When heated it is decomposed into 
 cyanogen and mercury, and this affords the most con- 
 venient method of making cyanogen. When heated rap- 
 idly there is formed a considerable quantity of a black 
 substance called paracyanogen, which has the same per- 
 centage composition as cyanogen, but is not volatile. This 
 is probably a polymeric variety of cyanogen. The 
 cyanide unites with the cyanides of other metals, form- 
 ing double cyanides, of which those corresponding to 
 the following formulas are good examples : K 2 Hg(CN) 4 , 
 MgHg(CN) 4 , etc. These are soluble in water, and, as a 
 rule, crystallize well. 
 
 Fulminating mercury is an explosive compound much 
 used in the manufacture of gun-caps. It is made by dis- 
 solving mercury in nitric acid and adding alcohol. Its 
 explosion consists in a sudden breaking down into nitro- 
 gen, carbon dioxide, and mercury. The composition of 
 the compound is represented by the formula C 2 N 2 O 2 Hg. 
 It would lead too far to discuss its constitution a sub- 
 ject which has occupied the attention of some of the 
 most celebrated chemists. 
 
 Mercurous Nitrate, HgNO 3 , is made by treating nitric 
 acid with an excess of mercury. This salt is easily de- 
 composed, forming difficultly soluble basic salts, some of 
 which are of complicated composition. The simplest is 
 that which has the composition HgOH.HgNO 3 . This 
 is easily explained on the assumption that the nitrate 
 
 Hg(NO 3 ) 
 has the formula I The decomposition by water 
 
 is then represented by the equation 
 
 a + H 2 = 
 Hg(N0 3 ) Hg(OH) 
 
 This may possibly be regarded as a slight argument in 
 favor of the doubled formula for the mercurous com- 
 pound. It is, however, far from conclusive, as the salt 
 may be considered to be made up as represented in this 
 
COMPOUNDS OF SALTS OF MERCURY AND AMMONIA. 633 
 
 (OH 
 
 formula, ON < OHg , according to which it is a derivative 
 (OHg 
 
 of the acid NO(OH) 3 , corresponding to phosphoric acid, 
 PO(OH) 3 . 
 
 Mercuric Nitrate, Hg(NO 3 ) 2 , is formed by treating mer- 
 curic oxide with an excess of nitric acid and evaporating, 
 when the salt can, under favorable circumstances, be 
 obtained in crystals. It is easily decomposed by water, 
 with formation of a basic salt which is insoluble in 
 water. It has the composition represented by the for- 
 mula NO 3 -Hg-O-Hg-O-Hg-NO 3 + H 2 O. 
 
 Compounds formed by Salts of Mercury with Ammonia. 
 Attention has already been called to the compounds 
 formed by the salts of copper with ammonia, and the 
 statement was made that this power to combine with am- 
 monia and with ammonium salts is very common among 
 the salts. Some of these compounds which are formed 
 by the salts of mercury are of special interest. When 
 mercuric chloride is treated with ammonia a white pre- 
 cipitate is formed, the composition of which is repre- 
 sented by the formula HgClNH 2 . The formation takes 
 place according to the equation 
 
 HgCl 2 + 2NH 3 = HgClNH 2 + NH 4 C1. 
 
 The simplest view in regard to the constitution of this 
 compound is that it is mercuric chloride in which a 
 chlorine atom is replaced by the group NH 2 , or amide, 
 
 Cl 
 as represented in the formula Hg<^-rr . According to 
 
 this view the compound is called mercuric chloramide. 
 It is known as white precipitate, or, to distinguish it from 
 another similar compound, as infusible white precipitate. 
 It is also possible that the constitution of the compound 
 
 (H 2 
 should be represented by the formula N< Hg, according 
 
 (a 
 
 to which it is ammonium chloride in which two atoms of 
 hydrogen are replaced by an atom of bivalent mercury. 
 The similar compound referred to is formed by adding a 
 solution of mercuric chloride to a boiling solution of am- 
 
634 INORGANIC CHEMISTRY. 
 
 monium chloride containing ammonia, and separates out 
 on cooling. It has the composition Hg(NH 3 Cl) 2 , and is 
 to be regarded as ammonium chloride, in two molecules 
 of which two hydrogen atoms are replaced by an atom 
 of bivalent mercury, as represented in the formula 
 
 This is called mercuric diammonium chlo- 
 
 ride, or fusible white precipitate. Mercurous compounds 
 corresponding to both the above-mentioned derivatives 
 of mercuric chloride are known. The first, or mercurous 
 chloramide, Hg 2 NH 2 Cl, is the black substance formed 
 when mercurous chloride is treated with ammonia : 
 
 2HgCl + 2NH 3 = Hg 2 NH 2 Cl + NH 4 C1. 
 
 The second, or mercurous ammonium chloride, HgNH 3 Cl, 
 is formed by treating calomel with ammonia gas. 
 
 Similar compounds are obtained from the other salts of 
 mercury. In a solution of mercurous nitrate ammonia 
 forms a black precipitate, known in pharmacy as Mercu- 
 rius solubilis Hahnemanni. It may be regarded as derived 
 from ammonium nitrate by replacement of two hydrogen 
 atoms by two atoms of mercury, as represented in the 
 formula Hg 2 H 2 N.NO 3 . 
 
 Reactions which are of Special Value in Chemical 
 Analysis. As has been seen in the account already given 
 of the conduct of the compounds of mercury, mercurous 
 and mercuric compounds conduct themselves quite dif- 
 ferently. Among the most characteristic reactions of 
 mercurous compounds are the following : 
 
 Sodium or potassium hydroxide gives a black precipi- 
 tate of mercurous oxide. 
 
 Ammonia gives a black precipitate which is a compound 
 of mercurous oxide and ammonia. 
 
 Hydrochloric acid and soluble chlorides form mercurous. 
 chloride, and the precipitate turns black when treated 
 with ammonia. 
 
 The reactions with stannous chloride, potassium iodide, 
 hydrogen sulphide, and ammonium sulphide have been 
 explained. 
 
 The principal reactions of the mercuric salts are the 
 
GALLIUM. 635 
 
 following : Sodium or potassium hydroxide produces a 
 yellow precipitate of mercuric oxide. 
 
 The reactions with hydrogen sulphide, stannous chlo- 
 ride, and potassium iodide have been described above. 
 
 ELEMENTS OF FAMILY III, GROUP B: 
 GALLIUM INDIUM THALLIUM. 
 
 General. All the elements of this group are rare. 
 Gallium forms compounds which in composition are 
 analogous to those of aluminium, in which it is trivalent, 
 as in the compounds Ga01 3 , Ga 2 O 3 , Ga(NO 3 ) 3 , Ga 2 (SO 4 ) 3 , 
 etc. On the other hand, it also forms compounds in which 
 it is bivalent, as in the chloride GaCl 2 and GaO. Indium 
 also forms compounds in which it is trivalent, and a few 
 in which it is bivalent. Thallium, like gold, is trivalent 
 and univalent in its compounds. 
 
 Gallium, Ga (At. Wt. 69.38). This element is found in 
 some varieties of zinc blende, and was discovered in that 
 which occurs at Pierrefitte, in France. It owes its name 
 to the Latin name of France, Gallia. Like scandium, it 
 is of special interest for the reason that its properties 
 were foretold by Mendeleeff some years before it was dis- 
 covered, and it was described by him under the name of 
 eka-aluminium, as scandium was described under the 
 name of eka-boron. The metal has a bluish- white color, 
 and does not lose its lustre in the air. It does not de- 
 compose water. Even when heated to a high tempera- 
 ture in oxygen it only becomes covered with a thin layer 
 of oxide. It dissolves in hydrochloric acid and in po- 
 tassium hydroxide with evolution of hydrogen. 
 
 Compounds of Gallium. The chlorides of gallium, 
 GaCl 2 and GaCl 8 , are formed by treating the metal with 
 chlorine. With water both give basic chlorides. That 
 formed from gallons chloride, GaCl 2 , is completely con- 
 verted into gallium oxide by further action of water. 
 Gallic sulphate, Ga 2 (SO 4 ) 3 , is decomposed by boiling water, 
 a basic salt being formed. With ammonium sulphate it 
 forms a double sulphate of the formula NH 4 Ga(SO 4 ) 2 -f- 
 12H 2 O, which is perfectly analogous to ammonium alum 
 (see Alums, p. 576). 
 
636 INORGANIC CHEMISTRY. 
 
 Indium, In (At. Wt. 112.99). Indium was discovered in 
 a variety of zinc blende found at Freiberg, in Germany. 
 The discovery was made by means of the spectroscope, 
 and as the spectrum of the element contains two very 
 characteristic blue lines it was called indium. It has 
 since been found in other varieties of zinc blende, but 
 always in small quantity. It is a soft, white, lustrous 
 metal. It does not undergo change in contact with the 
 air at ordinary temperatures, but when heated it takes 
 fire and burns, forming the oxide. It does not decom- 
 pose water even at the boiling temperature. 
 
 Compounds of Indium. With chlorine indium forms 
 the compound InCl 3 , which is volatile at a high tempera- 
 ture. The specific gravity of the vapor is that required 
 for the formula InCl 3 . The sulphate is easily formed 
 by dissolving the metal in sulphuric acid. It combines 
 with ammonium sulphate to form the double sulphate 
 ]S T H 4 In(SO 4 ) 2 -f 12H 2 O, analogous to alum. The corre- 
 sponding potassium salt does not contain the same num- 
 ber of molecules of water of crystallization. 
 
 Thallium, Tl (At, Wt. 202.61). This element was dis- 
 covered in the flue-dust of a sulphuric-acid factory in the 
 Harz mountains, by the aid of the spectroscope. As it 
 colors the flame a beautiful green it was called thallium, 
 from the Greek $orAAo, which signifies a green branch. 
 It has since been found in a number of varieties of iron 
 pyrites and copper pyrites. It is a soft, bluish-white 
 metal, and has a lustre like lead. It is oxidized when 
 heated to a sufficiently high temperature in the air. 
 
 Compounds of Thallium. When thallium is treated 
 with chlorine it is converted into thallous chloride, T1C1 ; 
 and when this is treated under water with chlorine it is 
 converted into the trichloride or thallic chloride, T1C1,. 
 Thallous chloride, further, is formed as a caseous pre- 
 cipitate when hydrochloric acid is added to a solution of 
 a thallous salt. When exposed to the light it turns violet. 
 Thallous hydroxide, Tl(OH), is formed by the action of 
 the metal on water in the air, and by treating a solution 
 of the sulphate with barium hydroxide. It is easily 
 soluble in water, and the solution has an alkaline reac- 
 
COMPOUNDS OF THALLIUM. 037 
 
 tion. Thallic hydroxide, T1(OH) 3 , is formed by treating 
 a solution of thallic chloride with potassium hydroxide. 
 When dried it loses water and forms the compound 
 T1O.OH, analogous to metaboric acid, BO. OH, and meta- 
 aluminic acid, A1O.OH. Thallium forms an insoluble 
 chloride, T1C1, which turns violet when exposed to the 
 light, and in this respect it resembles silver ; the forma- 
 tion of the soluble hydroxide, Tl(OH), which has an 
 alkaline reaction, is highly suggestive of the alkali 
 metals ; while the formation of the hydroxide, TIO(OH), 
 shows that thallium is allied to aluminium. An exami- 
 nation of its salts shows that those in which it is univa- 
 lent resemble the salts of the alkali metals, while those 
 in which it is trivalent resemble the salts of aluminium. 
 -Thallous sulphate, T1 2 SO 4 , and the thallous phosphates, 
 H 2 T1PO 4 , HT1 2 PO 4 , and T1 3 PO 4 , are isomorphous with 
 the corresponding salts of potassium. Thallous sulphide 
 is thrown down as a black powder when hydrogen sul- 
 phide is passed into a solution of a thallous salt. Thai' 
 lie sulphate, T1 2 (SO 4 ) 3 , combines with sulphates of the 
 alkali metals, forming salts of the general formula 
 T1M(SO 4 ) 2 ; but these do not crystallize like the alums. 
 On the other hand, thallous sulphate, T1 2 SO 4 , combines 
 with aluminium sulphate, and other similar sulphates, 
 forming salts perfectly analogous to the alums, and in 
 these the thallium takes the place of the alkali metal, as, 
 for example, in the salts 
 
 TlAl(SO 4 ) f + 12H a O, TlFe(SO 4 ), + 12H 2 O, etc. 
 
CHAPTER XXX. 
 
 ELEMENTS OF FAMILY IV, GROUP B : 
 GERMANIUM TIN LEAD. 
 
 General. Of the elements of this group germanium is 
 extremely rare. It was discovered by Winkler in 1885. 
 Tin and lead, on the other hand, have long been known, 
 and are extensively used. All form two series of com- 
 pounds, in one of which they are bivalent and in the 
 other quadrivalent. The general formulas of some of 
 the principal compounds of the first series are as fol- 
 lows: 
 
 MC1 2 , MO, M(OH) 2 , M(NO 3 ) 2 , MSO 4 , MCO 3 , etc. 
 
 The general formulas of some of the principal compounds 
 of the other series are as follows : 
 
 MC1 4 , M0 2 , M(OH) 4 , MO(OH) 2 , etc. 
 
 These elements have already been referred to at the 
 close of Chapter XXII. (see p. 425), and attention was 
 then called to the resemblance between them and carbon 
 and silicon. In this connection it will be well to repeat 
 what was there said. Of the three elements of the 
 group, germanium and tin are more acidic in character 
 than lead. They combine with chlorine in two propor- 
 tions, forming the chlorides GeCl 3 , SnCl a , PbCl a , GeCl 4 , 
 SnCl 4 , PbCl 4 . With oxygen they combine to form the 
 compounds GeO a , SnO 2 , and PbO 2 . Stannic oxide, SnO 2 , 
 and lead peroxide, PbO 2 , form salts with bases, and these 
 have the composition represented by the general formulas 
 M 2 SnO 3 and M 2 PbO 3 , and are therefore analogous to the 
 silicates, carbonates, and titanates. On the other hand, 
 further, salts are known which are derived from the 
 oxide PbO. These have the general formula M 2 PbO a> . 
 
 (638) 
 
 
GERMANIUM TIN. 639 
 
 and are to be regarded as salts of an acid Pb(OH) 2 . 
 These salts are not stable, and are not easily obtained. 
 Most of the derivatives of lead are those in which it 
 plays the part of a base-forming element. Notwith- 
 standing the marked analogy between some of the com- 
 pounds of tin and those of the members of the silicon 
 group, it appears, on the whole, advisable to treat of this 
 element in company with lead, which it also resembles in 
 many respects. 
 
 GERMANIUM, Ge (At. Wt. 71.93). 
 
 Germanium is the third element the properties of 
 which were foretold by Mendeleeff by the aid of the 
 periodic law. As it occurs in the silicon group he called 
 it eka-silicon. It was discovered in a silver ore occur- 
 ring at Freiberg, Germany. The name has, of course, 
 reference to the country in which it was discovered. It 
 acts mostly as a base-forming element, being perhaps 
 more like tin than any other one metal. It forms the 
 two chlorides GeCl 2 and GeCl 4 , and the corresponding 
 fluorides GeF 2 and GeF 4 ; but preferably it forms those 
 compounds in which the element is quadrivalent. The 
 fluoride forms double salts resembling the fluosilicates, 
 as, for example, pot assium fluogermanat e, K 2 GeF 6 . 
 
 TIN, Sn(At. Wt. 118.15). 
 
 General. The compounds of tin with which we gen- 
 erally have to deal belong, with the exception of stan- 
 nous chloride, to the series in which the metal is quad- 
 rivalent, and in this series it acts as an acid-forming 
 element. The chloride, SnCl 4 , corresponds to the chlo- 
 rides of carbon and silicon, CC1 4 and SiCl 4 . Unlike 
 these elements, however, it does not form a compound 
 with hydrogen. 
 
 Occurrence. Tin occurs almost exclusively as tin 
 stone or cassiterite in nature. This is the dioxide, SnO 2 , 
 corresponding to carbon dioxide, CO 2 ; silicon dioxide, 
 SiO 2 ; titanium dioxide, TiO 2 ; etc. It also occurs in 
 small quantities in company with gold as metallic tin, 
 
640 INORGANIC CHEMISTRY. 
 
 and in a variety of pyrites of the formula Cu 4 SnS 4 -)- 
 FeSnS 4 , known as stannite. 
 
 Metallurgy. The ores are roasted for the purpose of 
 getting rid of the sulphur and arsenic, and the oxide is 
 then heated with coal in a furnace. After the reduction 
 is complete the tin is drawn off and cast in bars. This 
 tin is impure, and when again slowly melted, that which 
 first melts is purer. By letting it run off as soon as it 
 melts the comparatively difficultly fusible alloy remains 
 behind, and the tin is thus rendered much purer. The 
 commercial variety of tin known as Bonca tin is the 
 purest. It receives its name from Banca, in the East 
 Indies, where it is made. Block-tin is made in England, 
 and is also comparatively pure. 
 
 Properties. Tin is a white metal, which in general ap- 
 pearance resembles silver. It is soft and malleable, and 
 can be hammered out into very thin sheets, forming the 
 well-known tin-foil. Its specific gravity is 7.3. At 200 
 it is brittle, and at 228 it melts. At ordinary tempera- 
 tures it remains unchanged in the air. It dissolves in 
 hydrochloric acid, forming stannous chloride, SnCl 2 ; in 
 sulphuric acid, forming stannous sulphate, SnSO 4 , sul- 
 phur dioxide being evolved at the same time. Ordinary 
 concentrated nitric acid oxidizes it, the product being a 
 compound of tin, oxygen, and hydrogen, known as meta- 
 stannic acid, which is a white, powder insoluble in ni- 
 tric acid and in water. It is dissolved by a hot solution 
 of potassium hydroxide which forms potassium stannate, 
 K 2 Sn0 3 . 
 
 Applications. It is used in making alloys, of which 
 bronze (see p. 591), soft solder, and britannia metal are 
 the most important. It is used also for protecting other 
 metals, as in the tinware vessels in such common use, 
 which are made of iron covered with a layer of tin. 
 Copper vessels are also frequently covered with tin. 
 
 Alloys. Bronze has already been treated of under 
 Copper. It is made of copper, tin, and zinc. Soft solder 
 is made of equal parts of tin and lead, or of two parts of 
 tin and one of lead. Britannia metal is composed of nine 
 parts of tin and one of antimony. Tin amalgam is made 
 
STANNOUS CHLORIDE. 641 
 
 by bringing tin and mercury together, and is used in the 
 silvering of mirrors. 
 
 Stannous Chloride, SnCl 2 , is formed by dissolving tin 
 in hydrochloric acid, and if the solution is concentrated 
 enough the compound crystallizes oat. The crystals have 
 the composition SnCl 2 -j- 2H 2 O. This is the commercial 
 product known as tin salt. It is very easily soluble in 
 water, but if the solution is dilute it becomes turbid in 
 consequence of the formation of the insoluble basic salt, 
 
 2SnCl a + 3H 2 O = 2Sn<j + H 2 O + 2HCL 
 
 This same decomposition takes place if the solution is 
 allowed to stand in contact with the air. Under these 
 circumstances a part of the stannous chloride is converted 
 into stannic chloride by oxidation : 
 
 3SnCl 2 + H 2 O + O = SnCl 4 + 2Sn(OH)Cl. 
 
 "When the crystals are heated they melt at about 40, and 
 if heated higher they undergo partial decomposition, 
 forming the oxide and hydrochloric acid : 
 
 SnCl 2 + H 2 O = SnO + 2HC1. 
 
 Stannous chloride has a marked tendency to combine 
 with chlorine and to pass over into stannic chloride. 
 This power has already been shown in its action upon 
 mercuric chloride and upon mercurous chloride. It re- 
 duces the former, first to mercurous chloride, and it then 
 abstracts the chlorine from this, leaving metallic mercury : 
 
 2HgCl 2 + SnCl 2 = 2HgCl + SnCl 4 ; 
 2HgCl + SnCl 2 = 2Hg + SnCl 4 . 
 
 Stannous chloride is an excellent mordant, and is ex- 
 tensively used by the dyers. It unites with other chlo- 
 rides, forming double chlorides of the general formula 
 SnCl 2 .2MCl or M 2 SnCl 4 , which are analogous to the salts 
 of an unknown stannous acid, H 2 SnO 2 . 
 
642 INORGANIC CHEMISTRY. 
 
 Stannic Chloride, SnCl 4 , formed by treating tin with 
 chlorine, is a colorless liquid which boils at 114, and 
 the specific gravity of its vapor is that required by a 
 compound of the formula SnCl 4 . In the air it gives off 
 fumes in consequence of the action of moisture. It has 
 long been known by the name spiritus fumans Libavii, 
 which has reference to its fuming quality and to the fact 
 that it was discovered by Libavius. It combines with- 
 water, forming a number of crystallized hydrates. When 
 its solution in water is boiled stannic acid is precipitated : 
 
 SnCl 4 + 3H 2 O = H 2 SnO 3 + 4HC1. 
 
 Probably the normal acid, Sn(OH) 4 , is first formed, and 
 this then breaks down with loss of water to form the 
 ordinary acid : 
 
 SnCl 4 + 4H 2 O = Sn(OH) 4 + 4HC1 ; 
 Sn(OH) 4 == SnO(OH) 2 + H 2 O. 
 
 Stannic chloride is used as a mordant. With other 
 chlorides it forms the chlorostannates of the general for- 
 mula M 2 SnCl 6 or SnCl 4 .2MCl, of which the ammonium 
 salt, (NH 4 ) 2 SnCl 6 , or pink salt, is the best known. This 
 is manufactured for use in making cotton -prints. Similar 
 fluorine compounds, HIQ fluostannates, are also easily ob- 
 tained, as K 2 SnF 6 , Na 2 SnF 6 , etc. 
 
 Stannous Hydroxide, Sn(OH) 2 , is not known. When a 
 solution of stannous chloride is treated with potassium 
 carbonate a precipitate of the composition H 2 Sn 2 O 3 is 
 formed, which is a derivative of the hydroxide Sn(OH) 2 , 
 as is shown by the formula HO-Sn-O-Sn-OH, which 
 probably expresses its constitution. It readily loses 
 water and passes over into stannous oxide, SnO, which is 
 a black powder. 
 
 Stannic Hydroxide, Sn(OH) 4 , is perhaps formed when a 
 solution of stannic chloride in water is boiled. The pre- 
 cipitate obtained has, however, the composition H 2 SnO 3 , 
 and this is known as stannic acid. Stannic acid is pre- 
 cipitated also by treating a solution of a stannate with 
 
METASTANNIC ACID-STANNOUS OXIDE. 643 
 
 just enough of an acid to effect decomposition. The de- 
 composition with hydrochloric acid takes place as repre- 
 sented in the equation 
 
 Na 2 SnO 3 + 2HC1 = 2NaCl + H 2 SnO 3 . 
 
 The compound thus obtained is insoluble in water, but 
 is easily soluble in hydrochloric, nitric, and sulphuric 
 acids, and in the caustic alkalies. With the alkalies it 
 forms stannates, as sodium stannate, Na 2 SnO 3 , and potas- 
 sium stannate, K 2 SnO 3 . The former is made on the large 
 scale, and is known as preparing salt. 
 
 Metastannic Acid. When tin is treated with concen- 
 trated nitric acid it is converted into a white powder which 
 is insoluble in water and in acids, but nevertheless seems 
 to be a hydroxide of tin of the same composition as stannic 
 acid. This is known as metastannic acid. With alkalies it 
 forms salts which in properties and composition are en- 
 tirely different from the stannates. They are known as 
 the metastannatcs. Two sodium salts are known, which 
 differ in composition. One of these has the composition 
 Na 2 Sn 6 O n , the other is Na 2 Sn 9 O 19 . They are probably 
 derived from acids analogous to the polysilicic acids, 
 which may be called poly stannic acids. The question as 
 to the composition of metastannic acid is an open one. 
 When heated it is converted into the oxide SnO 2 . When 
 treated with concentrated hydrochloric acid it is con- 
 verted into a compound containing chlorine which is 
 soluble in water, and the solution contains stannic chlo- 
 ride. When this solution is treated with sulphuric acid 
 stannic sulphate is thrown down, and when the solution 
 is boiled this salt is decomposed, leaving stannic acid. 
 
 Stannous Oxide, SnO, is formed from the correspond- 
 ing hydroxide when this is heated in a current of carbon 
 dioxide. When a solution of stannous chloride is treated 
 with caustic soda stannous hydroxide is first precipitated, 
 and this dissolves in an excess of the caustic soda. When 
 the solution is boiled the salt contained in it is decom- 
 posed, and black stannous oxide is thrown down in crys- 
 talline form. When heated in the air stannous oxide 
 takes fire, and is converted into stannic oxide. 
 
644 INORGANIC CHEMISTEY. 
 
 Stannic Oxide, SnO 2 , as has been stated, is the princi- 
 pal form in which tin occurs in nature. The mineral is 
 known as cassiterite or tin-stone. It is found at Corn- 
 wall, England ; on the East Indian islands Banca and 
 Biliton ; and in Malacca. It crystallizes in the tetragonal 
 system, and is generally colored from brown to black. 
 It is formed as a white powder by burning tin in the air, 
 and by heating the different varieties of stannic hydrox- 
 ide. It is infusible, and is not acted upon by concen- 
 trated hydrochloric or nitric acid. Concentrated sul- 
 phuric acid, however, converts it into a gelatinous liquid 
 from which stannic oxide is precipitated by water. 
 
 Stannous Sulphide, SnS, which is formed when hydro- 
 gen sulphide is passed into a solution of stannous chlo- 
 ride, is a brownish-black powder. When treated with 
 the soluble sulphides and sulphur, or with a soluble 
 polysulphide, it dissolves, forming a sulphostannate as 
 represented in these equations : 
 
 SnS + (NH 4 ) 2 S + S - (NH 4 ) 2 SnS 3 ; 
 SnS + (NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 3 . 
 
 Stannic Sulphide, SnS 2 . This compound is obtained 
 in crystallized form by heating together tin-filings, sul- 
 phur, and dry ammonium chloride ; and in amorphous 
 form by treating a solution of stannic chloride with hy- 
 drogen sulphide. In the former case it is a golden- 
 yellow crystalline substance ; in the latter a yellow 
 powder. The crystalline variety is known as mosaic gold. 
 "When heated to a high temperature it is converted into 
 stannous sulphide. The precipitated variety is easily 
 dissolved by concentrated hydrochloric acid, and con- 
 verted into metastannic acid by concentrated nitric acid. 
 The crystallized variety is not soluble in hydrochloric 
 acid, and is affected but slightly by nitric acid. The 
 crystallized variety, or mosaic gold, is used as a pigment, 
 particularly for bronzing. Stannic sulphide dissolves 
 easily in the soluble sulphides, forming sulphostan- 
 nates : 
 
 SnS, + (NH,),S = (NH,),SnS, ; 
 
 SnS, + K,S = K.SnS,. 
 
SULPHOSTANNATES, ETC. REACTIONS. 645 
 
 The sulphostannates are perfectly analogous in com- 
 position to the stannates, differing from them simply by 
 containing three atoms of sulphur in place of the three 
 atoms of oxygen. Although the comparison has been 
 repeatedly made before in this book, it is perhaps not 
 superfluous again to call attention to the analogy between 
 the stannates, sulphostannates, chlorostannates, and 
 fluostannates, all of which are easily obtained, and are 
 well-characterized compounds. The formulas of some 
 representatives of the four classes are here placed side 
 by side : 
 
 K 2 Sn0 3 K 2 SnS 3 K 2 SnCl 6 K 2 SnF 6 
 
 Na 2 SnO 3 Na 2 SnS 3 Na 2 SnCl 6 Na 2 SnF 6 
 
 When a solution of a sulphostannate is treated with 
 an acid a yellow precipitate is formed. This becomes 
 darker, and if filtered off and dried it forms a gray mass. 
 It is thought that this may be sulphostannic acid, 
 H 2 SnS 3 . It breaks down readily into hydrogen sulphide 
 and stannic sulphide. 
 
 Stannous and Stannic Salts. With acids, tin forms a 
 few salts, but they are not particularly well characterized, 
 and the stannic salts are easily decomposed by water. 
 Stannous sulphate, SnSO 4 , is formed by dissolving tin in 
 warm concentrated sulphuric acid. Stannic sulphate, 
 Sn(SO 4 ) 2 , is precipitated when sulphuric acid is added 
 to a dilute solution of stannic chloride. It is decomposed 
 by hot water, forming stannic acid. 
 
 Reactions which are of Special Value in Chemical Analy- 
 sis. The reactions with hydrogen sulphide, and the con- 
 duct of the precipitated sulphides when treated with 
 ammonium sulphide or polysulphide, are constantly 
 utilized for analytical purposes when tin is present. 
 
 The reducing action of stannous compounds serves to 
 distinguish them from stannic compounds. 
 
 The conduct of solutions of stannous and stannic com- 
 pounds towards caustic alkalies has been explained above. 
 
 The carbonates of the alkali metals precipitate the hy- 
 droxides, and these do not dissolve in an excess of the 
 carbonate. Metallic zinc precipitates the tin from a so- 
 
G46 INORGANIC CHEMISTRY. 
 
 lution of a tin compound as a spongy mass. If the pre- 
 cipitation is carried on in a platinum vessel the plati- 
 num is not colored by it. By this means tin can be dis- 
 tinguished from antimony, which is reduced under the 
 same circumstances, but is deposited as a black coating 
 upon the platinum. 
 
 LEAD, Pb (At. Wt. 205.36). 
 
 General. The basic properties of lead are stronger, 
 and its acidic properties weaker, than those of tin. Its 
 principal compounds are those in which it acts as a base- 
 forming element. The compounds in which it is quad- 
 rivalent, such as PbO 2 , are comparatively unstable, and 
 when treated with acids they readily pass over into the 
 compounds of the series in which the lead is bivalent. 
 Thus lead oxide itself readily gives up oxygen when 
 treated with acids, and yields salts of bivalent lead. 
 
 Forms in which Lead occurs in Nature. Lead occurs in 
 nature as the sulphide, PbS, which is known as galenite 
 or galena. Other natural compounds of the metal are 
 the carbonate, PbCO 3 , known as cerussite ; the phos- 
 phate, Pb 3 (PO 4 ) 2 ; the chromate, PbCrO 4 , or crocoisite ; 
 and the molybdate, PbMoO 4 , or wulfenite. 
 
 Metallurgy. Most of the lead in the market is ob- 
 tained from the sulphide, and as most of the sulphide 
 contains silver both metals have to be considered in the 
 treatment of the ore. Under the head of Silver (which 
 see) reference was made to the methods ' by which this 
 metal was separated from the lead after both have been 
 separated from their compounds. Here it will only be 
 necessary to show how the metals are extracted together 
 from the ore. This is accomplished in one of two ways : 
 
 (1) By heating the sulphide with iron, when the latter 
 combines with the sulphur, forming iron sulphide, while 
 the lead is set free. 
 
 (2) By roasting the sulphide until it is partly con- 
 verted into lead oxide and lead sulphate, and then heat- 
 ing the mixture without access of air, when two reactions 
 take place, which are represented in these equations : 
 
 PbS + 2PbO = 3Pb + SO 2 ; 
 PbS + PbS0 4 = 2Pb + 2SO a . 
 
PROPERTIES OF LEAD. 647 
 
 The lead is thus set free, and the sulphur is driven off as 
 sulphur dioxide. 
 
 Properties. Lead is a bluish-gray metal, with a high 
 lustre. It is very soft-, and not very strong ; melts at 
 325, and has the specific gravity 11.37. At a high tem- 
 perature it is converted into vapor. When heated in 
 contact with the air it becomes covered with a layer of 
 oxide, as can be seen in a vessel containing the molten 
 metal. In the air, at ordinary temperatures, it is tar- 
 nished in consequence of the formation of a suboxide of 
 the composition Pb 2 O. The formation of this compound 
 can be observed very readily by cutting a piece of lead 
 with a knife. At first the freshly- cut surface has a high 
 lustre, but this soon grows dim and acquires a bluish 
 color. Pure water acts upon lead when air has access 
 to it, and some of the lead dissolves. If the water con- 
 tains salts in solution, such as calcium carbonate, gyp- 
 sum, etc., or if it contains carbon dioxide, it acts only 
 very slightly upon the metal ; further, water which con- 
 tains organic matter in a state of decomposition dis- 
 solves lead with comparative ease. Nitric acid dissolves 
 Jead readily; as, however, lead nitrate is insoluble in 
 nitric acid, it is necessary to use comparatively dilute 
 acid. Concentrated sulphuric acid dissolves lead to 
 some extent, and therefore a little lead sulphate is some- 
 times contained in commercial sulphuric acid. When 
 such a solution is diluted with water the sulphate is pre- 
 cipitated. Hence, whenever the commercial acid is di- 
 luted with water it becomes turbid, and after standing 
 for a time it becomes clear, as the lead sulphate settles. 
 Hydrochloric acid acts only slightly upon lead. Acetic 
 acid dissolves the metal very readily. It is precipitated 
 in metallic form from a solution of one of its salts by 
 metallic zinc. The formation is sometimes very beauti- 
 ful, especially if the zinc is suspended in the solution. 
 It is called the " lead tree," or Arbor Saturni. The 
 action consists in a replacement of the lead by the zinc. 
 After the action is complete all the lead is deposited as 
 metallic lead, and the zinc has entered into its place, 
 
648 INORGANIC CHEMISTRY. 
 
 forming a salt which remains in solution. Thus, if lead 
 nitrate is used, zinc nitrate is in the solution. 
 
 Applications. Lead is extensively used for a variety 
 of purposes, as, for example, for making sulphuric-acid 
 chambers, for evaporating-pans for alum and sulphuric 
 acid, for shot, for water-pipes, and in making alloys. 
 The use of lead water-pipes is a matter of much im- 
 portance from the sanitary point of view, as is evident 
 from the statements above made concerning the action of 
 water upon the metal. Ordinary drinking-water acts under 
 most circumstances only very slightly upon lead, and not 
 enough is dissolved to be dangerous to those using the 
 water. At the same time circumstances may at any 
 time arise which will increase the solvent power of the 
 water, and thus cause serious results ; and it would un- 
 doubtedly be better if the use of such pipes could be 
 entirely avoided in cases in which the water is to be 
 used for drinking purposes. 
 
 Lead Chloride, PbCl 2 , is formed when hydrochloric 
 acid or a soluble chloride is added to a cold solution of 
 a lead salt, and appears as a white precipitate. It is 
 soluble in hot water, and is deposited in the form of long, 
 needle-shaped crystals when the solution cools. It occurs 
 in nature in small quantity, as cotunnite. 
 
 Lead Tetrachloride, PbCl 4 , is a liquid that forms a 
 crystalline mass at 15. It is decomposed into the 
 dichloride and chlorine at the ordinary temperature. 
 At 105 this takes place with explosion. It is formed by 
 passing chlorine into concentrated hydrochloric acid 
 containing lead chloride. After saturation, ammonium 
 chloride is added, when the double salt, PbCl 4 . 2NH 4 C1, 
 separates in yellow crystals. Ice-cold concentrated sul- 
 phuric acid decomposes it into ammonium sulphate, 
 hydrochloric acid, and lead tetrachloride, which collects 
 as a yellow oil below the sulphuric acid. 
 
 Lead Iodide, PbI 2 , is a yellow substance which crystal- 
 lizes from water in beautiful lustrous laminae. It is pre- 
 cipitated when potassium iodide is added to a solution 
 of a lead salt. It dissolves in potassium iodide and 
 forms a salt of the formula PbI 2 .KI or KPbI 3 . There 
 
OXIDES OF LEAD. 649 
 
 is also a hydrogen compound of the formula H 2 PbI 4 , 
 that is formed by dissolving lead iodide in hydriodic 
 acid. 
 
 Lead Hydroxide, Pb(OH) 2 , is not known, but when a lead 
 salt is treated with a soluble hydroxide a compound close- 
 ly related to this hydroxide is formed. This varies some- 
 what in composition, according to the method by which 
 it is made ; but it is usually either HO-Pb-O-Pb-OH or 
 HO-Pb-O-Pb-O-Pb-OH. 
 
 Oxides of Lead. Lead forms four distinct compounds 
 with oxygen, the formulas and names of which are as 
 follows : lead suboxide, Pb 2 O ; lead oxide, PbO ; lead ses- 
 quioxide, Pb 2 O 3 ; and lead peroxide, PbO 2 . 
 
 Lead Suboxide, Pb 2 O. This compound, as has been 
 stated, is formed when lead is exposed to the air. It has 
 been made in pure condition, and is a black powder. 
 When treated with acids it yields salts, and lead sepa- 
 rates. Thus with hydrochloric acid lead chloride is 
 formed, as represented in the equation 
 
 Pb 2 + 2HC1 = PbCl 2 + H 2 O + Pb. 
 
 Lead Oxide, PbO. This compound is formed by heat- 
 ing lead nitrate, and is then left behind in the form of a 
 yellow powder. If heated to melting it solidifies, form- 
 ing a yellowish or reddish mass known as litharge. This 
 is formed in large quantity in the process of separating 
 silver from lead. It will be remembered that, in order 
 to remove the last portions of lead from silver, the alloy 
 is melted and air blown upon it in this condition. Un- 
 der these circumstances, the lead is converted into the 
 oxide while the silver remains unchanged. The lead 
 oxide thus formed is the litharge found in the market. 
 The oxide can be obtained in crystallized form by heat- 
 ing a solution of the ordinary oxide in dilute caustic soda 
 or caustic potash ; and by boiling lead hydroxide with a 
 quantity of caustic alkali insufficient to dissolve it. In 
 the powdered condition lead oxide attracts carbon diox- 
 ide from the air. With acids it forms salts which in 
 some respects resemble those of barium and strontium. 
 With the strongest bases it forms salts similar to those 
 
650 INORGANIC CHEMISTRY. 
 
 formed by zinc. This is seen in the solubility of the 
 hydroxide in sodium and potassium hydroxide, which is 
 due to the formation of compounds known as plumbites : 
 
 Pb(OH) 2 + 2KOH = Pb(OK) 2 + 2H 2 O. 
 
 Heated with silicon dioxide it forms a silicate which is 
 easily fusible. Lead oxide is used extensively in the 
 manufacture of flint glass for optical purposes, as was 
 described under Glass (which see). It also finds appli- 
 cation in glass painting and porcelain painting, and is 
 used in the manufacture of lead salts, particularly " sugar 
 of lead," which is an acetate. 
 
 Lead Sesquioxide, Pb 2 O 3 , is obtained by bringing to- 
 gether lead acetate and caustic soda, and treating the 
 solution with sodium hypochlorite ; and is thrown down 
 as a reddish-yellow powder. The constitution of the 
 compound is unknown. It has been suggested that it 
 may be a lead salt of plumbic acid, as represented in the 
 
 formula PbO<Q>Pb, plumbic acid being, as will be 
 
 seen, 
 
 Lead Peroxide, PbO 2 , is formed by treating minium or 
 red lead with dilute nitric acid. Minium has the compo- 
 sition, Pb 3 O 4 . When treated with nitric acid, a part dis- 
 solves as lead nitrate, and lead peroxide remains behind, 
 as represented in the equation : 
 
 Pb 3 O 4 + 4HNO 3 = PbO a + 2Pb(N0 3 ) 2 + 2H 2 O. 
 
 The peroxide is formed in general by the action of oxidiz- 
 ing agents upon the lower oxides of lead. One of the 
 most convenient methods for making it consists in treat- 
 ing lead acetate with a filtered solution of bleaching- 
 powder. It is a dark-brown powder, insoluble in water. 
 When ignited it loses half of its oxygen, and it gives up 
 its oxygen readily to other substances. Towards hydro- 
 chloric acid it acts like manganese dioxide, giving lead 
 chloride and chlorine according to the equation 
 
 PbO a + 4HC1 = PbCl 2 + 2H 2 + C1 2 . 
 
RED LEAD. 651 
 
 It appears probable that the tetrachloride is first 
 formed, and that this then breaks down into the dichlo- 
 ride and chlorine. When the peroxide is treated in the 
 cold with hydrochloric acid it dissolves, and when this 
 solution is heated it gives off chlorine. Further, when it 
 is treated with caustic alkalies lead peroxide is thrown 
 down. 
 
 Lead peroxide dissolves in concentrated caustic potash 
 
 OTC 
 
 and forms a salt of the formula K 9 PbO 3 , or PbO<oK' 
 
 analogous to potassium stannate, K 2 SnO 3 , silicate, 
 K 2 SiO 3 , carbonate, K 2 CO 3 , etc. Other salts derived from 
 the acid PbO(OH) 2 are known, and are called plumbates. 
 Red Lead, Minium. When lead oxide is heated gently 
 in the air it takes up oxygen, and is converted into the 
 red compound known as minium or red lead. The com- 
 mercial article of this name varies in composition, but 
 approximates to that represented by the formula 
 Pb 3 O 4 , and, if the oxide is slowly heated, the amount of 
 oxygen taken up is that required to form a compound of 
 the above formula. Red lead varies in color from red 
 to yellowish, according to the method of preparation. 
 When heated, it becomes dark, but the red color appears 
 again on cooling. When heated to a high temperature, 
 it loses oxygen and yields lead oxide : 
 
 Pb,0 4 = 3PbO + O. 
 
 When treated with dilute nitric acid, lead nitrate is 
 formed, and lead peroxide is left undissolved. As re- 
 gards the relation existing between minium and the 
 other oxides of lead, no positive statement can be made. 
 The evidence points to the conclusion that it is a chemi- 
 cal compound and not a mixture. Dilute acetic acid 
 does not dissolve it, while this acid does dissolve the 
 monoxide. It has been suggested that it is a lead salt 
 of normal plumbic acid, Pb(OH) 4 , as represented in the 
 
 formula 
 
 'O 
 O 
 
 >Pb 
 
 o> pb 
 
 As partial experimental evidence 
 
652 INORGANIC CHEMISTRY. 
 
 in support of this view, the fact may be mentioned that a 
 compound similar to red lead is formed, when a solu- 
 tion of potassium plumbate is treated with a solution 
 of lead oxide in potassium hydroxide. In solution, the 
 potassium salt probably has the constitution repre- 
 
 roH 
 
 OH 
 
 sented by the formula Pb -j Q-^-. When this is treated 
 
 OK 
 
 with lead oxide the corresponding lead salt should be 
 formed. Red lead is used as a pigment, and sometimes 
 in place of litharge when an oxide of lead is needed: as 
 in the manufacture of glass, as a flux in the manufacture 
 of porcelain, etc. 
 
 Lead Sulphide, PbS This has already been referred 
 to as the principal compound from which lead is ob- 
 tained. The natural variety is called galena or galenite. 
 It is formed in the laboratory as a black precipitate, when 
 hydrogen sulphide is passed into a solution of a lead salt. 
 When heated in the air, as in the roasting of galenite, 
 the sulphur passes off as sulphur dioxide, and the lead 
 is converted into oxide. Concentrated hydrochloric acid 
 dissolves it. Concentrated nitric acid converts it into the 
 sulphate. When hydrogen sulphide is conducted into a 
 weak acid solution of lead chloride, a compound contain- 
 ing lead, sulphur, and chlorine is precipitated, the com- 
 position of which is approximately that represented by 
 the formula 3PbS.PbCl 2 , and this has a red or a yellow 
 color, according to the conditions. 
 
 Lead Nitrate, Pb(NO 3 ) 2 . The nitrate is easily made 
 by dissolving lead, lead oxide, or carbonate in nitric acid. 
 The salt crystallizes well, and is easily soluble in water. 
 It is difficultly soluble in dilute nitric acid, and insoluble 
 in concentrated nitric acid, resembling in this respect 
 barium nitrate. It is decomposed by heat, giving nitro- 
 gen peroxide, NO 2 , and lead oxide. 
 
 Lead Carbonate, PbCO 3 . The carbonate occurs in na- 
 ture as cerussite, crystallized in forms which are the 
 same as those of barium carbonate, and of that variety of 
 calcium carbonate known as aragonite. It can be ob- 
 
LEAD CARBONATE. 653 
 
 tained by adding lead nitrate to a solution of ammonium 
 carbonate, but, when solutions of lead salts are treated 
 with the secondary carbonates of the alkali metals, pre- 
 cipitates of basic carbonates are always obtained. When 
 an excess of sodium carbonate is added to a solution 
 of lead nitrate, the precipitate has the composition 
 HO-Pb-0-CO-O-Pb-O-CO-O-Pb-OH, or 3PbO.2CO 2 
 + H 2 O. The salts usually obtained are more complicated 
 than this, but the relations between them and lead oxide 
 and carbonic acid are of the same kind. Basic lead car- 
 bonate is prepared and used extensively, under the name 
 of white lead, as a pigment. It is manufactured by differ- 
 ent methods. The principal ones are the following : 
 
 (1) The Dutch Method. This consists in exposing sheets 
 of lead wound in spirals to the action of vinegar, air, 
 and carbon dioxide from decaying organic matter. The 
 spirals of sheet lead are placed in earthenware vessels, 
 on the bottom of which, but not in contact with the lead, 
 the vinegar is placed. The vessels thus arranged are 
 placed in beds of horse manure. In consequence of de- 
 composition, which is set up in the manure, carbon diox- 
 ide is given off slowly, and enough heat is generated to 
 start the action upon the lead. The chemical changes 
 involved in the process are, mainly, the formation of a 
 basic acetate of lead, and the subsequent decomposition 
 of this by carbon dioxide, forming a basic carbonate, 
 and leaving the acetic acid free to act upon a further 
 quantity of lead. 
 
 (2) The French Method. In this method a solution of 
 basic lead acetate is prepared by treating a solution of 
 the neutral salt with lead oxide. This is then decom- 
 posed by passing carbon dioxide into it, when a basic 
 carbonate is thrown down. The carbon dioxide is gen- 
 erally made by burning coke. 
 
 (3) The English Method. This is a modification of the 
 Dutch method, and differs from it chiefly in the replace- 
 ment of manure by spent tan in a state of fermentation, 
 and the use of dilute acetic acid in place of vinegar. 
 There is less risk of discoloration in consequence of 
 the formation of sulphuretted hydrogen, but the fermen- 
 
654 INORGANIC CHEMISTRY. 
 
 tation takes place more slowly, and the whole process, 
 therefore, requires a longer time. 
 
 The composition of white lead is not always the 
 same. That prepared by precipitating a solution of basic 
 lead acetate with carbon dioxide has the composition 
 Pb(OH) 2 .3PbCO 3 ; and that prepared by the Dutch 
 method has the composition Pb(OH) 2 .2PbCO 8 ; or these 
 may be expressed structurally by the formulas 
 
 Pb< o 
 t -' 
 
 and 
 
 Pb< OH OH 
 
 An objection to white-lead paint is that it turns dark 
 under the influence of hydrogen sulphide. It also turns 
 yellow in consequence of the action of some substance 
 contained in the oil with which the lead carbonate is 
 mixed. 
 
 Lead Sulphate, PbSO 4 , occurs to some extent in nature. 
 It is formed by adding sulphuric acid or a soluble sul- 
 phate to a solution of a lead salt, and by oxidation of 
 lead sulphide. Like barium sulphate, it is practically 
 insoluble in water. As stated above, it is somewhat 
 soluble in concentrated sulphuric acid, and it is there- 
 fore always found in the concentrated acid of commerce. 
 Nitric acid and hydrochloric acid dissolve it in consider- 
 able quantity. It dissolves further quite readily in solu- 
 tions of some ammonium salts, as in ammonium tartrate 
 and acetate. When heated to redness it is partly decom- 
 posed with loss of sulphur trioxide. 
 
 Reactions which are of Special Value in Chemical Analy- 
 sis. The reactions of lead salts with the soluble hy- 
 droxides, with sulphuric acid, hydrochloric acid, hydro- 
 gen sulphide, soluble carbonates, potassium chromate 
 and dichromate, are the ones which are principally used 
 in analysis. All of these have been treated of in this 
 chapter, with the exception of those with potassium 
 chromate and dichromate, which will be taken up in the 
 chapter on Chromium (which see). In anticipation it 
 
LANTHANUM CERIUM. 655 
 
 may be said that the reactions are based upon the fact 
 that lead chromate, PbCrO 4 , like barium chromate, is 
 insoluble in water. 
 
 The elements of Family V, Group A, are vanadium, 
 columbium, didymium, and tantalum. As they are 
 closely related to the members of Group B, of the same 
 family, they were treated of at the end of Chapter XVIII. 
 in connection with the members of the phosphorus 
 group. Among them the one which is least known is 
 didymium. This in turn is more or less closely related 
 to two other elements of nearly the same atomic weight 
 which occur in Families III and IV. These are lantha- 
 num and cerium. A few words in regard to these three 
 rare elements will suffice for the present purpose. 
 
 LANTHANUM, CERIUM, DIDYMIUM. 
 
 These three elements occur together in several rare 
 minerals of Norway, as cerite, gadolinite, and allanite. 
 Cerite is a silicate of the three metals, and its composi- 
 
 Ce.) 
 tion is represented by the formula La 4 V (SiO,) 3 + 3H a O. 
 
 Di. j 
 
 It is probably a mixture of three isomorphous silicates. 
 The principal constituent is cerium silicate, Ce 4 (SiO 4 ) 3 . 
 The perfect separation of the constituents of the mineral 
 is a very difficult operation. 
 
 Lanthanum, La (At. Wt. 137.59), forms an oxide of the 
 formula La 2 O 3 , analogous to that of aluminium. Its 
 chloride also is analogous to that of aluminium, and has 
 the composition LaCl s ; and in all its salts it acts as a 
 trivalent element. 
 
 Cerium, Ce (At. Wt. 139.1), forms two series of com- 
 pounds, in one of which it is trivaient, resembling lan- 
 thanum and the other members of the aluminium group ; 
 while in the other series it is quadrivalent, resembling 
 silicon and the other members of the silicon group. 
 The formulas of some of the principal members of the 
 first series are as follows : 
 
 CeCl 3 , Ce 2 3 , and Ce 2 (S0 4 ) 3 . 
 
656 INORGANIC CHEMISTRY. 
 
 Some of the principal members of the second series are 
 represented by the formulas 
 
 CeF 4 , CeO 2 , Ce(NO 3 ) 4 , and Ce(SO 4 ) a . 
 
 Didymium, Di (At. Wt. 142.1), has already been re- 
 ferred to on page 351 in connection with the members 
 of Family Y, Group A, which it resembles in some re- 
 spects. In most of its compounds it is, however, triva- 
 lent, forming compounds, of some of which the following 
 are the formulas : 
 
 DiCl s , Di 2 3 , Di(N0 3 ) 3 , Di 2 (S0 4 ) 3 , Di 2 (CO 3 ) 3 , etc. 
 
 Praseodymium, Pr, and Neodymium, Nd. While the 
 name didymium is still given above, and this substance 
 dealt with as though it were an element, as it was at first 
 held to be, it has been shown by Auer von Welsbach 
 that it consists of two very similar elements to which he 
 has given the names praseodymium and neodymium.. When 
 the double nitrate of ammonium and didymium is re- 
 peatedly recrystallized it is separated into two salts, one 
 of which is green, and the other rose-colored. When the 
 nitrate or oxalate of one of these new elements is ignited 
 it forms a black oxide, while from the other is formed an 
 oxide of a different color. The element that gives green 
 salts is called praseodymium, and the other neodymium. 
 The atomic weights of these elements are nearly the 
 same, but they have not yet been accurately determined. 
 
CHAPTER XXXI. 
 
 ELEMENTS OF FAMILY VI, GROUP A : 
 CHROMIUM MOLYBDENUM TUNGSTEN URANIUM. 
 
 General. At the end of Chapter XIY., in connection 
 with the elements of the sulphur group, the four ele- 
 ments which form the subject of this chapter were briefly 
 referred to, for the reason that in some respects they 
 resemble sulphur. As was there stated, this resem- 
 blance " is seen mainly in the formation of acids of the 
 formulas H 2 CrO 4 , H 2 MoO 4 , H 2 WO 4 , and H 2 UO 4 ; and the 
 oxides CrO 3 , MoO 3 , WO 3 , and UO 3 ." Further, it was 
 stated that " when the acids of chromium, molybdenum, 
 tungsten, and uranium lose oxygen, they form com- 
 pounds which have little or no acid character. The lower 
 oxides of chromium form salts with acids, and these bear 
 a general resemblance to the salts of aluminium, iron, 
 and manganese. The chromates lose their oxygen quite 
 readily when acids are present with which the chromium 
 can enter into combination as a base-forming element." 
 " Molybdenum and tungsten do not form salts of this 
 character : indeed they seem to be practically devoid of 
 the power to form bases. Uranium, on the other hand, 
 forms some curious salts which differ from the simple 
 metallic salts which we commonly have to deal with. 
 These are the uranyl salts which are regarded as acids, 
 in which the hydrogen is either wholly or partly replaced 
 by the complex UO 2 , which is bivalent. Thus, the nitrate 
 has the formula UO 2 (NO 3 ) 2 , the sulphate (UO 2 )SO 4 , etc. 
 These salts are derived from the compound UO 2 (OH) 2 , 
 acting as a base, whereas the compound has also dis- 
 tinctly acid properties." That member of the group the 
 compounds of which are most commonly met with in 
 the laboratory and in the arts is chromium, and this will 
 receive principal attention here. 
 
 (657) 
 
658 INORGANIC CHEMISTRY. 
 
 CHROMIUM, Or (At. Wt. 51.74). 
 
 General. This element forms three series of com- 
 pounds, in which it appears to be respectively bivalent, 
 trivalent, and sexivalent. Of these the members of the 
 series in which it is trivalent are most stable under ordi- 
 nary circumstances. Some of the principal members of 
 the first series, or the chromous compounds, are repre- 
 sented by the formulas 
 
 CrCl 2 , Cr(OH) 2 , CrSO 4 , CrCO 3 . 
 
 Of the second series, or the chromic compounds, some 
 of the principal members are : 
 
 CrCl 3 , Cr 2 3 , Cr 2 (SO 4 ) 3 , Cr(NO 3 ) 3 , KCr(SO 4 ) 2 + 12H 2 O. 
 
 And, finally, the members of the third series are derived 
 from the oxide CrO 3 , and they are for the most part salts 
 of the acid of the formula H 2 CrO 4 , known as chromic acid, 
 or of an acid of the formula H 2 Cr 2 O 7 , known as dichromic 
 acid, which is closely related to chromic acid. 
 
 When exposed to the air the chromous compounds 
 are converted into chromic compounds, and they are in 
 general readily converted into chromic compounds by 
 the action of oxidizing agents, as cuprous and mercurous 
 compounds are converted into cupric and mercuric com- 
 pounds. If the oxidation takes place in acid solution 
 the limit is reached when a chromic salt is formed. If, 
 however, the action takes place in the presence of a 
 strong base the limit is reached in the formation of a 
 chromate. Thus, suppose chromous oxide to be treated 
 with an oxidizing agent in the presence of sulphuric acid, 
 the final product would be chromic sulphate, as repre* 
 sented in the following equations : 
 
 CrO + H 2 S0 4 = CrS0 4 + H 2 O ; 
 
 2CrS0 4 + H 2 S0 4 + O = Cr 2 (SO 4 ) 3 + H 2 O. 
 
 On the other hand, if the oxidation takes place in the 
 presence of caustic potash the final product is potassium 
 chromate, as shown in the following equation : 
 
 CrO + 2KOH + O a = K 2 CrO 4 + H 2 O. 
 
CHROMIUM. 659 
 
 When a chromate is treated with an acid it tends to 
 pass back to a compound of the chromic series, and the 
 change involves the giving up of oxygen. Thus when 
 potassium chromate is treated with sulphuric acid in the 
 presence of something which has the power to take up 
 oxygen, potassium and chromium sulphates are formed, 
 and oxygen is given up, thus : 
 
 2K 2 CrO 4 + 5H 2 SO 4 = 2K 2 SO 4 -f Cr 2 (SO 4 ) 3 + 5H 2 O -f 3O. 
 
 All these relations will be more fully taken up in the 
 paragraphs which treat of the individual compounds. 
 
 Forms in which Chromium Occurs in Nature. The 
 principal form in which chromium occurs in nature is 
 the mineral chromite, also known as chromic iron and 
 chrome iron ore. This has the composition FeCr 2 O 4 , 
 and, as will be pointed out below, it is probably analo- 
 gous to the spinels (see p. 572), being an iron salt of the 
 acid CrO.OH, which may be called metachromous acid. 
 
 CrO O 
 This view is represented by the formula ^ Q*^>Fe. 
 
 It occurs also in the mineral crocoisite, which is lead 
 chromate, PbCrO 4 . The name chromium is derived from 
 the Greek jpcSyua', meaning color ; and the element is so 
 called because most of its compounds are colored. 
 
 Preparation. The metal is obtained by the electroly- 
 sis of chromic chloride ; by decomposing the chloride 
 by means of sodium in the form of vapor ; and by treat- 
 ing the chloride with zinc. 
 
 Properties. Chromium is a light-gray, crystalline, lus- 
 trous, metallic-looking substance ; or it consists of mi- 
 croscopic, lustrous rhombohedrons of a tin-white color. 
 It is very hard, and difficultly fusible. When heated in 
 the air it is oxidized very slowly, but in the flame of the 
 oxyhydrogen blowpipe it burns, forming chromic oxide, 
 Cr 2 O 3 . It is easily dissolved by hydrochloric acid. Cold 
 sulphuric acid does not dissolve it ; the hot acid does. 
 Nitric acid does not affect it. When treated with salts 
 of potassium which easily give up their oxygen, as the 
 chlorate and nitrate, it is converted into potassium 
 chromate. 
 
660 INORGANIC CHEMISTRY. 
 
 Chromous Chloride, CrCl 2 , is formed by dissolving the 
 metal in hydrochloric acid, and by carefully heating 
 chromic chloride in a current of hydrogen. It forms 
 white crystals, which dissolve in water, giving a blue 
 solution. This solution takes up oxygen very readily 
 from the air, and the compound is converted into others 
 which belong to the chromic series. The other chro- 
 mous compounds act in a similar way. 
 
 Chromic Chloride, CrCl 3 . This compound is made in 
 solution by dissolving chromic hydroxide, Cr(OH) 3 , in 
 hydrochloric acid. This solution has a dark-green color. 
 When evaporated to a sufficient extent crystals of the 
 composition CrCl 3 + 6H 2 O are deposited. If these are 
 heated in the air they undergo decomposition just as 
 aluminium chloride does, and the product left behind is 
 chromic oxide : 
 
 2CrCl 3 + 3H 2 O = Cr 2 O 3 + 6HC1. 
 
 If, however, the crystallized chloride is heated in an 
 atmosphere of chlorine or hydrochloric acid, the water 
 is given off, and the anhydrous chloride, which has a 
 beautiful reddish violet color, is formed. This dissolves 
 in water and forms a green solution. But if the dry 
 chloride thus obtained is sublimed, it is deposited in 
 lustrous laminae of the same color ; and this variety 
 is insoluble in water and acids, and is only slowly 
 acted upon by boiling alkalies. This insoluble, crystal- 
 lized variety of the chloride is obtained also by the same 
 method as that used in making aluminium chloride, that 
 is, by passing a current of chlorine over a heated mix- 
 ture of carbon and chromic oxide. Although it is called 
 insoluble, it passes gradually into solution by boiling 
 with water. Further, when a very minute quantity of 
 chromous chloride is mixed with it, it dissolves easily, 
 and forms a green colored solution. 
 
 Chromic chloride unites with other chlorides, as alu- 
 minium chloride does, and forms double chlorides, 
 analogous to the chlor-aluminates. Examples of these 
 are the compounds of the formulas CrCl 3 .KCl, or 
 
CHROMIC HYDROXIDE. 661 
 
 KCrCl 4 ; CrCl 3 .2KCl, or K 2 CrCl 6 ; and CrCl 3 .3KCl, or 
 K 3 CrCl 6 . 
 
 Chromous Hydroxide, Cr(OH) 2 , is formed as a brown- 
 ish-yellow precipitate by adding caustic potash to a solu- 
 tion of chromous chloride. It easily gives up hydrogen, 
 and is converted into chromic oxide : 
 
 2Cr(OH), = Cr.O. + H 2 O + H,. | , 
 
 Chromic Hydroxide, Cr(OH) 3 . When ammonia is added 
 to a solution of a chromic salt, a light-blue voluminous 
 precipitate, which has the composition Cr(OH) 3 -|- 2H 2 O, 
 is formed. When this is filtered off and dried in a 
 vacuum it loses the water and leaves the hydroxide. 
 This is readily converted by heat into a compound of the 
 formula CrO.OH, and finally into chromic oxide, Cr 2 O 3 . 
 The green precipitates formed in solutions of chromic 
 salts by sodium and potassium hydroxides always con- 
 tain some of the alkali metal in combination. 
 
 Chromic hydroxide, like aluminium hydroxide, dis- 
 solves in the soluble hydroxides, and forms salts known 
 as chromites, which are derived from the acid CrO.OH. 
 Thus with potassium hydroxide the action takes place 
 as represented in the equation 
 
 /OH 
 
 Crf- OH + KOH = Cr^Xi^ + 2H 2 O. 
 X)H 
 
 If the solution containing potassium or sodium chro- 
 mite is boiled, the salt is decomposed and chromic hy- 
 droxide precipitated, though the precipitate thus formed 
 always contains some of the alkali metal in combination. 
 It will be noticed that in this respect aluminium and 
 chromium conduct themselves differently towards the 
 alkaline hydroxides. 
 
 It has already been stated that chromite, (CrO.O) 2 Fe, 
 is regarded as an iron salt of the same order as the po- 
 tassium salt referred to. 
 
662 INORGANIC CHEMISTRY. 
 
 Another hydroxide formed by heating potassium di- 
 chromate and boric .acid together has the composition 
 represented by the formula Cr 2 O(OH) 4 or Cr 4 O 3 (OH) 6 . 
 This is known as Guignet's green. The relation between 
 the normal hydroxide and these compounds is shown 
 by means of the equations 
 
 2Cr(OH) 3 = Cr 2 O(OH) 4 + H 2 O ; 
 4Cr(OH) 3 = Cr 4 O 3 (OH) 6 + 3H 2 O. 
 
 Chromic Oxide, Cr 2 O s , is formed by igniting the hy- 
 droxides, and is most readily prepared by heating a 
 mixture of potassium dichromate and sulphur. The 
 sulphur is oxidized, and with the potassium forms 
 potassium sulphate, while the chromic acid is reduced 
 to the form of the oxide Cr 2 O 3 . It can be obtained in 
 crystals. As ordinarily obtained it is a green powder, 
 which after ignition is almost insoluble in acids. It is 
 dissolved, however, by treatment with fusing potassium 
 sulphate. The oxide colors glass green, and is used in 
 painting porcelain. 
 
 Chromic Sulphate, Cr 2 (SO 4 ) 3 , is made by dissolving the 
 hydroxide in concentrated sulphuric acid when it is de- 
 posited in purple crystals of the composition Cr 2 (SO 4 ) 3 -f- 
 15H 2 O. If the solution of this salt is boiled, the so- 
 lution becomes green, and crystals cannot be obtained 
 from it. But by standing for some time the green solu- 
 tion becomes reddish purple again, and yields the crys- 
 tallized salt. Other salts of chromium act in the same 
 way. They exist in two varieties, one of which crystal- 
 lizes and is reddish purple in color, while the other does 
 not crystallize and is green. The crystallized salts are 
 converted into the uncrystallized green salts by boiling, 
 and the green salts are converted into the crystallized 
 salts by standing. 
 
 Chrome- Alums. Chromic sulphate, like aluminium sul- 
 phate, combines with other sulphates, such as potassium, 
 sodium, and ammonium sulphates, and forms well-crys- 
 tallized salts, which are closely analogous to ordinary 
 
CHROMIC ACID AND THE CHROMATES. 663 
 
 alum. They all contain twelve molecules of water, as 
 represented in -the formulas below : 
 
 Chrome-Alum, KCr(SO 4 ) 2 -f 12H O 
 
 Sodium Chrome-Alum, . . . NaCr(SO 4 ) 2 + 12H 2 ~O 
 Ammonium Chrome-Alum, . (NH 4 )Cr(SO 4 ) 2 + 12H 2 O 
 
 The potassium compound which is commonly called 
 chrome-alum is made by adding a reducing agent, such 
 as alcohol or sulphur dioxide, to a solution of potas- 
 sium dichromate containing sulphuric acid. If the solu- 
 tion is heated it turns green, and crystals cannot be ob- 
 tained from it. But on standing for a considerable time 
 its color changes, and reddish-purple crystals of the 
 alum are deposited. This change can be facilitated by 
 putting some crystals of the salt in the concentrated 
 green solution. The action of reducing agents upon po- 
 tassium dichromate will be treated of farther on. The 
 salt finds application in dyeing and tanning. 
 
 Chromic Acid and the Chromates. It has already been 
 stated that when chromium compounds belonging to the 
 chromous and chromic series are oxidized in the pres- 
 ence of bases they are converted into chromates. These 
 salts are derived from an acid of the formula H 2 CrO 4 , 
 which is unknown, as it breaks down spontaneously into 
 chromium trioxide, CrO 3 , and water, when it is set free 
 from its salts, just as carbonic and sulphurous acids break 
 down respectively into carbon dioxide and water, and 
 sulphur dioxide and water. The starting-point for the 
 preparation of the chromates and the compounds re- 
 lated to them is chromic iron. This is ground fine, inti- 
 mately mixed with a mixture of caustic potash and lime, 
 and then heated in shallow furnaces in contact with the 
 air. ' Under these circumstances oxidation is effected by 
 the oxygen of the air. The iron is converted into ferric 
 oxide, and the chromium gives, with the calcium and 
 potassium, the corresponding chromates, CaCrO 4 and 
 K 2 CrO 4 . When the mass is treated with water these 
 salts dissolve, and ferric oxide remains undissolved. By 
 treating the solution with potassium sulphate the cal- 
 cium salt is converted into the potassium salt, and thus 
 
664 INORGANIC CHEMISTRY. 
 
 all the chromium appears in the form of potassium ehro* 
 mate, the changes referred to are represented in the fol- 
 lowing equations : 
 
 2(Cr0 2 ) 2 Fe + 8KOH + 7O = 4K 2 CrO 4 + Fe 2 O 3 + 4H 2 O : 
 2(CrO 2 ) 2 Fe + 4CaO + 7O = 4CaCrO 4 + Fe 2 O 3 ; 
 OaCr0 4 + K 2 SO 4 = K 2 CrO 4 + CaSO 4 . 
 
 As potassium chromate is easily soluble in water, and 
 therefore difficult to purify, it is converted into the 
 dichromate, which is less soluble and crystallizes well. 
 The change is easily effected by adding the necessary 
 quantity of a dilute acid. If nitric acid is used the re- 
 action is represented by the following equation : 
 
 2K 2 Cr0 4 + 2HN0 3 = K 2 Cr 2 O 7 + 2KNO 3 + H 2 O. 
 
 The salt thus obtained is manufactured on the large 
 scale and is the starting-point for the preparation of 
 other chromium compounds. 
 
 Potassium Chromate, K 2 CrO 4 , formed as above de- 
 scribed, is a light-yellow crystallized substance which is 
 easily soluble in water. It is isomorphous with potas- 
 sium sulphate. Acids convert it into the dichromate, as 
 just stated. 
 
 Potassium Dichromate, K 2 Cr 2 O7. This salt forms large 
 red plates, which are triclinic. It is soluble in ten parts 
 of water at the ordinary temperature, and is much more 
 soluble in hot water. When heated, it at first melts 
 without undergoing decomposition ; at white heat, how- 
 ever, it is decomposed, yielding the chromate, chromic 
 oxide, and oxygen : 
 
 2K 2 Cr 2 O 7 = 2K 2 Cr0 4 + Cr 2 O 3 + 3O. 
 
 It undergoes a similar change, but much more readily, 
 when heated with concentrated sulphuric acid. In this 
 case, however, the chromic oxide forms chromic sulphate 
 with the acid, and this forms chrome-alum with the po- 
 tassium sulphate : 
 
 K 2 Cr 2 O 7 + 4H 2 SO 4 = 2KCr(SO 4 ) 2 + 4H 2 O + 3O. 
 All the oxygen in the chromate in excess of that required 
 
POTASSIUM DICHROMATE. 665 
 
 to form the alum and water is given off. This also is 
 the character of the action towards reducing agents in 
 general. One molecule of the dichromate gives three 
 atoms of oxygen. With sulphur dioxide the action is 
 that represented in the equation 
 
 K 3 O 2 O 7 -f 4H 2 SO 4 + 3SO 2 = 2KCr(SO 4 ) 2 + 3H 2 SO 4 -f H 2 O. 
 
 Or, one molecule of the dichromate converts three mole- 
 cules of sulphur dioxide, SO 2 , into three molecules of 
 sulphuric acid, H 2 SO 4 . 
 
 The action with alcohol will be understood by the aid 
 of the following equation, which represents the action of 
 oxygen in general upon alcohol : 
 
 C 2 H 6 + O = C 2 H 4 + H,0. 
 
 Alcohol Aldehyde 
 
 Each molecule of alcohol requires one atom of oxygen 
 to convert it into aldehyde. Therefore, one molecule of 
 the dichromate oxidizes three molecules of alcohol to 
 aldehyde : 
 
 K 2 Cr 2 7 + 4H 2 SO 4 + 3C 2 H 6 O = 2KCr(SO 4 ) 2 + 3C 2 H 4 O + 7H 2 O. 
 
 Concentrated hydrochloric acid is oxidized by the di- 
 chromate, and chlorine is evolved : 
 
 K 2 Cr 2 7 + 14HC1 = 2KC1 + 2CrCl 3 + 7H 2 O + 601. 
 
 Here two atoms of chlorine are required to form potas- 
 sium chloride with the potassium, and six to form chro- 
 mic chloride with the chromium ; and the eight hydro- 
 gen atoms in combination with this chlorine combine 
 with four atoms of oxygen of the dichromate, leaving 
 three more to oxidize hydrochloric acid. Consequently 
 one molecule of the dichromate sets free six atoms of 
 chlorine : 
 
 3O + 6HC1 = 3H 2 O + 6C1. 
 
 When the dichromate in solution is treated with po- 
 tassium hydroxide, its color changes to yellow, in con- 
 sequence of the formation of the chromate, the action 
 taking place as represented in this equation : 
 
 K 2 Cr 2 O 7 + 2KOH = 2K 2 CrO 4 -f H 2 O. 
 
6 GO INORGANIC CHEMISTRY. 
 
 Potassium dichromate finds extensive use in the arts 
 and in the laboratory as an oxidizing agent. With gela- 
 tine it forms a mixture which is sensitive to light, which 
 turns it dark, and makes it insoluble. This fact is made 
 the basis of a number of photographic processes. The 
 dichromate is used, further, in dyeing. 
 
 Chromium Trioxide, CrO 3 , crystallizes out on cooling 
 when either the chromate or the dichromate is treated 
 in concentrated solution with concentrated sulphuric acid. 
 This is a beautiful red substance, which crystallizes in 
 needles. When dissolved in water it forms a solution 
 from which, by neutralization, the chromates can be ob- 
 tained. When heated alone it gives off half its oxygen, 
 and is converted into chromic oxide : 
 
 2Cr0 3 = Cr 2 3 + 3O ; 
 
 and when heated with sulphuric acid it gives chromic 
 sulphate and oxygen : 
 
 2CrO 3 + 3H 2 SO 4 = Cr 2 (SO 4 ) 3 + 3H 2 O + 3O. 
 
 It is an extremely active oxidizing agent, disintegrating 
 most organic substances with which it is brought in con- 
 tact. 
 
 Relations between the Chromates and Bichromates. 
 The fact that chromium trioxide with water gives chro- 
 mic acid, which is a dibasic acid, whose salts in general 
 resemble those of sulphuric acid, leads to the belief that 
 the structure of chromic acid should be represented by 
 a formula similar to that of sulphuric acid, thus : 
 
 O O 
 
 HO-S-OH HO-Cr-OH 
 
 6 b. 
 
 or or 
 
 0,S(OH), 2 Cr(OH), 
 
 Just as sulphuric acid by loss of water is converted into 
 disulphuric acid or pyrosulphuric acid, so chromic acid 
 is converted into dichromic acid, and in all probability 
 the relation between the chromates and dichromates is 
 
RELATIONS OF THE CHROMATES AND BICHROMATES. 667 
 
 the same as that between the sulphates and disulphates, 
 as represented by the equations 
 
 9O t < 
 
 CrO,< 
 
 But, as has been stated, neither chromic acid nor di- 
 chromic acid is known, as they break down into chromium 
 trioxide and water when set free from their salts. The 
 conversion of potassium chromate into the dichromate 
 by treatment with an acid is represented as follows : 
 
 .OK Prft . 
 
 '<OK HNO, _ Cr ><OH 
 
 N S " OrO.<g| + 
 
 CrO,< K 
 
 Sodium Chromate, Na 2 CrO 4 , and Sodium Dichromate, 
 Na 2 Cr 2 O7, are made in the same way as the potassium 
 compounds. They are both deliquescent as ordinarily 
 made. It has, however, recently been shown that a so- 
 dium dichromate which is not deliquescent can be made ; 
 and at present it is largely manufactured instead of the 
 somewhat more expensive potassium salt. 
 
 Barium Chromate, BaCrO 4 , like the sulphate, is insol- 
 uble in water, and is precipitated when a solution of a 
 barium salt is brought together with a soluble chromate 
 or dichromate. It is soluble in hydrochloric acid and 
 nitric acid, but not in acetic acid. The strontium salt is 
 somewhat soluble in water, and easily in hydrochloric,. 
 nitric, and acetic acids. 
 
668 INORGANIC CHEMISTRY. 
 
 Lead Chromate, PbCrO 4 , occurs in nature, and is formed 
 as a beautiful yellow precipitate when a solution of a 
 lead salt is treated with a solution of a chromate or di- 
 chromate. It is used as a pigment under the name 
 chrome yellow. When heated to a high temperature it 
 gives off some of its oxygen. In contact with oxidizable 
 substances it gives off its oxygen very easily, and it is 
 used in the laboratory in some cases instead of cupric 
 oxide in the analysis of organic compounds (see p. 594). 
 Treated with dilute potassium or sodium hydroxide in 
 insufficient quantity to dissolve it, it turns red in conse- 
 quence of the formation of a basic chromate which is 
 known as chrome red. The action takes place as repre- 
 sented in the equation 
 
 CrO,<>Pb 
 
 2KOH = 
 
 This then loses water, and forms the salt CrO 2 <Q~pi >O, 
 
 which is chrome red. 
 
 Lead chromate dissolves completely in the caustic 
 alkalies in consequence of the formation of chromates 
 and plumbites. 
 
 Silver chromate, Ag 2 CrO 4 , is formed as a red precipitate 
 when a chromate is treated with a silver salt. 
 
 Potassium trichromate t K 2 Cr 3 O ]0 , and potassium tetra- 
 cJiromate, K 2 O 4 O 13 , are formed by treating the dichromate 
 with nitric acid : 
 
 3K 2 Cr 2 O 7 + 2HNO 3 = 2K 2 Cr 3 O 10 + 2KN0 3 + H 2 O ; 
 2K 2 Cr 2 0, + 2HN0 3 - K 2 Cr 4 O 13 + 2KNO 3 + H 2 O. 
 
 The acids from which these salts are derived bear to or- 
 dinary chromic acid relations similar to those which the 
 polysilicic acids bear to ordinary silicic acid. 
 
 Chromium Oxychloride, Chromyl Chloride, CrOzCla, 
 is analogous to sulphuryl chloride, SO 2 C1 2 , and is to be 
 
ANALYTICAL REACTIONS OF CHROMIUM. 669 
 
 regarded as derived from chromic acid by the replace- 
 ment of the hydroxyls by chlorine : 
 
 CrO.<g|; OrO.<g|. / -; 
 
 It is formed by treating a mixture of sodium chloride 
 and potassium dichromate with sulphuric acid. Prob- 
 ably the sulphuric acid sets free hydrochloric acid and 
 chromic acid simultaneously, and these then act upon 
 each other as represented in the equation 
 
 It is a dark -red-colored liquid, which boils at 116 with- 
 out decomposition. With water it gives chromic acid 
 and hydrochloric acid, and with an alkaline hydroxide it 
 gives the corresponding chromate : 
 
 182 = Cr < 2HCL 
 
 Reactions which are of Special Value in Chemical Anal- 
 ysis. The reactions of chromic salts with the alkaline 
 hydroxides have been explained. With the soluble car- 
 bonates they give precipitates which consist mainly of 
 the hydroxide, though carbonic acid is to some extent 
 in combination in them. These precipitates are soluble 
 in a large excess of the carbonate. 
 
 The solution of chromic oxide in an alkali is green, 
 but by oxidizing agents it is turned yellow in consequence 
 of the formation of a chromate. 
 
 Hydrogen sulphide does not precipitate chromium from 
 its salts. Ammonium sulphide precipitates chromium hy- 
 droxide, the reaction being the same as in the case of 
 aluminium (which see). 
 
 Chromates give with barium and lead salts yellow pre- 
 cipitates (see Barium Chromate and Lead Chromate). 
 
670 INORGANIC CHEMISTRY. 
 
 When heated with sulphuric acid, the chromates are 
 decomposed, with evolution of oxygen. The action of 
 hydrochloric acid upon the chromates was explained 
 above. 
 
 In neutral or slightly acid solutions of chromates, 
 hydrogen sulphide and ammonium sulphide act as re- 
 ducing agents, and precipitate the hydroxide mixed with 
 sulphur : 
 
 K 2 Cr 2 O 7 + 3H a S+2HCl = 2KC1 + 2Cr(OH) 3 +3S+ H 2 O. 
 
 With borax or microcosmic salt chromium compounds 
 form green beads, both in the oxidizing and reducing 
 flames. 
 
 MOLYBDENUM, Mo (At. Wt. 95.26). 
 
 General. Molybdenum is of interest on account of the 
 variety of its compounds. It forms four compounds with 
 chlorine, the formulas of which appear to be MoCl a , 
 MoCl 3 , MoCl 4 , and MoCl 5 . With oxygen also it forms 
 four compounds, but, while the first three are analogous 
 to the first three chlorine compounds above mentioned, 
 the last differs from the last chlorine compound. In it 
 the element is sexivalent. The formulas of the oxygen 
 compounds are MoO, Mo 2 O 3 , MoO 2 , and MoO 3 . The 
 last of these is analogous to chromium trioxide, as it 
 forms salts with bases analogous to the chromates, and 
 known as the molybdates. 
 
 Occurrence and Preparation. Molybdenum occurs in 
 nature principally as molybdenite, which is the sulphide 
 MoS 2 , and as wulfenite, which is lead molybdate, PbMoO 4 . 
 It occurs also, but in smaller quantity, as molybdenum 
 trioxide, MoO 3 . It is obtained in free condition by heat- 
 ing the oxides or chlorides in a current of hydrogen. 
 
 Properties. It is a very hard silver-white metal of 
 specific gravity 8.6. It is apparently infusible when pure. 
 When heated for a considerable time in the air it is con- 
 verted into the trioxide. It dissolves in concentrated 
 nitric acid and in aqua regia. 
 
 Chlorides. When molybdenum is heated for a long 
 time in a current of chlorine it is finally completely con- 
 verted into the pentachloride, MoCl 5 . When the penta- 
 
OXIDES OF MOLYBDENUM. 671 
 
 chloride is heated to 250 in a current of hydrogen it is 
 converted into the trichloride, MoCl 3 . And when the 
 trichloride is heated in a current of carbon dioxide it is 
 converted into the dichloride and the tetrachloride : 
 
 2MoCl 3 = MoCl a + MoCl 4 . 
 
 Oxides. The final product of the action of oxygen or 
 oxidizing agents upon molybdenum is the trioxide, MoO 3 . 
 This is obtained comparatively easily when molybdenite, 
 MoS 2 , is mixed with pure sand and roasted, then treated 
 with ammonia, which forms ammonium molybdate. This 
 salt is crystallized, and afterwards decomposed by nitric 
 acid. The oxide is a white crystallized substance which 
 is difficultly soluble in water. When a solution of the 
 trioxide is treated with reducing agents, as sodium-amal- 
 gam, lower oxides are formed, and the reduction stops 
 at molybdic oxide, Mo 2 O 3 . This is a black compound. 
 The dioxide, MoO 2 , is a dark-blue substance which is 
 formed by gently heating the metal or molybdic oxide, 
 Mo 2 O 3 , in the air, and also by heating the corresponding 
 hydroxide, Mo(OH) 4 . 
 
 Molybdenum monoxide, MoO, is formed by treating a 
 solution of the dichloride with hot caustic potash. Like 
 molybdic oxide it is black. 
 
 Molybdenum trisulphide, MoS 3 , is precipitated as a red- 
 dish-brown substance when a moderately concentrated 
 solution of a molybdate is treated with hydrogen sul- 
 phide. When heated it is converted into the disulphide, 
 MoS 2 , which is identical in composition with the molyb- 
 denite found in nature. 
 
 Molybdic Acid and the Molybdates. Molybdenum tri- 
 oxide, MoO 3 , combines readily with bases, forming salts 
 which in composition are analogous to the chromates. 
 When ammonium molybdate is treated with moderately 
 dilute nitric acid free molybdic acid crystallizes out. 
 This has the composition represented by the formula 
 
 roH 
 
 f OTT 
 H 2 MoO 4 + H 2 O, and perhaps the constitution MoO -j QTT, 
 
 [OH 
 
672 INORGANIC CHEMISTRY. 
 
 analogous to that of tetrahydroxyl-sulphuric acid. Mo- 
 lybdic acid not only forms compounds with bases, but 
 also with acids. The latter are, however, not salts, but 
 complex acids which in turn form salts with bases, some 
 of which are extremely complicated. There are a num- 
 ber of molybdates known of the general formula M 2 MoO 4 , 
 in which M represents a univalent element, but the acid 
 further shows a marked tendency to form more complex 
 salts, derivatives of polymolybdic acids. Thus the fol- 
 lowing sodium salts are known : Na 2 MoO 4 , Na 2 Mo 2 O 7 , 
 Na 2 Mo 3 10 , Na 2 Mo 4 13 , Na 2 Mo 8 O 25 , Na 2 Mo 10 O 31 , and 
 Na,Mo 7 O 24 . The relations between the acids from which 
 these salts are derived, and molybdic acid, H 2 MoO 4 , will 
 readily be seen by the aid of the following equations : 
 
 2HJMo0 4 = H 2 Mo 2 O 7 + H 2 O ; 
 3H 2 Mo0 4 = H 2 Mo 3 10 + 2H 2 O ; 
 4H 2 Mo0 4 = H 2 Mc 4 13 + 3H 2 O ; 
 8H 3 MoO 4 = H 2 Mo 8 2B + 7H 2 O ; 
 10H 2 Mo0 4 = H 2 Mo 10 31 + 9H 2 ; 
 7H 2 Mo0 4 = H 6 Mo 7 O 2 , + 4H 2 O. 
 
 Lead Molybdate, PbMoO 4 , occurs in nature, as has been 
 stated, and the mineral is known as wulfenite. It can be 
 obtained artificially by melting together sodium molyb- 
 date, lead chloride, and sodium chloride ; or by treating 
 a solution of sodium molybdate with a solution of lead 
 nitrate. If the reagents are pure the artificially prepared 
 salt is white, while the natural variety is always yellow 
 or red. 
 
 Phospho-molybdic Acid. Among the best known and 
 most frequently met with compounds of molybdic acid 
 with acids is that which it forms with phosphoric acid, 
 known as phospho-molybdic acid. When a solution of 
 ammonium molybdate in an excess of nitric acid is added 
 in excess to a solution of phosphoric acid or a phosphate, 
 a yellow precipitate is formed. This is ammonium phos- 
 pho-molybdate, which, when dried, has the composition 
 represented by the formula 12MoO 3 .(NH 4 ) 3 PO 4 . This is 
 insoluble in water and in dilute acids, and also in a nitric- 
 
TUNGSTEN. 673 
 
 acid solution of ammonium molybdate. On account of 
 the properties mentioned, this salt furnishes a valuable 
 means of detecting phosphoric acid and of precipitating 
 it from its solutions. When the salt is treated with aqua 
 regia it is decomposed, and from the solution formed 
 a compound of the composition H 3 PO 4 .llMoO 3 -|- 12H 2 O 
 crystallizes out. 
 
 TUNGSTEN, W (At. Wt. 183.43). 
 
 General. Like molybdenum, tungsten forms a large 
 variety of compounds. With chlorine it forms four, of 
 which the formulas are WC1 2 , WC1 4 , WC1 6 , and WC1 6 . 
 With oxygen, however, it forms but two compounds, and 
 these are represented by the formulas WO 2 and WO 3 . 
 The trioxide forms salts with bases which are analogous 
 to the molybdates, and, like molybdic acid, tungstic acid 
 forms complicated salts which are derived from poly- 
 tungstic acids. Further, tungstic acid combines with 
 other acids, forming very complex acids. 
 
 Occurrence and Preparation. Tungsten occurs in na- 
 ture as tungstates. The principal one is the iron salt, 
 which always, however, contains some manganese. This 
 is known as wolframite, and has the composition repre- 
 sented by the formula FeWO 4 . Calcium tungstate, 
 CaWO 4 , or scheelite, and lead tungstate, PbWO 4 , or 
 stolzite, are also found in nature, but in smaller quantity 
 than wolframite. The element is prepared by reducing 
 the chlorides or oxides in a current of hydrogen. 
 
 Properties. Tungsten forms lustrous, steel-colored 
 laminae, or a black powder. It is very hard and difficultly 
 fusible, and has the specific gravity 19.129. It is not 
 changed by contact with the air at ordinary temperatures. 
 At higher temperatures it combines with oxygen and 
 forms the trioxide, WO 3 . Nitric acid and aqua regia 
 convert it into the trioxide. It is used in the manufac- 
 ture of steel, as the addition of from 8 to 9 per cent of it 
 makes steel extremely hard. 
 
 Chlorides. When tungsten is heated in a current of 
 chlorine it is converted into the hexachloride, WC1 6 . The 
 other chlorides are formed by heating this in hydrogen. 
 
674 INORGANIC CHEMISTRY. 
 
 Oxides. Tungsten trioxide, WO 3 , is found in small 
 quantity in nature, and is formed from wolframite by a 
 number of methods. When a solution of a tungstate is 
 boiled with an acid the trioxide is precipitated as a yel- 
 low powder. Under the influence of sunlight it turns 
 greenish. It is insoluble in water. In alkalies it dis- 
 solves, forming the tungstates. When heated in a current 
 of hydrogen, lower oxides are formed ; and a blue com- 
 pound thus obtained appears to have the composition 
 represented by the formula 2WoO 3 -[- WO 2 or W 3 O 8 . 
 The dioxide, WO 2 , is obtained by further reduction. 
 This is a brown powder. 
 
 Tungstic Acid and the Tungstates. When the required 
 quantities of tungsten trioxide and potassium carbonate 
 are brought together in solution, or are melted together, 
 potassium tungstate, K 2 WO 4 , is formed. If a solution of 
 this salt is treated with a strong acid at the ordinary 
 temperature, a white precipitate of the composition 
 H 2 WO 4 -f-H 2 O is formed. This is tungstic acid, analogous 
 to crystallized molybdic acid. If the solutions are hot 
 the precipitate has the composition H 2 WO 4 . Among the 
 complex salts derived from polytungstic acids are the 
 following : Na 2 W 2 O 7 , Na 4 W 3 O n , Na 10 W ]2 O 41 , etc. The re- 
 lations between the polytungstic acids and the ordinary 
 variety of the acid, H 2 WO 4 , will be readily understood. 
 The salt, Na 10 W 12 O 41 -j- 28H 2 O, is known as sodium para- 
 tungstate. It is manufactured on the large scale by heat- 
 ing together wolframite and calcined sodium carbonate. 
 Inflammable substances impregnated with a solution of 
 the salt burn with great difficulty, and it is used to pro- 
 tect various articles from fire. The salts derived from 
 the acid, H 2 W 4 O 13 , are called metatung 'states. 
 
 Silico-tungstic Acids. Among the most interesting of 
 the complex compounds formed by the combination 
 of tungstic acid with other acids are those known as 
 the silico-tungstic acids. When sodium paratungstate, 
 Na 10 W 12 O 41 , is boiled in solution with precipitated gelati- 
 nous silicic acid, the latter dissolves, and from the solution 
 a salt of the composition Na 8 W 12 SiO 42 + 7H 2 O crystal- 
 lizes. This is soluble in one fifth its weight of water, 
 
URANIUM. 675 
 
 and the solution has the remarkably high specific gravity 
 3.05. The acid from which the salt is derived is known 
 as silico-tungstic acid. Its composition is represented 
 by the formula 4H 2 O.12WO 3 .SiO 2 ; and it may be re- 
 garded as made up of a polytungstic acid, H^W^O^, 
 in combination with one molecule of silicon dioxide, 
 H 8 W 12 4 ,SiO, 
 
 UBANIUM, U (At. Wt. 237.77). 
 
 General. Uranium has stronger basic properties than 
 either molybdenum or tungsten ; and it differs from 
 chromium in the fact that the trioxide forms salts with 
 acids. These salts are the uranyl salts which are derived 
 from the hydroxide, UO 2 (OH) 2 or H 2 UO 4 . The cor- 
 responding compounds of chromium, molybdenum, and 
 tungsten are acids. Uranium also forms salts in which 
 it acts as a quadrivalent element, as U(SO 4 ) 2 . While the 
 hydroxide, UO 2 (OH) 2 , forms salts with acids, it also forms 
 salts with the strongest bases. These are analogous in 
 composition to the dichromates, and have the general 
 formula M 2 U 2 O 7 . With chlorine, uranium forms the 
 compounds UC1 3 , UC1 4 , and UC1 5 ; and with oxygen the 
 following : U0 2 , U 3 O 8 , UO 3 , and UO 4 . 
 
 Occurrence and Preparation. Uranium occurs in na- 
 ture chiefly in the form of the mineral known as pitch- 
 blende or uraninite, which consists of the oxide,U 3 O 8 , mixed 
 with a number of other substances in smaller or larger 
 quantities. When this is finely powdered and treated 
 with concentrated nitric acid, uranyl nitrate, UO 2 (NO 3 ) 2 , 
 is obtained, and, by igniting this, the trioxide, UO 3 , is 
 left behind. In order to isolate the metal, the oxide thus 
 obtained is mixed with charcoal and treated with chlo- 
 rine, when the tetrachloride, UC1 4 , is formed. This is 
 then reduced by heating it with sodium under a cover of 
 the molten chlorides of potassium and sodium. 
 
 Properties. Uranium has the color of nickel and the 
 specific gravity 18.4. When heated to redness it is oxid- 
 ized superficially. It dissolves in dilute acids with 
 evolution of hydrogen. 
 
676 INORGANIC CHEMISTRY. 
 
 Chlorides. When chlorine acts upon finely divided 
 uranium, the two combine to form the tetrachloride, UC1 4 . 
 When this is heated in hydrogen it loses a part of its 
 chlorine and forms the trichloride, UC1 3 ; and when the 
 tetrachloride is treated with chlorine it is partly converted 
 into the pentachloride, UC1 5 . The tetrachloride is the 
 most stable form. 
 
 Oxides. The oxide of uranium which is formed as the 
 last product of the action of oxygen on uranium or the 
 other oxides when these are heated in the air is that which 
 has the composition U 3 O 8 , which is also the composition 
 of the natural variety. When this is treated with nitric 
 acid, however, it. is converted into uranyl nitrate, 
 TJO 2 (NO 3 ) 2 , which is a derivative of the trioxide, UO 3 ; 
 and when the nitrate is ignited, the trioxide is left behind. 
 By reduction with hydrogen the trioxide is converted into 
 the dioxide, UO 2 ; and when either the dioxide or the tri- 
 oxide is heated in the air, the product obtained is the 
 oxide U 3 O 8 . As will be pointed out below, this is re- 
 garded as a uranium salt of uranic acid. 
 
 Uranous Salts. In the uranous salts, uranium acts as 
 a quadrivalent element, replacing four atoms of hydro- 
 gen, as, for example, in the sulphate, which has the com- 
 position U(SO 4 ) 2 . But few salts of this order are known. 
 
 Uranyl Salts. As already explained, the uranyl salts 
 
 OH 
 
 are derivatives of the hydroxide UO 2 <QTT, and are 
 
 formed by the action of acids, as represented in the 
 equations below : 
 
 uo .<81 + HO >S0 ' = UO,<g>SO, + 2H,0; 
 
 TTO , , HO.NO, TTO .O.NO, 
 U0 5 <OH 
 
 They are derived from the acids by replacing the hydro- 
 gen by uranyl, UO a , which is bivalent. Uranyl nitrate, 
 TJO 2 (NO 3 ) 3 , is easily obtained, as above described, and 
 crystallizes well in lemon-yellow prisms. Uranyl sul- 
 phate, UO 3 (SO 4 ), is formed by treating the nitrate with 
 
URANATES. 677 
 
 sulphuric acid. It combines with ammonium sulphate, 
 forming the salt UO 2 (SO 4 ) + (NH 4 ) 2 SO 4 + 2H 2 O, which 
 is difficultly soluble in water and crystallizes in lemon- 
 yellow prisms. 
 
 Uranates. When a uranyl salt is treated with a soluble 
 hydroxide a precipitate is formed which is a salt of an 
 acid, H 2 U 2 O 7 , which may be called diuranic acid, as in 
 composition it is analogous to dichromic and disulphuric 
 acids. 
 
 Sodium diuranate, Na 2 U 2 O 7 , is a fine yellow powder, and 
 is manufactured and sold under the name uranium yellow, 
 being used as a pigment for coloring glass, etc. Ammo- 
 nium diuranate, (NH 4 ) 2 U 2 O 7 , is also manufactured on the 
 large scale. When it is treated with a solution of am- 
 monium carbonate it dissolves, and from the solution a 
 salt of the composition UO 2 (CO 3 ) + 2(NH 4 ) 2 CO 3 crystal- 
 lizes out. The solubility of ammonium diuranate in am- 
 monium carbonate is utilized in analysis. The oxide of 
 the formula U 3 O 8 above referred to may be a uranous 
 salt of uranic acid as represented by tbe formula 
 
 1 
 
 UO,<r\ 
 
 k 
 uo,<<>j 
 
 Many uranium salts exhibit in solution a beautiful 
 fluorescence. 
 
CHAPTER XXXII. 
 
 ELEMENTS OF FAMILY VII, GROUP A : 
 MANGANESE (Mn, At. Wt. 54.57). 
 
 General. At the close of Chapter XII (which see), 
 which treated of the elements of Family VII, Group 
 B, or the chlorine group, reference was made to man- 
 ganese, and attention was called to the fact that in some 
 respects it resembles chlorine. The resemblance is 
 seen in the formation of an oxide, Mn 2 O 7 , and an acid, 
 HMnO 4 , analogous to perchloric acid. On the other 
 hand, in many of its compounds it plays the part of a 
 base-forming element, and in this capacity it forms two 
 series of compounds, known as the manganous and the 
 manganic compounds. In the former the element ap- 
 pears to be bivalent, and in the latter trivalent. The 
 formulas of some of the principal manganous com- 
 pounds are : 
 
 MnCl 2 , Mn(OH) 2 , MnO, Mn(N0 8 ) 2 , MnSO 4 , MnCO 3 , etc. 
 
 The formulas of some of the principal manganic com- 
 pounds are : 
 
 Mn(OH) 3 , Mn a O 3 , MnO(OH), Mn 2 (SO 4 ) s , KMn(SO 4 ) 3 + 12H 2 O, etc. 
 
 These two series of compounds are analogous in composi- 
 tion to the chromous and chromic compounds, but, while 
 the chromic compounds are more stable than the chro- 
 mous compounds, the manganous compounds are more 
 stable than the manganic compounds. By contact with 
 the air, the manganous are not as a rule converted into 
 the manganic compounds. 
 
 Corresponding to chromic acid, there is a manganic 
 acid, H 2 MnO 4 ; and, further, there is the permanganic 
 acid already mentioned, of the formula HMnO 4 . An anal- 
 ogous compound of chromium, perchromic acid, HCrO 4 , 
 
 (678) 
 
MANGANESE. 679 
 
 is believed to be formed when hydrogen peroxide is 
 added to an aqueous solution of chromic acid, but this is 
 by no means certain. Manganic acid and its salts are 
 very unstable, and are readily converted into perman- 
 ganic acid and the permanganates. On the other hand, 
 perchromic acid, if it exists at all, is spontaneously de- 
 composed, yielding ordinary chromic acid. To sum up, 
 then, both chromium and manganese form four classes 
 of compounds. But, while chromium under ordinary 
 circumstances preferably forms chromic salts and salts 
 of chromic acid, manganese preferably forms manganous 
 salts and salts of permanganic acid. 
 
 Manganese forms a number of oxides corresponding 
 to the formulas MnO, Mn 2 O 3 , Mn 3 O 4 , MnO 2 , and Mn 3 O 7 . 
 Of these probably the one of the composition Mn 3 O 4 
 and perhaps that of the composition Mn 2 O 3 are com- 
 pounds of the others, as will be pointed out further on. 
 
 Forms in which Manganese occurs in Nature. The 
 principal natural compound of manganese is the black 
 oxide or pyrolusite, MnO 2 . Besides this, however, there 
 are several compounds found in nature, the principal 
 ones being, braunite, Mn 2 O 3 , hausmannite, Mn 3 O 4 , man- 
 ganite, Mn 2 O 2 (OH) 2 , and rhodocroisite, which is the car- 
 bonate, MnCO 3 . 
 
 Preparation and Properties. The metal is isolated 
 from its oxides by heating them to a high temperature 
 with charcoal. It looks like cast iron, is brittle and 
 hard, and has the specific gravity 8. It easily becomes 
 oxidized in the air, decomposes warm water, and dis- 
 solves readily in dilute acids. It is used as a con- 
 stituent of some useful alloys, and imparts certain de- 
 sirable properties to iron, as will be pointed out when 
 that metal is taken up. 
 
 Manganous Chloride, MnCl 2 . This chloride is obtained 
 in solution by dissolving any one of the oxides or hy- 
 droxides or the carbonate of manganese in hydrochloric 
 acid with the aid of gentle heat. Its formation in the 
 preparation of chlorine from manganese dioxide and hy- 
 drochloric acid was referred to under Chlorine (which 
 see). When the solution is evaporated to the proper 
 
680 INORGANIC CHEMISTRY. 
 
 concentration, the salt crystallizes out in pink, mono- 
 clinic plates of the composition MnCl 2 -f- 4H 2 O. When 
 the crystallized salt is heated, it decomposes into the 
 oxide and hydrochloric acid, as so many other chlorides 
 do. It forms double chlorides, an example of which is 
 the ammonium compound of the formula MnCl 2 .2NH 4 Cl 
 or (NH 4 ) 2 MnCl 4 . 
 
 When manganic hydroxide, Mn(OH) 3 , is treated in the 
 cold with hydrochloric acid, a deep brown-colored solu- 
 tion is formed, which is believed to contain the trichlo- 
 ride, MnCl 3 . On standing, however, this solution gives 
 off chlorine slowly, and when heated it gives it off rap- 
 idly, and the color changes to pink when only mangan- 
 ous chloride is left in the solution. 
 
 Phenomena similar to those just mentioned are ob- 
 served when manganese dioxide is treated with hydro- 
 chloric acid in the cold, and it is believed that the 
 tetrachloride, MnCl 4 , is contained in the solution. 
 
 While the tetrafluoride itself has not been isolated, a 
 solution is obtained by treating manganese dioxide with 
 concentrated hydrofluoric acid which with potassium 
 fluoride gives a salt of the formula MnF 4 .2KF or 
 K 2 MnF 6 . 
 
 General Remarks Concerning the Oxides. The series of 
 oxides of manganese strongly suggests that of the oxides 
 of lead. Manganese, however, forms one oxide, the 
 heptoxide, Mn 2 O 7 , for which there is no analogue among 
 the compounds of lead. Placing the formulas of the 
 oxides of the two metals side by side, we have the follow- 
 ing table : 
 
 PbO MnO 
 
 Pb 2 3 Mn 2 3 
 
 Pb 3 O 4 Mn 8 O 4 
 
 PbO 2 MnO 2 
 
 Mn 2 7 . 
 
 Just as lead oxide when heated is converted into red- 
 lead, Pb 3 O 4 , so the other oxides of manganese are con- 
 verted into the oxide, Mn 3 O 4 , when heated. This has al- 
 ready been seen in the preparation of oxygen by heating 
 the dioxide. In the same way, the oxide, Mn 2 O 3 , loses 
 
OXIDES OF MANGANESE. 681 
 
 enough oxygen, and the lowest oxide, MnO, takes up 
 enough to form the same product : 
 
 3Mn 2 O 3 = 2Mn 3 O 4 + O ; 
 3MnO + O = Mn 3 O 4 . 
 
 When treated with energetic oxidizing agents in the 
 presence of the alkalies, all the oxides are converted into 
 manganates. On the other hand, if a manganate is re- 
 duced in the presence of an acid the tendency to form 
 inanganous compounds shows itself, and all oxygen pres- 
 ent in excess of that required to form the manganous 
 salt is given off. 
 
 Manganous Oxide, MnO, is formed by reducing one of 
 the higher oxides in a current of hydrogen. 
 
 Manganous Hydroxide, Mn(OH) 2 , is formed as a white 
 precipitate when a soluble hydroxide is added to a solu- 
 tion of a manganous salt. Suspended in the alkali, or 
 in contact with air, it absorbs oxygen, and is converted 
 into hydroxides corresponding to the higher oxides. 
 
 Manganous-manganic Oxide, Mn 3 O 4 , occurs in nature as 
 the mineral hausmannite. It is formed, as already stated, 
 by igniting the other oxides in contact with the air. 
 When heated with dilute nitric acid, it acts like the cor- 
 responding oxide of lead, giving manganous nitrate, and 
 leaving manganese dioxide : 
 
 Mn 3 4 + 4HN0 3 = 2Mn(NO 3 ) 2 + MnO, + 2H 2 O. 
 
 It breaks down in the same way with dilute sulphuric 
 acid. These facts make it appear probable that the ox- 
 ide is the manganous salt of normal manganous acid, 
 Mn(OH) 4 , just as minium or red lead is regarded as the 
 lead salt of normal plumbic acid, Pb(OH) 4 , (p. 651.) This 
 
 fg>Mn 
 
 view is expressed by the formula Mn i JJ . The de- 
 
 [ >Mn 
 
 composition with acids, as with sulphuric acid, would, 
 according to this, be represented thus : 
 
 Mn0 4 Mn a + 2H 2 SO 4 = 2MnSO 4 + Mn(OH) 4 . 
 The. hydroxide thus formed would then break down into 
 the dioxide and water. 
 
682 INORGANIC CHEMISTRY. 
 
 Manganic Oxide, Mn 2 O 3 , occurs in nature as the mm- 
 eral braunite, and it can be made from the other oxides 
 by igniting them in oxygen. A hydroxide related to 
 this, and having the composition MnO.OH, analogous to 
 the compounds of aluminium and chromium of the 
 formulas A1O.OH and CrO.OH, is found in nature, and 
 is known as manganite. The hydroxide, Mn(OH) 3 , is 
 formed when manganous hydroxide, Mn(OH) 2 , is ex- 
 posed in a solution of ammonia in contact with the air, 
 and forms a brownish black powder. 
 
 Manganese Dioxide, MnO 2 . This important compound 
 occurs in nature in very considerable quantities, and is 
 known as pyrolusite or the black oxide of manganese. 
 It is obtained artificially by gently igniting manganous 
 nitrate. A hydroxide derived from the dioxide is ob- 
 tained by treating a manganous salt in alkaline solution 
 with a soluble hypochlorite or chlorine or bromine. The 
 chief application of the dioxide is in the preparation of 
 chlorine, for which purpose it is used in large quantities. 
 It is also used for making oxygen, and for the purpose of 
 decolorizing glass. In the last process a small quantity 
 is added to the molten glass. This alone would give the 
 glass an amethyst color. Without it the glass would 
 be green. One color counteracts the other, and the glass 
 appears colorless. As regards the action of hydro- 
 chloric acid upon manganese dioxide, it has been sug- 
 gested, upon the basis of experimental investigations, 
 that the first product of the action is a compound of 
 the formula H 3 MnCl e , which is the chlorine compound 
 analogous to the oxygen acid, H a MnO 3 . The action is 
 supposed to take place as represented in the following 
 equation : 
 
 0=Mn=O 6HC1 = ~>Mn=Cl 2HO. 
 
 The suggestion is made, further, that it is this compound, 
 and not manganese tetrachloride, MnCl 4 , which breaks 
 down yielding chlorine, the action taking place thus : 
 
 H-(Cn >MnC1 * = MnC1 * + 2HC1 
 
MANGANITES. 683 
 
 The manganons chloride and some of the chlormangan- 
 ous acid then react, forming a compound which with 
 water undergoes decomposition. In regard to this sug- 
 gestion, it can only be said that as yet it is not sufficiently 
 supported by facts. The formation of the unstable 
 compound, H 2 MnCl 6 , appears highly probable, however, 
 in view of the conduct of so many other chlorides in 
 the presence of hydrochloric acid. 
 
 Manganites. There are some salts known as the man- 
 ganites, which are clearly derived from hydroxides re- 
 lated to manganese dioxide. Theoretically the simplest 
 hydroxides of this kind are those of the formulas 
 Mn(OH) 4 and MnO(OH) 2 . The salts are not, however, 
 derived from these, but from more complicated forms, as 
 H 2 Mn 2 O 6 and H 2 Mn 5 O u , the relations between which and 
 the simpler hydroxides are shown in the equations 
 
 2MnO(OH) 2 = Mn 2 O 3 (OH) 2 + H 2 O ; 
 5MnO(OH) 2 = Mn 5 O 9 (OH) 2 + 4H 2 O. 
 
 The potassium salt, K 2 Mn 5 O n , is obtained when carbon 
 dioxide is conducted into a solution of potassium man- 
 ganate. Further, a salt of the composition KH 3 Mn 4 O JO 
 is formed as a brown insoluble powder by boiling the 
 other manganites with potassium hydroxide or car- 
 bonate. 
 
 Weldon's Process for the ^Regeneration of Manganese 
 Dioxide in the Preparation of Chlorine. Under the head 
 of Chlorine (which see), Weldon's process was referred 
 to ; but as a satisfactory explanation of the working of the 
 process could not be given without dealing with some 
 rather complicated compounds of manganese, a fuller 
 account was postponed until these compounds should 
 be taken up. The object in view is to utilize the waste 
 liquors from the chlorine factories. When manganese 
 dioxide is treated with hydrochloric acid, as we have seen, 
 manganous chloride and chlorine are formed, according 
 to the equation 
 
 Mn0 2 + 4HC1 = MnCl 2 + C1 2 + 2H 2 O. 
 
684 INORGANIC CHEMISTRY. 
 
 The manganous chloride was of no special value until it 
 was shown that by a comparatively simple method it can 
 be converted into a compound which with hydrochloric 
 acid gives chlorine. When it is treated in solution with 
 lime the corresponding hydroxide is precipitated : 
 
 MnCl, + Ca(OH) 2 = Mn(OH) a + CaCl a ; 
 
 and when this hydroxide mixed with lime is allowed to 
 stand exposed to the air oxidation takes place, and a 
 compound CaMn0 3 or CaMn 2 O 6 is formed : 
 
 Mn(OH) 2 + Ca(OH) 2 + O = CaMnO 3 + 2H 2 O ; 
 2Mn(OH) 2 + Ca(OH) 2 + 2O = CaMn 2 O 5 + 3H 2 O. 
 
 These -compounds give chlorine when treated with hydro- 
 chloric acid. They may indeed be regarded as consisting 
 of lime and manganese dioxide, CaO. MnO 2 and CaO.2MnO 2 , 
 and the action of hydrochloric acid takes place thus : 
 
 CaO.Mn0 2 + 6HC1 =CaCl 2 + MnCl 2 +3H 2 O + C1 2 ; 
 Ca0.2Mn0 2 + 10HC1 = CaCl 2 + 2MnCl 2 + 5H 2 O + 2C1 3 . 
 
 In practice, the waste liquor is mixed with calcium 
 carbonate in order to neutralize the acid. After settling, 
 lime enough is added to precipitate the manganese as 
 hydroxide, and to form with this a mixture in molecular 
 proportions. Into this mixture steam and air are passed, 
 when the oxidation referred to takes place, and calcium 
 nianganite is formed. 
 
 Sulphides. When a solution of a manganous salt is 
 treated with ammonium sulphide, a flesh-colored pre- 
 cipitate, which is thought to be the hydrosulphide, is 
 formed. When this is exposed to the air it turns dark 
 in consequence of oxidation ; and if allowed to stand in 
 the liquid, if this is concentrated, it turns green and be- 
 comes crystalline. The product thus formed is man- 
 ganous sulphide, MnS. This also occurs in nature as 
 alabandite. A disulphide, MnS 2 , corresponding to 
 the dioxide is also found in nature, and is known as 
 hauerite. 
 
VARIOUS COMPOUNDS OF MANGANESE. 685 
 
 Manganous Cyanide, Mn(C!N") 2 , in combination with, po- 
 tassium or sodium cyanide as the compounds Mn(CN) 2 . 
 4KCN or K 4 Mn(CN) 6 , and Mn(CN) 2 .4NaCN or Na 4 Mn(CN) 6 , 
 is formed by treating solutions of manganous salts with 
 potassium or sodium cyanides. When exposed to the 
 air, or when the solutions are boiled, salts of the formulas 
 Mn(CN) 3 .3KCN or K 3 Mn(CN) 6 and Mn(CN) 3 .3NaCN or 
 Na 3 Mn(CN) 6 are formed. 
 
 Manganous Carbonate, MnCO 3 , is found in nature, and 
 is precipitated when a solution of a manganous salt is 
 treated with a soluble carbonate. 
 
 Manganous Sulphate, MnSO 4 , is formed by heating the 
 oxides of manganese with concentrated sulphuric acid. 
 If higher oxides than manganous oxide are used, oxygen 
 is given off : 
 
 MnO +H 2 S0 4 =MnS0 4 + H 2 O ; 
 
 Mn A + 2H 2 S0 4 = 2MnS0 4 + 2H 2 O + O ; 
 
 Mn0 2 + H 2 S0 4 = MnSO 4 + H 2 O + O. 
 
 It crystallizes at low temperatures with seven molecules 
 of water, and at ordinary temperatures with five, in this 
 respect resembling cupric sulphate (which see). * The 
 salt of the formula MnSO 4 + 7H 2 O forms bright-red 
 monoclinic prisms ; while that of the formula MnSO 4 -f- 
 5H 2 O forms pink triclinic crystals. Between 20 and 30 
 it forms monoclinic prisms with four molecules of water. 
 
 Manganic Sulphate, Mn 2 (SO 4 ) 3 , is formed when the oxide 
 Mn 3 O 4 , or the finely divided precipitated dioxide, MnO 2 , 
 is treated with sulphuric acid at not too high a tempera- 
 ture. It forms a dark green amorphous powder, which 
 is easily decomposed by heat and by water. With the 
 sulphates of the alkali metals it forms salts analogous 
 to the alums, as KMn(SO 4 ) 2 + 12H 2 O and NH 4 Mn(SO 4 ) 2 + 
 12H 2 O, in which manganese takes the place of aluminium. 
 This fact makes it appear probable that in the manganic 
 compounds manganese is trivalent, as aluminium prob- 
 ably is in its compounds. 
 
 Manganic Acid and the Manganates. When an oxide 
 of manganese is treated with an energetic oxidizing agent 
 in the presence of a strong base it is converted into a 
 
686 INORGANIC CHEMISTRY. 
 
 manganate, just as the oxides of chromium are converted 
 into chromates and the compounds of sulphur into sul- 
 phates. These three classes of compounds are analo- 
 gous as far as the composition is concerned, as shown by 
 the formulas 
 
 M 2 MnO 4 M 2 CrO 4 M 2 SO 4 
 
 Manganate Chromate Sulphate 
 
 The manganates are, however, quite unstable except in 
 alkaline solution, and when they decompose they form 
 the permanganates. Potassium manganate, K 2 MnO 4 , is 
 formed by fusing manganese dioxide with potassium hy- 
 droxide, when, if the air is not in contact with the mass, 
 the reaction takes place as represented in the equation 
 
 3MnO 3 + 2KOH = K 2 MnO 4 + Mn 2 O 3 + H,O. 
 
 It is also made by fusing the dioxide with potassium hy- 
 droxide and potassium chlorate, when this reaction takes- 
 place : 
 
 3Mn0 2 + 6KOH + KC1O 3 = 3K 2 MnO 4 + KG1 + 3H,(X 
 
 "When the mass obtained in either way is treated with 
 water a dark-green solution of the manganate is formed, 
 and by allowing this to evaporate at the ordinary tem- 
 perature in a partial vacuum, or in an atmosphere free 
 of oxygen, the salt is obtained in small crystals, which 
 are almost black. When a solution of a manganate i& 
 treated with an acid, the manganic acid is at once decom- 
 posed into permanganic acid and manganese dioxide : 
 
 3H 2 MnO 4 = 2HMn0 4 + MnO 2 + 2H 2 O. 
 
 The change of a manganate to a permanganate is ef- 
 fected simply by passing carbon dioxide into the solu- 
 tion, or by boiling or allowing the solution to stand in 
 the air. The change by means of carbon dioxide is rep- 
 resented by the equation 
 
 3K a MnO 4 + 2C0 3 = 2K a C0 3 + MnO 2 + 2KMnO 4 . 
 
POTASSIUM PERMANGANATE. 687 
 
 With water the change takes place thus : 
 
 3K 2 MnO 4 + 2H a O = 2KMnO 4 + MnO a + 4KOH. 
 
 The potassium hydroxide and the manganese dioxide 
 react upon each other to form a manganite of more or 
 less complicated composition. While the manganates 
 are decomposed by acids, forming permanganates, the 
 latter are decomposed by alkalies, forming manganates. 
 Thus, when a solution of potassium permanganate is 
 boiled with potassium hydroxide the color changes to 
 green, owing to the formation of the manganate : 
 
 2KOH + 2KMnO 4 = 2K 2 MnO 4 + H 2 O + O. 
 
 This change takes place readily in the presence of sub- 
 stances which have the power to take up oxygen ; but if 
 such substances are present the reduction goes further, 
 forming finally a manganite which is a derivative of the 
 hydroxide, MnO(OH) 2 . 
 
 Permanganic Acid and the Permanganates. The sim- 
 plest method of obtaining the permanganates is by de- 
 composition of the manganates, as described in the last 
 paragraph. 
 
 Potassium Permanganate, KMnO 4 , is manufactured on 
 the large scale by oxidizing manganese dioxide in the 
 presence of a base. Sometimes the oxidation is effected 
 by the oxygen of the air ; sometimes by the action of an 
 oxidizing agent, as potassium chlorate or nitrate. The 
 fundamental reaction in each case is that represented by 
 the equation 
 
 MnO 2 + 2KOH + O = K 2 MnO 4 + H 2 O. 
 
 As will be observed, it is a reaction of the same kind as 
 that involved in the conversion of a sulphite into a sul- 
 phate. Probably the first action of the hydroxide upon 
 the dioxide consists in the formation of the manganite, 
 K 2 MnO 3 , and this is then oxidized to the manganate. 
 When the solution of the manganate is treated with sul- 
 phuric acid a change similar to those referred to above 
 takes place, and the permanganate is formed. The salt 
 
688 INORGANIC CHEMISTRY. 
 
 is easily soluble in water, and is deposited from its solu- 
 tion in crystals, isomorphous with potassium perchlorate, 
 which appear nearly black, with a greenish lustre. Its- 
 solution in water has a purple or reddish-purple color, 
 according to the concentration. Yery concentrated solu- 
 tions appear almost black. The salt is used extensively 
 in the laboratory and in the arts as an oxidizing agent. 
 Its action will be readily understood from what has- 
 already been said in regard to the conduct of manganese 
 towards acids and towards alkalies. When the perman- 
 ganate undergoes decomposition in the presence of an 
 acid the manganese tends to form a manganous salt, and 
 all the oxygen present in excess of what is needed for 
 this purpose is given off. Thus the decomposition with 
 sulphuric acid takes place as represented in the equation 
 
 2KMn0 4 + 3H 2 S0 4 = 2MnSO 4 + K 2 SO 4 + 3H 2 O + 5O. 
 
 Therefore, when potassium permanganate is used as an 
 oxidizing agent in acid solution, two molecules of the 
 salt KMnO 4 give five atoms of oxygen. On the other 
 hand, when the action takes place in alkaline solution 
 the action reaches its limit in a manganite, which, for 
 purposes of calculation, may be regarded as having the 
 composition K 2 MnO 3 . The first change is from the per- 
 manganate to the manganate as represented in the equa- 
 tion 
 
 2KMn0 4 + 2KOH = 2K 2 MnO 4 + H 2 O + O, 
 and then the manganate loses another atom of oxygen, 
 2K 2 MnO 4 = 2K 2 MnO 3 + 2O. 
 
 Therefore, when the permanganate is used as an oxid- 
 izing agent in alkaline solution, two molecules of the 
 salt yield three atoms of oxygen. 
 
 The permanganates and maganates are valuable dis- 
 infecting agents, and the sodium salts are extensively 
 used for this purpose, under the name of Candy's liquid. 
 
 When a solution of barium permanganate is treated 
 with sulphuric acid, free permanganic acid is obtained 
 
REACTIONS OF MANGANESE COMPOUNDS. 689 
 
 in solution. It is extremely unstable, and decomposes 
 spontaneously when the solution is exposed to the light 
 or is heated. When dry potassium permanganate is 
 added to concentrated sulphuric acid, oily drops sepa- 
 rate and collect upon the bottom of the vessel. These 
 are manganese heptoxide, Mn 2 O 7 , which is formed thus : 
 
 2KMnO 4 + H 2 SO 4 = K 2 SO 4 + Mn 2 O 7 + H 2 O. 
 
 The compound bears to permanganic acid the relation 
 of an anhydride : 
 
 2HMnO 4 = Mn 2 O 7 + H,O. 
 
 It is extremely unstable, giving off oxygen with great 
 ease, and therefore acting as a powerful oxidizing agent. 
 As regards the constitution of the manganates and 
 permanganates, they are respectively regarded as anal- 
 ogous to the sulphates and perchlorates. Accord- 
 ingly manganic acid is represented by the formula 
 O 
 
 HO-Mn-OH or MnO 2 (OH) 2 , while permanganic acid is 
 
 6 
 
 o 
 
 represented by the formula O=Mn-OH or MnO 3 (OH). 
 
 b 
 
 Reactions which are of Special Value in Chemical 
 Analysis. The conduct of manganous salts towards 
 soluble hydroxides and towards soluble carbonates has 
 been described. The hydroxide is soluble in ammonia 
 and ammonium salts, but this solution turns brown when 
 exposed to the air and the manganese is gradually pre- 
 cipitated as the hydroxide Mn(OH) 3 . 
 
 The conduct towards ammonium sulphide has been 
 described. When oxidizing agents like hypochlorites, 
 chlorine, or bromine act upon manganous salts in solu- 
 tions in presence of soluble hydroxides, hydroxides cor- 
 responding to the dioxide MnO 2 , such as Mn(OH) 4 , 
 MnO(OH) 2 , are precipitated. Instead of the above- 
 
690 INORGANIC CHEMISTRY. 
 
 mentioned oxidizing agents, potassium permanganate may 
 be used. 
 
 The action of potassium permanganate and manganate 
 as oxidizing agents when used in alkaline and in acid 
 solutions has been described above. 
 
 Manganese is easily detected by heating the substance 
 under examination with nitric acid and lead peroxide, 
 when permanganic acid will be formed if manganese is 
 present, and its formation will be shown by the purple 
 color of the solution. 
 
 "With microcosmic salt and borax manganese gives an 
 amethyst-colored bead in the oxidizing flame, which 
 becomes colorless in the reducing flame. 
 
CHAPTER XXXIII. 
 
 ELEMENTS OF FAMILY 'VIII, SUB-GROUP A: 
 IRON COBALT NICKEL. 
 
 General. The three elements which form this group 
 are in many respects very similar, and their atomic 
 weights differ but little from one another. That of iron 
 (55.6) is nearly the same as that of manganese (54.57), 
 while cobalt and nickel have nearly the same atomic 
 weight. There is much in iron which suggests manganese. 
 It forms two series of compounds, the ferrous and ferric 
 compounds, which are analogous to the manganous and 
 manganic compounds. In the first series iron appears to 
 be bivalent, as shown in the formulas 
 
 FeCl a , Fe(OH) 2 , FeO, FeS, FeSO 4 , FeCO 3 , etc. 
 
 In the second series it appears to be trivalent, as indi- 
 cated in the formulas 
 
 FeCl s , Fe(OH) 3) Fe,0 s , Fe(NO,) s , Fe,(SO 4 ) a) etc. 
 
 Like chromium and manganese it also forms an acid 
 known as ferric acid, H 2 FeO 4 , which in composition is 
 analogous to chromic and manganic acids. The soluble 
 salts of this acid are, however, unstable, and on decom- 
 posing yield ferric hydroxide. Oxidizing agents readily 
 convert ferrous compounds into ferric compounds, and 
 reducing agents reconvert the latter into the former. 
 When exposed to the air most ferrous compounds are 
 oxidized to ferric compounds. The ferrous compounds 
 in which iron is bivalent are similar to the compounds 
 of the zinc group. The ferric compounds, however, in 
 which the iron is trivalent, are similar to the aluminium 
 compounds ; and in ferric acid it exhibits a resemblance 
 to chromium. Cobalt and nickel resemble iron in re- 
 
 (691) 
 
692 INORGANIC CHEMISTRY. 
 
 spect to their power to form two series of compounds 
 corresponding to the ferrous and ferric compounds. 
 Both elements preferably form compounds of the lower 
 series, examples of which are represented by the for- 
 mulas 
 
 CoCl a Co(OH) 2 CoO Co(NO 3 ) 2 CoSO 4 etc. 
 MCI, Ni(OH) 2 MO Ni(N0 3 ) a NiSO 4 etc. 
 
 Cobalt forms a few compounds corresponding to the 
 ferric series ; and nickel forms a hydroxide, of the for- 
 mula Ni(OH) 3 . While the power of cobalt to form com- 
 pounds in which it is trivalent is much weaker than that 
 of iron, it is stronger than that of nickel, the latter being 
 almost exclusively bivalent. In general terms, it may be 
 said that manganese forms a greater variety of compounds 
 than any other element except carbon. In the manganous 
 compounds it exhibits analogies with zinc, copper, and 
 some other bivalent elements; in the manganic compounds 
 it exhibits analogies with aluminium ; in manganic acid it 
 suggests sulphur and chromium ; and in permanganic 
 acid it suggests chlorine. In the following table some 
 of the analogies which are plainly discernible between 
 the elements mentioned, and iron, cobalt, and nickel, are 
 indicated. The formulas of those compounds which are 
 not easily obtained, and are exceptional, are put in 
 brackets : 
 
 MnSO 4 [Mn 2 (SO 4 ) 3 ] MnO a K 2 MnO 4 KMnO 4 
 
 [CrS0 4 ] Cr a (S0 4 )a CrO 3 K 2 CrO 4 [HCrO 4 ](?) 
 
 FeSO 4 Fe 2 (SO 4 ) 8 [K 2 FeO 4 ] 
 
 CoSO 4 Co(OH) 3 
 NiS0 4 [Ni(OH) 3 ] 
 A1 2 (S0 4 ) 3 
 ZnSO 4 
 
 SO 2 SO 3 K 2 SO 4 
 
 KC10 4 
 
 As regards the question whether the formula of the 
 simpler ferric compounds is to be written with two atoms 
 of iron in every case, it is in much the same state as the 
 question in regard to aluminic compounds. Is ferric 
 chloride FeCl 3 or Fe 2 Cl 6 ? A determination of the 
 
VALENCE OF IRON. 693 
 
 specific gravity of the vapor gave a result in accordance 
 with the larger formula, but this would not appear to be 
 sufficient evidence in view of the peculiar results ob- 
 tained with aluminium chloride. Considering the close 
 resemblance between ferric compounds and the com- 
 pounds of aluminium, it seems probable that, if alu- 
 minium is trivalent, iron is also trivalent in these com- 
 pounds. The ease with which ferric chloride forms com- 
 pounds with other chlorides suggests, further, that the 
 compound of the simpler formula FeCl 8 may combine 
 with another molecule of the same kind to form a double 
 
 /(Cl,)\ 
 chloride of the formula Fe,Cl. or Fef-(Cl,HFe. It may 
 
 be objected to this that it is not probable that such a 
 compound could be converted into vapor without under- 
 going decomposition, and, according to the one determi- 
 nation of the specific gravity, it does not appear to 
 undergo decomposition. What value to attach to this 
 objection it is impossible to say at present. In any case, 
 a further knowledge of the facts is needed before a final 
 conclusion can be reached. In the mean time it seems 
 to be justifiable to consider iron trivalent in ferric com- 
 pounds, as aluminium is considered trivalent in its com- 
 pounds, chromium in chromic compounds, and manga- 
 nese in manganic compounds. 
 
 That iron is bivalent in ferrous compounds is probable 
 from the analogy of these compounds with the distinctly 
 bivalent metals, like copper, zinc, etc. Further, a deter- 
 mination of the specific gravity of the vapor of ferrous 
 chloride gave a figure which indicated that the vapor 
 consisted of about an equal number of molecules of the 
 formulas FeCl 2 and Fe 2 Cl 4 , so that it appears that at a 
 lower temperature the compound has the formula Fe 2 Cl 4 , 
 and that the compound breaks down or dissociates, form- 
 ing the simpler compound. This subject requires further 
 investigation. In the mean time the simpler formula will 
 be used, as it probably represents the chemical molecule 
 or that smallest particle of the compound which comes 
 into play in chemical reactions. If iron is bivalent in 
 
694 INORGANIC CHEMISTRT. 
 
 ferrous compounds, then in all probability cobalt and 
 nickel are bivalent in their principal compounds. 
 
 IRON, Fe (At. Wt. 55.6). 
 
 Introductory. The importance of this metal to man- 
 kind can hardly be overestimated, and for many cen- 
 turies it has played a commanding part in the industries. 
 It requires little thought to convince one that without it 
 the earth would be quite a different place from what it 
 now is. In the earliest periods of history metals were 
 but little used, as but few of them are furnished ready 
 for use by nature. Stones were therefore first used, and 
 these were shaped into a variety of implements, many of 
 which still exist, and furnish evidence of the Stone Age. 
 After a time copper and tin were used in the form of an 
 alloy or bronze, as copper is found in nature in the free 
 condition. During this period, known as the Bronze Age, 
 stone implements gave way to those made of bronze. 
 Afterwards men learned to extract iron from its ores, 
 and the Iron Age was introduced ; and this has continued 
 up to the present, as nothing has since been found which 
 can advantageously take the place of iron. The sugges- 
 tion has been made that as it is less difficult to extract 
 iron from its ores than to make bronze, possibly iron 
 was used as early as bronze perhaps earlier, but that, 
 owing to the fact that iron easily rusts, implements of 
 this metal have disappeared, while those made of bronze 
 remain intact. 
 
 Forms in which. Iron occurs in Nature. Iron occurs in 
 small quantity native in meteorites, in the basalts of Bo- 
 hemia and Greenland, and in some gabbros. The iron 
 meteorites always contain nickel, and frequently small 
 quantities of other elements, as manganese and carbon. 
 Compounds of iron occur in enormous quantities, and 
 widely distributed in the earth. Among the more im- 
 portant are the following-named : hematite, Fe 2 O 3 ; mag- 
 netite, Fe 3 O 4 ; brown iron ore, Fe 4 O 3 (OH) 6 ; siderite, or 
 the carbonate, FeCO 3 ; pyrite, FeS 2 ; pyrrhotite, Fe 7 S 8 . 
 It is also contained in many silicates in small quan- 
 tity, and in consequence of the disintegration of the 
 
METALLURGY OF IRON. 695 
 
 constituents of rocks it is found in the soil, and in many 
 natural waters. In the vegetable kingdom it is always 
 found in chlorophyll, and in the animal kingdom always 
 in the blood. The compounds which are chiefly used 
 for the purpose of making iron, or the iron ores, are 
 magnetite, Fe 3 O 4 ; hematite, Fe 2 O 3 ; brown iron ore, 
 Fe 4 O 3 (OH) 6 ; and spathic iron, or siderite, FeCO 3 . 
 
 Metallurgy. The ores of iron, after they are broken 
 up, are first roasted, in order to drive off water from the 
 hydroxides ; to decompose carbonates ; to oxidize sul- 
 phides ; and, as far as possible, to convert the oxides 
 into ferric oxide, Fe 2 O 3 , which is the most easily re- 
 ducible of the oxides of iron. After the ores are pre- 
 pared in this way they are reduced by heating them with 
 carbon and fluxes in the blast- 
 furnaces, when the iron collects 
 in the molten condition under 
 the so-called slag at the bottom 
 of the furnace. Blast-furnaces 
 differ somewhat in construction, 
 but the essential parts are rep- 
 resented in Fig. 14. 
 
 The inner cavity of the furnace 
 is narrow at the top and bot- 
 tom, as is shown in the fig- 
 ure. Through pipes, known as 
 tuyeres, such as that represented 
 at the lower part of the left- 
 hand side of the figure, air is 
 blown into the furnace to facili- 
 tate the combustion. In modern 
 furnaces arrangements are made 
 above for carrying off the gases Fl - . 
 
 and utilizing them as fuel. The inner walls are built 
 of fire-bricks, and these are surrounded by ordinary 
 bricks, or stone-work. The furnaces vary in height 
 from 25 to 80 or 90 feet, an average height being about 
 45 feet. The reduction of the ores is accomplished bv 
 placing in the furnace alternating layers of coke or 
 charcoal, and the ores mixed with proper fluxes. The 
 
696 INORGANIC CHEMISTRY. 
 
 nature of the flux depends upon the ore. If this con- 
 tains silicon dioxide or clay, lime is added ; while, 
 if it contains considerable lime, minerals rich in silicic 
 acid are used, such as feldspar, clay-slate, etc. The 
 object of the flux is to form a slag in which the re- 
 duced iron collects, and by which it is protected from 
 oxidation. When the fire is once started in a blast- 
 furnace the operation of reduction is continuous until 
 the furnace is burned out. Alternate layers of ore and 
 flux and carbon are added, and, as the reduced iron col- 
 lects below, it is from time to time drawn off and allowed 
 to solidify in moulds of sand. The operation requires 
 close attention. The ores must be carefully studied, and 
 the nature and amount of flux regulated according 
 to the character of the ore as above stated. Then, 
 too, the temperature of the furnace is a matter of im- 
 portance, and must be watched, and regulated by means 
 of the blast. The reduction is largely accomplished 
 by carbon monoxide. In the lower part of the furnace 
 the fuel burns to carbon dioxide, but this comes in con- 
 tact with hot carbon, and is then reduced to the monox- 
 ide. The hot monoxide in contact with the oxides of 
 iron reduces these, and is itself converted into the diox- 
 ide. A large proportion of the carbon monoxide, how- 
 ever, escapes oxidation, and this is carried off from the 
 top of the furnace to the bottom by properly arranged 
 pipes, and is then utilized as fuel. A furnace lasts from 
 two to twenty years, and sometimes longer. 
 
 Varieties of Iron. The iron obtained as above de- 
 scribed is known as pig-iron or cast-iron. It is very 
 impure, containing carbon, phosphorus, sulphur, silicon, 
 etc. If, when drawn from the furnace, the iron is cooled 
 rapidly, nearly all the carbon contained in it remains in 
 chemical combination, and the iron has a silver-white 
 color. This product is known as white cast-iron. If the 
 iron cools slowly, most of the carbon separates as graph- 
 ite, and this being distributed through the mass gives it 
 a gray color. This product is known as gray cast-iron. 
 If the ore contains considerable manganese, this is re- 
 duced with the iron, and iron made from such ores and 
 
VARIETIES OF IRON. 697 
 
 containing manganese has the power to take up more 
 carbon than ordinary iron. This product, containing 
 from 3.5 to 6 per cent combined carbon, is known as 
 spiegel-iron. 
 
 All varieties of cast-iron are brittle, and easily fusible. 
 The gray iron fuses at a lower temperature than the 
 white, and is not as brittle ; it is therefore well adapted to 
 making castings. When cast-iron is treated with hydro- 
 chloric acid the carbon which is present in combined 
 form is given off in combination with hydrogen as hy- 
 drocarbons, some of which have a disagreeable odor. 
 This is, of course, the cause of the bad odor noticed in 
 dissolving ordinary cast-iron in acids. The uncombined 
 or graphitic carbon, on the other hand, remains undis- 
 solved. Owing to its brittleness, cast-iron cannot be 
 welded. When the carbon, silicon, and phosphorus 
 are removed the iron becomes tough and malleable, and 
 its melting-point is much raised. The product thus ob- 
 tained is known as wrought-iron. 
 
 Puddling. Wrought-iron is obtained from cast-iron by 
 the puddling process. The puddling furnace has a flat, 
 oval hearth, and low arched roof. The sides of the 
 hearth are lined with a layer of iron ore (oxide). Coal 
 is burned on a grate and the flame passes into the fur- 
 nace at one end and out at the other, thus coming in 
 contact with the roof and the charge of iron. By con- 
 tact with the flame, and by the heat radiated from the 
 roof, the cast-iron melts. The carbon and silicon are 
 removed from the molten cast-iron, partly by the oxy- 
 gen in the air or flame, but principally by the oxygen 
 in the iron ore, which is itself thus reduced to wrought- 
 iron. 
 
 Wrought-iron contains less than 0.6 per cent of car- 
 bon, and, as the percentage of carbon decreases, the. 
 malleability increases and the melting-point rises. The 
 melting-point of good wrought-iron is from 1900 to 
 2100. Small quantities of sulphur, phosphorus, silicon, 
 and manganese exert a very marked influence upoji its 
 properties. The process of welding consists in heating 
 
698 INORGANIC CHEMISTRY. 
 
 two pieces of iron to a high temperature, putting some 
 borax upon one of them, laying them together, and ham- 
 mering, when, as is well known, they adhere firmly to- 
 gether. The object of the borax is to keep the surfaces 
 bright, which it does by uniting with the oxide and form- 
 ing an easily fusible borate. 
 
 Bessemer Process. Molten cast-iron is poured into a 
 large vessel called the converter. The carbon and sili- 
 con are entirely oxidized and removed by means of a 
 blast of air forced through the metal from below. No 
 fuel is used, as the heat generated by the oxidation of 
 carbon and silicon is sufficient to raise the temperature 
 above 2100. The converter contains molten wrought- 
 iron after the oxidation. By addition of spiegel-iron & 
 product containing any desired percentage of carbon is 
 obtained. 
 
 Iron which contains more than a very small percent- 
 age of phosphorus is not adapted to the manufacture of 
 Bessemer steel in the ordinary way ; but it has been 
 found that, if the converters are lined with lime and 
 magnesia, such iron may be used. Under these circum- 
 stances the phosphorus is oxidized, and forms calcium 
 and magnesium phosphates, which are of value as fertil- 
 izers (see Calcium Phosphate). This process is known 
 as the Thomas-Gilchrist or the basic-lining process. 
 
 Siemens- Martin Furnace. This is simply a reversible 
 puddling furnace in which gas is used as fuel. The 
 gas is previously heated in a Siemens regenerative fur- 
 nace. 
 
 Steel and Wrought-iron. The product of the puddling 
 furnace is called wrought-iron ; while those formed in 
 the Bessemer process and in the Siemens-Martin fur- 
 nace are called steel. Bessemer steel often contains 
 less than 0.6 per cent of carbon, and Siemens-Martin 
 steel is the purest form of wrought-iron, containing less 
 carbon and silicon than the product of the puddling 
 furnace. 
 
 Tempering. When steel is heated and cooled sud- 
 denly, it is rendered extremely hard and brittle ; and 
 when hardened steel is carefully heated, and allowed to- 
 
PROPERTIES OF IRON. 
 
 cool slowly, it becomes very elastic. This process is 
 called tempering. 
 
 Properties of Iron. Pure iron is almost unknown. Of 
 the commercial varieties, it follows from what has been 
 said that wrought-iron is the purest. That which is 
 used for piano- strings is the purest iron which can be 
 bought ; it contains only about 0.3 per cent of impuri- 
 ties. Pure iron can be made in the laboratory by ignit- 
 ing the oxide or oxalate in a current of hydrogen, and 
 by reducing ferrous chloride in hydrogen. In larger 
 quantity it can be prepared by melting the purest 
 wrought-iron in a lime crucible by means of the oxy hy- 
 drogen flame. The impurities are taken up by the cru- 
 cible, and a regulus of the pure metal is left behind. 
 That made by reduction of the oxide or oxalate is, of 
 course, in finely divided, condition. If in its preparation 
 the temperature is kept as low as possible, the prod- 
 uct takes fire when brought in contact with the air ; 
 while if the temperature is high, the product has not 
 this power. Iron is white, and is one of the hardest 
 metals ; and its melting-point is higher than that of 
 wrought-iron. Pure iron is attracted by the magnet. 
 In contact with a magnet, or when placed in a coil 
 through which an electric current is passing, it becomes 
 a magnet ; but the purer it is the sooner it loses the mag- 
 netic power when removed from the magnet or the coil. 
 Steel, however, retains its magnetism. When heated to 
 a sufficiently high temperature iron burns, and forms 
 the oxide, Fe 3 O 4 . This takes place much more easily in 
 oxygen than in the air. In dry air iron does not under- 
 go change, but in moist air it rusts, or it becomes covered 
 with a layer of oxide and hydroxide, which is formed 
 by the action of the air, carbon dioxide, and water. 
 Water that contains salts in solution facilitates the 
 rusting. Various methods are adopted to protect iron 
 from this change, most of which are, however, purely 
 mechanical. A method which promises valuable results 
 is that invented by Barff, which consists in introducing 
 the iron into water vapor at a temperature of 650, when 
 it becomes covered with a firmly adhering layer of oxide. 
 
700 INORGANIC CHEMISTRY. 
 
 Iron dissolves in acids with evolution of hydrogen, 
 and generally with formation of ferrous salts : 
 
 Fe + 2HCl = FeCl 2 + H 2 ; 
 Fe + H 2 S0 4 = FeS0 4 + H Q . 
 
 When cold nitric acid is used, ferrous nitrate and am- 
 monium nitrate are the products ; if the acid is warmed, 
 ferric nitrate and oxides of nitrogen are formed. When 
 an iron wire which has been carefully polished is intro- 
 duced for an instant into red fuming nitric acid it can 
 afterward be put into ordinary nitric acid without under- 
 going change. It is said to be in the passive state ; and 
 the commonly accepted explanation of the phenomenon 
 is, that the wire is covered with a thin layer of oxide. 
 As, however, the passive condition is lost by contact 
 with an ordinary wire, the explanation does not appear 
 to be adequate. 
 
 Ferrous Chloride, FeOl a . When iron is dissolved in 
 hydrochloric acid without access of air, and the solution 
 evaporated, crystals of the composition FeCl 2 -f- 4H 2 O 
 are obtained. When heated for the purpose of driving 
 off the water, the crystallized compound decomposes. 
 The dry chloride can be obtained by heating iron in a 
 current of dry hydrochloric-acid gas. It is a colorless 
 mass, which deliquesces in the air, is volatile at a high 
 temperature ; and determinations of the specific gravity 
 of its vapor made at very high temperatures have shown 
 that its molecule under these conditions should be 
 represented by the formula FeCl 2 . At lower tempera- 
 tures the molecule appears to be more complex. The 
 evidence on this point is not conclusive. If allowed to 
 stand in contact with the air in hydrochloric-acid solu- 
 tion, it is changed to ferric chloride : 
 
 2FeCl 2 + 2HC1 + O = 2FeCl 3 + H 2 O. 
 
 If hydrochloric acid is not present, a basic chloride is 
 precipitated and ferric chloride is then in the solution : 
 
 4FeCl, + H 2 O + O = 2Fe < + 2FeCl 3 . 
 When treated with oxidizing agents in general, as nitric 
 
FERRIC CHLORIDE. 701 
 
 acid, potassium chlorate, potassium permanganate, etc., 
 it is converted into ferric chloride. 
 
 Ferrous chloride, like most other metallic chlorides, 
 combines with the chlorides of the strongest basic ele- 
 ments, forming double compounds. Those with potas- 
 sium and sodium chlorides have the formulas FeCl a . 
 2KC1 or K 2 FeCl 4 , and FeCl 2 .2NaCl or Na 2 FeCl 4 . It 
 combines also with other chlorides, such as those of 
 mercury and cadmium, forming similar salts. 
 
 A solution of ferrous chloride made by dissolving iron 
 in hydrochloric acid is used in medicine under the name 
 Liquor Ferri chlorati. It contains ten per cent iron. 
 
 Ferric Chloride, FeCl 3 . As stated in the last paragraph, 
 ferrous chloride is readily converted into ferric chloride 
 by oxidation. The simplest way to make a solution of 
 the ferric compound is to dissolve iron in hydrochloric 
 acid and pass chlorine into it to complete saturation. 
 The solution is decomposed by heating, especially if 
 dilute, yielding hydrochloric acid and an insoluble 
 oxychloride. The chloride can be obtained in yellow 
 crystals with six or twelve molecules of water. Like the 
 ferrous compound, it is decomposed into hydrochloric 
 acid and the oxide when heated. Anhydrous ferric 
 chloride is obtained by heating iron wire in dry chlorine. 
 It forms black, lustrous crystalline laminae, is volatile 
 &i a lower temperature than the ferrous compound, and 
 the specific gravity of the vapor is that required by a 
 compound whose molecule corresponds to the formula 
 Fe01 3 . "When treated with nascent hydrogen, ferric 
 chloride is converted into ferrous chloride : 
 
 FeCl s + H = FeCl 2 + HC1. 
 
 It combines with other chlorides, forming double chlo- 
 rides. A solution of ferric chloride is used in medicine 
 under the name Liquor Ferri sesquichlorati. 
 
 Cyanides. The compounds which iron forms with 
 cyanogen are of special interest. The simple com- 
 pounds, ferrous cyanide, Fe(CN) 2 , and ferric cyanide, 
 Fe(CN),, corresponding to the above-mentioned chlorides, 
 are not known : only double compounds of these with 
 
702 INORGANIC CHEMISTRY. 
 
 other cyanides are well known, and some of them are 
 manufactured on the large scale. When a solution of 
 potassium cyanide acts upon metallic iron or the 
 oxides of iron, a solution is formed from which the 
 salt known as potassium ferrocyanide or yelloiv prus- - 
 slate of potash crystallizes. This has the composition 
 K 4 Fe(CN) 6 -f- 3H 2 O, and may be regarded as made up 
 of a molecule of ferrous cyanide and four molecules 
 of potassium cyanide, as represented in the formula 
 Fe(CN) 2 .4KCN + 3H 2 O. When this salt is treated with 
 chlorine it is converted into potassium ferricyanide, or red 
 prussiate of potash, K 3 Fe(CN) 6 , which is to be regarded 
 as consisting of ferric cyanide and potassium cyanide, as 
 represented in the formula Fe(CN) 3 .3KCN. The trans- 
 formation is represented thus : 
 
 K 4 Fe(CN) 6 + 01 = K 3 Fe(CN) 6 + KC1. 
 
 From these two a number of other cyanogen compounds 
 are obtained. When treated in concentrated solution 
 with concentrated hydrochloric acid they yield the free 
 acids, and by treating them with solutions of different 
 metallic salts corresponding salts of these acids are ob- 
 tained. Among the most important of these derivatives 
 are the following : 
 
 Ferrohydrocyanic acid, Ferrihydrocyanic acid, 
 
 H 4 Fe(CN) 6 H 3 Fe(CN) 6 
 
 Potassium ferrocyanide, Potassium ferricyanide, 
 
 K 4 Fe(C]Sr)6 K 3 Fe(CN) 6 
 
 Sodium ferrocyanide, Sodium ferricyanide, 
 
 Na 4 Fe(ClSr) 6 Na 3 Fe(CN) 6 
 
 Barium ferrocyanide, 
 
 Ba 2 Fe(CN) 
 
 Ferric ferrocyanide, Ferrous ferricyanide, 
 
 Fe 4 [Fe(CN) 6 ] 3 Fe 3 [Fe(CN) 6 ] 3 
 
 Ferri- potassium ferrocyanide, 
 
 KFeFe(CN) 8 
 
 Potassium Ferrocyanide, K 4 Pe(CN) 6 + 3H 2 O. As stated 
 above, this salt can be made by treating iron or the 
 oxides of iron with a solution of potassium cyanide. On 
 the large scale it is manufactured by melting crude 
 potash or potassium carbonate, and gradually adding a 
 mixture of iron filings or turnings, and refuse animal- 
 
 
CYANIDES OF IRON. 703 
 
 matter, as claws, horns, hoofs, hair, etc. Or the potash is 
 melted with the animal substances and potassium cyanide 
 thus formed, and this treated in solution with ferrous 
 carbonate, when the ferrocyanide is formed. It forms 
 large yellow pyramids belonging to the tetragonal sys- 
 tem. At the ordinary temperature it dissolves in three 
 to four parts of water, and more easily in hot water. It 
 gives up its water of crystallization very easily. When 
 heated it is decomposed, forming potassium cyanide, 
 nitrogen, and a compound of iron and carbon : 
 
 K 4 Fe(CN) 6 = 4KCN + N 2 + FeC 2 . 
 
 Treated with concentrated sulphuric acid it undergoes 
 decomposition, giving as gaseous product carbon mon- 
 oxide, and this furnishes a good method for the prep- 
 aration of the gas : 
 
 K 4 Fe(CN) 6 + 6H 2 S0 4 + 6H 2 O = FeSO 4 + 2K 2 S0 4 
 
 + 3(NH 4 ) 2 S0 4 + 6CO. 
 
 With dilute sulphuric acid it gives hydrocyanic acid, and 
 forms at the same time a white insoluble compound of 
 the composition KFe(CN) 3 or Fe(CN) 2 .KCN : 
 
 2K 4 Fe(CN) 6 + 3H 2 SO 4 = 6HCN + 3K 2 SO 4 + 2KFe(CN) 3 . 
 
 Ferrohydrocyanic Acid, H 4 Fe(CN) 6 , formed as above 
 described, is a white crystalline substance, which is easily 
 soluble in water and alcohol. It takes up oxygen from 
 the air, and is converted into the ferric salt of the acid, 
 hydrocyanic acid being given off. The ferric salt is 
 the substance commonly called insoluble Prussian blue. 
 The relation of the salt to the acid is shown by the 
 formulas 
 
 H.Fe(CN). Fe,[Fe(CN).], 
 
 Ferro-hydrocyanic acid Ferric ferrocyanide, or 
 
 Prussian blue 
 
 Ferric Ferrocyanide, or Prussian Blue, Fe 4 [Fe(C]S")]3. 
 This compound is very readily formed by adding a solu- 
 tion of a ferric salt to a solution of potassium ferrocya- 
 nide, and appears as a dark-blue precipitate : 
 
 3K 4 Fe(CN) 6 + 4FeCl 3 = Fe 4 [Fe(CN)J 3 + 12KC1. 
 
704 INORGANIC CHEMISTRY. 
 
 It is obtained in pure condition by treating a solution of 
 a ferric salt with a solution of ferrohydrocyanic acid. 
 When a ferric salt is added to an excess of potassium 
 ferrocyanide a ferri-potassium salt, KFeFe(CN) 6 , is 
 formed. This is commonly called Prussian blue, and 
 the commercial article always contains some of it. It 
 is also known as soluble Prussian blue. When heated 
 with an alkaline hydroxide, Prussian blue is decom- 
 posed, the products of the action being the ferrocyanides 
 of the alkali metals and ferric hydroxide : 
 
 Fe 4 [Fe(CN) 6 ] 3 + 12KOH = 3K 4 Fe(CN) 6 + 4Fe(OH) 3 . 
 
 Potassium Ferricyanide, K 3 Fe(CN) 6 . This salt is 
 formed by treating the ferrocyanide, either dry or in 
 solution, with chlorine. It forms large, dark-red, mono- 
 clinic prisms. It dissolves in about three times its- 
 weight of water at the ordinary temperature, and is more 
 easily soluble in hot water. In alkaline solution it acts 
 as a strong oxidizing agent, on account of its tendency to 
 form the ferrocyanide. The character of the action is 
 indicated by the following equation : 
 
 6K 3 Fe(CN) 6 + 6KOH = 6K 4 Fe(CN) 6 + 3H 2 O + 3O. 
 
 Ferrihydrocyanic Acid, H 3 Fe(CN) 6 , is a crystallized 
 substance. 
 
 Ferrous Ferricyanide, Fe 3 [Fe(CN) 6 ] 2 , is commonly called 
 Turnbull's blue. It is formed by adding potassium 
 ferricyanide to a solution of ferrous sulphate, or any 
 ferrous salt : 
 
 3FeS0 4 + 2K 3 Fe(CN) 6 = Fe 3 [Fe(CN) 6 ] 2 + 3K 2 SO 4 . 
 
 Starting with ferrocyanic and ferricyanic acids, four 
 iron salts suggest themselves. These are ferrous and 
 ferric ferrocyanide, and ferrous and ferric ferricyanide. 
 The relations between them are indicated in the follow- 
 ing formulas : 
 
 Acid H 4 Fe(CN) 6 Acid, .... H 3 Fe(CN) a 
 
 (1) Ferrous salt, . Fe a Fe(CN)e (3) Ferrous salt, . Fe 3 [Fe(CN) 6 ] 
 
 (2) Ferric salt, . . Fe 4 [Fe(CN) 6 ] 3 (4) Ferric salt, . . FeFe(CN) 8 
 
IRON SALTS-NITROPRUSSIATES. 705 
 
 Of these (2) is Prussian blue and (3) is Turnbull's 
 blue. The commercial Prussian blue contains some 
 Turnbull's blue. The reason of this appears to be that 
 a part of the ferrocyanide of potassium used in the 
 preparation is oxidized by the ferric salt, and thus ferri- 
 cyanide of potassium and a ferrous salt come together. 
 When potassium ferrocyanide is added to a solution of 
 a ferrous salt, we should expect the formation of salt (1) 
 or ferrous ferrocyanide : 
 
 K 4 Fe(CN) 6 + 2FeCl 2 = Fe 2 Fe(CN) 6 + 4KC1. 
 
 But instead ol this a ferro-potassium salt, of the formula 
 
 K 2 FeFe(CN) 6 , is formed :' 
 
 K 4 Fe(CN) 6 + FeCl 2 = K 2 FeFe(CN) 6 + 2KC1. 
 
 This is a white powder, which is formed also when po- 
 tassium ferrocyanide is decomposed by dilute sulphuric 
 acid in the preparation of hydrocyanic acid. It is repre- 
 sented above by the formula KFe(CN) 3 , but taking the 
 method of formation into consideration the formula 
 K 2 FeFe(CN) 6 seems more probable. Ferric ferricyanide 
 is not known. When, however, a solution of a ferric 
 salt is added to one of potassium ferricyanide the solu- 
 tion turns dark brown, and perhaps contains this salt. 
 The composition of the salt is the same as that of ferric 
 cyanide, and possibly the two compounds are identical. 
 
 Nitroprussiates. When potassium ferrocyanide is 
 treated with nitric acid, potassium nitrate is formed. 
 When this is removed and the solution neutralized with 
 sodium carbonate, a salt known as sodium nitroprussiate 
 is obtained. This crystallizes very beautifully, and is 
 used to some extent in the laboratory. With soluble 
 sulphides it gives an intense violet color, but not with hy- 
 drogen sulphide. The composition of the salt is repre- 
 sented by the formula Na 2 Fe(CN) 6 (NO) + 2H 2 O. The 
 free acid corresponding to this salt, and also other salts 
 of the same acid, have been made. 
 
 Ferrous Hydroxide, Fe(OH) 2 , is formed when a soluble 
 hydroxide is added to a solution of a ferrous salt. It is 
 a white precipitate, but it is usually obtained as a green- 
 ish mass, as it is very easily oxidized by the oxygen of 
 
706 INORGANIC CHEMISTRY. 
 
 the air and that contained in the solutions. When al- 
 lowed to stand in contact with the air it turns a dirty 
 green, and finally brown, being converted into ferric hy- 
 droxide. When heated in the air it loses water, and 
 takes up oxygen, forming ferric oxide. 
 
 Ferrous Oxide, FeO, is formed by passing hydrogen 
 over ferric oxide heated to 300. It is a black powder, 
 which takes up oxygen from the air, and is converted into 
 the oxide Fe 2 O 3 . 
 
 Ferric Hydroxide, Fe(OH) 3 . This compound is formed 
 most readily by adding ammonia to a solution of a ferric 
 salt, when it appears as a voluminous brownish-red pre- 
 cipitate. When filtered, washed, and dried, its compo- 
 sition is not changed. If heated at 100, or if the solution 
 is boiled for some time, it loses water, and forms com- 
 pounds of the formulas FeO.OH, Fe 2 O(OH) 4 , etc. The 
 latter is derived from the normal hydroxide as repre- 
 sented in the equation 
 
 2Fe(OH) 3 = Fe 2 0(OH) 4 + H 2 O. 
 
 The mineral pyrosiderite is the hydroxide FeO.OH. 
 Brown iron ore is Fe 4 O 3 (OH) 6 ; and bog iron-ore is 
 Fe 2 O(OH) 4 . All of these are derivatives of the normal 
 hydroxide. The normal hydroxide differs from alumin- 
 ium hydroxide in the fact that it has no acid properties. 
 Therefore, if the two hydroxides are treated together with 
 a caustic alkali only the aluminium hydroxide dissolves. 
 The compound FeO.OH, corresponding to A1O.OH and 
 CrO.OH, yields salts under some circumstances. Thus 
 
 a calcium salt, -p 6 Q | *Q>Ca, is formed by heating together 
 
 ferric oxide and lime to a high temperature. In compo- 
 sition this is plainly analogous to the spinels. Magnetic 
 oxide of iron or magnetite is believed to be the corre- 
 sponding ferrous salt, -p, 6 Q *Q>Fe. Franklinite also is a 
 salt of the same order, containing zinc. It is essentially 
 a zinc salt, of the formula -p 6 Q*Q>Zn, but some of the 
 zinc is replaced by iron and manganese. 
 
OXIDES OF IRON. 707 
 
 Ferrous-Ferric Oxide, Fe 3 O 4 . As stated above, this 
 compound is regarded as analogous to the spinels, and 
 as the ferrous salt of the acidic hydroxide FeO.OH, as 
 represented in the equation (FeO.O) 2 Fe. It is found in 
 nature as the mineral magnetite, and loadstone, which 
 occurs in Sweden, Norway, and elsewhere. It is, further, 
 formed when iron is burned in oxygen, and when water 
 is passed over red-hot iron. Some of the magnetite 
 which occurs in nature has the power to attract iron, or 
 is magnetic. 
 
 Soluble Ferric Hydroxide is formed when a solution of 
 ferric chloride or ferric acetate is treated with ferric hy- 
 droxide, and the solution thus formed dialyzed (see page 
 421). The ferric salts pass through the membrane, and 
 the ferric hydroxide remains in solution in water, form- 
 ing a deep-red liquid. It is used in medicine. Small 
 quantities of salts cause the precipitation of ferric hy- 
 droxide from the solution. 
 
 Ferric Oxide, Fe 2 O 3 , is found in nature, and is known 
 as hematite, forming one of the most valuable ores of 
 iron. It can be made in the laboratory by igniting the 
 hydroxide. As hematite, it is a black, crystallized sub- 
 stance with a high lustre. Otherwise it has a red or a red- 
 dish-brown color. The oxide found in nature and that 
 which has been strongly ignited are very difficultly solu- 
 ble in acids. In the preparation of fuming sulphuric 
 acid by heating ferrous sulphate (see page 219) there is 
 left a residue of ferric oxide known as rouge, which is 
 used as a red pigment and as a polishing powder. A 
 specially fine variety of rouge for polishing is manufac- 
 tured by heating ferrous oxalate, FeC 2 O 4 , in contact with 
 the air. 
 
 Ferrous Sulphide, FeS, is formed by direct union of 
 iron and sulphur when the two are heated together. It 
 is manufactured by heating iron filings and sulphur to- 
 gether in a crucible. The pure compound is yellow and 
 crystalline. When heated in contact with the air it is 
 oxidized to ferrous sulphate, if the temperature is not 
 too high. At a higher temperature the products are 
 sulphur dioxide and ferric oxide. When a solution of a 
 
708 INORGANIC CHEMISTRY. 
 
 ferrous salt is treated with ammonium sulphide, ferrous 
 sulphide is precipitated as a black powder. When a 
 ferric salt is treated with ammonium sulphide it is re- 
 duced to the ferrous condition, and then ferrous sulphide 
 is precipitated : 
 
 Fe 2 (S0 4 ) 3 + (NH 4 ) 2 S = 2FeSO 4 +(NH 4 ) 2 SO 4 + S ; 
 2FeS0 4 + 2(NH 4 ) 2 S = 2FeS + 2(NH 4 ) 3 SO 4 . 
 
 The sulphide thus obtained oxidizes readily in the air, 
 and forms the sulphate. The compact variety is used 
 in making hydrogen sulphide (which see). 
 
 Ferric Sulphide, Pe 2 S 3 , is analogous to ferric oxide, 
 Fe 2 O 3 . It is formed artificially by heating iron and sul- 
 phur together in the proper proportions. 
 
 Just as there are salts derived from the hydroxide 
 FeO.OH, so there are salts which are derived from the 
 corresponding sulphide FeS.SH. The potassium, sodium, 
 and some other salts are obtained artificially. Chalco- 
 pyrite is apparently the cuprous salt SFe-S-Cu or 
 FeCuS 2 . 
 
 Ferrous Carbonate, FeCO 3 . This salt occurs in nature 
 as spathic iron or siderite. It crystallizes in forms 
 similar to those of calc spar or calcium carbonate CaCO 3 . 
 Like this, further, it dissolves in water which contains 
 carbon dioxide, and is therefore contained in natural 
 waters which come in contact with it. When a solution 
 of a ferrous salt is treated with a soluble carbonate a 
 white precipitate is formed, which is ferrous carbonate ; 
 but in contact with the air this is rapidly oxidized and 
 decomposed, leaving ferric hydroxide, which with car- 
 bonic acid does not form a salt. In this respect ferric 
 hydroxide acts like alurninic and chromic hydroxides, 
 and therefore when a soluble carbonate is added to a 
 solution of a ferric salt the hydroxide and not ferric car- 
 bonate is thrown down. 
 
 Ferrous Sulphate, FeSO 4 . This important compound 
 is manufactured on the large scale by the spontaneous 
 oxidation of pyrite in contact with the air, and by dis- 
 solving iron in sulphuric acid. It is frequently called 
 
FERROUS SULPHATE. 709 
 
 " green vitriol" (see p. 596), and more commonly " cop- 
 peras." Under ordinary conditions it crystallizes in 
 transparent, green, monoclinic crystals with seven mole- 
 cules of water, just as zinc sulphate, magnesium sulphate, 
 etc., do ; and when heated, six of these are given off 
 readily, while the last is given off with difficulty a fact 
 which makes it appear probable that the salt is a deriva- 
 tive of tetrahydroxyl- sulphuric acid, as represented in 
 
 ( (OH), 
 
 the formula OS-< O -ci . While it ordinarily crystal- 
 lizes in monoclinic crystals, it takes the rhombic form 
 if its supersaturated solution is touched with a crystal 
 of zinc sulphate. It also crystallizes in the triclinic sys- 
 tem with five molecules of water, like cupric sulphate, if 
 a crystal of the latter salt is placed in its concentrated 
 solution. The salt undergoes change when exposed to 
 the air, being converted into a compound containing 
 ferric sulphate, Fe 2 (SO 4 ) 3 , and ferric hydroxide, or more 
 probably a basic ferric sulphate, Fe 3 (SO 4 ) 3 (OH) 3 . 
 
 6FeS0 4 + 30 + 3H 2 O = 2Fe 3 (SO 4 ) 3 (OH) 3 , or 
 6FeSO 4 + 3O + 3H 2 O = 2Fe 2 (SO 4 ) 3 + 2Fe(OH) 3 . 
 
 The same change takes place when a solution of ferrous 
 sulphate is exposed to the air. When treated with oxid- 
 izing agents in the presence of sulphuric acid it is com- 
 pletely converted into ferric sulphate : 
 
 2FeS0 4 + H 2 S0 4 + O = Fe,(SO 4 ), + H,O. 
 
 Like other soluble ferrous salts it absorbs nitric oxide, 
 and when the solution of the unstable compound is heated 
 the nitric oxide is given off. Ferrous sulphate is used in 
 dyeing, in the manufacture of ink, etc.; and as a deodor- 
 izer. With sulphates of the alkalies ferrous sulphate 
 forms double salts, such as FeK 2 (SO 4 ) 2 + 6H 2 O, Fe(NH 4 ), 
 (SO 4 ) 2 -f- 6H 2 O, etc. These are not as easily oxidized as 
 the simple salt, and are convenient in the laboratory 
 when a pure ferrous salt is wanted. It is a fact worthy 
 of special notice, that while ferrous sulphate crystallizes 
 
710 INORGANIC CHEMISTRY. 
 
 with seven molecules of water, these salts contain only 
 six molecules, and all of this is easily given off when 
 the salts are heated. It appears, therefore, that these 
 double salts are formed from ferrous sulphate by re- 
 placing one molecule of the water by a molecule of some 
 sulphate. This is clear if ferrous sulphate and the other 
 salts are regarded as salts of tetrahydroxyl-sulphuric 
 acid. We should then have the relation between the 
 double sulphate and the simple ones as represented in 
 the formulas below : 
 
 
 H HO 
 H HO 
 
 O 
 
 OK 
 OK 
 
 Ferrous sulphate Potassium sulphate 
 
 Fe<>S<g>S<gf = FeK,(S0 4 ), 
 
 Ferrous-potassium sulphate 
 
 Ferric Sulphate, Fe 2 (SO 4 ) 3 . This salt, as stated in the 
 last paragraph, is formed by oxidation of ferrous sul- 
 phate. It is also formed by dissolving ferric oxide or 
 hydroxide in sulphuric acid. When the solution is 
 evaporated the salt remains behind as a white, anhydrous 
 mass. It readily forms basic salts, the composition of 
 which is not positively known. With the sulphates of 
 the alkali metals it forms double salts, which are per- 
 fectly analogous to alum, and are known as the iron alums ; 
 as, for example, FeK(SO 4 ) 2 + 12H. 2 O, Fe(NH 4 )(SO 4 ) 2 + 
 12H 2 O, etc. 
 
 Ferrous Phosphate, Fe 3 (PO 4 ) 2 , occurs in nature crystal- 
 lized with eight molecules of water as the mineral vivi- 
 anite. Both this salt and ferric phosphate, FePO 4 , are 
 insoluble and are formed when solutions of ferrous and 
 ferric salts are treated with sodium phosphate. 
 
 Ferric Acid, H 2 FeO 4 , is analogous in composition to 
 chromic and manganic acids. The acid itself is not known, 
 but its potassium salt is formed when iron or ferric oxide 
 is heated with saltpeter, or when chlorine is passed into 
 
REACTIONS OF IRON COMPOUNDS. 711 
 
 caustic potash containing ferric hydroxide in suspension : 
 
 Fe(OH) 3 + 2KOH + 3C1 = K 2 FeO 4 + H 2 O + 3HC1. 
 
 This, as well as the other ferrates, is unstable, the iron 
 tending to pass back into the condition of a ferric com- 
 pound. 
 
 Iron Bisulphide, PeS 2 , is not analogous to any oxygen 
 compound of iron. In it the metal appears to be quad- 
 rivalent. The disulphide occurs very widely distributed 
 and in large quantities in nature as the mineral iron 
 pyrites or pyrite, which crystallizes in the regular sys- 
 tem, and as marcasite, which crystallizes in the rhombic 
 system. It can be made artificially, and if crystallized 
 it appears in the form of pyrite. Its conduct under the 
 influence of heat has been repeatedly referred to in con- 
 nection with the roasting of iron and other ores. As 
 pyrite it has a golden-yellow color, and it has frequently 
 been taken for the precious metal by those not familiar 
 with it. The name " fool's gold," by which it is some- 
 times popularly known, suggests this fact. 
 
 Arsenopyrite, FeAsS, occurs in nature, and is a valu- 
 able source of the element arsenic ; for, as has been 
 stated (see p. 305), when it is heated it gives off arsenic, 
 and ferrous sulphide is left behind. 
 
 Iron Carbonyls. When finely-divided iron is alloweS 
 to cool in hydrogen gas to 80, and then treated with car- 
 bon monoxide, the issuing gas imparts a yellow color to 
 the flame of a Bunsen burner, and by passing it through 
 a heated glass tube a metallic mirror is formed at between 
 200 and 350. This mirror consists of iron. These 
 phenomena are due to the formation of volatile com- 
 pounds of iron with carbon monoxide, the principal one 
 having the composition Fe(CO) 6 . (See Nickel Carbonyl.) 
 
 Reactions which are of Special Value in Chemical 
 Analysis. Ferrous Compounds. The reactions of ferrous 
 compounds with the soluble hydroxides and carbonates, 
 ammonium sulphide, potassium ferricyanide, and with ox- 
 idizing agents have been explained above. With ammo- 
 nium salts ferrous chloride forms double salts, which are 
 soluble ; therefore, if ammonium chloride is added to a 
 
712 INORGANIC CHEMISTRY. 
 
 solution of the salt ammonia does not precipitate the 
 hydroxide. Further, ammonia does not completely pre- 
 cipitate the hydroxide from a solution of a ferrous salt, as 
 an ammonium salt is formed. By standing in the air, 
 however, these solutions containing the double salts are 
 oxidized, and ferric hydroxide is precipitated. The re- 
 actions with potassium cyanide will be understood from 
 what has been said concerning the compounds of ferro- 
 hydrocyanic and ferrihydrocyanic acids. 
 
 Ferric Compounds. The reactions of ferric compounds 
 with the soluble hydroxides and carbonates, ammonium sul- 
 phide, potassium ferrocyanide, and potassium ferricyanide 
 have been explained above. When hydrogen sulphide is 
 passed through a solution of a ferric salt, reduction to 
 the corresponding ferrous salt takes place, and sulphur 
 separates, which gives the solution a milky appearance : 
 
 Fe 2 (S0 4 ) 3 + H 2 S = 2FeS0 4 + H 2 SO 4 + S ; 
 2FeCl 3 + H 2 S = 2FeCl 2 + 2HC1 + S. 
 
 When a neutral solution of a ferric salt is treated with 
 suspended barium carbonate the iron is precipitated as the 
 hydroxide. 
 
 When a neutral solution of a ferric salt is treated with 
 acetate of potassium or sodium it turns dark red, in conse- 
 quence of the formation of ferric acetate which remains 
 in solution. When the solution is boiled the acetate 
 breaks down into acetic acid and ferric hydroxide, which 
 is precipitated : 
 
 FeCl 3 + 3NaC 2 H 3 2 = Fe(C 2 H 8 O 2 ) 3 + 3NaCl ; 
 Fe(C 2 H 3 2 ) 3 + 3H 2 = Fe(OH) 3 + 3C 2 H 4 O a . 
 
 When potassium sulphocy.anate, KCNS, is added to a solu- 
 tion of a ferric salt a blood-red color is produced. This 
 occurs even in extremely dilute solutions of ferric salts. 
 
 The borax bead is colored bottle-green in the reducing 
 flame, and brown-red to yellowish red in the oxidizing 
 flame, when treated with compounds of iron. 
 
 COBALT, Co (At. Wt. 58.49). 
 
COBALT. 713 
 
 General. As stated in the remarks introductory to the 
 iron group, cobalt, like nickel, preferably forms com- 
 pounds which are analogous to ferrous compounds. It r 
 however, forms a few which are analogous to ferric com- 
 pounds, its power in this direction being greater than 
 that of nickel. Its salts form a great variety of com- 
 pounds with ammonia, and these have been extensively 
 studied. 
 
 Occurrence and Preparation. Cobalt occurs in nature, 
 almost always in company with nickel. The principal 
 minerals containing it are smaltite, CoAs 2 , and cobaltite, 
 CoS 2 .CoAs 2 . In each of these a part of the cobalt is re- 
 placed by iron, and, generally, some nickel. By roasting 
 and melting the ores in blast-furnaces they are partly 
 purified. The product is dissolved in hydrochloric acid, 
 and treated with a small quantity of calcium hypochlorite 
 and hydroxide for the purpose of removing iron and 
 arsenic ; then with hydrogen sulphide to remove copper 
 and bismuth ; and finally with calcium hypochlorite 
 when cobaltic hydroxide, Co(OH) 3 , is precipitated. This 
 is readily converted into the oxide, Co 2 O 3 , from which 
 the metal can be prepared by heating it in a current of 
 hydrogen. It is also obtained by heating the oxalate to 
 a sufficiently high temperature. 
 
 Properties. Cobalt has a silver-white color, with a 
 slight cast of red. It is harder than iron, and melts at a 
 somewhat lower temperature ; is tenacious ; and has the 
 specific gravity 8.9. It dissolves in nitric acid. 
 
 Cobaltous Chloride, CoCl 2 , is formed by heating cobalt 
 in chlorine gas, and in solution by treating cobalt 
 carbonate with hydrochloric acid. From the solution 
 it crystallizes in dark-red prisms of the composition 
 CoCL, + 6H 2 O. The anhydrous salt is blue. When the 
 blue salt is treated with water it turns red, and when the 
 red salt is heated it turns blue. This difference in color 
 between the anhydrous and the hydrated salts is charac- 
 teristic of cobalt salts. If marks are made on paper 
 with a dilute solution of one of the salts the color is not 
 perceptible. If, however, the paper is held before a 
 fire, the salt loses water and turns blue, and as the blue 
 
714 INORGANIC CHEMISTRY. 
 
 is more intense than the red, it is visible. When the salt 
 becomes moist again it becomes invisible. This is the 
 basis for the preparation of the so-called sympathetic inks. 
 
 Cobaltous Hydroxide, Co(OH) 2 , is formed as a red pre- 
 cipitate when a soluble hydroxide is added to a cobalt- 
 ous salt, and the blue precipitate, which is first formed 
 and which is a basic salt, is allowed to stand. It is oxid- 
 ized by contact with the air, forming cobaltic hydroxide^ 
 which breaks down into cobaltic oxide, Co 2 O 3 . 
 
 Cobaltous Oxide, CoO, is formed when the correspond- 
 ing hydroxide is carefully heated without access of air. 
 "When heated in the air it is converted into cobaltous- 
 cobaltic oxide, Co 3 O 4 , which is analogous to ferrous-ferric 
 oxide. This is also formed when cobaltic oxide, Co 2 O 3 , 
 is heated in the air. 
 
 Cobaltic Hydroxide, Co(OH) 3 , is formed when calcium 
 hypochlorite is added to a solution of a cobaltous salt, 
 and is a black powder. When heated it is converted 
 into black cobaltic oxide, Co 2 O 3 . 
 
 Cobalt Sulphide, CoS, is the black precipitate which is 
 formed by adding ammonium sulphide to a cobaltous 
 salt. It is not soluble in dilute acids, and differs from 
 ferrous sulphide in this respect. Other sulphides of 
 cobalt are those of the formulas Co 3 S 4 and CoS 2 . The 
 former is found in nature, and is known as linnseite. 
 The latter occurs in combination with other sulphides, 
 as in cobaltite, CoS 2 .CoAs 2 . 
 
 Cyanides. Cobaltous cyanide is an insoluble dirty-red 
 compound which is formed when potassium cyanide is 
 added to a solution of a cobalt salt. It dissolves in an 
 excess of potassium cyanide, forming a double cyanide, 
 K 4 Co(CN) 6 , which is analogous to potassium ferrocya- 
 nide. When this solution is boiled the cyanide is oxi- 
 dized, forming a compound analogous to potassium 
 ferricyanide, thus : 
 
 2K 4 Oo(ON). + H 2 O + O = 2K 3 Co(CN) 6 + 2KOH. 
 
 This acts like the corresponding iron compound^ The 
 cobalt is not precipitated from it by ammonium sulphide 
 
COMPOUNDS OF AMMONIA WITH SALTS OF COBALT. 715 
 
 or sodium hydroxide. This conduct towards potassium 
 cyanide distinguishes cobalt from nickel salts. 
 
 Smalt. The beautiful pigment known by this name is 
 essentially a cobalt glass in which cobalt takes the place 
 of calcium. It is made by heating compounds of cobalt 
 with quartz and potassium carbonate. The glass thus 
 formed is powdered very finely and used as a pigment. 
 It does not change color in the sunlight, and is not af- 
 fected by acids nor by alkalies. 
 
 Compounds of Ammonia with Salts of Cobalt. In gen- 
 eral, when solutions of cobalt salts in ammonia are ex- 
 posed to the air they undergo oxidation, and complicated 
 salts are formed. In the case of cobaltous chloride, 
 the first product formed is one of the composition 
 Co(NH 3 ) 3 Cl s + H a O, which is known as dichro-cobaltic 
 chloride. At the same time another compound of the 
 composition Co(NH 3 ) 4 Cl 3 -|- H 2 O, known as praseo-cobaltic 
 chloride, is formed. If a solution of cobaltous chloride in 
 concentrated ammonia is allowed to stand longer than is 
 required to form the preceding compound, or if an oxid- 
 izing agent is used, the product has the composition 
 Co(NH 3 ) 6 Cl 3 , and is ^nown as purpureo-cobaltic chloride. 
 And, finally, by further action of oxidizing agents, luteo- 
 cobaltic chloride, Co(NH 3 ) 6 Cl 3 , is formed. It will be ob- 
 served that these compounds form a series, the members 
 of which differ from one another by NH 3 or a multiple : 
 
 Co(NH,) s Cl,; Co(NH,) 4 Cl,; Co(NH.) 6 Cl, ; Co(NH,).Cl, 
 
 They may be regarded as made up of cobaltic chloride 
 and different numbers of molecules of ammonia. In 
 regard to the first one, it is simplest to consider it as 
 analogous to the mercur-ammonium compounds (see 
 page 633). Accordingly, it is usually represented as 
 derived from three molecules of ammonium chloride by 
 the substitution of a trivalent atom of cobalt for three 
 hydrogen atoms. 
 
 When a solution of cobaltous chloride in ammonia, 
 which has become red by contact with the air, is treated 
 at the ordinary temperature with concentrated hydro- 
 chloric acid, a brick-red precipitate of roseo-cobaltic chlo- 
 
716 INORGANIC CHEMISTRY. 
 
 ride, which has the same composition as purpureo-co- 
 baltic chloride, Co(NH 3 ) B Cl 3 , with a molecule of water, 
 is formed. 
 
 If in the above reactions the nitrate or sulphate of 
 cobalt is used, nitrates and sulphates corresponding to 
 the chlorides mentioned are obtained. 
 
 Thus : 
 
 Hoseo-salts. Luteo-salts. 
 
 Co(NH 3 ) 6 01 3 Co(NH 3 ) 6 Cl 3 
 
 Co(NH 3 ) 6 (N0 3 ) 3 Co(NH 3 ) 6 (N0 3 ) 3 
 
 [Co(NH 3 )J a (S0 4 ) 3 [Co(NH 3 )J 3 (S0 4 ) 3 
 
 NICKEL, Ni (At. Wt. 58.24). 
 
 General. Nickel differs from cobalt in respect to the 
 difficulty with which it forms nickelic compounds or 
 those in which it is trivalent. At the same time it does 
 form an oxide of the composition Ni 2 O 3 , and the corre- 
 sponding hydroxide, Ni(OH) 3 . In all other compounds 
 it is bivalent, the compounds being analogous to ferrous 
 compounds. 
 
 Occurrence and Preparation. Nickel occurs native in 
 meteorites. The principal minerals containing it are 
 the arsenide, NiAs, known as niccolite, and the sulph- 
 arsenide, NiSAs or NiS 2 .NiAs 2 , known as gersdorffite. 
 From the ores the oxide is obtained in the same way 
 that cobalt oxide is obtained from its ores. This is then 
 pressed in the form of small cubes, mixed with charcoal 
 powder, and ignited. The commercial metal is generally 
 found in the form of these cubes. 
 
 Properties. Nickel is a white metal with a slight cast 
 of yellow. It is very hard, and capable of a high polish. 
 The metal in its ordinary condition is brittle, but after 
 treatment with a little magnesium it becomes very mal- 
 leable. Its specific gravity is 8.9, and it melts at a high 
 temperature. It is not changed in the air ; it dissolves 
 slowly in hydrochloric and sulphuric acids, and readily 
 in nitric acid. Like iron, it is magnetic. 
 
 Alloys. Alloys of nickel are extensively used. A.r- 
 gentan or German silver consists of copper, zinc, and nickel. 
 Various nickel alloys are used for making coins. The 
 5 and 3 cent pieces in the United States are made of an 
 
COMPOUNDS OF NICKEL. 717 
 
 alloy consisting of 25 per cent nickel and 75 per cent 
 copper. In Switzerland, and Belgium also, nickel coins 
 are used. 
 
 Other Applications of Nickel. Besides as a constituent 
 of important alloys, nickel is extensively used at present 
 in nickel-plating. Iron objects are covered with a thin 
 layer of the metal for the purpose of protecting them 
 from rusting. The plating is accomplished as silver- 
 plating and copper-plating are by means of electrolysis, 
 -a bath of nickel-ammonium sulphate being used. 
 
 Nickelous Chloride, NiCl 2 , crystallizes from aqueous 
 solution with six molecules of water, and the crystals are 
 green. When the water of crystallization is driven off 
 they become yellow. In general, nickel salts, with their 
 water of crystallization, are green, and in the anhydrous 
 condition they are yellow. 
 
 Nickelous Hydroxide, Ni(OH) 2 , is formed when a nickel 
 salt is treated with a soluble hydroxide, and is a green 
 insoluble substance. When heated it is converted into 
 the green oxide, NiO. 
 
 Nickelic Hydroxide, Ni(OH) 3 , is precipitated as a black 
 powder when a solution of a nickel salt is treated with 
 sodium hypochlorite. 
 
 Cyanides. When potassium cyanide is added to a so- 
 lution of a nickel salt, nickel cyanide, Ni(CN) a , is precipi- 
 tated as a greenish-white substance. With an excess of 
 potassium cyanide this forms the salt, Ni(CN) 2 .2KCN, 
 which, owing to the fact that nickelous salts are not 
 converted into nickelic salts by oxidation, does not 
 undergo change when boiled with potassium cyanide. 
 When hydrochloric acid is added to a solution of the 
 double cyanide, nickelous cyanide is precipitated. If 
 boiled with precipitated mercuric oxide the double cy- 
 anide is decomposed and nickel oxide is thrown down. 
 
 Nickel Carbonyl, Ni(CO) 4 . This interesting compound 
 is formed when finely-divided nickel, such as is obtained 
 by reducing nickel oxide by hydrogen at about 400, is 
 allowed to cool in a slow current of carbon monoxide. 
 A gas is formed which can easily be condensed, its 
 boiling-point being 43. At 25 it solidifies forming 
 
718 INORGANIC CHEMISTRY. 
 
 needle-shaped crystals. When the gas is passed through 
 a heated tube pure nickel is deposited. Advantage of 
 this fact is taken for the purpose of preparing pure 
 nickel on the large scale. Cobalt does not form a com- 
 pound of this kind. 
 
 Reactions of Cobalt and Nickel which are of Special 
 Value in Chemical Analysis. The reactions with the sol- 
 uble hydroxides have been explained. With ammonium 
 sulphide both give black sulphides, which are not easily 
 dissolved by dilute hydrochloric acid. From solutions 
 of the acetates hydrogen sulphide precipitates the sul- 
 phides. Nickel sulphide is slightly soluble in ammonium 
 sulphide, and the solution has a brownish-yellow color. 
 
 The action of the hypochlorites upon solutions of 
 nickel and cobalt salts has been explained above. The 
 reactions with potassium cyanide have also been ex- 
 plained. These furnish a good method for separating, 
 the two metals. 
 
 When a solution of potassium nitrite is added to a solu- 
 tion of a cobalt salt containing free acetic acid or nitria 
 acid, a precipitate of cobaltic potassium nitrite is formed. 
 This is a compound of cobaltic nitrite, Co(NO 2 ) 3 , and potas- 
 sium nitrite, of the composition Co(NO 2 ) 3 .3KNO 2 . The 
 formation involves oxidation of the cobaltous salt, and 
 this is effected by some of the nitrogen trioxide which is 
 set free. Thus with the chloride the action may be rep- 
 resented as follows : 
 
 CoCl 2 + 7KNO 2 + 2C 2 H 4 2 = 
 
 2KC1 + Co(N0 2 ) 3 .3KN0 2 + 2KC 2 H 3 O 2 + H 2 O + NO. 
 
 Nickel does not form a similar compound of a nickelic 
 salt, but simply forms a double nitrite, containing the 
 nickelous salt Ni(NO 2 ) 2 .4KNO 2 . 
 
 Cobalt compounds color the bead of microcosmic salt 
 blue both in the reducing and oxidizing flame. Nickel 
 colors it reddish brown in the oxidizing flame when hot, 
 and pale yellow when cold. In the reducing flame it ia 
 gray. 
 
CHAPTER XXXIV. 
 
 ELEMENTS OF FAMILY VIII, SUB GROUP B j 
 RUTHENIUM RHODIUM PALLADIUM. 
 
 ELEMENTS OF FAMILY VIII, SUBGROUP C 
 OSMIUM IRIDIUM PLATINUM. 
 
 General. Comparing the members of the three sub- 
 groups of Family VIII, with reference to their atomic 
 weights and specific gravities, we have the following re- 
 in a rkable table: 
 
 Fe 
 
 Co 
 
 Ni 
 
 At. Wt. 55.6 
 
 At. Wt. 58.49 
 
 At. Wt. 58.24 
 
 Sp. Gr. 7.8 
 
 Sp. Gr. 8.5 
 
 Sp. Gr. 8.8 
 
 Ru 
 
 Rh 
 
 Pd 
 
 At. Wt. 100.91 
 
 At. Wt. 102.23 
 
 At. Wt. 105.56 
 
 Sp. Gr. 12.26 
 
 Sp. Gr. 12.1 
 
 Sp. Gr. 11.5 
 
 Os 
 
 Ir 
 
 Pt 
 
 At. Wt. 189.55 
 
 At. Wt. 191.66 
 
 At. Wt. 193.41 
 
 Sp. Gr. 22.48 
 
 Sp. Gr. 22.42 
 
 Sp. Gr. 21.50 
 
 It will be observed that the atomic weights and specific 
 gravities of the members of each sub-group are approxi- 
 mately the same. But just as there is a gradual change 
 in the chemical conduct as we pass from iron to nickel 
 in the iron group, so a similar gradation of properties is 
 observed in the other two groups. As far as the variety 
 of compounds which the}' form is concerned, ruthenium 
 and osmium are more like iron than they are like rho- 
 dium and iridium. Further, rhodium and iridium re- 
 semble each other, as regards the variety of their com- 
 pounds, more closely than they resemble palladium and 
 platinum, and a similar resemblance is noticed between 
 palladium and platinum. These relations will appear 
 
 (719) 
 
720 INORGANIC CHEMISTRY. 
 
 more clearly if the formulas of some of the principal 
 compounds of the elements under consideration are 
 placed together in a table. 
 
 Ku and Os Bh and Ir Pd and Pt 
 
 Eu0 4 OsO 4 EhO 2 IrO a PdO a PtO, 
 
 HKuO 4 Kh a O 3 Ir 2 O 8 PdO PtO 
 
 H 2 Ku0 4 H a OsO 4 EhO IrO Pd 3 O 
 
 EuO a OsO 2 IrCl 4 PdCl 4 PtCl 4 
 
 Ku 2 O 8 Os 2 O 3 BhCl s IrCl 3 PdCl 2 PtCl, 
 
 EuO OsO IrCl a 
 
 EuCl 4 OsCl 4 
 
 EuCl 3 OsCl 3 
 
 EuCl a OsCl a 
 
 The elements ruthenium and osmium have a more acidic 
 character than the others ; just as iron has a more acidic 
 character than cobalt and nickel. Euthenium forms not 
 only ruthenious acid, H 2 EuO 4 , which is analogous to 
 ferric, chromic, and manganic acids, but also perru- 
 thenious acid analogous to permanganic acid. Osmium 
 forms osmious acid, H 2 OsO 4 , but apparently no peros- 
 mious acid. The highest known oxides are derived from 
 these elements. These are the tetroxides, EuO 4 and 
 OsO 4 , in which the elements appear to be octovalent. 
 Neither rhodium nor iridium forms acids. As far as their 
 oxides and chlorides are concerned, they suggest manga- 
 nese more than any other element, but their oxides have 
 only weak basic properties. Passing finally to the last 
 pair, palladium and platinum, we find that they have not 
 the power to form oxides of the general formula M 2 O 3 > 
 but that they act as the members of Family IV do either 
 as bivalent or quadrivalent elements. Palladium, to be 
 sure, forms a compound, the sub-oxide Pd 2 O, which is 
 like the oxide of silver, Ag 2 O, and in which it appears to 
 be univalent. In fact the members of Family VIII form 
 the connecting link between the members of Family VII 
 and those of Family I. In manganese, as we have seen, 
 a maximum of power is reached as far as the valence is 
 concerned. It forms compounds in which it appears to 
 
GENERAL IN REGARD TO THE PLATINUM METALS. 721 
 
 be septivalent, sexivalent, quadrivalent, trivalent, and 
 bivalent. When we pass to iron, however, we find that 
 it is not septivalent. In its most complex compounds 
 this element is sexivalent, as in ferric acid, H 2 FeO 4 , 
 but it acts preferably as a trivalent or a bivalent 
 element. Then, further, as we have seen, cobalt forms 
 a few compounds in which it is trivalent, but it is 
 generally bivalent, and nickel is scarcely ever trivalent. 
 In its compounds nickel resembles copper in the cupric 
 compounds, and copper is the next element in the order 
 of increasing atomic weights. 'But copper has an ad- 
 ditional power which allies it to the members of Group 
 A, Family I. It acts as a univalent element in the 
 cuprous compounds. 
 
 Now, in the same way, there is an increase in the 
 complexity of the compounds formed by the elements, as 
 we pass from zirconium, to niobium, to molybdenum, and 
 below manganese in Family VII we should expect to find 
 an element forming compounds which in general resem- 
 ble those of manganese, and leading up to the octovalent 
 element ruthenium. Considering the relations between 
 iron and ruthenium, one is tempted to suspect that 
 this unknown element may exhibit a valence of nine 
 in some unstable compounds. While ruthenium is oc- 
 tovalent in its highest oxide, it is also septivalent in 
 HEuO 4 , sexivalent in H 2 KuO 4 , quadrivalent in EuO 2 , 
 trivalent in Ru 2 O 3 , and bivalent in EuO. Rhodium, 
 however, is only quadrivalent, trivalent, and bivalent ; 
 and palladium is quadrivalent, bivalent, and univalent. 
 Just as nickel leads naturally to copper, so palladium 
 leads naturally to silver. 
 
 In regard to the series to which osmium, iridium, and 
 platinum belong, not as much is known as in regard to 
 the series just referred to, though the three elements 
 themselves have been carefully studied. There is here 
 observed the same falling off of valence power from os- 
 mium to platinum ; and* just as nickel leads to copper, 
 and palladium leads to silver, so platinum leads natu- 
 rally to gold in Family I. 
 
INORGANIC CHEMISTRY. 
 
 THE PLATINUM METALS. 
 
 The six elements of Sub-Groups B and C, Family 
 VIII, are generally grouped together and spoken of as 
 the platinum metals. They occur together in nature, 
 and almost always in alloys, into the composition of 
 which all enter. The chief constituent is platinum, 
 which is present to the extent of 50 to 80 per cent, and 
 over. The alloys occur in only a few localities, in the 
 Ural Mountains, in California, Australia, Borneo, and a 
 few other places, and form small pieces which are mixed 
 with sand and earth. They generally contain also gold, 
 iron, and copper. Palladium occurs, further, in a gold 
 ore which is found in Brazil. 
 
 Metallurgy. The process for obtaining the metals 
 from the ores is based mainly upon the following facts : 
 (1) Gold is soluble in dilute aqua regia, while platinum 
 requires concentrated aqua regia; (2) platinic chloride, 
 PtCl 4 , and iridium chloride, IrCl 4 , form, with ammonium 
 chloride, difficultly soluble compounds of the formulas 
 (NH 4 ) 2 PtCl 6 (PtCl 4 .2NH 4 Cl) and (NH 4 ) JrCl 6 (IrCl 4 .2NH 4 Cl). 
 When these compounds are ignited, they are completely 
 decomposed, and the metals are left behind. When, 
 therefore, platinum-ore has been freed as far as possible 
 from sand and earth, it is first treated with dilute aqua 
 regia, which removes the gold, and then with concen- 
 trated aqua regia, which dissolves the platinum together 
 with a little iridium, leaving an alloy of iridium and os- 
 mium. 
 
 When the solution thus obtained is treated with 
 ammonium chloride, both metals are precipitated ; and 
 when the precipitate is ignited, both metals are left be- 
 hind in the form of a spongy mass. This consists, how- 
 ever, almost wholly of platinum, the amount of iridium 
 being very small. 
 
 KUTHENIUM, Eu (At. Wt. 100.91). 
 
 Preparation. Euthenium is obtained from the residue 
 which is left undissolved when platinum-ore is treated 
 with concentrated nitro- hydrochloric acid. 
 
RUTHENIUM OSMIUM. 723 
 
 Properties. When heated in oxygen it burns and forms 
 the oxide, EuO 2 . It is insoluble in the strong acids, and 
 even in nitro-hydrochloric acid it is almost insoluble. 
 Owing to its power to form salts of ruthenious acid, it is 
 dissolved when heated with potassium hydroxide and an 
 oxidizing agent, such as saltpeter or potassium chlorate, 
 and afterwards treated with water. 
 
 Chlorides. When heated in chlorine it forms the di- 
 chloride, EuCl 2 , and some of the trichloride, EuCl 3 . The 
 tetrachloride, EuCl 4 , is known in combination with chlo- 
 rides of the alkali metals. 
 
 Oxides. When ruthenium is heated with potassium 
 hydroxide and saltpeter, potassium ruthenite, K 2 EuO 4 , is 
 formed. The acid from which this salt is derived is 
 plainly ruthenious acid, H 2 EuO 4 , and this is related to 
 the oxide, EuO 3 . Neither the acid nor the anhydride is 
 known, however. When the solution is treated with 
 chlorine the first product is potassium perruthenite, 
 KEuO 4 , which forms a dark green solution, and is iso- 
 morphous with potassium permanganate and potassium 
 perchlorate. By further treatment of the solution with 
 a rapid current of chlorine, ruthenium peroxide, EuO 4 , is 
 formed. This is a volatile crystalline solid, which ap- 
 parently is not acidic. It is easily reduced to the ses- 
 quioxide, Eu 2 O 3 ; and, if heated, it is decomposed with 
 explosion. 
 
 The oxides of the formulas EuO 2 , Eu 2 O 3 , and EuO are 
 not basic, and do not dissolve in acids. 
 
 OSMIUM, Os (At. Wt. 189.55). 
 
 Preparation. As stated above, this element is left un- 
 dissolved in the form of an alloy with iridium when plati- 
 num-ore is treated with concentrated nitro-hydrochloric 
 acid. In order to separate it from the iridium, advan- 
 tage is taken of the fact that it forms a volatile peroxide, 
 OsO 4 , similar to that formed by ruthenium, while iridium 
 does not. 
 
 Properties. The metal does not melt at the highest 
 temperatures reached artificially. It has the highest 
 specific gravity of all known substances ; is easily oxid- 
 
724 INORGANIC CHEMISTRY. 
 
 ized when in finely divided condition ; and is converted 
 either by the oxygen of the air or by nitric acid into 
 osmium peroxide, OsO 4 . 
 
 Chlorides. The dicliloride, OsCl 2 , and the tetraMoride, 
 OsCl 4 , are formed by treating the metal with chlorine* 
 The trichloride, OsCl 3 , is not known in free condition. 
 
 Oxides. The metal as well as the oxides forms the per- 
 oxide, OsO 4 , when heated in the air. This is also formed 
 by treating a heated mixture of sodium chloride and the 
 alloy of osmium and iridium with chlorine and water 
 vapor. It is commonly called osmic acid, though its acid 
 properties are very weak. Like ruthenium peroxide it 
 is volatile. It sublimes in colorless, lustrous needles, 
 and boils without decomposition at a temperature a little 
 above 100. It has an intense odor similar to that of 
 chlorine, and its vapor attacks the eyes and respiratory 
 organs somewhat in the same way that chlorine does. 
 It dissolves slowly in water, and reducing agents pre- 
 cipitate the metal from the solution. A solution of osmic 
 acid is used in microscopic work. When injected into 
 the tissues, the parts are hardened and colored. 
 Potassium osmite, K 2 OsO 4 , is formed when a solution of 
 the peroxide in potassium hydroxide is treated with a re- 
 ducing agent. It is easily decomposed in water solution. 
 The oxides OsO, OsO 2 , and Os 2 O 3 have neither acid 
 nor basic properties. 
 
 BHODIUM, Eh (At. Wt. 102.23). 
 
 Ehodium has no acid properties, and does not form a 
 peroxide corresponding to those of ruthenium and os- 
 mium. On the other hand, its oxide, Eh 2 O 3 , is basic. 
 The chloride EhCl 3 is readily formed, and it is doubt- 
 ful whether the di- and tetrachlorides have been made. 
 
 IRIDIUM, Ir (At. Wt. 191.66). 
 
 Preparation. The extraction of iridium with plati- 
 num and with osmium from platinum -ore was referred 
 to above. In order to separate it from platinum, advan- 
 tage is taken of the fact that it forms a trichloride, IrCl 3 , 
 
IRIDIUM. 725 
 
 which with ammonium chloride gives an easily soluble 
 double chloride. The reduction is accomplished either 
 by heating the tetrachloride for some time at 150, or by 
 treating the insoluble double chloride in water with 
 hydrogen sulphide or with sulphur dioxide. From os- 
 mium it is separated by treating with moist chlorine, 
 when, as stated above, the osmium is converted into the 
 peroxide, which being volatile passes over. The residue 
 contains the iridium in the form of the tetrachloride, and 
 this, when treated with potassium chloride, forms the 
 difficulty soluble chloriridate, K 2 IrCl 6 . 
 
 Properties. Iridium has a grayish- white color, and 
 resembles polished steel. Its specific gravity is nearly 
 the same as that of osmium, being 22.42. It is harder 
 and more brittle than platinum ; melts at a higher tem- 
 perature ; and unless it is finely divided it is not dis- 
 solved by nitro-hydrochloric acid. When heated with 
 potassium hydroxide and saltpeter it is converted into 
 the oxide. 
 
 Chlorides. When finely divided iridium is treated 
 with nitro-hydrochloric acid it is converted into the 
 tetrachloride, IrCl 4 . When the solution of the tetra- 
 chloride is heated it gives off chlorine, and the dicHoride, 
 IrCl 2 , is formed. The trichloride, IrCl 3 , is formed when 
 the metal is heated in chlorine gas. Both the tetrachlo- 
 ride and the trichloride form double salts with the chlo- 
 rides of the alkali metals. Those with the tetrachloride 
 have the general formula M 2 IrCl B , or IrCl 4 .2MCl ; while 
 those with the trichloride have the general formula 
 M 3 IrCl 3 , or IrCl,.3MCl. The latter are all soluble in 
 water ; of the former, the potassium salt, K 2 IrCl,, and 
 the ammonium salt, (NH 4 ) 2 IrCl 6 , are almost insoluble in 
 water. 
 
 Oxides. The oxides have neither acid nor basic prop- 
 erties. The one most easily obtained is the dioxide 
 IrO 2 . The hydroxides, Ir(OH) 3 and Ir(OH) 4 , are ob- 
 tained, the former as a black and the latter as a blue 
 precipitate, by treating the chlorides with potassium 
 hydroxide. 
 
726 INORGANIC CHEMISTRY. 
 
 PALLADIUM, Pd (At. Wt. 105.56). 
 
 Preparation. The chief source of palladium is a 
 Brazilian gold-ore. From this ore the metal can be ob- 
 tained by various methods, one of which consists in 
 melting it together with silver, and then treating it with 
 nitric acid, when the silver and palladium dissolve, and 
 the gold remains undissolved. The silver is precipitated 
 as chloride and the palladium as the cyanide, and when 
 the latter is ignited it is decomposed, leaving palladium. 
 
 Properties. Palladium resembles iridium and plati- 
 num in appearance. Its specific gravity is only about 
 half as great as that of platinum, being 11.5 ; it is more 
 easily fusible than platinum, and dissolves in nitric acid 
 and in hot concentrated sulphuric acid. The property 
 of palladium which has perhaps attracted most atten- 
 tion is its power to absorb hydrogen, and form 
 
 Palladium-Hydrogen. The formation of this com- 
 pound was referred to under Hydrogen (which see). The 
 combination takes place even at the ordinary tempera- 
 ture, but best at 100. If the me-tal is brought into 
 hydrogen at this temperature, it absorbs more than 900 
 times its volume, forming an alloy of the composition 
 Pd 2 H. This alloy has a greater volume and lower spe- 
 cific gravity than the palladium from which it is formed. 
 At 130 it begins to decompose under the atmospheric 
 pressure, but continued heating at a red heat is neces- 
 sary to decompose it completely. If allowed to lie in 
 contact with the air the hydrogen is oxidized to water. 
 Palladium-hydrogen acts as a strong reducing agent, 
 the hydrogen which it gives up being apparently in the 
 nascent or atomic condition. 
 
 Chlorides. When palladium is dissolved in concen- 
 trated nitro-hydrochloric acid it is converted into palladic 
 chloride, PdCl 4 , which with the chlorides of the alkali 
 metals forms double salts similar to those formed by 
 iridium tetrachloride, and, as we shall see, by platinic 
 chloride. The tetrachloride is decomposed by evapo- 
 ration of its solution, giving up chlorine and leaving pal- 
 ladious chloride, PdCl a , which crystallizes, and forms 
 
PLATINUM. 727 
 
 with the chlorides of the alkali metals double salts of the 
 general formula M 2 PdCl 4 , or PdCl 2 .2MCl. 
 
 Oxides. The point of chief interest presented by the 
 oxides is that in one of them, the suboxide, Pd 2 O, the 
 metal appears as a univalent element. The dioxide or 
 palladia oxide, PdO 2 , has neither acid nor basic proper- 
 ties. The monoxide, or palladious oxide, PdO, forms 
 unstable salts with acids, an example being the sulphate, 
 PdS0 4 + 2H a O. 
 
 PLATINUM, Pt (At. Wt. 193.41). 
 
 Preparation. A general idea of the method of pro- 
 cedure in extracting platinum from its ores was given 
 on p. 722. Thus prepared, however, it always contains 
 iridium, and for some purposes for which platinum is 
 used this is objectionable. In order to purify the metal 
 advantage is taken of the fact that iridium chloride can 
 be converted into a trichloride, which with ammonium 
 chloride forms an easily soluble double salt (see p. 
 725). The metal as obtained by igniting ammonium 
 platinic chloride forms a gray spongy mass known as 
 spongy platinum. When a solution of platinous chloride 
 is boiled with potassium hydroxide, and alcohol gradually 
 added, the salt is reduced, and the platinum is precipitated 
 as an extremely fine powder, known as platinum black. 
 When spongy platinum and platinum black are heated 
 to fusion by the oxyhydrogen flame they are converted 
 into the compact variety. 
 
 Properties. Platinum is a grayish-white metal re- 
 sembling polished steel ; it can be drawn out into very 
 fine wire ; it melts in the flame of the oxyhydrogen blow- 
 pipe, and when heated above its melting-point it is 
 volatile ; its specific gravity is 21.5. At white heat it 
 can be welded. It is not dissolved by nitric acid, hydro- 
 " chloric acid, nor sulphuric acid, but it dissolves in nitro- 
 hydrochloric acid, forming the acid, H 2 PtCl 6 . Fusing 
 alkalies, and particularly a mixture of caustic potash 
 'and saltpeter, act upon it; but the alkaline carbonates 
 do not. In contact with red-hot charcoal and silicon 
 dioxide a compound of silicon and platinum is formed. 
 
728 INORGANIC CHEMISTRY. 
 
 Finely divided platinum has to a remarkable extent the 
 power of condensing gases upon its surface. It ab- 
 sorbs, for example, 200 times its own volume of oxy- 
 gen, and other gases in a similar way. The oxygen 
 thus absorbed is in active condition, and if oxidizable 
 substances are brought in contact with it they are easily 
 oxidized. Thus when a current of hydrogen is allowed 
 to flow against a piece of spongy platinum it takes fire, 
 owing to the presence of the condensed oxygen in the 
 pores of the platinum. Similarly, when sulphur dioxide 
 and oxygen are allowed to flow together over spongy 
 platinum, or even the compact metal, the two gases unite 
 to form sulphur trioxide. 
 
 Applications of Platinum. The metal is of the great- 
 est value to the chemist on account of its power to resist 
 the action of high temperatures and of most chemical 
 substances. It is used in the laboratory in the form of 
 wire, foil, crucibles, evaporating-dishes, tubes, etc., etc. 
 From what was said above it cannot be used with 
 alkalies and saltpeter, nor with nitro-hydrochloric acid. 
 Platinum vessels, further, should not be placed upon 
 red-hot charcoal. Metallic salts which are easily re- 
 duced, such as those of antimony and bismuth, should 
 not be heated in platinum vessels, as the reduced ele- 
 ments, like silicon, form alloys with the platinum, and 
 these, as a rule, are easily fusible. In the concentration 
 of sulphuric acid on the large scale platinum stills are 
 used. The price of platinum is not as high as that of 
 gold, but much higher than that of silver. 
 
 Alloys of Platinum. The only alloy of platinum which 
 is of any special importance is that which it forms with 
 iridium. A small percentage of iridium diminishes the 
 malleability of platinum very markedly, and makes it 
 brittle ; it, however, increases its resistance to the action 
 of reagents. An alloy of 90 per cent platinum and 10 
 per cent iridium has been adopted by the French Gov- 
 ernment as the best material from which to make normal 
 meters. This alloy is very hard, as elastic as steel, more 
 difficultly fusible than platinum, entirely unchangeable 
 in the air, and is capable of a high polish. 
 
COMPOUNDS OF PLATINUM. 729 
 
 Chlorides. Like palladium, platinum forms two chlo- 
 rides, platinom chloride, PtCl 2 , and platinic chloride, PtCl 4 . 
 The latter is formed when platinum is dissolved in aqua 
 regia, and the solution evaporated to dryness. From its 
 solution in water it crystallizes with ten or five molecules 
 of water. It is soluble in alcohol as well as in water. 
 When the dry substance is heated for some time to 
 225-230, it is decomposed, yielding platinous chloride, 
 which is a grayish-green powder insoluble in water. 
 
 Chlorplatinic Acid, H 2 PtCl 6 , is formed by direct union 
 of platinic chloride with hydrochloric acid. It crystal- 
 lizes with six molecules of water, and forms a series of 
 salts called the chlorplatinates, to which reference has 
 already been made. Those most commonly met with 
 in the laboratory are the potassium salt, K 2 PtCl 6 , or 
 PtCl 4 .2KCl, and the ammonium salt, (NH 4 ) 2 PtCl 6 , or 
 PtCl 4 .2NH 4 Cl, both of which are difficultly soluble in 
 water, and are therefore precipitated when platinic chlo- 
 ride is added to solutions containing potassium or ammo- 
 nium chloride. The sodium salt is easily soluble in water. 
 Many other chlorplatinates are known, and many crys- 
 tallize well. Considering the similarity in composition be- 
 tween chlorplatinic acid and fluosilicic acid, the conclu- 
 sion seems justified that they have the same constitution. 
 The reasons which lead to the belief that the constitution 
 of fluosilicic acid is properly represented by the formula 
 fFl 
 
 I T^l 
 
 Si 4 v jj make it probable that the constitution of 
 chlorplatinic acid should be represented by a similar 
 formula, Pt 
 
 Platinous chloride like platinic chloride combines with 
 other chlorides to form double salts, the general formula 
 of which is M 2 PtCl 4 , or PtCl 2 .2MCl. 
 
 Cyanides. Platinum forms a number of beautiful 
 double cyanides derived from an acid of the formula 
 
730 INORGANIC CHEMISTRY. 
 
 H 2 Pt(CN) 4 or Pt(CN) 2 .2HCN, which should be called 
 cyanplatinous acid. It is analogous to the acid from 
 which the double chlorides of platinous chloride are 
 derived, H 2 PtCl 4 . These cyanplatinites are easily ob- 
 tained, and, as a rule, crystallize well and are beauti- 
 fully colored. The magnesium salt, MgPt(CN) 4 + 7H 2 O, 
 forms quadratic prisms, the side faces of which have a 
 green metallic lustre, while the end faces are deep blue. 
 Hydroxides and Oxides. When a solution of platinic 
 chloride is treated with sodium hydroxide, and afterward 
 with acetic acid, a white precipitate of platinic hydroxide, 
 Pt(OH) 4 + 2H 2 O, is formed, which when dried at 100 
 loses water and is converted into the brown hydroxide, 
 Pt(OH) 4 . This loses water when heated higher and is 
 converted into the oxide, PtO 2 . In a similar way plati- 
 nous hydroxide, Pt(OH) 2 , and platinous oxide, PtO, are 
 obtained from platinous chloride. Platinic hydroxide,, 
 Pt(OH) 4 , has acid properties, and forms a few salts of 
 the general formula M 2 PtO 3 , of which barium platinate, 
 BaPtO 3 , is the best known. Platinic acid, from which 
 these salts are derived, is plainly formed from platinic hy- 
 droxide by loss of one molecule of water, and bears to it 
 the same relation that ordinary silicic acid, H 2 SiO 3 , bears 
 to normal silicic acid, Si(OH) 4 . Further, platinic acid 
 and chlorplatinic acid appear to be analogous compounds; 
 and the latter may be regarded as derived from the 
 former by replacement of the three atoms of oxygen 
 by six atoms of chlorine, as shown in the formulas ] 
 
 H 2 Pt0 3 ; 
 
 Sulphides. There are two sulphides of platinum which 
 are analogous to the two oxides, PtO and PtO 2 . These 
 are platinous sulphide, PtS, and platinic sulphide, PtS 2 > 
 They are black insoluble compounds, which are precip- 
 itated when hydrogen sulphide or soluble sulphides are 
 added to solutions of platinous and platinic chlorides. 
 
THE PLATINUM BASES. 731 
 
 Compounds with Ammonia The Platinum Bases. Like 
 cobalt salts, the salts of platinum form a large number 
 of compounds with ammonia. When ammonia acts upon 
 a solution of platinous chloride a compound of the for- 
 mula PtCl 2 (NH 3 ) 2 is formed. This is the starting-point 
 for a series of compounds, as the bromide, PtBr 2 (NH 3 ) 2 ; 
 the nitrate, Pt(NO 3 ) 2 (NH 3 ) 2 ; the sulphate, PtSO 4 (NH 3 ) 2 ; 
 etc. There is another series beginning with the chlo- 
 ride, PtCl 2 (NH 3 ) 3 ; another beginning with the chloride, 
 PtCl 2 (NH 3 ) 4 . All the above are to be regarded as derived 
 from platinous chloride. Similarly there are other series 
 obtained from platinic chloride. The chlorides have 
 the formulas PtCl 4 (NH 3 ) 2 , PtCl 4 (NH 3 ) 3 , PtCl 4 (NH 3 ) 4 . It 
 seems probable that these salts are ammonium salts 
 in which a part of the hydrogen of the ammonium is 
 replaced by platinum. Thus the chloride PtCl a (NH 8 ), 
 
 probably has the constitution P^<NH 8 fT Although a 
 
 great deal of work has been done on these platino- 
 ammonium compounds, and much interesting informa- 
 tion in regard to them has been gained, the subject of 
 their constitution is still in an unsatisfactory state. 
 
APPENDIX T 
 
 CONTAINING SPECIAL DIRECTIONS FOR 
 LABORATORY WORK. 
 
 Introduction. In order to become familiar with the prin- 
 ciples of Chemistry it is absolutely necessary that the student 
 should devote a part of his time to work in the laboratory > 
 the more the better. It is, further, necessary that the labora- 
 tory work should be done with the greatest care. Every piece 
 of apparatus should be carefully constructed, the desk should 
 be kept clean and in good order; and no work should be 
 abandoned until the student is satisfied that he has seen all 
 there is to be seen, and that he has learned all that the work 
 can teach him. He must learn to use his own senses, and to 
 believe what he sees, and not simply "what the book says." 
 It sometimes happens that owing to the peculiar way in which 
 an experiment is performed results quite different from those 
 anticipated are obtained. Under these circumstances it is not 
 advisable to conclude at once that " the book must be wrong." 
 It may be; but the probabilities are against this explanation 
 of the discrepancy. Nothing is more instructive than well- 
 directed efforts to find the causes of difficulties. Such efforts, 
 more than anything else, develop the spirit of true scientific 
 inquiry. It is advisable for the student to carry on the work 
 for which directions are given below in connection with the 
 study of the book. A good plan to follow is to read a chapter 
 with care; then to perform the experiments which are intended 
 to illustrate that chapter, and, while doing the work, again to 
 read; and afterwards to write out an account of what has been 
 done, noting everything of importance exactly as it was ob- 
 served. If experiments are necessary to account for phe- 
 nomena not described in the book, these should be described; 
 and if a conclusion is reached in regard to these phenomena, 
 the evidence upon which the conclusion is based should be 
 clearly stated. It is only by patient work carried on in this 
 way that one can hope to reach a clear conception of the 
 science. But by such work the desired result will be reached. 
 
 (733) 
 
734 EXPERIMENTS TO ACCOMPANY CHAPTER I. 
 
 Progress will seem slow at first, as it always does in a new sub- 
 ject; but in time the ideas will begin to arrange themselves 
 systematically, order will come out of confusion, and in this 
 result the conscientious student will find a delightful reward 
 for his labor. Every great branch of knowledge is made up 
 of details which are bound together by certain broad govern- 
 ing principles. It is impossible to avoid these details. In 
 order to understand the governing principles the details 
 must be studied to some extent. They form the raw material 
 from which the science is constructed; without them the 
 science would be impossible. As well might one hope to learn 
 a language by studying its grammatical rules and avoiding the 
 details of the mere words, as to learn chemistry by studying 
 the laws and avoiding contact with the things to which these 
 laws have reference. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER I. 
 CHEMICAL CHANGE CAUSED BY HEAT. 
 
 Experiment 1 In a clean dry test-tube put enough white 
 sugar to make a layer to an inch thick. Hold the tube 
 in the flame of a spirit-lamp or a laboratory burner. What 
 evidence is furnished by this experiment that chemical change 
 may be caused by heat? What is left in the tube? Is it 
 soluble? Is it sweet ? Is it sugar? 
 
 Experiment 2. Half-fill the bulb of an arsenic-tube with 
 red oxide of mercury, or, if such a tube is not available, pro- 
 ceed as follows: From a piece of hard-glass tubing of about 
 6 to 7 millimeters (^ inch) internal diameter cut off a piece 
 about 10 centimeters (4 inches) long by making a mark 
 across it with a triangular file, and then seizing it with both 
 hands, one on each side of the mark, pulling and at the 
 same time pressing slightly as if to break it. Clean and dry 
 it, and hold one end in the flame of a blast-lamp until it 
 melts together. During the melting turn the tube constantly 
 around its long axis so that the heat may act uniformly upon 
 it. Put into the tube thus made enough red oxide of mer- 
 cury (mercuric oxide) to form a layer about 12 millimeters 
 (i inch) thick. Heat the tube as in the last experiment. 
 What change in color is noticed ? What is deposited upon 
 the glass in the upper part of the tube ? What evidence is 
 
CHANGES EFFECTED BY AN ELECTRIC CURRENT. 735 
 
 furnished by this experiment that chemical change can be 
 effected by heat ? 
 
 CHEMICAL CHANGES CAN BE EFFECTED BY AN ELECTRIC 
 
 CURRENT. 
 
 Experiment 3. To the ends of insulated copper wires con- 
 nected with two cells of a Bunsen's or Grove's battery fasten 
 platinum plates, say 25 mm. (1 inch) long by 12 mm. ( inch) 
 wide. Insert these platinum electrodes into water contained 
 in a shallow glass vessel about 15 cm. (6 inches) wide and 7 
 to 8 cm. (3 inches) deep, taking care to keep them separated 
 from each other. No action will take place, for the reason, 
 as has been shown, that water will not conduct the current, 
 and hence when the platinum electrodes are kept apart there 
 is no cucrent. By adding to the water about one tenth its 
 own volume of strong sulphuric acid it acquires the power to 
 convey the current. It will then be observed that bubbles 
 rise from each of the platinum plates. In order to collect 
 them an apparatus like that shown in Fig. 15 may be used. 
 
 h and o represent glass tubes which may conveniently be about 
 30 cm. (1 foot) long and 25 mm. (1 inch) internal diameter. 
 They are first filled with the water containing one tenth its vol- 
 ume of sulphuric acid, and then placed with the mouth under 
 water in the vessel J. The platinum elec- 
 trodes are now brought beneath the invert- 
 ed tubes. The b ubbles which rise from them 
 will pass upward in the tubes and the water 
 will be pressed down. Gradually the water 
 will be completely forced out of one of the 
 tubes, while the other is still half full of wa- 
 ter. The substances thus collected in the 
 tubes are invisible gases. After the first tube 
 is full of gas, place the thumb over its mouth 
 and remove the tube. Turn it mouth up- 
 ward, and at once apply a lighted match to it. 
 A flame will be noticed. The gas which was 
 contained in the tube is therefore capable of 
 burning. It cannot, therefore, have been 
 air. In the mean time the second tube will 
 have become filled with gas. Kemove this tube in the same 
 way and insert a thin piece of wood with a spark on it. The 
 spark will at once burst into flame, and the burning of the 
 
 FIG. 15. 
 
736 EXPERIMENTS TO ACCOMPANY CHAPTER I. 
 
 wood will take place more actively than it does in ordinary air y 
 as may be shown by withdrawing it and again inserting it inta 
 the tube. The gas in this tube, it will be noticed, does not 
 take fire. Without going into further details, it is clear from 
 the above experiment that when an electric current acts on 
 water two invisible gases are produced. A chemical change is 
 caused by an electric current. 
 
 MECHANICAL MIXTURES AND CHEMICAL COMPOUNDS. 
 
 Experiment 4. Examine carefully a piece of coarse-grained 
 granite; break off some of it, and separate the constituents. 
 How many are there? By what properties do you recognize 
 them ? Powder a small bit of one of the constituents, and 
 examine the powder with the microscope. Do you recognize 
 more than one kind of matter ? Mix the powder of the three 
 constituents, and see whether in the mixed powder there is 
 any difficulty in detecting the three kinds of matter with the 
 aid of the microscope. 
 
 Experiment 5. Mix a gram or two of powdered roll- 
 sulphur and an equal weight of very fine iron filings in a 
 small mortar. Examine a little of the mixture with a micro- 
 scope. Not only can we recognize the particles of iron and of 
 sulphur by means of the microscope, but we can also pick out 
 the pieces of iron by means of a magnet. The magnet attracts 
 the iron but not the sulphur, so that by passing the magnet 
 often enough through the mixture we can pick out all the 
 iron and leave all the sulphur. This separation is really a 
 mechanical separation. It is only a somewhat more refined 
 method of picking out than that used in the case of granite. 
 
 Experiment 6. Pass a small magnet through the mixture 
 above prepared. Unless the substances used are thoroughly 
 dry, particles of sulphur will adhere to the magnet, but even 
 then it will be seen that most of that which is taken out of 
 the mixture is iron. 
 
 The iron and sulphur can also be separated by treating the 
 mixture with a liquid known as carbon disulphide. Sulphur 
 dissolves in this liquid, but iron does not. So that when the 
 mixture is treated with it the- iron is left behind, and can 
 easily be recognized as such. 
 
 Experiment 7. Pour two or three cubic centimeters of 
 carbon disulphide on a little powdered roll-sulphur in a dry 
 test-tube. The sulphur dissolves. Treat iron filings in the 
 
VARIOUS EXAMPLES OF CHEMICAL ACTION. 737 
 
 same way. The iron does not dissolve. Now treat a small 
 quantity of the mixture with carbon disulphide. After the 
 sulphur is dissolved pour off the solution on a good-sized 
 watch-glass and let it stand. Examine what remains undis- 
 solved in the test-tube, and satisfy yourself that it is iron. 
 After the liquid has evaporated examine what is left in the 
 watch-glass and satisfy yourself that it is sulphur. Why are 
 you justified in concluding that the substance left in the test- 
 tube is iron and that left on the watch-glass is sulphur ? 
 
 Experiment 8. Make a fresh mixture of three grams each 
 of powdered roll-sulphur and fine iron filings. Grind them 
 together intimately in a dry mortar and put them in a dry 
 test-tube. Heat gradually until the mass begins to glow. At 
 first the sulphur melts and becomes dark-colored. It may 
 even take fire. But soon something else evidently takes place. 
 The whole mass begins to glow, and if you at once take the 
 tube out of the flame, the mass continues to glow, becoming 
 brighter. This soon stops; the mass grows dark and gradually 
 cools down. As soon as it reaches the ordinary temperature, 
 the tube should be broken and the contents put in a mortar. 
 A close examination will show that the mass does not look 
 like the mixture of sulphur and iron with which we started. 
 It has a bluish-black color, and is apparently homogeneous. 
 An examination with the microscope, the magnet, and car- 
 bon disulphide will prove that, while there may be a little 
 iron left, and possibly a little sulphur, most of the bluish- 
 black mass is neither iron nor sulphur, but a new substance 
 with properties quite different from those of iron and from 
 those of sulphur. 
 
 OTHER EXAMPLES OF CHEMICAL ACTION. 
 
 Experiment 9. Examine a piece of calc-spar or marble. 
 You see that it is made up of pieces of definite shape. It is, 
 as we say, crystallized. It is quite hard, though a knife will 
 cut it. Heated in a hard-glass tube, as in Experiment 2, it 
 does not melt, but remains essentially unchanged. It does 
 not dissolve in water. To prove this, put a piece the size of a 
 pea in a test-tube with pure water. Thoroughly shake, and 
 then, as heating usually aids solution, boil. Now pour off a 
 few drops of the liquid on a piece of platinum-foil or a watch- 
 glass, and by gently heating cause the water to evaporate. If 
 there is anything in solution, there will be a solid residue on 
 
738 EXPERIMENTS TO ACCOMPANY CHAPTER 1. 
 
 the platinum-foil or watch-glass. If not, there will be no 
 residue. Now treat a small piece of the substance with dilute 
 hydrochloric acid and notice what takes place. Bubbles of gas 
 are given off. After the action has continued for about a 
 minute, insert a lighted match in the upper part of the tube. 
 It is extinguished, and the gas does not burn. The gas formed 
 in this case is therefore plainly not identical with either one 
 of those obtained from water b}^ the action of the electric cur- 
 rent (see Experiment 3). It is what is commonly called car- 
 bonic-acid gas. As the action continues, the piece of calc-spar 
 or marble grows smaller and smaller, and finally disappears, 
 when there is a clear solution. The substance has dissolved 
 in the hydrochloric acid. In order to determine whether any- 
 thing else has taken place besides the dissolving, we shall have 
 to get rid 01 the excess of hydrochloric acid. This we can 
 easily do by boiling it, when it passes off in the form of vapor, 
 and then whatever is in solution will remain behind. For this 
 purpose put the solution in a small, clean porcelain evaporat- 
 ing-dish, and put this on a vessel containing boiling water, or 
 a water-bath. The operation should be carried on in a place 
 in which the draught is good, so that the vapors will not collect 
 in the working-room. They are not poisonous, but they are 
 annoying. The arrangement for evaporating is represented in 
 Fig. 16. 
 
 After the liquid has evaporated and the suoatance in the 
 
 evaporating-dish is dry, examine 
 it, and carefully compare its prop- 
 erties with those of the substance 
 which was put into the test-tube. 
 Its structure wilj be found not to 
 present the regularities noticed in 
 the original substance. It is much 
 softer, dissolv(nS in water, melts 
 when heated in a hard-glass tube. 
 It does not give off a gas when 
 treated with hydrochloric acid. 
 When exposed to the air it soon 
 
 FIG. 16. , . 
 
 becomes moist, and after a time 
 
 liquid. The experiment shows that when hydrochloric acid 
 acts upon calc-spar or marble the latter at least loses its own 
 properties. It might be shown that some of the hydrochloric 
 acid also loses its properties. In place of tha two wo get a new 
 
VARIOUS EXAMPLES OF CHEMICAL ACTION. 739 
 
 substance with entirely different properties. The two sub- 
 stances have acted chemically upon each other arid produced, 
 a chemical compound. In this case it was only necessary to 
 bring the substances in contact in order to cause them to act 
 chemically upon each other. It was not necessary to heat 
 them, as it was in the case of the iron and sulphur. 
 
 Experiment 1O. Bring together in a test-tube a small 
 piece of copper and some moderately dilute nitric acid. In a 
 short time action begins. The upper part of the tube becomes 
 filled with a dark, reddish-brown gas which has a disagreeable 
 smell. Do not inhale it, as when taken into the lungs it 
 produces bad effects. The solution becomes colored dark 
 blue, and the copper disappears. Examine this solution, 
 as in Experiment 9, and see what has been formed. What are 
 the properties of the substance found after evaporation of the 
 liquid? Is it colored? Is it soluble in water? Does it change 
 when heated in a tube ? Is it hard or soft ? Does it in any 
 way suggest the copper with which you started ? 
 
 Experiment 11. Try the action of dilute sulphuric acid 
 on a little zinc in a test-tube. A gas will be given off. Apply 
 a lighted match to it. Does the result suggest anything no- 
 ticed in an experiment already performed? After the zinc 
 has disappeared, evaporate the solution as in Experiment 9. 
 Carefully compare the properties of the substance left behind 
 with those of zinc. 
 
 Experiment 12. Hold the end of a piece of magnesium 
 ribbon about 20 centimeters (8 inches) long in a flame until it 
 takes fire; then hold the burning substance quietly over a 
 piece of dark paper, so that the light, white product may 
 be collected. Compare the properties of this white product 
 with those of the magnesium. Here again a chemical act has 
 taken place. The magnesium has combined with something 
 which it found in the air, and heat was produced by the com- 
 bination. The product is the white substance. 
 
 Experiment 13. In a small, dry flask (400 to 500 ccm.) 
 put a bit of granulated tin. Pour upon it 2 or 3 ccm. concen- 
 trated nitric acid. If no change takes place, heat gently, and 
 presently there will be a copious evolution of a reddish-brown 
 gas with a disagreeable smetl, (under what conditions has a 
 gas like this already been obtained ?) the tin will disappear, 
 and in its place will appear a white powder. Compare the 
 
740 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 properties of this white powder with those of tin. Why are 
 you justified in concluding that they are not the same thing ? 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 PREPAKATIOX OF OXYGEN. 
 
 Experiment 14. Make a small quantity of oxygen by 
 heating strongly an arsenic-tube half-filled with manganese 
 dioxide and fitted with a small rubber delivery-tube. 
 
 Experiment 15. Make some oxygen by heating a few 
 grams of mercuric oxide in a hard-glass tube closed at one 
 end and connected at the other end by means of a cork with 
 a bent glass tube. 
 
 Experiment 16. Arrange an apparatus as shown in Fig. 
 17. A represents a flask of 100 ccm. capacity. By means of 
 
 FIG. 17. 
 
 a well-fitting rubber stopper one end of the bent-glass tube B 
 is connected with it, and the other end, which should turn 
 upward slightly, is placed under the surface of the water in C. 
 In A put 4 to 5 grams (about an eighth of an ounce) potas- 
 sium chlorate, and gently heat by means of the lamp. Notice 
 carefully what takes place. At first the potassium chlorate 
 will melt, forming a clear liquid. If the heat is increased, 
 the liquid will appear to boil, and it will soon be seen that a 
 gas is given off, Now bring the inverted cylinder D filled 
 with water over the end of the tube, and let the bubbles 
 
MEASUREMENT OF THE VOLUME OF GASES. 741 
 
 of gas rise in the cylinder. After a considerable quantity of 
 gas has been collected in this way the action stops, the mass 
 in the flask becomes solid, and apparently the end of the pro- 
 cess is reached. But if the heat is raised again, gas will again 
 begin to come off, and in this second stage a larger quantity 
 will be collected than in the first. Finally, however, the end 
 is reached, and the substance left in the flask remains un- 
 changed, no matter how long heat may be applied. An ex- 
 amination of the gas collected will show that a piece of wood 
 will burn in it very readily. Explain the changes which have 
 taken place in this experiment. Calculate how much oxygen 
 can be obtained by heating 12 grams of potassium chlorate. 
 
 MEASUREMENT or THE VOLUME OF GASES. 
 
 In studying chemical changes it often becomes necessary to 
 measure the volume of a gas, and it is important to know what 
 precautions must be taken in such cases. For the purpose a 
 tube is used which is graduated by marks etched on the out- 
 side. These marks may either indicate the number of cubic 
 centimeters of gas contained in the tube, or the length of the 
 column of gas. In the latter case it is of course necessary to 
 determine what volume corresponds to a given length of the 
 column. The chief difficulty encountered in measuring gas 
 volumes is due to the fact that the volume varies with the 
 temperature and pressure. When the temperature of a gas is 
 raised one degree centigrade its volume is increased ^-fj part. 
 If, therefore, the volume of a gas at is V, at t its volume 
 V will be 
 
 This expression may also be written thus : 
 
 V' = F-f 0.00366* . F, or V = F(l + 0.003660. 
 From these we get the expressions 
 
 97Q V' V 
 
 V=^r-< and- F = 
 
 4- t 1 + 0.00366** 
 
 It is customary to reduce the observed volume of a gas to the 
 volume which it would have at 0. The correction is easily 
 
742 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 made by the aid of the above formula. Thus, if the volume 
 of a gas is found to be 250 cubic centimeters at 15, and it is 
 required to know what the volume would be if the temperature 
 were reduced to 0, the calculation is made thus : In this 
 case the observed volume V 9 is 250 cc. ; t, the temperature, 
 is 15. Substituting these values in the equation 
 
 273 V V 
 
 \/ _ _ of I/ ___ 
 
 " 273 + t 9 " 1 + 0.00366^' 
 
 we have 
 
 273 X 250 250 
 
 TF- or = 
 
 " 273 + 15 ' " 1 + 0.00366 X 15' 
 
 from which we get 236.99 as the value of F. 
 
 But the volume of a gas varies also according to the pres- 
 sure. If the pressure is doubled, the volume is decreased one 
 half ; and if the pressure is decreased one half, the volume 
 is doubled, and so on. In other words, the volume of a gas 
 varies inversely according to the pressure. Increase the pres- 
 sure two, three, or four times, and the volume becomes one 
 half, one third, or one fourth, and vice versa. If the gas has 
 the volume F at the pressure P, and at pressure P f the vol- 
 ume V't these values are found to bear to one another the 
 relations expressed in the equation 
 
 VP= V'P'. 
 
 The pressure is usually stated in millimeters, and reference is 
 to the height of a column of mercury which the pressure cor- 
 responds to. A gas contained in an open vessel, or in a vessel 
 over mercury or water, in which the level of the liquid inside 
 and outside the vessel is the same, is under the pressure of 
 the atmosphere. What that is we learn from the barometer. 
 As this pressure varies, it is necessary to read the barometer 
 whenever a gas is measured, and then to reduce the observed 
 volume to certain conditions which are accepted as standard. 
 If the gas is measured in a tube over mercury or water, and 
 the level of the liquid inside the tube is higher than that out- 
 side, the gas is under diminished pressure, the amount of 
 diminution depending on the height of the column of mer- 
 
MEASUREMENT OF THE VOLUME OF OASES. 743 
 
 cury or water in the tube. Thus, if the arrangement is as 
 
 represented in Fig. 18, the height 
 of the mercury column above the 
 level of the mercury in the trougli be- 
 ing 100 millimeters, and the pressure 
 of the atmosphere 760 millimeters 
 of mercury; then the gas in the tube 
 is plainly not under the full atmos- 
 pheric pressure, for the atmosphere 
 is supporting a column of mercury 
 100 millimeters high, and the pres- 
 sure actually brought to bear on the 
 gas corresponds to 760 100 = 660 
 mm. of mercury. Suppose that in 
 this case the volume of gas actually 
 FIG. is. measured is 75 cc. Call this V. 
 
 What would be the actual volume V under the standard 760 
 
 mm. ? We have seen that 
 
 V'P'. 
 Now, in this case P = 760, V = 75, and P ' = 660, There- 
 
 fore, 760 V = 75 X 660, or V= Q = 65.13. 
 
 In all cases it is necessary to make a correction similar to 
 this in dealing with the volumes of gases. The correction for 
 temperature and that for pressure may be made in one opera- 
 tion, the formula being 
 
 273 V'P' V'P' 
 
 T/ _____ _ f\-r* T/ 
 
 ' ' 
 
 760(273 + *)' ' 760(1 + 0.00366*)' 
 
 in which V = the volume of the gas at and 760 mm. pres- 
 sure ; V = the observed volume ; t = the observed temper- 
 ature ; P' = the pressure under which the .gas is measured. 
 Some of the most important ideas which have been intro- 
 duced into chemistry with a view of explaining the regularities 
 observed in the quantities of substances which act upon one 
 another chemically have their origin in observations on the 
 conduct of gases. It is therefore of the highest importance 
 that the student should familiarize himself with the meaning 
 of the expression, " the volume of a gas under standard condi- 
 tions." The presence of water vapor in a gas also influences 
 
744 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 its volume, and this must be taken into account in refined 
 work. The formula for making all the corrections required 
 in determining the volume of a gas is 
 
 FIG. 19. 
 
 
 760(273 + t) ' 
 
 or V= 
 
 V'(P'-d) 
 760(1 + 0.003660' 
 
 in which the letters F, V, P', and t 
 have the same significance as in the last 
 formula given, while a is the tension of 
 water vapor at t. 
 
 A convenient apparatus for meas- 
 uring gas volumes, which simplifies 
 the process, is that represented in 
 Fig. 19. It consists of two tubes con- 
 nected at the base by means of a piece 
 of rubber tubing, and containing water. 
 The tube A is graduated, the other 
 is not. The gas the volume of which 
 is to be measured is brought into the 
 tube A, with the narrow opening at 
 the top, and the other tube is then 
 placed at the side of the one con- 
 taining the gas, and its height ad- 
 justed so that the column of liquid 
 in both tubes is at the same level. 
 Under these circumstances, obviously, 
 the gas is under the atmospheric pres- 
 sure for which the necessary correction 
 must of course be made. It is also 
 necessary in this case to make the cor- 
 rections for temperature and the ten- 
 sion of aqueous vapor. It is, further, 
 sometimes convenient when the gas is 
 measured over water to transfer the 
 measuring-tube to a vessel containing 
 enough water to permit the immersion 
 of the tube to a point at which the level 
 of the liquid inside and outside of the 
 tube is the same. In this case the 
 conditions are the same as in the apparatus described in the 
 last paragraph. The arrangement is represented in Fig. 20. 
 
DECOMPOSITION OF POTASSIUM CHLORATE. 745 
 
 DETERMINATION 'OF THE AMOUNT OF OXYGEN LIBERATED 
 WHEN A KNOWN WEIGHT OF POTASSIUM G'HLORATE is_ 
 
 DECOMPOSED BY HEAT. 
 
 Experiment 17. To determine how much oxygen is given 
 off when a known weight of potassium chlorate is decomposed 
 by heat, proceed as follows: In a small dry hard-glass tube 
 about 10 cm. long and 5 to 7 mm. internal diameter, closed 
 at one end, weigh out on a chemical balance about 0.25 gram 
 dry potassium chlorate, first weighing the tube empty. In- 
 troduce just above the potassium chlorate a plug of asbestos 
 which has been ignited, then, by means of a 
 blast -lamp, soften the upper end of the 
 tube, and draw it out so that it has the 
 form shown in Fig. 21. Now weigh the tube 
 again. 
 Let 
 
 a = weight of tube empty ; 
 
 b = weight of tube with potassium chlorate ; 
 
 c = weight of tube with potassium chlorate 
 
 and plug. 
 
 Connect at A by means of a short piece of 
 rubber tubing with the measuring tube Fig. 19 
 B 
 
 FIG. 20. FIG. 21. 
 
 so that the ends of the two tubes are almost in contact with 
 each other, the measuring tube having been previously filled 
 with water to the zero point, and the top closed by means 
 of the stop-cock. Open the stop-cock, and now heat the po- 
 tassium chlorate gently at first, and gradually higher until no 
 more gas is given off. After the gas has stood for half an hour 
 to cool it down to the temperature of the air, adjust the two 
 tubes of the measuring apparatus so that the level of the wa- 
 ter in both is the same; read off the volume of gas. At the 
 same time read the barometer and thermometer; and now 
 make the corrections for pressure and temperature as above 
 directed. The weight of a liter or 1000 cc. of oxygen at 
 and 760 mm. pressure is 1.4290 grams. Knowing the volume 
 of oxygen obtained, calculate the weight of this volume. 
 
746 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 Remove the tube containing the product left after the decom* 
 position of the potassium chlorate, and weigh it. 
 
 Let 
 
 d = weight of tube after decomposition of potassium chlorate. 
 Now 
 
 b a = weight of potassium chlorate used ; 
 
 d (a -\- c 1} = weight of potassium chloride left. 
 
 Knowing further the weight of the oxygen obtained in the 
 
 decomposition, which we may call e, it is obvious from what has 
 
 been said that 
 
 d (a -(- c V) -\- e should be equal to b a t 
 
 and the weights should all be in accordance with the equation 
 
 KC10 3 = KC1 + 30. 
 Make all the calculations, and see how nearly the results 
 
 obtained agree with what is required by this equation. Should 
 
 the results not be satisfactory the first time, repeat the work. 
 
 The more carefully the work is done, the more nearly will the 
 
 results agree with the equation. 
 Experiment 18. Mix 25 to 30 grams (or about an ounce) 
 
 of potassium chlorate with an equal weight of manganese di- 
 oxide in a mortar. The sub- 
 stances need not be in the 
 form of powder. Heat the 
 mixture in a glass retort, and 
 collect the gas by displace- 
 ment of water in appropri- 
 ate . vessels, cylinders, bell- 
 glasses, bottles with wide 
 mouths, etc. It will also be 
 well to collect some in a 
 gasometer, such as is com- 
 monly found in chemical lab- 
 oratories, the essential features 
 of which are represented in 
 Fig. 22. It is made either of 
 metal or of glass. The open- 
 ing at d can be closed by 
 means of a screw cap. In 
 order to fill it with water, 
 open the stop-cocks and pour 
 FIG. 22. the water into the upper part 
 
 of the vessel after having screwed on the cap d. When it is 
 
PHYSICAL PROPERTIES OF OXYGEN. 
 
 747 
 
 full, water will flow out of the small tube e. Now close all 
 the stop cocks, and remove the cap d. The water will stay 
 in the vessel for the same reason that it will stay in a 
 cylinder inverted with its mouth below water. To fill the 
 gasometer with gas, put it over a tub or sink, and introduce 
 the tube from which gas is issuing into the opening at d. 
 The gas will rise and displace the water, which will flow out 
 at d. When full, put the cap on. To get the gas out of the 
 gasometer, attach a rubber tube to e, pour water into the 
 upper part of the gasometer, open the stop-cock a and that 
 at e, when the gas will flow out, and the current can be regu- 
 lated by means of the stop-cock at e. 
 
 The arrangement of the retort is shown in Fig. 23. 
 
 FIG. 23. 
 
 FIG. 24. 
 
 PHYSICAL PROPERTIES OF OXYGEN. 
 
 Experiment 19. Inhale a little of the gas from one of the 
 bottles. Has it any taste ? any odor ? any color ? 
 
 CHEMICAL PROPERTIES OF OXYGEN. 
 
 Experiment 2O. Turn three of the bottles containing oxy- 
 gen with the mouth upward, leaving them covered with glass 
 plates. Into one introduce some sulphur in a so-called de- 
 flagrating-spoon, which is a small cup of iron or brass at- 
 tached to a stout wire which passes through a metal plate, 
 usually of tin (see Fig. 24). In another put a little char- 
 
748 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 coal (carbon), and in a third a piece of phosphorus* about 
 the size of a pea. Let them stand quietly, and notice what 
 changes, if any, take place. Sulphur, carbon, and phos- 
 phorus are elements, and oxygen is an element. It will be 
 noticed that the sulphur and the carbon remain unchanged, 
 while some change is taking place in the vessel containing the 
 phosphorus, as is shown by the appearance of white fumes. 
 After some time the phosphorus will disappear entirely, the 
 fumes will also disappear, and there will be nothing to show 
 us what has become of the phosphorus. If the temperature 
 of the room is rather high, it may happen that the phosphorus 
 takes fire. If it should, it will burn with an intensely bright 
 light. After the burning has stopped, the vessel will be filled 
 with white fumes, but these will quickly disappear, and the 
 vessel will apparently be empty. What do these experiments 
 prove with reference to the action of oxygen on sulphur, car- 
 bon, and phosphorus at the ordinary temperature ? 
 
 Experiment 21. In a deflagrating-spoon set fire to a little 
 sulphur and let it burn in the air. Notice whether it burns 
 with ease or with difficulty. Notice the odor of the fumes 
 which are given off. Now set fire to another small portion, 
 and introduce it in a spoon into one of the vessels containing 
 oxygen. It will be seen that the sulphur burns much more 
 readily in the oxygen than in the air. Notice the odor of the 
 fumes given off. Is it the same as that noticed when the 
 burning takes place in the air ? 
 
 Experiment 22. Perform similar experiments with char- 
 coal. 
 
 Experiment 23. Burn a piece of phosphorus not larger 
 than a pea in the air and in oxygen. In the latter case the 
 light emitted from the burning phosphorus is so intense 
 fchat it is painful to some eyes. It is better to be cautious. 
 The phenomenon is an extremely brilliant one. The walls 
 of the vessel in which the burning takes place become cov- 
 
 * Phosphorus should be handled with great care. It is always kept 
 under water, usually in the form of sticks. If a small piece is wanted, 
 take out a stick with a pair of forceps, and put it under water in an 
 evaporating-dish. While it is under the water, cut off a piece of the 
 size wanted. Take this out by means of a pair of forceps, lay it for 
 a moment on a piece of filter-paper, which will absorb most of the 
 water; then quickly put it in the spoon. 
 
OXYGEN IS USED UP IN COMBUSTION. 749 
 
 ered with a white substance, which afterwards gradually dis- 
 appears. 
 
 What differences do you notice between the burning in the 
 air and in oxygen ? In the experiments is there any sulphur, 
 or carbon, or phosphorus left behind ? Do the experiments 
 furnish any evidence that oxygen takes part in the action ? or 
 that oxygen is used up ? 
 
 Experiment 24. Straighten a steel watch-spring,* and 
 fasten it in a piece of metal such as is used for fixing a 
 deflagrating-spoon in an upright position; wind a little thread 
 around the lower end, and dip it in melted sulphur. Set fire 
 to this, and insert it into a vessel containing oxygen. For a 
 moment the sulphur will burn as in Experiment 21; but soon 
 the steel begins to burn brilliantly, and the burning continues 
 as long as there is oxygen left in the vessel. Notice that in 
 this case there is no flame, but instead very hot particles are 
 given off from the burning iron. The phenomenon is of great 
 beauty, especially if observed in a dark room. The walls of the 
 vessel become covered with a dark reddish-brown substance, 
 some of which will also be found at the bottom in larger 
 pieces. 
 
 OXYGEN is USED UP IN COMBUSTION. 
 
 Experiment 25. Is the odor of the contents of the bottle in 
 which the sulphur was burned the same as before the experi- 
 ment ? Introduce a stick with a small flame on it successively 
 into the vessels used in burning sulphur, carbon, phosphorus, 
 and iron. Is oxygen present or not ? What evidence have 
 you on this point ? 
 
 Experiment 26. Fill a tube say 30 to 40 cm. (12 to 15 
 inches) long, and 2-J- to 3 cm. (1 to 1 inches) wide, with 
 oxygen, and arrange it in a vessel over water, as shown in 
 Fig. ^25. Now fasten a small stick of phosphorus to the end 
 of a wire and push it into the tube so that about to % inch 
 of the phosphorus is above the water and exposed to the 
 oxygen. At first no action will take place, but after a 
 
 * Old watch-springs can generally be had of any watch maker or 
 mender for the asking. A spring can be straightened by unrolling it, 
 attaching a weight, and suspending the weight by the spring. The 
 spring is then heated up and down to redness with the flame of a Bun- 
 sen burner. 
 
750 EXPERIMENTS TO ACCOMPANY CHAPTER II. 
 
 time white fumes will be seen to rise from the phosphorus, * 
 
 and the phosphorus will begin to melt. This 
 
 action will be accompanied by a diminution 
 
 of the volume of the oxygen, as will be shown 
 
 by the rise of the water. When the water has 
 
 risen so as to cover the phosphorus, shove the 
 
 stick up so that it is again just above the 
 
 surface of the water. Some of the oxygen 
 
 will again be used up. By working carefully, 
 
 and repeating this process as many times as 
 
 may be necessary, the oxygen can all be used 
 
 up without the active burning of the phos- 
 
 phorus. Usually, however, before the action 
 
 is completed, the temperature of the phos- 
 
 phorus becomes so high that it takes fire, 
 
 when there is a flash of light in the tube and 
 
 a sudden rise of the water, showing that the 
 
 gas is suddenly used up. 
 
 Experiment 27. Burn a steel watch-spring as directed in 
 Experiment 24, with the difference that the 
 spring is passed air-tight through a cork which 
 is fitted tightly into the neck of the bell-jar. 
 As the spring burns, the water will rise from 
 the vessel in which the bell-jar is standing, 
 and it is necessary to pour water into this ves- 
 sel. When the spring has burned near to- 
 the cork shove it through so that the burning 
 may continue. If the experiment is properly 
 performed the bell- jar will be nearly full of 
 water at the end. What does this prove ? 
 
 THE PRODUCTS OF COMBUSTION WEIGH MORE 
 THAN THE BODY BURNED. 
 
 FIG. 25. 
 
 Flo 
 
 Experiment 28. Weigh off about a gram 
 of magnesium ribbon in a porcelain crucible. 
 Heat over a Bunsen burner until the magnesium has turned 
 to a white substance (magnesium oxide). After cooling, weigh 
 again. Perform the same experiment with zinc, tin, and lead. 
 What conclusion are you justified in drawing? 
 
 Experiment 29. Over each pan of a large and rather 
 sensitive balance suspend a glass tube filled with pieces of solid 
 
PREPARATION OF HYDROGEN. 751 
 
 caustic soda. A balance that will answer the purpose very 
 well can be made of wood with metal bearings. It may con- 
 veniently be about 2? feet high, with a delicate beam about 
 3 feet long. The best tubes for the caustic soda are Argand 
 lamp-chimneys, around the bottom of which is tied a piece of 
 wire-gauze to prevent the caustic soda from falling out. On 
 one pan of the balance place a candle directly under one of 
 the caustic-soda tubes, so adjusted that the flame shall be not 
 more than ty to 3 inches below the bottom of the tube. By 
 means of weights placed on the other pan establish equilibrium. 
 Now light the candle. Slowly, as it burns, the pan upon 
 which it is placed will sink, showing that the products of com- 
 bustion which are partly absorbed by the caustic soda are 
 heavier than the candle was. While this is by no means an 
 accurate experiment, it is a very striking one, and proves be- 
 yond question that in the process of combustion matter is 
 taken up by the burning body. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER III. 
 
 PREPARATION OF HYDROGEN. 
 Experiment 3O. Kepeat Experiment 3 and examine the 
 
 Experiment 31. Throw a small piece of sodium * on water. 
 While it is floating on the surface apply a lighted match to it. 
 A yellow flame will appear. This is burning hydrogen, the 
 flame being colored yellow by the presence of the sodium, some 
 of which also burns. Make the same experiment with potas- 
 sium. The flame appears in this case without the aid of the 
 match. It has a violet color, which is due to the burning of 
 some of the potassium. The gas given off in these experi- 
 ments is either burned at once or escapes into the air. In the 
 case of the potassium it takes fire at once, because the action 
 takes place rapidly and the heat evolved is sufficient to set fire to 
 it; in the case of the sodium, however, the action takes place 
 more slowly, and the temperature does not get high enough to 
 set fire to the gas. In order to collect it unburned, it is only 
 necessary to allow the decomposition to take place so that the 
 
 * The metals sodium and potassium are kept under oil. When a 
 small piece is wanted take out .one of the larger pieces from the bottle, 
 roughly wipe off the oil with filter-paper, and cut off a piece the size 
 needed. It is not advisable to use a piece larger than a small pea. 
 
752 EXPERIMENTS TO ACCOMPANY CHAPTER III. 
 
 gas will rise in an inverted vessel filled with water. For this 
 purpose fill a good-sized test-tube with water and invert it in 
 a vessel of water. Cut off a piece of sodium not larger than a 
 pea, wrap it in a layer or two of filter-paper, and with the 
 fingers or a pair of curved forceps bring it quickly below the 
 mouth of the test-tube and let go of it. It will rise to the 
 top, the decomposition of the water will take place quietly, 
 and the gas formed, being unable to escape, will remain in the 
 tube. By repeating this operation in the same tube a second 
 portion of gas can be made, and so on until the vessel is full. 
 
 Examine the gas and see whether it acts like the hydrogen 
 obtained from water by means of the electric current. What 
 evidence have you that they are the same ? Is this evidence 
 sufficient to prove the identity of the two ? 
 
 The metals sodium and potassium disappear in these experi- 
 ments, and we get hydrogen. What becomes of the metals ? 
 and what is the source of the hydrogen ? If after the action 
 has stopped the water is examined, it will be found to contain 
 something in solution. It now has a peculiar taste, which we 
 call alkaline; it feels slightly soapy to the touch; it changes 
 certain vegetable colors. If the water is evaporated off, a 
 white substance remains behind, which is plainly neither 
 sodium nor potassium. In solid form or in very concentrated 
 solution it acts very strongly on animal and vegetable sub- 
 stances, disintegrating many of them. On account of this 
 action it is known as caustic soda, or, in the case of potas- 
 sium, as caustic potassa. 
 
 FIG. 27. 
 
 Experiment 32. Certain metals which do not decompose 
 water at ordinary temperatures, or which decompose it slowly, 
 decompose it easily at elevated temperatures. This is true of 
 
PREPARATION OF HYDROGEN. 
 
 "53 
 
 iron. If steam is passed through a tube containing pieces of 
 iron heated to redness, decomposition of the water takes place, 
 and the oxygen is retained by the iron, which enters into com- 
 bination with it, while the hydrogen is liberated. In this ex- 
 periment a porcelain tube with an internal diameter of from 
 20 to 25 mm. (about an inch) and a gas furnace are desirable, 
 though a hard-glass tube and a charcoal furnace will answer. 
 The arrangement of the apparatus is shown in Fig. 27. 
 
 Experiment 33. In a cylinder or test-tube put some small 
 pieces of zinc, and pour upon it some ordinary hydrochloric 
 acid. If the action is brisk, after it has continued for a min- 
 ute or two apply a lighted match to the mouth of the vessel. 
 The gas will take fire and burn. If sulphuric acid diluted 
 with five or six times its volume of water* is used instead of 
 hydrochloric acid, the same result will be reached. The gas 
 evolved is hydrogen. For the purpose of collecting the gas 
 the operation is best performed in a wide-mouthed bottle, in 
 
 FIG. 28. 
 
 FIG. 29. 
 
 which is fitted a cork with two holes (see .Fig. 28), or in a 
 bottle with two necks called a Wolffs flask (see Fig. 29). 
 Through one of the holes a funnel-tube passes, and through 
 the other a glass tube bent in a convenient form. 
 
 * If it is desired to dilute ordinary concentrated sulphuric acid with 
 water, the acid should be poured slowly into the water while the mix- 
 ture is constantly stirred. If the water is poured into the acid, the heat 
 evolved at the places where the two come in contact may be so great as 
 to convert the water into steam and cause the strong acid to spatter. 
 
754 EXPERIMENTS TO ACCOMPANY CHAPTER IIL 
 
 The zinc used is granulated. It is prepared by melting it 
 in a ladle, and pouring the molten metal from an elevation 
 of four or five feet into water. The advantage of this form 
 is that it presents a large surface to the action of the acids. 
 A handful of this zinc is introduced into the bottle, and 
 enough of a cooled mixture of sulphuric acid and water (1 
 volume concentrated acid to 6 volumes water) poured upon it 
 to cover it. Usually a brisk evolution of gas takes place at 
 once. Wait for two or three minutes, and then collect some 
 of the gas by displacement of water. When the action be- 
 comes slow, add more of the dilute acid. It will be well to 
 fill several cylinders and bottles with the gas, and also a gaso- 
 meter, from which it can be taken as it is needed for experi- 
 ments. 
 
 SOMETHING BESIDES HYDROGEN' is FORMED. 
 
 Experiment 34. After the action is over pour the contents 
 of the flask through a filter into an evaporating-dish, and 
 boil off the greater part of the water, so that, on cooling, the 
 substance contained in solution will be deposited. If the op- 
 eration is carried on properly, the substance will be deposited 
 in regular forms called crystals. It is zinc sulphate, ZnS0 4 , 
 formed by the replacement of the hydrogen of the sulphuric 
 acid by zinc. 
 
 PROBLEMS. How much zinc would it take to give 200 liters of hy- 
 drogen ? How much zinc sulphate would be formed ? How much hy- 
 drogen would be formed by the action of 50 grams of zinc on sulphuric 
 acid ? How much sulphuric acid would be used up ? 
 
 DETERMINATION OF THE AMOUNT OF HYDROGEN EVOLVED 
 WHEN A KNOWN WEIGHT OF ZINC is DISSOLVED IN SUL- 
 PHURIC ACID. 
 
 Experiment 35. This determination can be made by means 
 of an apparatus such as represented in Fig. 30. The bent 
 tube leading from the flask A is drawn out at B, and a plug 
 of glass-wool introduced below the constriction. The other 
 parts of the apparatus need no description. The flask should 
 have a capacity of about 40 to 50 cc. ; and the measuring tube 
 C should have a capacity of about 100 cc., and be graduated to 
 cc. 
 
AMOUNT OF HYDROGEN EVOLVED. 
 
 '55 
 
 " The experiment is conducted in the following manner : 
 D is filled with distilled water ; a piece of zinc weighing from 
 0.150 to 0.200 gram is placed in the flask; the pinch-cock E 
 is then opened, and the whole apparatus thus filled with 
 water. The apparatus is now examined in order to ascertain 
 if gas bubbles are lodged under the stopper F or in the glass- 
 wool. If so, they can usually be dislodged without difficulty. 
 If they persist, a few moments' boiling of the water in the 
 flask will eifect their complete removal. . . The eudiometer is 
 now placed over the outlet of the delivery-tube, and the 
 greater portion of the water remaining in D allowed to flow 
 through the apparatus. Sulphuric acid of the concentration 
 ordinarily employed in the laboratory (1 of H 2 S0 4 to 4 of H 3 0) 
 is poured into the reservoir D until it is nearly full. The 
 pinch-cock E is then opened, and the water which fills the 
 
 Fio. 30. 
 
 apparatus is displaced by sulphuric acid. The action of the 
 acid upon the metal may be facilitated by heat or by adding 
 some platinum scraps. When the action is over, the contents 
 of the flask are swept through the delivery-tube by again open- 
 ing the pinch-cock E. Finally, the measuring-tube is trans- 
 ferred to a cylinder of water, and the volume of the gas read 
 and corrected in the usual manner. If hydrochloric instead of 
 sulphuric acid has been used, which would be the case when 
 the metal employed is aluminium, a little caustic soda should 
 be added to the water in the cylinder to which the eudiometer 
 is transferred."* 
 
 * See Morse and Keiser, American Chemical Journal, vol. vi. p. 349. 
 
?5ti EXPERIMENTS TO ACCOMPANY CHAPTER III. 
 
 A liter of hydrogen at and 760 mm. weighs 0.089873 
 gram. How much does the hydrogen obtained in the experi- 
 ment weigh? How much ought to have been obtained? 
 How many cubic centimeters of hydrogen ought to have been 
 obtained ? 
 
 Try the same experiment, using tin- and hydrochloric acid. 
 The action takes place as represented in the equation 
 
 Sn + 2HC1 = SnCl a + H 2 . 
 
 It would be well, further, to try the experiment also with iron 
 and sulphuric acid, and with aluminium and hydrochloric 
 acid, and to calculate from the results the relation between 
 the weights of the four metals required to give equal volumes 
 of hydrogen, and the volumes of hydrogen given by, say, a 
 gram of each metal. The action between iron and sulphuric 
 acid takes place according to the equation 
 
 Fe + H 2 S0 4 = FeS0 4 + H a . 
 
 That between aluminium and hydrochloric acid is represented 
 by this equation : 
 
 Al + 3HC1 = A1C1 3 + 3H. 
 
 HYDROGEN is PURIFIED BY PASSING THROUGH A SOLUTION 
 
 OF POTASSIUM PERMANGANATE. 
 
 Experiment 36. Pass some of the gas, made by the action 
 of zinc on sulphuric acid, through a wash cylinder contain- 
 
 FIG. 31. 
 
 ing a solution of potassium permanganate; collect some of it, 
 and notice whether it has an odor. The apparatus should 
 
DIFFUSION. 
 
 75? 
 
 Fia. 32. 
 
 be arranged as shown in Fig. 31. The solution of potas- 
 sium permanganate is, of course, contained in the small cyl- 
 inder A, and the tubes so arranged that the gas bubbles 
 through it. 
 
 Has the gas any odor or taste or color ? 
 
 Experiment 37. Place a vessel containing hydrogen with 
 the mouth upward and uncovered. In a short time examine 
 the gas contained in the vessel, and see whether it is hydrogen, 
 What does this experiment prove with reference to the weight 
 of hydrogen as compared with that of the air ? 
 
 Experiment 38. Gradually bring a vessel containing hy- 
 drogen with its mouth upward 
 below an inverted vessel contain- 
 ing air, in the way shown in Fig. 
 32. After the vessel which con- 
 tained the hydrogen has been 
 brought in the upright position 
 beneath the other, examine the 
 gas in each vessel. Which one 
 contains the hydrogen ? 
 
 Experiment 39. Soap-bubbles filled with hydrogen rise in 
 the air. This experiment is best performed by connecting an 
 ordinary clay pipe by means of a piece of rubber tubing with 
 the delivery-tube of a gasometer filled with hydrogen. Small 
 balloons of collodion are also made for the purpose of showing 
 the lightness of hydrogen. 
 
 HYDROGEN PASSES READILY THROUGH POROUS VESSELS. 
 DIFFUSION. 
 
 9 
 
 Experiment 4O. Arrange an apparatus as shown in Fig. 
 S3. It consists of a porous earthenware cup, such as is used 
 in galvanic batteries, fitted wi'th a perforated cork connected 
 with a glass tube 2 to 3 feet long. The cork must fit air-tight 
 into the mouth of the cup, as well as the tube into the cork 
 This may be secured by shoving the cork into the cup until 
 its outer surface is even with the edge of the cup, and then 
 covering it carefully with sealing-wax. Put the lower end of 
 the glass tube through a cork into one neck of a Wolffs 
 bottle containing some water colored with litmus or indigo, so 
 that the end of the tube is above the surface of the water. 
 Through the other neck of the bottle pass a tube slightly bent 
 outward and drawn out at the end to a fine opening. This 
 
758 EXPERIMENTS TO ACCOMPANY CHAPTER III. 
 
 tube must also be fitted to the bottle by an air- tight cork, 
 
 and its lower end must be 
 below the surface of the 
 liquid. Now bring a bell- 
 jar containing dry hydro- 
 gen over the porous cup, 
 when the liquid will be 
 seen to rise in the short, 
 bent tube that dips be- 
 low the liquid, and be 
 forced out of it, some- 
 times with considerable 
 velocity. Withdraw the 
 bell-jar, and bubbles will 
 rise rapidly from the bot- 
 tom of the tube which 
 dips under the water, thus 
 showing that air is enter- 
 
 FIG. 33. 
 
 FIG. 34. 
 
 ing the bottle. This is due to the diffusion of the hydrogen 
 from the porous cup into the air. Explain all that you have 
 seen. 
 
 CHEMICAL PROPERTIES OF HYDROGEN. 
 
 Experiment 41. If there is no small platinum tube avail- 
 able, roll up a small piece of platinum-foil and melt it into 
 the end of a glass tube, as shown in Fig. 34. Connect the 
 burner thus made with the gasometer containing hydrogen, 
 and after the gas has been allowed to issue from it for a 
 moment, set fire to it. In a short time it will be seen that the 
 flame is practically colorless, and gives no light. That it is 
 
PRODUCT FORMED WHEN HYDROGEN IS BURNED. 759 
 
 hot can be readily shown by holding a piece of platinum wire 
 or a piece of some other metal in it. 
 
 Experiment 42. Into the flame of burning hydrogen in- 
 troduce a small coil of platinum wire. What change is ob- 
 served ? Introduce also a piece of magnesium ribbon. Explain 
 the difference between the two cases. What 
 becomes of the magnesium ? of the platinum ? 
 
 Experiment 43. Hold a cylinder filled 
 with hydrogen with the mouth downward. 
 Insert into it a lighted taper held on a bent 
 wire, as shown in Fig. 35. The gas takes fire 
 at the mouth of the vessel, but the taper is 
 extinguished. On withdrawing the taper and 
 holding the wick for a^ moment in the burn- 
 ing hydrogen, it will take fire, but on putting 
 it back in the hydrogen it will again be extin- 
 guished. Other burning substances should 
 be tried in a similar way. What conclu- 
 sions are justified by the last two experi- 
 ments ? FIG. 3t>. 
 
 PRODUCT FORMED WHEN HYDROGEN is BURNED. 
 
 Experiment 44. Hold a clean, dry glass plate a few inches 
 above a hydrogen flame. What do you observe? Kemove 
 what is deposited upon the plate, and hold the plate again 
 over the flame. Repeat this a number of times. What does 
 the substance deposited upon the plate suggest? Can you 
 positively say what it is ? 
 
 REDUCTION. 
 
 Experiment 45. Arrange an apparatus as shown in Fig. 
 36. The flask A contains zinc and dilute sulphuric acid; the 
 cylinder B a solution of potassium permanganate; the cylinder 
 concentrated sulphuric acid; and the tube D granulated 
 calcium chloride. The object of the potassium permanganate 
 is to purify the hydrogen; the object of the concentrated sul- 
 phuric acid and calcium chloride is to remove moisture from 
 the gas. In the tube E put a few pieces of the black oxide of 
 copper, or cupric oxide, CuO. After hydrogen has been pass- 
 ing long enough to drive all the air out of the apparatus 
 (about two or three minutes if there is a brisk evolution) heat 
 
760 EXPERIMENTS TO ACCOMPANY CHAPTER IV. 
 
 the oxide of copper by means of a flame applied to the tube. 
 What change in color takes place ? Try the action of nitric 
 
 Fio. 36. 
 
 acid on the substance before the action and after, and note 
 whether there is any difference. What appears in G ? Ex- 
 plain what you have seen. 
 
 Experiment 46. Try the experiment just described, using 
 ferric oxide, or oxide of iron, Fe 2 3 , instead of cupric oxide. 
 What is the common feature in the two reactions ? 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER IV. 
 COMPOSITION OF WATER. 
 
 Experiment 47. Arrange the apparatus shown in Fig. 36 
 with a straight tube instead of the bent tube E, and connect 
 this with a small bent tube containing calcium chloride, as 
 shown in Fig. 37. Weigh tube E empty, and after the cupric 
 
 u 
 
 FIG. 37. 
 
 oxide has been put into it. This gives the weight of the cupric 
 oxide. Weigh the tube F before the experiment. Now pro- 
 ceed as in Experiment 45. In this case all the water formed 
 by the action of the hydrogen on the cupric oxide will be 
 
COMPOSITION OF WATER. 761 
 
 absorbed by the calcium chloride in tube F. This tube will 
 therefore gain in weight, and as oxygen is removed from the 
 cupric oxide, tube E will lose in weight. After the reduction 
 is complete weigh tube E and tube F again. 
 
 Let x weight of tube E -f- cupric oxide before the ex- 
 periment; 
 
 y = weight of tube E -{ copper after the experiment. 
 Then x y = weight of oxygen removed from the cupric 
 
 oxide. 
 
 Let a = weight of tube F before the experiment, 
 and b = " " " after " " 
 
 Then b a = weight of water formed. 
 
 If the experiment is properly performed, it will be found 
 that the ratio -r - is very nearly -. Or the result may be 
 
 o ci> y 
 
 stated thus: In nine parts of water there are eight parts of 
 oxygen. 
 
 Experiment 48. The tubes in the apparatus used in Ex- 
 periment 3, or some other similar apparatus, should be 
 graduated. Let the gases formed by the action of the electric 
 current, as in Experiment 3, rise in the tubes, and observe 
 the volumes. It will be seen that when one tube is just full 
 of gas, the other, if it is of the same size, will be only half full. 
 On examining the gases the larger volume will be found to be 
 hydrogen, and the smaller volume oxygen. What are the 
 relative weights of equal volumes of hydrogen and oxygen ? 
 In what proportion by weight are the two gases obtained 
 from water in this experiment? How does this result agree 
 with that obtained in the preceding experiment ? Does this 
 experiment prove that water consists only of hydrogen and 
 oxygen ? 
 
 Experiment 49. Pass hydrogen from a generating- flask or 
 a gasometer through a tube containing some substance that 
 will absorb moisture ; for all gases made in the ordinary 
 way and collected over water are charged with moisture. 
 The calcium chloride should be in granulated form, not 
 powdered. After passing the hydrogen through the cal- 
 cium chloride, pass it through a tube ending in a narrow 
 opening, and set fire to it. If now a dry vessel is held over 
 the flame, drops of water will condense on its surface and run 
 
762 EXPERIMENTS TO ACCOMPANY CHAPTER IV. 
 
 down. A convenient arrangement of the apparatus is shown 
 in Pig. 38. 
 
 FIG. 38. 
 
 A is the calcium chloride tube. Before lighting the jet, 
 hold a glass plate in the escaping gas, and see whether water 
 is deposited on it. Light the jet before putting it under the 
 bell- jar ; otherwise, if hydrogen is allowed to escape into the 
 vessel, it will contain a mixture of air and hydrogen, and this 
 mixture, as we shall soon see, is explosive. 
 
 Experiment 50. Mix hydrogen and oxygen in the propor- 
 tions of about 2 volumes of hydrogen to 1 volume of oxygen, 
 in a gasometer. Fill soap-bubbles, made as directed in Ex- 
 periment 39, with this mixture, and allow them to rise in the 
 air. As they rise, bring a lighted taper in contact with them, 
 when a sharp explosion will occur. Great care must be taken 
 to keep all flames away from the vicinity of the gasometer 
 while the mixture is in it. This experiment is conveniently 
 performed by hanging up, about six to eight feet above the 
 experiment-table, a good-sized tin funnel-shaped vessel, with 
 the mouth downward. a Now place a gas jet or a small flame 
 of any kind at the mouth of the vessel. If the soap-bubbles 
 are allowed to rise below this apparatus they will come in con- 
 tact with the flame and explode at once.* What does this 
 experiment show ? Does it give any information in regard to 
 the composition of water ? 
 
 * The same apparatus may be used in experimenting with soap 
 bubbles filled with hydrogen. 
 
EUDIOMETRIC EXPERIMENTS. 763 
 
 EUDIOMETRIC EXPERIMENTS. 
 
 Experiment 51. The general method of studying the 
 combination of hydrogen and oxygen by means of the eudi- 
 ometer was described in the text (see p. 50). To what was 
 there said it need only be added that, in exploding the mix- 
 ture in the eudiometer, the latter should be held down firmly, 
 by means of a clarnp, against a thick piece of rubber cloth 
 placed on the bottom of the mercury-trough. In making the 
 measurements of the volume of the gases and the height of 
 the mercury column, care must be taken to have the eudiom- 
 eter in a perpendicular position. This can be secured by 
 means of plumb-lines suspended from the ceiling and reach- 
 ing nearly to the table, by which the position of the eudiom- 
 eter can be adjusted. 
 
 OXYHYDROGEN BLOW-PIPE. 
 
 Experiment 52. Hold in the flame of the oxyhydrogen 
 blow-pipe successively a piece of iron wire, a piece of a steel 
 watch-spring, a piece of copper wire, a piece of zinc, a piece 
 of platinum wire. 
 
 Experiment 53. Cut a piece of lime of convenient size 
 and shape, say an inch long by three quarters of an inch wide, 
 and the same thickness. Fix it in position so that the flame 
 of the oxyhydrogen blow-pipe will play upon it. The light 
 is very bright, but by no means as intense as the electric light. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER V. 
 
 ORGANIC SUBSTANCES CONTAIN WATER. 
 
 Experiment 54. In dry test-tubes heat gently various or- 
 ganic substances as a piece of wood, fresh meat, fruits, vege- 
 tables, etc. 
 
 WATER OF CRYSTALLIZATION. 
 
 Experiment 55. Take some of the crystals of zinc sul- 
 phate obtained in Experiment 34. Spread them out on a 
 layer of filter-paper, and finally press two or three of them 
 between folds of the paper. Examine them carefully. They 
 appear to be quite dry, and in the ordinary sense they are 
 dry. Heat them in a dry tube, when it will be observed that 
 water condenses in the upper part of the tube, while the 
 
764 EXPERIMENTS TO ACCOMPANY CHAPTER V. 
 
 crystals lose their lustre, becoming white and opaque, and at 
 last crumbling to powder. 
 
 Experiment 56. Perform a similar experiment with some 
 gypsum, which is the natural substance from which " plaster 
 of Paris " is made. 
 
 Experiment 57. Heat a few small crystals of copper sul- 
 phate, or blue vitriol. In this case the loss of water is accom- 
 panied by a loss of color. After all the water is driven off, 
 the powder left behind is white. On dissolving it in water, 
 however, the solution will be seen to be blue ; and if the solu- 
 tion is evaporated until the substance is deposited, it will 
 again appear in the form of blue crystals. 
 EFFLORESCENT SALTS. 
 
 Experiment 58. Select a few crystals of sodium sulphate 
 which have not lost their lustre. Put them on a watch- 
 glass, and let them lie exposed to the air for an hour or two. 
 They soon lose their lustre, and undergo the changes noticed 
 in heating zinc sulphate. 
 
 DELIQUESCENT SALTS. 
 
 Experiment 59. Expose a few pieces of calcium chloride 
 to the air. Its surface will soon give evidence of the presence 
 of moisture, and after a time the substance will dissolve in 
 the water which is absorbed. 
 
 PURIFICATION OF WATER BY DISTILLATION". 
 
 Experiment 6O. In an apparatus like that shown in Fig. 
 39 distil a dilute solution of copper sulphate or some other 
 
 FIG. 39. 
 
 colored substance. A slow current of cold water must be 
 
METHOD OF DUMA& 765 
 
 kept running through the condenser by connecting the lower 
 rubber tube with a water-cock. When the water is boiled in 
 the large flask, the steam passes into the inner tube of the 
 condenser. As this is surrounded by cold water, the steam 
 condenses and the distilled water collects in the receiver. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER VI. 
 
 It would be well in this connection to determine the specific 
 gravity of some substance in the form of vapor. The princi- 
 pal methods for this purpose are those of Dumas, Gay Lussac, 
 Hof mann, and Victor Meyer. That of Dumas, which consists 
 in measuring the volume and determining the weight of the 
 vapor under observation, is the most accurate. The method 
 of Hofmann is a modification of that of Gay Lussac. It con- 
 sists in weighing a small quantity of the liquid the specific 
 gravity of whose vapor is to be determined, and, after intro- 
 ducing the liquid in a minute glass vessel into a eudiometer 
 over mercury, heating the eudiometer and its contents by 
 passing steam through a jacket surrounding it and measuring 
 the volume of vapor formed. The method of Victor Meyer 
 is used very commonly, especially when it is required to 
 -determine the specific gravity of the vapor of a substance 
 which boils at a high temperature. 
 
 METHOD OF DUMAS. 
 
 Experiment 61. In this method the liquid to be vaporized 
 is brought into a small balloon like that shown in Fig. 40. 
 The dry balloon is first weighed, and a small quantity of 
 liquid then introduced by gently heating the balloon and pat- 
 ting the point of its stem into the liquid, when, on cooling, the 
 liquid rises and enough is easily brought into the balloon in 
 this way. The balloon is now placed (in the position shown in 
 Pig. 41) in a bath of water, oil, or paraffin, according to the 
 boiling-point of the liquid. The bath is heated 30-40 above 
 the boiling-point of the liquid under examination. The air 
 is thus driven out and the balloon is filled with the vapor. 
 When vapor no longer escapes, the point of the stem is closed 
 by melting it with a mouth blow-pipe. The balloon is then 
 cleaned, dried, and weighed. The temperature of the bath 
 and the height of the barometer are observed at the time the 
 balloon is closed. The point of the stem is broken off under 
 
766 EXPERIMENTS TO ACCOMPANY CHAPTER VI. 
 
 mercury, when the mercury rises and fills the balloon. By 
 pouring the mercury out into a graduated cylinder the ca- 
 pacity of the balloon is determined. The specific gravity of 
 the vapor is calculated by the aid of the formula 
 
 
 B + p)(l + 0.00366 X 
 
 vh v X 0.001293 
 
 in which 
 
 B = weight of balloon at t and h mm. ; 
 B 1 = " " " with vapor, at t t and h t mm. ; 
 
 v = capacity of the balloon in cubic centimeters; 
 0.001293 = weight of 1 cc, air at and 760 mm. ; 
 p = weight of air in balloon at t and h mm. 
 
 FIG. 40. 
 
 FIG. 41. 
 
 METHOD OF VICTOR MEYER. 
 
 Experiment 62. In this method a known 
 weight of substance is converted into vapor, and 
 the volume of vapor formed is determined by 
 measuring the volume of air which it displaces. 
 The apparatus consists of an outer cylindrical 
 vessel A, Fig. 42, and an inner vessel B, which 
 is connected with a tube 0. The vessel B has a 
 capacity of about 100 cc., and is about 200 
 mm. long. The tube C, with its funnel-shaped 
 end E, is about 600 mm. long. First, a small 
 quantity of some substance with a boiling-point ^, 
 high enough to secure the complete conversion FIG 42. 
 into vapor of the substance to be studied, is put in the 
 bottom of the vessel A, and a little ignited asbestos or dry 
 
OZONE. 767 
 
 mercury in the bottom of the vessel B. The substance in A 
 is now heated to boiling, and E is closed with a rubber stop- 
 per. After a time the temperature of the air in B is raised to 
 that of the vapor in A, and no more escapes from the tube D. 
 When this condition of equilibrium is reached, a small weighed 
 quantity of the substance under examination is dropped into 
 the vessel B, the stopper being removed from E and quickly 
 replaced. The substance is converted into vapor, and displaces 
 an equivalent volume of air, and this displaced air is collected 
 over water in the measuring-tube placed over the end of D. 
 When no more air escapes, the volume, is determined in the 
 usual way. The specific gravity of the substance is calculated 
 by the aid of the following formula: 
 
 _ s , (1 4- 0.00366 X Q760 
 " ' (B w)Vx 0.001293' 
 
 in which 0.001293 is the weight of 1 cc. air in grams at 760 
 mm. and 0; and, further, 
 
 S = weight of substance taken; 
 
 t = temperature of the room, or of the water in the measur- 
 ing apparatus; 
 B = height of barometer; 
 w = tension of aqueous vapor; 
 V = observed volume of air; 
 
 or, the formula can be simplified by division, when it takes 
 this form: 
 
 (1 + 0.00366 X 0587,780 
 (B-w)V 
 
 The above is the simplest form of apparatus used. To avoid 
 opening and shutting the vessel in order to introduce the sub- 
 stance, an arrangement has been devised for holding the sub- 
 stance below the stopper, until the proper temperature is 
 reached, and then releasing it without disturbing the stopper. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER VII. 
 OZONE. 
 
 Experiment 63. Put a few sticks of ordinary phosphorus 
 on the bottom of a good-sized bottle with a wide mouth, and 
 
768 EXPERIMENTS TO ACCOMPANY CHAPTER VIII. 
 
 partly cover the phosphorus with water. In a short time the 
 odor of ozone will be perceptible, and the gas can also be de- 
 tected by means of strips of paper which have been moistened 
 with a dilute solution of potassium iodide and starch-paste. 
 See whether such papers are changed in the air ? What is the 
 cause of the change? If convenient, examine the air in the 
 neighborhood of a frictional electrical machine, and see 
 whether it causes the papers to change color. 
 
 HYDROGEN DIOXIDE. 
 
 Experiment 64. Finely powder some barium dioxide, and 
 add some of it to dilute sulphuric acid. Filter from the pre- 
 cipitated barium sulphate, and with the solution try the fol- 
 lowing reactions : 
 
 Heat some in a test-tube. What takes place? Add to an- 
 other small portion a little of a dilute solution of potassium 
 permanganate. To another portion add a little finely pow- 
 dered manganese dioxide. What is given off ? To a dilute 
 solution contained in a small stoppered cylinder add a few 
 drops of a dilute solution of potassium dichromate, and quick- 
 ly add ether, and shake the cylinder thoroughly. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER VIII. 
 PREPARATION OF CHLORINE. 
 
 Experiment 65. Pour 2 or 3 cc. concentrated sulphuric 
 acid on a gram or two of common salt in a test-tube. A gas 
 will be given off which forms dense white fumes in the air and 
 has a sharp, penetrating taste and smell. This is hydrochloric 
 acid gas. 
 
 Experiment 66. Pour 2 or 3 cc. concentrated sulphuric acid 
 on a few grams of manganese dioxide in a test-tube. Heat, 
 and examine the gas given off. jQonvince yourself that it is 
 oxygen. 
 
 Experiment 67. Mix 2 grams manganese dioxide and 2 
 grams common salt. Pour 4 to 5 cc. dilute sulphuric acid on 
 the mixture in a test-tube. This experiment should be per- 
 formed under a hood in which the draught is good, as the gas 
 which is given off is not only disagreeable, but irritating to 
 the respiratory organs. Notice the color and odor of the gas. 
 [Does it support combustion ? Does it burn ?] 
 
PREPARATION OF CHLORINE. ?69 
 
 The best way to make chlorine is the following : Mix 5 parts 
 coarsely granulated manganese dioxide and 5 parts coai 
 ly granulated common salt. 
 Make a mixture of 12 parts 
 concentrated sulphuric acid 
 and 6 parts water. Let this 
 mixture cool down to the tem- 
 perature of the room, and then 
 pour it upon the mixture of 
 salt and manganese dioxide. 
 Gently heat on a sand-bath, 
 and a regular current of chlo- 
 rine will be given off. The gas 
 is collected by displacement of 
 air in a dry glass vessel. The 
 apparatus for the purpose is 
 arranged as shown in Fier. 43. 
 
 FIG 43. 
 
 The delivery-tube should reach 
 
 to the bottom of the collecting vessel, and the mouth of the 
 vessel should be covered with a piece of paper to prevent cur- 
 rents of air from carrying away the chlorine. As the gas col- 
 lects in the vessel the experimenter can judge of the quantity 
 present by means of the color. 
 
 Experiment 68. Collect six or eight dry cylinders or bot- 
 tles full of chlorine. Make the gas from about 30 grams of 
 manganese dioxide, using the other substances in the propor- 
 tions already stated. 
 
 (1) Introduce into one of the vessels containing chlorine a 
 little finely powdered antimony. 
 
 (2) Into a second vessel put a few pieces of heated thin 
 copper- foil. 
 
 (3) Into a third vessel put a piece of paper with some writ- 
 ing on it, some flowers, and pieces of cotton print. The sub- 
 stances used must be moist. 
 
 (4) Into a fourth vessel put a dry piece of the same cotton 
 print as that used in the previous experiment. 
 
 What conclusions do the results of the above experiments 
 justify as to the conduct of chlorine? 
 
 Experiment 69. Cut a piece of filter-paper about an inch 
 wide and six to eight inches long. Pour on this some ordinary 
 oil of turpentine previously warmed slightly. Introduce this 
 into one of the vessels of chlorine. A flash of flarne is noticed, 
 
and a dense black cloud is formed. The action in this case 
 is due to the great affinity of chlorine for hydrogen. Oil of 
 turpentine consists of carbon and hydrogen. The main action 
 of the chlorine consists in extracting the hydrogen and leaving 
 the carbon. The experiment is interesting chiefly in so far 
 as it illustrates the general tendency of chlorine to act upon 
 vegetable substances. 
 
 CHLORINE DECOMPOSES WATER IN THE SUNLIGHT. 
 
 Experiment 70. Seal the end of a glass tube about a metre 
 (or about a yard) long and about 12 mm. (| inch) 
 internal diameter. Fill this with a strong solu- 
 tion of chlorine in water. Invert it as shown in 
 Fig. 44, in a shallow vessel containing some of 
 the same solution of chlorine in water. Place 
 the tube in direct sunlight. Gradually bubbles 
 of gas will be seen to rise and collect in the up- 
 per end, and the color of the solution, which is 
 at first greenish yellow, like that of chlorine, 
 disappears. The gas can be shown to be oxygen. 
 
 CHLORINE HYDRATE. 
 
 Experiment 71. Conduct chlorine into a flask 
 containing water cooled down to about 2 or 3 
 Centigrade. If crystals are formed remove some 
 by filtering out-of-doors if the weather is cold. Expose some 
 of the crystals on filter-paper under a hood in the laboratory. 
 What changes have taken place ? 
 
 FORMATION OF HYDROCHLORIC ACID. 
 
 Experiment 72. Light a jet of hydrogen in the air and 
 carefully introduce it into a vessel containing chlorine. It 
 will continue to burn, but the flame will not appear the same. 
 A gas will be given off which forms clouds in the air. This 
 gas has a sharp, penetrating taste and smell. 
 
 Experiment 73. Half fill a small, wide-mouthed cylinder 
 over hot water with chlorine gas. Then fill it with hydrogen. 
 The direct sunlight must not shine upon the cylinder while 
 it contains the mixture. Turn it mouth upward and apply 
 a flame. 
 
PREPARATION OF HYDROCHLORIC ACID. 771 
 
 PREPARATION OF HYDROCHLORIC ACID. 
 
 - 
 
 Experiment 74. Arrange an apparatus as shown in Fig. 45i 
 
 FIG. 45. 
 
 "Weigh out 5 parts common salt, 5 parts concentrated sul- 
 phuric acid, and 1 part water. Mix the acid and water, tak- 
 ing the usual precautions; let the mixture cool down to the 
 ordinary temperature, and then pour it on the salt in the 
 flask. For the purposes of the experiment take about 20 
 grams of salt. Now heat the flask gently, and the gas will 
 be regularly evolved. Conduct it at first through water con- 
 tained in the two Wolff's bottles until what passes over is 
 all absorbed in the first bottle. The reason why gas at first 
 bubbles through all the bottles is, that the apparatus is full 
 of air, which is first driven out. When the air has been dis- 
 placed, the gas is all absorbed as soon as it comes in contact 
 with the water. After the gas has passed for ten to fifteen 
 minutes, disconnect at A. Notice the fumes. These become 
 denser by blowing the breath on them. Why? Apply a 
 lighted match to the end of the tube. Does the gas burn? 
 Collect some of the gas in a dry cylinder by displacement 
 of air, as in the case of chlorine. The specific gravity of the 
 gas being 1.26, the vessel must of course be placed with the 
 mouth upward. That the gas is colorless and transparent is 
 shown by the appearance of the generating flask, which is 
 filled with the gas. Insert a burning stick or candle in the 
 cylinder filled with the gas v Reconnect the gen erat ing-flask 
 with the series of bottles containing water, and let the pro- 
 cess continue until no more gas comes over. The reaction 
 represented in the equation 
 
 H 2 S0 4 = Na 2 S0 4 + 2HC1 
 
772 EXPERIMENTS TO ACCOMPANY CHAPTER IX. 
 
 is now complete. Disconnect the flask, and after it has 
 cooled down pour water on the contents ; when the substance 
 is dissolved filter it and evaporate to such a concentration 
 that, on cooling, the sodium sulphate is deposited. Pour off 
 the liquid and dry the solid substance by means of filter- 
 paper. Compare the substance with the common salt which, 
 you put in the flask before the experiment. What proofs 
 have you that the two substances are not the same? Heat 
 a small piece of each in a dry tube closed at one end. What 
 differences do you notice ? Treat a small piece of each in a 
 test-tube with sulphuric acid. What difference do you no- 
 tice ? If in the experiment we should recover all the sodium 
 sulphate formed, how much should we have ? Put about 50 
 cc. of the liquid from the first Wolff's bottle in a porcelain 
 evaporating-dish. Heat over a small flame just to boiling. Is 
 hydrochloric acid given off ? Can all the liquid be driven off 
 by boiling ? Try the action of the solution on some iron 
 filings. What is given off ? Add some to a little granulated 
 zinc in a test-tube. What is given off ? Add a little to 
 some manganese dioxide in a test-tube. What is given off ? 
 Add ten or twelve drops of the acid to 2 to 3 cc. water in a 
 test-tube. Taste the dilute solution. It has what is called a 
 sour or acid taste, the two terms being practically synony- 
 mous. Add a drop or two of a solution of blue litmus, 
 or put into it a piece of paper colored blue with litmus. 
 What change takes place? Litmus is a vegetable color pre- 
 pared for use as a dye. Other vegetable colors are changed 
 by hydrochloric acid. Steep a few leaves of red cabbage in 
 water. Add a few drops of the solution thus obtained to di- 
 lute hydrochloric acid. Is there any change in color ? The 
 color will be restored in each case by adding a few drops of a 
 solution of caustic soda. In what experiment has caustic 
 soda been obtained ? What relation does it bear to water ? 
 To the dilute solution of hydrochloric acid add drop by drop 
 a dilute solution of caustic soda. Is the acid taste destroyed ? 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER IX. 
 
 CHLORIC ACID AKD POTASSIUM CHLORATE. 
 
 Experiment 75. Dissolve 40 grams (or about 1J ounces) 
 caustic potash in 100 cc. water in a beaker-glass, and pass 
 chlorine into it. When chlorine passes freely through the 
 
PERCHLORIC ACID. 
 
 773 
 
 solution, thus indicating that it is no longer absorbed, stop 
 the action. After boiling filter the solution and allow it to 
 cool, when crystals of potassium chlorate will be deposited, 
 mixed with a little potassium chloride. Eecrystallize from a 
 little water. Filter off the crystals and dry them. What evi- 
 dence have you that the substance is potassium chlorate? Does 
 it give off oxygen when heated ? In a dry test-tube pour two 
 or three drops of concentrated sulphuric acid on a small crys- 
 tal of the substance. Do the same with a piece of potassium 
 chlorate from the laboratory bottle. Hold the mouth of the 
 test-tube away from the face. What is noticed in each case ? 
 Evaporate the solution from which the crystals of potassi- 
 um chlorate have been removed. On allowing it to cool 
 crystals will again be deposited. Take them out and recrys- 
 tallize them. Does this substance give off oxygen when 
 heated ? Does it give off a gas when treated with sulphuric 
 acid ? Is this gas colored ? Is it hydrochloric acid ? How 
 do you know that it is ? If the gas is hydrochloric acid, what 
 is the solid substance from which it is formed ? And what 
 is left in the test-tube ? 
 
 Experiment 76. Mix 10 
 grams fresh quick-lime with 
 20 cc. water. After the slak- 
 ing is over, pass chlorine into 
 it until the gas is no longer 
 absorbed. Put the powder 
 thus formed in a flask ar- 
 ranged as shown in Fig. 46. 
 Pour a mixture of equal parts 
 of sulphuric acid and water 
 slowly through the funnel- 
 tube. Collect by displacement 
 of air the gas given off. What 
 evidence have you that the. 
 gas is chlorine ? y IG . 40. 
 
 PERCHLORIC ACID. 
 
 Experiment 77. Make potassium perchlorate as follows : 
 Gently heat 50 to 100 grams potassium chlorate until after 
 having been liquid it becomes thick and pasty, and gas is not 
 given off without raising the temperature. After cooling, 
 break up the mass and treat it with cold water. This dissolves 
 
774 EXPERIMENTS TO ACCOMPANY CHAPTER X. 
 
 out the potassium chloride and leaves the perchlorate, which 
 can then be crystallized from hot water. After the crystallized 
 salt is dried it is decomposed by sulphuric acid. To effect 
 this decomposition, the finely powdered salt (10 parts) is 
 treated in a retort with 20 parts of pure sulphuric acid which 
 is free from nitric acid and diluted with -fa its volume of water. 
 The retort is connected with a receiver which can be well 
 cooled. The mixture is heated, and when the perchloric acid 
 begins to come over, the heat is so regulated that the tem- 
 perature does not rise above 140. When the mixture has 
 become colorless the operation is ended. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER X. 
 
 NEUTRALIZATION OF ACIDS AND BASES ; FORMATION OF 
 
 SALTS. 
 
 Experiment 78. Make dilute solutions of nitric, hydro- 
 chloric, and sulphuric acids (1 part dilute acid, such as is used 
 in the laboratory, to 50 parts water), 
 and of caustic soda and caustic pot- 
 ash (about 1 ^ram to 200 cc. of 
 water). Measure off about 20 cc. of 
 one of the acid solutions. Add a 
 few drops of a solution of blue 
 litmus. Gradually add to the meas- 
 ured quantity of acid sufficient di- 
 lute caustic soda to cause the red 
 color just to change to blue. As long 
 as the solution is red it is acid. 
 "When it turns blue it is alkaline. 
 At the turning-point it is neutral. 
 The operation is best carried on by 
 means of a burette, which is a gradu- 
 ated tube with an opening from 
 which small quantities can be poured. 
 A convenient shape is that repre- 
 sented in Fig. 47. At the lower end 
 is a small opening. The flow of the 
 liquid from the burette is controlled 
 by means of a small pinch-cock. It 
 will require some practice to enable the student to know ex- 
 
 FIG. 47. 
 
STUDY OF THE PRODUCTS FORMED. 775 
 
 actly when the red color disappears and the blue appears, but 
 with practice the point can be discerned with great accuracy. 
 Should too much alkali be allowed to get into the acid, add a 
 small measured quantity of the acid from another burette. 
 Having in one experiment determined how much of the solu- 
 tion of alkali is required to cause the red color to change to 
 blue in operating on a given quantity of the acid solution, try 
 the experiment again, using a different quantity of the acid 
 solution. If the results of several experiments with the same 
 acid and alkali are recorded, it will be found that there is 
 a definite ratio between the quantities of acid and alkali so- 
 lution required to neutralize one another. If, for example, 
 15 cc. of the alkali solution are required to neutralize 20 cc. of 
 the acid solution, 18 cc. of the alkali solution will be required 
 to neutralize 24 cc. of the acid solution, 30 cc. to neutralize 
 40 cc., etc. In other words, in order to neutralize a given 
 quantity of an acid, a definite quantity of an alkali is necessary. 
 Perform similar experiments with the other acids. Afterwards 
 carefully examine the numerical results. Suppose it should 
 require 15 cc. of the caustic-soda solution or 12 cc. of the 
 caustic-potash solution to neutralize 20 cc. of the hydrochloric- 
 acid solution. Compare the quantities of these alkali solu- 
 tions necessary to neutralize equal quantities of the other acids. 
 What conclusion is justified with reference to the act of neu- 
 tralization ? 
 
 STUDY OF THE PKODUCTS FOBMED. 
 
 Experiment 79. Dissolve about 10 grams caustic soda in 
 100 cc. water. Add hydrochloric acid slowly, examining the 
 solution from time to time by means of a piece of paper col- 
 ored blue with litmus. As long as the solution is alkaline it 
 will cause no change in the color of the paper. The instant 
 the point of neutralization is passed, the solution changes the 
 color of the paper to red; when exactly neutral, it will neither 
 change the blue to red, nor, if the color is changed to red 
 by means of another acid, will it change it back again. When 
 this point is reached, evaporate to complete dryness on the 
 water-bath, and see what is left. Taste the substance. Has 
 it an acid taste? Does it suggest any familiar substance? If 
 it is sodium chloride, how ought it to conduct itself when 
 treated with sulphuric acid? Does it conduct itself in this 
 way ? Satisfactory evidence can be given that the substance 
 
776 EXPERIMENTS TO ACCOMPANY CHAPTER XII. 
 
 is sodium chloride. It is not an acid nor an alkali. It is 
 neutral. 
 
 Experiment 80. Perform a similar experiment, using 
 dilute nitric acid and caustic soda. What evidence have you 
 that the product in this case is different from caustic soda ? 
 
 Experiment 81. Perform similar experiments with dilute 
 sulphuric acid and caustic soda; with sulphuric acid and 
 caustic potash; with nitric acid and caustic potash; with hy- 
 drochloric acid and caustic potash. Dry and examine the 
 product carefully in each case; and keep for future study what 
 is not used in these experiments. 
 
 FOR CHAPTER XI. 
 
 A large table of the Natural System of the Elements, like 
 that on page 151, should be hung up in a conspicuous place 
 in the laboratory. It would be well also to have such a table 
 pasted upon a cylinder which can be revolved on its axis, so 
 that the continuity of the system may be impressed upon the 
 mind. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XII. 
 PREPARATION OF BROMINE. 
 
 Experiment 82. Mix together 3.5 grams potassium bro- 
 mide and 7 grams manganese dioxide. Put the mixture into a 
 500 cc. flask; connect with a condenser (see Fig. 39). Mix 15 
 cc. concentrated sulphuric acid and 90 cc. water. After cool- 
 ing pour the liquid on the mixture in the flask. Gently heat, 
 when bromine will be given off in the form of vapor. A part 
 of this will condense and collect in the receiver. Perform this 
 experiment under a hood with a good draught. 
 
 HYDROBROMIC ACID. 
 
 Experiment 83. In a small porcelain evaporating-dish 
 put a few crystals, of potassium bromide. Pour on them a 
 few drops of concentrated sulphuric acid. The white fumes 
 of hydrobromic acid and the reddish-brown vapor of bromine 
 are noticed. Treat a few crystals of potassium or sodium 
 chloride in the same way. What difference is there between 
 the two cases ? 
 
 The preparation of hydrobromic acid may be shown in the 
 lecture-room as follows: 
 
HTDROBROMIO ACID. 777 
 
 Experiment 84. Arrange an apparatus as shown in Fig. 
 48. In the flask put 1 part red phosphorus and 2 parts water. 
 
 FIG. 48. 
 
 Let 10 parts bromine gradually drop into the flask from the 
 glass-stoppered funnel. Pass the gas through a U-tube 
 loosely packed with asbestos containing red phosphorus in 
 order to free the hydrobromic acid from bromine, which to 
 some extent passes over with it. Collect some of the gas in 
 water, and examine the solution. How does the gas act when 
 allowed to escape in the air ? Fill a cylinder with the gas in 
 the same way as was done with hydrochloric acid, and fill an- 
 other with chlorine. While covered with glass plates bring 
 their mouths together. Then withdraw the plates. What 
 change is observed ? What is this due to ? 
 
 Experiment 85. To a dilute solution of sodium hydroxide 
 add bromine water made by shaking up a little liquid bromine 
 in a bottle with water. What change takes place ? Add sul- 
 phuric acid until the liquid shows an acid reaction. What 
 takes place? The changes here referred to are perfectly anal- 
 ogous to those which would *take place if chlorine were used 
 instead of bromine. Shake a solution containing a little free 
 bromine with ether ; with chloroform ; with carbon disulphide. 
 What changes do you observe ? 
 
778 EXPERIMENTS TO ACCOMPANY CHAPTER XII. 
 
 IODINE. 
 
 Experiment 86. Mix about 2 grams of sodium or potas- 
 sium iodide and 4 grams manganese dioxide. Treat with a 
 little concentrated sulphuric acid in a one to two liter flask. 
 Heat gently on a sand-bath. Gradually the vessel will be filled 
 with the beautiful colored vapor of iodine. In the upper parts 
 of the flask some of the iodine will be deposited in the form 
 of crystals of a grayish-black color. 
 
 Experiment 87. Make solutions of iodine in water, in 
 alcohol, and in a water solution of potassium iodide. Use 
 small quantities in test-tubes. 
 
 Experiment 88. Dissolve a piece of potassium iodide the 
 size of a small pea in about 100 cc. water in a stoppered cylin- 
 der. Add enough carbon disulphide to make a layer about 
 an inch thick at the bottom of the cylinder. Shake the two 
 liquids together. Does the carbon disulphide become colored ? 
 Add a drop of chlorine water and shake again. What differ- 
 ence do you observe in the two cases ? Explain this. Try 
 the same experiment, using chloroform instead of carbon 
 disulphide. 
 
 IODINE CAN BE DETECTED BY MEANS OF ITS ACTION UPON 
 STAECH-PASTE. 
 
 Experiment 89. Make some starch-paste by covering a 
 few grains of starch in a porcelain evaporating-dish with cold 
 water, grinding this to a paste, and pouring 200-300 cc. boil- 
 ing-hot water on it. After cooling add a little of this paste 
 to a dilute water solution of iodine. The solution will turn 
 blue if the conditions are right. Now add a little of the paste 
 to a diluted water solution of potassium iodide. Is there any 
 change? Add a drop or two of a solution of chlorine in water. 
 Why the difference ? Will not chlorine water alone act this 
 way toward starch-paste ? 
 
 ACTION OF SULPHURIC ACID UPON POTASSIUM IODIDE. 
 
 Experiment 90. Bring a piece of potassium iodide the 
 size of a pea in a dry test-tube ; add one drop of water and 
 three or four drops of concentrated sulphuric acid ; the salt 
 becomes brown ; heat gently ; violet-colored vapor escapes, 
 and with it a gas with an odor like that of rotten eggs. At 
 
IODIC ACID PROPERTIES OF SULPHUR. 779 
 
 the same time a yellow coating appears on the inside of the 
 tube above the acid. Add five or six drops more of the acid 
 and continue to heat gently. The bad odor first noticed dis- 
 appears gradually, and another, quite different odor, irritating 
 to the throat is now perceptible. This is sulphur dioxide, 
 
 so, 
 
 IODIC ACID. 
 
 Experiment 91. Pass chlorine into a test-tube containing 
 iodine in suspension in water ; or add chlorine water. What 
 becomes of the iodine? 
 
 Experiment 92. Add chlorine water to a dilute solution 
 of potassium iodide, and note the successive changes. 
 
 Experiment 93. Dissolve iodine in caustic soda. Add 
 an acid to the solution. Explain the changes. 
 
 HYDROFLUORIC ACID. 
 
 Experiment 94. In a lead or platinum vessel put a few 
 grams (5-6) of powdered fluor-spar and pour on it enough 
 concentrated sulphuric acid to make a thick paste. Cover the 
 surface of a piece of glass with a thin layer of wax or paraffin, 
 and through this scratch some letters or figures, so as to leave 
 the glass exposed where the scratches are made. Put the 
 glass over the vessel containing the fluor-spar, and let it stand 
 for some hours. Take off the glass, scrape off the coating, 
 and the figures which were marked through the wax or paraf- 
 fin will be found etched on the glass. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XIII. 
 PROPERTIES OF SULPHUR. 
 
 Experiment 95. Distil about 10 grams roll sulphur from 
 an ordinary glass retort. What changes in color and in con- 
 dition take place ? Collect the liquid sulphur formed by the 
 condensation of the vapor in a beaker-glass containing cold 
 water. 
 
 Experiment 96. Treat some powdered roll sulphur with 
 carbon disulphide and filter. Does it all dissolve ? Try the 
 same experiment with flowers of sulphur. Does this all dis- 
 solve? Put the solutions together and allow to evaporate. 
 Examine the crystals deposited. Compare them with some 
 natural crystals of sulphur. See whether one of the crystals 
 will completely dissolve in carbon disulphide. 
 
rtfO EXPERIMENTS TO ACCOMPANY CHAPTER XIII. 
 
 Experiment 97. In a covered sand or Hessian crucible 
 melt about 25 grams of roll sulphur. Let it cool slowly, and 
 when a thin crust has formed on the surface make a hole 
 through this and pour out the liquid part of the sulphur. 
 What is left ? Compare with the crystals formed in the last 
 experiment. Lay the crucible aside, and in the course of a 
 few days again examine the crystals. What changes, if any, 
 have taken place? 
 
 Experiment 98. Add hydrochloric acid to a solution of 
 sodium thiosulphate. What takes place ? 
 
 Experiment 99. In a wide test-tube heat some sulphur 
 to boiling. Introduce into it small pieces of copper-foil or 
 sheet copper. Or hold a narrow piece of sheet copper so that 
 the end just dips into the boiling sulphur. 
 
 Experiment 100. Dissolve some sulphur in concentrated 
 caustic soda. In what form is the sulphur in the solution ? 
 
 HYDROGEN SULPHIDE. 
 
 Experiment 101. Arrange an apparatus as shown in Fig. 
 49. Put a small handful of the sulphide of iron, FeS, in the 
 
 FIG. 49. 
 
 flask, and pour dilute sulphuric acid upon it. Pass the 
 evolved gas through a little water contained in the wash cylin- 
 der A. Pass some of the gas into water. [What evidence 
 have you that it dissolves ?] Collect some by displacement of 
 air. Its specific gravity is 1.178. Set fire to some of the gas 
 contained in a cylinder. In this case the air has not free 
 
MANUFACTURE OF SULPHURIC ACID. 781 
 
 access to the gas, and the combustion is not complete. The 
 hydrogen burns to form water, while a part of the sulphur is. 
 deposited upon the inside walls of the cylinder. If there is 
 free access of air, the sulphur burns to sulphur dioxide and 
 the hydrogen to water. 
 
 Make a solution of the gas in water in the usual way. Put 
 some of this in a bottle and set it aside, and in the course of 
 a few days examine it again. Boil another portion for a time 
 in a test-tube, and note the changes. Pass a little of the gas 
 through concentrated sulphuric acid contained in a test-tube, 
 and note the changes. Moisten strips of paper with dilute 
 solutions of lead nitrate, copper sulphate, stannous chloride, 
 .antimony chloride, and mercuric chloride ; and expose these 
 papers in turn to the gas. What changes take place ? Eepeat 
 Experiment 90, and see whether one of the gases given off 
 produces similar changes. 
 
 Experiment 102. Pass hydrogen sulphide successively 
 through solutions containing a little lead nitrate, cadmium 
 nitrate, and arsenic prepared by dissolving a little white 
 arsenic, or arsenic trioxide, As 2 3 , in dilute hydrochloric acid. 
 What, action takes place in each case ? The formula of lead 
 nitrate is Pb(N0 3 ) 2 ; that of cadmium nitrate, Cd(N0 3 ) 2 ; and 
 that of the chloride of arsenic in solution is As01 3 . The cor- 
 responding sulphides are represented by the formulas PbS, 
 dS, and As 2 S 3 . 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XIV. 
 MANUFACTURE OF SULPHURIC ACID. 
 
 Experiment 103. The manufacture of sulphuric acid can 
 be illustrated in the laboratory by means of the apparatus 
 represented in Fig. 50. This consists of a large balloon flask 
 fitted with a stopper having five openings. By means of tubes 
 it is connected with three small flasks. One of these, a, con- 
 tains water for the purpose of providing a current of steam ; 
 another, c, contains copper-foil and concentrated sulphuric 
 acid, which give sulphur dioxide when heated ; and the third, 
 b, contains copper-foil and dilute nitric acid, which give oxides 
 of nitrogen, mainly nitric oxide, NO. When the nitric oxide 
 comes in contact with the air it combines with oxygen, form- 
 ing nitrogen trioxide and nitrogen peroxide; and when steam 
 and sulphur dioxide are admitted to the flask the reactions 
 
782 EXPERIMENTS TO ACCOMPANY CHAPTER XIV. 
 
 involved in the manufacture of sulphuric acid take place. 
 By means of a pair of bellows attached at d air is supplied. 
 If air is not forced in, the gases become colorless, owing to 
 
 FIG. 50. 
 
 complete reduction of the oxides of nitrogen to the form of 
 nitric oxide, NO, which is colorless. If steam is not admitted 
 the walls of the vessel become covered with crystals of nitro- 
 syl-sulphuric acid. This is, however, decomposed by an excess 
 of steam. 
 
 Experiment 104. Into a vessel containing ordinary con- 
 centrated sulphuric acid introduce small sticks of wood, pieces 
 of paper, and various other organic substances, and note the 
 result. The charring effect is particularly well shown by 
 adding the acid drop by drop to a concentrated solution of 
 sugar, or to molasses, and stirring. 
 
 Experiment 105. Sulphuric acid is detected in analysis 
 by adding barium cloride to its solution, when insoluble bar- 
 ium sulphate is formed. 
 
 H 2 S0 4 + Bad, = BaS0 4 + 2HC1. 
 
 Other insoluble sulphates are those of strontium and lead ; 
 and calcium sulphate is difficultly soluble. To a dilute solu- 
 tion of sulphuric acid or of any soluble sulphate, add in test- 
 tubes barium chloride, strontium nitrate, and lead nitrate. 
 
SULPHUROUS ACID AND SULPHUR DIOXIDE. 783 
 
 SULPHUROUS ACID AND SULPHUR DIOXIDE. 
 
 Experiment 106. Put eight or ten pieces of sheet copper, 
 one to two inches long and about half an inch wide, in a 500 
 cc. flask ; pour 15 to 20 cc. concentrated sulphuric acid on it. 
 On heating, sulphur dioxide will be evolved, The moment 
 the gas begins to come off, lower the flame, and keep it at 
 such a height that the evolution is regular and not too active. 
 Pass some of the gas into a bottle containing water. The 
 solution in water is called sulphurous acid. 
 
 Experiment 107. Pass sulphur dioxide into a moderately 
 dilute solution of potassium hydroxide, until the solution is 
 saturated. What is then contained in the solution? To a 
 little of it add hydrochloric acid. What takes place? 
 
 Experiment 108. Try the effect of heating concentrated 
 sulphuric acid with charcoal, and with sulphur. 
 
 Experiment 109. Collect by displacement of air some of 
 the gas made in Experiment 106. Does it burn ? or does it 
 support combustion? 
 
 Experiment 110. Pass some of the gas through a bent- 
 glass tube surrounded by a freezing mixture of salt and ice. 
 Tubes provided with glass stop-cocks are made for such pur- 
 
 FIG. 51. 
 
 poses. They generally have the form represented in Fig. 51. 
 If the tube is taken out of the freezing mixture, the liquid 
 sulphur dioxide changes rapidly to gas, if the tube is open. 
 
 Experiment 111. Burn a little sulphur in a porcelain 
 crucible under a bell-jar. Place over the crucible on a tripod 
 some flowers. In the atmosphere of sulphur dioxide the 
 flowers will be bleached. 
 
 SULPHUROUS ACID is A SEDUCING AGENT. 
 
 Experiment 112. To a dilute solution of potassium iodide 
 in a test-tube gradually add chlorine water until the solution 
 
784 EXPERIMENTS TO ACCOMPANY CHAPTER XV. 
 
 becomes clear and colorless. Now add a solution of sulphur- 
 ous acid. At first iodine is deposited, but on further addi- 
 tion of sulphurous acid it dissolves again. Explain all the 
 changes. 
 
 SULPHUR TRIOXIDE. 
 
 Experiment 113. Heat a little fuming sulphuric acid 
 gently in a test-tube. What takes place ? Put a little of the 
 acid (5-10 cc.) in a small dry retort provided with a glass 
 stopper and connect with a dry glass receiver. Heat the re- 
 tort gently, and keep the receiver cool. By means of a dry 
 glass rod take out some of the substance which collects in the 
 receiver and put it in water. Lay a little of it on a piece of 
 wood and on a piece of paper. 
 
 Experiment 114. Prepare finely divided platinum by 
 moistening some fine asbestos with a solution of platinic 
 chloride and heating to redness in a porcelain crucible. The 
 substance thus obtained is known as platinized asbestos, 
 Now arrange an apparatus so that both oxygen and sulphur 
 dioxide can be passed together through a tube of hard glass 
 as represented in Fig. 52. First pass the two dried gases 
 
 0- 
 
 S0 2 
 
 FIG. 52. 
 
 together through the empty tube and heat a part of the tube 
 by means of a burner. Is there any evidence of combination ? 
 Now stop the currents of the gases, let the tube cool down, 
 and introduce a small layer of the platinized asbestos. Pass 
 the dried gases over the heated asbestos. What takes place? 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XV. 
 PREPARATION OF NITROGEN. 
 
 Experiment 115. Place a good-sized stoppered bell- jar 
 over water in a pneumatic trough. In the middle of a flat 
 cork about three inches in diameter fasten a small porcelain 
 crucible, and place this on the water in the trough. Put in 
 
ANALYSIS OF AIR. 
 
 785 
 
 it a piece of phosphorus about twice the size of a pea, and set 
 fire to it. Quickly place the bell-jar over it. At first some 
 air will be driven out of the jar. The burning will continue 
 for a short time, and then gradually grow less and less active, 
 finally stopping. On cooling, it will be found that the 
 volume of gas is less than four fifths the original volume, for 
 the reason that some of the air was driven out of the vessel at 
 the beginning of the experiment. Before removing the 
 stopper of the bell-jar see that the level of the liquid outside 
 is the same as that inside. Try the effect of introducing suc- 
 cessively several burning bodies into 'the nitrogen, as, for 
 example, a candle, a piece of sulphur, phosphorus, etc. 
 
 Experiment 116. Place a live mouse in a trap in a bell- 
 jar over water. When the oxygen is used up the mouse will 
 die. After the animal gives plain signs of discomfort, it may 
 be revived by taking away the bell-jar and giving it a free 
 supply of fresh air. 
 
 Experiment 117. Pass air slowly over copper contained in' 
 a tube heated to redness and collect 
 the gas which passes through. Does 
 it act like nitrogen ? 
 
 Experiment 118. In a good-sized 
 Wolff's bottle provided with a safety- 
 funnel and delivery-tube as shown in 
 Fig. 53 put some copper-turnings 
 and pour upon them concentrated 
 ammonia, but not enough to cover 
 them. Close the delivery-tube by 
 means of a pinch-cock; and let the 
 vessel stand. What evidence of ac- 
 tion is there? After a time, force 
 some of the gas out of the bottle by 
 pouring water through the funnel, 
 and opening the delivery-tube. Does FIG. 53. 
 
 the gas act like nitrogen ? 
 
 ANALYSIS OF AIR. 
 
 Experiment 119. Arrange an apparatus as in Fig. 25. 
 Instead of a plain tube, use one graduated into cubic centi- 
 meters. Enclose 60 to 80 cc. air in the tube over water. 
 Arrange the tube so that the level of the water inside and 
 outside is the same. Note the temperature of the air and the 
 
786 EXPERIMENTS TO ACCOMPANY CHAPTER XV. 
 
 height of the barometer. Reduce the observed volume to 
 standard conditions. Now introduce a piece of phosphorus, 
 as in Experiment 26, and allow it to stand for twenty-four 
 hours. Draw out the phosphorus. Again arrange the tube 
 so that the level of the water inside is the same as that out- 
 side. Make the necessary corrections for temperature, pres- 
 sure, and the tension of aqueous vapor. It will be found that 
 the volume has diminished considerably, but that about four 
 fifths of the gas originally put in the tube is still there. If 
 the work is done properly, the volume of the gas left in the 
 tube will be to the total volume used as 79 to 100. In other 
 words, of every 100 cc. air used 21 cc. are absorbed by phos- 
 phorus, and 79 cc. are not. The gas absorbed is oxygen, 
 identical with the oxygen made from the oxide of mercury, 
 manganese dioxide, and potassium chlorate. The gas left 
 over has no chemical properties in common with oxygen. 
 Carefully take the tube out of the vessel of water, closing its 
 mouth with the thumb or some suitable object to prevent the 
 contents from escaping. Turn it with the mouth upward, and 
 introduce into it a burning stick. Does it support combus- 
 tion ? Is it oxygen ? 
 
 Experiment 120. Expose a few pieces of calcium chlo- 
 ride on a watch-glass to the air. It gradually becomes liquid 
 by absorbing water from the air. 
 
 Experiment 121. Expose some clear lime-water to the 
 air. It soon becomes covered with a white crust. A similar 
 change takes place if baryta- water is exposed in the same way. 
 Lime-water is made by putting a few pieces of quick-lime in 
 a bottle and pouring water upon it. The mixture is well 
 shaken up and allowed to stand. The undissolved substance 
 settles to the bottom, and with care a clear liquid can be 
 poured off the top. This is lime-water, which is a solution 
 of calcium hydroxide, Ca(OH) 3 , in water. Baryta- water is 
 a solution of a similar compound of the element barium. 
 When these solutions are exposed to nitrogen or oxygen, or to 
 an artificially prepared mixture of the two gases, no change 
 takes place. Further, if air is first passed through a solution 
 of caustic soda it no longer has the power to cause the forma- 
 tion of a crust on lime-water or baryta-water. 
 
 Experiment 122. Arrange an apparatus as shown in Fig. 
 54. The wash-cylinders A and B are half filled with ordi- 
 nary caustic-soda solution. The bottle C is filled with water. 
 
ANALYSIS OF AIR. 
 
 787 
 
 The tube D, which should be filled with water and provided 
 with a pinch-cock, acts as a siphon. Open the pinch-cock 
 and let the water flow slowly out of the bottle. As it flows 
 out air will be drawn in through the caustic soda in the wash- 
 
 Fio. 54. 
 
 cylinders. When the bottle is a quarter filled with air pour 
 some water in again until it is full. Then draw all the 
 water off. Now remove the stopper from the bottle, pour 
 in 20 to 30 cc. lime-water and cork the bottle. The crust 
 formed on the lime-water will now be hardly, if at all, per- 
 ceptible. There is, therefore, something present in the air 
 under ordinary circumstances which has the power to form a 
 crust on lime-water or baryta-water, and which can be re- 
 moved by passing the air through caustic soda. Thorough 
 examination has shown that this is the compound which 
 chemists call carlon dioxide, and which is commonly known as 
 carbonic acid gas. It is the substance which was obtained by 
 burning charcoal in oxygen. 
 
 Experiment 123. Into the bottle containing the air from 
 which the carbon dioxide has been removed hold a burning 
 stick or taper for a moment. Notice whether a crust is now 
 formed on the lime-water. Wood and the material from 
 which the taper is made contain carbon. Explain the forma- 
 tion of the crust on the lime-water after the stick of wood or 
 taper has burned for a short time in the vessel. 
 
 Experiment 124. Arrange an apparatus as shown in Fig. 
 55. The bottle A contains air. B contains concentrated sul- 
 phuric acid, C contains granulated calcium chloride, D is care- 
 
788 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. 
 
 fully dried and contains a few pieces of granulated calcium 
 chloride and air. Pour water through the funnel-tube into 
 Ay when the air will be forced through B and C and into D. 
 But in passing through B and G the moisture contained in it 
 
 FIG. 55. 
 
 will be removed, and the air which enters D will be dry. 
 After A has once been filled with water, empty it and fill it 
 again, letting the dried air pass into D. This operation may 
 be repeated indefinitely. The calcium chloride in D will not 
 grow moist. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. 
 PREPARATION AND PROPERTIES OF AMMONIA. 
 
 Experiment 125. To a little ammonium chloride on a 
 watch-glass add a few drops of a strong solution of caustic 
 soda, and notice the odor of the gas given off. Do the same 
 thing with caustic potash. Mix small quantities of ammo- 
 nium chloride and lime in a mortar, and add a few drops of 
 water. 
 
 Experiment 126. Mix 20 parts iron filings, 1 part potas- 
 sium nitrate, and 1 part solid potassium hydroxide, and heat 
 the mixture in a test-tube. Is there any evidence of the 
 formation of ammonia ? 
 
 Experiment 127. Arrange an apparatus as shown in Fig. 
 45. In the flask put a mixture of 100. grams slaked lime 
 and 50 grams ammonium chloride. Heat on a sand-bath. 
 
AMMONIA. 
 
 789 
 
 After the air is driven out, the gas will be completely ab- 
 sorbed by the water in the first Wolff's 
 flask if shaken from time to time. Dis- 
 connect the delivery-tube from the series 
 of Wolff's flasks, and connect with an- 
 other tube bent upward. Collect some of 
 the gas by displacement of air, placing the 
 vessel with the mouth doivmvard. (Why ?) 
 The arrangement is shown in Fig. 56. The 
 vessel in which the gas is collected should 
 be dry, as water absorbs ammonia' very 
 readily. Hence, also, it cannot be collected 
 over water. In the gas collected introduce 
 a burning stick or taper. Ammonia does 
 not burn in air, nor does it support com- 
 bustion. In working with the gas great care 
 must be taken to avoid inhaling it in any 
 quantity. After enough has been collected in cylinders to 
 'exhibit the chief properties, connect the delivery-tube again 
 with the series of Wolff's flasks, and pass the gas through the 
 water as long as it is evolved. 
 
 AMMONIA BURNS IN OXYGEN. 
 
 Experiment 128. Put a little of a concentrated solution 
 of ammonia in a flask placed upon a tripod. Heat gently and, 
 from a gasometer, pass a rapid current of oxygen through a 
 bent tube into the liquid. Apply a light to the mouth of the 
 vessel, when the ammonia will be seen to burn. 
 
 FIG. 56. 
 
 AMMONIA FORMS AMMONIUM SALTS WITH ACIDS. 
 
 Experiment 129. Put 100 cc. dilute ammonia solution in 
 an evaporating-dish. Try its effect on red litmus paper. 
 Slowly add dilute hydrochloric acid until the alkaline reaction 
 is destroyed and the solution is neutral. Evaporate to dry- 
 ness on a water-bath. Compare the substance thus obtained 
 with sal-ammoniac, or ammonium chloride. Taste. Heat 
 on a piece of platinum foil.* Treat with a caustic alkali. 
 Treat with a little concentrated sulphuric acid in dry test- 
 tubes. Do they appear to be identical ? Similarly sulphuric 
 acid and ammonia yield ammonium sulphate ; nitric acid and 
 ammonia yield ammonium nitrate ; etc. 
 
790 EXPERIMENTS TO ACCOMPANY CHAPTER XVL 
 
 Experiment 13O. Fill a dry cylinder with ammonia gas, 
 and another of the same size with hydrochloric acid gas. Bring 
 them together with their mouths covered. Quickly remove 
 the covers, when a dense white cloud will appear in and about 
 the cylinders. This will soon settle on the walls of the vessels 
 as a light white solid. It is ammonium chloride. Thus, from 
 two colorless gases we get a solid substance by an act of chemi- 
 cal combination. Heat is evolved in the act of combination. 
 
 COMPOSITION OF AMMONIA. 
 
 Experiment 131. This experiment should be performed 
 by a person experienced in the use of chemical apparatus. A 
 glass-tube, such as represented in Fig. 57, provided with a 
 glass stop-cock is needed. Fill this tube with chlorine free 
 from air over a saturated solution of sodium chloride. After 
 
 it is filled let it stand for some 
 time mouth downward in the 
 solution of sodium chloride to 
 let the liquid drip out of it. 
 Close the stop-cock and re- 
 move it from the solution. 
 Hold the tube mouth upward, 
 and pour a concentrated solu- 
 tion of ammonia into the fun- 
 nel-like projection above the 
 stop-cock, put in the glass 
 stopper, and now by slightly 
 opening the stop-cock let the 
 ammonia pass drop by drop 
 into the tube. Reaction be- 
 tween the chlorine and the 
 ammonia takes place, accom- 
 panied by a marked evolu- 
 tion of heat, and in a partly- 
 darkened room light is seen. 
 Great care must be taken not 
 to admit air with the am- 
 monia. After nearly all the 
 ammonia has passed in from 
 the funnel, pour into the 
 Fl - 57> funnel about two thirds as 
 
 much ammonia as has already been used, and let this 
 
PREPARATION AND PROPERTIES OF NITRIC ACID. 791 
 
 in gradually. Leave the stop-cock closed, and fill the 
 funnel with dilute sulphuric acid. . Fit a bent tube into a 
 cork, fill this tube with dilute sulphuric acid : put the 
 cork in the funnel, and the other end of the tube in a small 
 beaker containing dilute sulphuric acid, and, after immersing 
 the long tube in water of the ordinary temperature, open the 
 stop-cock. If the operation has been carried out as it should 
 be, the dilute acid will flow into the tube until it is two thirds 
 full, and will then stop. The residual gas is nitrogen. What 
 evidence in regard to the composition of ammonia is furnished 
 by this experiment ? 
 
 The arrangement of the apparatus in the last stage of the 
 experiment is shown in Fig. 57. 
 
 PREPARATION AND PROPERTIES OF NITRIC ACID. 
 
 Experiment 132 Arrange an apparatus as shown in Fig. 
 58. In the retort put 20 grams sodium nitrate (Chili salt- 
 
 FIG. 58. 
 
 peter) and 20 grams concentrated sulphuric acid. On gently 
 heating, nitric acid will distil over, and be condensed in the 
 receiver. After the acid is all distilled off, remove the con- 
 tents of the retort. Kecrystallize the substance from water, 
 and compare it with the sodium sulphate obtained in the 
 preparation of hydrochloric acid. (See Experiment 74.) In 
 the latter stage of the operation the vessels become filled with 
 a reddish-brown gas. The acid which is collected has a some- 
 what yellowish color. 
 
792 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. 
 
 Experiment 133. Mix together 400 grams concentrated 
 sulphuric acid and 80 grams ordinary concentrated nitric acid. 
 Pour the sulphuric acid into the nitric acid. Distil the mix- 
 ture from a retort arranged as in the preceding experiment, 
 taking care to keep the neck of the retort cool by placing 
 filter-paper moistened with cold water on it. Use the acid 
 thus obtained for the purpose of studying the properties of 
 pure nitric acid. 
 
 NITKIC ACID GIVES UP OXYGEN KEADILY, AND is HENCE A 
 GOOD OXIDIZING AGENT. 
 
 Experiment 134. Pour concentrated nitric acid into a 
 wide test-tube, so that it is about one-fourth filled. Heat 
 the end of a stick of charcoal of proper size, and, holding the 
 other end with a forceps, introduce the heated end into the 
 acid. It will continue to burn with a bright light, even, 
 though it is placed below the surface of the liquid. The 
 action is oxidation. The charcoal in this case finds the oxy- 
 
 Fio. 59. 
 
 gen in the acid and not in the air. Great care must be 
 taken in performing this experiment. The charcoal should 
 not come in contact with the sides of the test-tube. A large 
 beaker-glass should be placed beneath the test-tube, so that 
 in case it breaks the acid will be caught and prevented from 
 doing harm. The arrangement of the apparatus is shown in 
 Fig. 59. 
 
NITRATES. 793 
 
 The gases given off from the tube are offensive and poison- 
 ous. Hence this experiment as well as all others with nitric 
 acid should be carried on under a hood in which the draught 
 is good. 
 
 Experiment 135. Boil a little strong nitric acid in a test- 
 tube in the upper part of which some horse-hair has been in- 
 troduced in the form of a stopper. The horse-hair will take 
 fire and burn, and leave a white residue. Hold the test-tube 
 with a forceps over a vessel to catch the contents should the 
 tube break. 
 
 Experiment 136. In a small flask put a few pieces of 
 granulated tin. Pour on this just enough strong nitric acid 
 to cover it. Heat gently over a small flame. Soon action will 
 take place. Colored gases will be evolved, the tin will disap- 
 pear, and in its place will be found a white powder. This 
 consists mostly of tin and oxygen. (See Experiment 13.) 
 
 METALS DISSOLVE IN NITRIC ACID, FORMING NITRATES. 
 
 Experiment 137. Dissolve a few pieces of copper-foil in 
 ordinary commercial nitric acid diluted with about half its 
 volume of water. The operation should be carried on in a 
 good-sized flask and under an efficient hood. When the cop- 
 per has disappeared, pour the blue solution into an evaporat- 
 ing-dish, and evaporate down to crystallization. Compare the 
 substance thus obtained with copper nitrate. Heat specimens 
 of each. Treat small specimens with sulphuric acid. What 
 evidence have you that the two substances are identical ? 
 
 NITRATES ARE DECOMPOSED BY HEAT. 
 
 Experiment 138. Heat some potassium nitrate in a test- 
 tube. .Introduce a piece of wood with a spark on it. Heat 
 also lead nitrate, copper nitrate, and any other nitrates which 
 may be available. What difference do you observe between the 
 decomposition of potassium nitrate and that of lead nitrate ? 
 
 NITRATES ARE SOLUBLE IN WATER. 
 
 Experiment 139. Try the solubility in water of the ni- 
 trates used in the last experiment. 
 
794 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. 
 
 NITRIC ACID is REDUCED TO AMMONIA BY NASCENT 
 HYDROGEN. 
 
 Experiment 140 In a good-sized test-tube treat a few 
 pieces of granulated zinc with dilute sulphuric acid. What 
 is evolved? Prove it. Now add drop by drop dilute nitric 
 acid. The hydrogen ceases to be given off. Pour the con- 
 , tents of the tube into an evaporating-dish and evaporate the 
 liquid. Put the residue into a test-tube and add caustic-soda 
 solution, when the smell of ammonia will be noticed. Try the 
 action of the gas on red litmus-paper. Moisten the end of a 
 glass rod with a little hydrochloric acid, and hold it in the 
 tube. White fumes are seen. What are they ? Do the same 
 with nitric acid. What are the fumes in this case ? 
 
 NITROUS ACID. 
 
 Experiment 141. Melt 25 grams potassium nitrate in a 
 shallow iron plate and gradually add 50 grams metallic lead 
 cut in small pieces. Stir them together as thoroughly as 
 possible. After the mass is cooled down, break it up and 
 treat with water in a flask. The potassium nitrite will dis- 
 solve, while the lead oxide and unused lead will not dissolve. 
 Filter. Add a little sulphuric acid to some of the solution. 
 A colored gas will be given off. See whether a solution of 
 potassium nitrate acts in the same way. Treat with sulphu- 
 ric acid a little of the residue left after heating potassium 
 nitrate alone in a test-tube as in Experiment 138. 
 
 NITROUS OXIDE. 
 
 Experiment 142. In a retort heat 10 to 15 grams crystal- 
 lized ammonium nitrate until it has the appearance of boiling. 
 Do not heat higher than is necessary to secure a regular evo- 
 lution of gas. Connect a wide rubber tube directly with the 
 neck of the retort and collect the evolved gas over water, as 
 in the case of oxygen. It supports combustion almost as well 
 as pure oxygen. Try experiments with wood, a candle, and a 
 piece of phosphorus. 
 
OXIDES OF NITROGEN. 795 
 
 NITRIC OXIDE. 
 
 Experiment 143. Arrange an apparatus as shown in Fig. 
 60. In the flask put a few pieces of copper-foil. Cover this 
 with water. Now add slowly, waiting each time 
 for the action to begin, ordinary concentrated 
 nitric acid. When enough nitric acid has been 
 added gas will be evolved. If the acid is added 
 rapidly, it not unfrequently happens that the 
 evolution of gas takes place too rapidly, so that 
 the liquid is forced out of the flask through 
 the funnel-tube. This can be avoided by not 
 being in a hurry. At first the vessel becomes 
 filled with a reddish-brown gas, but soon the 
 gas evolved becomes colorless. Collect over 
 
 o 
 
 water two or three vessels full. The gas col- 
 lected is principally nitric oxide, NO, though FIG. eo. 
 it is frequently mixed with a considerable quantity of nitrous 
 oxide. 
 
 Experiment 144. Turn one of the vessels containing col- 
 orless nitric oxide with the mouth upward, and uncover it. 
 The colored gas is at once seen, presenting a very striking 
 appearance. Do not inhale the gas. Perform the experi- 
 ments with nitric oxide where there is a good draught. 
 
 Experiment 145. Pass nitric oxide into a concentrated 
 solution of ferrous sulphate. Afterwards heat the solution 
 and collect the gas. What do you conclude that the gas is? 
 
 NITROGEN" TRIOXIDE. . 
 
 Experiment 146. In a flask fitted with a safety-funnel and 
 a delivery-tube pour nitric acid of specific gravity 1.30-1.35 
 upon coarsely granulated arsenious oxide, As 2 3 . Heat gently, 
 and conduct the gases through a tube surrounded by a freez- 
 ing mixture, as in Experiment 110. 
 
 NITROGEN PEROXIDE. 
 
 Experiment 147. Admit a little air to nitric oxide con- 
 tained in a bell-jar over water, and let the ves'sel stand. 
 Almost immediately the color will disappear, showing that the 
 nitrogen peroxide formed is decomposed. Again admit air, 
 
796 EXPERIMENTS TO ACCOMPANY CHAPTER XVII. 
 
 and let the vessel stand. The same changes will be noticed as 
 in the first instance. If oxygen is used instead of air the 
 above changes can be repeated over and over again. Devise 
 an experiment for the purpose of determining whether the 
 nitric oxide is gradually used up or not. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVII. 
 
 PHOSPHORUS. 
 
 Experiment 148. [This, as well as the other experiments 
 with phosphorus, should be performed only by an experienced 
 person.] Arrange an apparatus as shown in Fig. 61. The 
 neck of the retort is somewhat drawn out and bent downward 
 and fitted air-tight by means of a cork to the wide glass tube 
 B. Some small pieces of ordinary phosphorus are now care- 
 fully slipped into the retort as much 
 as is obtained by cutting up two sticks 
 three to four inches long. The ap- 
 paratus is then adjusted as shown in 
 the figure, so that the end of the 
 tube B dips below the surface of the 
 water in the beaker C. The whole is 
 then allowed to stand for some hours. 
 The oxygen is absorbed from the air 
 contained in the vessel, and the water 
 rises in B. Without uncovering the 
 end of B t replace the water in C by 
 some that has a temperature of about 
 50. Now heat the retort gradually, 
 when the phosphorus will distil over 
 and condense in C in the molten condition. By lowering the 
 heat gradually at the end of the operation it can finally be 
 stopped without danger of breaking. 
 
 Experiment 149. Dissolve a little ordinary phosphorus in 
 carbon disulphide. Pour some of this solution upon a strip 
 of filter-paper, and let this hang in the air or wave it gently 
 in the air. After the carbon disulphide has evaporated the 
 phosphorus will take fire. 
 
 Experiment 150. Bring together in a porcelain crucible 
 or evaporating-dish a little phosphorus and iodine. It will be 
 seen that simple contact is sufficient to cause the two sub- 
 
 FIQ. 61. 
 
PHOSPHINE. 797 
 
 stances to act upon each other. Direct combination takes 
 place, and the action is accompanied by light and heat. 
 
 PHOSPHORUS ABSTRACTS OXYGEN FROM OTHER SUBSTANCES. 
 
 Experiment 151. Add a little of a solution of phosphorus 
 in carbon disulphide to a solution of copper sulphate. What 
 change takes place ? 
 
 Experiment 152. Put a few pieces of ordinary phosphorus 
 in a glass tube and seal it. Heat gradually to 300. Open 
 the tube and examine the product. See whether it takes fire 
 as readily as ordinary phosphorus does ; whether it dissolves 
 in carbon disulphide ; whether it melts easily when put in 
 water heated to between 45 and 50. 
 
 PHOSPHINE. 
 
 Experiment 153. Arrange an apparatus as shown in Fig. 
 <62. In the small flask B put about 5 grams caustic potash dis- 
 
 FIG. 62. 
 
 solved in 10-15 cc. water, and when the solution is cold add a 
 few small pieces of phosphorus the size of a pea. Pass hydrogen 
 for some time through the apparatus from the generating- 
 flask A until all the air is displaced ; then disconnect at D, 
 leaving the rubber tube, closed by the pinch-cock, on the 
 tube which enters the flask. Gently heat the contents of the 
 retort, when gradually a gas will be evolved, and will escape 
 through the water in C. As each bubble comes in contact 
 
798 EXPERIMENTS TO ACCOMPANY CHAPTER XVII. 
 
 with the air it takes fire, and the products of combustion ar 
 range themselves in rings, which become larger as they rise. 
 They are extremely beautiful, particularly if the air of the 
 room is quiet. Both the phosphorus and the hydrogen com- 
 bine with oxygen in the act of burning. Collect some of the 
 gas in a tube over water, and then place the tube mouth up- 
 ward. What difference is there between the burning of the 
 gas under these circumstances, and that noticed when the 
 rings are formed? Collect another tube full of the gas, and 
 let this stand for some time. Then open the vessel by taking 
 it out of the water. Has any change taken place in the gas ? 
 
 ARSENIC. 
 
 Experiment 154 Heat a small piece of arsenic on charcoal 
 in the flame of the blow-pipe. 
 
 ARSINE. 
 
 Experiment 155 Arrange an apparatus as shown in Fig. 
 63. Put some pure granulated zinc in the flask and pour 
 
 FIG. 68. 
 
 dilute sulphuric acid on it. The calcium-chloride tube serves 
 to dry the gas. When the air is all out of the vessel and the 
 hydrogen is lighted, add slowly a little of a solution of arsenic 
 trioxide, As 2 3 , in dilute hydrochloric acid. The appearance 
 of the flame will soon change. It will become paler, with a 
 slightly bluish tint, and give off white fumes. (See next ex- 
 periment.) 
 
MARSH 'S TEST FOR ARSENIC ANTIMONY, ETC. 799 
 
 MARSH'S TEST FOR ARSENIC. 
 
 Experiment 156. Into the flame of the burning hydrogen 
 and arsine produced in the last experiment introduce a piece 
 of porcelain, as the bottom of a small porcelain dish or a cru- 
 cible, and notice the appearance of the spots. Heat by means 
 of a Buusen burner the tube through which the gas is passing, 
 which should be of hard glass. Just beyond the heated 
 place there will be deposited a thin layer of metallic arsenic, 
 commonly called a mirror of arsenic. This deposit is due to 
 the direct decomposition of the arsine into arsenic and hydro- 
 gen by heat. [Compare ammonia, phosphine, and arsine with 
 reference to their stability.] 
 
 ANTIMONY. 
 
 Experiment 157. Heat a small piece of antimony on char- 
 coal in the blow-pipe flame. Try the action of dilute and of 
 concentrated hydrochloric acid, of dilute and of concentrated 
 nitric acid, and of a mixture of the two acids on a small piece 
 of antimony. 
 
 STIBINE. 
 
 Experiment 158. Stibine is made by the same method as 
 that used in making arsine. Make some, using a solution of 
 tartar emetic. Introduce a piece of porcelain in the flame, 
 and afterwards heat the tube through which the gas is passing. 
 Compare the antimony spots with the arsenic spots. Color? 
 Volatility? Conduct towards a solution of sodium hypochlo- 
 rite or hypobromite? 
 
 BISMUTH. 
 
 Experiment 159. Heat a piece of bismuth on charcoal in 
 the blow-pipe flame. See how it conducts itself towards hy- 
 drochloric acid; towards nitric acid. If a solution is obtained 
 in either case, add water to it. Explain what takes place. 
 
 PHOSPHORUS TRICHLORIDE. 
 
 The experiments with the chlorides of phosphorus must be 
 carried on under a hood or out-of-doors. 
 
 Experiment 160. Arrange an apparatus as shown in Fig. 
 64. The tube A is arranged so that it can be raised or low- 
 ered in the retort. Put 50 to 100 grams ordinary phosphorus 
 in the retort, taking precautions to prevent it from taking fire 
 
800 EXPERIMENTS TO ACCOMPANY CHAPTER XVIL 
 during the operation. This is best accomplished by fitting 
 
 FIG. 64. 
 
 corks in both openings of the retort; placing the retort in a 
 vessel of cold water; removing the cork from B, throwing in 
 a piece of phosphorus, and quickly putting the cork in. The 
 pieces must not be put in in too rapid succession. After all 
 the phosphorus is in the retort, adjust the apparatus as repre- 
 sented, placing the receiver D in a dish of cold water. Now 
 connect by means of the rubber tube E with an apparatus 
 furnishing chlorine, dried by means of concentrated sulphuric 
 acid and calcium chloride. As soon as the chlorine comes in 
 contact with the phosphorus action begins, and the product, 
 which is phosphorus trichloride, distils over into the receiver. 
 If the action is taking place too rapidly, the inside of the re- 
 tort will become covered with a coating of red phosphorus. 
 In this case raise the tube A a little and the red coating will 
 gradually disappear. If the tube is raised too high, not enough 
 heat is generated, and the trichloride in the retort is converted 
 into the pentachloride, which is deposited as a white coating. 
 By raising and lowering the tube ac- 
 cording to the indications, the retort 
 can be kept clear, and all the phos- 
 phorus converted into the trichloride. 
 This manipulation of the tube is much 
 facilitated by fitting into the cork a 
 somewhat larger tube, through which 
 the smaller one can pass easily; letting 
 this project about an inch and an half 
 above the cork and passing over it a 
 piece of rubber tubing of such size 
 that while the smaller tube moves through it readily, the two 
 form a gas-tight joint. This is shown in Fig. 65. After the 
 
 FIG. 65. 
 
PHOSPHORUS PENTACHLORIDE. 
 
 801 
 
 operation is finished, pour the liquid from the receiver into a 
 clean dry flask, and distil on a water-bath. Try the action of 
 a little of the compound on water. 
 
 PHOSPHORUS PENTACHLORIDE. 
 
 Experiment 161 Put the trichloride of phosphorus ob- 
 tained in the last experiment in a wide-mouthed 
 bottle surrounded by cold water. Through a 
 wide glass tube pass dry chlorine upon the 
 surface of the liquid, and as the action ad- 
 vances, and a solid begins to make its appear- 
 ance, stir the contents of the bottle. Con- 
 tinue the passage of the chlorine until the 
 product is a perfectly dry solid. The arrange- 
 ment of the bottle containing the trichloride, 
 and that of the delivery-tube, is shown in Fig. 
 66. The bottle is put in a larger vessel con- 
 taining cold water, which is renewed from 
 time to time during the process. 
 
 Try the action of a little phosphorus pentachloride on 
 water. In a large dry flask heat a little of the pentachloride. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XVIII. 
 
 PHOSPHORIC ACID. 
 Experiment 162. In a flask connected with an inverted 
 
 FIG. 66. 
 
 FIG 67. 
 
 condenser, as shown in Fig. 67, boil 10 to 15 grams of ordinary 
 phosphorus with 250 cc. ordinary commercial nitric acid. If 
 
802 EXPERIMENTS TO ACCOMPANY CHAPTER XVIIL 
 
 necessary, add more acid after a time. Boil gently until the 
 phosphorus disappears. Evaporate the solution to complete 
 dryness, so as to get rid of all the nitric acid. Dissolve a lit- 
 tle of the product in water, and add a few drops of the solu- 
 tion to a dilute solution of silver nitrate. What effect is pro- 
 duced? Heat some of the product gently in a porcelain 
 crucible, and from time to time take out a little, dissolve it in. 
 water, and try its action on silver nitrate. 
 
 Experiment 163. Try the action of ordinary sodium phos- 
 phate on silver nitrate. Heat a little of the salt in a porcelain 
 crucible to redness. After cooling, try the action of the salt 
 left in the crucible on silver nitrate. 
 
 ARSENIC ACID. 
 
 Experiment 164. Pass chlorine into water containing ar- 
 senic trioxide in suspension, until the oxide is dissolved. 
 Evaporate to crystallization. Into a dilute solution of the 
 product thus obtained, to which some hydrochloric acid is 
 added, pass hydrogen sulphide. Explain the changes. 
 
 KEDUCTION OF ARSENIC TRIOXIDE. 
 
 Experiment 165. In the bottom of a dry tube of hard glass 
 of the form represented in Fig. 68 put a minute piece of ar- 
 senic trioxide, and just above it a small bit of charcoal. 
 Heat gently. Explain the change. 
 
 SULPHIDES OF ARSENIC. 
 
 Experiment 166. Pass hydrogen sulphide into a di- 
 lute solution of arsenic trioxide in hydrochloric acid. 
 Filter off the precipitate, and try the action of ammoni- 
 um sulphide on some of it. 
 
 SULPHIDES OF ANTIMONY. 
 
 Experiment 167. Pass hydrogen sulphide into a so- 
 FIO. 68. j u ^ on O f antimonio acid made by treating antimony 
 with aqua regia and diluting with water. Pass hydrogen sul- 
 phide into a solution of antimony trichloride made by dis- 
 solving stibnite or antimony trisulphide in hydrochloric acid. 
 Try the action of ammonium sulphide on the precipitates 
 after filtering. 
 
 OXYCHLORIDES OF ANTIMONY. 
 
 Experiment 168. Treat a solution of antimony trichloride 
 with water. 
 
BASIC NITRATES OF BISMUTH CARBON. 803 
 
 BASIC NITRATES OF BlSMUTH. 
 
 Experiment 169. Dissolve a little bismuth in nitric acid 
 and evaporate. Add water. 
 
 BORON. 
 
 Experiment 170. Make a hot solution of 30 grams crystal- 
 lized borax in 120 cc. water. Add slowly 10 grams concen- 
 trated sulphuric acid. On cooling, the boric acid will crystal- 
 lize out. What evidence have you that the substance which 
 crystallizes out of the solution is not borax? Try the solu- 
 bility in alcohol of specimens of each. Is there any difference? 
 Treat a few crystals of borax with about 10 cc. alcohol ; pour 
 off the alcohol and set fire to it. Treat a few crystals of the 
 boric acid in the same way. What difference do you observe ? 
 Distil an aqueous solution of boric acid, and determine 
 whether any of the acid passes over with the water vapor. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XIX. 
 
 CARBON. BONE-BLACK FILTERS. 
 
 Experiment 171. Make a filter of bone-black by fitting a 
 paper filter into a funnel 12 to 15 mm. (5 to 6 inches) in di- 
 ameter at its mouth. Half fill this with bone-black. Pour a 
 dilute solution of indigo through the filter. If the conditions 
 are right the solution will pass through colorless. Do the 
 same thing with a dilute solution of litmus. If the color is 
 not completely removed by one filtration, heat and filter 
 again. The color can also be removed from solutions by put- 
 ting some bone-black into them and boiling for a time. Try 
 this with half a liter each of the litmus and indigo solutions 
 used in the first part of the experiment. Use about 4 to 5 
 grams bone-black in each case. Shake the solution frequently 
 while heating. 
 
 CHARCOAL ABSORBS GASES. 
 
 Experiment 172. Collect over mercury in glass tubes some 
 ammonia gas, and some carbon dioxide. Introduce into each 
 a piece of charcoal, which has been heated in a Bunsen- 
 burner flame in order to drive out gases which may be con- 
 tained in the pores. * 
 
 CARBON COMBINES WITH OXYGEN TO FORM CARBON DIOXIDE. 
 
 Experiment 173 Put a small piece of charcoal in a piece of 
 
 hard-glass tube. Heat the tube, and pass oxygen through it. 
 
804 EXPERIMENTS TO ACCOMPANY CHAPTER XIX. 
 
 Pass the gases into clear lime-water. Arrange the apparatus 
 as shown in Fig. 69. 
 
 FIG. 69. 
 
 A is a large bottle containing oxygen ; B is a cylinder con- 
 taining sulphuric acid; is a U-tube containing calcium 
 chloride ; D is the hard-glass tube containing the charcoal; 
 E is the cylinder with clear lime-water. Explain all that 
 takes place. 
 
 CARBON REDUCES SOME OXIDES WHEN HEATED WITH THEM. 
 
 Experiment 174 Mix together two or three grams pow- 
 dered copper oxide, CuO, and about one tenth its weight of 
 
 powdered charcoal ; heat in a hard- 
 glass tube, as shown in Fig. 70, or, 
 still better, use an arsenic-tube. 
 
 Pass the gas which is given off 
 into lime-water contained in a test- 
 tube. Is it carbon dioxide ? What 
 evidence have you that oxygen has 
 been extracted from the copper 
 oxide ? Compare the substance left 
 in the tube with metallic copper. 
 Treat both with nitric acid, with 
 sulphuric acid. 
 
 Experiment 175. Eepeat Experiment 165 with somewhat 
 larger quantities of the substances, and examine the gas 
 given off. 
 
 HYDROCARBONS. 
 
 Experiment 176. Make marsh-gas by heating in a retort a 
 
 FIG. 70. 
 
CARBON DIOXIDE. 
 
 805 
 
 mixture of 20 grams sodium acetate, 20 grams potassium hy- 
 droxide, and 30 grams slaked lime. Collect some of the gas 
 over water. Is it a combustible gas ? 
 
 Experiment 177. Make ethylene as follows : In a flask of 
 2 to 3 liters capacity put a mixture of 25 grams alcohol and 
 150 grams ordinary concentrated sulphuric acid. Heat to 
 160 to 170, and add gradually through a funnel tube about 
 500 cc. of a mixture of 1 part of alcohol and 2 parts of con- 
 centrated sulphuric acid. Pass the gas through three wash- 
 bottles containing, in order, concentrated sulphuric acid, 
 caustic soda, and concentrated sulphuric acid. Collect some 
 of the gas over water. Is it combustible ? 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XX. 
 CARBON DIOXIDE is FORMED WHEN A CARBONATE is 
 
 TREATED WITH AN ACID. 
 
 Experiment 178. In test-tubes add successively dilute hy- 
 drochloric, sulphuric, nitric, and acetic acids to a little sodium 
 carbonate. In each case pass the gas given off through lime- 
 water, and insert a burning stick in the upper part of each 
 tube. Perform the same experiments with small pieces of 
 marble. 
 
 PREPARATION AND PROPERTIES OF CARBON DIOXIDE. 
 Experiment 179. Arrange an apparatus as shown in Fig. 
 71. In the flask put some pieces of marble or limestone, 
 and pour ordinary hydrochloric acid 
 on it. The gas should be collected by 
 displacement of air, the vessel being 
 placed with the mouth upward. Col- 
 lect several cylinders or bottles full of 
 the gas. Into one introduce succes- 
 sively a lighted candle, a burning stick, 
 a bit of burning phosphorus. Into 
 another, if convenient, put a live mouse. 
 "With another proceed as if pouring 
 water from it. Pour the invisible gas 
 upon the flame of a burning candle. 
 Pour some of the gas from . one vessel 
 to another, and show that it has been ^ 
 transferred. Balance a beaker on a ~~ "~^ 
 
 good-sized pair of scales, and pour car- 
 bon dioxide into it. If the balance is at all sensitive, the pan 
 on which the beaker is placed will sink. 
 
806 EXPERIMENTS TO ACCOMPANY CHAPTER XX. 
 
 CARBON DIOXIDE is GIVEN OFF FROM THE LUNGS. 
 
 Experiment 180. Force the gases from the lungs through 
 
 some lime-water by means of an 
 apparatus arranged as shown in 
 Fig. 72. 
 
 FORMATION OF CARBONATES. 
 
 Experiment 181. Pass car- 
 bon dioxide into a solution of po- 
 tassium hydroxide to saturation. 
 Determine whether a carboiiate is 
 in solution or not. 
 
 Experiment 182. Pass car- 
 bon dioxide into 50 to 100 cc. clear 
 lime-water. Filter off the white 
 
 insoluble substance. Try the action of a little acid on it. 
 
 What evidence have you that it 'is a carbonate ? 
 
 Experiment 183. Pass carbon dioxide first through a little 
 
 water to wash it, and then into 50 to 100 cc. clear lime-water. 
 
 Continue to pass the gas for some time after the precipitate 
 
 is formed. The precipitate dissolves. Heat the solution. 
 
 What happens ? Explain these reactions. 
 
 PREPARATION AND PROPERTIES OF CARBON MONOXIDE. 
 
 Experiment 184 Put 10 grams crystallized oxalic acid and 
 50 to 60 grams concentrated sulphuric acid in an appropriate 
 flask. Connect with two Wolff's flasks containing a solution of 
 caustic soda, so that the gas evolved will bubble through the 
 solution. Heat gently. Collect some of the gas over water. 
 Set fire to some, and notice the characteristic blue flame. If 
 convenient put a live mouse in a vessel containing a mixture 
 of about equal parts of carbon monoxide and air. It will die 
 unless taken out. 
 
 CARBON MONOXIDE is A GOOD KEDUCING AGENT. 
 
 Experiment 185. Pass carbon monoxide over some heated 
 copper oxide contained in a hard-glass tube. Is the oxide re- 
 duced? How do you know? Is carbon dioxide formed? 
 What evidence have you ? Was the carbon monoxide used 
 free of carbon dioxide? If not, what evidence have you that 
 carbon dioxide is formed in this experiment ? 
 
 Experiment 186. Pass carbon dioxide over heated charcoal 
 in a hard-glass tube. What is formed ? 
 
COAL-GAS, ETC. 807 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXI. 
 COAL-GAS. 
 
 Experiment 187. Heat some bituminous coal in a retort 
 and collect over water the gases given off. Are these gases 
 combustible ? 
 
 OXYGEN BURNS IN AN ATMOSPHERE OF A COMBUSTIBLE GAS. 
 
 Experiment 188. Break off the neck of a good-sized re- 
 tort ; fit a perforated cork to the small end ; pass a piece of 
 glass tube through the cork, and connect by means 'of rubber 
 hose with an outlet for coal-gas. Fix the apparatus in position, 
 
 FIG. 73. 
 
 as shown in Fig. 73. Turn the gas on, and when the air is 
 driven out of the retort-neck, light the gas. The neck is now 
 filled with illuminating gas, and the gas is burning at the mouth 
 of the vessel. If now a platinum jet from which oxygen is 
 issuing is passed up into the gas the oxygen will take fire, and a 
 flame will appear where the oxygen escapes from the jet. 
 The oxygen burns in the atmosphere of coal-gas. 
 
 KINDLING TEMPERATURE OF GASES. 
 
 Experiment 189 Light a Bunsen burner. Bring down 
 upon the flame a piece of brass or iron wire-gauze. There is 
 no flame above the gauze. That the gas passes through un- 
 burned can be shown by applying a light just above the outlet 
 of the burner and above the gauze. The gas will take fire and 
 burn. By simply passing through the thin wire-gauze, then, the 
 gas is cooled down below its burning temperature, and does not 
 burn unless it is heated up again. Turn on a Bunsen burner. 
 
808 EXPERIMENTS TO ACCOMPANY CHAPTER XXL 
 
 Do not light the gas. Hold a piece of wire-gauze about one 
 and a half to two inches above the outlet. Apply a lighted 
 match above the gauze, when the gas will burn 
 above the gauze, but not below it. Here again 
 the heat necessary to raise the temperature of the 
 gas to the burning temperature cannot be com- 
 municated through the gauze. If in either of the 
 above-described experiments the gauze is held in 
 position for a time, it will probably become so 
 highly heated that the gas on the side where there 
 is no flame will be raised to the burning tempera- 
 ture. The instant that point is reached the 
 flame becomes continuous. 
 
 THE BLOW-PIPE AND ITS USES. 
 
 FIG ?4 The blow-pipe used in chemical laboratories is 
 
 constructed as shown in Fig. 74. 
 
 When used with the Bunsen burner it is best to slip into 
 the burner a brass tube ending above 
 in a narrow slit-like opening, as shown in 
 Fig. 75. The tube referred to, marked 
 
 FIG. 76. 
 
 a in the figure, reaches to the bottom of 
 the burner, and thus cuts off the supply 
 , of air which usually enters the holes at 
 FIG. 75. the base. The gas is now lighted, and 
 
 the current so regulated that there is a small flame about 1^- 
 to 2 inches long. The tip of the blow-pipe is placed on the 
 slit of the burner in the flame, as shown in Fig. 76. By blow- 
 ing regularly and not violently through the pipe the flame is 
 forced down in the same direction as the end-piece of the 
 blow-pipe, and the slant of the burner-slit. Under proper 
 
THE BLOW-PIPE AND ITS USES, 809 
 
 conditions the flame separates sharply into a central blue part 
 and an outer part of another color. The direction and lines of 
 division of the flame are indicated in Fig. 76. The outer 
 part of the flame marked o is the oxidizing flame ; the part 
 marked r is the reducing flame. 
 
 Experiment 190. Select a piece of charcoal about 4 
 inches long by 1 inch wide and 1 inch thick, with one surface 
 plane.* Near the end of the plane surface make a cavity by 
 pressing the edge of a small thin coin against it, and turning 
 it completely round a few times. Mix together equal small 
 quantities of dry sodium carbonate -and lead oxide. Put a 
 little of the mixture in the cavity in the charcoal, and heat it 
 in the reducing flame produced by the blow-pipe. In a short 
 time globules of metallic lead will be seen in the molten mass. 
 After cooling, scrape the solidified substance out of the cavity 
 in the charcoal. Put it in a small mortar, treat it with a little 
 water, and, after breaking it up and allowing as much as pos- 
 sible to dissolve, pick out the metallic beads. Is it malleable 
 or brittle? Is metallic lead malleable or brittle? Is it dis- 
 solved by hydrochloric acid ? Is lead soluble in hydrochloric 
 acid ? Is it soluble in nitric acid ? Is lead soluble in nitric 
 acid ? The action of the acids can be tried by putting the 
 bead on a small dry watch-glass and adding a few drops of the 
 acid. Does the substance act like lead? What has become 
 of the oxygen with which the lead was combined in the oxide ? 
 Is there any special advantage in having a support of charcoal 
 for this experiment? 
 
 Experiment 191. Heat a small piece of metallic lead on 
 charcoal in the oxidizing blow-pipe flame. Notice the forma- 
 tion of the oxide, which forms a coating or film on the char- 
 coal in the neighborhood of the metal. Is there any analogy 
 between this process and the burning of hydrogen ? In what 
 does the analogy consist? What differences are there between 
 the two processes? 
 
 Experiment 192 Repeat the experiments with arsenic, 
 antimony, and bismuth. Notice the colors of the films formed 
 on the charcoal. 
 
 Experiment 193. Melt into a bit of glass tubing a piece of 
 platinum wire 8 to 10 mm. (3 to 4 inches long) and bend the 
 end so as to form a small loop, as shown in Pig. 77. Heat the 
 
 * Pieces of charcoal prepared for blow -pipe work can be bought from 
 dealers in chemical apparatus, at small cost. 
 
810 EXPERIMENTS TO ACCOMPANY CHAPTER XXll. 
 
 loop in the flame of a Bunsen burner, and then dip it into 
 some sodium-ammonium phosphate (microcosmic salt). Heat 
 in the oxidizing flame of the blow-pipe until a clear glass bead 
 is formed in the loop. What changes have taken place ? and 
 what is the clear glass ? Bring a minute particle of a man- 
 
 FIG. 77. 
 
 ganese compound in contact with the bead, and heat again. 
 What change takes place ? Try the same experiment, using 
 successively a cobalt compound, a copper compound, and an 
 iron compound. Now, instead of using microcosmic salt, use 
 borax. Explain the changes in all the above-described experi- 
 ments. 
 
 CYANOGEN. 
 
 Experiment 194. Make potassium cyanide by heating po- 
 tassium ferrocyanide in an iron crucible. 
 
 Experiment 195. Make cyanogen by heating mercuric cy- 
 anide. Cyanogen is poisonous. Burn some of the gas. 
 
 Experiment 196. Make potassium cyanate from some of 
 the cyanide obtained in Experiment 194. This is done by 
 melting it in an iron crucible, and, while the mass is liquid, 
 adding about four times its weight of red lead, stirring during 
 the operation. After this the crucible should again be 
 put in the furnace for a little while, the metallic lead allowed 
 to settle, and the contents poured out on a smooth stone. 
 Break this up, and extract the cyanate with alcohol. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXII. 
 SILICON. 
 
 Experiment 197. Prepare sodium fluosilicate as directed 
 in the next experiment. Mix 3 parts of the dry salt with 1 
 part of sodium cut in pieces. Throw this mixture all at once 
 into a Hessian crucible heated to bright-red heat in a furnace. 
 Add immediately 9 parts granulated zinc, and a layer of sodium 
 chloride previously heated to drive off water. The crucible is 
 then covered, and the fire allowed to burn down. After cool- 
 ing, the regulus of zinc containing the silicon is separated from 
 the slag, washed with water, and treated with hydrochloric 
 
SILICON TETRAFLUORIDE AND FLUOSILICIC ACID. 811 
 
 acid. The zinc dissolves and leaves the silicon. This is again 
 washed with water and then heated with nitric acid, and 
 washed with water, when crystals of silicon, sometimes of 
 great beauty, are obtained. Try the effect of heating a little 
 of the silicon in the air. Try the action of acids and of 
 alkalies upon it. 
 
 SILICON TETKAFLUORIDE AND FLUOSILICIC ACID. 
 
 Experiment 198. Arrange an apparatus as shown in Fig. 
 78. A is a bottle of about 2 liters capacity, such as are com- 
 
 Ji 
 
 FIG. 78. 
 
 monly used for transporting acids. This is about two-thirds 
 filled with alternating layers of sand and powdered fluor-spar, 
 moistened with concentrated sulphuric acid. The bottle is 
 put in the deep sand-bath B, and connected by means of a 
 wide glass tube with the funnel C, which dips just below the 
 surface of the water in the large evaporating-dish D. The 
 sand-bath is now gently heated, when silicon tetrafluoride 
 passes over. Coming in contact with water, it is decomposed, 
 silicic acid being deposited -and fluosilicic acid passing into 
 solution. In order to prevent clogging, the gelatinous silicic 
 acid is from time to time removed from the mouth of the 
 funnel by means of a bent-glass rod. After the action is com- 
 plete, filter the solution. Take out one quarter, and to the 
 
812 EXPERIMENTS TO ACCOMPANY CHAPTER 
 
 rest slowly add a solution of sodium carbonate until the whole 
 just begins to show an alkaline reaction; now add the other 
 quarter of the acid, and filter. Explain all the reactions. 
 Heat a little of the dried salt in a covered platinum crucible. 
 What change takes place? What evidence have you that the' 
 change has taken place ? To a little of the salt in water add 
 a solution of potassium hydroxide. What change takes place ? 
 Dry the silicic acid formed in the first part of the experiment 
 by decomposition of the silicon tetrafluoride. 
 
 SILICIC ACID. 
 
 Experiment 199. Boil some of the silicic acid obtained in 
 the last experiment with sodium hydroxide. Treat some of 
 the solution with hydrochloric acid ; with ammonium chloride. 
 
 Experiment 200. Add some fine sand to about four times 
 its weight of a molten mixture of potassium and sodium car- 
 bonates, heated in a platinum crucible in the flame of the 
 blast-lamp. Continue the heating until no more sand is dis- 
 solved. Pour the molten mass out on a stone, and when 
 cooled break it up and treat it with water. 
 
 Experiment 201. Treat a little of the solution containing 
 sodium and potassium silicates, prepared in the last experi- 
 ment, with a little sulphuric or hydrochloric acid. A gelati- 
 nous substance will be precipitated. This is silicic acid. Some 
 of the acid remains in solution. By evaporating the solution 
 to dryness and heating for a time on the water-bath, all the 
 silicic acid is rendered insoluble. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXIV. 
 CHLORIDES, BROMIDES, AND IODIDES. 
 
 Experiment 202. Dissolve a small crystal of silver nitrate 
 in pure water. Add to a small quantity of this solution in a 
 test-tube a few drops of dilute hydrochloric acid. The white 
 substance thus precipitated is silver chloride, AgCl. To an- 
 other small portion of the solution add a few drops of a dilute 
 solution of common salt, or sodium chloride, NaCl. The 
 white substance produced in this case is also silver chloride. 
 Add ammonia to each tube. If sufficient is added the 
 precipitates will dissolve. On adding enough hydrochloric 
 acid to these solutions to combine with all the ammonia the 
 
HYDROXIDES. 813 
 
 silver chloride is again thrown down. On standing exposed to 
 the light both precipitates change color, becoming finally dark 
 violet. The reactions involved in the above experiments are 
 these : In the first place, when hydrochloric acid is added to 
 silver nitrate this reaction takes place : 
 
 AgNO, + HC1 = AgCl + HN0 3 . 
 
 When sodium chloride is added this reaction takes place : 
 AgN0 3 + NaCl = AgCl -f NaNO,. 
 
 In the first reaction nitric acid is set free ; in the second, 
 the sodium and silver exchange places. In addition to the 
 insoluble silver chloride, there is formed at the same time the 
 soluble salt, sodium nitrate. On adding ammonia the silver 
 chloride forms with it a compound which is soluble in water ; 
 and on adding an acid, the ammonia combines with it, leav- 
 ing the silver chloride uncombined and therefore insoluble. 
 
 Extensive use is made of insoluble compounds for the pur- 
 pose of detecting substances in analysis. The only insoluble 
 chlorides are those of silver, lead, and mercury. * If, there- 
 fore, on adding hydrochloric acid or a soluble chloride to a 
 solution, a precipitate is formed, the conclusion is justified 
 that one or more of the three metals silver, lead, or mercury 
 is present. By taking account of the differences in the 
 properties of these chlorides it is not difficult to decide of 
 which of them a precipitate consists. 
 
 HYDROXIDES. 
 
 Experiment 203. To some pieces of freshly-burnt lime 
 add enough cold water to cover it. The action which takes 
 place is represented by the equation 
 
 CaO + H 2 = Ca(OH),. 
 
 The process is known as slaking. 
 
 Experiment 204. To a small quantity of a dilute solution 
 of magnesium sulphate add a dilute solution of caustic soda. 
 The white precipitate is magnesium hydroxide. [Would you 
 
 * There are two chlorides of mercury. Only one of them, mercurous 
 chloride, is insoluble. 
 
814 EXPEE1MENTS TO ACCOMPANY CHAPTER XXIV. 
 
 expect this precipitate to be soluble in sulphuric acid ? in hy- 
 drochloric acid ? in nitric acid ?] The answers follow from 
 these considerations : When acids act upon hydroxides, salts 
 are formed ; magnesium sulphate is soluble, as is seen by the 
 fact that we started with a solution of this salt ; the only inso- 
 luble chlorides are those of silver, lead, and mercury ; all 
 nitrates are soluble. 
 
 When a solution of an iron salt is treated with sodium hy- 
 droxide a precipitate of iron hydroxide is formed : 
 
 FeCl s + 3NaOH = Fe0 3 H 3 + SNaOl. 
 
 Experiment 205. To a dilute solution of that chloride of 
 iron which is known as ferric chloride add caustic soda. 
 The reddish precipitate which is formed is ferric hydroxide. 
 [From the general statements made above, would you expect 
 this precipitate to be soluble in hydrochloric acid? in nitric 
 acid ? Try each. Is it soluble in sulphuric acid ?] 
 
 Experiment 206. Add to a solution of an aluminium salt 
 sodium hydroxide. After a precipitate is formed continue to 
 add the sodium hydroxide. Perform similar experiments with 
 a chromium and with a lead salt. Boil each of the solutions 
 obtained. Treat a solution of copper sulphate with sodium 
 hydroxide in the cold. Heat. 
 
 SULPHATES. 
 
 Experiment 207 Make a dilute solution of barium chlo- 
 ride, of lead nitrate, of strontium nitrate. To a small quan- 
 tity of each in a test-tube add a little sulphuric acid. [What 
 remains in solution ?] Make a somewhat concentrated so- 
 lution of calcium chloride. To this add sulphuric acid. 
 [What is in solution ?] Add more water, and see whether the 
 precipitate will dissolve. The formulas of the salts used in 
 the experiments are barium chloride, BaCl 2 ; lead nitrate, 
 Pb(N0 3 ) 2 ; strontium nitrate, Sr(N0 3 ) 2 . [Write the equations 
 expressing the reactions. ] If to the solutions of the salts any 
 soluble sulphate is added instead of sulphuric acid, the same 
 insoluble sulphates will be formed. The sulphates of iron, cop- 
 per, sodium, and potassium are among the soluble sulphates. 
 Make dilute solutions of small quantities of each of these, and 
 add them successively to the solutions of barium chloride, 
 
REDUCTION OF SULPHATES TO SULPHIDES. 815 
 
 lead nitrate, and strontium nitrate. The formula of iron sul- 
 phate is FeS0 4 ; of copper sulphate, CuS0 4 ; of sodium sul- 
 phate, Na 2 S0 4 ; and of potassium sulphate, K 2 S0 4 . Write the 
 equations representing the reactions which take place in the 
 above experiments. It need hardly be explained that the 
 action consists in an exchange of places on the part of the 
 metals. Thus, when the soluble salt iron sulphate, FeS0 4 , 
 is brought together with the soluble salt barium chloride, 
 BaCl 2 , the insoluble salt barium sulphate, BaS0 4 , and the 
 soluble salt iron chloride, FeCl 2 , are formed : 
 
 FeS0 4 + BaCl 2 = FeCl 2 + BaS0 4 . 
 
 REDUCTION OF SULPHATES TO SULPHIDES. 
 
 Experiment 208. Mix and moisten a little sodium sulphate 
 and finely-powdered charcoal. Heat the mixture for some 
 time in the reducing flame. After cooling scrape off the salt, 
 dissolve it in a few cubic centimeters of water, and filter 
 through a small filter. If the change to the sulphide has 
 taken place, sodium sulphide, Na 2 S, is in solution. A solu- 
 tion of a sulphide when added to a solution containing copper 
 gives a black precipitate of copper sulphide. Try this; also 
 try the action on the solution of the salt of copper of some of 
 the sulphate from which the sulphide was made. 
 
 CARBONATES. 
 
 Experiment 209. The formation of carbonates by the ad- 
 dition of soluble carbonates to solutions of salts of metals whose 
 carbonates are insoluble, is illustrated by the following experi- 
 ments: Make solutions of copper sulphate, iron sulphate, lead 
 nitrate, silver nitrate, calcium chloride, barium chloride. 
 Add to each a little of a solution of a soluble carbonate, as 
 sodium carbonate, potassium carbonate, ammonium carbon- 
 ate. Note the result in each case. Filter off all the pre- 
 cipitates and prove that they are carbonates. This may be 
 done by treating them with dilute acids, which decompose 
 them, causing an evolution of carbon dioxide, which can be 
 detected by passing a little .of it into lime-water. In some of 
 the cases mentioned the insoluble salts formed are basic car- 
 bonates, as, for example, those of copper and magnesium. The 
 salts of silver, calcium, and barium are the normal carbonates 
 Ag.CO a , BaC0 3 , and CaC0 3 . 
 
816 EXPERIMENTS TO ACCOMPANY CHAPTER XXV. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXV. 
 POTASSIUM SALTS. 
 
 Experiment 210. In preparing potassium iodide from 
 iodine and potassium hydroxide, proceed as follows : To 30 
 grams iodine use 15 grams hydroxide. Dissolve the latter in 
 100 cc. water. Add half this solution to the iodine in a por- 
 celain evaporating dish. Now slowly add the rest of the 
 liquid until the color disappears. Concentrate the liquid to 
 a syrupy consistence, add 1 gram finely-powdered charcoal, 
 mix, and evaporate to dryness. The residue is then heated 
 to redness in an iron vessel. After cooling extract with water. 
 
 Experiment 211. Potassium iodide can also be prepared 
 by the following method : Bring together in a capsule 200 
 grams water, 10 grams iron filings, and 40 grams iodine ; mix, 
 and heat gently. When the solution has become green, de- 
 cant, filter, and wash. Now heat the liquid nearly to boiling, 
 and gradually add a solution of 35 grams potassium carbonate 
 in 100 grams water. Filter, wash, and evaporate. 
 
 Experiment 212. Dissolve 50 grams potassium carbonate 
 in 500 to 600 cc. water. Heat to boiling in an iron or a silver 
 vessel, and gradually add the slaked lime obtained from 25 
 to 30 grams of good quicklime. During the operation the 
 mass should be stirred with an iron spatula. After the solu- 
 tion is cool, draw it off by means of a siphon into a bottle. 
 This may be used in experiments in which caustic potash is 
 required. 
 
 Experiment 213. Mix together 15 grams potassium nitrate 
 and 2.5 grams powdered charcoal. Set fire to the mass. 
 
 Experiment 214. Treat a quantity of wood ashes with 
 water. Filter, and examine by means of red litmus-paper. 
 Evaporate to dryness. What evidence have you that the 
 residue contains potassium carbonate ? 
 
 SODIUM SALTS. 
 
 Experiment 215. Make a supersaturated solution of sodi- 
 um sulphate by heating an excess of the salt with water at 33. 
 Filter the solution into small flasks and cork them. On re- 
 moving the corks and agitating the vessels, the salt will sud- 
 denly crystallize out. 
 
 Experiment 216. Pass carbon dioxide into a strong solu- 
 tion of ammonia (about 100 cc.) until it is no longer absorbed. 
 
CALCIUM SALTS. MAGNESIUM AND ITS SALTS. 817 
 
 A solution of acid ammonium carbonate is thus obtained. 
 Add this to a concentrated solution of sodium chloride as 
 long as a precipitate is formed. Filter off the precipitate, 
 and dry it by spreading it upon layers of filter-paper. Heat 
 some of the salt when dry, and determine whether the gas 
 given off is carbon dioxide or not. When gas is no longer 
 given off by heat, let the tube cool and examine the residue. 
 
 Experiment 217. Make ammonium sulphide thus : Divide 
 a given quantity of a solution of ammonia into two equal 
 parts. Saturate one half by passing hydrogen sulphide through 
 it, and then add the other half. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVI. 
 
 CALCIUM SALTS. 
 
 Experiment 218. Dissolve 10 to 20 grams of limestone or 
 marble in common hydrochloric acid. Filter, and evaporate 
 to dryness. Expose a few pieces of the residue to the air. 
 
 MAGNESIUM A$TD ITS SALTS. 
 
 ."Experiment 219. Make anhydrous magnesium chloride 
 thus : Dissolve 180 grams magnesia usta in ordinary hydro- 
 chloric acid ; shake the solution with an excess of magnesia 
 to remove iron and aluminium ; filter ; add 400 grams am- 
 monium chloride ; evaporate to dryness, keeping the mass 
 constantly stirred. The double salt thus formed must be 
 dried until a small specimen put in a test-tube is found not 
 to give off water when heated. The dry salt is then ignited 
 in a crucible placed in a furnace until ammonium chloride is 
 no longer given off, when the molten mass, which is anhydrous 
 magnesium chloride, is poured out on a stone and, after it is 
 broken up, it is put in a dry bottle provided with a good 
 stopper. 
 
 Experiment 220. Mix 6 parts anhydrous magnesium 
 chloride, 1 part of a mixture of sodium and potassium chlo- 
 rides, prepared by melting the two together and breaking up 
 after cooling, 1 part powdered fluor-spar, and 1 part sodium. 
 Throw this mixture all at once into a red-hot crucible in a 
 furnace, and cover the crucible. In a few moments a curious 
 sound is heard, and this indicates that the reaction is taking 
 place. Now take the crucible out of the furnace, and stir the 
 liquid in it with the aid of a clay pipe-stem. This causes the 
 
818 EXPERIMENTS TO ACCOMPANY CHAPTER XXVIII. 
 
 particles of the metal to collect in one large spherical mass. 
 After cooling, break the crucible, separate the metallic ball 
 from the slag, and wash it quickly with hydrochloric acid to 
 remove superficial impurities. If the slag is melted with a 
 quarter the weight of sodium that was used at first, a second 
 smaller piece of magnesium will be obtained. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVII. 
 ALUMINIUM CHLORIDE. 
 
 Experiment 221. Aluminium chloride is made thus : 
 Mix aluminic oxide with starch-paste ; form the mass into 
 small balls of the size of ordinary marbles ; ignite these in a 
 crucible in a furnace ; put them in a porcelain tube, and then 
 pass dry chlorine over them, at the same time heating the tube 
 to redness. The chloride will sublime in the front end of the 
 tube or in a receiver if the heat is sufficient. It can be puri- 
 fied by subliming it over heated iron or aluminium. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXVIII. 
 COPPER AND ITS SALTS. 
 
 Experiment 222. Cuprous chloride is best made as fol- 
 lows : Saturate a solution of 1 part sodium chloride and 2| 
 parts crystallized copper sulphate with sulphur dioxide. Fil- 
 ter, and wash with acetic acid. 
 
 SILVER AND ITS SALTS. 
 
 Experiment 223. Dissolve a ten or twenty-five cent piece 
 in dilute nitric acid. Dilute the solution to 200 to 300 cc. 
 with hot water. Add a hot solution of common salt until it 
 ceases to produce a precipitate. Filter off the white silver 
 chloride and wash with hot water. Dry the precipitate on 
 the filter, by placing the funnel with the filter and precipitate 
 in an air-bath heated to about 110. Eemove the precipitate 
 from the filter and put it into a porcelain crucible. Heat 
 gently with a small flame until the chloride is melted ; then 
 let it cool. Cut out a piece of sheet zinc large enough to 
 cover the bottom of the crucible, and lay it on the silver 
 chloride. Now add a little water and a few drops of dilute 
 
ZINC AND ITS SALTS TIN AND ITS COMPOUNDS. 819 
 
 sulphuric acid, and let the whole stand for twenty-four hours. 
 The silver chloride is reduced to silver, and zinc chloride is 
 formed : 
 
 Zn + SAgCl = ZnCl 2 + 2Ag. 
 
 Take out the piece of zinc and wash the silver with a little 
 dilute sulphuric acid, and then with water. Dissolve the 
 silver in dilute nitric acid and evaporate to dryness on the 
 water-bath, so that all the nitric acid is driven off. Dissolve 
 the residue in water, and put the solution either in a bottle 
 of dark glass or one wrapped in dark paper. 
 
 Experiment 224. To a solution of silver nitrate contain- 
 ing about 5 grams of the salt in 100 cc. water, add a few drops 
 of mercury, and let it stand. In a few days the silver will be 
 deposited in the form of delicate crystals. The formation is 
 called the "silver tree." 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXIX. 
 
 ZINC AND ITS SALTS. 
 
 Experiment 225. Heat a small piece of zinc on charcoal 
 in the oxidizing flame of the blow-pipe. The white fumes of 
 zinc oxide (philosopher's wool) will be seen, and the charcoal 
 will be covered with a film which is yellow while hot, but be- 
 comes white on cooling. 
 
 Experiment 226. Dissolve some zinc dust in a solution of 
 sodium hydroxide, and see whether hydrogen is given off. 
 
 MEKCURY AND ITS SALTS. 
 
 Experiment 227. Make a solution of mercurous nitrate 
 by treating at the ordinary temperature an excess of mercury 
 with nitric acid, which is not too concentrated ; and with this 
 solution study the conduct of mercurous salts. 
 
 Experiment 228. Heat some of the solution of mercurous 
 nitrate to boiling, then add a few drops of concentrated nitric 
 acid, and boil again. With the solution thus obtained study 
 the conduct of mercuric salts. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXX. 
 TIN AND ITS COMPOUNDS. 
 
 Experiment 229. Dissolve tin in hydrochloric acid and let 
 the product crystallize. 
 
820 EXPERIMENTS TO ACCOMPANY CHAPTER XXXI. 
 
 Experiment 230. Pass dry chlorine over granulated tin 
 contained in a retort connected with a receiver, using the ar- 
 rangement illustrated in Fig. 64, page 800. Kedistil the prod- 
 uct. Treat some of the liquid with water, and boil. 
 
 LEAD A^D ITS COMPOUNDS. 
 
 Experiment 231. Make specimens of lead chloride and 
 lead iodide, and crystallize them from water. 
 
 Experiment 232. Make lead sesquioxide by bringing to- 
 gether lead acetate and sodium hydroxide, and treating the 
 solution with a solution of sodium hypochlorite. 
 
 Experiment 233. Treat some red lead with dilute nitric 
 ?icid. Filter, wash, and treat the substance left on the filter 
 with hydrochloric acid. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXI. 
 CHROMIC ACID AND THE CHROMATES. 
 
 Experiment 234. Powder some chromic iron very finely. 
 Add 3 grams to a molten mixture of 3 grams each of potas- 
 sium carbonate, potassium hydroxide, and potassium nitrate, 
 heated in a porcelain crucible. After cooling treat the mass 
 with water. Potassium chromate is in the solution. 
 
 Experiment 235. To the solution of potassium chromate 
 obtained in the last experiment add nitric acid to decompose 
 the unacted-upon potassium carbonate, and give the solution 
 an acid reaction. The color will change from yellow to red. 
 The red color indicates the presence of the dichromate. 
 
 Experiment 236. Treat a solution of 10 to 20 grams 
 potassium dichromate with potassium hydroxide , until the 
 color becomes pure yellow, and evaporate to crystallization. 
 
 Experiment 237. Make a solution of potassium dichro- 
 mate saturated at the ordinary temperature. Pour into this 
 1 times its volume of ordinary concentrated sulphuric acid. 
 After the liquid cools, and the chromium trioxide separates, 
 filter with the aid of a filter-pump through glass-wool. 
 
 Experiment 238. To a solution of potassium dichromate 
 add some hydrochloric acid and a little alcohol. On boiling, 
 the alcohol is oxidized, and the solution now contains chromic 
 chloride. 
 
MANGANESE AND ITS COMPOUNDS PLATINUM. 821 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXII. 
 
 MANGANESE AND ITS COMPOUNDS. 
 
 Experiment 239. Make and crystallize some manganous 
 chloride by treating manganese dioxide with hydrochloric 
 acid. Also make some manganous sulphate by heating man- 
 ganese dioxide with sulphuric acid. Use these solutions for 
 the purpose of studying the conduct of manganous salts. 
 
 Experiment 240. In a small porcelain crucible heat to- 
 gether 5 grams manganese dioxide, 5 grams solid potassium 
 hydroxide, and 2J grams potassium chlorate. When the 
 mass has turned green, dissolve the contents in water and 
 neutralize most of the free alkali in the solution. Or pass 
 carbon dioxide through the solution without boiling. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXIII. 
 IRON AND ITS COMPOUNDS. 
 
 Experiment 241. Make ferric chloride by heating the 
 purest iron wire in a current of chlorine. Also make a solu- 
 tion by dissolving iron in hydrochloric acid, and oxidizing the 
 solution with nitric acid. 
 
 Experiment 242. Make ferrous sulphate by dissolving iron 
 in dilute sulphuric acid, and evaporating to crystallization. 
 Dissolve equivalent quantities of ferrous sulphate and ammo- 
 nium sulphate, and evaporate to crystallization. 
 
 Experiment 243. Dissolve ferric hydroxide in sulphuric 
 acid, and evaporate to dryness. 
 
 EXPERIMENTS TO ACCOMPANY CHAPTER XXXIV. 
 PLATINUM. 
 
 Experiment 244. Prepare a solution of platinic chloride 
 as follows : Heat platinum in a flask with concentrated nitric 
 acid, adding from time to time a few drops of hydrochloric 
 acid. After the metal is dissolved evaporate to dryness. Dis- 
 solve in water and filter. 
 
 CONCLUSION. 
 
 At the end of each chapter treating of the metallic or base- 
 forming elements there is given a list of such reactions as are 
 
822 INORGANIC CHEMISTRY, 
 
 of special value for analytical purposes, together with such ex- 
 planatory statements as seem called for. The student who is 
 engaged in analytical work will generally find these explana- 
 tions sufficient to enable him to keep his ideas clear in regard 
 to the reactions with which he is dealing, provided, at the 
 same time, he carefully studies the chapter to which the ex- 
 planations form an appendix. As an introduction to analyti- 
 cal work, a general study of chemical reactions is necessary, 
 and the fuller this is the better. There are many small books 
 in existence in which good directions are given for work of 
 this kind. The student is, however, advised to supplement 
 the book he may be using by such experiments as may suggest 
 themselves on reading the corresponding chapters in this 
 book. The directions there found will generally be quite suf- 
 ficient for the purpose, and it is therefore not considered 
 necessary to give more specific directions in this place. In 
 general, the more the student occupies himself in the labora- 
 tory with chemical substances the more rapidly will his chemi- 
 cal ideas grow. But it is necessary that he should avoid 
 working by "rule of thumb," and for this purpose constant 
 reference to some larger text-book in which the relations be- 
 tween the substances and the reactions he is dealing with are 
 discussed in a broad way is of the highest importance. 
 
APPENDIX II. 
 
 THE following tables have been prepared to furnish those 
 who may use this book as a reference-book in connection with 
 laboratory work with such information as they may need, but 
 which they can perhaps get only from works not always easily 
 available. It was not considered advisable to include tables 
 required in gas analysis and general quantitative analysis. 
 For part of the material I am indebted to that excellent 
 work, the Chemiker-Kalender, by E. Biedermann. In the 
 tables of solubilities the substances have been arranged 
 alphabetically as salts of acids, as for example the acetates, 
 chlorides, sulphates, and sulphydrates. 
 
 J. ELLIOTT GILPIK. 
 
824 
 
 APPENDIX II. 
 
 ATOMIC WEIGHTS. 
 (From " The Constants of Nature," by F. W. Clarke, 1897.) 
 
 
 H = l 
 2691 
 
 O = 16 
 27.11 
 
 Molybdenum. . . . 
 
 H = 1 
 95.26 
 
 Antimony 
 
 119 52 
 
 12043 
 
 Neodymium 
 
 139.70 
 
 
 9 
 
 ? 
 
 Nickel 
 
 58 24 
 
 
 74 44 
 
 75 01 
 
 Nitrogen 
 
 13 93 
 
 Barium . . . .... 
 
 136 39 
 
 137 43 
 
 
 189 55 
 
 Bismuth 
 
 206 54 
 
 208 11 
 
 Oxygen 
 
 15 88 
 
 
 10 86 
 
 10 95 
 
 Palladium . . . 
 
 105 56 
 
 Bromine . 
 
 79 34 
 
 79 95 
 
 Phosphorus .... 
 
 3079 
 
 
 111 10 
 
 111.95 
 
 Platinum 
 
 193 41 
 
 
 131 89 
 
 132.89 
 
 Potassium 
 
 38.82 
 
 Calcium. . .... 
 
 39 76 
 
 40 07 
 
 Praseodymium . 
 
 142 50 
 
 Cirbon 
 
 11 92 
 
 12 01 
 
 Rhodium 
 
 102 23 
 
 Cerium 
 
 139 10 
 
 140.20 
 
 Rubidium 
 
 84 78 
 
 Chlorine 
 
 35 18 
 
 35 45 
 
 Ruthenium . . . 
 
 100 91 
 
 Chromium 
 
 51 74 
 
 52 14 
 
 Samarium 
 
 149 13 
 
 Cobalt . . 
 
 58 49 
 
 58 93 
 
 
 43 78 
 
 Columbium 
 
 93 02 
 
 93 73 
 
 Selenium . . . . 
 
 78 42 
 
 CoDoer . 
 
 63.12 
 
 63.60 
 
 Silicon 
 
 28.18 
 
 Erbium . . 
 
 165 06 
 
 166 32 
 
 Silver 
 
 107.11 
 
 
 18.91 
 
 19.06 
 
 
 22.88 
 
 Gadolinium 
 
 155 57 
 
 156 76 
 
 Strontium . 
 
 86 95 
 
 Gallium 
 
 69 38 
 
 69 91 
 
 Sulphur 
 
 31 83 
 
 Germanium 
 
 71 93 
 
 72.48 
 
 Tantalum 
 
 181 45 
 
 
 9.01 
 
 9 08 
 
 
 126 52 
 
 Gold 
 
 195 74 
 
 197 23 
 
 Terbium 
 
 158 80 
 
 
 ? 
 
 ? 
 
 Thallium 
 
 202.61 
 
 Hydrogen 
 
 1 000 
 
 1 008 
 
 Thorium 
 
 230 87 
 
 Indium 
 
 112 99 
 
 113 85 
 
 Thulium 
 
 169 40 
 
 Iodine ... 
 
 125 89 
 
 126.85 
 
 Tin 
 
 118 15 
 
 Iridium 
 
 191 66 
 
 193 12 
 
 Titanium . 
 
 47 79 
 
 Iron 
 
 55 60 
 
 56.02 
 
 Tungsten 
 
 183 43 
 
 Lanthanum 
 
 137 59 
 
 138. '14 
 
 Uranium 
 
 237 77 
 
 Lead 
 
 205.36 
 
 206.92 
 
 
 50 99 
 
 Lithium 
 
 6 97 
 
 7 03 
 
 Ytterbium 
 
 171 88 
 
 Magnesium 
 
 24 10 
 
 24.28 
 
 Yttrium 
 
 88 35 
 
 Manganese 
 
 54 57 
 
 54 99 
 
 Zinc 
 
 64 91 
 
 Mercury.. 
 
 198.49 
 
 200.00 
 
 Zirconium . . 
 
 89.72 
 
 O = 16 
 
 95.99 
 
 140.80 
 
 58.69 
 
 14.04 
 
 190.99 
 
 16.00 
 
 106.36 
 
 31.02 
 
 194.89 
 
 39.11 
 
 143.60 
 
 103.01 
 
 85.43 
 
 101.68 
 
 15026 
 
 44.12 
 
 79.02 
 
 28.40 
 
 107.92 
 
 23.05 
 
 87.61 
 
 32.07 
 
 182.84 
 
 127.49 
 
 160.00 
 
 204.15 
 
 232.63 
 
 170.70 
 
 119.05 
 
 48.15 
 
 184.83 
 
 239.59 
 
 51.38 
 
 173.19 
 
 89.02 
 
 65.41 
 
 90.40 
 
APPENDIX II. 
 
 825 
 
 MELTING-POINTS AND BOILING POINTS OF THE 
 
 ELEMENTS. 
 
 Where the values of different observers vary, the highest and lowest 
 ill be given. 
 
 
 Melting-point. 
 
 Boiling-point. 
 
 Aluminium. ..... 
 Antimony 
 
 600-850 
 425-450 
 
 n.v. at white heat 
 1090-1700 
 
 
 189 6 
 
 186 9 
 
 Al'senic . . 
 
 under pressure at red heat 
 
 sublimes at 450 
 
 Barium 
 
 red heat 
 
 
 Beryllium 
 
 below 1000 
 
 
 Bismuth. . . ... 
 
 260-270 
 
 1090-1450 
 
 Boron 
 
 
 
 Bromine 
 
 7.3 
 
 59-63 
 
 Cadmium 
 
 310-320 
 
 760-860 
 
 CflBsiuni ... 
 
 26 5 
 
 
 Calcium 
 
 red heat 
 
 n v 
 
 
 infusible 
 
 
 Cerium 
 
 between Sb and Ag 
 
 
 Chlorine 
 
 solidifies at 102 
 
 33.6 
 
 Chromium 
 
 higher than platinum. 
 
 
 
 1500-1800 
 
 
 Copper 
 
 1000-1330 
 
 
 Fluorine 
 
 
 
 Gallium * 
 
 30.15 
 
 
 Germanium 
 
 about 900 
 
 
 Gold 
 
 1035-1250 
 
 
 Indium 
 
 176 
 
 red heat 
 
 Iodine 
 
 113-115 
 
 above 200 
 
 Indium 
 
 1950-2500 
 
 
 Iron pure. . 
 
 1500-1800 
 
 
 
 322-335 
 
 between 1450-1600 
 
 Lithium 
 
 180 
 
 
 Magnesium 
 Manganese 
 
 500-800 
 1600 
 
 about 1100 
 
 Jdercury 
 
 38 50 to 40 5 
 
 357 25 
 
 Molybdenum 
 Nickel 
 
 does not melt at a white heat 
 1450-1600 
 
 
 Nitrogen ., 
 
 solidifies at 203 at 60-70 mm. 
 
 1944 
 
 Osmium . ... 
 
 2500 
 
 
 Oxveen . 
 
 below 21] 5 at 9 mm 
 
 181 4 
 
 Palladium 
 
 1360-1950 
 
 
 Phosphorus . 
 
 44 2 
 
 287-290 
 
 Platinum 
 
 1460-2200 
 
 
 Potassium 
 
 58 62 5 
 
 667 731 
 
 Rhodium. . . 
 
 2000 
 
 
 Rubidium 
 
 38 5 
 
 
 Selenium . . . 
 
 217 
 
 bet 676 and 683 
 
 Silicon .... 
 
 about 1400 
 
 
 Silver 
 
 954-1040 
 
 
 Sodium 
 
 90 97 6 
 
 742-954 
 
 Strontium. . 
 
 red heat 
 
 
 Sulphur, rhombic, 
 mono. .. 
 ' ' amor. . . . 
 Tellurium 
 Tin 
 
 113-115 
 120 
 above 120 
 425-525 
 226 5-235 
 
 447-448.4 
 bet 1450 and 1600 
 
 Zinc 
 
 412-420 
 
 891 1040 
 
 
 
 
826 
 
 APPENDIX II. 
 
 ttj 
 
 53*3 
 
 -:-:dH -o 
 
 dgS 
 
 =>. sw 
 
 ^ I 
 
 Melting'- 
 
 OOi W 
 <M CO ~ 
 
 o 1 
 
 bfri 
 
 o 
 
 o o, 
 
 O 
 ? cS 
 
 W B 55 
 
 te 
 
 e 
 
 at 
 ce 
 
 . 
 
 ..fi| Iss^l 
 
 ll'" V 
 
 
 o .iJ 
 
 cid. 
 acid 
 
 a ^ ^ :: o 
 
 Sodium 
 Urauyl 
 
 
 2 1SJ 
 
 1 Sill 
 ilgfS^ 
 
APPENDIX II. 
 
 827 
 
 
 
 I 5 
 
 - 
 
 t-sfii. 
 
 S 
 
 88888 
 
 8 
 
 ; -5 
 
 888880' 'S | ^ 
 
 I* 
 
 J 
 
 O*T3 
 
 T-t 
 
 a 
 
 iri 
 :! I 
 
 I QJ ,Q 
 
828 
 
 APPENDIX II. 
 
 flj 
 
 *$t 
 
 ' * *" d 
 
 
 
 80^ 
 
 o 
 
 (D - 
 
 a 
 
 Sao - . 
 is^-*! 1 * 
 
 o'B 1o 
 
 o .3 a 
 
 O in 
 
 8 o^ 
 
 ^ ^^-r-J 
 
 10 
 
 t s I 
 
 , O TH O 
 
 ! O \ O 
 
 00 
 
 00 05 fc.'-JrH 
 
 JH ^ a>0'-'co 
 
 C^ ^^ *^ *H CO 
 
 ^^ ^D 03 
 
 . . > i-H T-H 
 
 O O P C5 
 
 5 8 SS5 
 
 I? 
 
 5 
 
 3 
 
 S 
 
 O q o 
 
 H?W W 
 
 * - 
 O PH 
 
 B8 
 
 O " 
 
 
 iS.'*5 ?" a 
 sill S? 
 
 s S-s ^ 
 
 si "-'S S ^ u. 
 
 O S3 OQ S ^ 
 
 S-IS3 : 
 
 M C2 rt 
 
 O <U .S .S t- 
 
 7='C - t- O 
 
 "^r o 
 
 c c^s ^ 
 
 3 a 
 
 o b^ 53 6 
 
 gWP,g*1 <1 <j X^ 
 
 111 
 
 a s "= o 
 s ii I 
 
APPENDIX II. 
 
 829 
 
 cohol, Acids, 
 lkalies, and 
 Other Sol- 
 vents. 
 
 s 
 
 ft 
 
 o S 
 
 03 O 
 
 
 i 
 
 -a i i* 
 
 ? 
 
 \t 
 
 r~ 
 
 i .Si : 
 
 W 
 
 SE S 
 
 13 ' 13 
 
 q q 
 
 as 
 
 e 
 
 Q 
 
 T3 03 
 
 O O 
 
 J Q 
 Q 
 
 o o 
 
 00 
 
 ww 
 
 - B .8J., 
 
 
 o fe 
 
 ?l 
 II 
 II 
 
 il 
 
 a 
 
 e 
 chlori 
 chlor 
 
 " o> fl> 
 
 O> Q^ I T3 r O 
 
 CJ ^ T3 0> - "^ 
 
 ?S"S ? 
 
 3-31 a !p?s 
 
 o -3 gg^ ggg 
 .Sig ^3^3-sS 
 BBS o -=> 
 
 III I 
 
 060 o 
 
 o 
 ^ 
 'C 
 
 o 
 
 1 
 
 
 3 
 c 
 
 osphorus 
 
 num chloride 
 ssium (Chlorid 
 
 II 
 
 B-3 
 
 T 
 
 GQ 
 
830 
 
 APPENDIX II. 
 
 & 
 
 
 
 ww 
 
 2 2 
 
 8O 
 O 
 C2 O 
 
 a :^ 
 
 s ~ 
 
 
 
 M ^l 
 
 G 
 O 
 
 I? 
 
 -3 
 
 
 
 
 
 3 
 
 S w 
 
 7" i 
 
 O 
 
 s 
 
 a 5iN 
 
 OQ .S "<*&> ^ 
 
 O j 
 
 ^ o 
 
 S 
 
 ^ eo 
 
 o o 
 
 OQ CO 
 
 
 O 
 
 -r^p? o9' 
 
 HLORIDES (cont'd) : 
 Silicon chloroform 
 
 . . . 
 
 
 
 
 
 . . . 
 
 
 
 
 
 
 
 
 S 
 
 s 
 
 
 
 
 
 
 
 c3 
 
 
 
 
 
 *r^ . 
 
 2 
 
 O 
 
 CJS 4 
 
 S 
 
 j 
 
 ! C 
 
 
 
 1 
 
 % 
 ?-Sb 
 
 Is 
 
 O rf - 
 
 <uS^ ^ 
 
 c 5- ? 'S 
 
 "C^^ o 
 
 0,0 j 3 
 
 ^ o o 7s 
 
 o .- o 
 
 o ^ S-^ 
 
 um 
 um 
 
 l 
 
 h 
 e 
 
 ^3 0^3 
 
 l',t 
 
 H 
 
 c^ eg 
 
 ori 
 ac 
 ri 
 lor 
 
 l 
 x 
 lo 
 h 
 e 
 
 m c 
 n h 
 c 
 
 Tha 
 Tun 
 Uram h 
 Vanadium c 
 Zinc chlorid 
 
 TE 
 
 - S^ 
 
 s 2 1 
 
 BD-3-gP 
 
 1a 
 Sl-SiS 
 
 gl^^ 
 
 Q 
 
APPENDIX II. 
 
 831 
 
 
 a co ~ ^' , " d a ' 
 
 13 1 
 
 ; _ 
 
 'p.l.S 
 
 
 
 
 
 !, 
 
 .s 
 
 O O O O GO 
 T-I 00-^QOTH 
 
 I I I I I 
 
 . 
 
 F 
 
 s 
 
 >> 
 
 
 
 s 
 ft 
 
 <S o' 
 
 ^ O 
 
 r5 
 
 
 tf 
 
 "Ili 
 
 00 
 
 
 5 A 5 
 
832 
 
 APPENDIX II. 
 
 
 So : s 
 
 1 O . 
 
 * rj 
 
 : 8 a S 
 
 oo ; o 
 01 g 
 
 
 .i 
 
 a 
 a. 
 
 8 . 
 
 T-l CU 
 
 W 
 
 1 
 
 I a 
 
 o o 
 
 g g 
 
 :H > 
 
 }>i i ( _ H H {> 
 
 I ooooo 
 
 ^-^ C^ X ^> '^7 
 
 o- 
 
 HH P 
 
 cont'd} 
 ydrox 
 
 i i : :S 
 
 4i- i 
 
 gill' 
 
 'Ha 
 
 Ms-ll 
 
 roxide 
 
 ydroxide. 
 ydroxide. 
 roxide 
 ydroxide. 
 
 de 
 
 S : 
 
 chlorite 
 pochlori 
 
 ! t ; 
 
 . [O 
 
 ' T3 oJ ' 
 
 03 .2 12 * 
 
 ^5 S o oi 
 
 O 3'~T3 
 
 11 1 
 11 tig 
 
 " CJ _rt o c 
 
 0^(X 
 
 : 
 
 mangan 
 
 ANGANAT 
 
 Potassiu 
 
APPENDIX II. 
 
 833 
 
 fl. 
 
 F 
 
 ~o . .0 - . 
 
 ' 
 
 g o 
 
 so 
 
 iti I* 
 
 ^ b 
 
 \a 
 
 <M O 
 
 A' ^ a S * 
 
 O g, . ^, 
 
 H 
 
 '!. 
 
 a 
 
 a. 
 UP 
 
 C o 
 
 s 
 "S 
 
 a 
 
 g co ^ 
 
 w 
 
 (53 
 
 ooo 
 
 C* CT. 
 
 + o 
 
 
 +1 ! 8 
 
 PQQO 
 
 Q> ' ' 
 
 l iig 
 
 ^ 
 
 ' P ^sg 
 
 ^ i-T! S C ^ H 
 
 a I 
 
 w -5 
 
 ^<j -^pppqpq 
 
 O 
 
 .2 g a 
 
 s 3 
 
 2 d 
 
 2 I 
 
 83 Oj 
 
 o o 
 
834 
 
 *g. 
 
 3S 
 <-r 
 
 Al 
 O 
 
 APPENDIX II. 
 
 * M&rl ^O 
 
 * 
 
 ft . 
 
 
 
 - .2 
 
 I* 
 
 g o 8 
 
 -H <M , S 
 
 
 OHCC 
 
 
 
 
 S oj 
 o-S 
 
 S w' 6 (n-S-c s g^^^S o SS ^^ g as 
 
 Hl|llftllli|S i &-s s Pi 
 
 r,SS>< 0tt rS^^a*ooa 'x 3 :, 5 'S 
 
 
 
 
 
 wcocJ^rf-3 hfiD 3Q.S 
 
 si 1111 1 II H| 
 
 !^{5 o P^S 
 
APPENDIX II. 
 
 835 
 
 acu no 
 
 a a " 
 
 99 oo* 
 
 : a -s 
 
 og.2 
 
 a o 
 
 o ^** 
 
 T3 
 
 | d 
 GO 
 
 13 
 
 d^ - 
 
 a i - 
 
 A G. O. 
 
 a a a 
 
 3 
 s^i 
 
 i*i 
 
 i 1 
 
 
 
836 
 
 APPENDIX II. 
 
 ... 
 
 ta 
 
 3 i 
 
 H 
 
 s ' II 
 
 * 7>-o 
 " .2 %<*. 
 
 c 
 
 fe 
 
 I' 
 
 c 
 
 a. 
 I? 
 
 I 
 
 *8 
 
 >& & 
 
 
 
 -*2 " oD 
 
 < o 
 
 ^3 Q^3 
 
 PH PQ P, 
 
 .2 a 
 
 
 1 
 
 H 
 
 
 
 ., 
 
 sill 
 
 PL, 
 
 ia 
 
 - 
 
 ,5 a 
 
 I* 
 It 
 
 i 
 
 OD OQ 
 
APPENDIX II. 
 
 837 
 
 fc 
 
 IN 
 
 .3 -3 
 
 
 1? 
 
 :oq oq 
 
 J3S 5 
 
 rn -4-j ~*j? gg ^ 
 
 oqqo 
 ggSS 
 
 3 its 
 o :o 
 
 5 jg" 
 
 - s t 
 
 ! 
 
 ?, 5 
 '3 S 
 
 . 
 1 
 
 I lit I 1 
 
 rf Ct t>^ r p3 *~^ rl 
 
 O 
 
 1 
 
 tl 
 
 -2 >> 
 
 ^^5 
 
 ."3 ^ 
 
 83 
 
 
 OQ 
 
 <JPH 
 
838 
 
 APPENDIX II. 
 
 Alcohol, Acids, 
 Alkalies, and 
 Other Sol- 
 vents. 
 
 S 8 i 
 
 o . a 
 
 i 
 
 7^ g 
 
 c 
 
 
 
 * 
 
 'o 
 
 go 
 
 O !<, 
 
 T-l TH <M 
 
 q O q 
 
 W W 
 
 
 
 1 
 
 
 + 
 
 1 
 
 H 
 
 S3 
 
 NH ffi 
 
 MJ r+4 
 
 | 
 
 CQ 
 
 M 
 
 
 f sp-sil f 
 
 3 2 ^5 rH Jd ~^ M 
 
 8*SB-M'a.l9i fi 
 
 ulph 
 ide 
 e 
 
 ocy 
 yau 
 
 l t 
 
 5s 
 
 ^3 
 
 ill 
 
 2 * d 
 
 QQ 
 
 
 " 
 
 
 33 v>2<2% 
 
 afli-si5lsa 
 
 2 tl5 iill1 
 
 ^..2^o'SS2-S^So 
 
 PnCCtSJ gPn^ g<lPn &4^ 
 
 otass 
 ver 
 c s 
 
 QQ 
 
 02 
 
 "^ * P-i a 
 
 "all 
 
 S o'S 
 E^ 
 
 H 
 
APPENDIX II. 
 
 839 
 
 WEIGHT OF 1 LITER OF GAS AT AND 760 MM. 
 PRESSURE IN LATITUDE 45. 
 
 Name. 
 
 Formulae. 
 
 Molecular Weight. 
 
 Weight of 1 Liter, 
 in Grains. 
 
 
 CH 
 
 2584 
 
 1 1611 
 
 Ammonia . . . * 
 
 NH 3 
 
 16 93 
 
 7607 
 
 Carbon dioxide 
 
 CO 2 
 
 43 68 
 
 1 962 
 
 " monoxide 
 
 CO 
 
 Clo 
 
 27.80 
 70 36 
 
 1.249 
 3 1617 
 
 Hydrochloric acid. . . . 
 
 HC1 
 
 36 18 
 
 1 625 
 
 
 H 2 
 
 2 
 
 089873 
 
 
 CH 4 
 
 15 92 
 
 0.7153 
 
 Nitric oxide 
 
 NO 
 
 29 81 
 
 1 339 
 
 Nitrogen 
 
 No 
 
 2786 
 
 1 252 
 
 
 NO 2 
 
 45 69 
 
 2053 
 
 Nitrous oxide 
 
 N 2 O 
 
 43 74 
 
 1 9655 
 
 Oxygen 
 
 Oo 
 
 31 76 
 
 1 4^9 
 
 Sulphur dioxide 
 
 SO 2 
 
 63 59 
 
 28575 
 
 
 
 
 
 Morley's values for oxygen and hydrogen were adopted, and the others 
 calculated on this basis. 
 
 APPROXIMATE COMPOSITION OF A NUMBER OF ALLOYS. 
 
 
 
 Pro] 
 
 jortion 
 
 sin 
 
 ivhic 
 
 h Con 
 
 ibim 
 
 ition t 
 
 akes 
 
 Place. 
 
 
 
 Cu 
 
 Zn 
 
 Sn 
 
 Ni 
 
 Ag 
 
 P 
 
 Cd 
 
 Bi 
 
 Pb 
 
 Al 
 
 Sb 
 
 Brflss ......... 
 
 66 
 
 33 
 
 
 
 
 
 
 
 3 
 
 
 
 Gun-metal 
 
 90 
 
 
 10 
 
 
 
 
 
 
 
 
 
 Bell metal 
 German silver 
 Silver coins . . . 
 
 77 
 60 
 10 
 
 20 
 
 23 
 
 20 
 
 90 
 
 
 
 
 
 
 
 Wood's metal 
 Rose's metal 
 
 
 
 13.5 
 25 
 
 
 
 
 10 
 
 49.8 
 50 
 
 26.7 
 25 
 
 
 
 Britannia metal 
 Pewter 
 
 7-8 
 1 
 
 2 
 
 77-90 
 2-20 
 
 
 
 
 
 
 
 
 10-15 
 
 Aluminium bronze 
 Art bronze . . 
 
 90 
 86 6 
 
 3 3 
 
 6.6 
 
 
 
 
 
 
 3 3 
 
 10 
 
 
 Soft solder 
 
 
 
 50 
 
 
 
 
 
 
 50 
 
 
 
 Babbitt metal 
 Nickel coins . . .. 
 
 3.7 
 
 75 
 
 
 88.9 
 
 95 
 
 
 
 
 
 
 
 7.4 
 
 Phosphor bronze .. 
 
 90 
 
 
 9 
 
 
 
 0.5 
 
 
 
 
 
 
840 APPENDIX II. 
 
 FREEZING-MIXTURES. 
 
 The mixtures given below will be found useful in cases where low 
 temperatures are desired. 
 
 Temperature falls 
 from to 
 
 1 pt. potassium sulphocyanate -f- 1 pt. water -|- 18 21 
 
 1 pt. sodium chloride + 3 pts. snow 21 
 
 3 pts. crystallized calcium chloride -j- 1 pt. snow 48.5 
 
 1 pt. snow -f- 1 pt. dilute sulphuric acid. ...... -f 5 41 
 
 1 pt. potassium chloride + 4 pts. water. ./ 11.8 
 
 3 pts. sodium nitrate + 4 pts. water -(- 13.2 5.3 
 
 Solid carbonic acid + ether 100 
 
 TABLE OF WEIGHTS AND MEASURES. 
 
 WEIGHT. 
 
 1 gram = 15.432 grains ; 
 
 1 grain = 0.06479 grams ; 
 
 1 ounce avoirdupois = 28.3495 grams ; 
 
 1 ounce troy = 31.1035 grams ; 
 
 1 pound avoirdupois = 0.45359 kilogram ; 
 
 1 pound troy = 0.37324 kilogram ; 
 
 1 kilogram = 2.2046 pounds avoirdupois. 
 
 LENGTH. 
 
 1 millimeter = 0.03937 inches ; 
 linch = 25.3995 millimeters; 
 
 1 meter 3.280899 feet ; 
 1 foot = 0.30479 meter ; 
 
 1 mile = 1609.31 meters. 
 
 VOLUME. 
 
 1 cubic centimeter 0.061027 cubic inches ; 
 1 cubic inch = 16 38,6 cubic centimeters ; 
 
 1 cubic centimeter = 0.03519 fluid ounce ; 
 1 fluid ounce = 28.41 cubic centimeters ; 
 
 1 liter = 1.76 pints ; 
 
 1 liter = 0.2201 gallon; 
 
 1 gallon = 4.543 liters. 
 
APPENDIX II. 
 
 841 
 
 COMPARISON OF THE TWADDELL SCALE WITH THE 
 BAUMft AND GAY-LUSSAC SCALES. 
 
 Tw. 
 
 B. 
 
 G.-L. 
 
 Tw. 
 
 B. 
 
 G.-L. 
 
 Tw. 
 
 B. 
 
 G.-L. 
 
 Tw. 
 
 B. 
 
 G.-L. 
 
 
 
 
 
 1.000 
 
 44 
 
 26.0 
 
 1.220 
 
 88 
 
 44.1 
 
 1.440 
 
 131 
 
 57.1 
 
 1.655 
 
 1 
 
 0.7 
 
 1.005 
 
 45 
 
 26.4 
 
 1.225 
 
 89 
 
 444 
 
 1.445 
 
 132 
 
 57.4 
 
 1.660 
 
 2 
 
 1.4 
 
 1.010 
 
 46 
 
 26.9 
 
 1.230 
 
 90 
 
 44.8 
 
 1.450 
 
 133 
 
 57.7 
 
 1.665 
 
 3 
 
 2.1 
 
 1.015 
 
 47 
 
 27.4 
 
 1.235 
 
 91 
 
 45.1 
 
 1.455 
 
 134 
 
 57.9 
 
 1.670 
 
 4 
 
 2.7 
 
 1.020 
 
 48 
 
 27.9 
 
 1.240 
 
 92 
 
 45.4 
 
 1.460 
 
 135 
 
 58.2 
 
 1.675 
 
 5 
 
 3.4 
 
 1.025 
 
 49 
 
 28.4 
 
 1.245 
 
 93 
 
 45.8 
 
 1.465 
 
 136 
 
 584 
 
 1.680 
 
 6 
 
 4.1 
 
 1.030 
 
 50 
 
 28.8 
 
 1.250 
 
 94 
 
 46.1 
 
 1.470 
 
 137 
 
 58.7 
 
 1.685 
 
 7 
 
 4.7 
 
 1.035 
 
 51 
 
 29.3 
 
 1.255 
 
 95 
 
 46.4 
 
 1.475 
 
 138 
 
 58.9 
 
 1.690 
 
 8 
 
 5.4 
 
 1.040 
 
 52 
 
 29.7 
 
 1.260 
 
 96 
 
 46.8 
 
 1.480 
 
 139 
 
 59.2 
 
 1.695 
 
 9 
 
 6.0 
 
 1.045 
 
 53 
 
 30.2 
 
 1.265 
 
 97 
 
 47.1 
 
 1.485 
 
 140 
 
 59.5 
 
 1.700 
 
 10 
 
 6.7 
 
 1.050 
 
 54 
 
 30.6 
 
 1.270 
 
 98 
 
 47.4 
 
 1.490 
 
 141 
 
 59.7 
 
 1.705 
 
 11 
 
 7.4 
 
 1.055 
 
 55 
 
 31.1 
 
 1.275 
 
 99 
 
 47.8 
 
 1.495 
 
 142 
 
 60.0 
 
 1.710 
 
 12 
 
 8.0 
 
 1.060 
 
 56 
 
 31.5 
 
 1.280 
 
 100 
 
 48.1 
 
 1.500 
 
 143 
 
 60.2 
 
 1.715 
 
 13 
 
 8.7 
 
 1.065 
 
 57 
 
 32.0 
 
 1.285 
 
 101 
 
 48.4 
 
 1.505 
 
 144 
 
 60.4 
 
 1 720 
 
 14 
 
 9.4 
 
 1.070 | 
 
 58 
 
 32.4 
 
 1.290 
 
 102 
 
 48.7 
 
 1.510 
 
 145 
 
 60.6 
 
 1.725 
 
 15 
 
 10.0 
 
 1.075 
 
 59 
 
 32.8 
 
 1.295 
 
 103 
 
 49.0 
 
 1.515 
 
 146 
 
 60.9 
 
 1.730 
 
 16 
 
 10.6 
 
 1.080 I 
 
 60 
 
 33.3 
 
 1.300 
 
 104 
 
 49.4 
 
 1.520 
 
 147 
 
 61.1 
 
 1.735 
 
 17 
 
 11.2 
 
 1.085 i 
 
 61 
 
 33.7 
 
 1.305 
 
 105 
 
 49.7 
 
 1.525 
 
 148 
 
 61.4 
 
 1.740 
 
 18 
 
 11.9 
 
 1.090 
 
 62 
 
 34.2 
 
 .310 
 
 106 
 
 50.0 
 
 1.530 
 
 149 
 
 61.6 
 
 1.745 
 
 19 
 
 12.4 
 
 1.095 
 
 63 
 
 34.6 
 
 .315 
 
 107 
 
 50.3 
 
 1.535 
 
 150 
 
 61.8 
 
 1 750 
 
 20 
 
 13.0 
 
 1.100 
 
 64 
 
 35.0 
 
 .320 
 
 108 
 
 50.6 
 
 1.540 
 
 151 
 
 62.1 
 
 1.755 
 
 21 
 
 13.6 
 
 1.105 ! 
 
 65 
 
 35.4 
 
 .325 
 
 109 
 
 50.9 
 
 1.545 
 
 152 
 
 62.3 
 
 1.760 
 
 23 
 
 14.2 
 
 1.110 ! 
 
 66 
 
 35.8 
 
 .330 
 
 110 
 
 51.2 
 
 1.550 
 
 153 
 
 62.5 
 
 1.765 
 
 23 
 
 14.9 
 
 1.115 
 
 67 
 
 36.2 
 
 .335 
 
 111 
 
 51.5 
 
 1.555 
 
 154 
 
 62.8 
 
 1.770 
 
 24 
 
 15.4 
 
 1.120 
 
 68 
 
 36.6 
 
 .340 
 
 112 
 
 51.8 
 
 1.560 
 
 155 
 
 63.0 
 
 1.775 
 
 25 
 
 16.0 
 
 1.125 
 
 69 
 
 37.0 
 
 .345 
 
 113 
 
 52.1 
 
 1.565 
 
 156 
 
 63.2 
 
 1.780 
 
 26 
 
 16.5 
 
 1.130 
 
 70 
 
 37.4 
 
 1.350 
 
 114 
 
 52.4 
 
 1.570 
 
 157 
 
 63.5 
 
 1.785 
 
 27 
 
 17.1 
 
 1.135 
 
 71 
 
 37.8 
 
 1.355 
 
 115 
 
 52.7 
 
 1.575 
 
 158 
 
 63.7 
 
 1.790 
 
 28 
 
 17.7 
 
 1.140 
 
 72 
 
 38.2 
 
 1.360 
 
 116 
 
 53.0 
 
 1.580 
 
 159 
 
 64.0 
 
 1.795 
 
 29 
 
 18.3 
 
 1.145 
 
 73 
 
 38.6 
 
 1.365 
 
 117 
 
 53.3 
 
 1.585 
 
 160 
 
 64.2 
 
 1.800 
 
 30 
 
 18.8 
 
 1.150 
 
 74 
 
 39.0 
 
 1.370 
 
 118 
 
 53.6 
 
 1.590 
 
 161 
 
 64.4 
 
 1.805 
 
 31 
 
 19.3 
 
 1.155 ! 
 
 75 
 
 39.4 
 
 1.375 
 
 119 
 
 53.9 
 
 1.595 
 
 162 
 
 64.6 
 
 1.810 
 
 32 
 
 19.8 
 
 1.160 
 
 76 
 
 39.8 
 
 1.380 
 
 120 
 
 54.1 
 
 1.600 
 
 163 
 
 64.8 
 
 1.815 
 
 33 
 
 20.3 
 
 1.165 
 
 77 
 
 40.1 
 
 1.385 
 
 121 
 
 54.4 
 
 1.605 
 
 164 
 
 65.0 
 
 1.820 
 
 34 
 
 20.9 
 
 1.170 
 
 78 
 
 40.5 
 
 1.390 
 
 122 
 
 54.7 
 
 1.610 
 
 165 
 
 65.2 
 
 1.825 
 
 35 
 
 21.4 
 
 1.175 
 
 79 
 
 40.8 
 
 1.395 
 
 123 
 
 55.0 
 
 1.615 
 
 166 
 
 65.5 
 
 1 830 
 
 36 
 
 22.0 
 
 .180 ! 
 
 80 
 
 41.2 
 
 1.400 
 
 124 
 
 55.2 
 
 1620 
 
 167 
 
 65.7 
 
 1.835 
 
 37 
 
 .22.5 
 
 .185 
 
 81 
 
 41.6 
 
 1.405 
 
 125 
 
 55.5 
 
 1.625 
 
 168 
 
 65.9 
 
 1.840 
 
 38 
 
 23.0 
 
 .190 ; 
 
 82 
 
 420 
 
 1.410 
 
 126 
 
 55.8 
 
 1.630 
 
 169 
 
 66 1 
 
 1.845 
 
 39 
 
 23.5 
 
 .195 
 
 83 
 
 42.3 
 
 1.415 
 
 127 
 
 56.0 
 
 1.635 
 
 170 
 
 663 
 
 1.850 
 
 40 
 
 24.0 
 
 .200 
 
 84 
 
 42.7 
 
 1.420 
 
 128 
 
 56.3 
 
 1.640 
 
 171 
 
 66.5 
 
 1.855 
 
 41 
 
 24.5 
 
 .205 
 
 85 
 
 43.1 
 
 1.425 
 
 129 
 
 56.6 
 
 1.645 
 
 172 
 
 66.7 
 
 1.860 
 
 42 
 
 25.0 
 
 210 
 
 86 
 
 43.4 
 
 1.430 
 
 130 
 
 56.9 
 
 1.650 
 
 173 
 
 67.0 
 
 1.865 
 
 43 
 
 25.5 
 
 .215! 
 
 87 
 
 43.8 
 
 1.435 
 
 
 
 
 
 
 
 The scale of the Banme hydrometer, for liquids lighter than water, is so adjusted 
 that the point to which the spindle \vould sink, in a solution of one part of sodium 
 chloride in nine parts of water, is marked 0, while the corresponding point, when 
 pure water is used, is marked 10. The instrument used for liquids heavier than 
 water has the at the point to which it sinks in pure water, and the 10 at the point 
 reached in a 10 per cent solution of sodium chlorine at 17.5. 
 
 The Twaddell hydrometer contains two hundred divisions for a difference in 
 specific gravity of one. The of Twaddell is the specific gravity of pure water and 
 = 1.000. The instrument is usually divided into five sections for convenience. The 
 200 mark would correspond to a specific gravity of 2.000. 
 
INDEX. 
 
 Acetone, 360 
 
 Acetylene, 368, 371, 374 
 
 Acid, acetic, 360 ; antimonic, 343 ; 
 arsenic, 336; arsenic, experiments 
 with, 802; arsenious, 337; boric, 
 354; bromic, 166; carbonic, 386 ; 
 chlorauric, 614; cbloric, 115 ; 
 chloric, experiments with, 772 ; 
 chlorous, 118; chlorplatinic, 730 ; 
 cblorsulphuric, 241 ; chromic, 
 652, 663 ; chromic, experiments 
 with, 820; cyanauric, 615 ; cyan- 
 aurous, 615 ; cyanic, 404; cyan- 
 platinous, 730; dichromic, 658 ; 
 disilicic, 422; disulphuric, 219; 
 dithionic, 208, 225; dititanic, 424; 
 diuranic, 677; ferric, 710; ferrihy- 
 drocyanic, 704 ; ferrohydrocy- 
 anic, 703 ; fluoboric, 353; fluo- 
 silicic, 415;fluosilicic, constitution 
 of, 416 ; fluosilicic, preparation 
 of, 811; fluotantalic, 350; fluo- 
 titanic, 417; fluozirconic, 417 ; 
 formic, 380, 391 ; fuming, sulphu- 
 ric, 219; hydrobromic, 163; hydro- 
 bromic, experiments with, 776 ; 
 hydrochloric, 43, 105; hydrochlo- 
 ric, experiments with, 771 ; hy- 
 drochloric, formation of, 770 ; 
 hydrochloric, preparation of, 771; 
 hydriodic, 169 ; hydriodic, reduc- 
 ing action of, 779; hydrocyanic, 
 402; hydrofluoric, 178; hydrofluo- 
 ric, constitution of, 179 ; hydro- 
 fluoric, experiments with, 779 ; 
 hydrosulphurous, 222; hypobro- 
 mous, 166; hypochlorous, 116 ; 
 hyponitrous, 282 ; hypophos- 
 phoric, 333; hypophosphorous, 
 333 ; hyposulphurous, 208, 221, 
 222; iodic,171; iodic, constitution 
 of, 176; iodic, experiments with, 
 779 ; manganic, 182, 685; meta- 
 boric, 355; metachromous, 659 ;^ 
 metantimonic, 296, 343 ; meta-* 
 phosphoric, 296, 331; metarsenic, 
 296, 336 ; metastannic, 643 ; 
 rnetavanadic, 350; molybdic, 671 ; 
 muriatic, 43; nitric, 43, 277 ; 
 
 nitric, experiments with, 792 ; 
 nitric, normal, 263; nitric, prepa- 
 ration of, 791; nitric, red fuming, 
 281; nitric, reduction of, 794 ; 
 nitrohydrochloric, 281 ; nitrosyl- 
 sulphuric, 212, 293 ; nitrous, 281; 
 nitrous, experiments with, 794 ; 
 nitrous, normal, 263; Nordhausen 
 sulphuric, 219; normal sulphuric, 
 210, 218; orthoantimonic, 343 ; 
 orthophosphoric,296,326 ; orthova- 
 nadic, 350; osmic, 724 ; pentathi- 
 onic, 208, 226 ; perbromic, 166 ; 
 perchloric, 118 ; perchloric, prep- 
 aration of, 773 ; periodic, 173 ; 
 periodic, constitution of, 174 ; 
 periodic, normal, 174 ; perman- 
 ganic, 182, 687; phosphoric, 326; 
 phosphoric, experiments with, 
 802; phosphoric, glacial, 331; phos- 
 phoric, "insoluble," 541; phos- 
 phoric, normal, 296 ; phosphoric, 
 "reverted," 541; phosphoric "sol- 
 uble, " 541 ; phospho-molybdic,672; 
 phosphorus, 332; platinic, 729 ; 
 prussic, 402 ; pyroantimonic, 297, 
 343; pyroarsenic, 297, 336 ; pyro- 
 carbonic, 388;pyroligneous, 360 ; 
 pyrophosphoric, 297, 330 ; pyro- 
 sulpharsenic, 342 ; pyrosulpharse- 
 nious, 342 ; pyrosulphuric, 219 ; 
 pyrotantalic, 351 ; pyrovanadic, 
 350; ruthenious, 723;selenic, 242; 
 selenious, 241; silicic, 420; silicic, 
 experiments with, 812 ; silicic, 
 insoluble, 422; silicic, normal, 420; 
 stannic, 642 ; sulphantimonious, 
 346; sulpharsenious, 340; sulpho- 
 carbonic, 406; sulphocyamc, 407; 
 sulphotelluric, 245 ; sulphoxyar- 
 senic, 342; sulphuric, 43, 209; sul- 
 phuric, experiments with, 782; 
 sulphuric, illustration of prepa- 
 tion, 782; sulphuric, solid, 219, 
 236; sulphurous, 220; sulphur- 
 ous, experiments with, 783; sul- 
 phydric, 200; telluric, 244; tellu- 
 rious, 244; tetrabortc, 354; tetra- 
 hydroxyl-sulphuric, 218; tetrathi- 
 843 
 
844 
 
 INDEX. 
 
 onic, 208, 226; thiocarbonic, 406; 
 thiosulphuric, 208, 223; triazoic, 
 277 ; trithionic, 208, 225; tungstic, 
 674; vanadic, 350; zirconic, 424 
 
 Acids, 61, 127; avidity of, 440; basi- 
 city of, 136; constitution of, 134; 
 dibasic, 138 ; monobasic, 138; no- 
 menclature of, 143; organic, 380: 
 pentabasic, 138 ; polymolybdic, 
 672; polysilicic, 422; polystannic, 
 643; polytungstic, 674; silico- 
 tungstic, 674; tetrabasic, 138; tri- 
 bisic, 138; trisilicic, 423 
 
 Acids of antimony and arsenic, con- 
 stitution of, 346 
 
 Acids of phosphorus, constitution 
 of, 335 
 
 Acids of sulphur, constitution of, 
 226 
 
 Acidum, phosphoricum, glaciale, 
 331 
 
 Affinity, chemical, 11, 435 
 
 Affinity, specific coefficient of, 442 
 
 Affinity, thermochemical study of, 
 438 
 
 Agate, 418 
 
 Air, 248, 251 
 
 Air, analysis of, 253, 785 
 
 Alabandite, 684 
 
 Alabaster, 539 
 
 Albite, 575 
 
 Alchemy, 4 
 
 Alcohol, 379 
 
 Aldehyde, formic, 379 
 
 Algaroth, powder of, 318, 347 
 
 Alkalies, 127 
 
 Alkali metals, 482 
 
 Alkaline earths, 527 
 
 Allanite, 655 
 
 Allotropy, 89 
 
 Alloys, approximate composition 
 of, 839 
 
 Allylene, 371 
 
 Alum, basic, 578; burnt, 578; insolu- 
 ble, 579; shale, 578 
 
 Alumina, aluminium oxide, 575 
 
 Alurninates, 571 
 
 Aluminium, 563; bronze, 566, 587; 
 chloride, 566; preparation of, 818; 
 hydroxide, 569; oxide, 575 ; sili- 
 cates, 579; sulphate, 576 ; sul- 
 phates, basic, 576 
 
 Alums, 576 
 
 Alunite, 578 
 
 Amalgamation process for silver, 
 602 
 
 Amalgams, 625 
 Amethyst, 418 
 
 Ammonia, 260, 266; composition of, 
 
 270; determination of composi- 
 tion of, 790; experiments with, 
 789; preparation of, 788 
 
 Ammoniacal liquor, 266, 365 
 
 Ammonia-process for soda, 512 
 
 Ammonium alum, 579; amalgam, 
 273, 625; carbamate, 522; carbon- 
 ate, 522 ; chloride, 266, 519; com- 
 pounds, metallic derivatives of, 
 274; compounds, structure of, 274; 
 hydrosulphide, 521; hydroxide, 
 270; nitrate, 522; nitrite, 267; mag- 
 nesium phosphate, 329, 560; 
 phospho-molybdate, 672 ; salts, 
 269, 517; salts, formation of, 789; 
 sulphide, 520; sulphide, prepara- 
 tion of, 817; sulphocyanate, 519; 
 uranate, 677 
 
 Ammonium sulphide group, 198 
 
 Ampere's law, 73 
 
 Anatase, 412, 423 
 
 Anhydrides, 172, 209 
 
 Anhydrite, 539 
 
 Annealing, 544 
 
 Antimony, 308; and its compounds, 
 experiments with, 799; blende, 
 345; oxy chlorides, 319, 347; oxy- 
 chloride, experiment with, 802; 
 pentachloride, 318 ; pentasul- 
 phide,346; pentoxide, 345; salts of, 
 343; sulphides, experiments with, 
 802 ; tetroxide, 345; trichloride, 
 318; trioxide, 343; trisulphide, 
 345 
 
 Antimony, constitution of acids of, 
 346 
 
 Antimonyl salts, 344 ; sulphate, 
 344 
 
 Apatite, 298, 529, 540 
 
 Aqua regia, 281 
 
 Aragonite, 536 
 
 Arbor Saturni, 647 
 
 Argentan, 716 
 
 Argentic nitrate, 608 
 
 Argentous chloride, 606 
 
 Argon, 259 
 
 Aristotle, four elements of, 8 
 
 Arsenic, 305 ; and its compounds, 
 experiments with,798; disulphide, 
 340; pentasulphide, 342 ; pentox- 
 ide, 340 ; sulphides, experiments 
 with, 802 ; trichloride, 317; triox- 
 ide, 338 ; trioxide, experiments 
 with, 802; trisulphide, 340 
 
 Arsenic, constitution of acids of, 
 346 
 
 Arsenides, 305 
 
 Arseniuretted hydrogen, 306 
 
 Arsenopyrite, 711 
 
INDEX. 
 
 845 
 
 Arsine, 306, 798 
 Arsonium compounds, 308 
 Atomic theory, 68; weights, 71, 77, 
 
 79, 824; weights and specific heat, 
 
 449 
 
 Atoms, 68, 74 
 Aurates, 614 
 Auric chloride, 613; compounds, 
 
 610; hydroxide, 614; oxide, 614 
 Aurous chloride, 613; compounds, 
 
 610; oxide, 614 
 Avidity of acids, 440 
 Avogadro's law, 73 
 
 Balance, 5 
 
 Banca tin, 640 
 
 Barff's process, 699 
 
 Barite, 547 
 
 Barium, 547; carbonate, 549; chlo- 
 ride, 547; chromate, 667; dioxide, 
 547; hydroxide, 547; nitrate, 549; 
 oxide, 547; peroxide, 547; phos- 
 phates, 550; platinate, 724; sul- 
 phate, 549; sulphide, 548 
 
 Bases, 61, 127, 131; acidity of, 138; 
 constitution of, 134; diacid, 138; 
 monacid, 138; nomenclature of, 
 144; triacid, 138 
 
 Basic-lining process, 698 
 
 Bauxite, 564, 571 
 
 Bell-metal, 591 
 
 Benzene, 372 
 
 Berthollet, 15 
 
 Beryll, 551 
 
 Beryllium, see Glucinum. 
 
 Bessemer process, 698 
 
 Bismuth, 311; basic nitrates of , 348, 
 803; dichloride, 319; dioxide, 348; 
 experiment with, 799; hydroxide, 
 347; nitrate, 311; oxychloride, 
 319, 349; pentoxide, 349; sub- 
 nitrate, 348; sulphate, 311; tri- 
 chloride, 319; trioxide, 347; tri- 
 sulphide.. 349 
 
 Bisrnuthyl salts, 348 
 
 Black ash, 511 
 
 Black-lead, 359 
 
 Bleaching by chlorine, 101, 103; by 
 sulphur dioxide, 235; Dowder, 
 116, 534 
 
 Block-tin, 640 
 
 Blow-pipe, 398; experiments with, 
 808 
 
 Blue vitriol, 595 
 
 Bog iron-ore, 706 
 
 Boiling-points, table of, 826 
 
 Bone-black, 362; filters, 363; ex- 
 periments with, 803 
 
 Boracite, 354 
 
 Borax, 352, 515 
 
 Boron, 351; experiments with, 803 
 nitride, 356; phosphate, 356; salts 
 of, 356; trichloride, 352; trifluo- 
 ride, 356 
 
 Boryl potassium tartrate, 356 
 
 Brass, 591, 617 
 
 Braunite, 679, 682 
 
 Bricks, 583 
 
 Brimstone, crude, 188; roll, 188 
 
 Britannia metal, 309, 640 
 
 Bromides, experiments with, 812 
 
 Bromine, 161; Chloride, 166; hy- 
 drate, 162; preparation of, 776; 
 water, 162 
 
 Bromoform, 375 
 
 Bronze, 591 
 
 Brookite, 412, 423 
 
 Brown iron ore, 694, 706 
 
 Bunsen burner, 400 
 
 Burning, 32 
 
 Butane, 369 
 
 Butter of antimony, 318 
 
 Butylene, 371 
 
 Biityrum antimonii, 318 
 
 Cadmium, 622 ; carbonate, 616 ; 
 chloride, 622; cyanide, 623; sul- 
 phate, 622; sulphide, 622 
 
 Caesium, 502 
 
 Calamine, 617 
 
 Calcium, 529; carbide, 545; carbon- 
 ate, 536; chloride, 530; chloride, 
 preparation of, 817; fluoride, 531; 
 hydroxide, 532 ; nitride, 545 ; 
 oxide, 531; phosphates, 540; sili- 
 cate, 542; sulphate, 538; sulphide, 
 545 
 
 Calc-spar, 536 
 
 Calomel, 626 
 
 Calorie, 37 
 
 Calorimeter, 37 
 
 Candle, standard, 395 
 
 Carbides, 368 
 
 Carbon, 357 ; amorphous, 360 ; 
 chemical conduct of, 366; diox- 
 ide, 380 ; dioxide, experiments 
 with, 805 ; dioxide, relation to 
 life, 385; disulphide, 404; experi- 
 ments on reducing power of, 804; 
 monoxide, 389; experiments with, 
 806; tetrachloride, 375 
 
 Carbonates, 386, 478 ; acid, 387 ; 
 basic, 387; experiments with, 815 
 
 Carbonyl chloride, 392 
 
 Carbon silicide, 424 
 
 Carborundum, 424 
 
 Carnallite, 483, 485 
 
 Carnelian, 418 
 
846 
 
 INDEX. 
 
 Cassiterite, 639 
 
 Cast-iron. 696; gray, 696; white, 696 
 
 Cast-steel, 698 
 
 Caustic soda, 506 
 
 Cavendish, 40 
 
 Celestite, 546 
 
 Cements, 545; hydraulic, 545 
 
 Cerite, 562, 655 
 
 Cerium, 413, 655 
 
 Cerussite, 646 
 
 Chalcedony, 418 
 
 Chalcocite, 589 
 
 Chalcopyrite, 708 
 
 Chalk, 529, 536 
 
 Chance process for recovery of sul- 
 phur, 512 
 
 Charcoal, 360; absorption of gases 
 by, 803; animal, 362; filters, 363; 
 kiln, 361 
 
 Chemical action, 10, 426; examples 
 of, 737 
 
 Chemical change, 2 ; caused by 
 electric current, 735; by heat, 734 
 
 Chemical reactions, causes of, 434; 
 kinds of, 431; ideal, 435 
 
 Chemistry, 3; organic, 358 
 
 Chili saltpetre, 507 
 
 Chlorates, 476 
 
 Chloraurates, 614 
 
 Chlorides, 99, 460 ; double, 465 ; 
 experiments with, 812 ; general 
 properties of, 463; of acids, 209 
 
 Chlorine, 96; dioxide, 121; experi- 
 ments with, 769; preparation of, 
 768; hydrate, 104; hydrate, pre- 
 paration of, 770 ; liquid, 104 ; 
 monoxide, 120; trioxide, 121 
 
 Chlormercurates, 622 
 
 Chloro-acids, 465 
 
 Chloroaluminates, 568 
 
 Chlorostannates, 642 
 
 Chloroform, 375 
 
 Chlorophyll, 695 
 
 Choke-damp, 384 
 
 Chlorplatinates, 730 
 
 Chromates, 663; experiments with, 
 820 
 
 Chrome-alums, 662 
 
 Chrome-red, 668 
 
 Chrome-yellow, 668 
 
 Chromic chloride, 660; compounds, 
 658; hydroxide, 661; iron, 659; 
 oxide, 662; sulphate, 662 
 
 Chromite, 573, 657, 659 
 
 Cromium, 657; oxychloride, 668; 
 trioxide, 666 
 
 Chromous chloride, 660; com- 
 pounds, 658; hydroxide, 661 
 
 Chromyl chloride, 668 
 
 Chrysoberyl, 573 
 
 Cinnabar, 624, 631 
 
 Clay, 581 
 
 Cleveite, 585 
 
 Coal, 364; anthracite, 364; bitumi- 
 nous, 364; gas, 394; experiments 
 with, 807; oil, 358; tar, 365 
 
 Cobalt, 713; cyanides, 714; sulphide, 
 714 
 
 Cobaltic hydroxide, 714; oxide, 714 
 
 Cobaltite, 713 
 
 Cobaltous chloride/ 713; cobaltic 
 oxide, 714; hydroxide, 714; oxide 
 714 
 
 Coke, 362, 394 
 
 Columbite, 350 
 
 Columbium, 351 
 
 Combination, 11, 24; direct, 431 
 
 Combining weights, 17, 71 
 
 Combustion, 34; experiments on, 
 749; Lavoisier's explanation of, 34 
 
 Compound, chemical, 10, 12, 736 
 
 Condy's liquid, 688 
 
 Constitution, 80 
 
 Copper, 587; basic carbonate, 387 ; 
 hydride, 592; metallurgy of, 589 ; 
 native, 589; peroxide, 595; plat- 
 ing, 600; ruby, 589; suboxide, 595 
 
 Copperas. 709 
 
 Corrosive sublimate, 627 
 
 Corundum, 575 
 
 Cotunnite, 648 
 
 Crocoisite, 646, 659 
 
 Cryolite, 177, 466, 503; constitution 
 of, 180 
 
 Cupric arsenite, 598; carbonates, 
 598; chloride, 592; cyanide, 599; 
 hydride, 592; hydroxide, 594; ni- 
 trate, 598; oxide, 594; sulphate, 
 595; sulphide, 600; sulphocyanate, 
 599 
 
 Cuprite, 594 
 
 Cuprous chloride, 592; chloride, 
 preparation of, 818; cyanide, 599; 
 hydride, 592; hydroxide, 593; 
 iodide, 593; oxide 594; sulphide, 
 599; sulphocyanate, 599 
 
 Cyanides, 401, 407 
 
 Cyanogen, 401; experiments with, 
 810 
 
 Cyanplatinites, 730 
 
 Dal ton, 15, 68 
 Datholite, 354 
 Davy, 96 
 Deacon's process, 
 Decay, 381 
 
 Decomposition, 11, 24; direct, 432; 
 double, 25 
 
INDEX. 
 
 847 
 
 Decrepitation, 505 
 
 De la Bastie glass, 544 
 
 Deliquescence, 58 
 
 Deliquescent salts, 764 
 
 Developers in photography, 607 
 
 Dialyser, 421 
 
 Dialysis, 421 
 
 Diamond, 359 
 
 Diaspore, 570 
 
 Di-chlor-ethane, 373; methane, 375 
 
 Dichro-cobaltic chloride, 715 
 
 Dichromates, experiments with, 820 
 
 Didirniuni, 351, 656 
 
 Diffusion, 45; experiments on, 757 
 
 Dimorphism, 190 
 
 Di-sodiuin pyroantimonate, 515 
 
 Dissociation, 61, 314, 443; of dis- 
 solved substance, 453 
 
 Distillation, 67; destructive, 360; 
 dry, 360; experiment on, 764 
 
 Dolomite, 380, 555 
 
 Double chloride, 465; union, 372 
 
 Double salts, 319 
 
 Drummond light, 55 
 
 Dumas, his study of combustion of 
 hydrogen, 49; method for deter 
 minating the specific gravity of 
 vapors, 765 
 
 Earthenware, 582 
 
 Efflorescence, 58, 508 
 
 Efflorescent salts, 764 
 
 Eka-aluminium, 635 
 
 Eka-boron, 584 
 
 Eka-silicon, 639 
 
 Electrolysis, 443, 445 
 
 Electrolytic dissociation, 446 
 
 Electrolytic process for chlorine, 
 99 
 
 Electro-negative ions, 446 
 
 Electro-positive ions, 446 
 
 Electrotypes, 600 
 
 Elements, 8, 19, 21; acid-forming. 
 112; base-forming, 111, 455; melt- 
 ing-points and boiling-points, 825 
 of Aristotle, 8 ; replacing power, 
 83; symbols of, 19, 21 
 
 Emerald, 552 
 
 Emery, 575 
 
 Energy, chemical, 38; conservation 
 of, 6; stored up in plants, 385 
 
 Enstatite, 422 
 
 Epsom salt, 559 
 
 Equations, chemical, 23 
 
 Erbium, 562 
 
 Etching on glass, 179 
 
 Ethane, 369, 371 
 
 EtMops martialis, 631 
 
 Ethylene, 371, 374 
 
 Euchlorine, 121 
 Eudiometer, 50 
 
 Eudiometric method, 50; experi- 
 ments, 763 
 Euxenite, 562, 585 
 
 Faraday, 96 
 
 Feldspar, 480, 564 
 
 Ferment, nitrifying, 277 
 
 Fermentation, alcoholic, 381 
 
 Ferric chloride, 695 ; chloride, pre- 
 paration of, 821; ferrocyanide, 
 703; hydroxide, 706; soluble, 
 707; oxide, 707; sulphate, 710; 
 sulphate, preparation of, 821; 
 sulphide, 708 
 
 Ferrous ammonium sulphate, pre- 
 paration of, 821; carbonate, 708; 
 chloride, 700; ferricyanide, 704; 
 hydroxide, 705; oxide, 706; phos- 
 phate, 710; sulphate, 708; sul- 
 phate, preparation of, 821; sul- 
 phide, 707; ferric oxide, 707 
 
 Fire-damp, 373 
 
 Flame, 395; oxidizing, 397; reac- 
 tions, 525; reducing, 397 
 
 Flames, 395; causes of luminosity 
 of, 399; structure of, 397 
 
 Flint, 418 
 
 Flores zmci, 620 
 
 Fluocolumbates 351 
 
 Fluorides, constitution of, 179 
 
 Fluorine, 177 
 
 Fluor-spar, 177, 529 
 
 Fluosilicates, 415 
 
 Fluostannates, 642 
 
 Fluotantalates, 350 
 
 Fluotitanates, 417 
 
 Flux, 531 
 
 Fool's gold, 711 
 
 Formulas, constitutional, 81; molec- 
 ular, 79 
 
 Franklinite, 617, 706 
 
 Freezing-mixtures, 840 
 
 Fulminating mercury, 632 
 
 Fumaroles, 354 
 
 Gadolinite, 562, 584, 655 
 
 Gadolinium, 585 
 
 Gahnite, 573, 617 
 
 Galenite, 646 
 
 Gallium, 635 
 
 Gallic sulphate, 635 
 
 Gallons chloride, 635 
 
 Garnet, 542 
 
 Gases, kindling temperature of, 
 
 807; measurement of volume of, 
 
 741; weights of, 839 
 Gay Lussac tower, 214 
 
848 
 
 INDEX. 
 
 Germanium, 425, 639 ; chloride, 
 
 425; oxide, 425 
 German silver, 591, 716 
 Gersdorffite, 716 
 Glass, 542 
 Glauber, 105 
 Glauber's salt, 507 
 Glover tower, 214 
 Glucinum, 551; carbonate, 554; 
 
 chloride, 552; hydroxide, 553 
 Gold, 610; alloys, 613; amalgam, 
 
 626; dichloride, 613; metallurgy 
 
 of, 611; sulphide, 615 
 Graham, 47 
 Graphite, 359 
 Greenockite, 622 
 Green Vitriol, 709 
 Guignet's green, 662 
 Gun-metal, 591 
 Gunpowder, 493 
 Gypsum, 529, 538 
 
 Halogens, 160 
 
 Hardness, permanent, 538; tem- 
 pory, 538 
 
 Hard water, 538 
 
 Hauerite, 681 
 
 Hausmannite, 679, 681 
 
 Heat of combustion, 36; of decom- 
 position, 38; of neutralization, 
 440; specific law of, 449 
 
 Heavy spar, 547 
 
 Helium. 585 
 
 Hematite, 694 
 
 Hepar sulfur is, 491 
 
 Heptane, 369 
 
 Hexane, 369 
 
 Holmium, 562 
 
 Homologous series, 370 
 
 Homology, 370 
 
 Hornblende, 555 
 
 Hydrargillite, 569 
 
 Hydrates, 61, 144, 146 
 
 Hydraulic cement, 545; mining, 612 
 
 Hydrazine, 275 
 
 Hydrocarbons, 368, 804 
 
 Hydrochloric acid hydrate, 109 
 
 Hydrogen, 40 ; amount evolved 
 when a known weight of a metal 
 is dissolved in an acid, 754; burn- 
 ing of, 49, 106, 759; chemical 
 properties of, 758; dioxide, 91; 
 dioxide, experiments with, 768; 
 peroxide, 91; persulphide, 201; 
 preparation, 751; selenide, 204; 
 sulpliide, 193; sulphide, experi- 
 ments with, 780; sulphide group, 
 198; sulphide in chemical analy- 
 sis, 196; telluride, 205 
 
 Hydrogenium, 47 
 Hydrogen- valence, 156 
 Hydrometer table, 841 
 Hydrosulphides, 200, 474 
 Hydroxides, 61, 469; experiments 
 
 with, 814; formation from oxides, 
 
 236 
 
 Hydroxol, 237 
 Hydroxylamine, 265, 276 
 Hypochlorites, 476 
 Hyposulphite of soda, 508 
 
 Iceland spar, 536 
 
 Ice-machine, 268 
 
 Illuminating gas, 394 
 
 Illumination, 394 
 
 Indium, 636 
 
 Infusorial earth, 420 
 
 Ink, sympathetic, 714 
 
 lodic anhydride, 172 
 
 Iodides, experiments with, 812 
 
 Iodine, 167; bromide, 177; chloride, 
 
 176; experiments with, 778; pen- 
 
 tafluoride, 180; pentoxide, 172; 
 
 trichloride, 177 
 lodoform, 375 
 Ions, 446 
 Iridium, 724; chlorides, 725; oxides, 
 
 725 
 Iron, 691; carbonyls, 711; disul- 
 
 phide, 711; galvanized, 618; 
 
 metallurgy, 695 
 Isomerism, ^30 
 
 Kainite, 499, 555 
 Koaline, 564, 580 
 Kelp, 167 
 Kieserite, 555 
 
 Kindling temperature, 34; of gases, 
 395 
 
 Lamp black, 362 
 
 Lanthanum, 655 
 
 Lapis-lazuli, 581 
 
 Laughing gas, 284 
 
 Lavoisier's work, 5 
 
 Law of definite proportions, 14; of 
 indestructibility of matter, 6; of 
 multiple proportions, 15 
 
 Lead, 425, 646; its com pounds r 
 experiments with, 820; carbonate, 
 652; chloride, 648; chromate, 668; 
 hydroxide, 649; iodide, 648; 
 molybdate, 672; metallurgy of, 
 646; nitrate, 652; oxide, 649; pen- 
 cils, 359; peroxide, 650; sesqui- 
 oxide, 650; suboxide, 649; sugar 
 of, 650; sulphate, 654; sulphide, 
 652; tree, 647 
 
INDEX. 
 
 849 
 
 Le Blanc process for soda, 510 
 
 Lepidolite, 516 
 
 Lignite, 364 
 
 Lime, 531; chloride of, 116; slaked, 
 531 
 
 Liine-light, 55 
 
 Lime-stone, 529, 536 
 
 Lime-water, 532 
 
 Linnaeite, 714 
 
 Liquor Ferri chlorati, 701; Ferri 
 sesquichlorati, 701; Stibii muria- 
 tici, 319 
 
 Litharge, 649 
 
 Lithium, 516; carbonate, 517; chlor- 
 ide, 517; phosphate, 517 
 
 Liver of sulphur, 491 
 
 Loadstone, 707 
 
 Lunar caustic, 609 
 
 Luteo-cobaltic chloride, 715 
 
 Magnesia, 558; alba, 560; usta, 558 
 
 Magnesio ferrite, 573 
 
 Magnesite, 555 
 
 Magnesium, 555;. borate, 561; car- 
 bonate, 559; chloride, 557; chlor- 
 ide, preparation of, 817; hydrox- 
 ide, 558; phosphates, 560; prepa- 
 ration of, 817; oxide, 558; silicate, 
 561; silicide, 561; sulphate, 559 
 
 Magnetite, 694, 707 
 
 Malachite, 589, 598 
 
 Malaria, 259 
 
 Manganates, 685 
 
 Manganese, 182, 678; black oxide of , 
 682; dioxide, 682; hept-oxide, 689; 
 salts, preparation of, 821; tetra- 
 chloride, 680; tetrafluoride, 680; 
 trichloride, 680 
 
 Manganic oxide, 682; sulphate, 685 
 
 Manganite, 679, 682 
 
 Manganites, 683 
 
 Manganous carbonate, 685; chlo- 
 rides, 679; cyanide, 685; hydrox- 
 ide, 681 ; maganic oxide, 681 ; ox- 
 ide, 681; sulphate, 685; sulphide, 
 684 
 
 Marble, 529, 536 
 
 Marcasite, 711 
 
 Marl. 537, 581 
 
 Marsh gas, 368, 369, 373 
 
 Marsh's test' for arsenic, 310, 797 
 
 Mass action, 436 
 
 Mass, influence of, 436 
 
 Matches, 302; safety, 302 
 
 Matter, constitution of, 68; early 
 views regarding composition of, 
 7; law of indestructibility of, 6 
 
 Meerscbaum, 555 
 
 Melting-point, table of, 826 
 
 Mendelejeff, 148 
 
 Mercuric chloramide, 633; chloride, 
 627; compounds, 623; cyanide, 
 631; diammonium chloride, 634; 
 iodide, 629; nitrate, 633; nitrate, 
 preparation of, 819; oxide, 630; 
 sulphide, 631 
 
 Mercurias solubilis HaJmemanm, 
 634 
 
 Mercurous ammonium chloride, 633; 
 chloramide, 633; chloride, 626; 
 compounds, 623; iodide, 628; 
 nitrate, 632; nitrate; preparation 
 of, 819; oxide, 630; sulphide, 630 
 
 Mercury, 623; metallurgy of, 624 
 
 Metallic properties, 456 
 
 Metallurgy, 458 
 
 Metals, 111, 134, 456; properties, 
 459 
 
 Metastannates, 643 
 
 Metathesis, 25, 433 
 
 Metatungstate, 674 
 
 Metavanadates, 350 
 
 Meteorites, 694 
 
 Methane, 369, 373 
 
 Methyl alcohol, 360, 379 
 
 Meyer, Lothar, 148 
 
 Meyer, Victor, method for deter- 
 mining the specific gravity of va- 
 pors, 766 
 
 Mica, 542; 564 
 
 Microcosmic salt, 332; 523 
 
 Minerals, 457 
 
 Miner's safety lamp, 396 
 
 Minium, 651 
 
 Mixtures, mechanical, 12, 736 
 
 Molecular weights, 75, 77 
 
 Molecules, 73, 74, 427 
 
 Molybdates, 671 
 
 Molybdenite, 670 
 
 Molybdenum, 670; chlorides, 670; 
 oxides, 671 
 
 Mono-chlor-methane, 375 
 
 Mortar, 544 
 
 Mosaic gold, 644 
 
 Nascent state, 90 
 Neodymium, 351, 656 
 Neutralization, 127; experiments 
 
 on, 774 
 
 Newlands, 148 
 Newton's metal, 312 
 Niccolite, 716 
 Nickel, 716; alloys ,716; carbonyl, 
 
 717; cyanide, 717; plating, 717 
 Nickelic hydroxide, 717 
 Nickelous chloride, 717, hydroxide, 
 
 717; oxide, 717 
 Niobium, 350 
 
850 
 
 INDEX. 
 
 Nitrates, 475; basic, 280; formation 
 of, 793; formation of, in the soil, 
 261 
 
 Nitric oxide, 285; experiments with, 
 795 
 
 Nitrification, 277, 492 
 
 Nitre-compounds, 280 
 
 Nitrogen, 248; boride, 356; experi- 
 ments with, 785; iodides, 292; 
 pentoxide, 289; peroxide, 288; 
 peroxide, experiments with, 795; 
 preparation of, 784; relations to 
 life, 257; structure of compounds 
 of, 289; sulphide, 293; trichlo- 
 ride, 292; triiodide, 292;trioxide, 
 287; trioxide, experiments with, 
 795 
 
 Nitroprussiates, 705 
 
 Nitrosyl chloride, 281 
 
 Nitrous oxide, 283; experiments 
 with, 794 
 
 Non-metals, 112, 456 
 
 Octane, 369 
 
 Olefiant gas, 374 
 
 Olivine, 561 
 
 Opal, 418 
 
 Ores, 458 
 
 Orpiment, 305, 340 
 
 Orthite, 562 
 
 Orthoclase, 423 
 
 Osmium, 723; chlorides, 724; oxides, 
 724 
 
 Oxidation, slow, 35 
 
 Oxides, 39, 467; acidic, 172, 209; 
 basic, 172 
 
 Oxygen, 28; amount liberated from 
 a known weight of potassium 
 chlorate, 745; and acid properties, 
 141; burning of, 807; chemical 
 properties of , 747; physical prop- 
 erties of, 747; preparation of, 740; 
 valence, 156 
 
 Oxyhydrogen light, 55; blow-pipe, 
 54; blow-pipe, experiments with, 
 763 
 
 Oxysulphides, 406 
 
 Ozone, 85; experiments with, 767; 
 in the air, 88 
 
 Palladic chloride, 726; oxide, 727 
 
 Palladious chloride, 726; oxide, 727 
 
 Palladium, 726; hydrogen, 726; 
 suboxide, 727 
 
 Paracyanogen, 632 
 
 Passive state of iron, 700 
 
 Pattison's method of separating sil- 
 ver from lead, 602 
 
 Peat, 364 
 
 Pentane, 369 
 
 Periodates, 174 
 
 Periodic law, 148, 158, 429 
 
 Permanent white, 549 
 
 Permanganates, 687 
 
 Petalite, 516 
 
 Petroleum, 358, 369 
 
 Philosopher's wool, 819 
 
 Phlogiston theory, 33 
 
 Phosgene, 392 
 
 Phosphates, 479; acid, 329; neutral, 
 329; normal, 329; primary, 329; 
 secondary, 329; tertiary; 329 
 
 Phosphine, 302; experiments with, 
 797; liquid, 304 
 
 Phosphonium bromide, 304; chlo- 
 ride, 304; iodide, 304 
 
 Phosphoric anhydride, 334 
 
 Phosphorite, 298, 529, 540 
 
 Phosphorus, 298; anhydride, 334; 
 crystallized, metallic, 301 ; exper- 
 iments with, 796; iodides, 317; 
 oxychlorides, 336; pentachloride, 
 3J4; pentachloride, action on hy- 
 droxides, 316; pentachloride, 
 preparation of, 801; pentafluor- 
 ide, 317; pentoxide, 334; red, 301; 
 suboxide, 335; tetroxide, 335; tri- 
 bromide, 317; trichloride, 812; 
 trichloride, preparation of, 799; 
 trioxide, 334 
 
 Phosphorus, constitution of acids 
 of, 335 
 
 Phosphuretted hydrogen, 302 
 
 Photography, 607 
 
 Photometerj 395 
 
 Physical change, 2 
 
 Physics, 3 
 
 Pig-iron, 696 
 
 Pinchbeck, 591 
 
 Pink salt, 642 
 
 Pithcblende, 675 
 
 Placer mining, 611 
 
 Plaster of Paris, 539 
 
 Platinic chloride, 729; chloride, 
 preparation of, 821; hydrox : de, 
 730; oxide, 730; sulphide, 730 
 
 Platinous chloride, 729; hydroxide, 
 730; oxide, 730; sulphide, 730 
 
 Platinum, 727; alloys, 728; bases, 
 731; black, 727; cyanides, 729; 
 metals, 722; spongy, 727 
 
 Plumbago, 359 
 
 Plurnbates, 651 
 
 Plumbites, 650 
 
 Polymerism, 515 
 
 Polysulphides, 207 
 
 Porcelain, 582 
 
 Potash, 483, 500 
 
INDEX. 
 
 851 
 
 Potassium, 483; alum, 578; bro- 
 mide, 485; carbonate, acid, 500; 
 carbonate, extraition from, wood 
 ashes, 816; chlorate, 114, 494; 
 chlorate, preparation of, 772; 
 chloride, 485; chloriridate, 725; 
 chlorthorate, 417; chrornate, 664; 
 cyan ate, 498; cyanate, preparation 
 of, 810; cyanide, 496; cyanide, 
 preparation of, 810; dichromate, 
 664; diperiodate, 496; disulphite, 
 500; ferricyanide, 704; ferrocya- 
 nide, 401, 702; fluogermanate,639; 
 fluoride, 485; fluosilicate,115; fluo- 
 thorate, 417; fluotitanate, 412; 
 hydride, 485; hydrosulphide, 490; 
 hydroxide, 488; hydroxide, prep- 
 aration of, 818; hypochlorite, 114; 
 iodide, 485; iodide, preparation of, 
 816; manganate, 686; manganate, 
 preparation of, 821; rnesoperio- 
 date, 496; nitrate, 492; nitrite, 
 494; osinite, 724; oxide, 489; 
 perchlorate, 118, 495; perchlorate, 
 preparation of, 773; periodate, 
 496; permanganate, 183, 687; per- 
 inangate, preparation of, 821; per- 
 oxide, 489 ; perruthenite, 723; 
 phosphates, 500 ; polysulphides, 
 491; ruthenite, 723; silicate, 501; 
 stannate, 643; sulphantimonate, 
 492; sulpharsenate, 492; sulph- 
 arsenite, 492; sulphate, 498; sul- 
 phate, acid, 499; sulphide, 490; 
 sulphite, 500; sulphite, acid, 500; 
 sulphocyanate, 498; tetrachro- 
 mate, 668; tichromate, 668 
 
 Praseo-cobaltic chloride, 715 
 
 Praseodymium, 351, 656 
 
 Preparing salt, 643 
 
 Priestley, 28 
 
 Printing-ink, 362 
 
 Propane, 369 
 
 Propylene, 371 
 
 Proust, 15 
 
 Prussian blue, 704 
 
 Puddling, 697 
 
 Purple of Cassius, 614 
 
 Purpureo-cobaltic chloride, 715 
 
 Pyrargyrite, 603 
 
 Pyrolusite, 679 
 
 Pyrophospbates, 331 
 
 Pyrosiderite, 706 
 
 Pyrrhotite, 694 
 
 Quartation. 613 
 Quartz, 410, 418 
 Quick-lime, 531 
 
 Raoult's method for determining 
 molecular weights, 521 
 
 Reaction, chemical, 24; endother- 
 niic, 94; exothermic, 94 
 
 Realgar, 305, 340 
 
 Red lead, 651 
 
 Reducing agent, 47 
 
 Reduction, 47; by carbon, 368; by 
 hydrogen, 759 
 
 Respiration, 383 
 
 Reversion of phosphates, 541 
 
 Rhodium, 724 
 
 Rhodocroisite, 679 
 
 Rinmann's green, 622 
 
 Roasting, 474 
 
 Rock-crystal, 418 
 
 Roseo-cobaltic chloride, 715 
 
 Rose's metal, 312 
 
 Rouge, 717 
 
 Rubidium, 502 
 
 Ruby, 575 
 
 Ruby copper, 594 
 
 Rupert's drops, 544 
 
 Ruthenium, 722; chloride, 722; ox- 
 ide, 723 
 
 Rutile, 412, 423 
 
 Safety lamp, 374 
 
 Sal ammoniac, 266, 519 
 
 Salt, common, 504 
 
 Saltpeter, 492; plantations, 492 
 
 Salts, 131, 139; acid, 139; basic, 138; 
 constitution of, 136; decomposi- 
 tion of by bases, 469; formation 
 of, 461; neutral, 137; nomen- 
 clature of, 144; normal, 137, 
 139 
 
 Samarium, 585 
 
 Samarskite, 585 
 
 Sand, 418 
 
 Sapphire, 575 
 
 Scandium, 584 
 
 Scheele, 28 
 
 Scheele's green, 598 
 
 Scheelite, 673 
 
 Schlippe's salt, 346, 506 
 
 Schweinfurt green, 598 
 
 Selenium, 203; acid chlorides of, 
 243; dioxide, 243 
 
 Serpentine, 423, 555 
 
 Siderite, 694, 702 
 
 Siemens-Martin furnace, 698 
 
 Silicates, 480 
 
 Silicides, 424 
 
 Silicon, 410; dioxide, 418; hexa- 
 chloride, 414; hydride, 412; mag- 
 nesium, 557; preparation of, 810; 
 tetrachloride, 413; tetrafluoride, 
 
652 
 
 INDEX. 
 
 414; tetrafluoride, preparation of, 
 811 
 
 Silver, 601; allotropic forms of, 605; 
 alloys of, 605 ; amalgam, 606 ; 
 borates, 609; bromide, 606; chlo- 
 ride, 606; chromate, 668; cyanide, 
 609; iodide, 606; metallurgy of, 
 602; nitrate, 608; octoborate, 
 609: oxide, 608; peroxide, 608; 
 suboxide, 608; sulphide, 608; sul- 
 phocyanate, 609; tree, prepara- 
 tion of, 819; triazoate, 608 
 
 Slaking, 532 
 
 Smalt, 715 
 
 Smaltite, 713 
 
 Smithsonite, 617 
 
 Soapstone, 555 
 
 Soda, calcined purified, 511; crude, 
 511; crystallized, 511; from cryo- 
 lite, 513 
 
 Soda-water, 383 
 
 Sodium, 502; alum, 579; amalgam, 
 625; ammonium phosphate, 523; 
 borate, 515 ; bromide, 506 ; car- 
 bonate, 509; carbonate, acid or 
 primary, 514; carbonate, prepar- 
 ation of, by ammonia process, 
 816 ; chloride, 504 ; chromate, 
 667; dichromate, 667; fluoride, 
 506 ; hydride, 504 ; hydroxide, 
 506; iodide, 506; nitroprussiate, 
 705 ; metaphosphate, 329, 515 ; 
 monoxide, 506; nitrate, 506; pan- 
 tungstate, 674 ; peroxide, 506 ; 
 phosphates, 514; potassium car- 
 bonate, 514; pyrophosphate, 330; 
 silicate, 516; stannate, 643; sul- 
 phantimonate, 506 ; sulphate, 
 507 ; sulphate, experiment on 
 supersaturated solution of, 816; 
 thiosulphate, 224, 508; uranate, 
 677 
 
 Solder, soft, 640 
 
 Solubilities, table of, 826 
 
 Solution, 62; as an aid to chemical 
 action, 64 
 
 Solvay process for soda, 512 
 
 Spathic iron, 694, 708 
 
 Specific heat and atomic weights, 
 449 
 
 Specific gravity of vapors, deter- 
 mination of, 765 
 
 Spectroscope, 525 
 
 Spectrum, banded, 526; continuous, 
 526 
 
 Sphalerite, 617 
 
 Spiegel iron, 697 
 
 Spinels, 572 
 
 Spiritus fumans Libavii, 642 
 
 Spitting of silver, 604 
 
 Spodumene, 516 
 
 Stahl, 33 
 
 Stalactites, 537 
 
 Stalagmites, 537 
 
 Stannic chloride, 642 ; chloride, 
 preparation of, 820; hydroxide, 
 642; oxide, 644; salts, 645; sul- 
 phate, 645; sulphide, 644 
 
 Stannite, 634 
 
 Stannous chloride, 641; chloride, 
 preparation of, 819; hydroxide, 
 642; oxide, 643; salts, 645; sul- 
 phate, 645; sulphide, 644 
 
 Stas, 15 
 
 Steel, 698 
 
 Stibine, 309, 799 
 
 Stibnite, 308, 345 
 
 Stolzite, 673 
 
 Strass, 543 
 
 Stromeyerite, 608 
 
 Stontianite, 546 
 
 Strontium, 546 ; chloride, 546 ; 
 hydroxide, 546 ; nitrate, 546 ; 
 oxide, 546; sulphate, 546 
 
 Sublimation, 519 
 
 Substitution, 41, 102 
 
 Superphosphate of lime, 541 
 
 Sulphantimonates, 346 
 
 Sulphates, 476; experiments with, 
 814 
 
 Sulphides, 191, 472 
 
 Sulphites, 2^2, 478; acid, 222; nor- 
 mal, 222 
 
 Sulpho-salts, 475 
 
 Sulphostannates, 644 
 
 Sulphur, 187 ; acid chlorides of, 
 239 ; auraturu, 346 ; dichloride, 
 202 ; dioxide, 233 ; dioxide, ex- 
 periments with, 783; experiments 
 with, 779; flowers of, 188; hept- 
 oxide, 233; hexiodide, 203; in- 
 soluble, 191; monochloride, 202; 
 sesquioxide, 233; soluble, 191; 
 stick, 188 ; tetrachloride, 202 ; 
 trioxide, 236; trioxide, experi- 
 ments with, 784; waters, 187 
 
 Sulphur auratum, 346 
 
 Sulphur,constitution of acids of ,226 
 
 Sulphuretted hydrogen, 193 
 
 Sulphuryl chloride, 240 
 
 Sulphuryl-hydroxyl chloride, 241 
 
 Sulphydrates, 200 
 
 Sylvite, 483, 485 
 
 Symbols of compounds, 22; of ele- 
 ments, 19, 21 
 
 Tachjdrite, 530 
 Talc,' 555 
 
INDEX. 
 
 853 
 
 Tantalite, 350 
 
 Tantalum, 350 
 
 Tartar, crude, 483; emetic, 344 
 
 Tellurium, 204; dioxide, 245; tetra- 
 chloride, 205; trioxide, 245 
 
 Tempering of glass, 544; of steel, 
 698 
 
 Tetra-chlor-ethane, 373 
 
 Tetra-chlor-metbane, 375 
 
 Tballic chloride, 636; hydroxide, 
 637; sulphate, 637 
 
 Thallium, 636 
 
 Thallous chloride, 636; hydroxide, 
 636; phosphate, 637; sulphate, 
 637; sulphide, 637 
 
 Theory, use and value of a, 70 
 
 Thermocheinical measurements, 
 value of, 439 
 
 Thermochemistry, 27 
 
 Thionyl chloride, 239 
 
 Thomas- Gilchrist process, 698 
 
 Thorite, 413 
 
 Thorium, 413; dioxide, 424; tetra- 
 chloride, 417; tetrafluoride, 417 
 
 Thulium, 562 
 
 Tin, 425 ; amalgam, 640 ; metal- 
 lurgy of, 639; salt, 641; stone, 
 639, 644 
 
 Titanium, 412; dioxide, 423; sul- 
 phate, 424; tetrachloride, 416; 
 tetrafluoride, 417 
 
 Titanyl sulphate, 424 
 
 Toluene, 372 
 
 Tri-chlor-methane, 375 
 
 Tridymite, 418 
 
 Triple linkage, 373; union, 373 
 
 Tungstates, 673 
 
 Tungsten, 673; chlorides, 673; 
 oxides, 674 
 
 Turnbull's blue, 705 
 
 Tuyeres, 695 
 
 Type metal, 309 
 
 Ultramarine, 581 
 Uranates, 677 
 Uraninite, 675 
 
 Uranium, 675; chlorides, 676; ox- 
 ides, 676; yellow, 677 
 Uranous salts, 676 
 Uranyl salts, 247, 676 
 
 Valence, 81, 428; variations of, 155 
 Vanadium, 350; chlorides, 350; 
 oxides, 350 
 
 Vein-mining, 611 
 
 Ventilation, 258 
 
 Vitriol, oil of, 211 
 
 Vitriols, 596 
 
 Volumes, law of combination by, 45 
 
 Water, 57; as a solvent, 62; con- 
 tamination of, by sewage, 67 ; 
 determination of composition of, 
 760; in organic substances, 763; 
 of crystallization, 57, 763; proofs 
 of composition of, 58; purifica- 
 tion of, 67; synthesis of, 59; 
 transformation into earth, 5 
 
 Water-gas, 43, 389 
 
 Water-glass, 501, 516 
 
 Waters, chalybeate, 66; efferves- 
 cent, 66; natural, 65; sulphur, 66 
 
 Weights, atomic, 19, 21; combin- 
 ing, 17, 18 
 
 Weights, length and volume, com- 
 parison of, 840 
 
 Weldon's process, 99, 682 
 
 White lead, 653 
 
 White precipitate, 633; fusible, 
 634; infusible, 633 
 
 Witherite, 547 
 
 Wolframite, 673 
 
 Wollastonite, 542 
 
 Wood's metal, 312 
 
 Wood spirits, 360 
 
 Work, chemical, 38 
 
 Wrought iron, 697 
 
 Wulfenite, 646 
 
 Xylene, 372 
 
 Yellow prussiate of potash, 401 
 Ytterbium, 581 
 
 Yttrium, 584 ; chloride, 584 ; 
 hydroxide, 584; oxide, 584 
 
 Zeolites, 542 
 
 Zinc, 617; alloys, 618; blende, 614; 
 carbonate, 621; chloride, 618; 
 dust, 617; experiment on burning 
 of, 819; hydroxide, 619; metal- 
 lurgy of; 617; oxide, 620; solu- 
 tion of, in sodium hydroxide, 819; 
 sulphate, 621 ; sulphide, 620 ; 
 white, 620 
 
 Zircon, 413, 424 
 
 Zirconium, 413; dioxide, 424; oxy- 
 fluoride, 417; tetrachloride, 417; 
 tetrafluoride, 417 
 
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 HENRY HOLT & CO., 29 WEST 230 ST., NEW YORK. 
 in, 1900 
 

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