GIFT OF ofessor F.T. Biolet.ti AMERICAN SCIENCE SERIES-ADVANCED COURSE INORGANIC CHEMISTRY BY IEA EEMSEN Professor of Chemistry in the Johns Hopkins University FOURTH EDITION, REVISED JtfEW YORK HENRY HOLT AND COMPANY 1895 .< :, r .. COPYRIGHT, 1889, . . * * b^" *' "" ' t HENRY 'HOLT & cd/ '"'*** * ** : *'* ^*l*-^4-i .*: PREFACE. 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 tha 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 M3G649 iu iT PEEFACE. 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. PEEFACE 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, December 6, 1889. CONTENTS. CHAPTER I. CHEMICAL AND PHYSICAL CHANGE EARLIEST CHEMICAL KNOWLEDGE LAW OP THE INDESTRUCTIBILITY OF MATTER LAW OF DEFINITE PROPORTIONS LAW OF MULTIPLE PROPORTIONS. PAGE 3!atter 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 Earliest 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 OF THE ACTION OF HYDROGEN ON OXYGEN. PAGBT. Burning of Hydrogen Method of Dumas Eudiometric Method Calculation of the Results obtained in exploding Mixtures of Hydrogen and OxygenDetermination of the 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 Oxy hydrogen 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, 5T CHAPTER VI. CONSTITUTION OF MATTER ATOMIC THEORY ATOMS AND MOLECULES CONSTITUTION VALENCE. Early ViewsThe Atomic Theory as proposed by Dalton Use and Value of a Theory Atomic Weights and Combining Weights Molecules A vogadro'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 Therniochemical Considerations, 85> CONTENTS. ix CHAPTER VIII. CHLORINE HYDROCHLORIC ACID. PAGE Historical Occurrence of Chlorine Preparation Weldon's Pro- cessProperties Different Kinds of Action Chlorine Hydrate and Liquid Chlorine Hydrochloric Acid Historical Study of the Action of Hydrogen upon Chlorine Preparation Proper- tiesChemical Action of Hydrochloric Acid, 96 CHAPTER IX. COMPOUNDS OF CHLORINE WITH OXYGEN AND WITH HYDROGEN AND OXYGEN. General Principal Reactions for Making Compounds of Chlorine with Hydrogen and Oxygen Properties Hypochlorous Acid Chlorous Acid Perchloric Acid General 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. Acid 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 Compound of Iodine and Bromine Fluorine Occurrence Properties Hydrofluoric Acid Constitution of Hydrofluoric Acid and the Fluorides Compound of Fluorine and 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 -OF FAMILY VI, GROUP B : SULPHUR SELENIUM TELLURIUM . Introductory Sulphur Occurrence Properties Uses of Sulphur Compounds of Sulphur with Hydrogen Hydrogen, Sul- phuretted Hydrogen Properties Action of Hydrogen Sul- phide upon Solutions of Salts, Uses in Chemical Analysis Hydrosulphides Hydrogen Persulphide Compounds of Sul- phur with Members of the Chlorine Group Selenium Occur- renceProperties Hydrogen Selenide Tellurium Occur- rence Properties Hydrogen Telluride, 185 CHAPTER XIV. COMPOUNDS OF SULPHUR, SELENIUM, AND TELLURIUM WITH OXYGEN AND WITH OXYGEN AND HYDROGEN. Introductory Sulphuric Acid Tetrahydroxyl Sulphuric Acid Normal Sulphuric Acid Sulphurous Acid Hyposulphurous Acid Thiosulphuric Acid Other Acids of Sulphur Consti- tution of the Acids of Sulphur Compound of Sulphur with Oxygen Sulphur Dioxide Sulphur Trioxide Acid Chlorides of Sulphur Thionyl Chloride Sulphuryl Chloride Chlor- sulphuric 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 Selenium Tellurious Acid Telluric Acid Oxides of Tellurium Sulphotelluric Acid Family VI, Group A, 206 CONTENTS. xi CHAPTER XV. NITROGEN THE AIR. PAGE Nitrogen General Occurrence of Nitrogen Preparation Prop- erties The Air Analysis of Air, 348 CHAPTER XVI. COMPOUNDS OP 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 Relations 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 Nitric Acid Red Fuming Nitric Acid Nitrous Acid Hyponitrous Acid- Nitrous Oxide Nitric Oxide Nitrogen Trioxide Nitrogen Pentoxide Structure of the Compounds of Nitrogen with Oxygen and Hydrogen Compounds of Nitrogen with the Ele- ments of the Chlorine Group Compounds of Nitrogen with the Members of the Sulphur Group, 260 CHAPTER XVII. ELEMENTS OP FAMILY V, 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 Occurrence Antimony Occurrence Properties Applications of Antimony Stibine Methods of distinguishing between Arsenic and Antimony Bismuth Occurrence Compounds 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 Chlo- rineDouble Salts, 294 CHAPTER XVIII. COMPOUNDS OP THE ELEMENTS OP THE PHOSPHORUS GROUP WITH OXYGEN AND WITH OXYGEN AND HYDROGEN. Introduction Phosphoric Acid, Orthophosphoric Acid Proper- ties Pyrophosphoric Acid Metaphosphoric Acid Phosphor- XH CONTENTS. PAGE ous Acid Hypophosphoric Acid Hypophosphorous Acid Phosphorus Pentoxide, Phosphoric Anhydride Phosphorus Trioxide or Phosphorous Anhydride Constitution of the Acids of Phosphorus Phosphorus Oxy chloride Arsenic Acid Arsenious Acid Arsenic Trioxide Arsenic Pentoxide Sulphides Arsenic Disulphide Arsenic Trisulphide Arsenic Pentasulphide Autirnouic Acid Antimony Trioxide Salts of Antimony Antimony Tetroxide Antimony Pentoxide Antimony Trisulphide Antimony Pentasulphide Constitu- tion of the Acids of Arsenic and Antimony Oxychlorides of Antimony Oxides of Bismuth Salts of Bismuth Bismuth Dioxide Bismuth Peutoxide Bismuth Trisulphide Bismuth Oxychloride Family V, Group A Vanadium Vanadic Acid Tantalum Niobium Didymium Boron General Oc- currence Boron Trichloride Boron Trifluoride Boric Acid Salts of Boron Nitrogen Boride, 322 CHAPTER XIX. CARBON 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 OF 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 Family V Compounds of the Members of the Silicon Group with Oxygen, and with Oxygen and Hydrogen Silicon Dioxide Properties Uses Silicic Acid Polysilicic Acids Titanium Dioxide Zirconium Dioxide Thorium Dioxide 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 Thermocheinical Study of Affinity 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 Relations between Specific Heat and Atomic Weights Exceptions to the Law of Specific Heats Raoult's Method for the Determination of Molecular Weights, . - ; . . . . ... .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 CONTENTS. PAGB the Metals Compounds of the Metals Formation of Salts in General The so-called Double Chlorides and similar Com- pounds of Fluorine, Bromine, and Iodine Different Chlorides of the same Metal Oxides Different Oxides of the same Metal Hydroxides Decomposition of Salts by Bases Sul- phides Hydrosulphides Sulpho -salts Nitrates Chlorates Sulphates Carbonates Phosphates Silicates, ...... 45] CHAPTER XXV. ELEMENTS OF FAMILY I, GROUP A: THE ALKALI METALS : LITHIUM SODIUM POTASSIUM RUBIDIUM CAESIUM AMMONIUM. General Potassium Occurrence Preparation Properties Po- tassium Hydride Potassium Fluoride, Chloride, Bromide, Iodide Properties Applications Potassium Hydroxide Potassium Oxide Potassium Hydrosulphide Potassium Sul- phide Potassium Nitrate Applications Gunpowder Potas- sium Nitrite Potassium Chlorate Potassium Perchlorate Potassium Periodate Potassium Cyanide Potassium Cyauate Potassium Sulphocyanate Potassium Sulphate Primary, or Acid, Potassium Sulphate Sulphites Phosphates Potas- sium Silicate Rubidium Caesium Sodium Occurrence Preparation Properties Applications Sodium Hydride Sodium Chloride Sodium Hydroxide Oxides Sodium Sul- phantimoniate 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-auti- inonate 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 Carbonate Reactions of the Members of the Sodium Group which are of Value in Chemical Analysis Flame Reactions and the Spectroscope, ..... 478 CHAPTER XXVI. ELEMENTS OF FAMILY II, GROUP A: BERYLLIUM MAGNESIUM CALCIUM STRONTIUM BARIUM [ERBIUM]. General Calcium Sub-GroupCalcium Occurrence Properties Calcium Chloride Calcium Fluoride Calcium Oxide Cal- cium Hydroxide Bleaching-powder Calcium Carbonate CONTENTS. XV PAGE Applications Calcium Sulphate Calcium Phosphates Cal- cium Silicate Glass Mortar Calcium Sulphide Strontium Occurrence and Preparation Properties Compounds of Strontium Barium Occurrence Properties Barium Chlo- rideBarium Hydroxide Barium Oxide Barium Peroxide or Dioxide Barium Sulphide Barium Nitrate Barium Sul- phate Barium Carbonate Phosphates of Barium Reactions which are of Special Value in Analysis Magnesium Sub- Group Beryllium Occurrence and Preparation Properties Compounds of Beryllium Beryllium Chloride Beryllium Hydroxide Beryllium Sulphate Beryllium Carbonate Weak Basic Character of Beryllium Magnesium Occurrence Preparation Properties Applications Compounds of Mag- nesium Magnesium Chloride Magnesium Oxide Magne- sium Sulphate Magnesium Carbonate Phosphates Borates Silicates Silicon Magnesium Reactions of Magnesium Salts which are of Special Value in Chemical Analysis Er- bium General, .523 CHAPTER XXVII. ELEMENTS OF FAMILY III, GROUP A : ALUMINIUM SCANDIUM YTTRIUM LANTHANUM YTTERBIUM. 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 Ammon- ium Alum, Ammonium-Aluminium Sulphate Aluminium Silicate Kaoline Clay Ultramarine Porcelain Earth- enware Reactions of Aluminium Salts which are of Special Value in Chemical Analysis Other Members of Family III, Group A Scandium Yttrium Ytterbium The Boron- Aluminium Group in General 559 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 Copper Hydride Cuprous Chloride Cupric Chloride Cuprous Iodide Cuprous Hydroxide Cuprous Oxide Cupric Hydroxide Cupric Oxide Other Oxides of Copper Cupric Sulphate Cupric Nitrate Cupric Arsenite Cupric Carbonates Cyanides of Copper Cuprous Sulphocyanate XVI CONTENTS. PAGE Cupric Sulphocyanate 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 Alloys of Silver Argentous Chloride Silver Chloride, Argentic Chloride Application of the Chloride, Bromide, and Iodide of Silver in the Art of Photography Silver Oxide Other Oxides of Silver Sul- phides of Silver Silver Nitrate, Argentic Nitrate Silver Cyanide Silver Sulphocyanate Borates of Silver Reactions which are of Special Value in Chemical Analysis Gold Gen- eral Forms in which Gold occurs in Nature Metallurgy of Gold Properties Alloys of Gold Chlorides of Gold Chlor- auric Acid Cyanauric Acid Auric Hydroxide, 583 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- cations Amalgams 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, 610 ELEMENTS OF FAMILY III, GROUP B : GALLIUM INDIUM THALLIUM. General Gallium Compounds of Gallium Indium Compounds of Indium Thallium Compounds of Thallium, 629 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. XVil 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 Iodide Lead Hydroxide Oxides of Lead Lead Suboxide Lead Oxide Lead Sesquioxide Lead Peroxide Red Lead, Minium Lead Sulphide Lead Nitrate Lead Carbonate Lead Sulphate Reactions which are of Special Value in Chem- ical Analysis Lanthanum Cerium Didymium, 632 CHAPTER XXXI. ELEMENTS OF FAMILY VJ, 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 Chrornate Po- tassium Dichromate Chromium Trioxide Relations between the Chromates and Bichromates Sodium Chrornate 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 Phosphomolybdic Acid Tungsten General Occur- rence and Preparation Properties Chlorides Oxides Tungstic Acid and the Tungstates Silicotungstic Acids Uranium General Occurrence and Preparation Chlorides Oxides Uranous Salts Uranyl Salts, 651 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 Mauganous Hy- droxide Manganous-manganic Oxide Manganic Oxide Manganese Dioxide Manganites Weldon's Process for the Regeneration of Manganese Dioxide in the Preparation of Chlorine Sulphides Manganous Cyanide Manganous Car- bonate Mauganous Sulphate Manganic Sulphate Manganic Acid and the Manganates Permanganic Acid and the Per- manganates Potassium Permanganate Reactions which are of Special Value in Chemical Analysis, 672 xviii CONTENTS. CHAPTER XXXIII. ELEMENTS OP FAMILY VIII, SUB-GROUP A : IRON COBALT NICKEL. PAGE General Iron Introductory Forms in which Iron occurs in Nature Metallurgy Varieties of Iron Steel Properties of Iron Ferrous Chloride Ferric Chloride Cyanides Potas- sium Ferrocyanide Ferro-hydrocyanic Acid Ferric Ferro- cyanide, or Prussian Blue Potassium Ferricyanide Ferri- hydrocyanic 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 Bisulphide Arsenopyrite Reactions which are of Special Value in Chemical Analysis Ferrous Compounds Ferric Compounds Cobalt General Occurrence and Prep- aration Properties 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, 685 CHAPTER XXXIV. ELEMENTS OP FAMILY VIII, SUB-GROUP B : RUTHENIUM RHODIUM PALLADIUM. ELEMENTS OP FAMILY VIII, SUB-GROUP C : OSMIUM IRIDIUM PLATINUM. General The Platinum Metals Ruthenium Properties Chlo- rides Oxides Osmium Preparation Properties Chlorides Oxides Rhodium Iridium Preparation Properties Chlorides Oxides Palladium Preparation Properties Palladium Hydrogen Chlorides Oxides Platinum Prep- arationPropertiesApplications of Platinum Alloys of Platinum Chlorides Chlorplatinic Acid Cyanides Hy- droxides and Oxides Sulphides Compounds with Ammonia, 712 APPENDIX CONTAINING SPECIAL DIRECTIONS FOR LABORATORY WORK. Introduction, 727 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, , 728 CONTENTS. XIX EXPERIMENTS TO ACCOMPANY CHAPTER U. 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, 734 EXPERIMENTS TO ACCOMPANY CHAPTER 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 745 EXPERIMENTS TO ACCOMPANY CHAPTER IV. Composition of Water Eudiometric Experiments Oxyhydrogen Blow-pipe 754 EXPERIMENTS TO ACCOMPANY CHAPTER V. Organic Substances contain Water Water of Crystallization Efflorescent Salts Deliquescent Salts Purification of Water by Distillation 757 EXPERIMENTS TO ACCOMPANY CHAPTER VI. Method of Dumas Method of Victor Meyer, 759 EXPERIMENTS TO ACCOMPANY CHAPTER VIL Ozone Hydrogen Dioxide, 761 EXPERIMENTS TO ACCOMPANY CHAPTER Vm. Preparation of Chlorine Chlorine decomposes Water in the Sun- light Chlorine Hydrate Formation of Hydrochloric Acid Preparation of Hydrochloric Acid, 762 EXPERIMENTS TO ACCOMPANY CHAPTER IX. Chloric Acid and Potassium Chlorate Perchloric Acid, .... 766 EXPERIMENTS TO ACCOMPANY CHAPTER X. Neutralization of Acids and Bases ; Formation of Salts Study of the Products formed, 768 FOR CHAPTER XI., 770 EXPERIMENTS TO ACCOMPANY CHAPTER XII. Preparation of Bromine Hydrobromic Acid Iodine Iodine can be detected by Means of its Action upon Starch-paste lodic Acid, . ... 770 XX CONTENTS. EXPERIMENTS'TO ACCOMPANY CHAPTER xin. PAGE Properties of Sulphur Hydrogen Sulphide, . . % , ; . . . 773 EXPERIMENTS TO ACCOMPANY CHAPTER XIV. Manufacture of Sulphuric Acid Sulphurous Acid and Sulphur Dioxide Sulphurous Acid is a Reducing Agent Sulphur Tri- oxide, 775 EXPERIMENTS TO ACCOMPANY CHAPTER XV. Preparation of Nitrogen Analysis of Air, 778 EXPERIMENTS TO ACCOMPANY CHAPTER XVI. Preparation and Properties of Ammonia Ammonia burns in Oxy- gen Ammonia forms Ammonium Salts with AcidsCompo- 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 N itric Oxide Nitrogen Trioxide Nitrogen Peroxide, 782 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, 790 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, 795 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, . 797 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, 799 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, 801 CONTENTS. xxi EXPERIMENTS TO ACCOMPANY CHAPTER XXH. PAGE Silicon Silicon Tetrafluoride and Fluosilicic Acid Silicic Acid, . 804 EXPERIMENTS TO ACCOMPANY CHAPTER XXIV. Chlorides, Bromides, and Iodides Hydroxides Sulphates Re- duction of Sulphates to Sulphides Carbonates, 806 EXPERIMENTS TO ACCOMPANY CHAPTER XXV. Potassium Salts Sodium Salts 810 EXPERIMENTS TO ACCOMPANY CHAPTER XXVI. Calcium Salts Magnesium and its Salts, 811 EXPERIMENTS TO ACCOMPANY CHAPTER XXVII. Aluminium Chloride, 812 EXPERIMENTS TO ACCOMPANY CHAPTER XXVIII. Copper and its Salts Silver and its Salts, 812 EXPERIMENTS TO ACCOMPANY CHAPTER XXIX. Zinc and its Salts Mercury and its Salts 813 EXPERIMENTS TO ACCOMPANY CHAPTER XXX. Tin and its Compounds Lead and its Compounds 813 EXPERIMENTS TO ACCOMPANY CHAPTER XXXI. Chromic Acid and the Chromates, 814 EXPERIMENTS TO ACCOMPANY CHAPTER XXXII. Manganese and its Compounds, ........ .... 815 EXPERIMENTS TO ACCOMPANY CHAPTER XXXIV. Platinum "... 815 Conclusion, , - . . , 816 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 wo 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) INORGANIC CHEMISTRY. Ghange. 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 chemical 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 electricity 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 off light, and giving no evidence of being warm. Kemove 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 no longer has the power to attract other pieces of iron ; and does not differ 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. 4 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 mere 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 a 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 burning or combus- tion which for a long time had been a subject of study. The explanation and the experiments which led to it will be taken up later. Suffice it to say here that the essen- tial feature of the work consisted in the fact that the substances which entered into action and those formed by 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 with 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 pl,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 shown that the total COMPOSITION OF MATTER. 7 amount of energy is unchangeable. One of the greatest discoveries in science is 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 mo- tion we can get a certain amount of heat. We know that a similar definite relation exists between heat and elec- trical 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 laio 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 simple 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 each other, and thus form more complex substaoces, 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 do 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 w^hich 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 would 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 call 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 with 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 able 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 possible to separate the three com- ponents 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 LA W 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 laiv of definite proportions has been drawn. This law may be stated thus : A chemical compound alivays 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 experiments ever performed in chemistry, showed that in the compounds which he worked on 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 cafbon 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 elements A and B form several compounds with each other, and we consider any fixed mass of A, tJien 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 respectively contain 7 parts of nitrogen and 8, 16, 24, 32 y and 40 parts of oxygen. The figures representing 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, potas- sium, and oxygen the proportions by weight are repre- sented in the following table : Chlorine. Potassium. Oxygen. 35.37 39.03 15.96 35.37 39.03 31.92 35.37 39.03 47.88 35.37 39.03 63.84 It will be observed that the ratio between the chlorine and potassium remains constant, but that the mass of oxygen varies regularly from 15.96 to 63.84 ; 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 35.37 parts Chlorine. 1 " t (i 35.37 parts Chlorine combine 79.76 " Bromine " 126.54 " Iodine 7976 126.54 39.03 39.03 39.03 Bromine. Iodine. Potassium.. 18 INORGANIC CHEMISTRY, 15.96 parts Oxygen combine with 65.1 parts Zinc. 15.96 .. 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, 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 known 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 be 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 which is required to express the weight of the oxygen which enters into the composition of the molecules, and that 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. Water 17.96 Carbon monoxide 27.93 Carbon dioxide. 43.89 Nitric oxide , 29.97 Nitrous oxide , 43.98 Sulphur dioxide 63.9 Sulphur trioxide 79.86 2 15.96 11.97 15.96 11.97 31.92 14.01 15.96 28.02 15.96 31.98 31.92 31.98 47.88 Composition. parts hydrogen, oxygen, carbon, oxygen, carbon, oxygen, nitrogen, oxygen, nitrogen, oxygen, sulphur, oxygen, sulphur, 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 27.93 times as much as the atom of hydrogen, and the 27.93 parts of matter are made up of 11.97 parts of carbon and 15.96 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 15.96 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 15.96 as great as that of the atom of hydrogen. The conclusion is therefore drawn that 15.96 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 15.96 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 HC1 ; 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 such compounds as are not gaseous, nor convertible into vapor there is, however, considerable doubt, as there is no general reliable method for the de- termination of molecular weights except the one described. 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- \H 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 dement. 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 O = O ; but this involves the conception that oxy- gen may act as a quadrivalent element ; and although there are some facts which give an air of plausibility to this conception, the evidence in favor of it is hardly suffi- cient to warrant the acceptance of the above formula. 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 3 O 2 + 2KI = 2KOH + I 3 . 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 hydrogen dioxide. If in a glass cylinder a layer of ether be poured upon a solution of hydrogen dioxide and a drop of a solution of potassium dichromate be 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 : Ag,0 + HA = Ag, + 11,0 + O,. So, also, it undergoes decomposition with ozone as rep- resented thus : 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 (0 2 , 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 2 2 = H 2 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. Afterwards, when it was learned that yellow crystals could be obtained from it by sub- jecting it to a low temperature, it was supposed to be a mixture of substances. In 1810 Davy studied these crystals and showed that they contained water, and a little later Faraday showed that the gas could be liquefied. For a long time it was supposed to contain oxygen, until, finally, all efforts to obtain oxygen from it having failed, it came to be looked upon as 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. Preparation. The problem to be solved in the prepa- ration of chlorine from common salt is how to separate (96) PREPARATION OF CHLORINE. 97 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 by heating. The method adopted consists in making hydrochloric acid, HC1, from sodium chloride, and then treating the hydrochloric acid with some substance which readily gives off oxygen. Owing to the strong affinity of oxygen for hydrogen, the two combine to form water and the chlorine is thus set free. The first change, that of sodium chloride to hydrochloric acid, is readily accomplished by treating salt with ordi- nary sulphuric acid a reaction which is 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 S0 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) 2HC1 + = H,O + C1 3 . 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 has not on the whole proved successful, and is practically supplanted by another, which is apparently more complicated, known as Weldon's process. In the laboratory the method employed consists in bringing hydrochloric acid in con- tact with manganese dioxide. The reaction is repre- sented thus : (1) Mn0 2 + 4HCl = MnCl 2 +2H 2 + 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 is probably 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 di- oxide is represented by the equation (1) above. Some recent investigations make it appear possible that the reaction is somewhat more complicated than it is here represented to be ; that the first product of the action of hydrochloric acid on manganese dioxide is a compound CHLORINE WELDOFS PROCESS. of the composition H 2 MnCl 6 ; and that this compound breaks down under the influence of heat thus : H 2 MnCl 6 = MnCl a + 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 given off. 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 3 + H a SO 4 = MnSO 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 2 , 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). The steps in- volved are (1) treatment of the waste-solutions contain- ing manganous chloride with calcium carbonate and lime ; and (2) treatment of the liquid thus obtained with steam and air. In this way a compound, calciiim manganite, CaMnO 3 , is obtained, which when treated with hydro- chloric acid undergoes the following reaction : CaMn0 3 + 6HC1 = CaCl a + MnCl 2 + 3H a O + Cl a . 100 INORGANIC CHEMISTRY. Properties. Chlorine is a greenish-yellow gas. It has a disagreeable odor, and acts upon the membranes lining the throat and nose, causing irritation and inflammation. The effect is much like that of a " cold in the head." Inhaled in concentrated form, i.e., not diluted with a great deal of air, it would cause death. It is much heavier than the air ; its specific gravity is 2.45 (air = 1), and as compared with hydrogen it is 35.37. A liter of chlorine gas, under standard conditions, weighs 3.167 grams. It is soluble in water and acts upon mercury, and therefore cannot be collected by displacement of either of these liquids. The most convenient way to collect it is by displacement of air. It can also be col- lected 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 a remarkably active substance, combining with or acting in some way upon most other substances even at ordinary temperature. This activity may be illustrated 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. Very marked changes will be observed at once. The antimony will take fire and a white substance will be formed. The reaction has been studied, and been found to take place in accordance with the following equation : Sb + 3C1 = SbCl 3 . The copper-foil also takes fire, and is converted into a, chloride as represented thus : Cu + 2C1 = CuCl 2 . 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 and which we have already studied under the name of Com- bustion. There is, however, this difference between re- actions in oxygen and in chlorine : the latter frequently take place at ordinary temperatures, whereas those in PROPERTIES OF CHLORINE., 101 oxygen require an elevation of temperature' to' start them. In both cases the gases combine : 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 be 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 the act of neutralization is a definite one, which 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 be 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 : HN0 3 +KOH =KN0 3 + H 2 O ; HN0 3 +NaOH = NaNO 3 + H 2 O ; 2HNO, + Ca(OH), = 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 = Naj3O 4 + 2H 2 O ; H 2 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 which differs from the acid in con- taining 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 O. 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 could 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 replacing the metal of the base by hydrogen is water ; (5) That the substance obtained from the acid by re- placing the hydrogen by a metal 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 replac- ing the hydrogen by a metal the substance has an alka- line 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 ftosc 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 S , po- tassium nitrate ; NaNO s , sodium nitrate ; Ca(XO,) t , cal- cium nitrate ; KjSO,, potassium sulphate ; XajSO^ 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 Adds 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 Here, as when the hydroxide is used, the acid is neutral- ized and the salt, sodium sulphate, Xa 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 nay be written thus : NaOH H) Na) ,HOH KaCl H) Na) Ha Nad The two reactions are thus seen to be of the same general character. That with the chloride does not take place 133 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 0; CaS + H 2 SO 4 = CaSO 4 +-H.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 SO 4 = CaSO 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, >ur 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 w T hich 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 = A1C1. + 3H 2 O. But it also neutralizes potassium hydroxide according to the equation A1(OH) 3 + 3KOH = Al(OK), + 3H,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 3 Sulphuric acid, (HO) 2 SO 2 Phosphoric acid, .,..-.. . (HO) 3 PO Carbonic acid, ... . . . (HO) 2 CO Metaphosphoric acid, .... (HO)PO a Nitrous acid, (HO)NO Arsenious acid, (HO) 3 As Hypochlorous acid, .... (HO)C1 Perchloric acid, . . . * : u |J . (HO)C10 3 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 H-O X 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), (0x A)"(OH) 2 HOH III. B"YOH) 3 HOH (O X A)'"(OH) 3 HOH In these general formulas B" means any bivalent metal, and B'" 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 a S), and (O X A) X// 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-O|-K = O a N-0-K + H-O-H ; or in this : O 2 N-|O-H + H|-O-K = 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 0; KOH + HN0 3 = KN0 3 + H 2 O. If only half the quantity of base which is required to neutralize the acid be 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 be added, half the base will remain unchanged. On the other hand, if an acid like sulphuric acid be 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 be added, the first reaction takes place ; but if the act of neutralization be 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. ; tetrabasic acids, like pyrophosphoric acid, H 4 P 2 O 7 ; pentabasic acids, like peri- odic acid, H B 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), 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 possible. 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 < ^ + H 2 O; Zn(OH) 2 + 2HC1 = ZnCl a + 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 ( NO, BiJ OH + HNO 3 = Bi^ OH + H 2 O; OH OH Bij OH ( N0 3 OH + 2HNO 3 = Bi^ NO, + 2H,O; (OH (OH ( OH ( NO 3 i^ OH + 3HNO 3 = Bi^ NO, + 3H 2 O. (OH (NO, The salts Bi | QH\ and Bi 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) Acid 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 hydrogen 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 ./OH") * s ca ^ e( i the mono-nitrate; that of the formula Bi ! XTT , 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 : CO + 2H,O. r ^x^-"- AJLV ^ ru, ^^^ < OH UU< 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,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 will be 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 hjdrobromic acid, HBr ; hydriodic acid, HI ; and hydrocyanic acid, HCN. 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, and others containing chlorine which form compounds with chlorides in much the same way that the oxygen acids form compounds with oxygen bases, and the compounds formed are analo- gous to ordinary salts, only they contain sulphur or chlo- rine or some other element in place of oxygen. Thus, there is a compound of arsenic and sulphur of the compo- sition H 3 AsS 4 , known as sulpharsenic acid, which is analo- 142 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 w r ell- 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 PtO 3 + 2KOH = K 2 PtO 3 + H 2 O. 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 Q 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 under. Thus there are the following examples : Hy- pochlorous acid, HC1O ; hyposulphurous acid, H 2 SO 2 ; hyponitrous acid, HNO : 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) s , and ferrous hydroxide the composition Fe(OH), ; cuprous hy- droxide is Cu(OH), and cupric hydroxide Cu(OH) Q , etc. These compounds are sometimes called hydrates, and there are some good reasons for using this name, as will be more fully shown in the next 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 be 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 3 + 10H 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 can 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 cMorites; those of nitrous acid, nitrites ; those of sulphurous acid, sulphites; etc., etc. It will be noticed that the terminal 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 ) 3 ; 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 phospTwrates, but for the sake of euphony and convenience these names are shortened to the above forms. 146 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 hydrochlorates, 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 a + H 2 ; while zinc and chlorine act thus : Zn + C1 2 = ZnCl 2 . 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 hydrocMorate is unnecessary. The name hydrate to which reference was made in the last paragraph 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 2 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 35.37, 79.76, and 126.54 respectively. It will be seen that the atomic weight of bromine, 79.76, is approximately the mean of those of chlorine and iodine. We have = 80.95. A 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 = 31.98, Se = 78.87, and Te = 125. We have here 3L98 125 Other groups are those of phosphorus, 30.96, vanadi- um, 51.1, and arsenic, 74.9 : lithium, 7.01, sodium, 23, and potassium, 39.01 : ^U 23.01. (147) 148 INORGANIC CHEMISTRY. In 1863-64 J. A. E. Newlands called attention to the fact that if all the elements be 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 be 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. Mendelejeff 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 law is a fundamental law of chemistry. Arrangement of the Elements. Mendelejeff 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. MENDELEJEFF'S TABLES OF THE ELEMENTS. 149 CO oo'os oc'i- H si II II 1 si II II 1 ^ II s i ' s" 2 .~: 1 S : 8' 1 n n 8 II II 1 ~- 91 1 d > - i- l~ II o 8 ~ II OS 1 ^ 1 S ^ o 1 II II h 1 II 1 1 1 1 . g? g *?" 1 1 > II co II os " h? cT OQ * II s OS II II U II i II II ^ H 10 eo 00 - i" 1 o S TABLE ill o II fc II Q. Q9 II II $ s II J3 i a 02 II 5 II e II 13 1 1 . ^ GO i at M . . ' ^ II rt * eo" w "* W O CO oo o a S II 1 Q II O " SO,. From 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 VII 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, Mendelejeff 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 Mendelejeff 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-CHLORINE BROMINE IODINE. 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 elements the elements of this group are univalent, as shown in such compounds as- NaCl, KBr, CaCl 2 , 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. 79.76). 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 there is some bromine in the form of sodium bromide, NaBr, and in some places it occurs as potassium bromide, KBr. 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 is prepared from the bromides in the same way that chlorine is made from the chlorides : by first treating with sulphuric acid, thus liberating hy- drobromic acid, and then treating with manganese diox- ide, or, better, by mixing the bromide with manganese dioxide and treating the mixture with sulphuric acid. The complete reaction is represented by the equation 2NaBr + MnO 2 + 2fl 2 SO 4 = Na 2 S0 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- brcmic acid from the bromide, and this would act upon the manganese dioxide thus : MnO 2 + 4HBr = MnBr 2 + 2H 2 O + Br 2 . But sulphuric acid would act upon manganous bromide, MnBr 2 , thus : MnBr a -f H 2 S0 4 = MnSO 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 be 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 fipoofiio?, a stench). Its properties are, in general, very much like those of chlorine. It acts violently upon organic substances ; at- tacking the skin, and the membranes lining the passages of the throat and lungs, in much the same way as chlorine. Wounds caused by the liquid coming in contact with the skin are painful and serious, and it must therefore 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 water. At a low temperature bromine forms with water a compound in every way analogous to chlorine hydrate, viz., 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 6 , yields the products" C 6 H 5 Br, C 6 H 4 Br 2 , C 6 H a 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 2 O = or 2Br, + 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 1886 amounted to over 400,000 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 S0 4 = 2H 2 O + SO 2 + Br a . 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 + SO,. 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 M were used, the reaction would be represented in this way : H7DROBROMIC 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 Br + HHO = PO 3 H 3 + 3HBr, Tir. TTTTn Br HHO or PBr 3 + 3H,O = 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 + 5H,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 . Hypobromous 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 2 = KBr + KBrO + H 3 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 2 + 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 = HBrO 4 + Cl. 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. 126.54). 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 3 + 2H,SO 4 = K 2 SO 4 + MnSO 4 + 2H 2 O + 1 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 Q . A con- siderable quantity of iodine is now 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 i^ 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 a SO 4 : H a SO 3 + 1 2 + H 2 O = H,SO 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 + 2HI = 2H 2 + SO, + 1, ; H 2 S0 4 + 6HI = 4H 2 + S + 3I 2 ; H 2 SO 4 + SHI = 4H 2 O + SH 2 + 4I 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 = 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 + IC1. Iodine monochloride is also decomposed by water, giv- ing iodic acid, hydrochloric acid, and free iodine : 10IC1 + 6H 2 = 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 = SKI + 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 : 172 INORGANIC CHEMISTRY. Its salts have the general formula MIO 3 , though it also 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 5 . 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 5 ; 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 3 ; CO 2 + H 2 O = H 2 CO 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 a + 2KOH = K 2 Si0 3 + H 2 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 Q 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 : A1A + 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 0. 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 + 01, = Na 2 H 3 IO e + 2NaCl. The salt Na 2 H 3 IO fl 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 : 2Ag 2 H 3 IO 6 + 2HNO 3 = 2AgNO 3 + 4H 2 O + 2AgIO 4 . 174 INORGANIC CHEMISTRY. The acid when separated from its solutions is a crys- tallized solid which has the composition H 5 IO 6 . When heated it undergoes decomposition, losing water and oxygen, and yielding iodic acid. It cannot, however, be 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 , Ag.1,0., ZnJ.0,,. 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 is 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 : OH OH OH OH OH 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 6 IO 6 can also be imagined to break down into H 3 IO 5 and water thus : O OH OH _ T OH : OH OH CONSTITUTION OF PERIODIC ACID. 175 The salt Ag 3 IO 6 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 : O O OH OH 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) 7 , 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 5 one molecule of water be taken, an acid of the formula H 4 I 2 O 9 would be formed ; or if two molecules of the acid H 5 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 : O O OH = (HO) 2 2 I-0-I0 2 (OH) 2 + H 2 ; OH OH 21 ^ O OH = (HO) 4 OI-0-IO(OH) 4 + H 2 C OH OH A salt of the formula Ag 4 I 2 O 9 , and another of the for- mula Zn 4 I Q O n , are known. The former is derived from the acid H 4 I 2 O 9 , the latter from the acid H 8 I 2 O n , by re- placement of the eight atoms of hydrogen by four biva- lent atoms of zinc. 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) 7J 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) 6 by loss of water : , = I0 2 (OH) + 2H 3 0; or r OH OH (O I-! OH = II O +2H 2 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 O=I=O 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. 177 chloric acid, and free iodine, as stated under lodic Acid (p. 171). If the passage of chlorine over iodine be 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. 19.06). 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 3 .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 up to quite re- cently, when its isolation was effected by passing an electric current through liquid hydrofluoric acid contained in a platinum vessel. Properties. Fluorine is the most active element at ordinary temperatures. It is a faintly greenish-yellow gas. It acts upon most substances. Thus, it decomposes water, yielding ozone and hydrofluoric acid; it combines directly with hydrogen at the ordinary temperature ; it 178 INORGANIC CHEMISTRY. combines with sulphur, phosphorus, iron, etc., with evo- lution 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, HP. 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 SO 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 dissolve 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, + 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 , and this perhaps represents the substance with which we ordinarily have to deal as hydrofluoric acid. It is quite possible that, at considerably lower temperatures than the ordinary, hydrochloric acid may have the formula H 2 C1 2 . 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 by the formula xF-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 5 . This is a liquid which is formed by treating silver fluoride, AgF, with iodine. Water decomposes it, forming iodic acid : IF 5 + 3H 2 = 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. 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 wiih Hydrogen. HF(H 2 F a ) HC1 HBr HI Compounds tvith Oxygen. C1 2 CIA CIO. Compounds tvith Hydrogen and Oxygen. HC1O HBrO HC10 a HC1O 3 HBrO 3 HIO 3 HC10 4 HI0 4 (H 5 I0 6 ) Compounds ivith One Another. ClBr IC1, IBr IC1 3 ,IF B 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 + C1 2 = 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. 182 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 Q 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 be dissolved in water it undergoes partial decomposition, which is complete if the solution be boiled, or if carbon dioxide be passed through it. The change consists in the transformation of manganic acid into per- manganic acid, HMnO 4 : 3H 2 MnO 4 = 2HMn0 4 + MnO 2 + 2H 2 O ; or 3K 2 Mn0 4 + 2H 2 O = 2KMnO 4 + MnO 2 + 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 be 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 hydrochloric acid potas- sium permanganate forms potassium chloride, manganous chloride, 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 : 2KMn0 4 + 6HC1 = 2KC1 + 2MnCl 2 + 3H 2 O + 5O. Similarly, with sulphuric acid manganous sulphate is formed, thus : 2KMnO 4 -f 3H 2 SO 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 7 , stands in the relation of an anhydride to permanganic acid. In water solution it passes over into the acid : Mn a 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 a SO 4 + Mn 2 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 II O=Mn-O-H, in which the manganese is septivalent. 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 2 Te. Of oxygen and sulphur there are other hydrogen compounds, as hydrogen dioxide, H 2 O 2 , and an analogous compound of sulphur, but it is probable that in these the valence towards hydrogen is 2, as in the more stable compounds, as was pointed out under Hydrogen Dioxide (v/hich see). It appears that 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,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 INORGANIC CHEMISTRY. forms the compound SI 6 , in which it is sexivalent. 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 6 . Towards oxygen the three elements of the sulphur group are quadrivalent and sexivalent, as seen in the compounds SO 2 , SeO 2 , TeO 2 , and SO 3 , SeO 3 , and TeO 3 . Of course, it is possible that in these oxygen compounds 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 V . 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 sexivalence 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 SO 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 jfinally 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. SULPHUR, S (At. Wt. 31.98). 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 accomplished by piling the ore in such a way as to. leave passages for air, cover- ing the piles with earthy matter to prevent free access of air and then setting fire to them. A part of the sul- phur burns, and the heat thus furnished melts the rest of the sulphur. The molten sulphur runs down to the bottom of the pile, and by a proper arrangement is drawn off from time to time. If the pile of ore were not covered up, and the oxygen allowed free access, the sul- phur would burn up, yielding the gas sulphur dioxide. The "crude brimstone" obtained in the manner de- scribed 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, as in making 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 3 ; 3MnO 2 = Mn 8 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 solutions of sulphur. The latter belong to the rhombic system. They are octahedrons with a rhombic base, which is also the form of the sulphur found in nature. The former are honey-yellow needles. A careful 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 be 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 be 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 O 3 + 2HC1 = H,S 2 O 3 + 2NaCl ; H,S,0 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 depositing 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 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 power to combine with oxygen, however, it differs markedly from oxvgen itself, and it also 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-form- ing elements if heated with them to a sufficiently high temperature. Its affinity for most other elements is less than that of oxygen. Thus, its compound with hydro- gen is decomposed very much more readily into its ele- ments than water is ; and the sulphides or its compounds with metals are decomposed by heating them 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 has a strong affinity for 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 S0 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 i& 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 = Na 2 S + H 2 + O; (2) Na 2 S + 4S = Na 2 S 6 ; (3) Na 2 S +S -f 30 = Na 2 S 2 O 3 . Expressing these reactions in one equation we have 6NaOH + 12S = 2Na 2 S 6 + 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. Attention has already been called to the fact that the specific gravity of the vapor of sulphur varies with the temperature, and is such as to indicate that at tempera- tures not far above the boiling point the molecule con- sists of six atoms, while at the temperature 800 and higher the molecule consists of two atoms. This sug- gests the question whether the molecule of sulphur m the solid form may not be more complex than the con- dition represented by the symbol S 6 . We have no means, of answering this question with any certainty at present. COMPOUNDS OF SULPHUR WITH HYDROGEN. 193 Uses 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 SULPHUR WITH HYDROGEN. 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 2 + 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 SO 4 + SHI = H 2 S + 4H 2 O + 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 + 30 = H 2 + SO 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 ol 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 staliie. 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 3 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 SO 4 + 2H = 2H 2 O + SO 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 + Cl a = 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 0. 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 replacement of the hydrogen by the metals 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 + 2HN0 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 sul- 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 : FeSO 4 + K 2 S = K 2 SO 4 + FeS. Here no sulphuric acid is formed, but, instea'd 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 : FeS0 4 + (NH 4 ) 2 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 pliides 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 s S s + 3(XH 4 ),S = 2(NH,) 3 AsS, 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 : 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 (b) 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 O K 2 S 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 2 S 2 (?). 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 + 4O = K 2 S0 4 . But for reasons which will be more fully considered 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 S 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 chlorine were simply re- placed by oxygen, 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 = S0 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, which breaks down into sulphur and iodine by exposure to the air. SELENIUM, Se (At. Wt. 78.87). Occurrence. Selenium occurs only in small quantity in 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 S 1 OH + HO \ S< OH = (HO) 2 OSSO(OH) 2 + 4H 2 O OH HOJ An acid formed in this way would have the formula DISULPHURIC ACID. 219 H 4 S 2 O 8 , and the salt K 3 HS 3 O 8 is tlie tertiary potassium salt of this acid. Disulphuric Acid, Pyrosulphuric Acid, H a S 2 O 7 . This compound, which is also known by the names fuming sidphuric 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 : It is made by distilling ferrous sulphate which is not perfectly dry : 4FeS0 4 + H 2 = 2Fe 2 O 3 + 2SO 2 + H 2 S 2 O 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 : I8> SO > = Or the formula of the acid may be written thus : O 2 S-OH 6 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 7 - M 2 S0 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 O 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 tliat the solution contains the acid which is formed by the action of sulphur dioxide on water, thus : S0 2 + H 2 O = H 2 SO 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 : SULPHUEOUS ACID. 221 H 2 Se0 3 + 2SO a + H 2 O = Se + 2H 2 SO 4 ; or H 2 Se0 3 + 2H 2 SO S = 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 : 2HIO 3 + 5H 2 S0 3 = H 2 O + 5H 2 SO 4 + I 2 . If sulphurous acid be 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 SO 2 , 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,S. Treated with zinc alone the reduction is not carried as far as this, the reduction-product being hyposulphurous acid, H 2 SO 2 : Zn + 2H 2 SO 3 = ZnSO 3 + H 2 SO 2 + H,O. Sulphurous acid when heated in a sealed tube breaks down into sulphuric acid and sulphur : 3H,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 2 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 S0 3 + O = Na 2 S0 4 ; Na 2 S0 3 + S = Na 2 S 2 3 (or Na 2 SO 3 S). 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 S0 3 + 2HC1 = 2NaCl + H 2 O + SO, ; with sulphuric acid thus : Na 2 S0 3 + H 2 S0 4 = Na 2 S0 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. Hyposulphurous Acid, H 2 SO 2 . 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 sulphurous acid by treating with zinc, when this reac- tion takes place : Zn + 2H 3 S0 3 = ZnS0 3 + H 2 SO 2 + H 2 O. The reduction is in all probability a secondary action, due to the hydrogen liberated from the sulphurous acid by the action of the zinc : Zn + H 2 S0 3 = ZnSO 3 + 2H ; H 2 S0 3 + 2H = H 2 S0 2 + H 2 0. 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 : 2H.SO. = H.8,0, + H.O ; . =80. + S + H.O. It will be seen that thiosulphuric acid bears to hypo- sulphurous acid a relation similar to that which disul- phuric acid bears to sulphuric acid. Just as acid sul- phates are converted into disulphates when heated, so the hyposulphites, all of which have the composition MHSO a , break down into thiosulphates and water : 2MHSO, = M 2 S 2 3 + H 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 one atom of oxygen has been replaced by 224 INORGANIC CHEMISTRY. sulphur, and hence the name thiosulphuric acid is ap- propriate, whereas the name hyposulphurous acid sug- gests at once a compound similar to sulphurous acid, but containing less oxygen, and is therefore inappropriate. Sodium thiosulphate is formed together with the penta- sulphide by the action of sulphur upon sodium hy- droxide : 6NaOH + 128 = 2Na 2 S 5 + Na 2 S 2 O 9 + 3H 2 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 re- placement of three atoms of sulphur by three of oxygen : Na 2 S 5 + 30 = Na 2 S 2 O 3 + 38. Sodium thiosulphate is formed, further, by the action of iodine upon a mixture of sodium sulphide and sodium sulphite : Q .Na /Na k + 2NaI. ^ Q Na ' O 3 b\ 8< Na X Na 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 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 3 = 3Na a SO 4 + Na 2 S 6 . 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 a + 2SO a = 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 : H 2 S 2 6 = H 2 S0 4 + SO,. So, too, when a salt of the acid is heated it breaks down, forming a sulphate and sulphur dioxide : KJ3.0. = K 5 SO. + SO,. Trithionic Acid, H 2 S 3 O 6 , 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 + 2S = 2K 2 S 3 O 6 + K 2 S 2 O 3 + 3H 2 O. The sodium salt is formed by treating a mixture of so- dium sulphite and sodium thiosulphate with iodine : 226 INORGANIC CHEMISTRY. nQ .Na /Na U 3 b< TVfl O<4/ * + 21 = q X q> + 2NaL or* a ^a b(J,b\ SO 8< Na X Na When heated in solution or dry, the potassium salt is decomposed, yielding the sulphate, sulphur dioxide, and sulphur : K,S 3 0. = K.SO, + S0 2 + 8. Similarly, when treated with acids, decomposition takes place with evolution of sulphur dioxide, separation of sulphur, and formation of sulphuric acid : H 3 S 3 6 = H 2 S0 4 + SO, + S. Tetrathionic Acid, H 2 S 4 O , 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 : Pentathionic Acid, H 2 S 5 O 6 , is formed by the action of hydrogen sulphide upon a solution of sulphur dioxide in water. Constitution of the Acids of Sulphur. The existence of the oxide SO 3 and of the acid SI 6 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) 7 , and the analogous facts presented under Sul- CONSTITUTION OF THE ACIDS OF SULPHUR. 22? 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- ? TOE sented by the formula H-Q-S-O-H or S J ^ H . Of ii ( x o LO 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 INORGANIG CHEMISTRY. o o II II Y-O-S-O-X in one case, and by X-O-S-O-Y 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- (01 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- il I rangements is this, O=S-O-H, in which the sulphur is II O O II 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. 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 Sx ONa, the conduct of which is such as to show that in it the group E 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 E a com- (E pound of the formula Sx OE is obtained, in which it to, appears that one of the groups is in combination with sulphur and the other with oxygen. Now, it is possible (Cl by starting with the compound S < Cl, which is made, as to we shall see, by replacing one oxygen atom of sulphur dioxide by two chlorine atoms, to introduce in the place of the two chlorine atoms two groups, OE, and thus ob- (OE tain the compound S -< OE, which plainly has the same to composition as the one represented by the formula (E S-< OE, but it has a different constitution. It was found to, 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 E 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 ( TT I ( H by the formula O=S=O or S-< OH, which may also be 6 *> TT written O 2 S < QJJ. At the same time it appears quite pos- sible that a sulphurous acid of the other constitution, OTT , 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. it II 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 ; ll II O O O I! 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 : ,Na < Na S-Na Na +2I = ^_ _ Na +2NaI; O=S-O-Na II 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, 232 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 O O i + Na-S-S-O-Na = NaO-S-S-S-S-O-Na + NaI. & & 6 & 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, and S0 3 , or sulphur di- oxide and sulphur trioxide. Besides these two it also forms a sesquioxide, S 2 O 3 , and a heptoxide, S 2 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 referred 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 = C0 3 + 2S0 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 2.24. 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 = 2S0 3 + S. When conducted into solutions of bases or of carbonates the corresponding sulphites are formed : SULPHUR TRIOXIDE. 235 2KOH + SO, = K a SO 3 + H 3 O ; K a CO 3 + S0 a = K a SO 3 + CO ,. Under certain conditions, as when the gas is passed into a hot solution of an alkali carbonate, a salt of the general formula M a S a O 5 is formed. This bears to the sulphite the same relation that the pyrosulphate bears to the sulphate : 2KHS0 3 =:K a S a 5 + H a O; and 2KHSO 4 = K a S a 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 '-herefore 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 Similarly, the acid sulphates of the alkali metals yield the corresponding normal sulphates and sulphur trioxide : 2NaHS0 4 = Na 2 S0 4 + SO 3 + H,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 a 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 -|- 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 : oH gg 0=8=0 + H*0 = Off H * I OH [OH 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 PIC 1 !-- H thus, ^ a Se < QTT , a few compounds of this kind I ing 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 : HOH r OH , HOH rp, J OH i + HOH = Te ] OH + HOH [OH OH The potassium salt of the acid is formed by melting together tellurium dioxide and potassium carbonate. K 2 C0 3 + TeO 2 = K 2 TeO 3 + CO 2 . If the salt thus formed is dissolved in water and nitric acid added to the solution, tellurious acid is thrown down: K 2 Te0 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 = TeO 2 (OH) a + 2H a O. Heated higher, to about 160, the acid is decomposed into tellurium trioxide and water : TeO a (OH) 3 = TeO 3 + H Q O. Although most of the tellurates are simple salts of the acid H a 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 : 4Te(OH). = Te t O n (OH), + HH.O. Oxides of Tellurium. Tellurium dioxide is formed by burning tellurium or by oxidizing it with nitric acid ; and, further, by the decomposition of tellurious acid by heat. It crystallizes 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, GROUP 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 CrO 4 + 10HC1 = 4KC1 + 2CrCl 3 + 5H 2 O + 3O. 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 2 (SO 4 ) 3 + 5H 2 O + 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. NITKOGEN, N (At. Wt. 14.01). 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 O 6 . The hydrogen compound is entirely different in character from those of chlorine and sulphur, for, while these are acid, the hydro- gen compound of nitrogen 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 $GDTIKO$, 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 ytreir, 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, an< I 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 3 + 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 1.257 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 is in solution in water be 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 wafor, 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 FBTTTTTFO- P art ^ ^ e 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-oVinr ^ * 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 Sckeele were made by burning sub- stances in air contained in closed vessels. They con- cluded that the air is made up of -J- oxygen and -J 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 a . 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. By other methods which need not be considered here it can be shown that there are many other substances in the air besides those men- tioned. Among them are ammonia, ozone, hydrogen di- oxide, and organic matters of various kinds, including a large variety of germs the presence of which can be de- tected 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 phos- phorus is inserted into a measured volume of air con- tained in a graduated glass tube over water or mercury, a diminution in volume will take place slowly. If, in the course of a few hours, the volume is again measured, 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, 254 INORGANIC CHEMISTRY. 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 a&, 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 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 ANALYSIS OF AIR. 255 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 be 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- 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 256 INORGANIC CHEMISTRY. 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. Water vapor is lighter than air. When air which is charged with water vapor comes in contact with cooler air, the vapor condenses and falls as rain. A great deal of attention has been given to the accu- rate estimation of the quantity of carbon dioxide in the air. The method employed for the purpose is similar to that employed for the purpose of estimating the quantity of water ; it consists in drawing a known volume of air over something which has the power to absorb the carbon dioxide, and then determining the in- crease in weight of the absorbing substance. Potassium or sodium hydroxide is well adapted to this. An appa- ratus has been constructed in which barium hydroxide, Ba(OH) 2 , is used as the absorbent. When carbon dioxide is passed through a solution of this substance insoluble barium carbonate, BaCO 3 , is thrown down according to the equation Ba(OH) 2 -}- CO 2 = BaCO 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 suimmer than in the country and in winter ; and greater in warm countries than in cold. It is as essential to the lifet 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, ozone, organic matters, etc., though there is no difficulty in de- termining 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 combination 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 withdrawn 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 rec'ent very elaborate experi- ments that the plants have the power to take up from 268 INORGANIC CHEMISTRY. the air a part of the nitrogen which they need. Whethei they take it up directly, or it is first taken up by the soil and converted into some compound which the plants 0, N=0, >0, 0=N=0, >0. N=O 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 in it the ni- H 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- 0=^=0 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 + C1 9 = NHC1, + HC1 ; NHC1 9 + C1 2 = NCI, + 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 8 . 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, but the compound need not here occupy attention. In compo- sition it corresponds to nitric oxide, NO. 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 nitrogen tri- oxide : 280, + H,0 + N.O. + 20 = 2SO,0 + HA qO ^ U f> c S .\ OH ; /OH POOH po / -L W- - - + H.O. \OH POOH 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 POJ Cl -f- 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 hydrochloric 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. 30.96). Occurrence. The name phosphorus is derived from the Greek 0c3?, light, and fiepeiv, 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 it is converted into primary calcium phosphate, CaH 4 (PO 4 ) 2 , which is soluble in water, and at the same time calcium sulphate which is difficultly soluble is formed. The reaction is represented as follows : Ca 3 (P0 4 ) 2 + 2H 2 SO 4 = CaH 4 (P0 4 ) ? + 2CaSO 4 . When any primary phosphate, MH 2 PO 4 , is heated it is converted into the corresponding metaphosphate, MPO 3 : MH 2 PO 4 = MPO 3 + H 2 O. The transformation in the case of primary calcium phos- phate is represented by this equation : CaH.(P0 4 ), = Ca(PO s ), + 2H 2 O. When calcium metaphosphate is mixed with charcoal, and the mixture distilled, two-thirds of the phosphorus is set free and passes over : 3Ca(PO 3 ) 2 + IOC = 4P + Ca 3 (PO 4 ), + 10CO. In this way one-third of the phosphorus passes back to the form of tertiary calcium phosphate. If, however, enough sand, SiO 2 , be added to form calcium silicate with the calcium, all the phosphorus is set free : 2Ca(PO 3 ) 2 + 2SiO 2 + IOC = 2CaSiO 3 + 10CO + 4P. The phosphorus passes over in the form of vapor, and is collected under water. The crude phosphorus thus 300 INORGANIC CHEMISTRY. 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. It is then cast into sticks in glass or copper tubes. In this form it comes into the market. At the time of the last report available there were manufactured in one year about 2500 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 disulp hide, 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 3 O 6 , 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 a 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 oft 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 P 2 H 4 , corresponding to hydrazine ; and a solid of the formula P 4 H 2 . 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 + PH 3 ; 2H 3 P0 2 = H 3 P0 4 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 3 = 3NaJS0 4 +Na f S.. So also sodium sulphite yields sodium sulphate and sodium sulphide : 4Na 2 SO s = 3Na 2 SO 4 + Na 2 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 calcium phosphide with water or hydrochloric acid. Assuming that calcium phosphide has the composition represented by the formula Ca 3 P 2 , the reaction with hydrochloric acid takes place accord- ing to the equation Ca 3 P 2 + 6HC1 = 3CaCl 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; PH 4 I + KOH = PH ? + KI + H 2 O. 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 + HI The products are called respectively phosphonium chlo- 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. ARSENIC, As (At. Wt. 74.9). 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 a 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 comes 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 chlorine group, with 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 a O 3 + 6H 2 = 2AsH 3 + 3H 2 O ; H 3 As0 4 + 4H 3 = AsH 3 + 4H 2 O. That involved in the second method mentioned is : As 2 Zn 3 + 3H 2 SO 4 = 2AsH 3 + 3ZnSO 4 . N It is a colorless gas with a peculiar and very unpleas- ant odor. 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. 119.6). \ 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 combinatior. 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 O 4 + 40 = 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 & . 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 = 3S0 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. heat. 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- r " 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 npt 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. 207.3). 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 ; Rose'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, BiCl 2 . 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 + 181 = Te OH + 4HC1 ' 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, + 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=Px 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 : and when it is brought together with a polysulphide, as Na 2 S 5 , 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 has 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 5 = PC1 3 + Cl a . PHOSPHORUS PENTACHLORIDE. 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 TTTTO ( OIL PC1 5 + *"JX = P 1 OH + 2HC1. (Cl, 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 6 + 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, and another, Sb 3 O 5 , and, finally, two oxides of bismuth, Bi 2 O 2 and Bi 2 O 5 . Phosphorus, further, forms the oxide P 2 O 3 . The table below con- tains the formulas of the above-mentioned compounds systematically arranged : Bi 2 2 P a O 3 As a O 3 Sb a O 3 Bi 2 O 3 Sb a o 4 P 2 5 As 2 O 6 Sb a 6 BiA 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 : Normal phosphoric acid Orthophosphoric acid As(OH) s As )(OH), + H '; Normal arsenic acid Orthoarsenic acid (321) 332 INORGANIC CHEMISTRY. Sb(OH) 5 Sb { ( Normal antiinonic acid Orthoantimonic acid (OH) s 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 Bi |oH + Normal bismutm'c acid Bismuthic acid P(OH) 5 = P j OH + 2H2 ; Metaphosphoric acid As(OH) & = As | Q 2 H + 2H 2 O ; Metarsenic acid Sb(OH) 6 = Sb | Q 2 H + 2H 3 O. 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 + H,0; COMPOUNDS OF THE PHOSPHORUS GROUP. 323 or 2H 3 P0 4 = H 4 P,0 7 + H 2 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), 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 P^H,OP^H, (H (H Unknown Hypophosphorous Phosphorous Phosphoric acid acid acid 324 INOEGANIC CHEMISTRY. The oxidation of hydrogen sulphide takes place sim ilarly, as has been shown : {g. 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 6 gives POC1 3 and P 2 O 5 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 : (01 (01 (OH ?-{ Cl , P^ OH, Pi OH. (OH (OH (OH {01 /QTTX 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 3 and P 2 S 6 , and which are, there- fore, analogous to the two oxides of phosphorus, P 2 O 3 P 2 O 6 . 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 a S 3 + 6H a O = 2H 3 PO 3 + 3H a S ; and the pentasulphide is converted by water into phos- phoric acid : p a S 6 + 8H 3 = 2H 3 PO 4 + 5H a S. Arsenic forms with sulphur several compounds, the principal of which are the disidphide, As a S 2 , the trisul- phide, As a S 3 , and the pentasulphide, As 2 S 5 . The principal sulphides of antimony are those of the formulas Sb 2 S, and Sb 2 S 6 , and of bismuth those of the formulas Bi 2 S, and Bi a 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 - P 2 S 3 As 2 S, Sb a S 8 Bi a S 3 P 2 S 6 As a 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 : and starting with arsenic acid, the following possibilities suggest themselves : (OH (OH (OH (OH (SH ^ OH , SAs^ OH , SAs^ OH , SAs^ SH, SAs^ SH . (OH (OH (SH (SH (SH 326 INORGANIC CHEMISTRY. While none of these compounds are 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 {S crrr, corresponding to the oxygen compound As j QTT , which in turn is derived from arsen- ious acid by loss of one molecule of water. ( OH (O As |oI = A So, too, we have : 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 6 , also known as phosphoric anhydride, and when this is treated with water it is converted into phosphoric acid : PA + 3H,0 = 2H,PO,. 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 = 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 a TeO 3 + 2HNO 3 = 2KN0 3 + H 2 Te0 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 PO 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 3 (P0 4 ) 2 + 2H 2 S0 4 = 2CaS0 4 + CaH 4 (PO 4 ) 2 . The phosphate thus formed is, as was seen in studying the method of extracting phosphorus from bone-ash, soluble in water, and the calcium is not easily precipitated from it. By evaporation and addition of sufficient sul- phuric acid and alcohol the precipitation can be effected, and a solution of phosphoric acid thus obtained. This acid is not pure, as there are substances in bone-ash which are not removed by the method described. 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 -j 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 4 . This salt has the formula OPx ONa , (ONH. and is called ammonium sodium phosphate. Another phosphate commonly met with is ammonium magnesium (ONH. phosphate, OP-( O - > 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 a P(ONa) + H a O. (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 : OP OP ONa ONa ONa OP ONa o + H.O. OP \ ONa (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 )PO 4 = H 2 NaPO 4 + NH 8 ; 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 4 )MgPO 4 = HMgPO 4 + NH 3 ; 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) 8 , 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, HXa 2 PO 4 , are heated. Metaphosphoric Acid, HPO 3 . This acid is formed by dissolving phosphorus pentoxide, P 3 O 6 , in cold water : It is also formed by heating phosphoric acid to 400 : H 3 PO 4 = HPO 3 + H 2 0. 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- phoricum 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 S , . . . . . Metaphosphoric acid. HXO 3 , . . . , . Nitric acid. HC10 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 : HP0 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 O = 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 phosphorus 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 P0 4 + PH 3 . HYPOPHOSPHORIC AND HTPOPHOSPHOROUS 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 PO 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 2 PO 3 , 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. It has been sug- gested that this acid has the constitution represented by OP(OH), the formula I . There is, however, no experi- mental evidence for or against this view. 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 S0 4 + H 3 P0 2 = 2H 2 S0 3 + H 3 PO 4 ; H 2 S0 3 + H 3 P0 2 = S + H 2 O + H 3 P0 4 . When heated, also, it forms phosphoric acid and phos- phine just as phosphorous acid does : 2H 3 PO 2 = H 3 PO 4 + 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 for the purpose of abstracting the elements of water from organic com- pounds, or as a dehydrating agent, as a substance that acts in this way is called. Phosphorus Trioxide, or Phosphorous Anhydride, P 2 O 3 , is formed by burning phosphorus in such a way that the air does not have free access to it. This can be accom- plished by putting a piece of phosphorus in a glass tube CONSTITUTION OF THE ACIDS OF PHOSPHORUS. 335 drawn out to a fine opening and, by means of a tube attached to the other end, drawing air over the phos- phorus and warming it gently. In this way not enough air can get access to the phosphorus to convert it into the pentoxide. The trioxide 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. Like the pentoxide it acts readily upon water, forming with it phosphorous acid : P,0 3 + 3H 2 O = 2H 3 PO 3 . 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- /Q-K tion is represented by the formula O=P^-O-H , which MX-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 M)H ( TT OP j /QJJ\ - 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 the formula O=P~O-H. \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 + O = POC1 3 ; PC1 5 + H 2 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 3 + 3H 2 + 20 = 2H 3 As0 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, HAs0 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 ARSENIOVS 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 AsO 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 i 2H 3 AsO 3 S + 6HC1 = 2AsCl 3 + 6H Q O + 2S ; I 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, but salts related to it are obtained by treating arsenic trioxide with bases. Thus, when it is treated with potassium hydroxide the salt KAsO a , or potassium meta-arsenite, is formed : As a O 3 + 2KOH = 2KAsO a + H a 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 roasting natural arsenides, particularly arsenical 'py- rites, FeAsS. The products of roasting this compound are ferric oxide, Fe 2 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 TBIOXIDE. 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 9 O, + 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 3 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 3 + 2KOH = 2KAs0 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 , s As-} SH = As] Q^ + HJS; (SH just as the acid HAsQ a is derived from the normal oxy- (OH gen acid As < OH , by loss of one molecule of water : (OH OH = As] XTT 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 a + 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 nitrite, when nitrogen trioxide is set free : KN0 2 + HC1 = HNO, + KC1 ; 2HN0 2 = N 2 O 3 + 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 nitrogen trioxide escapes as a gas. Besides salts of the acids H 3 AsS 3 and HAsS 2 , there are others derived from the more complex acid H 4 As 2 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 a S(SNH 4 ) 4 . Arsenic Pentasulphide, As 2 S 5 , 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 the reaction takes place thus : The ammonium salt formed in this way is, however, de- composed thus : (NH 4 ),As 2 S 7 = (NH.XAsS, + (NH,)AsS 3 . 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 3 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 r 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 3 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 s , 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 2 (S0 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 . ( 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 : 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 < 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 j Q-TT- . 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- ' 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 oxychlorides. 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 SbO 4 = Sb.O. + 3H 2 O ; Sb A = Sb 2 o 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) 3 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 ihe 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 crystals 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 olack or gray. The sulphide dissolves in solutions of metallic sulphides, forming salts of sulphantimonious acid, either SbS.SH or Sb(SH) 8 . Antimony Pentasulphide, Sb 2 S 5 , is formed by passing hydrogen sulphide into a solution of antimonic 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 6 + 8H S ; 2Na,SbS 4 + 6HC1 = 6NaCl + Sb 2 S B + 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 6 + 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 -I /QJJ\ and Sb j /-QJJ\ 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 O 3 S | QTT , that of selenious acid is represented by the formula OSe j QJJ . Oxychlorid.es 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 = SbOCl + 2HC1. If, however, hot water is used, the composition of the product approximates to that represented by the formula Sb 4 O 6 Cl 2 . This complex mixture of oxychlorides is known as the "Potvder of Algaroth" It may be regarded as derived from the simple oxychloride by loss of anti- mony trichloride, thus : 5SbOCl = Sb 4 O 6 Cl a + 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,), + 3KOH = Bi(OH) 3 + 3KNO, ; Bi(OH) 3 = BiO.OH + H 2 ; 2Bi(OH) 3 = Bi 2 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 ^ I \ 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 a . 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 a , and this with the potassium hydroxide breaks down, yielding the dioxide : 2BiCl 3 + SnCl, = 2BiCl 3 + SnCl 4 ; 2BiCl a + 4KOH = Bi,0 9 + 4KC1 + 2H a 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 oxychloride, is thrown down as a white powder when a solution con- taining bismuth chloride is treated with water : BiCl 3 + H a 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. 51.1). This element occurs in nature in the form of vanadates or salts of yanadic acid, H 3 YO 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 Y 2 O 5 . With chlorine, vanadium forms the compounds YC1 2 , YC1 3 , and YC1 4 ; with oxygen, the compounds Y 2 O, Y 2 O 2 , Y 2 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,(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. 182). 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 6 , and TaI 5 . 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 3> 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 3 O,. With oxygen it forms Ta a O 4 and Ta a O 5 . The latter forms the tantalates with bases. When tantalum pentachloride is decomposed with water it forms the acid H 4 Ta 3 O, or pyrotantalic acid : 2TaCl 5 + 7H a O == Ta 2 3 (OH) 4 + 10HCL The tantalates are derived from the meta-acid HTaO 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 19 + 5H a 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 5 , and a bromide and fluoride corresponding to the latter chlo- ride. The fluoride readily forms jluocolumbates, 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 is an extremely rare element. In some of its compounds it shows a resemblance to the members of this group. It forms, for example, an oxide of the formula Di 2 O 5 . On the other hand, it seems to be more closely related to cerium and lanthanum, which are also very rare elements, occurring associated with didymium. It will be further treated of in connection with lanthanum and cerium. BORON, B (At. Wt. 10.9). 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 the sodium salt or borax. From borax and the other borates the acid can easily be obtained, as will be shown pres- ently. This has the composition B(OH) 3 . When heated it loses all its hydrogen in the form of water, and boric oxide or boron trioxide, B 2 O 3 , is left : 2B(OH) 3 = B.O. + 3H,0. By heating the oxide with potassium the element boron- is obtained in amorphous form. By melting aluminium and boron trioxide together at a high temperature, the latter is reduced, and the boron thus formed is dis- solved in the molten aluminium, from which, on cooling, it is deposited in crystals. One of the chief difficulties en- countered 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. Boron, by which is meant the amorphous form, 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 carbonates of potassium or sodium, it forms borates of the corresponding metals. Boron Trichloride, BC1 3 . 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 2 O 3 + 3C + 6C1 3 = 4BC1 3 + 3CO 2 . 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 O = 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 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 = CaSO 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 = B(OH) 3 + 3HBF 4 . The first action which we should expect is the formation of normal boric acid, thus : BF 3 + 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 tlie 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 fumaroles 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 a B 4 O 7 + H 2 SO 4 + 5H 2 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 2 O 3 . These changes are represented in the equations follow- ing: (OH B^ OH (OH ,OH (OH , Q {OH = BJ +H a O; / ATT ( Ul y M)H B XDH / Bf \ B(OH VH \A | CTT f\ OH - >O + 5M (J - BfOH \OH w _ ,OH B - \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, + 5CuO = Na,Cu 5 (BO 8 ) 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 is utilized for the purpose of determining their nature 356 INORGANIC 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 it. 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 - - . This is the salt represented -- , 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 a O = 2B(OH) 3 + 2NH,. ft CHAPTER XIX. CARBON AND ITS SIMPLER COMPOUNDS WITH HYDRO- GEN 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, B(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 believed to be due to the decomposition of submarine animals continued through ages, consists of a large number of compounds which contain only carbon 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 cel- lulose. Most products of animal life contain 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 moun- tain ranges. Limestone, marble, and chalk are varieties of calcium carbonate. Uncombined, the element occurs pure in two very dif- DIAMOND GRAPHITE. 359 fereut forms in nature : (1) As diamond ; and (2) as graphite, or plumbago. Before considering the evidence which leads to the conclusion that diamond and graphite are only modifica- tions of the same element, and that while closely related to each other they are also equally closely related to charcoal, it will be best to study separately the proper- ties of each of these three substances. 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. 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. 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 ivood spirits or methyl alcohol, pyroligneous or acetic acid, acetone, etc. The carbon, however, is foj* 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, and which make 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 iavorable, 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. 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-black, 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 w r ater, 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, as 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 oxperiment is known, the weight of the 366 INORGANIC 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 ease 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 difference is not known, but, as in the cases of phosphorus and sulphur, it will probably be found to be in the different 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. They are tasteless, inodorous, and infusible. When heated without access of air, they remain un- changed, unless the temperature is very high, when the diamond swells up and is converted into a mass resembling coke a change which is connected with a rearrangement of the particles in an irregular 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 oon 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 + C0 2 = CaC0 3 + H 2 O; Calcium Carbon Calcium hydroxide dioxide carbonate Ba(OH) 2 + C0 2 = BaCO, +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 2 . 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 the 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 2As 2 O 3 + 30 = 4As + 3CO a . 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, and it has been shown that these relations are simple. But carbon is not known in the form of gas, and cannot even be melted. It is therefore, of course, impossible to de- termine 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. Di- rect combination of carbon and hydrogen cannot easily be effected. When the carbon pencils connected with a powerful battery, as in the production of the electric arc- light, are surrounded by an atmosphere of hydrogen the two elements combine to some extent to form the com- pound acetylene, C 2 H 2 . 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. The compounds which make up petroleum are hydrocarbons NUMBER OF HYDROCARBONS. 369 -which have probably been formed by decomposition of organic matter without free access of air. Finally, when wood or coal is heated, hydrocarbons are given off, and compounds 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. Its formation is probably due to the decomposition of submarine animals. In the earth it contains both gases and liquids. When it is brought into the air the gases and the liquids which are volatile 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 R , 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 1Q , 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 a H 6 , Ethane; C 3 H 8 , Propane ; C 4 H 10 , Butane ; C 6 H 12 , Pentane ; C 6 H 14 , Hexane ; C,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 t and such a series is known as an homologous series t Carbon is distinguished from all other elements by its power to form homologous series. Other elements form similar series/ jbut 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 HC1O 2 H 2 SO 3 HNO 2 HC1O 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, i 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 6 + 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 H unite, the resulting compound will have the structure H H represented by the formula H-C-C-H. In a similar 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 Seriasof 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 a H 4 , examples of which are Ethylene, C,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 a H a ; Allylene, C,H 4 . 372 INORGANIC CHEMISTRY. Another series begins with benzene, C 6 H 6 . Some of the members of this series are Benzene, C 6 H 6 ; Toluene, C 7 H 8 ; Xylene, C 8 H 10 . 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 ll +1=1 CH, Cl CH 2 C1 Similar reasons have led to the conclusion that in acetylene, C 2 H 2 , the carbon atoms are held together in MARSH-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-chlor-ethane : HC=CH + 401 = C1 2 HC-CHC1 2 (C 2 H 2 C1 4 ) ; HfeCH + 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. The sulphur is extracted from the compounds, and the carbon and hydrogen combine, as represented in the equation CS 2 + 2H 2 S + 8Cu = CH 4 + 4Cu 2 S. The gas is met with in coal-mines, and is known to the miners as fire-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, Oleflant Gas, C 2 H 4 . This Iiydrocarbon is formed by heating a mixture of ordinary alcohol and concentrated sulphuric acid. The reaction is represented thus: C 2 H 6 = H 2 + C 2 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. 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 replaces the hydrogen, atom by atom. Thus, when it acts upon marsh gas, the following reac- tions take place : CH 4 + 01, = CH 3 C1 + HC1 ; CH 3 C1 + C1 2 = CH 2 C1 2 + HC1 ; CH 2 C1 2 + C1 2 = CHC1 3 + HC1 ; CHC1 3 + C1 2 = CC1 4 + HC1. The four products are known respectively as morvo-cTdor- methane, di-cMor-methane, tri-cJdor-methane, and tetra-chlor- methane, or carbon tetracMoride. 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) the hydrogen atoms appear to be transformed into hydroxyl 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 : S H H OH OH Hydrogen Unknown Sulphurous Sulphuric Sulphide acid acid (376) RELATIONS BETWEEN COMPOUNDS OF CAEBON. 377 OH (OH (OH (OH N^OH, ON \ OH. (OH (OH Hydroxyl- Yields hypo- Yields nitrous Yields nitric amine nitrous acid acid acid The last three products, if formed, break down, losing water, and forming respectively hjponitrous, 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 (OH (OH (OH H, ON-{ H , ON-} OH, ON-{ OH. H (H (H (OH If nitrous acid were formed by the breaking down of (OH ihe compound ON-< OH, its structure would probably ( H (O 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. (OH (OH OP^OH, OP^OH. (H Hypophosphor- ous acid Phosphorous acid (OH 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 4 OH 5 But the tendency of the compounds containing twcr 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 : r OH C 10H [OH OH OH OH OH (H CJH + c |L H,0; = CJg + 2H.O. The products actually obtained, therefore, are as follow : RELATIONS BETWEEN COMPOUNDS OF CARBON. 379 H fH ,-rr ,TT E , CJ H > CJH, CJO , CJ. H [OH 1 ( H Methane Methyl Formic Formic Carbon alcohol aldehyde acid 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 vhich 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 < 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 ( TT H i ' C < OH or O=C is the simplest representative of a < 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 a , 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 nitrogen trioxide from a nitrite, and sulphur dioxide from a sulphite. In all these cases the products are volatile, and therefore, when a non- volatile acid is added to the salts, decomposition takes place. Nitrites and siilphites do not yield the corresponding acids, but these break down into water and the anhy- drides : j 2NaN0 2 + H 2 S0 4 = Na 2 SO 4 + 2HNO 2 ; |2HN0 2 =N f O, +H 2 0. ( Na 2 S0 8 + H 2 S0 4 = Na 2 S0 4 + H 2 SO 8 ; JH 2 S0 3 = S0 2 +H a O. j Na 2 C0 3 + H n SO 4 = Na 2 SO 4 + H 2 CO 3 ; |H 2 C0 3 =C0 2 +H 2 0. 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 3 + 2HC1 = 2NaCl + CO 2 + H 2 O ; 382 INORGANIC CHEMISTRY. that between nitric acid and sodium carbonate in this way : Na 2 CO 3 + 2HN0 3 = 2NaNO, + CO 2 + 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 : CaC0 3 + 2HC1 = CaCl 2 + CO 2 + 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 INORGANIC 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 have been shown to be due, not to the carbon dioxide, but 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 as 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 'tware 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 salt, 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 3 O. 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, = CaC0 3 + H 2 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 = CaCO 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 + S0 2 = K 2 SO 3 + H 2 O ; Ca(OH) 2 + S0 2 = 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,, 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 2 CO 3 , 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 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 : nn /OH , (HO)Cu(OH) _ np .OCuOH U0 < OH + (HO)Cu(OH) - < OCuOH ~ M * U ' 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 : ,X)MgOH CJ >0 2 Mg co< ^OMgOH 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 : Such a salt is formed by loss of water from the primary sodium salt : 2HNaC0 3 = Na 2 C 2 O 6 + H 2 O. 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, + 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 ^oncentraied sulphuric acid. The change which takes place is represented thus : 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 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. With proper precautions, however, there seems to be no good reason why it should not be used, although it is 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 cMoride, 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,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 - O . The relations here are similar to (OH those met with in phosphorous and sulphurous acids, 392 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 : H op ( OH OH u \ OH Formic acid Carbonic acid Sulphurous acid Sulphuric acid (H (OH OP I OH OP 4 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 2 = 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 co 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 5 , 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. IHuminating Gas, Coal-gas. Illuminating gas is gen- erally 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 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. When we say that the illuminat- ing power of a gas is fourteen candles, we mean 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 we 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. ii. 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 i eduction. The principal parts of d 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 exhaustively, 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 the aid of the oxygen of the air admitted to the interior of the flame. COMPOUNDS OF CARBON WITH NITROGEN AND WITH SULPHUR. Cyanogen, CaNa. 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 knowTi as potassium ferro-cyanide, or yettoiv prussiate of potash, K 4 Fe(CN) 6 -f- 3H 2 O, is formed. When this is simply heated it is decomposed, yielding potassium cyanide : K 4 Fe(CN) 6 = 4KCN + FeC a + N 2 . It is an easy matter to make the mercury salt, Hg(CN),, from the potassium salt. By heating mercuric cyanide it breaks up, yielding metallic mercury and cyanogen gas: Hg(CN) 2 = just as mercuric oxide yields mercury and oxygen when heated : HgO -- Hg + O. Cyanogen (from Kvaros, 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 ON , ._ CO,H bos. ^ +4H.O = ;_ + 2NH, 2 J TT The compound T 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, Prussic Acid, HON. 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 CL 'The product NH 4 CN is ammonium cyanide. HYDROCYANIC ACID. 403 Hydrocyanic acid is a volatile liquid, boiling at 26.5, and solidifying at 15. 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_ 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 oH OH the formula C -< -- 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 : ! HOH [81 + HOH=C * + 3HCl; C-l H |H OH OH 404 INORGANIC CHEMISTRY. , 1 H C 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 cannot be separated from its salts, as it breaks down at once into carbon dioxide and ammonia in presence of water : CONH + H 2 O = NH 3 + C0 2 . 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 DISULPBIDE. 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. 5. 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 2 + 3O 2 = CO 2 + 2SO 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 a + 2H 2 = C0 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 hy drosulphide 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 a 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 t CS, + 3H 2 O = Na,CO 3 + 3H 3 S. When a sulphocarbonate is treated with cold dilute Irydrochloric 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 3 CS 3 = CS 2 + 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 g , C S j OH , 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, + CS 2 + 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 che cyanogen group, plays the same part that the atom chlorine plays in the chlorides : H(CN) HC1 K(CN) KC1 Hg(CN), 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=NT. 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. The only member of the group which combines with hydrogen is silicon, and this forms but one compound with it, corre- sponding in composition to marsh-gas. This is silicon hydride, SiH 4 . 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 . Two of them, further, form compounds analogous to hexa- chlor-ethane, C a Cl e , and to tetra-chlor-ethylene, C 2 C1 4 : Si 2 Cl 6 Si 2 Cl 4 Ti 2 Cl. Ti 2 Cl 4 All the elements of the group form oxygen compounds analogous to carbon dioxide. They are : SiO 2 , TiO 2 , Zr0 2 , Ce0 2 , ThO 2 . 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 a Si0 3 , M a TiO 3 , M a ZrO 3 . 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). 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 element in nature. There are extensive mountain-ranges consisting almost entirely of the dioxide, SiO 2 , in the form known as quartz or quartzite. Other ranges are made up of silicates, which are compounds formed by the com- bination of silicon dioxide and bases. The clay of the valleys 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 2 , which is most easily obtained in the form of sand, is decomposed by heating it with potassium or magnesium, and sili- con is thus set free. When magnesium is used the SILICON. 411 action is violent, and besides the silicon there is formed a compound of silicon and magnesium. It is also formed by the action of potassium on silicon tetrachloride. The best way to make it is by heating together potassium fluosilicate, K 3 SiF,, sodium, and zinc : K a SiF 6 + 4Na = 4NaF + 2KF + 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 iluo- 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 is con- verted into crystallized silicon by continued heating 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 aqiia regia. It dissolves, however, in potassium hydroxide, forming potassium silicate, in this case acting like an acid-forming element : Si + 2KOH + H 2 == K 2 SiO 3 + 2H,. 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. 48). 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, FeTiO 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 TETRACHLORIDE. 413 Zirconium, Zr (At. Wt. 90.4). 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. 232). This element occurs prin- cipally in the mineral thorite, which is essentially a sili- cate 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 GROUP WITH THOSE OF THE CHLORINE 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, + 2C + 2C1 2 = 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 iogether 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,CO S . Silicon Hexachloride, Si 2 Cl 6 , is formed when silicon tetrachloride is heated with silicon : 3SiCl 4 + Si = 2Si 2 Cl 6 . 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,0,(OH), . 81,01, + 6H,O = Si,(OH), + 6HC1. SJ,(OH). = SiA(OH), + 2H.O. 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, + H 2 S0 4 = CaS0 4 + 2HF ; Si0 2 + 4HF = 2H 2 + SiF 4 . FLUOSILICIG ACID. 415 The tetrafluoride escapes as a colorless gas, which forms thick clouds in moist air on account of the action of water upon if. 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 = 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 8 . 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 Fluosilicic 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 re- placement of the three oxygen atoms by six fluorine 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 : or , 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 why 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, + 4HCL Titanium also forms with chlorine the compounds Ti a Cl 6 . and Ti a Cl 4 . THORIUM TETRAFLUORIDE. 417 Titanium Tetrafluoride, 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 3 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, : ZrCl 4 + H 2 = ZrOCl, + 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, + 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 3 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 OF 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 : K 2 C0 3 + Si0 2 = K 2 Si0 3 + CO, ; 2KOH + Si0 3 = K 2 Si0 3 + H 2 0. 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 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 with 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 a Si0 3 : Si(OH) 4 = OSi(OH) 2 + H 2 0. 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), = SiO, + 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 dicdyser, 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 Z> 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 : OH Its composition is, therefore, H 2 Si 2 O 6 , which may be TRISILICIC ACIDS. written O 3 Si a (OH) a . 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 a (OH) 6 + H a O. The structure of this acid is expressed by the formula The well-known mineral serpentine is appa- rently the magnesium salt of this acid. It is represented by the formula Mg 3 Si 2 O 7 . 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 8 Si 3 O 10 : 3Si(OH) 4 = H,Si,0 10 + 2H,0. 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 I (OH), Si^r ~lg si (OH)s II 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 the for- mula OK Si Si Si O O o o (XA1 424 INORGANIC CHEMISTRY. Titanium Dioxide, TiO 2 . 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 4 CO 3 + Si0 2 = K 2 Si0 3 + C0 2 ; K 2 C0 3 + Ti0 3 = K a Ti0 3 + C0 2 . 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 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 ) 2 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 CO 3 + ZrO 2 = K 2 ZrO 3 + CO 2 . The sodium salt of normal zirconic acid, Na 4 ZrO 4 , has also been obtained. The dioxide forms salts with acids, among which two of the sulphates are of special interest. One has the composition ZrSO 5 , and the other Zr 3 (SO 6 ) 2 . The former is to be regarded as the salt of the acid OS(OH) 4 , formed FAMILY IV, GROUP B. 425 by replacing the four hydrogen atoms by one atom of zirconium ; the other is the salt of normal sulphuric acid, S(OH) 6 , formed by replacing all the hydrogen by zirconium. 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. 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 first two form chlorides of the formulas GeCl 4 and SnCl 4 , while lead forms only the lower chloride, PbCl 2 . 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 law 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 quantities of any ele- ment which combine with a fixed quantity of another or others bear simple relations to one another. To account for the laws of definite and multiple pro- portions the Atomic Theory has been 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 a con- clusion 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 INOEGANIG CHEMISTRY. the extent to which the compound 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 how 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. Valence, 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 the 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 peri- CHEMICAL ACTION. 429 odic function of these tveights. This is the great periodic law of chemistry, which makes a systematic classification of the elements according to their atomic weights and their properties possible, and which is so full of sugges- tion 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 CIA H,S S0 3 (SA) H 3 P PA H 4 Si Si0 2 (SiA) 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 3 PO 4 , and H 4 Si0 4 . Here the remarkable fact is ob- 430 INORGANIC CHEMISTRY. served, that the number of hydrogen atoms in each of 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) a , P(OH) 5> 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 : 2KC10 3 = KC1 + KC10 4 + 2 ; 4K 2 SO 3 = K 2 S + 3K 2 SO 4 ; 4H 3 PO 3 = PH 8 + 3H 3 PO 4 . The highest forms are therefore evidently most stable. Turning to the compounds which the elements of Families IV, Y, VI, and VII 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 IV, 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 : Cl a + H 2 = Cl + Cl + H + H ; d -f 01 + 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- l sented by the formula N-j 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 + HG1 = NH 3 + H + 01; NH 3 + 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 ; that of water into hydrogen and oxygen by the electric current ; 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 2HgO = Hg + Hg + O + O ; 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 + Cl a . Similar remarks 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 nitrogen trioxide 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 come under our notice are : The action of sulphuric acid upon potas- sium 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 SO 4 = K 2 S0 4 + H 2 O ; CaCl 2 + H 2 SO 4 = CaSO 4 + 2HCL The Cause of Chemical Reactions. The prime cause of . chemical reactions is something which we think of as an AN IDEAL CHEMICAL EEACTION. 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, AE and CD act upon each other in the gaseous condition to form two compounds B C 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 AE 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 E. 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 of C for I?. 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 4 , 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. Considerable attention has been given to the study of mass action of late, and the result is to show that in reactions generally this kind of action comes prominently into play. The law has been established that chemical action is proportional to the active mass of each substance taking part in the change. It would appear from this that the decomposition of two compounds 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 reacting 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 : Bad, + H a 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 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, arid most of them are undoubtedly not true. 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 has 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,, Cl,] = 2[H, Cl] - [H, H] - [Cl, Cl] = 44,000 c. [H,, Br,] = 2[H, Br] - [H, H] - [Br, Br] = 16,880 c. [H,, I,] = 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- 4:4:0 INORGANIC CHEMISTRY. 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 nitric acid acts upon sodium sulphate in solution, several changes are possible, as represented in the equations (1) Na 2 SO 4 + HN0 3 = NaHSO 4 + NaNO 3 ; (2) Na 2 SO 4 + 2HNO 3 = H 2 SO 4 + 2NaNO 9 ; (3) 2Na a S0 4 + 4HN0 3 = Na,SO 4 + 2NaNO 3 + H 2 SO 4 + 2HNO 8 . 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 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 soda combines with the nitric acid and one-third with 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 are as given in the following table : Acids. Avidity. Nitric acid, 100 Hydrochloric acid, 100 Hydrobromic acid, 89 Hydriodic acid, 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 ^ molecule of orthophosphoric acid. 442 INORGANIC CHEMISTRJ. 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 from other points of view. 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 affinity of the acids. (4) The simplest method of all is the electrical. This consists in determining the conducting power of solutions of different dilutions. In this way figures are obtained which bear to one another the same relations 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, at the same time prevents 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 always 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 would 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 would 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 would 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 would be decom- posed in a given mass at a given temperature and pres- sure would 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 would not exist, and total decomposition could be effected at a lower temperature than that required for total decomposition in the process of dissociation proper. Dissociation also takes place in some solutions as the temperature is raised, and the phenomenon is in princi- ple the same as the dissociation of gases, and it is to be explained in the same way that is, by assuming that the motion of the atoms and the molecules in different parts of the solution is different, and that there is constant de- composition and re-formation of the compound. Electrolysis. One of the first chemical phenomena which we studied was that of the action of the electric current upon water, causing its decomposition into hy- drogen and oxygen. Phenomena of this kind are quite common. The decomposition of a chemical compound by an electric current is called electrolysis. In the de- composition of water the hydrogen appears at the nega- tive pole of the battery and the oxygen at the positive pole. So, too, when hydrochloric acid is decomposed in this way, the hydrogen is evolved from the negative pole and the chlorine from the positive pole. In a similar way, it has been found that when salts are decomposed by electrolysis the metal is deposited at the negative pole and the acid or its decomposition-products at 446 INORGANIC CHEMISTRY. the positive pole. There is undoubtedly some close con- nection between chemical phenomena and electrical, but what the connection is is not known. It was at one time held that those elements which in electrolysis appear at the negative pole do so because they are charged with positive electricity, and that those which appear at the positive pole do so because they are charged with nega- tive electricity. The elements were therefore divided into the electro-positive and the electro-negative. Those elements which we call acid-forming are electro-negative in this sense, while hydrogen and the base-forming ele- ments are electro-positive. The electrolysis of chemical compounds is not generally a simple decomposition into two compounds. Thus, when copper sulphate, CuSO 4 , is decomposed, the copper is deposited at the negative pole ; but no such compound as SO 4 is formed at the positive pole. This, if formed, breaks down into oxygen and sulphur trioxide, and the latter with water forms sul- phuric acid. Both oxygen and sulphuric acid are there- fore liberated at the positive pole. The changes in- volved may be represented thus : Cu +SO 4 ; S0 4 = S0 3 +0; S0 3 + H 2 = H 2 S0 4 . 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, has been re- ferred to. 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 SPECIFIC HEAT AND ATOMIC WEIGHTS. 447 same number of degrees. The quantity of heat required to raise the temperature of a certain weight of a sub- stance 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 sub- stance. Thus, from what was said above, the specific heat of mercury is -fy, or, in decimals, 0.03332. 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 : Element. Specific Heat. Atomic Weight. Silver, 0.0570 107.66 Zinc, 0.0955 65.1 Cadimum, .... 0.0567 111.7 Copper, 0.0952 63.18 Tin, 0.0562 117.4 An examination of this table will show that the atomic weights are inversely proportional to the specific heats. We have 107.66: 65.1 111.7 : 63.18 107.66 : 117.4 0.0955 0.0952 0.0562 0.0570 ; 0.0567 ; 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 can occasion no sur- prise. 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 448 INORGANIC CHEMISTRY. of the atomic iveighi into the specific heat is a constant. Thus, in the above cases : 107.66 X 0.057 =6.14; 65.1 X 0.0955 =: 6.22 ; 111.7 X 0.0567 = 6.33 ; 63.18 X 0.0952 = 6.01 ; 117.4 X 0.0562 = 6.51. 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 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 DETERMINATION OF MOLECULAR WEIGHTS. 449 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 carbon, boron, and silicon form exceptions to the law of specific heats as this law has been stated above. At or- dinary 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 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 point, 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 Method 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 450 INORGANIC CHEMISTRY. 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 Raoult. 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 Eaoult, as will readily be seen. 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 (451) 452 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, 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. 453 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 H, 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 IV, 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 YI, 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 cobalt ; (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 4:54 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 + 110 = Fe 2 3 + 4SO 2 ; 2Fe(OH) 3 = Fe 2 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. 455 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 compounds of the met- als 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 the acids of nitrogen ; or the ni- trates and nitrites. e. Compounds with the acids of chlorine, bromine, and iodine ; or the chlorates, bromates, iodates, hypochlorites, etc. /. Compounds with the acids of sulphur, selenium, and tellurium ; or the sulphates, sulphites, etc. g. Compounds with carbonic acid ; or the carbonates. h. Compounds with the acids of phosphorus, arsenic, and antimony ; or the phosphates, arsenates, etc. i. Compounds with silicic acid ; or the silicates. j. Compounds with boric acid ; or the borates. Of the almost infinite number of compounds belong- ing to the classes above referred to, only a compara- tively small number will be treated of in this book. It is more important to become acquainted with the general methods of preparation and the general properties of 456 INORGANIC CHEMISTRY. these compounds than to learn details concerning many individual members of each class. Only those com- pounds will be considered which illustrate general principles, or which, owing to some application, happen to be of special interest. 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, to consider briefly 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. 457 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 +C1 2 = 2Fe + 3C1 2 = 2FeCl, ; 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 + 30 + 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 meted ivith 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 S0 4 = ZnSO 4 + H 2 . 458 INORGANIC CHEMISTRY. Other common examples are those represented in the follow equations : Fe + H 2 S0 4 = FeS0 4 + H a ; Zn+2HCl = ZnCl 2 + H 2 . 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 : MgCO a + H 5 SO, = MgSO. + H,0 + CO, ; and calcium chloride is made by dissolving calcium car- bonate in hydrochloric acid : CaCO, + 2HC1 = CaCl;+ H 2 O + CO 2 ; etc. 4. By treating a salt of an insoluble acid ivith 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. 459 the silicic acid is partly precipitated, as we have seen Na 3 Si0 3 + 2HC1 + H 2 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 : CuS0 4 + Bad, = BaSO 4 + CuCl a . 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 oxychloride : SbCl 3 + H 2 O = SbOCl + 2HC1. 460 INORGANIC CHEMISTRY. The chlorides of the most strongly marked metals, like potassium, sodium, etc., are apparently not changed 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 Ca01 2 + 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 a + 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 SO 4 = M 2 SO 4 + 2HC1 ; MC1 2 + H 2 S0 4 = MS0 4 + 2HC1 ; etc. THE SO-CALLED DOUBLE CHLORIDES. 461 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 chloro-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 + 2MCl, or M 2 PtCl 6 . Antimony chlo- ride, combines with three molecules of potassium chlo- ride forming the compound SbCl 3 + 3KC1, or K 3 SbCl t . 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..KC1, MC1..2KC1, and * There are a few exceptions to this rule, but apparently not more than two or three in several hundred cases. 462 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. K 2 MC1 4 K 2 MC1 5 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. 463 reduction. The change can generally be effected by means of nascent hydrogen : SnCl 4 + 2H = SnCl, + 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, + 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, + SnCl, = 2HgCl + SnCl 4 . In this case mercuric chloride is changed to mercurous chloride by the action of stannous chloride. The latter has such a strong affinity for chlorine 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. 464 INORGANIC CHEMISTRY. "When magnesium is burned it is converted into mag- nesium oxide : M g + O = MgO. When lead nitrate is heated it gives off oxygen and an oxide of nitrogen, and lead oxide is left behind : Pb(NO 3 ) 2 = PbO + 2N0 2 + O. When calcium carbonate is heated it yields calcium oxide and carbon dioxide : CaCO 3 = CaO + CO 3 . When aluminium hydroxide, A1(OH) 3 , 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. 465 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 = Ca(OH) 2 . 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 : MgS0 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 466 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 8 + 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. 467 chromium, and lead. The salt-like compounds thus formed ars generally quite unstable. The precipitation and subsequent solution of the hydroxides of the three metals named take place thus : A1C1, + 3NaOH = A1(OH) 3 + 3NaCl; Al(OH), + 3NaOH = Al(ONa), + 3H a O ; CrCl, + 3NaOH = Cr(OH) 8 +3NaCl; Cr(OH) 3 + 3NaOH = Cr(OKa) 3 + 3H a O ; Pb(NO,) f + 2NaOH = Pb(OH), + 2NaNO, ; Pb(OH) a + 2NaOH = Pb(ONa) f + 2H,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 : 2AgNO s + 2KOH = Ag 2 O + H,O + 2KNO, ; HgCl, + 2NaOH = HgO + H a O + 2NaCl. It is probable that the first product is the hydroxide, and that this breaks down into the oxide and water : 2AgN0 3 + 2KOH = 2AgOH + 2KNO, ; HgCl, + 2NaOH = Hg(OH) 2 + 2NaCl ; Hg(OH), = HgO + H,0. 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 : CuS0 4 + 2NaOH = Na 2 SO 4 + Cu(OH) a ; 468 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 hydrogen 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. 469 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 a + H a & It is therefore obvious that the first reaction cannot take place. If, however, a soluble sulphide, as sodium or ammon- 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 a + (NH 4 ),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 + 40 = BaS + 4CO ; and by the formation of copper sulphide by heating copper sulphate in a current of hydrogen : CuSO 4 + 4H 2 = CuS + 4H 2 O. 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 sulphur 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 470 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 = CuSO 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 + 4O = K 2 SO 4 ; K 2 S + 4S = K 2 SS 4 , or K 2 S 5 . 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 ; 2 = CuO+.H a O. SULPHO-SALTS. 471 The only hydrosulpliides 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 a S = KSH + H 3 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,S = 2K 3 AsS 4 ; SnS 2 + KJ3 = 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 hydrogen displaced by the 473 INORGANIC CHEMISTRY. metal acts upon the nitric acid itself, reducing it and forming different products, according to the circum- stances. 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(NO 3 ) a = PbO + 2N0 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 hypocMorites 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. 473 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 SO 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 = CuSO 4 4- H ; H 2 SO 4 + 2H = SO, + 2H 8 O ; or by these : Cu + H 2 SO 4 = CuO + SO, + H a O ; CuO + H 2 S0 4 = CuS0 4 + H,0. 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. 474 INORGANIC CHEMISTRY. When heated with charcoal in the reducing flame of the blow-pipe, sulphates are reduced to sulphides : K 2 SO 4 + 4C = K 2 S + 4CO, or K 3 S0 4 + 2C = K 2 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, = Na 2 SO 3 + H 2 O. All sulphites are decomposed by the common acids, sulphur dioxide being given off : Na 2 S0 3 + H 2 SO 4 = Na 2 SO 4 + H 2 + 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 SO 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 + C0 2 = K 2 CO 3 + H,O. The formation of calcium carbonate takes place in the same way, but the carbonate formed is insoluble : Ca(OH) 2 + C0 2 = CaC0 3 + H 2 O. PHOSPHATES. 475 If in either case the action is continued, tiie normal carbonate first formed is converted into acid carbonate : K a CO 3 + CO a + H,O = 2KHC0 3 ; CaC0 3 + CO, + H,0 = Ca ) All 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 the best-known illustration of this fact : 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 a 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 476 INORGANIC CHEMISTRY. sodium phosphate, HNa 2 PO 4 , lose water when heated, and yield pyrophosphates : 2HNa 2 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 2 NaPO 4 lose water when heated, and yield metaphosphates : H 2 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 8 + 4H,0. 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 2 = CaSiO 3 ; Na 2 CO 3 + SiO 2 = Na 2 Si0 3 + CO 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. 477 Many silicates, which are not acted upon by strong acids, are decomposed by fusing 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 a 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 CAESIUM 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 0, 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, 23, is very nearly half the sum of those of lithium, 7.01, and potassium, 39.03. We have 7.01 39.03 So, also, that of rubidium, 85.2, is approximately half the sum of those of potassium, 39.03, and csesium, 132.7. 39.03 + 132.7 ^ = 85.87. (478) POTASSIUM. 479 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, . . . 7.01 0.594 180. Sodium, ... 23. 0.972 95.6 Potassium, . . 39.03 0.865 62.5 Eubidium, . . 85.2 1.52 38.5 Caesium, . . . 132.7 ? ? 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. 39.03). 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 480 INORGANIC CHEMISTRY. heating to a high temperature a mixture of potassium carbonate and carbon : K 2 C0 3 + 20 = 2K + 300. 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 iron retort, and at the end of the operation the retort 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. 481 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 j Br^mid!' KBr ' Of these salts the onl J L Iodide, ia 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 carnattite. 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 = 5KI + KIO 3 + 3H 2 O. By evaporating the water and heating the residue with very finely powdered charcoal, the iodate is decomposed 482 INORGANIC CHEMISTRY. into iodide and oxygen. The reduction of the iodate takes place in accordance with the equation : 2KI0 3 + 3C = 2KI + 3C0 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 more iodine is added as has already been used in its preparation, f erroso-f erric 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 8 + 4K 2 C0 3 + 4H 2 = SKI + Fe 3 (OH) 8 + 4CO 2 . 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 1J 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. 483 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 in 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. 461 and 462). Those containing fluorine have been studied most fully. Good examples are the fol- lowing : fF F As , or AsF 6 .KF ; F 2 K F ,orSbF 6 .2KF; F.K F F F , or AsF 6 .2KF ; KK F,K ,orBF 3 .KF; Si 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. 484 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 a C0 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. 485 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 out of the laboratory is for the pur- pose of 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,O + H 2 . With water it forms the hydroxide, with a marked evo- lution of heat : K,O + H,0 = 2KOH. Potassium also forms other oxides of which the peroxide 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 486 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 2 0. 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 = KSH + KOH. This reaction is analogous to that of water on the oxide . K 2 O + H 2 O = 2KOH. POLTSULPHIDES OF POTASSIUM. 487 The polysulphides 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 a 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 : 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 + 4O = 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\ a tetrathiosulphuric acid, H 2 SS 4 (S 2 S< C vrTJ, or hydrogen pentasulphide, H 2 S.. 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 a CO, + 8S = 2K 2 S 3 + K,S 2 O 3 + 3CO 2 . 488 INORGANIC CHEMISTRY. If the mass is ignited, of course the thiosulphate is de- composed, forming the sulphate and pentasulphide : 4K 2 S 2 3 = 3K 2 S0 4 -f K 2 S 6 . 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 S 5 = 2K 3 AsS 4 + 3H 2 S ; 6KSH + As 2 S 3 + 2S = 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 -|- 4JH 2 O, said potas- sium sulpharsenite, KAsS 2 + 2^H 2 O, are also easily made. The latter is plainly the analogue of the metarsenite, KAsO 2 metarsenious acid being derived from normal arsenious acid by elimination of one molecule of water : H 3 As0 3 = HAsO 2 + H 2 0. 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 489 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, + KC1 = KNO 3 + NaCl. 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 fact 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 = CO, + CO + N, + K a 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 : 2KX0 3 + 3C + S = 3CO, + N, + K 2 S. 490 INORGANIC CHEMISTRY. For this reaction the constituents should be mixed in the proportions : Saltpeter, ..... .'.,'... . 74.83 Charcoal, . . . V . . . V, , ". 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 -f CaCl 2 . POTASSIUM PERCHLORATE. 491 Potassium chlorate crystallizes in lustrous crystals of the monocliriic 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. 492 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 5 , 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.1,0. + 5H,0. 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 w r ith 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. 493 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 : Fe(CN) 2 + K 2 CO 3 = FeCO 3 + 2KCN. But the ferrous carbonate breaks down under the influ- ence of heat into ferrous oxide and carbon dioxide FeC0 3 = FeO + CO, ; 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) fl + K 2 CO S = 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 has a strong affinity for oxygen when in the molten condition, 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_j metallic elements, forming the double cyanides. These hare generally a composition analogous to that of the double chlorides and of similar compounds. Thus, silver cyanide forms the .compound AgCy.KCy or AgKCy,, in which the cyanogen group CN is represented by the symbol Cy, as is customary ; ferrous cyanide forms the 494 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,0 = K 2 CO 3 + (NH 4 ) 2 CO 3 . Acids do not set cyanic acid free from the cyanate, as, if formed, it is at once decomposed by water, thus : CNOH + H 2 = CO 2 + 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 SCX. This salt occurs in com- bination with others in nature, particularly in the mineral PRIMARY, OR ACID, POTASSIUM SULPHATE. 495 Icainite, 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 a 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 : 2KHSO, = K 2 S a 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 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 a O 8 . It is to be regarded as derived from two molecules of normal sulphuric acid by elimina- tion of four molecules of water : 496 INORGANIC CHEMISTRY. The acid probably lias the constitution represented thus : (HO) 1 OS80(OH) 1 = H 4 S,O e . Sulphites. When sulphur dioxide is passed into a solution of potassium carbonate until carbon dioxide ceases to escape, potassium sulphite, K 2 SO 8 , is formed. If the gas is passed to saturation the product is the pri- mary or acid sulphite, KHSO 8 . If the solution of the carbonate is hot and concentrated, the product is the disulphite, K 2 S 2 O 5 , 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 lese easily soluble in water than the normal salt. The dry RUBIDIUM CESIUM. 49? salt gives off carbon dioxide and water easily when heated, and is converted into the normal salt : 2KHC0 3 = K 2 C0 3 + C0 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. ; 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. 476). 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 iL 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. KUBIDIUM, Eb (At. Wt. 85.2). CESIUM, Cs (At. Wt. 132. 7). 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). Caesium was found in the Diirkheim mineral water, and was recognized by two characteristic blue lines, and the name caesium was given to it on this account (from ccesius^ 498 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. 23). 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 in 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. 499 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 AIF 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 a CO 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 method for the preparation of sodium on the large scale has recently been introduced 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 : 6NaOH + FeC, = 2Na 3 CO 3 + 6H + 2Na + Fe. The preparation of sodium is a problem of great im- portance to the world, for if this metal could be prepared cheaply the important metal aluminium could also be prepared cheaply. Both are supplied by nature prac- tically in unlimited quantities, but they are held so firmly in combination that it is a comparatively difficult matter to isolate them. 500 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, out 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 Eeichenhall 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. 501 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. The methods by which hydrochloric acid and chlorine are obtained have already been fully discussed. In the preparation of hydrochloric acid by the usual method sodium sulphate is necessarily formed. The methods for making the principal sodium compounds 602 INORGANIC CHEMISTRY. from the chloride will be taken up below. Salt is neces- sary to the life of man and many other animals. The role played by it in the animal economy is not under- stood, but it is found generally distributed throughout the body in small quantity. 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 C0 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 . The hydrosulphide and the sulphides of sodium are made just as the potassium compounds are, and resemble them very closely. Sodium Sulphantimonate, Na 3 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 iodide. 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. 503 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 + MgSO 4 = Na f 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 + 10H 2 O is present ; while if the solu- tion is heated above this point the compound breaks 504 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 exposed 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 2 SO 3 = Na 2 S + Na 2 S 2 O 3 . SODIUM CARBONATE. 505 It is also made by boiling a solution of sodium sulphite and adding sulphur : Na a SO, + S = Na 2 S 8 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 picturS has been developed. The action consists in the formation of salts in which both sodium and silver are contained, and which are soluble in water. It will be taken up more in detail under Silver (which see). Sodium Carbonate, Na a 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 3 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 + H,O, or (HO),C(ONa) 2 , separates. This, however, loses water 506 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 33 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 Le 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 SO 4 = Na,SO 4 + 2HC1. (2) The sodium sulphate thus obtained is heated with charcoal, which reduces it to sodium sulphide : Na,SO 4 + 2C = Na 2 S + 2CO 2 . SODIUM CARBONATE LE BLANC PROCESS. 507 (3) The sodium sulphide is heated with calcium car- bonate, when sodium carbonate and calcium sulphide are formed : Na,S + CaCO 3 = Na 2 CO 3 + CaS. The conversion of the sulphate into the carbonate is, therefore, expressed by the equation Na 2 SO 4 + 2C + CaC0 3 = Na 3 CO 3 + CaS + 2CO a . 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 -f- 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 + 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- 508 INORGANIG 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 2 + H 2 + = Na 2 SO 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 2 + CaH 2 S a + H 3 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 = HNaCO 3 + NH 4 C1. When the acid carbonate thus obtained is heated, it gives SODA FROM CRYOLITE. 509 off carbon dioxide, and is converted into the normal salt thus : 2HNaC0 3 = Na a CO 3 + CO a + H a O. The carbon dioxide given off is passed into ammonia, and thus again obtained in the form of acid ammonium carbonate : NH 3 + H a 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 drjness and heated, yields magnesia and hy- drochloric acid : MgCl 3 + H 2 O = 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 8 AlF 6 + 3CaCO 3 = 3CaF, + 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 AlO 3 + 3CO, + 3H,O = 3Na a 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. 510 INORGANIC CHEMISTRY. Mono-Sodium Carbonate, Primary Sodium Carbonate, HNaCOs. 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 2 C0 3 + C0 2 + 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, KWaCO 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 PO 4 + CO, + H 2 = 2HNa 2 PO 4 + Na 2 CO 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. t SODIUM BORATE. 511 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 polymerlsm ; 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 O 7 + 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, -f- 10H 2 O. By adding the required quantity of sodium hydroxide to a solution of borax, and evaporating to crystallization, sodium metaborate, NaBO 2 -f- 4H 2 O, is obtained : Na 2 B 4 O 7 + 2NaOH = 4NaBO a + H a O. The metaborate is decomposed when its solution is ex- posed to the action of the air. It is thus converted by the carbon dioxide into sodium carbonate and borax, or sodium tetraborate. 512 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. 7.01). 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. 513 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 514 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 : 2S0 4 1 ^ a = Na 2 S 2 7 + H 2 0. 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 ( ONa ONH 4 = PO^ OH + NH 3 ; OH (OH ONa OH = PO 2 ONa + H 2 O. OH 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. 515 so that the hydrogen forms water. This is illustrated in the decomposition of ammonium nitrate and nitrite : NH 4 NO, = N 2 O + 2H 2 ; and N 2H0. 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 sol 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 : 516 INORGANIC CHEMISTRY. The thiocarbamate afterwards breaks down when heated, forming the sulphocjanate and hydrogen sul- phide : OS - = CNSNH 4 + H 2 S. The salt, like so many other ammonium salts, causes a marked lowering of temperature when dissolveu 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, (]SrH 4 ) 2 S. This compound is ex- tensively used in chemical analysis for the purpose of J precipitating those sulphides which are soluble in dilute ! hydrochloric acid (see p. 198 and p. 469). 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. 517 power 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 : KE 3 + H 2 S = NH 4 HS. If only half this quantity of the gas is passed, the pro- duct is the sulphide : 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 ) a 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 ),S + O = 2NH, + H 2 + S. The sulphur set free in this way combines with the undecomposed ammonium sulphide, forming the com- pounds (NH,),S,, (NH 4 ),S a> (NH 4 ),S <( and (NH 4 ),S 8 . 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. 518 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 NO 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, CN"H 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 j This is the salt of an acid, CO -! Qjr 2 , known as car- bamic acid. When the carbamate is dissolved in water, it is converted into the carbonate : --roi NH < J JONH 4 . 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 : ONH 4 po j OH REACTIONS OF THE MEMBERS OF THE SODIUM GROUP. 519 It is easily decomposed into ammonia, water, and car- bon dioxide. Sodium-ammonium Phosphate, HNaNTELPO*. 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 + 4H 2 O. The changes which the anhydrous salt undergoes when heated were described on page 514. 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 ; 2XaP0 3 + M 2 = Na 2 M 2 P 2 O 7 ; NaPO, + M 2 = NaM 2 P0 4 . Many of these double phosphates and pyrophosphates are colored, and, like the double borates (see p. 512), 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. Yery few salts of the members of the sodium 520 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 ) 2 PtCl 6 . The cor- responding salt of sodium is easily soluble. Perchloric Acid, HC1O 4 , forms difficultly soluble potas- sium perchlor ate, 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 : KC1 + H 2 (C 4 H 4 6 ) = KH(C 4 H 4 6 ) + 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 < O e ) + K,C0 3 = 2K,(C 4 H 4 O.) + CO, + H,O. If, to the solution of the neutral salt, hydrochloric acid is added, the acid salt is again formed and precipitated : K,(C 4 H,0.) + HC1 = KH(C.H ,0,) + 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. 521 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 l>e 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. Ever}- light has its char- 522 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 tubes. Through one of the tubes 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 ob- served through the other tube, which is provided with lenses for the purpose of magnifying the spectrum. By means of a third tube, 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 quan- tities of some elements, and, since it was devised, several new elements have been discovered through its aid ; as, for example, csesium, rubidium, thallium, indium, gal- lium, and others. * For an account of the spectroscope and its uses, the student should consult some work on physics. The principles involved in its construe- 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: BERYLLIUM 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. Beryllium 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 Kb Cs 7.01 23. 39.03 85.2 132.7 Be 9.08 Mg 23.94 Ca 39.91 Sr 87.3 Ba 136.9 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 : (523) 524 INORGANIC CHEMISTRY. a b Lithium, Potassium, Sodium, and Bubidium, Caesium. Beryllium, 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, 23.94, is nearly half the sum of those of beryllium, 9.08, and calcium, 39.91. We have 9.08 + 39.91 ~T~ = <5 ' So, also, that of strontium, 87.3, is approximately half the sum of those of calcium, 39.91, and barium, 136.9 : 39 " 91 + 136 " 9 = 88.45. In the calcium group the specific gravities increase in the order of the atomic weights : At. Wt. Sp. Gr. Calcium, .... 39.91 1.57, Strontium, . . . 87.3 2.5, Barium, .... 136.9 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(NO 3 ) 2 , MSO 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. Beryllium sulphate, BeSO 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. 525 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 beryllium are insoluble in water. The solubility of the hydroxides increases as the atomic weight increases. Beryllium 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. 39.91). 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, 526 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 2 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 Q -|- 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. 490); 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 + 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 + 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. 527 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, + H 3 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 alcohol 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 much the same part at elevated temperature in facilitating reactions that water plays at ordinary tem- peratures. 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 : -s not complete in an at- 528 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 : OH OH OH OH OH CALCIUM HYDROXIDE. 529 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 : OH OH , - CaC L . OH OaU ' 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 530 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 alkaliss, 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) 3 + 4C1 = 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 {01 OC1 or CaOCl u . 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. 531 Ca(C10) 3 + 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 + 2HC1O = 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(ClO) 2 + H 2 S0 4 = CaSO 4 + 2HC1O ; CaCl 2 +H 2 SO 4 = CaSO 4 + 2HCl; 2HC1 + 2HC10 = 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, + 2HOC1 + CO 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, + CO a = CaCO 3 + C1 9 . 532 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(C10 3 ) a + 2CaCl a . This fact is taken advantage of, as has been shown, for the purpose of making calcium chlorate, and from this- potassium chlorate (see p. 490). 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 ita 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. 533 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- aute 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- 534 INORGANIC CHEMISTRY. cium hydroxide, the carbonate is precipitated ; and, if the current of gas is continued long enough, the carbonate is redissolved. It appears, therefore, that calcium car- bonate is soluble in water which 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 take 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, form- ing lime, or calcium oxide, and carbon dioxide, was re- ferred to on p. 464. Applications. Calcium carbonate is used, in the arts, for a great many purposes, as in the manufacture of glass ; as a flux (see p. 527) 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 U O. This is perhaps derived directly from the normal acid S(OH) 6 , having the con- CALCIUM SULPHATE. 535 stitution represented by the formula (HO) 4 SCa. The salt of the formula CaSO 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." 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 : CaS0 4 + 2H,0 = (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 CaSO. + Na 3 CO 3 = Na,SO 4 + CaCO,. 536 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 ) 2 . (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. 537 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 make 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,(PO.), + 2H 2 SO, = CaH 4 (P0 4 ), + 2CaSO ( . 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 : 2HNa a PO 4 + 3CaCl 3 + 2NH 3 = Ca 3 (PO 4 ) 2 + 4NaCl -j- 2NH 4 C1. 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. 538 . INORGANIC CHEMISTRY. The reaction takes place thus : 4HNa 2 PO 4 + 4CaCl a = CaH 4 (PO 4 ) 2 + Ca 3 (P0 4 ) 2 + SNaOl. On standing, the primary acts upon the tertiary salt, forming the secondary phosphate thus : CaH,(PO,)i + Ca 3 (PO,), = 4HCaP0 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 3 (P0 4 ) 2 + 4HC1 = 2CaCl 2 + H 4 Ca(PO 4 ) 2 ; 2CaCl 2 + H 4 Ca(PO 4 ) 2 + 4NH 3 = 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(P0 4 ) 2 =HCaPO 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. 539 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 gloss 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 540 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 Kupert'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 tempering. 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, 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, carbonate of cal- cium 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 STRONTIUM. 541 kept moist in consequence of the reaction above men- tioned. It is, however, difficult to see why the presence of a little extra moisture 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 hydravlic 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. STRONTIUM, Sr (At. Wt. 87.3). 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 542 INORGANIC CHEMISTRY. 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, Ba (At. Wt. 136.9). 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 -f 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. COMPOUNDS OF BARIUM. 543 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. 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 3 , is added to a solution of barium hydroxide, a precipitate is formed which has the composition BaO, -)- 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- 544 INORGANIC CHEMISTRY. soluble barium sulphate is formed and the hydrogen dioxide remains in solution : Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 O,. 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 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 ; while with barium peroxide the reaction takes place thus: BaO a + 2HC1 = BaCl, + H a O 2 . 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 : 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- PHOSPHATES OF BARIUM. 545 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 uoite and form insoluble compounds with me- 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 2 + 2HNa 2 P0 4 4- 2NH 3 = Ba 3 (PO 4 ) 3 + 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 2 + 4HNa 2 PO 4 = BaH 4 (PO 4 ) 2 + Ba 3 (PO 4 ) 2 546 INORGANIC CHEMISTRY. 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 (P0 4 X + Ba s (PO,) 2 = 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 6 , in solutions of barium salts. This is 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. BERYLLIUM. 547 Ammonium oxalate, (NH 4 ) a C 2 O 4 , produces precipitates of the oxalates in solutions of calcium, barium, and strontium. Only the calcium salt is insoluble in dilute acetic acid. Potassium dicliromate, K 2 Cr 3 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. BERYLLIUM, Be (At. Wt. 9.08). Occurrence and Preparation. The principal form in which the element beryllium occurs in nature is in the mineral beryl, which is a silicate of aluminium and beryl- lium of the formula Al a Be 3 (SiO 3 ) 6 . Emerald has the 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 beryllium, made by those who have prepared it in dif- ferent ways, differ somewhat from one another, evidently in consequence of the fact that it has not generally 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 potas- sium hydroxide, forming in all probability a beryllate of the composition Be(OK) 2 : Be + 2KOH = Be(OK) 2 + H a . 548 INORGANIC CHEMISTRY. The specific heat of beryllium is 0.425, and this, mul- tiplied by the atomic weight, 9.08, gives the product 3.86, instead of a figure near 6.4, as would be expected accord- ing to the law of Dulong and Petit. On the other hand, the analysis and the determination of the specific gravity of the vapor of beryllium chloride have shown that it has the composition represented by the formula BeCl 2 , in which the atomic weight of beryllium is 9.08. It appears, therefore, that beryllium, like carbon, boron, and silicon, is an exception to the law of specific heat, at least at the ordinary temperatures. Compounds of Beryllium. The compounds of beryl- lium differ in many respects from those of the group cal- cium, 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 beryllium oxide. It shows- a marked tendency to form basic salts. Beryllium Chloride, BeCl 2 , is formed by the action of chlorine on beryllium, and more easily by treating a. mixture of beryllium oxide and carbon with chlorine, the reaction being similar to that employed in making sili- con chloride (seo 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 beryllium is 9.08. 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 regarded as conclusive in favor of the atomic weight 9.08. Beryllium Hydroxide, Be(OH) 2 , is thrown down as a precipitate when a soluble hydroxide is added to a solu- tion of a beryllium salt : BeS0 4 + 2NaOH = Be(OH) 2 COMPOUNDS OF BERYLLIUM. 549 It is a white, gelatinous mass, which is soluble in potas- sium and sodium hydroxides and in ammonia, so that, after precipitation from beryllium salts by these reagents, it redissolves. This solution is due to the formation of beryllates of the formula Be(OM) a : Be(OH), + 2NaOH = Be(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 \vhich 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. Beryllium Sulphate, BeSO 4 , is formed by dissolving beryllium hydroxide in dilute sulphuric acid, and has the composition BeSO 4 -|- 4H 2 O when crystallized from water. When a solution of this salt is heated with be- ryllium hydroxide, basic salts are formed, of which the following are examples : Be 2 SO 5 and Be 3 SO 6 . The first of these is to be regarded as derived from a hydroxide of the formula HO Be O Be-OH, by neutralization with sulphuric acid, as represented in the formula O<-p^_Q>SO 2 ; the second from a hydroxide of the formula HO-Be-O-Be-O-Be-OH, by neutralization with sulphuric acid, as represented in the formula Q Be-0) Be, or possibly rk>Be. Beryllium Carbonate, BeCO 3 . When a slight excess of sodium carbonate is added to a solution of beryllium 550 INORGANIC CHEMISTRY. sulphate, a basic carbonate of the formula Be 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-Be-O-Be-O-Be-OH by neu- tralization with carbonic acid, as represented in the Be-0) formula X5 B e VCO. < Be-0 j Weak Basic Character of Beryllium. The power of beryllium 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 ; beryllium, 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 beryllium 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 beryllium it is complete, as it is with the similar metal MAGNESIUM. 551 magnesium. In general, the more acidic the element the more easily is its chloride decomposed in this way. MAGNESIUM, Mg (At. Wt. 23.94). 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 2 ; kieserite is magnesium sulphate, MgSO 4 -f- 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 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 652 INORGANIC CHEMISTRY. 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 beryl- lium. 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 beryllium. Its hy- droxide does not form salts with the hydroxides of potassium and sodium. On the other hand, its chlo- 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 beryllium. 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 beryl- lium 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 -|- 6H 2 O. When this MAGNESIUM CHLORIDE. 553 crystallized salt is heated for the purpose of driving off the water, it is completely decomposed in accordance with the following equation : MgCl, + 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 a .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- dium, and ammonium, forming crystallizing compounds of the formulas KMgCl 3 , NaMgCl 3 , and NH 4 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. Iu seems probable that the latter is analogous to the po- OT^ tassium compound of beryllium of the formula Be < Q-g-. It corresponds to an oxygen compound of the formula, O-TT- Mg<^-j^, 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- pound would have the formula ^g<(cn-K' The exist- ence of two double chlorides of magnesium and potas- 554 INORGANIC CHEMISTRY. slum is suggested by what has been said regarding compounds of this kind (see p. 461). One of these 01 would be represented by the formula Mg < / \TT , the other by the above formula. Both are known. Further, magnesium bromide forms the salt K 2 MgBr 4 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 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 fOH should be written (HO) 2 MgSO,, or OS -j , in which it appears as a derivative of the acid SO(OH) 4 . The salt MgSO 4 + 7H 2 O (or H 2 MgSO 5 + 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- MAGNESIUM CARBONATE. 555 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 beryllium sulphate (see p. 549). The constitution of the double sulphate of magnesium and potassium is probably that expressed in the for- mula n >Mg. Magnesium Carbonate, MgCO 3 . Like beryllium, mag- nesium shows a marked tendency to form basic salts 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 (MgO) 3 (OH) 2 (CO) 2 . It is derived from three molecules of magnesium hydroxide and two of carbonic acid, as is co< 0-Mg-OH more clearly shown in the formula X>Mg . The CO< 0-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. 556 INORGANIC CHEMISTRY. 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 + 3H 2 O and MgCO 3 + 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.OH = Mg(OH) 3 + (NH,),SO,. 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 ; 2MgHP0 4 = Mg 2 P 2 7 +H 2 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 1B + MgCl 2 . The borate, Mg 3 B H O 15 , is derived ERBIUM. 557 from the acid, H 6 B 8 O 1B , 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 a SiO 4 . It is the neutral salt of normal silicic acid. Serpentine is derived from the acid, O < gj/Q jjK an d has the composition Mg 3 Si 2 O 7 + 2H 2 O. Silicon-Magnesium, Mg-.-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 silicon-magne- sium. When treated with hydrochloric acid it gives silicon hydride, SiH 4 , and hydrogen : Mg 2 Si + 4HC1 = 2MgCl 3 + SiH 4 . The liberation of hydrogen is due to the presence of an excess of magnesium. Reactions of Magnesium Salts which are of Special 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. 166). 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 166, 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 558 INORGANIC CHEMISTRY. 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 the next chapter. 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, A1 2 0, CHAPTER XXVII. ELEMENTS OF FAMILY III, GROUP A : ALUMINIUM SCANDIUM YTTRIUM LANTHANUM- YTTERBIUM. 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 : MC1 3 , M(OH) S , M(N0 8 ) 3> M.(S0 4 )., M,(CO S ) S> 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. 27.04). Occurrence. Aluminium is an extremely important element in nature and in the arts. It occurs very (559) 560 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 -|- 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 most uses to which iron is put, and for most 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 thau 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. 561 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 laws 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 an opinion upon them. Until recently the commercial preparation of aluminium has appeared to be intimately connected with that of the commercial preparation of sodium ; but, if the latest method is as good as is claimed, this is no longer the cas6. 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, 562 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 Avith 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 2 O 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. 563 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 ai 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 A1 3 C1 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 ,C1 C1-)A1-A1^-C1 . On the other hand, in a compound / \ Cl 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 itself, 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 564 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 , N (01,)K so it is possible that aluminium chloride may have the constitution represented by the formula A1^-(C1 2 )-)A1, \c\y 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- mulas A1C1 3 .KC1 and A101,.NaCl, or probably Al-Cl \C1 2 )K /Cl and A1^-C1 . The fluoride forms two compounds \Cl 2 )Na with potassium fluoride and two with sodium fluoride. ALUMINIUM HYDROXIDE. 565 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,)K and A1^-(F 2 )K . The tri-sodium fluoaluminate is the min- \F,)K 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 : Aid, + 3NH 4 OH = Al(OH), + 3NH 4 C1. 566 INOEGANIC 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) 3 . When heated under proper conditions it loses water, and forms the compound A1O 2 H : A1(OH), = A10.0H + H 2 0. 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 O 3 + 3H 2 O. 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 : 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. 567 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 A1,0(OH) 4 . 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) 3 , 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(OlS T a) 3 + SCO, + 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,AlF 6 + 3CaO = NaAlO, + Na f 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- 568 INORGANIC CHEMISTRY. agent used dissolves the precipitate. This action is the same in character as that which takes place in the case of beryllium, 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), + 3H,O, and Al(OH), + 3NaOH = Al(ONa), + 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 might 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.ONa is 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 A-,Q *Q>Mg, or Mg(AlO 2 ) 2 . Chrysoberyll is the corresponding beryllium salt Be(AlO 2 ) 2 ; and gahnite is the zinc salt Zn(AlO 2 ) 2 . These salts are extremely stable, differing markedly in this ALUMINATES. 569 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. 495. 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> r Fe ( Cr *) 2 - 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 a ) 2 ; and there is also a compound of magnesium, TQ *Q>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, . . . *Q>Ca JBarium aluminate, . . . */ Q>Ba spinel, ....... i!8:o >M g Chrysoberyll, ..... AiaO >Be Gahnite, Chromite, ...... CrO'8 >Ee Magnetite, ...... FeO.O >Fe Magnesio-ferrite, .... There is a highly instructive analogy between the aluminates and the double chlorides and other similar 570 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 : KA1O 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 l Al^Cl \Cl,)Na. The compounds of the same order as cryolite have their analogues in such oxygen compounds as Al(ONa) 3 , etc., as is shown by the following formulas : Na 3 A10 8 ; Na 3 AlF 6 ; /ONa /( F *)Na Al^ONa ; Alf (F 2 )Na . \ONa \F f )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. 571 discussing the subject of double chlorides in general, three series of such salts may be looked for, correspond- ing to the formulas XC1 2 )M AlfCl , A1HC1 2 )M, and Alf (Cl f )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. 57*3 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 SQ probably be represented by the formula Al \ Cr k ' 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 : ^}Q'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 the ordinary aluminium salt, 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(SO 4 ) 2 + 12H 2 O or K 2 S0 4 + A1 2 (SO 4 ) 3 + 24H 2 O. The most rational view which has been expressed re- garding this compound is that it has the constitution ALUMS. 573 S0 2 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 ) 2 + 12H 2 ; LiAl(SO 4 ) 2 + 12H 2 O ; (NH 4 )A1(S0 4 ), + 12H 2 0; CsAl(SO 4 ) 2 + 12H 2 O ; TlAl(SO 4 ) a + 12H 2 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(S0 4 ) 2 + 12H 2 ; 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 574 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, a,nd 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 ),. ALUMINIUM SILICATE. 575 When the solution is boiled the salt contained in it is decomposed, forming ordinary alum and another basic alum which is insoluble : 3[K,(A1,0)(SO,)J = K(A10),(SO,) S + 3KA1(SO 1 ) 1 + 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, (NH 4 )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 576 INORGANIC CHEMISTRY. what has 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 which 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 -f- 4H 2 O, according to which the substance is the salt of normal silicic acid, Si(OH) 4 . Other specimens have the compo- oH sition HAlSi0 4 + H,O, or Si Q\ AI , HQ . When O/ CLAY ULTRAMARINE. 577 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-lazidi 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. 578 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 Jieated 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, with 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. 579 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 3 (SO 4 )3 + 3(NH 4 ) 2 S + 6H 2 O = 3(NH 4 ) 2 SO4 + 3H 2 S + 2A1(OH),. Soluble carbonates precipitate aluminium hydroxide for the same reason that the soluble sulphides do. The re- 580 INORGANIC CHEMISTRY. action between aluminium sulphate and sodium carbon- ate takes place thus : Al 2 (SO 4 ) 3 +3Na 2 CO 3 +3H 2 0=3Na 2 S0 4 +3C0 2 +2Al(OH) 8 . OTHER MEMBERS or FAMILY III, GROUP A. Scandium, Sc(At. Wt. 43.97). This element was discov- ered only recently 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, Sc 2 (SO 4 ) 3 ; a double sulphate, KSc(SO 4 ) 3 ; etc. It is of special in- terest for the reason that its properties were foretold by Mendelejeff 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 Mendelejeff, the latter described the properties of an ele- ment then unknown, and which he called ekaboron, which 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. 88.9), like scandium, is found in gadolinite, euxenite, and some other rare minerals. The element itself has not been isolated. Its chloride, YC1 3 -f- 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, Y. 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. 581 Ytterbium, Yb (At. Wt. 172.6). 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 (S0 4 ) 3 , etc. 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- 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, beryllium, and boron. The base- forming elements thus far considered form the principal 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 582 INORGANIC CHEMISTRY. 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. Having 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-SILVERGOLD. 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 trivalent, 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. 63.18). 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, (583) 584 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 2 CuBr CuBr 3 Cu 2 O 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. 585 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-ClCu 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 586 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 + SO 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. 587 the preparation of nitrogen, as was stated in speaking oi 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 imth 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 588 INORGANIC CHEMISTRY. nickel. The proportion of copper varies from 40 to 60 per cent ; that of zinc from 19 to 44 per cent i and that of nickel from 6 to 22 per cent. Copper Hydride, CuH. This compound is made by treating a solution of hypophosphorous acid or of hydro- sulphurous acid with a solution of copper sulphate. It is thrown down as a yellow precipitate which gradually becomes darker. At 60 it decomposes into copper and hydrogen, and treated with hydrochloric acid it yields cuprous chloride and hydrogen : CuH + HC1 = CuCl + H 2 . As will be seen, it is less stable than the hydrogen com- pounds of potassium and sodium, and has a different composition. Cuprous Chloride, CuCl, is formed by heating cupric chloride, CuCl 2 ; by passing hydrochloric acid over highly heated copper ; and by treating cupric chloride with re- ducing agents, as, for example, with stannous chloride, SnCl 2 , or sulphurous acid. The action with these two reagents takes place as represented in the equations 2CuCl 2 + SnCl 2 = 2CuCl + SnCl 4 ; 2CuCl 2 + S0 2 + 2H 2 O = 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, 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. 589 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 2 + 2H 3 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 ^jj s >CuCl 2 . 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, Col. 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 : 2CuSO 4 + 4KI = 2K 2 SO 4 + 2CuI + 1 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 a 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 s O 3 (OH) 2 or 4Cu 2 O.H 2 0, is easily made by adding sodium hydroxide to a solution of 590 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 ruby 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 a] lowed 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 SO 4 = CuSO 4 + H 2 + 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) 2 + 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 the 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. 591 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 -|- 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 SO 4 = CuSO 4 + S(X + 2H a 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 proportion 592 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 (OH) 2 or [g>Cu OS -j QTT , 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 + 7H 2 O Beryllium sulphate, BeSO 4 + 7H 2 O Ferrous sulphate (green vitriol), . . . FeSO 4 + 7H 2 O Nickel sulphate, . NiSO 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, etc. CUPRIC SULPHATE. 593 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 + 6H a O and (NH 4 ) 2 SO 4 .CuSO 4 + 6H 2 O, and probably have the consti- tution represented by the general formula qo OM bO 2 < o X>Cu. so 2 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 : CuS0 4 .5NH 3 ; CuS0 4 .4NH 3 .H 2 O ; CuSO 4 .3NH 3 .2H 2 O ; CuS0 4 .2NH 3 .3H,O; CuS0 4 .NH,.4H,O ; CuSO 4 .5H a O. 594 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(WO 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 blue precipitate is formed, which has the composition CuCO. 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. 595 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),. 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(CX) 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 2 S.Fe a S 3 or CuFeS 2 . It 596 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 oxidation, 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. 597 SO 2 <>Cu + 2NH 3 + 2H a O - X> S0 2 S0 2 <~>Cu Cu< OH If the action is carried farther, the basic salt dis- solves, forming the compound referred to under Cupram- monium 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) e 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. 107.66). 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 598 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- tinson's 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. 599 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 3 + 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,Cl f + 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 C0 3 = 2Ag + CO 2 + O + 2NaCl ; 600 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. 601 laboratory, or to evaporate their solution, as in the preparation of the caustic alkalies. 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 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 602 INORGANIC CHEMISTRY. 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 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 2 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 COMPOUNDS OF SILVER. 6C3 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 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 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 a S.Ag 2 S, and pyrargyrite, 3Ag 2 S.Sb 2 S 3 . 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 604 INORGANIC CHEMISTRY. 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 3 + 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 not 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 16 , 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 GOLD. 605 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 a 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, HNa 2 PO 4 , gives a precipitate of the neutral salt Ag 3 PO 4 , which is yellow. Potassium ferrocyanide precipitates white silver ferro- cyanide, Ag 4 Fe(CN) 9 . 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. 196.7). 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 3 O Au 8 O 3 Those of the first series are called aurous compounds, 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 606 INORGANIC CHEMISTRY. 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 MA1O 2 , 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 -f- 4H 2 O and HAuBr 4 -f- 5H 2 O, which are plainly the acids from which the chlcr- aurates and the brom-aurates are derived. Besides the compounds of gold in which the element is univalent and those in which it is trivalent, there are, further, a few in which it is bivalent, as the chloride AuCl 2 , and the bromide, AuBr 2 . 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 METALLURGY AND PROPERTIES OF GOLD. 607 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. The gold obtained as above 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 dilut- ing, and treating with a reducing agent, when metallic gold is precipitated. Thus when ferrous sulphate is used the following reaction takes place : 3FeS0 4 + Audi, = 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 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. 608 INORGANIC CHEMISTRY. 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, AuCl 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 forms the dichloride, AuCl 2 . This, when treated with a little water, breaks down into auric chloride and aurous chloride, and by further treat- ment with water the latter yields auric chloride and gold. By filtering and evaporating to dryness, auric chloride is obtained in the anhydrous condition. It can also be obtained in crystallized form, the crystals hav- ing the composition AuCl 3 -f- 2H 2 O. When anhydrous auric chloride is heated to 185, it loses chlorine and is converted into aurous chloride, AuCl. This, as stated above, 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 pre- cipitate, known as the purple of Cassius, 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 2 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 AURIC HYDROXIDE. 600 /Cl Au^--Cl , being similar to that of the acid from which \C1,)H /Cl potassium chlor-aluminate is derived. A1^-C1 . The 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, EAu(CN) 4 , with silver nitrate, which gives the silver salt, and then decomposing this with hydrochloric acid. Ihe 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 salts known as the aurates are obtained. In composition these are analogous to the meta-aluminates. Potassium anrate, for example, has the composition KAuO 3 . The analogy between some of the compounds of aluminium and those of gold is shown in the following table : A1A Au,0 3 Al(OH), Au(OH), A1C1 3 AuCl, /Cl /Cl A1^-C1 Au^-Cl \(C1 2 )K \(C1 2 )K 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, MS0 4 , MCO 3 , 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, Hgl Hgl, Hg,0 HgO HgNO, Hg(NO s ) 2 , etc. Just as the first member of Group A, Family II, beryl- lium, 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 B, Family II, zinc, is (610) ZINC. 611 acidic, while the other members of the group are not. Beryllium hydroxide and zinc hydroxide dissolve in caustic alkalies, forming beryllates 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. 65.1). 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 a 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 FeZn has been stated (see p. 592), 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 >co zinc sulphate give mainly the compound Zn<^ ZnCo. CADMIUM, Cd (At. Wt. 111.7). 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 are 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 112. 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 + 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. 617 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. MEKCURY, Hg (At. Wt. 199.8). 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 618 INORGANIC CHEMISTRY. mercurous 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 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. 619 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. M'ercury 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 \vater, 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,Hg. The product is extremely voluminous, and swells up during the reaction, so that it occupies under favorable 620 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 O + SO 2 = 2HgCl + H 2 S0 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. 621 mercurous chloride is treated with nitric acid it is con- verted into mercuric chloride and mercuric nitrate, as represented in the equation 6HgCl + 8HNO, = 3Hg(NO 3 ) 2 + 3HgCl, + 2NO + 4H 2 O. If hydrochloric acid is present in sufficient quantity the action takes place thus : 3HgCl + 3HC1 + HNO 3 = 3HgCl a + 2H 2 O + NO. Further, the conversion of mercurous nitrate into mer- curic nitrate is represented by the equation 3HgNO 8 + 4HXO, = 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,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, HgSO 4 + 2NaCl = HgCl, + 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,. It is easily reduced to mercurous chloride by 622 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 B ; or Hg^J^ HgPb, 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 2 + 2Pb(NO 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 + 01 ,. RED LEAD. 645 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 a PbO 3 , or PbOPb 0> Pb As partial experimental evidence 646 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 OTTT sented by the formula Pb 4 ^^. When this is treated LOK with lead oxide the corresponding lead salt should be formed. Eed lead is used as a pigment, and in place of litharge whenever 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(N"O 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. 647 tainecl 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, + 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 viuegar. There is less risk of discoloration in consequence of the formation of sulphuretted hydrogen, but the fermen- 648 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 s ; or these may be expressed structurally by the formulas Pb< ro ( ~' u <> 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 a red heat it is partly decomposed 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. 649 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 4 ) 3 -f- 3H 2 O. ttj 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. 138), 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 3 , and in all its salts it acts as a- trivalent element. Cerium, Ce (At. Wt. 141.2), forms two series of com- pounds, in one of which it is trivalent, 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 O 3 , and Ce 2 (S0 4 ) 3 . 650 INOMOANIC 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,0 3 , Di(N0 3 ) w Di a (S0 4 ) s) Di a (CO a ),, etc. An oxide of the formula Di 2 6 appears to have been made. Should further investigation confirm this, there would be some basis for classifying didymium with the elements of Family V, Group A, in which it falls accord- ing to its atomic weight. CHAPTER XXXI. ELEMENTS OF FAMILY VI, GROUP A : CHROMIUM-MOLYBDENUM-TUNGSTEN-URANIUM. General. At the end of Chapter XIV., 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 8 , 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 w r hich 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. (651) 652 INORGANIC CHEMISTRY. CHROMIUM, Cr (At. Wt. 52.45). General. This element forms three series of com- pounds, in which it appears to be respectively bivalent, trivalent, and sexivaleiit. 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,. Of the second series, or the chromic compounds, some of the principal members are : CrCl 3 , Or A, Cr,(S0 4 ) s) Cr(NO 3 ),, KCr^O.), + 12H.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 = CrSO 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 3 = K 2 Cr0 4 + H 2 O. CHROMIUM. 653 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 Cr0 4 + 5H 3 SO 4 = 2K a 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. 569), being an iron salt of the acid CrO.OH, which may be called metachromous acid. This view is represented by the formula r^o'o^Fe. It occurs also in the mineral crocoisite, which is lead chromate, PbCrO 4 . The name chromium is derived from the Greek ^pcSyw^, 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,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. 654: 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 laming of the same color ; and this variety is insolutJIcT in watc r 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. 655 KCrCl 4 ; CrCl s .2KCl, or K 2 CrCl 5 ; and CrCl 3 .3KCl, or K 3 CrCl a . 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 : 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 -f- 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^Xir + 2H 2 O. 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- um salt referred to. 656 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 20(OH) 3 = Cr 2 0(OH) 4 + H 2 O ; 4Cr(OH) 3 = Cr 4 3 (OH) 6 + 3H 2 O. Chromic Oxide, Cr a O 8 , 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 + 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. 657 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, CxO 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 658 INOEOANIC CHEMISTRY. all the chromium appears in the form of potassium chro- 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(Cr0 2 ) 2 Fe + 4CaO + 7O = 4CaCrO 4 + Fe 2 O 3 ; OaCrO 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 CrO 4 + 2HNO 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,OA + 4H 2 SO 4 = 2KCr(S0 4 ) 2 + 4H 2 O + 3O. All the oxygen in the chromate in excess of that required POTASSIUM DICHROMATE. 659 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,Cr a O 7 -f 4H 2 SO 4 + 3SO, = 2KCr(SO 4 ) 2 + 3H 3 SO 4 + 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 3 O. 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 O 2 O 7 + 4H 2 SO 4 + 3C 2 H 6 = 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,O 7 + 14HC1 = 2KC1 + 2CrCl 3 + 7H 2 O + 6C1. 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 : 30 + 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 8 O. 660 INORGANIG 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 : 2CrO 3 = Cr 2 3 + 3O ; and when heated with sulphuric acid it gives chromic sulphate and oxygen : 2Cr0 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 Dichromates. 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 6 or or 0,S(OH), O s 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 DICHROMATES. 661 the same as that between the sulphates and disulphates, as represented by the equations OH +H 8 0; ~r. 3< OH 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 : r OK r OK Cr 'Pb This then loses water, and forms the salt 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, K 2 Cr 3 O ]0 , and potassium tetra- chromate, K 2 Cr 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 + 2KNO 3 + H 2 O ; 2K 2 Cr 2 7 + 2HN0 3 = K 2 Cr 4 O J3 + 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 Oxy chloride, Chromyl Chloride, CrO 2 Cl 2r is analogous to sulphuryl chloride, SO 2 01 3 , and is to be ANALYTICAL REACTIONS OF CHROMIUM. 663 regarded as derived from chromic acid by the replace- ment of the hydroxyls by chlorine : CrO.<|; CrO, Na a Mo,O IO , Na 2 Mo 4 O 13 , Na 2 Mo 8 O 25 , Na 2 Mo 10 O 31 , and Na 6 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 : 2H 2 Mo0 4 = H 2 Mo 2 7 + H 2 O ; 3H 2 Mo0 4 = H 2 Mo 3 10 + 2H 2 O ; 4H 2 Mo0 4 = H 2 Mo 4 13 + 3H 2 O ; 8H 2 Mo0 4 = H 2 Mo 8 O 2B + 7H 2 O ; 10H 2 Mo0 4 = H 2 Mo ]0 31 + 9H 2 ; 7H 2 Mo0 4 = H 6 Mo 7 24 + 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. 667 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 a O crystallizes out. TUNGSTEN, W (At. Wt. 183.6). General. Like molybdenum, tungsten forms a large variety of compounds. With chlorine it forms four, of which the formulas are W T C1 2 , WC1 4 , W T C1 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 FeW T O 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 S . 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. 668 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 -|-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 ]0 W 12 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 -f- 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 -j- 7H 2 O crystal- lizes. This is soluble in one fifth its weight of water, URANIUM. 669 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 8 W M O 40 , in combination with one molecule of silicon dioxide, H 8 W 12 40 .SiO, UBANIUM, U (At. Wt. 239.8). 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 : UO 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 3 (NO 3 ) 2 , is obtained, and, by igniting this, the trioxide, UO S , 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. 670 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, UO 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 SO * = U0 2 S0 2 + 2H 2 0; TJO OH , HO.N0 2 _ uo O.N0 2 , OH o < HO.NO, - 2< O.N0 2 + 2 ^' They are derived from the acids by replacing the hydro- gen by uranyl, UO 2 , which is bivalent. Uranyl nitrate, UO 2 (NO 3 ) 2 , is easily obtained, as above described, and crystallizes well in lemon-yellow prisms. Uranyl sul- phate, UO 2 (SO 4 ), is formed by treating the nitrate with URANATES, 671 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 TJ 2 O 7 , which may be called diuranic acid, as in n<"ra position 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 ydlow, 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 the formula uo,Mn view is expressed by the formula Mn 4 . The de- [6> Mn composition with acids, as with sulphuric acid, would, according to this, be represented thus : MnO 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. 676 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 ot 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 recently been suggested, upon the basis of experimental investi- gations, that the first product of the action is a com- pound of the formula H 2 MnCl 6 , which is the chlorine compound analogous to the oxygen acid, H 2 MnO 3 . The action is supposed to take place as represented in the following equation : O=Mn=O + 6HC1 = ~>Mn-Cl 2 + 2H 2 O. 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 : 2 = MnCl 2 + 2HC1 + 01,. MANGANITES. 677 The manganous 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 5 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) a + H 2 O ; 5MnO(OH) 2 = Mn 5 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 10 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 io ; 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, + C1 2 + 2H 2 0. 678 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 2 + Ca(OH) 2 = Mn(OH) 3 + CaCl 2 ; and when this hydroxide mixed with lime is allowed to stand exposed to the air oxidation takes place, and a, compound CaMnO 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 & + 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.MnO, + 6HC1 = CaCl 2 +MnCl 2 +3H 2 O + C1 2 ; CaO.2Mn0 2 + 10HC1 = CaCl 2 + 2MnCl 2 + 5H 2 O + 2Cl a . 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 manganite 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. 679 Manganous Cyanide, Mn(CN) 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),.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),.3NaCN or Na,Mn(CN) 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 2 3 + 2H 2 S0 4 = 2MnS0 4 + 2H 2 O + O ; MnO 2 + H 2 SO 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 ) + 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 680 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,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 2 + 2KOH = K 2 MnO 4 + Mn 2 O 3 + H 2 O. It is also made by fusing the dioxide with potassium hy- droxide and potassium chlorate, when this reaction takes place : 3MnO 2 + 6KOH + KC1O 3 = 3K 2 MnO 4 + KC1 + 3H 2 O. 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 is treated with an acid, the manganic acid is at once decom- posed into permanganic acid and manganese dioxide : 3H 2 MnO 4 = 2HMnO 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 2 Mn0 4 + 2C0 2 = 2K a C0 3 + MnO 2 + 2KMnO 4 . POTASSIUM PERMANGANATE. 681 With water the change takes place thus : 3K a MnO 4 + 2H,O = 2KMnO 4 + MnO, + 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 + 2KMn0 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, + 2KOH + O = K 2 MnO 4 + H a 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 682 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. Very 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 a MnO 8 . 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 Condy's liquid. When a solution of barium permanganate is treated with sulphuric acid, free permanganic acid is obtained REACTIONS OF MANGANESE COMPOUNDS. 683 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 S0 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 2 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). 6 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- 684 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 l)orax 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 (55.88) is nearly the same as that of manganese (54.8), 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 Fed,, 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 3 , Fe(OH) 3 , Fe,O 3 , Fe(NO,),, Fe a (SO 4 ) 3 , 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- (685) 686 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. MC1 2 M(OH) 2 MO M(N0 3 ) 2 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 which are exceptional, are put in brackets : MnSO 4 [Mn 2 (SO 4 ) 3 ] MnO 2 K 2 MnO 4 KMnO 4 [CrSO 4 ] Cr 2 (SO 4 ) 3 CrO 3 K 2 CrO 4 [HCrO 4 ](?) FeSO 4 Fe 2 (SO 4 ) 3 [K 2 FeO 4 ] CoS0 4 Co(OH) 3 NiS0 4 [Ni(OH) 3 ] A1 2 (S0 4 ) 8 ZnSO 4 SO 2 SO 3 K 2 S0 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. 687 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 Fed, may combine with another molecule of the same kind to form a double chloride of the formula Fe.Cl. or Fe<-(Cl.)-^Fe. It may X (C1 2 )/ 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 a 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 688 INORGANIC CHEMISTRY. ferrous compounds, then in all probability cobalt and nickel are bivalent in their principal compounds. IKON, Fe (At. Wt. 55.88). 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. 689 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 by placing in the furnace alternating layers of coke or charcoal, and the ores mixed with proper fluxes. The 90 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. 691 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 spiegd-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 ivrought-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 upon its properties. The process of welding consists in heating 692 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- iroii after the oxidation. By addition of cast-iron a 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 limestone, such iron may be used. Under these circumstances the phosphorus is oxidized, and with the limestone forms calcium phosphate, which is of value as a fertilizer (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 Wrouglit-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. 693 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 wrrought-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 which 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. 694 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 2 . 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, FeCl 2 . 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 2 + H 2 O + O = 2Fe < + 2FeCl 3 . When treated with oxidizing agents in general, as nitric FERRIC CHLORIDE. 695 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 Fed,. 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 cJdorati. 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 at 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 FeClg. When treated with nascent hydrogen, ferric chloride is converted into ferrous chloride : FeCl 3 + 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) 3 , corresponding to the above-mentioned chlorides, are not known : only double compounds of these with 696 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- siate of potash crystallizes. This has the composition K 4 Fe(CN) 6 + 3H Q 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 .3KCK 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(CN) 6 K 3 Fe(CN) 9 Sodium ferrocyanide, Sodium ferricyanide, Na 4 Fe(CN) 6 Na 3 Fe(CN) Barium ferrocyanide, Ba a Fe(CN) Ferric ferrocyanide, Ferrous ferricyanide, Fe 4 [Fe(CN) G ] 3 Fe 8 [Fe(CN) 6 ] a Ferri- potassium ferrocyanide, KFeFe(CN) e 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. 697 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, + 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 SO 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). + 3H 2 SO 4 = 6HCN + 3K,SO 4 + 2KFe(CN) 3 . Ferrohydrocyanic Acid, H 4 Fe(CN;, 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 4 [Fe(CN) 6 ] 8 Ferro-hydrocyanic acid Ferric ferrocyanide, or Prussian blue Ferric Ferrocyanide, or Prussian Blue, Fe4[Fe(CN)e]s. 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) 6 ], + 12KC1. 698 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)J 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) e + 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,Fe(CN) e (1) Ferrous salt, . Fe 2 Fe(CN). (3) Ferrous salt, . Fe 3 [Fe(CN) 6 ] 3 (2) Ferric salt, . . Fe 4 [Fe(CN) 6 ] 3 (4) Ferric salt, . . FeFe(CN) 6 IRON SALTS-NITROPRUSSIATES. 699 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 ot this a ferro-potassium salt, of the formula K 3 FeFe(CN) 6 , is formed : K 4 Fe(CN) 6 + Fed, = 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) 5 (NO) + 2H,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 700 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 ars 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, f> r\r\>Q** * s 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 e Q*Q>Zn, but some of the zinc is replaced by iron and manganese. OXIDES OF IRON. 701 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 702 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 ) 2 SO 4 . The sulphide thus obtained oxidizes readily in the air, ^nd forms the sulphate. The compact variety is used in making hydrogen sulphide (which see). Ferric Sulphide, Fe 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 aluminic 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. 703 " green vitriol" (see p. 592), 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, ^tc., 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 -,-, . While it ordinarily crystal- (o > 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 O = 2Fe 3 (SO 4 ) 3 (OH) 3 , or 6FeS0 4 + 30 + 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 2 (SO 4 ) 3 + H 2 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 ) 2 (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 704 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 : 1 es< ,0 ^0 H H HO HO \J ,OK ^OK Ferrous sulphate Potassium sulphate ( ) \ ~\TT = FeK s (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.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. 705 caustic potash containing ferric hydroxide in suspension : Fe(OH) s + 2KOH + 3C1 = K 3 FeO 4 + H,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. 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 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. 706 INORGANIC CHEMISTRY. 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 2 . When potassium sulphocyanate t 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. 58.74). 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, 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. COBALT. 707 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 oC! 3 + 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 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- 708 INORGANIC CHEMISTRY. 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 linn^eite. 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 Co(CN) 6 + H a 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 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. COMPOUNDS OF AMMONIA WITH SALTS OF COBALT. 709 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 a variety of complicated salts are formed. Among those which have been best studied are the chlorides, which may be briefly mentioned here as examples. When a solution of cobaltous chloride in ammonia is exposed to the air, the first product formed is one of the composition Co(NH 3 ) 3 Cl 3 + H 2 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-coboltic 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 ) 5 Cl 3 , and is known 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 3 ) 3 Cl s Co(NH 3 ) 4 Cl 3 Co(NH 3 ) 5 Cl 3 Co(NH 3 ) 6 Cl 3 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 an- alogous to the mercur-amnionium compounds (see page 627). Accordingly, it is usually represented as derived from three molecules of ammonium chloride by the re- placement of three hydrogen atoms by a trivalent atom of cobalt. It must be said, however, that our knowledge of the constitution of these compounds is very limited. When the 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- 710 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 of roseo- salts and luteo-salts we have examples as below : Roseo-salts. Luteo-salts. Co(NH 3 ) 5 01 3 Co(NH 3 ) 6 Cl 3 Co(NH 3 ) & (N0 3 ) 3 Co(NH 3 ) 6 (N0 3 ) 3 [Co(NH 3 ) 5 ],(S0 4 ) 3 [Co(NH 3 )J 2 (S0 4 ) 3 NICKEL, Ni (At. Wt. 58.56). G-eneral. 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 3J 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 ar& 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 always 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 when it contains a small quantity of magnesium or phosphorus it becomes very malleable. 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 sul- phuric acids, and readily in nitric acid. Like iron, it is magnetic. Alloys. Alloys of nickel are extensively used. Ar- gentan or German silver consists of copper, zinc, and nickel. COMPOUNDS OF NICKEL. 711 Various nickel alloys are used for making coins. The 5 and 3 cent pieces in the United States are made of an 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) 2 , is precipi- tated as a greenish-white substance. With an excess of potassium cyanide this forms the salt, Ni(CN),.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. 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 712 INORGANIC CHEMISTRY. 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 nitric 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 + 7KN0 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 is gray. CHAPTER XXXIV. ELEMENTS OF FAMILY VIII, SUB GROUP B ; RUTHENIUM RHODIUM PALLADIUM. ELEMENTS OF FAMILY VIII, SUB GROUP 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- markable table : Pe At. Wt. 55.88 Sp. Gr. 7.8 Bu At. Wt. 103.5 Co At. Wt. 58.74 Sp. Gr. 8.5 Rh At. Wt. 104.1 Ni At. Wt. 58.56 Sp. Gr. 8.8 Pd At. Wt. 106.2 Sp. Gr. 12.26 Sp. Gr. 12.1 Sp. Gr. 11.5 Os IT Pt At. Wt. 191 At. Wt. 192.5 At. Wt. 194.3 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 they 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 Dalladium and platinum. These relations will appear (713) 714 INORGANIG CHEMISTRY. more clearly if the formulas of some of the principal compounds of the elements under consideration are- placed together in a table. Ru and Os Rh and Ir Pd and Pt EuO 4 OsO 4 EhO 2 IrO 2 PdO 2 PtO 2 HEu0 4 Eh 2 3 Ir 2 O 8 PdO PtO H 2 Ku0 4 H 9 OsO 4 EhO IrO Pd 2 O Eu0 2 Os0 2 IrCl 4 PdCl 4 PtCl. Eu a O 8 Os 2 O 8 EhCl s IrCl, PdCl a PtCl, EuO OsO IrCl 3 EuCl 4 OsCl 4 EuCl 8 OsCl 8 EuCl 2 OsCl 3 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 IY 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. 715 be septivalent, sexivalent, quadrivalent, trivalent, and. bivalent. "When we pass to iron, however, we find that, the septivalence is gone. 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 YII 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 HKuO 4 , sexivalent in H 2 RuO 4 , quadrivalent in RuO 2> trivalent in Ru 2 O 3 , and bivalent in RuO. 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. 716 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 sancl 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, Ku (At. Wt. 103.5). Preparation. Kuthenium is obtained from the residue which is left undissolved when platinum-ore is treated with concentrated nitro-hydrochloric acid. RUTHENIUM OSMIUM. 717 Properties. When heated in oxygen it burns and forms the oxide, Hud),. It is insoluble in acids, and even in nitro-hydrochloric acid it is almost insoluble. Owing to its power to form salts of ruthenious acid, it is dis- solved 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 a O s , and EuO are not basic, and do not dissolve in acids. OSMIUM, Os (At. Wt. 191). 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- 718 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 dichloride, OsCl 2 , and the tetrachloride, 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. KHODIUM, Eh (At. Wt. 104.1). Rhodium has no acid properties, and does not form a peroxide corresponding to those of ruthenium and os- mium. On the other hand, its oxide, Rh 2 O 3 , is basic. The chloride KhCl 3 is readily formed, and it is doubt- ful whether the di- and tetrachlorides have been made. IRIDIUM, Ir (At. Wt. 192.5). 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, Ir01 3 , IRIDIUM. 719 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 is not dissolved by nitro-hydrochloric acid unless it is finely divided. 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 tetracMoride, IrCl 4 . When the solution of the tetra- chloride is heated it gives off chlorine, and the dicJdoride, 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,, or IrCl 4 .2MCl ; while those with the trichloride have the general formula M 3 IrCl 3 , or IrCl 3 .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. 720 INORGANIC CHEMISTRY. PALLADIUM, Pd (At. Wt. 106.2). 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 metal 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. 721 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 pcdladic 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 2 O. PLATINUM, Pt (At. Wt. 194.3). Preparation. A general idea of the method of pro^ cedure in extracting platinum from its ores was given on p. 716. 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. 719). 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 8 . 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. 722 INORaANIC 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 nitre-hydrochloric acid. Platinum vessels, further, should not be placed upon red-hot charcoal. Metallic salts which are easily re- duced should not be heated in platinum vessels, such as those of antimony and bismuth, 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. 723 Chlorides. Like palladium, platinum forms two chlo- rides, platinous chloride, PtCl a , 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 3 PtCl e , 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 n make it probable that the constitution of chlorplatinic acid should be represented by a similar roi Cl formula, Pt-^ (C1)H . Platinous chloride like platinic chloride combines with other chlorides to form double salts, the general formula of which is M,PtCl 4 , or PtCl 3 .2MCl. Cyanides. Platinum forms a number of beautiful double cyanides derived from an acid of the formula 724 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,Pt0 8 ; OH; OH 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. 725 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,(NH 3 ) a is formed. This is the starting-point for a series of compounds, as the bromide, PtBr a (NH 3 ) 2 ; the nitrate, Pt(NO 3 ) 9 (NH 3 ) a ; the sulphate, PtSO 4 (NH 3 ) a ; etc. There is another series beginning with the chlo- ride, PtCl 3 (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 2 (NH 3 ) 2 probably has the constitution Pt < 3 - Another com- pound of the same composition, but of the constitution X^TT \TTT O1 Pt< 3 3 , suggests itself, and there are two salts of this composition known. Although a great deal of work has been done on these platino-ammomum com- pounds, arid much interesting information in regard to them has been gained, the subject of their constitution is still in an unsatisfactory state. APPENDIX 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. (727) 72S 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 J- to -J 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 From a piece of 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 laboratory burner until it melts to- gether. 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 mercury (mer- curic oxide) to form a layer about 12 millimeters ( 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 familiar substance does it suggest? Introduce a splinter of wood upon the end of which there is a spark. What do you observe? Is there any dif- ference between the burning of the wood in the air and in the CHANGES EFFECTED BY AN ELECTRIC CURRENT. 729 tube ? What evidence is 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 current. 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 A. The platinum elec- trodes are now brought beneath the invert- ed tubes. The bubbles 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. Remove 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. 730 EXPERIMENTS TO ACCOMPANY CHAPTER 1. wood will take place more actively than it does in ordinary air, as may be shown by withdrawing it and again inserting it into 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 disulphide of carbon. 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 disulphide of carbon on a little powdered roll-sulphur in a dry test-tube. The sulphur dissolves. Treat iron filings in the VARIOUS EXAMPLES OF CHEMICAL ACTION. 731 same way. The iron does not dissolve. Now treat a small quantity of the mixture with bisulphide of carbon. 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 disul- phide of carbon 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 small 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 732 EXPERIMENTS TO ACCOMPANY CHAPTER L 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 of 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 substance 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 will be found not to present the regularities noticed in the original substance. It is much softer. It dissolves in water. It melts when heated in a tube. It does not give off a gas when treated with hydrochloric acid. When exposed to the air it soon 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 prop- erties. It might be shown that some of the hydrochloric acid also loses its properties. In place of the two we get a new VARIOUS EXAMPLES OF CHEMICAL ACTION. ?33 substance with entirely different properties. The two sub- stances have acted chemically upon each other and 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 smell, (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 734 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. PREPARATION OF OXYGEN. Experiment 14. Make some oxygen by heating to redness 10 to 15 grams (about half an ounce) of manganese dioxide in an iron tube closed at one end and connected at the other end by means of a cork with a bent-glass tube. Experiment 15. Make some oxygen by heating a few grams of mercuric oxide in a 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 com. capacity. By means of FIG. 17. a good-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 OASES. 735 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 oan be obtained by heating 12 grams of potassium chlorate. MEASUREMENT OF 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 -^ part. If, therefore, the volume of a gas at is V, at t its volume V will be F+ J_ F or F' This expression may also be written thus : V = F+0.00366/. V or V = V(l + 0.003660- From this we get the expression V v= 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 736 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 is 250 cc.; t, the temperature, is 15. Substituting these values in the equation F = 1 + 0.00366^ we have 250 F= 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' the vol- ume F', 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. 737 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 trough 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 FlG - 18 - measured is 75 cc. Call this V. What would be the actual volume V under the standard 760 mm. ? We have seen that VP= V'P'. Now, in this case P = 760, V = 75, and P ' = 660. There- fore, 760 F= 75 X 660, or F = 75 X 660 760 = 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 V V'P 760(1 + 0.003660' 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 738 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 V- ~ 760(1 + 0.00366^' in which the letters F, F', 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 FIG. 19. 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. 739 DETERMINATION OF THE AMOUNT OF OXYGEN LIBERATED WHEN A KNOWN WEIGHT OF POTASSIUM CHLORATE is DECOMPOSED BY HEAT. Experiment 17. In order 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 glass tube about 10 cm. long and 8 to 10 mm. internal diameter, closed at one end, weigh out on a chemical balance about 0.2 gram dry potassium chlorate, first weighing the tube empty. Introduce just above the potassium chlorate a plug of glass-wool, then soften the tube in a flame, and draw it out so that it has the form shown in Fig. 21, the plug of glass- wool being at the constricted part of the tube. 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 th3 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 Aveight of a liter or 1000 cc. of oxygen at and 760 mm. pressure is 1.4298 grams. Knowing the volume of oxygen obtained, calculate the weight of this volume. 740 EXPERIMENTS TO ACCOMPANY CHAPTER II. Eemove 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 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 -f c V) + e. should be equal to b a, 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 Fl - 22. the water into the upper part of the vessel after having screwed on the cap d. When it is PHYSICAL PROPERTIES OF OXYGEN. 741 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- 742 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 deflagra ting-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 that 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. 743 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. OXYGEX 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^ 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 i 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 eenerally 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. 744 EXPERIMENTS TO ACCOMPANY CHAPTER II. FIG. 25. 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. 26. 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. 745 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 2J 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 2 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 30. Repeat Experiment 3 and examine the gases. 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 name 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. 746 EXPERIMENTS TO ACCOMPANY CHAPTER III. gas will rise in an inverted vessel filled with water. For 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. 747 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 Wolff's 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. 748 EXPERIMENTS TO ACCOMPANY CHAPTER III. 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. 749 " 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 effect 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 3 S0 4 to 4 of H,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. 750 EXPERIMENTS TO ACCOMPANY CHAPTER III. A liter of hydrogen at and 760 mm. weighs 0.089578 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 2 . That between aluminium and hydrochloric acid is represented by this equation : Al + 3HC1 = A1C1, + 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. 751 FIG 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 th,e mouth upward and uncovered. In a short time examine J&\e 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. Experiment 40. 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 e 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 Wolff's 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 752 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 bubbles will be seen to appear rapidly at the end of the long straight tube, and the liquid will rise in the other tube, 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- Fio. 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. 753 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. Wliat conclu- sions are justified by the last two experi- ments ? FIG. 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? Eemove 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 C 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 :s a brisk evolution) heat 754 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 FIG. 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. 755 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 ^and tube F again. Let x = weight of tube E + cupric oxide before the ex- periment; y = weight of tube E -f- 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 I a = weight of water formed. If the experiment is properly performed, it will be found that the ratio -7 - 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 756 EXPERIMENTS TO ACCOMPANY CHAPTER IV. down. A convenient arrangement of the apparatus is shown in Fig. 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 names 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. 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. EUDIONETEIC EXPERIMENTS. 757 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. Put them in a dry tube and heat them gently, when it will be observed that water condenses in the upper part of 758 EXPERIMENTS TO ACCOMPANY CHAPTER V. the tube, while the 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 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 60. 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 DUMAS. 759 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 when it is required to determine the specific gravity of the vapor of a substance which boils at a high temper- ature. 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 Fig. 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 760 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 n _ (-g, ~ B +p) . (1 + 0.00366 . '*,) . 760 v . 0.001293 . A, in which B = weight of balloon at t and li mm. ; B l = " " " with vapor, at t and h l 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 FTO 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. 761 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: (1 + 0.00366 . t) . 760 ' (B - iv)V . 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: n g (l + 0.00366 . t) 587,780 (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 762 EXPERIMENTS TO ACCOMPANY CHAPTJSB 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. Convince 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. 763 The best way to make chlorine is the following : Mix 5 parts coarsely granulated manganese dioxide and 5 parts coarse- 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 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 10 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 w r ith 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 flame is noticed, 764 EXPERIMENTS TO ACCOMPANY CHAPTER VIII. 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. CHLORIDE 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. 765 PREPARATION OF HYDROCHLORIC ACID. Experiment 74. Arrange an apparatus as shown in Fig. 45. 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 three Wolffs 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. 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 SNaCl + H 2 S0 4 = Na,S0 4 + 2HC1 766 EXPERIMENTS TO ACCOMPANY CHAPTER IX. is now complete. After the flask has cooled down, pour wa- ter on the contents; and when the substance is dissolved fil- ter 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 placing it upon folds of filter-pa- per. 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 ac\d taste, the two terms being practically synony- mous. Add a drop or two of a solution of Uue litmus, or put into it a piece of paper colored Hue 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 AND 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. 76? 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. Becrystallize 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 ? FIG. 46. PERCHLORIC ACID. Experiment 77. Make potassium perchlorate a^ 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 768 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 OP 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 gram to 200 cc. of water). Measure off about 20 cc. of each of the acid solutions. Add a few drops of a solution of blue litmus. Gradually add to each of the meas- ured quantities 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- Em. 47. STUDY OF THE PRODUCTS FORMED. 769 actiy 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 PRODUCTS FORMED. 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 770 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 or 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: HYDROBROXIC ACID. 971 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 con- taining red phosphorus in order to free the hydrobromic acid from bromine, which to some extent passes over with it. Col- lect 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 hydro- chloric acid, and fill another with chlorine. AVhile covered with glass plates bring their mouths together. Then with- draw 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 ? 772 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 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 crys- tals 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 STAKCH-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. 773 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, S0 2 . 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 completely dissolve in carbon disulphide. 774 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 Fio. 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. 775 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 ? Kepeat 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 AsCl 3 . The cor- responding sulphides are represented by the formulas PbS, CdS, 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, I, 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 776 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. 777' 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- Fio. 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 REDUCING AGENT. Experiment 112. To a dilute solution of potassium iodide in a test-tube gradually add chlorine water until the solution 778 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 gase& O- 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. 779 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 780 EXPERIMENTS TO ACCOMPANY CHAPTER XV. height of the barometer. Keduce 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) 2 , 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. 781 The tube D, which should be filled with water and provided with d 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- FIG. 54. cylinders. When the bottle is filled with air pour some water in again so that it is about a quarter full. Draw this water off as before. 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 carbon 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- 782 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 A 9 when the air will be forced through B and C and into D. But in passing through B and C the moisture contained in it Fia. 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 oil 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. 783 After the air is driven out, the gas will be completely ab- sorbed by the water in the first Wolff's flask. Disconnect the delivery-tube from the series of Wolff's flasks, and connect with another tube bent upward. Collect .some of the escaping gas by displacement of -air, placing the vessel with the mouth doivn- ward, as the gas is much lighter than air. The arrangement is shown in Fig. 56. The vessel in which the gas is collected should bo 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. 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. 734: EXPERIMENTS TO ACCOMPANY CHAPTER XVI. Experiment 130. Fill a 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. 785 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. Recrystallize 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. 786 EXPERIMENTS TO ACCOMPANY CHAPTER 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. NITRIC 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. 787 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 1 3. ) METALS DISSOLVE IN NITRIC ACID, FORMING KITRATES. 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. 788 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. 789 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 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 vessel stand. Almost immediately the color will disappear, showing that the nitrogen peroxide formed is decomposed. Again admit air, 790 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 By 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- FIG. 61. PHOSPHINE. 791 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 flask B put about 5 grams caustic potash dissolved FIG. 62. in 10-15 cc. water, and ivlien 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 792 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. 63. dilute sulphuric acid on it. When the air is all out of the vessel and the hydrogen is lighted, add slowly a little of a so- lution 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 experiment. ) MARSH '8 TEST FOR ARSENIC ANTIMONY, ETC. 793 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 Bunsen burner the tube through which the gas is passing, which should be of hard glass. Just in front of 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 794 EXPERIMENTS TO ACCOMPANY CHAPTER XVII. during the operation. This is best accomplished by fitting Fio. 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, tfye two form a gas-tight joint. This is shown in Fig. 65. After the FIG. 65. PHOSPHOR US PENT A CHL OBIDE. 795. 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 XVHL 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 7% EXPERIMENTS TO ACCOMPANY CHAPTER XVIII. 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 m 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- 10.68. i u tj on O f antimonic 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. 797 BASIC NITRATES OF BISMUTH. Experiment 169. Dissolve a little bismuth in nitric acid and evaporate. Add water. BORON. Experiment 17O. 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. 798 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 ; C 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 tube to which is fitted an outlet tube, as shown in Fig. 70. 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. Repeat 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.* 799 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 ~~ , . , , FIG. 71. 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. 800 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 carbonate is in solution or not. Experiment 182. Pass car- bon dioxide into 50 to 100 cc. clear FlG>72> 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. Heat the contents of the flask gently. Collect some of the gas over water. Set fire to some, and notice the characteristic blue flame. Put a live mouse in a vessel con- taining a mixture of about equal parts of carbon monoxide and air. It will die unless taken out. CARBON MONOXIDE is A GOOD EEDUCING 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. 801 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. KINDLIXG 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. 802 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. Flo 74 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. n 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 l 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. 803 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 a is the oxidizing flame ; the part marked r is the reducing flattie. 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 Fig. 77. Heat the * Pieces of charcoal prepared for blow-pipe work can be bought from dealers in chemical apparatus, at small cost. 804 EXPERIMENTS TO ACCOMPANY CHAPTER XX2I. 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. 805 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 TETRAFLUORIDE 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- Fio. 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 fluosiKoic 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 806 EXPERIMENTS TO ACCOMPANY CHAPTER XXIV. 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. Mix together some fine sand and about four times its weight of a mixture of potassium and sodium carbonates. Heat in a platinum crucible in the flame of the blast-lamp until the mass is thoroughly melted. 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, NaCJ. 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. 807 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 -f NaCl = AgCl + 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) 2 . 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, rnercurous chloride, is insoluble. 808 EXPERIMENTS 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 + 3NaCl. 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 ; Jead 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. 809 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, Bad,, the insoluble salt barium sulphate, BaS0 4 , and the soluble salt iron chloride, FeCl 2 , are formed : FeS0 4 + BaCl 2 = FeCl a + 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- ble sulphide when added to a solution containing copper gives a black precipitate of copper sulphide. Try this ; also try the action on 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,, BaC0 3 , and CaC0 3 . 810 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. 811 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 oil 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 AND 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. Xowtake the crucible out of the furnace, and stir the liquid in it with the aid of a clay pipe-stem. This causes the 812 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- COS S WORKS ON SCIENCE. Allen's Laboratory Exercises in Elementary Physics. By CHARLES R. ALLEN, Instructor in the New Bedford, Mass., High School. Pupils' Edition: x -\- 209 pp. I2mo. 8oc. net. 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