tm m LIBRARY CALIFORNIA COLLEGE PHARMACY COLLEGE OF PHARMACY CHEMISTRY DEPARTMENT OUTLINES OF CHEMISTRY THE MACMILLAN COMPANY NEW YORK BOSTON CHICAGO SAN FRANCISCO MACMILLAN & CO., LIMITED LONDON BOMBAY CALCUTTA MELBOURNE THE MACMILLAN CO. OF CANADA, LTD TORONTO OUTLINES OF CHEMISTRY A TEXT-BOOK FOR COLLEGE STUDENTS BY LOUIS KAHLENBERG, PH.D. PROFESSOR OF CHEMISTRY AND DIRECTOR OF THE COUR8B IX CHEMISTRY IN THE UNIVERSITY OF WISCONSIN REVISED California College of Pharmac; gctk THE MACMILLAN COMPANY 1920 -r; iijlit* re* All rialit* reserved COPYRIGHT, 1909 AND 1915, BY THE MACMILLAN COMPANY. Set up and electrotyped. Published September, 1909. NortoootJ J. 8. Gushing Co. Berwick & Smith Co. Norwood, Mass., U.S.A. PREFACE TO THE FIRST EDITION THIS book is intended to represent one year's work of chemistry in college. It should be used in connection with a course of experimental lectures and laboratory exercises. The matter has been selected so as to meet the needs of those that can devote but one year to the study of chemistry, and also to serve as a suitable basis for future work in the case of students who desire to pursue the subject further. In writing the book, the author has naturally had in mind the needs of his own students, over six hundred in number, who are pre- paring for careers in chemistry, pharmacy, medicine, engineer- ing, or agriculture, or who desire a course in chemistry for work in other natural sciences or as a means of general culture. In the first five chapters, experimental work has been placed in the foreground, and all reference to atomic and molecular theo- ries has been purposely avoided in order that the student may properly be impressed with the fundamental facts and laws, which are independent of the theories, though they serve as a foundation for the latter. In the sixth chapter, these funda- mental laws are then reviewed, and the atomic and molecular theories are presented as views growing out of the experimental facts. The nomenclature is then also introduced, and the reactions which so far have been written in words are expressed by means of chemical symbols. This offers an excellent oppor- tunity for reviewing the experimental work of the foregoing chapters. While the teacher is somewhat inconvenienced by thus postponing the introduction of the atomic theory and the use of formulation till the student has at least a fair stock of carefully selected facts upon which to found the theory, it really pays to make the exertion, for thus greater interest is created and the student sees the facts and theoretical viewa 42130 vi PREFACE in their proper relations. He becomes a clear, logical thinker^ and does not look upon the atomic and molecular theories as something arbitrary, metaphysical, and well-nigh incomprehen- sible. The method here adopted is not new. It is essentially the same in principle as that followed by Bunseii and many other successful teachers of chemistry. Throughout the book, the endeavor has been to convey the salient facts in as simple and direct a manner as possible, developing cardinal principles, and carefully keeping the dis- tinction between facts and theories in mind. The aim has been to enlist the interest of the student in the study of chem- istry, and to this end the historical development of certain aspects of the subject has been presented as far as space would permit. The most important technical applications and processes have constantly been emphasized, though they have been in- troduced in connection with the description of the various elements and compounds rather than as special chapters. On the other hand, it has been thought best to treat the subjects of thermochemistry and solutions and electrolysis in special chapters, after a sufficient number of fundamental facts have been acquired by the student, so that he is in a position to comprehend the more difficult relationships which these topics involve. Only the essential parts of chemical theory which can be comprehended by college students who are beginning the study of chemistry have been presented. My own experience would indicate that fully as much has been given as they can well digest at this stage of their work. In touching upon contro- verted points, the aim has been to present both sides of the question involved. I have felt that the teacher should not entirely avoid mooted questions even during the first year of work in chemistry, for by so doing the impression is conveyed that all matters are in a settled state, and thus a powerful stimulus toward further study and inquiry is lost. On the whole, however, the presentation of the subject has been along rather well established, conservative lines. The dominant idea has been to select with care what the student needs, what he can reasonably be asked to comprehend, at his stage of advance- ment, and to present this in a clear, simple, and direct manner, PREFACE v i\ taking the trouble to repeat and to emphasize here and there in order to secure the desired end. My best thanks are due to Dr. J. H. Walton for suggestions and reading of proof, also to Messrs. C. W. Hill, D. Klein, F. C. Krauskopf, and W. G. Wilcox for reading proof sheets of some of the chapters. Additional suggestions or corrections to be used in preparing further editions will be welcomed from others. LOUIS KAHLENBERG. MADISON, WISCONSIN, June 8, 1909. PREFACE TO THE SECOND EDITION IN preparing this new edition, the entire book has been gone over with special care so as to bring the subject matter up to date and improve the presentation wherever possible. As a special aid to the student, a set of review questions has been added to each chapter. These questions are intended to cover the matter presented in the respective chapters, but they by no means exhaust the subject. On the contrary, the questions will doubtless serve to suggest many others to the teacher. The wide popularity which the first edition of this book has enjoyed among students and teachers has been gratifying to the publishers and the author, and it is hoped that the new edition will similarly commend itself to a still larger circle. The aim has been to keep the presentation as simple, clear, and direct as possible, giving as much theory as is necessary for a thorough comprehension of the important facts to be inculcated, yet always bearing in mind that the book is essentially for first year college students whose interest must be awakened, and who must not be overwhelmed with too much detailed and ab- struse matter. To the many kind friends who have so gener- ously aided by means of helpful suggestions the author desires to express his best thanks. LOUIS KAHLENBERG. MADISON, WISCONSIN, February, 1915. CONTENTS CHAPTER I THE SCOPE OP CHEMISTRY AND ITS RELATIONS TO OTHER SCIENCES CHEMICAL CHANGE, ELEMENTS, AND COMPOUNDS PAGE Physical and Chemical Changes Definite Proportions Solutions and Chemical Compounds Chemical Elements Compounds Types of Chemical Change Conservation of Mass Conserva- tion of Energy Cause of Chemical Change Chemical Affinity Factors affecting Chemical Change . . . . . 1 CHAPTER II HYDROGEN History Occurrence Preparation Properties Uses Hydrogen Equivalents of the Metals 13 CHAPTER m OXYGEN History Occurrence Preparation Properties Combustion in the Air Kindling Temperature and Temperature of Combustion Heat of Combustion Different Stages of Oxidation Law of Multiple Proportions Role of Oxygen in Respiration Oxy- hydrogen Blowpipe Detonating Gas Combustion of Oxygen in Hydrogen Earlier Views of Combustion . . . .26 CHAPTER IV WATER Occurrence Preparation Natural Waters Potable Water Min- eral Water Composition Gay-Lussac's Law of Combination of Gases by Volume Properties of Water Super-cooled Water Change of Freezing-point with Pressure Principle of Le Cha- telier Crystalline Nature of Ice Compounds with Water Water as a Solvent , . 38 ix X CONTENTS CHAPTER V HYDROCHLORIC ACID AND CHLORINE MM Preparation and Properties of Hydrochloric Acid -Composition and Chemical Behavior of Hydrochloric Acid Occurrence, History, and Properties of Chlorine Uses of Chlorine Some Compounds of Chlorine with Oxygen Law of Reciprocal Proportions . . 51 CHAPTER VI THE LAWS OF COMBINING WEIGHTS AND COMBINING VOLUMES AND THE ATOMIC AND MOLECULAR THEORIES Retrospect Laws of Definite, Multiple, and Reciprocal Proportions Combining Weights and Chemical Equivalents Chemical Sym- bols Atomic Theory of Matter Difference between Theory and Law Law of Combination of Gases by Volume Avogadro's Hypothesis Molecular Weight Determinations Determination of Atomic Weights Law of Dulong and Petit Other Methods of Choosing Atomic Weights from the Combining Weights Law of Isomorphism Table of Atomic Weights Interpreta- tion of a Chemical Formula Valence and Structural Formulae Nomenclature Chemical Equations Retrospect Phe- nomena of the Nascent State ....... 60 CHAPTER VII OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE History, Occurrence, and Preparation of Ozone Relation between Ozone and Oxygen A llotropy -Properties of Ozone History, Occurrence, and Preparation of Hydrogen Peroxide Properties of Hydrogen Peroxide Formula of Hydrogen Peroxide Uses of Hydrogen Peroxide Ozonic Acid ...... 92 CHAPTER THE HALOGENS The Halogen Family Compounds of Chlorine with Oxygen Hypo- chlorous Acid and Hypochlorites Chloric Acid and Chlorates Perchloric Acid and Perchlorates Nomenclature and General Relations Occurrence, Preparation, and Properties of Fluorine Hydrofluoric Acid Occurrence, Preparation, and Properties of Bromine Hydrobromic Acid Oxy-acids of Bromine Bromic Acid and Bromates Uses of Bromine and its Compounds History and Occurrence of Iodine Preparation of Iodine CONTENTS xi PAGB Properties of Iodine Uses of Iodine Hydriodic Acid Oxide of Iodine Oxy-acids of Iodine Compounds of the Halogens with Each Other General Relations of the Halogens to One Another 101 CHAPTER IX ACIDS, BASES, SALTS, HYDROLYSIS, MASS ACTION, AND CHEMICAL EQUILIBRIUM Acids Bases Salts Older View of the Process of Salt Formation Acid- and Base-forming Elements Other Views of Solutions of Acids, Bases, and Salts Basicity of Acids Acid Salts Acidity of Bases Basic Salts Normal Salts Acidirnetry and Alkalimetry Indicators Hydrolysis Mass Action Chemi- cal Equilibrium Additional Illustrations of Chemical Equi- librium and the Operation of the Law of Mass Action Strength of Acids and Bases . 127 CHAPTER X NITROGEN, THE ATMOSPHERE, AND THE ELEMENTS OF THE HELIUM GROUP History and Occurrence of Nitrogen Preparation and Properties of Nitrogen The Air The Elements of the Helium Group . . 146 CHAPTER XI COMPOUNDS OF NITROGEN WITH HYDROGEN AND WITH THE HALOGENS History and Occurrence of Ammonia Preparation and Properties of Ammonia Hydrazine Hydroxylamine Hydrazoic Acid Compounds of Nitrogen with the Halogens 158 CHAPTER XII OXY-ACIDS AND OXIDES OF NITROGEN History, Occurrence, and Preparation of Nitric Acid Properties of Nitric Acid Nitrogen Pentoxide Nitric Oxide Nitrogen Dioxide and Tetroxide Nitrous Acid Nitrogen Trioxide Hyponitrous Acid Nitrous Oxide General Considerations , 170 CHAPTER XIII SULPHUR, SELENIUM, AND TELLURIUM Occurrence and Preparation of Sulphur Properties of Sulphur Uses of Sulphur Crystals and Crystal Systems Hydrogen Xii CONTENTS PAGE Sulphide Poly-sulphides and Hydrogen Persulphide Compari- son of Hydrogen Sulphide with Water Compounds of Sulphur with the Halogens Sulphur Dioxide and Sulphurous Acid Sulphur Sesquioxide Sulphur Trioxide and the Contact Process of making Sulphuric Acid Sulphuric Acid and the Lead Cham- ber Process Properties of Sulphuric Acid Hydrates of Sul- phuric Acid - Pyrosulphuric Acid Thiosulphates Persulphates Polythionic Acids Thionyl Chloride Sulphuryl Chloride, Selenium Compounds of Selenium Tellurium Compounds of Tellurium General Considerations 185 CHAPTER XIV CARBON AND SOME OF ITS TYPICAL COMPOUNDS Occurrence and Allotropic Forms of Carbon Chemical Behavior of Carbon Carbon Dioxide Properties of Carbon Dioxide Physiological Effects of Carbon Dioxide Relations of Carbon Dioxide to Plant and Animal Life Early Work on Carbon Dioxide Carbon Monoxide Properties of Carbon Monoxide Physiological Effects of Carbon Monoxide Carbon Bisulphide Cyanogen Hydrocyanic Acid Cyanates and Sulphocyanates 221 CHAPTER XV HYDROCARBONS AND ADDITIONAL COMPOUNDS OF CARBON Hydrocarbons General Behavior of Hydrocarbons Halogen Sub- stitution Products Alcohols Phenols Aldehydes Organic Acids Esters Ethers Ketones Carbohydrates Fermen- tation and Enzymes Starch and Dextrine Cellulose Nitro- benzene, Aniline, and Coal Tar Dyes Alkaloids Proteins . 244 CHAPTER XVI ILLUMINATING GAS AND FLAMES Illuminating Gas Flame Luminosity of Flame Structure of Flame Davy Safety Lamp 279 CHAPTER XVn THERMOCHEMISTRY General Remarks Calorimeters Laws of Thermochemistry Thermochemical Equations Thermochemical Data Tables Uses of Thermochemical Data 29 CONTENTS xiii CHAPTER XVIH SILICON AND BORON AND THEIR IMPORTANT COMPOUNDS PAQH Occurrence, Preparation, and Properties of Silicon Silicon Dioxide Silicic Acids Action of Water on Silicates Decomposition of Silicates in the Laboratory Hydrogen Silicide Compounds of Silicon with the Halogens Esters of Silicic Acid Silicon Carbide Titanium Zirconium Thorium Occurrence, Preparation, and Properties of Boron Boric Acid and its Salts Other Compounds of Boron 306 CHAPTER XIX PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH Occurrence and Preparation of Phosphorus Properties and Allo- tropic Forms of Phosphorus Uses of Phosphorus, Matches Compounds of Phosphorus with the Halogens Oxides and Acids of Phosphorus Formulae of the Acids of Phosphorus Com- pounds of Phosphorus with Sulphur Occurrence, Preparation, and Properties of Arsenic Arsine Compounds of Arsenic with the Halogens Oxides and Oxy-acids of Arsenic Occurrence, Preparation, and Properties of Antimony Stibine Compounds of Antimony and Sulphur Occurrence, Preparation, and Prop- erties of Bismuth Halogen Compounds of Bismuth Oxides of Bismuth Bismuth Salts of Oxy-acids Bismuth Trisulphide General Considerations of the Group Vanadium, Columbium, and Tantalum 321 CHAPTER XX CLASSIFICATION OF THE ELEMENTS THE PERIODIC SYSTEM . . 355 CHAPTER XXI , THE ALKALI METALS Occurrence, Preparation, and Properties of Potassium Potassium Hydride Compounds of Potassium with the Halogens Potas- sium Hydroxide Potassium Oxide Potassium Chlorate Po- tassium Nitrate Potassium Cyanide Potassium Carbonate Potassium Silicate Potassium Fluosilicate Potassium Phos- phates Potassium Sulphate Potassium Sulphite Sulphides of Potassium Tests for Potassium Rubidium and Caesium Occurrence, Preparation, and Properties of Sodium Sodium Chloride Oxides and Hydroxides- of Sodium Sodium Carbon- ate Sodium Nitrate Phosphates of Sodium Sodium Sul- xiv CONTENTS phate Sodium Sulphite Sodium Thiosulphate Sodium Sili- cate Sodium Cyanide Sodium Borate Lithium and its Compounds The Alkali Metals as a Group Spectrum Analysis Ammonium Salts Detection of Ammonium Salts . . . 362 CHAPTER XXII THE ALKALINE EARTH METALS Occurrence, Preparation, and Properties of Calcium Calcium Oxide Cement Calcium Sulphate Calcium Sulphite Calcium Sulphide Calcium Fluoride' Calcium Chloride Bleaching Powder Calcium Phosphate Calcium Carbide Calcium Phosphide Calcium Cyanamide Calcium Silicide Calcium Silicate Glass Occurrence, Preparation, and Properties of Strontium Strontium Compounds Occurrence, Preparation, and Properties of Barium Compounds of Barium Detection of the Alkaline Earth Metals Radium and Radio-activity . . 394 CHAPTER XXIII THE METALS OF THE MAGNESIUM GROUP Glucinum Occurrence, Preparation, and Properties of Magnesium Magnesium Oxide Magnesium Carbonate Magnesium Chloride Magnesium Sulphate Magnesium Phosphates Magnesium Ammonium Arsenate Tests for Magnesium Occurrence, Prepa- ration, and Properties of Zinc Zinc Oxide Zinc Carbonate Zinc Chloride Zinc Sulphate Zinc Sulphide Analytical Tests for Zinc Salts Occurrence, Preparation, and Properties of Cadmium Cadmium Compounds Occurrence, Preparation, and Properties of Mercury Amalgams Compounds of Mercury Oxides of Mercury Halides of Mercury Mercuric Cyanide Nitrates of Mercury Mercuric Fulminate Sulphates of Mer- cury Mercuric Sulphide Compounds of Mercury Salts with Ammonia Physiological Properties of Mercury Compounds Tests for Mercury General Remarks 414 CHAPTER XXIV SOLUTIONS, ELECTROLYSIS, AND ELECTRO-CHEMICAL THEORIES Nature and Kinds of Solutions Absorption of Gases by Liquids Solutions of Liquids in Liquids Solutions of Solids in Liquids Degrees of Saturation Solid Solutions Precipitation Col- loidal Solutions Boiling Points of Solutions Use of Boiling Points of Solutions in Molecular Weight Determinations The Freezing Points of Solutions Discussion of Molecular Weights CONTENTS XV PAGE Determined in Solutions Osmosis and Osmotic Pressure Elec- trolysis Electrolytic Theories Electric Batteries Electro- chemical Series of the Metals 432 CHAPTER XXV COPPER, SILVER, AND GOLD Occurrence, Metallurgy, and Properties of Copper Alloys of Copper Oxides of Copper Halides of Copper Cyanides of Copper Copper Salts of Oxy-acids Sulphides of Copper Analytical Tests for Copper Occurrence, Metallurgy, and Properties of Silver Oxides of Silver Halides of Silver Uses of Silver Halides in Photography Silver Nitrate Silver Nitrite Silver Sulphate Silver Carbonate Silver Phosphate Silver Sul- phide Silver Cyanide Silver Plating Silver Fulminate Analytical Tests for Silver Occurrence, Metallurgy, and Proper- ties of Gold Gold Alloys Compounds of Gold Analytical Tests for Gold '.460 CHAPTER XXVI THE METALS OF THE EARTHS Occurrence, Preparation, and Properties of Aluminum Uses of Alu- minum Aluminum Oxide Aluminum Hydroxide Aluminum Chloride Aluminum Sulphide Aluminum Sulphate Alums Aluminum Silicates Analytical Tests for Aluminum Gal- lium Indium Thallium and its Compounds The Rare-Earth Elements . . . . . . . . . . . .482 CHAPTER XXVII LEAD AND TIN Germanium Occurrence, Metallurgy, and Properties of Tin Uses of Tin Chlorides of Tin Oxides of Tin Sulphides of Tin Analytical Tests for Tin Occurrence, Metallurgy, and Proper- ties of Lead Uses of Lead Oxides of Lead Halides of Lead Lead Nitrate Lead Acetate Lead Sulphate Lead Arse- nate Lead Carbonate Analytical Tests for Lead . . . 498 CHAPTER XXVIII CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM Occurrence, Preparation, and Properties of Chromium Chromic Oxides and Hydroxides Chromous Compounds Chromic Salts XVI CONTENTS PAGI Chromates, Bichromates, and Chromium Trioxide Chromyl Chloride Analytical Tests for Chromium Molybdenum Tungsten Uranium 512 CHAPTER XXIX MANGANESE Occurrence, Preparation, and Properties Oxides Salts of Manga- nese Manganates and Permanganates Uses of Permanganates Analytical Tests for Manganese 523 CHAPTER XXX IRON, NICKEL, AND COBALT Occurrence of Iron Metallurgy of Iron Cast Iron Wrought Iron Steel Properties of Iron Oxides and Hydroxides of Iron Chlorides of Iron Sulphides and Sulphates of Iron Ferrous Carbonate Cyanides of Iron Blue Printing Other Com- pounds of Iron Analytical Tests for Iron Occurrence, Prepa- ration, and Properties of Nickel Nickel Oxides and Hydroxides Salts of Nickel Nickel Carbonyl Occurrence, Preparation, and Properties of Cobalt Oxides and Hydroxides of Cobalt Other Cobalt Compounds Analytical Tests for Cobalt and Nickel 531 CHAPTER XXXI THE METALS OF THE PLATINUM FAMILY Occurrence Extraction of Platinum from the Ores Ruthenium Rhodium Palladium Osmium Iridium Platinum Ana- lytical Tests for Platinum . . 552 INDMX 559 LIST OF ILLUSTRATIONS no. PAG 1. Tube used in demonstrating that weight remains constant during chemical changes .......... 10 2. Electrolysis of water . . . . . . . . .14' 3. Preparation of hydrogen by action of sodium on water ... 15 4. Preparation of hydrogen by action of steam on heated iron . . 15 5. Preparation of hydrogen by action of sulphuric acid on zinc . . 16 6. Transferring hydrogen from one jar to another .... 17 7. Diffusion of hydrogen ......... 18 8. Formation of water when hydrogen burns in the air . . . 19 9. Singing flame 20 10. A candle will not burn in hydrogen 20 11. Oxidation of copper when heated in the air ..... 21 12. Reduction of hot copper oxide by hydrogen 21 13. Cylinder for compressed gases 22 14. Apparatus for determining hydrogen equivalents of metals . . 22 15. Burning of an iron wire in oxygen ....... 28 16. Burning of sulphur in oxygen ........ 28 17. Oxyhydrogen blowpipe 33 18. Combustion of oxygen in hydrogen 34 19. Lavoisier's apparatus to show that mercury unites with oxygen when calcined 36 20. Distillation 39 21. Demonstration of volumetric relations between oxygen, hydrogen, and steam 43 22. Desiccator 48 23. Composition of hydrochloric acid gas . . . ... .52 24. Electrolysis of hydrochloric acid 53 25. Synthesis of hydrochloric acid gas by volume . . . . .54 26. Burning of arsenic in chlorine 56 27. Action of chlorine on water in sunlight 57 28. Ozone apparatus 92 29. Preparation of fluorine . . . 108 30. Preparation of hydrobromic acid 113 31. Sublimation of iodine in the laboratory . . . . 117 32. Sublimation of iodine on commercial scale 118 33. Titration 136 34. Preparation of nitrogen from the nir 147 35. Oxidation of nitrogen by means of the electric spark . . .150 xvii xviii LIST OF ILLUSTRATIONS FIG. 36. Volumetric composition of ammonia gas ..... 160 37. Decomposition of ammonia by the electric spark .... 161 38. Burning ammonia mixed with oxygen ...... 162 39. Oxidation of ammonia by use of a platinum spiral . . . 162 40. Preparation of nitric acid ........ 171 41. Heating sodium in nitric oxide ....... 176 42. Commercial distillation of .sulphur ...... 186 43. Crystal of rhombic sulphur . . . . . . . . 187 44 to 54. Crystals of the isometric system ...... 190 55 to 60. Crystals of the tetragonal system ...... 191 61 'to 67. Crystals of the hexagonal system ...... 192 68 to 71. Crystals of the orthorhombic system ..... 193 72 to 74. Crystals of the monoclinic system ..... 193 75 to 76. Crystals of the triclinic system . ..... 194 77. When sulphur burns in oxygen the volume remains unchanged . 199 78. Bleaching of flowers by means of sulphur dioxide . . . 200 79. Sulphuric acid by the contact process . . . . . . 202 80. Diagram of a sulphuric acid factory ...... 206 81 and 82. Crystal forms of diamond ....... 221 83. Acheson graphite furnace ........ 224 84. Typical arc furnace for experimental work ..... 224 85. Absorption of ammonia gas by charcoal ..... 225 86. Kipp apparatus .......... 230 87. Siphoning carbon dioxide from one jar to another . . . 232 88. Taylor's carbon bisulphide furnace ..... . 239 89. Yeast cells ........... 250 90. Acetic acid organisms ... ...... 255 91. Lactic acid organisms ......... 258 92. Formulae of dextro and Isevo lactic acids : 260 93. Polariscope ........... 261 94. Crystals of dextro and Isevo tartaric acid ..... 263 95. Grains of potato starch . . ....... 272 96. Grains of wheat starch . . . . ..... 272 97. Grains of corn starch ......... 272 98. Potato starch grains in polarized light ...... 272 99. Manufacture of coal gas ........ 280 100. Gases burn in the flame of a candle ...... 282 101 to 103. Demonstration of the reverse flame . . . . . 283 104. Burning oxygen in coal gas . . . . . . . . 284 105. Enriching carbon monoxide gas ....... 284 106. Principle of the Bunsen burner ....... 285 107. Zones of the flame of a candle ....... 287 108 and 109. A flame will not pass through a wire gauze . . . 287 110. Davy safety lamp .......... 288 111. Calorimetric apparatus ......... 291 112. Combustion bomb and calorimeter ( ..... 292 113. Right and left quartz crystals ....... 308 LIST OF ILLUSTRATIONS x ix FIG. PAGE 114. Crystal of tridymite 308 115. Dialyser 310 116. Making hydrofluosilicic acid 314 117. Retorts for making phosphorus . . . . . . . 322 118. Electric furnace for making phosphorus 323 119. Making phosphine from phosphorus and caustic alkali . . 326 120. Phosphine from calcium phosphide 327 121. Marsh test for arsenic ......... 336 122. Curve of atomic weights and atomic volumes. (L. Meyer.) . 359 123. Hopper-shaped crystal of sodium chloride 375 124. Acker process of making caustic soda 376 125. Salt cake furnace, Le Blanc soda process 377 126. Revolving black ash furnace ........ 378 127. Solubility curve of sodium sulphate ...... 381 128. Spectroscope 385 129. Spectra of some common elements 386 130. Tube for examining spectra of gases 387 131. Spectra of gases 388 132. Absorption of spectrum of blood ....... 389 133. Making metallic calcium ........ 395 134. Common limekiln . . .397 135. Glass pots, open and closed form 404 136. Solubility curve of magnesium chloride ..... 416 137. Iron flask for shipping mercury . . . . . . . 423 138. Solubility curves of various salts ....... 435 139. Making colloidal silver 439 140 and 141. Explanation of osmosis 441 142. Demonstration of osmotic pressure 443 143. Simple osmometer 443 144. Pfeft'er's osmotic apparatus . . . . . . . . 444 145. Grotthus's theory of electrolysis . 449 146. Electrolysis according to the theory of electrolytic dissociation . 451 147. An electric battery .......... 455 148. Measuring the voltage of a cell ....... 455 149. Gravity battery 456 150. Electrolytic production of aluminum . . . . . 483 151. Blastfurnace 532 152. Bessemer converter 536 153. Solubility curve of ferric chloride 540 154. Pyrite crystal . . . .541 155. Dobereiner's lamp - . . . 556 OUTLINES OF CHEMISTRY CHAPTER I THE SCOPE OP CHEMISTRY AND ITS RELATIONS TO OTHER SCIENCES CHEMICAL CHANGE, ELEMENTS, AND COMPOUNDS OUR own bodies, and the various objects that surround us, constitute the subject of study of the natural sciences. The investigation of the things that make up the universe as we know it, is conducted by means of our senses, either aided or unaided. For the sake of classifying our knowledge, we are wont to distinguish between the biological sciences, which deal with living things, and the so-called physical sciences of astron- omy, geology, physics, and chemistry. Astronomy, which deals with the heavenly bodies, is nevertheless closely related to the sciences of physics and chemistry, though obviously not to biology. But the study of living things and the life history of the earth and the processes that are continually going on on its surface is inseparably linked with the subjects of physics and chemistry. The latter sciences may indeed be regarded as basal in character. The study of matter that is, anything which occupies space comes within the scope of these two sciences. Viewed in this light, biology, astronomy, and geol- ogy merely present special complex phases and combinations of physics and chemistry. Physical and Chemical Changes. The changes which any object may undergo are either superficial or deep-seated in character. Thus, if a stick of sulphur be thrown or whirled through the air, the character of the sulphur is not altered, though the sulphur has undergone change of position through expenditure of mechanical energy upon it. Energy is any- thing which does work or is capable of doing work. Energy itself is measured by the amount of work it has done or is B 1 2 OUTLINES OF CHEMISTRY capable of doing. Indeed, as energy is always measured in terms of work, the two are often regarded as synonymous. Work is equal to the force multiplied by the distance through which the force acts, a force being defined as that which causes or modifies motion, the latter being a change of place. The motion might have been imparted to the sulphur by means of the muscles or by a contrivance in which the energy was furnished by gravity, heat, light, electricity, magnetism, etc. These agencies are consequently capable of doing work ; that is, they represent forms of energy. As long as the sulphur remains sulphur, no matter through what motions or other alterations, like contraction, expansion, electrification, change of temper- ature, pulverization, liquefaction, or vaporization, it may go, the change in question is called a physical change, and the study of such changes in all their various phases belongs to the ' subject of physics. But if, for example, we burn the sulphur in the air we obtain a gas of a pungent odor which may be con- densed with the aid of pressure and lowering of the tempera- ture to a colorless, mobile liquid. This is quite unlike sulphur in all its various properties, and we consequently say that a new substance has been formed. The process of forming a new substance is called a chemical change. Any process in which given substances disappear and new ones are formed is chemical, and the study of such deep-seated processes in all their various phases is the subject with which the science of chemistry is concerned. It would thus seem fairly easy to dis- tinguish between chemical and physical processes. Indeed, in general, such a distinction can readily be made on the basis of what has just been said. But whether new substances have been formed must be decided from the properties of the material; and there must consequently be some definite way of telling whether an alteration of substance has occurred or not. It is evident at once that the term substance must be clearly defined. For our present purpose, it will suffice to say that a substance is matter which is perfectly homogeneous throughout, considered without respect to shape or amount. Thus sulphur, iron, and water are substances. Many things which are apparently homogeneous in character are not so in reality. Thus, the atmosphere on closer study is found to be a mixture of nitrogen, oxygen, carbon dioxide, and other gases ; sea water is found to THE SCOPE OF CHEMISTRY 3 consist of water together with various saline substances ; brass is made up of copper and zinc in proportions that may vary to a considerable extent in different samples. If we pulverize a piece of roll sulphur and grind it together with iron filings in a mortar as intimately as possible, a fine grayish powder results which has the outward appearance of homogeneity. On closer inspection, however, with the aid of a microscope perchance, this powder appears heterogeneous; in other words, it is merely a physical mixture Indeed, it is very easy to separate the iron from the sulphur, for by passing a magnet through the mixture the iron will adhere to the magnet, and the sulphur will be left behind. We could also separate the iron from the sulphur in the mixture by treating the latter with carbon disulphide, which liquid dissolves the sulphur and leaves the iron unaltered. The mixture of iron filings and sul- phur represents a typical physical mixture. It is obviously heterogeneous in character, the proportion of iron and sulphur in the mixture may be varied at will, and the iron and sulphur may readily be separated from each other by simple means. If now we heat some of the mixture of pulverized sulphur and iron filings in a test tube, we observe that at a certain temperature the contents of the tube begin to glow. As we take it out of the flame the glowing nevertheless increases, and the contents of the tube become hotter. After a time the glow- ing becomes weaker, and gradually ceases as the material cools. It is evident that by raising the temperature of the mixture of iron filings and sulphur to a certain point, a change was inaugu- rated, which on taking away the source of heat nevertheless continued, giving off additional heat and light. On examining the contents of the tube after it has cooled to room temperature, we find a black mass, quite unlike either the sulphur or iron in appearance. We can no longer detect heterogeneity in it even with the aid of the microscope. The magnet is unable to extract iron from this material, and carbon disulphide will not alter it in any way. A few drops of hydrochloric acid poured upon it evolve a malodorous gas called hydrogen sulphide, which is not formed when a simple mixture of iron filings and sulphur is moistened with that acid. We clearly have formed a new substance by heating the sulphur and iron together. It is called ferrous sulphide, and results from simple union of sulphur and 4 OUTLINES OF CHEMISTRY iron at elevated temperatures. It has been found that ferrous sulphide contains 63.52 per cent of iron and 36.48 per cent of sulphur, and that it always has exactly this composition no matter by what methods it has been formed. This is in fact a characteristic of all chemical compounds. We may express this fact by saying that every definite chemical compound always contains the same ingredients in the same proportion by weight. This is the law of definite proportions. A law, as the word is used in science, is a general statement summarizing what has actu- ally been found to be true in a large number of individual cases that have been carefully investigated. Other typical examples of chemical change are the rusting of iron, the combustion of coal or wood, the decomposition of water by electrolysis, the formation of quicklime from lime- stone by the agency of heat, the change of carbon dioxide and water into starch by sunlight in the green leaf of the plant, and the darkening of a photographic plate when exposed to light. In all these cases new substances are formed, and the actions are accompanied by changes in temperature, volume, outward appearance, and other specific properties which char- acterized the original substances before the change occurred. It is the province of chemistry to study such changes in all their various aspects. This involves a close study of the com- position and specific properties of the substances before and after the chemical change, which is commonly termed the chem- ical reaction, has taken place. But, in addition, a study of the conditions that must obtain in order that the reaction may begin and proceed, and an investigation of the various energy changes that accompany the reaction, also fall within the field of chemistry. Thus we have various branches of chemistry. So analytical chemistry seeks to determine the qualitative arid quantitative composition of substances by tearing them apart or analyzing them ; synthetic chemistry seeks to build up more complex substances from simpler ones ; thermochemistry concerns itself with the thermal changes accompanying chemical reac- tions ; electrochemistry is concerned with electricity as an agent in producing chemical changes, or as an accompaniment of chemical phenomena ; photochemistry treats of the relations of light to chemical changes. In the crust of the earth, in the atmosphere, in natural waters, in the bodies of plants and ani- THE SCOPE OF CHEMISTRY 5 mals, chemical changes are continually going on. Upon these all life on the globe depends. Every breath we breathe, every move we make, every thought we think, is accompanied by chemical changes and their concomitant physical phenomena as above briefly mentioned. The importance of the study of chem- istry, therefore, is clearly apparent, and it is also evident why there must needs be many special and applied lines of this sub- ject, which seek to investigate certain special fields. Thus we have agricultural chemistry, pharmaceutical chemistry, physio- logical chemistry, food chemistry, industrial chemistry, etc., the province of each of which is indicated sufficiently by the name itself. From what has been stated, it would seem a fairly simple mat- ter to distinguish a chemical change from a purely physical one, but this is by no means always easy. Suppose a block of ice and one of common salt be placed in contact with each other ; we note that the salt and ice gradually disappear, forming a brine. Evi- dently the brine has quite different properties from those of either the salt or the ice. Moreover, there was a marked change of temperature, in this case a cooling effect, as the salt and ice acted on each other. Furthermore, a contraction ensued, for the volume of the brine is less than the sum of the volumes of the blocks of ice and salt. Again, as a block of ice and one of par- affine, or one of salt and one of paraffine, for example, do not act on each other at all when brought into contact, it is clear that the action between ice and salt takes place because of the specific nature of the substances. Furthermore, it has been found that below 22 C. ice and common salt no longer act on each other, just as iron and sulphur do not act on each other at ordinary temperatures. Raise the temperature sufficiently in each case, and at a certain definite point action begins. Thus, the interaction of ice and common salt apparently bears all the earmarks of a genuine chemical change. This is indeed true except in one particular which has not yet been mentioned, namely, it is possible to vary the composition of the brine grad- ually, by adding common salt to it till a point of saturation is reached. Even then the brine will still take up somewhat more salt gradually if the temperature of the whole is slowly raised. The brine is termed a solution of common salt in water. It results from the action of salt and water on each other. The 6 OUTLINES OF CHEMISTRY water used may be liquid, or in form of ice above - 22 0. A distinction is commonly made between solutions and chemical compounds. In a solution, the relative amounts of the ingredients that it contains may be varied -gradually within certain limits, as we have seen in the case of the brine. In a chemical compound, the constituents cannot thus be varied in amount. Not many years ago, chemists spoke of solutions as chemical combinations according to variable proportions, and this term is indeed indic- ative of the real relation that they bear to definite chemical compounds which follow the law of definite proportions. Brine, then, is not a mere physical mixture, and it is conse- quently not to be classed with such mixtures as that of sulphur and iron filings rubbed together in a mortar, which represents a typical physical mixture. In chemistry we frequently have to deal with (1) physical mixtures, (2) solutions (i.e. com- pounds according to variable proportions), and (3) definite chemical compounds. As further typical examples of solutions may be mentioned, solution of sugar in water, of camphor in petroleum oil, of ether in alcohol, of carbon disulphide in olive oil. The subject of solutions clearly forms an important part of chemistry, and it will consequently be considered more fully later. Chemical Elements. A careful study of all substances known has revealed the fact that there are about eighty which it has been impossible to decompose into simpler substances thus far. These substances are regarded as elementary in char- acter. They are termed the chemical elements. Whether a substance is an element or not is thus determined by experi- ment. As new methods of experimental attack are discovered, substances that are now regarded as elements may prove to be complex and consequently capable of synthesis. Thus at one time lime and caustic potash were regarded as elements, whereas now we know that lime contains calcium and oxygen, and caustic potash consists of potassium, hydrogen, and oxygen. Sir William Ramsay found that the emanations from radium show the spectra of helium, argon, and neon, and this is by many regarded as a case of synthesis of the latter gases from the products of the decay of radium. Again, Ramsay claims to have obtained spectroscopic traces of lithium by the action of radium emanation upon copper sulphate solutions, though THE SCOPE OF CHEMISTRY 7 Mme. Curie's investigations do not substantiate his results. Thus it is evident that some of the substances we now term elements may prove to be composite. It is also obviously im- possible to state just how many elements there are, for it is uncertain whether some substances are elementary or complex in character. The following is an alphabetical list of the chemical elements as commonly recognized at present. CHEMICAL ELEMENTS Aluminum Europium Mercury Silicon Antimony Fluorine Molybdenum Silver Argon Gadolinium Neodymium Sodium Arsenic Gallium Neon Strontium Barium Germanium Nickel Sulphur Bismuth Glucinum Nitrogen Tantalum Boron Gold Osmium Tellurium Bromine Helium Oxygen Terbium Cadmium Hydrogen Palladium Thallium Csesium Indium Phosphorus Thorium Calcium Iodine Platinum Thulium Carbon Iridium Potassium Tin Cerium Iron Praseodymium Titanium Ohlorine Krypton Radium Tungsten Chromium Lanthanum Rhodium Vanadium Cobalt Lead Rubidium Xenon Columbium Lithium Ruthenium Ytterbium Copper Lutecium Samarium Yttrium Dysprosium Magnesium Scandium Zinc Erbium Manganese Selenium Zirconium It will be observed that the list contains a goodly number of common, well-known substances. Notably, it appears that all the metals are elements. Again, there are substances in the list which are not metals, like sulphur, chlorine, bromine, iodine, oxygen, hydrogen, phosphorus, etc. The elements may be divided into two groups ; namely, the metals and non-metals. It is difficult to draw a sharp line between these groups, however, for elements like arsenic, antimony, and tellurium clearly rep- resent transitions between the metals and non-metals. Such transition elements are sometimes called metalloids. Some of the elements are gases, others are liquids, and still others are solids, under ordinary conditions of temperature and pressure. Whether an element is a solid, liquid, or gas OUTLINES OF CHEMISTRY is determined entirely by the conditions of temperature and pressure to which it is subjected. Less than half of the elements mentioned in the table enter into the composition of ordinary objects. The solid crust of the earth, also called the lithosphere, makes up about 93 per cent of all known terrestrial matter, while the ocean represents about 7 per cent, and the atmosphere only 0.03 per cent, of the total. The following table, by F. W. Clarke, gives an estimate of the relative amounts of the elements contained in the litho- sphere and the ocean. The third column of the table gives a total average including the atmosphere. AVERAGE COMPOSITION OF LITHOSPHERE, OCEAN, AND ATMOSPHERE LITHOSPHERE (93 PEE CENT) OCEAN (7 PER CENT) AVERAGE INCLUDING THE ATMOSPHERE 47.07 28.06 7.90 4.43 3.44 2.40 2.43 2.45 0.22 0.40 0.20 0.07 0.11 0.11 0.09 0.07 0.03 0.02 0.50 85.79 0.05 0.14 1.14 0.04 10.67 0.002 2.07 0.008 0.09 49.78 26.08 7.34 4.11 . 3.19 2.24 2.33 2.28 0.95 0.37 0.19 0.21 0.11 0.11 0.09 0.07 0.03 0.02 0.02 0.48 Iron ...... Calcium . ... Potassium Carbon Phosphorus All other elements . . 100.00 100.00 100.00 The following table gives the approximate amounts of the elements found in the human body: THE SCOPE OF CHEMISTRY 9 AVERAGE ELEMENTARY COMPOSITION OF THE HUMAN BODY Oxygen 66.0 per cent Carbon 17.6 per cent Hydrogen , . .10.1 per cent Nitrogen 2.5 per cent Calcium . . . . . . 1.5 per cent Phosphorus . . . . . . . .1.0 per cent Potassium . . . . . . . . 0.4 per cent Sodium 0.3 per cent Chlorine 0.3 per cent Sulphur 0.25 per cent Magnesium 0.04 per cent Iron 0.004 per cent Silicon, Fluorine, Iodine, etc., in traces. Compounds. Most substances are non-elementary in charac- ter ; that is, they are combinations of two or more elements. Such substances are consequently termed compounds. They may be formed by direct union of the elements with one another under proper conditions ; as, for instance, sulphur may unite with iron to form ferrous sulphide. Again, limestone, which is carbonate of calcium, decomposes at a high tempera- ture, forming two simple substances, lime and carbon dioxide, the latter being a gas. Further, by action of two compounds on each other, two other compounds may result. As an example, when common salt and nitrate of silver are brought together in aqueous solution, silver chloride, a substance in- soluble in water, and sodium nitrate, a soluble substance, are formed. This latter change is termed double decomposition or metathesis. In the three cases cited we have, indeed, the three types of chemical changes ; namely, (1) the direct union of two or more substances to form a single compound, (2) the breaking up of a compound into simpler ones, and (3) the interaction of substances with one another to form new substances. Like elements, compounds may also assume the solid, liquid, or gaseous state, according to the conditions of temperature and pressure that obtain. However, by no means all compounds are capable of assuming these three states, for many readily decom- pose when an attempt is made to liquefy them or volatilize them by means of heat. 10 OUTLINES OF CHEMISTRY . Compounds which contain different elements are, of course, different in character. The same is true of compounds that contain the same elements, though in different proportions by weight. For a long time it was thought that one compound could differ from another only because it contained either differ- ent elements, or the same elements in different proportions. However, we now have knowledge of a large number of com- pounds that are quite different substances, and yet they con- tain the same elements in exactly the same proportion by weight. Such compounds are called isomers, and the difference between them is explained by the different manner in which the elements are combined in these substances, in other words, by the difference in inner structure or constitution of the compounds. Conservation of Mass. Investigations have shown that when chemical changes take place, the weight of all the substances before the reaction is equal to the weight of all the substances after the reaction has taken place. In other words, in any chemical change the total weight remains the same. As at any place on the surface of the earth weight and mass are propor- tional to each other, we may say that during chemical changes, the total mass of the reacting substances remains constant. This is simply the law of conservation of mass, which applies to chemi- cal as well as to physical changes. It is sometimes called the law of conservation of matter. It is the outcome of experi- mental investigations, the most careful of which were conducted by having chemical changes go on in sealed glass vessels, which, together with their contents, were weighed before and after the sub- stances they contained had reacted chemically on one another. Figure 1 shows a common type of such sealed glass tubes. The sub- stances are introduced into each limb, and the tube is then sealed by drawing off the end as shown. After the whole has been very accurately weighed, the contents are allowed to act on each other by inclining or shaking the tube. After the action has ceased and the whole has cooled to room temperature, the tube is carefully weighed again. H. Landolt has performed THE SCOPE OF CHEMISTRY 11 many careful experiments of this nature in recent years. His results show that if there is any change of weight, it lies very near the limit of experimental error. That is, it is so slight as to be quite negligible for all ordinary purposes. Conservation of Energy. Like mass, energy also cannot be created or destroyed. It can simply be transformed. Thus, for example, electricity may be converted into heat, mechanical energy, or chemical energy ; and again, each of these latter may be converted back into electricity. When coal burns, to be sure, new substances are being formed, but in addition chemical energy is being converted into heat and light. When water is decomposed by means of the elec- tric current, electrical energy is being converted into chemical energy. When lime is produced at the high temperature of the limekiln, heat is transformed into chemical energy. When starch is formed in the sunlight in the green leaf of the plant, light is converted into chemical energy. The Cause of Chemical Change. As to the cause why certain substances act on each other to form new substances under given conditions and other substances do not, we are quite ignorant. Thus, we cannot tell why a piece of sulphur will burn when heated in the air and a piece of platinum or gold will not. We know that, in the act of burning, the sulphur unites with the oxygen of the air, and therefore we explain this by saying that sulphur and oxygen have a specific attraction for each other. This specific attraction, which is regarded as the cause of chemical union, is commonly called chemical affinity. Thus, the fact that platinum and gold do not burn when heated in the air is explained by saying that these elements have too slight a chemical affinity for oxygen. The word affinity means relationship. It was adopted at a time when it was thought that substances that are similar are more prone to unite chemically with one another than those which are dissimilar. While it is true, as we shall see, that substances of similar characteristics do frequently unite chem- ically, nevertheless, as a rule, substances that are unlike in character react more energetically with one another. So, for instance, while metals do form chemical compounds with metals, yet they react much more energetically with non-metals and thus form stabler compounds. 12 OUTLINES OF CHEMISTRY Factors affecting Chemical Change. In order that a chemi cal change may take place, it is first of all necessary that the substances that are brought together be of the right kind ; that is, they must be of such a specific nature that they will react. According to the preceding paragraph, we should say, the sub- stances must have chemical affinity for one another. When we study any substance as to its power to react with other bodies to form new substances, we are investigating the chemical properties of the substances. Intimate contact of the substances that are to react is always necessary. From this fact we con- clude that chemical affinity acts at insensible distances. Again, temperature is a great factor in promoting chemical change; in- deed, in most of the changes studied in the laboratory, tempera- ture is second only to chemical affinity itself in determining whether chemical action will take place or not. Electricity, light, pressure, concussion, various forms of vibration, contact ivith other substances which often need to be present only in relatively minute quantity, are also frequently important factors in determin- ing whether a chemical change will proceed or not. Furthermore, the relative amounts of the reacting substances brought into contact also affect the rate of a chemical change and the extent or degree of completion to which it will proceed. REVIEW QUESTIONS 1. What is a science? 2. Name five of the more fundamental natural sciences and state of what each treats. 3. Distinguish between chemical change and physical change, and give six examples of each. 4. What is a substance, as that term is used in chemistry? 5. Distinguish between a physical mixture, a solution, and a chemical compound, giving three examples of each. What is a chemical element ? 6. Name the important phenomena that may accompany chemical change. 7. Make a list of the most important metallic elements that occur in the earth's crust. Make a similar list of the nonmetallic elements. 8. What elements occur in the human body to the extent of 1 per cent or more ? 9. Mention the three types of chemical change, giving illustrations. 10. State the law of conservation of mass ; also the law of conserva- tion of energy. Illustrate by an example in each case. 11. What is meant by the term "chemical affinity "? 12. Prepare a list of the more important factors or agencies that may inaugurate chemical change or affect its rate of progress. CHAPTER II HYDROGEN History. It was known to Paracelsus (1493-1541) that an inflammable gas is produced when dilute acids act on certain metals ; but the English physicist Cavendish (1731-1810) was the first to isolate hydrogen and recognize it as a special gas. In 1766 he prepared hydrogen by the action of either hydro- chloric or sulphuric acid on zinc, iron, or tin, and described the characteristic properties of the gas. Hydrogen is an essential constituent of water, and derives its name from the Greek words meaning water and to generate. Occurrence. Hydrogen is perhaps the most widely distributed element in the universe. It occurs in very large quantities in the sun, where it is heated to incandescence owing to the high temperature that obtains. It is found in all fixed stars and nebulae that have been examined by means of the spectroscope. On the earth it occurs only in small amounts in the free state. The atmosphere contains only about 0.005 per cent of uncom- bined hydrogen by volume. In the gases emitted from vol- canoes, oil wells, and some natural salt deposits, notably those at Stassfurt, Irvdrogen is found in the free state. It occurs further in the gases resulting from certain forms of fermentation, in the gases emitted from living plants, and in the intestinal gases of human beings and animals. In meteoric iron, and in various minerals, hydrogen has also at times been found as an occlusion. While hydrogen exists only in small quantities in the free state on the earth, in combination with other elements it is found in very large quantities. Thus, 11.19 per cent of the weight of water consists of hydrogen. It forms an essential part of all plants and animals, in which it occurs chiefly in combination with the elements oxygen, carbon, and nitrogen. In petroleum, natural gas, and marsh gas it occurs combined with carbon. It is an essential constituent of all acids. 13 14 OUTLINES OF CHEMISTRY Preparation. WMn an electric current is passed through water acidulated wit'i sulphuric acid (Fig. 2), both hydrogen and oxygen are produced, two volumes of the former and one volume of the latter appearing at the opposite plates used as electrodes. This method is an excellent one for preparing very pure hy- drogen. The process itself is, however, some- what complex in nature and will receive special attention later (see Elec- trolysis). When metallic sodium acts on water, hydrogen and caustic soda are formed. The sodium may be introduced into a test tube which has been filled with water and in- verted in a basin as shown in Fig. ' 3. The metal, being lighter than water, rises in the tube, and as the hydrogen is generated it forces the water out of FlG 2 the tube. The metal melts owing to the heat generated during the reaction and floats in form of a globule on top of the water. We may express what takes place by writing : Water -f- Sodium = Hydrogen + Caustic Soda. The latter substance is dissolved in the water after the change has taken place. It may be obtained as a white solid by boiling the solution till all the water has evaporated. The caustic soda solution turns red litmus blue, has an "alkaline" HYDROGEN 15 taste,- and feels slippery to the touch. Caustic soda consists of three elements, sodium, oxygen, and hydrogen. It is also called sodium hydroxide. It is a strong alkali, that is, a sub- stance which is able to neutralize acids and thus form salts. Potassium acts on water like sodium, only much more vigor- ously. The metal in this case catches tire and burns with a brilliant flame. Frequently the action is so violent as to result in explosions. Lithium, rubid- ium, csesium, barium, strontium, and calcium also act on water at room temperature, forming hydrogen and the hydroxide of the metal employed. It is therefore evident that all of these metals cannot be kept in contact with the air, which always contains some moisture. They are kept under hydrocarbon oils, like kerosene, with which they do not react. Magnesium decomposes water at room temperatures, but very slowly indeed. If, however, a magnesium salt is dis- solved in the water, the action goes on much more rapidly. Magnesium salts aid the action by dissolving the magnesium hydroxide formed, which would inclose the metal in a pro- FIG. 3. FIG. 4. 16 OUTLINES OF CHEMISTRY tecting film. On boiling water magnesium acts quite readily, forming hydrogen and the hydroxide of the metal. Zinc or iron when heated to redness in a tube will decompose steam, yielding hydrogen and an oxide of the metal employed (Fig. 4). Furthermore, by similarly passing steam over red-hot carbon, hydrogen and carbon monoxide are formed. This latter process is used in making water gas (which see). By boiling zinc in aqueous caustic potash solution, hydT\,g>o^ and potassium zincate result. The latter is a salt which remains in solution; thus: Caustic Potash + Zinc = Potassium Zincate + Hydrogen. Caustic soda acts like caustic potash. Aluminum acts in a manner similar to zinc. In this case an aluminate instead of a zincate is formed and remains in solution. By heating zinc dust or scrap iron with slaked lime, hydrogen is liberated, and an oxide of the metal used is simultaneously formed. This method is frequently used for preparing hydrogen in large quantities for industrial purposes. By far the commonest way of preparing hydrogen in the laboratory is by treating zinc with dilate sulphuric acid. The apparatus used for this purpose is shown in Fig. 5. In this reaction there is formed, besides hydrogen, a white salt Fia. 5. HYDROGEN 17 called zinc sulphate. It remains in solution, and may be obtained in form of crystals by evaporating the solution to a small bulk and allowing it to cool. We may express the change thus : Sulphuric Acid + Zinc = Hydrogen -f Zinc Sulphate. Instead of sulphuric acid, dilute hydrochloric acid or acetic, acid may be used. Furthermore, iron may be substituted for the zinc, in which case hydrogen and corresponding salts of iron are formed. Hydrogen thus obtained is never quite pure. The impurities present in ordinary zinc and iron, such as carbon, arsenic, sulphur, and phosphorus, combine with some of the hydrogen, and the relatively small amounts of the resulting gases contaminate the larger portion of hydrogen which remains. These impurities may be removed by passing the gas through appropriate absorbents. Ordinary cast iron usually contains so much of the impurities mentioned, notably of carbon, that, when treated with an acid, the hydrogen liberated is con- taminated sufficiently to have a very disagreeable odor. Properties. Hydrogen is the lightest of all known sub- stances. It is a colorless, odorless, tasteless gas. At and 760 mm. barometric pressure, namely, under standard conditions, one liter weighs 0.08987 gram. It is 14.388 times lighter than the air ; in other words, its specific gravity with respect to air is 0.0695. Because of the light- ness of hydrogen, jars containing the gas are held bottom upward. Figure 6 shows how hydrogen may be transferred from one jar A into another jar B. At and below 241, its critical tem- perature, hydrogen may be liquefied by subjecting it to pres- sure. At 241 a pressure of 20 atmospheres will liquefy the gas; but at 252.5 the vapor tension of liquid hydrogen is practically one atmosphere ; that is to say, the liquid boils at FIG. 6. 18 OUTLINES OF CHEMISTRY H the last-named temperature. Liquid hydrogen is clear ancl colorless, like water, and has a specific gravity of about 0.07. Solid hydrogen may be obtained by evaporating liquid hydrogen in a partial vacuum. The melting point of the solid, which consists of white crystals, is 259. Its power to refract light is 6.5 times greater than that of air. On account of its light- ness, hydrogen diffuses very rapidly, and readily passes through porous sub- stances like unglazed por- celain, brick, mortar, and paper. The rate of diffu- sion of gases is inversely proportional to the square roots of their densities ; hence, air diffuses only V0.0695 or 0.2636 time as fast as hydrogen. The rapid diffusion of hydrogen may be demonstrated by means of the apparatus shown in Fig. 7. When the unglazed porcelain cup A, which contains air, is surrounded with hydrogen gas, which is passed into the inverted vessel B by means of the tube C, the hydrogen diffuses into the porous cup A much faster than the air diffuses out. FIG. 7 A . ,, A pressure is consequently produced in A which is connected with the Wolf bottle D containing water ; and this pressure forces the water out of the tube U in form of a fountain. Hydrogen is but slightly soluble in water, for only 19 vol- umes are absorbed by 1000 volumes of water at 15. Certain HYDROGEN 19 solids absorb hydrogen in notable quantities. Freshly ignited charcoal absorbs about twice its volume of hydrogen. Palla- dium absorbs 500 volumes, platinum 49 volumes, iron 19 vol- umes, gold 46 volumes, copper 4.5 volumes, nickel 17 volumes, aluminum 2.7 volumes, lead 0.15 volume. At red heat, pal- ladium may even absorb as much as 900 volumes, according to Graham. This power of solids to absorb gases is sometimes termed adsorption, or occlusion. The amount absorbed depends upon the specific nature of the solid and also of the gas ; and as the absorption is accompanied with changes of temperature and of volume, it is clear that the phenomenon is akin to the process of solution. Furthermore, hydrogen passes through iron and platinum tubes when these are hot. This is readily explained by the fact that these metals absorb the gas. The most notable characteristic of hydrogen is its inflamma- bility. It burns readily in the air or in oxygen, and the product formed is water. This can easily be shown by holding a cold bell jar over a burning jet of hydrogen (Fig. 8). The water FIG. 8. formed condenses in drops on the sides of the jar. At ordinary temperatures, hydrogen and oxygen do not act on each other appreciably, but the action takes place when the -gases are heated to the kindling temperature, which is about 615, ac- cording to V. Meyer. The hydrogen flame is colorless. To show this the gas must be burnt from a platinum jet, which is not affected during the process, so that particles of foreign matter do not get into the flame and color it. When a glass 20 OUTLINES OF CHEMISTRY tube is held over a hydrogen flame, as shown in Fig. 9, the column of air in the tube is set in vibration, thus producing the phenomenon known as the singing flame. The hydrogen flame is very hot, which is evident from the fact that platinum, which fuses above 1700, will melt in it. The burning of one gram of hydrogen develops about 34.5 large calories of heat, which is enough heat to raise the temperature of 345 grams of water from to the boiling point or to melt 431 grams of ice. Hydrogen is not poisonous, but animals would suffocate in the gas for lack of oxygen, without which they cannot live. Hydrogen will also not support ordinary combustion. Thus, when a lighted candle is thrust into a jar of hydrogen (Fig. 10), the gas at the mouth of the jar , which contains a weighed quantity of a metal, say magnesium. The upper end of A is then filled with dilute acid, which by carefully opening the cock O is allowed to flow down upon the metal. The cock is closed before all the acid has passed the cock so as to avoid admitting air into the graduated tube. After the acid has dissolved the Fia. 13. FIG. 14. HYDROGEN 23 metal completely, the level of the liquid in the tube A is adjusted so that it is the same as that in the beaker The volume of the hydrogen is noted, the temperature and barometric pressure are taken, and from these data the volume of the hydrogen under standard conditions is computed. Knowing this and the weight of one liter of hydrogen, the weight of the hydrogen liberated may readily be calculated. The result would be the weight of hydrogen displaced from the acid by the given weight of magnesium, and from this the amount of magnesium required to liberate 1 gram of hydro- gen can easily be found. An experiment of this kind yields the result that it requires 12.16 grams of magnesium to liberate 1 gram of hydrogen. Similarly, it has been found that 23.00 grams of sodium, or 39.10 grams of potassium, or 9.03 grams of aluminum, or 27.9 grams of iron, or 59.5 grams of tin, or 32.7 grams of zinc are required to set free 1 gram of hydrogen. The quantities mentioned are called the hydrogen equivalents of the re- spective metals ; or sometimes they are simply spoken of as the chemical equivalents. It is evident that the amounts of the vari- ous metals that are chemically equivalent to 1 gram of hydrogen are very different. The chemical equivalents of the elements are of great importance, and they will be referred to again later. When each of the metals above mentioned acts upon dilute hydrochloric acid, it is evident, from even a rough observation, that the rate with which the different metals liberate hydrogen varies greatly. Arranging these metals in the order of rapidity with which they react with dilute acid, WQ have : potassium, sodium, magnesium, aluminum, zinc, iron, and tin, the action being strongest in the case of potassium, and weakest in the case of tin. This gives us an idea of the relative affinity or chemi- cal attraction that exists between these metals and the dilute aqueous solution of the acid used, or rather between the metals and that part of the aqueous acids with which the displaced hydrogen was combined. By measuring accurately, at con- stant tempeiature, the rate of the liberation of hydrogen per minute when one and the same area of the different metals acts on samples of the same dilute acid solution, the relative affini- ties of the metals for the acid may be determined ; for, other factors being constant, the rate with which a chemical reaction pro- ceeds is proportional to the chemical affinity that comes into play. 24 OUTLINES OF CHEMISTRY REVIEW QUESTIONS 1. Where is hydrogen found in nature? 2. Name six compounds that contain hydrogen. 3. Give four general methods for preparing hydrogen in the laboratory. 4. What are the most important properties of hydrogen, and what practical use is made of hydrogen gas ? 5. Give two illustrations of the adsorption of hydrogen by solids. . 6. What is formed when hydrogen burns in the air? What are the characteristics of the hydrogen flame ? 7. What is a hydride ? Give three illustrations. 8. Give an illustration of the process of reduction ; also one of the process of oxidation. 9. What are the products formed when dilute sulphuric acid acts on zinc? What becomes of these products? Similarly, what are the prod- ucts formed when aluminum acts on sodium hydroxide and what be- comes of them in the experiment ? 10. How test a hydrogen generator to ascertain if it is free from air? Why is it necessary to do this ? 11. Define the term "hydrogen equivalent," and illustrate by two examples. 12. If the metal used in determining the hydrogen equivalent con- tained two per cent of insoluble material, would the hydrogen equivalent found by the method shown in Fig. 14 be too high or too low? 13. How many grams of hydrogen could be obtained by the action of 200 grams of iron on an acid? What volume would this gas occupy at C. and 760 mm. pressure? What volume would it occupy at 20 C. and 744 mm. pressure ? 14. How many grams of sodium would be required to liberate five liters of hydrogen measured over water at 22 C. and 750 mm. pressure? 15. What volumes of hydrogen and oxygen form the most explosive mixtures of these two gases? 16. How do you account for the disagreeable odor of the gas that is liberated when hydrochloric acid acts upon cast iron? How could you demonstrate that hydrogen gas is quite odorless ? 17. When and by whom was hydrogen gas first isolated? How was this gas prepared by its discoverer? 18. What explanation can you give of the following facts : sodium liberates hydrogen from water at room temperatures; magnesium liber- ates that gas from boiling water ; red hot iron liberates it from steam ; and gold does not decompose water at all. 19. What peculiar property does metallic palladium manifest toward hydrogen gas, especially when red hot ? 20. How do you explain the fact that a candle will not burn in an atmos- phere of hydrogen? 21. If materials like fats, oils, and wood will not burn in an atmosphere HYDROGEN 25 of hydrogen, why is it nevertheless not feasible to use the latter gas as a fire extinguisher? 22. Though animals die when left in an atmosphere of hydrogen gas, how could you prove that the latter is not poisonous ? 23. Explain how the density of hydrogen may be found by ascertain- ing how fast the gas flows from a very small orifice. 24. What causes the loud report when a mixture of oxygen and hydrogen gases is brought in contact witl\ a flame or a spark? Give a detailed explanation of each step in the process. CHAPTER III OXYGEN History. Oxygen was discovered in 1774 by Joseph Priestley, who liberated the gas by heating the red oxide of mercury. It was independently discovered in 1773 by Scheele, but he did not publish his work till 1775. Lavoisier, who found the dis- covery of the gas of particular interest in connection with his studies of the process of combustion, named the element oxygen, from the Greek words meaning acid and to generate. He found that the union of oxygen with such elements as sulphur, nitrogen, and arsenic produced substances that were sour to the taste, and in general behaved like other well-known acids. His con- clusion was that oxygen is an essential constituent of all acids, but later -work has shown this to be erroneous. Occurrence. Oxygen is the most abundant element on the earth. The atmosphere contains about 21 per cent of free oxy- gen by volume. Water contains 88.88 per cent of oxygen by weight, and the rocks of the earth's crust contain from 44 to 48 per cent. It is present in all animals and plants, in which it occurs in combination with hydrogen and carbon, and also with hydrogen, carbon, and nitrogen. Preparation. (1) When liquid air is allowed to evaporate, the nitrogen, which is more volatile than the oxygen, passes off first, and thus a considerable portion of the oxygen is left in the container, approximately free from nitrogen. (2) By electroly- sis of water, acidified with sulphuric acid, two volumes of hydro- gen and one volume of oxygen are produced. (3) By heating red oxide of mercury, this compound is decomposed, yielding oxygen and mercury; similarly, the oxide of silver may be decomposed by heat. Again, the peroxides of manganese, lead, and barium give off a portion of their oxygen on heating them. The peroxides of these metals also evolve oxygen when heated with sulphuric acid. (4) Certain salts rich in oxygen give off their oxygen content either in part or entirely upon being heated. OXYGEN 27 Thus, saltpeter yields oxygen and potassium nitrite on ignition, and potassium chlorate when heated yields oxygen and potassium chloride. The latter method is very commonly used for preparing oxygen for laboratory purposes. One hundred grams of potassium chlorate yield about 39 grams of oxygen. In the process of heating potassium chlorate, potassium perchlorate first forms, and this upon further heating breaks down into oxygen and potassium chloride. (5) By treating a solution of hydrogen peroxide, acidified with sulphuric acid, with potassium perman- ganate or potassium bichromate, oxygen is evolved. This method is very convenient for laboratory purposes. (6) When bleaching powder acts on peroxide of hydrogen, oxygen is evolved. (7) Barium oxide when heated in the air to about 500 takes on oxygen, forming barium peroxide. The latter on being heated up to 1000 parts with half of its oxygen, forming the original barium oxide, and the process can then be repeated. This is known as Brin's process. It will be seen that it is a con- venient method of preparing oxygen from the air. It is used for preparing oxygen for commercial purposes. (8) The green leaves of plants in the sunlight decompose carbon dioxide and water, forming starch and oxygen. Large quantities of oxygen are thus supplied to the atmosphere. Properties. Oxygen is a colorless, odorless, tasteless gas. It is 1.10 times as heavy as air. One liter under standard conditions (0 and 760 mm.) weighs 1.4290 grams. Its power to refract light is only 0.8616 time that of air. The gas may be liquefied at and below 119, its critical temperature. At 119 a pressure of fifty atmospheres is required to liquefy oxygen. This pressure is consequently the critical pressure. Liquid oxygen is a light blue, mobile liquid which boils at 182.5 under atmospheric pressure. It is attracted by a mag- net. At 182.5 the specific gravity of the liquid is 1.1315. By means of liquid hydrogen, Dewar froze oxygen to a pale blue, snowlike solid, whose melting point is 227. Oxygen is slightly soluble in water. At and atmospheric pressure 100 volumes of water dissolve four volumes of oxygen, while at 15, 3.4 volumes of the gas are absorbed. Oxygen may consequently be collected over water. Chemically, oxygen is a very active substance combining directly with all known elements, the only exceptions being 28 OUTLINES OF CHEMISTRY FIG. 15. fluorine and the gases of the argon group, namely, helium, neon, argon, krypton, and xenon. The compounds of the elements with oxygen are called oxides. At ordinary tempera- tures, oxygen unites but slowly with most substances. Thus, the rusting of iron consists of a slow union with oxygen of the air. Sodium is oxi- dized quite rapidly on exposure to air or oxygen at room temperature, while in the case of wood, charcoal, or sulphur, the union with oxygen at ordinary tempera- tures proceeds very slowly indeed. How- ever, at elevated temperatures all of these substances combine readily and vigorously with oxygen, with concomitant evolution of heat and light. This process is termed combustion. All chemical processes which proceed with the evolution of light and heat may, in general, be called cases of combustion ; ordinarily, however, the term is applied to union with oxygen. In an atmosphere of the latter gas, iron will burn with brilliant scintillations (Fig. 15) and evolution of much heat. The product formed is an oxide of iron of a reddish brown color. Phosphorus burns brilliantly in oxygen, forming phosphoric oxide, consisting of white fumes which condense on the sides of the container. On moistening this white solid with water, a solution of phos- phoric acid is formed. This solution is sour and turns blue litmus red. Carbon burns in oxygen to carbon dioxide ; sul- phur burns to sulphur dioxide (Fig. 16). These gases, too, form acids when treated with water. The oxides of phosphorus, carbon, and sulphur are consequently acid -forming oxides. They are also spoken of as acid anhydrides; that is, the acids minus water. Sodium when burned in oxygen forms a white powder, called sodium oxide, which readily dissolves in water, yielding a solu- tion which is alkaline to the taste, turns red litmus blue, and irtt FIG. 16. OXYGEN 29 feels slippery to the fingers. It is an alkali or base, and is capable of reacting with acids, forming salts whose aqueous solutions have no effect on litmus, i.e. they are neutral. Potassium and calcium also burn readily in oxygen, forming the oxides of potassium and calcium. These are white caustic substances which resemble the oxide of sodium. The oxide of calcium is ordinary lime. The oxides of potassium and calcium are caustic alkalies. Other oxides, like those of zinc, iron, and lead, are insoluble in water. They are consequently tasteless and do not affect litmus. The oxides of most metals can be formed by direct union with oxygen. Some metals, like gold and platinum, do not burn in oxygen, but their oxides may be formed indirectly by double decomposition. On heating such oxides, they yield the metal and oxygen. Combustion in the Air. The combustion of substances in the air yields precisely the same products as combustion in oxy- gen. Indeed, the process of burning substances in the air is in all respects, except in brilliancy, rapidity, and vigor, like that of burning them in oxygen. As the oxygen of the air is di- luted with four times its volume of nitrogen, which latter gas is rather inert in character, it is quite natural that combustion in the air should go on less vigorously than in oxygen. The total energy liberated as heat is the same, however, whether the oxidation of a substance takes place rapidly in pure oxygen, or less rapidly in the air, or extremely slowly at ordinary tempera- tures in the air. Kindling Temperature and Temperature of Combustion. In order to burn a substance in oxygen, it must be heated to a certain minimum temperature at which it will burst into flame. This temperature, which is very different for different sub- stances, is called the kindling temperature. Thus, phosphorus catches fire at a much lower temperature than sulphur, and the latter ignites at a lower temperature than wood. The highest temperature attainable during the process of combustion of a substance is sometimes called the temperature of combustion. It varies greatly with the nature of the sub- stance. It is higher in pure oxygen than in air, and higher in compressed oxygen than in that gas at atmospheric pressure. The temperature of combustion is generally very much higher than the kindling temperature. 30 OUTLINES OF CIIKMISTKY Heat of Combustion. The heat evolved during the combus tion of a substance is called its heat of combustion. As above stated, it is the same whether the combustion goes on rapidly or slowly, though the maximum temperature reached during the process of combustion varies greatly under different conditions. The unit of heat is the calorie. The small calorie is the amount of heat required to raise 1 gram of water 1 degree ; UK l;u _;(! calorie is 1000 times as large, i.e. it is the amount of heat required to raise iOOO grams of water 1 degree in tern perature. It is very important to ascertain the heat of com- bustion of various substances, not only for purely scientific purposes, but also for the determination of the relative value of fuels and certain classes of food. Heats of combustion will consequently receive special consideration in the chapter on thermochemistry. Different Stages of Oxidation. While it is true that combus- tion in the air or in oxygen is essentially the same process, except as to rapidity, it not infrequently happens that when a sul .stance is burnt in an excess of oxygen, more of the latter enters into the oxides formed than when the burning proceeds in the air. Thus, when iron is oxidized by heating it in the air, a black oxide is formed which is magnetic in character, and which consists of 72.38 per cent iron and 27.62 per cent oxy- gen ; whereas when iron is burned in oxygen, there is formed mainly a reddish brown oxide of iron which is practically non- magnetic, and which contains 69.96 per cent iron and 30.04 per cent oxygen. By carefully heating the latter oxide in a cur- rent of hydrogen at 500 a black oxide may be obtained which consists of 77.75 per cent iron and 22.25 per cent oxygen. Writing the composition of these oxides of iron, the only ones known, in form of a table, we have as follows : l'i i: (Y.x-r |I;MN PKR CENT OXYGEN PARTS OXYGEN TO 77.75 PARTS IKON (1) 7U.:is ('_) 09.90 (3) 77.7.", 27.62 30.04 22.25 29.67 33.38 22.25 In the third column are placed the amounts of oxygen by weight combined with one and the same amount of iron; namely, OXYGEN 31 77.75 parts. The latter figure was chosen simply for conven- ience, as it represents the percentage of iron in the oxide poor- est in oxygen. Now, inspecting the table, we see that 29.67:22.25 = 4:3, and that 33.38 : 22.25 = 3 : 2. This means that in these three different oxides of iron the amounts of oxygen that are combined with one and the same amount of iron are simple, rational multiples of one another. This being the case, had we calculated the amounts of iron combined in these oxides with one and the same amount of oxygen, we should have found that these amounts of iron are also simple, rational multiples of one another. Again, there are five different oxides of lead known. These as follows: (1) lead suboxide, a black substance formed rhen lead is heated at its melting point in the air; (2) lith- irge, a yellow powder formed when lead is very strongly heated in air ; (3) lead sesquioxide, an orange-yellow powder formed 'hen bleaching powder acts on litharge dissolved in caustic >tash ; (4) red lead or minium, a bright red powder, which be obtained by heating litharge in the air at a temperature lot above 450; and (5) lead peroxide, a brown powder, which be prepared by treating red lead with dilute nitric acid, percentage composition of these oxides is as follows : NAME PEE CEST LEAD PER CEMT OXTOEK PAKT LEAD TO 3.72 I'AKTH OXYOEX (1) Lead suboxide . . . (2) Litharge 96.28 92.82 3.72 7.18 96.28 48.14 (3) Lead sesquioxide . . (4) Red lead (5) Lead peroxide . . . 89.61 90.65 86.60 10.39 9.35 13.40 32.09 36.11 24.07 In the last column we have the amount of lead combined ith 3.72 parts of oxygen in each of the oxides. Comparing figures in the last column we note as follows : (1) and (2) (1) and (3) (1) and (4) O > and (5) 96.28:48.14 = 2:1, 96.28 : 32.09 = 3:1, JS.3tf.ll =8:3, 96.28:24.07 = 4:1. 32 OUTLINES OF CHEMISTRY Thus we see that in the five oxides of lead, the amounts of lead combined with one and the same amount of oxygen are simple multiples of one another. Obviously the amounts of oxygen which in these oxides are combined with one and the same amount of lead are also simple multiples of one another. Law of Multiple Proportions. These results of the quantita- tive study of the composition of the oxides of iron and lead are typical of a large number of similar cases. It has been found to be general, that whenever two elements form more than one compound with each other, the amounts by weight of the one that are united with one and the same weight of the other are simple rational multiples of one another. This is the law of multiple proportions. It was discovered by John Dalton about 1806. Many careful analyses of various compounds have since yielded results confirming this law, which is of fundamental importance in chemistry. As we proceed, we shall meet numerous addi- tional instances illustrating the law of multiple proportions. Role of Oxygen in Respiration. Oxygen is necessary for all animal life. If the oxygen supply is cut off from an animal, it soon dies from suffocation. Pure oxygen may be inhaled with- out evil effects for a while. An animal placed in oxygen shows invigoration by its more lively movements; but after a while febrile symptoms appear, and a reaction sets in which may cause death. The air as it enters the lungs is virtually oxygen diluted with four times its volume of nitrogen. It is the oxygen only that is absorbed by the membranes of the lungs. Furthermore, only 4 to 5 per cent of the oxygen contained in the air is thus absorbed in the process of respiration. The ex- haled air contains water and also about 3 to 4 per cent of car- bon dioxide, gained from the body. The oxygen from the air passes through the membranes of the lungs, into the blood, where it is taken up by the blood corpuscles. The latter contain hemoglobin, a crystalline sub- stance which unites with oxygen, forming oxyhemoglobin, which has a red color, giving a bright appearance to arterial blood. As oxyhemoglobin attached to the blood corpuscles, the circulation carries the oxygen to all parts of the body, where it is given off, entering into various combinations with the tissues. As the blood is thus deprived of oxygen, carbon dioxide, which is formed during the oxidation of the tissues, is OXYGEN 83 taken up and carried to the lungs, where it is exhaled and ex- changed for oxygen. The blood deprived of a portion of its oxygen and laden with carbon dioxide is so-called venous blood. It is dark in color instead of bright red. On discharging its carbon dioxide and taking on oxygen, it is converted into so-called arterial blood, which is bright red. All of these pro- cesses go on much more rapidly and vigorously in an atmos- phere of pure oxygen than in air. It is for this reason that animals succumb in oxygen ; they are destroyed by the too rapid changes. On the other hand, if the supply of oxygen is unduly diminished, the transformations described, which are necessary for life, cannot go on and the animal dies of suffoca- tion. As stated above, pure oxygen may be breathed for a time; it is frequently administered to patients who are suffering de- pression because of difficulty experienced in breathing. Fishes derive their supply of oxygen by means of their gills from the oxygen dissolved in the water. In the respiration of plants, carbon dioxide is taken up by the green leaves in the sunlight, and oxygen is exhaled. In the leaf, starch is simultaneously formed, as carbon dioxide and water act on each other with elimination of oxygen. Thus, while animals are using up oxygen in breathing and are giving off carbon dioxide, plants are taking up the latter gas and re- turning oxygen to the air. Oxyhydrogen Blowpipe. When a jet of hydrogen is burned in the air, a high temperature is developed ; this may be FIG. 17 further increased by burning the jet in oxygen, or by supplying oxygen to the jet of hydrogen as it burns. The oxyhydrogen blowpipe (Fig. 17) is an arrangement for securing very high temperatures. As a rule the burner is made of brass. Hydro- gen passes in as shown and issues at the tip, where the jet is 34 OUTLINES OF CHEMISTRY lighted. Oxygen is then passed in as indicated, and thus the gases do not mix except in the jet itself. In this way explosions are avoided. The oxyhydrogen flame readily fuses platinum or silica, and is used in working such refractory materials. When the jet is directed against a piece of lime, the latter is heated to incandescence, producing a very intense white light, known as Drummond's lime light. This is used at times in projection lanterns, and for signaling purposes where a very intense light is required. Detonating Gas. We have seen that when water is decom- posed by electrolysis, two volumes of hydrogen and one volume of oxygen are produced. A mixture of these two gases in the proportions mentioned is highly explosive when ignited, for water is formed which, by the intense heat generated, is at once converted into steam, thus producing the explosion. The ex- plosive character of oxyhydrogen gas may be demonstrated in a harmless way by making soap bubbles filled with the gas and then igniting them. Not too large a quantity of the gas should be exploded at once in a room, for the report is very loud and may rupture the eardrum. Combustion of Oxygen in Hydrogen. It has been mentioned that a jet of hydrogen will burn in an atmosphere of oxygen, developing intense heat. It is equally possible to burn- a jet of oxygen in an atmosphere of hydrogen (Fig. 18). The hydrogen is first lighted at the mouth of the cylinder, and a jet of oxygen is then introduced. It ignites and continues to burn in the atmosphere of hydrogen as shown. The fact that either of these two gases may be burned in an atmosphere of the other shows the real nature of combustion, which consists of a chemical union of the two gases. The product formed is, of course, water in either case. Earlier Views of Combustion. That the combus- IG ' 18t tion of substances in the air is a process of oxida- tion was not recognized till Lavoisier showed it to be true by experiment. Before Lavoisier, the view prevailed that when a substance is burned a subtile principle flies out of it. This action dates back to antiquity. It was probably suggested by H ft/ OXYGEN 35 the rising of the smoke of ordinary fires. It was Georg Ernst Stahl (1660-1734), professor of medicine at the University of Halle, who first formulated a definite theory of combustion. He called the subtile principle, which he assumed flies out of bodies on burning them, phlogiston, which means that which is com- bustible. So, for instance, when mercury is heated in the air to 500 a red powder results, which, according to Stahl's view, would be dephlogisticated mercury. Similarly he looked upon other oxides as bodies that had been deprived of phlogiston. Anything that was combustible contained phlogiston. Thus carbon was considered very rich in phlogiston. By heating, for example, dephlogisticated lead (yellow oxide of lead) with carbon, the latter would give off phlogiston to the yellow pow- der and thus change it back to lead. In general, what we now term oxidation was regarded as dephlogistication, and what we call reduction was regarded as a process of taking on phlo- giston. The phlogistic theory dominated chemistry in the eighteenth century ; and, indeed, many chemical changes, and among them rather complicated ones, could in a way be ex- plained by means of the theory. In fact, Cavendish, Priestley, and Scheele adhered to the phlogistic theory. It was known to the adherents of the phlogistic view that when metals are calcined by heating them in the air, the result- ing powder is heavier than the original metal. In fact, this was known even a hundred years before the phlogistic theory was promulgated ; but it was not regarded as an especially vital fact in forming a correct view of combustion. It was not an age of careful quantitative experimentation, and the value of facts established by accurate measurements was frequently not seri- ously considered. And so it was that when Lavoisier pointed out that metals grow heavier when burned in the air, and argued that this means that something is added to the metal rather than subtracted from it during the process, his argument did not meet with favor, even on the part of the discoverers of oxy- gen themselves. The adherents of the phlogistic view argued that the fact that substances increase in weight when burned could not serve to prove that something, namely phlogiston, might not also fly out of the substances during the process of combustion. In order to explain the fact that substances grow heavier when burned, some of the followers of Stahl even 36 OUTLINES OF CHEMISTRY suggested that phlogiston might be a substance of negative weight. Antoine Laurent Lavoisier (1743-1794), the founder of mod- ern chemistry, laid great stress upon the increase in weight of substances during combustion, and when oxygen was discovered by Scheele and Priestley he actually demonstrated that it is this gas which unites with bodies when they are burned in the air. Thus, he heated a quantity of mercury in a retort (Fig. 19) in contact with air for twelve days. The end of the FIG. 19. retort opened into a bell jar, the opening of which was shut off from the outer air by means of mercury, as shown. The total volume of the air in the retort and bell jar was about one liter. After the apparatus had cooled, it was found that a diminution of volume of the air in the apparatus amounting to about 170 cc. had taken place. From the calcined mercury, which he col- lected, he obtained 160 cc. oxygen by heating ; and thus he showed by synthesis and analysis the real nature of calcined mercury. He further demonstrated that carbon unites with the oxygen of certain metallic oxides when heated, and that thus the metal itself is prepared by subtraction of oxygen from the calcined metal rather than by the addition of phlogiston to it. The views of Lavoisier were stoutly opposed by the follow OXYGEN 37 ers of the phlogistic theory. However, facts began to increase in favor of Lavoisier's explanations, arid when Cavendish showed that water is formed when hydrogen and oxygen unite chemically, the former's views soon triumphed. Whereas the followers of phlogiston regarded the metals and other combusti- ble elements as compound bodies containing phlogiston, we now look upon them as simple bodies capable of uniting with oxy- gen under proper conditions. REVIEW QUESTIONS 1. State how oxygen occurs in nature. 2. Make a list of five different methods of preparing oxygen in the laboratory, stating which of these methods is most commonly employed. 3. What are the important properties of oxygen? 4. What is combustion? What phenomena accompany this pro- cess? Give an illustration. 5. Define: kindling temperature, oxide, phlogiston, heat of com- bustion. 6. How does combustion in the air differ from combustion in oxygen? Illustrate. 7. State the law of multiple proportions and illustrate it by means of an example. 8. What is the relation of oxygen to animal life? To plant life? 9. What led to the overthrow of the phlogiston theory ? 10. Why does painting an iron bridge or polishing a stove prevent rusting ? 11. How many tons of oxygen are there in ten tons of water? 12. Compare the chemical and physical properties of hydrogen and oxygen ; also the methods of preparing these gases. 13. What important views did Lavoisier introduce into chemistry, anc upon what experiments were these views founded ? 14. Who discovered oxygen ? How did he prepare this gas ? CHAPTER IV WATER Occurrence. Water is found in oceans, lakes, and rivers, in the soil and in the atmosphere. It occurs in the solid, liquid, and vapor states. As snow and ice it covers the vast fields of the polar regions, the highest mountain peaks, and, during the winter, large areas of the temperate zones. Falling in form of rain, snow, and hail, water permeates the soil and forms springs, lakes, and rivers that carry it to the sea. In the atmosphere, it exists as vapor which by condensation may form fogs and clouds. The amount of aqueous vapor that the air may hold varies with the temperature. One million liters of air satu- rated with water vapor at contain 4800 grams of water, while at 20 and at 30 this amount of air will take up 17,000 and 29,840 grams of water, respectively. Ordinarily, air is saturated with water vapor to but two thirds of its capacity. When the moisture content of the air reaches but four tenths of its capacity, the air feels dry, whereas it requires nearly double this amount of humidity to cause the sensation of damp- ness. In all plants and animals, water is found in relatively large quantities. Usually organisms are made up of over fifty per cent of water. Many minerals, salts, and manufactured products contain water more or less loosely bound. Preparation. Water is formed not only when hydrogen and oxygen gases unite, but also when hydrogen acts on vari- ous oxides at high temperatures, and when compounds contain- ing hydrogen are oxidized. It forms during the process of the oxidation of the tissues of organic beings, and together with carbon dioxide is exhaled by animals. All natural waters are, chemically speaking, impure. Rain water is the purest of nat- ural waters, but even this contains air, dust, and not infre- quently estimable amounts of nitrites and nitrates of ammonium. All water that has been in contact with the soil contains some of the ingredients of the latter in solution. On evaporating 38 WATER 39 off the water, these dissolved ingredients, which are in the main salts of various kinds, are left behind as a residue. The amount of material taken up from the soil by water varies very greatly ^vith the nature of the soil. Thus from soil formed mainly from the disintegration of granite rocks, relatively small amounts of material are dissolved, whereas, from limestone soils large quantities enter into solution. By distilling natural waters, they may be freed from the dissolved, non-volatile in- gredients. In this way pure water may be obtained. The process consists of boiling the water in a retort and condensing the steam formed (Fig. 20). In this process the condenser is, FIG. 20. of course, always dissolved to a slight extent. The material of which it is constructed is somewhat soluble in water. Thus glass condensers are always somewhat attacked by water, though not sufficiently so to make the distilled water unfit for ordinary purposes. When a very pure water is desired, a block tin, or, still better, a platinum condenser, is used. On boiling water, the dissolved gases it contains are almost completely ex- pelled. Distilled water tastes flat; whereas natural waters, which contain air, have a refreshing taste. Natural Waters. The solid ingredients in natural waters vary greatly in character and amount. In oceanic waters there is about 3.5 per cent of solid material, of which 2.7 per cent consists of common salt, and the remainder mainly of chlorides and sulphates of magnesium, calcium, and potassium, together with smaller amounts of the bromides and carbonates of these CALIFOKNIA COLLEGE ^f PHARMACY . 40 OUTLINES OF CHEMISTRY metals. Some thirty elements occur in oceanic waters, most of them in very minute amounts. The water of the Dead Sea contains 22.8 per cent of saline matter and that of Great Salt Lake in Utah 23.04 per cent. Fresh water as we find it in rivers and many lakes usually contains from 0.005 to 0.15 per cent of solid material, and deep well water averages from 0.01 to 0.4 per cent. " The amount of salts contained in the waters of springs and wells varies greatly with the character of the strata of the earth's crust with which the water has been in contact. Sandstone and granitic material is less attacked by water than soils rich in the carbonates of lime and magnesia; consequently, springs and wells in limestone regions contain much more solid material in solution than those where sand- stone and granitic rocks abound. Rain water is really distilled water ; though as it falls through the atmosphere it gathers dust and dissolves the atmospheric gases. If water is gathered during a shower, that which falls after a time is much purer than that which first falls to earth. This is due to the fact that the air is fairly well washed during the earlier part of a copious rainfall. Waters containing a large amount of calcium salts are called hard waters. They do not form a lather with soap, and do not soften vegetables properly when these are boiled in such water. Furthermore, these waters produce a hard sediment consisting mainly of carbonate of lime which clogs up cooking utensils, boilers, and pipes. The purification of hard waters will be considered in connection with the salts of calcium. Potable Water. For ordinary drinking purposes, water should be colorless, odorless, tasteless, and free from materials that may prove to be deleterious to health. The mineral in- gredients commonly found in waters from springs, wells, brooks, rivers, and lakes are not injurious to the system. It is when these sources are contaminated by sewage, which very fre- quently gets into them, that the waters become dangerous to health ; for the organic animal and vegetable material in de- composing develops products which may be injurious, and often affords a place for the growth of bacteria that cause disease. For this reason, any water that clearly shows that it has been contaminated by sewage is pronounced dangerous to health. It is clear that a bacteriological examination ought to accom WATER 41 pany a chemical examination of a drinking water, for injurious organisms may be present in water even though the sewage contamination be so slight that a chemist would pronounce the water fit to drink. As common salt and organic matter and its decomposition products, especially nitrites and nitrates of ammonium, characterize sewage, the determination of the amounts of these ingredients forms the chief task of the chem- ist in analyzing a potable water. The air dissolved in natural waters renders it refreshing. As boiling kills the disease germs in water, it is frequently resorted to, especially in large cities, in cases of epidemics caused by contaminated water. The process of boiling expels the gases dissolved in the water and renders it insipid to the taste. Thus wholesome drinking water is not at all chemically pure water. The latter is not even common in chemical laboratories, for ordinary distilled water, though free from non- volatile ingredients, still contains carbon dioxide, air, and riot infrequently ammonium salts in solution. These impurities are not of consequence, however, for ordinary purposes. Very frequently, river and lake water must be used for drinking purposes, even though it is somewhat contaminated by sewage. These waters must then be subjected to purifica- tion, which commonly consists of filtration through beds of sand and exposure to the air, the oxygen of which being ab- sorbed by the water, oxidizes the organic material to simpler products that are comparatively harmless to the human system. The filters, of course, must be renewed from time to time, for the organic material collects in them and thus they may after a while themselves become a source of contamination. On a small scale, the Pasteur-Chamberland water filter is entirely efficient in freeing water from suspended matter and bacteria. This filter consists of unglazed porcelain, generally in form of a tube closed on one end, which is attached to the ordinary water cock. The water thus filters through the pores of this un- glazed porcelain under the pressure of the waterworks system Mineral Water. Waters containing special mineral ingre- dients or dissolved gases are frequently used for medicinal pur- poses. Among the mineral waters are distinguished : (1) bit- ter waters that are rich in magnesium salts ; (2) chalybeate waters that contain iron salts ; (3) sulphur waters which con 42 OUTLINES OF CHEMISTRY tain hydrogen sulphide ; (4) carbonated waters which are charged with carbon dioxide so that they effervesce ; (5) lithia waters containing lithium salts. Thermal waters are those which have a higher temperature than the surrounding atmos- phere. They frequently also contain special mineral ingredi- ents which are considered valuable for therapeutic purposes. Composition. Chemically pure water is a compound of oxygen and hydrogen, two volumes of the latter uniting with one volume of the former to form water. By weight water consists of 88.864 per cent oxygen and 11.136 per cent hydro- gen. Knowing that by the electrolysis of water two volumes of hydrogen and one volume of oxygen are produced, and hav- ing found the weight of a liter of hydrogen and that of a liter of oxygen, the composition of water by weight can readily be computed. When hydrogen is passed over copper oxide heated to a dull redness, the oxide is reduced to metallic copper and water is formed. Consequently, by heating a known amount of dry copper oxide in a tube, in a current of dry hydrogen, arid col- lecting and weighing the water formed, and also weighing the metallic copper obtained, the percentage composition of water may be computed. Obviously, the loss of weight of the copper oxide represents the oxygen that has entered into combination in the water formed ; and the difference between the weight of the latter and the oxygen given off by the copper oxide is the weight of the hydrogen in the water produced. This method of determining the composition of water was used by Dulong and Berzelius in 1819. It was also employed by Dumas in 1842 with greater refinements. The researches on the composition of water have yielded the result that for each gram of hydrogen, water contains 7.94 grams of oxygen ; that is to say, the ratio of hydrogen to oxygen in water is nearly 1 to 8. Gay-Lussac's Law of Combination of Gases by Volume. When two volumes of hydrogen and one volume of oxygen unite chemically, and the water formed is not allowed to condense to the liquid state, it is found that the steam obtained occupies two volumes, measured, of course, at the same temperature and pressure as the oxygen and hydrogen. To demonstrate this, the apparatus of Hoffman (Fig. 21) is convenient. The inner WATER 43 long eudiometer tube A is filled with mercury and then in- verted in the mercury bath B. Thus, a Torricelli vacuum is formed in A, whose upper end is provided with a pair of plati' FIG. 21. 44 OUTLINES OF CHEMISTRY num terminals, across which an electric spark may be passed by connecting with the induction coil C. The eudiometer tube A is placed inside of the larger tube D, which is filled with steam from the boiler .E. By this means the eudiometer tube is heated to the boiling point of water. If now a mixture of two volumes of hydrogen and one volume of oxygen is intro- duced into the eudiometer tube A, the volume carefully noted, and then the mixture is exploded by passing the electric spark, the resulting water va.por will be found to have two thirds of the volume of the mixture of the oxy hydrogen gas introduced, when the level of the mercury in the eudiometer has been restored to its original place. Therefore, two volumes of hydrogen and one volume of oxygen unite to form tivo volumes of water vapor. This very simple relation is typical of the volume relations in general which have been found to obtain when gases combine chemically. Expressed in general terms we may say : When gases combine chemically with one another, the volumes that unite bear a simple relation to one another ; and if the product formed be gaseous, its volume also bears a simple relation to the volumes of the original gases that have entered into combination. This law was discovered by Joseph Louis Gay- Lussac, professor of chemistry at the Sorbonne, Paris. We shall meet with further specific illustrations of this law, which is known as the law of Gay-Lussac of combination of gases by volume. It is of great importance in chemistr}^ as will appear in the succeeding chapters. Properties of Water. In thin layers, pure water is colorless, but in deep layers it has a greenish blue color. This explains' the beautiful hue of the waters of the sea and many lakes. River waters are commonly brownish in color, due to the humus material which they contain from the soils through which they have coursed. The freezing point of water is taken as the zero of the centigrade scale, and the boiling point under a pressure of 76 cm. of mercury is taken as the 100 point of that scale. At and below 360 C water may be condensed to a liquid ; above this point, which is the critical temperature', water is a gas which cannot be condensed to a liquid even though very high pressures be applied. Like liquids in general, water is but slightly compressible. Thus, by placing a liter of water at 20 under a pressure of two WATER 45 atmospheres its volume is diminished only by 0.046 of a cubic centimeter. The volume of a given weight of water varies with the temperature. Water expands in volume when heated above 4, and also when cooled below that temperature to its freezing point. Water, therefore, has its maximum density at 4. Most substances show a continuous diminution in volume when cooled. The fact that water expands when cooled below 4 is therefore a very exceptional behavior. At 4 a cubic centimeter of water weighs one gram. Water at its maximum density is commonly taken as the standard liquid with which the weights of equal volumes of other liquids and solids are compared. In other words, water at 4 is the standard of comparison of the specific gravities of liquids and solids. At 100 the volume of water is about 4.3 per cent greater than at 0. The amount of heat required to raise the temperature of one gram of water one degree is called a calorie (cal.) ; it is the unit used in the measurement of heat. It requires 80 cal. to trans- form one gram of ice at to water of the same temperature ; i.e. the latent heat of fusion of ice is 80 cal. To convert one gram of water at 100 into vapor of the same temperature re- quires 537 cal., which is the so-called latent heat of evapora- tion of water. The specific heat, the latent heat of fusion, and the latent heat of evaporation of water are very high indeed, as compared with similar constants of most other substances. When water freezes it expands, and the ice at occupies 1.0908 times the volume of the water at the same temperature. This behavior of water is again unusual, for most substances contract during the act of congealing, thus forming a solid that is heavier than the liquid. The fact that water increases in volume as it solidifies is an important factor in the disintegra- tion of rocks, for the force exerted by water in freezing is enormous. The bursting of frozen water pipes and other con- tainers in winter is also due to the expansion of water in freez- ing. But the fact that ice is lighter than water is of further importance in nature; for were it not for this, many of our lakes and rivers would freeze to the bottom in winter, and thus fishes and other organisms in these waters would be destroyed. The huge masses of ice that would accumulate in winter also would 46 OUTLINES OF CHEMISTRY materially reduce the temperature for the remainder of the year. Supercooled Water. Ice melts at 0, but when water is cooled to 0, it does not necessarily freeze. In fact, water may be cooled several degrees below zero and still be liquid. Water in this condition is said to be supercooled, or in a metastable condition. If water thus supercooled is brought in contact with a piece of ice, the whole mass freezes to a solid. If super- cooled water is cooled still further, a point is finally reached at which it will congeal without being touched with ice. Super- cooled water may be kept in the liquid condition for a long time. Sometimes shaking, jarring, or brisk stirring induces freezing of supercooled water, but this is not necessarily the case. The lower the temperature of the metastable water, the more likely is jarring to induce ice formation. However, touch- ing supercooled water with ice, always causes freezing. The freezing point and the melting point of water are the same ; namely, 0. This is the temperature at which ice and water are in equilibrium with each other at ordinary pressure. Raise the temperature above and all ice disappears ; cool below in the presence of ice and the whole mass freezes; i.e. liquid water disappears. Similarly, the freezing or melting point of any solid is an equilibrium temperature at which the solid and liquid can exist side ly side in contact with each other without change. Change of Freezing Point with Pressure. The freezing point of ice is altered by change of pressure. Since water ex- pands on congealing, an increase of pressure on its surface would make it more difficult for ice to form. In other words, we should then have to cool water under pressure to a lower temperature in order to freeze it ; or what comes to the same thing, ice under pressure melts at a lower temperature than at ordinary pressure. Substances which do not expand, but con- tract as they congeal, act just the opposite from water in this respect when put under pressure ; i.e. increase of pressure causes them to freeze at a higher temperature, the increase of pressure aiding contraction which accompanies the solidification in these cases. These instances of the alteration of the freezing point of substances with increase of pressure are illustrations of a far* WATER 47 reaching principle which may be stated as follows : When chemical or physical equilibrium exists, and one of the factors upon which it depends is altered, a change is produced which opposes the first alteration. This is known as the principle of Le Chatelier, who first enun- ciated it. Thus increase of pressure upon any solid or liquid tends to diminish its volume. When ice and water exist in equilibrium at and the pressure is increased, the ice melts, which process is accompanied with a diminution in volume, which has a tendency to lessen the pressure. In the case of in- crease of pressure upon solid and liquid tin in equilibrium with each other at the melting point of tin, the liquid tin will congeal, for thus contraction is brought about and consequently the pressure exerted upon the tin is lessened. The principle of Le Chatelier is of far-reaching importance, and we shall have further examples of it later. Crystalline Nature of Ice. When water solidifies, it tends to take on regular forms. This is evident from the frost on the windows, from the shapes of snowflakes, and the radial structure of ice. The needles that form as ice congeals tend to arrange themselves so as to form angles of 60. These forms are most perfect perhaps in the case of snowflakes, which as they fall on a still day are frequently quite large and perfect. Water crystallizes in the hexagonal system, which is one of the six systems into which all known c^stals may be divided (see Crystal Systems). Not only do crystals exhibit outward regularity of form, but they also show different degrees of hardness, tenacity, refrangibility, light absorbing power, etc., in different directions. We therefore distinguish crystalline substances, which show these characteristics, from amorphous substances, which do not have regularity of form and which exhibit the same properties irrespective of the direction through them. Ice is a typical crystalline substance, while glass is a typical amorphous substance. Amorphous means without form. Crystalline substances commonly have a definite melt- ing point and definite solubility, while amorphous substances do not. Thus glass has no definite temperature at which it melts. It softens when heated and gradually passes through all stages 48 OUTLINES OF CHEMISTRY of gradations to the liquid state on further heating. Not so with water, for it has a sharp melting point at 0. Many definite chemical compounds tend to form crystals; and since the same compound tends to take on the same shape under given conditions, the study of crystallography is of value to the chemist in aiding him in purifying and identifying sub- stances. However, many definite chemical compounds have never been obtained in the crystalline condition. Compounds with Water. Many salts, like copper sulphate, Glauber's salt, and Epsom salt, crystallize with water. The water in these salts is spoken of as water of crystallization. On exposure to the air, some of these salts lose a portion of this water of crystallization and become opaque or crumble to a powder. They are said to effloresce. Other salts, like calcium chloride, have such a strong attraction for water, that on ex- posure to the air they take on water from the air and finally become completely dissolved. They are said to deliquesce. Substances that have attraction for water are also called hygro- scopic. Concentrated sulphuric acid, phosphorus pentoxide, calcium chloride, lime, and caustic potash are strongly hygroscopic. Gases passed over these are deprived of their moisture, and many solids left with them for a time in a confined space are dried or desiccated. A typical form of desiccator is shown in Fig. 22. The strongly hy- groscopic substance is placed in the bottom of the vessel, and the substance to be dried is placed on the support in the upper part of the apparatus. Through the cock the air can be exhausted from the apparatus and thus the drying process be aided still further. Such desiccators are frequently used in chemical work, for many substances like glass, porcelain, and even metallic utensils attract moisture and form a film of it on their surface. This film varies in thickness and weight with the degree of humidity of the atmosphere. In accurate quantitative experiments it is FIG. 22. WATER 49 necessary to eliminate this film of moisture, and for this purpose desiccators are commonly used. If permissible, the objects are heated to drive off moisture, and then cooled in the desiccator. If heating is not permissible in a given case, the substance is introduced into the desiccator and kept there for a longer time, frequently in a vacuum. It is evident that in a desic- cator, the drying material used must have a greater affinity for water than the substance to be dried. When water simply adds itself to another compound, the product is commonly termed a "hydrate. Such hydrates are quite common ; thus ferric chloride forms several hydrates with water, which follow the laws of definite and multiple proportions. When oxides unite with water, or when a metal like sodium acts on water, crowding out a portion of its hydrogen, the product formed is commonly termed an hydroxide. In these cases it is always possible to regard the resulting substance as water in which a portion of the hydrogen has been replaced by another element. So when lime and water act on each other they unite and form slaked lime, which is hydroxide of calcium. Caustic potash, which may be formed by the action of potas- sium on water, with concomitant evolution of hydrogen, is potassium hydroxide. Water as a Solvent. Many substances are soluble in water. Indeed, of so many is this the case that water has at times been termed a universal solvent. There are, however, very many compounds that are not soluble in water. In general, the ordinary acids, alkalies, and salts used in the chemical lab- oratory are soluble in water to a greater or lesser degree. The rocks of the earth's crust are all soluble to some extent, though to a very slight degree in some cases ; yet this slight solubility of rocks is of the highest importance to plants whose rootlets are thus able to take up mineral matter needed for their economy and growth. In geological transformations, such as the weathering of rocks, the formation of soils, and the deposition of ores, this slight solubility is nevertheless the determining factor, without which these processes could not proceed. Fats, waxes, oils, and kindred substances are generally not soluble in water. Yet many of these have some degree of 50 OUTLINES OF CHEMISTRY attraction for water, which is shown by the fact that they are often slightly hygroscopic. And again, in the bodies of plants, and particularly in those of animals, fatty material is very closely associated with tissues which are rich in water. So that although fats are generally not soluble in water to speak of, yet in many cases there is evidence that some degree of attraction between them and water does exist. Solutions will receive further consideration later. REVIEW QUESTIONS 1. Describe three different methods of demonstrating that water consists of hydrogen and oxygen and nothing else. 2. What impurities are found in rain water? In well water? In ocean waters? 3. What impurities are removed from a natural water by the pro- cess of distillation ? Why are natural waters never pure ? 4. (a) Mention the essential characteristics which every good drinking water ought to possess. (6) What tests must be made to ascertain whether a water is fit to drink? (c) Why is it well to boil contaminated water before drinking it ? 5. What impurities are removable from water by the process of nitration ? 6. Mention five different kinds of mineral waters. 7. How much hydrogen could be prepared from 50 kilograms of water? Give two methods that are essentially different in character by means of which the hydrogen could be liberated from the water. State also what volume the hydrogen would occupy under standard conditions. 8. State the Law of Gay-Lussac of the combination of gases by vol- ume, and show how the composition of water, by volume, illustrates this law. 9. Make a list of the more important physical properties of water. 10. Define the following terms : water of crystallization, efflores- cence, deliquescence, hydrate, hydroxide. Give an example in each case. 11. Name some of the more important substances that are soluble in water. What substances are not soluble in water? Why is the slight solubility of the rocks of the earth's crust nevertheless of great impor- tance in the economy of nature? 12. What is the difference between an unsaturated, a saturated, and a supersaturated solution? 13. How may a supersaturated solution of common salt in water be prepared ? CHAPTER V HYDROCHLORIC ACID AND CHLORINE Preparation and Properties of Hydrochloric Acid. When con- centrated sulphuric acid is poured upon common salt, an effer- vescence ensues, a gas being evolved which is colorless, has a pungent odor, and is neither combustible nor a supporter of combustion. This gas has a very sour taste, and produces suffocation when inhaled in quantity. It reddens moist blue litmus paper, and is very soluble in water. At one volume of water will absorb 503 volumes of the gas, while at room tem- perature about 450 volumes are thus absorbed. This gas, which was at first called " spirit of salt," was discovered by Johann Rudolf Glauber in 1658. It is hydrochloric acid. Priestley called it " marine acid air " ; he collected the gas over mercury. Hydrochloric acid is sometimes emitted during volcanic erup- tions. It also occurs in the gastric juice of man and other animals. In normal condition the human gastric juice contains about 0.33 per cent of hydrochloric acid, which is essential in the process of digestion. Hydrochloric acid comes in the market as a solution of the gas in water. It also goes by the name of muriatic acid. On dissolving pure hydrochloric acid gas in distilled water, a color- less solution is obtained. However, much of the commercial muriatic acid is colored yellowish by the impurities, especially salts of iron, that it contains. The attraction between hydrochloric acid gas and water is so great that the gas fumes strongly in the air. This is due to the fact that it condenses moisture from the air in drops, which consist of an aqueous hydrochloric acid solution. When the gas is conducted into water, heat is evolved. The thermal change accompanying the solution of any substance is termed the heat of solution (see Thermochemistry). Aqueous solutions of hydrochloric acid are heavier than water. Thus, a solution of specific gravity 1.024 at 15 contains 5 per cent hydrochloric E 51 52 OUTLINES OF CHEMISTRY acid, while solutions having the specific gravities 1.049, 1.100, 1.152, and 1.200 contain 10, 20, 30, and 40 per cent, respec- tively. A solution which is saturated with hydrochloric acid at 15 contains 42.9 per cent and has a specific gravity of 1.212. The usual " pure," commercial, concentrated hydrochloric acid is about 38 per cent strong and has a specific gravity of 1.19. It fumes strongly when exposed to the air. - On boiling a saturated solution of hydrochloric acid, the gas is in part expelled, and finally a 20.2 per cent solution with a boiling point of 110 is obtained. At ordinary pressure, this solution distills over without change of composition. The same strength of solution is finally obtained when a dilute solu' tion is boiled. In this case water is mainly expelled until the solution reaches a strength of 20.2 per cent, when it distills over without decomposition. The final acid thus obtained at different pressures, however, has a slightly different composition. Pure, dry hydrochloric acid gas may be condensed to a liquid at 10 under a pres- sure of 40 atmospheres. Under atmospheric pressure the liquid, which is colorless, boils at - 84 and freezes at about - 110. Composition and Chemical Behavior of Hydrochloric Acid. When metallic sodium is introduced into pure hydrochloric acid gas, the metal burns in the gas, forming common salt and hydrogen. This fact shows that hydrogen is one of the con- stituents of hydrochloric acid. The right limb of the apparatus (Fig. 23) is filled with pure, dry hydrochloric acid gas. The press P, which fits securely on the top of the glass tube, contains metallic sodium. When the latter metal is pressed out into the tube A, by turning the screw of the press, the sodium and hydrochloric acid react and form common salt and hydrogen with concomitant evolu- tion of light and heat. If the level in the F JG> 23. limbs A and B is kept constant by pouring HYDROCHLORIC ACID AND CHLORINE 53 mercury into B as required, it will be seen that when furthei addition of sodium no longer produces any action in A, the hydrogen obtained occupies just one half of the volume of the original hydrochloric acid gas. Hydrochloric acid is a very powerful acid and acts strongly on many metals, hydrogen being liberated and chlorides of the metals being formed during the reaction. When a concentrated aqueous solution of hydrochloric acid is subjected to electrolysis {Fig. 24), equal volumes of hydrogen and FIG. 24. a greenish yellow gas, chlorine, appear. Carbon electrodes are used in this electrolysis, lor platinum would be attacked by the chlorine. This apparatus, designed by Lothar Meyer, differs from that in Fig. 2, because chlorine when collected over an aqueous hydrochloric acid solution under pressure is quite appreciably absorbed, so that the volume of the chlorine gas would be diminished. When equal volumes of dry chlorine and dry hydrogen con- tained in the two parts of the strong tube (Fig. 25) are allowed to mix by opening the stopcock, and the mixture is then ex- posed to diffused daylight, hydrochloric acid is formed, and neither hydrogen nor chlorine is left uncombined. Moreover, 54 OUTLINES OF CHEMISTRY the volume of the hydrochloric acid gas formed is exactly equal to that of the hydrogen plus chlorine. That is, equal volumes of hydrogen and chlorine unite to form hydrochloric acid without change of volume, which is demonstrated by the fact that when the stopper at the lower end of the tube (Fig. 25) is removed under mercury after the hydrochloric acid has formed, neither gas escapes nor mercury enters the tube. In direct sunlight or on exposure to a strong magnesium flash light the union takes place with explosive violence. Thus it is that one volume of hydrogen unites with one volume of chlorine to form two volumes of hydro- chloric acid gas. This is an- other example illustrating the law of Gay-Lussac of combina- tion of gases by volume. It has been found that one volume of chlorine is 35.45 times as heavy as an equal volume of hydrogen. From this and the fact that equal volumes of hydrogen and chlorine unite to form hydrochloric acid, it is evident that by weight, 1 part of hydro- gen unites with 35.45 parts of chlorine to form hydrochloric acid. According to H. Sainte-Claire Deville, hydrochloric acid gas is partially decomposed into hydrogen and chlorine when heated to temperatures of 1300 or above. Enormous quantities of -hydrochloric acid are manufactured as a by-product of the Le Blanc soda process (which see). Occurrence, History, and Preparation of Chlorine. Chlorine occurs in nature only in combination with other elements. The chlorine-bearing compounds are chiefly common salt, the chloride of sodium, and the chlorides of potassium, magnesium, and cal- cium. Chlorine is also found in the native chlorides of lead, copper, and silver. In combination with hydrogen, it occurs in the gastric juice, and as sodium chloride and potassium chlo- ride it forms an essential constituent of the bodies of animals. FIG. 25. HYDROCHLORIC ACID AND CHLORINE 55 It is also an important constituent of plants, in which it is probably mainly combined with potassium. Chlorine was first prepared in the free state by Scheele in 1774, who treated manganese dioxide with hot hydrochloric acid. He called the gas " dephlogisticated hydrochloric acid," for at that time hydrogen was regarded as the phlogiston of Stahl. However, in 1785 Berthollet, who belonged to the anti- phlogistic school, called chlorine "oxidized hydrochloric acid." He was of the opinion that chlorine contained oxygen, and this view prevailed till 1807 ; when, on the basis of their researches, Gay-Lussac and Thenard showed the gas to be a simple sub- stance, i.e. an element. The gas was named chlorine by Sir Humphry Davy. The name comes from the Greek, meaning greenish yellow. We have seen that chlorine is one of the products of the elec- trolysis of hydrochloric acid. The simplest way of preparing chlorine is by treating hydrochloric acid with an oxidizing agent, whose oxygen unites with the hydrogen of the hydrochloric acid, thus forming water and setting the chlorine free. As such an oxidizing agent, manganese dioxide is commonly em- ployed. Chlorine may be formed by treating manganese diox- ide with the aqueous solution of hydrochloric acid and heating gently ; or by mixing common salt with manganese dioxide and treating the mixture with sulphuric acid. In the latter case, the sulphuric acid acts on the sodium chloride forming hydrochloric acid, which then acts upon the manganese dioxide as before. In these processes manganous chloride is also formed. In place of manganese dioxide, other oxidizing agents, like po- tassium dichromate, potassium chlorate, red lead, or bleaching powder, may be employed. In preparing chlorine by subtract- ing the hydrogen from the hydrochloric acid by an oxidizing agent, the oxygen of the air may be employed. By passing a mixture of air and hydrochloric acid at about 400 over porous bricks which have been soaked with copper sulphate solution, chlorine is liberated. The method is called the Deacon process and is used on a commercial scale. In this process cupric chlo- ride is formed, and this is decomposed into cuprous chloride and chlorine. The cuprous chloride is then again converted into cupric chloride, which suffers decomposition as before, and so on. 56 OUTLINES OF CHEMISTRY Properties of Chlorine. Chlorine is a greenish yellow gas which is 2.6 times as heavy as air. It has a very disagreeable odor, attacks the mucous membranes strongly, giving rise to a cough, and causes death by suffocation. At it may be liquefied by means of six atmospheres of pressure. The criti- cal temperature of the gas is 146, and the critical pressure is 84 atmospheres. Thus, at ordinary temperatures chlorine is a condensable vapor. Under atmospheric pressure it becomes a liquid at 34, its boiling point. Liquid chlorine has a golden yellow color. At 102 it freezes, forming yellow crystals. Liquid chlorine is now obtainable in the market in lead-lined steel flasks (Fig. 13). In this form it is shipped for use in laboratories and various industrial plants. Chemically, chlorine is a very active element, combining at ordinary temperatures with evolution of light and heat with sodium, phosphorus, arsenic, antimony, and many other metals when these are introduced into an atmosphere of the gas in the form of powder or very thin sheets. In all such cases chlorides form by direct union of the chlorine with the other ele- ment. An apparatus for burning arsenic in chlorine is shown in Fig. 26. The cork fits loosely. When the test tube containing the powdered arsenic is raised, FlG 26 the latter falls into the bottle and unites with the chlorine with evolution of light. Chlorine does not act directly on carbon or nitrogen ; but chlorides of these elements may be formed by the indirect methods of double decomposition, as will appear later. Chlo- rine and hydrogen have a strong affinity for each other. A jet of hydrogen will burn in an atmosphere of chlorine, or a jet of chlorine in an atmosphere of hydrogen. In either case hydro- chloric acid is formed as the product. A lighted taper or gas jet will continue to burn in chlorine, forming hydrogen chloride and carbon, which forms dense clouds of soot. Similarly, when a strip of filter paper moistened with turpentine is introduced into an atmosphere of chlorine, hydrochloric acid is formed, much soot escapes in dense clouds, and the paper instantly catches fire. HYDROCHLORIC ACID AND CHLORINE 57 Chlorine is soluble in water. At 10 1 volume of watei absorbs about 3 volumes of chlorine, and at 50 about 1.5 volumes. The solution is commonly called chlorine water. When it is exposed to sunlight, the chlorine gradually unites with the hydrogen of the water, forming hydrochloric acid and oxygen. . By filling a retort (Fig. 27) with chlorine water and exposing it to sun- light, the solution becomes colorless, hydrochloric acid being formed and oxygen liberated. The latter gas collects in the bulb, as shown in Fig. 27. By tilting the retort, this gas may be brought into the neck of the vessel and tested with a glowing splint. Because chlorine thus unites with the hydrogen of water and sets oxygen free, which in turn is capable of oxidizing sub- stances, chlorine is spoken of as a powerful oxidizing agent. Upon this power to set free oxygen from water depends the bleaching action of chlorine. When moist flowers, green leaves, colored cloth, and paper on which marks have been made with ordinary ink are introduced into an atmosphere of chlorine, they are bleached ; that is, the color is destroyed. Moisture is essential to have the bleaching take place; for the chlorine unites with the hydrogen of the water, forming hydrochloric acid and setting oxygen free. The latter then unites with the coloring matter and destroys it. Printer's ink consists largely of carbon, which at ordinary temperatures is not attacked either by oxygen or chlorine ; it consequently is not bleached. It should further be stated that fabrics dyed with some of the aniline dyestuffs also retain their color, even when treated with chlorine water. By the action of chlorine on water, some hypochlorous acid is always formed. Other Uses of Chlorine. The oxygen liberated when chlorine acts upon water as explained is very destructive to organic life ; for this reason chlorine is used as a disinfectant. Fungi and disease germs are rapidly destroyed by the action of chlorine. Chlorine is also sometimes used in extracting gold from its ores. By direct union with the metallic gold, the chloride of 58 OUTLINES OF CHEMISTRY that metal is formed; and this salt being readily soluble in water, can then be leached out of the ores. Some Compounds of Chlorine with Oxygen. Chlorine does not form compounds with oxygen by direct union of the two gases. However, by the indirect method of double decom- position, compounds of oxygen and chlorine may be obtained. These compounds are gases which readily decompose. The compounds of oxygen and chlorine will be considered in Chapter VIII. Here only two of these compounds will be mentioned briefly. Chlorine Monoxide. When dry chlorine is passed over red oxide of mercury in the cold, a pale yellow gas is formed, which does not have the greenish tint of the chlorine and readily decomposes with explosive violence, even when moderately heated. At 5 it may be condensed to a liquid of orange- yellow color, which readily explodes in sunlight or on slight heating, at times even on pouring it from one dish to another. The gas is soluble in water. One volume of water absorbs about 200 volumes of chlorine monoxide gas at 0. This substance is an oxide of chlorine, and consists of 35.45 parts of chlorine to 8 parts of oxygen by weight. It is called chlorine monoxide. Chlorine Dioxide. By carefully treating powdered potassium chlorate with concentrated sulphuric acid added in very small quantities at a time, a heavy, deep yellow gas is evolved which has a very disagreeable odor, attacks mercury, and is readily soluble in water. It is very unstable, exploding in the sun- light, or when heated by means of the electric spark or a hot iron rod. In the cold, it may be condensed to a liquid of dark ?ed color, which is of a highly explosive nature. This compound is an oxide of chlorine, which contains 35.45 parts of chlorine and 32 parts of oxygen by weight. It is called chlorine dioxide or chlorine peroxide. Thus, in the case of these two oxides of chlorine we have another illustration of the law of multiple proportions ; for in the monoxide 35.45 parts of chlorine are united with 8 parts of oxygen by weight, whereas in the peroxide 35.45 parts of chlorine are united with 4 times 8 parts of oxygen. The Law of Reciprocal Proportions. We have learned that in water hydrogen and oxygen are united in the proportions of HYDROCHLORIC ACID AND CHLORINE 69 1 part of hydrogen to 8 parts of oxygen by weight. In hydro- chloric acid 1 part of hydrogen is united with 35.45 parts of chlorine by weight. In chlorine monoxide we find that 35.45 parts of chlorine are united with 8 parts of oxygen by weight ; and in chlorine peroxide 35.45 parts of chlorine are united with 4 times 8 parts of oxygen. Thus we see that the proportions by weight in which hydrogen and oxygen combine, and in which hydrogen and chlorine com- bine, also determine the proportions in which chlorine and oxygen combine with each other. This is an illustration of a general law which holds in all chemical combinations. It may be stated thus in general terms : If three elements, A, B^ and (7, are able to unite to form chemical compounds with one another, the proportions by weight with which A and B unite to form the compound AB, and the proportions in tvhich A and C unite, also determine the proportions in which B and C unite with each other. This law is known as the law of reciprocal proportions. It was discovered by Jeremias Benjamin Richter, and is one of the fundamental laws of chemical combination by weight. In the further study of chemistry, the student will meet numerous illustrations of this law. REVIEW QUESTIONS 1. Describe the common way of preparing hydrochloric acid gas. 2. About how much hydrochloric acid gas will a liter of water absorb at ordinary temperature, 15 C. ? What different names are given to an aqueous solution of hydrochloric acid gas ? What properties has the gas ? 3. How many liters of hydrogen will be required to unite with 25 liters of chlorine to form hydrochloric acid gas? What would be the volume of the latter? What law is illustrated by these facts? 4. Describe two different methods of proving that hydrochloric acid consists of hydrogen and chlorine, and that these are present in the pro- portion of 1 part of hydrogen to 35.5 parts of chlorine by weight. 5. Why is the method of preparing chlorine by acting upon hydro- chloric acid with any of the following substances one and the same in principle : manganese dioxide, potassium chlorate, red lead, potassium dichromate? Give a method of preparing chlorine which is essentially different in principle. 6. Mention the properties of chlorine. What use is made of chlorine? 7. Name two oxides of chlorine and state what law these illustrate. 8. What is the law of reciprocal proportions ? Illustrate. 9. How much chlorine is there in a barrel of common salt containing 280 Ib. net? How much hydrochloric acid could be made from this salt? CHAPTER YI THE LAWS OF COMBINING WEIGHTS AND COMBINING VOLUMES AND THE ATOMIC AND MOLECULAR THEORIES Retrospect. In the preceding chapters we have found that certain general laws regulate the quantities in which the chemical elements combine to form compounds. The laws governing the combination of the elements by weight are as follows : (1) The Law of Definite Proportions. A chemical compound always contains the same elements in the same proportions by weight. No matter when, where, or by what process hydro- chloric acid, for example, is formed, it always contains only the elements hydrogen and chlorine in the proportions of 1 gram of hydrogen to 35.45 grams of chlorine. Water always consists of hydrogen and oxygen united in the proportions of 1 gram of hydrogen to 8 grams of oxygen. Common salt is made up of 23 parts of sodium to 35.45 parts of chlorine by weight ; and similarly every other chemical compound always has exactly the same invariable qualitative and quantitative com- position. The law of definite proportions, it will be recalled, was discovered by Lavoisier. (2) The Law of Multiple Proportions. When any two ele- ments, A and B, form more than one compound uith each other, the amounts of B that unite with one and the same weight of A are simple rational multiples of one another. Thus iron and sulphur form ferrous sulphide, which consists of 28 grams of iron to every 16 grams of sulphur ; and pyrite or fool's gold, a native sulphide of iron, always contains 28 grams of iron to every 32 (i.e. 2 times 16) grams of sulphur. Again, in chlorine monox- ide, every 35.45 grams of chlorine are united with 8 grams of oxygen. In chlorine peroxide, every 35.45 grams of chlo- rine are united with 32 (i.e. 4 times 8) grams of oxygen; and in chlorine heptoxide (which see) every 35.45 grams of chlorine 60 FUNDAMENTAL LAWS AND THEORIES 61 are combined with 56 (i.e. 7 times 8) grams of oxygen. In the oxides of lead the proportions of lead and oxygen by weight are as follows : (a) In the black oxide, (6) In the yellow oxide, (me cases the symbols are derived from the Latin names of the elements in the manner mentioned. So the symbol for silver, argentum, is Ag ; for sodium, natrium, Na ; for mer- jury, hydrargyrum, Hg ; for lead, plumbum, Pb ; etc. A com- pete list of all the names and symbols of the elements will be fiven later. Having thus adopted the symbols for the elements, the com- >ounds would naturally be designated by simply writing side >y side the symbols of the elements that occur in the com- >ounds. But since the elements always enter into combina- ;ion in the ratio of their combining weights, it is easy to have 66 OUTLINES OF CHEMISTRY the symbol of a compound stand for both its qualitative and its quantitative composition. This is readily accomplished by letting the symbol of each element stand for not only the verbal name of that element, but also for its combining weight. Thus the symbols H and Cl would stand not only for hydro- gen and chlorine, but also for 1 part of hydrogen by weight, and 35.45 parts of chlorine by weight, respectively. And so the symbol HC1 stands not only for hydrochloric acid, but it also tells us that in that compound 1 part of hydrogen is com- bined with 35.45 parts of chlorine by weight. Likewise, the symbol for common salt, NaCl, denotes that this compound con- sists of sodium and chlorine united in the proportions of 23 parts of sodium to 35.45 parts of chlorine by weight. Again, the symbol HO, which was formerly used for water, denoted that this compound consists of hydrogen and oxygen united in the proportions of 1 to 8 by weight. The symbol of ferrous sulphide, FeS, denoted that in this compound iron and sulphur are present and in the proportion of 28 parts of iron to 16 parts of sulphur by weight. The symbol FeS 2 , the 2 being used as a subscript to the S, was the symbol for pyrite, and denoted that in it 28 parts of iron are combined with 2 times 16 parts of sulphur, i.e. two combining weights of sulphur. In general, whenever more than one combining weight of an element enters into the compound, that fact is indicated by the appropriate figure used as a subscript. So, for instance, the symbol for red lead is Pb 3 O 4 , indicating that in that compound 3 combin- ing weights (3 x 103.5) of lead are united with 4 combining weights (4 x 8) of oxygen. This mode of designating chemical compounds by having the symbols stand for equivalent or com- bining weights was in vogue for many years ; and with a slight modification it is still in use at present. The nature of this modi- fication lies merely in the fact that we do not always designate the chemical equivalent by the symbol of the element, but fre- quently the symbol stands for some other simple multiple of the chemical equivalent, for reasons that will presently be explained. The Atomic Theory of Matter. The fact that the chemical elements always unite in definite proportions by weight in accordance with the three laws of definite, multiple, and recip- rocal proportions, finds a ready explanation in a simple as- sumption as to the nature of matter. If we assume that each FUNDAMENTAL LAWS AND THEORIES 67 elementary substance is made up of extremely minute, ultra- microscopic, indivisible particles, termed atoms (from the Greek meaning indivisible), and that these atoms are of exactly the same weight and also otherwise alike in the case of any one element, but different in weight and other properties in the case of different elements, and that chemical compounds are formed by the union of the atoms of the various elements with one another, the experimental facts expressed in the laws of definite, multiple, and reciprocal proportions are readily ex- plained. So, for instance, by this hypothesis, hydrogen would be considered as made up of minute particles, atoms, all of the same weight and otherwise also alike. Chlorine would similarly be regarded as composed of atoms which are of the same weight and otherwise alike among themselves, but quite different in weight and other properties from the atoms of hydrogen or those of any other element. Each element would similarly be composed of atoms that are alike in weight and otherwise, but different in weight and other characteristics from the atoms of all other elements. Since the atoms of each element are assumed to be indivisible, in forming compounds a whole number of atoms of one element must always unite with a whole number of atoms of another element or elements. Consequently the proportions by weight in which, for instance, two elements A and B unite with each other to form a compound AB, are proportional to the atomic weights of A and B ; in other words, the combining weights of the elements are proportional to the atomic weights. So, for example, 1 gram of hydrogen unites with 35.45 grams of chlorine to form 36.45 grams of hydrochloric acid, consequently, in the light of the atomic theory, 1 grain of hydrogen must contain a definite whole number of atoms of hydrogen, and similarly the 35.45 grams of chlorine must contain a definite whole num- ber of atoms of chlorine. If we let x represent the actual weight of one of the hydrogen atoms, and n the number of atoms of hydrogen in 1 gram of hydrogen, then xn equals 1 gram. Similarly, if we let y represent the weight of 1 atom of chlorine and n' the number of atoms of chlorine in 35.45 grams of that gas, then yn r equals 35.45 grams. We may consequently write the relation xn : yn f : : 1 gram : 35.45 grams. 68 OUTLINES OF CHEMISTRY It is clear that we might deduce a similar equation in the case of the union of any two or more of the chemical elements. Let us now view this equation more closely. It contains four un- known quantities : namely, the atomic weight of hydrogen x, the atomic weight of chlorine ?/, the number of atoms of hydrogen n, in 1 gram of hydrogen, and the number of atoms of chlorine n 1 , in 35.45 grams of chlorine. Of course, neither of these values can be ascertained from the equation as it stands. If now we arbitrarily choose some definite value for either x or ?/, say we assume with Dalton the atomic weight of hydrogen as 1, the x will disappear from our equation as an unknown quantity ; still we should have the three unknown quantities n, n', and y present in the equation. There is no way of determining how many atoms of hydrogen unite with how many atoms of chlo- rine in forming hydrochloric acid, and so it is customary to make the simplest possible assumption here, namely, that 1 atom of hydrogen, unites with 1 atom of chlorine in forming a particle of hydrochloric acid. On the basis of this assump- tion, it becomes clear that 1 gram of hydrogen would contain as many atoms of hydrogen as 35.45 grains of chlorine contain atoms of chlorine ; or in other words, in our equation n equals n'. Since we have assumed that x equals 1, and also that n equals n', the equation becomes 1 : y : : 1 gram : 35.45 grams ; whence y equals 35.45. That is, the atomic weight of chlorine is 35.45, if the atomic weight of hydrogen is assumed to be 1, and it is further assumed that in forming hydrochloric acid 1 atom of hydrogen unites with 1 atom of chlorine. Similarly, the proportions by weight in which hydrogen and oxygen unite to form water, namely, 1 to 8 (if we were to make the simplest assumption, as Dalton did, that in forming water 1 atom of hydrogen unites with 1 atom of oxygen), lead to the conclusion that the atomic weight of oxygen is 8. And this value was assigned to it by Dalton, though it is not the one used at present, as will be explained shortly. However, if we thus take the atomic weight of hydrogen as 1, that of oxygen as 8, and that of chlorine as 35.45, then since in chlorine monoxide every 35.45 grams of chlorine are combined with 8 grams of oxygen, we should have in this compound 1 atom of chlorine united with 1 atom of FUNDAMENTAL LAWS AND THEORIES 69 oxygen. In the chlorine peroxide again, in which every 35.45 grams of chlorine are combined with 32 grams of oxygen, we should have 1 atom of chlorine combined with 4 atoms of oxygen. Thus on this basis the formulae for hydrochloric acid, water, chlorine monoxide, and chlorine peroxide would be HC1, HO, CIO, and C1O 4 , respectively. In these formulae the symbols of the elements simply stand for the combining or equivalent weights. Thus by introducing the assumptions : (1) that the elements are made up of atoms, (2) that the atomic weight of hydrogen is 1, (3) that in water 1 atom of hydrogen is united with 1 atom of oxygen, and (4) that in hydrochloric acid 1 atom of chlorine is united with 1 atom of hydrogen, we simply arrive at the conclusion that the combining weights are the relative atomic weights ; the lowest combining weight of an element found in any compound into which it enters being, of course, taken as the atomic weight. This system of using the equiva- lent weights as the atomic weights was employed by many chemists during the first half of the 19th century. Thus, in the atomic theory, the law of definite proportions finds a ready explanation, for each compound would always contain the same relative number of atoms of each of the ele- ments of which it is composed. The law of multiple propor- tions is readily explained by the theory, since according to it a fixed number of atoms of one element can only combine with one atom or some other whole number of atoms of another ele- ment. And finally, the law of reciprocal proportions also is easily accounted for by the theory, since according to it the weight of an atom of any one element is constant and different from that of any other element, and combination can only take place by whole numbers of atoms, from which it follows that the proportion by weight in which an element occurs in one compound will be either the same as, or some multiple of, the proportion in which it occurs in any other compound. It should thus be borne in mind that the modern atomic theory of matter is based upon the weight relations that obtain when the elements unite chemically, and these weight relations are expressed in the three fundamental stoichiometrical laws. It is of interest to note that, though Dalton promulgated the modern atomic theory in 1802, the basis for that theory has been furnished by chemists of three different nations ; for the TO OUTLINES OF CHEMISTRY law of definite proportions was discovered by Lavoisier and Proust, the law of multiple proportions by Dalton, and the law of reciprocal proportions by Richter. The atomic conception of matter was not original with Dalton ; indeed, its origin dates back to the times of classical Greece. Democritus, Epicurus, and Leucippus taught that matter is made up of indivisible particles or atoms, whereas according to the doctrine of Anaxagoras, matter is infinitely divisible. However, the atomic conception of matter of the Greeks was a mere metaphysical speculation, not founded upon experimental facts. It was Dalton who first used the concep- tion of the atomic nature of matter in explaining actual facts established by experiments, and to him consequently we rightly ascribe the origin of the modern atomic theory. Difference between Theory and Law. The student must always clearly bear in mind the distinction between a theory on the one hand, and facts and laws on the other hand. Facts are the results of actual observation and experiment. When a large number of similar facts have been found and these are ex- pressed in a general statement, the latter is a law. Thus we actually find that the composition of salt, water, lime, sal am- moniac, etc., is constant, no matter when or where prepared. We have here a series of facts. If now we formulate this into the general statement, that a chemical compound always has the same composition, we have a law. A theory or hypothesis, however, is neither a fact nor a general statement of fact, it is merely an assumption made for the purpose of correlating, explain" ing, or accounting for facts that have been collected and formulated into laws. So the atomic theory is a theory which enables us to correlate and comprehend better the facts expressed in the stoichiometrical laws. A theory, however, not only enables one to see facts in their relations and thus satisfies the natural craving of the human intellect for a better comprehension of things observed, but it also suggests new avenues of inquiry and experimentation by means of which further facts may be acquired. A theory is thus a powerful stimulus to scientific research, and is conse- quently of almost inestimable value. On the other hand, it is to be remarked that theories by implication also suggest that it is useless for actual inquiry to proceed in certain directions, and that certain things are impossible, when after all they are quite possible. FUNDAMENTAL LAWS AND THEORIES 71 and thus a theory may be a bar to progress. Carefully ascertained facts formulated into laws constitute the unchangeable, the eternal part of science. Theories and hypotheses on the other hand are the changeable, the ephemeral part of science ; for the views we entertain concerning the relationship of natural phe- nomena frequently change as new facts are acquired. A theory is a cord by which the precious pearls of truth are held together, but when the pearls found become too numerous or too heavy so that the old cord can no longer hold them together, it must be discarded, and the pearls must ultimately be arranged on a new cord of adequate length and strength. Thus theories and hypotheses are frequently discarded in science. In fact, the pathway of the progress of science is strewn with defunct theories. As we continue our considerations, we shall see that the atomic theory, simple and even crude as it seems, has been in a high degree useful in correlating even facts other than those upon which it is actually based, and has suggested many new avenues of further experimental inquiry. It has thus ful- filled in a high degree the function of a theory. The Law of Combination of Gases by Volume. It will be re- called that Gay-Lussac discovered the law that when gases com- bine to form chemical compounds the volumes of the gases that enter into combination bear a simple relation to one another ; and if the product formed be gaseous, its volume also bears a simple relation to the volumes of the original gases. This law was estab- lished at about the time when Dalton formulated the atomic theory. In viewing the fact that 1 volume of hydrogen and 1 volume of chlorine unite to form hydrochloric acid gas, in the light of the atomic theory of Dalton, according to which hydrogen and chlorine are made up of atoms and 1 atom of the one unites with 1 atom of the other to form one particle of hydrochloric acid, it follows at once that 1 volume of hydrogen must contain exactly as many atoms of hydrogen as the same volume of chlorine contains atoms of the latter ele- ment ; for were this not the case, there would be either hydro- gen or chlorine left uncombined when exactly equal volumes act on each other chemically. Thus the idea was naturally formed that equal volumes of gases under the same conditions of temperature and pressure contain the same number of atoms. Th-is is, of course, not a law, but simply an hypothesis evolved 72 OUTLINES OF CHEMISTRY to explain the law of Gay-Lussac of combination of gases by volume. To the hypothesis in the form stated, Berzelius inter- posed a serious objection. Thus he called attention to the fact that when 1 volume of hydrogen and 1 volume of chlorine unite, 2 volumes of hydrochloric acid are formed ; and if 1 atom of hydrogen unites with 1 atom of chlorine, there will of course be formed as many particles of hydrochloric acid as there are particles of hydrogen, or what amounts to the same thing, as there are particles of chlorine. Therefore, if equal volumes of hydrogen, chlorine, and hydrochloric acid contain the same number of atoms or particles, the hydrochloric acid formed by the union of equal volumes of hydrogen and chlorine ought to occupy the same volume as the original hydrogen ; that is, it ought to occupy one half of the volume that it actually does occupy. To hold the volume hypothesis, a scheme consequently had to be proposed whereby 1 atom of hydrogen would unite with 1 atom of chlorine and form 2 particles of hydrochloric acid ; for only then the actual volume relations that obtain when the latter substance is formed by the union of hydrogen and chlorine would be accounted for. Such a scheme was pro- posed in 1811 by Amadeo Avogadro, who was then professor of physics at the University of Turin. He made the bold assumption that the particles of hydrogen gas really are double atoms, that is, that they consist of 2 atoms joined together, and that the particles of chlorine gas are similarly made up each of 2 chlorine atoms. These double atoms of hydrogen and double atoms of chlorine he called molecules, and then stated the hypothesis as follows : Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules. This is known as AvogadrcTs hypothesis. It is very important in chemistry. Thus, Avogadro considered the molecule of hydrogen as H 2 , and the molecule of chlorine as C1 3 ; and when these react with each other we should have H 2 + C1 2 =2HC1. 1 molecule -f 1 molecule = 2 molecules. 1 volume + 1 volume = 2 volumes. On this basis, there would be twice as many molecules of hydrochloric acid formed as there were molecules of hydrogen or molecules of chlorine, and consequently one would expect FUNDAMENTAL LAWS AND THEORIES 73 the hydrochloric acid formed to occupy twice the volume of the original hydrogen ; or, what is the same, twice the volume of the chlorine. Let us now review the volume relations that obtain when hydrogen and oxygen combine to form water vapor. We have by experiment 2 volumes of hydrogen -f 1 volume of oxygen = 2 volumes of water vapor. If, now, we desire to hold Avogadro's hypothesis, we clearly must assume the molecules of hydrogen, oxygen, and water so constituted that : 2 molecules of hydrogen + 1 molecule of oxygen = 2 molecules of water vapor. But in connection with the synthesis of hydrochloric acid, it was already assumed that the molecule of hydrogen consists of 2 atoms (i.e. that it is H 2 ); we must consequently adhere consist- ently to this assumption wherever hydrogen gas is considered. Now if we assume that the oxygen molecule is made up of 2 atoms of oxygen (i.e. is O 2 ), the volume relations in the case of the synthesis of water are readily explained; the water mole- cule then, however, must be considered as composed of 2 atoms of hydrogen and 1 atom of oxygen. In form of an equation we should have 2 volumes of hydrogen + 1 volume of oxygen = 2 volumes of water vapor. 2 molecules of hydrogen -f 1 molecule of oxygen = 2 molecules of water vapor. i.e. 2H 2 + O 2 =2H 2 O. Thus it appears that if we hold Avogadro's hypothesis, and assume with him, that in hydrogen gas and oxygen gas there are molecules that consist of 2 atoms of these respective ele- ments, we are bound to conclude that the molecule of the very common compound water is not made up of 1 atom of hydro- gen united with 1 atom of oxygen, but rather of 2 atoms of hydrogen united with 1 atom of oxygen. We should thus have to assign to water the formula H 2 O instead of HO. Further, since by weight 1 part of hydrogen unites with 8 parts of oxy- gen, and since we assume with Dalton that the atomic weight of hydrogen is 1, we consequently must assume the atomic weight of oxygen as equal to 16 instead of S. U OUTLINES OF CHEMISTRY All this seemed to many chemists of the early part of the nineteenth century as a set of rather violent changes to make. In fact, Avogadro's molecules, his double atoms as they were frequently termed in the literature, were not regarded seriously by many able chemists for nearly half a century ; they simply continued to work with the tables of equivalent or combining weights in which the value for oxygen was 8. One of the reasons why Avogadro's hypothesis was laid aside for a time was that in the study of the ammonium salts certain apparent contradictions were met. These will be considered when those salts are discussed. Suffice it here to say that Avogadro's hypothesis has gained general acceptance and is now commonly regarded as of vital consequence in chemistry. Thus, while the atomic theory is based upon the weight relations that obtain when substances unite chemically, the molecular theory came into being as a consequence of the acceptance of Avogadro's hypothesis, which in turn grew out of Cray-Lussac's law of the combination of gases by volume. Avogadro's hypothesis is further supported by the fact that all gases contract and expand alike under the same changes of temperature and pressure. Molecular Weight Determinations. If equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules, it is clear that the weights of equal volumes of gases are to one another as the mo- lecular weights of the gases. To fully appreciate Avogadro's hypothesis, the student must bear in mind that on the basis of the molecular theory a gas consists of molecules that are en- tirely remote from one another. So, for example, if the mole- cules of a gas were, say by pressure, all crowded together so that they touched one another, they would occupy but a small portion of the volume originally occupied by the gas. In other words, the actual volume of a gas consists largely of space or interstices between the molecules, as it were. Bearing this in mind, it is clear that equal volumes of gases may well contain equal numbers of molecules, though the individual molecules of each of the gases may occupy very different volumes. Ac- cording to Avogadro's hypothesis, the volume of any gas, at constant temperature and pressure, depends not upon the size, weight, or kind of molecules it contains, but solely upon the number of molecules present. Hence at constant temperature and FUNDAMENTAL LAWS AND THEORIES 75 pressure, the volumes of any two gases are to each other as the number of molecules in the volumes. Thus, for instance, 10 liters of any gas contain 10 times as many molecules as 1 liter of the same gas or any other gas. Again, take any vol- ume of hydrogen, say 1 liter, and compare its weight with the weight of the same volume of some other gaseous substance, say chloroform vapor, at the same temperature and pressure. Now since by Avogadro's hypothesis each of these volumes contains the same number of molecules, which we shall call n, then we should have The wt. of 1 liter of Chloroform Vapor : wt. of 1 liter Hydrogen : : mn : 2 w, where ra'is the molecular weight of chloroform, and 2 that of hydrogen. From the equation, , 9 wt. of liter of Chloroform Vapor wt. of liter of Hydrogen That is, the molecular weight of chloroform is equal to twice its vapor density as compared with hydrogen. Hence the rule for finding molecular weight of any gas : find the density of the gas with respect to hydrogen, and multiply the result by 2. The volume occupied by 2 grams of hydrogen under standard conditions, and 760 mm. pressure, is 22.38 liters. The same volume is occupied by 32 grams of oxygen, 70.9 grams of chlorine, 18 grams of water vapor, and in short, by the molec- ular weight in grams of any gaseous substance whatever, under standard conditions. Consequently 22.38 liters is termed the molecular volume of all gases. We may also state that to find the molecular weight of any gas, we simply need to determine the weight of 22.38 liters of that gas at and 760 mm. pressure, and the result is the molecular weight in grams. The student should assure himself that this really comes to the same thing as finding the density of the gas with respect to hydrogen and multiplying the result by 2. The molecular weights of substances that can be obtained in the vapor state can thus readily be determined. But there are liquid and solid substances that cannot be vaporized without decomposition, and so their vapor densities cannot be deter- mined. In the case of substances that can be dissolved, it is possible to make molecular weight determinations by studying the freezing point, boiling point, or vapor pressure of the solu- 76 OUTLINES OF CHEMISTRY tion. This will be explained later in connection with the subject of solutions. Determination of Atomic Weights. The proportions by weight in which the elements combine with one another are determined by very exact chemical analyses of the compounds containing the elements in question ; or, if the latter will unite directly, by ascertaining the weights of the elements that enter into combination, when compounds are thus formed by synthesis. The values so found are the combining weights. Dalton took the combining weight of hydrogen as equal to unity, and expressed the combining weights of the other elements on this basis. We now, for reasons already stated above, take the com- bining weight or chemically equivalent weight of oxygen as equal to 8, oti which basis that of hydrogen equals 1.008. We have seen that Gay-Lussac's law of combination of gases by volume led to Avogadro's hypothesis, which in turn led to the idea of molecules, and to the conception that a molecule of oxygen consists of 2 atoms. This further made it necessary to adopt for water the formula H 2 O, instead of HO, and con- sequently for oxygen the atomic weight 16 instead of 8. Thus the experimental fact that 2 volumes of hydrogen unite with 1 volume of oxygen to form 2 volumes of water vapor really determined us in choosing the atomic weight of oxygen as 16 instead of 8 ; in other words, the vapor density of water has really been used in deciding whether we should use 8 or some multiple of that figure as the atomic weight of oxygen. Similarly, the vapor density of substances has in many other cases been used in choosing the atomic weights, the combining weights being known. Thus, in hydrochloric acid hydrogen and chlorine are combined in the proportions of 1 to 35.45 by weight. By volume, on the other hand, we have, 1 volume of hydrogen uniting with 1 volume of chlorine to form 2 volumes of hydrochloric acid. We have seen that this volume relation led Avogadro to distinguish between the atom of hydrogen, H, and the molecule of hydrogen, H 2 , and also between the atom of chlorine, Cl, and the molecule of chlorine, C1 2 . We have also noted, that having assumed each of the molecules of hydrogen and chlorine to consist of 2 atoms, the composition of the molecule of hydrochloric acid could still be expressed by the simple formula, HC1 ; and thus the volume FUNDAMENTAL LAWS AND THEORIES 77 relations that obtain when hydrochloric acid is formed from the elements could be explained in the light of Avogadro's hypothe- sis, and 35.45 be regarded as the atomic weight of chlorine Here again, the vapor density of hydrochloric acid gas has been the determining factor in choosing the atomic weight of chlorine as 35.45 rather than some multiple thereof. In marsh gas, which consists of hydrogen and carbon, 1 gram of hydrogen is combined with every 3 grams of carbon. The combining weight or chemical equivalent of carbon is therefore 3. How now proceed to ascertain the atomic weight of carbon ? Marsh gas is 8 times heavier than hydrogen ; whence 22.38 liters of marsh gas weigh 16 grams. From what has been stated above, 16 is therefore the molecular weight of marsh gas. But in 16 grams of marsh gas there are 4 grams of hydrogen, which is 4 times the atomic weight of hydrogen in grams. The molecule of marsh gas therefore contains 4 atoms of hydrogen. Now if in the case of this compound, which of all the compounds of hydrogen with carbon contains the least amount of carbon by weight as compared with the hydrogen, we assume that there is but 1 carbon atom to the 4 hydrogen atoms, we must ascribe to the carbon atom the weight of 12. Thus it is clear that the vapor density of marsh gas has not only fixed its molecular weight, but has also led us to choose the atomic weight of carbon as 12, rather than as some other multiple of 3, the combining weight. Further, we find that by thus taking 12 as the atomic weight of carbon, the com- position of other compounds into which that element enters 'can readily be expressed. Of compounds of carbon with oxygen, the one that contains the least carbon as compared with the oxygen is carbonic acid gas. In this 3 grams of car- bon are combined with every 8 grams of oxygen. The gas is 22 times heavier than hydrogen, that is 22.38 liters of carbonic acid gas weigh 44 grams, and the molecular weight of this sub- stance is consequently 44. Since it contains the least carbon of any of the known compounds of oxygen and carbon, it would be natural to hold that it contains but 1 atom of carbon. Now if the atomic weight of carbon be 12, that of oxygen 16, and the molecular weight of carbonic acid gas 44, we have (since in 44 grams of carbonic acid gas there are 12 grams of carbon and 32 grams of oxygen) in the carbonic acid molecule 78 OUTLINES OF CHEMISTRY 1 atom of carbon united with 2 atoms of oxygen, which is ex- pressed by the formula CO 2 . Since this compound contains but two elements, it is a so-called binary compound. The names of all binary compounds, i.e.. compounds consisting of but two elements, end in ide. Since carbonic acid gas contains twice as much oxygen as the lower oxide of carbon, carbon monoxide, and as we have assigned to the former the formula, CO 2 , indicating that the molecule contains 2 atoms of oxygen, it is called carbon dioxide. In carbon monoxide there are 3 grams of carbon combined with every 4 grams of oxygen, and carbon monoxide is 14 times heavier than hydrogen ; its mo- lecular weight is consequently 28. The atomic weights 12 for carbon and 16 for oxygen would consequently lead us to write the formula for carbon monoxide as CO. The name carbon monoxide is given to the compound because it contains but 1 atom of oxygen in its molecule. When carbon is burned in oxygen, the carbon dioxide formed is of exactly the same volume as the original oxygen, as will appear later. In other words : 1 volume of oxygen yields 1 volume of carbon dioxide. Accepting Avogadro's hypothe- sis, there are then as many molecules of carbon dioxide formed as there are molecules of oxygen. The process may be ex- pressed by the equation, C + 2 = C0 2 , which, like all so-called chemical equations, simply expresses the march of the reaction. The sign = stands for yields. Some use the sign > instead of = . The latter is, however,, more frequently employed. When carbon monoxide is burned in oxygen, 2 volumes of the former unite with 1 volume of the latter to form 2 vol- umes of carbon dioxide. Assuming Avogadro's hypothesis, there must consequently be formed as many molecules of carbon dioxide as there were molecules of carbon monoxide. These relations are expressed by the simple equation : 2CO + 2 =2C0 2 . The above illustrations may suffice to indicate how the vapor densities of substances have been employed in choosing the atomic weights, the combining weights having been ascertained by careful quantitative analytical or synthetical experiments. FUNDAMENTAL LAWS AND THEORIES 79 When Avogadro put forth his hypothesis in 1811, it was by no means at once generally accepted. Indeed, it was not till the vapor densities of a very considerable number of substances had become known, that the value of the hypothesis was really recognized. It was Charles Gerhardt, professor of chemistry at the University of Montpellier, who in 1842 used the vapor densities of substances as a guide in determinining their for- mulae and in choosing the atomic weights from the equivalents or combining weights, which were at that time in almost general use. But it was Auguste Laurent, professor of chem- istry at the University of Bordeaux, who in 1846 grasped the great value of Avogadro's hypothesis and paved the way for its general acceptance. He distinguished clearly between atomic and molecular weights, defining the molecule as the smallest weight of any substance that can exist by itself, and the atom as the smallest weight of a substance that can enter into combination. But there are elements which do not enter into compounds that can be vaporized, and consequently the atomic weights of such elements cannot be chosen from the combining weights by means of the vapor density, as described. This is particu- larly true of many of the metals. In determining the atomic weights of the latter, Berzelius simply took the number of parts by weight of the metal that united with 16 parts by weight of oxygen as the atomic weight of the metal. In case a metal formed more than one oxide recall the oxides of lead, for instance Berzelius assumed the one most commonly found as containing 1 atom of the metal to 1 atom of oxygen, and then computed the formulae of the other oxides accord- ingly. When there was but one oxide known, as in the case of zinc, for instance, he assumed that the molecule consisted of 1 atom of the metal to 1 atom of oxygen. Thus he proceeded on the basis of simplicity, guarding himself by assigning similar formulas to substances that exhibit similar chemical properties. Gerhardt, however, considered it likely that the molecules of the oxides of the metals are similar to the molecule of water in construction, and consequently contain 2 atoms of metal to 1 atom of oxygen, which, of course, led him to adopt atomic weights for the metals which were just half of the values adopted by Berzelius. This led to considerable discus- sion. But in 1858 Stanislao Cannizzaro, then professor of 80 OUTLINES OF CHEMISTRY chemistry at Genoa, brought order into the confusion that had arisen by pointing out that the specific heats of the elements in the solid state may be employed with great advantage in choosing the true atomic weights from the combining weights. The Law of Dulong and Petit. Cannizzaro recalled a simple relation, discovered by Dulong and Petit of Paris in 1819, be- tween the atomic weight of an element and its specific heat. This relation is that the product of the specific heat of an element in the solid state and its atomic weight is constant. This law, which is known as the law of Dulong and Petit, may also be expressed by saying that the atoms of the elements have the same heat ca- pacity. The experimental researches of Victor Regnault, the great French physicist (1810-1878), added many new data to confirm this law ; but, like the hypothesis of Avogadro, its value was not clearly recognized till Cannizzaro showed how useful it is in fixing the atomic weights of many of the ele- ments. The product of the atomic weight and the specific heat of an element in the solid state is approximately 6.4. The follow- ing table gives the specific heats of a number of elements in the solid state : ELEMENT SYMBOL ATOMIC WEIGHT SPECIFIC HEAT ATOMIC HEAT Lithium Li 6.94 0.941 6.53 Sodium Na 23.00 0.293 6.74 Magnesium Mg 24.32 0.245 5.95 Aluminum Al 27.1 0.214 5.80 Phosphorus P 31.0 0.202 6.26 Sulphur S 32.07 0.203 6.50 Potassium K 39.10 0.166 6.49 Iron Fe 55.85 0.112 . 6.26 Copper Cu 63.57 0.095 6.04 Zinc Zn 65.37 0.093 6.08 Silver Ag 107.88 0.057 6.15 . Platinum Pt 195.2 0.0325 6.34 Gold Au 197.2 0.0324 6.40 Mercury Hg 200.6 0.0333 6.68 Lead Pb 207.1 0.0315 6.52 Glucinum Gl 9.1 0.42 3.82 Boron B 11.0 0.24 2.64 Carbon (Graphite) C 12.0 0.200 2.40 Silicon Si 28.3 0.177 5.01 FUNDAMENTAL LAWS AND THEORIES 81 As the specific heat of a substance is a function of the tem- perature at which it is determined, one would clearly not ex- pect the product of the atomic weight and the specific heat to yield exactly the same value. An inspection of the table shows that the atomic heat is generally about 6, except in the case of the last four elements, where it varies greatly from that value. Now it is found that the specific heats of glucinum, boron, car- bon, and silicon increase greatly with rise of temperature, finally becoming nearly constant. So the specific heat of graphite is 0.160 at 10 ; 0.199 at 60 ; 0.445 at 600 ; and 0.460 at 900. The specific heat of silicon is 0.20 at 200 and 0.203 at 300; that of glucinum is 0.617 at 400, and 0.620 at 500. Thus at higher temperatures these elements approximately obey the law of Dulong and Petit, though at room temperatures they appar- ently are exceptions. By means of the law of Dulong and Petit the atomic weight of an element may be found by dividing the atomic heat, ap- proximately 6.4, by the specific heat; that is, / 4 Atomic Weight Specific Heat This method can obviously not be used for determining atomic weights with accuracy ; but it is of great value in choosing the true atomic weight from the combining weights, and it is in this way that it was employed with great success by Canniz- zaro. The latter thus showed that in nearly all cases the values that Berzelius had assigned to the atomic weights of the metals were the correct ones ; and that only in a few instances, like potassium, sodium, and silver, did the oxides have two atoms of the metal in the molecule, though Gerhardt had assumed this in all cases. Cannizzaro also showed that wherever volatile metallic compounds were known, the choice of the atomic weights from the vapor density of these agreed with the values deduced from the specific heats. Other Methods of choosing the Atomic Weights from the Com- bining Weights. Many substances form crystals. Crystals are solids bounded by plane faces which are the outcome of a regular internal structure. Substances which do not crystal- lize, that is, are non-crystalline, are called amorphous. All crystals may be classified into six crystal systems (which see). 82 OUTLINES OF CHEMISTRY In 1819, Eilhard Mitscherlich discovered that chemical com- pounds which are similar in character crystallize in the same, forms. This is the law of isomorphism, for isomorphous sub- stances are such as crystallize in the same form. Whenever compounds are isomorphous, they are chemically analogous ; and so if the formula of one compound has been determined, the formulae of other compounds that are isomorphous with it are readily deduced by analogy. Ismorphism may conse- quently be used in choosing the atomic weights from the com- bining weights. It was so employed by Berzelius. Thus, the sulphate of magnesium is isomorphous with that of zinc. If the atomic weight of the latter metal has been fixed as 65.5, with the aid of the law of Dulong and Petit, then the amount of magnesium that is required to replace 65.5 parts of zinc by weight in the sulphate, namely 24.32, is the atomic weight of magnesium. In this way isomorphism has been of great use in atomic weight determinations. It must be applied with great care, however, for it is true there are many cases where com- pounds are chemically dissimilar and yet possess like crystal forms. For this reason, isomorphism is not so reliable a guide as the specific heat or vapor density method, and is only em- ployed when other methods cannot be used. In choosing the atomic weights from the combining weights, the principle of simplicity and analogy was employed with much success by Berzelius. As already stated, in the case of metals, like zinc and magnesium, that form but one compound with oxygen, he assumed that the oxide contains 1 atom of the metal to 1 atom of oxygen. Since the atomic weight of oxygen was taken as 16, the atomic weight of the element combined with oxygen could readily be found. Further, Ber- zelius sought to assign analogous formulae to compounds that are actually analogous in their chemical behavior. By this method he consequently chose the atomic weight of the ele- ment in question in accordance with the formulas assigned. Finally, the arrangement of the atomic weights of the ele- ments in the so-called periodic system (which see) has in some cases influenced the choice of the atomic weights. In summary, then, the methods of choosing the atomic weights from the combin- ing weights are: (1) the vapor density, (2) the specific heats in the solid state, (3) isomorphism, (4) the principle of simplicity FUNDAMENTAL LAWS AND THEORIES 83 and chemical analogy, and (5) the periodic system. Concerning the last two of these it should be said at this juncture that the manner of their application cannot be elucidated before more substances have been studied. Table of Atomic Weights. The following is a table of the atomic weights of the elements as adopted by the International Commission on Atomic Weights : INTERNATIONAL ATOMIC WEIGHTS ELEMENT SYMBOL ATOMIC WEIGHT ELEMENT SYMBOL A'lOMIC WEIGHT Aluminum Al 27.1 Neodymium Nd 144.3 Antimony Sb 120.2 Neon Ne 20.2 Argon A 39.88 N ickel Ni 58.68 Arsenic As 74.96 Niton (radium Barium Ba 137.37 emanation) Nt 222.4 Bismuth Bi 208.0 Nitrogen N 14.01 Boron B 11.0 Osmium Os 190.9 Bromine Br 79.92 Oxygen O 16.00 Cadmium Cd 112.40 Palladium Pd 106.7 Caesium Cs 132.81 Phosphorus P 31.04 Calcium Ca 40.07 Platinum Pt 195.2 Carbon C 12.00 Potassium K 39.10 Cerium Ce 140.25 Praseodymium Pr 140.6 Chlorine Cl 35.46 Radium Ra 226.4 Chromium Cr 52.0 Rhodium Rh 102.9 Cobalt Co 58.97 Rubidium Rb 85.45 Columbium Cb 93.5 Ruthenium Ru 101.7 Copper Cu 63.57 Samarium Sa 150.4 Dysprosium Dy 162.5 Scandium Sc 44.1 Erbium Er 167.7 Selenium Se 79.2 Europium Eu 152.0 Silicon Si 28.3 Fluorine F 19.0 Silver Ag 107.88 Gadolinium Gd 157.3 Sodiurn N! 23.00 Gallium Ga 69.9 Strontium Sr 87.63 Germanium Ge 72.5 Sulphur S 32.07 Glucinum Gl 9.1 Tantalum Ta 181.5 Gold. Au 197.2 Tellurium Te 127.5 Helium He 3.99 Terbium Tb 159.2 Holmium Ho 163.5 Thallium Tl 204.0 Hydrogen H 1.008 Thorium Th 232.4 Indium In 114.8 Thulium Tm 168.5 Iodine I 126.92 Tin Sn 119.0 Iridium Ir 193.1 Titanium Ti 48.1 Iron Fe 55.84 Tungsten W 184.0 Krypton Kr 82.92 Uranium U 238.5 Lanthanum La 139.0 Vanadium V 51.0 Lead Pb 207.10 Xenon Xe 130.2 Lithium Li 6.94 Ytterbium Lutecium Lu 174.0 (Neoytterbium) Yb 172.0 Magnesium Mg 24.32 Yttrium Yt 89.0 Manganese Mn 54.93 Zinc Zn 65.37 Mercury Hg 200.6 Zirconium Zr 90.6 Molybdenum Mo 96.0 84 OUTLINES OF CHEMISTRY Interpretation of a Chemical Formula. A chemical formula expresses (1) what elements occur iu the compound, (2) the relative weights in which these elements occur in the compound, and (3) the weight of 22.38 liters of the vapor of the com- pound under standard conditions. In case the compound can- not be converted into the vapor state, the formula is derived from a study of the freezing or boiling point of its solution, the crystalline form, the specific heat, or from its chemical behavior and analogy to other compounds. These facts are thus recorded in the formula. So the formula of carbonic acid gas CO 2 tells us that this compound consists of carbon and oxygen in the proportions of 12 parts of the former to 32 parts of the latter by weight. It also tells that the weight of 22.38 liters of the gas under standard conditions is 44 grams, or that the gas is 22 times as heavy as hydrogen. Thus we see that chemical formulae are a system of shorthand writing, as it were, for they express in a small space the salient facts known about a compound. Valence and Structural Formulae. For hydrochloric acid we have developed the formula HC1. In this compound 1 atom of hydrogen is combined with 1 of chlorine. In water H 2 O, on the other hand, 2 atoms of hydrogen are combined with 1 of oxygen. The power which an atom of one element has to unite with one or more atoms of other elements is called its valence. Thus in hydrochloric acid, hydrogen and chlorine each have a valence of one. Hydrochloric acid is said to be a saturated compound, for it will unite with neither more hydro- gen nor more chlorine. Hydrogen always has a valence of one, it is consequently called a univalent element or a monad. The number of hydrogen atoms, or other univalent atoms, with which an atom of a given element combines determines the valence of the latter. In water, we have 2 hydrogen atoms united with 1 oxygen atom. Oxygen, consequently, has a valence of 2; i.e. it is a bivalent element or dyad. The formula for water is also written H O H to indicate that each of the atoms of hydrogen is bound to oxygen, which idea may be deduced from the fact that when sodium acts on water only half of the hydrogen of the latter is displaced, and that the other half of the hy- drogen is set free when the resulting sodium hydroxide FUNDAMENTAL LAWS AND THEORIES 85 is treated with zinc, as indicated by the equations (compare Chapter II)- H-O-H + Na = NaOH + H. Water + sodium = sodium hydroxide -f- hydrogen. 2 NaOH + Zn = Na 2 O 2 Zn + H 2 . Sodium hydroxide + zinc = sodium zincate -f hydrogen. The formula H O H is therefore the structural formula for water. It is clear that it is derived from reactions that water will undergo, and.is consequently merely a brief way of express- ing these changes. Structural formulae are often used in chem- istry, particularly in connection with the compounds of carbon. It is not to be thought that such a formula expresses the actual conditions that exist within the molecule itself; it is rather simply a concise expression of the reactions which the compound in question will undergo with other chemical compounds. In chlorine monoxide C1 2 O chlorine is univalent and oxy- gen is bivalent. The structural formula of the compound is Cl O 01 ; we may consider it as water in which the hydro- gen atoms are replaced by chlorine. In chlorine dioxide C1O 2 the oxygen is bivalent, and the chlorine has a valence of four; i.e. it is quadrivalent, the formula being O = 01 = O. Oxygen is always bivalent except in very rare cases. It is thus clear that the number of oxygen atoms, or other atoms of known valence, with which an atom of another element combines, may also serve to ascertain the valence of the latter. In carbon dioxide CO 2 carbon is quadrivalent ; thus, O = = O In carbon monoxide CO carbon is bivalent; thus, C = O ; or sometimes it is considered as quadrivalent, two combining powers or bonds being free or unsaturated, thus, = C = O. This at once brings us to the question whether the valence of an element is always the same or not. There has been considerable dispute over this question, but now it is quite generally held that the valence of an ele- ment may vary in different compounds. The highest valence which an element exhibits in any known compound is called its maximum valence. The valence of an element may vary from one to eight, though in case of most elements it varies but slightly. As already stated, hydrogen is always univalent ; oxygen is almost CALIF OR NIA COLLfcfii 86 OUTLINES OF CHEMISTRY always bivalent; and carbon may practically always be con sidered as quadrivalent, though in some compounds it is biva lent and even trivalent. Chlorine is always univalent toward hydrogen, while toward oxygen it may be either univalent, bivalent, quadrivalent, or heptavalent. The valences of the various elements will be taken up in connection with the description of each element, for the subject cannot be con- sidered fully except in connection with actual illustrations. Nomenclature. The names of the metallic elements end in um, lik'e lithi?/m, sodium, baritm, etc., except, in case of some metals that have been known for a very long time, which retain their old names, as iron, lead, gold, silver, etc. The elements selenium and tellurium are not metals. They were thought to be such when discovered, on account of their outward properties, hence the ending um in their names. Substances containing but two elements are called binary compounds ; their names end in ide. Thus, common salt NaCl is sodium chloride / magnesia is magnesium oxide MgO ; lime is calcium oxide CaO, etc. In some cases the suffix ide is used in connection with compounds containing more than two ele- ments. In all these, however, two or more of the elements act as a group, i.e. a unit or radical, so called, which may pass from compound to compound ; thus, sodium hydroxide NaOH contains the OH group, which is called the hydroxyl group. This OH group passes from one compound to another as a unit, and when elements or other groups are combined with this group, the compounds formed are termed hydroxides. So we have calcium hydroxide Ca(OH) 2 formed when calcium acts on water, Ca + 2 H 2 O = Ca(OH) 2 + H 2 ; or when lime, calcium oxide CaO, is treated with water, i.e. is slaked, thus : CaO + H 2 = Ca(OH) 2 . In the hydroxyl group OH one of the two combining powers or bonds of oxygen is satisfied by hydrogen, the other bond being free. The group is therefore univalent, and we may write it thus : O H. When two elements form more than one compound with each other (which is frequently the case), a name indicating the nura- FUNDAMENTAL LAWS AND THEORIES 87 Der of atoms of the one element that are united with the other is given to each compound. Thus we have carbon monoxide CO, carbon dioxide CO 2 , sulphur dioxide SO 2 , sulphur trioxide SO 3 , phosphorus trichloride PC1 3 , phosphorus pentachloride PC1 6 , lead sesquioxide Pb 2 O 3 . The ending ous is frequently used for one compound and the ending ic for another compound richer than the former in one of the ingredients. Thus SO 2 or sulphur dioxide is also called sulphurous oxide, and SO 3 or sul- phur trioxide is also called sulphuric oxide. Again, PC1 3 is phosphorus trichloride or phosphorous chloride, and PCl g is phosphorus pentachloride or phosphor^*? chloride. When more than two compounds are formed by two elements, the endings ous and ic are retained and the prefixes proto, Jiypo or sub, and per are added as required. Thus litharge PbO is lead mon- oxide, lead protoxide or plumbic oxide ; black oxide of lead Pb 2 O is lead swooxide or plumbous oxide ; Pb 2 O 3 is lead sesqui- oxide ; minium or red lead Pb 3 O 4 is the proto-sesquioxide of lead, i.e. PbO Pb 2 O 3 ; brown oxide of lead PbO 2 is lead dioxide or lead peroxide. The prefix per stands for the highest oxida- tion stage in the case of oxides, for the highest chlorination stage in the case of chlorides, etc. Chlorine monoxide or pro- toxide C1 2 O is also called %/>ochlorous oxide. Water H 2 O is hydrogen protoxide or monoxide, or hydrogen hydroxide, or hydroxyl hydride. The prefix "hypo is rarely used in case of binary compounds. In ternary compounds, that is, those that are made up of three elements, somewhat similar designations are employed, which will be explained when compounds of this character are con- sidered. Chemical Equations. Retrospect. Chemical compounds may be designated by means of symbols or formulae, as we have seen, and chemical changes may be indicated by writing equa- tions in which these formulae are used instead of the names of the compounds. Reviewing the work on hydrogen, oxygen, and chlorine, and writing the principal chemical changes that have been studied in form of chemical equations, we have as follows : (1) Preparation of Hydrogen Na + H 2 O = NaOH + H. Sodium -f- water = sodium hydroxide + hydrogen 88 OUTLINES OF CHEMISTRY H 2 S0 4 + Zn = ZnS0 4 + H 2 . Sulphuric acid + zinc = zinc sulphate + hydrogen, 2KOH + Zn = K 2 O 2 Zn + H 2 . Potassium hydroxide -f zinc = potassium zincate 4- hydrogen. 3KOH + Al = K 8 8 A1 + 3 H. Potassium hydroxide + aluminum = potassium aluminate+ hydrogen. 2 H 2 O (on electrolysis) = 2 H 2 + O 2 . 3Fe + 4H 2 O = Fe 3 O 4 + 4 H 2 . Mg +H 2 MgO + H 2 . (2) Preparation of Oxygen HgO (on heating) = Hg + O. Mercuric oxide (on heating) = mercury + oxygen. Ag 2 O (on heating) = 2 Ag + O. Argentic oxide (on heating) = silver + oxygen. KC1O 3 (on heating) = KC1 + 3 O. Potassium chlorate (on heating) = potassium chloride + oxygen. 3 MnO 2 (on ignition) = Mn 8 O 4 + O 2 . Manganese dioxide (on ignition) = manganese proto-sesquioxide + oxygen. KNO 3 (on heating) = KNO 2 + O. Potassium nitrate or saltpeter (on heating) = potassium nitrite + oxygen. (3) Oxidations 2H a + O 2 =2H 2 O. Mg+ O = MgO. Cu+ O = CuO. phosphorus pentoxide. C + O 2 =CO 2 . 3 Fe + 2 O 2 = Fe 3 O 4 . 2 Fe + 3 O = Fe a O 8 , ferric oxide or sesquioxide of iron S + O g = SO 2 . FUNDAMENTAL LAWS AND THEORIES 89 (4) Reductions CuO + H 2 = Cu + H 2 O. Fe 3 O 4 + 4 H a = S Fe + 4 H 2 O. (5) Preparation of Chlorine 4HCl-fMnO 2 = MnCl 2 manganous chloride. (6) Reactions of Chlorine H 2 + C1 2 =2HC1 + 0. P + 3C1 = PC1 3 , phosphorus trichloride. P + 5 Cl = PC1 6 , phosphorus pentachloride. Sb + 3 Cl = SbCl 3 , antimony trichloride. G io H i6 + 8 C1 2 = 16 HC1 + 10 C. turpentine. Phenomena of the Nascent State. When a dilute solution of sulphuric acid is acting on zinc, the hydrogen liberated will reduce many substances like potassium permanganate, potas- sium bichromate, or saltpeter, if these are added directly to the mixture in the generator. The reduction will not take place if the hydrogen is passed through solutions of these salts contained in a separate vessel. The explanation of this as commonly given is that at the moment of liberation, the hydrogen is in the so-called nascent state, i.e. in an atomic con- dition represented by H, whereas afterwards it passes over into the molecular condition H 2 . While in the nascent state the hydrogen is more active than in the molecular state, and it con- sequently effects many reductions. Similarly we may have nascent oxygen O, as compared with molecular oxygen O 2 . Cases of this kind will be mentioned later. 90 OUTLINES OF CHEMISTRY REVIEW QUESTIONS 1. What is the difference between a law and a theory, as these terms are used in science ? 2. State the three fundamental laws of stoichiometry, illustrating each by means of a concrete example. 3. Upon what facts is the atomic theory of matter based? 4. What is meant by the term " combining weights"? Illustrate by means of an example. 5. What is the hypothesis of Avogadro ? Upon what facts is it based ? 6. Why is the symbol for hydrogen H 2 , and for chlorine C1 2 ? 7. Define atomic weight, also molecular weight. 8. Why is the formula for water written H 2 0? State fully all the facts which this formula represents. What three important facts does every chemical formula represent ? 9. What is meant by the term "specific heat"? State the law of Dulong and Petit. For what purpose is it used in chemistry? 10. A new element has a specific heat of 0.0328; what is its atomic weight ? 11. What is the significance of the volume 22.38 liters? The weight of a liter of a certain gaseous substance is 2.1 grams, under standard con- ditions ; what is its molecular weight ? 12. State the law of isomorphism. Who discovered it? 13. What are the different methods used in choosing the atomic weights from the combining weights? 14. What is the valence of the first element in the compounds repre- sented by the following formulas: C0 2 ; P 2 5 ; NH 3 ; Fe 2 3 ; S0 3 ; NaCl; BaCl 2 ; CuO; BiCl 3 ; S0 2 ; PC1 5 ; NO; Pb0 2 ; C1 2 7 . 15. How much oxygen may be prepared from 100 kilograms of silver oxide? What volume would this gas occupy at standard conditions? How much greater would its volume be at 10 C and 760 mm. pressure ? 16. What chemical elements generally have names ending in um ? Give six illustrations. Mention four cases that do not follow the rule. 17. Using the formula for water, explain what is meant by valence and structural formula. 18. Write the chemical equation expressing the action of metallic potassium upon water and explain in detail what the equation means. 19. Name two compounds illustrating the use of the endings ous and ic in chemical nomenclature. 20. Write the equation expressing the reaction by means of which Joseph Priestley first prepared oxygen. Why can not oxygen be pre- pared in an analogous manner by heating cupric oxide ? 21. Distinguish between nascent hydrogen and molecular hydrogen. Use examples. 22. Write the equations expressing the change when each of the fol- lowing is burned in oxygen : Mg, C, S, Cu, P. FUNDAMENTAL LAWS AND THEORIES 91 23. By means of a chemical equation express the electrolytic decom- position of water. 24. Write the equation expressing the action of metallic calcium on water. 25. Write the names of the compounds represented by the following formulas: Mn 3 4 , FeCl 2 , FeCl 3 , BaO, Ba0 2 , Pb 2 3 , PbO, NaOH. CHAPTER VII OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE History, Occurrence, and Preparation of Ozone. When a f fic- tional electrical machine is operated, there is observed in its neighborhood a peculiar characteristic odor, which is sometimes described as similar to the odor of chlorine, burnt sulphur, or garlic. The observation that this smell is produced when electric sparks are passed through oxygen was made in 1785 by. Van Marum, who had constructed an especially powerful machine. The same odor is noticed whenever electric sparks pass through the air, as, for instance, from an induction coil, or when objects are struck by lightning. In 1840 Christian Schonbein, professor at the University of -Basel, showed that when water is electrolyzed the oxygen obtained always con- tains some of this odoriferous substance, which he named ozone, meaning a smell. From the fact that ozone is produced when electric sparks pass through pure, dry oxygen, it is clear that the substance consists of oxygen. By means of the silent electrical discharge ozone is produced in larger quantities. For this purpose an apparatus like that in Fig. 28 is commonly employed. The apparatus is blown of one piece of glass. The OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE 93 outside of the tube A is coated with tin foil, as is also the inside of the tube B, as indicated. Dry oxygen is passed through the apparatus as shown, and when the tin foil coatings are connected with the poles of an induction coil, ozone issues at 0. In this way about 5 to 8 per cent of the oxygen is con verted into ozone. By liquefying oxygen and ozone by means of liquid air, a liquid is obtained which upon slow evaporation leaves a very dark blue liquid consisting of about 86 per cent ozone and 14 per cent oxygen. Besides being formed by means of electrical discharges and in the electrolysis of water, ozone is produced in chemical reactions, notably when moist phosphorus slowly oxidizes in the air ; also generally when oxygen is rapidly evolved, as by heating potassium chlorate, or when potassium permanganate is treated with strong sulphuric acid. Further, ozone is formed in very small quantities when hydrogen burns in oxygen. By the action of fluorine on water, oxygen containing up to 15 per cent of ozone is formed. Relation between Ozone and Oxygen, Allotropy. As already stated, ozone is produced from oxygen. By passing ozone through a red-hot tube, it is again converted into oxygen. Under standard conditions, 22.38 liters of ozone weigh 48 grams. The molecular weight of ozone is consequently 48; and since the atomic weight of oxygen is 16, the formula of ozone is O 3 . The change of oxygen to ozone is expressed by the following equation : 3 O 2 (plus energy) = 2 O 3 . The energy that must be added to oxygen to convert it into ozone may be obtained from the silent electric discharge, or from chemical changes, as we have seen. When ozone is heated, the reaction is reversed. We have here then a reversi- ble reaction. This fact may be expressed thus : 3 O 2 (plus energy) ^ 2 O 3 , where the arrows are used instead of the usual sign of equality. In forming ozone, oxygen shrinks from 3 volumes to 2 vol- umes and simultaneously a considerable amount of energy is absorbed. Ozone is called an allotropic form of oxygen. The property which some elements possess of occurring in two or more forms is called allotropy. 94 OUTLINES OF CHEMISTRY Ozone is a much more powerful oxidizing agent than oxygen. Many of the reactions which take place in oxygen only at higher temperatures proceed readily in ozone at room temperatures. Properties of Ozone. In thick layers ozone gas has a bluish color. Inhaled in quantity, it attacks the mucous membranes and produces headache. Liquid ozone is indigo-blue in color, and boils at 119 under atmospheric pressure. The liquid is strongly magnetic. On warming, it is liable to explode, due to sudden change of the substance to ordinary oxygen. Ac- cording to Ladenburg, 1000 volumes of water dissolve 10 vol- umes of ozone. It acts slowly on water, forming oxygen and hydrogen peroxide (which see), and the solubility in water may be due to this fact. The chief chemical property of ozone is its oxidizing power. It will bleach litmus, indigo, and other dyestuffs, the colors being destroyed by oxidation. Ozone destroys disease germs and other minute organisms, and, consequently, it is used as a germicide in sterilizing drinking water. Ozone is soluble in turpentine, also in oil of cinnamon and other similar oils. In solutions, ozone is also a powerful oxidizing agent. On account of its oxidizing power, it causes many oils to thicken and become resinous. Ozone rapidly oxidizes such substances as silver, lead, arsenic, phosphorus, and sulphur, to their highest stages of oxidation. It is the most powerful oxidizing agent known. It acts on potassium iodide solutions, liberating iodine, thus : 2 KI + H 2 O + O 3 = 2 KOH + O 2 + I 2 . Iodine turns starch paste blue, and so when a strip of paper saturated with starch paste plus a solution of potassium iodide is exposed to ozone, the paper turns deep blue in color. This is a common test for ozone. However, it must be used with proper care ; for, as we shall see, there are other things besides ozone that turn starch potassium iodide paper blue. The above reaction may be used in estimating the amount of ozone in a given sample of oxygen, by determining the quantity of iodine set free. Ozone is the only gaseous oxidizing agent that will blacken a bright silver foil, and consequently this test is used in detecting ozone in presence of other oxidizing gases. OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE 95 History, Occurrence, and Preparation of Hydrogen Peroxide. - In 1818 Thenard prepared a compound of hydrogen and oxy- gen containing twice as much oxygen as there is in water. He treated barium dioxide with hydrochloric acid, thus : BaO 2 + 2 HC1 = BaCl 2 + H 2 O 2 . Both the barium chloride and hydrogen peroxide (which is also called hydrogen dioxide or hydroperoxide) remain in solu- tion. Hydrogen peroxide may also be prepared by adding barium dioxide to cold, dilute sulphuric acid : Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 O 2 ; or by passing carbon dioxide through water and gradually adding barium dioxide in small amounts : BaO 2 + CO 2 + H 2 O = BaCO 3 + H 2 O 2 . Barium sulphate and barium carbonate are insoluble in water, and hence may be removed by filtration ; and thus a filtrate, which is an aqueous solution of hydrogen peroxide, may be obtained. When ozone acts on water, hydrogen peroxide is produced: H 2 O + O 3 = H 2 O 2 + O 2 . Hydrogen peroxide occurs in very small amounts in the air, and this is probably due to the fact that ozone has been pro- duced, which in turn has acted on the moisture in the air. It is consequently very doubtful whether ozone itself occurs in air. It should be stated here that the occurrence of hydrogen peroxide in the air has been questioned by some chemists, the claim being made that the strong- oxidations observed may very well be caused by oxides of nitrogen which are present in the atmosphere. Hydrogen peroxide may also be formed by treating cold, dilute hydrochloric acid with sodium peroxide : 2 HC1 + Na 2 2 = 2 NaCl + H 2 O 2 . Both the sodium chloride and hydrogen peroxide remain in solution. Instead of the peroxide of barium or sodium, that of potassium or strontium may be used. By distilling an aqueous solution of hydrogen peroxide in a partial vacuum, the water passes off first, leaving hydrogen peroxide in the retort. 96 OUTLINES OF CHEMISTRY On heating a 3 per cent solution of hydrogen peroxide on the water bath to temperatures below 70, in a retort from which the air has been exhausted so as to create a partial vacuum, a 45 per cent solution may readily be obtained without loss. On continuing the distillation further, nearly pure hydrogen perox ide passes over between 84 and 85 at 68 mm. pressure. Properties of Hydrogen Peroxide. Pure hydrogen peroxide is a colorless, sirupy liquid, which, like water, has a bluish hue in thick layers. At its specific gravity is 1.458. It boils at 69 under 26 mm. pressure, and at 84 to 85 under 68 mm. pressure. It forms colorless prismatic crystals which melt at - 2. Hydrogen peroxide slowly decomposes into water and oxygen on standing. In the sunlight the decomposition proceeds more rapidly. By warming hydrogen peroxide the rate of decom- position is increased ; and at 100 the evolution of oxygen becomes so rapid as to cause explosion. It is, therefore, neces- sary to distill hydrogen peroxide in a vacuum, so that it will not need to be heated to a temperature at which violent decom- position sets in. Solutions of hydrogen peroxide have a peculiar bitter, disa- greeable taste. Concentrated solutions act on the skin. The aqueous solutions on the market usually contain about 3 per cent hydrogen peroxide, though 30 per cent solutions are also now placed on sale. The latter are kept in small bottles coated with paraffin on the inside; for in contact with glass the solu- tion soon suffers decomposition on account of the fact that alkali is dissolved from the glass. In contact with platinum black, manganese dioxide, or finely divided silver, gold, or carbon, hydrogen peroxide is decomposed into oxygen and water even at room temperatures and in dilute solutions. The action is more rapid at higher temperatures. All these cases are illustrations of catalytic or contact action. Hydrogen peroxide is 2, strong oxidizing agent. It will act on black sulphide of lead and convert it into lead sulphate, which is a white salt : PbS + 4 H 2 O 2 = PbSO 4 + 4 H 2 O. Potassium iodide in solution is oxidized thus : 2 KI + H 2 O 2 = 2 KOH + I 2 . OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE 97 For this reason, starch potassium iodide paper may be used to detect the presence of hydrogen peroxide. The action goes on much more slowly than in the case of ozone ; but the addition of a little ferrous sulphate hastens the action very markedly, so that the test is really a sensitive one. In presence of ozone, which also liberates iodine from potassium iodide, this test for hydrogen peroxide can, of course, not be used. Hydrogen peroxide does not oxidize a bright silver foil as ozone does, and thus the latter may be detected in presence of the former. In contact with blood, meat, and the mucous membranes, hydrogen peroxide decomposes. The oxygen thus liberated destroys germs by oxidizing them, hence the use of hydrogen peroxide in medicine as a gargle and an antiseptic. When lime water is treated with hydrogen peroxide solution, a precipitate of calcium peroxide is formed : - Ca(OH) 2 + H 2 O 2 = CaO 2 + 2 H 2 O. The action on the hydroxide of barium or strontium is similar. All of these peroxides may be regarded as hydrogen peroxide in which the hydrogen is replaced by metals. When hydrogen peroxide solution is slightly acidified with sulphuric acid, and a few drops of potassium bichromate solu- tion and some ether are added, and the mixture is then shaken, an indigo -blue compound is formed which dissolves in the ether and so finally collects in the light ethereal layer on standing. This reaction is used as a test for hydrogen peroxide. The nature of the blue compound is not known with certainty, though it is probably perchromic acid. Wliile hydrogen peroxide is an oxidizing agent, it may also at times act as a reducing agent, in which case ordinary oxygen gas is evolved. So the oxides of metals like silver, gold, and plati- num suffer reduction to the metallic state when treated with hydrogen peroxide, thus : Ag 2 + H 2 2 = 2 Ag + H 2 + O r We see that hydrogen peroxide in such cases loses one atom of oxygen which unites with oxygen of the metallic oxides and escapes as ordinary oxygen gas. Lead peroxide is changed to lead monoxide : PbO 2 + H 2 O 2 = PbO + II 2 O + 2 . Added to a potassium permanganate solution acidified with sul 98 OUTLINES OF CHEMISTRY phuric acid, hydrogen peroxide reduces the permanganate with liberation of oxygen and formation of a solution of potassium sulphate and manganous sulphate, which is nearly colorless : - This reaction is used in the quantitative determination of the strength of hydrogen peroxide solutions ; for if a certain volume of a potassium permanganate solution of known strength is just decolorized by a known volume of a hydrogen peroxide solu- tion, the strength of the latter can readily be computed from the data given in the above equation. It would seem rather peculiar that hydrogen peroxide, which is a good oxidizing agent, may also serve in effecting reductions. It must be borne in mind, however, that it only reduces com- pounds that are rich in oxygen which is readily set free. The explanation of the reduction is that when compounds like potas- sium permanganate, or oxides of silver, gold, lead, etc., are brought in contact with hydrogen peroxide, the tendency to form the ordinary oxygen molecule O 2 , that is, the attraction of oxygen for oxygen, is so great that the compounds mutually reduce each other. Formula of Hydrogen Peroxide. Thenard, the discoverer of hydrogen peroxide, determined that it consists of 16 parts of oxygen to 1 part of hydrogen by weight. The simplest for- mula one could assign to the compound would therefore be HO, the atomic weight of oxygen being 16. However, the fact that water H 2 O and oxygen are formed when hydrogen peroxide decomposes, is much better indicated by adopting the formula H 2 O 2 for the latter substance. The vapor density of hydrogen peroxide cannot well be determined because the sub- stance is so unstable, and so the weight of 22.38 liters of its vapor under standard conditions is unknown. Its molecular weight has, however, been found to be 34, from a study of the freezing point of its aqueous solution. The fact that hydrogen peroxide decomposes into water and TJ\ oxygen has led Kingsett to ascribe to it the formula T [ /O = O, in which it will be seen that one oxygen atom is regarded as a tetrad and the other as a dyad. From a study of the index of refraction of the substance, Briihl has on the other hand sug- gested that both oxygen atoms are tetrads and that the formula OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE 99 should be written, H O = O H. As" a rule chemists regard both atoms of oxygen as bivalent, writing the structural for- mula of hydrogen peroxide, H O O H. A structural formula expresses not only the qualitative and quantitative composition of a substance and its molecular weight, but it also indicates its chemical behavior. This is accomplished by arranging the relative position of the atoms in the formula so as to indicate what chemical changes the compound will undergo. Uses of Hydrogen Peroxide. As already stated, hydrogen peroxide is used in medicine as a germicide. As such it has the distinct advantage that, after it has acted, only water re- mains, which is harmless. The usual 3 per cent solution on the market is also called dioxogen ; it frequently is diluted further as required. It is kept in brown bottles, in a cool place, and is generally very slightly acidified, which greatly reduces the rate of its decom- position by neutralizing the alkali that is taken up from the glass of the bottle. Hydrogen peroxide is manufactured on a large scale, and most of it is employed as a bleaching agent. Thus, delicate silks, ostrich feathers, ivory, hair, and sponges are bleached with hydrogen peroxide. It is used to change dark-colored living hair to lighter color. It is also employed similarly in changing the color of furs. In these bleaching processes, hy- drogen peroxide is used because it is a mild agent, which does not injure .these animal tissues as much as other bleaching agents do. Hydrogen peroxide is also used in photography to remove the last traces of sodium thiosulphate from the pho- tographic plates, after the latter have been "fixed." In ana- lytical chemistry it is frequently employed as an oxidizing agent. Ozonic Acid. Baeyer and Villiger have described an oxide of hydrogen which contains still more oxygen than hydrogen peroxide. This compound, to which the formula HO 2 or H 2 O 4 has been assigned, has been called ozonic acid, because it may be regarded as formed by the addition of ozone to water : 3 +H 2 = H 2 4 . Ozonic acid has not yet been isolated, but Baeyer and Villiger regard the peroxide of potassium K 2 O 4 , for instance, as a salt of ozonic acid, the two potassium atoms having replaced the hydrogen atoms. 100 OUTLINES OF CHEMISTRY REVIEW QUESTIONS 1. Give two essentially different methods by means of which ozone may be prepared, and state the important properties of this substance. 2. Under standard conditions, how much ozone could be prepared from 2.5 liters of oxygen? What would be the weight of the ozone in grams? 3. Define "allotropy." 4. Why should the formula of ozone be written 3 ? 5. What are the important properties of ozone? What use may be made of ozone? 6. What action has ozone upon potassium iodide? Upon water? Upon metallic silver? Write the appropriate chemical equation in each case. 7. How determine experimentally whether oxygen or ozone is the more active chemically? 8. Describe two methods of making hydrogen peroxide. Write the chemical equations expressing the changes that take place. 9. How prepare pure hydrogen peroxide from sodium peroxide ? 10. What use is made of hydrogen peroxide ? Upon what property of this Compound does its use depend? 11. Mention three catalytic agents that will decompose hydrogen peroxide. 12. How distinguish between ozone and hydrogen peroxide? 13. Why is the formula of hydrogen peroxide written H 2 2 ? 14. Explain the action of hydrogen peroxide upon potassium iodide- starch paper. 15. Mention some of the different structural formulas that have been proposed for hydrogen peroxide and point out the balance of the oxygen atoms in each formula. 16. Explain the use of the prefix per in the name hydrogen peroxide. Why may this substance also be called hydrogen dioxide ? 17. Why does hydrogen peroxide attack the mucous membranes and "set the teeth on edge"? 18. How much per cent of pure hydrogen peroxide does the ordinary commercial solution of hydrogen peroxide contain? Why should such solutions be kept in a cool place and not exposed to the light ? 19. Compare the colors of liquid oxygen, ozone, water, and hydrogen peroxide. Explain. 20. How much pure hydrogen peroxide may be prepared from 25 pounds of barium dioxide ? 21. Give an illustration showing that hydrogen peroxide, while com- monly an oxidizing agent, may yet also act as a reducing agent. 22. State briefly the history of the discovery of ozone. 23. By whom was hydrogen peroxide first prepared ? What method was used in the preparation of this compound? Why is not this method used at present? Write the appropriate chemical equations. CHAPTER VIII THE HALOGENS The Halogen Family. The elements that belong to thia group are flu.or.irie, chlorine, bromine, and iodine. Of these chlorine is the most common and the most abundant in nature. Its properties have already been discussed. Fluorine, bromine, and iodine form with hydrogen the compounds hydrogen fluoride or hydrofluoric acid HF, hydrogen bromide or hydro- bromic acid HBr, and hydrogen iodide or hydriodic acid HI. These compounds are analogous to hydrogen chloride or hydro- chloric acid HC1. By replacing the hydrogen of these hydro- halogen acids by means of sodium, the sodium salts, sodium fluoride NaF, sodium chloride NaCl, sodium bromide NaBr, and sodium iodide Nal are formed. These salts are quite similar to one another ; and as common salt is a member of the group, the elements fluorine, chlorine, bromine, and iodine have been termed the halogens, meaning salt formers. This must not be taken to mean that all salts contain one of these four ele- ments, for such is not at all the case. With the exception of fluorine, the halogens unite with oxygen and hydrogen to form certain acids. Chlorine and iodine also unite with oxygen to form oxides. Furthermore, the halogens form compounds with one another, with the metals, and with many other elements. We shall now take up the compounds which chlorine forms with oxygen and hydrogen, after which the remaining halogens and their principal compounds will be considered. Compounds of Chlorine with Oxygen. There are three of these compounds, namely, chlorine monoxide C1 2 O, chlorine dioxide C1O 2 , and chlorine heptoxide C1 2 O 7 . These are all very unstable substances, decomposing readily into chlorine and oxygen. They are not formed by direct interaction of chlorine and oxygen. Chlorine monoxide is formed when chlorine acts on cold mer curie oxide : 2 HgO + 2 C1 2 = HgO - HgCl 2 + C1 2 O. 101 102 OUTLINES OF CHEMISTRY It is a brownish yellow gas, which may be condensed to a liquid boiling at +5. The substance, especially when liquefied, is highly explosive. It detonates when heated or subjected to concussions; but in the sunlight it soon decomposes into chlorine and oxygen without explosion. Chlorine dioxide is formed when potassium chlorate is treated with concentrated sulphuric acid. The reaction may be re- garded as taking place in two steps, thus : (1) KC1O 3 + H 2 SO 4 = KHSO 4 + HClO r chloric acid (2) 3HClO 3 =HClO 4 -j-H 2 O + 2ClO 3 . perchloric acid Chlorine dioxide is also called chlorine peroxide. It is a yellow gas which may be condensed to a liquid, boiling at +9.9. Solid chlorine dioxide melts at 79. The substance is very explosive. Its odor resembles that of chlorine. In the sunlight it slowly decomposes into the elements. It is a powerful oxidizing agent. Sugar mixed with potassium chlorate bursts into flame when touched with a drop of concentrated sulphuric acid ; for thus chlorine peroxide is liberated, which at once attacks the sugar violently. Phosphorus introduced into chlorine peroxide gas at once takes fire. When the gas is touched with a red-hot iron, it explodes. Chlorine heptoxide is formed by the action of phosphorus pentoxide on perchloric acid. The action simply consists of the elimination of a molecule of water from two molecules of perchloric acid : 2 HC10 4 = H 2 + C1 2 7 . Chlorine heptoxide is a colorless oil which bo'ls at 82. On percussion it explodes with violence, also when brought in con- tact with a flame. It is therefore a dangerous substance to handle, and great care must be exercised in distilling it. Hypochlorous Acid and Hypochlorites. When chlorine mon- oxide acts on water a solution of hypochlorous acid is formed : C1 2 O + H 2 O = 2 HOC1. Hypochlorous acid is known only in solution and in form of its salts. THE HALOGENS 103 When caustic potash solution is treated with chlorine at room temperatures, the following change occurs : 2 KOH + Cl a = KOC1 + KC1 + H 2 O. potassium potassium hypochlorite chloride A perfectly analogous change occurs when chlorine acts on cal- cium hydroxide, slaked lime : 2 Ca(OH) 2 + 2 C1 2 = Ca(OCl) 2 + CaCl 2 + 2 H a O. calcium hypochlorite The product is bleaching powder or so-called chloride of lime. It consists of calcium hypochlorite Ca(OCl) 2 and calcium chloride CaCl 2 . The formula of bleaching powder is, however, best expressed thus : Ca^ ^ , , for the substance really con- tains no calcium chloride, since it lacks the hygroscopicity of the latter salt. Furthermore, alcohol will not extract calcium chloride from bleaching powder, though calcium chloride is soluble in alcohol. By treating calcium hypochlorite with very dilute, cold nitric acid HNO 3 , hypochlorous acid is liberated, thus : Ca(OCl) 2 + 2 HN0 3 - Ca(NO 3 ) 2 + 2 HOCL Hypochlorous acid readily loses oxygen and passes over into hydrochloric acid, especially in the sunlight : HOC1=HC1 + 0. The nascent oxygen thus liberated readily oxidizes substances like coloring matters, and hence hypochlorous acid is an oxidiz- ing and bleaching agent. Hypochlorous acid has twice the bleaching power possessed by chlorine water containing the same amount of chlorine, as is evident from the following equations : = 2HC1+0 2 . Hypochlorous acid readily decomposes in sunlight into oxy- gen and hydrochloric acid. Concentrated solutions readily form chloric acid and hydrochloric acid : 3 HC1O = 2 HCi + HClOg. 104 OUTLINES OF CHEMISTRY Hypochlorites are also oxidizing agents. They part with their oxygen and pass over into chlorides. Thus calcium hypochlo- rite slowly forms calcium chloride and oxygen : Ca(OCl) 2 = CaCl 2 + O 2 . Upon this fact depends the bleaching action of bleaching powder, and also its disinfecting action, for the oxygen liberated destroys organic matter. Javelle water is made by treating bleaching powder with sodium carbonate. By treating bleaching powder with sulphuric acid all the chlorine is liberated, thus : The chlorine liberated then acts upon water, forming hydro- chloric acid and oxygen, the latter destroying the coloring matter to be bleached. Hence, in using bleaching powder in practice it is generally treated with an acid. Chloric Acid and Chlorates. When a solution of potassium hypochlorite is heated, the following change occurs : 3 KC1O = 2 KC1 + KC1O 3 . potassium chlorate Potassium chlorate may be formed directly by saturating a hot solution of caustic potash with chlorine, thus : 6 KOH + 3 C1 2 = 5 KC1 + KC1O 3 + 3 H 2 O. As potassium chlorate KC1O 3 is much less soluble in water than potassium chloride KC1, the former readily crystallizes from a hot saturated solution on cooling. By treating potassium chlorate with dilute sulphuric acid, chloric acid, whose composition is represented by the formula HClOg, is liberated :- 2 KC1O 3 + H 2 SO 4 = K 2 SO 4 + 2 HC1O 3 . This reaction is perfectly analogous to that of making hydro- chloric acid : 2 KC1 + H 2 SO 4 = K 2 SO 4 + 2 HC1. Chloric acid is known only in solution and in form of its salts. Its anhydride C1 2 O 5 is not known at all. Chloric acid solutions, forming thick, colorless sirups of specific gravity 1.25, have been obtained. They contain 40 per cent of the free acid and corre- spond approximately to the formula, HC1O 3 -4- 7 H 2 O. THE HALOGENS 105 Attempts to concentrate the solution farther always result in decomposition of the chloric acid into chlorine, oxygen, and perchloric acid. The sirupy solution of chloric acid oxidizes linen, wood, paper, and other organic material very rapidly, with evolution of light and heat. The aqueous solutions of the acid are much more stable than those of hypochlorous acid; still, on standing perchloric acid is formed in them, especially in the sunlight. The salts of chloric acid, namely the chlorates, are much more stable than the hypochlorites. Perchloric Acid and Perchlorates. When potassium chlorate is melted, it gives off oxygen slowly and then becomes nearly solid, forming potassium chloride and potassium perchlorate : 4 KC1O 3 = 3 KC1O 4 + KC1. The potassium perchlorate KC1O 4 is much less soluble in water than potassium chloride, hence the latter salt may readily be separated from the former. Sodium perchlorate NaClO 4 is found in small amounts in Chili saltpeter. On heating potassium perchlorate, it gives up all of its oxygen, passing over into potassium chloride. Hence, the formation of potassium perchlorate is really an intermediate step in the making of oxygen by heating potassium chlorate. By treating potassium perchlorate with strong sulphuric acid, perchloric acid is formed : KC1O 4 + H 2 SO 4 = KHSO 4 + HC1O 4 . Perchloric acid HC1O 4 may also be produced by heating chloric acid, or by exposing the latter to sunlight : 3 HC1O 3 = C1 2 + 2 O 2 + H 2 O + HC1O 4 . Perchloric acid, prepared by carefully distilling a mixture of potassium perchlorate and sulphuric acid in a partial vacuum, is a colorless, very corrosive liquid which fumes strongly in the air. It has a specific gravity of 1.782 at 15.5 and a boiling point of about 40 at 60 mm. pressure. It is the most stable of the oxy-acids of chlorine ; still, it cannot be kept long even in the dark, for after a few days decomposition with violent explosion occurs. The acid is a dangerous product. In con- tact with the skin it produces wounds that are painful and very slow to heal. A few drops put on paper, wood, etc. causes these substances to burst into flames, while a drop of 106 OUTLINES OF CHEMISTRY the acid on charcoal produces a violent explosion. These phenomena occur because perchloric acid is very rich in oxy- gen, with which it parts readily, thus producing violent oxida- tion accompanied with sudden liberation of much heat. The anhydride of perchloric acid is chlorine heptoxide C1 2 O 7 , which, as has been stated, is produced by abstracting water from perchloric acid by treatment with phosphorus pentoxide. Nomenclature and General Relations. The following table presents the formulae and names of the compounds of chlorine with oxygen and hydrogen : HC1, hydrochloric acid. KC1, potassium chloride. HC1O, hypochloroMS acid. KC1O, potassium hypochlori'te. (HC1O 2 , chlorous acid). KC1O 2 , potassium chlorite. HC1O 3 , chloric acid. KC1O 3 , potassium chlorate. HC1O 4 , perchloric acid. KC1O 4 , potassium perchlorate. Chlorous acid HC1O 2 is not known in the free state ; but its salts, like potassium chlorite KC1O 2 , are known. The latter, for instance, is formed together with potassium chlorate when chlo- rine dioxide acts on caustic potash : 2 KOH + 2 C10 2 = KC10 2 + KC1O 3 + H 2 O. The above table presents an interesting series of compounds. Beginning with hydrochloric acid and its salt potassium chlo- ride, each member of the series contains one atom of oxygen more than the preceding. Hydrochloric acid is a very stable compound ; but hypochlorous acid is very unstable. On the other hand, chloric acid is more stable than hypochlorous acid, and perchloric acid is the most stable of the three known oxy- acids of chlorine. The salts of these acids, obtained by replac- ing the hydrogen of the acid by a metal, are much more stable than the corresponding acids. Such salts form articles of com- merce. Their uses will be considered more fully later. The names given the oxy-acids of chlorine and their corre- sponding salts afford an excellent illustration of the system of naming a series of acids of increasing oxygen content and the salts which they form. From the table it appears that HC1O 3 is called chloric acid and its salts chlorates ; the acid which is richest in oxygen, HC1O 4 , is called perchloric acid, and its salts perchlorates ; the acid containing less oxygen than chloric acid, THE HALOGENS 107 namely HC1O 2 , is termed chlorous acid and its salts the chlorite* ; whereas the acid containing still less oxygen, HC1O, is called hypochlovous acid and its salts hypochlorite*. Finally, HC1, which contains no oxygen at all, is termed hydrochloric acid, which distinguishes it sufficiently from HC1O 3 , chloric acid. This method of naming acids and their corresponding salts is generally applied in chemistry whenever a similar series of com- pounds is found. The io acid forms the ate salt ; the ous acid forms the ite salt ; the hypo . . ous acid forms the hypo . . ite salt ; and the per . . ic acid forms the per . . ate salt. Numerous other illustrations of this will be met in our further considerations. Occurrence, Preparation, and Properties of Fluorine. This element is widely distributed in nature. It occurs in large quantities, but always in combination with other elements. It is chiefly found combined with calcium as fluorspar, calcium fluoride CaF 2 , which crystallizes in octahedra and in cubes like common salt. In Greenland, fluorine is found in the mineral cryolite, which is a fluoride of sodium and aluminum, the com- position of which is expressed by the formula (NaF) 3 . A1F 8 . In many minerals and siliceous rocks fluorine occurs in small quantities, in combination with calcium and other metals. Fluorides also are found in small quantities in sea water, in many mineral waters, in the ashes of plants, and in the teeth and the bones of animals. Fluorspar has been known for a very long time. It melts at red heat, and has been used as a flux in metallurgical processes as early as the fifteenth century. It used to be called fluate of lime. The name "fluorine" comes from the use of fluorspar as a flux. Fluorine was not isolated till 1886, when Henri Moissan pre- pared it by passing the electric current through dry, liquid hydrofluoric acid HF, in which potassium hydrogen fluoride KHF 2 had been dissolved, in order to have the liquid conduct electricity. The solution was placed in a tube made of platinum (Fig. 29), the stoppers being made of fluorspar. The electrodes were made of an alloy of platinum and iridium. The apparatus was kept at 23 C., and the fluorine was collected in a platinum tube, the ends of which were closed with transparent plates of fluorspar. The difficulty in isolating fluorine lies in the fact that the element combines so readily with other elements. Moissan found later that perfectly pure fluorine attacks glass 108 OUTLINES OF CHEMISTRY FIG. 29. but very slowly indeed, so that the gas may be collected in glass vessels. It has also been demonstrated that a copper vessel may be used instead of one of platinum in pre- paring fluorine. Fluorine is a gas of a light, greenish yellow color and a strong pun- gent odor. It may be condensed to a liquid which boils at -187. By chilling the liquid with liquid hydrogen, it freezes, the white crystals formed melting at - 223. Fluorine gas is 19 times as heavy as hydrogen. Its molec- ular weight is therefore 38 ; and since its atomic weight is 19.0, the formula of fluorine is F 2 . The atomic weight of fluorine has been determined from the analysis of calcium fluoride. Fluorine is the most active of all the elements. It acts on water, yielding ozone and hydrofluoric acid : 3 H 2 + 3 F 2 = 6 HF + O 8 . It unites with hydrogen with great violence in the dark at ordinary temperatures, and even at 253 solid fluorine still acts with explosive violence on liquid hydrogen, according to Dewar and Moissan. Most of the non-metallic elements unite directly with fluorine at ordinary temperatures with evolution of heat and light. Iron, lead, barium, strontium, calcium, so- dium, and potassium are acted upon by fluorine at ordinary temperatures ; magnesium, aluminum, manganese, nickel, and silver burn in fluorine when slightly heated. At ordinary temperatures gold and platinum are not attacked, but between 300 and 400 they are converted into fluorides. Copper is acted upon at ordinary temperatures, a coating of cuprous fluoride being formed at once on the metal, which is thus pro- tected from further action. Oxygen, chlorine, nitrogen, and THE HALOGENS 109 argon do not unite with fluorine. Organic substances generally burn in fluorine gas. Hydrochloric acid gas is decomposed by fluorine with explosive violence: 2 HC1 + F 2 = 2 HF + C1 2 . Dry glass is but very slowly attacked by fluorine, but in pres- ence of hydrofluoric acid or water, even in traces, glass is rapidly destroyed. Hydrofluoric Acid. When calcium fluoride is treated with sulphuric acid, hydrofluoric acid is formed : CaF 2 + H 2 S0 4 = CaS0 4 + 2 HF. The experiment is carried on in a platinum or lead dish, for hydrofluoric acid acts upon glass or porcelain. The process of making the acid is perfectly analogous to that of preparing hydrochloric acid. To obtain hydrofluoric acid which is anhy- drous, i.e. free from water, potassium hydrogen fluoride KHF t is heated to redness in a platinum retort : KHF 2 = KF + HF. Hydrofluoric acid is a liquid whose boiling point is + 19.4. Solid hydrofluoric acid melts at 92.3. When perfectly dry, the liquid does not act upon glass. In presence of moisture, however, glass is rapidly attacked, fluorides and water being formed. Glass consists essentially of the silicates of sodium and calcium, Na 2 SiO 3 and CaSiO 3 . When hydrofluoric acid acts upon these, the following changes occur : CaSiO 3 + 6 HF = CaF 2 + SiF 4 + 3 H 2 O. Na 2 SiO 3 + 6 HF = 2 NaF + SiF 4 + 3 H 2 O. The compound silicon tetrafluoride SiF 4 is a gas, and so es- capes. Calcium fluoride is soluble in acids, so that when glass is attacked by hydrofluoric acid, it is dissolved. Use is made of this fact in the chemical analysis of glass and other silicates, also in etching glass. In the latter process the glass is first coated with paraffin ; the design is traced in the paraffin coat- ing so as to expose the portions of the glass to be etched, and the whole is then treated either with the fumes of hydrofluoric acid or with an aqueous solution of the latter. When the par- affin is finally removed, the design is found etched into the sur- 110 OUTLINES OF CHEMISTRY face of the glass. This process is used in marking graduations on glass utensils, thermometers, etc. Because hydrofluoric acid attacks glass, it is kept in rubber or wax bottles. It is very soluble in water, and fumes in contact with moist air. The concentrated solution boils at 120, and contains about 36 to 38 per cent of the anhydrous acid. Hydrofluoric acid is a dangerous substance, for it is very poisonous. When inhaled it produces death. In contact with the skin it produces swellings, pains, and wounds that are very slow to heal. At 100 hydrogen fluoride is about ten times as heavy as hydrogen ; this leads to the molecular formula HF. But at 25 the vapor of hydrofluoric acid is nearly 20 times as heavy as hydrogen, which leads to the formula H 2 F 2 . The acid is prone to form acid salts like KHF 2 and NaHF 2 . The other hydro- halogens do not form analogous compounds. Occurrence, Preparation, and Properties of Bromine. Like chlorine, bromine does not occur in nature except in combina- tion with other elements. Bromine is generally found in nature with chlorine in salt deposits. And just as chlorine occurs mainly in form of sodium chloride, so bromine occurs chiefly as sodium bromide. Bromine is widely distributed in nature, but it is not found anywhere in very large quantities. In sea water, sodium bromide . and magnesium bromide are found. Together these constitute from 0.3 to 1.3 per cent of the residue obtained by evaporating the water. In the Stass- furt salt beds, bromine occurs as magnesium bromide. The salt wells of West Virginia, Ohio, and Michigan furnish most of the bromine used in the United States. Here the element occurs as sodium bromide together with common salt. On evaporating the brine the sodium chloride is first deposited, it being less soluble than sodium bromide. From the mother liquor sodium bromide, together with some common salt, is obtained by further evaporation. In the year 1910 the United States produced 245,437 pounds of bromine valued at $41,684. Bromine was discovered in 1826 by Balard, who prepared it from the residue obtained by evaporating sea water. The method of preparing bromine is the same as that of pre- paring chlorine, thus : - 2 = 2NaHSO 4 +MnSO 4 +2H 2 O-|-Br 9 , THE HALOGENS 111 Chlorine will readily replace bromine, and so this method ma} be employed in making bromine : 2 NaBr + C1 2 = 2 NaCl + Br 2 . MgBr 2 + C1 2 = MgCl 2 + Br 2 . This method is used in manufacturing bromine in Michigan and at Stassfurt. Bromine is the only non-metallic element which is a liquid at ordinary temperatures. It is dark reddish brown in color, boils at 59, and has a specific gravity of 3.188 at 0. At - 7.5 it freezes to a dark brown solid, and at 98 it crystallizes from carbon disulphide in carmine red needles. At room tem- peratures bromine vaporizes readily. It irritates the eyes and the mucous membranes of the mouth and throat and has an extremely disagreeable odor ; whence its name bromine, mean- ing a stench. In contact with the skin it produces wounds that are painful and difficult to heal. Bromine dissolves in water. The solution has the color of bromine and is known as bromine water. At room temperature 20 the saturated solution contains about 3 per cent bromine. On cooling the solution to about a hydrate of the composi- tion Br 2 -f 10 H 2 O separates out. This, however, readily decomposes when warmed to room temperature. In its chemical behavior, bromine closely resembles chlorine. With metals and a large number of other elements it unites directly, forming bromides. Thus arsenic and antimony wi-.l burn in bromine, which also reacts vigorously with phosphorus and sulphur. On the other hand, it does not unite with carbon or oxygen directly. It acts violently on potassium ; but dry sodium may even be heated with bromine up to 200 before appreciable action begins. Bromine turns starch paste yellow. It bleaches like chlorine, only much more slowly. The bleach- ing action depends upon the fact that bromine, like chlorine, acts upon water, liberating oxygen, which attaoks organic coloring matters, thus: The atomic weight of bromine is 79.92, and as its vapor is about 79.5 times heavier than hydrogen, its molecular weight is 159.84, and its molecular formula is Br 2 . 112 OUTLINES OF CHEMISTRY Hydrobromic Acid. Hydrogen bromide, or hydrobromic acid HBr, may be formed by direct union of hydrogen with bromine, which occurs when hydrogen charged with bromine vapor is ignited. By treating sodium bromide with sulphuric acid hydrogen bromide is formed, just as hydrogen chloride forms when common salt is similarly treated, thus : NaBr + H 2 SO 4 = NaHSO 4 + HBr. However, in this case a portion of the hydrobromic acid liber- ated at once reacts with some of the sulphuric acid, forming bromine, water, and sulphur dioxide. That is to say, some of the hydrobromic acid reduces sulphuric acid : H 2 SO 4 + 2 HBr = SO 2 + 2 H 2 O + Br 2 . Thus, pure hydrobromic acid cannot be obtained by treating sodium bromide with sulphuric acid, for the product contains free bromine and also sulphur dioxide. Pure hydrogen bromide is formed when phosphorus tribro- mide PBr 3 or phosphorus pentabromide PBr g is acted upon by water, thus : PBr 3 + 3 H 2 O = H 3 PO 3 + 3 HBr. phosphorous acid PBr 5 + 4 H 2 = H 3 P0 4 + 5 HBr. phosphoric acid Phosphorous and phosphoric acids are not volatile, but hydro- bromic acid is, and so the latter can readily be separated from the former. The apparatus used for making hydrobromic acid is shown in Fig. 30. The flask F contains red phosphorus covered with a little water. Bromine is gradually added by opening the cock (7. Phosphorus bromide is formed, which is at once decomposed by the water, yielding hydrobromic acid. The latter generally contains some bromine vapor, which is removed by allowing the gas to pass over pumice covered with moist red phosphorus in the U-tube A. Hydrobromic acid is a colorless gas of strong pungent odor. It may be condensed to a liquid which boils at 64.9 under 738.2 mm. pressure. It forms colorless crystals which melt at 88. It fumes strongly in the air, and is very soluble in THE HALOGENS 113 water, one volume of the latter absorbing about 600 volumes oi hydrobromic acid gas at 10. Hydrobromic acid is a strong acid which readily attacks many metals, forming bromides of the metals and liberating hydrogen, thus : Mg + 2 HBr = MgBr 2 + H 2 . Zn + 2HBr=ZnBr 2 + H 2 . In general, the chemical behavior is like that of hydrochloric acid. Like the chlorides of the metals, the bromides are gen- erally soluble in water ; and just as the chloride of silver AgCl FIG. 30. is insoluble, so the bromide of silver AgBr is also insoluble. Further, the bromide of lead PbBr 2 and mercurous bromide HgBr are difficultly soluble like the corresponding chlorides. When hydrobromic acid is treated with chlorine, hydrochloric acid and bromine are produced : 2HBr + Cl a =2HCl + Br a . On boiling, a strong aqueous solution of hydrobromic acid becomes weaker, and a weak solution becomes stronger, till finally a solution containing from 47.4 to 47.8 per cent of hydrobromic acid is formed. This solution then distills over unchanged in concentration at 752 to 762 mm. pressure. How- ever, by distilling it at other pressures its strength is changed. 114 OUTLINES OF CHEMISTRY Thus at 16 mm. pressure the acid that distills over contains 51.6 per cent HBr. Hydrobromic acid is 40.45 times heavier than hydrogen. Its molecular weight is therefore 80.9. By weight it contains 1.008 grams of hydrogen to every 79.92 grams of bromine. From these data, its formula is HBr. When dry hydrobromic acid gas is treated with metallic sodium in an apparatus like that used in investigating the composition of HC1 (Fig. 23), it is found that the hydrogen liberated occupies one half of the volume of the hydrobromic acid taken. At very high temperatures hydrobromic acid gas decomposes into hydrogen and bromine. This reaction, which is a rever- sible one, like the decomposition of hydrochloric acid gas and of water in the gaseous state at very high temperatures, is another typical case of dissociation, and may be represented thus : The term dissociation is only applied to reversible reactions in which a compound is decomposed into products which may again unite to form the original compound as the pressure, temperature, or amount of material contained in unit of volume is varied . We shall have occasion to refer to other instances of dissociation. Oxy-acids of Bromine. There are no oxides of bromine known, and but two ^oxy-acids have thus far been prepared. They are hypobromous acid HBrO and bromic acid HBrO 3 , the latter being known in aqueous solution only. Hypobromous acid and its salts, the hypobromites, are prepared in a manner analogous to the preparation of hypochlorous acid and hypochlorites. Thus, by action of bromine water on mer- curic oxide, hypobromous acid HBrO results, just as hypochlo- rous acid is formed when chlorine water acts on mercuric oxide. The changes are expressed as follows : 2 Br 2 + H 2 O + HgO = HgBr 2 + 2 HBrO. 2 C1 2 -h H 2 O + HgO = HgCl 2 + 2 HC1O. While hypochlorous acid is known only in aqueous solutions, hypobromous acid may be isolated by distillation in a partial vacuum at 40. The solution of the acid in water is straw-yel- low in color. The acid readily decomposes into hydrobromic THE HALOGENS 115 acid and oxygen, and is, therefore, like hypochlorous acid, a strong oxidizing and bleaching agent. When bromine acts on a cold solution of caustic alkali, hypo- bromites are formed. The process is analogous to the formation of hypochlorites, thus : 2 KOH + Br 2 = KBr + KBrO + H 2 O. Hypolromites, like hypochlorites, are unstable compounds, readily giving up oxygen. The hypobromite solutions yield bromates readily, especially at higher temperatures: 3KBrO = 2KBr+KBr0 3 ; or in warm caustic potash solution, bromine at once forms potassium bromate, thus : 3 Br 2 + 6 KOH = 5 KBr + 3 H 2 O + KBrO 3 . Bromic Acid and Bromates. When silver bromate AgBrO 3 is treated with bromine and water, bromic acid is formed : 5 AgBrOg + 3 Br 2 + 3 H 2 O = 5 AgBr + 6 HBrO 3 . It is also formed when dilute sulphuric acid acts on barium bromate : Ba(Br0 3 ) 2 + H 2 SO 4 = BaSO 4 + 2 HBrO 8 ; or when chlorine is passed into bromine water, thus : Br 2 + 6 H 2 O + 5 C1 2 = 10 HC1 + 2 HBrO 3 . Bromic acid is very similar to chloric acid in its behavior. At 100 the aqueous solution decomposes, yielding oxygen and bromine. The pure anhydrous acid has not been prepared. On heating potassium bromate, it yields oxygen and potas- sium bromide, without, however, first forming a potassium per- bromate. In this respect the behavior of potassium bromate KBrO 3 differs from that of potassium chlorate. By melting potassium bromide with potassium chlorate, po- tassium bromate results : IKClOg + KBr = KBrO 3 + KC1. We thus see that under these conditions the bromate is more stable than the chlorate, and the chloride more stable than the bromide. 116 OUTLINES OF CHEMISTRY Uses of Bromine and its Compounds. Bromine is used in the manufacture of dyestuffs from coal-tar products. In medicine potassium bromide is used as a sedative. In photography silver bromide is used in the sensitized plates. History and Occurrence of Iodine. This element is a solid at room temperatures. It was discovered in 1812 by Courtois, who evolved it from the ashes of seaweeds. It forms beautiful violet vapors, whence its name iodine, meaning violet-colored. Indeed it was the color of the vapor that led to the discovery of iodine. The substance was then examined by Sir Humphry Davy and by Gay-Lussac ; the latter in 1815 established its elementary character. Like bromine, iodine always occurs in nature associated with chlorine. It has been reported by Wanklyn that the water from the spring, Woodhall Spa, near Lincoln, Nebraska, con- tains iodine in minute quantities ; but with this singular excep- tion, iodine has always been found in combination with other elements, chiefly with sodium, potassium, magnesium, and cal- cium in form of iodides and iodates. In sea water it occurs in extremely minute quantity. Seaweeds, particularly those grow- ing in deeper waters, like the genera Fucus and Laminaria, assimilate iodine and store it up in their bodies. The ashes of such seaweeds are termed kelp in Scotland and varech in Nor- mandy, and from these iodine is prepared. However, the chief source of iodine at present is the crude Chili saltpeter, or caliche NaNO 3 , in which iodine occurs mainly as sodium iodate NaIO 3 . The amount of iodine in caliche is, however, only about 0.2 per cent. Besides occurring in seaweeds, iodine is found in many sponges, oysters, and other sea animals, in cod-liver oil, in some fresh-water plants, in coal, and in the thyroid glands of animals. In combination with silver, copper, and lead it occurs as iodides, though these minerals are rare. Many mineral springs contain minute amounts of iodine, and in deposits of common salt the element generally occurs in small quantity. Thus it is evident that iodine is quite widely distributed in nature, though it is nowhere present in large amounts. Preparation of Iodine. From the ashes of seaweeds, iodine is liberated by treatment with sulphuric acid and manganese dioxide, or by passing chlorine through the solution, which con- THE HALOGENS in tains the iodine mainly in the form of sodium iodide. The equations expressing the changes that occur are as follows : (2) It will be observed that equation (1) is perfectly analogous to the process of making chlorine or bromine from chlorides or bro- mides by treatment with sulphuric acid and manganese dioxide. Further, equation (2) is analogous to the process of making bromine from a bromide by treatment with chlorine. On the coasts of France and Scotland the seaweeds are gathered, dried, and burned, the latter process being carried on in closed retorts so that no iodine is lost by volatilization. The charcoal remaining after the ash has been leached out of it is sim- ilar to animal charcoal. It readily absorbs odors and is used as a deodorant. The iodine is then liberated from the solution by means of one of the processes just men- tioned. It is purified by volatilizing it and condens- ing the vapor, which forms crystals. The process of FlG 31 vaporizing a solid without melting it and condensing the vapor to the solid state is called sublimation. To get pure iodine the latter is mixed with potassium iodide, and the mix- ture is heated so as to volatilize the iodine, which is condensed on cool surfaces in form of crystals. In this way bromine and chlorine remain behind, in combination with potassium. Figure 31 shows a simple laboratory apparatus for subliming iodine, and Fig. 32 represents an arrangement for resubliming raw iodine on a commercial scale. From Chili saltpeter, in which iodine occurs as sodium iodate together with smaller amounts of sodium iodide and magnesium iodide, iodine is prepared by treatment with sodium bisulphite. The reaction which takes place is as follows : 118 OUTLINES OF CHEMISTRY 2 NalOg + 5 NaHSO 3 = 2 Na 2 SO 4 + 3 NaHSO 4 + H 2 O -t- 1 2 . The iodine is thus obtained in precipitated form from the aque- ous solution. It is allowed to settle and is then collected and purified by sublimation. The quantity of iodine produced annually from Chili saltpeter is about 300 tons, which is somewhat more than half of the total pro- duction. Of recent years, the process of obtaining iodine from seaweeds has been improved, the best method consist- ing of lixiviating the seaweeds without previous charring. In this way much less iodine is lost, and the remains of the seaweeds are worked up into algin, which is like gelatine. Thus the method of preparing iodine from seaweeds has again become profitable. Properties of Iodine. Iodine is a grayish black, lustrous solid which crystallizes in plates that belong to the rhombic system. From solutions in alcohol or hydriodic acid beautiful crystals may be obtained. Iodine really has a metallic luster. It is brittle and may be pulverized readily. Its specific gravity is 4.95 at 17. It melts at 116.1, forming a reddish brown liquid which boils at 184. Iodine volatilizes perceptibly, though slowly, at room temperatures. Its vapors are violet- colored, but when dense they appear very dark and opaque. Its odor reminds one of that of chlorine and bromine, but it is much less intense. It colors the skin brown and exerts an irritating and corrosive action upon it. In water it is but very slightly soluble, about 1 part in 5000. However, water containing potassium iodide or hydriodic acid readily dissolves iodine. These solutions are brown, as is also the solution of iodine in alcohol, which is called tincture of FIG. 32. THE HALOGENS 119 iodine. Iodine furthermore dissolves in hydrocarbon oils, chlo- roform, and carbon disulphide. With the latter it forms beau- tiful violet solutions, which fact is frequently used in detecting iodine in chemical analysis. We have already learned that iodine turns starch paste blue. This is used as a test for iodine and also for starch in analytical chemistry. Uses of Iodine. The solution of iodine in alcohol, tincture of iodine, is used as a counter irritant in medicine. Iodine is also administered internally in form of potassium iodide as a specific in certain diseases, particularly those which, like goiter, are caused by disturbances in the thyroid gland. The latter normally contains iodine in form of an organic compound known as thyroiodine. This also occurs in the thyroid glands of animals, particularly in sheep, from which source it is mainly obtained. It is administered as a specific for goiter and myxoadema. Iodine is also used in the manufacture of iodoform, iodocrol, and other iodine preparations. These are used principally as antiseptics in healing wounds. In synthetic chemistry, hydriodic acid and compounds of iodine with carbon and hydrogen are frequently employed. Hydriodic Acid. There is but one compound of hydrogen and iodine known ; namely^ hydriodic acid or hydrogen iodide. Its composition and vapor density are represented by the for- mula HI. It may be prepared by passing a mixture of hydrogen and iodine vapor through a red-hot tube containing platinum in a finely divided state, thus: The reaction is, however, incomplete since it is a reversible one, hydriodic acid decomposing readily into iodine and hydrogen. By treating potassium iodide with sulphuric acid, we cannot obtain hydriodic acid ; for the latter reduces sulphuric to sul- phurous acid far more readily than does hydrobromic acid: 2 KI + 3 H 2 SO 4 = 2 KHSO 4 + SO 2 + 2 H 2 O + 21. However, by treating potassium iodide with hot, concentrated phosphoric acid, hydriodic acid may be obtained, thus : KI + H 8 PO 4 = KH 2 PO 4 + HI. The acid is best prepared by decomposition of phosphorus iodide by water: 120 OUTLINES OF CHEMISTRY PI 3 + 3H 2 = H 3 P0 3 +3HI; or by simply having red phosphorus and iodine act on each other in presence of water. In this way phosphorus iodide is formed and then decomposed into phosphorous acid and hydrogen iodide: p + 3 I + 3 H 2 O == H 3 PO 3 + 3 HI. Hydrogen iodide may also be obtained by the action of iodine on hydrogen sulphide H 2 S (which see), or by the action of hydrogen sulphide upon cuprous iodide suspended in water, thus: = 2HI-|-S. These methods are quite similar to those by means of which pure hydrobromic acid can be obtained. Hydrochloric acid may also be prepared by similar methods ; but it is not at all necessary to resort to these in this case, since this acid is much more stable than hydrobromic or hydriodic acid. It does not reduce sulphuric acid, and can therefore readily be prepared by the action of the latter on common salt. At room temperatures hydrogen iodide is a colorless gas, which, like hydrochloric and hydrobromic acids, fumes strongly in the air and is very soluble in water. At 60, 485 volumes of hydriodic acid gas are absorbed by 1 volume of water. At a solution of specific gravity 2.0 may be obtained which contains about 90 per cent hydrogen iodide. On distillation, a solution of hydrogen iodide behaves like the corresponding solution of hydrogen bromide and hydrogen chloride. On boiling, a concentrated solution becomes weaker, and a weak solution becomes more concentrated, till finally a liquid is obtained which contains 57 per cent HI. This boils at 127 at 774 mm. and distills over without change of composition. ' On changing the pressure, however, the composition of the dis- tillate is changed. The distillation of the hydriodic acid must be conducted in a current of hydrogen to prevent decomposi- tion of the acid. Pure hydriodic acid may be condensed to a colorless liquid which boils at 34.1. Solid hydriodic acid forms colorless crystals which melt at 50.8. The vapor of hydriodic acid THE HALOGENS 121 is 62.92 times heavier than hydrogen, whence its moleculai weight is 126.84, which corresponds fairly well to the formula HI; for the atomic weight of iodine is 126.92, and the calcu- lated molecular weight for the formula HI is 127.92. Hydriodic acid forms iodides and hydrogen when treated with many metals. These salts are as a rule soluble in water, the exceptions being the iodides of silver, mercury, and lead. Hydriodic acid is a powerful reducing agent, which comes from the fact that it readily gives up its hydrogen to oxidizing agents. The decomposition of hydriodic acid proceeds more rapidly in the light and at higher temperatures. Its reducing power is frequently used in chemistry, especially in the investi- gation of the compounds of carbon. Oxide of Iodine. But one oxide of iodine is known. Its com- position is expressed by the formula I 2 O 5 . It is the anhydride of iodic acid, and is prepared by heating the latter to 170, thus : 2HI0 3 =H 2 + I 2 6 - When the oxide is dissolved in water, the acid is regenerated. Iodine pentoxide is a white crystalline solid, which decomposes into its elements at 300 ; thus it is much more stable than the oxides of chlorine. Oxy-acids of Iodine. There are three oxy-acids of iodine known, namely hypoiodous acid HIO, iodic acid HIO 3 , and periodic acid HIO 4 . A dilute solution of hypoiodous acid may be prepared by shaking together mercuric oxide, iodine, and water: * HgO + 2 I 2 + H 2 = 2 HIO + HgI 2 . The method is thus similar to the preparation of HC1O and HBrO. When iodine is introduced into cold caustic alkali solutions, a colorless liquid having bleaching power results, due to the .formation of hypoiodites, thus : 2 NaOH + T 2 = NalO + Nal + H 2 O. Hypoiodites are, however, extremely unstable, readily passing over into iodates, especially on warming, thus: Iodic acid HIO 3 is perfectly analogous to chloric and bromic 122 OUTLINES OF CHEMISTRY acids. It is, however, much more stable than the latter. It may be formed by treating barium iodate with sulphuric acid : Ba(I0 3 ) 2 + H 2 S0 4 = BaS0 4 + 2 HIO 8 ; or by oxidation of iodine either by means of chlorine or nitric acid, thus : I + 3 H 2 + 5 Cl = 5 HC1 + HIO 3 31 + 5 HNO 3 = 5 NO + H 2 O + 3 HIO 8 . nitric acid nitric oxide lodic acid readily gives up oxygen and is consequently a good oxidizing agent. Thus in contact with hydriodic acid, both acids are decomposed, yielding water and iodine : The salts of iodic acid are called the iodates. The potassium and sodium salts readily dissolve in water ; but, in general, the salts of other metals are sparingly soluble. On being heated, the iodates behave like the bromates. The iodates of sodium and potassium yield iodides and oxygen, whereas other iodates decompose into oxides of the metal, iodine, and oxygen. The iodates of potassium and sodium readily unite with one or two additional molecules of iodic acid, forming acid salts. Thus, KIOo - HIO q and KIO, 2 HIOo are known. The chlorates and 66 66 bromates do not thus add on chloric and bromic acid. Periodic acid is formed by the action of iodine upon an aque- ous solution of perchloric acid : 2 HC10 4 + 4 H 2 + I 2 = Cl a + 2 (HIO 4 . 2 H 2 O). Periodic acid has the composition corresponding to the- formula HIO 4 2 H 2 O, or as it is often written, H 5 IO 6 . The acid of the formula HIO 4 has never been obtained. Periodic acid forms colorless, transparent, deliquescent, prismatic crystals that melt at 133, while at 140 they are entirely decomposed, forming water, oxygen, and iodine pentoxide, thus : 2H 6 I0 6 = 5H 2 0-r0 2 + I 2 5 . Periodic acid is a strong oxidizing agent. Periodates are gener- ally difficultly soluble in water. Sodium periodate may readily be prepared by the interaction of chlorine, sodium hydroxide, and sodium iodate, thus : C1 2 + 3 NaOH + NaIO s = 2 NaCl + Na 2 H 3 IO 6 . THE HALOGENS 123 By heating barium iodate, the periodate of barium may be ob tained, thus : 5 Ba(I0 3 ) 2 = Ba,(IO.), + 4 1 2 + 9 O 2 . It will thus be seen that while periodic acid and the periodates are analogous to perchloric acid and the perchlorates, the fact that periodic acid has the composition HIO 4 2 H 2 O or H 5 TO 6 leads to a more complicated series of salts than we have in case of the perchlorates. Compounds of the Halogens with Each Other. By passing chlorine over iodine, a dark reddish brown liquid not unlike bromine in appearance is formed. It is very volatile and has an exceedingly pungent odor. It is about 8 times as heavy as water and boils at about 101, during which process it suffers partial decomposition. Its composition corresponds to the formula IC1 ; it is iodine monochloride. Two modifications of this compound have been described, the one melting at 24.7, and the other at 13.9. Iodine monochloride does not turn starch paste blue. By contact with water it is decomposed : 3 H 2 O + 5 IC1 = HI0 3 + 5 HC1 + 2 I 2 . When iodine is treated with chlorine in excess, or when iodine monochloride is further treated with chlorine, yellow, needle-like crystals are formed having the composition ICl g . They are iodine trichloride. They may be purified by sublima- tion at ordinary temperatures. On heating, they decompose into chlorine and iodine monochloride, but on cooling the tri- chloride forms again, thus : Water dissolves iodine trichloride. The solution has great germicidal power and is consequently used as an antiseptic. With bromine, iodine forms a crystalline compound, iodine monobromide, of the composition IBr. It has chemical proper- ties similar to those of iodine monochloride. It melts at 36. A compound of iodine with fluorine, iodine pentafluoride IF 6 , is also known. It is formed by direct union of the elements. It is a colorless liquid which boils at 97 and solidifies at 8. On heating it to 400, it suffers decomposition. Water decom poses it into hydrofluoric and iodic acids. 124 OUTLINES OF CHEMISTRY General Relations of the Halogens to One Another. Fluorine, chlorine, bromine, and iodine increase in atomic weight in the order named. With increasing atomic weight their melting points and boiling points rise, their specific gravities increase, and their color becomes more intense. Thus, with increasing atomic weight, we have here an increase in the degree of con- densation of matter, as it were. While the physical properties thus show a regular change with increasing atomic weight, the chemical properties also exhibit regularity of change. So the affinity for hydrogen is greatest in the case of fluorine, and least in the case of iodine. The general chemical activity of the halogens diminishes as the atomic weight increases. So fluorine is by far the most active element of the group, and iodine the least active. However, for oxygen iodine has a much greater affinity than fluorine, which unites with oxygen neither directly nor indirectly. In- deed, in case of the oxygen compounds of the halogens the sta- bility increases with the atomic weight of the halogen, being greatest in the iodine compounds. It is interesting to note that the atomic weight of bromine, 79.92, is approximately equal to one half the sum of the atomic weights of chlorine, 35.46, and iodine, 126.92. We shall meet more such groups of three elements in which a similar relation holds. The consideration of the relations between the atomic weights of the elements and their physical and chemical proper- ties has led to a classification of the elements known as the periodic system, which will be considered when more of the elements have been studied. REVIEW QUESTIONS 1. Name the halogens. Why are they so called? Why are they grouped together? 2. How do the properties of the halogens vary with their atomic weights ? 3. What is the weight of 22.38 liters of each of the halogens in the gaseous state under standard conditions ? 4. What is the action of chlorine upon a cold solution of potassium hydroxide? Upon a hot concentrated solution of the latter? Write the appropriate equations. Write similar equations expressing the action of bromine and iodine respectively on cold and hot solutions of sodium hydroxide. THE HALOGENS 125 5. What is bleaching powder? How is it prepared? Write the equation. Upon what properties of bleaching powder do its uses depend? 6. Measured under standard conditions, how many liters of hydro- bromic acid gas could be prepared from 8 liters of hydrogen ? How much would the hydrobromic acid gas weigh? 7. Give two general methods for preparing chlorine, bromine and iodine. What difficulty would be met in preparing fluorine by these methods ? 8. Write the equation expressing the action of manganese dioxide on a mixture of common salt and sulphuric acid. What actions would take place if potassium bromide and sodium iodide, respectively, were substituted for the common salt ? Write the. equation in each case. 9. Name three products that are formed when chlorine acts upon water. How do the other halogens act upon water? 10. Compare the action of the four halogens upon hydrogen. Write the equations. 11. How is hydrochloric acid commonly prepared? Why cannot hydrobromic and hydriodic acids be prepared similarly ? How can these three hydrohalogen acids be prepared from the phosphorus halides? Write the appropriate equation in each case. 12. Complete the equation in each of the following cases, if an action occurs : C1 2 + NaBr ; Br 2 + NaCl ; C1 2 + Nal ; Br 2 + Nal ; I 2 + KBr; I 2 + CaCl 2 ; NaCl + F 2 . 13. In which cases may the halogens replace one another in chemical compounds ? 14. Given the formulas of the following oxides : Pb0 2 , Fe 2 3 , Ag 2 0, CuO, Cr0 3 , write the formulas of the corresponding chlorides and fluorides. 15. Given the formulas of the following chlorides: NaCl, CrCl 3 , CaCl 2 , SnCU, PCls, write the formulas of the corresponding oxides. 16. What is an oxygen acid? Give the names and formulas of all of the oxygen acids of chlorine, and the corresponding sodium salts. Do the same for bromine and iodine. 17. Give the names of the following compounds : KI0 4 , Ca(C10 2 ) 2 , HI0 2 , KBr0 3 , NaCIO, Ca - Cl - CIO. 18. Write the equation expressing the action of hydrochloric acid upon hypochlorous acid. 19. What is the chief use of hydrofluoric acid? Illustrate the action by means of an equation. 20. What use is made of bromine in the arts? How is iodine used? In what liquids is the latter soluble? How may bromine and iodine be distinguished from each other? 21. What interesting relation exists between the atomic weights of chlorine, bromine, and iodine ? 22. What is thyroiodme? 126 OUTLINES OF CHEMISTRY 23. Write the equation expressing the action of each of the following upon hydrogen sulphide : chlorine, bromine, iodine. 24. What is iodine trichloride ? How may it be formed ? Write the equation. 25. Give an illustration showing that hydriodic acid is a powerful reducing agent. Write the equation. CHAPTER IX ACIDS, BASES, SALTS, HYDROLYSIS, MASS ACTION, AND CHEMICAL EQUILIBRIUM Acids. In connection with the study of oxygen it was found that this element readily unites with non-metals like phos- phorus, sulphur, and carbon, forming oxides which, when dissolved in water, have a sour taste and redden blue litmus. These oxides are consequently called acidic oxides or acid-form- ing oxides, for with water they form acids. Thus when sulphur burns in oxygen sulphur dioxide results : S + O 2 = SO 2 . When conducted into water, sulphur dioxide unites with the water, forming sulphurous acid H 2 SO 3 , thus: SO 2 + H 2 O = H 2 SO 3 . It is possible to form a similar acid of higher oxygen content. By passing sulphur dioxide mixed with oxygen over red-hot, finely divided platinum, a higher oxide of sulphur, namely sulphur trioxide SO 3 , is formed. This is a white crystalline solid which greedily unites with water, forming sulphuric acid H 2 SO 4 . The changes may be expressed as follows : 2SO 2 + O 2 = 2 SO 8 . S0 3 + H 2 = H 2 S0 4 . Similarly, when phosphorus is burned in oxygen phosphorus pentoxide P 2 O 5 is formed, which readily unites with water, forming phosphoric acid H 3 PO 4 : P 2 6 + 3 H 2 = 2 H 3 P0 4 . Again, when carbon is burned in oxygen, carbon dioxide CO 2 is produced, which, when dissolved in water, forms carbonic acid H 2 CO 3 , thus : C0 2 + H 2 = H 2 C0 3 . Carbonic acid H 2 CO 3 has not been isolated ; it exists merely in solution. The combination of water and carbon dioxide is very 127 128 OUTLINES OF CHEMISTRY \ weak, and carbonic acid is but slightly sour to the taste, red- dens blue litmus slowly, and acts in all respects much more feebly than the other acids just mentioned. It is a good exam- ple of a weak acid. It will be recalled that the fact that oxides like the above yield sour or acidic substances originally led to the idea that it is the oxygen that imparts these acidic characteristics to the com- pounds. Indeed, oxygen received its name in accordance with this notion. However, we have seen that the halogens form a series of compounds with hydrogen, namely, HF, HC1, HBr, and HI, which are all pronounced acids, in that they are sour to the taste, redden litmus, and attack metals like magnesium, zinc, and iron, evolving hydrogen and yielding compounds consisting of the halogen and the metal employed. These latter products are salts of the metal. When chlorine was discovered, it was looked upon as an oxide (as oxidized hydrochloric acid) ; for it was well known that chlorine is an acid-forming substance, and consequently it was thought that it must contain oxygen, which was regarded as the essential element in every acid. In fact, it was not till iodine was discovered that the correct view of chlorine as an element was really established. For this reason the discovery of iodine and the proof that it is elementary in character was of great importance in the development of the idea of the real nature of an acid. Thus, from the study of the hydrohalogens, which are all pronounced acids, came the true notion that hydrogen, and not oxygen, is the essential constituent of every acid. An acid is a compound containing hydrogen which may be replaced by a metal, the product formed being a salt. Acids commonly have sour taste and redden blue litmus ; but we shall learn of acids that are so weak that they do neither of these things ; yet they are nevertheless acids, because they contain hydrogen which may be replaced by a metal, the product formed being a salt. While it is thus true that there are very pronounced acids that contain no oxygen, still it must be stated that after all by far the great majority of acids do contain oxygen, and often it is by the addition of the latter element that the acidic character is produced. Since phosphoric acid H 3 PO 4 , sulphuric acid H 2 SO 4 , sulphur- ous acid H 2 SO 3 , and carbonic acid H 2 CO 3 , may be formed by ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 129 the addition of water to the oxides P 2 O 5 , SO 3 , SO 2 , and respectively, in accordance with the equations already givei? above, these oxides are called the anhydrides of the respective acids. When an acid acts on a metal like zinc or magnesium, hydro- gen is evolved and a salt is formed, thus : Zn + H 2 SO 4 = ZnSO 4 + H 2 . Mg + 2 HC1 = MgCl 2 + H 2 . A salt is thus one of the products of the interaction of an acid and a metal. It is possible to form salts by other means, how- ever, as will be shown below. Bases. We have seen in connection with the study of hydrogen that when this element is liberated by the action of sodium or potassium on water, caustic soda or caustic potash results, thus : H 2 + Na = NaOH + H. H 2 O + K = KOH + H. The solutions of sodium hydroxide and potassium hydroxide are alkaline in character. They turn red litmus blue, and when treated with an acid they become neutral ; that is, they do not affect either red or blue litmus. When these hydroxides are treated with an acid, they are said to be neutralized. In this process the acid is, of course, also neutralized. The interaction of an alkaline hydroxide with an acid is a mutual act, resulting in the neutralization of both compounds. On evaporating the neutral solution it is found that a salt has been formed. The reaction or neutralizing sodium hydroxide with hydrochloric acid may be expressed thus : NaOH + HC1 = NaCl + H 2 O. When, for example, potassium hydroxide is neutralized with sulphuric acid, the following change takes place : 2 KOH + H 2 SO 4 = K 2 8O 4 + 2 H 2 O Hydroxides of metals which are thus capable of reacting with acids, forming salts and water, are called basic hydroxides or bases. Hence a base is an "hydroxide or oxide of a metal which will react with an acid, forming (1) a neutral substance called a salt, and (2) water. 130 OUTLINES OF CHEMISTRY Elements which are thus capable of uniting with the hydroxyl radical OH to form bases are called base-forming elements, while those that form acids by union with hydrogen are called acid- forming elements. Hydroxides of some of the elements, how ever, are capable of acting as bases toward more acidic hydrox- ides, and as acids toward hydroxides that are more basic than themselves. This will be more evident as we proceed in our considerations. Salts. From what has been stated in connection with the consideration of acids and bases, the nature of salts is already sufficiently characterized. Thus, a salt is a neutral compound resulting as a product of the interaction of an acid with a base ; or a salt is a neutral compound which is formed when the hydrogen of an acid is replaced by a metal. So it appears that a salt may be formed in the following ways : (1) By the neutralization of an acid with a base, as, for exam- ple, sodium sulphate Na 2 SO 4 is formed when sodium hydroxide and sulphuric acid act on each other : 2 NaOH + H 2 SO 4 = Na 2 SO 4 + 2 H 2 O. (2) By the action of a metal on an acid, thus : 2 Na + H 2 S0 4 = Na 2 S0 4 + H a . (3) It is possible, however, to form a salt by the direct action of two oxides, a base-forming oxide or basic oxide, and an acid- forming oxide or acidic oxide, on each other, thus : (4) It is also possible to form a salt by direct union of a base- forming element with an acid-forming element, thus : Na + Cl = NaCl. In all cases,' however, one may think of the salt as derived from some acid whose hydrogen has been replaced by the metal. So, though in (3) the sulphate of sodium was made by the union of sodium oxide and sulphur trioxide, that is sulphuric anhydride, one may think of the product Na 2 SO 4 as derived from H 2 SO 4 in which the hydrogen is replaced by sodium. Likewise, all sulphates may be regarded as similarly derived from sulphuric acid. The latter may in turn be looked upon as hydrogen sulphate, i.e. a salt in which hydrogen is the basic ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 131 element; indeed, any acid may be regarded as a salt of hydrogen, Thus, hydrochloric acid HC1 is the chloride of hydrogen. Sodium chloride NaCl, whether made by the direct union of the elements as in (4) above, or by the action of sodium hydroxide upon hydrogen chloride, may yet be regarded as derived from HC1 in which the hydrogen is replaced by sodium. Similarly, all fluorides may be regarded as derived from HF, the hydrogen of which has been replaced by the metal. All iodides may similarly be considered as derived from HI, all chlorates from HC1O 3 , etc. Older View of the Process of Salt Formation. Salts were formerly considered as the result of the union of a basic oxide with an acidic oxide. Thus, sodium sulphate was regarded as sodium oxide Na 2 O plus sulphur tri oxide SO 3 , and the formula of the salt was written Na 2 O SO 3 . Similarly, calcium carbon- ate was regarded as made up of calcium oxide CaO, plus carbon dioxide CO 2 , and the formula of calcium carbonate was conse- quently CaO CO 2 . Again, ferrous sulphate was considered as ferrous oxide FeO plus sulphur trioxide, thus : FeO SO 3 . This was the dualistic way of writing which was in vogue dur- ing the former half of the last century, and it is not to be denied that it had many advantages. So these formulse indicated at once that the salts can be formed by direct union of the acidic and basic oxides ; and since it is true that many of these salts when strongly heated decompose into the basic and acidic oxides, the formulse also in a simple way represented this fact. For instance, on heating, calcium carbonate yields calcium oxide and carbon dioxide : CaCO 3 =CaO + CO 2 ; and ferrous sulphate decomposes thus : According to the older view the process of solution of a metal like zinc or iron in sulphuric acid consisted of two steps. First, when the zinc was introduced into dilute sulphuric acid, the metal was oxidized to zinc oxide ZnO, hydrogen being simultaneously liberated from the water; and second, zinc oxide would then combine with the sulphur trioxide, forming zinc sulphate ZnO SO 3 . Sulphuric acid was regarded as SO 3 dissolved in water. This dualistic way of writing the formulse 132 OUTLINES OF CHEMISTRY of salts was strongly defended by Berzelius, and it was only through the development of the study o the compounds of carbon and of electrochemistry that the present method of ex- pressing the formulse was finally adopted. In the study of some of the complicated silicates, however, the old way of writing is still frequently employed -with distinct advantages, as will be seen when the compounds of silicon are discussed. The dualistic formulae of Berzelius were, moreover, also based upon electrochemical ideas. Acid- and Base-forming Elements. In general, the acid- forming elements are the non-metals, and the base-forming ele- ments are the metals. Oxygen, sulphur, nitrogen, phosphorus, carbon, the halogens, etc., are acid-forming elements; and potassium, sodium, magnesium, zinc, lead, copper, etc., are base- forming elements. However, as stated above, an acidic element may act as a basic element toward a still more acidic element ; and a basic element may act as an acidic element toward a still more basic element. Of this we have already had illustrations. Thus while zinc acts as a base in zinc sulphate ZnSO 4 , in which compound sulphur and oxygen form the acid radical SO 4 , in potassium zincate K 2 ZnO 2 , formed thus, Zn(OH) 2 + 2 KOH = K 2 ZnO 2 + 2 H 2 O, zinc is a part of the acid radical ZnO 2 . So toward the acid group SO 4 zinc acts as a base, while toward the strongly basic potassium the zinc forms a part of the acidic group ZnO 2 . Again, in sodium iodide Nal, sodium is the basic and iodine the acidic element ; whereas in iodine chloride IC1, the iodine plays the role of base toward the more acidic chlorine. Further, in phos- phorus trichloride PC1 3 , phosphorus is the basic and chlorine the acidic element, whereas in phosphates, like sodium metaphos- phate NaPO 3 , phosphorus plays the role of an acidic element. Additional examples will readily occur to the reader, and many others will be met in our further study. The distinction between acid- and base-forming elements is thus not a sharp one ; nevertheless, from what has been stated, the difference between acids and bases can generally be made without difficulty. Other Views of Solutions of Acids, Bases, and Salts. In recent years the attempt has been made to define acids, bases, and salts ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 133 on the basis of the behavior of their solutions toward the elec- tric current. In this attempt the boiling and freezing points of dilute solutions have also been a prime consideration. This study has further led to another way of regarding the act of neutralization and the resulting salt solutions. A considera- tion of these views of acids, bases, and salts in solution will be taken up later in connection with the subjects of solutions and electrolysis. Basicity of Acids; Acid Salts. An acid which contains one replaceable hydrogen atom in its molecule is called a monobasic acid; one that contains two replaceable hydrogen atoms is called a dibasic acid ; etc. There are also tribasic, tetrabasic, and pent a- basic acids. For example, HC1, HBr, HI, HC1O, HBrO, HIO, HC1O 3 , HBrO 3 , HIO 3 , HC1O 4 , HNO 3 (nitric acid), are all monobasic acids. With a univalent basic element like potas- sium or sodium, for instance, each of these monobasic acids forms but one salt, thus : KC1, Nal, KBrO, KC1O 4 , NaNO 8 . Sulphuric acid H 2 SO 4 and carbonic acid H 2 CO 8 are dibasic acids, since they contain two atoms of replaceable hydrogen per molecule. In the neutralization of a dibasic acid the opera- tion may take place in two steps, thus : (1) H 2 SO 4 (2) NaHS0 4 +NaOH = Na 2 S0 4 + H 2 O. (1) H 2 CO 3 +KOH =KHCO 3 + H 2 O. (2) KHC0 3 +KOH =K 2 CO 3 + H 2 O. The salt NaHSO 4 still contains replaceable hydrogen, i.e. acidic hydrogen. It is consequently an acid salt, as contrasted with Na 2 SO 4 , in which all the hydrogen has been replaced. The latter salt is a neutral or normal salt. Acid sodium sulphate NaHSO 4 is also called sodium bisulphate, for it contains twice as much acid radical per same amount of sodium as does the normal salt Na 2 SO 4 , which is also called bisodium sulphate, as well as simply sodium sulphate. Similarly, KHCO 3 is potas- sium bicarbonate, and K 2 CO 3 is bipotassium carbonate, or simply potassium carbonate. Thus an acid salt is one which still contains hydrogen that is replaceable by a metal. 134 OUTLINES OF CHEMISTRY The ability of an acid to form acid salts shows that the acid is not monobasic. Thus, for instance, the fact that hydro fluoric acid forms acid salts like KHF 2 and NaHF 2 would argue in favor of the view that the acid is dibasic in character and has the formula H 2 F 2 , rather than the simple formula HF. In phosphoric acid H 3 PO 4 and periodic acid H 6 TO 6 we have an example of a tribasic and a pentabasic acid, respectively. The hydrogen atoms of these acids can be replaced step by step, thus forming a series of salts which grow less and less acid in character. For instance, in case of phosphoric acid, we may form the three salts KH 2 PO 4 , K 2 HPO 4 , and K 3 PO 4 , which are called primary, secondary, and tertiary potassium phosphate, respectively. We may also call these salts monopotassium phos- phate, dipotassium phosphate, and tripotassium phosphate. The commonest salt of the three is dipotassium phosphate, and so it is generally referred to simply as potassium phosphate. In the case of the pentabasic periodic acid, we have a still greater range of possibility of formation of acid salts ; indeed, the greater the number of replaceable hydrogen atoms the molecule of an acid contains, i.e. the greater its so-called basicity, the large', is the number of acid salts that it is able to form. Acidity of Bases. A base like KOH or NaOH, which is an hydroxide of a univalent metal, is called a monoacid base, for it is capable of reacting with an acid, and thus forming one mole- cule of water and replacing one atom of acidic hydrogen. A base like Ca(OH) 2 is called a diacid base, for it is capable of reacting with an acid, forming two molecules of water and a salt in which the bivalent metal replaces two atoms of acidic hydro- gen. Similarly, we may have triacid, tetraacid, and pentaacid bases like antimonous hydroxide Sb(OH) 3 , stannic hydroxide Sn(OH) 4 , and antimonic hydroxide Sb(OH) 5 . The so-called acidity of a base consequently is determined by the valence of the metal of the base ; or what really comes to the same thing, by the number of basic hydroxyl groups the molecule of the base contains. Basic Salts. Just as in the case of acids that contain more than one hydrogen atom to the molecule it is possible to secure acid salts by replacing these hydrogen atoms step by step, so in the case of bases that contain more than one basic OH group it is possible to neutralize those groups step by step by means of an acid, thus forming a series of basic salts. For example, ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 135 (1) Sb(OH) 3 + HC1 = Sb(OH) 2 Cl + H 2 O. (2) Sb(OH) 2 Cl + HC1 = Sb(OH)Cl 2 + H 2 O. (3) Sb(OH)Cl 2 + HCl = SbCl 3 + H 2 0. Thus Sb(OH) 2 Cl and Sb(OH)Cl 2 are basic salts, for they still contain an excess of base that is not yet neutralized. The compound Sb(OH) 2 Cl readily splits off water, thus : Sb(OH) 2 Cl = SbOCl + H 2 0. The salt SbOCl is called antimony oxy chloride; it is also clearly a basic salt, for it is still capable of further neutraliza- tion with an acid. So with hydrochloric acid it undergoes the following change: SbOCl + 2 HC1 = SbCl 3 + H 2 O. A basic salt is one that contains an excess of base, which may still be neutralized with an acid. Basic salts are frequently met with, and they often split off water, as was exemplified above. Normal Salts. Normal salts are those that contain neither hydrogen that is replaceable by a metal nor an excess of base that may still be neutralized by an acid. Thus NaCl, K 2 SO 4 , CaCO 3 , are normal salts of hydrochloric, sulphuric, and carbonic acids, respectively. From what has been said it is evident that a monobasic acid can form only normal salts with monoacid bases ; whereas with polyacid bases it may form basic salts as well as normal salts. Again, polybasic acids may form either normal salts, acid salts, or basic salts with polyacid bases; whereas with monoacid bases they can form only acid salts and normal salts. Acidimetry and Alkalimetry. The fact that it always takes a definite amount of acid and a definite amount of base to exactly neutralize each other is. used in the quantitative esti- mation of acids and markedly alkaline bases in processes that are called acidimetry and alkalimetry. A given volume of a solution of an acid of known strength will always neutralize a perfectly definite volume of a given solution of an alkali. If we place the acid of known strength in the burette A (Fig. 33) and an alkaline solution of unknown strength, e.g. of sodium hydroxide, in the burette B, then run out say 20 cc. into the dish D and find that it is necessary to add 23.6 cc. of the acid in A to just make the solution in the dish D neutral to litmus, 136 OUTLINES OF CHEMISTRY it is possible to compute the strength of the sodium hydroxide solution from the data at hand. Any solution of known strength is called a standard solution. In working with acid solutions of known strength normal solutions are frequently used. A normal solution of an acid is one which contains 1.008 grams of replaceable hydrogen per liter of solution. Thus a normal solution of hydrochloric acid would contain 36.468 grams of pure HC1 per liter, for in this amount there are 1.008 grams of replaceable hydrogen. In the case of hydrobromic acid, a liter of normal solution would contain 80.928 grams of HBr; in the case of sulphuric acid a liter of normal solution would contain 49.043 grams of H a SO 4 , etc. Sometimes solutions of one half, one tenth, or one twentieth, etc., of the strength of normal solutions are em- ployed; these contain the cor- responding amounts of re- placeable hydrogen per liter. Solutions of an acid may be prepared which contain some multiple of 1.008 grams of replaceable hydrogen or some aliquot part thereof per liter. These are then called twice normal, normal, half normal, fiftieth normal, etc., as the case may be. A normal solution of an alkali is one that will just neutralize a normal solution of an acid volume for volume ; consequently a normal solution of an alkali is a solution that contains the chemical equivalent of 1.008 grams of replaceable hydrogen per liter. Al- kaline solutions may be made up as normal solutions or as some multiple or fractional part of normal, just as in the case FIG. 33. ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 137 D the acid solutions. So, for instance, a normal solution of sodium hydroxide NaOH contains 40.008 grams of pure NaOH per liter, and a tenth normal solution contains one tenth as much per liter. The process of exactly neutralizing an acid solution with an alkaline solution to ascertain the strength of one of them in terms of the other, as illustrated in. connection with Fig. 33, is called titration. If, in the instance cited above, the acid solu- tion was normal, then the 20 cc. of NaOH solution required would just be equal to 23.6 cc. of normal solution. In other words, the NaOH solution would be ' normal, and contain 23 6 x 40.00 grams of pure NaOH, or 47.20 grams per liter. Z\j Indicators. Litmus may serve to indicate the acidity or alkalinity of a solution. There are, however, many other coloring matters that change their hue on being treated with an acid and an alkali in succession. Such substances may con- sequently also serve as indicators. So phenolphthalein is color- less in acid solution, pink in a faintly alkaline solution, and a beautiful purplish red in strongly alkaline solution. Methyl orange is red when acid, straw yellow when alkaline, and orange colored when neutral. Congo red is red when alkaline and blue when acid ; that is, it acts in just the opposite way that litmus does. Turmeric paper, that is paper soaked in a decoc- tion of turmeric root, is yellow when neutral or acid, but turns brown when moistened with an alkaline solution. There are still other indicators in use, but those mentioned are the ones commonly employed in the laboratory. Hydrolysis. When phosphorus trichloride is brought into contact with water, violent action ensues, heat is liberated, and phosphorous and hydrochloric acids are formed, thus : PC1 3 + 3 H 2 O = P(OH) 3 + 3 HC1. The action is complete and irreversible. No such change takes place, for example, when sodium chloride NaCl is treated with water. The solution in this case remains quite neutral, and the salt may be recovered by evaporating off the water. In phosphorus trichloride we have a compound which in contact with water readily passes over into the much more stable com- 138 OUTLINES OF CHEMISTRY pounds hydrochloric and phosphorous acids. In the forma- tion of hydrochloric acid, the great affinity of hydrogen foi chlorine comes into play, and in the formation of phosphorous acid the strong affinity of phosphorus for oxygen exerts itself. We thus see that all three chlorine atoms in PC1 3 are exchanged for OH groups taken from three water molecules, the remain- ing hydrogen atoms of which unite with the chlorine atoms to form hydrochloric acid. When a substance is thus decomposed by water, the process is termed hydrolytic decomposition, or hydrolysis. Very many salts suffer hydrolysis in water to a slight extent, others are not decomposed by water, and still others are com- pletely hydrolyzed. Thus we have seen that PC1 3 , which is a salt of an extremely weak basic element, phosphorus, with a very strong acidic element, chlorine, is completely decomposed by hydrolysis. On the other hand, sodium chloride, a salt of the strongly basic sodium with the strongly acidic chlorine, is not hydrolyzed. In general, salts of weak bases with strong acids are more or less hydrolyzed when brought into contact with water. The same is true of salts of strong bases with very weak acids. So cupric sulphate CuSO 4 , ferric chloride FeCl 3 , and in general all the ordinary salts of the heavy metals, are somewhat hydrolytically decomposed when dissolved in water. This is made evident, for instance, by the fact that all these solutions react acid toward litmus. On the other hand, salts like sodium carbonate Na 2 CO 3 and borax Na 2 B 4 O 7 , which contain a very weak acid radical united to the strongly basic sodium, are also hydrolyzed to some extent. This is evidenced by the fact that their solutions react alkaline toward litmus. Even sodium bicarbonate NaHCO 3 is thus hydrolyzed, and its solutions react alkaline toward litmus, though the salt still contains hydrogen that is replaceable by a metal and consequently is an acid salt. Thus it is that the normal salts need not necessarily yield solutions that are neutral toward indicators, for many of them are hydrolyzed. Indeed, neutrality toward indicators is met only in dealing with solutions of normal salts of strong acids with strong bases. Thus, the chlorides, sulphates, nitrates, and bromides of sodium, potassium, calcium, and magnesium in solution are neutral toward indica- tors ; whereas the corresponding salts of iron, copper, lead, and mercury have acid reactions in solution, and the carbonates, silicates, and borates of sodium and potassium yield alkaline so- ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 130 tutions. From this it is evident that the processes of acidimetry and alkalimetry described above can be used only in estimat- ing the strength of solutions of fairly pronounced acids and alkalies. If now we take cupric chloride CuCl^ and dissolve it in water, we find that the solution is distinctly acid toward litmus and other indicators, showing that the salt has suffered decomposi- tion to a certain degree. The hydrolysis which has taken place has resulted in the liberation of a slight amount of hydrochloric acid. One may indicate the change, at least in part, thus : CuCI a + H a O^Cu(OH)Cl+ HC1 ; meaning that only a small percentage of the salt is thus hydro- lyzed in solutions that contain say ten per cent or more of the salt. Such solutions remain clear, the basic cupric chloride forming with the cupric chloride CuCl 2 a compound which consists of a soluble, though slightly basic cupric chloride. On dilution of the solution of cupric chloride, the hydrolysis proceeds further and a basic cupric chloride finally forms, which is richer in base and poorer in chlorine than that in the stronger solutions. This basic salt is difficultly soluble and gradually separates out in form of a precipitate. The more dilute the solution, then, the more does the reaction proceed in the direction indicated by the upper arrow in the equation. On the other hand, on con- centrating the solution the reaction proceeds from right to left; i.e. it reverses. We have here then a case of reversible hydrolysis. Mass Action; Chemical Equilibrium. Thus we see, in the instance just mentioned, that the more water there is added to the cupric chloride solution, the greater is the extent of the hydrolysis. The water consequently influences the process to proceed from left to right (see the equation) ; but it is to be borne in mind that it is the mass of water present relative to the amount of cupric chloride in the solution that really deter- mines the direction of the reaction. The amount of matter contained in unit of volume is commonly called the concentration. In the case of a solution, the concen- tration is the amount of dissolved substance contained in unit of volume of the solution. Very commonly the concentration 140 OUTLINES OF CHEMISTRY of a solution is expressed by stating the number of gram- molecules contained in one liter of the solution. By a gram-molecule is meant the molecular weight in grams. Thus a gram-molecule of hydrochloric acid is 36.46 grams HC1 ; a gram-molecule of potassium iodide, 166.02 grams KI. We may state the facts elucidated in the case of the hydroly- sis of cupric chloride by saying that the extent of the hy- drolysis is determined by the relative concentrations of the cupric chloride and the water in the solution ; i.e. by the relative masses of the substances that are acting on each other. The amount of undecomposed cupric chloride, water, basic cupric chloride, and hydrochloric acid which the solution contains at any particular temperature and concentration of the cupric chloride solution is perfectly definite ; and the amounts of these four ingredients are said to be in chemical equilibrium with one another. The equilibrium of this system may be disturbed by changing the concentration of any one of the four ingredients that make up the solution. Thus to abstract water from the solution causes the action to proceed in the direction of the lower arrow, CuCl 2 + H 2 O ^ Cu(OH)Cl + HC1 ; while to add water causes the action to go in the direction of the upper arrow. To abstract cupric chloride from the system causes the equilibrium to be displaced in the direction of the lower arrow, for this practically amounts to the same thing as adding more water relatively. Addition of cupric chloride causes the opposite effect. Addition of hydrochloric acid to the system causes the action to proceed in the direction of the lower arrow ; abstracting hydrochloric acid causes the equilibrium to be displaced in the direction of the upper arrow. Addition of basic cupric chloride to the system causes the equi- librium to be changed in the direction of the lower arrow, while removal of basic cupric chloride from the system causes the reaction to proceed in the direction of the upper arrow. It is obvious that if either the hydrochloric acid or the basic cupric chloride were taken from the system as fast as formed, the re- action would go to completeness from left to right and with increased rapidity. What has been thus presented is really a special case of a ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM general law, termed the law of mass action, which may be stated thus : The speed or rate of any chemical change is proportional to the active mass, that is, the molecular concentration of each substance engaged in the reaction. This is universal and holds for all chemical changes, whether they are reversible or not. In case of reversible reactions, the law holds for the change from right to left as well as from left to right, and hence the final chemical equilibrium reached is also determined by the law of mass action. One can best comprehend this by thinking of the equilibrium as reached when the rate of speed of the for- ward action just equals that of the reverse action. Chemical equilibrium is commonly regarded as dynamic rather than static in character. Additional Illustrations of Chemical Equilibrium and the Oper- ation of the Law of Mass Action. In the first chapter it was stated that the factors which determine whether a chemical change will go on or not are : (1) the right substances must be brought into contact, i.e. chemical attraction or chemical affinity must exist between the substances that are to react ; (2) the temperature must be properly chosen ; (3) the pressure is of consequence, particularly when a gas enters into the change ; (4) the concentrations of the active substances must be con- sidered. All of these factors determine not only whether the change will proceed at all or not, but they also influence the rate with which the action proceeds and consequently affect the final equilibrium reached. Now it is clear that it is with factor (4), above mentioned, that the law of mass action is concerned. There are many reactions which are, so far as we know, irreversible ; that is, they go to completion in one direction. We have already seen that the hydrolysis of phosphorus chlo- ride is of this class. The combustion of calcium, magnesium, or sodium in oxygen, the decomposition of potassium chlorate into potassium chloride and oxygen, the neutralization of po- tassium hydroxide by hydrochloric acid, the burning of sugar to water and carbon dioxide, are further examples of this kind. In these reactions, the chemical affinity factor, namely (1) above, is really the determining one ; i.e. its influence overshadows ali the other factors very greatly, and so the action goes on in one 142 OUTLINES OF CHEMISTRY direction to completion and is irreversible. The cases of irre versible reactions are after all then fairly clearly distinguished, for as a rule they belong to one of two categories ; namely : (1) they represent the formation of very stable compounds directly from the elements, in which processes powerful affinities come into play ; or (2) they represent the decomposition of relatively unstable or complicated compounds into simpler and much stabler ones. The burning of barium to barium oxide is a typical illustra- tion of the first class; the decomposition of sugar or nitroglycer- ine by heat illustrates the second class. In speaking of irre- versible reactions, it must be borne in mind that the term does not mean that the original substances taken cannot be got back by roundabout means. So while the burning of magne- sium to magnesium oxide is an irreversible reaction, it is yet possible to get back the metallic magnesium and also the oxygen that it contains. In this sense, of course, all chemical reactions could be reversed, for matter cannot be destroyed, but simply transformed. What we mean by an irreversible reaction, in the sense in which the term is here used, is a reaction that cannot be re- versed entirely or in part by simply altering the temperature, pressure, or the concentration of the substances concerned in the reaction. In the irreversible reactions the factors of temperature, pressure, and concentration cannot reverse the process. But in many chemical changes the affinities that come into play in fix- ing the direction in which the change will go on are so well balanced that changes of temperature, pressure, or concentration suffice as determining factors in altering the direction the reaction takes. Such reactions are consequently reversible. This class of reactions is very large indeed. It is therefore evident that at constant temperature and pressure the effect of concentration is of vast importance in case of reversible reactions, for in these it de- termines the direction of the change and consequently the final equilibrium. On the other hand, in the irreversible re- actions the concentration changes can only affect the rate of the reaction. Thus, in the burning of magnesium ribbon the final product is MgO, and whether the action proceeds in oxygen at atmospheric pressure, or in compressed oxygen, only affects the rate of the combustion, not the character or the ACIDS, BASES, sAi/rs, CHEMICAL EQUILIBRIUM 143 amount of the final product. But when, for instance, chlorine acts on water in diffused light, we have a case of a reversible reaction, thus: HO + C1HOC1 + HC1. All four ingredients are finally in equilibrium at any given temperature and pressure. By increasing either the relative amount of water or chlorine the change progresses somewhat more from left to right ; the reverse happens by increasing the relative concentration of either the hydrochloric acid or hypo- chlorous acid. By diminishing the concentration of either the water or chlorine or both, the reaction proceeds from right to left. Decreasing the concentration of either the hypochlorous acid or hydrochloric acid or both causes the change to proceed from left to right. If we were to abstract say the hypochlo- rous acid as fast as it forms, the reaction would complete itself from left to right. Now, in the sunlight hypochlorous acid undergoes decomposition, thus: 2 HOC1 = O 2 + 2 HC1. Therefore as the oxygen escapes and only hydrochloric acid remains in the solution, we have (when chlorine acts on water in sunlight) a reaction which goes on to completion. This reaction is complete because of the removal of one of the in- gredients ; namely, the hypochlorous acid. Again, when sulphuric acid acts on common salt in moder- ately dilute solution, 10 to 20 per cent for instance, an equi- librium is established which may be expressed thus : NaCl + H 2 S0 4 ^NaHSO 4 + HCL The action is reversible, for it is possible to displace the equilib- rium in the one direction or the other by changing the concen- tration of the substances that enter into the change. Now, when concentrated sulphuric acid is poured on sodium chloride and the mass becomes warm, the reaction will complete itself from left to right; for the hydrochloric acid is volatile, and under the conditions of the experiment it can escape and so get out of the field of action. This does not necessarily mean that the sulphuric acid is stronger than the hydrochloric acid and so drives the latter out, as was formerly supposed. It will 144 OUTLINES OF CHEMISTRY be observed that the determining factor is rather the volatility of the hydrochloric acid, which takes it out of the reacting mass. Indeed, it is possible to displace the hydrochloric acid from common salt by boiling it with strong solutions of much weaker acids than hydrochloric acid, provided that the acid so employed is non-volatile. So, for example, it is possible to evolve hydrochloric acid from salt by employing boric acid, which, as we shall learn, is a very weak yet practically non- vola- tile acid. Whenever liquids act on liquids, or solids act on liquids, forming a product which is gaseous and so escapes, the reaction proceeds practically to completion. The same is true whenever in such cases a solid forms which is insoluble, i.e. is practically not acted upon, and so is thrown out of the field of action. Thus, for instance, when sodium sulphate acts on barium chloride we have the following change taking place : BaCl 2 + Na 2 SO 4 = BaSO 4 + 2 NaCl. This goes practically to completion because the barium sulphate formed is very difficultly soluble, and nearly all of it drops out of the field of action as a precipitate. In the case of gases we frequently have instances of rever- sible changes. At red heat, water vapor partially decomposes into hydrogen and oxygen; at still higher temperatures, the reaction progresses further in the sense mentioned, whereas on cooling it again is reversed. The process of thus decomposing a substance on heating it is called dissociation. It was studied particularly by Henri Saint Claire Deville. We shall consider cases of the dissociation of gases more carefully later. Strength of Acids and Bases. The relative strengths of acids has been a favorite subject of study with chemists. By having, let us say, tenth normal solutions of hydrochloric, sulphuric, and acetic acids each separately act on a piece of zinc (the pieces being arranged so as to expose the same area of zinc to each acid) and estimating the volume of hydrogen liberated by each acid per minute, it is possible to compare the relative strengths of the acids. The apparatus for this purpose might be arranged as in Fig. 14. In each tube is placed a piece of zinc of the same size and shape. The whole apparatus is then filled with water, the same quantity being used in each case ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 145 The acids are then introduced in chemically equivalent amounts from above by means of the stopcocks, care being taken to admit no air through the cocks. The volumes of hydrogen evolved per minute may then readily be read. We should thus be estimating the strengths of these acids by their rate of action upon zinc. It is, of course, possible to use other characteristic activities of acids as a basis of estimating their strength. Similarly, strengths of alkalies might be compared by measuring the rate with which they transform a fat into soap (which see). REVIEW QUESTIONS 1. What is formed when each of the following compounds acts upon, water: S0 2 , P 2 5 , K 2 0, C0 2 , CaO, Na 2 0? Write the equation in each case and state what conclusion may be drawn as to the nature of (a) the metallic oxides, (b) the non-metallic oxides. 2. Define : acid, base, salt, neutralization, normal salt, acid salt, basic salt, anhydride, and give an example of each. 3. Mention four general methods of preparing a salt. Give an equa- tion for the preparation of calcium chloride by each method. 4. How many grams H 2 S0 4 will neutralize 250 grams KOH ? 5. Define : normal solution, indicator, hydrolysis. Give illustrations. 6. Explain the fact that a solution of NaHS0 4 reacts acid toward litmus while a solution of NaHC0 3 reacts alkaline. 7. State how an aqueous solution of each of the following reacts towards litmus and explain the action: A1C1 3 , K 2 C0 3 , FeCl 3 , KN0 3 , NH 4 C1, CuS0 4 , NaCl, Na 2 B 4 7 . In general what normal salts react neutral toward litmus? What normal salts react alkaline? What nor- mal salts show an acid reaction? 8. Given the following normal salts: Bi(N0 3 ) 3 ; SbCl 3 ; FeCl 3 ; write the formulas of two basic salts which each could form. 9. Write the formulas for the anhydrides of the oxy-acids of chlorine. 10. What is meant by the term "reversible reaction"? Illustrate. 11. In the following reaction, state what would be the effect of increas- ing (a) the amount of water, and (6) the amount of nitric acid in the solution : Bi(N0 3 ) 3 + 2 H 2 O^Bi(OH) 2 N0 3 + 2 HN0 3 . 12. If dry common salt is treated with hot, concentrated sulphuric acid, it is completely converted into hydrochloric acid and sodium sul- phate ; but, if a twenty per cent solution of common salt is treated with sulphuric acid, the reaction is incomplete. Explain these facts from the standpoint of chemical equilibrium. 13. How compare the strengths of hydrochloric, sulphuric, and acetic acids ? CHAPTER X NITROGEN, THE ATMOSPHERE, AND THE ELEMENTS OF THE HELIUM GROUP History and Occurrence of Nitrogen. In 1772 Dr. Rutherford, professor of botany at Edinburgh, found that when animals are confined in an air-tight space, the air they breathe becomes incapable of supporting combustion or respiration. After treat- ing such air with caustic potash solution to absorb the carbon dioxide, then called " fixed air," he showed that the remaining gas supported neither life nor combustion. A lighted candle thrust into the gas, for instance, was immediately extinguished. He called this residual gas " mephitic air." Priestley burned carbon in a confined volume of air, and then treated the latter with limewater ; thus the carbon dioxide formed during the com- bustion was absorbed, and a residual gas was obtained, which he called " phlogisticated air." He found that one fifth of the vol- ume of atmospheric air can thus be converted into " fixed air " and absorbed by caustic lime. But he did not regard the " phlo- gisticated " air he had prepared as a constituent of the atmos- phere. It was Scheele (1777) who first showed that there are two different gases in the air. Lavoisier was the first to con- sider mephitic or phlogisticated air as an element. He called it azote, because of its inability to support life. The name nitro- gen was given to the gas by Chaptal, because it forms an essen- tial constituent of niter or saltpeter. Cavendish showed that nitrogen obtained from air is essentially a simple body which is somewhat lighter than ordinary air ; and, indeed, till 1894 the residual gas thus prepared was regarded as pure nitrogen. Sir William Ramsay and Lord Rayleigh showed that the gas remain- ing after the oxygen and carbon dioxide have been removed from the air consists of 98.814 per cent nitrogen and 1.186 per cent other gases, which, unlike nitrogen, will not unite with oxygen or with red-hot magnesium. This notable observation led to the discovery of the new elements of the helium group. 140 NITROGEN, AIR, AND THE HELIUM GROUP 147 About 80 per cent of the volume of atmospheric air consists of nitrogen in the free state. In combination with carbon, hydrogen, and oxygen, nitrogen forms an essential constituent of the bodies of all plants and animals. It is found especially in the blood, muscles, nerves, seeds, and, in general, in all tissues that are concerned in movement or reproduction. When plants and animals die and their bodies decay, their nitrogen content passes over into simpler compounds, namely, ammonia, nitrites, and nitrates (which see). Thus it is that in all soils nitrogen is present in the form of nitrates and ammonium salts. It also occurs in all refuse matter of plant or animal origin, like barn- yard manure, guano, sewage, etc. In coal, which represents the remains of plants of the carboniferous age, nitrogen is found in combination with hydrogen, carbon, and oxygen. In minute quantities, nitrogen also occurs in granitic rocks, in meteoric iron, and in steel. In Chili saltpeter, consisting chiefly of sodium nitrate, nitrogen occurs in large quantities. Preparation and Properties of Nitrogen. Nitrogen which is approximately 99 per cent pure may be prepared from the air by removing the oxygen from the latter. This is generally accom- plished by heating in the air some elementary substance which will read- ily combine with oxygen, forming an oxide that is either a non-volatile solid or that can readily be removed by absorption in some liquid. Thus, when phosphorus is burned in a little dish resting on water under a bell jar (Fig. 84), phosphorus pentoxide is formed, which is a solid that is readily absorbed by water, forming phosphoric acid : P 4 + 5 O = 2 P 2 O 5 , and FIG. 34. Again, air may be passed over red-hot copper, when the latter unites with the oxygen, forming cupric oxide CuO, which is non-volatile, thus leaving the nitrogen. The oxygen may also be removed from the air by shaking the latter with an alkaline solution of pyrogallic acid, which readily absorbs oxygen, and 148 OUTLINES OF CHEMISTRY which is frequently used for this reason in gas analysis. Left in contact with moist yellow phosphorus, the air is also deprived of its oxygen even at room temperatures. This fact is often used in determining the amount of oxygen in a given sample of gas. By cooling air to 182 the oxygen liquefies, leaving the nitrogen in form of a gas. Pure nitrogen cannot very well be prepared from atmospheric air, for the gases of the helium group, with which it is always contaminated, are chemically very inert, and hence difficult to remove. Pure nitrogen is prepared from compounds in which it occurs. Thus, by treating ammonia NH 3 with chlorine, nitrogen and hydrochloric acid are formed, the latter uniting with some of the ammonia (which should be present in excess) to form ammonium chloride, which dissolves in water. The reactions may be expressed thus : 2NH 3 + 3C1 2 = 6HC1+N 2 . NH 3 + HC1 =NH 4 C1. The simplest way of preparing pure nitrogen consists of heating ammonium nitrite NH 4 NO 2 , either in pure form or in strong aqueous solution. The compound when thus treated decom- poses into water and nitrogen : NH 4 N0 2 =2H 2 + N 2 . Frequently ammonium nitrite is not at hand, and a mixture of sodium nitrite and either ammonium chloride or ammonium sulphate is employed. By the interaction of the sodium nitrite and the ammonium salt employed, ammonium nitrite is formed, which on heating decomposes into water and nitrogen. When, for instance, sodium nitrite and ammonium sulphate are em- ployed, the reaction is as follows : 2 NaN0 2 + (NH 4 ) 2 S0 4 = Na 2 SO 4 + 4 H 2 O + 2 N 2 . By heating ammonium bichromate (NH 4 ) 2 Cr 2 O 7 , or a mixture of ammonium chloride and potassium bichromate, nitrogen is formed, thus : K 2 Cr 2 7 + 2 NH 4 C1 = 2 KC1 + (NH 4 ) 2 Cr 2 O 7 , and (NH 4 ) 2 O 2 O 7 = Cr 2 O 3 + 4 H 2 O + N 2 ; or, by combining the two equations, K 2 Cr 2 7 + 2 NH 4 C1 = 2 KC1 + O 2 O 3 + 4 H 2 O + N r NITROGEN, AIR, AND THE HELIUM GROUP 149 When oxides of nitrogen are passed over red-hot copper, cupric oxide and nitrogen are formed, for example : When urea CO(NH 2 ) 2 is oxidized by means of hypochlorous or hypobromous acids or their salts, nitrogen is formed. So, for instance, with potassium hypobromite the reaction is : CO(NH 2 ) 2 + 3 KBrO = 2 H 2 O + 3 KBr + CO 2 + N a . The potassium hypobromite solution as usually prepared con- tains an excess of caustic potash, Avhich at once absorbs the carbon dioxide, forming potassium carbonate, which dissolves in water: 2 KOH + CO 2 == K 2 CO 3 + H 2 O. In estimating the quantity of urea in urine, which often needs to be done in medical practice, these reactions are used.- Nitrogen is a colorless, odorless, tasteless gas, which is 0.9672 time as heavy as air. It may be liquefied and solidified. Liquid nitrogen is colorless, and boils at 195.5 at atmospheric pres- sure. The critical temperature is 146, at which it requires a pressure of 35 atmospheres to liquefy the gas. Liquid nitro- gen has the specific gravity 0.80 at its boiling point. Solid nitrogen is a white, crystalline substance melting at 214; its specific gravity is 1.0265 at 252.5. Nitrogen is less soluble in water \han oxygen. At 10, 1000 cc. of water dis- solve 16.1 cc. of nitrogen, while at 0, 20.34 cc. are absorbed. At ordinary temperatures, nitrogen is a very inert element chemically. At higher temperatures it unites with lithium, boron, silicon, magnesium, barium, strontium, or calcium to form nitrides. Lithium burns readily in nitrogen, and even unites slowly with that gas at ordinary temperatures, forming the nitride Li 3 N. Magnesium at red heat absorbs nitrogen greedily, forming Mg 3 N 2 . In general, nitrogen is trivalent in the nitrides. When nitrogen and oxygen are mixed and subjected to the action of the electric spark (Fig. 35), nitrogen and oxygen unite to form an oxide of a brown color. Its formula is NO 2 ; at room temperatures it is N 2 O 4 . Hydrogen and nitrogen when mixed and similarly sparked yield small amounts of ammonia, which is a nitride of hydrogen having the composition NH 3< 150 OUTLINES OF CHEMISTRY FIG. 35. Due to electrical disturbances in the atmosphere, especially during thunder storms when lightning flashes from cloud tc cloud or to earth, small amounts of ammonia and oxides of nitrogen are formed. The atomic weight of nitrogen is 14.01; and since at and 760 mm. pressure 22.38 liters of nitrogen weigh 27.98 grams, the molecule contains 2 atoms and the mo- lecular formula is N 2 . This is also in harmony with the composition of ammonia and of the oxides of nitrogen by volume, as will appear later. In compounds nitrogen is either triv- alent or pentavalent. Its atomic weight was determined by r Stas, who ascertained the proportion by weight in which nitrogen exists in silver nitrate and in ammonium chloride. The Air. As already stated above, the air consists of about one fifth oxygen and four fifths nitrogen by volume. That these gases are not chemically bound to each other but simply mixed is evident from the following facts : (1) When the air is cooled, the oxygen condenses to a liquid first, leaving the nitrogen in form of a gas ; or when liquid air is boiled, the nitrogen distills off first, leaving nearly pure liquid oxygen behind. (2) The composition of the air, though nearly con- stant, varies somewhat at different times and places, the oxy- gen content commonly varying from 20.9 to 21.0 per cent. (3) Water will dissolve air to some extent. When the water is then deprived of this air by boiling, the air expelled from the water is richer in oxygen and poorer in nitrogen than ordinary air. Thus in air expelled from water the oxygen content is 35.1 per cent and the nitrogen is 64.9 per cent; whereas in ordinary air the corresponding figures are 20.96 and 79.04 per cent, respectively. (4) Air made by mixing four volumes of nitrogen and one of oxygen behaves like ordinary air. During NITROGEN, AIR, AND THE HELIUM GROUP 151 the preparation of the mixture there is neither a change of volume nor of temperature. The amount of oxygen and nitrogen in the air may be deter- mined by passing air freed from carbon dioxide and moisture over red-hot copper and collecting and weighing the nitrogen, which is not absorbed by the copper. The oxygen is determined by the increase of weight of the copper, which has united with the oxygen of the air passed over it. This is the method em- ployed by Dumas and Boussingault in 1841. Another method consists of mixing a carefully measured volume of air with a known excess of hydrogen and exploding the mixture by means of an electric spark. In this way the oxygen completely unites with hydrogen to form water whose volume is extremely small relatively. And so from the diminution of the gaseous volume after the explosion and the known relation of the volumes of hydrogen and oxygen in water, the amount of oxygen in the air may readily be computed. As a result of the average of numerous analyses of air, it has been found that the atmosphere consists essentially 0/21 volumes of oxygen to 79 volumes of nitrogen, or of '23*2 per cent oxygen and 76.8 per cent nitrogen by weight. Usually the composition of differ- ent samples of air varies from these figures by only one-tenth of a per cent. A liter of air at and 760 mm. pressure weighs 1.2933 grams. That the ratio of oxygen to nitrogen in air is so nearly constant is due to the fact that while animals are con- tinually using up oxygen in respiration, plants are on the other hand giving off oxygen to the air. Furthermore, the atmos- phere is so vast that the ordinary processes of combustion make scarcely a preceptible impression upon its oxygen content. Besides oxygen and nitrogen, the air always contains water vapor, ammonia, hydrogen, nitric acid, carbon dioxide, dust particles of organic as well as inorganic nature, and various bacteria and other microbes. All of these constituents are, how- ever, quite variable in amount. In the neighborhood of cities, sulphur dioxide and hydrocarbon gases have also been found in the air. The amount of water vapor in the air varies greatly with the locality and the temperature. Air saturated with moisture at contains 4.87 grams of water vapor per cubic meter, while at 20 it contains 17.157 grams. As stated in connection with the consideration of water, the air is usually 152 OUTLINES OF CHEMISTRY saturated to only about two thirds of its capacity. The amount of moisture in the air is best found by passing a given volume of it through sulphuric acid and phosphorus pentoxide and determining the increase in weight of these drying agents. In normal country air or air over the sea, there are about 3 volumes of carbon dioxide in every 10,000 volumes. In city air, the carbon dioxide content is often from 6 to 7 volumes per 10,000. In closed rooms where the air is contaminated by respiration and combustion of illuminating gas or oil, the carbon dioxide content may run as high as 6 to 8 times the latter amount. Air containing more than 7 volumes of carbon dioxide is con- sidered harmful for continuous breathing. The carbon dioxide in the air is determined by passing a known volume of the latter through baryta water and weighing the barium carbonate formed. The reaction that takes place is : Ba(OH) 2 + CO 2 =BaCO 3 + H 2 O. City air contains more carbon dioxide than country air, mainly because of large amounts of fuel consumed in cities, and because in the country the carbon dioxide is taken from the air to a con- siderable extent by plants. Ammonia occurs in the air in very minute and variable amounts hardly exceeding from 0.5 to 1 gram per 10,000 grams of air. It arises as a decomposition product of organic matter and is not present in the air in the free state, but is commonly combined with nitrous and nitric acids as nitrites and nitrates. The latter acids are formed, as already mentioned, when light- ning discharges in the air. The ammonium salts are washed from the atmosphere during rains. Thus they get into the soil and serve as an important nitrogen supply for plants. The latter get their nitrogen from this source or from manures. Leguminous plants, like peas, beans, and the various varieties of clover, are able to get their nitrogen supply from little nodules which are produced on their roots by certain species of bacteria, which get nitrogen directly from the atmosphere that circulates in the porous soil. These nodules may contain up to five per cent of nitrogen. Many plants are incapable of assimi- lating nitrogen in form of ammonia. The latter must first be oxidized. This is brought about by bacterial action in the soil. The amount of nitric acid in the air is small and very variable. R'din \\ NITROGEN, AIR, AND THE HELIUM GROUP 153 water has been observed to contain 0.14 part of nitrogen as nitrates per million parts of water on the average in some localities. As has already been stated, it is doubtful whether ozone is normally present in the air. The effects observed on starch potassium iodide paper may well be due to hydrogen peroxide or higher oxides of nitrogen. The hydrogen content of the air varies considerably. Ray- leigh found it to be 0.003 per cent by volume. Dewar isolated 0.001 per cent hydrogen from liquid air ; while Gautier claims to have found as much as 0.02 per cent. The hydrogen gets into the atmosphere from volcanic gases, and as a product of bacterial action. During the process of the decay of animal arid vegetable matter there are also still other gases produced, which enter the atmosphere. These are, however, soon oxi- dized, especially in the presence of sunlight. The particles of solid matter in the air frequently carry bacteria. As a rule the bacteria found in the air are harmless, though pathogenic organisms do get into the air, especially in the sick room and in crowded cities. Dry weather and winds increase the amount of dust in the air, and also the number of organisms that cling to dust particles. The spores of molds and microbes producing fermentation and putrefaction are practically always present in the air. By filtering the latter through plugs of cotton, dust and microbes may be removed from the air. Normal air contains but 4 or 5 microorganisms per liter. The waters of rivers and inland lakes contain from 5,000 to 20,000 organisms per cubic centimeter, whereas the soil con- tains about 5 times the latter number per cubic centimeter. Thus, it is clear that the air is relatively free from organisms. The latter get into the air chiefly from the dry soil, or dry ob- jects, as dust is carried from them by currents of air. Dust particles act as nuclei for the condensation of moisture in the formation of fogs. The air that is exhaled by animals and human beings con- tains, besides carbon dioxide, organic material ; and it is chiefly the latter which gives rise to headache and general depression that one experiences in crowded rooms. The decomposition products of this organic matter give rise to unpleasant odors which are frequently m^t in crowded, poorly ventilated rooms. 154 OUTLINES OF CHEMISTRY Liquid air is now produced on a commercial scale. The methods employed are founded upon the principle that by sub- jecting a gas to very high pressure and then allowing it to escape through a small orifice, the remaining gas is cooled, due to the heat absorbed in expansion. Thus, air compressed to about 200 atmospheres (i.e. 3000 pounds to the square inch) is cooled to room temperature by means of cold water, and this air is then allowed to escape from the long tube in which it is con- tained, through an orifice the size of which is controlled by means of a needle valve. The air thus enters another chamber which surrounds the first tube. The outflow is regulated so that in this second chamber the pressure of the air is about 20 atmospheres. In thus coursing from the first chamber into the second against a pressure of 20 atmospheres work is done, and the heat required to do this work is taken from the tube contain- ing the highly compressed air. The apparatus is carefully in- sulated from the surroundings by means of wool. After thus continuing to feed the apparatus compressed air for a few hours, the temperature in the inner tube becomes so low that the air liquefies and can then be drawn off. It is turbid in appear- ance, due to the solid particles of carbon dioxide and water that it contains. These may be filtered off. The filtrate is a clear liquid of bluish hue. Liquid air rapidly changes its compo- sition, since nitrogen evaporates faster than oxygen. Liquid air boils at about 190. The boiling point of nitrogen is 190.5 and that of oxygen is 182.5. After a time, nearly all the nitrogen has evaporated, leaving practically only oxygen behind. The latter is put on the market in steel cylinders as compressed oxygen. The Elements of the Helium Group. It was found by Lord Ray- leigh that a liter of nitrogen prepared from air weighs 1.2572 grams and that the same volume of nitrogen prepared from chem- ical compounds weighs 1.2521 grams. This led Rayleigh and Ramsay to investigate the composition of air more carefully, with the result that they discovered in it the new element argon in 1894. Argon may be prepared by passing air over heated copper to take out the oxygen, and then over hot magnesium or lithium to absorb the nitrogen. Or air may be mixed with an excess of oxygen and subjected to the action of the electric spark, the gas being kept over caustic potash solution to absorb the oxides of NITROGEN, AIR, AND THE HELIUM GROUP 155 nitrogen formed. In the latter method, the excess of oxygen may finally be absorbed by passing the gas over heated copper or by treatment with an alkaline solution of pyrogallic acid. About 0.9 per cent of the air, by volume, consists of argon. The gas has properties similar to those of nitrogen ; but argon has thus far not been obtained in combination with other ele- ments. It is very inert, chemically, whence its name argon, meaning inactive. The boiling point of argon is 186 and the melting point 189. It is more soluble in water than nitro- gen. At room temperatures about 40 cc. of argon are dissolved by 1 liter of water. The gas has consequently been found in all natural waters. Argon is 19.95 times as heavy as hydrogen ; its molecular weight is consequently 39.9. As it combines with no other elements whatever, its atomic weight cannot be ascertained by the usual means. The molecular heat of gases containing two atoms to the molecule is approximately 5 Cal. ; in the case of mercury vapor, which contains but one atom in the molecule, the molecular heat is but 2.5 Cal. Now it has been found that to heat 39.9 grams of argon one degree requires 2.5 Cal., con- sequently argon, like mercury, contains but one atom in its molecule. The molecular weight and atomic weight of argon are consequently the same, namely 39.9. That argon is a sim- ple substance is supported by the fact that it has a constant boiling point, and that by shaking the gas with water the dis- solved portion is identical with the undissolved portion. Argon is extraordinarily stable, and since it has not been decomposed into anything simpler, it must be regarded as an element. Helium, neon, krypton, and xenon, four additional new ele- ments, were later also discovered in the air by Ramsay and Travers. Helium was known to exist in the sun, whence its name. In 1895 Ramsay prepared helium by heating the mineral cleveite with sulphuric acid. The element exhibits a characteristic t yellow line in its spectrum. This line had previously been observed in the spectrum of the sun by Lockyer, who ascribed .t to an element which he called helium, then unknown on the iarth. Helium does not unite with any other element. Its molecular weight is 3.99 and its atomic weight is the same. A.bout 1.4 cc. of helium are dissolved by 100 cc. of water at room temperature. Helium was liquefied in 1908 by Professor 156 OUTLINES OF CHEMISTRY Kammerlingh Onnes of the University of Leiden. Its boiling point is 268.7, and its specific gravity in the liquid state is 0.15. This gas is the most difficult one to liquefy, and for sev- eral years it had resisted all attempts to condense it to a liquid. We may now say that all known gases have been liquefied. In the air helium occurs to the extent of 1 to 2 volumes per million volumes. Ramsay found neon, krypton, and xenon in the argon pre- pared from air. Helium and neon are dissolved in the liquid argon, from which they are expelled, together with argon, as the temperature rises. The residue Ramsay subjected to further fractional distillation, and so separated krypton and xenon from each other. By cooling a mixture of neon and helium with liquid hydrogen, neon solidifies, while the helium remains in the gaseous state. Neon, krypton, and xenon are inert gases that combine with no other elements; they are mono-atomic. Their atomic weights as determined from their densities are : Neon . . . 20.2 Krypton . . . 82.92 Xenon . . . 130.2 Their boiling points are as follows : Neon ... 243 (approximately) Krypton . . . -152 Xenon . . . -109 Krypton melts at 169 and xenon at 140. The following table gives the amounts of the gases of the helium group contained in one cubic meter, i.e. 1000 liters, of air : Helium . . 0.0015 liter . . = 0.00027 gram Neon . . 0.015 liter . . = 0.01339 gram Argon . . 9.4 liters . . =16.76 grams Krypton . . 0.00005 liter . . = 0.00018 gram Xenon . . 0.000006 liter . . = 0.00003 gram NITROGEN, AIR, AND THE HELIUM GROUP 157 REVIEW QUESTIONS 1. Compare oxygen and nitrogen as to (a) their occurrence in nature, (6) their physical properties, (c) their chemical properties. 2. Give four reasons for regarding the air as a mixture and not a chemical compound. 3. Make a list of all of the gases that normally occur in the air, stating their amounts in per cent by volume. 4. State of what use each of the constituents of the air is in plant and animal life. 5. How do plants obtain their supply of nitrogen? What is a legu- minous plant? How do such plants aid in the assimilation of nitrogen? 6. Outline the essential features of a process of preparing argon from the air. 7. What property do the rare gases of the atmosphere have in common ? CHAPTER XI COMPOUNDS OF NITROGEN WITH HYDROGEN AND WITH THE HALOGENS History and Occurrence of Ammonia. Ammonia is by far the most important compound of nitrogen with hydrogen. Up to 1774 it was known only in its aqueous solution, which Glauber called " spiritus volatilis salis armoniaci," and which was later named spirits of hartshorn and spirits of sal ammoniac. Am- monia was prepared by treating sal ammoniac, that is, ammo- nium chloride, with lime or some other alkali, whence the name spirits of sal ammoniac. It may also be obtained by heating hoofs and horns of animals out of contact with the air, whence the term spirits of hartshorn. The process of thus heating substances out of contact of the air in a retort and decompos- ing them into other products is called destructive distillation. Priestley discovered ammonia gas in 1774 by evolving it from lime and ammonium chloride and collecting it over mercury. He called it " alkaline air " ; for the gas turns red litmus blue and acts in other ways like a strong alkali. As stated in con- nection with nitrogen, ammonia occurs in small amounts in the atmosphere in form of ammonium salts, particularly as ammo- nium carbonate. It is a product of the decomposition of all vegetable and animal matter, and hence is found in all natural waters and soils. In the form of salts, mainly nitrate and nitrite, it occurs in rain water. Its occurrence in soils is im- portant, for it is a fertilizer. In the neighborhood of volcanoes, particularly those of Tuscany, ammonia occurs in the form of sulphate and chloride of ammonium. Ammonium chloride used to be prepared in Egypt in the oasis near the temple of Jupiter Ammon, from the soot obtained by heating camel's dung which was used as fuel; thus the salt received its name sal ammoniac, from which comes the term ammonia. Preparation and Properties of Ammonia. When a mixture of nitrogen and hydrogen is subjected to the silent electrical 158 AMMONIA AND OTHER NITROGEN COMPOUNDS 159 discharge as in making ozone, small amounts of ammonia are formed by direct union of the elements, thus : The common method of preparing ammonia is by heating am- monium chloride with slaked lime : 2 NH 4 C1 + Ca(OH) 2 = CaCl 2 + 2 H 2 O + 2 NH 3 . Any other ammonium salt may be used instead of the chloride, and other alkalies, like sodium or potassium hydroxide, may serve in place of lime, which, however, is the cheapest. By the reduction of nitrates and nitrites with nascent hydro- gen, ammonia may be produced, thus : KNO 3 + 8 H = KOH + 2 H 2 O + NH 3 . KNO 2 + 6 H = KOH + H 2 O + NH 8 . By the dry distillation of nitrogenous animal and vegetable material ammonia is formed. So by heating coal (which rep- resents the remains of vegetation of the carboniferous age) out of contact with air, as is done in the manufacture of illuminat- ing gas from coal, ammonia is produced. The coal gas formed is passed through water, which readily dissolves the amnxonia, and it is from these ammoniacal liquors of the gas works that the ammonia of commerce is almost entirely obtained at present. From these liquors the gas is expelled by heating with slaked lime. The ammonia so liberated is passed into sulphuric acid, and the sulphate of ammonium thus formed is purified by re- crystallization. From this pure salt, pure ammonia and other ammonium products are in turn prepared. By heating organic nitrogenous products with strong alkalies, ammonia is produced. This is frequently used in ascertaining the amount of nitrogen in organic substances, particularly as they occur in fertilizers, sewage, drinking water, etc. In this process a strong solution of caustic potash and potassium per- manganate is frequently employed. The latter substance is a powerful oxidizing agent and so aids in the destruction of the organic material. Animal matter when heated with fuming sulphuric acid is decomposed, the nitrogen being converted into ammonia, which unites with the sulphuric acid, forming ammonium sulphate. This process (known as 160 OUTLINES OF CHEMISTRY method) is of importance in the chemical analysis of nitroge- nous organic substances. Ammonia is a colorless gas of a strong, peculiar, penetrating odor. It is 0.59 time as heavy as air. It may be condensed to a liquid which boils at 32.5. It has also been obtained in form of white crystals that melt at 78. The specific gravity of liquid ammonia, taken under pressure at 0, Is 0.6233. In water the gas is extremely soluble At 0, 1 volume of water absorbs 1148 volumes of ammonia, while at 16 and 50, only 764 volumes and 306 volumes, respec- tively, are absorbed. On boiling an aqueous solution of ammonia, the gas is completely expelled, which fact is frequently used in laboratories for preparing ammonia gas. On account of its solubility in water, ammonia gas is collected over mercury, or simply by displace- ment of air, the vessel in which gas is to be collected being supported with the bottom upward, for the gas is but little more than half as heavy as air. The weight of a liter of ammonia gas at and 760 mm. is 0.7635 gram, and since the gas consists of 82.27 per cent nitrogen and 17.73 per cent hydrogen by weight, its formula is NH 3 . By electrolyzing an aqueous ammonia solution, to which some common salt has been added to make the solution conduct better, three volumes of hydrogen are obtained to one volume of nitrogen. The apparatus used for this purpose is the same as that shown in Fig. 2. Again, when a given volume of ammonia gas (Fig. 36) is treated with a concentrated solution of potassium hypobromite, nitrogen is formed which occupies half the volume of the original ammonia. Care must be taken not to admit air into the tube during the experiment. We thus see that 2 volumes of ammonia yield 1 volume of nitrogen, while by electrolysis 3 volumes of hydrogen and 1 volume of nitrogen were obtained from ammonia. Con- sequently, these volume relations may be expressed thus : 3 volumes hydrogen + 1 volume nitrogen = 2 volumes ammonia gas. We have here another excellent confirmation of the law of Gay-Lussac of the combination of gases by volume. By FIG. 36. AMMONIA AND OTHER NITROGEN COMPOUNDS 161 Avogadro's hypothesis, equal volumes of gases contain an equal number of molecules, hence : 3 molecules hydrogen 4- 1 molecule nitrogen = 2 molecules ammonia or 3H 2 + N 2 = 2NH 3 . While it is true that a mixture of 3 volumes of hydrogen and 1 volume of nitrogen when subjected to the electric spark yields small amounts of ammonia, it is also the case that when the latter gas is thus treated it is partly decomposed into nitro- gen and hydrogen. The reaction is thus a reversible one : If none of the gases are removed, an equilibrium is finally slowly reached, which is the same in each case, the gases con- sisting of 2 per cent ammonia and 98 per cent of uncombined nitrogen and hydrogen. So if ammonia gas con- tained in the closed limb of the appa- ratus shown in Fig. 37 is treated with the electric spark, the volume of hy- drogen plus nitrogen formed will be nearly twice that of the original volume of ammonia. If, however, the ammonia is removed (absorbed by sulphuric acid, for example) as fast as it is formed, the reaction completes itself from left to right as would be expected, according to the law of mass action. When ammonia is oxidized, as, for instance, by passing it over hot copper oxide, the latter is reduced and the only products formed are water and nitrogen, thus showing that the gas is composed of hydrogen and nitrogen. By thus oxidizing a definite volume of ammonia and weighing the water and nitrogen formed, the percentage composition of ammonia has been determined. The results of such analyses have already been given above. On account of its hydrogen content ammonia will burn. The action proceeds in oxygen, but not in air. Thus, when FIG. 37. 162 OUTLINES OF CHEMISTRY in a flask (Fig. 38) strong ammonia water is heated till ammonia is copiously evolved, and oxygen is then conducted into the gas, the mixture when lighted will burn at the mouth of the flask. The products are mainly water and nitrogen, though ni- trous and nitric acids are also formed to a slight extent. These acids unite with the excess of ammonia to form ammonium nitrite and nitrate. The latter salts are more copiously formed when a heated spiral of platinum wire (Fig. 39) is hung into a mixture of oxygen and ammonia gases. The platinum continues to glow, and white fumes form which consist of the salts men- tioned. The platinum here acts as a catalytic agent. Ammonia water is lighter than water. The saturated solu- tion at 14 C. contains 36 per cent NH 8 and has a specific gravity of 0.8844. It is sold as a con- centrated ammonia, and may be diluted to any other strength desired. Ammonia unites directly with acids, forming salts, thus : FIG. 38. = NH 4 C1. 2NH 3 +H 2 S0 4 =(NH 4 ) 2 S0 4 . NH 3 +HN0 3 =NH 4 N0 3 . We may regard these salts as derived from the acids by the replacement of each hydrogen atom by the group NH 4 . So we may also consider that the group NH 4 plays the role of an atom of a univalent metal, like Na or K. For this reason, NH 4 FIG. 39. is called ammonium, the ending urn being used to indicate that chemically it is analogous to a metal. When ammonia dis- AMMONIA AND OTHER NITROGEN COMPOUNDS 163 solves in water, much heat is evolved, and we may consider that the addition product formed is NH 4 OH, thus : NH 3 +H 2 O = NH 4 OH. The latter has not been isolated ; but the aqueous solutions act as though this substance were contained in them. So, for example, when ammonia water is neutralized by hydrochloric or sulphuric acid, the action may be expressed thus : NH 4 OH + HC1= NH 4 C1 + H 2 O. 2 NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2 H 2 O. At dull red heat all ammoniu'm salts are volatilized. This is a very important fact in chemical analysis. In many cases the salts are simply broken up into ammonia and the free acid, thus : In such instances, which are typical cases of dissociation, the actions are reversible, the products again uniting as the tem- perature is lowered. The vapor of ammonium chloride con- tains all three products named in the above equation ; which was demonstrated by Henri Saint Claire Deville, who separated hydrochloric acid and ammonia from these vapors by diffusion, making use of the fact that ammonia, the lighter gas, diffuses more rapidly than hydrochloric acid. By means of chlorine, ammonia is decomposed : 2 NH 3 + 3 C1 2 = 6 HC1 + N 2 . When excess of ammonia is present, the latter at once unites with the hydrochloric acid formed and produces ammonium chloride. Ammonia acts on many metals. Thus, sodium and potassium when heated act on ammonia as follows : 2 Na + 2 NH 3 = 2 NaNH 2 + H 2 . 2 K + 2 NH 3 = 2 KNH 2 + H a . The compounds formed are sodium amide and potassium amide. They are decomposed by water into ammonia and the hydrox- ide of the metal, so, for instance : KNH Q + H 2 = KOH 164 OUTLINES OF CHEMISTRY The nitrides of lithium NLi 3 and of magnesium N 2 Mg 3 be regarded as ammonia in which the hydrogen atoms are re- placed by the respective metals. These nitrides may be formed by igniting the metals in ammonia. On treating the nitrides with water, ammonia and the metallic hydroxides are formed, thus : NLi 8 + 3 H 2 O = 3 LiOH + NH 3 . N 2 Mg 3 + 6 H 2 = 3 Mg(OH) 2 + 2 NH 3 . The fact that ammonia water will dissolve metals like zinc, copper, and silver and also their oxides, is used in cleaning tarnished metallic articles, for the action of ammonia is not as drastic as that of an acid, and, furthermore, the ammonia readily evaporates after use. Liquid ammonia is a great solvent. In this respect it is similar to water, for it dissolves many salts, forming with some of them addition products that are analogous to hydrates that water forms with salts. Thus, with water copper sulphate forms the compound CuSO 4 5 H 2 O, in which the water is spoken of as water of crystallization. Similarly with dry ammonia copper sulphate forms the compound CuSO 4 5 NH 3 , in which the ammonia may be called ammonia of crystallization. The properties of liquid ammonia solutions have been investigated by E. C. Franklin in recent years. Liquid ammonia is much like water in that it has a high specific heat, 1.02 between and 20, and a high latent heat of vaporization. It takes 316 cal. to vaporize 1 gram of liquid ammonia at 0. This fact is used in the artificial ice machines^ in which liquid ammonia is evaporated in tubes which are sur- rounded with concentrated calcium chloride solution. The evaporation of the ammonia requires heat, which is abstracted from the brine; and this cold brine is then distributed by means of a system of tubes to the places where refrigeration is re- quired. The ammonia is again liquefied by compression with a powerful pump, and so it can be used over and over again in a closed system of tubes. The calcium chloride brine also circu- lates in a closed system of tubes, the brine after becoming warmed being returned and again chilled. That a considerable lowering of temperature can be produced by the evaporation of ammonia may be shown by a simple experiment. When a con- AMMONIA AND OTHER NITROGEN COMPOUNDS 1G5 centrated aqueous solution is placed in a flask standing upon a board that is wet with a few cubic centimeters of water, and a strong current of air is passed into the solution by means of a pair of bellows, the evaporation of the ammonia proceeds so rapidly that in a few minutes the cold produced is sufficient to freeze the flask to the board. In ammonia NH 3 nitrogen is trivalent, whereas in ammonium salts the element is quinquivalent. Indeed, in most of its com- pounds nitrogen may be considered as having a valence of either three or five. Ammonium salts are readily detected by the fact that ammonia is evolved when they are treated with caustic alkali. The ammonia gas is easily distinguished by its odor and by the fact that it turns red litmus paper blue. When present in very small quantities, as in drinking water, ammonium salts are detected by means of a solution of mer- curic iodide HgI 2 in potassium iodide. This solution is made strongly alkaline by addition of caustic potash and is then known as Nesslers reagent. When added to a very dilute solution of an ammonium salt a yellow color is produced. In stronger solutions of ammonium salts a dark brown color or a precipitate is formed. Nessler's reagent is of great importance in analyzing potable waters, sewage, and the like. Hydrazine. The composition of hydrazine or diamide is ex- pressed by the formula H 2 N-NH 2 . This compound was dis- covered by Curtius in 1887. It may be made by the oxidation of urea, thus: NH 2 -CO-NH 2 + O = H 2 N-NH 2 + CO 2 ; or by the reduction of hyponitrous acid, thus : H-O-N -N-O-H + 6 H = H 2 N-NH 2 + 2 H 2 O. It forms white crystals melting at 1. Liquid hydrazine boils at 113. Its specific gravity at 15 is 1.013. It is miscible with water in all proportions and forms a hydrate H 2 N-NH 3 (OH), which melts at - 40 and boils at 120. Its 166 OUTLINES OF CHEMISTRY specific gravity is 1.03. Like ammonia, hydrazine is a strong base. Its solutions have a corrosive action on cork and rubber, and even on glass, especially at higher temperatures. With acids hydrazine forms salts. The action is similar to the formation of ammonium salts. A large number of compounds derived from hydrazine by replacing one or more of its hydrogens by hydrocarbon radicals are known in organic chemistry. One of these, namely, phenyl hydrazine (C 6 H 5 )HN-NH 2 , has been of special importance in the synthesis and investigation of sugars. Hydroxylamine. This compound, which was discovered by Lossen in 1865, and prepared in the pure state by Lobry de Bruyn in 1891, may be formed by the action of nascent hydro- gen either on nitric acid or nitric oxide, thus : HN0 3 + 6 H = 2 H 2 + NH 2 (OH). NO + 3H = NH 2 (OH). It may be considered as ammonia NH 3 in which one hydrogen atom has been replaced by the OH group. The group NH 2 is called the amido or amine group, whence the name hydroxyl- amine. It consists of white hygroscopic needles that melt at 33. The boiling point is 58 at 22 mm. and 70 at 60 mm. pressure. In water hydroxylamine dissolves readily. It has basic properties, showing alkaline reaction toward indicators, and forming crystalline salts with acids, thus : NH 2 OH + HC1 = NH 3 (OH)C1. 2 NH 2 OH 4- H 2 S0 4 = (NH 3 OH) 2 SO 4 . These salts may be regarded as ammonium salts in which one hydrogen has been replaced by OH. Hydroxylamine is, how- ever, a much weaker base than ammonia or hydrazine. On heating hydroxylamine or its compounds, decomposition sets in, which, on account of sudden evolution of gas, may take place with explosive violence. The reducing power of hydrox- ylamine is characteristic. By treating a hot alkaline solution of a cupric salt with hydroxylamine, red cuprous oxide is at once formed, thus : 4 CuO + 2 NH 2 (OH) = N 2 O + 3 H 2 O -f 2 Cu 2 O. The reaction will take place even when hydroxylamine is present merely as 1 part in 100,000. AMMONIA AND OTHER NITROGEN COMPOUNDS 167 Hydroxylamine readily decomposes into ammonia, nitrogen, and water, thus : 3 NH 2 OH = NH 3 + N 2 + 3 H 2 O. In organic chemistry hydroxylamine is of importance, because with aldehydes and ketones (which see) it forms compounds known as oximes. Hydrazoic Acid. This compound has the composition ex- pressed by the formula N 3 H. It is also called hydronitric acid, triazoic acid, or azoimide. It was discovered by Curtius in 1890. It may be made by passing nitrous oxide over sodium amide at 200, and then treating the resulting sodium hydrazo- ate with dilute sulphuric acid. The reactions are as follows: NaNH 2 + N 2 = H 2 O + NaN 3 . 2 NaN 8 + H 2 SO 4 = Na 2 SO 4 + 2 HN 3 . By carefully distilling the aqueous solution produced, a solution of the free acid in water may be obtained. The pure acid boils at 37. It is a colorless liquid with a disagreeable, penetrating odor. When inhaled, it irritates the mucous membranes. It explodes with violence, forming nitrogen and hydrogen with liberation of much heat, thus : = 3N 2 +H a . It is a monobasic acid, and in this respect it is similar to the hydrohalogen acids. Its salts are also unstable and liable to explode with violence. It is of interest to note that the one hydride of nitrogen NH 3 is alkaline and the other N 3 H is acid. The two will combine to form a salt, thus : The empirical formula of NH 4 (N 8 ), ammonium hydrazoate, is N 4 H 4 . Compounds of Nitrogen with the Halogens. With the halo- gens nitrogen forms extremely unstable compounds. Nitrogen trichloride. NC1 3 is formed by the action of chlorine upon ammonium chloride, thus : NH 4 C1 H- 3 C1 2 = 4 HC1 + NC1 8 . The compound may be prepared by the electrolysis of an aque- ous solution of ammonium chloride, the chlorine liberated acting 168 OUTLINES OF CHEMISTRY on the solution according to the above equation. Nitrogen trichloride is a thin, yellowish, oily liquid of specific gravity 1.65. It is an extremely dangerous substance to deal with, for it explodes with great violence when heated or brought into con- tact with substances like turpentine or phosphorus, or when exposed to sunlight. Often the explosion occurs spontaneously, which makes the danger of working with it very great indeed. It has a pungent odor, and its fumes irritate the mucous mem- branes. It is soluble in hydrocarbons and carbon disulphide, the solutions thus formed being yellow in color and compara- tively harmless. At 71 nitrogen trichloride boils and may be distilled, though the danger incurred in the operation is extremely great. By concentrated hydrochloric acid or aqueous ammonia solution, nitrogen trichloride may be decomposed, thus : NC1 3 + 4 HC1 = NH 4 C1 + 3 C1 2 . NC1 3 + 4 NH 4 OH = 3 NH 4 C1 + 4 H 2 O -f- N 2 . Nitrogen trichloride was discovered by Dulong in 1811. In working with the substance he was so unfortunate as to lose an eye and three fingers in consequence of an explosion. Nitrogen tribromide is a red, oily, highly explosive substance formed by the action of potassium bromide on nitrogen chloride. The substance is believed to have the composition represented by the formula NBr 3 . Nitrogen iodide is formed when iodine is treated with a con- centrated aqueous ammonia solution, or when an alcoholic solu- tion of iodine is mixed with strong aqueous ammonia. The compound is a brown powder having the composition N 2 H 3 I 3 , probably I 3 N = NH 3 . It is not explosive when wet; but when dry it is very explosive, a touch with a feather sufficing to cause it to explode with detonation. By treatment of silver hydrazoate AgN 3 with a solution of iodine in ether, triazoiodide IN 3 may be formed : AgN 3 +I 2 =AgI + IN 8 . It is a yellow powder of a very penetrating odor, and is ex- tremely explosive. AMMONIA AND OTHER NITROGEN COMPOUNDS 169 REVIEW QUESTIONS 1. What is the composition of ammonia? State its chief physical and chemical properties. 2. Write the equation expressing the action of slaked lime upon ammonium chloride. This action will proceed to completion, whereas if a dilute solution of ammonium chloride is treated with lime, the action does not proceed to completion. Explain these facts. In general, how may ammonia be liberated from any ammonium salt ? 3. From what source are the ammonium salts of commerce obtained? Explain fully. 4. What gases are produced when ammonia is decomposed by elec- trolysis? What are the respective volumes of these gases? What law does this illustrate ? 5. Under what conditions will nitrogen and hydrogen unite directly to form ammonia ? Write the equation and discuss it from the standpoint of chemical equilibrium. 6. What use is made of ammonia and of ammonium compounds? 7. In testing a drinking water why does a chemist always test for the presence of ammonium compounds? 8. Mention three halides of nitrogen. How may they be prepared? What is their chief characteristic ? CHAPTER XII OXY-ACIDS AND OXIDES OF NITROGEN THREE oxy-acids and five oxides of nitrogen are known. These are nitric acid HNO 3 , nitrous acid HNO 2 , hyponitrous acid H 2 N 2 O 2 , nitrogen pentoxide or nitric anhydride N 2 O 5 , nitrogen peroxide NO 2 or N 2 O 4 , nitrogen trioxide or nitrous anhydride N 2 O 3 , nitric oxide NO, and nitrous oxide N 2 O. In the consideration of these compounds nitric acid will be taken up first, for from it the other oxy-acids and oxides named are generally prepared. History, Occurrence, and Preparation of Nitric Acid. Nitric acid was known to the alchemists under the name of aquafortis. Up to the seventeenth century the acid was prepared by heat- ing a mixture of saltpeter, copper sulphate, and alum according to directions given by the alchemist Geber, who probably lived in the ninth and tenth centuries. In this process copper sulphate and alum yield sulphuric acid, which unites with the potassium of the saltpeter, thus setting nitric acid free. Pre- pared by this method, the acid was impure. In 1650 Glauber- prepared nitric acid by treating saltpeter with sulphuric acid, and this method is in vogue to the present day. Though Lavoisier studied nitric acid and showed that it contained oxy- gen, he did not ascertain the real nature of the acid. In 1784 Cavendish demonstrated the nature of the acid by preparing it by passing an electric spark through air. In this way nitrogen peroxide is formed, which in contact with water yields nitric acid (see below). It has already been stated that nitric acid and nitrates occur in small amounts in the atmosphere. In the soil and in natural waters nitrates occur as the final product of the decomposition and oxidation of animal and vegetable matter. The chief source of nitric acid is Chili saltpeter or sodium nitrate. Nitric acid gets its name from the fact that it is commonly pre- pared from niter, saltpeter. 170 OXY- ACIDS AND OXIDES OF NITROGEN 171 By treating a nitrate like potassium or sodium nitrate with concentrated sulphuric acid, nitric acid is liberated, thus : NaNO 3 + H 2 SO 4 = NaHSO 4 + HNO 3 . In the laboratory the sodium nitrate is generally placed in a glass retort (Fig. 40), sulphuric acid is added, and the mixture FIG. 40. gently heated, when nitric acid distills over. On a commercial scale Chili saltpeter is treated with sulphuric acid in a cast-iron retort, and the nitric acid formed is condensed in bottles of stone- ware that contain a little water, the last bottle being connected with a tower filled with coke over which water trickles so as to dissolve the acid vapors that still remain uncondensed. Of late stoneware pipes are frequently employed instead of the bottles. In this way an aqueous solution is obtained which contains about sixty per cent nitric acid and has a specific gravity of 1.37. By using dry sodium nitrate and concentrated sulphuric acid, the nitric acid obtained has a specific gravity of 1.53 and is practically free from water. On heating the pure acid, as in the process of distillation, it decomposes in part, thus : 4 HN0 3 = 2 H 2 + 4 NO 2 + O 2 . The nitrogen dioxide forms reddish brown fumes that dissolve in the nitric acid. This solution, which fumes strongly in the 172 OUTLINES OF CHEMISTRY air, is termed red fuming nitric acid. Its specific gravity is about 1.54. When the electric spark passes through air, brown fumes are formed which are nitrogen dioxide. These in contact with water form nitric acid, thus : 3 NO 2 + H 2 O = 2 HNO 3 + NO. This may readily be shown by means of the apparatus in Fig. 35. Sparks from an induction coil are passed through the air between the platinum points in the glass globe. After a time the gas in the globe appears brownish in color. On shaking the gas with water, and testing with blue litmus paper, the presence of acid is demonstrated. Many attempts have been made to use this process for the profitable production of nitric acid on a commercial scale. These have been unsuccessful till recently, for the amount of nitric acid produced was too small as compared with the electric power that had to be expended, even when water power was available for running dynamos. Of late, however, the process has been perfected by subjecting the electric arc formed to the action of a powerful electro-mag- netic field. In this way arcs produced by means of large alter- nating current dynamos are obtained in form of disks over six feet in diameter, through which air is passed. Thus in this process, which is used in Norway and is known as the Birkelund and Eyde process, nitric oxide NO is formed at the very high temperature of the arc. It is the high temperature secured by means of the electric arc, and not an electrical effect, that causes the oxygen and nitrogen to unite. When the nitric oxide is then treated with air and water in a tower filled with moist coke, nitrogen dioxide and nitric acid form, thus : 2 NO + O 2 5> 2 NO 2 . H 2 O + 3 NO 2 ;j NO + 2 HNO S . All of the reactions involved in the process are reversible, so that in order to have them go to completion from left to right as far as possible, the products formed are rapidly removed by condensation and solution. The dilute nitric acid thus obtained is neutralized with lime, and the calcium nitrate formed is sold as a fertilizer. This process of making nitric acid is of special importance, because the large demands made upon the deposits OXY-ACIDS AND OXIDES OF NITROGEN 173 of Chili saltpeter annually will ere long exhaust this source of supply, though new deposits have been found of recent years in the same region. About one and a half million tons of Chili saltpeter have been used annually in recent years. The salt is used as a fertilizer to a large extent, but it is also employed in making nitric acid, which is used in the manufacture of explo- sives, dyestuffs, medicinal chemicals, nitrates of metals, etc. Upwards of 100,000 tons of nitric acid are used annually in the chemical industries of the world. Properties of Nitric Acid. Pure nitric acid is a colorless liquid which boils at 86 with partial decomposition, as stated above. It is a monobasic acid whose composition is expressed by the formula HNO 3 . On distilling the acid under diminished pressure, this decomposition may be avoided, and this is actually done in the manufacture of pure nitric acid. On cooling, nitric acid forms colorless crystals that melt at 42. The acid that is sold in the market as concentrated nitric acid is a 68 per cent solution. It has a constant boiling point, which is 120.5, and a specific gravity of 1.414 at 15. The composition of this con- stant boiling solution changes when it is distilled under dimin- ished pressure (compare hydrochloric acid) ; the solution is consequently not regarded as a chemical compound. Nitric acid is a powerful acid which fumes in the air. In aqueous solutions it is much more stable than when pure. The concentrated acid is rather an unstable substance. It is slowly decomposed in sunlight to a slight extent, the yellow color de- veloped being due to the formation of nitrogen dioxide, which remains in solution. At about 280 nitric acid decomposes, practically completely, into nitrogen dioxide, water, and oxy- gen. Concentrated nitric acid has a very corrosive action on the skin, producing painful wounds that are slow to heal. More dilute solutions color the skin yellow, due to the forma- tion of nitro products. The effect upon wool, linen, silk, and other organic substances is similar. When nitric acid is neutralized with bases, nitrates are formed. These salts are all readily soluble in water. Nitric acid does not attack gold or platinum. When it attacks other met- als, they are either oxidized, as is the case with tin, or converted into nitrates, as is more frequently the case. So zinc, copper, iron, magnesium, when treated with nitric acid, are converted 174 OUTLINES OF CHEMISTRY into the corresponding nitrates. There is, however, no concom- itant evolution of hydrogen, as when these metals are attacked with hydrochloric acid, for instance ; for the hydrogen at once attacks the nitric acid, reducing it commonly to nitric oxide NO and water. With metals like zinc, iron, and magnesium, the temperature and concentration of the acid and resulting solutions may be regulated so as to secure a very gradual reduc- tion of the nitric acid, the products being successively nitrogen dioxide NO 2 , nitrous acid HNO 2 , nitric oxide NO, nitrous oxide N 2 O, nitrogen N 2 , hydroxylamine NH 2 OH, and ammonia NH g . As nitrogen dioxide and nitrous acids are readily reduced to nitric oxide, the latter is generally formed when nitric acid acts on a metal, thus : 3 Zn + 8 HNO 3 = 3 Zn(NO 3 ) 2 + 4 H 2 O + 2 NO. Nitric acid is a powerful oxidizing agent and will convert many of the non-metals into their highest oxidation products with ease. Thus, when heated with nitric acid, phosphorus is oxidized to phosphoric acid, sulphur to sulphuric acid, carbon to carbon dioxide. A glowing stick of charcoal thrust into concentrated nitric acid continues to burn brightly. While neither nitric nor hydrochloric acid alone attacks gold or platinum, these metals are readily dissolved in a mixture of nitric and hydrochloric acids. This mixture, since it dis- solves gold, the " king of metals," is called aqua regia. The action depends upon the fact that nitric acid oxidizes hydro- chloric acid, one of the products formed being chlorine, which attacks gold. Aqua regia was known even in the days of alchemy, for Geber dissolved gold in a solution of ammonium chloride in nitric acid. The action of concentrated nitric and hydrochloric acids on each other may be represented thus 3 HC1 + HNO 3 = 2 H 2 O + NOC1 + Cl a . The compound NOC1 is called nitrosyl chloride. It occurs here as one of the products of the reaction, but it does not attack gold. Nitrogen Pentoxide. When nitric acid is treated with phos- phorus pentoxide, nitrogen pentoxide or nitric acid anhydride is formed. The action consists of the subtraction of water from nitric acid : 2 HN0 3 + P 2 5 = 2 HP0 8 -f N 2 6 . OXY-ACIDS AND OXIDES OF NITROGEN 175 In this process, pure nitric acid is carefully mixed with about an equal weight of phosphorus pentoxide in the cold, and the sirupy mass obtained is carefully distilled. Nitric anhydride may also be formed from silver nitrate and chlorine, thus : 4 AgN0 3 + 2 C1 2 = 4 AgCl + 2 N 2 O 5 + O 2 . It was by this method that the substance was discovered by Deville in 1849. Nitrogen pentoxide consists of colorless pris- matic crystals that melt at 30, forming a dark yellow liquid. The latter boils at 50 with concomitant partial decomposition. It is very unstable, readily giving off a portion of its oxygen, thus : 2N 2 6 = 4N0 2 +0 2 . The decomposition goes on slowly, though spontaneously, at ordinary temperatures. When rapidly heated, the decomposi- tion proceeds with explosive violence. The substance cannot be kept long in any case. Dissolved in water, nitric anhydride N 2 O g yields nitric acid. Nitric Oxide. Nitric oxide NO, discovered by Priestley in 1772, is formed by the action of copper, silver, mercury, and many other metals upon a solution of nitric acid of about 30 to 35 per cent : 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO. The temperature should be kept low during the reaction, as otherwise nitrous oxide N 2 O and nitrogen are apt to form. Nitric oxide is also conveniently produced by the action of fer- rous chloride or sulphate on nitric acid in presence of hydro- chloric or sulphuric acid, thus : HNO 3 + 3 FeCl 2 + 3 HC1 = 2 H 2 O + 3 FeCl 8 + NO. 2 HN0 3 + 6 FeS0 4 + 3 H 2 .SO 4 = 4 H 2 O + 3 Fe 2 (SO 4 ) 3 +2 NO. The gas is colorless, but on coming in contact with the oxygen of the air it immediately turns brown, due to the formation of nitrogen dioxide : 2NO + O 2 =2NO 2 . It is consequently necessary to expel the air from the apparatus before collecting the gas, which may be done over water since the latter dissolves nitric oxide but slightly. Nitric oxide is a colorless, neutral gas which is 1.039 times 1T6 OUTLINES OF CHEMISTRY as heavy as air. Its critical temperature is 94, and its criti- cal pressure 71.2 atmospheres. Under atmospheric pressure the liquid, which is colorless, boils at -150. When solidified, nitric oxide forms colorless crystals that melt at 167. At one volume of water absorbs 0.075 volume of the gas, and at 20, 0.05 volume. Nitric oxide is the most stable of the oxides of nitrogen. A lighted candle or burning sulphur will be extinguished when introduced into the gas. On the other hand, burning mag- nesium or phosphorus will continue to burn in the gas with great brilliancy. On heating metallic sodium or iron in nitric oxide (Fig. 41), these metals are oxidized, and the nitrogen which remains occupies just one half of the volume of the nitric oxide, thus : 4Na + 2 NO = 2 Na 2 O + N 2 . (solid) (2 volumes) (solid) <1 volume) 3Fe + 4 NO = Fe 3 O 4 + 2 N 2 . (solid) (4 volumes) (solid) (2 volumes) Knowing the specific gravities of nitric oxide and nitrogen, and the fact that 2 volumes of nitric acid yield 1 volume of nitrogen, it follows that in nitric oxide 14 grams of nitrogen are combined with every 16 grams of oxygen. As nitric oxide is 15 times heavier than hydrogen, its molecular weight is 30. The for- mula of nitric oxide is therefore NO. Nitric oxide may be used to detect the presence of free oxy- gen in a mixture of gases on account of its ability to form brown fumes NO 2 with oxygen. In solutions of ferrous salts nitric oxide dissolves readily, forming a dark brown liquid. From these solutions the gas is expelled by heating. It is probable that the solutions contain the unstable compound FeSO 4 -NO. This reaction is a delicate test for nitrates, for, as we have seen, ferrous salts in presence of free acid readily reduce nitrates to NO, which then gives the brown color with the excess of the ferrous salt. Nitrogen Dioxide and Tetroxide. Nitrogen dioxide is pro- duced by heating nitrates of the heavy metals : OXY-ACIDS AND OXIDES OF NITROGEN 177 2 Pb(N0 3 ) 2 = 2 PbO + 2 + 4 N0 2 . 2 Cu(NO 3 ) 2 = 2 CuO + O 2 + 4 NO 2 . When oxygen acts on nitric oxide, nitrogen dioxide is formed : 2 NO + 2 = 2 N0 2 . When the electric spark is passed through a mixture of oxygen and nitrogen, nitrogen dioxide forms: As stated under nitric acid, this reaction takes place slowly and is ordinarily very incomplete. Concentrated nitric acid oxi- dizes nitric oxide to nitrogen dioxide, consequently the latter is formed when metals like copper or tin are acted upon by strong nitric acid even out of contact with the air. At ordinary temperatures, nitrogen dioxide is a gas of a dark reddish brown color. When chilled with a freezing mixture, consisting of ice and common salt, the dark brown nitrogen dioxide becomes much lighter in color and condenses to a pale yellow liquid. At 30 this liquid congeals, yielding colorless crystals that melt at 10, thus forming a colorless liquid which is fairly stable even at 0. On gently warming this liquid, it assumes a greenish yellow hue. At about 10 it is yellow in color, at 15 it is orange colored, and at higher temperatures it becomes still darker, till at 26, its boiling point, the color becomes a dark reddish brown. On lowering the temperature, these changes occur in the reverse order. At 2 the vapor is 38 times as heavy as hydrogen, while at 140 the vapor is only 23 times as heavy as hydrogen. At 26, therefore, the molecular weight would be 76, and at 140, 46. Now the formula NO 2 corresponds to a molecular weight of 46, con- sequently at 140 the gas is NO 2 . But the double formula N 2 O 4 corresponds to a molecular weight of 92, so that at 26 the gas has more nearly the formula N 2 Q 4 . The vapor density decreases gradually as the temperature is raised, and all these facts are best explained by assuming that at low temperatures the molecules are N 2 O 4 , which are colorless, and that these decompose gradually, with rise of temperature, into brown molecules of NO 2 , thus: N 2 4 ;2N0 2 . (1 vol.) (2 vols.) ITS OUTLINES OF CHEMISTRY At 26, therefore, the dissociation will have progressed to the extent of about 34 per cent, as may be computed from the densities above given ; while at 140 the dissociation is practi- cally complete. When nitrogen dioxide acts on water in the cold, nitrous and nitric acids result, as already mentioned in connection with nitric acid. On passing nitrogen dioxide through a red-hot tube, it is decomposed into oxygen and nitric oxide. The action is reversible, thus : 2 N0 2 ^ 2 NO + 2 . Nitrogen dioxide is a poisonous gas having a corrosive action on the mucous membranes. It is also a strong oxidizing agent, and consequently will support the combustion of many sub- stances. Nitrous Acid. When potassium nitrate is heated, it loses a portion of its oxygen and is converted into potassium nitrite, thus : 2 KN0 3 = 2 KN0 2 + O 2 . The nitrite may also be formed by heating saltpeter with lead or copper, thus : KN0 3 + Pb' = PbO + KN0 2 . KNO 3 + Cu = CuO + KNO 2 . Potassium nitrite KNO 2 is a salt of nitrous acid HNO 2 . The latter has never been prepared except in dilute solutions at low temperatures. On attempting to isolate nitrous acid from a nitrite by treating with sulphuric acid, the following reaction occurs : 2 KN0 2 + H 2 S0 4 = K 2 S0 4 + H 2 O + NO + NO 2 . It is possible that at first HNO 2 is set free, which undergoes decomposition into nitrogen trioxide, that is, nitrous acid an- hydride N 2 O 8 , and water, thus : 2 HN0 2 = H 2 + N 2 3 . The latter then decomposes into nitric oxide NO and nitrogen dioxide NO 2 , thus : N 2 3 = NO + N0 2 . OXY-ACIDS AND OXIDES OF NITROGEN 179 By dissolving nitrogen trioxide (see below) in water at 0*\ a 'blue solution is obtained which is commonly regarded as a solution of nitrous acid HNO 2 . This solution readily evolves nitric oxide with concomitant formation of nitric acid, thus : 3 HN0 2 = H 2 + 2 NO + HNO 3 . Thus it is evident that nitrous acid is very unstable. Its salts, however, are fairly stable. They may be formed by neutraliz- ing aqueous solutions of nitrous acid with bases, or by reduc- ing nitrates. In rain water and frequently in contaminated drinking water, nitrites are present. Nitrous acid may act as a reducing agent, for it will take up oxygen and form nitric acid. Thus it will reduce a potassium permanganate solution as follows : 2 KMn0 4 + 3 H 2 SO 4 + 5 HNO 2 = K 2 S0 4 + 2 MnS0 4 + 3 H 2 O + 5 HNO 3 . On the other hand, toward substances that will take up oxygen, nitrous acid plays the role of an oxidizing agent. So with hydriodic acid, the following reaction occurs : 2 HN0 2 + 2 HI = 2 NO + 2 H 2 O + I 2 . Nitrous acid is of importance in the study of carbon compounds, and in the preparation of aniline dyes. Nitrites are readily distinguished from nitrates, for nitrites evolve the charac- teristic brown nitrogen dioxide fumes when acidified with sulphuric acid. Furthermore, a dilute solution of a nitrite acidulated with sulphuric acid will turn starch potassium iodide paper blue ; compare the last equation above. Nitrogen Trioxide. Nitrogen trioxide or nitrous anhydride N 2 O 3 readily decomposes into NO and NO 2 . When equal volumes of the latter gases are mixed and cooled to 21 in a tube, nitrous anhydride, a deep blue liquid, is formed. It slowly decomposes even at 21, but at its boiling point the decomposition is more rapid. The dissociation may be indi- cated thus : Hyponitrous Acid. By reducing sodium or potassium nitrate or nitrite with nascent hydrogen formed by the action of 180 OUTLINES OF CHEMISTRY sodium amalgam on the aqueous solution, a salt of hyponitrous acid may be obtained thus : 2 KNO 2 + 4 H = K 2 N 2 O 2 + 2 H 2 O. When the potassium hyponitrite is treated with sulphuric acid, the hyponitrous acid liberated is decomposed into nitrous oxide and water, thus : K 2 N 2 O 2 + H 2 SO 4 = K 2 SO 4 + H 2 O + N 2 O. The reaction is not reversible, and so hyponitrous acid can- not be obtained by dissolving N 2 O in water. Free hyponitrous acid may be obtained by first making the silver salt Ag 2 N 2 O 2 and decomposing this by means of hydrochloric acid, thus forming insoluble silver chloride AgCl and the free acid. The latter may also be obtained by the oxidation of hydroxylamine NH 2 OH by means of nitrous acid, thus : NH 2 OH + HN0 2 = H 2 N 2 2 + H 2 O. The anhydrous acid forms transparent crystalline plates which are highly explosive. On exploding, the acid decomposes into water, nitrogen and oxygen ; while on slow decomposition at room temperatures in aqueous solution nitrous oxide and water are formed. The aqueous solution, however, is more stable. Nitrous Oxide. On heating ammonium nitrate, nitrous oxide N 2 O and water are produced, thus : NH 4 N0 3 =2H 2 + N 2 0. A mixture of ammonium chloride and saltpeter may be sub- stituted for the ammonium nitrate, for thus potassium chloride and ammonium nitrate are formed, and the latter then decom- poses on heating as represented above. Nitrous oxide is a colorless, neutral gas which is 1.52 times as heavy as air. It has a sweetish odor and taste, and when inhaled it produces peculiar symptoms that frequently are accompanied by fits of hysterical laughing; whence its name, laughing gas. On continued inhalation, it produces insensi bility, and hence the gas has been used as an anaesthetic in dental operations. The gas cannot take the place of oxygen in respiration, however, and if inhaled for a long time death results. OXY-ACIDS AND OXIDES OF NITROGEN 181 Nitrous oxide is much more soluble in cold than in warn? water; so at 0, 1.30 volumes are absorbed by 1 volume ol water, while at 25 only 0.59 volume is absorbed. For this reason the gas is collected over warm water. Nitrous oxide boils at 89.5. The solid melts at 102.7. It may be obtained in the market in compressed form in steel cylinders. A glowing splinter burns in nitrous oxide as in pure oxygen. Similarly phosphorus and sulphur burn in nitrous oxide as in oxygen, the products in all cases being oxides and free nitrogen. Nitrous oxide is, however, readily distinguished from oxygen by the fact that nitric oxide and oxygen form red fumes NO 2 , which does not take place when nitrous oxide and nitric oxide are mixed. When mixed with an equal volume of hydrogen, nitrous oxide explodes on ignition with an electric spark, and the volume of nitrogen formed is equal to that of the original hydrogen, thus : H 2 + N 2 = H 2 + N 2 . (1 volume) (1 volume) (liquid) (1 volume) On heating nitrous oxide with sodium (Fig. 41), the following reaction takes place : 2 Na + N 2 = Na 2 + N 2 . (solid) (1 volume) (solid) (1 volume) The volume of nitrogen formed equals that of the nitrous oxide. From this fact and the specific gravities of the gases involved, it follows that the formula of nitrous oxide is N 2 O. While nitrous oxide is a good oxidizing agent, exhibiting many of the properties of oxygen, it is not as energetic as the latter. So metals will not rust in contact with moist N 2 O as in contact with moist oxygen. A feebly burning piece of sulphur or phosphorus will be extinguished in nitrous oxide, though when these substances are burning strongly, they con- tinue to burn brilliantly in the gas. General Considerations. In ammonia, nitrogen has a valence of three, thus : 182 OUTLINES OF CHEMISTRY In ammonium salts it has a valence of five, thus : Cl In hydroxylamine, in hydrazine, in hydrazoic acid, and in nitrous oxide, nitrogen is trivalent, thus : - ,H H H N \ = J N = N NN-N/ ; X N X ; \ o / / >TJ H H' (hydroxylamine) (hydrazine) (hydrazoic acid) (nitrous oxide) At first sight one might be inclined to the view that nitrogen is univalent in N 2 O, and that this compound is analogous to water, the two hydrogen atoms of which are replaced by nitro- gen atoms. However, the ease with which nitrous oxide parts with its oxygen and forms free nitrogen speaks for the above formula. If in nitrous oxide the oxygen is replaced by the bivalent group = NH, called the imide group, hydrazoic acid results. In the salts of hydroxylamine and of hydrazine, nitrogen is quinquivalent as in the ammonium salts. In nitrogen pentoxide and nitric acid nitrogen has a valence of five, thus : N=O ^O ;>0, and N=0 N^Q N) - H In nitrogen dioxide, nitrogen is tetravalent, thus : O = N = O. In nitrogen tetroxide, formed by chilling nitrogen dioxide, nitrogen has been regarded as quinquivalent, thus : s~\ s~\ This formula is not generally accepted, however. OXY-ACIDS AND OXIDES OF NITROGEN 183 In nitrogen trioxide and nitrous acid nitrogen is trivalent, thus : and N OH. In nitric oxide, nitrogen is bivalent, thus: Finally in hyponitrous acid (N-O-H) 2 , as in N 2 O, nitrogen has at times been considered as univalent. However, the com- pound NOH is not known. Attempts to isolate it have always yielded (NOH) 2 , the constitution of which is best represented by regarding nitrogen as trivalent, thus : HO-N = N-OH. From this compound, water readily splits off, forming nitrous oxide thus : N = N NO/ ; The valence of nitrogen thus exhibits a relatively great range of variation in different compounds. We may regard all the oxides and oxy- acids of nitrogen as derived from hypothetical hydroxides by successive elimination of water, as the following table shows, in which all the compounds except those in the first column are known : 2 N(OH) 5 minus 4 H 2 O = 2 HNO 3 ; 2 HNO 3 minus H 2 O = N 2 O 6 . 2 N(OH) 4 minus 4 H 2 O = N 2 O 4 ; N 2 O 4 yields 2 NO 2 . 2 N(OH) 3 minus 2 H 2 O = 2 HNO 2 ; 2 HNO 2 minus H 2 O = N 2 O 3 . 2N(OH) 2 minus2H 2 O = 2NO; ......... 2 NOH ..... N 2 O 2 H 2 ; N 2 O 2 H 2 minus H 2 O = N 2 O. There is also a striking similarity between the oxy-acids of nitrogen and their salts on the one hand, and the oxy-acids of the halogens and their salts on the other hand. This similarity, which is evident from the following table, is not a mere simi- larity of formulae, for the compounds themselves exhibit anal- ogies in their crystal forms, solubility in water, and general 184 OUTLINES OF CHEMISTRY stability on heating and on treatment with reagents. Only known compounds are included in the table. N 2 . . . H 2 N 2 2 . . . Na 2 N 2 2 C1 2 O . . . HC1O . . . NaClO. NO N 2 O 8 . . . HNO 2 .... NaNO 2 NaClO 2 . N0 2 C10 2 N 2 6 . . . HN0 3 NaN0 3 I 2 O 5 . . . HC1O 3 . . . NaClO 3 . C1 2 O 7 . .HC1O 4 ... NaClO 4 . REVIEW QUESTIONS 1. Make a list of the oxy-acids of nitrogen, giving their formulas, also of their anhydrides. What other oxides of nitrogen are there, and what is the valence of nitrogen in each of the nitrogen compounds you have listed? 2. Write the equation for the preparation of nitric acid from Chili saltpeter and compare it with the equation for the preparation of hydro- chloric acid from common salt. 3. If solid saltpeter is treated with hot, concentrated nitric acid, the change proceeds to completion ; why is this not true when a twenty per cent solution of saltpeter is treated with sulphuric acid ? 4. How much nitric acid could be prepared from 100 tons of NaN0 3 ? 5. What are the two main properties of nitric acid ? Illustrate each by means of an example, writing the appropriate chemical equation. 6. How may nitric acid be prepared from the air and water? Equations. 7. Make a list of the uses of nitric acid. 8. Why is hydrogen not liberated when nitric acid acts on metals? Write the equation showing the action of nitric acid on copper. 9. How could one form ammonia from nitric acid? What inter- mediate products might be prepared in this process? 10. What principle of chemical equilibrium is illustrated in making hydrochloric acid, nitric acid, and ammonia? 11. What is aqua regia? Write the equation. Upon what property of nitric acid does its efficiency in aqua regia depend? What is the active agent in aqua regia ? 12. What is meant by "fixation of atmospheric nitrogen"? How may this be accomplished? Why is it of importance commercially? 13. How distinguish between oxygen and laughing gas? Between the latter and nitric oxide ? 14. Show the resemblance of the structural formulas of bromic acid and nitric acid, also of potassium nitrite and potassium chlorite. What is the valence of nitrogen and of the halogen in each of these cases? CHAPTER XIII SULPHUR, SELENIUM, AND TELLURIUM Occurrence and Preparation of Sulphur. Sulphur has been known since ancient times, for it occurs in nature in the uncom- bined state, especially in the vicinity of active or extinct volca- noes. Thus in Italy, Sicily, Spain, Poland, Egypt, Iceland, California, the Yellowstone Park, China, and India, sulphur is found native. As a result of volcanic action sulphur probably is formed by the reduction of hydrogen sulphide H 2 S by sul- phur dioxide SO 2 , thus : 2H 2 S + S0 2 =2H 2 + 3S. Sulphur also occurs in sedimentary deposits, where it is formed as a product of the decay of certain bacteria and algae which are able to store up this substance in their organisms in form of minute particles. This sulphur originates from deposits of gypsum, from which it is liberated as hydrogen sulphide as the result of cellulose fermentation. This hydrogen sulphide is then taken up by the algse and bacteria, which convert it into sulphates ; but in this process they store up a reserve stock of free sulphur in their bodies. The sulphur which is found in sedimentary deposits then really occurs from the oxidation of hydrogen sulphide through the action of these organisms, thus : H 2 S + 0=H 2 + S. Some of the sedimentary sulphur is, however, probably also formed by direct oxidation of hydrogen sulphide by the oxygen of the air. Especially rich sedimentary deposits of sulphur occur in Texas and Louisiana, where by means of superheated steam the sulphur in the lower strata is melted and forced up to the surface in the liquid state. On account of this rich deposit of sulphur, the amount produced in the United States in 1910 was 255,534 long tons, which is about half of the world's annual production of sulphur. 185 186 OUTLINES OF CHEMISTRY Sulphur further occurs as hydrogen sulphide in the waters ol sulphur springs and in the air near active volcanoes, where sul- phur dioxide is also frequently found. In combination with metals, sulphur occurs as sulphides, as in galenite PbS, pyrite FeS 2 , zinc blende ZnS, cinnabar HgS, and copper pyrite CuFeS 9 FIG. 42. It is also found in form of sulphates of various metals. Thus fer- rous sulphate FeSO 4 , lead sulphate PbSO 4 , heavy spar BaSO 4 . are found in nature ; but above all, gypsum CaSO 4 2 H 2 O and anhydrite CaSO 4 are found in very extensive deposits. Thfc amount of gypsum produced in the United States alone in 1910 was 2,379,057 tons. Sulphur occurs in small quantities in com- SULPHUR, SELENIUM, AND TELLURIUM 187 bination with other elements in nearly all plant and animal tissues, for it is a constituent of albumen. So it is found par- ticularly in muscles, hair, nails, hoofs, and horns. In urine sulphur is found as sulphates. In some plants, like mustard, onions, garlic, and skunk cabbages, it enters into odoriferous compounds that have an irritating action on the mucous mem- branes and the skin. The preparation of sulphur from the native deposits consists of melting it out of contact of the air and thus freeing it from the gypsum, calcium carbonate, sand, etc., with which it iscom- "nonly contaminated. Thus, a raw material about 90 per cent pure is obtained, which is placed in cast-iron retorts and distilled (Fig. 42). The vapors enter brick chambers, where they are condensed on the cold walls in form of fine powder which is placed on the market as flowers of sulphur. As the walls finally become hot the sulphur melts and collects on the bottom of the chamber, where it is drawn off from time to time and cast into sticks in moist, wooden, slightly conical molds. In this form it is called roll sulphur or brimstone. Sulphur is also prepared by heating pyrites FeS 2 and condensing the product. It is further prepared from the waste liquors of the Le Blanc soda process (which see), and from the sulphide of iron secured as a by-product in purifying illuminating gas. Properties of Sulphur. Native sulphur and roll sulphur form lemon-yellow crystals, of specific gravity 2.06, belonging to the orthorJiombie system (Fig. 43). When heated, this rhombic sulphur melts at 114.5 to a mobile, light yellow liquid, which on further heating to 160 becomes dark brown and viscous. In the neighborhood of 200 the viscosity is so great that the vessel in which the sulphur is contained may be turned bottom upward without causing the sulphur to run out. On still further heat- ing, the viscosity of the liquid diminishes, but its color remains dark brown. At 400 the liquid is quite mobile, and at 450 it boils, emitting a heavy, dark brown vapor. When sulphuji- is melted in a crucible and the mass is allowed to cool till a crust forms over the top of the liquid, and the latter is then poured out through a hole punctured in the crust^ 188 OUTLINES OF CHEMISTRY it is found that the walls of the crucible are lined with needle- like, almost colorless crystals of sulphur that belong to the monoclinic system. These crystals of monoclinic sulphur melt at 119. They have a specific gravity of 1.96. On standing they very slowly change to crystals of the orthorhombic sys- tem. The rhombic crystals are thus the stable ones at ordinary temperatures, whereas the monoclinic crystals are stable at high temperatures. The temperature at which the transition from the one form to the other takes place is .96. 5; at this point both rhombic and monoclinic sulphur remain side by side in equilibrium with each other without change. Below the tran- sition point all passes over into rhombic sulphur, while slightly above that point all is converted into the monoclinic variety. When sulphur heated almost to the boiling point is poured into cold water, an elastic mass is formed which is called plas- tic sulphur. After a few days it loses its plasticity and be- comes hard, but for a while it remains non-crystalline, that is, amorphous. This amorphous sulphur is practically insoluble in all solvents ; however, it very gradually passes over into rhom- bic sulphur. This is soluble in carbon disulphide to the extent of about 40 parts in 100 at room temperature. On evapora- tion, it may again be obtained from this solution in rhombic form. Rhombic sulphur is also soluble to a slight extent in liquids like alcohol, ether, turpentine, fats, and linseed oil. Flowers of sulphur dissolve only partially in carbon disulphide. They are a mixture of amorphous and rhombic sulphur. Substances which are able to crystallize in tivo different systems are called dimorphous. This property is not uncommon. In passing from the monoclinic to the rhombic form, sulphur slowly evolves heat. Precipitated sulphur, or milk of sulphur, is prepared by add- ing an acid to a polysulphide like K 2 S 5 : K 2 S 5 + 2 HC1 = 2 KC1 + H 2 S + 4 S. Thus formed, it is a grayish white powder, which is used in medicine. Precipitated sulphur is soluble in carbon disulphide. Sulphur is thus a polymorphous substance. The ability of an element to occur in different forms has been called allotro- pism, and so the different forms of sulphur are sometimes called the allotropic forms of sulphur. Their existence has been SULPHUR, SELENIUM, AND TELLURIUM 189 explained by assuming that the molecules of the different modi fications contain a different number of atoms, similar to the case of ox}rgen and ozone. It is doubtful, however, whether the cases are similar. Sulphur is insoluble in water and is devoid of taste and smell. In contact with moist air it very slowly oxidizes super- ficially and passes into solution as sulphuric acid. Sulphur combines with many metals and non-metals, forming sulphides. Heated together with iron or copper, for instance, the union takes place with evolution of light and heat. In the air and in oxygen, sulphur burns to sulphur dioxide SO 2 , which in pres- ence of platinum black will take on more oxygen and form SO 3 . The atomic weight of sulphur is 32.07. Investigations of the vapor density of sulphur show that at diminished pressure and low temperatures the molecular formula of sulphur is S 8 , whereas at 800 to 1000 the density corresponds to the formula S 2 . There is a gradual decomposition of the molecules from S 8 to S 3 as the temperature rises. In the neighborhood of 2000 the S 2 molecules are further largely dissociated into monatomic molecules S. Uses of Sulphur. Sulphur is used in the manufacture of sulphuric acid and sulphur dioxide, the latter being used as a bleaching and disinfecting agent. Sulphur is also used in making black gunpowder, fireworks, vulcanized caoutchouc, and hard rubber. In medicine it is employed as a specific. Crystals and Crystal Systems. Many substances are able to assume the crystalline state. Crystals are generally formed by allowing liquids to congeal or solutions to evaporate to a point at which the dissolved substances separate out. Crystals may, however, also be formed when vapors condense, as in the sublimation of iodine or sulphur ; or they may form gradually from amorphous, solid substances, as in the case of the conver- sion of amorphous sulphur to rhombic sulphur. There are many substances which, like sulphur, are known in both the crystalline and amorphous states; others have never been found in crystalline condition, like cellulose and dextrine ; whereas still others, like water, are always crystalline when solid. Crystalline substances are said to have crystallizing power, whereas those substances that are only known in amor- phous form are said to be devoid of crystallizing power. 190 OUTLINES OF CHEMISTRY FIG. 44. FIG. 45. FIG. 46. FIG. 47. FIG. 48. FIG. 49. FIG. 50. FIG. 61. FIG. 52. / > ^t T<\ FIG. 53. FIG. 54. SULPHUR, SELENIUM, AND TELLURIUM 191 do not know of what this tendency to form crystals really con- sists, much less are we able to measure or compare quantita- tively the crystallizing power of various substances. The most striking external characteristic of a crystal is its regularity of form. A study of crystals has led to the conclu- sion that a crystal is a solid bounded by plane faces which are the outcome of a regular internal arrangement of the molecules. So FIG. 55. FIG. fiff FIG. 57. FIG. 58. FIG. 59. FIG. 60. the hardness, color, index of refraction, crushing strength, resist- ance to corrosion by chemical agents, etc., may vary as different directions in one and the same crystal are considered. It has been found that all known crystals may be classified into six crystal systems, according to their symmetry. All crystals whose faces may be referred to a system of three axes of equal length and at right angles to one another are said to belong to the isometric or regular system. Some com- mon forms are shown in Figs. 44 to 54. These crystals may have nine so-called planes of symmetry, a plane of symmetry being a plane which cuts a crystal in two halves that are to each other as an object is to its reflection in a mirror. Many 192 OUTLINES OF CHEMISTRY substances crystallize in the regular system. Among these are common salt, alum, fluorspar, galena, pyrite, garnet, diamond, gold, silver, mercury, and copper. Crystals whose planes may be referred to a system of three axes, of which but two are of equal length but all at right angles to one another, are said to belong to the tetragonal or quadratic system, in which there are five planes of symmetry possible. Figures 55 to 60 show some common forms of crys- FIG. 64. FIG. 65. FIG. 66. FIG. 67. tals of this system as they occur in rutile, titanium dioxide TiO 2 ; in cassiterite, stannic oxide SnO 2 , the most important ore of tin; and in calomel HgCl. In the hexagonal system, the forms may be referred to four axes, three of which are of equal length, lie in the same hori- zontal plane, and bisect one another in a point so as to form six angles of sixty degrees each. The fourth axis is either longer or shorter than the others, and runs through their point of intersection at right angles to the horizontal plane, which bisects the vertical axis. In this system there are seven pos- sible planes of symmetry. Figures 61 to 67 show some typical SULPHUR, SELENIUM, AND TELLURIUM 193 crystals of the hexagonal system. To it belong the crystal forms assumed by many important substances, like water H 2 O, quartz SiO 2 , calcium carbonate CaCO 3 , Chili saltpeter NaNO 3 , and calcium phosphate Cti 3 (PO 4 ) 2 . The so-called rhombohedral FIG. 68. FIG. FIG. 70. FIG. 71. division of the hexagonal system in particular has many repre- sentatives. It has sometimes been termed a separate system, the trigonal system. Crystal forms that can be referred to a system of three axes, all of which are at right angles to one another but of unequal lengths, are said to belong to the orthorhombic or rhombic sys- tem, in which there are but three possible planes of symmetry. FIG. 72. FIG. 74. Figures 68 to 71 exhibit some typical rhombic forms. In this system crystallize many substances, like sulphur, iodine, olivine Mg 2 SiO 4 , saltpeter KNO 3 , heavy spar BaSO 4 , and magnesium sulphate MgSO 4 7 H 2 O. In the monoclinic or monosymmetric system the forms are referred to a system of three axes all of which are of unequal length. The two axes that lie in the vertical plane bisect each other at right angles, and the third axis is bisected at the point 194 OUTLINES OF CHEMISTRY of intersection of the other two, but it does not make a right angle with the plane in which the other two axes lie. The angle which it makes with that plane varies in different crystals. In this system there is but one plane of symme- try. Figures 72 to 74 show some representative morioclinic forms. Many compounds crystallize in this system, among which are monoclinic sulphur, gypsum, feldspar, cane sugar, Glauber's salt Na 2 SO 4 - 10 H 2 O, copperas FeSO 4 7 H 2 O, and borax Na 2 B 4 O 7 10 H 2 O. Finally, in the triclinic or asymmetric system the forms are referred to three unequal axes bisecting one another in a point at angles that are unlike and not right angles. In this system there is no symmetry whatever. Figures 75 and 76 show some triclinic forms. Copper sulphate CuSO 4 5 H 2 O, plagioclase feldspar NaAlSi 3 O 8 (albite), and CaAl 2 Si 2 O 8 (anorthite) crystallize in the triclinic system. Under the same conditions a chemical substance always crystal- lizes in the same system. Most substances crystallize in but one system. However, we have, seen that under different conditions one and the same substance may crystallize in two different systems. This prop- erty is called dimorphism. Thus, sulphur may form rhombic or monoclinic crystals ; calcium carbonate CaCO 8 may form hex- agonal or rhombic crystals ; iron pyrites may form isometric or rhombic crystals. These substances are consequently dimor- phous. Substances that have similar chemical composition generally crystallize in the same system and exhibit the same forms. This is the law of isomorphism, discovered by Eilhard Mitscherlich. So, for instance, the carbonates CaCO 3 , FeCO 8 , MgCO 3 are rhombohedral ; the chlorides NaCl, KC1, NH 4 C1 are isometric. Hydrogen Sulphide. This is by far the most important com- pound which sulphur forms with hydrogen. The elements unite directly with each other at higher temperatures, forming the compound whose composition and vapor density are repre- SULPHUR, SELENIUM, AND TELLURIUM 195 sen ted by the formula H 2 S. So when a current of hydrogen is passed over heated sulphur in a tube, H 2 S is formed; also when certain sulphides are similarly heated in a current of hydrogen, thus : The common way of preparing the gas consists of treating fer- rous sulphide FeS (made by heating sulphur and iron together) with either dilute sulphuric or hydrochloric acid ; FeS + H 2 S0 4 = FeS0 4 + H 2 S. FeS + 2 HC1 = FeCl 2 + H 2 S. Instead of ferrous sulphide, which is the cheapest, other sul- phides might be employed. The gas may also be prepared by reduction of sulphuric or sulphurous acids with nascent hydio- gen: H 2 SO 3 + 6 H = 3 H 2 O + H 2 S. In nature hydrogen sulphide occurs in sulphur springs, vol- canic gases, and wherever organic matter is decomposing, as in sewer gas, in the intestinal gases, and in some pathological cases in urine. Hydrogen sulphide is a colorless gas which is 1.19 times as heavy as air. It boils at 62 and melts at 86. The gas has a very disagreeable odor, being that of rotten eggs, in which it is contained. Hydrogen sulphide is a very poisonous gas and overcomes persons and animals suddenly, in which respect it resembles hydrocyanic acid. Inhaled in small amounts, hydrogen sulphide produces headache and at times vomiting. The gas is combustible, burning with a blue flame to water and sulphur dioxide : 2 H 2 S + 3 O 2 = 2 H 2 O + 2 SO 2 . In an insufficient amount of oxygen, the products are, in part, water and sulphur : 2 H 2 S + 2 = 2 H 2 + 2 S. In water, hydrogen sulphide is but slightly soluble, about 3 volumes being absorbed by 1 volume of water at ordinary temperature and pressure. On boiling this solution, all the gas escapes. On standing exposed to the air, the gas in the 196 OUTLINES OF CHEMISTRY solution is gradually oxidized to water and sulphur which separates out in the form of a precipitate. When chlorine, bromine, or iodine act on hydrogen sulphide, the latter is decomposed, sulphur being liberated and hydro- halogen being formed, so for instance : H 2 S + I 2 = 2 HI + S. The aqueous solution of hydrogen sulphide is feebly acid toward litmus, and in many ways it deports itself like a weak acid. So it will react with metals even at room temperature, forming sulphides and hydrogen, thus : 2 Ag + H 2 S = Ag 2 S + H 2 . Pb + H 2 S = PbS + H 2 . Furthermore it reacts with many basic oxides and hydroxides, thus : PbO + H 2 S = PbS + H 2 0. 2 NH 4 OH + H 2 S = (NH 4 ) 2 S + 2 H 2 O. KOH + H 2 S = KSH + H 2 O. 2 KOH + H 2 S = K 2 S + 2 H 2 O. The sulphides of sodium and potassium show a strong alkaline reaction toward indicators. They are salts of a very weak acid with a strong base, and hence are decomposed by water by hydrolysis. The reaction, which is reversible, may be written thus : When passed through a red-hot tube, hydrogen sulphide is decomposed to hydrogen and sulphur. It thus parts readily with its hydrogen, and is consequently a good reducing agent, as is evident, for instance, from the fact that it will reduce sulphuric or nitric acid, thus : H 2 SO 4 + H 2 S = 2 H 2 O + SO 2 + S. 2 HN0 3 + 3 H 2 S = 4 H 2 O + 2 NO + 3 S. Hydrogen sulphide is a very important reagent in chemical analysis, for while the sulphides which it forms with metals like sodium, potassium, calcium, and magnesium are soluble in water, other sulphides like those of iron, zinc, and nickel are not soluble in water, but soluble in dilute acids, and still other sulphides like those of arsenic, copper, and lead are insoluble SULPHUR, SELENIUM, AND TELLURIUM 197 both in water and dilute acids. A very careful study of these and similar properties of the sulphides of the metals has led to a system by means of which the metals can be detected and separated when they occur together. Polysulphides and Hydrogen Persulphide. When sulphur is added to a solution of sulphide of potassium, sodium, calcium, ammonium, etc., it dissolves, forming polysulphides. Thus, with K 2 S sulphur may form compounds varying in composition from K 2 S to K 2 S 5 according to the amount of sulphur dissolved. When such a persulphide is gradually added to a very dilute solution of hydrochloric acid, a thick, yellow oil of disagree- able odor separates out which has the composition H 2 S 5 , no matter what the sulphur content of the poly sulphide was, thus: 2 K 2 S 3 + 4 HC1 = 4 KC1 + H 2 S + H 2 S 6 . 4 Na 2 S 2 + 8 HC1 = 8 NaCl + 3 H 2 S + H 2 S 6 . Hydrogen persulphide bleaches organic dyestuffs. It reacts with iodine, forming hydriodic acid and sulphur. It gradually decomposes into hydrogen sulphide and sulphur on standing. Comparison of Hydrogen Sulphide with Water. It is evident that hydrogen sulphide and water possess many points of analogy. Thus the one is H-S-H and the other H-O-H. With the univalent metals they form hydrosulphides MSH and hydroxides MOH, respectively; furthermore, the corresponding sulphides M 2 S, and oxides M 2 O, are also known. With ele- ments of higher valence, analogous sulphides and oxides are formed. Thus we have FeS and FeO, P 2 O 6 and P 2 S 6 , Sb 2 O g and Sb 2 S 3 , etc. Again, just as oxygen and hydrogen form a peroxide H 2 O 2 , so sulphur and hydrogen form a persulphide, which, to be sure, has the composition H 2 S 5 . We shall later see further points of resemblance between oxygen and sulphur in their chemical behavior. The two elements indeed belong to the same family group. Compounds of Sulphur with the Halogens. Fluorine unites directly with sulphur to form sulphur hexafluoride SF 6 , which consists of white crystals that melt at 55. The substance boils but slightly above its melting point. The gas is colorless, odorless, tasteless, and practically as indifferent toward othei reagents as nitrogen. 198 OUTLINES OF CHEMISTRY When dry chlorine is passed over molten sulphur in a tubu- lated retort, sulphur monochloride S 2 C1 2 , boiling at 138, is formed. It is a fuming yellowish red liquid of suffocating odor. Its specific gravity is 1.7. It dissolves sulphur readily^ solutions containing over 60 per cent sulphur being obtainable. For this reason sulphur monochloride is used in preparing vulcanized rubber. Water decomposes sulphur monochloride, thus : 2 S 2 C1 2 + 2 H 2 O = 4 HC1 + SO 2 + 3 S. Sulphur dichloride SC1 2 is formed when sulphur monochlo- ride is saturated with chlorine in the cold. It is an oil of reddish brown color and specific gravity 1.6. It readily decomposes at 64, yielding sulphur and sulphur monochloride. It is also de- composed by water, thus: 2 SC1 2 + 2 H 2 O = 4 HC1 + SO 2 + S. Sulphur fetrachloride SC1 4 is formed by saturating sulphur dichloride with chlorine at temperatures below 25. The substance forms crystals which melt at 30. It readily dissociates above 22, the decomposition being practically complete at +6. With water it reacts violently, thus: SC1 4 + 2 H 2 = S0 2 + 4 HC1. With bromine, sulphur forms sulphur monobromide S 2 Br 2 , a brownish red liquid which congeals at 46 and boils at about 200, accompanied by partial decomposition. With iodine, sulphur forms sulphur monoiodide S 2 I 2 , consist- ing of dark grayish crystals melting at 60, and also sulphur hexaiodide SI 6 , which forms dark crystals that readily decom- pose on standing, yielding free iodine. Sulphur Dioxide and Sulphurous Acid. When sulphur is burned in the air or in oxygen, the following reaction takes place : S + 2 = S0 2 . The resulting sulphur dioxide occupies the same volume as the oxygen, which may be demonstrated by means of the apparatus of Victor Meyer shown in Fig. 77. The sulphur is burned in oxygen, with which the flask has been filled. On cooling, the manometer indicates that the volume of the gas in the appara- tus has not changed. SULPHUR, SELENIUM, AND TELLURIUM 199 Sulphur dioxide is a colorless gas of suffocating odor. It is 2.21 times heavier than air. It may readily be condensed to a liquid at ordinary pressure by cooling to 10. Under a pres- sure of about two atmospheres it may be liquefied at room tem- peratures. The liquid boils at 8, and the solid melts at 76. Sulphur dioxide will- not support combustion ; never- theless, at higher temperatures many metallic oxides unite vigorously with it with evolu- tion of light, thus : Pb0 2 + S0 2 = PbS0 4 . FIG. 77. Besides being produced by the burning of sulphur, sulphur dioxide is formed by heating sulphides of certain metals in the air ; thus, pyrite acts as follows : 2 FeS 2 + 11 O = Fe 2 O 3 + 4 SO 2 . In the laboratory, sulphur dioxide is commonly made by heating copper turnings with concentrated sulphuric acid : 2 H 2 SO 4 Cu = 2 H 2 O CuSO SO 2 . It may also be formed by heating concentrated sulphuric acid with carbon or sulphur : 2 H 2 S0 4 + C = 2 H 2 + C0 2 + 2 SO 2 . 2 H 2 S0 4 + S = 2 H 2 + 3 SO 2 . When dilute sulphuric acid acts on sulphites, sulphur dioxide is formed ; also when metallic oxides are heated with sulphur: NaHSO 3 + H 2 SO 4 = NaHSO 4 + SO 2 2 MnO 2 + 48=2 MnS + 2 SO 2 . H 2 O. 2 CuO + 2 S = Cu 2 S + SO 2 . In the presence of water, sulphur dioxide bleaches many organic coloring matters. Figure 78 shows the bleaching of flowers by sulphur dioxide evolved by burning sulphur. This bleaching OUTLINES OF CHEMISTRY does not depend upon the oxidation of the dyes, but rathei upon their union with the sulphur dioxide, for on warming some of the articles thus bleached their color may be restored. In other cases, the bleaching action depends upon the subtrac- tion of oxygen from the sub- stances. Sulphur dioxide is used to bleach silk, wool, straw, and other fibers that would be de- stroyed by means of chlorine. It is also used as an antiseptic and disinfectant, for it is a powerful germicide. For these purposes it may now be obtained in liquid form in tin cans. About 50 volumes of sulphur dioxide are dissolved by 1 volume of water at 15, while at 40 but 18.8 volumes are thus absorbed. From the solution all of the sul- FlG - 78 ' phur dioxide may be expelled by boiling. The solution reacts acid and behaves as though it contained sulphurous acid H 2 SO 3 , but this substance has never been isolated, thus : SO H 2 O = H 2 SO 3 . With bases, sulphurous acid forms salts called sulphites, thus : - H 2 SO 3 NaOH = NaHSO H 2 O. H 2 SO 3 + 2 NaOH = Na 2 SO 3 + 2 H 2 O. H 2 SO 3 + Ca(OH) 2 = CaSO 3 + 2 H 2 O. Sulphurous acid is dibasic in character. Both the acid and the normal sulphites of the alkali metals are soluble in water, but other normal sulphites are sparingly soluble. From sul- phites, sulphur dioxide may readily be regenerated by addition of sulphuric or hydrochloric acid. This fact is used in the detection of sulphites in chemical analysis. Sulphur dioxide is a reducing agent, which property is pos- sessed in a still greater degree by its aqueous solutions. This is because sulphurous acid is able to take up additional oxy- gen readily, thus passing over into sulphuric acid. Even the SULPHUR, SELENIUM, AND TELLURIUM 201 oxygen from the air slowly converts sulphurous acid into sul- phuric acid in solution, thus : 2H 2 SO 3 +O 2 =2H 2 SO 4 . Chlorine, bromine, or iodine rapidly change sulphurous acid into sulphuric acid, thus : H 2 SO 3 + H 2 O + C1 2 = H 2 SO 4 + 2 HCL H 2 S0 3 + H 2 + I 2 = H 2 S0 4 + 2 HI. Sulphur Sesquioxide. This compound has the composition S 2 O 3 . It may be prepared by treating molten sulphur trioxide SO 3 with pulverized sulphur. The product consists of bluish green crystals. With fuming sulphuric acid it forms a blue solution. Water decomposes the sesquioxide into sulphuric acid and sulphur. Sulphur Trioxide and the Contact Process of making Sulphuric Acid. Sulphur trioxide SO 3 is formed by heating sulphates of many of the heavy metals, thus : Fe 2 (S0 4 ) 8 = Fe 2 3 +3S0 3 . Oxygen unites but very slowly with SO 2 to form SO 3 , in spite of the fact that the union is accompanied with considerable evolution of heat. However, when a mixture of sulphur diox- ide and oxygen is passed over finely divided platinum, the union readily takes place, the action being practically complete at 450. In this process, the platinum remains unchanged. It acts as a contact or catalytic agent. In place of finely divided platinum, ferric oxide or chromic oxide will also serve. The residues of the oxides obtained by roasting pyrites are some- times used for this purpose. The sulphur dioxide obtained by burning sulphur or roasting native sulphides, generally pyrites, is mixed with air in such proportion that there is present a large excess of oxygen beyond what is needed to produce sul- phur trioxide according to the equation : 2.S0 2 +0 2 ^2S0 3 ; for this, reaction is a reversible one and the presence of the excess of oxygen, according to the law of mass action, displaces the equilibrium toward the right. The temperature should be held at about 400 to 450, for at higher temperatures the sul- phur trioxide dissociates, that is, the action reverses. The CALIFORNIA COLLE6E of PHARMACY 202 OUTLINES OF CHEMISTRY gases should be purified. It is especially necessary that they be freed from dust and from arsenic. The latter is generally present in the gases and is removed by means of steam. Both the residues from roasting pyrites, and platinized asbestus are used at present in thus preparing sulphur trioxide by what is known as the "contact process." The bulk of this sulphur trioxide formed is used in making sulphuric acid, and to this end it is absorbed in sulphuric acid of 97 to 98 per cent strength. The strength of the acid is regulated by addition of water. Enormous quantities of sulphuric acid are now pre- pared annually by the contact process, both in Europe and FIG. 79. America; and this method, the success of which on a commer- cial scale is due to the labors of Knietsch (1901), has to a large extent displaced the lead chamber process for making sulphuric acid, at least so far as making concentrated sulphuric acid is concerned. On a small scale, in the laboratory, sulphur trioxide can readily be made by means of the apparatus shown in Fig. 79. Sulphur dioxide from a generator and oxygen from a tank are passed into the wash-bottle JB; the mixed gases then pass through the drying tube T, filled with pumice soaked in sul- phuric acid, and finally enter the tube containing the asbestus, which contains finely divided platinum heated to 400. The SO 3 formed is condensed in the receiver. Sulphur trioxide is also formed by heating fuming sulphuric SULPHUR, SELENIUM, AND TELLURIUM 203 acid or warming concentrated sulphuric acid with phosphorus pentoxide, or by heating sodium or potassium pyrosulphate, thus : H 2 S 2 O 7 = H 2 SO 4 + SO 3 . H 2 S0 4 + P 2 5 = S0 3 + 2 HP0 3 . K 2 S 2 7 = K 2 S0 4 4-S0 3 . Sulphur trioxide forms long, colorless, prismatic crystals that melt at 14.8, forming a colorless, mobile liquid that boils at 46. At 20 the specific gravity is 1.97. Below 27 sulphur trioxide forms sulphur hexoxide S 2 O 6 , the crystals of which look like long-fiber asbestus and melt at 50. On further heat- ing, it passes over into vapors that are identical with those of SO 3 , i.e. it dissociates into SO 3 , which on cooling yields a liquid boiling at 46. Sulphur trioxide has great affinity for water. It fumes strongly in the air, and unites with water with great avidity and liberation of much heat which forms steam, causing a hissing noise as the substance is brought into contact with water. It is dangerous to bring large quantities of sulphur trioxide into contact with water at once, for the heat liberated causes explosions. At temperatures above 600 sulphur triox- ide dissociates into sulphur dioxide and oxygen, the reaction being practically complete at 1000. Sulphuric Acid and the Lead Chamber Process. Sulphuric acid H 2 SO 4 has been known for a long time. The alche- mists prepared it by heating ferrous sulphate, green vitriol FeSO 4 7 H 2 O, hence the name oil of vitriol. This process was described by Basil Valentine in 1450, who also prepared the acid by burning sulphur in presence of saltpeter. In 1746 Roebuck, in England, made use of the principle of the latter, method by burning sulphur mixed with saltpeter in closed leaden chambers in presence of moisture which absorbed the gases, forming sulphuric acid. By admitting more air to the chambers, and burning more sulphur in them, additional sul- phuric acid was formed, and so on. This process was the beginning of what is to the present day known as the lead chamber process of the manufacture of sulphuric acid. In its essence the method consists of oxidizing sulphurous acid H 2 SO 3 to sulphuric acid H 2 SO 4 , by means of nitric a< id and its decomposition products. 204 OUTLINES OF CHEMISTRY In practice, the manufacture of sulphuric acid by the lead chamber process involves: (1) The burning of sulphur to sul- phur dioxide, either by using sulphur or commonly by roasting native sulphides like pyrite FeS 2 , copper pyrite, CuFeS 2 , gale- nite PbS, zinc blende ZnS ; (2) the oxidation of the sulphur dioxide in presence of water by means of nitric acid and nitro- gen dioxide, one of its decomposition products ; (3) the oxi- dation of the nitric oxide NO formed by the reduction of the nitric acid and NO 2 ; and (4) the concentration of the sul- phuric acid obtained. In the roasting of the native sulphides mentioned, the latter are heated in a current of air, whereby sulphur dioxide and the oxides of the metals result. The nitric oxide is oxidized to NO 2 by means of oxygen of the air. We may write the chemical changes involved as follows : (1) S + 2 =S0 2 . (2) 3 SO 2 + 2 H 2 O + 2 HNO 3 = 3 H 2 SO 4 + 2 NO. (3) 2 NO + H 2 + 3 O = 2 HNO 3 , and (4) NO + O = NO 2 . (5) S0 2 + H 2 + N0 2 = H 2 S0 4 + NO. Thus it will be seen that when nitric acid acts on sulphur diox- ide in presence of moisture (equation 2), sulphuric acid and nitric oxide result. The latter is then oxidized by oxygen from the air, in part to nitric acid (equation 3), and in part to nitrogen dioxide (equation 4). The nitric acid so formed then reacts with more sulphur dioxide, according to equa- tion (2), and the nitrogen dioxide oxidizes sulphurous acid according to equation (5), the nitric oxide NO formed in both cases being again oxidized by oxygen, and. in turn reduced by sulphurous acid with concomitant formation of sulphuric acid, and so on. While the above equations may be used to represent what occurs in the manufacture of sulphuric acid, the actual process is no doubt of more complicated character. It has been studied by various investigators, among whom George Lunge holds that a compound HO -SO 2 -O(NO), nitrosyl sulphuric acid, is formed in the chambers during the process, and that this compound is then decomposed by water with resulting formation of sulphuric SULPHUR, SELENIUM, AND TELLURIUM 205 acid HO SO 2 OH. The reactions involved in this explanation are : (1) S0 2 + HN0 8 = HO - S0 2 0(NO). (2) The nitrosyl sulphuric acid is then again decomposed by water., according to equation (2), and so on. In nitrosyl sulphuric acid we have the univalent -N=O group, which takes the place of one of the hydrogen atoms in sulphuric acid. Now, in the ordinary manufacture of sulphuric acid, when things are running properly, the formation of nitrosyl sulphuric acid, which consists of colorless crystals known as "chamber crys tals," is not observed. It is only when the supply of water is deficient that these crystals are actually formed, for they are decomposed by water, as stated above. Although there is dif- ference of opinion as to what actually occurs in the details of the sulphuric acid manufacture, the changes in which process are undoubtedly rather. complicated, it nevertheless is certain that by this process sulphurous acid is completely and economi- cally converted into the end product, sulphuric acid. The oxides of nitrogen can be used over and over again, though of course there is always some loss of the latter that must be replenished. The accompanying Fig. 80 shows in diagrammatic form the essentials of a lead chamber sulphuric acid factory. In the furnaces F, the pyrites and other native sulphides are roasted in a current of air. The sulphur dioxide thus produced contains dust carried along mechanically, which deposits in a special long dust flue in which the gas is also mixed with air in proper proportion. The gases, which are at a temperature of about 300, then pass into the Glover tower 6r. This is a structure about 10 meters high and 3 meters in diameter, lined inside with sheet lead and filled with acid proof stones, over which dilute sulphuric acid containing oxides of nitrogen in solution contin- ually trickles from the reservoir on top of the tower. This acid is derived from the Gay-Lussac tower and from the chambers, and contains also some nitric acid, which has been added to replace the oxides of nitrogen that are inevitably lost during the process of manufacture. As the hot gases from the furnaces come into contact with this sulphuric acid of the Glover tower, 206 OUTLINES OF CHEMISTRY SULPHUR, SELENIUM, AND TELLURIUM 207 they are gradually cooled till they attain a temperature of about 70 when they reach the top. At the same time, the acid is heated up and thus concentrated, water being lost which is car- ried off with the gases in form of steam. Again, practically all of the oxides of nitrogen are carried off by the gases, which when they leave the tower pass into the first lead chamber laden with oxides of nitrogen and water vapor. The acid which flows from the bottom of the Glover tower contains only traces of oxides of nitrogen and is about 80 per cent strong. There are com- monly three lead chambers, so connected that the gases enter at the top of each and pass out at the bottom. In these cham- bers, which often have a volume of 1000 cubic meters each, the reactions above' mentioned take place. In the first and second chambers, water vapor is added to the gases. This is done either by blowing in steam from the boiler, or by forcing water into the chambers in form of a spray.- In the third chamber the gases are cooled, and they then pass (charged with oxides of nitrogen regenerated during the formation of sulphuric acid in the chambers) into the bottom of the Gay-Lussac tower. This is lined with lead and filled with coke over which 80 per cent sulphuric acid continually trickles from the tank at the top of the tower L. This 80 per cent acid is obtained from the reser- voir at the bottom of the Glover tower, from which place it is forced through a lead pipe P to the top of the Gay-Lussac tower. In the latter the 80 per cent acid dissolves practically all the oxides of nitrogen, and the residual gases, consisting mainly of nitrogen, leave the top of the tower and pass into a large chimney which keeps up a sufficient draught. The acid drawn from the bottom of the Gay-Lussac tower is thus strongly charged with oxides of nitrogen. It is the purpose of this tower to preserve these oxides. This acid, together with some of the chamber acid, is used again in the Glover tower as already explained. The acid produced in the chambers is known as "chamber acid." It is about 60 to 70 per cent strong, i.e. of specific gravity of about 1.5 to 1.6. The acid may be further concentrated by evaporation in leaden pans to 78 per cent. Stronger acid than this attacks lead too much, and so the 78 per cent acid must be further concentrated by evaporation either in cast-iron, glass, or platinum vessels. The chamber acid is commonly used directly in the manufacture of so-called " superphosphate " fertilizers, VVti OUTLINES OF CHEMISTRY and the acid from the bottom of the Glover tower is employed in the Le Blanc soda process. The concentrated sulphuric acid on the market has a specific gravity of from 1.83 to 1.84, and hence contains from 93 to 98 per cent of H 2 SO 4 . In making concentrated sulphuric acid, the contact process already described obviously has distinct advan- tages, and it is fast taking the place of the lead chamber method. The latter will, however, very likely continue to serve to pre- pare the more dilute acid, for which purpose it is well adapted. The amount of sulphuric acid produced in the world annually is over four million tons. The material is used in making soda, aniline dyes, fertilizers, and explosives like gun cotton, nitro-powder, and dynamite. Again, it is used in storage batteries, in converting starch to sugar in the glucose industries, in refining petroleum, in making alum, copper sulphate, and many other sulphates that are used in medicine and in the arts. Properties of Sulphuric Acid. Sulphuric acid is a colorless, odorless, heavy, oily liquid of specific gravity 1.8384 at 15. It has a very great affinity for water, with which it unites with great evolution of heat. For this reason the acid, when it is to be diluted with water, must always be poured gradually into an excess of water. It is dangerous to proceed in the reverse manner, that is to pour the water into the acid, for the great amount of heat suddenly liberated is very apt to lead to explo- sions throwing the acid out of the container. On account of its powerful affinity for water, sulphuric acid exercises a destruc- tive action upon all plant and animal tissues, for it abstracts hydrogen and oxygen from them in proportions to form water, thus leaving a dark brown or black, charred mass behind. So wood, sugar, cork, muscular tissues, etc., are charred by sul- phuric acid. When sulphuric acid is mixed with water a very appreciable contraction occurs ; thus 500 cc. sulphuric acid mixed with 500 cc. water yield a mixture that has a volume of 971 cc. On account of its affinity for water, concentrated sul- phuric acid is very often used as a drying agent in various chemical operations, particularly in drying certain gases that are not affected by the acid. The commercial sulphuric acid commonly contains lead sul- phate, arsenic, and oxides of nitrogen. By distilling it from retorts of platinum it may be purified. When pure anhydrous SULPHUR, SELENIUM, AND TELLURIUM 209 sulphuric acid (that is, H 2 SO 4 , also called the monohydrate because it is H 2 O-SO 3 ) is heated, it begins to fume at about 150 because of the escape of SO 3 . At 338 the acid boils and the distillate contains 1.5 per cent water. This 98.5 per cent acid thus has a constant boiling point and cannot be further concentrated by fractional distillation. At 85. mm. pressure, pure H 2 SO 4 boils without decomposition at 145-146. The monohydrate H 2 SO 4 melts at +10. The crystals are colorless and may be freed from adhering sulphuric acid by means of a properly constructed centrifugal machine. Sulphuric acid is a very strong dibasic acid. It is capable of forming acid sulphates, like NaHSO 4 , and normal sulphates, like Na 2 SO 4 . As it is also non-volatile except at comparatively high temperatures, it is very often used in liberating other acids from their salts. Besides acting as an acid, sulphuric acid may also play the role of an oxidizing agent toward many substances. So by means of hydrogen it may be reduced to sulphurous acid. When the metals act on sulphuric acid, the hydrogen liberated reduces the acid when the latter is used in concentrated form, sulphates and sulphurous acid being formed simultaneously. The sulphurous acid formed may, of course, be reduced still further. Gold and platinum do not act on sulphuric acid. The other metals react with it under certain conditions, forming sulphates. Dilute sulphuric acid acts readily on some metals, like zinc and magnesium, at room temperatures liberating the hydrogen, as was pointed out when the latter element was studied. Upon other metals, like copper and lead, for instance, sulphuric acid acts but slightly. Even hot, fairly concentrated sulphuric acid, as we have seen, does not attack lead much. This is due in part to the fact that the lead sulphate formed is difficultly soluble in sulphuric acid and so forms a protective coating on the lead. On the other hand, copper acts on hot concentrated sulphuric acid, forming copper sulphate and sul- phur dioxide. By means of hydrobromic or hydriodic acid, sulphuric acid is readily reduced to sulphurous acid and to hydrogen sulphide. The sulphates are all soluble in water except the sulphate of barium. The sulphates of lead, stron- tium, and calcium are sparingly soluble in water. As a rule, sulphates are not as soluble as chlorides and nitrates. Sul- phates of the alkalies are quite stable at high temperatures. 210 OUTLINES OF CHEMISTRY Sulphates of the heavy metals decompose at high temperatures, yielding oxides of the metals and sulphur trioxide. Hydrates of Sulphuric Acid. Pure H 2 SO 4 is commonly called the monohydrate, as stated above. When one molecule of water is added to it, it forms crystals of the composition H 2 SO 4 .H 2 O or H 4 SO 5 , which melt at 8. These are called the dihydrate. By a further addition of a molecule of water a trihydrate H 2 SO 4 -2H 2 O or H 6 SO 6 , also called orthosulphuric or normal sulphuric acid, is formed. It is evident that it may be regarded as S(OH) 6 , in which sulphur is combined with six hydroxyl groups. The trihydrate does not form crystals, except at ver}- low temperatures. Its existence is largely based upon the fact that it represents the composition of the compound formed when sulphuric acid and water react with maximum contraction of volume. There are no salts of either H 4 SO 5 or H 6 SO 6 known. In all its salts sulphuric acid is distinctly dibasic. Pyrosulphuric Acid. When sulphur trioxide is dissolved in pure sulphuric acid, pyrosulphuric acid or disulphuric acid H 2 S 2 O 7 is formed. It consists of crystals that melt at 36, and is sometimes called solid sulphuric acid. This acid fumes strongly in the air. The fuming sulphuric acid of commerce consists of sulphuric acid containing varying amounts of sul- phur trioxide in solution. An acid containing 10 to 20 per cent of additional SO 3 in solution used to be called Nordhausen sulphuric acid. It was prepared by Basil Valentine at Erfurt in 1450 by heating partially dehydrated sulphate of iron. From pyrosulphuric acid, sulphur trioxide may readily be pre- pared by heating. The so-called "oleum" of commerce con- sists of about 80 per cent SO 3 and 20 per cent H 2 SO 4 . It is used industrially. The salts of pyrosulphuric acid are called the pyrosulphates. They are readily prepared by heating acid sulphates, thus : KHSO 4 ^ K 2 S 2 O 7 + H 2 O. The water escapes as vapor. On moistening the pyrosulphate with water, the acid sulphate is again obtained, so that the above reaction is reversible. Thiosulphates. When a solution of a sulphite is boiled with sulphur, a thiosulphate results : - SULPHUR, SELENIUM, AND TELLURIUM 211 We may look upon this salt as sodium sulphate in which one oxygen atom is replaced by a sulphur atom, whence the name thiosulphate. Sodium thiosulphate is used in photography, and in commerce it is frequently called hyposulphite of soda or "hypo." These names are not in accord with chemical usage, since the salt is not a salt of an acid containing less oxygen than sulphurous acid H 2 SO 3 . By treating a thiosulphate with hydrochloric acid, the chloride of the metal, sulphur, sulphur dioxide, and water are formed, thus : Na 2 S 2 3 + 2 HC1 = 2 NaCl + S + SO 2 + H 2 O. Thiosulphuric acid H 2 S 2 O 3 is not known in the free state. Its salts are very common, but attempts to isolate the acid fail because it decomposes into the products indicated by the above equation. Per sulphates. By electrolyzing a concentrated solution of acid potassium sulphate, potassium persulphate KSO 4 is readily obtained. Sodium persulphate may be similarly prepared. It is used in photography. Persulphuric acid HSO 4 is unstable. It may be prepared by dissolving its anhydride, S 2 O 7 , sulphur peroxide, in water, thus : S 2 7 + H 2 = 2 HSO 4 . Sulphur peroxide or heptoxide wa& formed by Berthelot by the action of the silent electric discharge on a mixture of sulphur dioxide and oxygen. It is unstable, and but little is known about it. Persulphuric acid is formed to a slight extent in the lead storage cells, in which sulphuric acid of specific gravity 1.2 is commonly used. Polythionic Acids. Polythionic acids contain more than one sulphur atom. Of these thiosulphuric acid H 2 S 2 O 3 is the simplest example. The following acids are known : Thiosulphuric Acid H 2 S 2 O 3 , forms thiosulphates, like Na 2 S 2 O 3 . Dithionic Acid H 2 S 2 O 6 , forms' dithionates, like Na 2 S 2 O 6 . Trithionic Acid H 2 S 3 O 6 , forms trithionates, like Na 2 S 3 O 6 . Tetrathionic Acid H 2 $ 4 O 6 , forms tetrathionates, like Na 2 S 4 O 6 . Pentathionic Acid H 2 S 6 O 6 , forms pentathionates, like Na 2 S 5 O 6 . With the exception of thiosulphuric acid (which is known only in form of salts), the free acids are known only in aqueous 212 OUTLINES OF CHEMISTRY solutions ; and even in these they readily decompose. The salts, however, are as a rule quite stable. Thionyl Chloride. Thionyl chloride SOC1 2 is formed when phosphorus pentachloride acts on sulphur dioxide, or on a sulphite, thus : PC1 5 + SO 2 = POC1 3 + SOC1 2 . 2 PC1 5 + K-jSOg = 2 POC1 3 + 2 KC1 + SOC1 2 . It is a colorless liquid of very pungent odor. It fumes in the air and is readily decomposed by water, thus : SOC1 2 +H 2 = S0 2 + 2HC1. Thionyl chloride boils at 78. Its specific gravity at is 1.676. It may be regarded as SO 2 with one oxygen atom replaced by two chlorine atoms. Sulphuryl Chloride. This compound is made by the action of equal volumes of chlorine and sulphur dioxide on each other in sunlight, or in presence of a little camphor, thus : SO 2 +C1 2 = SO 2 C1 2 . It may be regarded as SO 3 with one oxygen atom replaced by two chlorine atoms. It is a colorless liquid of very pungent odor. It boils at 70, and has a specific gravity of 1.66 at 20. In contact with the air it fumes strongly. By addition of one gram-molecule of water to one gram-molecule of sulphuryl chloride, chlorsulphonic acid is formed, thus : SO 2 C1 2 + H 2 O = SO 2 Cl OH + HC1. Chlorsulphonic acid SO 2 -C1-OH may be regarded as sulphuric acid SO 2 (OH) 2 with one OH group replaced by chlorine. With more water, chlorsulphonic acid decomposes, thus : S0 2 .C1-OH + H 2 = S0 2 (OH) 2 + HC1. Selenium. This element belongs to the rarer elements, for though it is fairly widely distributed in nature, it generally occurs in extremely small quantities. It has been found in the free state in Mexico ; but it occurs mainly in combination with metals like lead, copper, iron, silver, and thallium. Not infre- quently it is present in small amount in pyrites, and so in roasting the latter the selenium is oxidized to selenium dioxide which is carried into the dust flues of sulphuric acid factories. SULPHUR, SELENIUM, AND TELLURIUM 213 In these flues there is also deposited some free selenium, for the latter forms when hot sulphur acts on selenium dioxide. This gets into the lead chambers, where it is reduced to selenium by the action of sulphur dioxide, and so accumulates in the slime at the bottom of the chambers. In 1817 Berzelius dis- covered selenium in the slime of the lead chambers at Gripsholm. He named the element selenium, from the Greek word mean- ing moon, because of its similarity to tellurium, which is named from tellus, the earth. There are three varieties of selenium : (1) a red amorphous precipitate which dissolves in carbon disulphide arid separates from the latter solution in form of (2) red monoclinic crystals fusing at 170-180, which are also soluble in carbon disul- phide ; and (3) a bluish gray, metallic form which crystallizes in the hexagonal system and is insoluble in carbon disulphide. This metallic form conducts electricity slightly, which property may be increased tenfold by exposure to light. The conduc- tivity depends on the intensity of the light. The metallic form has a specific gravity of 4.8, melts at 217, and boils at 680. The atomic weight of selenium is 79.2, and at high temperatures the molecular weight corresponds to the formula Se 2 . Compounds of Selenium. These are similar to the compounds of sulphur. So hydrogen selenide may be formed by treating ferrous selenide with hydrochloric acid : FeSe + HC1 = FeCl 2 + H 2 Se. The compound H 2 Se is a gas that has the smell of horseradish and is more poisonous than hydrogen sulphide. The aqueous solution deposits selenium on exposure to the air or to oxygen. With the exception of the selenides of the alkalies, the com- pounds of the metals with selenium are difficultly soluble in water. With chlorine, selenium forms selenium monochloride Se 2 Cl 2 and selenium tetrachloride SeCl 4 . The former is a dark, brownish yellow oil and the latter a light yellow crystalline solid. Selenium dioxide SeO 2 is a solid formed by burning selenium in the air. It is the only oxide of selenium known. It forms long white prismatic crystals that sublime at about 300, 214 OUTLINES OF CHEMISTRY When sulphur and selenium dioxide are heated together, sul phur dioxide and selenium are formed : S + SeO 2 = SO 2 + Se. By oxidizing selenium with nitric acid, selenious acid H 2 SeO 3 is produced. By means of sulphur dioxide, selenious acid is reduced to selenium : H 2 SeO 3 + 2 SO 2 4- H 2 O = 2 H 2 SO 4 + Se. In this way the element is formed in the lead chambers of the sulphuric acid factories. When SeO 2 and SeCl 4 react with each other, they form SeOCl 2 , selenyl chloride : SeO 2 + SeCl 4 = 2 SeOCl a . The compound melts at 10 and boils at 179. Selenic acid H 2 SeO 4 is formed by oxidation of selenious acid by means of chlorine : H 2 Se0 3 + H 2 + C1 2 ^ 2 HC1 + H 2 SeO 4 . The action is reversible, for selenic acid is able to liberate chlorine from hydrochloric acid. Selenic acid is thus a more powerful oxidizing agent than sulphuric acid. The latter oxidizes hydrobromic acid, but not hydrochloric acid. Pure selenic acid is a solid melting at 62. The 95 per cent solution is a thick, oily liquid not unlike sulphuric acid in appearance. When hydrogen sulphide is passed into a solution of seleni- ous acid, selenium sulphide SeS is precipitated. It is yellow in color and does not dissolve in ammonium sulphide. Tellurium. Tellurium is one of the rare elements. It has been found in the free state, and also in the form of tellurides in combination with gold, silver, lead, and bismuth. It occurs in Colorado, California, Hungary, Brazil, and the Liparian Islands. It is a brittle, crystalline, silvery white substance having metallic luster. In precipitated amorphous form it is a black powder. In metallic form it conducts heat and elec- tricity like other metals. It has a specific gravity of 6.26 and melts at 455. Its atomic weight is 127.5; and at 1400, its boiling point, the vapor density corresponds to the formula Te 2 . Tellurium was discovered in 1782 by Miiller von Reich- ens te in, whose work was confirmed by Klaproth in 1798. Tin; latter called the element tellurium, from tellus, earth. SULPHUR, SELENIUM, AND TELLURIUM 215 Compounds of Tellurium. By the action of hydrochloric acid upon zinc telluride ZnTe, hydrogen telluride H 2 Te is formed : ZnTe + 2 HC1 = ZnCl a + H 2 Te. The product is generally contaminated with some hydrogen, which is liberated simultaneously. Hydrogen telluride is a colorless, poisonous gas of disagreeable odor. It is combusti- ble and fairly soluble in water. Its aqueous solutions when in contact with oxygen or air gradually deposit tellurium. When conducted into solutions of metallic salts, tellurides of the metals are in general precipitated. Such tellurides may also be prepared by heating metals with tellurium. With chlorine, tellurium forms tellurium dichloride TeCl 2 and tellurium tetrachloride TeCl 4 . These are formed when chlorine is passed over hot tellurium. If the chlorine is in large excess, the tetrachloride is formed ; if less chlorine is used, the dichloride forms together with some tetrachloride. The di- chloride is a black crystalline substance melting at 175 and boiling at 324. The tetrachloride forms white, shining crys- tals that melt at 224 and boil at 380". Both chlorides are decomposed by water. It is to be noted that the dichloride TeCl 2 is not analogous to the lower chloride of sulphur, which is S 2 C1 2 . Tellurium dibromide TeBr 2 and tetrabromide TeBr 4 have also been prepared. Tellurium diiodide TeI 2 and tellurium tetraiodide TeI 4 are also known. When sulphur trioxide acts on tellurium, tellurium sulphur trioxide TeSO 3 , a red amorphous solid, forms, which on heating is decomposed into sulphur dioxide and tellurium monoxide TeO. The latter is a black, amorphous substance, which on heating yields tellurium dioxide TeO 2 and tellurium. When heated in the air, tellurium is oxidized to tellurium dioxide TeO 2 . This is a white crystalline powder which is volatile at red heat (i.e. at higher temperatures than tellurium itself) and difficultly soluble in water. By means of nitric acid, tellurium may be oxidized to tellurous acid H 2 TeO 8 . This is a feeble acid that forms a white powder which is slightly soluble in water. On heating, it decomposes into water and tellurium dioxide. With strong bases it forms both acid and normal tellurites, like KHTeO 8 and K 2 TeO 3 . However, towards strong acids it behaves like a base. The salts thus formed 216 OUTLINES OF CHEMISTRY may be considered as derivatives of Te(OH) 4 , that is, H 2 TeO 3 .H 2 O. So, for instance, tellurium sulphate Te(SO 4 ) 2 has been prepared. Moreover, the salt TeCl 4 may be retained in aqueous solutions in presence of an excess of hydrochloric acid. Being both a weak base and also a weak acid, the salts that tel- lurous acid forms with either bases or acids are not very stable. This is generally the case with substances that do not have pro- nounced chemical characteristics. On fusing together barium nitrate and tellurium dioxide, barium tellurate may be formed: Ba(NO 3 ) 2 + TeO a = BaTeO 4 + 2 NO 2 . By decomposing barium tellurate with the calculated quantity of sulphuric acid, barium sulphate, which is insoluble in water, and telluric acid H 2 TeO 4 , which remains in solution, result: BaTeO 4 -f H 2 SO 4 = BaSO 4 + H 2 TeO 4 . The latter may also be prepared by first making potassium tel- lurate K 2 TeO 4 , by fusing either tellurium or tellurium dioxide with potassium carbonate and potassium nitrate, or by passing chlorine into an alkaline solution of potassium tellurite. The potassium tellurate is then converted into the barium salt by means of barium chloride, thus: K 2 Te0 4 + BaCl 2 = 2 KC1 + BaTeO 4 ; and the barium tellurate is then decomposed by dilute sul- phuric acid as before. From the aqueous solution, telluric acid separates in form of monoclinic crystals of the composition H 2 TeO 4 -2 H 2 O or Te(OH) 6 . On heating these, H 2 TeO 4 forms, which loses water at 160, yielding tellurium trioxide TeO 3 , an orange-yellow, crystalline substance that unites with water ex- tremely slowly and decomposes into tellurium dioxide and oxygen on ignition. While telluric acid forms tellu rates with the alkalies and other metals, its resemblance to sulphuric acid and selenic acid is extremely slight. Like tellurous acid, telluric acid may act as a base toward strong acids. General Considerations. Oxygen is commonly considered as forming with sulphur, selenium, and tellurium a natural family group of elements. We have already seen that fluorine, chlorine, bromine, and iodine form such a group in which SULPHUR, SELENIUM, AND TELLURIUM 217 fluorine is rather less closely related to chlorine, bromine, and iodine, than these three are to one another. Now, the relation of oxygen is similarly less close to sulphur, selenium, and tel- lurium. From oxygen to tellurium we have a gradation of physical properties, as the following table shows : NAME COLOR ATOMIC WEIGHT SPECIFIC GRAVITY MELTING POINT BOILING POINT Oxy'en blue 16 1 124 (at - 181) -181.4 Sulphur Selenium yellow red or 32.07 1.96 to 2.0 114.5 450 metallic 79.2 4.8 217 680 Tellurium black or metallic 127.5 6.3 455 1400 All of these elements exhibit allotropism. Toward hydrogen these elements are bivalent, forming com- pounds of the type H 2 X, thus : H 2 O, H 2 S, H 2 Se, H 2 Te. The stability of these compounds decreases as the atomic weight of the elements in question increases. Sulphur, selenium, and tellurium form compounds with oxygen, whose composition is represented by the types XO 2 and XO 3 . In the former, that is SO 2 , SeO 2 , and TeO 2 , the elements are tetravalent ; whereas in the latter, namely, SO 3 , SeO 3 , and TeO 3 , the elements in question are hexavalent, which is the high- est valence they are capable of exhibiting. Again, in the acids of the type H 2 XO 3 , namely H 2 SO 3 , H 2 SeO 3 and H 2 TeO 3 , and in those of the type H 2 XO 4 , namely H 2 SO 4 , H 2 SeO 4 , and H 2 TeO 4 , we plainly have striking analogies. In the com- pounds H 2 XO 3 , the elements are quadrivalent, thus: x- =o X)-H /O-H \O-H / O-H In the compounds H 2 XO 4 , the elements are hexavalent, thus : O-H O-H \ O-H )-H 'O-H O-H OUTLINES OF CHEMISTRY Toward the halogens, sulphur, selenium, and tellurium are bivalent and quadrivalent, while in some oxy-halogen deriva- tives they are hexavalent, thus : ' C1 Cl < / Se < ci Cl /Cl /Cl //Cl i ci /C\ ci' e ^ci C1 \C1 \C1 /Cl Se^O , \C1 The halogen compounds of selenium are more stable than those of sulphur, and those of tellurium are more stable than the selenium halides. One may regard the very unstable com- pound C1 2 O as analogous to C1 2 S, Cl 2 Se, and Cl 2 Te. Thus, it is apparent that as the atomic weight of the elements of the oxygen group increases, their affinity for halogen increases. This is also evident from the fact that in the halogen com- pounds, sulphur is readily replaced by selenium, and the latter is in turn replaced by tellurium. Sulphur, selenium, and tellurium, like chlorine, bromine, and iodine, form a group of three, in which the atomic weight of the middle element is very nearly equal to one-half the sum of the atomic weights of the other two, thus : S Te 1(32.07 + 127.5)= 79.78; whereas, Se = 79.2. In spite of the relationships noted, it should be borne in mind, however, that tellurium after all shows some decided points of departure in its chemical behavior from sulphur and selenium, so that the closeness of relationship between the latter elements and tellurium has repeatedly been called into question. SULPHUR, SELENIUM, AND TELLURIUM 219 REVIEW QUESTIONS 1. Where is sulphur found and in what forms? What are the allo- tropic forms of sulphur ? 2. How is sulphur refined, and for what purposes is it used? 3. Make a list of the six crystal systems and mention at least two substances that crystallize in each system. What is dimorphism? State the law of isomorphism. 4. How much hydrogen sulphide may be prepared from 85 kilograms of ferrous sulphide ? What volume would this gas occupy under standard conditions ? 5. What are the chief properties of hydrogen sulphide? Illustrate each of its important chemical properties by means of an appropriate equation. 6. Complete the following equations and state what property of hy- drogen sulphide is illustrated in each case : NH 4 OH + H 2 S = Ag + H 2 S = H 2 SO 4 + H 2 S = CuS0 4 + H 2 S = H 2 S + I = H 2 S + S0 2 = 7. Explain the apparent contradiction of the following reactions : HgCl 2 + H 2 S = HgS + 2 HC1. ZnS + 2 HC1 = ZnCl 2 + H 2 S. 8. 'What is the formula of sulphur monochloride ? Why? How may this substance be prepared ? For what purpose is it used ? 9. When 12 liters of hydrogen sulphide, measured under standard conditions, are burned to water and sulphur dioxide, what volume of oxy- gen will be required, and what is the volume of the sulphur dioxide formed ? 10. What is a sulphite? Give five examples. How do sulphites act when treated with hydrochloric or sulphuric acid ? Illustrate by means of an equation. 11. How does bleaching with sulphur dioxide differ from bleaching with chlorine or hydrogen peroxide ? 12. What is sulphurous acid? How may it be changed to sulphuric acid by action of the halogens? Write appropriate equations. 13. How would you make sulphuric acid from the elements? Write the equations. 14. By what two processes is sulphuric acid prepared commercially? Write the equations. 15. What three important properties does sulphuric acid possess? Which of these is illustrated in the charring of wood and sugar ? 16. What is sulphuric acid used for? How does this substance act when poured into water? Why? 17. What is sodium thiosulphate ? How may it be prepared? Of what use is it? 220 OUTLINES OF CHEMISTRY 18. What two other elements are analogous to sulphur? Why? 19. How much sulphuric acid could be prepared from 350 tons of sulphur ? 20. Why is oxygen classified with sulphur, selenium, and tellurium? 21. What is the valence of sulphur in hydrogen sulphide ? In sulphur dioxide? In sulphuric acid? In sulphur trioxide? In sulphur mono- chloride ? 22. Assuming the air to consist of 20 per cent oxygen and 80 per cent nitrogen by volume, what will be the resulting volume if 1 gram of sul- phur is burned in 10 liters of air? How much of the final volume is due to each of the gases present? CHAPTER XIV CARBON AND SOME OF ITS TYPICAL COMPOUNDS Occurrence and Allotropic Forms of Carbon. Carbon occurs in the free state in nature as diamond and graphite, which are crystalline in character. It is also known in amorphous form as charcoal, coke, soot, lampblack, bone black, etc., resulting from the charring of animal and vegetable matter, and various compounds of carbon. Diamond, graphite, and amorphous carbon are the allotropic forms of the element. Carbon is a most important constituent of all plants and animals. Large quantities of carbon are found in form of coal, which represents the remains of vegetation of past geological ages. Coal is consequently not pure carbon. Indeed, the amount of free carbon in different kinds of coal varies con- siderably. In natural gas and petroleum, carbon occurs in combination with hydrogen. In form of carbon dioxide, carbon is found in the air and in many natural waters. As carbonates, especially calcium carbonate and magnesium carbonate, carbon is found in huge masses widely distributed in various parts of the earth. Calcium carbonate is commonly found in form of chalk, limestone, and marble ; whereas calcium magnesium carbonate, or dolomite, also called magnesian limestone, occurs very frequently and covers extensive areas of the earth's crust, often forming mountains. Diamond. The diamond is a crystalline form of carbon, and belongs to the regular or FIG. si. FIG. 82. isometric system, Figs. 81 and 82. It is colorless when pure, but frequently it is dark-colored or black, in which form it is known as carbonado, or bort. Some- times diamonds are colored blue, green, yellow, or red by small amounts of foreign substances. The diamond is very hard. It will scratch all other minerals. For this reason diamond 221 222 OUTLINES OF CHEMISTRY dust, usually in form of bort, is used by lapidaries in cutting and polishing. Drills of carbonado, which sometimes occurs in pieces as large as one's fist, are used in boring rocks, and glaziers make use of the diamond in cutting glass. Diamonds are found in Brazil, South Africa, India, and Australia, and sometimes also in the United States. They are usually cor- roded, and so their brilliant luster does not appear till the outer layer is removed by the lapidary, who, by grinding suitable artificial faces on the diamonds, brings out their highly prized brilliancy. The diamond has a specific gravity of 3.5 to 3.6 and a refractive index of 2.416 to 2.43. That it consists of carbon only, appears from the fact that on combustion in oxy- gen the sole product formed is carbon dioxide. Diamonds of microscopic dimensions have been prepared artificially by Mois- san, by dissolving carbon in molten iron and then chilling the same suddenly. Under the great pressure thus produced in the interior of the iron, graphite and some very small diamonds crystallize out. The diamond is not attacked by acids. It is a non-conductor of electricity, and becomes electrically charged when rubbed with a cloth. On heating it highly out of contact with the air, it may be changed to graphite ; whereas when heated highly in oxygen, it burns with great brilliancy to car- bon dioxide. The diamonds found are usually small ; only rarely do they weigh as much as 20 grams. The largest diamond known was found near Pretoria, South Africa. It is called the Cullinan, and weighed about 600 grams when found. Graphite. Graphite crystallizes in the monoclinic system in plates that simulate hexagonal forms. It is widely distributed in nature in form of small flakes or granules in various granite rocks. It is black or grayish black, having metallic luster. It is very soft, and may readily be crushed to a fine powder, which is frequently employed as a lubricant. Graphite is also used in making "lead pencils." At one time it was thought that graphite contained lead, hence the name plumbago, by which graphite is also known. Graphite conducts electricity, and in its artificial forms it is used as electrodes for electric arcs and in electro-chemical work, especially in making chlorine and caus- tic soda from common salt. Graphite powder is employed in electrotyping, in making stove polish, etc. The specific gravity CARBON AND SOME OF ITS TYPICAL COMPOUNDS 228 of graphite varies from 1.8 to 2.5. It burns with great difficulty even in oxygen, and the natural varieties leave from 3 to 20 per cent ash, which usually consists of silicates of various bases. Graphite is very refractory and not readily attacked by chemi- cal reagents, for this reason it is frequently employed together with fire clay in making crucibles. A mixture of concentrated nitric acid and potassium chlorate converts graphite into gra- phitic acid, which consists of small yellow crystals that explode on heating, leaving a mass of finely divided carbon. Graphitic acid consists of 56 per cent carbon, 2 per cent hydrogen, and 42 per cent oxygen. On treating graphite with concentrated nitric acid, and then igniting the material strongly, various samples show a different behavior. Thus, graphite found in the State of New York greatly increases in volume wlren so treated, whereas Siberian graphite is not so affected at all. Ceylon and Siberia furnish most of the natural graphite. Arti- ficial graphite is formed when carbon dissolved in molten iron crystallizes out, and when coke is heated to very high tempera- tures in the electric furnace out of contact of the air and then allowed to cool slowly. By the latter process, known as the Acheson process, much graphite of excellent quality is made at Niagara Falls. By first grinding up the coke, mixing it with a binder, like coal tar or black strap molasses, and molding it into desired forms and baking these in ovens, carbons for bat- teries, electric arc lights, and other purposes are obtained. These carbons may readily be converted into graphite by heat- ing them in the electric furnace as above stated. Artificial graphite is now much employed in the arts in place of natural graphite. Electric furnaces are either resistance furnaces or arc furnaces. In the former the heat is generated by passing a strong current of electricity through conducting material which is thus heated to high temperatures. In the arc furnaces, the high tempera- tures are obtained by means of the electric arc ; in principle, the arrangement is similar to that of the arc lamp. The Acheson graphite furnace is a form of resistance furnace, Fig. 83. The molded sticks of carbon are piled between two carbon terminals which are about two feet square and thirty feet apart. The whole is then covered with a thick layer of granular carbon and carborundum (which see), and a current 224 OUTLINES OF CHEMISTRY of 3000 amperes at a pressure of 220 volts is turned on which is gradually changed to 9000 amperes at 80 volts at the end of FIG. 83. the twenty-four hours during which the furnace is run. The whole is then allowed to cool, and the carbon is found to be converted to graphite. The heat is largely generated by the current as it passes through the granular carbon lodged in the interstices between the carbon sticks. Smaller resistance furnaces of different types are used in the arts and also for experimental purposes. Figure 84 shows a typical arc furnace for experimental work. The poles are of carbon, and the walls of the furnace are usually built of quicklime or mag- nesia. In the case of larger furnaces, this material is then reen forced on the outside by bricks made of crude mag- nesia. We shall have occa- sion to refer to electric fur- naces again as we proceed, * IG - 84t for they are now used for various purposes that require high temperatures. Amorphous Carbon. Amorphous carbon is formed when animal or vegetable matter, coal, or various compounds of car- bon are heated while access of air is excluded entirely or in part. Thus, by heating wood in this way charcoal is produced. CARBON AND SOME OF ITS TYPICAL COMPOUNDS 225 Similarly, bone black is made by heating bones in closed iron cylinders ; coke is formed by heating coal ; and blood charcoal is produced by heating blood till it chars. On burning hydro- carbons, like kerosene or turpentine, lampblack, or soot, is deposited on a cold object held in the smoky flame. Lamp- black is nearly pure carbon. All of these forms of amorphous carbon differ according to the source from which they are obtained. Charcoal contains about 85 per cent carbon. It is very porous and when freshly heated it has the power of absorbing many gases. So freshly ignited charcoal absorbs from 50 to 100 times its own volume of ammonia, hydrogen sulphide, bromine vapor, etc., which are again given off 011 heating the charcoal. This process of absorption consists of condensation of the gases on the sur- face of the charcoal and is called adsorption. Figure 85 illustrates th^ adsorption of ammonia (which has been collected over mercury) by means of charcoal. Charcoal is frequently used as a deodorant of vaults, cisterns, etc., because it absorbs the noxious gases. Bone black acts similarly. The latter is much used to decolorize various solutions. So in the refining of sugar, bone black serves to remove the brown coloring matter. Bone black contains from 70 to 80 per cent calcium phosphate, and only from 8 to 12 per cent carbon, together with some calcium carbonate and calcium sulphate. By repeated treatment with acids and water, nearly pure carbon may be prepared from bone black. A very pure amorphous carbon may be prepared by charring sugar. All animal charcoal contains some nitrogen, which is very tenaciously held. Very finely divided, dry charcoal has a strong affinity for oxygen ; indeed, it may catch fire on being thrown from its container into the open air, pyrophoric carbon. The various forms of amorphous carbon differ in their physical nature, and in the chemical character and the amounts of the impurities they contain. 226 OUTLINES OF CHEMISTRY Lampblack is much used as a pigment in paints, India ink. etc., whereas coke is used for fuel and in the reduction ot ores, particularly iron ore. In the production of coke, the gaseous products, tar, etc., should be saved and not allowed to burn and go to waste as is still frequently done. Similarly man}' valuable volatile products are lost when charcoal is made by heating wood in a pile covered with sod and earth in the old- fashioned way. By heating wood and coal in retorts, these gaseous and other volatile products may be saved and used. The process of thus heating and decomposing substances in retorts is called destructive or dry distillation. Coal. Coal is found in large deposits in various parts of the earth. It represents the plant remains of various geologi- cal periods from the carboniferous to the tertiary. Coal has been formed by the gradual abstraction of carbon, hydrogen, and oxygen, mainly in form of water, marsh gas, etc., from the vegetable remains. Much marsh gas (which see) is found associated with coal. There are many varieties of coal, which are commonly roughly divided into two great classes, namely, soft coal and hard coal. Hard coal, or anthracite, contains much less volatile matter than soft coal, which is also called bitu- minous coal. The latter burns with a sooty flame, and evolves more hydrocarbon gases than anthracite when heated. It is consequently used in manufacturing illuminating gas (which see). The amount of carbon in charcoal, coke, or anthracite is generally in the neighborhood of 95 per cent, whereas soft coal contains about 80 per cent carbon. Chemical Behavior of Carbon. The various forms of carbon above mentioned are practically not at all attacked by chemi- cal reagents at room temperatures. When heated in the air, carbon burns ; that is, it unites with oxygen. When a suffi- cient supply of oxygen is at hand, the product formed is car- bon dioxide, a gas whose composition corresponds to the formula CO 2 . When an insufficient amount of oxygen is furnished, some carbon monoxide is also formed. The latter substance is also a gas ; its composition is expressed by the formula CO. The very careful work of Dumas and Stas has shown that 3 parts of carbon by weight unite with 8 parts of oxygen by weight to form carbon dioxide gas. Under standard conditions CARBON AND SOME OF ITS TYPICAL COMPOUNDS 227 the weight of 22.4 liters of carbon dioxide is nearly 44 grams. The molecular weight of carbon dioxide is consequently 44. However, 44 parts of carbon dioxide by weight contain 12 parts of carbon and 32 parts of oxygen. Now, since we have previously chosen 16 as the atomic weight of oxygen, it is clear that the molecule of carbon dioxide contains 32-f-16 or 2 atoms of oxygen. The question now arises, how many carbon atoms there are in a molecule of carbon dioxide ? To answer this question really involves choosing the atomic weight of carbon from the combining weight. A study of all of the gaseous compounds of carbon that are known has revealed the fact that in no case do these contain less than 12 grams of carbon in 22.4 liters under standard conditions ; that is to say, there is no compound into which carbon enters whose molecule con- tains less than 12 parts of carbon by weight. For this reason, the atomic weight of .carbon is taken as 12, rather than some fraction or multiple thereof. Having thus determined upon 12 as the atomic weight of carbon, it is obvious from what has been said that the molecule of carbon dioxide contains one atom of carbon and two atoms of oxygen, and its formula is therefore CO 2 . The specific heat of carbon increases with the temperature, becoming nearly constant in the neighborhood of 1000. Between 900 and 1000 it is about 0.46, which fact also yields approximately 12 as the atomic weight of carbon according to the law of Dulong and Petit. When carbon is burned in oxygen, the volume of the car- bon dioxide formed is the same as that of the original oxygen, measured, of course, under the same conditions of temperature and pressure. Figure 77 represents an apparatus for demon- strating this fact. Carbon on a platinum spoon is burned in the flask which is filled with oxygen. After the whole has again cooled to room temperature, the mercury manometer, attached as shown, indicates that there has been no change of volume. Therefore we have : C + 2 = C0 2 (1 volume) (1 volume) which is quite in harmony with what we should expect on the basis of Avogadro's hypothesis. It will be recalled that when sulphur is similarly burned in oxygen the volume of the SO a formed is also the same as that of the oxygen. 228 OUTLINES OF CHEMISTRY Carbon does not unite directly with hydrogen except at very high temperatures such as are produced by the electric arc, and even then the union takes place with difficulty. Very many compounds of hydrogen and carbon are known, however, for they may be prepared by indirect methods. Similarly carbon does not combine directly with the halogens, except in the case of fluorine. However, by indirect methods compounds of car- bon with the halogens may be formed fairly readily. With sulphur carbon unites directly at high temperatures. At the temperature of the electric furnace, carbides of calcium, alumi- num, silicon, and boron may be formed, and iron generally contains some carbon which is present in form of carbide of iron. With nitrogen carbon does not unite directly, though by indirect means it is quite possible to effect the union of these two elements. In general, carbon is rather inert chemically, though its com- pounds, when once formed, frequently have a very considerable degree of stability. We shall see that the carbon atom has a great tendency to unite with other carbon atoms, which often results in building up of large and complex molecules. At high tempera- tures carbon is generally bivalent, whereas at lower tempera- tures it is quadrivalent. Sometimes carbon is regarded as quadrivalent in all of its compounds. This view would neces- sitate that a goodly number of carbon compounds be regarded as unsaturated. The number of carbon compounds known is very great, so that it is customary to treat these in a separate division of chemistry, namely, organic chemistry. The term organic chemistry, as used at present, signifies that branch of the science which deals with the compounds of carbon, or with the hydrocarbons and their derivatives ; for all the carbon compounds may be considered as derived from compounds of carbon and hydrogen by substituting other elements in place of the hydrogen. Since carbon is an essential constituent of all living beings, the study of the carbon compounds is closely associated with the study of the chemistry of the products that are formed in or- ganic beings. Hence the name organic chemistry. Indeed, till 1828 it was quite generally held that compounds that are produced by the life process could not be prepared artificially in the laboratory, but Wohler's synthesis of urea (which see) CARBON AND SOME OF ITS TYPICAL COMPOUNDS 229 showed that such syntheses are quite possible ; and now many of the products that are formed by the metabolic processes in plants and animals have been prepared in the laboratory. Carbon Dioxide. As already mentioned, there are two oxides of carbon known, namely, carbon monoxide CO, and carbon dioxide CO 2 . Carbon dioxide is the final oxidation product of carbon. It always results when carbon is burned in a suffi- cient amount of air or oxygen. It is moreover exhaled as a product of respiration of both plants and animals. It will be recalled that every 1000 volumes of air contain 3 volumes of car- bon dioxide. This gas is also contained dissolved in all natural waters. Many spring waters, like those at Saratoga, Colorado Springs, Selters, Vichy, and Narzan, are so highly charged with carbon dioxide, under pressure, that the gas escapes with effer- vescence when the pressure is released. In volcanic regions, like those of Italy, South America, and Java, large quantities of carbon dioxide issue from fissures in the earth's crust. In fermentation and in the decay of animal and vegetable matter carbon dioxide is always formed. It is thus evident that the air is continually receiving carbon dioxide from quite a variety of sources. Carbon dioxide maybe produced by the oxidation of carbon : C + 2 =C0 2 . In the process of fermentation of sugar by means of yeast, alcohol and carbon dioxide result. Thus the fermentation of glucose may be expressed by the following equation : C 6 H 12 6 =2C 2 H 6 OH + 2C0 2 . glucose alcohol carbon dioxide However, the simplest way of obtaining carbon dioxide is by the action of an acid upon a carbonate, like calcium carbonate CaCO 3 , or sodium carbonate : CaCO 3 + 2 HC1 = CaCl 2 + H 2 O + CO 2 . Na 2 CO 3 + H 2 SO 4 = Na 2 SO 4 + H 2 O + CO 2 . It will be recalled that when an acid acts on a sulphite, H 2 SO 3 is not formed, for it at once decomposes, yielding H 2 O and SO 2 . Similarly, when an acid acts on a carbonate, we do not get H 2 CO 3 , carbonic acid, but H 2 O and CO 2 . The latter is clearly carbonic acid anhydride. It is probable that the compound 230 OUTLINES OF CHEMISTRY H 2 SO 3 , sulphurous acid, exists in aqueous solutions of SO 2 ; and it is also likely that aqueous solutions of CO 2 contain H 2 CO 3 , carbonic acid. While an aqueous solution of any of the ordinary acids will decompose carbonates, owing to the weak- ness of carbonic acid and the volatility of carbon dioxide, cal- cium carbonate, in form of marble, and hydrochloric acid are com- monly employed, because these materials are cheap, and the cal- cium chloride is readily soluble in water. A Kipp apparatus (Fig. 86) is very convenient for evolving carbon dioxide from marble and hydrochloric acid. Just as sulphites are salts of sulphurous acid H 2 SO 3 , so car- bonates may be regarded as salts of the dibasic acid H 2 CO 3 . We have then normal carbonates, like Na 2 CO 3 , and acid carbonates, or bicarbonates, like NaHCO 3 . Basic carbonates of many metals, like zinc, copper, lead, etc., are also known. When carbon dioxide is conducted into clear limewater, a white precipitate of calcium carbonate is formed, thus: Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O. On continuing to conduct in more carbon dioxide the calcium carbonate again dissolves, the solution becoming clear. In this process calcium bicarbonate is formed : CaC0 3 + H 2 + C0 2 = Ca(HC0 3 ) 2 . On boiling the solution, however, carbon dioxide is expelled and calcium carbonate is reprecipitated : Ca(HCO 3 ) 2 = CaCO 3 + H 2 O + CO 2 , The fact that limewater, or baryta water Ba(OH) 2 , is rendered turbid by carbon dioxide is commonly used as a means for detecting the latter, though in some cases such a test requires further confirmation. Breathing into limewater renders the FIG. 86. CARBON AND SOME OF ITS TYPICAL COMPOUNDS 231 latter turbid, due to the formation of a precipitate of calcium carbonate, and this demonstrates the presence of carbon diox- ide in the breath. Many natural waters contain lime and magnesia in solution in form of the bicarbonates, Ca(HCO 3 ) 2 and Mg(HCO 3 ) 2 . When such waters are boiled, carbon dioxide escapes and the normal carbonates, CaCO 3 and MgCO 3 , are precipitated. Water containing carbonate or bicarbonate of calcium in solu- tion is called hard water. Since boiling decomposes the bicar- bonate and thus removes some of the calcium salts in form of a precipitate of CaCO 3 , it is clear that after boiling the water has lost some of its "hardness." Thus it is common to speak of temporary hardness of water, which can be removed by boil- ing, and permanent hardness, which remains even after boiling. The carbonates of the alkalies, like sodium carbonate Na 2 CO 3 , and potassium carbonate K 2 CO 3 , may be fused without de- composition. But the carbonates of other metals are com- monly decomposed on ignition, yielding carbon dioxide and the oxide of the metal, thus : CaCO 3 = CaO + CO 2 . SrCO 3 = SrO + CO 2 . Properties of Carbon Dioxide. Carbon dioxide is a colorless gas, of slightly acid taste and a feeble, agreeably pungent smell. It is 1.529 times as heavy as air. Being so heavy, it may be readily poured or even siphoned from one jar to another (Fig. 87). It neither supports combustion nor respiration. At and below 31, its critical temperature, carbon dioxide may be liquefied by pressure. The liquid boils at 79 under atmos- pheric pressure. Solid carbon dioxide readily forms when the liquid is allowed to evaporate rapidly in the air, for thus much heat is absorbed and the remaining liquid is chilled below the melting point. The solid looks like snow, and melts at 57. In the air it evaporates without first passing into the liquid state, which is quite natural, since its boiling point under at- mospheric pressure is much lower than its melting point. The large amounts of carbon dioxide evolved by fermentation in the brewing industries are now collected, washed, and pumped into cylinders made of mild steel (Fig. 13.) In this process the car- bon dioxide is chilled and put into the cylinders under pressure. 232 OUTLINES OF CHEMISTRY It is used for making soda water, which consists of water charged with carbon dioxide under pressure. When the pressure is re- leased, a portion of the gas escapes, for under atmospheric pressure and room temperature, water dis- solves only about its own volume of the gas. The solution has a faintly acid reaction toward lit- mus. The name "soda water" comes from the fact that sodium bicarbonate NaHCO 3 , also popu- larly called bicarbonate of soda or simply soda, is at times used in preparing carbonated water. Solid carbon dioxide is fre- quently used for securing low temperatures. For this purpose it is often employed mixed with ether in order to secure more rapid evaporation. Thus tem- peratures as low as 80 may be secured, and in a partial vacuum even 100 may be reached. Natural carbon dioxide, as it issues from the earth, is also frequently bottled in steel cylin- ders as described, and placed on the market. Thousands of tons of carbon dioxide are thus sold annually. Carbon dioxide is also used as a fire extinguisher. The reason why it does not burn or support combustion is that it contains all the oxygen it can hold, and retains this very tenaciously. At high temperatures potassium will burn brilliantly in carbon dioxide, for potassium is under these conditions able to rob carbon dioxide of its oxygen, thus : 3 CO a + 4 K = C + 2 K 2 CO 3 . Here again we see the relative character of the process of combustion. Physiological Effects of Carbon Dioxide. In pure carbon dioxide, living beings will die for lack of oxygen, just as they would succumb in nitrogen, for example. Air containing over 20 per cent of carbon dioxide may also finally produce death FIG. 87. CARBON AND SOME OF ITS TYPICAL COMPOUNDS 238 when breathed ; for in respiration carbon dioxide is given off and oxygen is taken up from the air, which processes are greatly impeded by a high content of carbon dioxide in the air. Upon the mucous membranes, carbon dioxide has a stimulating action which is agreeable and refreshing, for this reason car- bonated drinks are highly esteemed. Water from which all carbon dioxide has been expelled by recent boiling tastes flat, as already stated. Carbon dioxide colors the blood dark brown, and so when air containing even but 6 per cent of carbon diox- ide is breathed continuously drowsiness results. The anaes- thetic effects of carbon dioxide are similar to those of nitrous oxide. These effects may be counteracted by taking the person affected into fresh air that contains only the usual amount of carbon dioxide, for thus the blood again is able to get its needed oxygen, and so the intoxication gradually disappears. The necessity of good ventilation of buildings, especially sleep- ing rooms and audience halls, is consequently apparent. The evil effects of re-breathing carbon dioxide are great in them- selves, but the exhalations aside from their carbon dioxide con- tent are much more injurious to health. Relations of Carbon Dioxide to Plant and Animal Life. All animals exhale carbon dioxide, which is produced during the life process as a result of slow oxidation of their tissues. The heat of the animal body is a result of the oxidation that is con- tinually going on while the animal lives. Now the carbon diox- ide thus exhaled enters the atmosphere, and is again taken up by the green leaves of plants in which in the sunlight in presence of the chlorophyll, water and carbon dioxide react with each other, forming compounds containing carbon, hydrogen, and oxygen, notably starch (CgH^Og)^, and free oxygen, the latter being exhaled. Thus, carbon dioxide is decomposed and starch is formed, which may again serve as food for animals, in whose bodies it is oxidized to carbon dioxide and water. And so the cycle, often spoken of as the carbon cycle, repeats itself over and over. The energy is furnished by the sun's rays, which pro- duce the decomposition of carbon dioxide and water into starch and oxygen in the green leaf of the plant. The starch and free oxygen formed contain more energy than the carbon dioxide and water ; and this excess of energy in the starch and oxygen is again given off in the animal body, when carbon dioxide and 234 OUTLINES OF CHEMISTRY water are formed as a result of the oxidation of the starch. In connection with the study of the element nitrogen, we have seen that here too a somewhat similar cycle occurs. Early Work on Carbon Dioxide. J. B. van Helmont (1577- 1644) showed that the same gas is formed during alcoholic fermentation, during the action of acids on chalk, and during the combustion of carbon. He demonstrated that this gas is also found issuing from fissures in volcanic regions, and that it is contained in the waters of mineral springs. Indeed, it was he who first used the term gas. He called carbon dioxide " gas sylvestre " or " gas carbonum." He was aware of the fact that carbon dioxide does not support combustion or respiration. Stephen Hales was the first to collect gases over water by dis- placement, as is still in vogue, while Priestley taught how to use mercury for this purpose instead of water. The fact that carbon dioxide is absorbed by caustic alkalies was discovered by Joseph Black in 1757. He called the gas " fixed air " be- cause it is thus absorbed or fixed by caustic alkalies. Black demonstrated that soluble salts are formed when carbon dioxide acts on caustic potash or soda, and that an insoluble precipitate results when the gas is conducted into limewater. He also found that carbon dioxide is liberated when limestone is strongly heated, as in the process of making lime. McBride showed that the gas is liberated during putrefaction. Priestley demonstrated its presence in the air, and Lavoisier proved that carbon dioxide is formed during respiration and the reduction of metallic oxides by charcoal. The latter also showed that the gas contains only carbon and oxygen, while to the work of Berzelius, Dumas, Stas, and Roscoe we owe the careful deter- mination of the percentage composition of carbon dioxide. Carbon Monoxide. Carbon monoxide occurs in gases that issue from volcanoes. It is a constituent of illuminating gas and particularly of so-called water gas. Furthermore, it often occurs in the gases issuing from blast furnaces in which iron ores or other metallic oxides are being reduced. The gas may readily be formed by passing carbon dioxide over red-hot carbon, thus : The reaction is reversible. At 1000 it is nearly complete in CARBON AND SOME OF ITS TYPICAL COMPOUNDS 235 i lie sense of the upper arrow, while at 450 it is practically completely reversed. Carbon monoxide is commonly formed in ordinary coal fires, where the carbon dioxide passes upward through red-hot layers of coal. The blue flame so frequently observed in a coal fire is due to the combustion of carbon mon- oxide. This gas is always formed to some extent when carbon is burned in an insufficient supply of oxygen. In this case, however, it is never pure, being always associated with carbon dioxide. The latter may be readily removed by passing the gases through caustic potash solution, which absorbs the car- bon dioxide, but not the carbon monoxide. When air is passed through beds of incandescent coke or coal, in furnaces of special type, the issuing gas consists of 28 to 30 per cent carbon monoxide, 63 per cent nitrogen, and smaller amounts of carbon dioxide. This gas is known as producer gas and is extensively used for fuel, being readily obtained. The nitrogen and carbon dioxide greatly dilute the gas and diminish its heating power. When steam is passed over carbon heated to 1000 to 1400 C. (see apparatus, Fig. 4), carbon monoxide and hydrogen, a mix- ture which is known as water gas, is produced, thus : C + H 2 = CO + H 2 . This water gas is frequently made on a large scale and used as a fuel. In America it is also often used for illuminating pur- poses, in which case, since it burns with a non-luminous flame, it must be enriched or " carbureted " by the addition of the vapors of hydrocarbons that are rich in carbon (see illuminat- ing gas). Carbon monoxide is produced when many metallic oxides are heated with excess of carbon, thus : ZnO + C = Zn + CO. Indeed, it was by means of this reaction that carbon monoxide was first observed (1776). Carbon monoxide is further formed by heating carbonates with carbon or zinc dust, thus : CaCO 3 +C = MgCO 8 + Zn = ZrO + MgO + CO. 236 OUTLINES OF CHEMISTRY The gas is also formed by heating many organic acids with concentrated sulphuric acid, which abstracts water from the organic acids and so decomposes them. The resulting carbon monoxide generally contains carbon dioxide, which may, how- ever, be removed by means of caustic potash, as already de- scribed. In the laboratory carbon monoxide is often prepared by heating oxalic acid with sulphuric acid, thus : (COOH) 2 + H 2 SO 4 = H 2 SO 4 H 2 O + CO 2 + CO. When formic acid is used, carbon monoxide only is obtained: HCOOH + H 2 S0 4 = H 2 S0 4 - H 2 + CO. Frequently, when larger quantities of carbon monoxide are required for experimental work, the gas is prepared by heating together potassium ferrocyanide with ten times its weight of concentrated sulphuric acid in a flask of relatively large capacity. The reaction which occurs is as follows : 6 H 2 S0 4 + K 4 Fe(CN) 6 + 6 H 2 O = 2 K 2 SO 4 + 3 (NH 4 ) 2 SO 4 + FeSO 4 + 6 CO. Properties of Carbon Monoxide. Carbon monoxide is a colorless, odorless, tasteless gas which is 0.9672 time as heavy as air. At and below 139.5 it may be condensed to a liquid by pressure. The critical pressure is 35.5 atmos- pheres. The liquid boils at 190, and the white, snowlike solid melts at 207. In the air and in oxygen, the gas burns with a rather small blue flame, forming carbon dioxide, thus : 2CO + O 2 =2CO 2 . (2vols.) (1vol.) (2vols.) It is found by exploding a mixture of carbon monoxide and oxygen that 2 volumes of the former and 1 volume of the latter form 2 volumes of carbon dioxide, as indicated in the above equation. From this fact and the one that 22.4 liters of carbon monoxide under standard conditions weigh 28 grams, it follows that the formula for the gas is CO. The composition of carbon dioxide must, of course, be known as a result of independent experiment. In the volumetric relation of the combination of carbon monoxide and oxygen, we have another excellent illustration of the law of Gay-Lussac of combination of gases by volume, which law serves as the main support <.f A.vogadro's hypothesis. CARBON AND SOME OF ITS TYPICAL COMPOUNDS 237 Carbon monoxide is a strong reducing agent. It is able to abstract oxygen from many metallic oxides at higher tempera tures, thus : CuO + CO = Cu + CO 2 . Fe 2 8 + 3 CO = 2 Fe + 3 CO 2 . In sunlight carbon monoxide unites with chlorine, forming an addition product, phosgene COC1 2 ; while 011 heating sul- phur vapor with carbon monoxide, carbon oxysulphide COS is formed. Phosgene, or carbonyl chloride, boils at + 8 and is readily decomposed by water: COC1 2 + H 2 = C0 2 + 2 HCL Carbon oxysulphide is a colorless, inflammable gas with an odor like that of hydrogen sulphide. When burned, the prod- ucts formed are carbon dioxide and sulphur dioxide : COS + 3O= CO 2 + SO 2 . With nickel and iron, carbon monoxide forms the carbonyl compounds Ni(CO) 4 and Fe(CO) 5 . The chemical behavior of carbon monoxide is readily ex- plained by regarding it as an unsaturated compound, the two free bonds of the carbon atom enabling the formation of the various addition products to take place. On the other hand, it must be remembered that while carbon monoxide is a strong reducing agent, it may itself in turn be reduced by still stronger reducing agents like magnesium and aluminum, whose oxides at the high temperatures at which the reaction takes place are very stable and non- volatile, thus : 3 CO + 2 Al = A1 2 O 3 -|- 3 C. CO + Mg = MgO + C. In these reactions, then, carbon monoxide is compelled to play the role of an oxidizing agent, which again shows us that oxida- tion and reduction are processes that are relative in character. Carbon monoxide is readily absorbed by an ammoniacal or hydrochloric acid solution of cuprous chloride at room tempera- tures. The latter solution is much used in estimating carbon monoxide in gas analysis. From these solutions carbon mo- noxide may be expelled by heating. 238 OUTLINES OF CHEMISTRY Physiological Effects of Carbon Monoxide. Carbon monoxide is a very poisonous gas, and is all the more dangerous because it is odorless, and so does not betray its presence till it has already produced toxic effects. The gas unites with the hemo- globin of the blood, forming an addition product which is bright red in color and very stable. This fact was discovered in 1826 by Piorry. The carbon monoxide hemoglobin is much more stable than the oxyhemoglobin. Air containing as low as one- twentieth of one per cent carbon monoxide exerts toxic effects. These manifest themselves as headache, unconsciousness, con- vulsions, and finally death, which is caused by about 100 cc. of pure carbon monoxide for every 10 kilograms of weight of the person or animal inhaling the substance. The resuscitation of persons poisoned by inhaling carbon monoxide is effected by means of fresh air, in very mild cases. In serious cases, oxy- gen must be supplied, preferably under pressure of from one and one half to two atmospheres. As carbon monoxide is commonly produced in coal stoves, it is necessary to provide suitable draught to burn the gas completely, or at any rate to carry off the products of combustion so that they cannot escape into the room. Water gas, which, as we have seen, contains about 50 per cent carbon monoxide, is doubtless more poisonous than ordinary illuminating gas made by heating coal. Carbon Bisulphide. When sulphur vapor is passed over charcoal or coke heated to redness, carbon bisulphide CS 2 is formed. It is also called carbon disulphide. The carbon is heated in a tall iron cylinder, and the sulphur vapors are passed upward through the hot coal, the product being conducted off in tubes and condensed to a liquid. The heating is now gen- erally accomplished by means of electricity. The continuous process thus devised by E. R. Taylor represents a very great improvement over older methods. Figure 88 represents the Taylor furnace. Pieces of coke are placed between the elec- trodes .27, which are supplied with a strong alternating current. The heat produced melts and volatilizes the sulphur S, which is continuously supplied through B. The coke is renewed through Q. The tower is filled with charcoal which is intro- duced from above through D. The carbon disulphide vapors pass off through A to the condensers. This furnace is 40 ft. high and 16 ft. in diameter. CARBON AND SOME OF ITS TYPICAL COMPOUNDS 239 Carbon bisulphide forms a colorless, volatile liquid of not un- pleasant ethereal odor when pure. As it usually comes in the market, it is slightly yellowish in color and has a disagree- able odor, which is due to impurities. These readily form on standing, espe- cially in presence of moisture, because of slight decompo- sition of the carbon disulphide. The latter has a specific gravity of 1.262 at 20;. its vapor is 2.68 times as heavy as air. Carbon di- sulphide is an ex- tremely inflammable liquid and is con- sequently dangerous to handle. The vapors catch fire in the air when heated to but 232. Mixed with air, its vapors are explosive. Car bon disulphide burns with a blue flame, forming car- bon dioxide and sulphur dioxide : B FIG. 88. Carbon disulphide is a good solvent for fats, oils, iodine, rubber, and sulphur. It is used as a solvent for fats and oils on a large scale, also for vulcanizing rubber. It is further employed in exterminating ants, lice, and other insect pests. When in- 240 OUTLINES OF CHEMISTRY haled, its vapors act as an anaesthetic, large quantities producing intoxications and serious disturbances of the nervous system. Carbon disulphide CS 2 is analogous to carbon dioxide CO 2 . Just as CO 2 forms carbonates with oxides of alkalies, so CS 2 forms analogous compounds, namely trithiocarbonates, with sul- phides of alkalies : CaO+CO 2 =CaCO 3 . CaS + CS 2 = CaCS 3 . K 2 S + CS 2 =K 2 CS 3 . By treating a trithiocarbonate with a dilute acid, trithiocar- bonic acid H CSo is liberated as an unstable oil which decom- & o poses readily into CS 2 and H 2 S, thus showing great analogy to H 2 C0 3 : H 2 C0 3 =H 2 + C0 2 . H 2 CS 3 =H 2 S + CS 2 . Cyanogen. Under ordinary conditions, carbon and nitrogen do not combine with each other ; but at high temperatures in presence of carbonates of the alkalies or oxides of the alkaline earth metals, cyanides are formed. These are compounds con- taining carbon, nitrogen, and the alkali metal employed. Thus, nitrogen passed over a mixture of carbon and fused potassium carbonate yields potassium cyanide : K 2 C0 3 + 3 C + N 2 = 2 KCN + CO + CO a . In this reaction, metallic potassium is formed, which then unites with the carbon and nitrogen to form potassium cya- nide. Whenever any carbon compound containing nitrogen is heated with metallic potassium, potassium cyanide results, which fact is used as a test for nitrogen in organic compounds. When calcium oxide is employed, the cyanide of calcium is formed : CaO + 3 C + N a = Ca(CN) 2 + CO. On passing ammonia over carbon heated to redness, ammonium cyanide is produced : 2NH 3 +C=NH 4 CN + H 2 . The cyanides are salts derived from hydrocyanic acid HCN (which see)-. On heating mercuric cyanide Hg(CN) 2 , it decom- poses, yielding mercury and cyanogen, thus: Hg(CN) 2 =Hg+(CN) r CARBON AND SOME OF ITS TYPICAL COMPOUNDS 241 This reaction is quite similar to that of the decomposition of mercuric oxide by heat. Cyanogen is an extremely poisonous, colorless gas of sharp characteristic odor. It may be condensed to a colorless liquid that boils at 21. The solid melts at 35. The gas burns with a beautiful purple flame. In water, the gas is readily soluble, also in alcohol. Cyanogen gets its name from the fact that it enters into a number of com- pounds that are blue in color. Hydrocyanic Acid. Hydrocyanic acid, or Prussic acid, has the composition HCN. It is formed when potassium cyanide is treated with hydrochloric acid : KCN + HC1 = KC1 + HCN. It may also be prepared by treating potassium ferrocyanide K 4 Fe(CN) 6 -f- 3 H 2 O with dilute sulphuric acid, thus : 2 K 4 Fe(CN) 6 + 3 H 2 SO 4 = 6 HCN + 3 K 2 SO 4 + K 2 Fe Fe(CN) 6 . The potassium ferrocyanide is prepared by heating animal refuse, like blood, hoofs, horns, etc., with iron and potassium carbonate. The product, when purified, forms beautiful lemon- yellow crystals of the composition K 4 Fe(CN) 6 + 3 H 2 O. It is known also as the yellow prussiate of potash. Hydrocyanic acid is an extremely poisonous, colorless, mobile liquid, which smells like bitter almonds. It boils at 26. The colorless crystalline solid melts at 11. Hydrocyanic acid is a very weak acid which is readily soluble in water and alcohol. Its salts are also very poisonous. One twentieth of a gram of hydrocyanic acid is sufficient to cause death in case of a human being. The best antidote consists of a 3 per cent solution of hydrogen peroxide, which acts thus : H 2 2 + 2 HCN = H 2 NOC - CONH 2 . The latter compound is called oxamide. Hydrocyanic acid is used to kill insects that infest shrubs and fruit trees. It is also at times used in medicine. It affects chiefly the respiratory organs. Cyanates and Sulphocyanates. On heating potassium cya- nide with lead oxide, potassium cyanate KCNO is formed : KCN + PbO = Pb + KCNO. 242 OUTLINES OF CHEMISTRY The free cyanic acid HCNO is a liquid which readily decom poses into ammonia and carbon dioxide when treated witli water : H 2 O + HCNO = N H 3 + CO 2 . Fused with sulphur, potassium cyanide forms potassium sulphocyanate KCNS : KCN + S = KCNS. The free acid HCNS is extremely unstable. With ferric salts potassium sulphocyanate forms ferric sulphocyanate Fe(CNS) 3 , which is blood-red in solution, and serves as a delicate test for ferric salts. The fact that cyanides pass over into cyanates and sul- phocyanates shows that cyanides are really unsaturated in character. REVIEW QUESTIONS 1. What are the different allotropic forms of carbon? How do we know that the different substances you have mentioned consist of carbon only? 2. How is artificial graphite produced? Mention some of its uses. What use is made of the other forms of carbon you have named in answer to question 1. 3. Discuss the occurrence, kinds, and supposed origin of coal. 4. Describe three methods of preparing each of the two oxides of carbon, giving the equations expressing the changes that take place. 5. How much carbon dioxide could be prepared from 37 grams of carbon ? What volume would this gas occupy under standard conditions ? What was the volume of the oxygen consumed in preparing the carbon dioxide ? 6. What are the chief chemical characteristics of carbon? 7. What use is made of carbon dioxide ? Upon what property of this substance does each use mentioned depend ? 8. What is a carbonate? Give five illustrations, and write the equation showing how each reacts when treated with hydrochloric and sulphuric acid respectively. Compare this with the action of these acids upon the corresponding sulphites. 9. What is a bicarbonate? Give two examples. Are bicarbonates of practical importance? Why? How do carbonates and bicarbonates behave when heated ? 10. Why does the percentage of carbon dioxide in the atmosphere not increase continually due to the breathing of animals and the combustion of fuel? CARBON AND SOME OF ITS TYPICAL COMPOUNDS 243 11. What is meant by the carbon cycle ? 12. What is the valence of carbon in sodium carbonate? In carbon monoxide ? What use is made of the latter gas ? 13. What is water gas? Write the equation expressing its mode of formation. What is producer gas, and how is it made ? 14. How much carbon disulphide could be made from two tons of sulphur ? How could this be done ? What use could be made of the carbon bisulphide? Why is the latter a dangerous substance? 15. Describe a way of making Prussic acid. What other name has the substance ? What are its characteristics and uses ? 16. Give the formulas of the following and write one characteristic reaction in which each occurs : potassium ferrocyanide, potassium sulphocyanate, mercuric cyanide. 17. State the volume relations in each of the following equations : C0 2 +C=2CO; S+O 2 =S0 2 ; 2NO+0 2 = 2N0 2 ; 2 H 2 S + 3 2 = 2 S0 2 + 2 H 2 0; 2 NH 3 + 3 C1 2 = 6 HC1 + N 2 . 18. How much barium carbonate can be made from 18 grams of barium hydroxide ? 19. How much calcium sulphate may be formed by treating 25 pounds of calcium carbonate with sulphuric acid, and how much acid would this require? 20. Write the equation showing how nitric acid acts when used as an oxidizing agent, and then use this equation to show the oxidation of car- bon to carbon dioxide by means of nitric acid. CHAPTER XV HYDROCARBONS AND ADDITIONAL COMPOUNDS OP CARBON Hydrocarbons. Hydrocarbons are compounds of hydrogen and carbon. They are very numerous, nearly three hundred being known. The simplest hydrocarbon is marsh gas, or methane CH 4 . It issues from the decaying vegetable matter in ditches and marshes on warm summer days, hence its popular name, marsh gas. Methane, together with hydrogen, is a prime constituent of natural gas, which is found in the coal and oil regions of Indiana, Ohio, Pennsylvania, and other states. Methane may be prepared artificially by a process to be described later. Petroleum is essentially a mixture of hydro- carbons. The chief petroleum fields are located in Pennsyl- vania, New York, Ohio, Indiana, Kentucky, Kansas, Texas, Colorado, California, and Canada. Less extensive deposits are found in Russia near the Caspian Sea and Black Sea, in China, India, and Japan. The hydrocarbons in American petroleum practically all belong to the so-called paraffin series, of which methane is the first and simplest member. The fol- lowing table gives a number of hydrocarbons of this series : NAME FORMULA BOILING POINT Methane .... CH 160 Etliane Vxll 4 C H fl - 93 Propane . . . V ^2 IA 6 C H Q - 45 Butane .... v - y 3 i '*8 C II + 1 Pentane ^4 10 ^5^12 37 Hexane 69 Heptane Octane ^6 14 C 7 H 1C ^8 18 98 125 C ft H 149.5 It will be noted that the boiling points increase with the car- bon content, and that the first three substances are gases at 244 HYDROCARBONS AND THEIR DERIVATIVES 245 ordinary temperatures. The difference in composition between any two adjacent members is CH 2 , and any series of compounds in which this is the case is called an homologous series. The series above given is the so-called normal paraffin series. The general formula for any of its members is C n H 2w + 2 . Of this series, compounds containing up to sixty carbon atoms are known. The higher members are solids. Thus pentadecane C 15 H 32 melts at +10; eicosane C 20 H 42 melts at 36.7; heptacosane C 27 H 56 melts at 59.5 ; and hexacontane C 60 H 122 melts at 102. In the members of this series, carbon is quadrivalent, and the compounds are called saturated hydrocarbons. When crude petroleum is placed in a retort and subjected to fractional distillation the following fractions are obtained : Cymogene (mainly butane), B. P. about C. Bbigolene (butane and pentane), B. P. about 16. Petroleum ether (pentane and hexane), B. P. about 50 to 60. Gasoline (hexane and heptane), B. P. about 70 to 90. Naphtha (heptane and octane) B. P. about 90 to 120. Benzine (octane and nonane), B. P. about 110 to 140. Kerosene (nonane to heptadecane), B. P. about 150 to 300. Naphtha is also sometimes called ligroin. Above 300 heavy oils pass over which are used as lubricating oils. At still higher temperatures, vaseline is obtained, which consists mainly of C 19 H 40 to C 21 H 44 . Finally, from the residue in the retort paraffin is separated out at low temperatures. Paraffin is a white waxlike substance consisting of C 21 H 44 to C 32 H 66 and melting from 45 to 70 C. according to composition. Petroleum ether is used as a solvent, also in making illumi- nating gas, and as an anaesthetic. For the latter purpose it must be specially purified. Sometimes a mixture of the first four products named in the last table is called petroleum ether. Gasoline, naphtha, and benzine are also used as solvents and frequently as fuels in stoves and engines. They are also em- ployed in gas manufacture. Benzine is frequently used in paints and varnishes as a substitute for turpentine. Kerosene is used as a fuel and for purposes of illumination. There are different grades of kerosene. These vary as to color and fire test or flash point; i.e. the temperature at which evaporation is sufficient so that the vapors may be lighted in the air. Kero- sene having a flash point of 110 F. is safe for use in lamps; 246 OUTLINES OF CHEMISTRY this is the standard flash test fixed by law in most of the states, but sometimes a test of 150 F. is required. In purify- ing kerosene it is washed with sulphuric acid, then with an alkali, and finally with water. Paraffin lubricating oils are now very much used particularly in gas and gasoline engines. Vaseline is used in ointments of various kinds. In crude form, it serves as cup grease and axle grease. Paraffin is used in making candles and chewing gum, also in making paper and fabrics waterproof, in insulating wires and electrical apparatus of various kinds, in manufacturing matches, and in the laundry. Crude oils from the Texas and California fields are used for the preservation of railroad ties, which are saturated with the oils before using. Hydrocarbons may be prepared in several ways. The com- mon method of making methane is by heating sodium acetate with lime or caustic soda : CHgCOONa + NaOH = Na 2 CO 3 + CH 4 . The higher hydrocarbons may be prepared in a similar way by using salts of acids of higher carbon content. Thus to prepare propane, sodium butyrate is heated with a caustic alkali : C 3 H 7 COONa + NaOH = Na a CO 8 -+ C 3 H 8 . Hydrocarbons may also be formed by treating carbides of metals with water. Thus aluminum carbide and water yield meth- ane : A1 4 C 3 + 12 H 2 = 4 A1(OH) 3 + 3 CH 4 . The carbide of aluminum may be prepared by heating oxide of aluminum with carbon in the electric furnace, when the follow- ing reaction occurs : 2 A1 2 O 3 + 9 C = 6 CO + A1 4 C 3 . Other simple methods of preparing hydrocarbons consist of treating halogen substitution products with nascent hydrogen or with sodium : CH 3 I + 2 H = HI + CH 4 . methyl iodide methane 2 CH 3 I + 2 Na = 2 Nal + C 2 H 6 . ethane Hydrocarbons act neither as acids nor as bases, and thus they differ materially from the hydrides of such elements as nitro- gen, sulphur, and the halogens. HYDROCARBONS AND THEIR DERIVATIVES 247 Bthylene C 2 H 4 , also called olefiant gas, is the first of an homologous series of unsaturated hydrocarbons whose general formula is C w H 2re . Ethylene may be prepared by heating alco- hol with concentrated sulphuric acid, which abstracts water from the alcohol : C 2 H 5 OH + H 2 S0 4 = H 2 S0 4 -H 2 + C 2 H 4 . The relation, of ethylene to ethane C 2 H 6 is readily seen from the following graphic formulae : H H H H H H H-C-C-H, C = C, or - C - C -. H H H H H H ethane ethylene When ethane is treated with bromine, substitution takes place : C 2 H 6 + Br 2 = C 2 H 5 Br + HBr ; ethylene, however, reacts as follows : C 2 H 4 + Br, = C 2 H 4 Br 2 . That is, . ethylene bromide C 2 H 4 Br 2 , a colorless liquid, is formed, showing that ethylene is unsaturated, the two free bonds manifesting themselves in the fact that the hydrocarbon is able to unite with two atoms of bromine by simple addition. This is characteristic of all of the members of the ethylene series, and is often expressed by the so-called double bond, I I C = C , which, however, is not a source of additional strength. It simply indicates that the compound is unsatu- rated, i.e. is capable of forming addition products, as above illustrated. Ethylene gas burns with a luminous flame. Acetylene C 2 H 2 is the lowest member of an homologous series of hydrocarbons of the general formula C w H 2ra _ 2 . Acety- lene gas may be produced by passing ethylene through red-hot tubes : C 2 H 4 = C 2 H 2 + H 2 . It is also formed to some extent when a Bunsen burner burns below, that is, has "struck back." However, the gas is best prepared by the action of calcium carbide on water : CaC ?i + 2 H 2 = Ca(OH) 2 + C 2 H a . 248 OUTLINES OF CHEMISTRY Acetylene burns with a very bright, luminous flame. It is conse- quently often used for illuminating purposes. Acetylene is still less saturated than ethylene, which fact is expressed by I I the formula HC CH, or more frequently by H C = C H. I I The triple bond, or acetylene bond, indicates the unsaturated condition of the compound, i.e. its ability to unite with four additional atoms of halogen, for instance. Benzene C 6 H 6 (not to be confounded with the petroleum product, benzine, which is entirely different) is a colorless, mobile liquid of specific gravity 0.8799 at 20, boiling at 80. The 'compound forms crystals that melt at 6. Benzene is obtained as the light oil from coal tar. It is the first member of an homologous series of hydrocarbons of the general formula C n H 2w _ 6 . It burns with a luminous flame, and is an excellent solvent for fats, resins, iodine, sulphur, and phosphorus. From benzene many very important substances like carbolic acid, aniline, the coal tar dyes, and many medicinal and aromatic substances are derived. Because of the aromatic odor of the hydrocarbons of the benzene series and their derivatives, they are commonly called the aromatic series. The hydrocarbons of the paraffin series and their derivatives are also known as the fatty series, for the fats commonly belong to this series. Naphthalene C 10 H 8 forms shining leaflets that melt at 79. It occurs in coal tar and gives the latter its peculiar odor. In form of moth balls naphthalene is used to protect woolen goods from moths. Naphthalene is closely related to benzene. General Behavior of Hydrocarbons. All hydrocarbons burn, and when the combustion is complete, the products formed are car- bon dioxide and water. G-aseous hydrocarbons, or the vapors of all light hydrocarbon oils, are inflammable. With oxygen or air, they form mixtures that explode when brought in the neigh- borhood of a flame. The larger the carbon content of the molecule of a hydrocarbon, the more luminous is the flame with which it burns (see luminosity of flames). On the whole, hydrocarbons are rather inert substances chemically. They are practically insoluble in water, but they are miscible with one another in all proportions. They dissolve fats, oils, ether, alcohol, carbon disulphide, and many other sub- stances of kindred character. HYDROCARBONS AND THEIR DERIVATIVES 249 Halogen Substitution Products. When methane is treated with chlorine in the sunlight, the hydrogen atoms may be re- placed one after another by chlorine as indicated by the follow- ing equations : CH 4 +C1 2 =CH 3 C1 + HC1. methyl chloride CH 3 C1 + C1 2 = CH 2 C1 2 + HC1. dichlormethane CH 2 C1 2 + C1 2 = CHClg + HC1. chloroform CHClg + C1 2 = CC1 4 + HC1. carbon tetrachloride Chloroform, a heavy colorless liquid boiling at 61, is very important as an anaesthetic. The corresponding bromine com- pound, bromoform CHBr 3 , is also known. It is a liquid boiling at 151; while iodoform CHTg, a yellow crystalline solid melt- ing at 119, is used in dressing wounds. Carbon tetrachloride boils at 76. It rs not inflammable like gasoline, though like the latter it dissolves fats readily. Hence carbon tetrachloride is often used as a solvent, particularly for cleaning clothes, being less dangerous to handle than volatile hydrocarbons. We have already seen that unsaturated hydrocarbons may form simple addition products with halogens. Alcohols. On treating methyl chloride with caustic potash, the following reaction occurs : CHgCl + KOH = KC1 + CHgOH. methyl alcohol The radical CH 3 is called methyl. There are many similar hydrocarbon radicals, which are also called alkyl radicals. So, for instance, we have : ethyl C 2 H 5 , in ethyl chloride C 2 H 6 C1 ; propyl C 3 H 7 , in propyl iodide C 3 H 7 I ; phenyl C 6 H 5 , in phenol C 6 H 5 OH. These radicals may pass from one compound to an- other precisely as the radical ammonium NH 4 does in ammonium compounds. Now C HgOH or methyl hydroxide is methyl alcohol. It is also called wood alcohol, or spirit of wood, for it may be obtained as one of the products of the dry distillation of wood. We note that methyl alcohol is an hydroxide. Indeed, we may regard it as water with one hydrogen atom, replaced by methyl, or as 250 OUTLINES OF CHEMISTRY sodium hydroxide with the sodium -atom replaced by methyl All alcohols are hydroxides of alkyl radicals, and the general for- mula of an alcohol is R OH where R stands for an alkyl radical. Thus, C 2 H 6 OH, or ethyl alcohol, is ordinary alcohol; C 3 H 7 OH is propyl alcohol; C 4 H 9 OH, butyl alcohol; and so on up the homol- ogous series, the higher members of which, like C 16 H 33 OH, cetyl alcohol, and C 30 H 61 OH, melissyl alcohol, are waxlike solids. Ethyl alcohol, or ordinary alcohol, is also called spirit of wine, for it may be obtained by distilling wine. When thus pre- pared it contains other aromatic substances from the wine. Alcohol is contained in all fermented liquors. It may readily be prepared by fer- mentation of glucose with yeast, thus : C 6 H 12 6 =2C 2 H 6 OH glucose alcohol FK;. 89. Figure 89 shows a common form of yeast cells as they appear under the microscope. Pure alcohol boils at 78 and solidifies at about 130. Beers contain from 3 to 5 per cent alcohol, wines from 8 to 20 per cent, and brandy, whisky, and rum from 45 to 65 per cent. By adding wood alcohol or other poisonous substances to ethyl alcohol, the latter is made unfit for use as a beverage, and is said to be "denatured." Usually about 10 volumes of wood alcohol and half a volume of benzene are added to 100 volumes of 90 per cent alcohol to make so-called denatured alcohol. The latter may be used as fuel or for purposes of manufacturing, without the payment of duty. When ethyl alcohol is treated with phosphorus trichloride, the following change takes place : 3 C 2 H 6 OH + PC1 8 = P(OH) 3 + 3 C 2 H 5 C1. Thus chlorine is substituted for the OH group. All alcohols undergo a similar change when treated with either phosphorus chloride, bromide, or iodide. The halogen takes the place of the HYDROCARBONS AND THEIR DERIVATIVES 251 hydroxyl, and phosphorous acid is formed. Indeed, this fact is used to ascertain whether the OH group is present in a com- pound. The number of such groups may also be determined thus by the number of halogen atoms that enter the molecule. Water itself reacts perfectly analogously with phosphorus trichloride, as is evident from the following equations : 3 H 2 + PC1 8 = P(OH) 3 + 3 HC1. 3 C 4 H 9 OH + PC1 3 = P(OH) 3 + 3 C 4 H 9 C1. When treated with sodium, the hydrogen of the OH group of an alcohol is replaced, thus : 2 C 2 H 6 OH + 2 Na = 2 C 2 H 5 ONa + H 2 . sodium alcoholate This is perfectly analogous to what happens when water is treated with sodium : 2 H 2 + 2 Na = 2 HONa + H a . Just as we have hydroxides of the metals which contain more than one hydroxyl group, like Ca(OH) 2 , Bi(OH) 3 , Sn(OH) 4 , so we also have alcohols that contain two or more hydroxyl groups. Thus, we have : CH 2 -OH I CH-OH CH 2 -OH I CH 2 -OH | CH-OH glycol, CH OH glycerine, and | mannite. | CH-OH CH 2 -OH | CH-OH CH-OH I CH 2 -OH Glycol and glycerine are rather viscous liquids that mix with water in all proportions. They are slightly sweet. Moreover, the sweet taste increases as we go up the series from glycol to mannite. Erythrite, which contains four carbon atoms and four hydroxyl groups, and arabite, which contains five carbon atoms and five Ir^droxyl groups, are also well known, though they are of no practical importance. These alcohols are called polyhydric alcohols. They may act towards acids like bases, forming salts which are called esters (which see). Mannite is 252 OUTLINES OF CHEMISTRY a beautifully crystalline substance which readily dissolves in water, but not in hydrocarbons. It is closely allied to the sugars (which see). Taken internally, mannite acts as a mild purga- tive. All the polyhydric alcohols are soluble in water, and from erythrite up they are solids under ordinary conditions. Phenols. The hydroxyl derivatives of benzene C 6 H 6 are called phenols. The simplest of these is C 6 H 5 OH. It is called phenol, or more commonly carbolic acid. It is really not an acid, though it does exhibit acidic properties to some extent. Thus with caustic alkalies it forms phenolates : C 6 H 5 OH + KOH = C 6 H 5 OK + H 2 O. potassium phenolate Alcohols do not form alcoholates when treated with caustic alkalies. It will be recalled that it is necessary to treat alcohols with metallic sodium or potassium to form the corresponding alcoholates. Carbolic acid crystallizes in long needles that melt at 42 and turn pink when exposed to the air. In water it is but spar- ingly soluble, about 1 part dissolves in 15 parts of cold water ; but in alcohol and many other organic liquids it dissolves much more copiously. Carbolic acid has a characteristic odor and is very poisonous, whence its use as a disinfectant and antiseptic. When brought in contact with the skin it exerts a corrosive action. Among the phenols containing more than one hydroxyl group hydroquinone C 6 H 4 (OH) 2 and pyrogallol C 6 H 3 (OH) 3 , also called pyrogallic acid, are of importance as developers in photography (which see). Further, in the incomplete combus- tion and dry distillation of wood there are formed, along with other products, phenols, notably guajacol C 6 H 4 (OCH 3 )OH and kreosol C 6 H 3 (CH 3 )(OCH 3 )OH, a mixture of which is called creosote. These give the smoke a penetrating odor and anti- septic value that is used in preserving meats, sausages, and fish. Aldehydes. On careful oxidation of alcohols, aldehydes are formed, which are compounds containing two hydrogen atoms less than the alcohol from which they may be obtained, whence the name alcohol dehydrogenatum, abbreviated aldehyde. So when methyl alcohol is partially oxidized by means of potassium ' ., HYDROCARBONS AND THEIR DERIVATIVES 253 permanganate, or incomplete combustion, the following change takes place : -r- H I CH 8 OH + = H 2 + C = O. I H The product, formaldehyde, is a gas which may be condensed to a liquid that boils at 21. It has a penetrating, suffocating odor, acts strongly on the mucous membranes, and is a powerful antiseptic. In water it dissolves up to about 40 per cent, and it is this aqueous solution that is sold under the name of forma- line. It is used as a disinfectant, antiseptic, or preservative in various strengths as required. All the aldehydes are very active chemically, being specially strong as reducing agents. We may regard formaldehyde as carbon dioxide with one oxygen atom replaced by two hydrogen atoms (see formula above). By some it is thought that formaldehyde is the first product formed when water and carbon dioxide act upon each other in the green leaf of the plant in the sunlight, forming starch and liberating oxygen. Just as the hydrocarbons, their halogen substitution products, and the alcohols form homologous series, so the aldehydes form similar series. Thus, we have formaldehyde HCHO, acetic aldehyde CH 3 CHO, propionic aldehyde C 2 H 5 CHO, etc. ; further, benzole aldehyde C 6 H 5 CHO, also called oil of bitter almonds, and its homologues. When the hydrogen atoms of the methyl group in acetic alde- hyde are replaced by chlorine, CC1 3 -CHO is formed. This is trichloraldehyde, or chloral. It readily unites with water, form- ing the white crystalline addition product CC1 3 CH(OH) 2 , chloral hydrate, which is so much used in medicine as a soporific. The general formula of an aldehyde is R C = O, in which R I H represents either hydrogen or an alkyl radical. Organic Acids. On further oxidation of aldehydes, acids are produced. In this process one molecule of aldehyde takes up an additional atom of oxygen. The reactions in the formation of formic and acetic acids from formic and acetic aldehydes may be represented as follows : CALIf-OHNIA COLLtGE of PHARMACY . 254 OUTLINES OF CHEMISTRY H-C = O H-O= H-C = O; I I H OH formic aldehyde formic acid CH 3 - C = O + O = CH 3 - C = O. I I H O-H acetic aldehyde acetic acid In general, the reaction may be represented thus : R_C = O +O= R - C = O, I I H OH aldehyde organic acid R representing either hydrogen or any organic radical. The last formula above given in the general formula of an organic acid. The characteristic group which it contains, namely, C = O I OH or COOH, is called the carboxyl group. The hydrogen in this group is replaceable by metals or radicals, just as, for instance, the hydrogen in nitric acid may be thus replaced. Only a few typical organic acids can be mentioned here. Formic acid H-COOH occurs in red ants and stinging nettles. It is a colorless liquid boiling at 101. It is soluble in water in all proportions, has a pungent odor, and blisters the skin. With bases it forms the formates, thus : HCOOH + NaOH = HCOONa + H 2 O. sodium formate The acid is consequently monobasic, only one hydrogen atom being replaceable by a metal. When heated in closed vessels to 160, formic acid yields carbon dioxide and hydrogen : HCOOH = CO 2 + H 2 . Sodium formate may be obtained by passing carbon monoxide over heated caustic soda, and from the formate the free acid may be obtained by means of sulphuric acid, thus : NaOH f CO = HCOONa; HCOONa + H 2 SO 4 = NaHSO 4 + HCOOtt. Acetic acid CH 3 COOH occurs in combination with organic radicals in many odoriferous plant oils. It is formed as one of HYDROCARBONS AND THEIR DERIVATIVES 255 the products of the dry distillation of wood. It is made on a large scale in vinegar factories, the process depending on (1) the formation of alcohol by fermentation of sugar pro- duced from the starch in grain, and (2) the oxidation of this alcohol to acetic acid, which is brought about by Mycoderma aceti, "mother of vinegar," a bacterium shown in Fig. 90. In practice, the dilute alcohol, 8 to 15 per cent, is allowed to trickle over beech wood shavings con- tained in a barrel. Thus the alcohol is thoroughly exposed to the oxygen of the air, and the acetic-acid-forming bacteria cause the oxidation to take place. The oxidation process may be represented thus : FIG. 90. C 2 H 5 OH 2 = H 2 CH 3 COOH. When alcohol is treated with oxygen alone, this process does not occur ; however, the spores of the acetic-acid-forming bac- teria are commonly present in the air, and so various alcoholic solutions like beer and wines get sour because of the oxidation of the alcohol to acetic acid. Cider slowly ferments, forming alcohol, which is then similarly converted into acetic acid. Vinegar obtained as above described is a solution containing essentially 4 to 10 per cent acetic acid. Anhydrous acetic acid may be obtained from this solution by forming sodium acetate and then treating the latter salt with sulphuric acid and dis- tilling. The reactions are : 2 CHgCOOH (in solution) CHgCOONa (dry salt) Na a C0 H 2 O CO + 2 CHCOONa. H 2 SO 4 = NaHSO 4 + CH 3 COOH. (glacial acetic acid) The process is then quite similar to the preparation of hydro- chloric acid. Pure acetic acid boils at 118. It forms crystals that mel^ at 16.5, so that it is easy to cause it to solidify on a cold day, 256 OUTLINES OF CHEMISTRY whence the popular name glacial acetic acid. The acid has a pungent odor and corrodes the skin. It is a monobasic acid, only one hydrogen atom being replaceable by a basic element or radical. The salts formed are called acetates. Among the homologues of acetic acid are propionic acid C 2 H 5 COOH, butyric acid C 3 H 7 COOH, which smells like rancid butter (in which it occurs), palmitic acid C 15 H 31 COOH, and stearic acid C 17 H 35 COOH. The last two are solids at ordinary temperatures. They occur in fats and oils of plant and animal origin, commonly together with oleic acid C 17 H 33 COOH, an unsaturated acid belonging to another homologous series. These acids occur in fats and oils as esters (which see), not as free acids. Benzoic acid C 6 H 6 COOH is the first of an homologous series of acids derived from benzene. It occurs in gum benzoin and Peru balsam ; and a derivative of benzoic acid, hippuric acid C 6 H 6 .CO-NH.CH 2 .COOH, occurs in the urine of herbivorous animals. Benzoic acid crystallizes in shining flakes that melt at 121.5. The boiling point is 249, but the acid sublimes readily even at 100. In alcohol, ether, and hot water it readily dissolves, while in cold water it is but sparingly soluble, 1 part in 400. Its odor is very characteristic, and its vapor when inhaled irritates the throat and nasal passages, causing coughing and sneezing. It is much used as an antiseptic and preserva- tive, commonly in the form of sodium benzoate C 6 H 5 COONa. The free acid forms readily by oxidation of benzaldehyde. Oxalic acid is the simplest dibasic organic acid. Its formula is HoC 2 O 4 , or COOH I COOH that is, it consists of two carboxyl groups. It is readily formed by oxidizing sugar, sawdust, glycol, fats, and many other organic substances with strong nitric acid. On a commercial scale, sawdust is employed. The normal sodium salt may be obtained by passing carbon dioxide over sodium at 350, thus: 2 CO 2 + 2 Na= (COONa) 2 , i.e. Na 2 C 2 O 4 . This salt is also formed when sodium formate is quickly heated to 250, thus : - 2 HCOONa = H 2 + Na a C a O 4 . HYDROCARBONS AND THEIR DERIVATIVES 257 The calcium salt is difficultly soluble in water, and this fact is often used in analytical chemistry, in detecting and estimating calcium. When calcium oxalate is heated, there is first formed calcium carbonate and carbon monoxide, after which the car- bonate decomposes to calcium oxide and carbon dioxide or, further heating, thus : CaC 2 O 4 = CaCO 3 + CO, CaCOg = CaO + CO 2 ; or CaC 2 O 4 = CaO + CO + CO 2 . This behavior is typical of the oxalates of many other metals. Oxalic acid is oxidized to carbon dioxide and water by oxidizing agents like potassium permanganate, for example, thus : COOH I + O = 2 CO 2 + H 2 0. COOH It is evident, then, that oxalic acid is a good reducing agent. In the market the acid is commonly sold in form of its white crys- talline hydrate H 2 C 2 O 4 -2H 2 O. By heating this, the anhy- drous acid may be obtained, which on further heating passes over into carbon dioxide and formic acid : (COOH) 2 = CO 2 + HCOOH. Just as there are homologous series of monobasic organic acids, so there are homologous series of the dibasic acids. So we have malonic acid CH 2 (COOH) 2 , succinic acid (CH 2 COOH) 2 , etc. Hydroxycarboxylic acids. The simplest of these is glycolic acid, or liydroxyacetic acid, CH 2 OHCOOH, which may be pro- duced by careful oxidation of glycol : CH 2 OH CH 2 OH I +20=1 + H 2 0. CH 2 OH COOH glycol glycolic acid Glycolic acid is monobasic, only the hydrogen atom in the carboxyl group being replaceable by a basic element or radical. The other OH group is alcoholic in character. This compound is therefore both an alcohol and an acid. Its homologue, C 2 H 4 -OH-COOH, is lactic acid. It is of great practical 258 OUTLINES OF CHEMISTRY importance. In its pure form it is a thick, colorless, odorless, hygroscopic liquid of pronounced acidic properties. The acid is monobasic, and forms the lactates with bases, thus : C 2 H 4 . OH COOH + KOH = H 2 O + C 2 H 4 OHCOOK. Lactic acid is readily miscible with water, alcohol, and ether in all proportions. Like other hydroxyacids, it cannot be dis- tilled, for it decomposes into various simpler products like CO, H 2 O, CH 3 CHO, etc., on heating. Lactic acid is the acid that causes the acidity of sour milk, whence the name lactic acid. Lactic acid is produced by a special form of fermentation caused by lactic-acid-forming bacteria. These are shown in Fig. 91, together with a few yeast cells to indicate approxi- mately the relative size of the organisms. By the action of these bacteria, starch and sug- ars are converted into lactic acid, which is consequently formed in many liquids that contain starch, sugars, or kin- dred organic substances. So, for example, lactic acid occurs in sauerkraut, in sour pickles, in fermented beet juice, and at times in the contents of the alimentary tract. The lactates of strontium and of silver are used in medicine. Lactic acid also occurs in muscular tissues, but this acid is not quite identical with that in sour milk. So the acid in the muscles, sarcolactic acid, has the power to rotate the plane of polarized light ; that is, it possesses optical activity, a property not exhibited by lactic acid produced by fermentation, which is consequently termed inactive lactic acid. Sarcolactic acid turns the plane of polarized light to the right. When the fungus Penicillium glaucum is allowed to grow in solutions of fermentation lactic acid, the latter also acquires the power to turn the plane of polarized light to the right, i.e. it becomes dextroactive, and is in every way identical with sarcolactic acid. Now when the bacillus Acidi Icevolactici feeds upon FIG. 91. HfDROCARBONS AND THEIR DERIVATIVES 259 solutions of fermentation lactic acid, the latter acquires the power to turn the plane of polarized light to the left, that is, it becomes Icevoactive. Fermentation lactic acid is conse- quently a mixture of equal parts of dextro and Isevo lactic acid, which accounts for its optical inactivity. The optical activity is produced as described, because one organism feeds on the dextro variety of the acid, whereas the other organism lives upon the leevo variety. The formula of lactic acid is H I H- C- H I H - C - OH. I c = o O- H This, like all other chemical formulae, has been established by a study of the methods of synthesizing the compound, and by its deportment toward various reagents. It will be observed that lactic acid has one carbon atom, the center one, whose four bonds are satisfied by four different elements or radicals. Such a carbon atom is called an asymmetric carbon atom. All carbon compounds that are optically active possess at least one asymmetric carbon atom in the molecule; though, as we have seen, the pos- session of such an asymmetric carbon atom does not necessarily make the compound optically active, for the substance under consideration may be a mixture of equal parts of the dextro and Isevo varieties, as in the case of fermentation lactic acid. In general, for every dextroactive compound there is a Isevo- active compound that rotates the plane of polarized light in the opposite direction to the same degree. In order to repre- sent the difference between dextro and Isevo compounds by means of formulse, Le Bel and van't Hoff simultaneously and independently of each other (1874) proposed to represent the carbon atom by a regular tetrahedron the corners of which in- dicate the four valences. Figure 92 shows the two formulse for dextro and Isevo lactic acids. It will be observed that the formulse are alike, but not superposable ; that is, they are to each other as the right hand OUTLINES OE CHEMISTRY is to the left hand, or as an object is to its image in the mirror The study of the composition of carbon compounds by the aid of formulae thus represented in three dimensions has been carried on with marked success, and has in. recent years, also been extended to the investigation of compounds of other elements, notably those of nitrogen. The branch of chemistry which seeks to further inquiry by the use of formulae in three dimensions is called stereo-chemistry. Compounds that contain the same elements in the same pro- portions by weight are said to be isomers. So acetic acid CHgCOOH and methyl formate HCOOCH 3 are isomers, for HO COOH Cti COOH FIG. 92. they contain the same elements in the same proportions by weight, and yet they are quite different compounds. Dextro and Isevo lactic acids also contain the same elements in the same proportions by weight, and in addition they are in every way identical in their chemical behavior. The difference be- tween them lies simply in their behavior toward polarized light and they are consequently called optical isomers, or stereo-iso- mers, and the property they thus exhibit is called stereo-isom- erism. When an aqueous solution of formaldehyde HCHO is evapo- rated, an amorphous white substance (HCHO) 2 , paraformalde- hyde, separates out. On careful heating, this may be transformed into metaformaldehyde (HCHO) 3 , a crystalline compound melt- ing at 171, which on being heated to 140 with much water is again decomposed into formaldehyde. Other aldehydes exhibit similar properties. The process of forming larger molecules by simple aggregation of two or more molecules is called poly- merization, and the new products formed are said to be polymers. Polymerization is quite a common process. It will be noted that polymers contain the same elements in the same propor- HYDROCARBONS AND THEIR DERIVATIVES 261 tions by weight as the simple compounds from which they have sprung; they are consequently isorners as above defined. When, however, the molecular weight of two isomeric compounds is identical, they are said to be metameric, in contradistinction to polymeric, which latter term is applied when the isomerism depends upon difference in. molecular weight, as already ex- plained. The optical activity of substances is examined by means of a polariscope, also called a polarimeter, a common form of which FIG. 93. is shown in Fig. 93. The substance to be tested, commonly in the form of a liquid or solution, is placed in a tube between the polarizer and analyzer; that is, the nicol prisms which consti- tute the vital part of the instrument. The degree of rotation is read off on the scale attached to the analyzer, the yellow sodium flame being usually employed as a source of monochro- matic light. The rotation is proportional to the length of the tube and the concentration of the solution or other liquid employed. By concentration is meant the number of grams of active substance contained in one cubic centimeter. The 262 OUTLINES OF CHEMISTRY specific rotatory power of any substance is the number of degrees of rotation it exhibits in a tube 1 decimeter long when 1 gram of active substance is contained in 1 cc. Many important substances of commerce, notably sugars, essential oils, and other com- pounds obtained from plants, possess optical activity. The polariscope serves in detecting the presence of such substances, and also in estimating their amounts when present. So, for example, in estimating the amount of sugar in the juice of sugar beets, the polariscope enables one to obtain very rapid and accurate results. Malic, tartaric, and citric acids are important fruit acids. Malic acid occurs in sour apples, in mountain ash berries, and in many other fruits. It is monohydroxysuccinic acid : H I HO - C - COOH I H-C-COOH I H From the formula it is evident that the acid is dibasic, and that its molecule contains an asymmetric carbon atom. The acid is optically active. This is also true of tartaric acid, which is dihydroxysuccinic acid: H I HO - C - COOH I HO - C - COOH I H This acid occurs in grapes as the acid potassium salt C 4 H 5 O 6 K, which is difficultly soluble in water and is commonly known as cream of tartar. In its pure form this salt is perfectly white ; but in its crude state it commonly has a brownish red appear- ance from the coloring matter of the grapes. In this crude state it is called argol. Sodium potassium tartrate C 4 H 4 O 6 NaK is called Rochelle salt. It is used as a mild purgative. The acid sodium salt and the normal sodium and potassium salts of tartaric acid are readily soluble in water. They are all optic- ally active. The common variety of tartaric acid is dextrc- HYDROCARBONS AND THEIR DERIVATIVES 263 FIG. 94. active. It forms beautiful monoclinic prisms. The dextro and laevo varieties crystallize in forms that are to each other as an object is to its im- age in the mirror; that is, they are enantiomor phous forms. The fact that dextro and Ise vo tartaric acids also exhibit dex- tro and leevo char- acter in their crystal forms was discovered by Louis Pasteur. It was his first notable scientific discovery. Figure 94 shows crystals of dextro and Isevo tartaric acid. In lemons and other citrus fruits citric acid abounds. This is a strong tribasic hydroxyacid whose composition is repre- sented by the following formula : H I H - C - COOH I HO- C- COOH. I H - - COOH I H It crystallizes with one molecule of water in beautiful rhombic prisms. These melt at 100 and lose their water at 130. This acid forms three series of salts, for there are three re- placeable hydrogen atoms in the molecule. Thus we have: C 6 H 6 O 7 K 3 , the normal or tertiary potassium citrate; C 6 H 6 O 7 K 2 , secondary potassium citrate ; and C 6 H 7 O 7 K, primary potassium citrate. Malic, tartaric, and citric acids were discovered by the great Scheele. He prepared them by treating fruit juices with lime, thus obtaining the calcium salts, and then decom- posing these with sulphuric acid, forming calcium sulphate, which is difficultly soluble, and the free acids. The latter remained in the solutions, from which they were obtained by evaporation. 264 OUTLINES OF CHEMISTRY Esters. Esters, or ethereal salts, are formed by replacing the hydrogen of an acid by an alkyl radical. Just as metallic hy- droxides react with acids forming salts and water, so alcohols, which are hydroxides of alkyl radicals, react with acids to form esters and water: NaOH + HCOOH = HCOONa + H 2 O. CH 3 OH + HCOOH = HCOOCHg + H 2 O. The compound HCOONa is sodium formate, whereas HCOOCH 3 is methyl formate. The latter is a typical ester. In making esters, it is common practice to add a dehydrating agent to take up the water formed; usually hydrochloric acid gas is em- ployed, for it has great affinity for water, and can readily be removed afterwards because of its volatility. Ethereal salts of either inorganic or organic acids may be formed. Thus, we have such compounds as methyl iodide CH 3 T, ethyl nitrate C 2 H 5 NO 3 , ethyl nitrite C 2 H 5 NO 2 , methyl hydrogen sulphate CH 3 HSO 4 , amyl acetate CH 3 COOC 5 H n , propyl tartrate C 4 H 4 O 6 (C 3 H 7 ) 2 , etc. Many of the simpler esters are mobile liquids of pleasant aromatic odor, which can readily be dis- tilled and purified. The general formula of an ester of an organic acid is RCOOR/, where R and R' represent alkyl radicals that may or may not be alike. Esters occur in flowers, fruits, and other parts of plants, and impart to these their characteristic perfumes or flavors. So amyl acetate, or banana oil, smells like bananas; methyl butyrate C 3 H 7 COOCH 3 is known as pineapple oil; methyl salicylate C 6 H 4 (OH)COOCH 3 is tha oil of winter- green {Gaultheria procumbens). Again, many of our well-known fats and oils are esters. So olive oil and cotton-seed oil are essen- tially trioleine, an ethereal salt in which the oleic acid radical is united with the glycerine radical which acts as base, thus : (C 17 H 33 COO) 3 'C 3 H 5 . Similarly in beef tallow, tripalmitine (C 15 H 31 COO) 3 -C 3 H 5 and tristearine (C 17 H 35 COO) 3 - C 3 H 5 are found together with some trioleine. In fact, these three esters in which glycerine is the base make up the fats. Trioleine is a liquid at ordinary temperatures, whereas tristearine and tripalmitine are solids. In mutton tallow, which is hard, tristearine predomi- nates, whereas in beef tallow there is more tripalmitine and tri- oleine. In lard, trioleine is still more abundant, which is indicated by the soft, pasty consistency of the material. HYDROCARBONS AND THEIR DERIVATIVES 265 On treating esters with caustic alkalies, decomposition oc- curs, thus : CH 3 COOC 2 H 5 + NaOH = CH 3 COONa + C 2 H 5 OH. A similar action occurs when esters are boiled simply with water : CH 3 COOC 2 H 5 + H 2 O 5 CH 3 COOH + C 2 H 6 OH. In this case we. have a typical instance of hydrolysis. The action is incomplete, being reversible. When fats are boiled with caustic soda, glycerine and the sodium salts of stearic, palmitic, and oleic acids result. The latter are soaps. The reactions may be expressed thus : (C 15 H 31 COO) 3 C 3 H 5 +3NaOH = 3C 16 H 31 COONa+C 3 H 6 (OH) 3 . (C 17 H 36 COO) 8 C 3 H s +3NaOH = 3C 17 H 36 COONa + C 3 H 6 (OH> 8 . (C 17 H 83 COO) 3 C 3 H 6 + 3 NaOH = 3 C 17 H 33 COONa+ C 3 H 3 (O H) 3 . In each case the sodium salts are solids soluble in water. By addition of common salt brine these soaps are precipitated or "salted out," while the glycerine remains dissolved. The potas- sium salts of mixtures of these higher fatty acids are soft soap, and the sodium salts are hard soap. The real nature of fats became known through the investigations of the French chem- ist Chevreul. The process of decomposing any ester by means of an alkali is called saponification. This process is essentially the same in nature, no matter which ester is decomposed. When a soap solution is mixed with hard water, a curdy white precipitate is formed ; this is the calcium soap or calcium salt of the fatty acid. Writing the reaction for sodium oleate, Castile soap, and calcium sulphate solution, we have : 2 C 17 H 33 COONa + CaSO 4 = Na 2 SO 4 + (C 17 H 33 COO) 2 Ca. The latter compound is the insoluble calcium soap. The cleansing power of soap depends upon the fact that soap is soluble in water due to its content of sodium or potassium, and that the soap also coalesces with fats because of the large fatty radical it contains. Thus soap loosens the greasy mate- rial from the skin or fabrics, forming an emulsion which may be washed away. 266 OUTLINES OF CHEMISTRY Butter fat consists of about 92 per cent of a mixture of trl oleine, tripalmitine, and tristearine, and about 7.7 per cent of tributyrine, together with smaller amounts of other glycerides that give butter its characteristic flavor. Butter usually con- tains from 12 to 15 per cent water, also minor amounts of salt, casein, and milk sugar. The butter fat in butter amounts to from 82 to 85 per cent. Nitroglycerine C 3 H 5 (NO) 3 is an ester made by adding glyc- erine to a well-cooled mixture of nitric and sulphuric acid. The latter serves as a dehydrating agent. Nitroglycerine is a colorless, odorless, viscous liquid which explodes violently when heated rapidly to 200, or when jarred mechanically. Mixed with infusorial earth, nitroglycerine forms a mass called dynamite, which may be transported with far less danger. On treating nitroglycerine with caustic soda, sodium nitrate and glycerine are formed. Ethers. Just as the alcohols are the hydroxides, so the ethers are the oxides of alkyl radicals. The relations between water, alcohol, and ether on the one hand and metallic hydroxides and oxides on the other hand are illustrated by the following formulae : H 2 O water, NaOH sodium hydroxide, Na 2 O sodium oxide, (C 2 H 5 )OH alcohol, (C 2 H 5 ) 2 O ether. Ethers are readily formed by treating alcohols with a dehydrating agent like sulphuric acid. Ordinary ether is ethyl oxide. It is made by carefully running alcohol into concentrated sulphuric acid at 140-145. The reaction is 2 C 2 H 6 OH = (C 2 H 5 ) 2 O + H 2 O, the water being taken up by the sulphuric acid. For the reason that sulphuric acid is used in the manufacture of ether, the latter is often termed sulphuric ether. It, however, does not contain sulphuric acid, for the latter remains behind as the ether distills off. Ethyl ether is a mobile liquid of agreeable odor. It boils at 35 and its specific gravity at is 0.736. It is very inflamma- ble. It is used as an anaesthetic and also as a solvent for fats, oils, and kindred substances. By using other alcohols, ethers of various composition may be prepared. Sulphur ethers are compounds in which the oxygen of ordi- nary ethers is replaced by sulphur. Sulphur ethers are vola- tile, inflammable liquids with an extremely nauseating odor. HYDROCAHBONS AND THEIR DERIVATIVES 267 They may be regarded as derived from H 2 S by replacing the hydrogen atoms by alkyl groups; thus, we have methyl sul- phide (CH 3 ) 2 S, ethyl sulphide (C 2 H 5 ) 2 S, etc. Ketones. Acetone is the simplest representative of a class of organic compounds called ketones. It may be prepared by the dry distillation of sodium acetate : 2 CH 3 COONa = Na 2 CO 3 + (CH 3 ) 2 CO. acetone By heating the sodium salts of other acids, other ketones may similarly be formed. The general formula of a ketone is |>c=o. Acetone is a colorless, mobile liquid of agreeable odor. Its specific gravity is 0.792 at 20. It boils at 56.5 and is misci- ble in all proportions with water or alcohol. Both ketones and aldehydes contain the carbonyl group C = O. In ketones two radicals are combined with the CO group, whereas in alde- hydes one hydrogen atom and one radical are combined with this group. Acetone is important as a solvent. It commonly occurs in wood alcohol, being one of the products of the destructive distillation of wood. Acetone is also used in making iodoform. Carbohydrates. The carbohydrates are compounds that con- sist of carbon, hydrogen, and oxygen, the last two elements being present in the same proportions as in water, whence the name carbohydrates. This is one of the most important groups of compounds, for it contains the substances that are found in greatest abundance in the vegetable world, namely, (1) the sugars, (2) the starches, (3) the celluloses, (4) dextrine and gums. The most important sugars may be divided into two groups, the monoses, having the empirical formula C 6 H 12 O 6 , and the bioses, having the formula C 12 H 22 O n . The monoses are not altered by dilute acids, whereas the bioses are converted into monoses by the action of dilute acids. Among the important monoses are (1) glucose C 6 H 12 O 6 or CH 2 OH (CH OH) 4 CHO, also known as grape sugar or dextrose; and (2) Icevulose 6 H 12 O 6 or CH 2 (OH).(CHOH) 3 CO.CH 2 .OH, also called fructose or fruit sugar. From the formulae, it is evident that these sugars are polyalcohols and that in addition glucose is an aldehyde and Isev.ulose a ketone. 268 OUTLINES OF CHEMISTRY Glucose or dextrose occurs in the juice of grapes (whence its name, grape sugar) and in many other sweet fruits. It is also found in honey, in the blood, and in the urine of diabetic patients. Solutions of dextrose turn the plane of polarized light to the right, whence the name dextrose. A Isevo variety has been made in the laboratory. Dextrose may be prepared from cane sugar by the action of dilute acids. The process, which is called inversion of cane sugar, may be represented thus : CI H O U + H 2 = 6 H 12 6 + C 6 H 12 6 . cane sugar dextrose laevulose On evaporation, Isevulose remains in solution, while dextrose separates out. Dextrose crystals obtained from aqueous solu- tions have the composition C 6 H 12 O 6 + H 2 O. They melt at 86, whereas the anhydrous substance melts at 146. The action of yeast readily converts dextrose to alcohol and carbon dioxide, thus : C 6 H 12 6 = 2 C 2 H 6 OH + 2 CO 2 . Solutions of glucose readily reduce hot, alkaline solutions of cop- per sulphate and Rochelle salt, known as Fehling's solution, caus- ing a red precipitate of cuprous oxide Cu 2 O to form. This test is much used in practice, particularly in testing the urine of diabetics. Glucose is only about three fifths as sweet as cane sugar. Glucose is manufactured in large quantities from starch by boiling the latter with dilute sulphuric acid, the reaction being C 6 H 10 5 + H 2 = C 6 H 12 6 . starch glucose The acid is finally removed by treating with calcium carbonate and filtering off the calcium sulphate formed. In the United States large quantities of glucose are thus made annually from starch obtained from corn. Laevulose, or fructose, commonly occurs with glucose in fruits and in honey. As already stated, it is formed together with glucose, by inverting cane sugar. Crystals of Isevulose may be obtained from alcoholic solutions. They melt at 95. Solu- tions of Isevulose turn the plane of polarized light to the left, whence the name laevulose. Pure Isevulose may be obtained HYDROCARBONS AND THEIR DERIVATIVES 269 from inulin (a starch that occurs in dahlia roots and many other plants) by boiling with dilute sulphuric acid, thus : C 6 H 10 5 +H 2 = C 6 H 12 6 . inuline lisvulose Fructose also reduces Fehling's solution, and by means of yeast it may be converted into alcohol and carbon dioxide. Among the most important bioses are (1) sucrose, or cane sugar, (2) maltose, or malt sugar, (3) lactose, or milk sugar Each of these has the empirical formula C 12 H 22 O n . Cane sugar, also called saccharose or sucrose, C^H^Ojj, is very widely distributed in the vegetable kingdom. Sugar cane contains from 15 to 20 per cent sucrose, and sugar beets in some cases contain up to 20 per cent. It is further found in the juice of the sugar maple, in sorghum, in the flowers of plants, and in many varieties of nuts. Cane sugar crystallizes in beautiful monoclinic prisms (rock candy). It is very solu- ble in water ; 1 part dissolves in one third its weight of water at room temperatures. At 160 sucrose melts, and on cooling it solidifies to an amorphous glassy mass called barley sugar, which on long standing again becomes crystalline. On heat- ing cane sugar somewhat over 200, water is given off and a brown substance called caramel is formed. As already stated under dextrose, sucrose yields dextrose and Isevulose on hy- drolysis, the biose being thus split into two monoses. A. cane sugar solution does not reduce Fehling^s solution. It does not readily ferment with yeast. The latter, however, contains an enzyme called invertase, which inverts cane sugar, and the fructose and glucose thus formed are then converted into alcohol and carbon dioxide by fermentation. Nearly ten million tons of sugar are produced annually from sugar cane and sugar beets. The process of preparing sugar commercially consists of expressing the juice from the cane or beets and boiling this with about a 1 per cent solution of milk of lime to neutralize acids, coagulate vegetable albuminous matter, and prevent fermentation. The solution so obtained is then treated with carbon dioxide to remove the excess of lime, and filtered through bone black to decolorize it. It is then concentrated by evaporation in a partial vacuum in so- called vacuum pans, which are heated with steam. The crys- tals which are deposited on cooling are freed from the adhering 270 OUTLINES OF CHEMISTRY brown mother liquor by centrifugal force in the " centrifugals." The latter consist of sieves rapidly rotated by machinery. The mother liquor is hurled through the meshes of these sieves, and the crystals remain behind almost dry. The drying is com- pleted with the aid of heat. The mother liquor, from which crystals will no longer form because of the presence of various impurities, is called molasses. The residue of the cane or beets from which the juice has been extracted is called the begasse. It is made into paper or used as fuel. With lime, cane sugar forms calcium sucrate C 12 H 22 O n -CaO, which is soluble in water. The analogous strontium sucrate C 12 H 22 O n ' SrO is used as a means of recovering further crystal- lizable sugar from molasses. Cane sugar solutions rotate the plane of polarized light to the right. The specific rotatory power of cane sugar for sodium light is 4- 66.5, while that of dextrose is -f 52.7, and of Igevu- lose 93. Consequently, solutions of invert sugar, which contain equal parts of dextrose and Isevulose, always exhibit Isevorotatory power. Maltose. C 12 H 22 O n + H 2 O crystallizes with one molecule of water, as its formula indicates. It may be formed by the action of diastase, a ferment contained in malt, upon starch. Maltose is of importance as an intermediate product in the manufacture of alcohol from starch. Maltose is strongly dextrorotatory, its specific rotatory power being + 137. On boiling it with dilute mineral acids, dextrose only is produced, showing that maltose is an anhydride of dextrose. Unlike cane sugar, maltose reduces Fehling's solution, and readily ferments with yeast. Lactose C 12 H 22 O n + H 2 O, milk sugar, is found in milk. Like maltose, its crystals contain one molecule of water. It is not as sweet as cane sugar, and is much less soluble in water; 1 part dissolves in 6 parts of water. Lactose is dextrorotatory, its specific rotation being 4- 52.5. On boiling with dilute min- eral acids, it yields glucose and another monose called galactose C 6 H 12 O 6 , thus : - C 12 H 22 O n + H 2 = C 6 H ]2 6 + C 6 H 12 6 . lactose glucose galactose Lactose reduces Fettling" s solution. But the reduction is less rapid than when glucose is used. Pure yeast does not produce fermen- tation of milk sugar ; but ordinary yeast acts upon milk sugar HYDROCARBONS AND THEIR DERIVATIVES 271 as it does upon cane sugar. The products formed are alcohol and lactic acid. Cow's milk contains nearly 5 per cent milk sugar. Fermentation and Enzymes. Under fermentation are classed a large number of chemical processes that are produced directly or indirectly by organisms commonly termed ferments. These organisms are yeasts, molds, or fungi, and bacteria. They se- crete complex compounds called enzymes, the chemical nature of which is not yet understood. They seem to be closely related to the albumins and peptones (which see). In the presence of these enzymes, which are also called unorganized ferments, fer- mentation takes place, each enzyme causing its own particular chemical change, which commonly progresses at room tempera- tures, being accomplished with evolution of heat and frequently with liberation of gas. As a rule the action is checked by either raising or lowering the temperature materially, also by the addition of various poisons. The action of the enzymes is often termed catalytic, for a small amount of enzyme may produce a large amount of change without being itself seemingly affected. Among the common enzymes are : zymase, contained in yeast cells, producing alcoholic fermentation ; malt diastase, contained in malt, converting starch into malt sugar ; invertase, contained in ordinary yeast, inverting cane sugar and lactose; emulsin, contained in bitter almonds ; pepsin, contained in the stomach juices, aiding in the digestion of albumins ; trypsin, contained in the intestinal juices, aiding in so-called tryptic digestion. Starch and Dextrine. Starch (CgH^Og)^ is a carbohydrate found in large quantities in all plants, particularly in tubers like potatoes, and in grains like corn, rye, oats, wheat, rice, etc. Starch is insoluble in water and occurs as a nodular deposit of varying sizes and forms in different plant cells. Figure 95 shows grains of potato starch 'as they appear under the microscope. Figure 96 shows grains of wheat starch, and Figure 97 shows grains of corn starch. Figure 98 shows potato starch grains as they look in polarized light. Starch is prepared mainly from potatoes and corn, the former being commonly used in Europe and the latter in America. The potatoes or grains are ground so as to break up the plant cells and lay the starch granules bare. These are then washed out with water, forming a thin milky paste which is passed through fine sieves that retain the 272 OUTLINES OF CHEMISTRY pulp. The starch is allowed to settle, the water is drained off, and the remaining material is dried at low temperatures. Though starch is insoluble in cold water, its granules swell up and burst when treated with boiling water, and the contents of FIG. 95. FIG. 96. the cells, the granulose, dissolves, whereas the cell walls remain undissolved and can be filtered off. From the nitrate, the gran- ulose or soluble starch may be precipitated by adding alcohol. Starch paste is prepared by grinding starch with a little cold FIG. 97. FIG. 98. water and then adding boiling water with constant stirring. Thus a semi-transparent, gelatinous mass results, which is used as an adhesive. It is also employed in the laundry for stiffen- ing clothes, the hot sad iron converting the starch into dextrine, HYDROCARBONS AND THEIR DERIVATIVES 273 which covers the linen, imparting to it a characteristic luster. Starch has the empirical formula C 6 H 10 O 5 . Its molecular weight is unknown, but it is probably some rather high multi- ple of the formula given. Starch is an exceedingly important article of food. Wheat flour, for instance, consists of about 70 per cent starch, 10 per cent gluten, a sticky nitrogenous sub- stance closely allied to albumins like white of egg, and minor amounts of water, sugar, and inorganic material. Gluten, like starch, is valuable as food material. On heating starch to about 210 it is converted into dextrine (C 6 H 10 O 6 ), which is also obtained as an intermediate product in the conversion of starch to glucose by means of dilute sul- phuric acid. Dextrine is a colorless, amorphous substance that has a strong adhesive property. It is consequently frequently used as a cheap gum. Cellulose. This carbohydrate has the same empirical formula as starch (Cgll^Og)^. Its molecular weight is unknown, but it is probably rather high. Cellulose is very widely distributed in nature, constituting the material out of which the cell walls of plants are made. Thus wood, cotton, linen, hemp, flax, etc., when freed from the mineral matter they contain, are almost pure cellulose. Filter paper after extraction with hydrochloric and hydrofluoric acids and washing with water is a nearly pure form of cellulose. Cellulose is insoluble in water but soluble in a solution of copper in strong ammonia water, which is known as Schweitzer's reagent. From this solution cellulose may be pre- cipitated by means of acids in the form of a gelatinous mass that may be filtered off and dried to an amorphous powder. Cellu- lose gradually dissolves in concentrated sulphuric acid. On diluting the solution and boiling, dextrine is formed, which is finally converted into glucose. Thus wood may be changed to glucose, from which in turn alcohol may be obtained by fermen- tation. Paper consists essentially of thin sheets of fibers of cellulose matted together. On heating cotton with a mixture of nitric and sulphuric acids, nitrates of cellulose are formed. This action is similar to the formation of glycerine nitrate, i.e. nitroglycerine- The composition of cellulose nitrates varies according to the length of time that the nitric acid has acted upon the cellu- lose, and the concentration of the acid. Cellulose hexanitrate 274 OUTLINES OF CHEMISTRY 6 O 4 i y called nitrocellulose, or gun cotton. It does not dissolve in a mixture of alcohol and ether. Gun cotton looks like ordinary cotton, but it is not as soft to the touch. It burns rapidly and quietly, producing no smoke. By means of fulminating mercury it can be made to explode violently. It is used as an explosive, frequently together with nitroglycerine. Gun cotton and nitroglycerine are made into threads with the aid of acetone and vaseline and are thus used as smokeless gun- powder. Celluloid consists of camphor and gun cotton. It is not explosive, but it burns readily. The tetra- and petranitrates of cellulose C 12 H 16 (NO 3 ) 4 O 6 and C 12 H 15 (NO 3 ) 5 O 5 readily dis- solve in a mixture of alcohol and ether. This solution is known as collodion. It is used in photography for making films, and in surgery for protecting wounds from the air, for when collodion solution is poured out in thin layers, the alcohol and ether evaporate, leaving a tough, amorphous, transparent layer of nitrocellulose. On treating nitrocellulose with caustic soda, cellulose and sodium nitrate are formed, showing that nitrocellu- lose is a nitrate of cellulose. Nitrobenzene C 6 H 5 NO 2 is a light yellow oil formed by the action of nitric acid on benzene. Its boiling point is 208. It is used as a perfume for laundry soaps, under the name oil of mirbane. By reducing nitrobenzene with nascent hydrogen, aniline C 6 H 5 NH 2 is formed. Aniline is a base which forms salts like ammonia. Thus, C 6 H 5 -NH 2 -HC1 is aniline hydro- chloride. Aniline is a nearly colorless liquid which boils at 189. On exposure to the air, it soon turns brown. Aniline may be regarded as ammonia in which one hydrogen atom has been replaced by the phenyl group, C 6 H 5 . Many similar sub- stituted ammonias are known; they are called amines. So we have methyl amine CH 3 -NH 2 , ethyl amine C 2 H 5 -NH 2 , propyl amine C 3 H 7 -NH 2 , etc. In general, their chemical behavior is similar to that of ammonia. The aniline dyes are deriva- tives of pararosaniline HO C(C 6 H 4 NH 2 ) 3 and rosaniline HO C(C 6 H 4 NH 2 ) 2 - C 6 H 3 CH 3 - NH 2 . The hydrochloride of the latter is the dyestuff magenta. Many aniline dyes are known. They represent very many beautiful colors. The ani- line dyestuff industry is an important one. It will be noted that the dyes are all derived from benzene, which is obtained from coal tar, whence the name coal-tar dye. HYDROCARBONS AND THEIR DERIVATIVES 275 Alkaloids are complex basic substances that occur in plants. These compounds are so named because they form salts with acids, thus playing the role of alkalies. The base pyridine C 5 H 5 N is a colorless, odoriferous liquid boiling at 115, which is found in coal tar and in the products of the dry distilla- tion of green bones. Quinoline C 9 H 7 N occurs with pyridine. Many of the plant alkaloids are related to these two bases. Nicotine C 10 H 14 N 2 is found in tobacco. It is a poisonous liquid, boiling at 247. Atropine C 17 H 23 O 3 N occurs in nightshade, Atropa belladonna. It forms very poisonous crystals which melt ut 115.5. It is used by oculists to cause expansion of the pupil of the eye. Cocaine C 17 H 21 O 4 N is found in coca leaves. It consists of crystals that melt at 98. It is used as a local anaesthetic. Quinine C 20 H 24 O 2 N 2 and cinchonine C 19 H 22 O 2 N 2 are found in Peruvian bark, which also contains other alkaloids. Quinine is usually administered in form of the sulphate. It is a specific for malarial fever. Morphine C 17 H 19 O 3 N, narco- tine C 22 H 23 O 7 N, and codeine C 18 H 21 NO 3 occur in opium, which is the dried sap of the unripe seed pods of the opium poppy. Morphine is a crystalline powder which was first isolated in 1806. It induces sleep. As a rule it is administered in form of the hydrochloride. Strychnine C 21 H 22 O 2 N 2 and brucine C 23 H 26 N 2 O 4 occur in nux vomica. These alkaloids are very poison- ous, causing death accompanied by muscular contraction and rigor. Strychnine forms crystals that melt at 265. They are nearly insoluble in water. Strychnine compounds have a very bitter taste. In minute quantities they are often prescribed in medi- cine as a tonic. Proteins. These are very important compounds, consisting of carbon, hydrogen, nitrogen, oxygen, and sulphur, which make up a large share of the bodies of animals and also play an important role in plants. Exclusive of water, fats, and mineral constitu- ents, the animal matter consists of proteins, formerly called proteids. Without proteins as food, animals will finally die. The chemical structure of the proteins is very complicated, On the average, proteins have about the following com- position: . 276 OUTLINES OF CHEMISTRY Carbon 50.5 to 54.5 per cent Hydrogen 6.5 to 7.3 per cent Nitrogen 15.0 to 17.7 per cent Oxygen 21.0 to 24.0 per cent Sulphur - 0.3 to 2.3 per cent Phosphorus 0.4 to 0.8 per cent Nucleoproteins may contain from 5 to 6 per cent phosphorus. Among some of the common protein substances may be mentioned the albumins, as they occur for instance in serum, in eggs, in milk, in muscular tissues, and in leguminous and other seeds. Albumins are coagulated by heat and by acids. Albuminoids are closely related to albuminous bodies. Among the albuminoids are gelatine, elastine, and keratine. Elastine enters into the composition of connective tissues, while kera- tine is the main substance found in hoofs, horns, feathers, hair, nails, and the epidermis. By the action of pepsin albumins are converted into peptones, which process involves an addition of the elements of water to the molecule. The molecular weight of the protein molecule as determined by the freezing point method (which see) is approximately 15,000. Proteins are consequently regarded as colloidal sub- stances. The fact that they diffuse very slowly in solutions also accords with this view. The idea that proteins are very complicated bodies comes from the fact that a large number of products may be obtained from them by treatment with various reagents like acids, alkalies, oxidizing agents, etc. The study of proteins has of recent years been prosecuted with special success by Professor Emil Fischer of the University of Berlin, to whom we also owe much of our knowledge of the constitu- tion of sugars. When protein substances, like meat, fish, etc., putrefy, pto- maines are often formed. These are basic substances which act like the alkaloids in many respects. They are poisonous. Among them are putrescine C 4 H 8 (NH 2 ) 2 and cadaverine C 6 H 10 (NH 2 ) 2 . Illness from eating spoiled meat is commonly due to ptomaine poisoning. HYDROCARBONS AND THEIR DERIVATIVES 277 REVIEW QUESTIONS 1. What is a hydrocarbon? Give five illustrations. What are the ultimate products of combustion of a hydrocarbon? Compare the chemical properties of hydrocarbons with : (a) the hydrohalogens, (6) ammonia, hydrazine, and hydrazoic acid, (c) hydrogen sulphide. 2. Give the methods for preparing two typical hydrocarbons, writing the equations. 3. Mention the more important products that are made from petro- leum. Where is petroleum found and how is it refined ? 4. What is meant by the term " halogen substitution product"? Illustrate by giving three examples of halogen substitution products and the equations showing their formation. 5. What are the formulas and uses of the following : chloroform, iodoform, carbon tetrachloride, ether, carbolic acid? 6. What is an alcohol? Give the names and formulas of five alco- hols. How is ordinary alcohol made and how may it be converted into vinegar ? 7. What is denatured alcohol? About what is the alcoholic content of beer, wine, whisky? 8. What is the chemical nature of glycerine? Mention five other substances that belong in the same general class of compounds as glycer- ine. Give the formulas of each of these and state the characteristic prop- erties of this class of compounds. 9. What is hydroquinone ? Pyrogallic acid? Creosote? What are these used for? How are they related to: (a) alcohol, (6) phenol, (c) glycerine? 10. What is an aldehyde? Give two illustrations and the equations showing their formation ? 11. What is formaline ? What use is made of it ? 12. What is an organic acid? Write the formulas of the following organic acids : formic, acetic, propionic, butyric, lactic, oxalic, benzoic, stearic, oleic, tartaric, malic, citric. State where each of these occurs, or how it may be prepared. Mention also the chief properties of each of these acids. 13. What is meant by optical activity? Give six examples of opti- cally active substances and state what they have in common that causes le optical activity. 14. To what class of substances do fats belong? How may a fat converted into a soap ? What other product is formed simultaneously ? r rite the equations showing the change of olive oil into castile soap. 15. Why is soft water superior to hard water for washing purposes? rive a chemical equation showing how hard water acts on soap. Dis- the cleansing power of soap. 16. What is oil of wintergreen, banana oil, pineapple oil? To what ineral class of substances do these belong? 278 OUTLINES OF CHEMISTRY 17. What is nitroglycerine, and how does it differ from dynamite? 18. Write the general formula for : (a) an alcohol, (6) an aldehyde, (c) an organic acid, (d) a ketone, (e) an ester, (/) an ether, (g) a sulphur ether, and give a concrete example of each. 19. What is a carbohydrate? Classify the important carbohydrates. 20. What is a sugar? Classify the important sugars. State how the sugars you have mentioned react with Fehling's solution. 21. What is fermentation? What is an enzyme? Give an illustra- tion of each. 22. What is each of the following: smokeless powder, celluloid, collodion, nitrobenzene, aniline, quinine, morphine? 23. What is meant by the term "protein"? Give an illustration. What is the cause of so-called ptomaine poisoning ? 24. How much carbon dioxide and how much water may be obtained from the combustion of a pound of tartaric acid? CHAPTER XVI ILLUMINATING GAS AND FLAMES Illuminating Gas. Illuminating gas is often produced by the destructive distillation of soft coal. During this process there are formed coal gas, coal tar, and water containing ammonia and other products in solution. Coal gas consists mainly of hydrogen, carbon monoxide, and methane, together with a small amount of higher hydrocarbons known as illumi- nants. Impurities like nitrogen, carbon dioxide, and hydrogen sulphide are also present in small quantities. The tar from the gas works is distilled, and the benzene obtained therefrom is used in manufacturing dyestuffs and many medicinal prep- arations. The remaining tar is then employed for roofing, making artificial asphalt, black varnishes, etc. The coke from the gas retorts serves as fuel. The ammoniacal gas liquors serve as a source of ammonium salts. These liquors are neu- tralized with sulphuric acid, and the ammonium sulphate so obtained is used as a fertilizer or converted into ammonia water by treatment with lime and absorption of the gas lib- erated. The coal gas is treated with lime to remove carbon dioxide and sulphides, particularly hydrogen sulphide, which is always present. In the process of manufacture, coal gas passes from the retort R (Fig. 99), through the condensers (7, in which the tar and ammoniacal liquors are condensed, into the scrubber S, where it is washed with water. The gas then goes through layers of lime or oxides of iron, or both, which are contained in the purifiers P. This removes the hydrogen sulphide. Finally the gas enters the holder H, from which it is distrib- uted through the mains. Not all illuminating gas is coal gas. Much illuminating gas is made in America from petroleum. By allowing the latter to come into contact with the walls of a chamber lined with fire bricks heated to about 1000 C. the heavy hydrocarbons are 279 280 OUTLINES OF CHEMISTRY ILLUMINATING GAS AND FLAMES 281 decomposed into simple hydrocarbons that contain fewer car- bon atoms in the molecule, and are consequently gaseous at ordinary temperatures. This process is called " cracking " the petroleum oils. Pintsch gas is made entirely by this process. It has very high illuminating power and is used for lighting railway cars, for which purpose it is compressed, with moderate pressure, in steel cylinders. A special burner must be used with this gas to prevent smoking of the flame. Oil gas thus made is rich in benzene and the hydrocarbons of the ethylene and acetylene series. These give it its high illuminating power, and are consequently called illuminants. It is now common practice to increase the illuminating power of coal gas by " enriching " the gas by means of petroleum oils added directly to the coal in the retorts, or by adding oil gas to the coal gas. It will be recalled that water gas consists of hydrogen and carbon monoxide, and consequently burns with a flame that is nearly non-luminous. To make it serve for illuminating pur- poses, it is enriched, or "carbureted," by means of oil gas. To accomplish this the water gas is passed through heated chambers into which petroleum oil is run and cracked as above stated, thus furnishing the necessary illuminants. Sometimes the oil gas is made separately and then added to the water gas in proper proportion. Enriched water gas is much used in the large cities of the United States. Coal gas was first used for house illumination in 1792, by William Murdock in London, where it was employed for street lighting in 1812. Three years later it was also used for this purpose in Paris. However, the fact that a combustible gas is evolved from coal when it is heated, was discovered as early as the year 1680 by Becher. The illuminating power of a gas is usually expressed in candle power. The gas is burned from a burner consuming 5 cu. ft. per hour, and this light is compared with that of a standard candle by means of a photometer. The composition of gases used for illuminating purposes varies considerably. A general idea of the composition of illuminating gas may, however, be obtained from the fol- lowing table, from the reports of the Massachusetts State Gas Inspectors, which presents average results in per cent 'oy volume : 282 OUTLINES OF CHEMISTRY COAL GAS CARBURETTED WATER GAS OIL GAS Candle power .... 17.5 25. 65. Illuminants 5.0 16.6 45.0 34.5 19.8 38.8 Hydrogen 49.0 32.1 14.6 l Carbon monoxide . . . Nitrogen 7.2 3.2 26.1 2.4 1.1 Oxygen Carbon dioxide .... 1.1 3.0 Flame. Flames are produced by burning gases. Whenever a continuous stream of one gas issues into an atmosphere of another gas upon which it acts chemically, producing a suffi- cient rise of temperature and a certain degree of luminosity, we have at the surface of contact of the gases the phenomenon of flame. Solids like charcoal and coke burn practically with- out flame. In the kerosene lamp the wick draws up the oil and the latter is converted into gas by heat, and thus it is really the gas that burns in the flame. Similarly in the flame of a candle it is gas that burns ; for the wax or tallow is melted by the heat, and the liquid is drawn up by the wick, being converted to gas before it is burned. This is well illustrated by the experiment represented in Fig. 100, in which some of the unconsumed gas from the inner part of the flame is conducted off by means of the glass tube, at the end of which it has been lighted. Oxygen is said to support the combustion of a gas that burns in the air. Obviously, however, combustion is really the mutual interaction of the two gases. So, for instance, coal gas can be burned in the form of a flame in the air, or air can be burned in the form of a flame in an atmosphere of coal gas. This is well illustrated by the experiment shown in Figs. 101, 102, and 103. 1 This figure represents ethane and not hydrogen. The latter is not present in oil gas. FIG. 100. ILLUMINATING GAS AND FLAMES 283 Illuminating gas is passed into the apparatus (Fig. 101), through the tube at the left, and then the gas is ignited at the end of the other tube at the right, as shown. Now the lid on top of the apparatus is opened (Fig. 102), whereupon the draft created causes the flame to strike inward and burn inside of the large tube, as shown in the figure. This is now a jet of air burning in an atmosphere of illuminating gas. The gas issuing from the opening created by removing the lid may be lit as shown, and thus we have the upper flame of illuminating gas burning FIG. 101. FIG. 102. FIG. 103. in the air, and at the same time the inner flame of air burning in the atmosphere of the gas. This is commonly spoken of as the reverse flame. That it is actually air that is burning in the illuminating gas may be demonstrated by passing a small gas flame up into the inside of the flame of air burning in the illuminating gas, as shown in Fig. 103 ; the small gas jet burns in the inner cone of the flame. It will be recalled that a jet of hydrogen will burn in oxygen and a jet of oxygen in an atmosphere of hydrogen ; and that a jet of chlorine will burn in hydrogen and a jet of hydrogen in an atmosphere of chlorine. These experiments also demon- strate further the relative nature of combustion. The latter experiment further shows that we may have flames due to chemical changes other than oxidation. 284 OUTLINES OF CHEMISTRY FIG. 104. The fact that ox}^gen will burn in coal gas is readily demon- strated by the experiment illustrated in Fig. 104. Chlorate of potassium heated in a deflagrating spoon till oxygen is evolved freely is introduced into the atmosphere of coal gas as shown. The oxygen burns brilliantly, the flame being colored purple by the presence of potassium. The gas should be ignited at the upper orifice, as shown in the figure, before in- troducing the deflagrat- ing spoon. Luminosity of Flame. Flames are either lumi- nous or non-luminous. Thus, hydrogen burns with a non-luminous flame. This may, how- ever, be made luminous by blowing fine solid particles, like soot, for example, into the flame. These fine particles be- come heated to incandes- cence and emit light. Similarly, a platinum wire introduced into the hydrogen flame soon becomes hot and gives off light. Carbon monoxide burns with a blue flame, which can readily be made luminous by the introduction of fine particles of carbon. Figure 105 shows a convenient apparatus for accomplishing this. Carbon monoxide is passed through the ap- paratus, the cock at the left being open, and ignited. We now have the characteristic blue flame. If the cock at the right is then opened and the one at the left closed, thus causing the gas to pass through the cotton saturated with benzene contained in the right limb, the flame becomes luminous, for the heavy hydrocarbon vapors carried into the FIG. 105. ILLUMINATING GAS AND FLAMES 285 A B C flame lose their hydrogen at the high temperature of the flame, thus setting particles of carbon free which are heated to incan- descence and so give the flame luminosity. By proper manipu- ^tion of the two cocks, the flame may be rendered luminous or non-luminous alternately. That all luminous gas flames, in- cluding those of oil lamps and candles, contain carbon particles in suspension is readily de- monstrated by the fact that a layer of soot is deposited upon glass or porcelain held in such flames. According to Lewes, acetylene is formed in all luminous flames, and it would seem probable that the luminosity is due to the pres- ence of this gas, which breaks up, yielding carbon particles that then become heated to incandescence. At any rate, it is certain that the lumi- nosity of the flame of illumi- nating gas is due to the presence of benzene and gases of the ethylene and acetylene series, which are commonly called the illumi- nants. It is well known that at high temperatures these decompose, yielding acetylene as one of the products. When a jet of illuminating gas is mixed with air and then ignited, a non-luminous flame is obtained. This is the principle used in the Bunsen burner., Figure 106 shows a simple apparatus for illustrating this principle. Gas issues through the opening of the small tube, over which the larger tube is placed as shown. The gas passing upward creates a current of air which enters the larger tube, as shown by the arrows, and mixes with the gas in the large tube at the end of which the mixture is lighted, yielding Air FIG. 106. 286 OUTLINES OF CHEMISTRY a blue, practically non-luminous flame. It is commonly held that this non-luminosity is due to the oxygen, which causes complete combustion of the carbon particles. The Bunsen burner is used in various forms in the laboratory. In principle, gas stoves and furnaces are all Bunsen burners. By supplying compressed air by means of a bellows or other- wise, the blowpipe or blast flame is produced. This additional supply of air insures more rapid and more complete combus- tion, and consequently a higher temperature is obtained. The ordinary form of blast lamp is similar to the oxy hydrogen lamp, except that coal gas and air are used instead of hydrogen and oxygen. The flame of an alcohol lamp is non-luminous because alcohol contains some oxygen, and this, together with that supplied by the air, is sufficient to secure complete combustion of the par- ticles of carbon. For similar reasons ether burns with a prac- tically non-luminous flame. As the carbon content of com- pounds increases, however, the flames with which the compounds burn become luminous, and even sooty. It will be recalled that when calcium oxide was introduced into the oxyhydrogen flame, the lime was heated to a tempera- ture at which it emitted a brilliant white light. The brilliant light formed when magnesium or phosphorus burn in the air or in oxygen is similarly caused by the incandescence of the particles of MgO and P 2 O 5 that are formed during the com- bustion. In the Welsbach light this principle is used by hanging over the flame of a Bunsen burner a mantle consisting of a network made of 99 per cent thorium dioxide and 1 per cent cerium dioxide. Thus a strong, brilliant, white light is produced with a relatively low consumption of gas. Other oxides will also serve, but they do not yield nearly as good results in practice. The oxides mentioned, when used in the proportion indicated, have been found to give the strongest light. The Structure of Flame. Taking as a typical luminous flame that of a candle (Fig. 107), we see that it exhibits three distinct zones. The inner zone A is dark and non-luminous. It consists of the gases produced by the decomposition and volatilization of the material of the candle drawn up by the wick, which fact may be demonstrated by the experiment ILLUMINATING GAS AND FLAMES 287 shown in Fig. 100. This zone is also of relatively low tem- perature. The next zone B is brilliantly luminous. Here partial combustion of the expanding gases is going on. The ethylene and other hydrocarbon gases lose their Irydrogen, probably forming at first acetylene, which in turn loses hydro- gen and yields carbon particles whose incandescence gives the luminosity to this zone of the flame. Finally, the outer fringe C is practically non- luminous, for here more oxygen of the air comes into contact with the hot gases, thus causing more complete combustion. This latter zone is the hottest part of the flame. In the Bunsen flame (Fig. 106) the luminous zone is absent. We have here only an inner greenish blue cone O and an outer practically non- luminous mantle A. In the inner FlG - 107 - cone the essential processes are the combustion of hydrogen to water, and of carbon to carbon monoxide. The inner zone consequently contains an excess of reducing' gases and is termed the reducing flame, whereas the outer zone contains an excess of oxygen and is called the oxidizing flame. Many of the metals are oxidized when intro- duced into the oxidizing flame. Again, when oxides, like those of lead for instance, are placed in the inner zone they are re- duced. Blowpipe flames exhibit the same general structure as the Bunsen flame. That the lower part of the inner cone of the latter is relatively low in temperature is demonstrated by the fact that a match head may be placed in it for some time without taking fire. The outer fringe and cone near B are the hottest parts of the FIG. 108. the tip flame. of the 288 OUTLINES OF CHEMISTRY FIG. 109. Davy Safety Lamp. When a wire gauze is held over a Bunsen burner, and the gas is then lighted on the upper side of the gauze, the flame burns on that side and does not pass through the gauze to the lower side (Fig. 108). Again, if a wire gauze is pressed down upon a Bunsen flame, the flame does not pass through to the upper side of the netting, but only partially con- sumed gas makes its appearance there (Fig. 109). These phenomena are due to the fact that the wire netting lowers the temperature of the gases below the kindling point ; that is, the temperature at w^hich the gases take fire in the air. If the gauze should become very hot, the flame will pass through, of course. Upon the principle that a wire netting is thus able to intercept a flame as explained, Sir Humphry Davy devised the miner's safety lamp (Fig. 110). This consists of an oil lamp having a tight-fitting chimney of wire gauze. When this lamp is lighted and taken into a mine where fire damp, methane jDH 4 , is present, the flame is not communi- cated through the gauze to the explosive mixture, though to be sure the latter may get into the chimney through the gauze and burn there or cause small, harmless explo- sions. These serve to warn the miner of the presence of the dangerous gases. The safety lamp is consequently very useful ; neverthe- less, explosions do still occur in mines because currents of air arising from blasting opera- tions may blow fine coal dust into the lamp and so enable the flame to communicate itself to the fire damp on the outside of the gauze. After such explosions have occurred, the carbon dioxide (called choke damp by the miners) formed is dangerous also, because it does not support respiration and so gives rise to suffocation. FIG. 110. ILLUMINATING GAS AND FLAMES 289 REVIEW QUESTIONS 1. What is the difference in composition and method of preparation of water gas and coal gas? What are the products of the coal gas in- dustry? 2. What is a flame? Upon what does the illuminating power of a flame depend? What is meant by the term " enriching a gas " ? 3. Explain the principle of a Bunsen burner. 4. What determines the temperature of a flame? Give illustrations substantiating your answer. 5. Explain the principle of the Davy safety lamp. Why do explo- sions in mines occur in spite of this safety device? What is fire damp? Choke damp ? 6. Explain the action of the following when used to extinguish fires : (a) water, (6) carbon dioxide, (c) carbon tetrachloride, (d) sand, (e) a blanket, (/) saleratus. 7. Why should water not be used to extinguish gasoline or oil fires? What would you do to extinguish such fires? CHAPTER XYII THERMOCHEMISTRY General Remarks. All chemical changes are accompanied by either an evolution or an absorption of heat. In most of the ordinary chemical processes heat is liberated. These are consequently called exothermic changes, to distinguish them from endothermic changes, or reactions in which heat is absorbed. Endothermic changes are by no means uncommon. In fact many reactions occur with absorption of heat, particularly at higher temperatures. It must be borne in mind that physical changes as well as chemical reactions are generally accompanied by thermal changes. Thus, in melting ice or vaporizing water heat is absorbed, while in freezing water or in condensing vapor heat is liberated. Similarly, whenever a solid is con- verted into a liquid, or a gas is formed from a solid or liquid, heat is absorbed so far as the physical change is concerned; and heat is liberated when the reverse action takes place. The amount of heat required to convert 1 gram of a given solid into liquid of the same temperature is termed the latent heat of fusion. And the amount of heat necessary to change 1 grain of a liquid into vapor of the same temperature is called the latent heat of vaporization. When a piece of zinc is dissolved in hydrochloric a-cid, the solid zinc disappears and becomes part of the liquid, and simul- taneously a gas, hydrogen, is evolved. Both of these processes, the liquefaction of the solid and the liberation of the gas, con- sidered as physical processes, would proceed with absorption of heat. However, the action of hydrochloric acid on zinc pro- ceeds with disengagement of heat, which fact can readily be demonstrated by means of a thermometer placed in the acid. It is therefore evident that the thermal change observed is equal to the heat developed by the chemical interaction, minus the heat required for the liquefaction of the metal and the conversion of the hydrogen into the gaseous state. It is at 290 THERMOC H E MISTR Y 291 present impossible to determine just how much energy the last- named processes represent when they occur at room tempera- ture, and so it is also impossible to tell how much heat the actual chemical part of the change evolves. All chemical changes are similarly accompanied by physical changes of some kind. The thermal effect of the latter must be taken into consideration ; or, at any rate, if this effect cannot be evaluated and subtracted, as is frequently the case, the physi- cal state of the substances before and after the reaction must be mentioned. Calorimeters. Thermal changes are measured by means of calorimeters. A thermometer is introduced into a known weight of water contained in a cylindrical dish, the calorimeter, which is preferably made of platinum. The apparatus is so arranged that the heat evolved by the chemical reaction is communicated to the calorimeter water. Knowing the initial and final temperature of the latter, and multiplying the weight of the water by the num- ber of degrees of temperature change, the number of calories of heat evolved is obtained. A large calorie is the amount of heat necessary to raise 1000 grams of water 1 degree ; it is commonly designated by Cal. A small calorie is 0.001 of a large calorie and is desig- nated by cal. In technical work in England and America another heat unit known as the British thermal unit is frequently used. A British thermal unit is the amount of heat required to raise the temperature of 1 pound of water 1 degree Fahrenheit; it is designated by B. T. U. During calorimetric measurements care must be taken to pre- vent loss of heat by radiation, or the exact amount of heat lost 292 OUTLINES OF CHEMISTRY by radiation must be ascertained, and a proper correction made therefor in the final result. Figure 111 shows an ordinary calorimetric apparatus in cross section. The calorimeter itself should not have less than 500 cc. capacity. In Fig. 112 a com- bustion calorimeter is represented. The sub- stance to be burned is placed in the steel bomb, which is lined with platinum, gold, or porcelain enamel. The bomb is then filled with oxygen under 20 atmospheres pressure and finally immersed in the water of the calorimeter. The ig- nition is effected by means of a small wire in the bomb, heated by an electric current. Thus the combustion proceeds almost instan- taneous^, and the heat is communicated to the calorimeter water and measured in the usual way. Laws of Thermo- chemistry. Inasmuch as energy can neither be created nor de- stroyed, it is evident that if no heat be lost, the heat evolved during a chemical change is always exactly equal to the heat that is absorbed when the reaction is reversed. This law was pointed out in 1783 by Lavoisier and Laplace, who regarded it as self- evident. In 1840 G. H. Hess, professor at the University of St. Peters- FIG. 112. THERMOCHEMISTRY 298 burg, showed that the thermal change accompanying any chemical reaction depends on the initial and final condition of the substances* involved, and is independent of the intermediate changes that may occur during the reaction. Thus the total amount of heat evolved when a gram of carbon is burned to CO 2 is the same whether the combustion proceeds in one step, or whether CO is first formed, and this is then oxidized to CO 2 . This law of Hess really follows from the law of conservation of energy. It is of great importance in thermochemical measurements, for it enables many determinations to be made indirectly that could not be carried out by direct means. So it is practically im- possible to determine the amount of heat evolved when carbon is burned to CO, for some CO 2 always forms when this is at- tempted. But it is quite possible to find the heat developed when carbon is burned to CO 2 , and also that evolved when CO is burned to CO 2 ; and the difference between these two experi- mental results is the heat evolved when carbon is burned to CO. Thus : - C(solid) + O 2 (gas) = CO 2 (gas) + 97.65 Cal. CO(gas) + O(gas) = CO 2 (gas) + 68. Cal. Therefore, C(solid) + O(gas) = CO(gas) + 29.65 Cal. These are typical thermochemical equations. For instance, the first one states that when 12 grams of carbon and 32 grams of oxygen unite, 44 grams of carbon dioxide are formed and 97.65 Cal. of heat are liberated. All other thermochemical equations are interpreted similarly. According to the law of Lavoisier and Laplace, it is evident that if carbon dioxide is to be decom- posed into carbon and oxygen, energy to the amount of 97.65 Cal. is absorbed during the process per every 44 grams of CO 2 . At first it appears peculiar that the combustion of carbon to CO yields but 29.65 Cal., whereas the combustion of CO to CO 2 evolves 68 Cal. But it must be remembered that in the first step, when solid carbon passes into CO, much energy is absorbed in the process of vaporizing the carbon, which doubt- less accounts for the fact that we get but 29.65 Cal. when car- bon is burned to CO. It is evident that when furnaces are run so that fuel is but partially burned, i.e. so that a considerable proportion of the carbon is merely oxidized to monoxide, a 294 OUTLINES OF CHEMISTRY large proportion of the energy that might have been gained as heat is wasted. The development of the subject of thermochemistry is mainly due to the work of Julius Thomsen, who was professor at the University of Copenhagen, and Marcellin Berthelot, who was professor at the University of Paris. In 1853 the former stated that every simple or complex change of a purely chemical nature is accompanied by an evolution of heat ; and in 1879 Berthelot an- nounced that every change accomplished without the intervention of extraneous energy tends to produce a substance or substances in the formation of which the greatest amount of heat is disengaged. This is now commonly termed Berthelot's law of maximum work. It is true that under ordinary conditions those re- actions generally take place that evolve the greatest amount of heat; so in dissolving metals in acids, in neutralizing the latter with bases, in displacing one metal by another in solution, in the combustion of carbonaceous substances, etc., heat is evolved. Nevertheless, Berthelot's law, for which he contended strongly, does not hold rigidly ; for, as already remarked, at very high temperatures many endothermic reactions proceed readily. Furthermore, at ordinary temperatures many changes like the interaction of ice and salt proceed spontaneously with absorp- tion of heat ; though here doubtless the amount of heat required for the liquefaction of the ice and salt is greater than that evolved by the action of the salt on the ice, whence the cooling effect observed. Thermochemical Equations. As already stated, it is cus- tomary to express the thermal accompaniment of a chemical change for the molecular weight in grams of the substances in- volved. Thus in making lead iodide from lead and iodine we have : [Pb] + 2 [I] = [PblJ + 39.8 Cal. indicating that when 207.1 grams of lead and 2 x 126.92 grams of iodine unite, 460.94 grams lead iodide are formed and 39.8 Cal. are simultaneously liberated. The brackets indicate that the substances are in the solid state. When liquids come into consideration parentheses are used, and in the case of gases, both brackets and parentheses are omitted. Thus, [P] yellow + 3 Cl = (PC1 8 ) + 76.6 Cal., THERMOCHEMISTRY 295 means that when 31 grams of solid yellow phosphorus react with 3 x 35.46 grams of gaseous chlorine to form 137.38 grams of liquid phosphorus chloride, 76.6 Cal. are liberated. Thermo- chemical equations must .not be confounded with the ordinary chemical equations. The latter indicate the direction the chem- ical change takes and specify the nature and amounts of the substances formed, whereas thermochemical equations are energy equations. For example the last equation states that the energy represented in 31 grams of solid phosphorus plus the energy in 106.38 grams of gaseous chlorine is equal to the en- ergy in 137.38 grams of liquid phosphorus trichloride plus 76.6 Cal., at room temperature, i.e. about 18 C. All other thermochemical equations are to be interpreted similarly. It should be stated that the use of brackets and parentheses to in- dicate solids and liquids respectively has been proposed but re- cently. It is a simple form of designation which will probably be generally adopted. Different allotropic forms of an element contain different amounts of energy. Thus when 31 grains of red phosphorus are converted into phosphorus trichloride, we have : [P] red + 3 Cl = (PC1 3 ) + 49.34 Cal.; therefore from the last two equations it follows that the conver- sion of yellow phosphorus to red proceeds with liberation of 27.26 Cal. thus: - [P] yellow = [P] red +27.26 Cal. . Since thermochemical equations are energy equations, they may be transformed like any algebraic equation, for instance: (1) (Hg) + 2 Cl = [HgClJ + 53.3 Cal. (2) (Hg) + 2 Cl - 53.3 Cal. = [HgClJ. ( 3 ) (Hg) + 2 Cl - [HgClJ = 53.3 Cal. ( 4 ) ( H g) + 2 Cl - [HgClJ - 53.3 Cal. = zero. Equation (1) indicates that when liquid mercury and gaseous chlorine unite to form solid mercuric chloride, 53.3 Cal. are lib- erated. Equation (2) indicates that if solid mercuric chloride were transformed into liquid mercury and gaseous chlorine, 53.3 Cal. would be absorbed. Equation (3) states that the energy in 200 grams mercury plus that in 2 x 35.46 grams 296 OUTLINES OF CHEMISTRY chlorine is greater than that contained in mercuric chloride try 53.3 Cal., and equation (4) expresses the same fact. The total energy contained in any substance is an unknown quantity, for we have no way of robbing a substance of all of its energy and measuring the same. A certain quantity of energy may, however, be obtained from substances; this is the available or free energy. It varies according to the nature of the changes to which a substance is subjected. So by burning phosphorus in excess of oxygen more heat is developed than by burning it in excess of chlorine: 2[P] + 5 O = [P 2 O 5 ] + 370 Cal. 2[P] + 10 Cl = 2[PC1 6 ] + 218.4 Cal. It is clear that thermochemistry can deal only with available energy. Definitions. The heat of solution is the thermal change ac- companying the solution of a substance in so much solvent that the addition of more solvent causes no further appreciable ther- mal change. The heat of solution is commonly stated per gram- molecule of dissolved substance, thus: [NaCl] + (aq) = (NaCl aq) - 1.3 Cal. indicates that when 1 gram-molecule of sodium chloride is dis- solved in much water (100 to 400 gram-molecules of water, which is indicated by aq in all thermochemical equations) there is formed the dilute solution NaCl aq, and 1.3 Cal. are absorbed. The heat of dilution is the thermal change accompanying the dilution of a given solution with a definite amount of pure solvent, usually so much that the addition of further solvent does not cause any appreciable change of temperature. The heat of reaction is a general term used to express the thermal change that accompanies any chemical reaction. The heat of formation of a chemical compound is the thermal change accompanying the formation of that compound from the ele- ments. The term is also sometimes used to indicate the ther- mal change that accompanies the formation of a compound from other compounds, or from elements and compounds. When so used, it is necessary to specify from what substances the compound whose heat of formation is under consideration has been formed. THERMOCHEMISTRY 297 The heat of neutralization is the heat liberated when an acid is neutralized by a base. The heat of combustion is the heat evolved when a substance is completely burned. In all cases the thermal change is computed per gram-molecule. Thermochemical Data. These generally consist of tables of heats of formation, solution, neutralization, and combustion. From what has been stated, tables of this kind will readily be understood. The thermochemical data of nearly all of the important substances have been determined by Thomsen and Berthelot. When the heat of formation of a compound in solution is known, the heat of formation in the anhydrous condition may be found by subtracting the heat of solution, carefully consid- ering the sign of the latter. By making use of the law of Hess, the heat of formation of any compound may be computed from the heat of any reaction involving that compound, providing the heats of formation of the other compounds in the reaction are known. In this way the heat of formation of a compound from the elements may be found indirectly, even though it has not been synthesized. Thus, let the heat of formation of cane sugar be required. Its heat of combustion found experiment- ally is: [C 12 H 22 O n ] + 12 O 2 = 12 CO 2 + 11(H 2 O) + 1353 Cal. Again by experiment it has been found that [C] + O 2 = CO 2 + 97.65 Cal., and H 2 + O = (H 2 O) + 68.4 Cal. It is clear then that 12 CO 2 when formed from 12 [C] and 12 O 2 will liberate 12 x 97.65 Cal. ; and similarly 11 (H 2 O) rep- resents a heat of formation of 11 x 68.4 Cal. Thus, 12 CO 2 and 11 (H 2 O) together represent a liberation of 12 x 97.65 + 11 x 68.4 or 1924.2 Cal. We may conceive of 12 [C] and 11 H 2 as oxidized in one step to 12 CO 2 and 11 (H 2 O) when 1924.2 Cal. are liberated ; or we may think of the operation as going on in two steps : (1) the oxidation of 12 [C] and 11 H 2 to sugar (i.e. to [C 12 H 22 O n ]), and then (2) the oxida- tion of the latter to 12 CO 2 and 11 (H 2 O). Now, since the complete oxidation evolves 1924.2 Cal. and step (2) evolves 298 OUTLINES OF CHEMISTRY 1353 Cal., it is evident that step (1), which is the formation of sugar from the elements, proceeds with an evolution of 1924.2 - 1353 or 571.2 Cal. In this computation 97.65 Cal. rep- resents the heat of combustion of amorphous carbon. The heat of combustion of diamond is 94.3 Cal. If the latter value be employed in the problem selected, the heat of formation from the elements will obviously be 571.212x3.35, or 531.3 Cal. The value 3.35 Cal. clearly represents the difference in energy between amorphous carbon and diamond. From the foregoing illustration, the value of the heats of formation of compounds in computing the thermal accompani- ments of chemical changes is evident. The following tables, giving the thermochemical data of a few of the most important compounds, will serve to illustrate how such results are usually presented: THERMOCHEMISTRY 299 TABLE 1 HEATS OF FORMATION VALUES ARE EXPRESSED IN LARGE CALORIES. THE SUBSTANCES NAMED ARE IN THE USUAL STATE AT 15 C. COMPOUND FORMED FROM GASEOUS LIQUID SOLID DISSOLVED HF H,F 38.5 45.7 50.3 HC1 H, Cl 22.0 39.3 HBr H, (Br) 8.6 28.6 HI H, [I] -6.1 13.1 H 2 O H 2 , 55.3 68.4 69.8 HA H 2 ,0 2 > 45.3 HA (H 2 0), -23.1 H 2 S H 2 , [S] rhombic 2.7 7.3 NH 3 N,H 3 12.0 16.6 20.4 PH 3 ' yellow [P], H 3 4.3 AsH 8 cry st. [As], H 3 -44.1 SbH 3 [Sb], H 3 -86.8 C 2 H 2 diamond [C 2 ], H 2 -58.1 C 2 H 4 [C 2 ], H 4 -14.6 C 2 H 6 [C 2 ], H 6 23.3 CH 4 [C], H 4 17.3 SiH 4 cryst. [Si], H 4 -6.7 Ob 2 -30.7 HC10 H, Cl, O 31.65 HC1O 3 H, Cl, 3 24.0 HC10 4 H, Cl, O 4 18.3 38.6 HBr0 3 H, (Br) 0, 12.3 HI0 3 H, [I], 3 57.9 55.7 HI0 4 H, [I], 4 47.6 SO 2 rhombic [S], O 2 71.0 78.8 S0 3 rhombic [S] , O,j 103.3 142.5 H 2 SO 4 H 2 , [S], 4 189.9 210.9 H 2 S 2 3 H 2 , [S], 3 141.7 N 2 O N 2 , -17.4 NO N, -21.5 N 2 3 N 2 ,0 3 6.8 NO 2 N,0 2 -7.7 N 2 4 N 2 ,0 4 -2.6 - N 9 O, N,, O, 13.1 oq Q 25 HN0 3 * 2'. 5 H, N, 3 34.4 41.5 42.2 0.0 48.8 PA yellow [P 2 ],O 5 . 370.0 406.0 H 3 P0 4 H 3 , [P], 4 304.1 306.8 As 2 O 3 [As 2 ], 3 154.7 147.0 Sb 2 3 [SbJ, 3 166.9 Bi 2 3 [Bi 2 ], 3 139.2 B 2 3 amorph. [BJ, O 3 272.6 _____ SiO 2 aq cryst. [Si], O2, aq. 179.6 CO amorph. [C], O 29.4 CO diamond [C], O 26.1 C0 2 amorph. [C], O 2 97.65 103.25 300 OUTLINES OF CHEMISTRY TABLE 1 Continued COMPOUND FORMED FROM GASEOUS LIQUID SOLID DISSOLVED C0 2 diamond [C], O 2 94.3 99.9 PC1 3 yellow [P], C1 3 69.3 76.6 PC1 5 yellow [P], C1 5 109.2 AsCl 3 [As], C1 3 71.7 SbCl 3 [Sb], Cl, 91.4 BiCl 3 [Bi], Cl, 96.6 CC1 4 diamond [C], C1 4 68.5 75.7 SiCl 4 cryst. [Si], C1 4 121.8 128.1 SnCl 4 [Sn], C1 4 122.2 129.8 _ (CN) 2 diamond [C 2 ], N" 2 -73.9 -68.5 67.1 HCN diamond H, [C], N" -30.5 -24.8 24.4 CS 2 diamond [C], [S 2 ] rhombic -25.4 -19.0 TABLE 2 HEATS OF FORMATION OF SOME METALLIC COMPOUNDS AT 15 C. VALUES ARE EXPRESSED IN LARGE CALORIES COMPOUND FORMED FROM SOLID DIS- SOLVED COMPOUND FORMED FROM SOLID DIS- SOLVED KOH [K] , 0, H 103.2 116.5 NH 4 C1 N, H 4 , Cl 75.8 71.9 NaOH [Na], O, H 101.9 111.8 CaC] 2 [Ca], C1 2 169.8 187.2 LiOH [Li], 0, H 112.3 118.1 ZiiCl 2 [Zn], C1 2 97.2 112.8 NH 3 aq NH 3 , ?i(H 2 O) 20.3 A1C1 3 [Al], C1 3 161.0 237.8 MgO [Mg], 143.4 FeCl 2 [Fe], C1 2 82.1 100.0 CaO [Ca], 131.5 149.6 NiCl 2 [Ni], C1 2 74.5 93.7 Ca(OH) 2 [Ca], 2) H 2 214.9 217.9 CoCl 2 [Co], C1 2 76.5 94.8 SrO [Sr], 128.4 157.7 HgCl (Hg), Cl 31.4 BaO [Ba], O 124.2 158.7 HgCl 2 (Hg),Cl 2 53.3 50.0 BaO 2 [Ba], 2 12.1 AgCl [Ag], Cl 29.4 MnO [Mn], 90.9 AuCl [Au], Cl 5.8 Mn0 2 [Mn], O 2 125.3 AuCl 8 [Au], C1 8 22.8 27.3 FeO [Fe], 65.7 PtCl 4 [Pt], C1 4 59.8 79.4 Fe 8 4 [Fe 3 ], 4 270.8 NaBr [Na],(Br) 85.8 83.9 ZnO [Zn], 84.8 Nal [Na], [I] 69.1 70.3 CuO [Cu], 37.2 Na 2 C0 3 [Na 2 ], [C],0 3 269.9 275.4 Cu 2 O [Cu 2 ], 40.8 NaHC0 3 [Na],H,[C],O a 227.0 223.7 PbO [Pb], 50.3 Na 2 SO 4 [Na 2 ], [S], 4 328.4 329.0 PbO 2 [Pb], 2 63.4 NaHS0 4 [Na],H,[S],0 4 267.8 266.6 HgO (Hg), 20.1 KN0 3 [K], N, 3 119.5 111.0 KSH [K],[S],H 62.3 63.1 KClOg [K], Cl, 8 95.0 85.0 CaS [Ca], [S] 89.6 KBrOa [K], (Br),0 3 84.1 74.3 SrS [Sr], [S] 97.4 KlOg [K], [I], 3 124.5 117.4 BaS [Ba], [S] 98.3 KCN [K], [C], N 29.8 26.8 FeS [Fe], [S] 24.0 KMn0 4 [K], [Mn], 4 195.0 184.8 CuS [Cu], [S] 10.1 AgN0 3 [Ag], N, 8 28.7 23.3 Ag 2 S [Ag 2 ], [S] 3.3 CuSO 4 [Cu], [S], 4 1828 '198.4 NaCl [Na], Cl 105.6 101.2 BaSO 4 [Ba], [S], 4 338.1 THERMOCHEMISTRY 301 TABLE 3 HEATS OF COMBUSTION OF SOME CARBON COMPOUNDS VALUES ARE GIVEN IN LARGE CALORIES. SUBSTANCES ARE IN THEIR USUAL STATE AT 15 C. COMPOUND FORMULA HEAT OF COMBUSTION HEAT OF FORMATION DIAMOND = [C] Methane CH 4 213.5 16.5 Ethane C 2 H 6 370.5 22.1 Propane 529.2 25.4 Benzene CA 787.8 -9.1 Methyl alcohol CH 3 OH 170.6 61.7 Ethyl alcohol C 2 H 5 OH 325.7 69.9 Glycerine C 3 H 5 (OH) 3 397.2 161.7 Acetic acid CH 3 COOH 209.4 117.2 Oxalic acid (COOH) 2 60.2 196.7 Stearic acid C 18 H 36 O 2 2677.8 227.6 Starch C 6 H 10 O 5 " 684.9 225.9 Dextrine C 6 H 10 5 667.2 243.6 Cellulose C 6 H 10 5 680.4 230.4 Cane sugar C 12 H 22 O n 1353.0 531.3 Milk sugar C 12 H 2211 1351.4 537.4 Malt sugar C 12 H 22 O n 1350.7 538.1 Dextrose C 6 H ]2 6 677.2 302.6 Laevulose C 6 H 12 6 675.9 303.9 Urea CO(NH 2 ) 2 152.2 77.5 TABLE 4 HEATS OF COMBUSTION OF VARIOUS OTHER ORGANIC SUBSTANCES SUBSTANCE HEAT OF COMBUSTION PER 1 GRAM SUBSTANCE Butter Animal or vegetable fats and oils, average Caseine ....... White of egg ...... Egg yolk Peptone Gluten Muscular tissues ..... Fibrin Hemoglobin ..... o 9.2 Cal. 9.5 Cal. 5.6 Cal. 5.7 Cal. 8.1 Cal. 5.3 Cal. 6.0 Cal. 5.7 Cal. 5.5 Cal. 5.9 Cal. 302 OUTLINES OF CHEMISTRY TABLE 5 HEATS OF NEUTRALIZATION VALUES GIVEN IN LARGE CALORIES. (THE SOLUTIONS CONTAINED 1 GRAM EQUIVALENT OF ACID OR BASE IN Two LITERS. SOME o* THE BASES USED AND SULPHATES FORMED ARE INSOLUBLE.) BASES HClaq HN0 3 aq CH 8 COOH aq lHS0 4 aq HCNaq KOHaq 13.7 13.8 13.3 15.7 3.0 NaOH aq 13.7 13.7 13.3 15.85 2.9 NH 4 OH aq ' 12.45 12.5 12.0 14.5 1.3 * Ca(OH) 2 aq 14.0 13.9 13.4 15.6 3.2 * Sr(OH) 2 aq 14.0 13.9 13.3 15.4 3.1 $ Ba(OH), aq 13.85 13.9 13.4 18.4 3.2 Mg(OH) a aq 13.8 13.8 15.6 J Fe(OH) 2 aq 10.7 9.9 12.5 Ni(OH) 2 aq 11.3 13.1 $ Co (OH) 2 aq 10.6 13.3 Zn(OH) 2 aq 9.8 9.8 8.9 11.7 Cu(OH) 2 aq 7.5 7.5 6.2 9.2 Uses of Thermochemical Data. From Table 1 it appears that the heat of formation of the hydrohalogens diminishes as the atomic weight of the halogens increases, hydriodic acid even having a negative heat of formation. We have seen that the stability of these compounds diminishes in the same way, hydro- fluoric acid being the stablest and hydriodic acid the least stable. On the other hand, it will be recalled that iodic and periodic acids are more stable than chloric, bromic, and perchloric acids, and, indeed, Table 1 shows that the heats of formation of iodic and periodic acids are higher than those of the corresponding compounds of the other halogens. Water is a much stabler compound than hydrogen sulphide, as is borne out by the great difference in their heats of formation. Ammonia, phosphine, arsine, and stibine diminish in stability in the order named, which is precisely what one would expect from their heats of formation, which also diminish in the same order. Marsh gas and ethane, it will be observed, have positive heats of forma- tion, whereas the unsaturated compounds ethylene and acety- lene are formed with absorption of heat. The formation of ozone from oxygen takes place with absorption of much energy, as the negative heat of formation of ozone indicates. From these illustrations and from others with which Tables 1 and 2 THERMOCHEMISTRY 303 are replete, it appears that thermochemical data offer a means of comparing the relative stability of compounds. Both Thomsen and Berthelot had hoped that thermochemical data would offer a means of exact measurement of chemical attractions, but this has not been realized. Thermochemical data are complicated by the fact that they also represent the energy concomitants of physical changes which invariably accompany chemical reactions, and which cannot be evaluated, as already explained. Moreover, it must be borne in mind that, in speak- ing of the stability of a substance, it is really necessary to specify toward what agencies such stability is being considered. Thus, a substance A might be much stabler than another substance B towards the decomposing action of heat, whereas towards the action of electricity, light, or the inroads of various reagents, A might be less stable than B. So, for instance, carbon tetra- chloride, with its heat of formation +68.5, ought to be less stable than silicon tetrachloride, whose heat of formation is + 121.8. While this is substantiated by the fact that silicon tetrachloride may readily be obtained by passing chlorine over hot silicon, whereas carbon tetrachloride cannot be similarly obtained, it must also be borne in mind that when treated with water, silicon tetrachloride is at once decomposed into hydro- chloric and silicic acids, whereas carbon tetrachloride remains unchanged under the same treatment. However, here the fact that the heat of formation of silicic acid by far exceeds that of carbonic acid no doubt is a determining factor. By means of the electric current neither SiCl 4 nor CC1 4 can be decomposed, whereas common salt, which per gram equivalent has over five times as high a heat of formation as carbon tetrachloride, is nevertheless easily decomposed by electrolysis (which see). Thus it is clear that great care must be exercised in using thermo- chemical data in arguing as to the relative stability of compounds. The value of fuels depends upon their heat-giving power; that is, their heat of combustion. And so it is clear that the heats of combustion of wood, coal, and various liquid and gas- eous fuels is of utmost practical importance. In the animal body the foods consumed are digested, assimilated, and finally slowly oxidized and eliminated in the form of carbon dioxide and water in the case of carbohydrates and fats, and in the form of carbon dioxide, water, urea, and other nitrogenous products 304 OUTLINES OF CHEMISTRY in the case of nitrogenous foods. Therefore the heats of com- bustion of foodstuffs have sometimes been considered in deter- mining the value of various foods. In such a procedure great care must again be exercised ; for foods that have nearly the same heat of combustion are frequently of quite different value, because they are not all digested and assimilated with equal readiness. Compare, for example, the heats of combustion of starch and cellulose in Table 3 ; the values are nearly the same, and yet the food value of the substances to an animal is very different. An inspection of the heats of combustion in Table 3 shows that analogous substances of the same carbon and hydrogen content have approximately the same heats of combustion, in spite of their differences in structure. Nevertheless, differences in structure do yield corresponding differences in heats of com- bustion. This matter has been studied in some detail, espe- cially by Stohmann. Adjacent members of homologous series on the average show a difference of about 158 Cal. for CH 2 . The heat of combustion of carbon compounds is approximately an additive property. In Table 4 are given the heats of com- bustion of a few additional important substances. It will be observed that the heats of neutralization of differ- ent bases by different acids, Table 5, are approximately the same in the case of the strong bases and strong acids. This will be discussed in connection with the subject of electrolytic dissociation. In general, Table 5 shows that when a given acid is neutralized, the heat thus developed by bases that are known to be closely related chemically is approximately the same. So when hydrochloric acid is neutralized by sodium or potassium hydroxide, the heat of neutralization is 13.7 Cal. When the same acid is neutralized by ferrous, cobaltous, or nickelous hydroxide, the heat developed is about 10.8 Cal. THERMOCHEMISTRY 305 REVIEW QUESTIONS 1. Define the following terms: exothermic change, endothermie change, latent heat of fusion, latent heat of vaporization, calorie, British thermal unit, heat of reaction, heat of combustion. 2. What is the law of Hess? Give an example of the law. 3. How do thermochemical equations differ from ordinary chemical equations? Illustrate by means of a concrete example. 4. What is Berthelot's law of maximum work? Discuss its validity. 5. What is meant by each of the following: heat of solution, heat of dilution, heat of neutralization, heat of formation? Give an illustra- tion in each case. 6. How proceed to find the heat of formation of starch? Explain in detail. 7. Of what use is thermochemistry : (a) in the study of the value of fuels, (6) in the study of the value of foodstuffs ? Illustrate by means of concrete instances. CHAPTER XVIII SILICON AND BORON AND THEIR IMPORTANT COMPOUNDS Occurrence, Preparation, and Properties of Silicon. Next to oxygen, silicon is the most abundant element found in the earth's crust, constituting more than one fourth of ^ the latter. Silicon does not occur in the free state. It is always found in combination with other elements, especially with oxygen as silica, and with oxygen and various metals as silicates. Quartz, quartzite, flint, and the white sands of the seashore and the deserts are nearly pure silicon dioxide; whereas clays are largely composed of silicates. Silicon was first prepared in pure form in 1823 by Berzelius, who heated potassium silicofluoride with metallic potassium : K 2 SiF 6 + 4 K = 6 KF + Si. The element may also be obtained by heating finely powdered quartz sand with magnesium powder : SiO 2 + 2 Mg = 2 MgO + Si. In this case magnesium silicide Mg 2 Si is generally also formed ; but by means of hydrochloric acid the silicon can readily be freed from this silicide and also from the oxide of magnesium. Silicon may also be obtained by heating sodium or aluminum in a current of silicon tetrachloride vapor, thus : SiCl 4 + 4Na = 4NaCl+Si. 3 Si01 4 + 4 Al = 4 A1C1 8 + 3 Si. On a large scale, silicon is now manufactured at Niagara Falls by heating together quartz sand and coke in the electric furnace, thus : SiO 2 +2 C = 2 CO + Si. Silicon is run out of the electric furnaces into molds. It thus forms " pigs " that weigh from 600 to 800 pounds. The material varies in purity from 90 to 97 per cent, though silicon SILICON AND BORON 307 over 99 per cent pure has thus been prepared. Silicon is sold in car lots at about $120 per ton. It is mainly used in the steel industry as a reducing agent. Hundreds of tons of 90 per cent silicon are used annually in manufacturing steel. It is very likely that silicon will be used for many other purposes in the near future. Silicon is either crystalline, or an amorphous brown powder. In the latter form it is commonly obtained by the first three methods above described. Amorphous silicon burns when highly heated in the air, the product being silicon dioxide SiO 2 . Since the latter is practically non-volatile, its accumulation hinders the securing of complete oxidation of all the silicon. Under a layer of common salt, amorphous silicon may be melted, and on cooling it becomes crystalline. Silicon may also be obtained in crystalline form by dissolving amorphous silicon in molten zinc ; on cooling, silicon separates out in form of crystals. The zinc may be removed with hydrochloric acid. Silicon crystallizes in the isometric system, forming dark gray shining plates or rods, which in reality consist of octahedra that have grown together so as to form twin crystals. The specific gravity of silicon is 2.49. It is so hard that it will scratch glass. The crystalline variety conducts electricity, though rather poorly. The amorphous powder is a non-conductor. Like graphite, crystalline silicon is hard to oxidize by heating it in the air or in oxygen. Hydrofluoric acid attacks it but slowly ; nitric and hydrofluoric acids act on it more rapidly. Fluorine reacts with silicon even at ordinary temperatures with evolution of light and heat : Si -f 4 F = SiF 4 . Hot solutions of caustic potash dissolve silicon : 2 KOH + H 2 + Si = K 2 Si0 8 + 2 H 2 . The atomic weight of silicon is 28.3. It has been determined by analyzing its compounds with the halogens. Silicon is quad- rivalent in all of its compounds, the formulae of which conse- quently are analogous to those of the compounds of carbon. Indeed, silicon and carbon bear many resemblances in their chemical behavior, and while carbon is exceedingly important in the organic world, silicon plays a similar role in the inorganic realm. 308 OUTLINES OF CHEMISTRY Silicon Dioxide, Silica. This is by far the most important compound of silicon. Its formula is SiO 2 . It is silicic acid anhydride. In the form of quartzite, it often forms mountains. It is the chief constituent of sandstones, and sand is largely silica. In crystalline form it occurs as quartz and amethyst, and also, though rarely, as tridymite. In the amorphous form it is found as agate, opal, flint, carnelian, and chalcedony, which frequently contain water in combination. Pure silicon dioxide is colorless, but many of the varieties found in nature are colored by vari- ous impurities. Thus smoky quartz is discolored with organic matter, rose quartz with manganese, carnelian with oxide of iron, etc. Quartz crystallizes in the hexagonal system. Its crystals occur in two forms that are non-superposable (Fig. 113); that is, they are to each other as the right hand is to the left. These crystals rotate the plane of polarized light passed through them paral- lel to the main axis. The degree of rotation is propor- tional to the thickness of the layer traversed, and the deviation is either dextro or laevo according to the crystal used. This property makes quartz useful in certain kinds of optical in- struments, particularly in certain types of polariscopes. Tri- dymite also crystallizes in the hex- agonal system. It usually occurs in prismatic plates (Fig. 114). Quartz is brittle and very hard. It is consequently used as an abra- sive material in grinding glass, metals, etc. Glued on paper, it forms sandpaper. Its specific gravity is 2.6. It requires the temperature of the oxyhydrogen flame to melt quartz. When thus heated, it forms a viscous liquid that can be drawn out and worked like glass. In the electric furnace, the liquid may be boiled and evaporated. In recent years flasks, crucibles, evaporating dishes, and other utensils have been made of quartz glass. These have the great advantage that they will not break when subjected to sudden and very great differences Fia. us. FIG. 114. SILICON AND BORON 309 of temperature. This is due to the 'fact that quartz changes its volume but very slightly with alterations of temperature. The coefficient of expansion of quartz between and 1000 is only 0.0000007 on the average, being less than that of any other known substance. A white-hot quartz crucible may be quenched in cold water without injuring the dish. Silica constitutes about 40 per cent of the ash of the feathers of birds. It is also found in egg albumin, in the hair of ani- mals, and in various crustaceans. Diatomic or infusorial earth consists of the siliceous remains of minute organisms called diatoms or infusoria. The stalks of grasses, cereals, field horsetails, bamboo, and other canes contain notable amounts of silica, which is in combination with other elements and aids in giving the stalks stability. Sometimes over half of the ash of these stalks consists of silica. Besides being used as an abrasive material, silica is employed in the manufacture of glass and in making mortar, cement, and porcelain. Silicic Acids. Silicon dioxide is the anhydride of a series of silicic acids. These may all be considered as composed of silica and water in various proportions. They may all be referred to orthosilicic acid Si(OH) 4 , which is well known in the form of salts, though it has not been prepared in the pure state. The acid is probably present in the gelatinous precipitate formed when silicon tetrachloride or tetrabromide is treated with water : SiCl 4 + 4 H 2 = 4 HC1 + Si(OH) 4 . By losing a molecule of water, orthosilicic acid passes over into metasilicic acid H 2 SiO 3 . From two molecules of ortho- silicic acid by elimination of one, two, and three molecules of water the following acids, commonly known as disilicic acids, are formed: H 6 Si 2 O 7 , H 4 Si 2 O 6 , H 2 Si 2 O 5 . From three molecules of the ortho acid by loss of two and four molecules of water the trisilicic acids H 8 Si 3 O 10 and H 4 Si 3 O 8 are formed. None of these polysilicic acids have been isolated. Their existence is simply vouchsafed by the fact that salts of these acids occur in nature, or have been made in the labora- tory. The mineral olivine Mg 2 SiO 4 (Fig. 70) is a salt of 310 OUTLINES OF CHEMISTRY H 4 SiO 4 ; sodium silicate, or water glass, Na 2 SiO 3 is a salt oi H 2 SiO 3 ; serpentine Mg 3 Si 2 O 7 is a salt of H 6 Si 2 O 7 ; and the feld- spars, orthoclase AlKSi 3 O 8 and albite AlNaSi 3 O 8 , are salts of H 4 Si 3 8 . When silica is fused with sodium carbonate, sodium silicate is formed : Na 2 C0 3 + Si0 2 = Na 2 Si0 3 + CO 2 . Sodium silicate is soluble in water and is known as water glass, as is also the silicate of potassium K 2 SiO 3 , which may be made similarly. The silicates of metals other than the alkalies are very slightly soluble in water. On treating a solution of sodium or potassium silicate with a mineral acid, silicic acid is set free : Na 2 SiO 3 + 2 HC1 = 2 NaCl + H 2 SiO 3 . If the solution is concentrated, the silicic acid is precipitated in the form of a jelly. If dilute solutions are used and the water glass is poured into an excess of hydrochloric acid, no precipi- tate forms. From this clear solution, the sodium chloride and excess of hydrochloric acid may be removed by dialysis. The apparatus required for the purpose is called a dialyser, a common form of which is shown in Fig. 115. A parchment paper or animal bladder is se- curely tied over one end of a cylinder into which the solution is poured. The whole is then im- mersed in a larger outer dish of water as shown in the figure. The so- dium chloride and hy- drochloric acid pass through the septum into the outer liquid, while the silicic acid remains behind in the inner vessel. By renewing the water in the outer dish from time to time, practically all of the chlorides can be removed from the inner vessel, which then contains only a solution of silicic acid. This may be concentrated FIG. 115. SILICON AND BORON 311 by careful evaporation to about 10 per cent, if not quite all of the chlorides have been removed, or to about 1 per cent if practically all the chlorides have been taken out. If attempts are made to concentrate to a greater extent or to preserve the solution for a long time, the silicic acid largely separates out in form of a gelatinous mass, which is termed a hydrogel, the clear solu- tion from which the latter has been formed being termed a hydrosol. The solution of silicic acid obtained by dialysis as described is also commonly called a colloidal solution. This term was introduced by Thomas Graham to denote solutions of non-crystalline bodies which do not pass through membranes used in dialysis experiments. Thus Graham found that, like silicic acid, substances such as glue, gums, albumin, ferric hydroxide, etc., which are non-crystalline, do not pass through parchment or animal membranes as readily as crystalline sub- stances. He consequently made two classes of substances : colloids, which do not pass through membranes on dialysis ; and crystalloids, which do make their way through such septa readily. Though this distinction is still frequently made, it really cannot be held in the light of more recent experiments ; for it is quite possible to separate crystalline substances from each other by this process. It is even possible to effect the separation of crystalline from non-crystalline substances by having the latter pass through the septum and the crystal- line substances remain behind. It all depends upon the nature of the septum chosen and the character of the substances under consideration. So when cane sugar and camphor, both crystal- line substances, are dissolved in pyridine, and the solution is separated from pure pyridine by means of a vulcanized caout- chouc membrane, such as the dentists use as "rubber dam," camphor passes through and sugar remains behind. Again, when copper oleate, a non-crystalline substance, and cane sugar together in pyridine solution are similarly subjected to dialy- sis, the copper oleate passes through the rubber membrane, and the crystalline sugar remains behind. Finally, if to a solu- tion of collodion in alcohol and ether, copper oleate is added and this solution is then (by means of a rubber membrane) separated from a mixture of alcohol and ether such as is used in making up the collodion copper oleate solution, the copper oleate passes through the septum and the nitrocellulose remains behind. 312 OUTLINES OF CHEMISTRY As both copper oleate and nitrocellulose are non-crystalline in character, we have here a case of the separation of two non-crystalloids, that is, in Graham's language, of two colloids, by dialysis. On drying gelatinous silicic acid, it loses water and finally forms a white amorphous powder which must be heated in the blast to expel all traces of moisture. Action of Water on Silicates. Silicates are difficultly soluble in water, yet the earth's crust is continually being worn away by the solvent action of rain water upon the siliceous geological deposits. Rocks like granite, gneiss, schists, shales, and slates are continually being washed away by the solvent action of water, slight though it be. Thus a gradual leveling process is going on which is aided by the action of wind and the disintegrating effects of alternate freezing and thawing. So the silicates are gradually dissolved, and the more resistant quartz grains remain behind as sand. This, however, finds its way into the sea and other depressions filled with water, where the sand grains are frequently gradually cemented together with calcium carbonate or oxides of iron, thus forming so-called sandstones. As silicic acid is a very weak acid, which is evident from the fact that its solutions neither react toward litmus nor have any taste, we should expect solutions of silicates to contain these salts, largely in a state of hydrolytic decomposition ; and such is actually the case. Decomposition of Silicates in the Laboratory. This is effected b}' fusing the pulverized silicate with sodium carbonate or a mixture of this salt and potassium carbonate. In this way sodium silicate is formed which is soluble in water. The other bases present may generally be readily dissolved with the aid of hydrochloric acid. Silicates may also be decomposed with hydrofluoric acid, or with this and hydrochloric -or sul- phuric acid. Thus the silicon is volatilized in form of SiF 4 , and the bases remain as chlorides or sulphates. Silicates may further be decomposed by heating with calcium carbonate and ammo- nium chloride and then extracting the mass with water. In the latter process calcium silicate is formed, and the bases are converted into chlorides. Hydrogen Silicide. When magnesium silicide Mg 2 Si is SILICON AND BORON 313 treated with hydrochloric acid, hydrogen silicide or silico methane SiH 4 is formed : Mg 2 Si + 4 HC1 = 2 MgCl 2 + SiH 4 . The colorless gas so obtained always contains hydrogen and some silicoethane Si 2 H 6 . Pure SiH 4 does not take fire in the air except under diminished pressure. Silicoethane, however, ignites spontaneously on exposure to the air, and it is this gas whose presence causes SiH 4 to burn in contact with air at ordinary pressure. Silicon tetrahydride may be liquefied at 11 under a pressure of 50 atmospheres. In chlorine gas SiH 4 takes fire. On burning silicon hydride in the air, the products formed are water and silica ; the latter forms white smoke. Silicoethane Si 2 H 6 boils at 4- 52. Compounds of Silicon with the Halogens. Silicon tetrafluoride SiF 4 is formed by treating silicon with fluorine, or more readily by treating silica with hydrofluoric acid or a mixture of fluor- spar CaF 2 and sulphuric acid, thus : CaF 2 + H 2 S0 4 = 2 HF + CaSO 4 , and Si0 2 + 4 II F = 2 H 2 + SiF 4 ; or 2 CaF 2 + SiO 2 + 2 H 2 SO 4 = 2 CaSO 4 + 2 H 2 O + SiF 4 . Silicon tetrafluoride is a colorless gas of very pungent odor. It boils at - 65, and the solid melts at 77. The tetra- fluoride is always formed when hydrofluoric acid acts on sili- cates, and it is consequently produced when that acid is used in etching glass. Water decomposes silicon tetrafluoride : 3 H 2 + 3 SiF 4 = H 2 Si0 3 + 2 H 2 SiF 6 . The silicic acid formed separates out as a gelatinous precipitate, while the hydrofluosilicic acid H 2 SiF 6 remains in solution. The latter may be concentrated to some extent by evaporation. The concentration must be carried on in a platinum dish, because hydrogen fluoride is formed during the process, and so glass or porcelain dishes would be attacked. Pare H 2 SiF 6 is not known, for on attempting to concentrate its solutions beyond a certain point the acid breaks up, yielding hydrogen fluoride and silicon tetrafluoride. In making fluosilicic acid the silicon tetra- 314 OUTLINES OF CHEMISTRY fluoride generated in a flask by the reaction above mentioned is conducted into water by means of a tube whose lower end dips in mercury (Fig. 116), so that the gelatinous silicic acid formed will not stop the end of the tube. As the gas rises from the mercury, clouds of silicic acid are formed in the water. Hydrofluosilicic acid is a strong acid. It readily decomposes carbonates and hy- droxides of the metals, form- ing the fluosilicates. The latter are decomposed by heat, yielding fluorides of the metals and silicon tetra- fluoride. The silicofluorides are commonly soluble in water, insoluble, and the potassium salt is The barium salt is sparingly soluble. Silicon tetrachloride SiCl 4 is formed by heating silicon in a current of chlorine, or by passing chlorine over a heated mix- ture of carbon and silica, thus : Si + 2 Cl 2 =SiCl 4 , or SiO 2 + 2 C + C1 4 = 2 CO + SiCl 4 . The product is a liquid of pungent odor. It boils at 59, has a specific gravity of 1.52 at 0, and solidifies at 89. Water at once decomposes it : SiCl 4 + 4 H 2 O = 4 HC1 + H 4 SiO 4 . On heating silicon in a current of hydrochloric acid gas, silicon chloroform SiHCl 3 may be obtained. This boils at 34, has a specific gravity of 1.3, and, like silicon tetrachloride, it is at once decomposed by water. Bromine and iodine compounds of silicon, analogous to the chlorine compounds, have been prepared by similar methods. Silicon tetrabromide boils at 153 and melts at 12, silicon tetraiodide SiI 4 forms octahedra that melt at 120 and boil at 290. SILICON AND BORON 315 Esters of Silicic Acid. Methyl silicate (CH 3 ) 4 SiO 4 , boiling at 121, and ethyl silicate (C 2 H 6 ) 4 SiO 4 , boiling at 165, are also known. They are formed by the action of alcohols on silicon tetrachloride, thus : SiCl 4 + 4 CH 3 OH = 4 HC1 + (CH 3 ) 4 SiO 4 . Water decomposes the esters to alcohol and silicic acid. Silicon Carbide, Carborundum, SiC. This substance is formed in the electric furnace by heating together silica or quartz sand, carbon, and common salt to about 3500. The following reaction occurs : SiO 2 + 3 C = 2 CO + SiC. Silicon carbide forms hexagonal plates that commonly have a dark greenish blue color. The substance is not attacked by acids; not even hydrofluoric acid makes inroads upon it. It may readily be decomposed, however, by fusion with caustic alkalies. Carborundum has a specific gravity of 3.2, and is extremely hard, being next to the diamond in hardness. It is consequently used as an abrasive material. Grinding wheels, whetstones, etc., made of carborundum are in common use. Titanium, Zirconium, and Thorium. These are quadrivalent metallic elements whose compounds are analogous to those of silicon. The elements are steel-gray, brittle metals. Titanium (Ti 48.1) is found in nature as the dioxide TiO 2 , in form of rutile, brookite, and anatase. The element is widely distributed, but occurs nowhere in large quantities. It is also met in titaniferous iron ores, which are in the main ferrous titanate FeTiO 3 . It also occurs together with zircon in certain silicates. Zirconium (Zr 90.6) is found chiefly in the mineral zircon, which forms tetragonal crystals of the composition ZrSiO 4 , from which Klaproth prepared the dioxide ZrO 2 in 1789. Moissan prepared the metal by heating the oxide with carbon in the electric furnace. Though in its compounds cerium is also more frequently quadrivalent, it will nevertheless be discussed in connection with lanthanum and other rare-earth elements (which see). Thorium (Th 232.4) was found in thorite ThSiO 4 2 H 2 O by Berzelius, in 1828. Thorium salts are now prepared from 316 OUTLINES OF CHEMISTRY monazite found in North Carolina. Welsbach light mantles consist of 99 per cent thoria ThO 2 and 1 per cent ceria CeO 2 (see under cerium). Thorium compounds are radio-active (see radium). Occurrence, Preparation, and Properties of Boron. This ele- ment occurs in nature in the form of boric acid and its salts, called borates. . Of the latter borax, the sodium salt, and borocalcite and colemanite, which are calcium salts, are the most im- portant. The methods of preparing boron are analogous to those of making silicon. So boron may be prepared by reduction of its oxide by means of potassium, sodium, magnesium, or aluminum or by passing the vapors of boron chloride over heated sodium. An amorphous and a crystalline variety of boron are known. The former results when the oxide B 2 O 3 is reduced with potas- sium, or when borax is heated with magnesium powder. Amorphous boron is a brown powder. On being heated in the air it burns, forming the oxide B 2 O 3 and the nitride BN. Sulphuric or nitric acid and other oxidizing agents convert boron into boric acid. When fused with caustic alkalies or their carbonates, borates result. Amorphous boron dissolves in molten aluminum, and on cooling it crystallizes out in tetragonal crystals, which are transparent and generally some- what colored, due to impurities. These crystals are nearly as hard as the diamond. They are less readily attacked by reagents than the amorphous variety. Boron is trivalent in all of its compounds. Its atomic weight is 11. While boron resembles silicon and carbon in many respects, the formulae of its compounds, owing to its trivalence, are analogous to those of the compounds of the phosphorus group and to those of aluminum. The latter element and boron really belong to the same family, though aluminum is a pronounced metal and shows but slight acid-forming properties, while just the opposite is true of boron. The latter really occupies a somewhat lone position amongst the chemical elements. Boric Acid and its Salts. Boric acid H 3 BO 3 occurs in vol- canic regions, particularly in Tuscany, where it issues from the earth in jets of steam. These jets, which contain only small amounts of boric acid, are called soffioni, whereas the hot SILICON AND BORON 317 i springs from which the jets issue are termed fumaroles. The vapors are condensed in small natural or artificial basins sur- rounding the fumaroles, and the boric acid is finally obtained by evaporation to the point at which the acid crystallizes out, the heat necessary being furnished by the hot springs. The presence of boric acid in these steam jets is due to the fact that boric acid may be volatilized with water vapor. In the Caucasus Mountains and in some of the hot springs of Cali- fornia, boric acid issues from the earth in a similar manner. 'Much boric acid is also prepared from borax Na 2 B 4 O 7 -10 H 2 O, particularly in Nevada and California. A hot, concentrated solution of borax is treated with either hydrochloric or sulphuric acid, and on cooling boric acid crystallizes out. The reaction is : Na 2 B 4 O 7 + 5 H 2 O + 2 HC1 = 2 NaCl + 4 H 3 BO 3 . Boric acid crystallizes in shining white scales which are "soapy" to the touch. At 18, 100 parts of water dissolve 3.9 parts of boric acid, whereas at 100, 33 parts of the acid are thus dissolved. This fact makes it simple to recrystallize boric acid from its aqueous solutions. The acid is quite weak. It affects litmus but slightly, and its taste is not sour but simply astringent, and somewhat bitter. Solutions of boric acid turn turmeric paper reddish brown. To bring out this color the paper must be dried when very dilute solutions are used. This test for boric acid is a very delicate one. When the paper reddened by boric acid is treated with caustic alkali, a black stain is produced, which further serves to characterize boric acid. On treating boric acid with alcohol and sulphuric acid, a volatile ester, ethyl borate, is formed, which when ignited burns with a characteristic green flame. This also serves as a test for boric acid. Boric acid is often used in medicine and surgery as an antiseptic. It is also employed in making certain glazes for pottery, and it is still sometimes used as a preserva- tive for meat, fresh fish, milk, and other foods. The latter practice is to be condemned, because the substance is injurious to health. At 100 boric acid loses water and so forms metaboric acid HBO 2 , which on further heating to 140 passes over into pyroboric acid or tetraboric acid H 2 B^O 7 . The latter on igni- 318 OUTLINES OF CHEMISTRY tion forms the trioxide or boric anhydride B 2 O 8 , which fuses at a high temperature and congeals to a glassy mass on cooling. When treated with water it forms boric acid. Salts of the acid H 3 BO 3 are not known, but the esters like (CH 3 ) 3 BO 3 and (C 2 H 6 ) 3 BO 3 are well known. Metaborates like NaBO 2 have been formed, but they are unstable. By far the most important salt of boric acid is borax Na 2 B 4 O 7 10 H 2 O. It is the sodium salt of tetraboric acid H 2 B 4 O 7 , which may be considered as 4 H 3 BO 3 minus 5 H 2 O. Borax is found in the borax lake of California and in certain marshes of that state and Nevada. It also occurs in Thibet, Ceylon, and Bolivia. Large quantities of borax and boric acid are prepared from colemanite Ca 2 B 6 O n 5 H 2 O, which is found in California and Oregon. The amount of borax produced annually from the deposits in the United States is in the neigh- borhood of 50,000 tons. Borax solutions have a slightly alkaline reaction toward indicators, which is explained by the fact that boric acid is weak and its salts are somewhat decomposed by hydroly- sis. At 100, 100 parts of water dissolve 201.4 parts of Na 2 B 4 O 7 10 H 2 O, whereas at 10 only 4.6 parts are thus dis- solved. Borax crystallizes in large monoclinic prisms from solutions below 50 ; above that temperature the crystals formed are octahedra of the composition Na 2 B 4 O 7 5 H 2 O. The salt comes in the market in both forms. When borax is heated, it swells up because of loss of water in the form of steam. A clear liquid is finally obtained which solidifies to borax glass. The latter when molten dissolves many metallic oxides, and these solutions have colors charac- teristic of the metals they contain, a fact that is often used in chemical analysis, and in making glazes and enamels for pottery. Borax is used in the laundry for softening water, and to increase the gloss of starch in ironing. It is further employed as a flux in welding and brazing. metals, as a mordant in dyeing fabrics, as an antiseptic in medicine, and as a preservative. It ought not to be used as a preservative for foods. Other Compounds of Boron. Boron hydride BH 3 is a gas formed by the action of magnesium boride upon hydrochloric acid. SILICON AND BORON 319 Boron nitride BN is a white solid formed by the direct union of nitrogen with boron when heated. Water vapor decom- poses it at high temperatures, forming boric acid and ammonia. Boron trifluoride BF 3 is a colorless, pungent gas made by the action of hydrofluoric acid on boron trioxide, or by heating the latter with fluorspar : 3 CaF 2 + 2 B 2 O 3 = Ca 8 B 2 O 6 + 2 BF 3 . Boron trichloride BC1 3 is a colorless liquid of pungent odor. It boils at 18.2 and is decomposed by water into hydrochloric and boric acids : BC1 3 + 3 H 2 O = 3 HC1 + H 3 BO 3 . Boron carbide B 6 C is obtained as an extremely hard solid by heating boron with carbon in the electric furnace. Boron sulphide B 2 S 3 forms small white crystals obtained by heating boron and sulphur together. The sulphide is decom- posed by water with violence, thus : B 2 S 3 + 6 H 2 O = 2 B(OH) 3 + 3 H a S. REVIEW QUESTIONS 1. Compare the physical and chemical properties of silicon and carbon. 2. How may silicon be prepared? Give three methods and the chemical equation illustrating each. 3. What practical use is made of silicon? 4. Discuss the occurrence of silica and silicates in nature ? 5. What is water glass? Give the reaction showing its preparation. What use is made of water glass ? 6. What is quartz? What uses are made of quartz? How is it affected by hydrofluoric acid? Write the equation. 7. Classify the silicic acids. Mention five important silicates that are formed in nature and give their formulas. Mention two different methods by means of which these silicates may be decomposed chemically. 8. Define : dialysis, hydrogel, hydrosol, colloid, crystalloid, and give an example of each. 9. How does water act upon the silicates of the earth's crust? 10. Compare the properties of hydrogen silicide with those of marsh gas. 320 OUTLINES OF CHEMISTRY 11. What is carborundum, and how is it made? 12. What other elements are commonly classed with silicon? Why? What use is made of these elements or their compounds ? 13. How show that silica has acidic properties? Equation. What compound of silicon is analogous to : (a) potassium carbonate, (6) car- bon dioxide? 14. How may a soluble substance be prepared from silicon dioxide? How may a volatile substance be formed from it ? 15. What are the most important forms in which boron is found in nature ? 16. How prepare boric acid from borax. Write the equation. Com- pare this method with that of making : (a) silicic acid, (6) oleic acid, (c) hydrochloric acid, (d) nitric acid, (e) hydrogen sulphide. Do the reactions involved in all these cases come under the same principle of chemical equilibrium? Explain. 17. What are the uses of borax and boric acid? 18. Give two tests for boric acid. 19. How much pure crystallized borax would be required to prepare 100 pounds of boric acid ? ' 20. Many metallic oxides give characteristic colors to borax beads. Explain the chemical changes involved, and state what practical use is made of them. 21. What color does cupric oxide impart to a borax bead when heated (a) in the oxidizing flame, (6) in the reducing flame? Explain. CHAPTER XIX PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH Occurrence and Preparation of Phosphorus. Phosphorus does not occur in the free state in nature because of its great affinity for oxygen. It is widely distributed in the form of phosphates, particularly as calcium phosphate Ca 3 (PO 4 ) 2 , or apatite 3 Ca 3 - (PO 4 ) 2 + Ca(ClF), though it is at times also found as waveliite 2 A1 2 (PO 4 ) 2 + A1 2 (OH) 6 + 9 H 2 O, vivianiteFe 3 (PO 4 ) 2 + 8 H 2 O, and pyromorphite 3 Pb 3 (PO 4 ) 2 + PbCl 2 . In iron ores, phos- phorus occurs as phosphates of iron and calcium, and these are obtained from the slags of blast furnaces. Calcium phosphate is found in many rocks and in all fertile soils. Phosphorus is also an essential ingredient of plant and animal tissues. It is specially necessary in the development of the seeds of plants, hence its importance in the soil, from which the phosphates are taken up by the roots of plants. The ash of bones consists of 80.85 per cent calcium phosphate. In the brain, nerves, blood, albumen, and muscles, phosphorus plays an important role. It occurs here in complex compounds with carbon, hydrogen, nitrogen, oxygen, and sulphur, the nervous tissues being espe cially rich in a compound called lecithine C 44 H 90 NPO 9 . The urine and excreta of animals always contain phosphates. Phosphorus was first prepared in 1669 by the alchemist Brandt, in Hamburg, who evaporated urine and heated the residues mixed with sand to high temperatures. The process was kept a secret, but was soon discovered by Boyle in Eng land and Kunkel in Germany. Gahn showed that calcium phosphate is abundant in bones (1769), and two years later Scheele developed a method for preparing phosphorus from bone ash. Thus calcium sulphate and phosphoric acid ar formed by means of the following reaction : Ca 3 (PO 4 ) 2 + 3 H 2 SO 4 = 3 CaSO 4 + 2 H 3 PO 4 . The calcium sulphate is insoluble, while the phosphoric acic remains in solution and is drained off. This solution is evapo 321 322 OUTLINES OF CHEMISTRY rated to dryness after coke, charcoal, or sawdust have been added, and the mass is then transferred to retorts and heated. In this way water is driven off first, finally carbon monoxide, hydrogen, and phosphorus are formed, the latter being con- densed and collected under water. An older process consists of first forming monocalcium phosphate, which under the name of superphosphate is used as a fertilizer, thus : C ' J 3( po 4 ) 2 + 2 H 2 SO 4 = 2 CaSO 4 + CaH 4 (PO 4 ) 2 . This is then heated to form calcium metaphosphate Ca(PO 3 ) 2 : CaH 4 (P0 4 ) 2 = 2 H 2 + Ca(P0 8 ) 2 . Finally, by mixing the calcium metaphosphate with sand and coke or charcoal, and heating the mixture in earthenware re- torts, of which a number are placed in a furnace, the phospho- rus is obtained and condensed as before. The reaction is : 2 Ca(PO 3 ) 2 + 2 SiO 2 + 10 C Figure 117 shows an arrange- ment of retorts for making phosphorus. By using the electric fur- nace, phosphorus is now being prepared in a simpler way, the process being a continuous one. Calcium phosphate is thoroughly mixed with carbon and silica in pulverized form, and this mixture is heated to a high temperature in the electric furnace. Figure 118 shows the arrangement. The charge is fed in continuously on top by the conveyor, the cal- cium silicate slag is tapped off at the bottom, and the phos- phorus vapors issue from the pipe in the upper part of the furnace and are condensed un- der water. The reaction is : 2 CaSiO + 10 CO + 4 P. 1 I 1 1 1 1 1 1 1 FIG. 117. Ca 3 (PO 4 ) 2 + 3 SiO 2 + 5 C = 3 CaSiO 8 + 5 CO + 2 P. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 323 Thus the silica lays hold of the calcium oxide as it were, form- ing calcium silicate, and the oxygen is taken away from the phosphorus by the carbon at the high temperature, carbon monoxide being formed. Phos- phorus when first condensed as described is contaminated with sand, carbon, and other matter, from which it must be freed. This is accomplished by melting it under water and straining it, also under water, of course, through canvas sacks. It is then redistilled from retorts made of iron, and cast into sticks in glass or tin molds kept in cold water. These sticks are com- FlG> 118 . monly half an inch in diameter and 7.5 inches long, so that nine sticks make approximately a pound of phosphorus. Phosphorus is shipped immersed in water in tin cans. Properties and Allotropic Forms of Phosphorus. The phos- phorus obtained by the methods above described is known as yellow or white phosphorus. It is a pale yellow, translucent, waxlike solid, which in a high state of purity is nearly color- less. In the cold it is brittle, somewhat above room tempera- tures it has the consistency of wax, at 44 it melts under water, and at 269 it boils under atmospheric pressure. Yellow phos- phorus is practically insoluble in water, but it may be dissolved to some extent in alcohol, ether, benzene, and various ethereal oils and fats. It is copiously soluble in carbon disulphide, from which it may be obtained in rhombic dodecahedra of the iso- metric system (Fig. 46) by evaporating off the solvent out of contact with the air. When exposed to the air, phosphorus slowly oxidizes, during which process the oxidation products form fumes, and emit a faint light that is visible in the dark. From the latter phe- nomenon phosphorus derives its name. By such slow oxida- tion phosphorus gradually forms a solution of hypophosphoric 324 OUTLINES OF CHEMISTRY acid which has reducing properties. During this oxidation at room temperatures, ozone and ammonium nitrite are also formed from the air. At about 35 phosphorus catches fire in the air. It must consequently be kept under water. If a little of the solution of phosphorus in carbon disulphide is poured upon filter paper and allowed to evaporate, the finely divided phos- phorus remaining on the paper takes fire spontaneously. Phosphorus is very poisonous, 0.1 gram being a fatal dose for adults. Employees in match factories are apt to suffer from phosphorus poisoning, which manifests itself in enlargement of the liver and necrosis of the jawbones. Phosphorus should always be handled with a forceps and with great care, for phos- phorus burns are dangerous and very slow to heal. When yellow phosphorus is heated from 250 to 300 in closed vessels out of contact of the air, it is gradually converted into red phosphorus, an allotropic form of phosphorus which was discovered by Schrotter in 1845. The reaction is accom- panied with evolution of heat, and is never quite complete, be- ing reversible. When red phosphorus is heated to 260 in a current of carbon dioxide or nitrogen, and the vapors are con- densed under water, the yellow variety is again obtained. Light acting on yellow phosphorus slowly produces some of the red variety, so that ordinary sticks of phosphorus often have a reddish brown outer appearance. Red phosphorus is also called amorphous; it does not emit light in the dark. It may be heated to about 200 in the air without taking fire and con- sequently need not be kept under water. It is insoluble in carbon disulphide and other solvents that dissolve yellow phos- phorus. Moreover, red phosphorus is not poisonous ; and, in general, it is much less active than yellow phosphorus, which contains more energy. The specific gravity of red phosphorus is 2.25. By careful heating, it may be sublimed. The atomic weight of phosphorus is 31. Its valence is either three or five. The vapors of red and yellow phosphorus are identical. The density of the vapor corresponds to the formula P 4 . Uses of Phosphorus, Matches. A small portion of the phos- phorus produced is used for poisoning rats and other vermin. Most of it is used in making matches. The annual production of phosphorus amounts to over 3000 tons. Flint, steel, and tinder were still used to light fires at the beginning of the nine- PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 325 ceenth century. In 1812 the first matches made their appear, ance. They were invented by Chancel, and consisted of sticks dipped in molten sulphur which was afterwards covered with sugar mixed with potassium chlorate. To light such a match its head was brought in contact with concentrated sulphuric acid, which was commonly absorbed in asbestus and kept in a bottle. Thus chloric acid was liberated, and this set the sul- phur and sugar on fire. In 1827 friction or lucifer matches came into use. These had a head consisting of potassium chlorate, antimony sulphide, and glue. They were set on fire by rubbing them vigorously on sandpaper. Phosphorus matqhes appeared in the market in 1832. They contained a little phosphorus in place of the sulphide of antimony, which caused them to ignite more readily. Soon potassium nitrate came into use in matches in place of potassium chlorate, which is apt to cause explosions. At present the oxidizing agents in matches are red lead Pb 3 O 4 , lead peroxide PbO 2 , or manganese peroxide MnO 2 . In making matches the ends of the well-dried sticks are first dipped into paraffine. Afterwards they are dipped into the igniting mixture, consisting of phosphorus stirred into a solution of glue or dextrine, to which the oxidiz- ing agents are added, together with some coloring matter like lamp-black, chalk, or ultramarine to form a paste of proper con- sistency. Safety matches were invented by Bottger in 1848. They had a head of potassium chlorate and antimony trisul- phide like the lucifer matches, but it contained enough glue so that the match ignited with great difficulty on ordinary sur- faces. However, by rubbing these matches on a surface con- taining red phosphorus, which was glued on the box, they would ignite very readily. These safety matches, which are often called Swedish matches, for they were first placed on the market in large quantities in Sweden, are now in common use. The use of the ordinary match that will ignite by friction on any surface is prohibited by law in some countries. The modern safety matches commonly have a head consisting of potassium chlorate, potassium bichromate, powdered glass, and glue or dextrine ; and the friction surface on the box contains antimony trisulphide, red phosphorus, manganese dioxide, and glue. The purpose of the powdered glass in the head is to increase the friction, the heat from which raises 326 OUTLINES OF CHEMISTRY the temperature so that the phosphorus unites vigorously with the oxygen of the oxidizing agents, thus setting the match on fire. Compounds of Phosphorus with Hydrogen. Three compounds of phosphorus and hydrogen are known. They are called phosphines or phosphureted hydrogen. Their composition cor- responds to the formulae : PH 3 , a gas : P 2 H 4 , a liquid ; and P 4 H 2 , a solid. Gaseous phosphine PH 3 is prepared by heating phosphorus in a concentrated solution of caustic potash out of contact with the air. The reaction is : P 4 + 3 KOH + 3 H 2 O = 3 KH 2 PO 2 + PH 8 . In addition, there is always some hydrogen and P 2 H 4 formed. The vapors of the latter are spontaneously inflammable in the air. The experiment is conducted with the apparatus shown in Fig. 119. The small flask is filled half full of caustic potash 0* FIG. 119. solution, and the remaining air is displaced by conducting in illuminating gas or hydrogen through the small tube at the left, which is then closed. On applying heat, phosphine forms and catches fire, forming white smoke rings as it issues from the mouth of the delivery tube, which is kept under warm water PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 327 to prevent its clogging by phosphorus that mi S ht distU1 Ver and solidif y in the end of the tube. If the phos- phine formed is first passed through alcohol or hydrochloric acid, the P 2 H 4 is removed and the gas PH 3 is then no longer spontaneously inflam- mable in the air. In a simpler manner, phosphine may be obtained by treating calcium phosphide with water or dilute FIG. 120. hydrochloric acid (Fig. 120) thus : Ca 8 P 2 + 6 H 2 O = 3 Ca(OH) 2 + 2 PH 3 , or Ca 3 P 2 + 6 HC1 = 3 CaCl a + 2 PH 8 . In these reactions, smaller amounts of the solid and liquid hydrides of phosphorus, P 4 H 2 and P 2 H 4 , are also obtained by secondary reactions. Phosphides of magnesium, zinc, and iron similarly yield phosphine with hydrochloric acid. By heating phosphorous or hypophosphorous acid, phosphine is produced, thus : 4 H 8 P0 8 = 3 H 3 P0 4 + PH 3 , or phosphorous acid phosphoric acid 2H 3 PO 2 = H 3 PO 4 +PH 3 . hypophosphorous acid phosphoric acid When phosphoniurn iodide PH 4 I is treated with caustic alka- lies, phosphine is formed, thus : PH J + NaOH = Nal + H 2 O + PH 3 . Water also decomposes phosphonium iodide : PH 4 I + H 2 = HI + H 2 + PH 3 . Gaseous phosphine is colorless. It boils at 85 and solidi- fies at 133. The gas has the odor of rotten fish and is very poisonous. Heated to about 100 in the air it burns and forms water and phosphoric acid. Phosphine is but slightly soluble in water. Alcohol dissolves it more copiously. With the hydrohalogens phosphine forms phosphonium com- pounds, which are analogous to ammonium salts. Phosphonium iodide, the best known of the phosphonium compounds, is pre- pared by the following reaction : PH 3 4- HI = PH 4 I ; which is analogous to 328 OUTLINES OF CHEMISTRY Phosphonium iodide is a very unstable, colorless, crystalline salt, which is decomposed by water into phosphine and hydri- odic acid, as stated above. Oxygen acids do not form phos- phonium salts with phosphine. Liquid phosphine P 2 tT 4 * s analogous to hydrazine N 2 H 4 . It is a colorless liquid of specific gravity 1.01 at 15. It boils at 57 and is insoluble in water. Solid phosphine P 4 H 2 is a yellow, flocculent powder which is devoid of odor and taste. It does not dissolve in water. At about 160 it takes fire in the air. Compounds of Phosphorus with the Halogens. Phosphorus forms compounds with all of the halogens. These have the general formula? PX 3 and PX 5 . The chlorides PC1 3 and PC1 5 are the most important. Phosphorus trichloride PC1 3 is formed when chlorine is passed upon phosphorus in a retort. The action proceeds readily with liberation of heat, the product being a colorless liquid of pungent odor. In the pure state phosphorus trichlo- ride boils at 76 and solidifies at 115. Its specific gravity is 1.613 at 0. Water decomposes it : PC1 3 + 3 H 2 O = 3 HC1 + P(OH) 3 . Phosphorus pentachloride PC1 5 is formed by treating phos- phorus trichloride with chlorine, or by passing an excess of chlorine upon phosphorus in a retort : PC1 3 + Cl a = PC1 5 , or The product is a light yellow, finely crystalline solid which can- not bo melted under atmospheric pressure, for the temperature at which its vapor tension equals atmospheric pressure lies be- low the melting point of the compound. Under the pressure of its own vapor in a sealed tube, phosphorus pentachloride may be melted at 148. When heated, phosphorus pentachlo- ride decomposes into phosphorus trichloride and chlorine : PC],5PC],+ CI r At 300 this dissociation is nearly complete. The action is reversible, as indicated. It is quite similar to the dissociation of ammonium chloride by means of heat: f HC1. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 329 With water, phosphorus pentachloride forms hydrochloric acid and phosphorus oxychloride : PC1 5 + H 2 = 2 HC1 + POC1 3 . The latter is a colorless liquid of specific gravity 1.712 at 0. It boils at 107.5 and melts at 1.8. On further treatment with water, the oxychloride also decomposes, yielding hydrochloric acid and phosphoric acid : POC1 3 + 3 H 2 O = 3 HC1 + H 8 PO 4 . Phosphorus trifluoride PF 3 is a colorless gas. It boils at - 95 and congeals at 160. The pentafluoride PF 6 melts at 83 and boils at 75. These compounds are decom- posed by water like the analogous chlorides, but more slowly. Phosphorus oxyfluoride POF 3 is a gas which may be liquefied at - 50. Phosphorus tribromide PBr 8 is a colorless liquid boiling at 172. Its specific gravity is 2.925 at 0. Phosphorus penta- bromide PBr 5 forms yellow crystals, which on heating disso- ciate into bromine and phosphorus tribromide. Phosphorus triodide PI 8 forms dark red, prismatic crystals melting at 61. Phosphorus pentaiodide is not known, but a diphosphorus tetraiodide P 2 I 4 is known. It forms orange-yellow crystals which melt at 110. On treatment with water, both the bromides and iodides of phosphorus are decomposed into the hydrohalogen acids and oxygen acids of phosphorus. From phosphorus pentabro- mide, phosphorus oxybromide POBr 3 may be obtained in a manner analogous to the formation of POC1 3 . The treatment of phosphorus tribromide or triodide with water affords excellent methods for making pure hydrobromic and hydriodic acids, as already stated. Oxides and Acids of Phosphorus. The following oxides of phosphorus are well known : phosphorus trioxide P 2 O 3 ; phosphorus tetroxide P a O 4 ; and phosphorus pentoxide P 2 O 5 . Of these the latter is the most important by far. The trioxide is a white crystalline solid melting at 22.5. It is obtained together with the pentoxide by burning phosphorus in an in- sufficient? amount of oxygen. The tetroxide is a white solid formed, together with red phosphorus, by heating the trioxide 330 OUTLINES OF CHEMISTRY in a sealed tube to 440. Phosphorus pentoxide P 2 O 5 is formed when phosphorus is burned in the air or in oxygen. It is a light white powder which unites with water with great avidity, forming metaphosphoric acid, thus: Phosphorus pentoxide is the best drying agent known. Its action on water is accompanied with evolution of much heat and a hissing noise resembling that accompanying the quench- ing of hot iron. In union with different amounts of water, phosphorus pen- toxide forms three acids, thus . P 2 O 6 + H 2 O = 2 HPO 3 (metaphosphoric acid), P 2 O 5 + 2 H 2 O = H 4 P 2 O 7 (pyrophosphoric acid), P 2 O 5 + 3 H 2 O = 2 H 3 PO 4 (orthophosphoric acid). By union with two or six molecules of water phosphorus tri- oxide forms two acids, thus : 2 P 2 O 3 + 2 H 2 O = 4 HPO 2 (metaphosphorous acid), 2 P 2 O 3 + 6 H 2 O = 4 H 3 PO 3 (phosphorous acid). There are also known hypophosphoric acid H 4 P 2 O 6 and hypo- phosphorous acid H 3 PO 2 . The former is prepared by allowing sticks of phosphorus to oxidize slowly in contact with moist air, under which conditions phosphoric and phosphorous acids are also formed to some extent. The acid is tetrabasic, and consequently is able to form four kinds of salts by successive replacement of the hydrogen atoms. Hypophosphorous acid H 3 PO 2 may be liberated from its barium salt by action of sul- phuric acid, thus : 8 P + 3 Ba(OH) 2 + 6 H 2 O = 2 PH 3 + 3 Ba(H 2 PO 2 ) 2 , and Ba(H 2 PO 2 ) 2 + H 2 SO 4 = BaSO 4 + 2 H 3 PO 2 . It is a monobasic acid, forming crystals that melt at 17.4. It is a strong reducing agent. On being heated, it yields phos- phine and phosphoric acid. Orthophosphoric Acid. This compound is also called simply phosphoric acid. Its composition is expressed by the formula H 3 PO 4 . It may be considered as derived from the hypothet- ical pentahydroxide of phosphorus P(OH) 5 by loss, of one molecule of water. Pure phosphoric acid is prepared by action PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 33l of phosphorus pentoxide on water or by the oxidation of phos- phorus by means of nitric acid. Phosphoric acid is also made by the action of sulphuric acid upon calcium phosphate. The calcium sulphate formed simultaneously, being insoluble, is readily removed and the clear solution containing the phos- phoric acid is then evaporated. It commonly still contains some calcium salts which may be precipitated .by means of alcohol. Solutions of pure phosphoric acid may be evaporated to a thick, colorless sirup of specific gravity 1.88, from which upon cooling a crystalline mass is obtained, which melts at 42 c The crystals are deliquescent and dissolve in water with great readiness. The solutions are strongly acidic in character, The acid is not poisonous. Phosphoric acid is tribasic and con- sequently is able to form three classes of salts, the primary, sec- ondary, and tertiary phosphates, for instance : - H 3 P0 4 + KOH = KH 2 P0 4 4- H 2 O. H 3 P0 4 + 2 KOH = K 2 HP0 4 + 2 H 2 O. H 3 PO 4 + 3 KOH = K 3 PO 4 + 3 H 2 O. The tertiary phosphates are the normal or neutral salts ; whereas the secondary and primary salts still contain one and two hydro* gen atoms respectively in the molecule. The hydrogen atoms need not all be replaced by the same metal or radical. Thus we have sodium ammonium hydrogen phosphate NaNH 4 HPO 4 , which is also known as microcosmic salt. Magnesium ammonium phosphate MgNH 4 PO 4 forms white insoluble crystals ; it is of importance in analytical chemistry. Solutions of the secondary salts have an alkaline reaction, being to some extent deconir posed by hydrolysis. The tertiary salts are much more hy- drolyzed by water; indeed they are stable only as solids, and are obtained by evaporating the acid to dryness with the proper amount of alkali. These salts are not decomposed by heat, whereas both the secondary and primary phosphates lose water on being heated ; so, for instance : P0 4 ^:Na 4 P 2 7 + H 2 O ; and sodium pyrophosphate NaH 2 PO 4 ^NaPO 3 + H 2 O. sodium inetaphosphate 332 OUTLINES OF CHKMIS'LKY Thus secondary phosphates yield pyrophosphates, and primary phosphates yield metaphosphates, on heating. Conversely, on treatment with water the pyrophosphates gradually pass back into secondary phosphates, and the metaphosphates into pri- mary phosphates. Microcosmic salt and magnesium ammonium phosphate lose ammonia as well as water on being heated, thus: NaNH 4 HPO 4 = H 2 O + NH 3 + NaPO 3 . 2 MgNH 4 P0 4 = H 2 + 2 NH 3 + Mg 2 P 2 O 7 . Pyrophosphoric acid H 4 P 2 O 7 is formed by heating phosphoric acid to about 250 till a sample neutralized with ammonia and tested with silver nitrate solution yields a white precipitate. The white precipitate is silver pyrophosphate Ag 4 P 2 O 7 , whereas the phosphate of silver Ag 3 PO 4 is yellow. The formation of pyrophosphoric acid takes place thus : 2H 3 P0 4 =H 4 P 2 7 + H 2 0. The aqueous solutions of pyrophosphoric acid are fairly stable, the acid passing over into orthophosphoric acid but slowly. The presence of sulphuric or nitric acids hastens the change. Though the molecule of pyrophosphoric acid contains four hydrogen atoms, but two kinds of pyrophosphates are known. These correspond to the types K 4 P 2 O 7 and K 2 H 2 P 2 O 7 . By the color of the silver salt, pyrophosphoric acid is readily distinguished from orthophosphoric acid. From metaphos- phoric acid, pyrophosphoric acid is distinguished by the fact that it does not coagulate albumen like the former. Metaphosphoric acid HPO 3 is made by heating phosphoric acid to 400: H 3 P0 4 =H 2 + HP0 3 ; or by treating phosphorus pentoxide with water; P 2 6 + H 2 = 2HP0 8 ; or by heating ammonium phosphate; (NH 4 ) 2 HP0 4 = 2 NH 3 + H 2 + HPO 3 . The acid is a glassy, semitransparent mass which is also called glacial phosphoric acid. In contact with water, it slowly passes over into phosphoric acid, the action being hastened by boiling. The acid is monobasic and is analogous to nitric, chloric, and PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 333 bromic acids. Solutions of glacial phosphoric acid coagulate albumen and give white precipitates with the chlorides of barium or calcium, which behavior is different from that of solutions of pyrophosphoric acid. Phosphorous acid H 3 PO 3 forms as one of the products of the slow oxidation of phosphorus in moist air. It is best prepared by treating phosphorus trichloride with water and driving off the hydrochloric acid formed simultaneously, by heating to 180. The acid forms very hygroscopic crystals that melt at 70. On heating, it decomposes into phosphoric acid and phosphine : 4 H 3 PO 3 = 3 H 3 PO 4 + PH 3 . At the high temperature at which the reaction takes place, the phosphoric acid formed passes over into metaphosphoric acid, and the phosphine burns with a green flame. Though phos- phorous acid has three hydrogen atoms in the molecule, it is only dibasic. Its salts correspond to the type Na 2 HPO 3 , the third hydrogen atom not being replaceable by a metal. Formulae of the Acids of Phosphorus. The following struc- tural formulae of the oxy-acids of phosphorus will serve to impress their relationships further : O H orthophosphoric \M3-H acid. p/O-H r \O-H , . pyrophosphoric metaphosphoric \ Q _ H acid. O H phosphorous O-H acid. \ \ P-O-H O hypophosphoric / p //O-H acid. \ O - H hypophosphorous acid. )-H 334 OUTLINES OF CHEMISTRY It will be seen that the dibasic character of phosphorous acid is expressed by connecting the non-replaceable hydrogen atom directly with the phosphorus, Similarly the monobasic char- acter of hypophosphorous acid is indicated by connecting the two non-replaceable hydrogen atoms directly with phosphorus. Compounds of Phosphorus with Sulphur. With sulphur, phosphorus unites directly, forming a series of compounds: P 4 S 3 , P 2 S 3 , P 8 S 6 , and P 2 S 5 . The action of yellow phosphorus upon hot sulphur is violent; the sulphides are consequently made by using red phosphorus. Phosphorus pentasulphide P 2 S 6 forms yellow crystals which melt at 275. The liquid boils at 518. With potassium sulphide it forms potassium sulphophosphate : - P 2 S 5 + 3K 2 S = 2K 3 PS 4 . With phosphorus pentachloride phosphorus sulphochloride PSCL results : P 2 S 5 -h 3 PC1 5 = 5 PSC1 3 . The latter compound is a colorless liquid of specific gravity 1.168" ats?0. It boils at 125, and decomposes upon treatment with water : PSC1 3 + 4 H 2 O = 3 HC1 + H 2 S + H 8 PO 4 . Occurrence, Preparation, and Properties of Arsenic. Arsenic is very widely distributed in nature in minute quantities. It rarely ocpurs in the uncombined state, being found in larger quantities in combination with sulphur, as in realgar As 2 S 2 and orpiment As 2 S 3 . It is also found combined with oxygen, as in arsenolite As 2 O 3 , and with iron and sulphur and cobalt and sulphur, as in arsenical pyrites or mispickel FeAsS and cobaltite CoAsS. Arsenic is commonly prepared by heating mispickel or by reducing arsenolite with carbon. The reactions are : FeAsS=FeS + As. 2As 2 O 3 +6C= As 4 +6CO. Arsenic is volatile ; it sublimes, and is readily condensed. Arsenic is steel-gray in color, has a bright metallic luster, and is very brittle. Its specific gravity is 5.73 at 15. On heating, it volatilizes without melting ; but under pressure it may be melted at about 480. At 450 its vapor tension equals PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 335 atmospheric pressure. Heated in the air, it burns, the fumes having a garlic-like odor and the flame a pale lavender color ; these are characteristic of arsenic. Between 5t)0 and 860 the vapor of arsenic is about 150 times as heavy as hydrogen. Hence the molecular 'weight is approximately 300; and since the atomic weight of arsenic is 75 as determined from the analysis of the chloride, the molecular formula of arsenic is As 4 . Between 1600 and 1700, Victor Meyer found the vapor of arsenic to be only 75 times as heavy as hydrogen, which leads to the molecular formula As 2 . The valence of arsenic is either three or five, and the formulae of its compounds are consequently analogous to those of nitrogen and phosphorus. Arsenic burns to As 2 O 3 in the air when heated to 180. It combines directly with many elements like chlorine, bromine, sulphur, and some of the metals. When boiled with nitric acid or aqua regia, arsenic is oxidized to arsenic acid H^AsO 4 . Besides the metallic form, of arsenic above described, this element may be obtained as yellow crystals by rapidly cooling its vapor. The crystals resemble ordinary phosphorus in that they dissolve in carbon bisulphide. Arsenic itself does not act as a poison, for it is not taken up by the animal system. Its insoluble sulphides also are not especially toxic in character. However, all other compounds of arsenic, notably arsine AsH 3 , arsenious, oxide As 2 O 3 , halogen compounds, and salts of arsenious and arsenic acids are very poisonous. From 0.1 to 0.4 gram of arsenious oxide is sufficient to cause death. The antidote for arsenic is freshly precipitated ferric hydroxide. Arsine, Arseniureted Hydrogen, AsH 3 . Arsine is a colorless gas. It was discovered by Scheele in 1755. It melts at 113.5 and boils at 55. It is analogous to ammonia NH 3 and phosphine PH 3 . It is commonly prepared (1) by the action of hydrochloric or sulphuric acid upon the arsenide of zinc or sodium, or (2) by introducing compounds of arsenic in a flask containing zinc and hydrochloric or sulphuric acid. -y The reactions involved in these processes are typified by the follow- ing equations : - (1) Zn 3 As 2 4- 6 HC1 = 3 ZnCl 2 + 2 AsH 3 . AsNa 3 + 3 H 2 S0 4 = 3 NaHSO, +. AsH 8 . (2) As 2 O 3 + 12 H = 3 H 2 O + 2 AsH 8 . 336 OUTLINES OF CHEMISTRY The odor of arsine is very disagreeable, resembling that of garlic. Arsine is extremely poisonous and great care must con- sequently be exercised in experimenting with it. Arsine does not unite with water or with acids ; it thus exhibits much less basic properties than ammonia or phosphine. Ignited in the air, arsine burns with a pale lavender flame, forming water and arsenious oxide : 2 AsH 3 -f 3 O 2 = 3 H 2 O + As 2 O 3 . On being heated, the gas readily dissociates into arsenic and hydrogen : So when dry arsine is passed through a tube heated to dull redness, the reaction just given takes place, the arsenic con- densing in form of a metallic mirror in the colder parts of the tube. Since solutions of all arsenic compounds when intro- duced into a flask containing zinc and hydrochloric or sulphuric acid yield arsine, a simple and very efficient method of testing arsenic, known as Marsh's test, has been devised. The appara- tus is shown in Fig. 121. Pure zinc and hydrochloric acid are FIG. 121. introduced into the flask. The calcium chloride in the tube serves to dry the gases evolved. After all air has been ex- pelled, the hydrogen is lighted and the solution to be tested for arsenic is poured down the funnel tube. If arsenic is present, the flame will acquire the characteristic pale lavender color, and dark spots of metallic arsenic will be deposited upon a white porcelain dish held in the flame. If the tube, which should be of hard glass, is heated as shown, a mirror of metallic PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 337 arsenic will deposit on the sides of the tube just beyond th flame. Both the mirror and the spots are soluble in sodium hypochlorite or bleaching powder solution. It is to be noted that compounds of antimony under like treatment yield similar spots and mirrors ; these are, however, not soluble in hypo- chlorites. Moreover, the arsenical mirror is more volatile than that of antimony. The former may be converted into yellow sulphide of arsenic and the latter into red sulphide of antimony by means of hydrogen sulphide. When conducted into a solution of silver nitrate, arsine pre- cipitates metallic silver, thus : 2 AsH 3 + 12 AgNO 3 + 3 H 2 O = As 2 O 3 + 12 HNO 3 + 12 Ag. Since the corresponding antimony hydride, stibine SbH 3 , does not reduce silver nitrate solutions thus, this reaction may be used to distinguish between arsine and stibine. Compounds of Arsenic with the Halogens. Of these com- pounds arsenic trichloride AsCl 3 is the most important. There are also known : the trifluoride AsF 3 , boiling at 63 and melting at 8.5; the tribromide AsBr 3 , -melting at 31 and boiling at 221 ; the triodide AsT 3 , melting at 140, as well as iodides of the formulae AsI 2 and AsI 5 . Arsenic trichloride is formed by conducting chlorine upon powdered arsenic contained in a retort, or by the action of hydrochloric acid upon arsenic trioxide. It is a colorless, fum- ing liquid of specific gravity 2.205 at 0. It boils at 129 and solidifies to a crystalline mass at 18. It is very poisonous. Water decomposes it : 2 AsCl 3 + 3 H 2 O = As 2 O 3 + 6 HC1. By addition of concentrated hydrochloric acid, the hydrolysis may be reversed. It will be recalled that this cannot be done in the case of the analogous chloride of phosphorus. Oxides and Oxy-acids of Arsenic. Two oxides, the trioxide As 2 O 8 and the pentoxide As 2 O 5 , are known ; and the corre- sponding acids, arsenious acid H 3 AsO 3 and arsenic acid H 3 AsO 4 , are of importance. Arsenic trioxide As 2 O 3 , also called "white, arsenic" or com- monly simply " arsenic'' is the commonest, and by far the most important, of all the compounds of arsenic. It is found in nature CALIF Oh NIA COLLEGE 338 .OUTLINES OF CHEMISTRY arid is formed when arsenic burns in the air or in oxygen. Ar- senic trioxide is manufactured on a commercial scale by roasting arsenical pyrites in the air. Jii this process iron oxide remains as a non-volatile residue, sulphur dioxide escapes, and the arse- nious oxide condenses as a white powder upon the brick walls of the 'chambers. It is purified by resublimation. Each year approximately 4000 tons of arsenious oxide are produced in the United States. From two to three times this amount is annually produced in Europe. On heating arsenic trioxide, it gradually forms an ; amorphous glassy mass, which after a time becomes white, crystalline, and opaque. Below 200 the crystals formed are octahedra of the regular system, whereas above that tem- perature crystallization in monoclinic forms takes place. At 800 P the 'vapor density of arsenious oxide corresponds to the formula (As 2 O 3 ) 2 , whereas at about 1800 the density of the gas leads to the simple formula As 2 O 3 , the double molecules having been dissociated. Arsenic trioxide is readily reduced to arsenic by heating it with carbon, or cyanide of potassium. Its conversion to arsine has already been mentioned. In water rt dissolves but slightly. Hydrochloric acid dissolves it, form- ing arsenic trichloride. The trioxide has a sweetish, disagreeable taste. It is a strong poison. It is used as rat poison, also in taxidermy, in calico printing, in the manufacture of certain kinds of glass, in the preparation of many other compounds of arsenic, and in medi- c'ine. Freshly precipitated ferric hydroxide forms an insoluble compound with arsenious oxide and is consequently used as an antidote in cases of poisoning. Arsenious acid H 3 AsO 3 has not been isolated. It probably exists in the aqueous solutions of arsenious oxide. Its salts, the arsenites, are known. Among these may here be mentioned silver arsenite Ag 3 AsO 3 and copper hydrogen arsenite, or Scheele's green, CuHAsO 3 . Salts of meta-arsenious acid HAsO 2 are also" known, like KAsO 2 and Pb(AsO 2 ) 2 . Paris green, also called Schweinfurt green, is a double salt of cupric arsenite &nd cupric. acetate Cu 3 As 2 O 6 -Cu(C 2 H 3 O 2 ) 2 .: It is used as a poison for potato bugs and other insects. Arsenic acid H 3 AsO 4 is readily produced by oxidation of arsenious acid. Scheele prepared arsenic acid in 1776 by passing chlorine into arsenic trioxide suspended in PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 339 water ; nitric acid or a mixture of nitric and hydrochloric acids serves equally well. The reaction in the former case is : As 2 3 + 2 C1 2 + 5 H 2 = 2 H 3 AsO 4 + 4 HC1. The acid forms rhombic, deliquescent prisms or plates of the composition 2 H 3 AsO 4 + H 2 O. At 100 the water of crystalli- zation passes off. At about 180 the acid loses water, passing over into pyroarsenic acid H 4 As 2 O 7 , which on being heated still further again loses water, forming meta-arsenic acid HAsO 3 . So far then the behavior is entirely similar to that of phos- phoric acid, though in contact with water pyro- and meta- arsenic acids at once form arsenic acid. On further ignition of meta-arsenic acid, water is again split off and arsenic pentoxide As 2 O 5 is formed, thus : 2 HAsO 3 = H 2 O + As 2 O 5 . It will be recalled that metaphosphoric acid cannot thus be decomposed into P 2 O 5 and water. Furthermore, phosphorus pentoxide is very stable when heated, whereas arsenic pentoxide decomposes upon ignition into arsenic trioxide and oxygen : As 2 O 5 = As 2 O 3 + O 2 . The salts of arsenic acid are quite analogous to those of phosphoric acid. Thus, there are primary, secondary, and tertiary arsenates, also pyroarsenates and meta-ar senates. In contact with water, however, all the salts form orthoarsenates at once. Sulphides of Arsenic. Three sulphides of arsenic are known, namely : the disulphide As 2 S 2 , the trisulphide As 2 S 3 , and the pentasulphide As 2 S 5 . Arsenic disulphide A s 2 S 2 occurs in nature as realgar, in red monoclinic prisms. It is also manufactured by fusing sulphur and arsenic together. Thus made, it forms a dark red, glassy substance, which in pulverized condition is sometimes used as a pigment in paints. A mixture of 1 part arsenic disulphide, 12 parts saltpeter, and 3.5 parts sulphur when ignited makes white Bengal fire. Arsenic trisulphide As 2 S 3 occurs in nature in short rhombic prisms as orpiment. It was formerly used as a pigment. It is readily obtained as a lemon-yellow precipitate by passing 340 OUTLINES OF CHEMISTRY hydrogen sulphide into a solution of arsenic trioxide in hydro chloric acid : 2 AsCl 3 + 3 H 2 S = As 2 S 3 + 6 HC1. On heating the precipitate with concentrated hydrochloric acid, it may be redissolved ; that is, the reaction just given may be reversed. Arsenic trisulphide may also be obtained by fusing sulphur and arsenic together in the right proportions. In am- monium sulphide, arsenic trisulphide is soluble, forming ammo- nium sulpharsenite (NH 4 ) 3 AsS 3 : As 2 S 3 + 3 (NH 4 ) 2 S = 2 (NH 4 ) 3 AsS 8 . In solution of yellow ammonium sulphide, that is, in ammonium sulphide containing an excess of sulphur, arsenic trisulphide dissolves as ammonium sulpharsenate (NH 4 ) 3 AsS 4 : AsjjSg 4- 3 (NH 4 ) 2 S + 2 S = 2 (NH 4 ) 3 AsS 4 . On treatment with hydrochloric acid the sulpharsenites and sulph- arsenates are decomposed : 2 (NH 4 ) 3 AsS 3 + 6 HC1 = As 2 S 3 + 6 NH 4 C1 + 3 H 2 S. 2 (NH 4 ) 3 AsS 4 + 6 HC1 = As 2 S 6 + 6 NH 4 C1 4- 3 H 2 S. Arsenic pentasulphide As 2 S 6 , made by means of the reaction just given or by melting together sulphur and arsenic in proper proportions, is a yellow solid which may be sublimed when heated out of contact with the air. Occurrence, Preparation, and Properties of Antimony. Anti- mony (stibium) is sometimes, though rarely, found in nature in the uncombined state. When thus found, it occurs in rhombohedral crystals. The mineral stibnite Sb 2 S 3 , found in Hungary and Japan, is the chief source of antimony, though the latter also occurs combined with sulphur in many native sulphides of lead, copper, silver, iron, and arsenic. Native oxide of antimony, senarmontite Sb 2 O 3 , forming white octa- hedra of the regular system, is also known. Stibnite was known in ancient times. The Chaldeans manufactured vari- ous articles out of metallic antimony, and the alchemists frequently used the metal. Antimony is prepared by heating stibnite with iron, thus : Sb 2 S 8 4- 3 Fe = 3 FeS + 2 Sb. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 341 It is also made by roasting stibnite in the air, and reducing the tetroxide thus formed, by means of carbon. The reactions are as follows : Sb 2 S 3 + 5 2 = 3 S0 2 + Sb 2 4 . Sb 2 4 + 4 C = 4 CO + 2 Sb. To free the antimony ' regulus " so obtained from iron, lead, copper, etc., it is fused with a little sulphur or saltpeter. Thus the impurities are converted to sulphides or oxides, which float on top and can be removed. Antimony free from arsenic and other metals may be obtained by reducing pure sodium metantimoniate NaSbO 3 . Antimony is a hard, brittle, silvery-white metal having a high metallic luster. It can readily be ground to powder. At 625 it melts, and on cooling it forms rhombohedral crystals. Its boiling point is approximately 1400, and its specific gravity is 6.75. In the air it remains practically unchanged, but when strongly heated it burns with a bluish white flame to Sb 2 O 3 or Sb 2 O 4 . Introduced into an atmosphere of chlorine, it takes tire and burns to SbCl 5 . It dissolves in hot concentrated sulphuric acid, also in aqua regia, but nitric acid converts it into Sb 2 O 3 or antimonic acid H 3 SbO 4 . Hydrochloric acid acts slowly on antimony, liberating hydrogen. The latter gas is also formed by the action of steam on antimony at high temperatures. The atomic weight of antimony is 120.2. The vapor density leads to a molecular weight of approximately 290, which rep- resents a formula lying between Sb 2 and Sb 3 . The valence of antimony is either three or five. Its compounds consequently have formulae analogous to those of nitrogen, phosphorus, and arsenic. The latter is a clcise relative of antimony. Metallic antimony is much used in alloys, particularly in type metal and britannia metal. Type metal consists of approxi- mately 25 per cent antimony, 25 per cent tin, and 50 per cent lead. The presence of antimony in alloys makes them hard. Furthermore, antimony expands as it congeals (resembling water in this behavior) and consequently fills molds perfectly, thus yielding sharply defined castings. Hydrogen Antimonide, Stibine, SbH 3 . This compound is analogous to ammonia, phosphine, and arsine. It is quite simi- lar to the latter and is prepared by similar methods. So, .for OUTLINES OF CHEMISTRY instance, by treating an alloy of magnesium and antimony or zinc and antimony with dilute hydrochloric or sulphuric acid, stibine is formed. Again, by introducing a solution of any antimony compound into a flask in which zinc is being acted upon by hydrochloric or sulphuric acid, stibine results, which in this case is mixed with hydrogen. Stibine is a colorless gas of peculiar odor, reminding one somewhat of that of hydrogen sulphide. The odor is distinctly different from that of arsine. Stibine melts at 88 and boils at 17. The gas readily dissociates into antimony and hydro- gen, thus : 2 SbH 3 = 2 Sb + 3 H 2 . The change begins at 150. Even when diluted with hydro- gen, stibine is largely decomposed when passed through a tube heated to 150, yielding a deposit of antimony in the form of a mirror, which is insoluble in hypochlorites. Thus, in the apparatus used for making Marsh's test for arsenic, antimony compounds would yield a similar mirror; but the latter is readily distinguished from arsenic by the method described under arsine. The dissociation of stibine is practically com- plete at 200, at which temperature arsine remains unchanged. Stibine is moderately poisonous. Water dissolves about four times its own volume of the gas at room temperature. In the air or in oxygen, when ignited, stibine burns with a bluish white flame, forming water and Sb 2 O 3 .- Conducted into a silver nitrate solution, stibine is decomposed, the antimony being precipitated as silver antimonide SbAg 3 . When pure or when diluted with hydrogen, stibine may be kept unchanged ; but the presence of even small amounts of oxygen in the gas leads to the deposition of some of the antimony. Compounds of Antimony with the Halogens. Of these, anti- mony trichloride SbCl 3 and antimony pentachloride SbCl 5 are of most importance. Antimony trichloride is formed by the action of chlorine on antimony or of hydrochloric acid on antimony sulphide : 2 Sb + 3 C1 2 = 2 SbCl 3 . Sb 2 S 3 + 6 HC1 = 3 H 2 S + 2 SbCl 3 . The antimony trichloride is purified by distillation. It is a colorless crystalline mass which at ordinary temperatures is PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 343 soft, reminding one of the consistency of butter, hence it goes by the name of butter of antimony. It melts at 73 and boils at 223. At 26 its specific gravity is 3.064. Its vapor is 229 times as heavy as hydrogen, which fact leads to the formula SbCl 3 . It is deliquescent and has caustic properties. Antimony trichloride is used as a mordant, also in medicine and in burnishing metals, notably gun barrels, to which it imparts a brown hue. Antimony trichloride may be dissolved in water containing hydrochloric acid. But when treated with water alone, antimony trichloride is decomposed into hydro- chloric acid and insoluble oxychlorides, the composition of which varies according to the temperature and relative amount of water used. Two oxychlorides of antimony, SbOCl and (SbOCl) 2 Sb 2 O 3 , are well known as white crystalline powders. They are formed thus : (1) SbCl 3 + H 2 O = SbOCl -f 2 HC1. (2) 4 SbCl 3 + 5 H 2 O = (SbOCl) 2 - Sb a O 8 + 10 HC1. The second reaction takes place in hot solutions. The com- pound (SbOCl) 2 'Sb 2 O 3 , or Sb 4 O 5 Cl 2 , was used by the Italian physician Victor Algarotus, and is consequently known as the powder of algaroth. Antimony pentachloride is prepared by burning antimony in an excess of chlorine or by conducting chlorine upon antimony trichloride. It is a fuming liquid of yellow color. At 6 its crystals melt. It can only be distilled in a partial vac- uum, for on heating it readily dissociates into chlorine and the trichloride. With water it forms crystalline hydrates, SbCl 5 -H 2 O and SbCl 6 -4H 2 O. Antimony pentachloride is decomposed by hot water. It readily gives off part of its chlorine, and is consequently used in organic chemistry in chlorinating substances. It will be observed that while antimony pentachloride forms crystalline hydrates with water, the latter decomposes the chlorides of phosphorus at once. Antimony trifluoride SbF 3 forms deliquescent rhombic crystals that are not decomposed by cold water. With ammonium sul- phate it forms a compound that is used as a mordant. Antimony pentafluoride SbF 5 is an amorphous gummy mass, It readily enters into the formation of double salts. 344 OUTLINES OF CHEMISTRY Antimony tribromide SbBr g forms white rhombic crystals that melt at 94. The salt boils at 275, and is decomposed by water. Antimony triiodide SM 3 forms three different varieties ot crystals. The common red crystals melt at 171. The boiling point is 430. Antimony pentiodide SbI 5 is a dark brown, crystalline mass of melting point 79. It is unstable. Oxides and Oxy-acids of Antimony. There are three oxides of antimony: antimony trioxide Sb 2 O 3 , antimony tetroxide Sb 2 O 4 , and antimony pentoxide Sb 2 O 6 . The trioxide acts mainly as a base, though toward very strong bases, like caustic potash and soda, it is also able to act as an acid. The tetroxide exhibits neither acid nor basic properties, whereas the pentoxide acts solely in an acid-forming capacity. Antimony trioxide is found in nature as senarmontite. It is formed by burning antimony in the air or by oxidizing the metal with nitric acid. The oxide is white and may be sub- limed. It crystallizes in octahedra or rhombic prisms, being dimorphous. At 1560 the density of its vapor corresponds to the formula Sb 4 O 6 , nevertheless it is commonly called the tri- oxide. It is possible that at higher temperatures it would dissociate into Sb 2 O 3 like the corresponding oxide of arsenic. In water and nitric or sulphuric acid, antimony trioxide is practically insoluble, while in hydrochloric or tartaric acid, or in acid potassium tartrate or caustic alkalies, it dissolves, thus : Sb 2 3 + 6 HC1 = 2 SbCl 3 + 3 H 2 0. Sb 2 O 3 + 2 KOH = 2 KSbO 2 + H 2 O. Sb 2 3 + 2 (C 4 H 4 6 )HK = 2 (C 4 H 4 O 6 )SbO . K + H 2 O. The salt KSbO 2 is potassium metantimonite. It is plainly a salt of metantimonious acid HSbO 2 , which may be considered as derived from antimonious acid H 3 SbO 3 by loss of a molecule of water. The salt (C 4 H 4 O 6 ) SbO K is potassium antimonyl tartrate or tartar emetic. It contains the univa-lent antimonyl group, Sb = O, which is frequently found in other antimony salts. Tartar emetic has been known for a long time. The salt crystallizes with half a molecule of crystal water, a part of which escapes on exposure to the air. The salt is still some- PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 345 times used in medicine. Antimony salts were formerly fre- quently prescribed by physicians. These compounds gained in prominence through the work of Basil Valentine, who in the fifteenth century, published his book on " The Triumphal Chariot of Antimonium." The compounds (C 4 H 4 O 6 ) AsO - K potassium arsenyl tartrate and (C 4 H 4 O 6 ) BO K potassium boryl tartrate are analogous to tartar emetic. On treating tartar emetic with dilute sulphuric acid, the hydrate H 3 SbO 3 separates out as a precipitate, which, how- ever, loses water and forms metantimoiiious acid HSbO 2 , i.e. SbO- OH. The basic properties of antimony are shown in its salts, in which either Sb(OH) 3 or SbO OH act as bases. Thus, there are known antimony nitrate Sb(NO 3 ) 3 , antimony sulphate Sb 2 (SO 4 ) 3 , and the halogen salts like SbCl 8 ; further, when these salts are acted upon by water, oxy-salts or basic salts are produced, which may be considered as derived from SbO OH. So antimonyl nitrate SbO NO 3 and antimonyl sulphate (SbO) 2 SO 4 are known, and antimony oxychloride and tartar emetic, already mentioned, belong in this category. Antimony tetroxide is a white powder obtained by burning antimony in oxygen or by heating the trioxide in the air. In water it is insoluble, while boiled with cream of tartar it is converted into tartar emetic and nietantimonic acid, thus : Sb 2 O 4 + (C 4 H 4 O 6 )HK = (C 4 H 4 O 6 )SbO K + HSbO 3 . The tetroxide is also obtained on igniting antimony pentoxide : Sb 2 5 = Sb 2 4 4-0. Antimony tetroxide may be regarded as the antimonyl salt of nietantimonic acid, which is HSbO 3 . The antimonyl salt would have the formula (SbO) - SbO 3 . Antimonic acid H 3 SbO 4 is formed as an insoluble white powder by the action of concentrated nitric acid upon antimony, or by the action of water on antimony pentachloride. Salts of this acid and also of its dehydration products, pyro- and metan- timonic acids are known. So on fusing antimony with potas- sium nitrate, there is formed with explosive violence potassium metantimonate KSbO 3 , which on being heated with water passes into solution as potassium antimonate KH 2 SbO 4 . On fusing 346 OUTLINES OF CHEMISTRY potassium metantiraonate with caustic potash, the pyroantimo nate K 4 Sb 2 O 7 results : 2 KSbO 3 + 2 KOH = K 4 Sb 2 O 7 + H 2 O. Potassium pyroantimonate is decomposed by water : K 4 Sb 2 7 + 2 H 2 = 2 KOH + K 2 H 2 Sb 2 O 7 . When the latter salt is added to a solution of a sodium salt, sodium pyroantimonate Na 2 H 2 Sb 2 O 7 is precipitated. This is practically the only sodium salt known that does not dissolve in water readily. Antimonic acid and its dehydration products are then quite analogous to those of the corresponding phosphorus and arsenic compounds. Antimony pentoxide Sb 2 O 5 is a yellow powder obtained by heating antimonic acid to 275 At higher temperatures it is decomposed, yielding the tetroxide and oxygen. With strong bases it forms salts. It is soluble in hydrochloric acid. Compounds of Antimony with Sulphur. It has already been mentioned that antimony trisulphide Sb 2 S 3 is found in nature as stibnite. Precipitated from solutions of antimony salts by means of hydrogen sulphide, antimony trisulphide is an orange- red powder, which is insoluble in dilute hydrochloric acid, but soluble in concentrated hydrochloric acid, with concomitant evolution of hydrogen sulphide. In ammonium sulphide it dissolves, yielding ammonium sulphantimonite, thus : Sb 2 S 3 + 3 (N II 4 ) 2 S = 2 (NH 4 ) 3 SbS s . The latter is decomposed by hydrochloric acid : 2 (NH 4 ) 3 SbS 3 + 6 HC1 = 6 NH 4 C1 + Sb 2 S 3 + 3 H 2 S. In yellow ammonium sulphide, antimony trisulphide dissolves more readily, yielding ammonium sulphantimonate : Sb 2 S 3 + 3 (NH 4 ) 2 S + S 2 = 2 (NH 4 ) 3 SbS 4 . On treating the latter with hydrochloric acid, antimony penta- sulphide Sb 2 S 5 is obtained : 2 (NH 4 ) 3 SbS 4 + 6 HC1 = 6 NH 4 C1 + Sb 2 S 5 + 3 H 2 S. Antimony pentasulphide may also be obtained by treating anti monic acid with hydrogen sulphide, thus : 2 H 3 SbO 4 4- 5 H 2 S = Sb 2 S 5 + 8 H 2 O. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 347 It is a powder of golden yellow color, hence it is called sulphur auratum. On being heated, it gives off sulphur and forms the trisulphide. In soluble sulphides of the metals it dissolves, forming sulphantimonates. Thus with sodium sulphide it forms Na 3 SbS 4 + 9 H 2 O, which is known as u Schlippe's salt." Antimony pentasulphide is used in making red vulcanized caoutchouc. The trisulphide is used in making matches, also as a pigment. Antimony cinnabar, kermes mineral, a mixture of the trisulphide and trioxide of antimony, is used in medicine. Occurrence, Preparation, and Properties of Bismuth. This element, though not abundant or widely distributed in nature, has been known since the fifteenth century, when it was referred to by Basil Valentine, who, on account of its brittleness, re- garded it as a half metal. Bismuth generally occurs in the free state in nature, and is almost always fairly pure. Sometimes it is found as the sulphide, bismuth glance Bi 2 S 3 , more rarely as the oxide, bismuth ocher Bi 2 O 3 . The sulphide is roasted to oxide, which is then reduced with charcoal. The bismuth so obtained, or the native bismuth, is refined by fusing it with saltpeter or soda plus a little potassium chlorate. Thus, arsenic and other impurities, consisting mainly of lead, iron, antimony, copper, sulphur, etc., are oxidized and removed as a slag that floats on the surface. Bismuth is a white, brittle metal having a high metallic luster and a slightly reddish sheen, which readily distinguishes it from antimony. Bismuth is crystalline. Its crystals belong to the rhombohedral division of the hexagonal system. Its specific gravity is 9.82. It melts at 269, and may be distilled in a vacuum at about 995. It is a rather poor conductor of heat and electricity, as compared with other metals. The atomic weight of bismuth is 208, and its valence is commonly either three or five ; so that the formulse of its compounds are analogous to those of nitrogen, phosphorus, arsenic, and anti- mony. Nevertheless, bismuth is more pronouncedly basic in character than these, and consequently it is to be grouped with the metals. In the air bismuth remains practically unchanged. On ignition in the air it burns with a bluish- white flame ; the prod- uct formed is a yellow powder, the trioxide, Bi 2 O 3 . In nitric acid, bismuth may readily be dissolved, forming the nitrate 348 OUTLINES OF CHEMISTRY Bi(NO 3 ) 3 ; likewise when the metal is treated with sulphuric acid, the sulphate, Bi 2 (SO 4 ) 3 , is formed. Hydrochloric acid scarcely attacks bismuth. The latter does not combine with hydrogen. Bismuth is used in pharmaceutical preparations. It is also used in making alloys that have a low melting point. Of these the following are frequently used : Rose's metal, consisting of 1 part tin, 1 part lead, and 2 parts bismuth, melts at 93.8; Newton's metal, consisting of 3 parts tin, 5 parts lead, and 8 parts bismuth, melts at 94.5; and Wood's metal, which consists of 1 part tin, 2 parts lead, 1 part cadmium, and 4 parts bismuth, melts at 60.5. On changing from the liquid to the solid state bismuth expands even more than antimony. It is consequently also employed, like the latter, in alloys for stereotyping and other purposes where castings of sharp outline are required. Halogen Compounds of Bismuth. In these compounds bis- muth is always trivalent. Bismuth chloride BiCl 3 is made by the action of chlorine upon bismuth, or by dissolving the latter in nitro-hydrochloric acid. It may also be obtained by dissolv- ing the trioxide, Bi 2 O 3 , in hydrochloric acid. The salt con- sists of white crystals melting at 227, and boiling at about 445. It is soluble in hydrochloric acid solutions, from which it is precipitated in the form of bismuth oxychloride BiOCl : BiCl 3 + H 2 O = BiOCl + 2 HC1. Bismuth fluoride BiF 3 is a grayish powder formed by the action of hydrofluoric acid on bismuth trioxide. On treatment with much water, bismuth oxyfluoride BiOF is formed. Bis- muth bromide BiBr 3 forms orange-colored crystals melting at 215 and boiling at 453. With water they yield bismuth oxybromide BiOBr. Bismuth iodide BiI 3 consists of dark brown or black crystals of metallic luster, melting at 439. On boiling with water they are decomposed, yielding red crys- tals of bismuth oxyiodide BiOI. Halogen compounds of bismuth in which the element has a valence of live have not been prepared, but a dichloiide of the formula (BiCl 2 ) 2 has been described as a white powder formed by heating bismuth with mercurous chloride. Oxides of Bismuth. Bismuth trioxide Bi 2 O 3 , which is formed as a yellow powder when the metal is burned in the air, is the PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 349 most important of the oxides. It acts only as a base, forming salts which may be considered as derived from either Bi(OH) 3 or BiO - OH. Bismuth dioxide Bi 2 O 2 is obtained as a dark brown precipi- tate by pouring a solution containing stannous chloride and bismuth chloride into caustic potash solution. Bismuth tetroxide Bi 2 O 4 is a reddish yellow powder formed by heating the pentoxide to about 165. Bismuth pentoxide Bi 2 O 5 is an unstable brown powder ob- tained by passing chlorine into caustic potash solution contain- ing bismuth trioxide in suspension. On being heated, it forms the tetroxide. With hydrochloric acid it forms bismuth tri- chloride and chlorine : Bi 2 5 + 10 HC1 = 5 H 2 + 2 BiCl 3 + 2 C1 2 . Bismuth Salts of Oxy-acids. The salts of bismuth with the halogens have already been described. With sulphuric acid bismuth forms bismuth sulphate Bi 2 (SO 4 ) 3 , which on treat- ment with water yields the oxysulphate or bismuthyl sulphate (BiO) 2 SO 4 , thus:- Bi 2 (SO 4 ) 3 + 4 H 2 O = (BiO) 2 SO 4 + 2 H 2 SO 4 + 2 H 2 O. With nitric acid, bismuth forms the nitrate Bi(NO 3 ) 8 , which crystallizes in triclinic forms with five molecules of water. The salt is decomposed into basic nitrates by treatment with water. The composition of these basic nitrates varies with the temperature and the relative amounts of water and normal nitrate used in preparing them. Thus a white powder, bismuth oxynitrate BiO NO 3 , is known. On boiling this salt with water, a more basic salt of approximately the composition BiO NO 3 + BiO OH is obtained which is used as a cosmetic and antiseptic under the name bismuth subnitrate. Furthermore, it is very often prescribed in medicine in cases of dysentery and other disturbances of the digestive tract. In the treatment of dis- eases of the skin, particularly in cases of acute inflammations, it is also frequently employed. All salts of bismuth may be regarded as derived from the two basic hydroxides Bi(OH) 3 and BiO OH. Tha univalent radi- cal Bi = O, bismuthyl, is analogous to the antimonyl radical Sb = O. The tendency to form oxy-salts or basic salts is very characteristic of bismuth and also of antimony. 350 OUTLINES OF CHEMISTRY Bismuth Trisulphide Bi 2 S 3 occurs in nature as bismuth glance. It may also be obtained as a very dark brown or black precipitate by passing hydrogen sulphide into a solution of a salt of bismuth : 2 BiCl 3 + 3 H 2 S = 6 HC1 + Bi 2 S 3 . It is insoluble in ammonium sulphide solution, also in solutions of the sulphides of the alkalies. This behavior distinguishes it from the sulphides of arsenic and antimony, which readily dis- solve in alkali sulphides as sulpho-salts. On heating a precipi- tate of bismuth trisulphide suspended in a solution of an alkali sulphide to 200, the compound becomes crystalline. Bismuth trisulphide may also be obtained by melting together sulphur and bismuth in proper proportions. A compound of the composition Bi 2 S 2 , bismuth disulphide, has also been described as consisting of steel-gray needles formed by melting sulphur and bismuth together in the pro- portions represented by the formula. General Considerations of the Group. Nitrogen, phosphorus, arsenic, antimony, and bismuth form another natural group of elements. Their atomic weights increase in the order named, and their physical properties show a corresponding gradation of changes, as is evident from the following table : ELEMENT ATOMIC WEIUIIT CO 1.0 K SPECIFIC GRAVITY MELTING J'OINT BOILING POINT Nitrogen, N Phosphorus, P 14.01 3J..O colorless yellow or red 0.885 (liquid) 1.8-2.3 -210.5 4-44.4 -194.4 + 278.0 Arsenic, As 75.0 i gray, lustrous 5.7 500 (approx.) 450 (approx.) Antimony, Sb 120.2 white, lustrous 6.8 625 1500 (approx.) Bismuth, Bi 208.0 reddish white 9.8 268 1600 (approx.) The chemical properties of the members of the group also present an interesting series of changes as the atomic weight increases. The compounds with hydrogen have the formula RH 3 . So we have ammonia NH 8 , phosphine PH 3 , arsine AsH 3 , PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 351 and stibine SbH 8 . The stability of these compounds dimin- ishes in the order named, a hydride of bismuth being unknown. Ammonia has strong basic properties ; these are also exhibited by phosphine, but to a lesser degree, in the phosphonium salts. But arsine and stibine are no longer able to unite with acids to form salts. Hydrazine (NH 2 ) 2 has its analogue in liquid phos- phine (PH 2 ) 2 , while analogous hydrides of arsenic and anti- mony are unknown. Furthermore, hydrazoic acid HN 3 and solid phosphine P 4 H 2 stand alone, no analogous compounds of the group being known. In general, as the atomic weight of the elements of this group increases, the affinity for hydrogen decreases. Just the reverse is true of the affinity of nitrogen, phosphorus, arsenic, antimony, and bismuth for the halogens. The halogen III V compounds have the general types RX 3 and RX 5 . Thus, we have the following series of the halogen compounds : HALOGEN COMPOUNDS OF THE NITROGEN GROUP NF 3 (?) NC1 3 NBr 3 NI 3 + NH 3 PF 3 PF 5 PC1 3 PC1 S PBr 3 PBr 5 P 2 T 4 PI 3 r AsF 3 AsCl 3 AsBr 3 As 2 I 4 AsI 3 SbF 3 SbF 5 SbCl 3 SbCl 5 SbBr 3 SbI 3 SbI 5 BiF 3 BiCl 3 BiBr 3 BiI 8 While the halogen compounds of nitrogen are so unstable as to be explosive in character, the phosphorus halides possess a considerable degree of stability, which increases as we pass to corresponding compounds of arsenic, antimony, and bismuth in the order named. The phosphorus halides are at once decom- posed by water completely. The arsenic halides suffer sucli hydrolysis more slowly, and even incompletely if but little water is used, while the halides of antimony and bismuth are but partially decomposed by water, forming oxy-salts that are fairly stable. These oxy-salts generally have the formula ROX, like SbOCl, etc., though on treatment with boiling water they form more basic salts because of further hydrolysis. The affinity of the elements of this group toward oxygen and sulphur also diminishes as the atomic weight increases With oxygen we have the following compounds : - 852 OUTLINES OF CHEMISTRY N 2 O NO NA (N0 2 ) N 2 5 . PA (P0 2 ) 2 PA As 2 O 3 As 2 O 4 Sb 2 O 3 (Sb0 2 ) 2 Sb 2 5 (BiO) 2 Bi 2 3 (Bi0 2 ) 2 BiA The oxy-acids are as follows, those in parentheses being known only in the form of salts : H 2 X 2 2 HNO 2 HN0 3 H 3 PO 2 H 3 P0 3 H 3 P0 4 H 4 P 2 7 HP0 3 (HAsO 2 ) H 3 As0 3 H 3 AsO 4 H 4 As 2 O 7 HAsO 3 H8bO 2 H 3 Sb0 3 H 3 SbO 4 H 4 Sb 2 O 7 HSbO 3 .(HBi0 3 ) The sulphides are commonly of the general type R 2 S 8 or R 2 S 5 . They are given in the following table : (N 2 S 2 ) 2 FA As 2 S 2 As 2 S 3 As 2 S 6 Sb 2 S 3 Sb 2 S 6 Bi 2 S 2 BiA The sulphides of arsenic and antimony unite with ammonium sulphide and other alkaline sulphides to form sulpho-salts, while the sulphide of bismuth does not do this. Bismuth is already a fairly pronounced metal. It is, after all, not very closely related to the other members of the group. Nitrogen, too, it must be admitted, stands rather remotely related to the other members of the group. Its relation to the latter is more like that of fluorine toward the other halo- gens, or like oxygen toward sulphur, selenium, and tellurium. Vanadium, Columbium, and Tantalum. These are rare metal- lic elements which form compounds that are analogous to those of the phosphorus group, though in some respects these three elements are also similar to aluminum, iron, chromium, and tungsten. Vanadium (V 51.0) was discovered by Del Rio in 1801 in vanadinite PbCl 2 3 Pb g (VO 4 ) 4 , but was characterized defi- PHOSPHORUS,' ARSENIC, ANTIMONY, AND BISMUTH 353 nitely as an element in 1830 by Sefstrom. In 1867 Roscoe prepared vanadium by heating its dichloride, VC1 2 , in a current of hydrogen. Thus made, vanadium is a crystalline, grayish powder of specific gravity 5.8, which readily burns to V a O 5 in oxygen. Other oxides are V 2 O, V 2 O 2 , V 2 O 3 , V 2 O 4 . The chlorides are VC1 2 , VC1 8 , VOC1 3 , VC1 4 . The vanadates are salts of vanadic acid H 3 VO 4 . As in the case of the phos- phates, there are ortho-, meta-, and pyrovanadates ; of these the metavanadates, like NaVO 3 , are met most frequently. Columbium (Cb 93.5) is also called niobium, No. It occurs together with tantalum (Ta 181.0) in columbites and tantalites. These metals may be prepared like vanadium. Moissan prepared tantalum by reduction of its oxide Ta a O fi with carbon in the electric furnace. Tantalum is a steel-gray, malleable metal which melts at about 2300. It is now being used to make filaments for incandescent electric lamps. As the metal conducts better than carbon, tantalum filaments must be made longer than carbon filaments to obtain the necessary amount of electrical resistance. REVIEW QUESTIONS 1. From what sources is phosphorus obtained? 2. Mention the allotropic forms of phosphorus, stating their char- acteristics. 3. How may yellow phosphorus be prepared from phosphate rock? Give the equations. 4. Explain the chemistry of: (a) friction matches, (b) safety matches. 5. What compounds does phosphorus form with hydrogen? When these compounds are burned in oxygen, what products result? 6. What compound of phosphorus is analogous to : (a) ammonia, (6) nitric acid, (c) nitrogen chloride, (d) nitric acid anhydride ? 7. How may the halides of phosphorus be prepared? Write the equations. How does water act on these halides? Equations. 8. What is the formula of ortho phosphoric acid ? How much of this can be prepared from 50 kilograms of normal calcium phosphate? 9. Define the following and give an example of each, writing the appropriate formula : phosphate, phosphite, pyrophosphate, meta- phosphate, hypophosphite. 10. What is microcosmic salt? What changes does it undergo on heating? Equation. 11. What is bone ash? What is it used for? 354 OUTLINES OF CHEMISTRY 12. By means of structural formulas show the relation between: ortho, meta and pyro phosphoric acids; ortho and meta boric acids; ortho and meta silicic acids. 13. What is the commonest and most important compound of ar- senic? Describe its properties and mention its uses. 14. What property have all arsenic compounds in common? De- scribe arsenic and state how it may be prepared, writing the equation. 15. What is Paris green? How may it be prepared and what is it used for? 16. Write the names and formulas of the hydrogen compounds of phosphorus, arsenic, and antimony that are analogous to ammonia. How may these compounds be prepared ? Discuss the relative stability of these compounds. 17. Write the names and formulas of the chlorides of nitrogen, phos- phorus, arsenic, antimony, and bismuth, and compare their properties. 18. Give the names of the following compounds and also the names and formulas of the corresponding compounds of phosphorus : As 2 03, SbCl 5 , H 3 As0 4 , HSb0 3 , As 2 S 3 , Sb 2 S 5 , (NH 4 ) 3 AsS 4 , (NH 4 ) 3 SbS 3 . 19. By means of chemical equations, show that antimonious oxide may act either as an acid or a base. 20. Write the equations expressing the successive steps that take place when a suspension of white arsenic in water is subjected to the Marsh test. 21. How distinguish between arsenic and antimony mirrors resulting from the Marsh test ? 22. Mention some of the uses of the metals antimony and bismuth. Why are they sometimes called half-metals? 23. Explain the solubility of the sulphides of arsenic and antimony in yellow ammonium sulphide. Write the equations. Why does the sulphide of bismuth not dissolve in yellow ammonium sulphide? 24. What is bismuth subnitrate? What use is made of it? Given 500 grams of metallic bismuth, how much bismuth subnitrate could you prepare from it? Write all the equations. 25. What rare elements belong to the phosphorus group? Why? 26. How many liters of hydrogen would be required to make 60 liters of arsine, measured under the same standard conditions of temperature and pressure? How much zinc would be necessary to prepare this hydrogen? 27. Give the valence of each element in each of the following com- pounds: KC10 4 , Na 2 HP0 4 , HP0 3 , H 3 P0 3 , Bi 2 S 3 , KSb0 2 , PbSi0 3 , NaHC0 3 , KI0 3 , (NH 4 ) 3 AsS 4 , SbOCl, POC1 3 . CHAPTER XX CLASSIFICATION OF THE ELEMENTS THE PERIODIC SYSTEM WE have already seen that the chemical elements may be divided into metals and non-metals, although a sharp line of division between these two groups does not exist. It would be natural to classify the elements according to their physical and chemical properties. Experience has shown that the properties of the elements are closely related to their atomic weights. As early as 1817 Dobereiner called attention to the fact that the atomic weight of strontium, 87.62, is approxi- mately the arithmetical mean of the atomic weights of barium, 137.37, and calcium, 40.09. These three metals are very sim- ilar in character, forming the group of the alkaline earth metals. A number of other elements that are closely related also form similar groups of three, or so-called triads, which is evident from the following cases : Chlorine, 35.46 ; Bromine, 79.92 ; Iodine, 126.92. 35.46 + 126.92 2~~ - = i.iy. Sulphur, 32.07; Selenium, 79.2; Tellurium, 127.5. 32.07 + 127.5 = 7973 Phosphorus, 31.0; Arsenic, 75.0 ; Antimony, 120.2. 31.0+120.2 -- _ ~= 75 - 6 - Lithium, 6.94 ; Sodium, 23.00 ; Potassium, 39.10. 6.94 + 89.10 = 23Q2 2 In 1875 Lenssen attempted to arrange all the elements known in such groups of three. 355 356 OUTLINES OF CHEMISTRY In 1864 Newlands pointed out a relation which he termed the law of octaves. He arranged the elements in the order of the magnitude of their atomic weights, and thus found that the eighth element has properties similar to the first, no matter from which element we begin to count. He did not work out a complete classification, however. In 1869 Dimitri Mendeleeff, and practically simultaneously Lothar Meyer, arranged all of the elements in a table which is known as the periodic system of the elements. In slightly modified form, this table has remained the best classification of the elements to the present day. Referring to the table on page 339 it will be noted that the elements are arranged horizontally in the order of magnitude of their atomic weights, in nine groups, which are numbered from zero to VIII. The symbols in the horizontal series are so written that similar elements appear in the same vertical column. At the head of each column is indicated the valence of the elements towards oxygen and also towards hydrogen. So from left to right the maximum valence of the elements towards oxygen increases from zero to eight, while the valence towards hydrogen is greatest in group IV, i.e. in the middle of the table. Beginning with lithium and passing horizontally to fluorine, the elements show a gradation from .powerfully basic to strongly acidic properties. The second horizontal series, beginning with sodium and ending with chlorine, also shows the same phenom- enon. The first and second series are consequently two com- plete short periods. In the third horizontal series, beginning with potassium and ending with manganese, iron, nickel, and cobalt, we do not have a complete change from strongly basic to strongly acidic properties; but by continuing in the fourth horizontal series from copper to the right we do pass from the more basic elements to the acidic bromine. Consequently the third and fourth horizontal series together are said to form one long period. Again, taking the fifth series, beginning with the basic element rubidium and ending with the metals ruthenium, rhodium, and palladium, and then continuing in the sixth series from silver to iodine, we have a complete change from strongly basic to markedly acidic properties. For this reason, the fifth and sixth horizontal series together are a second long period. CLASSIFICATION OF THE ELEMENTS 357 .2 T3 OK OK MM 5 saiaag saoinaj 68 si as flS Z s 358 OUTLINES OF CHEMISTRY The seventh, eighth, ninth, and tenth horizontal series taken together are sometimes considered as the fifth period, and the eleventh series as a sixth period. These latter periods, to be sure, are incomplete. When helium, neon, argon, krypton, and xenon were discov- ered, the question as to their place in the periodic system arose. As these gases do not combine with anything, their valence is zero, and so these elements have been placed in a zero group at the head of the system. It will be noted that the periodic system contains many blank spaces. These represent elements yet to be discovered. When Mendeleeff first published his table, the spaces now occupied by the elements scandium, gallium, and germanium were vacant. He boldly predicted that these metals would be found; and from the known characteristics of the neighboring elements in the table, he foretold the approximate atomic weights and also described in some detail what the physical and chemical prop- erties of these metals would be. He called the elements ekaboron, ekaalu minium, and ekasilicon. When some years later scandium, gallium, and germanium were discovered, their atomic weights and other properties proved to be those of the metals foretold by Mendeleeff. This brilliant achievement of the great Russian chemist attracted special attention to the value of the periodic system, which has since served to stimu- late inquiry in various lines. Some of the elements did not seem to fit properly into the table and so the question arose whether their atomic weights had really been correctly determined. This led to more accu- rate atomic weight determinations of a number of elements. It will be noted that according to the size of their atomic weights tellurium and iodine ought to change places in the table ; and this ought also to occur in the case of argon and potassium. Considering the properties of these- elements, how- ever, such changes are not to be thought of for a moment, for it would take iodine from the column of the halogens, in which it certainly belongs, and place it in the sixth group with oxygen and sulphur. Similarly, potassium must remain in group I with the other alkali metals. Redeterminations of the atomic weight of tellurium have shown that this element, indeed, has a slightly higher atomic weight than iodine ; and so the anomalies CLASSIFICATION OF THE ELEMENTS 359 o CVJ m 360 OUTLINES OF CHEMISTRY mentioned remain unexplained. The position of hydrogen in the system is also uncertain. This element does not seem to fit into the table. Furthermore, group VIII is peculiar as compared with the other groups. It contains three groups of three elements each, though to be sure the elements in each of these groups have approximately the same atomic weight. A number of the rare-earth elements (which see) have atomic weights that are not widely different from one another, and for these there do not appear to be suitable places in the system. Stated in words, the so-called periodic law is that the physical and chemical properties of the elements are periodic functions of their atomic weights. An illustration of this is given in Fig. 122, in which the atomic weights are represented as abscissas and the atomic volumes (i.e. the atomic weights divided by the specific gravities) are represented as ordinates. The trend of the curve shows the periodicity, similar elements appearing in similar positions on the curve, which was first published by Lothar Meyer. The periodic system does not represent a sharp quantitative relationship between the atomic weights and the properties of the elements. Some of its anomalies have already been men- tioned. In spite of these imperfections, the periodic system offers a useful means of classifying the elements, which will in our further considerations be grouped accordingly. It will be noted that in considering the non-metals the natural families that have been studied really represent the essential elements of certain groups of the periodic system. The reader will comprehend the significance of the periodic system much better after having studied the physical and chemical peculiarities of the metals which still remain to be considered. REVIEW QUESTIONS 1. What relation exists between the atomic weights of chlorine, bro- mine, and iodine ? Mention three other groups of three elements each which also exemplify this relation. 2. What is the law of octaves ? Who discovered it? 3. What is the first complete classification of the chemical elements called? 4. State the so-called periodic law in words. Mention the names of the two men who discovered this law. CLASSIFICATION OF THE ELEMENTS 361 5. How many vertical columns are there in the periodic system of the elements and what do the numbers of these columns correspond to? What is meant by a complete short period ? By a complete long period ? 6. Given a chart of the periodic system, point out the complete short periods, also the complete long periods. 7. Criticize the position of the following elements in the periodic system : manganese, iodine, tellurium, argon, nickel, hydrogen. 8. Of what use has the periodic system been in the advancement of chemistry? CHAPTER XXI THE ALKALI METALS THE alkali metals are potassium, sodium, lithium, rubidium, and caesium. Of these potassium and sodium are by far the most abundant and important. Lithium is also found in fair quantities, but rubidium and caesium occur only in very small amounts, and they will consequently receive less consideration here. None of the metals of this group are found in the free state in nature. They always occur in the form of salts, which is due to the fact that their affinity for oxygen, the halogens, sulphur, and other non-metals is very great. The compounds which the alkali metals form with the non-metals are in general simple, very stable, and well characterized. These metals are univalent in all of their salts, hydroxides, and oxides. The hydroxides of the alkali metals are the most powerful bases known. The solutions of the hydroxides are very alkaline and caustic, so that they are commonly called the caustic alkalies. These alkalies are non-volatile, and are consequently termed the fixed alkalies, in contradistinction to ammonium and other very basic groups which may under proper conditions be volatilized. So far as its general chemical behavior is concerned, ammonium (NH 4 ) is closely allied to the alkali metals, and it will conse- quently be advantageous to refer to the chemistry of the ammonium compounds in this chapter. Occurrence, Preparation, and Properties of Potassium. Potas- sium is very widely distributed as a constituent of silicates like potassium feldspar and certain forms of mica. Inasmuch as soils are produced by the disintegration of rocks by the process of weathering, all soils contain potassium in the feldspathic constituents they have derived from rocks. Plants take up potassium salts from the soil, and in the ashes of plants the potassium is found as potassium carbonate, commonly called potash. By treating ashes with water and filtering, the potas- THE ALKALI METALS 363 slum carbonate is obtained by evaporating the filtrate. Since human beings and animals get their food supply directly or indirectly from plants, it is not strange that potassium is found in all animal tissues and secretions, like muscles, bones, blood, urine, albumen, eggs, milk, etc. Oceanic water contains about 0.04 per cent potassium, while the earth's crust contains about 2.45 per cent of the element. Though potassium is thus widely distributed, it occurs in large quantities in but few places. The ehief deposits of potassium salts are found in G-ermany, notably at Stassfurt, where they occur as layers twenty to thirty meters thick, covering strata of native common salt. Potassium occurs here mainly as carnallite KC1 MgCl 2 6 H 2 O and kainite MgSO 4 - KC1 3 H 2 O, but also as sylvite KC1. Associated with sodium nitrate, potassium nitrate is found to some extent in Peru and Chili. Metallic potassium was first prepared by Sir Humphry Davy, who in 1807 electrolyzed molten caustic potash. At present it is prepared commercially by electrolysis of either potassium chlo- ride or potassium hydroxide, though formerly it was largely made by heating potassium carbonate with carbon: K 2 CO 3 + 2 C = 2 K + 3 CO. In this process the potassium passes off as vapor which is con- densed and kept under petroleum oils. Potassium is a silvery white metal, which has a bright metal- lic luster, and is soft as wax at ordinary temperatures. Below it becomes hard and brittle. Its specific gravity is 0.865 at 15. It melts at 62.5 and boils at 667. Its vapors are green. The atomic weight of potassium is 39.10; and its molecular weight is the same, the vapors being about twenty times as heavy as hydrogen. Potassium reacts vigorously with water, evolving hydrogen and forming potassium hydroxide. The heat generated during the action is generally so great as to set the hydrogen and some of the potassium on fire, thus giving rise to explosions. A freshly cut surface of potassium at once becomes blurred because of reaction with the moisture of the air. The metal is consequently kept under petroleum oils. Potassium hydride KH is formed by passing hydrogen over potassium at 360. It consists of white needlelike crystals, that catch fire on exposure to the air. Water decomposes 364 OUTLINES OF CHEMISTRY the compound, forming potassium hydroxide and hydrogen, thus : 2 KH + 2 H 2 = 2 KOH + 2 H 2 . With carbon dioxide it readily forms potassium formate : KH + CO 2 = HCOOK. Compounds of Potassium with the Halogens. Of these salts potassium chloride KC1 is the commonest. In nature it occurs as sylvite KC1, and also in carnallite MgCl 2 KC1 6 H 2 O, as already stated. Potassium chloride crystallizes in cubes which melt at 730 ; at higher temperatures it is converted into vapor. With many other salts it unites to form double salts, examples of which we have in carnallite and kainite. Water dissolves potassium chloride readily. The salt is not soluble in liquid hydrochloric acid ; hence the addition of hydrochloric acid to a concentrated aqueous solution of potassium chloride causes a precipitate of the latter to form. Potassium bromide KBr is made by the action of bromine on potassium hydroxide : 6 KOH + 3 Br 2 = KBrO 3 + 5 KBr + 3 H 2 O. The potassium bromate simultaneously formed is reduced by heating the product with carbon. Potassium bromide forms cubical crystals that melt at 715. It is used in medicine and in the process of preparing silver bromide for photographic plates. Potassium iodide KI is prepared by the action of iodine upon potassium hydroxide, the process being analogous to that de- scribed for making the bromide. The following method is also used for making the iodide : Iodine is mixed with iron filings under water, when a solution of a compound Fe 3 T 8 (that is, FeI 2 + 2 FeI 3 ) is formed. This when treated with potassium carbonate yields potassium iodide, which remains dissolved, and an hydroxide of iron which is insoluble and can be filtered off. Carbon dioxide gas is also given off during the change, which is : Fe 3 I 8 + 4 K 2 C0 3 + 4 H 2 = 8 KI + Fe 3 (OH) 8 + 4 CO 2 . Potassium iodide crystallizes in cubes, on evaporation of the nitrate. The salt melts at 625. It is more copiously soluble THE ALKALI METALS 365 in water than the bromide. Its aqueous solutions readily ac- quire a yellow color, due to the separation of free iodine formed by the action of oxygen and carbon dioxide of the air upon the salt. Solutions of potassium iodide readily dissolve additional iodine. These solutions are frequently employed in analytical chemistry. Potassium iodide is often used in medicine, also in photography. Potassium fluoride KF is formed by treating potassium hydroxide or carbonate with hydrofluoric acid. With the latter it readily forms the double compound KF HF. Potas- sium fluoride is a deliquescent white salt, forming cubical crystals of the composition KF + 2 H 2 O. The solutions attack glass. All of the halides of the alkali metals form double salts with salts of. many other metals. The halides of the alkalies may all be prepared by the action of the caustic alkalies upon the hydrohalogen acids, or by the direct union of the metals with the free halogens. The meth- ods employed in preparing the bromide and iodide of potassium, as above described, are used because of the difficulty of making pure hydrobromic and hydriodic acids, to which fact attention has already been called. Potassium Hydroxide KOH. This compound is also called caustic potash and potassium hydrate. It is prepared by treat- ing potassium carbonate with slaked lime in vessels of iron or silver, for caustic alkalies attack glass or porcelain. The reaction is : K 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2 KOH. The calcium carbonate is insoluble, which fact really forms the basis of the process. On evaporating the clear nitrate, caustic potash is obtained. The latter is also made in large quantities by the electrolysis of solutions of potassium chloride. In this process the electric current enters the solution by a carbon plate dipping into it, and leaves the solution by a mercury sur- face also submerged in it, but not in contact with the carbon. Thus as the current passes, chlorine is liberated on the carbon and is conducted off in pipes and used for making bleaching powder ; at the same t^'me, potassium is liberated on the mer- cury, in which it dissolves. This solution of potassium in 366 OUTLINES OF CHEMISTRY mercury is called potassium amalgam. Water acts on it slowly, forming potassium hydroxide and hydrogen, leaving the mer- cury behind. The aqueous solution on evaporation yields solid caustic potash. The essential reactions of the process are, first, by electrolysis, 2 KC1 = 2 K + C1 2 ; and when the amalgam is acted upon by water, 2 K H- 2 H 2 O = 2 KOH + H 2 . Potassium hydroxide is a hard, brittle, white solid, which is deliquescent and very soluble in water with evolution of heat. The solution is very caustic, having a corrosive and disinte- grating action upon animal and vegetable tissues. It is the most powerful of the ordinary bases, and consequently gen- erally decomposes the salts of other bases. Caustic potash commonly comes into the market cast in sticks which contain about 80 per cent of the compound KOH and 20 per cent water. Caustic potash readily absorbs carbon dioxide, forming potassium carbonate. As a drying agent and an absorbent for carbon dioxide, potassium hydroxide is much used in chemical laboratories, though sodium hydroxide, which is cheaper, is often employed in its place when it will do just as well. Caustic potash is used in making soft soaps. Potassium Oxide K 2 O may be prepared (1) by melting potas- sium and potassium hydroxide together, or (2) by heating potassium nitrate with potassium ; the reactions are : (1) 2 KOH + 2 K = 2 K 2 + H 2 . (2) 2 KN0 3 + 10 K = 6 K 2 O + N a . The oxide is a white, unstable powder. With water it unites, yielding potassium hydroxide. Exposed to the air, it absorbs oxygen, forming potassium peroxide KO 2 , which is a yellow powder. With water this yields oxygen, hydrogen peroxide, and potassium hydroxide : 4 K0 2 + 6 H 2 = 4 KOH + 4 H 2 O 2 + O 2 . Peroxide of potassium is also formed together with the oxide when potassium is burned in the air or in oxygen. Potassium Chlorate KC1O 3 may be obtained by passing chlorine into a hot solution of potassium hydroxide, as already THE ALKALI METALS 367 described. By electrolyzing a solution of potassium chloride, chlorine and potassium hydroxide form at the opposite elec- trodes ; and by stirring the hot solution, the chlorate thus forms and crystallizes out. Often solutions of the more soluble sodium chlorate are thus prepared, by means of which potassium chlorate is precipitated from potassium chloride solutions. Nearly all the potassium chlorate of commerce is now made elec- trolytically . Potassium chlorate crystallizes in the monoclinic system. About six parts of it dissolve in 100 parts of water at room temperature. The salt melts at 350 and yields oxygen at a slightly higher temperature. It is used for making oxygen, also in manufacturing matches, fireworks, and explosives. The ease with which the salt gives up oxygen is shown by mixing two or three grains of it with a grain of sulphur or of red phosphorus in a mortar. As the substances are pressed to- gether by means of the pestle, there is an explosion. By heating the chlorate, potassium perchlorate KC1O 4 is produced as a first decomposition product, thus : 8 KC1O 3 = 5 KC1O 4 + 3 KC1 + 2 O 2 . It forms rhombic crystals and is less soluble than either the chlorate or chloride, and consequently it may readily be sepa- rated from these by fractional crystallization. At 400 the perchlorate decomposes into chloride and oxygen : KC10 4 = KC1 + 2 2 . Potassium bromate KBrO 8 and potassium iodate KIO 3 are analogous to the chlorate. The methods of their preparation have already been mentioned under potassium bromide and iodide. Potassium Nitrate KNO 3 , also called saltpeter, is widely dis- tributed in soils in small quantities, being formed wherever organic substances decay. It was formerly produced on a large scale by allowing refuse of nitrogenous organic bodies to decay in presence of potassium salts. At present potassium nitrate is made by treating hot, saturated solutions^of Chili saltpeter NaNO 3 with potassium chloride, thus : NaNO q + KC1 = KNO, + NaCl. 368 OUTLINES OF CHEMISTRY The sodium chloride formed, being far less soluble than potas- sium nitrate, is precipitated, and from the clear supernatant solution potassium nitrate is readily obtained in form of crys- tals on cooling. The product is further purified by recrystal- lizing. At 0, 100 parts of water dissolve 13 parts of KNO 3 , while' at 100, 247 parts of the salt are dissolved. Potassium nitrate crystallizes in rhombic prisms, which change into rhombohedra at about the melting point of the salt, 339. When heated above its melting point, potassium nitrate gives off oxygen, forming potassium nitrite KNO 2 . The latter salt is more readily formed by heating the nitrate with lead or iron, which take up the oxygen, forming oxides. Potassium nitrate is used as a fertilizer, as a preservative for meat, as an oxidizing agent in the laboratory, and as an ingredi- ent of fireworks and gunpowder. Most of it is used in making gunpowder, which consists of a mixture of 75 parts saltpeter, 13 parts charcoal, and 12 parts sulphur. When it is ignited, the following reaction takes place : 2 KN0 3 + S + 3 C = K 2 S + 3 CO 2 + N a . Thus a large volume of gas is suddenly liberated, and this causes the explosion. The pressure of the gases produced at 2200, the temperature of the discharge, is about 96,000 pounds per square inch. When discharged under pressure, as in a gun, the chemical reaction is somewhat more complicated than the one above given, potassium carbonate, sulphate, and thiosur* phate being also produced in notable amounts. Besides this, products of partial combustion remain suspended in tile air in a finely divided state, as smoke, after the discharge. Black potvder, as it is called, is being displaced more and more by smokeless powder, which see. Potassium Cyanide KCN has already been mentioned in con- nection with cyanogen. Besides being made as there described, it is prepared on a large scale by heating potassium ferrocyanide K 4 Fe(CN) 6 with potassium carbonate : K 4 Fe(CN) 6 + K 2 C0 3 = 5 KCN + KCNO + Fe + CO 2 . The potassium cyanide thus prepared always contains some potassium cyanate KCNO, which, in case a pure product is required, is reduced by means of charcoal or zinc. On evapo- THE ALKALI METALS 369 ration of the clear solution, potassium cyanide is obtained in white deliquescent lumps that readily dissolve in water. Moisture decomposes the salt somewhat, yielding caustic potash and hydrocyanic acid. In presence of carbon dioxide, potas- sium carbonate is formed together with hydrocyanic acid. Hence the odor of the latter is ever present with potassium cyanide. This salt is extremely poisonous. It is a powerful reducing agent, readily passing over into the cyanate, a white salt which also dissolves in water.. Potassium cyanide also unites with sulphur to form potassium sulphocyanate KCNS, as already mentioned. This salt is crystalline, deliquescent, and consequently readily soluble in water. Potassium cyanide is used in very large quantities for extract- ing gold from its ores. It is also employed in photography and in gold and silver electroplating, which see. Potassium Carbonate K 2 CO 3 was formerly prepared as potash by leaching out wood ashes. The molasses residues of the beet sugar industry and the fat of sheep's wool also contain potas- sium salts, from which potash is obtained. The bulk of the po- tassium carbonate on the market is made from potassium sulphate and chloride obtained from the Stassfurt. deposits. The method employed is analogous to the Le Blanc process for making sodium carbonate (which see). In addition, however, potas- sium carbonate is now made at Neustassfurt by passing carbon dioxide into magnesium carbonate MgCO 3 3 H 2 O suspended in a solution of potassium chloride, thus : 2 KC1 + 3(MgCO 3 . 3 H 2 O) + CO 2 = 2(MgC0 8 KHC0 8 - 4 H 2 O) + MgCl 2 . When the double carbonate thus formed is properly heated, magnesium carbonate and potassium carbonate are obtained ; the former, being insoluble, is filtered off, and the filtrate, upon evaporation, yields potassium carbonate. The latter melts at about 840. At 0, 100 parts of water dissolve 83 parts of the salt, while at 20, 112 parts of the salt are thus dissolved. From concentrated solutions, monoclinic crystals, 2 K 2 CO 8 4- 3 H 2 O, may be obtained. The solutions have a strongly alka- line reaction, for the salt is one of a weak acid with a powerful base, and is consequently appreciably hydrolyzed. In the mar- ket, potassium carbonate is often called pearlash. It is used 370 OUTLINES OF CHEMISTRY in making hard glass (potash glass), soft soap, and many salts of potassium. By passing carbon .dioxide into solutions of potassium car- bonate, potassium bicarbonate KHCO 3 is formed. It is less soluble than the carbonate, but more so than sodium bicarbon- ate, and hence cannot, like the latter, be obtained by the Solvay process (which see). Heating potassium bicarbonate, even in aqueous solutions, decomposes it, yielding the carbonate, carbon dioxide, and water. Potassium Silicate K 2 SiO 3 is made by fusing silica with the carbonate : K 2 CO 3 + SiO 2 = K 2 SiO 8 + CO 2 . Thus a glassy, deliquescent mass is obtained, which dissolves in water, yielding a thick sirupy solution popularly called potassium water glass. The solutions commonly contain other potassium silicates besides K 2 SiO 3 . Its uses are the same as those of the cheaper sodium water glass (which see). Potassium Fluosilicate K 2 SiF 6 is formed as an amorphous, translucent precipitate, when solutions of a potassium salt are treated with hydrofluosilicic acid. It is soluble in about 800 parts of water at 20. Potassium Phosphates. The existence of three phosphates, the primary, KH 2 PO 4 , the secondary, K 2 HPO 4 , and the tertiaiy, K 3 PO 4 , has already been mentioned. They are white salts, readily soluble in water. Their general characteristics have been sufficiently described. Potassium Sulphate K 2 SO 4 is found at Stassfurt in schonite K 2 SO 4 MgSO 4 3 H 2 O, from which it is obtained by treatment with potassium chloride in solutions, thus : K 2 SO 4 - MgSO 4 + 2 KC1 = 2 K 2 SO 4 + MgCl 2 . The magnesium chloride is very soluble and hence remains in solution, while the less soluble potassium sulphate is precipi- tated. Potassium sulphate is also obtained by the action of sulphuric acid on potassium chloride. It crystallizes in rhombic forms without water of crystallization. Its melting point is 1080. At room temperature 10 parts of the salt dissolve in 100 parts of water. It is used as a fertilizer, also in making potassium alum, hard glass, and potassium carbonate. THE ALKALI METALS 371 On treating potassium sulphate with sulphuric acid, acid potassium sulphate KHSO 4 is formed. This is very soluble in water and melts at about 200, forming water and potassium pyrosulphate K 2 S 2 O 7 , which, on further heating, decomposes into the normal sulphate and sulphur trioxide. The latter at higher temperatures readily unites with many metallic oxides ; hence the practice of fusing refractory metallic oxides and many minerals with potassium bisulphate KHSO 4 to convert the bases into soluble sulphates. Potassium Sulphite K 2 SO 8 is formed by passing sulphur dioxide into a solution of potassium carbonate till carbon diox- ide is no longer formed. The salt crystallizes in monoclinic prisms with two molecules of water. On saturating a solution of potassium carbonate or sulphite with sulphur dioxide, acid potassium sulphite or potassium bisul phite KHSO 3 , crystallizing in needles, is obtained. Sulphides of Potassium. Potassium sulphide K 2 S is made by fusing potassium sulphate with charcoal: K 2 S0 4 + 4C = K 2 S + 4 CO. It is a flesh-colored, crystalline mass that readily dissolves in water. The oxygen of the air acts on the solutions, gradually forming potassium thiosulphate K 2 S 2 O 3 , thus : - 2 K 2 S + H 2 O + 2 2 = 2 KOH + K 2 S 2 O 8 . By saturating a caustic potash solution with hydrogen sul- phide, potassium sulphydrate KSH is formed : KOH + H 2 S = KSH + H 2 O. Its solutions are alkaline, and upon evaporation with caustic potash they yield potassium sulphide: KSH + KOH = K 2 S + H 2 0. Solutions of K 2 S or KSH will readily dissolve sulphur, forming a series of compounds known as polysulphides. Thus K 2 S 3 , K 2 S 4 , and K 2 S 5 have been obtained. On treatment with acids the polysulphides yield hydrogen sulphide and sulphur. By fusing potash with sulphur out of contact with air, a mass of liver-brown color, known as liver of sulphur, is obtained. It consists of a mixture of polysulphides of potassium together with potassium sulphate and thiosulphate. 372 OUTLINES OF CHEMISTRY Tests for Potassium. To detect the presence of potassium in small quantities, the spectroscope (which see) is employed. In addition, the fact that acid potassium tartrate KHC 4 H 4 O 6 , potassium silicofluoride K 2 SiF 6 , and potassium platinic chloride K 2 PtCl 6 are difficultly soluble in water serves to determine whether potassium salts are present in a given solution or not The reactions involved, written for potassium chloride, are as follows : KC1 + H 2 C 4 H 4 6 = KH . C 4 H 4 6 + HC1. 2 KC1 + H 2 SiF 6 = K 2 SiF 6 + 2 HC1. 2 KC1 + PtCl 4 = K 2 PtCl 6 . Potassium platinic chloride is soluble in about 100 parts of water; but its solubility is much less in alcoholic solutions, of which fact the analytical chemist avails himself. Rubidium and Caesium. These metals have the atomic weights 85.45 and 132.81 respectively. They were discovered in 1860 by Robert Bunsen by means of the spectroscope, in the residues obtained by evaporating Durkheim mineral water. The spectrum of rubidium contains certain characteristic red lines, while that of caesium exhibits striking blue lines. Hence Bunsen named the elements rubidium (red) and caesium (blue) ; their symbols are Rb and Cs. Salts of both of these metals occur widely distributed in nature, though in extremely minute quantities; they are commonly associated with salts of potassium. Carnallite contains about 0.025 per cent of rubid- ium, and it has been estimated that from the Stassf urt salts that are used as fertilizers more than 200 tons of rubidium are dis- tributed annually over the soils, from which plants absorb it. Thus in the ash of sugar beets, tobacco, coffee, and tea, rubidium is frequently met. The methods that are used for preparing potassium also serve for making rubidium. The chief source of rubidium salts is the Stassf urt deposit. Caesium is much rarer than rubidium. The mineral pollux found on the island of Elba is a silicate of aluminum and caesium. It contains about 34 per cent of caesium oxide. Metallic caesium was first prepared in 1881 by Setterberg, who electrolyzed the cyanide CsCN. Caesium may also be ob- tained bv heating its oxide or carbonate with magnesium. The salts of both rubidium and caesium are in general analogous THE ALKALI METALS 373 to those of potassium, with the exception that the former elements are also able to form halides containing three or jive halogen atoms . In such compounds rubidium and caesium consequently exhibit valences of three and five, respectively. The hydroxide of rubidium is a stronger base than that of potassium, while the hydroxide of caesium is the most powerful base known. Occurrence, Preparation, and Properties of Sodium. The chief compound of sodium is sodium chloride, which is widely distributed in nature in large quantities. Thus, it occurs in oceanic waters, while many salt seas and lakes are practically saturated solutions of common salt. Mineral springs are often rich in sodium salts, which also occur in huge deposits as chlo- ride, nitrate, and borate, in various parts of the globe. Cryolite, an aluminum sodium fluoride, is found in Greenland, and albite or soda feldspar, a silicate of sodium and aluminum, is widely distributed in nature. Just as land plants contain potassium, so sea plants contain sodium, which is found in their ashes as carbonate. From the soil, sodium gets into plants and then into animal organisms, where it occurs in the blood and the various tissues and secretions. Metallic sodium was first prepared by Sir Humphry Davy in 1807 by electrolysis of molten sodium hydroxide. In this way it is now prepared on a large scale. Thus, sodium is deposited at one pole, and the hydroxyl liberated at the opposite pole at once decomposes, yielding water and oxygen. The yield is only about 40 per cent, for some of the metal con- tinually reacts with the water, liberating hydrogen and forming sodium hydroxide. The methods described for making potassium may also be used for preparing sodium. Sodium has properties similar to those of potassium. At room temperature the metal is soft like wax, while at 20 it is hard. It has a bright, silvery luster. Its melting point is 95.5, and its boiling point 742. The vapor of sodium is colorless, though in thick layers it appears violet. The vapor is 12 times as heavy as hydrogen, whence the molecular weight is 24. The atomic weight is 23, and the valence is one. The molecules of so- dium, like those of potassium, consist of but one atom. Sodium decomposes water, like potassium, but the action is not as violent as in the case of the latter. The specific gravity of sodium is 0.974 at 15. The metal is used in making sodium cyanide and 374 OUTLINES OF CHEMISTRY sodium peroxide, also in the manufacture of complex organic compounds. In the laboratory, it is frequently employed as a reducing agent. Like potassium, it is kept under petroleum oil ; but sodium is also shipped in tightly soldered tinned iron boxes. Sodium dissolves in mercury, forming sodium amalgam, which when treated with water yields mercury, sodium hydroxide, and hydrogen. The latter is liberated much less rapidly than when sodium alone acts on water, hence the amalgam is often em- ployed in effecting reductions that are to proceed slowly. With potassium, sodium forms alloys, 1 part sodium to 2 to 10 of potassium, that are liquid at room temperature and have the appearance of mercury. If but little potassium is alloyed with much sodium, that is, if the proportions mentioned are reversed, the alloys formed are brittle solids. With hydrogen, sodium forms sodium hydride NaH, which is similar to the hydride of potassium. Sodium Chloride NaCl, common salt, is the chief source of sodium and its compounds. Large deposits are found at Stassfurt and Reichenhall, in Germany, at Wieliczka in Galicia, at Cheshire in England, at Syracuse in New York, in Michigan, Kansas, Texas, Utah, California, as well as in Asia and Africa. In some localities the salt is mined in solid form, and again it is fre- quently obtained from brines by evaporation. When the brine is dilute, as in case of sea water, it is concentrated by allowing it to trickle over a large surface of twigs, thus giving better op- portunity for evaporation to take place. The concentrated solu- tion thus obtained is then generally evaporated by means of arti- ficial heat ; though in hot climates, solar heat is often relied upon entirely. Brine is also evaporated in shallow basins either by means of artificial heat or the heat of the sun. Again, so-called vacuum pans, in which brines are evaporated under diminished pressure, are frequently employed. Ordinary salt is not pure. It contains small amounts of sodium sulphate and chlorides of calcium and magnesium. The latter are deliquescent, and their presence in common salt causes it to attract moisture from the air. Pure common salt is obtained by conducting hydrochloric acid gas into a saturated solution of common salt. The latter is thus thrown out of solution, for it is less soluble in hydro- chloric acid than in water. THE ALKALI METALS 375 Each year the United States produces about 5,000,000 tons of sodium chloride, which is approximately one fourth of the world's annual output. Sodium chloride crystallizes in cubes, which when obtained from aqueous solutions form groups of hollow pyramids, that is, are hopper-shaped (Fig. 123). They contain occluded water which escapes on heating, causing the salt to decrepitate. At 10, sodium chloride crystallizes with two mole- cules of water, forming monoclinic prisms. Sodium chloride melts at about 800, at which temperature it also volatilizes appreciably. At FIG. 123. inn room temperature. 1UO parts ot water dissolve 36 parts of salt, and at 100, 39 parts. The salt is therefore about as soluble in hot as in cold water. In practically all solvents other than water, common salt is insoluble. Human beings and many animals cannot live without sodium chloride, which is found distributed throughout their systems in small amounts, though its function is not known. It 'has been estimated that the quantity of salt used annually by a i uman being amounts to about one tenth of his weight. Common salt is used in enormous quantities in making hydrochloric acid, chlorine, and nearly all of the sodium com- pounds. Sodium bromide NaBr and sodium iodide Nal are more solu- ble in water than the chloride; they are also more volatile. In general, these halides and sodium fluoride are much like the corresponding salts of potassium. Oxides and Hydroxide of Sodium. When sodium is burned in oxygen, a mixture of the two oxides, Na 2 O and Na 2 O 2 is formed. Sodium oxide Na 2 O is also obtained as a gray mass by heat- ing the hydroxide with sodium. When treated with water, it dissolves, forming a solution of the hydroxide. Sodium peroxide Na 2 O 2 is a white powder formed by pro- longed heating of sodium or sodium oxide in oxygen or in the air at about 300. At high temperatures it decomposes into the oxide and oxygen. With water it yields sodium hydroxide and oxygen, though in the cold it may in part be dissolved. 376 OUTLINES OF CHEMISTRY With dilute acids it yields hydrogen peroxide ; hence it may be regarded as the latter substance whose hydrogen atoms are replaced by sodium. Sodium peroxide is prepared on a large scale as an oxidizing and bleaching agent. It is shipped in air- tight tin cans. Sodium hydroxide NaOH, caustic soda, is made by the same methods as caustic potash, namely, by treating the carbonate with slaked lime or by electrolysis of solutions of the chloride, using a mercury cathode. It is also made by the Acker process, which consists of electrolyzing fused sodium chloride, using a FIG. 124. carbon anode and a cathode of molten lead. The chlorine is conducted off and used in making bleaching powder, while the sodium forms an alloy with the lead. A jet of steam directed on this alloy reacts with the sodium, forming hydrogen, which burns at once, and sodium hydroxide, which, being molten, is drawn off into proper containers. Figure 124 shows the essen- tial features of the apparatus. THE ALKALI METALS 377 The properties of sodium hydroxide are much like those ol potassium hydroxide. The former is much cheaper, however, and is used in place of potassium hydroxide whenever possible. FIG. 125. Large quantities of sodium hydroxide are used in making soap and in " softening water." It is also used in the paper industry and in making carbolic acid, oxalic acids, and many other products. Sodium Carbonate Na 2 CO 3 , popularly called soda or sal soda, is manufactured from common salt on a very large scale, for it is essential in making glass, soap, caustic soda, and other sodium compounds. Sodium carbonate is found in nature in Wyoming, California, Mexico, Egypt, and in the ashes of marine plants. It is manufactured by the Le Blanc process, the Solvay process, and the electrolytic process. The Le Blanc soda process was introduced in France by Le Blanc in 1791. It is based upon three chemical reactions. First* common salt is treated with sulphuric acid in equivalent quantity ; thus hydrochloric acid and " salt cake," consisting of sodium chloride and sodium acid sulphate, are formed : HC1. 2 NaCl + H 2 S0 4 = NaCl + NaHSO 4 This salt cake is then transferred to the hearth of a furnace, (Fig, 125) and heated ; thus sodium sulphate and more hydro- chloric acid are produced : NaCl + NaHSO 4 = Na 2 SO 4 + HCL The hydrochloric acid formed is absorbed in water and placed on the market. The second step consists in mixing the sodium sul- phate formed with charcoal and calcium carbonate, and heating 378 OUTLINES OF CHEMISTRY the mixture in a rotating cylindrical furnace (Fig. 126) which has a central flue through which the hot gases of the furnace pass. Thus sodium sulphate is reduced to sulphide : Na 2 S0 4 + 40 = Na 2 S + 4 CO ; and sodium sulphide reacts with calcium carbonate, yielding sodium carbonate and calcium sulphide : Na a S + CaCOg = CaS + Na 2 CO 3 . The third step consists in treating the mixture of calcium sul- phide and sodium carbonate; called black ash, with water. Thus sodium carbonate dissolves and calcium sulphide remains behind, as an insoluble residue. By evaporation of the clear solution FIG. 126. crystals of the composition Na 2 CO 3 H 2 O are obtained, which are heated to drive off the water, thus forming Na 2 CO 3 , calcined soda. When this is recrystallized from water at room temper- atures, crystals of the composition Na 2 CO 3 ' 10 H 2 O are formed. This is the washing soda or crystallized soda of commerce. Sodium carbonate made by the Le Blanc process generally con- tains small amounts of chloride, sulphate, and sulphite. The sulphur contained in the calcium sulphide, also called tank waste, is recoveredby a process developed in 1888 by Chance in England. It consists of diluting the waste with water and arranging it in a series of cylinders through which carbon THE ALKALI METALS 379 dioxide from a limekiln is run. There is first formed calcium carbonate and calcium sulphydrate, thus : 2 CaS + C0 2 + H 2 = CaCO 3 + Ca(SH) 2 . By further action of carbon dioxide, this sulphydrate is decom- posed, yielding hydrogen sulphide : Ca(SH) 2 + C0 2 + H 2 = CaCO 3 + 2 H 2 S. The hydrogen sulphide set free in one cylinder reacts with the material in the next, and so on through the series of cylinders; thus, H 2 S = Ca(SH) 2 . Finally, the hydrogen sulphide is burned to water and sulphur, thus, 2H 2 S + O 2 = 2H 2 O + 2 S; or it is burned completely to water and sulphur dioxide, and then made into sulphuric acid. About 90 per cent of the sul- phur in the tank wastes may be recovered. The Solvay process, also called the ammonia soda process, was perfected by the Belgian chemist Solvay in 1863. It is based upon the fact that sodium bicarbonate NaHCO 3 is relatively sparingly soluble in water. The process consists of conducting ammonia and carbon dioxide into a cold, saturated solution of common salt. In this way ammonium bicarbonate is formed, which reacts with sodium chloride, causing a precipitate of so- dium bicarbonate to separate out : NaCl + NH 4 HCO 3 = NH 4 C1 + NaHCO 3 . The sodium bicarbonate is placed on the market as such, or it is heated and thus converted into carbonate, 2 NaIIC0 3 = Na 2 C0 3 + CO 2 + H 2 O. The ammonium chloride remains in solution, from which, by heating with lime or magnesia, ammonia is again regenerated. Thus, the waste products are chlorides of calcium or magnesium. On heating magnesium chloride, magnesium oxide and hydro- chloric acid are formed : MgCl 2 + H 2 O = MgO + 2 HC1. 380 OUTLINES OF CHEMISTRY In this way the magnesia can be recovered and used over again. In Germany and the United States most of the sodium carbon- ate is made by the Solvay process, while in England the Le Blanc process still predominates. It is evident that a purer product is obtained by the Solvay process. The electrolytic process consists of making caustic soda by one of the electrolytic methods described and then treating the solution with carbon dioxide. At ordinary temperatures sodium carbonate crystallizes in monoclinic prisms of the composition Na 2 CO 3 10 H 2 O, which effloresce. At 60 this compound melts in its crystal water and on continued heating it yields a deposit of the composition Na 2 CO 3 2 H 2 O, which when dried in the air readily forms Na 2 CO 3 H a O. At 100 the salt may be completely dehydrated. At 15, 100 parts of water dissolve 55 parts of Na 2 CO 3 , at 38, 138 parts, and at 100, 100 parts. The solution deposits Na 2 CO 3 7 H 2 O at 50. Solutions of sodium carbonate have a strong alkaline reaction ; for, like potassium carbonate, the salt is hydrolyzed. Sodium carbonate melts at red heat, forming a clear liquid. Sodium bicarbonate NaHCO 3 , or sodium acid carbonate, is soluble in about 10 parts of water at 20. Its solutions have an alkaline reaction, for the salt is decomposed by hydrolysis to a slight extent. When warmed, the solutions give off car- bon dioxide, the carbonate being formed. Sodium bicarbonate is used in medicine, in baking powder, as saleratus, and in mak- ing soda water and other effervescent beverages. Sodium Nitrate NaNO 3 occurs in large quantities in Chili and Peru as Chili saltpeter, which also contains other salts of sodium, notably the iodate, sulphate, and chloride. The salt crystallizes in the rhombohedral division of the hexagonal system and melts at 318. In general, its chemical behavior is like that of potas- sium nitrate. Sodium nitrate is used as a fertilizer, also in the manufacture of potassium nitrate and nitric acid. The salt is somewhat hygroscopic, which unfits it for use in gunpowder. By heating sodium nitrate with lead or iron, sodium nitrite NaNO 2 is formed. This salt is much used in the coal-tar dye- stuff industry. Phosphates of Sodium. The most important of these is the secondary or disodium phosphate Na 2 HPO 4 12 H 2 O. This is THE ALKALI METALS 381 commonly called simply sodium phosphate. It is prepared by the action of phosphoric acid on sodium hydroxide or carbonate. The crystals effloresce. At 20, 100 parts of water dissolve 9.3 parts of the salt. The solution has a slightly alkaline reaction. On adding another equivalent of caustic soda to disodium phos- phate solution and evaporating to dryness, the tertiary sodium phosphate Na 3 PO 4 12 H 2 O is obtained. In aqueous solution, this does not exist, being hydrolyzed into disodium phosphate and sodium hydroxide. Primary sodium phosphate NaH 2 PO 4 -H 2 O ^ ou i Rft ; K OU bj. 1 4 f 30 fe ? - - . ~- .^ .. ji 2fe SO d . / ( y ( * y .0^ / b/o f 10 ] s x / \ ^ V -lc< 10 20 30 40 50 60 70 80 90 10 FIG. 127. Temperature in degrees C. has an acid reaction in solutions. On heating, the salt loses water and forms sodium metaphosphate NaPO 3 . Sodium Sulphate Na 2 SO 4 - 10 H 2 O, also called Glauber's salt, crystallizes in large monoclinic crystals that effloresce on expos- ure to the air. It is manufactured as one of the products of the Le Blanc soda process, also by Hargreave's process, which consists of passing air, sulphur dioxide, and steam over sodium chloride at high temperatures, thus : 2 NaCl + H 2 + O -f- SO 2 = Na 2 SO 4 + 2 HC1. At Stassfurt the salt is formed by the action of magnesium sulphate upon sodium chloride at low temperatures, when the sodium sulphate is deposited, while the magnesium chloride simultaneously forced remains in solution: MgSO 4 + 2 NaCl = Na 2 SO 4 + MgCl 3 . 382 OUTLINES OF CHEMISTRY The salt was first prepared by Glauber, whence its name. The solubility of Glauber's salt Na 2 SO 4 10 H 2 O increases with rise of temperature to 32.4, beyond which it decreases, for at higher temperatures the salt loses water, becoming anhydrous Na 2 SO 4 , which is less soluble than the decahydrate. Figure 127 shows the solubility curve, which has a sharp point of inflection at 32.4. This point is really the intersection of the solubility curve of the decahydrate and that of the anhydrous salt. Glauber's salt melts in its water of crystallization at 32.4. When completely melted, the solution so obtained may be cooled to room tempera- ture and even lower. This is a so-called supersaturated solu- tion, for it contains more salt than would be taken up at the lower temperature in presence of an excess of the solid salt. Indeed, if a crystal of the solid Na 2 SO 4 10 H 2 O is introduced into this supersaturated solution, the whole of it at once solidi- fies. Many other substances form similar supersaturated solu- tions. Glauber's salt simply furnishes an excellent illustration. Sodium sulphate occurs in many mineral waters. As Glauber's salt it is used as a purgative. On treating sodium sulphate with an equivalent quantity of sulphuric acid, acid sodium sulphate or sodium bisulphate NaHSO 4 H 2 O is obtained. It becomes anhydrous above 50 and melts at about 300. Its behavior is similar to that of potas- sium bisulphate. Sodium Sulphite Na 2 SO 3 7 H 2 O is formed by passing sulphur dioxide into a concentrated solution of sodium hydroxide or carbonate to neutrality. On saturating a strong solution of the sulphite with sulphur dioxide, sodium bisulphite NaHSO 3 is formed. This is frequently used as a source of sulphur dioxide in bleaching fabrics of silk, wool, etc. Sodium Thiosulphate Na 2 S 2 O 3 5 H 2 O is made by boiling sul- phur in a solution of sodium sulphite. It is also, though wrongly, called hyposulphite of soda or hypo. Its solution serves in photography as a " fixing bath," for it dissolves the excess of silver bromide on the photographic plate after the latter has been exposed to the light and developed. Sodium thiosulphate absorbs chlorine and is used as so-called " antichlo- rine," thus: 2 Na 2 S 2 3 + C1 2 = 2 NaCl + Na 2 S 4 O 6 . The salt is also used in chemical analysis. THE ALKALI METALS 388 The sulphides of sodium and sodium hydrosulphide are analo- gous to those of potassium and need no special description. Sodium Silicate, or sodium water glass, is made by fusing silica with sodium hydroxide or carbonate, or with' Glauber's salt and carbon. It is also prepared by boiling silica with concentrated solutions of sodium hydroxide under pressure. Sodium silicate may be obtained in form of monoclinic crystals of the composi- tion Na 2 SiO 3 8 H 2 O. Water glass comes in the market as a thick sirupy solution containing various sodium silicates, the composition averaging about Na 2 Si 4 O 9 . It is used in laundry soaps as a " filler," in making fabrics and wood fireproof, in the production of artificial stone, as a preservative for wood, and as a cement for glass, asbestus, mineral wool, etc. Sodium Cyanide NaCN is very similar to potassium cyanide, and is used for the same purposes. It is made commercially by the action of ammonia gas upon a mixture of carbon and metallic sodium, and is extremely soluble in water. Sodium Borate Na 2 B 4 O 7 10 H 2 O, or borax, has been described under boron. Lithium and its Compounds. Lithium is found in the miner- als, lepidolite or lithia mica, spodumene, triphylite, and a few others of rare occurrence. In very minute quantities lithium salts are also found widely distributed in soils, from which they pass into plants, a number, like tobacco and beets, being particu- larly prone to store up lithium. Many mineral waters contain lithium, though generally in rather small amounts. Lithium compounds are readily detected by means of the spectroscope: for they exhibit two characteristic red lines. To the Bunsen flame lithium salts impart a characteristic red color. Metallic lithium may be obtained by electrolyzing the fused chloride or a concentrated solution of the latter in pyridine. It has a very low specific gravity, 0.59, and hence floats even on petroleum oils. In general, its chemical behavior is similar to that of sodium, only less vigorous. The atomic weight of lithium is 6.94, and its valence is 1. Lithium chloride LiCl is deliquescent and extremely soluble in water. On the other hand, lithium carbonate Li 2 CO 3 is very slightly soluble, for 100 parts of water dissolve but 0.77 part at 15. Lithium phosphate Li 3 PO 4 is also but slightly soluble, 1 part in 2500 of water. This slight solubility of the 884 OUTLINES OF CHEMISTRY carbonate and phosphate distinguishes lithium sharply from the other alkali metals and shows its similarity to the alkaline earth metals. The chloride of lithium dissolves in pyridine, whereas the chlorides of all the other alkali metals are insoluble in that liquid, which fact is used in separating lithium from the other members of the group. Commercially, lithium carbonate is the most important of the lithium salts. With uric acid C 5 H 4 N 4 O 3 , lithium forms a moderately soluble salt, lithium urate C 5 H 2 N 4 O 3 Li 2 , upon which fact is based the administration of lithium carbonate in medi- cine, in cases of gout, which is caused by deposits of uric acid in the joints and muscles. The Alkali Metals as a Group. The alkali metals form a natural group, the properties of whose members exhibit a regu- lar variation with the atomic weights. The relations that obtain may be seen from the following table : ELEMENT ATOMIC WEIGHT MELTING POINT .SPECIFIC VOLTME Lithium I>i 6.94 1800 11 9 Sodium Net . 23 00 97 6 237 Potassium, K . . . . Rubidium, Ub . . . . Cesium Cs . . 39.10 85.45 13281 62.5 38.5 26 5 45.2 . 56.2 707 The gradation of chemical properties has already been men- tioned. It is evident that as the atomic weight increases the chemical activity of the members of the group also increases. Spectrum Analysis. A little of any sodium compound in- troduced into the Bunsen flame colors it a bright yellow. Similarly potassium compounds produce a violet and lithium compounds a characteristic red color. These colored flames may serve to detect the presence of small amounts of these metals. At the high temperature of the Bunseri flame, the salts of these and other metals are decomposed, and the incan- descent vapors emit light of definite wave length, i.e. of definite color, in each individual case. When sodium and potassium are both present in the flame, the latter appears yellow, for that color completely masks the delicate violet of the potassium THE ALKALI METALS 385 flame. The latter, however, can be detected by viewing the yellow flame through a blue glass or through a layer of a solu- tion of indigo, in which case the yellow light is absorbed and the violet color due to the potassium becomes apparent. Robert Bunsen sought to work out a system of detecting the presence of metals by noting the color which they impart to the flame. Later Bunsen and Kirchhoff (1859) passed the composite light through a glass prism, and thus the different colors could be FIG. 128. seen side by side instead of superimposed, for light of different color possesses different refrangibility. The instrument de- signed by Bunsen and Kirchhoff for producing and inspecting such spectra is called a spectroscope. Figure 128 shows a simple form of the instrument with the cover removed from the prism. The colored light from the flame enters a narrow adjust- able slit at $, and a lens or set of lenses in C gathers the rays and transmits them to the prism. After passing the latter, they enter the telescope T, which can be moved so as to catch the rays. Thus, if the flame is colored by sodium alone, the observer sees 386 OUTLINES OF CHEMISTRY E ~O O cog. CO Na Li K Rb Cs Ca Sr Ba Tl In A 9 1 on 1 irtft IA! >iri icU! ' | O ^H ^ ^j[ T w T"iL T w I TrT R ( d Oranj je Yellow Green Blue Initgo Vitlet i i i i i ' iii i i i l i i i 1 i i i i i iii 1 i i i i ii i i i iii i ii ii ! i i i i i ii i i i i i ! ! FIG. 129. THE ALKALI METALS 387 simply one yellow line, the image of the slit magnified, of course, by the lenses of the telescope. If now potassium be added to the flame, the yellow line due to the sodium remains, but there appear in addition two characteristic red lines to the left of the sodium line, while to the ex- treme right of the latter a violet line is found. The two red lines and the violet line are due to potassium. They always appear when potassium is present and are located in the same relative position to the sodium line and to one another. In order to observe more accu- rately the relative positions of these lines a scale is re- quired. In the tube IF a photographic scale is placed which is illuminated by a candle or other luminous flame as shown. Light from this scale strikes the prism at such an angle as to be reflected through the telescope to the eye. Thus, the scale and the spectrum are seen together. Every incandescent gas produces its own char- acteristic lines in the spectrum, and this fact constitutes the basis of spectrum analysis. Figure 129 exhibits the spectra of a number of common elements, the color of the lines being evident from the solar spectrum given kV at the head of the figure. Very minute quantities of many substances can still be detected by means of the spectroscope; so one three-millionth of a milligram of sodium will still show its characteristic yellow line in the spectrum. When a solid body is heated to incandescence in the flame, the spectrum observed is not a band spectrum, that is, it is not made up of a series of definite lines or bands, but consists of a continuous spectrum like that produced when sunlight passes through the slit. It is well known that sunlight passed through a prism is decomposed, yielding a continuous spectrum consisting of the colors of the rainbow. The light of the sun thus yields a continuous spectrum; this is, however, crossed by numerous dark lines called the Fraun- hofer lines. These dark lines are produced by light from the incandescent gases in the sun passing through the sun's atmos- phere, whicli absorbs the light emitted by the gases, thus leav- ing black lines wherever colored lines would have been. The 388 OUTLINES OF CHEMISTRY Fraunhofer lines, then, are absorption spectra. When, for exam- ple, sodium light is passed through the slit of the spectroscope, we get its characteristic yellow line ; but if the light is passed through an atmosphere of sodium vapor before entering the slit, we see a dark line just where the yellow line was before. This is because sodium vapor absorbs the light produced by 800 700 600 450 400 Rfd OrangeiYellow Gr$en 1 Blue Indie io Violet 1 1 i i I 1 He ! i 1 J i i i i i i Vapor II FIG. 131. incandescent sodium vapor, and thus we have the dark line, the so-called reversed spectrum. Now the Fraunhofer lines are similar lines, for the sun's atmosphere contains the vapors of the very substances whose incandescent gases near the sun are sending' out light. The principle established by Bunsen and Kirchhoff is that the vapors of a substance absorb the same light which the incandescent gases of the substances emit. Thus it has been possible to analyze the sun and other heavenly bod- ies that emit light of their own by means of the spectroscope. The result has been that it has been established that these bodies THE ALKALI METALS 389 contain practically the same elements that are found on the earth ; though in some cases fewer and in other cases more lines have been found in the spectra of celestial bodies than correspond to known terrestrial elements. By means of the spectroscope, a number of new elements have been discovered, among which are rubidium, caesium, thallium, indium, gallium, helium, neon, crypton, and xenon. Many metals whose salts cannot be vaporized in the Bunsen tiame are heated in the electric arc. Spectra so obtained are called spark spectra. Again, gases are inspected spectroscop- ically by introducing them into a tube (Fig. 130), provided with platinum or aluminum electrodes, and then exhausting with a vacuum pump till the gas has a pressure of but a fraction of a millimeter. When such a tube is connected with an induction coil, the gas emits light and can thus be examined with the spectroscope in the usual way. Figure 131 gives the spectra of a few simple gases. All colored solutions have their own characteristic absorption spectra. So if light from a Welsbach burner is passed through a colored solution and then through the slit of the spectroscope, it will appear that the continuous spectrum is crossed by black absorption bands characteristic of the solution. For example, blood yields the bands shown in Fig. 132, and thus it is clear that the spectroscope may be used to detect the presence of olood. Ammonium Salts. In their general behavior these are simi- ^a,r to the salts of sodium and potassium. All ammonium salts are volatile, however, when heated, decomposing into ammonia and the acid, or into other products of greater or less complex- ity. So when ammonium chloride NH 4 C1 is heated, it dissoci- ates into ammonia and hydrochloric acid; ammonium nitrite NH 4 NO 2 similarly yields nitrogen and water; ammonium ni- trate NH 4 NO 3 yields nitrous oxide and water; and ammonium oxalate (NH 4 ) 2 C 2 O 4 decomposes into ammonia, water, carbon monoxide, and carbon dioxide. Ammonium salts may be ob- RED YELLOW GREEN BLUE FIG. 132. 890 OUTLINES OF CHEMISTRY tained by neutralizing solutions of the acids with ammonia dis solved in water and evaporating to dryness. They may fre- quently also be formed by direct union of ammonia with the acid in question. When treated with an hydroxide of an alkali metal, ammonium salts are decomposed, ammonia being liberated. It has already been stated that ammonium chloride is ob- tained as a by-product in the manufacture of coal gas. Am- monium chloride, or sal ammoniac NH 4 C1, crystallizes in cubes and dissolves readily in water with absorption of heat. At 0, 100 parts of water dissolve 28 parts of the salt, and at 100, 73 parts. The salt has a very sharp, salty taste. Ammonium chloride is used in medicine, in making certain kinds of electric batteries, in the dyestuff industry, and, in general, as a source of ammonia. Ammonium bromide and ammonium iodide are deliquescent salts that also crystallize in cubes. They slowly decompose on exposure to the air. Ammonium sulphate (NH 4 ) 2 SO 4 is a very common ammonium salt, being generally manufactured by neutralizing the gas liquors with sulphuric acid. It crystallizes in the rhombic system, and is soluble in about 1.5 parts of cold water. The salt is used as a source of ammonia for making other compounds, and also as a fertilizer. Over 600,000 tons of this salt are manufactured annually. On electrolysis of a concentrated solution of ammonium bisul- phate NH 4 HSO 4 , ammonium persulphate (NH 4 ) 2 S 2 O 8 is pro- duced. The latter forms monoclinic crystals that separate out from the solution. It serves as an oxidizing agent. Ammonium sulphide (NH 4 ) 2 S is obtained by passing hydro- gen sulphide into a solution of ammonia in water, till the latter is half saturated, thus : 2NH 3 + H 2 S = (NH 4 ) 2 S. On passing the hydrogen sulphide into the solution till it is completely saturated, ammonium hydrosulphide NH 4 SH results : (NH 4 ) 2 S + H 2 S = 2 NH 4 SH. The sulphide (NH 4 ) 2 S may be obtained in form of colorless needles which readily give off ammonia and thus pass over into the hydrosulphide NH 4 SH, which also forms colorless crystals THE ALKALI METALS 391 that dissociate into ammonia and hydrogen sulphide even at room temperatures. At 50 this dissociation is nearly complete. Almost invariably aqueous solutions of ammonium sulphide or hydrosulphide are prepared for laboratory use as above stated. These are important in analytical chemistry. Ammonium sul- phide solutions are colorless when freshly prepared. They have a disagreeable odor, due to the fact that by hydrolysis both ammonia and hydrogen sulphide are liberated. On standing in the air, the solutions turn yellow on account of oxidation. The sulphur thus liberated remains in solution, forming a yellow polysulphide. By dissolving sulphur in ammonium sulphide solutions, a series of polysulphides may be obtained which are analogous to the polysulphides of the alkali metals. The solu- tions of these polysulphides are commonly termed yellow ammo- nium sulphide. It serves in analytical chemistry for dissolving the sulphides of arsenic, antimony, tin, gold, and platinum, with which it forms ammonium sulpho-salts. Ammonium nitrate NH 4 NO 3 forms rhombic prisms that are isomorphous with potassium nitrate. It melts at about 160, and on further heating, it decomposes into water and nitrous oxide. In water it dissolves readily with absorption of heat, and it is consequently sometimes used, mixed with ice, to pro- duce low temperatures. The salt is also employed in explosives in place of potassium nitrate. Ammonium nitrite forms deli- quescent crystals that readily decompose into water and nitro- gen, even in aqueous solutions, on being heated to 70. Ammonium carbonate (NH 4 ) 2 CO 3 + H 2 O is obtained in form of a crystalline precipitate by passing carbon dioxide into a concentrated solution of commercial ammonium carbonate. The latter consists of a mixture of acid ammonium carbonate NFI 4 HCO 3 and ammonium carbamate NH 2 CO 2 NH 4 , and is obtained by heating either ammonium sulphate or ammonium chloride with calcium carbonate. The sublimate formed is a hard white mass. The normal salt, (NH 4 ) 2 CO 3 , readily loses ammonia and passes over into the acid salt, NH 4 HCO 3 , which forms crystals that decompose into ammonia, carbon dioxide, and water at 60. In aqueous solutions, the salt loses carbon dioxide, thus forming the normal carbonate. Detection of Ammonium Salts. These salts are characterized by their volatility and the fact that ammonia is evolved by 392 OUTLINES OF CHEMISTRY treating them with caustic alkalies. With ammonium chloride, platinic chloride forms ammonium platinic chloride (NH 4 ) 2 PtCl 6 , which, like the analogous potassium salt, is sparingly soluble. Acid ammonium tartrate C 6 H 5 O 6 . NH 4 is precipitated from concentrated solutions of ammonium salts by means of tartaric acid. REVIEW QUESTIONS 1. Name the alkali metals. Why are they so called? Who dis- covered them? How? 2. Describe metallic potassium. Discuss the occurrence of potas- sium in : (a) sea water, (6) the earth's crust, (c) plants and animals. 3. Why is potassium hydroxide called caustic potash? Mention five other caustic alkalies, giving their formulas. 4. Explain what is meant by the term "fixed alkalies." 5. Make a list of the halides, nitrates, sulphates, sulphides, oxides, hydroxides, phosphates, and arsenates of sodium, potassium, and ammo- nium, giving the formula of each compound and arranging them so as to show their chemical relationships. 6. How demonstrate that a carrot contains potassium? How show that a seaweed contains sodium? 7. What is the maximum amount of potassium carbonate that could be prepared from 100 tons of carnallite? Give the chemical equations illustrating the steps in the process. 8. What use is made of each of the following compounds : potas- sium iodide, potassium carbonate, potassium cyanide, potassium chlorate, potassium nitrate, potassium bisulphate. 9. Characterize the compounds of rubidium and caesium. How were these elements discovered, and by whom? 10. Where is sodium chloride found ? By what processes may sodium chloride be changed into sodium carbonate? Write the appropriate equations showing the steps in these processes. 11. Discuss the uses of common salt. 12. How prepare caustic soda from common salt ? Equations. How much pure caustic soda could be prepared from 250 kilograms of common salt? 13. What is saleratus? Write the equation, expressing what happens when it is used in making biscuits. 14. How may sodium silicate be made? Equation. What use is made of this compound ? 15. Where does lithium occur in nature? Mention the chief charac- teristics of lithium compounds. What use is made of 'lithium carbonate and lithium citrate? THE ALKALI METALS 393 16. Why is the spectroscope of special value in detecting the alkali metals? 17. Explain fully what evidence we have that there is sodium in the sun. 18. In what respects are ammonium salts similar to those of the alkalies? In what respects are they different? 19. How detect the presence of an ammonium salt when it, is mixed with salts of the alkalies? 20. In making a cream of tartar baking powder, how much baking soda is required to every five pounds of cream of tartar ? CHAPTER XXII THE ALKALINE EARTH METALS THE metals of the alkaline earths are calcium, strontium, and barium. They form another natural group of closely related elements. These metals never occur in the free state in nature. They are harder than the alkali metals, have higher atomic weights, and do not melt below red heat. They act on water, yielding hydroxides and hydrogen, though the action is less vigorous than in the case of the alkali metals. The hydroxides formed are alkaline and rather sparingly soluble in water, the solubility increasing as the atomic weight of the metal increases. On heating the hydroxides, they lose water, forming the oxides, which are white, earthy powders that give an alkaline reaction with moist litmus paper. This dehydration is accomplished most readily in the case of calcium hydroxide, and least readily in the case of barium hydroxide ; that is, the stability of the hydroxides increases with the atomic weight of the metal. In all their compounds calcium, strontium, and barium are ii bivalent. The chlorides, bromides, and iodides, MX 9 , the ii ii nitrates, M(NO 3 ) 2 , and the acetates, M(C 2 H 3 O 2 ) 2 , are readily soluble in water ; whereas the sulphates, MS(X, phosphates, ii n ii n M 3 (PO 4 ) 2 , carbonates, MCO 3 , silicates, MSiO 3 , oxalates, MC 2 O 4 , and fluorides, MF 2 , are sparingly soluble in water. It will be recalled that the carbonate and phosphate of lithium are also but slightly soluble, and thus lithium in a way represents a transition between the alkali metals and those of the alkaline earths. The chlorides of the alkaline earth metals are more soluble than the nitrates. The solubility of the sulphates decreases as the atomic weight- increases. The insoluble car- bonates, sulphates, phosphates, and silicates are specially charac- teristic of this group. They are of importance in nature, in the arts and industries, and in chemical analysis. The acid car- bonates or bicarbonates are much more soluble than the carbonates. 394 THE ALKALINE EARTH METALS 395 Occurrence, Preparation, and Properties of Calcium. This metal occurs very widely distributed and often in enormous masses as carbonate in form of marble, chalk, or limestone. It is further found as sulphate in form of gypsum and anhydrite, as phosphate, as fluoride or fluorspar, and as an essential constit- uent of many silicates. Nearly all natural waters contain calcium sulphate and bicarbonate. The bones and teeth of animals consist mainly of calcium phosphate together with some carbonate and small amounts of fluoride. In eggshells, coral, and the shells of molluscs and various crustaceans the carbon- ate of calcium predominates. Calcium salts are also found in plants. These are the sulphate, oxalate, phosphate, and car bonate, as well as salts of various complex organic acids. Similarly calcium compounds are distributed throughout the bodies of animals in small quantities. Metallic calcium may be obtained by heating calcium iodide with sodium, or by heating an excess of calcium oxide with carbon or calcium carbide in the electric furnace. The best way to prepare the metal is by electroly- sis of the molten chloride. The chloride is placed in a carbon con- tainer (Fig. 133), the walls of which serve as anode (see elec- trolysis). The cathode is of iron or copper. After the electrolysis has started, the heat developed by the current is sufficient to keep the salt in molten condition. The metal separates out at the cathode, and, being light, rises to the top of the molten chloride. By slowly raising the cathode as the elec- trolysis proceeds, a rough stick of calcium is obtained, for the metal adheres to the electrode. Metallic calcium may now be purchased at less than a dollar a pound. i .'Calcium is a silver-white metal of specific gravity 1.85. It crystallizes in the hexagonal system and is fairly hard, tough, and malleable. It may leadily be worked in a lathe FIG. 133. 396 OUTLINES OF CHEMISTRY It decomposes water, and consequently is kept under petroleum, or more frequently simply in air-tight containers of glass or tinned iron. Calcium melts at about 760, and at that tem- perature catches fire in the air, burning to the oxide, CaO, and the nitride, Ca 3 N 2 . The latter is a yellow powder which is de- composed by water, yielding the hydroxide and ammonia. Calcium was formerly described as a yellow metal. The yellow color was due to the presence of calcium nitride as an impurity. With hydrogen, calcium readily forms calcium hydride CaH 2 , a white powder that acts more vigorously on water than the metal itself. In general, calcium is very active chemically, for it unites with all the elements except those of the argon group. Calcium Carbonate CaCO 3 is the most abundant of all the cal- cium compounds. In an impure form, as limestone, it forms mountains and strata of vast extent and great thickness. Dolo- mite is essentially a calcium, magnesium limestone of the com- position MgCO 3 CaCO 3 . It generally contains silica, iron, alumina, and other impurities. Marl consists of limestone mixed with clay. Chalk is fairly pure calcium carbonate. In a crystalline state calcium carbonate occurs as marble in many localities. In pure form it occurs as calcite, particularly as Ice- land spar in Iceland. The stalactites and stalagmites found in many caves consist of calcium carbonate. Calcium carbonate crystallizes in the hexagonal system as calcite, commonly form- ing rhombohedra (Figs. 65 and 66) or scalenohedra (Figs. 64 and 67), and also in the orthorhombic system as aragonite. The hexagonal form is the stable one at ordinary temperatures ; at higher temperatures the rhombic form is stabler. So from a cold solution calcium carbonate deposits in hexagonal, and from a hot solutio-n in rhombic, crystals. Calcium carbonate is prac- tically insoluble in water; but in water charged with carbon dioxide it dissolves fairly readily. The solution contains cal- cium bicarbonate Ca(HCO 3 ) 2 , in all probability. On boiling such solutions, carbon dioxide is expelled and the normal carbonate is precipitated. Waters containing calcium salts in solution are said to be "hard. The hardness that can be dis- pelled by boiling, as just mentioned, is termed temporary hard- ness, as compared with permanent hardness which is produced by the presence of calcium salts other than the bicarbonate, and consequently persists even on boiling. THE ALKALINE EARTH METALS 397 Limestone and marble are used as building stones, in glass manufacture, in the reduction of iron ores, and in making lime, Portland cement, sodium carbonate, and many other products Much calcium carbonate is also used as chalk and whiting. Mixed with linseed oil, calcium carbonate forms putty. Calcium Oxide CaO, lime, is formed by heating calcium car- bonate above 600. To obtain the pure oxide, pure calcium carbonate is strongly ignited ; whereas on a commercial scale, lime is prepared by heating limestone in limekilns (Fig. 134). Lime is a white, amorphous, porous solid. It may be melted in the electric furnace. Lime unites with water with evolu- tion of heat, forming a powder, slaked lime or calcium hydroxide Ca(OH) 2 . The unslaked oxide, CaO, is called quicklime or FIG. 134. caustic lime. Calcium hydrox- ide is somewhat soluble in water. At 15, 100 parts of water dissolve about 0.14 part of the hydroxide, while at 100 but half of that amount dis- solves. The solution is alka- line and is known as limewater, when clear. When it contains excess of hydroxide in suspen- sion, it is termed milk of lime. When exposed to the air, quicklime gradually absorbs moisture and carbon dioxide, and crumbles, being converted into calcium carbonate or air-slaked lime. Besides being used in making mortar for building purposes, lime is generally employed in chemical industries when a cheap base is required. Large amounts are used in purifying coal gas and sugar, in removing hair from hides, in making bleaching powder, sodium and potassium hydroxides, glass, and oxalic, tartaric, and citric acids. Lime is also used as a disinfectant, and limewater is frequently employed in medicine. Mortar consists of a pasty mass obtained by mixing sand, slaked lime, and water. After it has been applied, water dries out gradually and carbon dioxide is absorbed from the air, thus 398 OUTLINES OF CHEMISTRY forming hard calcium carbonate. This process is called the setting of the mortar. It will not take place while the mortar remains wet, and will commonly require a rather long time for its completion ; for after the outer layer has become trans- formed to carbonate, the deeper layers are partially protected from the air and so are altered but slowly. When thoroughly hardened, the sand grains are firmly fixed in the matrix of crystalline calcium carbonate, which adheres well to the brick or stone used. Lime mortar is unfit for use in places that are always wet, for it hardens only when dry. After hundreds of years, some calcium silicate is formed by interaction of the lime with the sand grains. Lime mortar has been in common use for a very long time. It was quite generally employed by the Romans in their buildings. Lime made from magnesian limestone, dolomite, consists of CaO and MgO. It slakes very slowly with cold water, owing to the presence of the magnesia. In cold weather, hot water is commonly employed in slaking this lime, which, however, makes very good mortar and hence is used in many localities. Cement. When limestone containing silicates of aluminum is heated in a kiln and the product ground to a powder, the latter forms a so-called hydraulic cement, for it will unite with water and form a hard, insoluble mass. The hardening process takes place uniformly, and relatively quickly throughout the mass even under water. The cement is very valuable, for it can be used in damp as well as in dry places. In some locali- ties, as near Milwaukee and Louisville, limestone containing a suitable amount of aluminum silicates for producing cement is found and made into cement. Such cements are called natural cements. They generally contain notable quantities of magnesia and other ingredients. The following table gives the composition of Louisville cement in per cent : SiO 2 . . . . . 21 1 CaSO 6 8 FeJojJ ' CaO . . ... 7.5 . . . 44.4 K 2 O 1 CO ... 0.8 11.2 MffO 7.0 PLO . 1.2 Portland cement is made by artificially mixing materials con- taining silica, lime, and alumina in proper proportions and then THE ALKALINE EARTH METALS 399 firing the mixture in a kiln. The hard mass or clinker thus obtained is ground to a fine powder and constitutes the so-called Portland cement. In practice, clay rich in silica, and calcium carbonate, are used in making this cement. Often marl or the slag from blast furnaces is employed. In all cases, it is neces- sary to determine the composition of the materials used by chemical analysis, so that they may be mixed in the proper proportions of silica, alumina, and lime. The following table gives the composition of Portland cement in per cent. The last column indicates the result of an analysis of a typical American Portland cement: i COMPOSITION OF PORTLAND CEMENT IN PER CENT SiO 2 . . . . . . . 20 to 25 22.5 CaO . . . . . . . 58 to 65 62.9 A1 2 8 . . . . . . . 5 to 10 traces to 5 6.4 3.1 Fe O a . . . . . 2 to 5 3.3 K 2 + Na 2 H 2 O + CO 2 . SO, . . . . traces to 3 . . . traces to 2 0.7 1.0 99.9 No definite chemical formula can be ascribed to Portland cement, the composition of excellent cements varying within the limits indicated in the first column above. Experience has shown that the ratio of the amount of lime to that of silica plus alumina and ferric oxide should fall between 1.8 and 2.2, being about 2 on the average. That is, _ CaO _ SiO 2 + A1 2 O 3 + Fe 2 O 3 should be greater than 1.8 and less than 2.2; for the cement whose analysis is given in the table above, the ratio is = 1.95. 22.5+6.4+3.3 The chemical changes involved in the making and hardening of cement have been the subject of much investigation and discussion. The clinker probably contains mainly tricalcium 400 OUTLINES OF CHEMISTRY silicate Ca 3 SiO 5 and calcium aluminate Ca 3 Al 2 O 6 . These when ground fine and then treated with water probably suffer decom- position thus : 2 Ca 8 SiO 6 + 9 H 2 O = 4 Ca(OH) 2 + (CaSiO 3 ) 2 . 5 H 2 O, and Ca 3 Al 2 6 + Ca(OH) 2 + 11 H 2 O = Ca 4 Al 2 O 7 . 12 H 2 O, so that the hardened cement contains the hydrated calcium silicate (CaSiO 3 ) 2 - 5 H 2 O and the hydrated basic calcium alu- minate Ca 4 Al 2 O 7 12 H 2 O. The hardening of the cement is supposed to be due mainly to the formation of the former compound. Cement powder has a greenish gray color. Its specific gravity is 3.1 to 3.2. The hardened mass has a drab color, resembling the rock found at Portland, England, whence the name of the cement. Mixed with crushed stone and water in proper proportions, Portland cement hardens into a mass called concrete, which wears as well as excellent stone. Often concrete is strengthened by imbedding in it rods of iron or steel. The product is called reenforced concrete. Plain and reenforced concrete are much used in modern structures, and the production of cement has greatly increased in recent years. Over 90,000,000 barrels of Portland cement are made in the United States each year, and the demand for the product is still growing. Calcium Sulphate CaSO 4 occurs in large quantities in nature as gypsum CaSO 4 2 H 2 O, which forms monoclinic crystals (Fig. 73) called selenite, and also as anhydrite CaSO 4 in rhom- bic forms. Alabaster is a granular, crystalline form of gypsum. Calcium sulphate occurs in soils and natural waters. At 0, 100 parts of water dissolve 0.19 part of calcium sulphate; at 35, 0.21 part. In nitric or hydrochloric acids and in many salt solutions calcium sulphate dissolves much more copiously. At about 110, gypsum loses water, forming (CaSO 4 ) 2 H 2 O, which is a white powder known as plaster of Paris. When this is mixed with water, a paste may be obtained which soon hardens. This hardening or " setting " is due to the fact that water is again taken up, a crystalline, coherent mass of gypsum being formed, thus : (CaSO 4 ) 2 H 2 O + 3 H 2 = 2 (CaSO 4 - 2 H 2 O). Plaster of Paris is much used for making casts, surgical band ages, and stucco. THE ALKALINE EARTH METALS 401 At about 200, gypsum loses all of its water, and is then said to be dead burned, for in this condition it unites with water but slowly and without hardening. Gypsum is often used as a fertilizer, land plaster. Its action probably depends on the fact that it reacts with the ammonium carbonate in soils, form- ing ammonium sulphate, which, being practically non-volatile, remains in the soil and is utilized by plants. Calcium Sulphite CaSO 3 is formed by passing sulphur dioxide into calcium hydroxide. The salt crystallizes in prisms of the composition CaSO 3 2 H 2 O, which are slightly soluble in water, 1 in 800. In aqueous solutions of sulphur dioxide, the salt dis- solves more copiously, and such solutions are used in paper mills in preparing wood pulp. Calcium Sulphide CaS is obtained by heating calcium sul- phate with charcoal : CaSO 4 + 4 C = CaS + 4 CO. With water the sulphide reacts thus : 2 CaS + 2 H 2 = Ca(OH) 2 + Ca(SH) 2 , the latter compound being soluble in water. Calcium sulphide is used in making luminous match safes, clock faces, etc., for, after exposure to sunlight, it emits a faint light which is visible in the dark. Barium sulphide BaS and strontium sulphide SrS serve similarly for making so-called luminous paint. Calcium Fluoride CaF 2 crystallizes in cubes, and is found in nature as fluorite. It is insoluble in water. It serves as a flux, and is used in making hydrofluoric acid and other fluorine com- pounds. Calcium Chloride CaCl 2 occurs at Stassfurt in tachhydrite CaCl 2 MgCl 2 12 H 2 O. It is obtained as a by-product in the Solvay soda process and in the production of ammonia by the action of lime on ammonium chloride. It is also made by the action of hydrochloric acid on calcium carbonate or hydroxide. From solutions it crystallizes in hexagonal prisms, CaCl 2 6 H 2 O. These melt in their crystal water at 29, and become a porous, anhydrous mass at 200. The anhydrous salt melts at 719. It is deliquescent, and is much used as a dry- ing agent in chemical laboratories. With ice and the hydrate, CaCl 2 6 H 2 O, temperatures as -low as 50 may be produced. 402 OUTLINES OF CHEMISTRY The anhydrous salt dissolves in water with liberation of heat. With alcohol, and also with ammonia, calcium chloride forms addition products, so that these substances must be dried with other agents like the oxide of calcium or barium. Calcium bro- mide CaBr 2 and calcium iodide CaI 2 are even more deliquescent than calcium chloride. Bleaching Powder or chloride of lime is made in large quanti- ties by passing chlorine into calcium hydroxide, i.e. slaked lime. The composition of the compound is expressed by the formula Ca(OCl)Cl, as explained under chlorine, where the reactions involved in its preparation and use are also de- scribed. Bleaching powder is slightly yellowish in color. It absorbs carbon dioxide and moisture from the air. Thus hypo- chlorous acid is formed, to which the odor of bleaching powder is due. Enormous quantities of bleach, as it is also called, are used in paper making and in the manufacture of cotton and linen goods. Calcium Phosphate Gu 3 (PO 4 ) 2 is found in nature as already stated, in apatite, and in the bones of animals. By treating a solution of calcium chloride with sodium ammonium phosphate, calcium phosphate is precipitated, thus : 2 Na 2 NH 4 PO 4 + 3 CaCl 2 = 4 NaCl + 2 NH 4 C1 + Ca 3 (PO 4 ) 2 . Calcium phosphate is practically insoluble in water, but in acids it dissolves readily, also in many solutions of salts, like chlorides or nitrates of the alkalies. As the latter are present- in soils, calcium phosphate is dissolved by their solutions and hence made available to plants. Calcium phosphate is necessary to plant life, and also to the life of animals, into whose bones it enters as a prime constituent. So-called superphosphate of lime, a fertilizer of great value, consists of calcium sulphate and primary calcium phosphate produced by the action of sulphuric acid on calcium phosphate, thus : Ca 3 (P0 4 ) 2 + 2 H 2 S0 4 = CaH 4 (P0 4 ) 2 + 2 CaSO 4 . The primary calcium phosphate CaH 4 (PO 4 ) 2 is soluble in water and hence is readily available to plants. Each year the United States produces about 3,000,000 tons of phosphate rock, most of which is used as fertilizer. The most important beds of this rock are in South Carolina and Florida. As this material is of prime importance in maintaining the fer- THE ALKALINE EARTH METALS 403 tility of the soil, its exportation has recently been forbidden by law. Calcium Carbide CaC 2 is made by heating lime or calcium carbonate with carbon in the electric furnace, thus : Pure calcium carbide is white, but the commercial article is dark in color, owing to the presence of impurities. The sub- stance yields acetylene when treated with water, as already stated, and hence large quantities of it are manufactured annually. Calcium Phosphide Ca 2 P 2 is formed by heating lime and phosphorus together, thus : 14 CaO + 14 P = 2 Ca a P 2 7 + 5 Ca 2 P 2 . The product is a brown solid, which on treatment with water yields calcium hydroxide and phosphine. Calcium Cyanamide CaN CN is formed by passing nitrogen over calcium carbide heated to white heat in an electric furnace, thus : CaC 2 + N 2 = CaN CN + C. Calcium cyanamide is used as a fertilizer, for in the soil its nitrogen gradually is converted into ammonia and nitrates, owing to the action of water and oxygen from the air. Calcium Silicide CaSi 2 is produced by heating lime with silicon in the electric furnace. It forms hexagonal crystals that react but slowly with water. Dilute acids decompose the silicide readily. Calcium Silicate CaSiO 8 is occasionally found in nature in monoclinic crystals as wollastonite. It occurs very frequently in complex silicates like mica, feldspar, garnet, hornblende, and many ethers. Calcium silicate may be formed by heating together silica and lime or calcium carbonate, also by adding sodium silicate to a solution of a calcium salt, thus : CaCl 2 + Na 2 SiO 3 = CaSiO 8 + 2 NaCl. Calcium silicate is practically insoluble in water ; but in hydro- chloric acid it dissolves, being decomposed into calcium chloride and silicic acicl. Glass is a mixture of the silicates of sodium and calcium. Sodium silicate is transparent and soluble in water, while 404 OUTLINES OF CHEMISTRY calcium silicate is neither transparent nor soluble in water. When sodium silicate is fused together with calcium silicate in proper proportion, a liquid results which on cooling forms ordinary glass. This is practically insoluble in both water and acids, except hydrofluoric acid. Ordinary glass varies somewhat in composition. It is a mixture of silicates of sodium and calcium which approximates the composition Na 2 O 3 SiO 2 + CaO 3 SiO 2 . This soda-lime glass, as it is also called, is used for windows and various ordinary vessels. It melts readily and is consequently relatively easy to work into desired forms. In making glass, finely ground and intimately mixed quartz sand SiO 2 , calcium car- bonate CaCO 3 , and soda Na 2 CO 3 are melted together in large pots of special fire clay. These pots are about 4 feet high and 4 feet in di- ameter ; Fig. 135 shows an open and a closed form. During this process carbon dioxide escapes. The temperature is raised after the mass has melted, to expel all gases. The scum and various impurities that have gathered on top of the molten material are removed mechanically. In making plate glass, the molten, viscous material is poured upon large iron or bronze plates and rolled to the desired thickness with hot iron rollers. The plates are afterwards ground and polished. For ordinary window glass, the glass is first blown into cylindrical forms which are cut open lengthwise and then flattened out. Heavy glass dishes and many other similar articles are made by pressing the plastic glass into suitable molds. Bottles are made by taking glass on the end of an iron pipe and blowing it into molds. All flasks, retorts, beakers, and many other thin glass dishes are blown into suitable molds. Glass tubing is made by blowing a small bulb on the end of an iron blowpipe, and then attaching the bulb at its lower end and drawing the plastic mass out into the form of a tube. All ylass articles must be annealed. This FIG. 135. THE ALKALINE EARTH METALS 405 process consists of cooling the glass gradually in suitable fur- naces. Glass suddenly cooled is under internal strain, and a slight scratch will cause it to break ; thus the "Prince Rupert's drop " of suddenly chilled glass is shattered by nipping off its tail. The strength and transparency of glass depends upon its amorphous nature, for when glass crystallizes it becomes opaque and brittle. Since the composition of glass may be varied to a considerable extent at will, it is of the nature of a solution; and indeed it is sometimes called a solid solution, resulting from the supercooling of the liquid mass. Since sodium silicate is soluble in water, and calcium silicate is readily decomposed by hydrochloric acid, and on the other hand glass is practically not attacked by water or hydrochloric acid, it is clear that in glass the silicates of sodium and calcium are chemically com- bined with each other, in spite of the fact that their relative proportions may gradually be varied to some extent. Soda-lime glass is unsuitable for laboratory utensils, for it is too easily attacked by chemical reagents, especially by alkalies. For this reason potash-lime glass, in the manufacture of which potassium carbonate is used instead of sodium carbonate, is commonly employed in making finer glassware, for it is much harder and less readily acted upon by reagents. This potash glass is also called Bohemian glass or hard glass ; it has a much higher melting point than ordinary soda glass, which is fre- quently called soft glass. Crown glass is a potash-lime glass. Flint glass is produced by melting together silica, potash, and lead oxide. It has a high specific gravity and also a high index of refraction, hence it is used in making lenses, prisms, and other optical instruments as well as ornaments. Cut glass is made by grinding and polishing flint glass. Jena glass, which is now much used for utensils, contains boric anhydride. Colored glass is made by adding small quantities of various metallic oxides to molten glass. In this way colored silicates are formed. Thus cobalt oxide colors glass blue ; chromic oxide produces a bright green shade ; manganese dioxide yields a violet color ; cuprous oxide or gold produces a red shade owing to the very finely suspended particles of these substances in the glass. The ordinary green bottle glass is colored by ferrous oxide, while the brown or brownish yellow glass is colored by ferric oxide. Window glass also contains small amounts of 406 OUTLINES OF CHEMISTRY iron, which produces the green color observed when viewing the glass 011 edge. Black or very dark glass is produced by large amounts of cobalt, iron, and other oxides. Bottle glass is commonly made from impure, cheap materials that contain relatively large amounts of iron and other oxides that color the glass. Enamel glass is a lead glass to which oxide of tin has been added, while milk glass contains calcium phosphate in suspension, which produces the characteristic white, opaque appearance. The following table gives the approximate percentage com- position of a few typical kinds of glass. Si0 2 Na 2 O K 2 CaO PbO Quartz glass .... 100 Common glass . . . 75,5 12.9 11.6 Bohemian or crown glass 70.8 18.3 10.9 Crystal or flint glass 53.5 13.8 32.7 The art of making glass is very old. It is supposed to have originated with the Phoenicians. The Egyptians practiced it long before the Christian era. From the thirteenth to the seventeenth century Venice was noted for its glass manufac- tures. The general introduction of the use of window glass dates from about the sixteenth century. Bohemian glass was placed in the market as early as the fifteenth century, though glass was manufactured in Germany and Bohemia even in the tenth century. Occurrence, Preparation, and Properties of Strontium. Stron- tium occurs in nature as celestite SrSO 4 and strontianite SrCO 3 . It is of much rarer occurrence than either calcium or barium. The metal, which has properties similar to those of calcium, may be obtained by electrolysis of the molten chloride, SrCl 2 . Metallic strontium acts on water, liberating hydrogen and forming strontium hydroxide. In the air, strontium is slowly oxidized. The metal is silver white and plastic like lead. It melts at about 800 and volatilizes readily at 950. Strontium Compounds. These are in general very similar to the compounds of calcium. They are prepared by treating the carbonate with acids. The native sulphate serves as the THE ALKALINE EARTH METALS 407 chief source of strontium salts. It is first reduced to sulphide by heating with charcoal and then treated with acids. Stron- tium salts impart a brilliant red color to the Bunsen flame. The spectrum contains a blue line, an orange line, and six bright lines in the red. Strontium carbonate is not as readily decomposed into strontium oxide SrO and carbon dioxide as calcium carbonate. With water, the oxide readily forms stron- tium hydroxide Sr(OH) 2 , which is more soluble than calcium hydroxide. With sugar, strontium hydroxide forms insoluble compounds, hence it is used in extracting sugar from molasses that will 110 longer yield crystals of sugar. Strontium dioxide SrO 2 is also known. Strontium chloride SrCl 2 is hygroscopic, but it does not pos- sess this property to as marked a degree as calcium chloride. The salt is isomorphous with calcium chloride, forming hexag- onal prisms of the composition SrCl 2 + 6 H 2 O. Strontium sulphate SrSO 4 is precipitated from solutions of strontium salts by adding a solution of a soluble sulphate. It dissolves in about 7000 parts of water. In alcoholic solutions it is much less soluble. On boiling strontium sulphate with sodium or potassium carbonate, the salt is transformed into the carbonate : SrSO 4 + Na 2 CO 3 = Na 2 SO 4 + SrCO 8 . Under like treatment, barium carbonate remains practically unchanged. Strontium nitrate Sr(NO 3 ) 2 forms octahedra or cubes when crystallized from concentrated solutions. .At low temperatures the hydrate, Sr(NO 3 ) 2 4 H 2 O, separates from solutions in mono- clinic crystals. These effloresce on exposure to the air. Though readily soluble in water, strontium nitrate is insoluble in alcohol. When heated, the nitrate is decomposed, yielding the oxide. Strontium nitrate is much used in fireworks and red Bengal lights. The latter usually consist of a mixture of about 50 parts stron- tium nitrate, 30 parts shellac, and 20 parts potassium chlorate. Besides being used in the sugar industry, strontium salts are sometimes prescribed in medicine. Occurrence, Preparation, and Properties of Barium. Barium occurs as witherite, the native carbonate, BaCO 8 , and as barite, or heavy spar, the native sulphate, BaSO 4 . These serve as sources for the preparation of barium compounds. Metallic CALIFORNIA COLLEGE of PHARMACY . 408 OUTLINES OF CHEMISTRY barium may be obtained as a silver- white metal by electrolysis of the molten chloride. In general, its properties are like those of strontium and calcium. Barium has a specific gravity of 3.78. It melts at about 850 and boils at 1150. Compounds of Barium. These are analogous to those of calcium and strontium. They are, however, poisonous in character. Barium oxide BaO is best prepared by heating the nitrate, for the carbonate requires even a higher temperature to decompose it than strontium carbonate. On treating the oxide with water, barium hydroxide Ba(OH) 2 is formed, which is more soluble than calcium hydroxide. The solutions of barium hydroxide are alkaline and yield a precipitate of barium carbonate BaCO 3 on treatment with carbon dioxide. Barium hydroxide is frequently used in place of limewater in testing for carbon dioxide. On heating barium oxide in the air or in oxygen, barium dioxide BaO 2 is formed. This interesting substance, dis- covered by Thenard and Gay-Lussac, is a grayish white powder which is insoluble in water. When barium dioxide is heated to high temperatures, it is converted into barium oxide and oxygen. This affords a method of preparing oxygen from the air, as already mentioned. Barium dioxide is used in prepar- ing hydrogen peroxide. By treating a solution of barium hydroxide with hydrogen peroxide a precipitate of the compo- sition BaO 2 8 H 2 O is obtained which loses water at 130, and yields barium dioxide. Barium chloride BaCl 2 2 H 2 O, prepared by treating either the carbonate or the sulphide with hydrochloric acid, forms rhombic prisms soluble in about 3 parts of water. The salt is not deliquescent. It is practically insoluble in alcohol. This salt is then less soluble than' the chloride of strontium, and the lat- ter is less soluble than calcium chloride. Barium chloride has a bitter taste, and is a fairly strong poison. Barium fluoride BaF 2 is practically insoluble in water, but readily soluble in acids. Barium bromide BaBr 2 and barium iodide BaI 2 are soluble in water, and also in alcohol. Barium nitrate Ba(N0 3 ) 2 forms anhydrous crystals of the regular system. It is soluble in about 12 parts of water, and hence is precipitated by adding sodium nitrate to a concentrated solution of barium chloride. It may also be obtained by the THE ALKALINE EARTH METALS 409 action of nitric acid on the carbonate, hydroxide, or sulphide. On heating, it decomposes, yielding the oxide, oxygen, and oxides of nitrogen. Barium nitrate is used in making green Bengal lights. Barium carbonate BaCO 3 occurs in rhombic crystals as with- erite, and is prepared commercially by heating the natural sul- phate or barium sulphide with sodium carbonate. The product consequently is contaminated with sodium carbonate, which it is practically impossible to remove by washing. Pure barium carbonate may be obtained by adding ammonium carbonate to a solution of barium chloride. It requires white heat (1500 C.) to decompose barium carbonate, but by heating it with carbon it may readily be reduced to barium oxide. Barium sulphate BaSO 4 is the chief source of barium. In nature it is found in compact masses or rhombic crystals of specific gravity 4.48. Barium sulphate is almost completely insoluble in water and dilute acids. Though alkaline carbon- ates in aqueous solution affect barium sulphate but slightly, the latter may readily be transformed to carbonate by fusion with sodium carbonate, thus : BaSO 4 + Na 2 CO 3 = Na 2 SO 4 + BaCO 3 . Barium sulphate is easily prepared by treating a solution of a barium salt with any soluble sulphate. The precipitate is usu- ally composed of very small crystals. These grow larger on standing and can then be filtered off readily. In analytical operations, the fact that barium sulphate is insoluble and readily prepared is much used in testing for barium and also for sul- phates. Barium sulphate is often used as a white pigment, known as permanent white. On heating barium sulphate with carbon, harium sulphide BaS is produced, which, like the sulphides of calcium and strontium, phosphoresces in the dark. This fact was discovered in 1603 by Casciorolo, a shoemaker of Bologna, who, in an at- tempt to make silver, had heated native barium sulphate with carbon, thus forming the sulphide of barium. In 1774 Scheele investigated barium sulphate and showed that it contains an earth quite different from lime. Barium carbonate was then soon prepared artificially, and its probable occurrence in nature was foretold, when indeed in 1783 native barium carbonate was 410 OUTLINES OF CHEMISTRY discovered in Scotland by Withering, after whom the mineral is called witherite. Barium receives its name from the Greek word meaning heavy. It was reported that native barium car- bonate also occurred at Strontium in Scotland. Investigations by Crawford in 1790 showed, however, that this mineral was still another carbonate. Thus strontium was discovered. It derives its name from the locality where the native carbonate, strontianite, was found. Detection of the Alkaline Earth Metals. These may readily be detected by means of the spectroscope, as already stated. In addition the following facts are frequently used in detecting calcium, strontium, and barium. By adding a calcium sulphate solution to a solution of a barium or strontium salt an insoluble sulphate is precipitated ; but by similar treatment solutions of calcium salts remain clear. Stron- tium sulphate solution precipitates barium sulphate from solu- tions of barium salts. Ammonium oxalate precipitates the oxa lates of barium, strontium, and calcium from solutions of their salts. The oxalates of barium and strontium are soluble in dilute acetic acid, while the oxalate of calcium is not. Calcium nitrate is soluble in alcohol, while strontium and barium nitrates are not. Barium chromate BaCrO 4 is insoluble in water, while calcium chromate is soluble. The sulphates of calcium and strontium are converted into carbonates by boiling with sodium carbonate solution, while barium sulphate is not changed by this treatment. Radium and Radio-activity. In 1896 Becquerel discovered that photographic plates, carefully protected from light by means of dark coverings, are nevertheless affected as though they had actually been exposed to light, when kept near ura- nium or its compounds. Furthermore, it was found that a well- insulated, charged electroscope would soon become discharged in the neighborhood of uranium or its compounds, showing that by the latter the air had been rendered a conductor of electric- ity. It was consequently concluded that these phenomena are due to peculiar rays emitted from uranium. These rays were called uranium rays or Becquerel rays, and the substances from which they issued were called radio-active. The mineral pitch- blende, uraninite, consisting essentially of a black oxide of uranium U 3 O 8 , together with various impurities in minor quan- THE ALKALINE EARTH METALS 411 titles, shows this radio-activity in a high degree; and M. and Mme. Curie found that the residues of this mineral after removal of the uranium were still very radio-active. This led them to further investigations, and in 1898 they announced that they had discovered in these residues a new element, radium, to which the radio-active phenomena are due. In pitchblende radium is found together with relatively very much larger quantities of barium, so that a ton of residues yielded about thirty pounds of barium chloride, from which by very many re- peated fractional crystallizations a few tenths of a gram of radium chloride were finally obtained. This compound was over a million times as radio-active as pitchblende. Radium chloride or bromide is less soluble than the corre- sponding barium salt, hence the possibility of effecting a tolerable separation from the latter by fractional crystallization. Radium compounds exhibit a spectrum similar to that of the alkaline earths. The metal radium itself has not yet been isolated. The analysis of the chloride has shown that 35.46 grams of chlorine are combined with 113.2 grams of radium, hence, re- garding radium as bivalent like the alkaline earth metals, to which it appears to be analogous, the atomic weight is 226.4. The symbol of radium is Ra. On account of the scarcity of the element, but few of its compounds have been studied. The chloride RaCl 2 , the bromide RaBr 2 , the nitrate Ra(NO 3 ) 2 , the carbonate RaCO 3 , and the hydroxide Ra(OH) 2 have been pre- pared. All of the compounds show radio-activity. In the dark they are luminous. They also emit heat continually. It has been estimated that a gram of radium in its compounds gives off heat at the rate of approximately 100 calories per hour. As the rays from radium compounds impinge upon a screen of ba- rium platinocyanide or zinc sulphide, they cause these to phos- phoresce; furthermore, radio-activity may be communicated from radium compounds to substances placed near them. So the walls of containers of radio-active substances acquire this property, and solids kept in contact with radium salts also become active. The rays emitted from radio-active substances are able to kill germs, destroy the germinating power of seeds, and act destructively on living tissues, so that experimentation with the purer and consequently more powerful radium compounds must be conducted with great care. Glass con- 412 OUTLINES OF CHEMISTRY tamers of radium salts slowly turn violet in color, while cloth is gradually disintegrated by the action of the rays. The rays emitted from compounds of radium are complex in character. This is shown by the fact that a portion of them (the a rays) is readily absorbed by metals, the air, etc., and is deflected with great difficulty by a magnetic field. Another portion (the ft rays) is readily deflected by a magnet, and is particularly active on photographic plates; while a third por- tion (the 7 rays) is not deflected by a magnet and has great penetrating power, passing even through a thickness of 30 cm. of iron. The study of these rays has largely been conducted by testing their ability to affect photographic plates and to dis- charge an electroscope. Radium compounds emit an emanation which may be con- densed at about 150, by the aid of liquid air. This emana- tion is therefore of the nature of a gas. Indeed, it may be passed from one tube to another like air. When examined in a vacuum tube by means of the spectroscope, the emanation from radium compounds after a time shows the spectrum of helium. For this reason it has been concluded that radium is slowly changing to helium. Furthermore, since radium compounds gradually lose their activity, and since uranium compounds gradually again acquire the ability of giving off emanations after having been deprived of the same, the assumption has been made that uranium is gradually changing into radium and that the latter is " decaying " into emanations* which in turn are transformed into helium. Thus, the study of radio-active phenomena has opened up the question of the possibility of the transmutation of the chemical elements. Thorium compounds also exhibit radio-activity and give off emanations. A particularly active constituent, known as tho- rium Jf, has been separated by chemical means from thorium, while a similar product, uranium X, has been obtained from uranium, and another, actinium J, from actinium. Actinium is a radio-active substance found in uranite by Debierne. It acts like thorium, but more intensely. The elementary char- acter of actinium and of polonium, which has also been found in pitchblende by Mme. Curie, still awaits confirmation. THE ALKALINE EARTH METALS 413 REVIEW QUESTIONS 1. Why are calcium, strontium, and barium called the alkaline earth metals? Make a list of the formulas of their oxides, hydroxides, car- bonates, sulphates, nitrates, phosphates, silicates, borates, oxalates. Indicate which of these compounds are soluble in water, and which are not soluble. 2. Compare the alkali metals with the alkaline earth metals in their physical properties and their chemical action on water. 3. Make a list of the forms in which calcium carbonate is found in nature, and state the use of each. 4. How is lime made from limestone? Equation. Write an equa- tion expressing what happens when mortar sets. What is the function of hair or wood fiber in mortar used for plastering walls ? 5. Besides its use for building purposes, what other uses are made of lime? 6. What are the chief ingredients in Portland cement? In what proportions should these be present? 7. In what respects does the setting of Portland cement differ from the setting of mortar? From the setting of plaster of Paris? How is plaster of Paris made ? 8. How prepare calcium sulphite ? What use is made of it ? 9. How much calcium chloride could be made from 12 pounds of pure chalk? Describe the properties of calcium chloride and compare it with the chlorides of strontium and barium. 10. What is superphosphate of lime fertilizer? Write the equation showing how it is formed from phosphate rock. 11. How is calcium carbide made? Write the reaction showing how acetylene is made from calcium carbide. How many liters of acetylene measured under standard conditions could be produced from one kilogram of calcium carbide? 12. Name four different kinds of glass, stating their similarities and differences in chemical composition. 13. How do strontium and barium occur in nature, and what use is made of their compounds? 14. Given a solution which contains the chlorides of calcium, barium, and strontium, describe two distinctly different methods of demonstrat- ing the presence of each of these three metals in the solution. Write equations. 15. Write chemical equations expressing the action of a barium chloride solution upon the solution of (a) copper sulphate, (b) ammonium oxalate, (c) sodium phosphate. 16. How much pure gypsum would be required to produce one thou- sand pounds of plaster of Paris ? 17. How is the element radium related to the alkaline earth metals ? Where does radium occur? What are some of its striking properties? What is meant by the term "radium emanation " ? CHAPTER XXIII THE METALS OF THE MAGNESIUM GROUP THE metals of this group are glucinum, magnesium, zinc, cad mium, and mercury. Of these, magnesium closely resembles the alkaline earth metals. Zinc, cadmium, and mercury, on the other hand, have a much higher specific gravity, and are much less readily oxidized; their affinity for oxygen diminishes in the }rder named. Glucinum, unlike the other members of this group, has a high melting and boiling point. It really is a transition element between the magnesium group and the metals of the earths. Glucinum, or Beryllium, Gl, or Be, At. Wt. 9.1, is found in nature as a constituent of the rather rare minerals beryl Gl 3 Al 2 Si 6 O 18 , phenacite Gl 2 SiO 4 , and chrysoberyl G1(A1O 2 ) 2 . When colored green by traces of chromium silicate, beryl is called emerald, which is used as a gem. Glucinum is a white, malleable metal having a specific gravity 1.8. It does not liberate hydrogen from water even on boiling. It dissolves in hydrochloric or sulphuric acid, while nitric acid attacks it but slightly. In caustic potash, glucinum dissolves readily, liberat- ing hydrogen. The metal is obtained by electrolysis of molten chrysoberyl, or by heating the oxide with magnesium powder. In its compounds, glucinum is always bivalent. The compounds have a sweetish taste, whence the name glucinum (sweet); the element is also called beryllium from beryl, in which it occurs. Glucinum oxide G1O was discovered in 1797 by Vauquelin. It is basic, but also possesses very weak acidic properties, form- ing soluble salts with caustic alkalies. From solutions of glucinum salts, ammonium hydroxide precipitates gltLCinum hydroxide G1(OH) 2 . Among the well-known salts are glucinum sulphate G1SO 4 , glucinum chloride G1C1 2 , and glucinum carbonate G1CO 3 . Basic salts, notably basic nitrates and sulphates, are also known. Occurrence, Preparation, and Properties of Magnesium. Mag- nesium is very widely distributed. It occurs in large quanti- 414 THE METALS OF THE MAGNESIUM GROUP 415 ties in dolomite MgCO 3 -CaCO 3 , which often forms mountains. It is also found in carnallite MgCl 2 KC1 6 H 2 O, in kieserite MgSO 4 -H 2 O, in schonite MgSO 4 K 2 SO 4 6 H 2 O, in kainite K 2 SO 4 -MgSO 4 - MgCl 2 - 6 H 2 O, in magnesite MgCO 3 , in soapstone or talc Mg 3 H 2 Si 4 O 12 , in meerschaum Mg 2 Si 3 O 8 -4 H 2 O, in serpen- tine Mg 3 Si 2 O 7 2 H 2 O, in hornblende Mg 2 CaFeSi 4 O 12 , in asbes- tus Mg 3 Si 2 O 7 2 H 2 O, and in many other complex silicates. Sea water contains magnesium chloride and sulphate. These salts also occur in many spring waters, which are called bitter waters. Magnesium salts are found in soils, being decomposition prod- ucts of rocks. Plants and animals contain magnesium, which appears in their ash as carbonate and phosphate. Guano con- tains magnesium ammonium phosphate. Metallic magnesium is prepared by electrolyzing molten carnal- lite in an iron pot, which also serves as cathode, a carbon anode being employed. The metal is silver-white, and when hot it may be drawn into ribbon or wire. At ordinary temperatures magnesium is but slightly malleable. Exposed to air, it grad- ually becomes covered with a thin, white layer of oxide. It melts at 633 and boils at 1100. Its specific gravity is 1.75. It may be sublimed in a vacuum, yielding hexagonal, prismatic crystals. In an atmosphere of hydrogen the metal may be dis- tilled. It acts but very slightly on pure water even at 100 ; but when heated in a current of steam, it takes fire. Dilute acids act vigorously on magnesium. In the air, the metal burns with a brilliant white light, yielding the oxide MgO, together with some nitride Mg 3 N 2 . Magnesium light has a powerful effect upon photographic plates and is consequently used in flash lights. Flash light powder consists of about 5 parts of powdered magnesium to 9 parts of potassium chlorate. Magnesium is also used in fireworks, for signal lights, and as a reducing agent in chemical operations. The metal was first obtained in pure form by Liebig and Bussy in 1830. Bunsen prepared it by electrolysis twenty-two years later. The atomic weight of magnesium is 24.32, and its valence is two. Magnesium Oxide MgO is a white powder, usually prepared by heating the carbonate. It is also called magnesia, magnesia usta, or calcined magnesia. In contact with water it does not dissolve, but it forms magnesium hydroxide Mg(OH) 2 , w r hich is 416 OUTLINES OF CHEMISTRY very sparingly soluble. Magnesium oxide is used in medicine. It is even more difficult to melt magnesia than lirne, conse- quently magnesia is used in making fire brick and in the con- struction of the electric furnace. Magnesium oxide does not have acidic properties. Magnesium Carbonate MgCO 3 occurs in nature as magnesite in hexagonal crystals that are isomorphous with calcite. On adding a solution of an alkaline carbonate to a solution of a magnesium salt, basic magnesium carbonate is precipitated, the composition of which varies according to the temperature and concentration of the solutions used. The ordinary magnesium carbonate of commerce, called magnesia alba, is prepared by precipitation. Its composition is approximately (MgCO 3 ) 3 - Mg(OH) 2 - 3 H 2 O. Though insoluble in water, magnesium carbonate, like calcium carbonate, dissolves in water charged with carbon dioxide. From such solutions crystals of the composition MgCO 3 -3 H 2 O (and at low temperatures MgCO 3 -5H 2 O) have been obtained. Magnesium carbonate is used in medicine, also as a cosmetic. Magnesium Chloride MgCl 2 6 H 2 O forms highly hygroscopic, monoclinic crystals. On heating these they completely decom- pose, thus : MgCl 2 - 6 H 2 O = MgO + 2 HC1 + 5 H 2 O. The anhydrous salt MgCl 2 may be obtained by heating magne- sium ammonium chloride MgCl 2 -NH 4 Cl-6 H 2 O. Figure 136 .20 u> e, 80 Jf o be - 60 S20 | MgCl 2 .I2H -40 -20 20 40 60 80 100 120 Temperature in degrees C. FIG. 136. 140 ISO THE METALS OF THE MAGNESIUM GROUP 417 presents the solubility curve of magnesium chloride. It will bfl seen from the figure that five different hydrates are known. Magnesium Sulphate MgSO 4 7 H 2 O, also known as Epsom salt arid bitter salt, forms large rhombic prisms that dissolve in about four parts of water. The solution has a disagreeable bitter taste. The salt is found in the waters of Epsom springs and many other mineral springs. It is used as a purgative, and is also employed in " loading " cotton goods and in making the sulphates of potassium and sodium. Occasionally it serves as a fertilizer. Magnesium Phosphates are in general similar to those of calcium. Magnesium ammonium phosphate MgNH 4 PO 4 is im- portant in analytical chemistry. The salt forms rhombic crystals that are insoluble in ammonia water. On ignition they yield magnesium pyrophosphate Mg 2 P 2 O 7 . Magnesium Ammonium Arsenate MgNH 4 AsO 4 is analogous to magnesium ammonium phosphate. On ignition it yields magnesium pyroarsenate Mg 2 As 2 O 7 . Tests for Magnesium. In testing for magnesium compounds, the following facts are of importance : Basic carbonate of magnesium is precipitated by carbon- ates of the alkalies. Hydroxides of the alkalies precipitate magnesium hydroxide ; but ammonium hydroxide does not precipitate magnesium hydroxide in presence of ammonium chloride, for the solutions of the latter dissolve magnesium hydroxide. In presence of ammonium chloride, clear am- moniacal solutions of magnesium salts are precipitated by sodium phosphate, magnesium ammonium phosphate being formed. Occurrence, Preparation, and Properties of Zinc. Zinc occurs in nature mainly as the carbonate, ZnCO 3 , in the mineral called smithsonite, calamine or zinc spar, and as the sulphide, ZnS, in zinc blende or black-jack. Other ores of zinc are franklinite Zn(FeO 2 ) 2 , gahnite Zn(AlO 2 ) 2 , also called zinc spinel, and red zinc ore ZnO. Most of the ores of zinc also contain some cadmium. In making metallic zinc, the ores are first roasted. Thus carbonates and sulphides are finally all transformed to zinc oxide. The latter is then reduced by heating it with carbon. The reactions involved are : 418 OUTLINES OF CHEMISTRY ZnCO 3 = ZnO + CO 2 . ZnS + 3 O = ZnO + SO 2 . ZnO + C = Zn + CO. The operation of reducing the zinc oxide is conducted in earthenware retorts. At the bright red heat developed, 1200 to 1300, zinc is converted into vapor, which is condensed in iron receivers connected with the retorts. Zinc dust, consist- ing of finely divided zinc plus 5 to 10 per cent of zinc oxide, is at first obtained on the sides of the condenser ; later, the metal condenses to a liquid and is run into molds. This crude zinc, called spelter, is contaminated with carbon, iron, arsenic, lead, and cadmium. It is purified by redistillation. Chemically pure zinc is made by heating pure, precipitated zinc carbonate with pure carbon, or still better by the electrolysis of pure zinc salts. Zinc crystallizes in the hexagonal system. It is bluish white, and has a bright metallic luster. Cast zinc has a specific gravity of 6.9 ; but hammered zinc, or zinc wire, is denser, having a specific gravity of 7.2. Zinc is rather brittle, but at 120 it becomes malleable and ductile. When heated to 200, it becomes so brittle that it may be pulverized. Zinc melts at 420 and boils at 918. Its atomic weight is 65.37 and its molec- ular weight is the same, for the vapor is 33.93 times as heavy as hydrogen. Zinc is always bivalent. On exposure to moist air, the luster of zinc is dimmed by the formation of a thin coating of white basic carbonate. At high temperatures, zinc burns in the air with a brilliant, bluish white flame. On water zinc does not act, but when steam is passed over heated zinc, the oxide and hydrogen are formed. Nearly all dilute acids act on zinc ; the purer the zinc, the less rapid is the action. In general, hydrogen is liberated when zinc acts on acids, though in some cases, like that of nitric acid for instance, the hydrogen is not set free, for it at once reduces some of the acid present. On heating zinc with concentrated sulphuric acid, sulphur dioxide and zinc sulphate are formed : Zn + 2 H 2 SO 4 = ZnSO 4 + SO 2 + 2 H 2 O. In hot caustic alkalies, zinc dissolves, as already stated, forming a zincate and hydrogen : Zn + 2 NaOH = Na 2 ZnO 2 + H 2 . THE METALS OF THE MAGNESIUM GROUP 419 When introduced into solutions of many of the salts of either lead, tin, copper, mercury, silver, platinum, or gold, zinc precipi- tates these metals, usually as finely divided powders, thus : CuSO 4 + Zn = ZnSO 4 + Cu. Zinc is used in sheet form for many purposes, such as making roofs, gutters, and architectural ornaments. Galvanized iron, so-called, consists of sheet iron which has been coated with zinc by dipping the thoroughly clean sheets of iron into highly heated molten zinc. The zinc coating prevents the iron from rusting. Much zinc is also used in making electrical batteries (which see), and in preparing many alloys, especially brass. Brass, which is an alloy of zinc and copper, was known long before zinc, for it was obtained by melting native copper ores that contained zinc. On the European continent zinc production on a commercial scale began about one hundred years ago ; though in England zinc was manufactured some fifty years earlier. Every year the United States produces about 360,000 tons of zinc, which is approximately one third of the world's annual output. Zinc Oxide ZnO, also called flores zinci or lana philosophica, is a white, bullsy powder obtained by burning zinc in the air or by heating the basic carbonate of zinc. When hot, it is yellow, but on cooling, it turns white. It is much used in white paints under the name zinc white. The oxide is also used in pharmacy for making ointments. Zinc hydroxide is obtained as an amor- phous precipitate by adding caustic alkalies to solutions of zinc salts. The hydroxide is soluble in an excess of the precipitant, forming a zincate. On heating the hydroxide, it loses water, forming the oxide. Native oxide of zinc is frequently colored red because of the presence of oxides of manganese. Zinc Carbonate ZnCO 3 occurs in nature, as already stated. It forms rhombohedral crystals that are isomorphous with cal- cite. On treating solutions of zinc salts with alkali carbonates, basic zinc carbonates are precipitated, whose composition varies according to the temperature and concentration of the solu- tions. These carbonates are approximately ZnCO 3 2 Zn(OH) 2 or (ZnCO 3 ) 2 3 Zn(OH) 2 . They are soluble in an excess of am- monium carbonate, but not in sodium or potassium carbonate. Zinc Chloride ZnCl 2 is made by burning zinc in chlorine, or by the action of hydrochloric acid on the carbonate, oxide, hy- 420 OUTLINES OF CHEMISTRY droxide, or the metal itself. It is a white, deliquescent mass, having caustic properties. It is soluble in alcohol as well as in water. On attempting to dehydrate the salt by heating it, hydrochloric acid is given off, as in the case of magnesium chlo- ride. On mixing concentrated zinc chloride solutions with zinc oxide, oxychlorides like Zn(OH)Cl are formed, the mass hard- ening in a manner not unlike the setting of plaster of Paris. Such zinc chloride and zinc oxide mixtures are used as cements in filling teeth. Zinc chloride is further employed in preserv- ing railroad ties, which, after being soaked in its solutions, do not rot readily. It is also used to cleanse the surface of metals in the process of soldering. In medicine zinc chloride is used as a caustic agent and as a disinfectant. In the laboratory it is used in connection with certain syntheses of organic com- pounds. Zinc bromide ZnBr 2 is analogous to the chloride. The iodide ZnI 2 is also very soluble ; it readily splits off iodine. The fluoride ZnF 2 is but sparingly soluble. Zinc Sulphate ZnSO 4 , white vitriol, is formed by the oxida- tion of zinc sulphide or by the action of sulphuric acid on the metal, oxide, or carbonate. It crystallizes in the rhombic sys- tem in colorless prisms of the composition ZnSO 4 7 H 2 O, which effloresce on exposure to the air. The crystals are ismorphous with MgSO 4 7 H 2 O. They are readily soluble in water. With sulphates of the alkalies, zinc sulphate forms double salts, like K 2 SO 4 ZnSO 4 6 H 2 O, which is analogous to scho- nite, K 2 SO 4 MgSO 4 6 H 2 O. Like other zinc compounds, zinc sulphate is moderately poisonous, upon which its use as an antiseptic depends. Zinc Sulphide ZnS is white when pure. Native sulphide of zinc, black-jack, is colored dark brown by ferric oxide and other impurities. It crystallizes in the regular system. On adding sulphides of the alkalies or of ammonium to a solution of a zinc salt, a white precipitate of zinc sulphide is formed. This is soluble in dilute mineral acids but not in acetic acid. Analytical Tests for Zinc Salts. The reactions that are used in testing for zinc salts are : Ammonium sulphide precipitates white zinc sulphide from solutions of zinc salts. The sulphide is insoluble in acetic acid but soluble in mineral acids. Alkaline hydroxides precipitate zinc hydroxide from solutions THE METALS OF THE MAGNESIUM GROUP 42 1 of zinc salts The precipitate dissolves in an excess of the reagent. Alkaline carbonates precipitate basic zinc carbonate, which is soluble in ammonium carbonate but not in sodium or potas sium carbonate. Zinc oxide is white when cold and yellow when hot. When moistened with cobalt nitrate solution and then strongly heated on charcoal, zinc oxide yields a green mass called Rinmann's green. Occurrence, Preparation, and Propertiesof Cadmium. Cadmium is a rather rare element. It frequently occurs in small quanti- ties in ores of zinc. It is also found, though rarely, as the native sulphide, CdS, known as greenockite, which is hexagonal and isomorphous with a rare, native zinc sulphide, wiirzite. Cadmium is commonly obtained in connection with the reduc- tion of zinc from its ores. The metal boils at 775 and hence passes over into the condensers before the zinc, which boils at 918. Cadmium is silver-white, malleable and ductile even at ordinary temperatures. It resembles tin in appearance, melts at 320, and has a specific gravity of 8.6. Its atomic weight is 112.4, and its molecular weight is the same, for its vapor is 55.65 times as heavy as hydrogen. In the air, cadmium is quite stable, though its luster becomes somewhat tarnished, due to the formation of a thin layer of oxide. When strongly heated in the air or in oxygen, it burns, forming a brown oxide. Cadmium is always bivalent. Dilute acids act on cadmium, liberating hydrogen. The element was discovered in 1817 by Stromeyer, who investigated a yellowish zinc oxide which contained no iron. Almost simultaneously, cadmium was discovered by Hermann. The metal is used in preparing standard cells for measuring electromotive forces. It is also used in certain alloys of low melting point, like Wood's metal, which has already been described. Cadmium Compounds. The oxide, CdO, is a brown powder which is readily reduced at higher temperatures by means of carbon or hydrogen. The hydroxide, Cd(OH) 2 , is white. It is formed by leaving the oxide in contact with water or by pre- cipitation from solutions. It is soluble in ammonium hydroxide. The chloride, CdCl 2 , may be obtained from solutions in the form of crystals of the composition CdCl 2 2 H 2 O, which effloresce, 422 OUTLINES OF CHEMISTRY The anhydrous salt melts at 540, and may be distilled at about 900. The bromide, CdBr 2 4 H 2 O, is soluble in alcohol as well as in water. The iodide, CdI 2 , is obtained by the direct union of iodine and cadmium in presence of water. This salt is soluble in alcohol, and is used in photography. The nitrate^ Cd(NO 3 ) 2 4 H 2 O, is a deliquescent salt obtained by the action of nitric acid upon either the metal, the hydroxide, or the car- bonate. The sulphate, 3 CdSO 4 8 H 2 O, forms monoclinic crys- tals that readily dissolve in water. The salt effloresces on exposure to the air. It will be observed that the sulphate is not analogous to zinc sulphate ZnSO 4 7 H 2 O and magnesium sulphate MgSO 4 7 H 2 O. Cadmium sulphate is used in treat- ing diseases of the eye. The sulphide, CdS, is obtained as a bright yellow precipitate by conducting hydrogen sulphide into a solution of a cadmium salt. Dilute acids do not affect it. It is used as a pigment in paints. The behavior of the hydroxide, carbonate, and sulphide, as above described, is used in testing for cadmium salts. Occurrence, Preparation, and Properties of Mercury. Mercury, or hydrargyrum, meaning silver water, is the only metal that is liquid at ordinary temperatures. It was known at least three hundred years before the Christian era. In nature it is sometimes found uncombined in small drops in the interstices of rocks. Its principal ore is cinnabar, the sulphide, HgS, which forms dark red, hexagonal, prismatic crystals. Mercury ores occur at Almaden in Spain, Idria in Austria, New Alma- den in California, in Prussia, Peru, Japan, and China. The metal is obtained by simply roasting the sulphide : The mercury vapors are condensed and collected. Sometimes the ore is heated with lime in iron retorts, when calcium sul- phide remains behind and mercury passes over in form of vapor and is condensed. Mercury is purified by redistillation after treatment with dilute nitric acid, ferric chloride solution, or dilute sulphuric acid plus potassium bichromate. Mechanical impurities are removed from mercury by nitration through cloth or chamois skin. Perfectly pure mercury is made by lib- erating the metal from pure mercury salts, and finally distilling the product, preferably in vacuo. THE METALS OF THE MAGNESIUM GROUP 423 Mercury is silver-white, and has a brilliant metallic luster. It melts at 39.4 and boils at 357. Its specific gravity is 13.59. Its atomic weight is 200.6, which is also its molecular weight, for mercury vapor is about 100 times as heavy as hydrogen. Mercury vapor has a very characteristic spectrum, exhibiting a bright yellow and a green line, besides a red, a blue, and three violet lines. The vapor conducts electricity and emits a bright, pale, greenish light rich in rays that affect photographic plates. The mercury lamp depends on the prin- ciple that mercury vapors emit light when they are conducting high tension electricity. The vapors of mercury are very poisonous, which is also true of compounds of mercury. Mercury crystallizes in the regular system. Solid mercury is malleable. It may be cut with tools, and beaten into sheets with a hammer. In the air mercury remains practically unchanged. By sulphuric or hydrochloric acids it is attacked but slightly. Nitric acid or hot, concen- trated sulphuric acid readily dissolves it. With sulphur or the halogens it combines at slightly elevated temperatures. Mercury is used in making thermometers, barometers, mir- rors, various amalgams in dentistry, and many compounds that find application in medicine. Large amounts of mercury are also employed in extracting gold and silver from their ores. The United States produces about 750 tons of mercury per year, which is about one fourth of the annual amount produced in the world. Mercury is shipped in iron flasks containing about 75 pounds each (Fig. 137). Amalgams. The alloys or combinations of mercury with other metals are called amal- gams. These are usually made by simply bringing the metals in contact with mercury, though they may also be obtained by electro- lyzing a salt, using a mercury cathode ; or frequently by introducing a metal into a solu- tion of a mercury salt, like the nitrate. In general, amalgams are of the nature of solutions of the metals in mercury. They are liquid when the mercury preponderates, and solid when relatively less mercury is used. 424 OUTLINES OF CHEMISTRY Sodium combines with mercury with evolution of light and heat. Sodium amalgam containing less than 1 per cent sodium is liquid ; a 1 per cent sodium amalgam is viscous, while amalgams containing 2 per cent or more of sodium are solid. Crystalline sodium amalgams corresponding to the formulse Na 3 Hg and NaHg* 6 have been isolated. Sodium amalgam is used as a reducing agent, as already mentioned under sodium. With the exception of iron and platinum, practically all the metals form amalgams with mercury. Even platinum may be amalgamated by electrolyzing a mercury salt with a platinum cathode. However, the readiness with which the various metals unite with mercury varies very greatly. So gold and silver dissolve with ease in mercury, which fact is used in extracting these metals from pulver- ized rocks with which they are mixed. Copper, cadmium, and tin also readily unite with mercury. The amalgam used in filling teeth usually consists of mercury mixed with an alloy containing essentially silver and tin, together with smaller amounts of other metals, among which are copper, cadmium, and gold. The amalgam is made into a stiff paste and intro- duced into the cavity in the tooth, where it soon hardens or sets without material change of volume. Tin amalgam is used in making mirrors. Amalgamated zinc is hardly acted upon by dilute acids; on the other hand, magnesium amalgam, which forms only on heating mercury and magnesium together, is a black incoherent mass which violently decomposes water with liberation of hydrogen. When sodium amalgam is treated with a concentrated solu- tion of ammonium chloride or other ammonium salt, a bulky mass called ammonium amalgam is obtained. It is supposed to contain ammonium NH 4 , dissolved in mercury. The large bulk is produced by hydrogen and ammonia gases that are set free. This so-called ammonium amalgam is also formed by electro- lyzing a solution of an ammonium salt, using mercury as a cathode. By some it is considered practically certain that the material contains NH 4 , dissolved in mercury ; by others this has been denied. At any rate, if ammonium amalgam exists at all, it is very unstable. Compounds of Mercury. Mercury forms two distinct series of compounds : the mercurous compounds, in which the metal THE METALS OF THE MAGNESIUM GROUP 425 is univalent ; and the mercuric compounds, in which it is bivalent. Oxides of Mercury. Mercurous oxide Hg 2 O is a dark brown, slightly greenish, unstable powder obtained as a precipitate by adding sodium or potassium hydroxide solution to a mercurous salt. It decomposes into mercuric oxide and mercury on ex- posure to light or when slightly heated. Mercuric oxide HgO may be obtained by prolonged heating of mercury in the air at about 360. It is made on a large scale by heating a mixture of mercuric nitrate and mercury. It is a bright red, crystalline powder also known as red precipitate. On heating mercuric oxide, it turns dark and gives off oxygen. Above 500 the de- composition into mercury and oxygen progresses rapidly. Red oxide of mercury is used in medicine in red precipitate oint- ment. The yellow variety is also used for similar purposes ; being finely divided, it is more active than the red. By adding caustic soda to a solution of a mercuric salt, mercuric oxide is obtained in form of a finely divided, amorphous, yellow powder. On heating, it behaves like the red oxide. Hydroxides of mercury are not known, probably because they are so unstable, decomposing at once into oxide and water. Halides of Mercury. Mercurous chloride HgCl, also called calomel, is obtained as a slightly yellowish, crystalline mass by heating mercury with mercuric chloride : Hg+HgC] 2 = 2HgCl. It may also be formed by passing sulphur dioxide into a hot solution of mercuric chloride: 2HgCl 2 + 2H 2 O + SO 2 = 2 HC1 + H 2 SO 4 + 2HgCl. Other reducing agents may be employed instead of sulphur dioxide; thus, with stannous chloride the reaction is, 2 HgCl 2 + SnCl 2 = SnCl 4 + 2 HgCl. By adding a soluble chloride to a solution of a mercurous salt, mercurous chloride is precipitated, thus ;. HgN0 3 + NaCl = NaN0 3 + HgCl. Mercurous chloride is not soluble in water. It may be sub- limed, and thus obtained in crystalline form. The precipitated variety is an amorphous powder. This salt is much used in medi- cine as a purgative and as a stimulant for the secretory organs 426 OUTLINES OF CHEMISTRY It should be kept in the dark, for on exposure to light it gradu ally decomposes somewhat, yielding mercury and mercuric chloride., which is a strong poison, thus : 2HgCl = Hg + HgCl a . Mercurous bromide HgBr is obtained as a white, insoluble powder by adding a soluble bromide to a mercurous salt. It may be obtained in crystalline form by sublimation, or by treating mercury with bromine water. Mercurous iodide Hgl is formed by triturating mercury and iodine together in presence of a little alcohol, or by precipitat- ing a solution of a mercurous salt with sodium or potassium iodide. It is a dark, yellowish green powder which readily decomposes into mercury and mercuric iodide, especially in the light, thus : 2 Hgl = Hg + HgI 2 . The salt is used in medicine. Mercuric chloride HgCl 2 , also called corrosive sublimate, or sublimate, is made on a large scale by heating mercuric sul- phate with common salt, thus : HgSO 4 + 2 NaCl = Na 2 SO 4 + Hg01 2 . In this process the mercuric chloride vapors formed are con- densed, thus yielding beautiful, rhombic, prismatic crystals that readily dissolve in water (1 in 15) and also in alcohol and in ether. The salt melts at 265 and boils at 307. Mercuric chloride may also be made by the action of hydrochloric acid on mercuric oxide, or by dissolving mercury in aqua regia. Mercuric chloride is reduced to mercurous chloride by reduc- ing agents as already stated above. An excess of stannoue chloride reduces mercuric chloride to metallic mercury : HgCl 2 + SnC1 2 = SnCl 4 + Hg. Mercuric chloride is a powerful poison. It is used as a disin- fectant in surgery, for cleansing wounds and washing the hands and surgical instruments. It is employed only in dilute solutions. The solutions have corrosive properties and a sharp, very disagreeable, "metallic" taste. Mercuric chloride is also employed in preserving anatomical specimens, herbaria, stuffed animals, wood, etc. With albumen it unites, forming insoluble compounds, hence the use of white of egg and milk in case of THE METALS OF THE MAGNESIUM GROUP 427 poisoning with corrosive sublimate. With chlorides of the alka- lies, mercuric chloride forms double salts, like HgCl 2 - KC1 H 2 O. Mercuric bromide HgBr 2 is similar to mercuric chloride. It is isomorphous with the latter and melts at 325. Mercuric iodide HgI 2 may be prepared by the direct union of iodine and mercury, or by adding potassium iodide to a solution of mercuric chloride, when a yellowish precipitate forms which soon becomes bright red. Though insoluble in water, it readily dissolves in alcohol, from which solutions it crystallizes in bright red, tetragonal pyramids. In aqueous potassium iodide solu- tions, mercuric chloride is readily soluble, forming a yellow solution containing the double salt HgI 2 2 KI. This solution may be concentrated till it has a specific gravity of over 3. It is known as Thoulet's solution and is used by mineralogists in determining the specific gravity of small pieces of minerals. Potassium mercuric iodide solution to which caustic potash has been added is Nessler's reagent, which is used in determin- ing ammonia in the analysis of potable waters. With ammonia such solutions give a brown precipitate having the composi- /Hg\ tion O/ yNH 2 -T. If the ammonia solutions are very di- / lute, only a yellowish brown coloration is observed. Mercuric Cyanide Hg(CN) 2 , the only cyanide of a heavy metal that is soluble in water, is obtained by the action of hydrocyanic acid on mercuric oxide. It forms tetragonal prisms, which when heated yield mercury and cyanogen. Nitrates of Mercury. Mercurous nitrate HgNO 3 - H 2 O forms monoclinic crystals, obtained by action of nitric acid upon an excess of mercury in the cold. On diluting its solutions with water, yellow basic salts separate out, like Hg(OH)HgNO 3 . This hydrolysis is counteracted by keeping a slight excess of nitric acid in the solutions. Mercuric nitrate Hg(NO 3 ) 2 is formed by dissolving mercuric oxide in nitric acid, also by the action of an excess of hot, con- centrated nitric acid on mercury. From the solutions, deli- quescent crystals 2 Hg(NO 3 ) 2 + H 2 O may be obtained. On dilution with water, these suffer hydrolysis, forming a series of basic salts whose composition varies according to the rela- tive amount of water present; so we have, for instance, 428 OUTLINES OF CHEMISTRY Hg(NO 3 ) 2 2 HgO + H 2 O. On boiling with much water, these basic salts are decomposed, yielding nitric acid and mercuric oxide. Addition of nitric acid, of course, reverses the re- action. Under the name Millon's reagent, mercuric" nitrate solution is used in testing for albumins, which are coagulated by it. Mercuric Fulminate HgC 2 O 2 N 2 , also called fulminating mer- cury, is made by the action of nitric acid on mercury in pres- ence of alcohol. It is a white powder, which is very explosive when dry. It is used in making percussion caps. The explo- sion of a small amount of the substance will cause gun cotton or nitroglycerine to explode. Sulphates of Mercury. Mercurous sulphate Hg 2 SO 4 is formed by treating an excess of mercury with sulphuric acid, or by precipitating a mercurous nitrate solution with sulphuric acid. The colorless crystals so obtained are sparingly soluble in water. The salt is used in making standard cells for compar- ison of electromotive forces. Mercuric sulphate HgSO 4 is made by digesting mercury or mercuric oxide with an excess of sulphuric acid. The salt so obtained is a white, crystalline mass. On treatment with water, it is hydrolyzed, forming sulphuric acid and a yellow basic salt HgO HgSO 4 , known as Turpeth mineral, which is insoluble. With sulphates of the alkalies, mercuric sulphate forms double salts like K 2 SO 4 HgSO 4 6 H 2 O. These are isomorphous with the analogous magnesium double salts. Mercuric Sulphide HgS is very stable. It is readily obtained by triturating mercury and moist sulphur together, by heating mercury with sulphur, or by conducting hydrogen sulphide into a solution of a mercury salt. Mercurous sulphide is not known with certainty; the black precipitate produced when a mercurous salt is treated with hydrogen sulphide or an alkaline sulphide is mercuric sulphide mixed with mercury. When made by precipitation, mercuric sulphide is a black, amorphous powder, which is insoluble even in concentrated acids on boil- ing. It is soluble in aqua regia, however. Black mercuric sul- phide may be sublimed, and thus converted into dark red rhombohedral crystals that are identical with cinnabar, which occurs in nature. Towards acids the red and black varieties act alike. Red sulphide of mercury serves as a pigment in THE METALS OF THE MAGNESIUM GROUP 429 paints under the name of vermilion. It has been used thus since ancient times. Compounds of Mercury Salts with Ammonia. When mercury salts are treated with ammonia, compounds are formed which may be regarded as ammonium salts in which one or more of the hydrogen atoms are replaced by mercury. So when mer- curous chloride is treated with ammonia water, a black insoluble powder, mercurous ammonium chloride, is formed : 2 HgCl + 2 NH 3 = (NH 2 Hg 2 ) 01 + NH 4 C1. Similarly, mercurous nitrate yields black mercurous ammonium nitrate (NH 2 Hg 2 )NO 3 plus mercury. n Mercuric ammonium chloride (NH 2 Hg)Cl, known as infu- sible white precipitate, is formed by adding ammonia to mer* curie chloride solution: HgCl 2 + 2 N H 8 = (NH 2 Hg)Cl + NH 4 C1. The so-called fusible white precipitate, mercuric diammonium chloride (NH 3 Cl) 2 Hg, is obtained by treating a boiling hot ammoniacal solution of ammonium chloride with mercuric chloride. Physiological Properties of Mercury Compounds. It has already been stated that mercury compounds are poisonous. When mercury is introduced into the system, it produces a peculiar taste in the mouth commonly described as metallic. Salivation follows, and the gums, teeth, liver, kidneys, and other organs frequently also become involved. In small doses mer- cury compounds stimulate the action of various glands; this furnishes the basis for the internal use of calomel. The intro- duction of mercury compounds into medicine dates back to Paracelsus, who lived in the first half of the sixteenth century. Tests for Mercury. Mercury is very readily detected in its compounds. By mixing any mercury compound with soda and heating, mercury is driven off, which condenses in drops in the cooler parts of the ignition tube. With iodine these drops form red iodide of mercury. In solutions of mercurous salts, chlorides produce a white precipitate of calomel which turns black on treatment with ammonia, as already stated. The fact that stannous chloride reduces mercuric chloride to 430 OUTLINES OF CHEMISTRY mercurous chloride, and that the latter is further reduced to mercury by adding more stannous chloride, is frequently used in testing for mercury. Further, bright copper when intro- duced into a solution of a mercury salt becomes coated with mercury. The sulphide of mercury is also very characteristic, as already mentioned. General Remarks. It will be observed that the compounds of magnesium and zinc bear close resemblances to each other, and also a fair resemblance to the compounds of cadmium. On the other hand, the mercury compounds do not resemble those of zinc, magnesium, and cadmium. To be sure, in the mercuric compounds, mercury is bivalent, and between this series and the compounds of the other metals mentioned, some analogies are apparent. It will be seen later that the compounds of mercury bear a closer resemblance to those of copper. Glucinum stands rather isolated, being, as has already been mentioned, a transi- tion element between this group and that of the earth metals. Thus it is evident that the metals of the magnesium group do not form as closely related a family as some of the others that have already been studied. REVIEW QUESTIONS 1. Why are beryllium, zinc, cadmium, and mercury grouped with magnesium? 2. What is beryl? What characterizes the solutions of beryllium compounds? 3. Mention ten compounds in which magnesium occurs in nature, writing the appropriate formulas. 4. How is metallic magnesium obtained, and what use is made of it ? Equations. 5. Write the formulas of magnesium carbonate, sulphate, chloride, nitrate, phosphate, and oxalate, and compare the solubilities of these salts in water with those of the corresponding compounds of : (a) the alkalies, (6) the alkaline earths, (c) zinc, cadmium, mercury. 6. How much magnesium oxide may be prepared from 30 grams of pure magnesium carbonate? How may this be done? Equation. 7. If a solution contains salts of barium, strontium, and magnesium, how demonstrate that the latter is present? 8. What are the principal ores of zinc ? By means of equations, show how the metal may be extracted from these ores. 9. What is zinc dust, spelter, brass, galvanized iron? THE METALS OF THE MAGNESIUM GROUP 431 10. State what use is made of each of the following: zinc oxide, zinc chloride, zinc sulphate. 11. If a solution contains a mixture of zinc and magnesium salts, how demonstrate the presence of these two metals in the solution? Equations. 12. How was cadmium discovered? Describe the properties of this metal and mention its uses. 13. Given a solution of cadmium nitrate and one of zinc nitrate, how distinguish between the two by means of a single reagent. Equations. 14. What property have all mercury compounds in common? 15. What two series of salts does mercury form? Give the formulas of five salts in each series. 16. In what form does mercury occur in nature? Where does it occur? 17. Describe mercury. 18. What is an amalgam? Mention ten different amalgams. How may each be formed? What use is made of each of these amalgams? 19. Mention three compounds of mercury that are used in medicine, stating the purpose for which they are employed. 20. How much corrosive sublimate may be prepared from 5 Ib. of mercury? How may this be done? Write equations. 21. What is the action on corrosive sublimate of : (a) stannous chloride, (b) hydrogen sulphide ? Equations. 22. What antidote is used in case of poisoning with mercury com- pounds ? Explain. 23. State the chemical nature and the use of each of the following : Thoulet's solution, Nessler's reagent, Millon's reagent, fulminating mer- cury, cinnabar, calomel. 24. Given a solution of mercuric chloride, describe five different ways of showing that mercury is present in the solution. Write the appropri- ate reactions. 25. Explain why mercury, which is so different outwardly from mag- nesium, zinc and cadmium, is nevertheless classified with these metals. 26. How much zinc sulphate can be prepared from 250 grams of pure zinc oxide? 27. Given 2000 grams of corrosive sublimate, how much calomel could be prepared from it? 28. How much mercury would 40 grams of copper precipitate from a solution of corrosive sublimate? 29. How much cadmium sulphide may be formed from 60 grams of cadmium chloride? 30. How much cadmium would 24 grams of magnesium displace from a solution of cadmium sulphate? CHAPTER XXIV SOLUTIONS, ELECTROLYSIS, AND ELECTRO -CHEMICAL THEORIES Nature and Kinds of Solutions. As stated in Chapter I, solutions were formerly regarded as chemical compounds accord- ing to variable proportions. This designation expresses the fact that the process of solution is accompanied by all the phe- nomena that are observed when chemical change takes place, and that the composition of the final homogeneous product obtained may gradually be varied at will within certain limits, as has already been explained. The term solution has of recent years received a somewhat broader meaning than it formerly had. So mixtures of .gases are sometimes spoken of as solu- tions. Glass, various alloys, isomorphous mixtures of crystal- line substances, gases absorbed by solids, etc., are frequently termed solid solutions. A solution is commonly defined as a homogeneous mixture, the constituent parts of which cannot be separated by mechanical means. The term mixture as here used is, however, not to be confused with a mere mechanical mixture like that of sulphur and iron filings ground together, as already remarked in Chapter I. The word mixture, as used in connection with a solution, indicates solely that the proportions of the ingredients may be varied arbitrarily, at least to some extent. Any gas will mix with any other gas or mixture of gases in all proportions, as would naturally be expected, for in a gas the molecules are relatively remote from one another. But gases will not be absorbed by liquids or solids in all proportions. Here the specific nature of the gas and that of the liquid or solid under consideration is of prime importance. The specific nature of liquids and solids is also the determining factor in fixing the solu- bility of liquids or solids in other liquids or solids. Moreover, in all cases of solution, the temperature is very important, being second only to the effect of the nature of the substances. Pressure is also a factor, which is particularly important when gases 432 SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 433 come into play, for these are highly compressible. On the other hand, the pressure factor is of less importance when liquids and solids only are used. Indeed, the factors mentioned are of the same relative importance in the process of solution as in chemical action. Absorption of Gases by Liquids. The amount of a gas absorbed by a liquid increases as the temperature is lowered and as the pressure is increased. A.t constant temperature, the amount of gas absorbed by a liquid is directly proportional to the pressure. This law was discovered in 1803 by Henry, whose name it bears. Henry's law holds only in case the gas is not very soluble in the liquid, like oxygen or hydrogen in water, or nitrogen in alco- hol. When gases are copiously soluble, like hydrochloric acid or ammonia in water, the law does not hold. Gases that follow Henry's law have a very low heat of solution, showing that but little affinity exists between the gas and the liquid. Gases that do not follow the law have a high heat of solution, indicat- ing the prominence of the affinity factor ; for the heat developed is frequently greater than that produced when the dry gas is liquefied by pressure. Thus we have : NH 8 + (aq) = (NH 3 aq) + 8.8 Cal., whereas the heat of liquefaction of 17 grams of ammonia is only 4.4 Cal. Solutions of Liquids in Liquids. Many liquids mix perfectly with one another in all proportions, forming solutions. Thus water will mix with glycerine, alcohol, or acetone ; ether with kerosene or fatty oils ; olive oil with linseed oil ; carbon disul- phide with ether, hydrocarbon oils, or fats. In general, liquids like hydrocarbon oils and their halogen substitution products, ethers, esters, carbon disulphide, and the various fatty arid oily products of plant or animal origin, are readily soluble in one another in all proportions. Liquids that are miscible in all pro- portions are called consolute liquids. Now, hydrocarbon oils and animal and vegetable oils, carbon disulphide, and carbon tetrachloride are practically insoluble in water. When such substances are added to water, the liquids separate in two distinct layers. When the specific gravities of the two non-miscible liquids are nearly alike arid the liquids are shaken together, a mixture of milky appearance is obtained, which on closer inspection is found to consist of small globules 434 OUTLINES OF CHEMISTRY of one liquid suspended in the other. This is an emulsion The greater the difference in specific gravity between the two liquids, the sooner will they again separate into two layers. The specific nature of the liquids also has to do with the length of time the globules remain in suspension, for the greater the ad- hesion between the liquids, the longer the emulsion lasts. In milk, the butter fat is present in the form of minute globules in suspension, that is, in emulsified condition. On standing, the fat gradually comes to the top, forming a layer of cream. By centrifugal force, the constituents of an emulsion can be rapidly separated, which fact is used in separating butter fat from milk by means of the cream separators in use in creameries. A few liquids like alcohol and acetone are consolute either with hydrocarbons and fats or with water. It is to be observed that, when alcohol or acetone is added to water, the resulting solution will then take up much more hydrocarbon or fatty oil than will pure water. A knowledge of this fact is often of practical value. Ether is soluble in water to a limited extent. When an excess of ether is added to water, two layers form, the upper one finally becoming a saturated solution of water in ether and the lower one a saturated solution of ether in water. The solubility in this case, as in all others, is a function of the temperature. Phenol C 6 H 6 OH and water are also only partially miscible at ordinary temperatures; however, at and above 68.4 the two liquids become perfectly consolute. Again, above 20 triethyl- araine N(C 2 H 5 ) 3 and water act like ether and water, that is, they are not consolute. But below 20 triethylamine and water are miscible in all proportions. In this case the miscibility increases as the temperature is lowered, though more frequently the op- posite is true as in the pair, phenol and water. Solutions of Solids in Liquids. Whether a given solid will dis- solve in a given liquid and to what extent, depends primarily upon the specific nature of the two substances, and also upon the tempera- ture. The pressure is also a factor, strictly speaking, but ordi- narily it is of little practical importance. A liquid that dissolves a solid adheres to or wets the latter. Nevertheless, in many cases a liquid will wet a solid without dissolving it appreciably. We may look upon adhesion as an unsuccessful attempt to form a solution, for here the attraction between solid and liquid is SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 435 insufficient to overcome their cohesions and thus form a solu- tion. Many solids have definite solubility in a given liquid, which varies with the temperature. Thus the ordinary salts and other 140 20 40 60 Temperature in degrees C. Fia. 138. 80 100 crystalline solids, like sugar or urea, have a definite solubility in water. On the other hand, substances like gelatine, gum arabic, 436 OUTLINES OF CHEMISTRY and many others of similar non-crystalline character do not have a definite solubility ; they are practically consolute with water. So when water is dropped on glue, the latter swells up, and as more and more water is gradually added all stages of plasticity are obtained, till a thick sirup, and finally a limpid liquid, results. In general, crystalline substances have sharply defined melting points as well as definite solubilities, whereas non-crystalline substances frequently lack definite melting points and solubilities, though this rule is not without its exceptions. The solubilities of solids in liquids vary very greatly with the nature of the substances brought together, as well as with the tem- perature. Some substances, like common salt in water, are about as soluble at lower as at higher temperatures. More frequently, the solubility increases with rise of temper- ature, as in the case of saltpeter in water. The number of grams of a substance dissolved by 100 grams of a given liquid at a certain temperature is called the solubility of that substance at the given temperature. Solubility curves are commonly constructed by charting temperatures as abscissas and solubilities as ordinates. Figure 138 shows a number of solubility curves of well-known salts. It will be observed that these by no means have the same trend. Even much greater varieties than these are met frequently. Degrees of Saturation. In determining the solubility of a solid, the usual practice is to shake the liquid with an excess of the solid till no more is dissolved. This process is conducted at the temperature at which the solubility is sought. The liquid finally obtained is said to be a saturated solution at that temperature. The strength of the saturated solution is found by analyzing a given weight of it. This frequently consists simply of evaporating the solution to dryness and weighing the residue. If the saturated solution thus obtained at a given temperature, say 30, is carefully decanted from the excess of solid substance so that none of the latter is present, and then this liquid is heated above 30, the solution is said to be unsaturated; for at this higher temperature it would take up more of the solid if some of the latter were introduced. On the other hand, if the saturated solution is cooled slightly below its temperature of saturation out of contact with the solid SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 437 it contains, the liquid usually remains clear. The solution is now said to be supersaturated, for when brought into contact with some of the solid (even traces will suffice), more of the latter will drop out and the solution will then become saturated at this lower temperature. This is another method of making a saturated solution. The degree of supersaturation that may be obtained by cooling saturated solutions as described varies in the case of different solutions. The phenomena of supersatu- rated solutions are analogous to those of the supercooling of liquids. So, out of contact with ice, water may be cooled several degrees below zero and still be a liquid. On touching this supercooled water with a piece of ice, however, the whole congeals. Unsaturated, saturated, and supersaturated solutions are similarly obtained in case of solutions of gases in liquids and of liquids in liquids. Solid Solutions. These have already been mentioned. They are commonly prepared by melting two solids together and al- lowing the liquid to congeal. Sometimes by crystallizing cer- tain substances from solution, crystals of variable composition, solid solutions, result. Again, gases are absorbed by solids to form so-called solid solutions. In these cases we may also have either restricted solubility, or miscibility in all proportions. Precipitation. The process of precipitation is the opposite to that of solution. Thus when a salt has been dissolved in water, the latter may be evaporated from the solution till the salt separates out. Again, common salt is soluble in water, but not in alcohol; hence by adding alcohol to salt brine, some of the salt is precipitated. Similarly, camphor is soluble in alcohol but not in water ; consequently camphor is precipitated by adding water to a solution of camphor in alcohol. Sodium chloride is insoluble in liquid hydrochloric acid, hence it is readily compre- hended why the salt is precipitated when hydrochloric gas or even a concentrated solution of it is added to salt brine. Cal- cium phosphate is soluble in nitric or hydrochloric acids but not in water; hence by neutralizing such an acid solution of calcium phosphate with ammonia or caustic soda, calcium phosphate is precipitated. These are all simple, typical illustrations. Precipitation may, however, also le caused by double decom- position in solutions. So silver chloride is precipitated thus : AgNO 3 + NaCl = NaN0 3 + AgCl. 438 OUTLINES OF CHEMISTRY This reaction occurs because under the conditions that obtain silver chloride and the solution of sodium nitrate are the stabler products. The precipitation does not take place simply because there is an opportunity for the insoluble silver chloride to form, it proceeds also because there is the chance to form the stable and soluble sodium nitrate. In other words, reactions of this kind must be considered as a whole. When silver nitrate solu- tion is shaken with carbon tetrachloride, silver chloride does not form. Clearly there is plenty of silver and chlorine to form silver chloride, but the other product would have to be a nitrate of carbon, which is unknown, probably because of its great instability. This lack of action between silver nitrate and carbon tetrachloride is not to be ascribed to the fact that the latter compound is a non- electrolyte, as is often done, for similar changes do occur in the best of insulators, as will be shown below. By heating carbon tetrachloride with silver nitrate and nitric acid to high temperatures in a sealed tube (method of Carius) silver chloride is indeed obtained, for here there are conditions under which carbon is oxidized. Thus carbon tetrachloride is de- stroyed, carbon dioxide, oxides of nitrogen, and silver chloride resulting as the stable products under the conditions to which the mixture was subjected. The method of precipitation by double decomposition in solutions is very commonly employed in chemistry. In all cases the explanation is similar to that of the precipitation of silver chloride. Liquids also may be thrown out of solution. So, many oils are insoluble in water but soluble in alcohol. By adding water to such alcoholic solutions, the oil separates out in the form of a layer. G-ases also may be liberated from their solutions in liquids. In many cases this can be done by simply heating the solution. But it may also be accomplished by adding some other appro- priate liquid, solid, or gas to the solution. So, for instance, by adding concentrated sulphuric acid to a strong solution of hydrochloric acid, hydrochloric acid gas is liberated ; similarly, ammonia is set free when concentrated caustic potash solution is added to a strong aqueous solution of ammonia. Colloidal Solutions. These may be made by dialysis, as al- ready explained. When an electric arc is formed between two rods of silver dipping into water (Fig. 139) a dark-colored liquid SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 439 results which is an extremely finely divided suspension of sil- ver. This is sometimes called a colloidal solution of silver. Similar colloidal solutions of other metals, like platinum or gold, can be formed. In general, colloidal solutions boil at the same temperature as the pure solvent. They vary from mere suspensions to true solu- . . -IT -i i FIG. 139. tions. In living beings, colloidal solutions play an important role. Many colloidal solutions are precipitated by slight additions of various salts, sugar, alcohol, etc. The statement sometimes made that only electrolytes precipitate colloidal solutions is erroneous. For instance, sugar added to colloidal solutions of ferric hydroxide causes precipitation. Boiling Points of Solutions. The boiling point of a liquid is the temperature at which its vapor tension just overcomes atmos- pheric pressure. It consequently varies with the pressure. At constant pressure, the boiling point of a liquid is changed by dissolving substances in it. Thus solutions are obtained which may have a higher or lower boiling point than the solvent, depending on the nature of the substances brought together. So water boiling at 100 and alcohol boiling at 78 yield solu- tions that boil between these temperatures. Formic acid boil- ing at 101 may yield an aqueous solution that boils at 107. Pyridine boiling at 115 and water may yield a solution boiling at 93. Use of Boiling Points of Solutions in Molecular Weight Determi- nations. When a substance that is practically non-volatile, as compared with the solvent liquid, is dissolved in the latter, the solution so obtained boils higher than the pure solvent. So, when sugar is added to water, the solution obtained boils above 100. The rise of the boiling point is proportional to the amount of solute added. Furthermore, the elevations of the boiling point are approximately the same for equimolecular quantities of solute each dissolved in an equal quantity of solvent. This offers a simple method of determining the molecular weights of sub- stances in solution. So 342 grams of sugar C 12 H 22 O n , 60 grams of urea CO(NH 2 ) 2 , 94 grams of glycerine C 3 H 5 (OH) 3 , 120 440 OUTLINES OF CHEMISTRY grams of magnesium sulphate, 159 grams of copper sulphate CuSO 4 , each dissolved in 5000 grams of water, will yield solu- tions that boil at 100.104. This rise of 0.104 in 5000 grams of water would be proportional to a rise of 50 x 0.804 or 5.20 in 100 grams of water. The value 5.20 is termed the molecular rise of the boiling point when water is used as the solvent. For other solvents this value is quite different. For instance, for ether it is 21.6, for alcohol 11.7, for carbon tetrachloride 41.0, for benzene 26.7. It is possible to compute these values from the absolute boiling point T and the latent heat of evaporation L of the solvent by means of the formula, 2 T 2 Molecular rise of boiling point = T 1UU /> The theoretical considerations upon which this formula is based will not be entered into here. To ascertain the molecular weight of any substance soluble in water, dissolve so much of the substance in 5000 grams of water till the solution boils at 100.104. The amount of solute added is its molecular weight in grams. In practice one would, of course, take a much smaller, convenient amount of water and solute to determine the boiling point of the solution, and compute the quantity of solute required to elevate the boiling point 5.2 per every 100 grams of water. Instead of the rise of the boiling point of a solution, the lowering of the vapor tension caused by the introduction of the solute may be used to determine the molecular weight of the solute. The Freezing Point of Solutions. The freezing point of a solution is lower than that of the pure solvent. The lowering of the freezing point is also about the same for equimolecular quantities of solutes each dissolved in an equal amount of solvent. Thus, in aqueous solutions, the molecular lowering of the freez- ing point is about 18.9 per 100 grams of water. The value varies for different solvents, but may be computed approximately from the absolute freezing point and the latent heat of fusion of the solvent according to theoretical considerations of van't Hoff, which will not be reproduced here, the formula being 9 77 2 Molecular lowering of the freezing point = \--. 100 L-, The method of determining molecular weights in solution from SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 441 the lowering of the freezing point is analogous to that of making such estimations from the rise of the boiling point. Discussion of Molecular Weights Determined in Solutions. The freezing and boiling point methods yield only approxi- mate results at best, and can be used with success only in dilute solutions. Special care must be exercised in interpret- ing the results obtained, for they are often anomalous. So, many substances like gelatine, dialyzed silicic acid, or soap when dissolved in \>ater hardly cause any rise of the boiling point. These substances consequently have enormous molecu- lar weights as found by these methods. Such substances are consequently frequently regarded as. highly associated or poly- merized. Again, many substances, among them a large number of the ordinary acids, salts, and bases, when dissolved in water and tested by the freezing or boiling point method, yield much lower molecular weights than the formulae usually ascribed to them indicate. These substances are by some chemists re- garded as dissociated in solution ; they seek to connect the abnormally low molecular weights observed with the fact that such solutions are frequently electrolytic conductors. So, for- instance, common salt, saltpeter, caustic potash, nitric acid, etc., yield molecular weights that are usually somewhat over half of those represented by the formulae of these compounds (see the theory of Arrhenius). Osmosis and Osmotic Pressure. If carbon disulphide (Fig. 140) be placed in a tube and a layer of water B be poured upon it, and again a layer of ether A be carefully poured upon the water, there will, after long standing, eventually be but two layers A' and B' (Fig. 141). The lower layer B f consists of carbon disul- phide and ether saturated with water, whereas the upper layer A f consists of water saturated with carbon disulphide and ether. The change which has FIG. 140. FIG. 141. 442 OUTLINES OF CHEMISTRY occurred is easily explained. Ether dissolves very readily in carbon disulphide, but much less readily in water. Again, carbon disulphide and water hardly dissolve each other at all. In Fig. 140 the water layer B dissolves ether and in turn the ethereal layer A also takes up some water. When the ether has gone into B till it touches the carbon disulphide (7, the latter extracts ether from the aqueous layer B. Thus the carbon disulphide layer Q becomes enriched with ether, whereas the aqueous layer becomes depleted in ether. The depletion is made good by a continuous supply of ether from A till the latter is exhausted and equilibrium is produced. Water charged with ether dissolves carbon disulphide to a greater extent than pure water does, and so the aqueous layer B always contains some carbon disulphide. Moreover, before equilibrium is reached, some of this carbon disulphide is ex- tracted from the layer B by the ethereal layer A. Thus, in the attempt to reach equilibrium, ether is passing from A through B into (7, and on the other hand carbon disulphide is traveling from through B into A. The former current is much the stronger, and so the layer A gradually disappears. The example just cited is a typical case of osmosis. The layer B is the septum which separates the liquids A and C. Whether the latter substances will pass through B or not is determined by the specific nature of B and also that of O and of A. Further- more, the specific nature of the septum B and that of the liquids A and C which bathe it also determines the direction which the major current will take. In all osmotic processes there is a major and a minor current, but in some cases the latter is so slight as to be almost negligible. Under these conditions the osmotic process appears to be one-sided and the septum used is said to be semipermeable. It is clear that if the layer B in Fig. 140 could be held rig- idly in place, pressure would result on the walls of the compart- ment O because of the influx of ether. Figure 142 shows an apparatus of glass containing mercury, carbon disulphide in O, and ether in A, while B is a slice of cork which is saturated with water, and tightly jammed into the tube. As the apparatus stands, the mercury rises in the limb at the right because the major current consists of the passage of ether from A through B into (7, as already explained. The pressure so produced is SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 443 called the osmotic pressure. It is clear thai this pressure is the result of differences in solubility. The apparatus shown in Figure 142 serves only to demonstrate the existence of osmotic pressure, not to measure its value, for the water-soaked cork will give way long before the maxi- mum pressure is reached. Osmosis has been studied for over ;i hundred years. The water, ether, carbon disulphide experiment was performed by L'Hermite in 1854. The importance of osmotic investigations in physiological processes was recognized very early, and physiologists have contributed largely to our knowledge of osmosis. When septa, like parchment paper, or pieces of animal bladder, are used to separate an aqueous sugar solution & from pure water W (Fig. 143), the main osmotic current is from the water through the membrane to the solution, for the latter extracts water from the water- soaked septum. At the same time much sugar passes through the septum into the water W. If, however, a precipitate of copper fer- rocyanide Cu 2 Fe(CN) 6 -# H 2 O is formed in the pores of the septum, almost no sugar passes from the cell into the outer liquid TT; that is to say, copper ferrocyanide is a semipermeable mem- brane, for it allows water to pas? through it, whereas it is nearly im- pervious to sugar. The precipitate of copper ferrocyanide is formed in the membrane by first placing about a three per cent potassium ferrocyanide solution in the cell and then immersing the latter for a time in a solution of copper FIG. 142. 444 OUTLINES OF CHEMISTRY sulphate. of about equal strength. Thus the precipitate forms in the pores of the septum : K 4 Fe(CN) 6 + 2 CuSO 4 = 2 K 2 SO 4 + Cu 2 Fe(CN) 6 . By forming this precipitate in the pores of a small unglazed porcelain cup (Fig. 144), the plant physiologist, Pfeffer, in 1877, measured the maximum osmotic pressure of dilute cane sugar solutions of several concentrations at a number of different temperatures. Figure 144 shows the cell, with manometer M attached, immersed in the large glass dish filled with water. T and T 1 are thermometers. In 1887 van't Hoff showed that Pfeffer's results indi- cate that the osmotic pres- sure is proportional to the absolute temperature, and that the osmotic pressure of the sugar solution is the same as the gas pressure that would be produced if the sugar were in the gase- ous state and confined in the same volume that the solution occupies. In other words, van't Hoff showed that when dilute aqueous cane sugar solutions are separated from water by means of copper ferrocyanide membranes, the osmotic pressures developed may be represented ly the gas equation PVRT, where P is the osmotic pres- sure measured, V the volume of the solution, T the absolute temperature, and R the usual gas constant. More recently, Morse, also using copper ferrocyanide membranes, showed by similar, though far more elaborate, experiments that aqueous cane sugar solutions do develop osmotic pressures that approxi- FIG. 144. SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 445 mately follow the gas laws. On the other hand, in 1906, Kah- lenberg found that vulcanized caoutchouc, so-called sheet rubber, acts as a semipermeable membrane when used to separate cane sugar solutions in pyridine from pure pyridine. Moreover, the osmotic pressures measured in pyridine solutions thus far indi- cate that here the gas laws do not hold at all. Experimental osmotic investigations are still being pursued at the pres- ent time, and it is to be hoped that they will cast more light upon the laws that regulate osmotic pressure in the various cases. The subject is by no means as simple as it was formerly regarded by physical chemists, for specific selective action on the part of the so-called semipermeable membrane comes strongly into play. Solutions having the same freezing or boiling point are fre- quently termed isosmotic or isotonic, for they are said to have the same osmotic pressure. The latter is sometimes computed on the basis of the assumption that the gas laws hold univer- sally for osmotic pressures of all solutions, provided that the membrane is semipermeable. This assumption is, however, not justified by known experimental facts. Indeed, in speaking of the osmotic pressure of a solution it is always necessary to specify what membrane separates that solution from the pure solvent. A more detailed consideration of osmotic processes belongs to the subject of physical chemistry. Electrolysis. There arc two kinds of conductors of electricity: (1) those that show no chemical change as electricity passes through them, and (2) those that conduct with concomitant chem- ical decomposition. To the conductors of the first class belong all metals, alloys, and graphite. A few other solids, like lead per- oxide, manganese peroxide, pyrite, and other native sulphides and arsenides also conduct slightly and probably without de- composition. Those conductors that are decomposed chemi- cally by the electric current are said to conduct electrolytically. They are called electrolytes or conductors of the second class. To this class belong many metallic oxides, hydroxides, sul- phides, chlorides, nitrates, sulphates, carbonates, and silicates, as well as other salts of metallic bases, when these compounds are in molten condition. Some of these substances even con- duct electrolytically to an appreciable degree before they are actually molten. So, for instance, at room temperatures a 446 OUTLINES OF CHEMISTRY block of rock salt is a non-conductor ; but on heating it strongly, it begins to conduct quite noticeably before it is actually melted. Again, the oxides of the earths are non- conductors at room temperatures, whereas at higher tempera tures they conduct, as is well known in the case of the glower of the Nernst lamp. Water, alcohols, ethers, acids, esters, aldehydes, ketones, mercaptans, sulphur ethers, hydrocarbons and their halogen substitution products, fats, oils, waxes, pitch, resin, caoutchouc, and the chlorides, bromides, iodides, sul- phides, oxides, and hydrides of non-metallic elements, when pure are practically all either exceedingly poor conductors or insulators, whether they are liquefied or not. On the other hand, solutions of many acids, bases, and salts in water, liquid ammonia, amines, liquid hydrocyanic acid, alcohols, esters, ketones, sulphur dioxide, etc., are very good electrolytes. This is particularly the case with the ordinary acids, salts, and bases in aqueous solutions; but it would be quite wrong to think that electrolytic conduction is confined to aqueous solutions, for some non-aqueous solutions conduct quite as well and even better than those in which water is the solvent. Furthermore, many salts when in the liquid condition or in solution are insulators. So the oleates, palmitates, and stearates of the heavy metals are insulators whether solid, molten, or dissolved, though, to be sure, they do not happen to dissolve in water. Pure tetra- chloride of tin SnCl 4 is an insulator at all temperatures from its freezing point to its critical temperature. The analogous chlorides, SiCl 4 , TiCl 4 , and CC1 4 , are also insulators. Since the days of Michael Faraday (1791-1867), who first investigated the subject, a large amount of data on electrolytic conductors has been gathered, and yet no way has been found to foretell with certainty whether a new compound will prove to be a conductor or not. The only way is to actually test the compound with the electric current itself. All highly rarefied gases conduct electricity ; and in some gases, like hydrochloric acid, for instance, this conductivity has been found to be accompanied with chemical decomposition. Electrolysis was discovered in 1800 by Nicholson and Carlisle, who decomposed water by means of the electric current. This new means of effecting chemical decomposition was studied by Berzelius and Hisinger, Sir Humphry Davy, and especially by SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 447 Faraday. The latter called substances that can be decomposed by the electric current electrolytes; while the process of such decomposition he named electrolysis. In electrolysis two plates, usually consisting of metal or conducting carbon, dip into the electrolyte. These plates Faraday termed the electrodes. The plate from which the positive current passes into the electrolyte he called the anode, and the other electrode, by which the cur- rent leaves the electrolyte, he called the cathode. It is a re- markable fact that the products of decomposition first appear right on the surface of the electrodes. As these products are eliminated or produced at the electrodes, the corresponding material in the electrolyte is continually moving towards the electrodes. The particles which thus move towards the elec- trodes during electrolysis Faraday termed the ions. Those particles that move towards the anode he termed anions, and those that move towards the cathode the cations. To Faraday, the simplest conceivable case of electrolysis consisted of two silver electrodes dipping into molten silver chloride. As a current passes through this electrolyte, silver is deposited on the cathode, and chlorine is simultaneously produced at the anode, which is thus at once attacked, forming silver chloride. Thus, while the cathode is increasing in thickness, the anode is wear- ing away. This process is, of course, accompanied by a con- tinuous movement of the cations, silver, toward the cathode, and of the anions, chlorine, toward the anode. If the anode consisted of carbon, the chlorine would appear as a gas. Now, why should molten silver chloride thus conduct electrolytically and molten tetrachloride of tin not? This is a question which we cannot yet answer, any more than we can tell why a piece of silver conducts electricity and a piece of sulphur is an in- sulator. In solutions, electrolysis is precisely the same as in molten electrolytes, though, to be sure, the products eliminated at the electrodes in solutions react with the electrolyte more frequently because of the more complex nature of the latter. So, on electrolysis of an aqueous sodium sulphate solution, hydrogen appears at the cathode and oxygen at the anode. This is because the sodium eliminated at the cathode at once reacts with the water, forming caustic soda and hydrogen, which is consequently a secondary product. If a mercury cathode is used, the sodium liberated dissolves in the latter and no hydro- 448 OUTLINES OF CHEMISTRY gen appears. At the anode SO 4 is liberated, which, however, at once reacts with water, yielding sulphuric acid and oxygen, which is consequently of secondary origin, like the hydrogen at the cathode. In 1833 Faraday demonstrated that the passage of the electric current through an electrolyte is always accompanied ly the ap- pearance of decomposition products at the electrodes, and that the amount of such decomposition is proportional to the current. Furthermore, he found that chemically equivalent amounts of substances are separated out from different electrolytes ly the same amount of current. These facts are commonly known as Faraday's law. They are fundamental in all electrolytic work. Faraday tested his law on molten electrolytes as well as on aqueous solutions, and its validity has since been confirmed by many careful investigations. In 1900 Kahlenberg showed that the law holds also in non-aqueous solutions. A current of 1 ampere will effect the deposition of a gram-equivalent of a substance in 96,540 seconds ; that is to say, whenever 96,540 coulombs of electricity pass through an electrolyte, a gram-equivalent of decomposition product is deposited at each electrode. The quantity 96,540 coulombs is known as the con- stant of Faraday's law ; it is now generally called a faraday. Electrolytic Theories. The Grotthus Theory. The first ex- planation of the process of electrolysis was made by Grotthus, in 1805. He assumed that each molecule of an electrolyte possesses electrical polarity similar to that of a magnet, and that these molecules are irregularly distributed throughout the electrolyte. On closing the electric circuit, the + poles of the molecules would be attracted by the electrode, and the poles of the molecules by the -f electrode. Thus the mole- cules in the electrolyte would all arrange themselves with their + poles directed toward the electrode, and their poles toward the + electrode. The molecules in actual contact with the electrodes were then conceived as decomposed by the attrac- tion of the -f- electrode for the part of the molecule, and of the electrode for the -f- part of the molecule. So with molten silver chloride as electrolyte between two platinum electrodes (Fig. 145), after closing the circuit, the arrangement would be as shown in line (1). A moment later silver would be depos- ; ted at and chlorine at A, and line (2) would represent the SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 449 condition of the remaining molecules. A general decomposition all along the line would then take place, so that the arrange- ment in line (3) would result. Then these molecules would turn on an axis, their poles heading toward the + electrode, llh= QJ2LJ2J1 W H-H H H-H (1) (2) (3) FIG. 145. and the + poles toward the electrode, as in line (1), after which the whole process would repeat itself. Faraday's View. This mechanical explanation of Grotthus held its own for over half a century, though Faraday protested that any explanation that the molecules were rent asunder by an attraction of the electrodes for the molecules was untenable. According to Faraday, all that can be maintained is that the electric current acts as an axis of force, ejecting the decomposi- tion products at the electrodes, which simply serve as doors for the exit of the products. Clausius's Theory. In 1856 Robert Clausius called special attention to the fact that the theory of Grotthus does not explain why a small electro-motive force suffices to send an appreciable current through an electrolyte, though, to be sure, such current might pass for but a short time. It was at the time when the kinetic theory of gases and the mechanical theory of heat were taking form. In these theories Clausius was interested. He consequently naturally assumed the mole- cules of an electrolyte to be in motion due to heat, just as the molecules of a gas are supposed to be in motion. He further stated that sometimes some of the molecules would collide in such a way that the basic and acid radicals of two different molecules would unite to form a new molecule, thus leaving a basic and an acid radical in an uncombined or free state for 450 OUTLINES OF CHEMISTRY a moment. At any instant, then, there would be a certain number of free positive basic radicals and free negative radi- cals, and these would be separated at the electrodes by electrical attraction, as Grotthus explained. However, Clausius did not assume that the molecules ever arranged themselves definitely in the electrolyte. He thought that the electric current merely directed the general trend of the decomposition within the electrolyte. Moreover, Clausius held that the free, or unde- composed, parts of an electrolyte would at any moment not amount to more than a small fraction of 1 per cent. This distinguishes Clausius's theory from that of Arrhenius. Arrhenius's Theory. The theory of Arrhenius, also known as the theory of electrolytic dissociation, and frequently nowa- days designated as the ionic theory, is founded upon a supposed connection between the vapor tensions, boiling points, or freez- ing points of dilute solutions of electrolytes, on the one hand, and their electrolytic conductivity on the other hand. In 1887 Arrhenius published a series of data intended to show that dilute solutions of electrolytes lower the freezing point to a much greater extent than similar solutions of non-electrolytes containing equimolecular quantities, and that consequently the solutions that conduct the current must contain relatively more dissolved molecules; that is, the latter must be dissociated. Because this dissociation is supposed to occur in electrolytes, it has been named electrolytic dissociation. Arrhenius assumes that when an acid, base, or salt is dissolved, yielding a solution that conducts the current, the molecules of the dissolved sub- stance are by the very act of solution decomposed into part molecules, which are charged with electricity, thus : H 2 SO 4 ^ H + HSO 4 or H + H + SO 4 . etc. In each case the dissociation from left to right is supposed to be complete in infinitely dilute solutions. At ^ny finite concen- SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 451 tration the solution contains a certain percentage of undissociatcd molecules and a certain percentage of charged part molecules. The latter according to Arrhenius are the only particles in the solution that are active in conducting the current. He calls these part molecules the ions, and conceives each gram-equiva- lent of cations charged with 96,540 coulombs of positive electric- ity and each gram-equivalent of anions charged with the same amount of negative electricity, postulating further that any conducting solution contains equivalent amounts of cations and anions. It will be observed that in the Arrhenius theory the word ion is used in a different sense from that proposed by Faraday, who regarded the ions as the substances that migrate toward the electrodes during actual electrolysis, not as part molecules charged with electricity which are at all times present in an electrolyte. The explanation of the passage of the elec- tric current through an electrolyte on the basis of Arrhenius's theory is simple, consisting merely of the movement of the free charged cations and anions toward the cathode and anode respectively, under the influence of the impressed electro-motive force (Fig. 146). This explanation is then essentially the same <9 G GCZJGG 6> E> Q) FIG. 146. as that given by Clausius. Indeed, the main difference between the Clausius and Arrhenius theories is that the latter assumes the presence of a very much larger percentage of dissociation, and claims this may be computed from either conductivity or freezing point or boiling point measurements of dilute solutions. According to Arrhenius's theory all the physical, chemical, and physiological properties of solutions that are electrolytes are 452 OUTLINES OF CHEMISTRY determined by the properties of the undissociated molecules of solute and those of the ions. The latter are the chief factor in dilute solutions where dissociation has often progressed to the extent of 80 per cent or more. Thus, the adherents of this theory hold that copper ions are blue, cobalt ions are red, MnO 4 ions are purple, sodium ions are colorless, etc., since aqueous solutions of copper salts are blue, those of cobalt are red, those of permanganates purple, and those of sodium chlo- ride, etc., colorless. Again, the sour taste and other acid properties of electrolytic solutions of acids would be due to hydrogen ions, and the alkalinity of caustic alkali solutions to hydroxyl ions. Indeed, in terms of the theory of electrolytic disso- ciation an acid would be defined as a substance capable of yielding hydrogen ions, while a base would be a substance capable of yield- ing hydroxyl ions. Moreover, the act of neutralization in dilute solutions would consist essentially of the union of hydrogen and hydroxyl ions to form undissociated water, thus: K, O~H + H, Cl = K, Cl + H 2 O. Since K and 01 appear on both sides of the equation, the latter might even be written : OH + H = H 2 0. In the chapter on thermochemistry it has been stated that the heat of neutralization of strong acids by strong bases is approxi- mately the same for all. This fact clearly would readily be explained by the above assumption that the neutralization in all cases consists essentially of a union of H and OH. On the other hand, it is to be noted that the heats of neutralization of some of the weaker acids, like oxalic and acetic acids, are nearly the same as that of hydrochloric acid ; though according to the theory, the latter is far more highly dissociated. Again, the fact that two solutions of neutral salts which form no precipitate when mixed also exhibit no change 'of temperature on being poured together (so-called law of thermoneutrality of Hess) is readily explained by Arrhenius's hypothesis, for there would be no action to occasion a thermal effect. However, when similar salt solutions that are non-electrolytes are mixed, there is also no thermal effect. Precipitation in electrolytic solutions by double decomposi- SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 453 tion is explained on the basis of the Arrhenius theory by saying that certain ions meet to form insoluble compounds. The pre- cipitation of silver chloride when silver nitrate and common salt solutions are brought together would be expressed as follows : + + Ag, NO 3 + Na, Cl = Na, NO 3 + AgCl. In other words, the chlorine ions and silver ions meet to form neutral silver chloride, which is difficultly soluble. Since all conducting chloride solutions are supposed to contain free chlorine ions, it is clear that these would all precipitate silver chloride from solutions of silver salts that contain silver ions. Similarly, the fact that barium sulphate is precipitated when a solution of a barium salt and one of a soluble sulphate are brought together finds a ready explanation, etc. At one time the view was even advocated that all chemical reactions, or at least those that take place instantaneously, occur only between free ions, that is to say, they occur only in electrolytes. Now the fact is that all chemical reactions known to occur in solutions that are electrolytes can be reproduced, as to type, in solutions that are the best of insulators. So, for instance, from copper oleate solutions in hydrocarbon oils, brown cupric chlo- ride is instantly precipitated by adding any of the following chlorides, also dissolved in the same hydrocarbon : HC1, PC1 3 , SriCl 4 , SiCl 4 , SbCl 3 , etc., though these solutions are all non- electrolytes. Lead will precipitate copper from insulating copper oleate solutions just as zinc precipitates copper from aqueous copper sulphate solutions. Furthermore, the colors exhibited, for example, by solutions of copper oleate, cobalt oleate, nickel oleate, etc., each dissolved in a hydrocarbon like benzene, toluene, etc., are entirely similar to the colors of aqueous electrolytic solutions of copper and cobalt salts respectively. In fact all the physical and chemical properties exhibited by salt solutions that are electrolytes can be duplicated in salt solutions that are insulators, except the phenomena of elec- trolysis themselves. But the distinguishing feature of Arrhenius's theory is the claim that there is a quantitative relation between lowering of the freezing point or elevation of the boiling point and the elec- trical conductivity of solutions. It must be stated that the numerous cases thus far adduced to support this claim are not 454 OUTLINES OF CHEMISTRY at all conclusive, for they show variations that are far beyond the limits of experimental errors. Furthermore, solutions of mag- nesium sulphate as well as those of all the other vitriols are good electrolytes, whereas according to the freezing points of their solutions they ought to be non-electrolytes. Again, soap solu- tions boil at practically the same temperature as water, yet they conduct electricity well. And so numerous other cases might be cited, the field of non -aqueous solutions being especially re- plete with such. It should be stated that the behavior of elec- trolytes is in general not in harmony with the law of mass action, as ought to be the case if Arrhenius's theory were tenable. A more detailed account of this interesting theory cannot be given here, for it belongs to the subject of physical chemistry. Suffice it to say that, taking all known facts into consideration, the theory of Arrhenius is in the author's opinion untenable. This view was also clearly voiced by Mendeleeff in the last edi- 'tion of his great work on the " Principles of Chemistry," when, referring to the theory of electrolytic dissociation, he said, " I do not consider the hypothesis in question to be in accordance with fact, and therefore refrain from giving a detailed exposition of it in this work." The reader will have no difficulty in compre- hending books that still use the nomenclature of the theory of electrolytic dissociation by remembering that the term ion as used in expressing chemical changes means the same as atom or radical, the atoms or radicals being simply thought of as charged with electricity, as already explained. The electron theory considers electricity itself to be material in character and to consist of corpuscles or electrons that weigh about 0.0005 as much as a hydrogen atom. This theory has developed from a study of radium rays and the discharge of electricity through rarefied gases. The electrons are considered to be negative electricity itself. Positive electrons appear to be much more difficult to isolate. J. J. Thomson has advanced a theory that the atoms of the various elements are composed entirely of electrons, and has shown that, on the basis of such an assumption, the properties of the elements would exhibit periodicity as indicated by the periodic system. The electron theory has thus far not proved to be of special value in chemistry. Electric Batteries are contrivances for converting chemical energy into electrical energy. So when a zinc plate and a SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 455 graphite plate are dipped into dilute sulphuric acid, as shown in Fig. 147, zinc dissolves and hydrogen is liberated on the graph- ite ; while at the same time an electric current passes through the solution from the zinc to the graphite and from the latter through the wire to the zinc. Some hydrogen also appears at the zinc, though in very small amounts if the zinc is pure or if it is amalgamated. The chemical action that takes place is: Zn + H 2 SO 4 = ZnSO 4 H 2 . FIG. 147. But the zinc dissolves at one plate and the hydrogen appears at the other, while in the middle of the electrolyte there is no visible change for some time. The ex- planation of the passage of the current through the electrolyte is the same as in the case of electrolysis. In a battery, however, no external electro-motive force is used to force the current through the electrolyte. Batteries develop electro-motive force of their own because of the chemical affinity that comes into play ; and the electrical energy developed by batteries comes from the energy of the chemical changes that take place while the battery is in action. The cell above mentioned will exhibit an E. M. F. of about 1.3 volts, which on closed circuit soon drops rapidly because of the accumulation of hy- drogen on the graphite, for this hydrogen- laden plate produces a counter E. M. F. Any two different conductors of the first class dipping into an electrolyte will show an E. M. F. when connected with a voltmeter as shown in Fig. 148. A complete considera- tion of electric batteries belongs to the sub- ject of electro-chemistry and cannot be entered into here. The ordinary batteries used for ringing doorbells consist of zinc and carbon dipping into a concentrated solution of ammonium chloride. As the battery acts, zinc is x: I \ 1 ;^ "^L > 1 1 \ 1 \ FIG. 148. 456 OUTLINES Ob CHEMISTRY dissolved, forming zinc chloride, while hydrogen is liberated on the carbon, thus : Zn + 2 NH 4 C1 + 2 H 2 O = ZnCl 2 + 2 NH 4 OH + H a . The ammonium hydroxide forms at the carbon. As these bat- teries are used only occasionally, and then only for a short time, this hydrogen ordinarily escapes while the battery is at rest, thus prolonging its life and usefulness. The dry batteries in use com- monly consist of zinc, ammonium chloride solution, and carbon, the latter being surrounded with coarsely powdered manganese dioxide which serves to oxidize the hydrogen liberated. These batteries are not perfectly dry, as their name would indicate, but they contain enough plaster of Paris to solidify their contents. Another battery, which frequently is used in telegraphic work, consists of zinc dipping into dilute sulphuric acid, and cop- per surrounded by a saturated solution of copper sulphate. The two solutions, being of dif- ferent density, are kept separate by gravity (Fig. 149). The battery must be kept in use all the time, however, or the solu- tions will diffuse into each other, and the copper sulphate solution will reach the zinc and react with it, forming zinc sulphate and copper, which by coating the zinc would spoil the battery. The E. M. F. of this battery, known as the blue cup battery, is about 1.1 volts. The ordinary storage battery consists of two lead plates, one of which is coated with lead peroxide, dipping into a solution of sulphuric acid of 1.2 specific gravity. As the battery acts, lead sulphate is formed. The complete change may be represented thus : Pb + H 2 S0 4 = PbS0 4 + 2 H, and PbO 2 + 2 H + H 2 SO 4 = PbSO 4 + 2 H 2 O ; or Pb + 2 H a SO i + PbO 3 = 2 PbSO 4 + 2 H a O. FIG. 149. SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 457 At the cathode lead dissolves, while at the anode the hydrogen liberated reduces the lead peroxide, which is then acted upon by sulphuric acid, forming lead sulphate. The E. M. F. of the storage cell is about 2 volts. When in use its voltage decreases. When the E. M. F. has run down to 1.8 volts, the battery should be recharged. This is done by passing a current from a dynamo through the cell in the opposite direction from that produced by the battery when in action. In the charging pro- cess the reactions just given are reversed; that is, lead is depos- ited at the cathode and lead peroxide is formed at the anode. Electro-chemical Series of the Metals. If in the battery zinc, dilute sulphuric acid, carbon (Fig. 147), the zinc be replaced successively by iron, copper, silver, and gold, the voltage will diminish. The order in which the metals are mentioned is the same as that in which they will replace one another in solutions of their salts. All of the metals may thus be arranged in a series, beginning with the most basic, and ending with the least basic ; or, as it is sometimes stated, beginning with the most electro-positive metal and ending with the least electro-positive. Such a series is called the electro-chemical series of the metals. The series varies somewhat for different solutions, but the usual order for the common metals is about as follows : K, Na, Ba, Sr, Ca, Mg, Al, Mn, Zn, Cd', Fe, Co, Ni, Pb, Bi, Sb, Sn, Cu, Hg, Ag, Pt, Au. This is also called arranging the metals in the order of their electrolytic solution tensions by those who think of the metals as having a tension or tendency to form ions in the sense of the theory of Arrhenius. The ease with which the metals are separated from electro- lytes by electrolysis is in the reverse order of that given above. REVIEW QUESTIONS 1. What is the difference between a solution and a mechanical mix- ture? Illustrate by giving five examples of each. . 2. In what respects does a solution differ from a chemical compound ? Give three examples of each by way of illustration. 3. What are the factors that determine the amount of a gas that will be absorbed by a given liquid? Give a concrete example. State Henry's law and discuss its exceptions. 4. Why is a solution of hydrochloric acid gas in water heavier than water and a solution of ammonia gas in water lighter than water? 5. What is meant by the term "solid solution"? Give three ex- amples of solid solutions. 458 OUTLINES OF CHEMISTRY 6. What are consolute liquids ? Give six examples. 7. What determines the solubility of one liquid in another? Of a solid in a given liquid ? 8. What is the difference between a solution and an emulsion ? 9. What is a solubility curve ? What fact does it show ? 10. Compare supercooled liquids with supersaturated solutions. Ex- amples. 11. Given a 20 per cent solution of cane sugar in water ; describe two essentially different methods of obtaining solid sugar from the solution. 12. Give three different methods by means of which solid salt may be obtained from a 15 per cent solution of common salt. Explain each case. 13. How does a solution of barium chloride react with one of sodium sulphate ? Write the equation. Compare this case of precipitation with those you have described in answer to questions 11 and 12. 14. Mention two distinctly different methods of obtaining hydro- chloric acid gas from a concentrated aqueous solution of that substance. 15. How form a solution of colloidal gold? Mention four other colloidal solutions. How do they differ from ordinary solutions? 16. Mention three methods of determining the molecular weight of grape sugar in solution. What results would be obtained if these methods were applied to : (a) a solution of common salt, (6) a solution of Epsom salt, (c) a soap solution ? Explain these peculiar results from the stand- point of the theory of Arrhenius. What other explanation may be given ? 17. Define osmosis and osmotic pressure, giving an illustration of each. 18. What are isotonic solutions ? 19. Distinguish between conductors of the first class and conductors of the second class and give ten illustrations of each. 20. Define the following terms : electrolyte, electrolysis, electrode, cathode, anode, ion, cation, anion. Give an example in each case. 21. What is Faraday 'slaw? 22. Explain the electrolysis of a copper chloride solution by: (a) the Grotthus theory, (6) the Clausius theory, (c) the Arrhenius theory. 23. State the essential features of the theory of electrolytic dissocia- tion. By what other name is this theory also known ? Using this theory, write the equations expressing the following reactions : (a) the neu- tralization of nitric acid by sodium hydroxide, (6) the precipitation of lead sulphate by pouring together solutions of lead nitrate and potassium sulphate, (c) the precipitation of silver iodide, by mixing solutions of potassium iodide and silver nitrate. 24. Upon what facts is the theory of Arrhenius based? What objec- tions are there to this theory? 25. What is the so-called electron theory? 26. What is an electric battery? Give four concrete illustrations. 27. What is the essential difference between a primary battery and a storage battery? What happens when a lead storage battery is charged? SOLUTIONS AND ELECTRO-CHEMICAL THEORIES 459 28. What is meant by the term electrochemical series of the metals? 29. How much copper would a current of one ampere precipitate from a copper sulphate solution in ten hours ? How much silver would be pre- cipitated by this current ? 30. How much zinc is required to precipitate ten grams of silver from a silver nitrate solution ? CHAPTER XXV COPPER, SILVER, AND GOLD COPPER, silver, and gold occur in nature in the uncombined state, and consequently have been known to man since earliest times. These metals have a high specific gravity, are very malleable and ductile, and most excellent conductors of heat and electricity. The atomic weights of the elements of this group are relatively high: Cu, 63.57; Ag, 107.88; Au, 197.2. Chemically these metals are rather inert, and their chemical activity decreases as the atomic weight increases. Like mer- cury, copper forms two series of compounds, the cuprous, in which copper is univalent, and the cupric, in which it is biva- lent. Silver is practically always univalent in its compounds. Gold, on the other hand, forms aurous compounds, in which the metal is univalent, and auric compounds, in which it is trivalent. The compounds of copper, silver, and gold in which these metals are univalent are analogous to those of the alkali metals, which are also univalent. On the other hand, copper in cupric compounds, in which it is bivalent, is analogous to mercury in mercuric compounds. Copper and gold are the only colored metals known. Copper, silver, and gold are used in all civilized countries in making coins. Occurrence, Metallurgy, and Properties of Copper. In the uncombined condition copper is found in large quantities near Lake Superior. It also occurs in the Urals, in Sweden, Japan, and China. Large amounts of chalcocite or copper glance Cu 2 S, chalcopyrite or copper pyrites Cu 2 S Fe 2 S 3 , and bornite (Cu 2 S) 3 Fe 2 S 3 occur in Montana, where they are mined and smelted for copper. Other important copper ores are ruby copper Cu 2 O, malachite CuCO 3 Cu(OH) 2 , and azurite Cu(OH) 2 2 CuCO 3 . Malachite occurs especially in Siberia. In extremely small amounts copper compounds are also some- times found in plants and animals. Thus, in plants growing in copper- bearing regions, copper is frequently met; similarly 460 COFFEE, SILVER, AND GOLD 461 in oysters ; also in the feathers of some birds, like those of the genus turacus. In extracting copper from its ores, the process is simple when the ores are oxides or carbonates, for then all that is nec- essary is to heat the ore with coke or coal in a blast furnace ; the reaction involved is : Cu 2 0-hC = 2Cu + CO. But if the ores contain sulphides, which is commonly the case, the process is much more difficult, for iron and other impurities, like lead, arsenic, and antimony, besides sulphur, must be elim- inated. For ores rich in copper a dry process is used, whereas for ores that have a low copper content the wet process is com- monly adopted. In the dry process, the ores are roasted to convert most of the sulphides into oxides. In this way, sulphur burns to sul- phur dioxide, and iron to iron oxides to a large extent. The mass is then mixed with carbon and silicates rich in silica, and heated in a blast furnace. Thus iron enters the slag as a sili- cate, which floats on top and is tapped off. In the process, copper oxides are reduced, the copper formed uniting with any sulphides still present, forming a heavy liquid which accumu- lates at the bottom of the iurnace. This liquid is run off, fre- quently into water so as to granulate the product, which is called copper matte, and consists mainly of copper and iron sulphides containing about 33 per cent copper. This copper matte is then roasted, and again melted with silicates and carbon. Thus more iron is removed as silicate, and copper and other ox- ides are reduced to metal containing 90 to 95 per cent copper. This product is then melted on the hearth of a furnace in a current of air, whereby metals like iron, arsenic, lead, and antimony are oxidized before copper, and these oxides either pass off as vapors or float on the surface and are removed. Ths copper is also slightly oxidized, and this oxide reacts with any cuprous sulphide left over, forming copper and sulphur dioxide, thus : Cu 2 S + 2 CuO = 4 Cu + SO 2 . Finally, by adding carbon, any further copper oxide is reduced to metal. Bessemer converters (which see) are sometimes used in effecting the oxidation of the ores or the copper matta 462 OUTLINES OF CHEMISTRY Air containing sand or other finely divided silicates is blown into these converters containing the molten matte. Thus the process above described is conducted in practically one opera- tion. In the wet process, ores low in copper, like oxidized pyrites residues from sulphuric acid factories, are roasted with common salt, thus forming cupric chloride, which is leached out with water. From the clear solution, copper is precipitated by scrap iron, thus : Or the ores are moistened with water and exposed to the air ; thus ferric sulphate gradually forms, in a solution of which copper is soluble. This solution is then treated with scrap iron to precipitate the copper. The supernatant solution, con- sisting essentially of ferrous sulphate, is again poured upon fresh ore, and so on. Electrolysis is used in copper refining. The impure copper from the above processes is melted and cast into thick plates. These serve as anodes, a copper sulphate solution is used as the electrolyte, and thin copper plates serve as cathodes. The copper is easily eroded from the anode and deposited on the cathode. As the process virtually consists of simply trans- ferring copper from anode to cathode, an E. M. F. below one volt is sufficient to keep the current flowing. The deposited copper is stripped off from the cathodes when they have become sufficiently thick. As the impure copper anodes are dissolving, the impurities, like copper sulphide, silver, gold, bismuth, and lead, fall to the bottom of the vats, forming a thick mud or slime, from which the two precious metals mentioned are sep- arated. Electrolytically refined copper is in a high state of purity. Its malleability, ductility, and electrical conductance are much greater than that of impure copper. The United States produces about 620,000 tons of copper each year, which is nearly two thirds of the world's annual output. Copper is a rather hard, tough, very ductile and malleable metal of reddish brown color and high metallic luster. It may be crystallized in regular jctahedra or cubes. In dry air it re- mains unchanged. In moist air it gradually becomes coated with a greenish basic copper carbonate known as verdigris. The specific gravity of copper is 8.94. Its melting point is 1084, COPPER, SILVER, AND GOLD 463 and its boiling-point 1500. Just below its melting point copper becomes so brittle that it may be pulverized. At red heat it may be welded. Molten copper absorbs gases like car- bon monoxide, hydrogen, and sulphur dioxide, which are given off with some violence on cooling. Copper is readily oxidized by heating it in the air. It imparts a green color to the Bunsen flame. In dilute sulphuric acid, copper is insoluble, but the hot concentrated acid attacks it, forming copper sulphate, sulphur dioxide, and water. In dilute nitric acid copper is readily soluble : 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO. Cold hydrochloric acid is practically without action on copper, but the hot concentrated acid attacks the metal somewhat, hydrogen and cuprous chloride being formed. Ammonia water slowly dissolves copper in contact with air. The deep blue solution thus obtained, known as Schweitzer's reagent, is a good solvent for cellulose. Many dilute acids also gradually act on copper, in presence of oxygen or air, which fact is important, since copper compounds are poisonous and copper is often used in making 1 cooking pots, kettles, and other utensils. Large quantities of copper are used for electric wires and cables, for covering the bottoms of ships, for coins, for roofing and archi- tectural ornamentation, for electroplating, and for a number of important alloys. Copper salts are important as germicides in agriculture. From solutions of its salts, copper is precipitated by magnesium, zinc, iron, phosphorus, and various other reduc- ing agents. Alloys of Copper. Most of the copper produced is used in the form of alloys. Pure copper does not make good castings, for on cooling it contracts and does not fill the molds. Many alloys of copper, however, are excellent for making castings, and they are constantly of very great importance. The alloys are made by melting the metals together in the proportions required. Copper readily forms alloys with most of the metals. Among the most important of these alloys are the following: Brass is a golden yellow alloy consisting of 2 parts copper and 1 part zinc. Other proportions may be used, however; thus 5 parts copper and 1 part zinc yield a reddish brass known as tombac, Dutch brass, or Dutch metal. Muntz metal contains 464 OUTLINES OF CHEMISTRY 3 parts copper and 1 part zinc. The larger the content of zinc, the paler the color of the brass obtained ; the varieties of brass on the market commonly contain from 18 to 50 per cent zinc. Brass is harder than copper. It is nevertheless malleable and ductile and makes good castings. It can also readily be worked in a lathe. German silver consists of about 80 to 95 per cent brass plus 5 to 10 per cent nickel. Bronzes are alloys of tin and copper. Gun metal consists of about 9 parts copper and 1 part tin ; bell metal of 3 parts copper and 1 part tin. Many bronzes also contain some zinc and lead. This is particularly the case with bronze used for statuary, which consists of from 80 to 90 parts copper, 3 to 8 parts tin, 1 to 3 parts lead, and 1 to 10 parts zinc. Bronze used in ma- chinery commonly contains rather more lead than this, and correspondingly less copper. Phosphor bronze is bronze to which from 0.5 to 3 per cent phosphorus has been added. It is particu- larly hard, and is often used in making machinery. Aluminum bronze consists of from 5 to 10 per cent aluminum and 90 to 95 per cent copper. It has a golden yellow color and is quite hard. Oxides of Copper. Cuprous oxide Cu 2 O is obtained as one of the products of the incomplete oxidation of copper when the latter is slightly heated in the air. By reducing hot, alkaline solutions of cupric salts with glucose, cuprous oxide is readily obtained; also by heating cuprous chloride with sodium car- bonate, thus : 2 CuCl + Na 2 CO 3 = Cu 2 O + CO 2 + 2 NaCl. Cuprous oxide is a bright red, crystalline powder which remains unchanged in the air. Dilute acids convert it into cupric salts and copper, thus : Cu 2 O + 2 HNO 3 = Cu(NO 3 ) 2 + H 2 O + Cu. Cu 2 O + H 2 SO 4 = CuSO 4 + H 2 O + Cu. Cuprous oxide is fused with glass to give the latter a red color. Cupric oxide CuO is a black powder formed by heating copper in oxygen or in the air. It is also obtained by ignition of the nitrate or carbonate. It readily gives up its oxygen when heated with carbon or carbon compounds ; hence the use of cupric oxide in the analysis of organic compounds. Cuprous hydroxide is not known with certainty. Cupric COPPER, SILVER, AND GOLD 465 hydroxide Cu(OH) 2 is an amorphous, blue precipitate formed when caustic alkali is added to a solution of a cupric salt. The precipitate readily loses water on boiling it in the solution from which it has separated, thus forming a black hydrated oxide of approximately the composition Cu(OH) 2 2 CuO. In ammonia water cupric hydroxide is soluble, yielding the deep blue solution called Schweitzer's reagent. When treated with acids, cupric hydroxide yields cupric salts. Halides of Copper. Cuprous chloride CuCl is formed by boil- ing cupric chloride CuCl 2 , with metallic copper and hydrochloric acid : CuCl 2 + Cu = 2 CuCl. The cuprous chloride separates out in the form of a white, finely crystalline powder, on dilution with cold water. Cuprous chlo- ride is insoluble in water and also in alcohol; but it dissolves in hydrochloric acid or ammonia. The hydrochloric acid solution absorbs carbon monoxide, forming a product which separates from the concentrated solutions in unstable, shining scales that prob- ably have the composition CO Cu 2 Cl 2 2 H 2 O. With ammonia, cuprous chloride forms the compound Cu 2 Cl 2 2 NH 3 . Thi? ammoniacal solution also has the property of absorbing gases. In many ways, cuprous chloride is analogous to mercurous chloride HgCl, silver chloride AgCl, and thallous chloride T1C1. Cupric chloride CuCl 2 is. made by burning copper in chlorine, also by dissolving the metal in nitre-hydrochloric acid, or by treating cupric oxide, hydroxide, or carbonate, with hydro- chloric acid. The anhydrous salt is brown. The aqueous solutions are bluish green, and from them green, rhombic, pris- matic crystals, CuCl 2 2 H 2 O, may be obtained. These readily dissolve in water or in alcohol. At 500 anhydrous cupric chloride melts, and on further ignition it may be converted into cuprous chloride. With ammonia, cupric chloride forms CuCl 2 -6NH 3 , a dark blue, unstable powder. From aqueous ammoniacal solutions of cupric chloride, dark blue crystals of the composition CuCl 2 4 NH 3 H 2 O may be obtained. At about 150 these compounds are converted into CuCl 2 2 NH 3 , which is a green powder. Cuprous bromide CuBr, like cuprous chloride, is a white, in- soluble powder. It may be melted without decomposition. Cuprous iodide Cul is the only iodide of copper known. It is 466 OUTLINES OF CHEMISTRY formed by heating copper and iodine together, also by adding potassium iodide to solutions of cupric salts : CuS0 4 + 2 KI = K 2 S0 4 + Cul + I. The cuprous iodide so formed is thus precipitated together with free iodine. Cuprous iodide is a white, insoluble, crys- talline powder. It may be melted at red heat. Cuprous fluo- ride CuF is also known. It is a bright red, insoluble powder. Cupric bromide CuBr 2 , like cupric chloride, is readily soluble in water, yielding green solutions. Cupric fluoride CuF 2 is also known, but cupric iodide CuI 2 is not known. Cyanides of Copper. When a solution of a cupric salt is treated with potassium cyanide, cupric cyanide Cu(CN) 2 is formed as a yellow unstable compound which at once decom- poses into cuprous cupric cyanide and cyanogen, thus : CuSO 4 + 2 KCN = Cu(CN) 2 + K 2 SO 4 , and 3 Cu(CN) 2 = Cu(CN) 2 . Cu 2 (CN) 2 + (CN) 2 . This reaction is frequently used in making cyanogen. Cuprous cyanide CuCN may be obtained as a white precipitate by adding potassium cyanide to a solution of copper sulphate saturated with sulphur dioxide. Potassium cyanide forms a soluble double cyanide with cuprous cyanide CuCN KCN. When potassium ferrocyanide is added to a solution of a cupric salt, copper ferrocyanide forms as a dark brownish red, hydrous, amorphous precipitate : 2 CuS0 4 + K 4 Fe(CN) 6 = Cu 2 Fe(CN) 6 + 2 K 2 SO 4 . Copper Salts of Oxy-acids. No cuprous salts of oxy- acids are known; in its salts with such acids, copper is always bivalent. The following are the most important of these salts : Copper sulphate CuSO 4 is obtained by heating copper with concentrated sulphuric acid. On a commercial scale it is formed by heating copper pyrites, or copper plus sulphur, in a current of air, and lixiviating the mass with water and evap- orating the clear solution, from which large, blue, triclinic crystals CuSO 4 5 H 2 O, known as blue vitriol, are obtained (Fig. 75). This is a very common copper salt. On exposure to the air the crystals effloresce. At 100 they lose about four fifths of their water, but it requires about 200 to dehydrate them completely. The anhydrous salt is a gray- COPPER, SILVER, AND GOLD 467 ish white powder, insoluble in alcohol. It is hygroscopic, and is consequently sometimes used in the laboratory as a drying agent. Blue vitriol is soluble in about 3 parts of water. It is used in large quantities in copperplating, as a mordant in dye- ing fabrics, as a source for the preparation of other copper com- pounds, and as a germicide, particularly in spraying plants to protect them from insect pests. Mixed with calcium hydroxide solution, copper sulphate forms the so-called Bordeaux mixture, which is often used in spraying fruit trees and other plants. With alkali sulphates, copper sulphate forms double salts, like K 2 SO 4 CuSO 4 6 H 2 O. These are monoclinic and isomorphous with the analogous magnesium salts. With ammonia, copper sulphate yields deep blue solutions from which dark blue, ortho- rhombic crystals of the composition CuSO 4 4NH 3 H 2 O are precipitated by adding alcohol. On heating this compound carefully, it may be converted into CuSO 4 2 NH 3 , a green pow- der which is called cuprammonium sulphate. It is regarded as ammonium sulphate in which two hydrogen atoms are replaced by a bivalent copper atom, thus : y^Hsv S0 4 < >Cu. N NH Basic copper sulphates, of the formulae CuSO 4 Cu(OH) 2 , CuSO 4 - 2 Cu(OH) 2 , and CuSO 4 3 Cu(OH) 2 , have also been prepared. Copper nitrate Cu(NO 3 ) 2 is formed by dissolving copper, cupric oxide, hydroxide, or carbonate in nitric acid. From the solutions the salt may be obtained in deliquescent crystals, Cu(NO 3 ) 2 6 H 2 O. On heating the nitrate, it yields cupric oxide. Copper Carbonates. Only basic carbonates of copper are known. On adding an alkaline carbonate to a solution of a copper salt, a green basic copper carbonate CuCO 3 - Cu(OH) 2 is formed, thus : 2 CuSO 4 + 2 Na 2 CO 3 + H 2 O = 2 Na 2 SO 4 + CuCO 3 - Cu(OH) 2 f CO 2 . This behavior is then analogous to that of zinc and mag- nesium salts when they are similarly treated with alkaline car- bonates, Malachite, which, as already stated, occurs in nature, 468 OUTLINES OF CHEMISTRY is CuCO 3 - Cu(OH) 2 , and verdigris, which forms in the ah on copper roofs, etc., is also basic copper carbonate. The mineral azurite forms monoclinic crystals of the formula, Cu(OH) 2 .2CuCO 3 . Copper Arsenite. Cu 3 (AsO 3 ) 2 2 H 2 O is a beautiful, green precipitate formed by adding potassium arsenite to a solution of copper sulphate. It is known as Scheele's green and is used as a pigment. Sulphides of Copper. Cuprous sulphide Cu 2 S is formed by heating cupric sulphide or a mixture of copper and sulphur in a current of h}^drogen, also by burning copper in sulphur vapors. It forms black crystals of the regular system. Cupric sulphide is obtained as a black precipitate by passing hydrogen sulphide into a solution of a copper salt. It is practically in- soluble in water and also in dilute acids. Analytical Tests for Copper. In testing copper, the following characteristic reactions are generally used: An excess of ammonia yields deep blue solutions with copper salts. Potassium or sodium hydroxide precipitates blue cupric hydroxide, which turns black on boiling. Hydrogen sulphide produces a black precipitate, insoluble in dilute acids. Potas- sium ferrocyanide produces a reddish brown precipitate of copper ferrocyanide. Alkaline carbonates precipitate basic copper carbonate. Hydrogen sulphide does not precipitate copper sulphide from solutions of potassium cuprous cyanide; whereas, in a similar solution of potassium cadmium cyanide, yellow cadmium sulphide is precipitated, and thus copper may be separated from cadmium. Occurrence, Metallurgy, and Properties of Silver. In nature, silver frequently occurs in the uncombined condition, usually scattered in quartz and other rocks in small amounts, though at times it is found in masses of a hundred pounds or more. Native silver generally is contaminated with other metals, like gold, copper, iron, and lead. It is frequently found in fairly well-developed c^stals of the regular system. The most important ores of silver are argentite or silver glance Ag 2 S, pyrargyrite or ruby silver ore Sb 2 S 3 3 Ag 2 S, proustite or light red silver ore As 2 S 3 3 Ag 2 S, and other complex sulphides of silver, copper, arsenic, and antimony. Horn silver AgCl and other halides of silver are occasionally found. Nearly all COPPER, SILVER, AND GOLD 469 galenite PbS contains some silver, and a considerable amount of silver is obtained from this source. The world's supply of silver comes mainly from the United States, Mexico, western South America, Australia, Germany, and Austria. The extraction of silver from its ores is carried on by several processes; these vary according to the nature of the ore and the locality. The method of obtaining silver from ores rich in lead consists of roasting the ores to get rid of the sulphur, and then heating the residues with carbon and suitable fluxes in a blast furnace. Thus the lead containing the silver in solution settles to the bottom of the furnace and is drawn off. To this molten lead, zinc is added, and the mixture is thoroughly stirred. Zinc is but slightly soluble in lead, but silver is much more greedily taken up by zinc than by lead. Therefore the silver enters the molten zinc, and this solution of silver in zinc floats on top, where it solidifies in the form of a foam, which is skimmed off and then treated with superheated steam at red heat. This oxidizes the zinc and leaves the silver in metallic form. The process of thus using zinc in separating silver from its solutions in molten lead is known as Parke's process. The amalgamation process consists of extracting silver from its ores by dissolving the metal in mercury and then distilling off the latter. In order to accomplish this, the silver in the ores must first be converted into the metallic form. This is done by roasting the ores with sodium chloride, thus forming silver chloride, which is then moistened with water and treated with iron. In this way ferrous chloride is formed, which remains dissolved, and metallic silver is precipitated. The latter is then extracted from the mass by means of mercury. Where fuel is expensive, as in some of the South American countries, the roasting pro- cess is not used, the ores first being finely ground and then mixed with sodium chloride, copper sulphate, water, and mer- cury. This thick pasty mass is mixed by driving mules through it, and shoveling it over occasionally. The sodium chloride and copper sulphate react, forming cupric chloride, which in turn converts the silver in the ores to silver chloride. The latter is then decomposed by mercury, forming mercurous chloride and silver, which is dissolved in the excess of mercury present. This process usually requires a week or more. The silver amalgam formed is washed and strained through canvas 470 OUTLINES OF CHEMISTRY From the solid thus remaining, the silver is obtained by dis tilling off the mercury. The reactions involved are in the main as follows: CuSO 4 + 2 NaCl =Na 2 SO 4 + CuCl 2 . CuCl 2 + Ag 2 S = CuS + 2 AgCl. 2 AgCl + 2 Hg = 2 HgCl + 2 Ag. Still other processes are also sometimes employed. So, when lead containing but little silver cools, the crystals that form are practically pure lead, and when these are removed, the remain- ing liquid is richer in silver. As this liquid cools further, more solid lead forms, after the removal of which the remaining solu- tion is still richer in silver. By repeating this process, known as Pattinson's process, an alloy of lead relatively rich in silver remains behind, from which the silver is obtained by cupella- tion. This process consists of heating the lead-silver alloy in a current of air in a muffle furnace ; the lead is thus oxidized to lead oxide, which melts and runs off or is volatilized, and the silver remains behind. Silver is refined by electrolysis or by the sulphuric acid method. The latter consists of dissolving the silver in boiling hot, con- centrated sulphuric acid and then diluting the solution with water. In this way, the noble metals, like gold, platinum, palladium, etc., remain behind, and the silver is obtained as silver sulphate in solution. From the latter the silver is pre- cipitated by means of iron and then cast into molds. The electrolytic process of refining silver is similar to that of pre- paring pure copper by electrolysis. The impure silver is used as anodes. A solution of silver nitrate acidified with nitric acid ssrves as the electrolyte, and thin silver plates are used as cathodes, on which silver is deposited in the form of shining crystals that are continually rubbed off by a mechanical con- trivance. These crystals of pure silver fall to the bottom of the vat, and after draining off the electrolyte and washing them, they are ready for further use. The electrolytic process of refining silver is much used at present. In the chemical lab- oratory pure silver is often made by heating pure silver chloride with soda, or by treating moist silver chloride with zinc. The United States produces about 4000 tons of silver each year, which is somewhat over half of the world's annual production. COPPER, SILVER, AND GOLD 471 Silver, argentum, is the most abundant of the noble metals. It has been known to man since earliest times, and has been used for utensils and coins for many centuries. It is a pure white metal of high metallic luster. Its specific gravity is 10.6. At 962 it melts, and at 2050 it boils and may be distilled. Silver is very tough, ductile, and malleable, and may be beaten into very thin sheets, which transmit blue light. It is the best conductor of heat and electricity known. Molten silver absorbs about 20 times its volume of oxygen, which is again liberated when the metal cools, thus causing " spitting." Silver is softer than copper, but harder than gold. Copper is often added to silver to secure greater hardness and durability. Sterling silver, which is used for coins, spoons, dishes, and many other useful articles, is an alloy containing 90 per cent silver and 10 per cent copper. It is said to be 900 fine, pure silver being called fine silver, or silver that has a fineness of 1000. Silver also readily forms alloys with zinc, lead, mercury, tin, gold, and other metals. When alloyed with tin and mercury it forms amal- gams which are used for filling teeth. In the air, silver remains unchanged, but ozone oxidizes it to peroxide AgO. Molten nitrates or hydroxides of the alka- lies do not attack silver, hence silver dishes are used in the laboratory in working with caustic alkalies or molten saltpeter. Hydrochloric acid attacks silver but slightly, a protective coating of silver chloride being formed. In hot concentrated sulphuric acid, silver readily dissolves, liberating sulphur dioxide : 2 Ag + 2 H 2 SO 4 = Ag 2 SO 4 + 2 H 2 O + SO 2 . In nitric acid, silver is readily soluble even in the cold : 3 Ag + 4 HNO 3 = 3 AgNO 3 + 2 H 2 O + NO. Nitric acid is the best solvent for silver. With sulphur, silver readily forms silver sulphide Ag 2 S, which is a black powder, insoluble in dilute acids. Silver sulphide forms as a black or dark brown coating on silver spoons when in contact with the contents of eggs, for the latter contain sulphur. With the halogens, silver readily unites even at ordinary temperatures. Oxides of Silver. Silver monoxide, commonly called simply silver oxide, Ag 2 O, is formed as an amorphous, dark brown pre- cipitate when potassium or sodium hydroxide is added to a silver 472 OUTLINES OF CHEMISTRY nitrate solution. Above 250, silver oxide decomposes into the metal and oxygen, a behavior which is similar to that of mer- curic oxide. Silver peroxide AgO forms as a black compound when ozone acts on silver. It is also formed on a platinum anode when silver nitrate solutions are electrolyzed. Silver suboxide Ag 4 O has also been prepared. Silver hydroxide is not known. Halides of Silver. These are formed by the action of halo- gens on silver, also by adding a halide to a solution of a silver salt, or by treating silver oxide with hydrohalogen acids. Silver chloride is formed as a curdy, white precipitate when hydrochloric acid or another soluble chloride is added to a solution of silver nitrate. It is insoluble in water and nitric acid, but it dissolves in concentrated hydrochloric acid, also in ammonia water, alkaline chlorides, sodium thiosulphate, or potassium cyanide. Silver chloride is white, but on exposure to light it turns violet and then brown or black, probably because of the separation of finely divided silver. The salt melts at 151. It is much used in analytical chemistry in the detection and estimation of silver and also of chlorine. Silver chloride readily absorbs ammonia, forming AgCl-3NH 3 . Silver bromide AgBr is obtained as a slightly yellowish pre- cipitate by adding a soluble bromide to a silver nitrate solution. If a soluble iodide is added instead of the bromide, silver iodide Agl, a yellow precipitate, is produced. Both the bromide and the iodide of silver are also practically insoluble in water and dilute acids. The solubility of silver bromide in ammonia is less than that of the chloride; and the iodide is least soluble in ammonia. When exposed to light, the bromide and iodide of silver are partially decomposed, like the chloride. Silver fluo- ride AgF is very soluble in water, and consequently differs greatly from the other halides of silver. It may be obtained by adding hydrofluoric acid to silver oxide. The crystals are deliquescent and correspond to the formula AgF H 2 O or AgF . 2 H 2 0. Use of Silver Halides in Photography. The fact that the halides of silver are decomposed when exposed to light forms the basis of modern photography. The sensitive dry plates or films are covered with a layer of gelatine containing finely divided silver bromide in suspension. When such a plate is COPPER, SILVER, AND GOLD 473 exposed to light in a camera, the silver bromide is acted upon more strongly in some places than in others, according to the intensity of the light reflected from the different parts of the object whose image is focused upon the plate. After being thus momentarily exposed to light, the plate exhibits no visible change. However, on now treating the plate with reducing solutions like ferrous salts, alkaline pyrogallic acid, or hydro- chinone, silver is deposited from the bromide more rapidly and densely on those places where the light has acted more intensely, and less rapidly where the action of the light has been less in- tense. Thus these reducing solutions, or developers, as they are called, develop the picture, which is a so-called negative, for it is dark where the object was light, and light where the object was dark. The developing process must be stopped when a good picture of sharply denned outline has been secured. This is done by rinsing the developing solution from the plate and introducing the latter into a solution of sodium thiosul- phate Na 2 S 2 O 3 , commonly called hyposulphite of soda or "hypo." Thus the developing process is arrested, and the sodium thiosulphate dissolves the remaining unaltered silver bromide from the plate, forming a soluble sodium silver thio- sulphate and sodium bromide, which are finally rinsed off : AgBr + Na 2 S 2 O 3 = NaBr + NaAgS 2 O 3 . The film or the plate now consists simply of silver imbedded in gelatine and is said to be " fixed, " for it may be exposed to the light without suffering further alteration thereby. The sodium thiosulphate solution is termed the "fixing bath." The pictures are printed by placing the negative upon paper covered with a sensitive film of silver chloride, bromide, or iodide, and then exposing this to the light. This yields a positive, for obviously the print is dark where the negative is light and vice versa. The prints must be " fixed " by the same process as the negative. Usually silver bromide papers are employed for printing. They are more readily acted upon by light than papers covered with a silver chloride film. Silver Nitrate AgNO 3 , also known as argentic nitrate and lunar caustic, is the most important silver compound, for from it nearly all the other compounds are directly or indirectly prepared. Silver nitrate is formed by dissolving silver in nitric acid. 474 OUTLINES OF CHEMISTRY It forms transparent, rhombic crystals which melt at about 200, and are extremely soluble in water, although they are an- hydrous and not at all deliquescent. At 0, 100 parts of water dissolve 122 parts of silver nitrate, while at 100 nearly ten times that amount of salt is dissolved. In contact with organic matter, like the skin, cloth, or paper, silver nitrate turns black on exposure to light, owing to separation of metallic silver. Hence the salt is used in indelible inks. Silver nitrate has strong caustic properties, and hence is used in cauterizing sores and removing warts. For this purpose it is commonly cast into sticks, which are either pure silver nitrate or a mixture of the latter with potassium nitrate. It is in the form of such caustic pencils that silver nitrate is termed lunar caustic or lapis infernis. Silver nitrate is frequently used in analytical chem- istry. It is poisonous, and its solutions have a disagreeable metallic taste. With ammonia, silver nitrate forms rhombic crystals of the composition AgNO 3 -2NH 3 . Silver Nitrite AgNO 2 is obtained as a sparingly soluble crystalline precipitate by adding potassium nitrite to a solution of silver nitrate. The salt is used in water analysis. Silver Sulphate Ag 2 SO 4 is commonly made by dissolving silver or silver carbonate in concentrated sulphuric acid. It forms rhombic crystals soluble in about 200 parts of water. In sulphuric acid it dissolves more readily, because of the for- mation of silver bisulphate AgHSO 4 . Silver Carbonate Ag 2 CO 8 is obtained as a yellowish precipi- tate by adding soluble carbonates to solutions of silver salts. It dissolves in water charged with carbon dioxide. On heating silver carbonate, it decomposes : Silver Phosphate Ag 8 PO 4 is a yellow, amorphous precipi- tate produced by adding sodium phosphate to a solution of a silver salt. Silver Sulphide Ag 2 S is obtained as a dark brown precipi- tate, when hydrogen sulphide acts on solutions of silver salts. It is insoluble in dilute acids. Silver Cyanide AgCN is formed by adding potassium cya- nide to a solution of silver nitrate. In excess of potassium COPPER, SILVER, AND GOLD 475 cyanide, silver nitrate is soluble, forming the double cyanide KAg(CN) 2 . The latter is very important, for its solutions are used in silver plating. Silver cyanide is stable in the light, and is used in pharmaceutical practice for preparing dilute solutions of hydrocyanic acid by the following reaction : AgCN + HCl = AgCl + HCN. The silver chloride being insoluble is filtered off. Silver Plating. This is commonly accomplished by elec- trolysis. The objects to be plated are first thoroughly cleaned, and then immersed in a bath consisting of a solution of potas- sium silver cyanide KAg(CN) 2 , in which they serve as cathodes. The anode consists of a thick plate of silver. As the current passes, the objects are coated with a dense, white, well-adhering deposit of silver which can afterwards be pol- ished. Aqueous silver nitrate will not serve as the electrolyte in electroplating, for from such solutions crystalline,. poorly adhering deposits are obtained. Silver mirrors are formed when a clean surface of glass is treated with an ammoniacal silver nitrate solution plus a reducing agent like formic aldehyde, glucose, or liochelle salt. Silver Fulminate AgONC is made like mercuric fulminate. It is even more explosive than the latter. Analytical Tests for Silver. Silver is very readily detected in its compounds. All silver salts of organic acids yield white metallic silver when ignited. Silver is precipitated from solu- tions of silver salts by copper, mercury, zinc, or iron. Solu- tions of silver salts yield insoluble silver chloride on treatment with hydrochloric acid. Silver chloride is soluble in ammonia, which is not the case with mercurous chloride or lead chloride. The last two salts are also insoluble in water, though lead chloride dissolves in boiling water. Silver chromate Ag 2 CrO 4 , a dark red precipitate, forms by adding a solution of either po- tassium chromate or bichromate to a soluble silver salt. Other characteristic precipitates are the carbonate, which dissolves in ammonium carbonate, and the phosphate, which is yellow ; these have already been described. In potassium cyanide or in sodium thiosulphate solutions, silver salts are soluble. The fact that most silver compounds are affected by light is also characteristic. 476 OUTLINES OF CHEMISTRY Occurrence, Metallurgy, and Properties of Gold. Gold is usually found in nature in the uneombined state in quartz veins and alluvial sands, though sometimes it occurs combined with tellurium as calaverite AuTe 2 , petzite (AuAg) 2 Te, and sylvanite AgAuTe 2 . Native gold always contains some sil- ver, frequently also copper, lead, iron, and other metals. Many samples of iron pyrites contain small amounts of gold. The chief localities are South Africa, Alaska, California, Colorado, Dakota, the Urals, and Australia. Alluvial gold-bearing sands are usually washed with water, which carries away the light, loose material, leaving the heavier particles of gold behind. The latter is gathered by treatment with mercury; and from the gold amalgam thus formed, the mercury is removed by distillation from iron retorts. From quartz, gold is separated by pulverizing the material in stamp- ing mills, and then running water bearing the finely divided ore over copper plates amalgamated with mercury. The latter thus catches and dissolves the gold particles. From time to time, the amalgam is removed and distilled to separate the gold from the mercury. The latter is used over and over again. Not all the gold is extracted from the ores by this amalgamation process, the " tailings " generally containing very finely divided particles of the metal. These are recovered by the cyanide process, which consists of treating the ore with a dilute solution of potassium cyanide (from 0.25 to 0.8 per cent), in which both gold and silver dissolve in presence of the oxygen of the air, thus : 4 Au + 8 KCN + 2 H 2 O + O 2 = 4 KOH + 4 K Au(CN) 2 . 4 Ag + 8 KCN + 2 H 2 O + O 2 = 4 KOH + 4 KAg(CN) 2 . These cyanide solutions are then subjected to electrolysis to obtain the gold, or the latter is precipitated by means of zinc, thus : 2 KAu(CN) 2 + Zn = K 2 Zn(CN) 4 + 2 Au. If electrolysis is used, iron anodes and lead cathodes are em- ployed. The latter are finally melted and cupelled, by heating them in a current of air in a muffle furnace. Thus lead is transformed to oxide and runs off, and gold remains. Aurife- rous pyrites ores are roasted, and then subjected to the chlorination process, which consists of treating the finely ground COPPER, SILVER, AND GOLD 477 ore with chlorine, thus forming soluble auric chloride AuCl 3 . From the solution, gold is then precipitated by means of ferrous sulphate, thus : 2 AuCl 3 + 6 FeSO 4 = 2 FeCl 3 + 2 Fe 2 (SO 4 ) 3 + 2 Au. Gold from any of the above processes usually contains silver from which it must be parted. The parting is accomplished by treating the alloy in form of foil or granulated pieces with hot concentrated sulphuric acid which dissolves the silver and leaves the gold behind as a dark powder. This is finally melted with fluxes, like potassium nitrate or borax, to remove any further base metals that may still be present. Gold may also be separated from silver by electrolytic methods. In the process of assaying gold ores, the latter are heated with lead and fluxes like borax or soda. Thus gold and silver finally accumulate, dissolved in lead. This lead alloy, or " button," is then heated in a muffle furnace in a small dish called a cupel, made of bone ash, which absorbs the lead oxide formed, leaving the alloy of gold and silver. This is then treated with hot nitric acid which dissolves the silver and leaves the gold behind. This method of parting is called quartation, for it was formerly believed that in order that the separation be com- plete, the alloy must not contain more than 25 per cent gold. As a matter of fact, the alloy may contain up to 35 per cent gold and still the separation be successful. Pure gold is a rather soft, bright yellow metal having a high luster. It is the most malleable and ductile of all the metals, arid conducts heat and electricity well. It may be beaten in to very thin leaves, which transmit green light. A grain of gold may be beaten into a leaf covering an area of about half a square meter. Gold may be obtained in regular octahedra or dodeka- hedra. Its specific gravity is 19.3. It melts at 1064, and may be volatilized in the electric furnace. Chemically, gold is rather inert. In the air or in oxygen, it remains unchanged even at high temperatures. Gold is attacked by fused caustic alkalies and nitrates of the alkalies, aurates being formed. It is soluble in aqua regia, but not in nitric, hydrochloric, or sulphuric acid. Chlorine attacks gold, but sulphur does not. In potassium cyanide, it readily dissolves, as already stated. On ignition, all compounds of gold are decomposed, leaving a residue of the metal. The atomic weight of gold is 197.2. 478 OUTLINES OF CHEMISTRY Gold Alloys. Gold forms alloys with many metals, like mer- cury, copper, silver, cadmium, tin, and lead. Pure gold is too soft for coinage, jewelry, and ornaments; hence it is alloyed with copper for these purposes. Indeed, of the gold alloys those with copper are by far the most important ones. They are harder, stronger, darker in color, and more readily fusible than pure gold. The gold coins of the United States consist of 1 part copper and 9 parts gold. It is customary to express the purity of gold in carats. Pure gold is called 24-carat gold ; a gold alloy consisting of 8 parts copper and 16 parts gold is 16-carat; 14-carat gold contains 10 parts copper and 14 parts gold. For many articles of jewelry, it is necessary to use at least as low as 14-carat gold to secure sufficient rigidity. Pure gold leaf in form of a wad is readily condensed to a solid piece by hammering, which fact is used by dentists in filling teeth with gold. Gold has been known since earliest times and has always been regarded as an article of great value. Each year the United States produces nearly 160 tans of gold, representing a value of about 196,500,000. This is approximately one fifth of the world's annual output. The yield of the Transvaal mines is now greater than that of America. The production of gold has increased greatly in recent years. Only about 26 tons of gold were produced annually on the average during the first fifty years of the last century. Since 1850 there has been a steady increase in the production of gold. Compounds of Gold. On account of its inertness, gold forms but few compounds. There are aurous compounds in which gold is univalent, and auric compounds in which it is trivalent. When dissolved in aqua regia or when treated with chlorine, auric chloride AuCL is formed. It consists of reddish-brown deliquescent crystals which, when heated to about 180, decompose into chlorine and aurous chloride AtiCl, which is white. On boiling aurous chloride with water, auric chloride and gold are formed : 3AuCl = AuCl 3 + 2Au. When auric chloride is evaporated with excess of hydrochloric acid, yellow prismatic crystals of chlorauric acid HAuCl 4 -4 H 2 () are formed. With sodium chloride, auric chloride forms COPPER, SILVER, AND GOLD 479 sodium chloraurate NaAuCl 4 or NaCl AuCl 3 , which is a salt of chlorauric acid. This salt is used in photography in dilute solu tions as a toning bath, for silver prints immersed in it are col- ored slightly yellowish, due to the fact that silver decomposes the salt and so deposits a very thin layer of gold. The analo- gous bromides and iodides of gold are also known. They are rather unstable. Aurous oxide Au 2 O is a dark violet powder formed by the action of caustic alkalies upon aurous chloride. When heated, it yields gold and oxygen. Auric oxide Au 2 O 3 is a brown powder. Auric hydroxide Au(OH) 3 is formed when caustic alkalies are added to auric chloride solutions. This hydroxide is soluble in excess of caustic alkalies, forming aurates like KAuO 2 : Au(OH) 3 + KOH = K AuO 2 + .2 H 2 O. By treating auric chloride or oxide with ammonia, a precipi- tate is formed which is called fulminating gold. It is very explosive when dry. Its composition is not known with cer- tainty. Aurous cyanide AuCN is a yellow, crystalline powder which is soluble in potassium cyanide, forming potassium aurous cya- nide KAu(CN) 2 . The latter is also obtained by dissolving gold in dilute potassium cyanide solutions in presence of air. When auric compounds are treated with potassium cyanide, potassium auric cyanide KAu(CN) 4 is formed. The solutions of these double cyanides of potassium and gold are used as baths in gold electroplating, which process is similar to silver plating. A gold plate serves as the anode, and the object to be plated is the cathode. Auric sulphide Au 2 S 3 is formed as brownish black precipitate when hydrogen sulphide is passed into cold solutions of auric salts. The precipitate usually contains free sulphur. Auric sulphide is soluble in alkali sulphides, sulpho- salts, like (NH 4 ) 3 AuS 3 , being formed. From hot solutions, aurous sul- phide Au 2 S is similarly precipitated. It has a steel-gray appearance and dissolves in water, from which it may be pre- cipitated by adding hydrochloric acid. Purple of Cassius consists of a finely divided, brownish pur- ple precipitate of gold formed by adding stannous chloride to 480 OUTLINES OF CHEMISTRY solutions of auric chloride. The powder is used in coloring glass and in painting porcelain ware. Analytical Tests for Gold. On ignition with soda on char- coal, gold compounds yield a globule of gold. Ferrous sulphate precipitates gold from solutions of its salts. Other reducing agents, like oxalic acid and sulphur dioxide, also precipitate gold in brown, pulverulent form. Gold is readily displaced from its solutions by many other metals like copper, zinc, and iron. The purple of Cassius reaction already mentioned is also characteristic. REVIEW QUESTIONS 1. Why do copper, silver, and gold occur in nature in the uncombined state? What other metals that you know of occur thus? What is meant by the terms noble metal, base metal ? 2. Mention three important ores of copper and tell where they are found, also how metallic copper is obtained from them. 3. How much copper would be necessary to produce 100 tons of each of the following : gun metal, bell metal, Dutch metal ? 4. What is the difference between brass and bronze? What is Ger- man silver? 5. What use is made of metallic copper? 6. How does metallic copper act toward : (a) nitric acid, (6) hot concentrated sulphuric acid, (c) dilute sulphuric acid? Equations. 7. What is a cupric compound? A cuprous compound? Give five illustrations of each. 8. What is the most common salt of copper? How is it prepared, and for what purposes is it used ? 9. Mention six different ways of showing that a copper nitrate solu- tion contains copper. Equations. 10. Mention four important silver ores and tell where the main silver mines of the world are located. 11. Explain two processes for refining silver. 12. Compare the action of silver, copper, and gold toward nitric acid, also toward sulphuric acid, writing the appropriate equations. 13. What is the most important compound of silver? Describe its properties and uses. 14. What is sterling silver? What use is made of it? How could you demonstrate that its composition is as you have stated? 15. What compounds of silver are used in photography? Upon what fact does their use depend? 16. What compounds of silver are insoluble in water and dilute acids? What silver compounds dissolve in water? 17. Explain the action of the fixing bath in photography. Equation. COPPER, SILVER, AND GOLD 481 18. What property have all photographic developers in common ? 19. How much metallic copper would be required to completely de- compose 20 grams of silver nitrate? Equation. 20. Why is silver plating done from a potassium silver cyanide solution, and not from one of silver nitrate ? 21. In a solution containing the chlorides of gold, copper, and zinc, how demonstrate that these metals are present? Equations. 22. What is the composition of gold coins? How would you prove this? Equations. 23. What metals would precipitate gold from an auric chloride solu- tion? Equations. How would a ferrous chloride solution act on one of auric chloride? Equation. 24. Where are copper, silver, and gold located in the periodic system of the elements ? Explain. CHAPTER XXVI THE METALS OF THE EARTHS OF the metals of the earths, aluminum is by far the most important. It has already been stated that boron is a trivalent element and that its compounds consequently bear some analogy to those of aluminum. Boron, however, exhibits but slightly basic properties, whereas aluminum is a basic element, though its hydroxide is also capable of acting as a weak acid in the formation of alum mates. The other metals of the earths are gallium, indium, thallium, scandium, yttrium, lanthanum, and ytterbium. With the exception of thallium, these are all quite rare. In their com- pounds, these rare-earth metals are trivalent like aluminum. Occurrence, Preparation, and Properties of Aluminum. Alu- minum is very widely distributed. It is the most abundant of the metals and enters into the composition of the earth's crust to the extent of nearly 8 per cent. Aluminum has a great affinity for oxygen, and so is never found in the uncombined state in nature. It occurs as an essential ingredient of practi- cally all the common siliceous rocks. It is found particularly in feldspars, micas, chlorite, granites, slates, and clays. Oxide of aluminum occurs as corundum, also in form of sapphire or ruby. Bauxite is a mineral consisting of the hydrated oxides of alumi- num and iron. Cryolite, which is found in Greenland, is a double fluoride of aluminum and sodium of the formula A1F 8 . 3 NaF. Though widely distributed, aluminum is not found in animals, and only in traces in some plants. Aluminum was formerly prepared by heating sodium alumi- num chloride NaCl A1C1 8 , or cryolite, with metallic sodium. It is now made in large quantities by electrolysis of a solution of aluminum oxide A1 2 O 3 in molten cryolite. The containing vessel (Fig. 150) is made of graphitic carbon, which also serves as cathode. The anode consists of sticks of carbon placed vertically. Thus, as the current passes, oxygen is evolved on 482 THE METALS OF THE EARTHS 483 the carbon anode and passes off. Aluminum is deposited at the bottom of the containing vessel and is tapped off from time to time. The heat generated by the current depositing the metal is sufficient to keep the electrolyte and the aluminum FIG. 160. in a molten state. As the aluminum is deposited, more oxide is placed on top of the electrolyte, as shown in the figure. Large quantities of aluminum are now prepared, especially where water power is available, as at Niagara Falls and other localities. Over 36,000 tons of aluminum are used in the United States each year. Approximately two thirds of this amount is produced in this country. Aluminum is a silver-white, lustrous metal of specific gravity 2.6 to 2.7; that is, it is about one third as heavy as iron. It melts at about 660. It is very ductile and malleable and is a good conductor of heat and electricity. Aluminum is about as hard as silver. The hammered or rolled metal is denser than when cast, which is also in general true of other metals. Sufficiently large pieces of aluminum ring like a bell when struck. At a temperature slightly below its melting point, aluminum becomes brittle and crumbles when shaken or touched. At still lower temperatures, it is again pliable and may be welded and readily worked into desired forms. It is very difficult to solder aluminum, and consequently it cannot very well be used for 'many purposes. Chemically aluminum is relatively inert as compared with the metals of the alkalies, alkaline earths, and magnesium. It oxidizes but very slowly on exposure to the air or oxygen ; for a thin, though resistant, coating of oxide forms on the metal, giving it a slightly bluish 484 OUTLINES OF CHEMISTRY hue. This film protects the aluminum from further oxidation and from attack by many acids. Thin pieces of aluminum, when strongly heated in the air or in oxygen, burn with a bril- liant light, forming the oxide and some nitride. The metal may also be burned in a current of superheated steam, oxide and hydrogen being formed. In caustic alkalies, aluminum dissolves, yielding aluminates and hydrogen. Hydrochloric acid readily acts on aluminum, forming hydrogen and alumi- num chloride, which is soluble. Nitric acid and dilute sul- phuric acid are practically without action on aluminum, which is, however, soluble in hot, concentrated sulphuric acid with concomitant evolution of sulphur dioxide. The atomic weight of aluminum is 27.1; the metal is always trivalent. Uses of Aluminum. On account of the fact that aluminum is a good conductor of electricity, it is used for electric cables and wires, especially at times when copper is relatively high in price. Aluminum is used for cooking utensils, and many other useful articles. In finely divided form, it is used in aluminum paints. In the form of leaf, it is employed in stamp- ing titles on the covers of books, for it does not blacken on exposure to the air as silver does. Aluminum is further em- ployed in removing oxides from iron, and thus denser castings are obtained. Aluminum alloys are also in use. Besides the aluminum-bronze already mentioned, magnalium, consisting of 10 to 25 per cent magnesium and 90 to 75 per cent aluminum, is coming into use. It is lighter and much harder than alu- minum and may be polished to a higher degree. With cadmium, aluminum forms a very tough alloy. Alloys with nickel, zinc, and other metals have also been studied. On account of the great affinity which aluminum has for oxygen, it is used in preparing other metals from their oxides by the Goldschmidt process, which consists of heating together a mixture of powdered aluminum and the oxide to be reduced. Thus the oxides of iron, nickel, chromium, etc., are readily reduced; for example, 3 Fe 3 O 4 + 8 Al =4 A1 2 O 3 + 9 Fe. Fe 2 O 3 + 2 Al = A1 2 O 3 + 2 Fe. 3 NiO + 2 Al = A1 2 O 3 + 3 Ni. Cr 2 ? + 2 Al = A1 2 3 + 2 Cr, THE METALS OF THE EARTHS 485 In practice the mixture of the oxide and aluminum powdei is placed in a crucible and ignited by means of a luse of magne- sium ribbon or a mixture of either aluminum or magnesium and barium peroxide. The reaction, when once started, continues with great evolution of heat. The temperature attained during the reaction when ferric oxide is reduced by aluminum is about 3000, which is quite sufficient to melt both the iron and the aluminum oxide formed. Upon this fact Goldschmidt has based his famous method of welding iron, which consists essentially of butting together the parts to be welded, surrounding the joint with an ignited mixture of oxide of iron, usually Fe 3 O 4 , and powdered aluminum. This mixure is called thermite. The heat developed by the reaction is sufficient to weld the iron securely. Thus car rails can be welded when in place, and many repairs on machinery, ships, etc., can conveniently be made. Thermite welding is consequently frequently used. Many metallic sulphides, when heated with aluminum powder, may similarly be reduced to metal, aluminum sulphide being simultaneously formed. Aluminum Oxide or Alumina A1 2 O 3 is found in nature as corundum, ruby, or sapphire in hexagonal crystals. Corundum is colorless, ruby is red due to the presence of a little chromium, and sapphire is blue because it contains a trace of cobalt. Im- pure corundum is called emery. It contains ferric oxide, and is used as an abrasive material on account of its great hardness, which is but slightly below that of the diamond. Aluminum oxide results in the Goldschmidt reduction process, above de- scribed, also when the hydroxide of aluminum is strongly ignited. The oxide thus obtained is practically not attacked by acids. After fusion with bisulphate of potassium or with caustic alkalies, it may be dissolved in water, for thus aluminum sulphate or aluminates, which are soluble, are formed. Aluminum Hydroxide A1(OH) 3 is found in nature as the mineral hydrargillit. Diaspore is a hydrated oxide A1 2 O 3 H 2 O, and bauxite is a hydrated oxide A1 2 O 3 3 H 2 O, containing ferric oxide. When an alkaline hydroxide is added to a solution of an aluminum salt, aluminum hydroxide is precipitated in the form of a gelatinous mass: A1C1 8 4- 3 KOH = 3 KC1 + A1(OH) 9 . 486 OUTLINES OF CHEMISTRY Aluminum hydroxide is also formed when an alkaline carbonate is added to a solution of an aluminum salt : 2 A1C1 3 + 3 Na 2 CO 3 = 6 NaCl + A1 2 (CO 3 ) 3 , and A1 2 (C0 3 ) 3 + 3 H 2 = 2 A1(OH) 3 + 3 CO 2 . The aluminum carbonate formed is at once completely decom- posed by hydrolysis, as indicated. Aluminum hydroxide may be dehydrated by heat. At about 100, A1 2 O 3 2 H 2 O, and at 300 A1 2 O 3 H 2 O is formed ; finally, on strong ignition all the water is driven off, thus leaving A1 2 O 3 . The fact that a carbonate of aluminum does not exist shows the feebly basic character of aluminum hydroxide. Indeed, the latter is made on a large scale by fusing bauxite with soda, extracting the resulting mass containing sodium alumi- nate with water and passing carbon dioxide into the solution, thus : (1) Na 2 C0 3 + A1 2 3 2 H 2 O = 2 NaAlO 2 + 2 H 2 O + CO 2 . (2) 2 NaA10 2 + 3 H 2 O + CO 2 = Na 2 CO 3 + 2 A1(OH) 8 . As the second equation indicates, sodium carbonate and alu- minum hydroxide are formed simultaneously. The aluminates are therefore rather unstable salts, showing that aluminum hydroxide is but a feeble acid. Just as the zincates are formed by dissolving zinc hydroxide in caustic alkalies, so the aluminates may be obtained by dissolving aluminum hydroxide in caustic alkalies, thus : Zn(OH) 2 + 2 NaOH = Na 2 ZnO 2 + 2 H 2 O. A1(OH) 3 + 3 NaOH = Na 3 AlO 3 + 3 H 2 O. A1(OH) 3 + NaOH = NaAlO 2 + 2 H 2 O. Like the zincates, the aluminates suffer hydrolysis to some extent : NaAlO 2 +2H 2 O^Al(OH) 3 + NaOH. Ammonia, being a weak base, does not react with aluminum hydroxide like sodium or potassium hydroxide. The aluminates are salts formed by replacing the hydrogen atoms of aluminum hydroxide by bases. When aluminum hydroxide loses one molecule of water, A1O OH or HA1O 2 , results which may be regarded as meta-aluminic acid, analo- gous to metaboric acid BO -OH or HBO 2 . Salts of meta- THE; METALS OF THE EARTHS 487 aluminic acid, that is meta-aluminates, are found in nature and are known as spinels. These minerals generally crystallize in regular octahedra. Thus we have spinel Mg(AlO 2 ) 2 , gahnite or zinc spinel Zn(AlO) 2 , iron spinel or pleonast Fe(AlO 2 ) 2 , and chrysoberyl G1(A1O 2 ) 2 . The latter is rhombic. The spinels are all very stable. They have been prepared artificially by heating the constituent oxides together, using boric anhydride as a flux. Insoluble aluminates may also be prepared by pre- cipitation, so, for instance : CaCl 2 + 2 NaAlO 2 = 2 NaCl + Ca(AlO 2 ) 2 . It will be recalled that calcium aluminate is one of the prod- ucts formed when cement sets. This aluminate hardens under water. When aluminum hydroxide is precipitated in a solution con- taining coloring matter, the latter commonly unites with the precipitate, and thus the supernatant liquor remains clear. Suspended matter is also dragged down with the precipitate. This fact is sometimes used in clarifying drinking 1 water. Pig- ments called lakes are made by dissolving dyestuffs in water together with an aluminum or tin salt, and then adding an alkali like soda ; thus precipitates are formed which consist of the coloring matter united with the hydroxide of aluminum or tin. Many dyestuffs do not unite directly with cotton fibers, which, it will be recalled, are practically cellulose. Aluminum hydroxide does unite with cotton fiber, and as dyestuffs in turn unite with aluminum hydroxide, the latter may act as a means of fixing the dye to the fabric. This is done by first dipping the fabric in a solution of an aluminum salt, then treating with steam, whereby aluminum hydroxide is formed which adheres to the fiber, and finally immersing the cloth in the dye, when the latter unites with the aluminum hydroxide, forming an insoluble compound which is thus fixed on the goods. Some- times salts of tin are similarly employed in fastening dyestuffs to fabrics. Substances that will serve this purpose are called mordants, from the Latin word meaning to bite. Besides alum and aluminum acetate, salts of tin, chromium, iron, and anti- mony are frequently employed as mordants. Not all dyestuffs require the use of mordants, for many unite with the fiber directly. 488 OUTLINES OF CHEMISTRY Again, though wool or silk more frequently unite directly with dyestuffs, they must, nevertheless, be treated with suitable mordants in many cases. It must be borne in mind that wool and silk are nitrogenous, organic compounds of animal origin, and are quite different chemically from cotton and linen, which consist mainly of cellulose. Aluminum Chloride A1C1 8 . By dissolving aluminum or its hydroxide in hydrochloric acid and evaporating the solution, deliquescent crystals, A1C1 3 6 H 2 O, form, which, on being heated, give off water and hydrochloric acid, leaving the oxide behind: 2 A1C1 8 6 H 2 = A1 2 3 + 6 HC1 + 3 H 2 O. Anhydrous aluminum chloride is made by heating aluminum in chlorine or by passing that gas over a heated mixture of car- bon and alumina : A1 2 O 3 + 3 C + 3 C1 2 = 2 A1C1 8 + 3 CO. The chloride sublimes and thus may readily be purified. It is very hygroscopic, and fumes in the air because hydrochloric acid is formed by hydrolysis when the salt is in contact with moisture. The chloride is used in the synthesis of organic compounds. With alkali chlorides, it forms double salts like AlClo'- NaCl and A1CL 3 KC1. Aluminum fluoride occurs in o o cryolite, which is an analogous double salt. Aluminum bromide AlBr 3 and aluminum iodide A1I 3 are colorless salts which are quite analogous to the chloride. Aluminum Sulphide A1 2 S 3 is formed by heating aluminum and sulphur together. It is a yellowish mass which is decom- posed by water, forming the hydroxide and hydrogen sulphide. This behavior is similar to that of magnesium sulphide. Be- cause of its complete hydrolysis by water, aluminum sulphide is not precipitated when ammonium sulphide is added to a solution of an aluminum salt. The precipitate consists of aluminum hydroxide. Aluminum Sulphate A1 2 (SO 4 ) 3 is formed by dissolving the hydroxide in sulphuric acid. On a large scale, it is prepared by treating kaolin or bauxite with sulphuric acid. On evapo- rating the solution, monoclinic crystals A1 2 (SO 4 ) 3 18 H 2 O may be obtained. The salt is readily soluble in water. On heating, it loses both water and sulphur trioxide, leaving THE METALS OF THE EARTHS 489 aluminum oxide. Aluminum sulphate is used as a mordant. It is also used in sizing paper. With the exception of blotting paper and filter paper, all papers are sized so that they will be smooth and not absorb ink. Rosin, which consists essentially of a complex organic acid, abietic acid C 44 H 64 O 5 , is dissolved in caustic soda, thus forming sodium abietate, or resin soap, which is added to the paper pulp, and this mixture is then treated with aluminum sulphate. Thus, sodium sulphate and insoluble aluminum resinate are formed. The latter acts as a binder, and gives the paper a smooth surface ; for under the hot rollers the resinate is melted and pressed upon the fibers. Alums. Aluminum sulphate forms double salts with ammo- nium sulphate and the sulphates of the alkali metals. These double sulphates are readily obtained by adding a solution of the alkali sulphate to one of aluminum sulphate and evapo- rating. Crystals are thus obtained which are regular octa- hedra. The compounds all correspond to the general formula M 2 SO 4 - A1 2 (SO 4 ) 3 - 24 H 2 O, or MA1(SO 4 ) 2 - 12 H 2 O, where M is either NH 4 , K, Na, Cs, Rb, Ag, or Tl. These double salts are less soluble than aluminum sulphate. They are called alums. Potassium alum K 2 SO 4 A1 2 (SO 4 ) 3 24 H 2 O is common alum. It is prepared on a large scale by calcining the mineral alunite K 2 SO 4 - A1 2 (SO 4 ) 3 - 4 A1(OH) 3 at about 500 and ex- tracting the mass with water, after having exposed the material to the action of the air and moisture for months. Alum is also made from clays, cryolite, and bauxite. Alunite occurs in large quantities in Hungary and in Italy near Rome. Alum dissolves in about 8 parts of water. Its solutions are astringent and have an acid reaction. Solutions of aluminum acetate (C 2 H 3 O 2 ) 3 A1 act similarly. When alum is heated, it melts, loses water, and finally also some sulphur trioxide, thus leaving a somewhat basic aluminum potassium sulphate behind. This is commonly called burnt alum. On further heating, it is converted into potassium sulphate and alumina. Ammonium alum (NH 4 ) 2 SO 4 A1 2 (SO 4 ) 3 24 H 2 O is also made on a large scale. It is cheaper than potassium alum. Sodium alum Na 2 SO 4 A1 2 (SO 4 ) 3 24 H 2 O is more soluble than either potas- sium or ammonium alum; moreover, it does not crystallize readily and is consequently not made commercially. Alum is used as a mordant, also as an astringent in medicine, especially 490 OUTLINES OF CHEMISTRY as a mouth wash. The acetate is to be preferred for the latter purpose. In the alums, the aluminum may be replaced by other trivalent elements like iron, chromium, manganese, gallium, and indium. The compounds so obtained are isomorphous with the alums. They are consequently also called alums, though they contain no aluminum. So we have ferric ammo- nium, alum (NH 4 ) 2 SO 4 . Fe 2 (SO 4 ) 3 24 H 2 O, potassium chrome alum K 2 SO 4 Cr 2 (SO 4 ) 3 24H 2 O, sodium manganese alum Na 2 SO 4 Mn 2 (SO 4 ) 3 24 H 2 O, etc. In general, the formula of an alum is M 2 SO 4 M 2 (SO 4 ) 3 - 24H 2 O, or M . M(SO 4 ) 2 . 12 H 2 O, where M is a univalent metal or radical, and M is a trivalent metal. Aluminum Silicates are found in enormous quantities in nature. They are also very widely distributed. Thus potash feldspar, orthoclase, K 2 O A1 2 O 3 6 SiO 2 , or KAlSi 8 O 8 , soda feldspar, albite, Na 2 O A1 2 O 3 6 SiO 2 , or NaAlSi 3 O 8 , lime feldspar, anor- thite, CaO A1 2 O 8 2 SiO 2 , or CaAl 2 Si 2 O 8 , occur in granitic rocks, together with quartz and micas. The latter are also silicates of aluminum, containing potassium, magnesium, cal- cium, and sometimes iron. The mineral disthen is a pure aluminum silicate Al 2 SiO 5 , which occurs in rhombic crystals ; it is also found in triclinic forms as andalusite. These two minerals are rather rare. By the action of water and carbon dioxide of the air upon feld spars, the latter lose their alkali content, which is dissolved and enters the soil as soluble silicates or carbonates, thus supplying potash and lime needed for plant growth. A hydrous aluminum silicate H 2 Al 2 (SiO 4 ) 2 2 H 2 O, called kaolin, remains as a white clay. The reactions of the weathering process of typical feld- spars are in the main as follows : 2 K AlSi 3 8 + 2 H 2 = K 2 Si 4 9 + H 2 Al 2 (SiO 4 ) 2 - H 2 O. CaAl 2 Si 2 O 8 +2 CO a + 3 H 2 O = Ca(HCO 8 ) 2 + Il 2 Al 2 (SiO 4 ) a -H a O. Ordinary clays contain ferric hydroxide, sand, and various sili- cates besides kaolin. Frequently calcium carbonate is also present. Clay containing a large amount of calcium carbonate is called marl. When mixed with water, clay forms a plastic mass that can readily be molded into various forms as desired. On drying THE METALS OF THE EARTHS 491 and heating this material to higher temperatures (i.e. "burn- ing " or " firing " it), it becomes hard, dense, and resistant with- out actually melting. These facts form the basis of making bricks, earthenware, and porcelain from clay. The rather im- pure clays are used for the manufacture of bricks, which get their color from 'the iron oxides contained in the clay. Red pottery is made from similar material. Light-colored stone- ware and pottery requires a clay relatively free from iron; while fine white porcelain necessitates kaolin that contains no iron. Earthenware is frequently porous. Sometimes this porous material is covered with a glaze ; and again, as porce- lain, the material may be made perfectly vitreous throughout. Porous earthenware is made by simply shaping clay containing but little fusible material and firing it, as in making bricks, flowerpots, etc. If such articles are to be glazed, they are covered with a layer of fusible silicates, which is fired on. In the case of the cheaper grades of stoneware, glazed bricks, etc., this is accomplished by simply throwing salt into the kiln. At the high temperature that obtains, the salt is volatilized and decomposed as it comes in contact with the earthenware; hydrochloric acid and a readily fusible sodium aluminum sili- cate being formed. The latter covers the surface of the arti- cles as a glaze. In making porcelain, the purest kaolin is mixed with finely pulverized feldspar and quartz. Of this mixture the dishes are shaped, and when they are finally fired in the kiln, the feldspar melts, fills the pores, and thus produces the translucent material known as porcelain. In actual practice, the articles are fired twice. After the first heating, they are dipped in water containing in suspension very finely ground material intended for the glaze. This material commonly con- sists of kaolin, together with enough finely ground feldspar to make a mixture that will fuse at a somewhat lower temperature than the ware of which the articles are composed. This fine material is taken up by the porous porcelain, which is then dried and fired again. As the feldspar melts, a smooth glaze is produced, and a mass that is vitreous throughout results. Ultramarine is an important blue pigment which is manu- factured in large quantities by heating together soda, sulphur, and clay, or, more frequently, sodium sulphate, carbon, and clay. The product, which is a silicate of sodium and aluminum 492 OUTLINES OF CHEMISTRY combined with sodium sulphides, is at first green and is called ultramarine green. This is also used as a pigment. On heat- ing ultramarine green in a current of air, it turns blue, forming ultramarine blue. Ultramarine violet results when the blue is heated in a current of hydrochloric acid gas to about 175, and ultramarine red is formed by the same process at about 145. Only the green and blue ultramarines have much value in practice. The constitution of ultramarine has caused much discussion, and has not yet been finally settled. It is probable that ultramarine blue is (Na 2 Al 2 Si 2 O 8 ) 2 . Na 2 S 2 . It is found in nature as lapis lazuli, from which the pigment was first prepared. Ultramarine is stable towards light, and the action of the air and water ; but acids readily destroy the color, liberating sul- phur and hydrogen sulphide. Ultramarine blue is much used in paints and laundry blue. It is also employed in removing the yellow tinge of sugar, paper pulp, linen and cotton fabrics, and in making cotton prints and wall paper. Analytical Tests for Aluminum. Aluminum is precipitated as the gelatinous hydroxide from solutions by either potassium, sodium, or ammonium hydroxides, carbonates, or sulphides. Aluminum hydroxide is soluble in potassium or sodium hydrox- ides, but -only slightly soluble in ammonium hydroxide. By means of hydrogen sulphide or carbon dioxide, solutions of aluminates are decomposed, yielding a precipitate of aluminum hydroxide. When alumina is moistened with a solution of cobalt nitrate and then strongly ignited, a blue mass Co(AlO 2 ) 2 results. This blue is known as Thenard's blue or cobalt ultramarine. Gallium. This is a rare metal whose atomic weight is 69.9. It was discovered in 1875 by Lecoq de Boisbaudran, who detected its presence in a sample of zinc blende from Pierrefitte in the Pyrenees, by means of the spectroscope. In 1869 Men- deleeff foretold the existence and described the properties of this element, which he called eka-aluminum. He was able to do this from the periodic system of the elements, in which the space now occupied by gallium was then vacant. Gallium receives its name from Gaul, the native land of its discoverer. Gallium is a hard, white metal which is stable in the air. It melts at 30, and its specific gravity is 5.9. In its compounds, it is trivalent and analogous to aluminum, though a chloride of THE METALS OF THE EARTHS 493 the formula GaCl 2 is also known. The spark spectrum of gal Hum exhibits two characteristic lines in the violet, through which the element was discovered. Indium, whose atomic weight is 114.8, is also a rare metal. It was discovered by Reich and Richter in 1863 by means of the spectroscope, in a sample of zinc blende from Freiberg. The metal receives its name from the bright indigo-blue line that characterizes its spectrum. Indium is a silver-white metal, softer than lead and very malleable. Its specific gravity is 7.4, and its melting point is 176. Like gallium, it is stable in the air. In its compounds, it is commonly trivalent like aluminum, though indium dichloride InCl 2 and monochloride InCl are also known. Thallium and its Compounds. Thallium is a metal whose physical properties are similar to those of lead. It was dis- covered in 1861 by means of the spectroscope by Sir William Crookes, in the mud at the bottom of the lead chambers of the sulphuric acid factory at Tilkerode in the Harz. Thallium compounds yield a very characteristic green line in the spec- trum, whence the name thallium, meaning a green branch. Crookesite, a selenide of copper, thallium, and silver contains about 17 per cent thallium. Many native sulphides of iron and copper contain small amounts of thallium, whence its appear- ance in the flues and lead chambers of sulphuric acid plants. In 1862 Lamy demonstrated that thallium is a metal. It oxi- dizes readily on exposure to air, and hence is kept under petro- leum oil or glycerine. The atomic weight of thallium is 204.0, and its valence is either 1 or 3. In thallous compounds, the metal is univalent and therefore analogous to the alkali metals. Thus we have T1OH, T1 2 O, T1F, T1C1, TIBr, Til, T1C1O 3 , T1NO 3 , T1 2 CO 3 , T1 2 SO 4 , T1 2 S, Tl 2 PtCl 6 , etc. Thallous chloride, bromide, and iodide are nearly insoluble in water, and so in this respect these com- pounds are similar to the corresponding ones of silver. Thai- lie chloride T1C1 3 is very soluble in water. In thallic compounds, thallium is trivalent. Thus we have T10-OH, Tl a 8 , T1C1 8 , T1(N0 8 ) 8 , T1 2 (SO 4 ) 3 , T1 2 S 3 , etc. Thallium readily dissolves in nitric or sulphuric acid, whereas, on account of the fact that thallous chloride is difficultly soluble, hydrochloric acid attacks the metal but slightly, it 494 OUTLINES OF CHEMISTRY being soon covered with a protective coating of the chloride. Thallium salts are readily recognized by means of the spectro- scope. From neutral or faintly acid solutions of thallous salts, hydrogen sulphide precipitates black thallous sulphide, which is insoluble in water and acetic acid, but it readily dissolves in sulphuric or nitric acid. The Rare-Earth Elements. These are a series of metals that form compounds which are in general analogous to those of aluminum. The rare-earth metals are, as the name implies, of very rare occurrence. They are found in the complex and very rare minerals monazite, gadolinite, euxenite, samarskite, orthite, cerite, yttrotantalite, hjelmite, and several others which occur mainly in the Scandinavian peninsula, Finland, Green- land, France, Bavaria, and the United States. In general, the elements are trivalent, like aluminum. Chemically the rare-earth metals deport themselves very much alike, which makes it extremely difficult to separate them. In general, their oxalates are insoluble, and their sul- phates are soluble and readily form double salts with the sul- phates of the alkalies. As the nitrates are decomposed by heat at different temperatures, this fact is used in separating the rare earths from one another. Fractional crystallization and fractional precipitation are also used to effect separations ; but these processes are laborious, and generally yield products that are, after' all, not quite pure. Scandium (Sc 44.1) was discovered in 1879 by Nilson and Cleve. Mendeleeff predicted the existence of this element in 1869. Compare gallium and germanium. He called it ekaboron, and described its properties. The element occurs in euxenite and gadolinite Volatilized in the electric arc, scan- dium chloride yields a bright characteristic spectrum of many lines. The oxide Sc 2 O 3 is a white, earthy powder. Yttrium (Y 89.0) occurs in the silicate gadolinite found at Ytterby, whence its name. Yttrium was discovered in 1843 by Mosander. Its oxide Y 2 O 3 is a white, earthy powder of very high melting point. The chloride YCJ 3 yields a bright spectrum containing many lines. Lanthanum (La 139.0) was found in cerite in 1839 by Mosander. The name lanthanum comes from the Greek, mean- ing hidden. Like yttrium, lanthanum may be prepared by THE METALS OF THE EARTHS 495 electrolysis of its molten chloride. Lanthanum is a white, malleable metal not unlike iron in appearance. When heated, it burns to oxide La 2 O 3 , a white powder which readily absorbs water, forming La(OH) 3 . Other compounds are La 2 (CO 3 ) 3 3H 2 0, La(N0 8 ) 8 6 H 2 0, La 2 (SO 4 ) 3 - 9 H 2 O, LaCl 3 . The latter is deliquescent and readily forms 2 LaCl 8 15 H 2 O ; it shows a characteristic spectrum of many lines. Ytterbium (neoytterbium) (Yb 172.0) occurs with scandium and yttrium in euxenite and gadolinite. Its compounds are in general like those of yttrium. In 1907 Urbain in Paris showed that the old ytterbium consists of two elements, neoyt- terbium (Yb 172.0) and lutecium (Lu 174.0). Auer von Welsbach, in Vienna, published the same discovery almost simul- taneously. He named the two new elements aldebaranium and cassiopeium, and found their atomic weights to be 172.90 and 174.23, respectively. As Urbain's work appeared first, his nomenclature will probably be adopted. Cerium (Ce 140.25) resembles lanthanum very much. It was discovered by Klaproth and also by Berzelius and Hisinger in 1803. The element is named from the planet Ceres, which had just been discovered. The mineral cerite contains about 60 per cent cerium. Cerium forms compounds in which the element is trivalent, like Ce 2 O 3 , Ce 2 (SO 4 ) 3 , and CeCl 3 , but it also forms compounds in which it is quadrivalent, like CeO 2 , Ce 2 (SO 4 ) 2 4 H 2 O, Ce(OH) 4 , CeF 4 . H 2 O. The latter class of compounds indicates that cerium is closely related to other quadrivalent elements, and consequently probably belongs in the same group as silicon, titanium, zirconium, and thorium. Cerium may be prepared by electrolysis of the molten chloride. It is a steel-gray, very malleable metal of specific gravity 7.0. It readily burns in the air, forming a yellow powder, CeO 2 . Ceric hydroxide Ce(OH) 4 is a red precipitate. In general, cerous compounds are colorless, while the eerie compounds are yellow, brown, or red. Cerium nitrate is now prepared from monazite sand found in North Carolina in considerable quan- tities. It is used together with thorium nitrate in making the mantles for Welsbach gas lights. The process of making these mantles consists essentially of knitting the mantle of cotton thread, and then dipping this into solutions of the nitrates of cerium and thorium. The mantle is then dried and calcined; 496 OUTLINES OF CHEMISTRY thus the thread is destroyed, and the nitrates are converted to oxides which adhere together. Such mantles are, of course, always frail, arid consequently must be handled with care. As already mentioned, experience has shown that mantles consist- ing of 1 per cent cerium oxide CeO 2 and 99 per cent thorium oxide ThO 2 give the best efficiency. The light emitted from such mantles, when heated to incandescence by a Bunsen burner, is a brilliant white. Mantles containing different pro- portions of ceria and thoria are less efficient, and those made of other earths give a light of poorer quality and also lower inten- sity. The Nernst lamp consists of a filament of earths heated to incandescence by passing an electric current through it. At ordinary temperatures, this filament is practically a non-con- ductor ; but on being heated, it conducts electricity and gives a bright light. Praseodymium (Pr 140.6) and neodymium (Nd 144.3) were for half a century regarded as but one element, didymium. But in 1885 Auer von Welsbach separated it into praseodym- ium, which forms leek-green salts, and neodymium, which forms rose-violet salts. These elements occur with cerium arid lan- thanum, which they resemble. The absorption spectra of solu- tions of salts of the didymiums are especially characteristic. Samarium (Sa 150.4) was discovered in Samarskite, a mineral from North Carolina, in 1878 by Delafontaine and also by Lecoq de Boisbaudran. Typical compounds are Sm 2 O 3 , SmCl 3 -6H 2 O, Sm 2 (SO 4 ) 3 8 H 2 O, Sm(NO 3 ) 2 - 6 H 2 O. These compounds are similar to those of lanthanum. Erbium (Er 167.7), terbium (Tb 159.2), thulium (Tm- 168.5), and dysprosium (Dy 162.5) are found associated with yttrium. Terbium also occurs in samarskite and in small quan- tities in gadolinite, which mineral also contains gadolinium (Gd 157.3). Europium (Eu 152.0) is another element about which but little is known. The position of a goodly number of these elements in the periodic system is still unsettled. It is also quite probable that some of the rare-earth elements, especially those which have been studied but little, will be separated further. THE METALS OF THE EARTHS 497 REVIEW QUESTIONS 1. Why is aluminum called an earth metal ? Where is it placed in the periodic system ? Why ? What other elements are classed with aluminum ? 2. What properties of aluminum make it valuable in the arts? 3. Discuss the occurrence of aluminum in the earth's crust, also in plants and animals. 4. What is the chief ore from which metallic aluminum is prepared? How is this accomplished ? 5. Describe how aluminum acts when treated with each of the fol- lowing reagents: caustic alkalies, nitric acid, hydrochloric acid, sul- phuric acid. Equations. 6. What is magnalium and what use is made of it ? 7. What is thermite? Describe its uses and write the equation expressing the change that takes place when it is used. 8. Compare the action of aluminum and zinc on caustic soda, also on hydrochloric acid, writing the equations. 9. What is the chemical nature of : ruby, corundum, spinel, emery. 10. Write an equation indicating how aluminum hydroxide may be prepared. How does this compound act : (a) when used in clarifying drinking water, (6) when employed in making pigments called lakes? 11. What is a mordant ? Give an illustration. 12. What use is made of aluminum sulphate in the paper industry? 13. What is an alum? Give the names and formulas of six different alums. What have they in common? 14. How does sodium alum act when treated with baking soda? Write the equation. What use is made of this fact ? ; 15. Give the names and formulas of three different feldspars and write the equations showing how in the process of weathering they pass over into kaolin. What is the latter used for? 16. What is the difference between clay and kaolin? Between clay and marl? Between clay and sand? 17. Tell how earthenware and porcelain are made. What other silicate industries are there? 18. How is ultramarine blue prepared? 19. By what tests could you show that alum contains aluminum? Equations. 20. Why is thallium so called? Who discovered this element? Write the formulas of thallous and thallic chloride. Describe the physi- cal properties of thallium. 21. What is meant by the term rare earths? Are any of the rare earths of practical importance? If so, state their use. 22. How much copper will be precipitated from a copper sulphate solution by 32 grams of metallic aluminum ? 23. From 100 Ib. of potassium alum how much aluminium oxide can be prepared ? 24. How demonstrate that clay contains aluminum ? CHAPTER XXVII LEAD AND TIN GERMANIUM, lead, and tin are three metallic elements that exhibit a valence of two or four in their compounds, which are consequently analogous to those of carbon and silicon. The hydroxides of these metals increase in basicity with rise in atomic weight. These hydroxides, however, also have weakly acidic properties, for towards strong bases they may act as acids. Germanium forms a compound with hydrogen, GeH 4 , but tin and lead do not. Germanium (Ge 72.5) is a very rare metal whose existence was foretold by Mendeleeff in, 1871. He called the element ekasilicon and described its properties and fixed its atomic weight at about 73. In 1886 Clemens Winkler discovered germanium in the silver ore argyrodite GeS 2 3 Ag 2 S, which occurs near Freiberg. The metal is also found in euxenite, samarskite, and confieldite. Germanium has properties which are practically the same as those predicted by Mendeleeff for ekasilicon. Germanium is a grayish white, brittle metal of specific gravity 5.47. It crystallizes in octahedra, and melts at about 900. From the formulae of the following compounds, the analogy of germanium to carbon and silicon is evident : GeH 4 , GeF 4 , GeCl 4 , GeHCl 3 , GeO 2 , GeOCl 2 , GeS 2 , K 2 GeF 6 . Besides these, Ge(OH) 2 and GeS are also known. Occurrence, Metallurgy, and Properties of Tin. Tin occurs mainly in cassiterite, or tin stone, SnQ 2 , which crystallizes in the tetragonal system (Fig. 59), and is usually colored brown or black by oxides of iron and manganese. The metal has been known for a long time, for though it does not occur in the free state, its oxide is readily reduced by means of carbon. Tin was alloyed with copper to make bronze, even in ancient times. The Latin name of tin is stannum, whence the symbol Sn. It was also called plumbum candidum to distinguish it from lead, plumbum nigrum. Tin was obtained from the 498 LEAD AND TIN 499 mines at Cornwall in the days of the Roman Empire. Tin ores also occur in Saxony, Peru, Australia, Alaska, and the islands of Billiton and Banca east of Sumatra. To extract tin from its ores, the latter are first roasted to expel any sulphur or arsenic present. The material is then treated with crude hydrochloric acid to remove iron, copper, etc., and lixiviated with water, after which the finely divided ore is mixed with carbon and heated in a furnace, thus : The molten tin collects at the bottom of the furnace and is drawn off and cast into bars. It is purified by remelting it and collecting that portion which fuses at the lowest tempera- ture, for this is freest from other metals. Banca tin is the purest; though German and English tin, block tin, also often is 98 to 99.9 per cent pure. The world's annual output of tin is about 130,000 tons, about two thirds of which comes from the East Indies. Tin is a silver- white, lustrous metal of specific gravity 7.3. It melts at 232, and boils at about 1600. Tin is crystalline in character, and its bars, when bent, give a low, peculiar, crac- kling noise known as tin cry, which is due to the friction of the crystalline particles moving over one another. Tin is very ductile, malleable, and so soft that it may readily be cut with a knife. At 100, the metal is still malleable, but this property decreases on raising the temperature further. At about 200, tin is brittle and may be powdered. On cooling, molten tin always congeals in crystalline form. At low temperatures, tin gradually changes to a gray, brittle variety. This change takes place most rapidly at 48, though it proceeds appre- ciably even at 15. It is the cause of the tin pest, which consists of the disintegration of tin organ pipes and tin roofs and gutters in Russia, where winters are very cold. Above 20, the ordinary malleable tin is the stable form, while below that temperature the gray, brittle variety is the stable modification. In the air tin remains practically unchanged. When strongly heated, it burns with a white flame to SnO 2 . In hot hydrochloric acid tin dissolves : Sn+2HCl=SnC 500 OUTLINES OF CHEMISTRY Hot concentrated sulphuric acid acts on tin, forming sulphui dioxide and stannous sulphate, SnSO 4 : Sn + 2 H 2 SO 4 = SnSO 4 + 2 H 2 O + SO 2 . Cold, dilute, nitric acid acts on tin, forming stannous nitrate, thus: 4 Sn + 10 HN0 3 = 4 Sn(NO 3 ) 2 + 3 H 2 O + NH 4 NO 3 . Concentrated nitric acid converts tin into metastannic acid (which see). Uses of Tin. Tin is used to a very large extent in making tin plate or tinned iron. This process consists of dipping thor- oughly cleaned sheet iron in molten tin. The ordinary tin cans and other tin utensils are made of iron covered with a layer of tin by this dipping process. Copper is also often tinned in the same way. Solder usually consists of 1 part tin and 1 part lead, but these proportions are often varied. An alloy of about 90 per cent tin, 8 per cent antimony, and 2 per cent copper is called Britannia metal. Pewter consists of 75 per cent tin and 25 per cent lead. Bronzes contain tin, copper, and sometimes also zinc, as stated under copper. Alloys of silver and tin are employed to make amalgams for filling teeth. Tin is recov- ered from old tin cans, etc., either by melting off the coating or more frequently by electrolysis. In this process, the cans placed in a wire basket are the anode, caustic alkali solution serves as the electrolyte, and an iron plate is used as the cathode. Much tin is saved in this way. Chlorides of Tin. Stannous chloride SnCl 2 is formed by dis- solving the metal in hydrochloric acid. From the solution, monoclinic crystals SnCl 2 -2H 2 O are obtained, which are very soluble in water. This solution is a strong reducing agent. For instance, it readily reduces mercuric chloride to calomel and even to metallic mercury: 2 HgCl 2 + SnCl a = 2 HgCl + SnCl 4 . HgCl 2 + SnCl 2 = Hg + SnCl 4 . Stannous chloride is hydrolyzed in solution. By treatment with much water a basic chloride of the composition Sn(OH)CJ is precipitated : SnCl 2 + H 2 O = Sn(OH)Cl + HC1. LEAD AND TIN 501 Stannous chloride is used as a reducing agent in the laboratory and as a mordant in dyeing fabrics. Stannic chloride SnCl 4 is made by the action of chlorine on tin or stannous chloride, or by treating stannic oxide or hydroxide with hydrochloric acid. It is a colorless, fuming liquid of specific gravity 2.28. It boils at 114, and congeals at 33. Tin tetrachloride is also known as spiritus fumans Libavii. With water, the chloride forms the hydrates SnCl 4 .3H 2 O, butter of tin, SnCl 4 .5H 2 O, and SnCl 4 - 8 H 2 O. These are soluble in water. Stannic chloride readily forms double salts with other chlorides, like SnCl 4 2 HC1 6 H 2 O, or H 2 SiiCl 6 .6H 2 0; SnCl 4 -2 KC1, or K 2 SnCl 6 ; SnCl 4 . 2NH 4 C1, or (NH 4 ) 2 SnCl 6 . The latter is called pink salt. It is used as a mordant, as is also the hydrate SnCl 4 5 H 2 O. Stannic chloride, on being boiled with water, yields a precipitate of stannic acid, thus : SnCl 4 + 3 H 2 O = 4 HC1 + H 2 SnO 3 . In hydrocarbons, carbon disulphide, and many other organic and inorganic liquids SnCl 4 is soluble in all proportions. Compounds like SnCl 4 -PCl 6 , SnCl 4 -2SCl 4 , SnCl 4 -2NOCl, SnCl 4 POC1 3 are also known. Indeed, stannic chloride enters into a large number of compounds. Tin tetrabromide SnBr 4 and tin tetraiodide SnI 4 are also known. Oxides of Tin. Stannous oxide SnO is formed as a black powder by heating stannous hydroxide out of contact with oxygen. When ignited in the air, it burns, forming stannic oxide SnO 2 . Stannous hydroxide is formed by adding sodium carbonate to a stannous chloride solution : SnCl a + Na 2 CO 3 + H 2 O = 2 NaCl + Sn(OH) 2 + CO 2 . In potassium or sodium hydroxide, stannous hydroxide is sol- uble, but not in ammonium hydroxide, thus : Sn(OH) 2 + 2 KOH = K 2 SnO 2 + 2 H 2 O. On boiling the solution, the potassium or sodium stannite is converted to stannate, with concomitant precipitation of tin, thus : 2 KSnO -}- HO = KSnO + 2 KOH + Sn. . 502 OUTLINES OF CHEMISTRY Stannic oxide SnO 2 , or stannic acid anhydride, is obtained bj burning tin in oxygen or in the air, or by igniting stannic acid. Stannic oxide is found in nature in quadratic crystals as cassit- erite, the principal ore of tin. Stannic oxide as artificially pre- pared is a white or slightly yellowish powder, which is insoluble in water, acids, or alkalies especially after strong ignition. Dur- ing ignition, the oxide turns darker in color, but it again changes to white on cooling. When fused with caustic alkalies, stannic oxide forms stannates, which are soluble, thus : Sn0 2 + 2 NaOH = Na 2 SnO 3 + H 9 O. Stannic hydroxide is formed as a gelatinous precipitate by boiling tin tetrachloride with water, or by adding ammonia to the solution. Stannic hydroxide or orthostannic acid Sn(OH) 4 easily splits off water, and forms SnO (OH) 2 or H 2 SnO 8 , which is called stannic acid. This also gradually loses water so that neither Sn(OH) 4 nor SnO (OH) 2 have really been definitely isolated. It will be observed that stannic acid H 2 SnO 3 is anal- ogous to H 2 CO 3 and H 2 SiO 3 . Stannic acid readily dissolves in concentrated sulphuric, nitric, or hydrochloric acid, and also in dilute solutions of caustic alkalies. With the latter it forms stannates, which dissolve. From these soluble stannates, the stannates of other metals may be prepared by precipitation ; for as in the case of the silicates and carbonates, only the stan- nates of the alkalies are soluble. Sodium stannate Na 2 SnO 3 is made by fusing cassiterite with caustic soda or by fusing tin with sodium carbonate and sodium nitrate. From its aqueous solutions, sodium stannate separates in monoclinic crystals Na 2 SnO 3 3 H 2 O. It is used as a mordant in calico printing, being termed preparing salt. Concentrated nitric acid converts tin into a white, insoluble powder known as metastannic acid, which is a hydrated oxide of tin, whose composition, like that of stannic acid, varies slightly according to temperature and other conditions of preparation, being probably H 2 SnO 4 or H 2 SnO 3 , thus : Sn + 4 HNO 3 = H 2 SnO 3 + H 2 O + 4 NO 2 . But metastannic acid is different from stannic acid, for it is insoluble in acids, and with alkalies it forms salts like K 2 Sn 5 O n 4 H 2 O and Na 2 Sn 5 O n 4 H 2 O, metastannates, which LEAD AND TIN 503 would indicate that the acid is dibasic and probably analogous to polysilicic acids. On fusion with caustic alkalies, metastannic acid yields the stannates which are identical with those of stannic acid. Sulphides of Tin. Stannous sulphide SnS is formed as a dark brown precipitate when hydrogen sulphide is conducted into a solution of a stannous salt. This precipitate is insoluble in dilute acids or solutions of the monosulphides of the alkalies ; but when boiled with the latter and sulphur, or when treated with alkaline polysulphides, soluble sulpho-stannates are formed, thus : SnS + S + K 2 S = K 2 SnS 8 . SnS + (NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 8 . These sulpho-salts are decomposed by hydrochloric acid, yield- ing precipitates of stannic sulphide SnS 2 , thus : K 2 SnS 3 + 2 HC1 = 2 KC1 + SnS 2 + H 2 S." (NH 4 \SnS 8 + 2 HC1 = 2 NH 4 C1 + SnS 2 + H 2 S. Stannic sulphide is a yellow, amorphous powder, which is also obtained by passing hydrogen sulphide into solutions of stannic salts. It is insoluble in dilute acids. On heating, it decom- poses into stannous sulphide and sulphur. In alkaline sulphides it dissolves, yielding sulphostannates. Concentrated hydro- chloric acid dissolves stannic sulphide, and concentrated nitric acid converts it into metastannic acid. By heating tin and sulphur together, stannous sulphide is formed. Stannic sulphide cannot be thus obtained, for it de- composes into sulphur and stannous sulphide at the high tem- perature reached during the progress of the reaction. However, by heating together sulphur, ammonium chloride, and finely divided tin, stannic sulphide is obtained as a golden yellow, crystalline mass which is used for bronzing, being called mosaic gold. Analytical Tests for Tin. When mixed with sodium car- bonate and ignited on charcoal, tin compounds are reduced, yielding bright globules of the metal. The behavior of stan- nous chloride toward mercuric chloride is often used to char- acterize tin. Furthermore, the reactions of the sulphides of tin, as given above, are of importance, as is also the fact that the hydroxides are precipitated by alkaline carbonates, and do 504 OUTLINES OF CHEMISTRY not dissolve in excess of the latter. Zinc precipitates spongy tin from solutions of tin salts. This may then be dissolved in hydrochloric acid and tested with mercuric chloride. In this way, tin may be distinguished from antimony. Occurrence, Metallurgy, and Properties of Lead. Lead, plumbum, is seldom found in the uncombined state in nature. It occurs chiefly as the sulphide PbS, which commonly crystal- lizes in cubes of the regular system, and is called galenite. The following ores are also found, though rarely and in smaller quantities: cerussite PbCO 3 , wulfenite PbMoO 4 , crocoisite PbCrO 4 , bouronite Cu 2 S Sb 2 S 3 2 PbS, anglesite PbSO 4 , pyro- morphite PbCl 2 Pb 3 (PO 4 ) 2 . Galenite, or galena, is found in fairly large quantities in the United States, Great Britain, Ger- many, Spain, and Australia. To obtain lead from galena, the latter is roasted so as to oxidize it in part to oxide and in part to sulphate, thus : 2 PbS -f 3 O 2 = 2 PbO + 2 SO 2 , and PbS + 2 2 = PbS0 4 . By then strongly heating this mixture of lead oxide, lead sul- phate, and unchanged lead sulphide, out of contact with the air, lead is obtained, thus : 2 PbO + PbS = SO 2 + 3 Pb, and PbSO 4 + PbS = 2 SO 2 + 2 Pb. Lead may also be prepared from galena by heating the latter with iron, thus : PbS + Fe = FeS + Pb. As the ferrous sulphide formed is much lighter than lead, it floats on top of the latter, and is run off like a slag. The mol- ten lead can readily be drawn off from below. The separation of lead from silver, which commonly accompanies lead, has already been described. The world's annual production of lead is about 1,300,000 tons. The United States furnishes about one third of this amount. Lead is a soft, malleable, bluish white metal, having a bright luster on freshly cut surfaces, which soon becomes dull on ex- posure to the air because of the formation of a film of oxide. The specific gravity of lead is 11.4. It melts at 327, and LEAD AND TIN 505 boils at about 1200 in vacuo. By heating lead in the air, the oxide is formed. Water in contact with the air acts on lead somewhat, forming lead hydroxide, which is slightly soluble. Waters containing calcium sulphate or bicarbonate act on lead, forming lead sulphate or carbonate. These salts are insoluble, and so form a coating that protects the metal from further action. Hence it is quite feasible to use lead pipes for conduct- ing drinking water. Hydrochloric and sulphuric acids act on lead but slightly, because the resulting products, lead chloride and lead sulphate, are sparingly soluble. In nitric or' acetic acid lead is readily soluble. Many other weak organic acids dissolve lead, hence it is unsuitable for cooking utensils. The atomic weight of lead is 207.10, and its valence is generally two, though in some compounds the element is quadrivalent. From solutions of its salts, lead is displaced by zinc, iron, and tin. When a strip of zinc is hung in a dilute lead acetate so- lution, a bulky, branching mass of lead, known as a lead tree, is formed thus : Zn + Pb(C 2 H 8 2 ) 2 = Zn(0 2 H 3 2 ) 2 + Pb. Uses of Lead. Lead has been known since earliest times. The Romans used it for water pipes, and Pliny distinguished between lead and tin. Much lead is used as lead pipes in plumbing. These pipes are readily made of hot, plastic lead by means of hydraulic pressure. In the sulphuric acid indus- try, lead is used for lining the chambers and making evaporat- ing dishes and pipes. It is also employed in other chemical operations for containers. Sheet lead further serves for roofs and gutters. Alloyed with tin, lead forms solder, pewter, and Britannia metal, which have already been described. It also enters into the composition of Rose's metal and Wood's metal, which are alloys of low melting points. Type metal is an alloy of lead with antimony, which has been mentioned in connection with the latter. Lead used for shot and bullets contains from 0.2 to 0.4 per cent of arsenic. Babbitt metal consists of 70 to 90 per cent lead alloyed with tin and antimony. It is used for bearings in machines. Large quantities of lead are also used in making storage batteries. Besides this, much lead is consumed in the production of various lead compounds. Oxides of Lead. The following five oxides of lead are 506 OUTLINES OF CHEMISTRY known: Pb 2 O, PbO, Pb 2 O 3 , Pb 3 O 4 , and PbO 2 . Lead suboxide Pb 2 O is a black powder, formed by the action of the air 011 lead, or by heating lead oxalate, thus : 2 PbC 2 4 = Pb 2 + 3 C0 2 + CO. Lead oxide, or plumbic oxide, PbO is formed by burning lead in the air. It results as a by-product in the separation of lead from silver by cupellation. Lead oxide is also made by calcin- ing the carbonate or nitrate. It is a yellow, amorphous pow- der ; but when fused and allowed to cool, it forms a crystalline mass. This, when pulverized, is termed litharge. It is used in making glass of high refracting power, in glazing pottery, deco- rating porcelain, and preparing many lead compounds. Lead hydroxide Pb(OH) 2 is formed by adding caustic alkali to a solution of a lead salt. Lead hydroxide is a strong base and is appreciably soluble in water, imparting a faintly alkaline reaction to the latter. With caustic alkalies, however, lead hydroxide exhibits acid properties, forming plumbites, thus : 2 NaOH + Pb(OH) 2 = Na 2 PbO 2 + 2 H 2 O. On boiling plumbic oxide with caustic alkalies, plumbites are also formed. Lead sesquioxide Pb 2 O 3 is an orange-yellow powder, obtained by the action of a hypochlorite on a solution of a plumbite. Red lead, or minium, Pb 3 O 4 is formed by heating plumbic oxide in the air to about 400. It is a bright red powder, which is generally contaminated with some litharge, just as the latter always contains small amounts of minium. Red lead is used as a pigment in paints. When treated with nitric acid, it reacts thus : Pb 3 4 + 4 HN0 3 = 2 Pb(N0 3 ) 2 + PbO 2 + 2 H 2 O. It is evident, then, that minium behaves like a mixture of 2 PbO and PbO 2 . However, it is probably a plumbate of lead, Pb 2 PbO 4 ; and as nitric acid acts upon it, lead nitrate and plumbic acid H 4 PbO 4 are formed. The latter, being very unstable, is decomposed at once into 2 H 2 O and PbO 2 . Lead peroxide PbO 2 is a brown powder, which is also formed by the action of hypochlorites upon lead salts in alkaline solu- tion, and at the anode during the electrolysis of lead salts. LEAD AND TIN 507 With sulphuric acid it forms lead sulphate and oxygen, and with hydrochloric acid lead chloride and chlorine, thus : 2 Pb0 2 + 2 H 2 S0 4 = 2 PbSO 4 + 2 H 2 + O a . Pb0 2 + 4 HC1 = PbCl 2 + 2 H 2 + C1 2 . In hot, concentrated solutions of caustic alkalies lead peroxide dissolves, forming plumbates, thus : 2 KOH + Pb0 2 = K 2 Pb0 3 + H 2 0. Plumbates are salts of metaplumbic acid H 2 PbO 3 > They are analogous to carbonates, silicates, and stannates. All oxides of lead when strongly ignited in the air are finally converted into litharge, thus : Pb 2 +0 = 2 PbO. Pb 2 3 = 2 PbO + O. Pb 3 O 4 = 3 PbO + O. Pb0 2 = PbO + O. Halides of Lead. Lead chloride PbCl 2 is a white salt obtained as a precipitate by adding a soluble chloride to a solution of a lead salt. It is sparingly soluble in cold water, but 100 parts of boiling water dissolve about 4 parts of the salt. On ignition in the air, lead chloride forms lead oxychloride Pb 2 OCl 2 . This oxychloride is also produced as Pb 2 OCl 2 H 2 O, Pattinson's white lead, by adding milk of lime to a hot solution of lead chloride: 2 PbCl- 2 + Ca(OH) 2 = Pb 2 OCl 2 - H 2 O + CaCl 2 . Lead tetrachloride PbCl 4 is a yellow oil of specific gravity 3.2, which congeals as a crystalline mass at 15. It is formed by dissolving lead peroxide in well -chilled concentrated hydro- chloric acid, or by passing chlorine into a cold solution of lead chloride in hydrochloric acid. On then adding ammonium chloride, the double salt PbCl 4 2 NH 4 C1 is obtained as a crys- talline precipitate, which when treated with concentrated sul- phuric acid at yields lead perchloride as a yellow oil. The latter readily decomposes into PbCl 2 and C1 2 . At 105 this decomposition proceeds with explosive violence. Lead bromide PbBr 2 and oxybromide Pb 2 OBr 2 are analogous to the corresponding chlorides. Lead iodide PbI 2 is a yellow precipitate formed by adding a 508 OUTLINES OF CHEMISTRY soluble iodide to a solution of a lead salt. It dissolves in about 200 parts of boiling water, from which it crystallizes in shining hexagonal scales. Like the iodide of mercury, it readily forms double salts with iodides of the alkalies. Lead Nitrate Pb(NO 3 ) 2 is formed by the action of nitric acid on lead, lead oxide, or lead carbonate. It forms octahedra which are soluble in about 2 parts of water. Lead Acetate Pb(C 2 H 8 O 2 ) 2 3 H a O is made by dissolving lead or lead oxide in acetic acid. The salt forms prismatic crystals that readily effloresce. They are very soluble in water, and the solution has a sweetish taste, hence the salt is called sugar Df lead, or saccharum saturni. Lead acetate solutions dissolve lead oxide or hydroxide, thus forming basic lead acetate Pb(C 2 H 3 O 2 ) 2 (PbO)^, dilute solutions of which (about 2 per cent) are known as lead water. These basic lead acetate solu- tions readily become milky, due to the formation of lead car- bonate with carbon dioxide of the air. Lead Sulphate PbSO 4 is a white crystalline precipitate formed by adding a soluble sulphate to a solution of a lead salt. It is practically insoluble in water and dilute sulphuric acid. But it dissolves appreciably in hydrochloric and nitric acids, and fairly readily in concentrated sulphuric acid. It is also soluble in caustic alkalies, in sodium thiosulphate, and in ammonium acetate and other ammonium salts of organic acids. Lead Sulphide PbS is found in nature as already stated. It is formed as a black precipitate when hydrogen sulphide is passed into a solution of a lead salt. In dilute acids it is in- soluble also in sulphides of the alkalies; but when boiled with concentrated hydrochloric acid, it forms lead chloride and hydrogen sulphide. When boiled with dilute nitric acid, lead sulphide forms lead nitrate, while concentrated nitric acid oxidizes it to sulphate. The latter change also takes place slowly when moist lead sulphide is exposed to the air. Lead Arsenate Pb 8 (AsO 4 ) a is formed by treating a lead acetate solution with sodium arsenate. Lead arsenate is a white pow- der that is but sparingly soluble in water. Like Paris green, lead arsenate is used for exterminating potato bugs and other insects that infest cultivated plants. Lead Carbonate PbCO 8 is found in nature, as already men- tioned. It is formed as a white precipitate by adding lead LEAD AND TIN 509 nitrate to a solution of ammonium carbonate. When a solution of a lead salt is precipitated with normal sodium or potassium carbonate, basic lead carbonates are obtained whose composi- tion varies with the concentration and temperature of the solu- tions used. White lead, which is used in paints, is a basic lead carbon- ate, commonly of the composition Pb(OH) 2 2 PbCO 3 . It is formed by the action of carbon dioxide on basic lead acetate. Thenard's method, also called the French method, consists of passing carbon dioxide into a solution of basic lead acetate formed by dissolving lead oxide in lead acetate solution. Thus basic lead carbonate is precipitated, while neutral lead acetate remains in solution and can be used over again. This is a rapid, direct process, but it yields a crystalline or granular product which does not possess the desirable covering properties of white lead produced by the Dutch process. The latter has been in use for nearly three hundred years, and consists of allowing the vapors of vinegar and carbon dioxide to act slowly upon large surfaces of lead. To effect this, the sheet lead is loosely rolled together in spirals, each of which is placed in a small earthenware pot into which vinegar, 4 to 5 per cent acetic acid, has been poured. The lead rests on a shelf so as not to be in contact with the vinegar. Lead gratings or other forms of lead castings are sometimes used instead of sheet lead spirals. The earthen pots, which are about 20 cm. high and 12 cm. in diameter, are loosely covered and set in horse manure or spent tan bark. The acetic acid thus vaporizes and acts upon the lead, forming basic lead acetate, which is then acted upon by the carbon dioxide liberated during the process of fermentation of the manure or spent tan, resulting in the formation of white lead. Several weeks are required to complete the action. The fermentation process not only furnishes the required carbon dioxide, but also produces the heat necessary to vaporize the vinegar. An amorphous white lead of superior covering power is obtained in this way. Many other processes are in use, and some of them yield white lead of excellent quality. They all depend upon the action of acetic acid and carbon dioxide upon lead. Electrolytic processes have also been proposed; they de- pend upon the fact that when the electric current is passed be- tween two lead plates dipping in a solution of sodium nitrate 510 OUTLINES OF CHEMISTRY or chlorate, and sodium bicarbonate, basic lead carbonate is precipitated. It is claimed that the product obtained is amorphous. White lead is ground with linseed oil and used as a paint. The oil gives the paint a yellowish tinge which is usually dispelled by adding a trace of blue or black pigment. Barium sulphate, lead sulphate, and calcium carbonate, particularly the former, are used as adulterants of white lead. Their presence can readily be detected. White lead turns dark when exposed to hydrogen sulphide, or when sulphide pigments are mixed with it. It is also poisonous, and care must be exercised in its manu- facture and use so that it will not get into the system. All lead compounds are poisonous ; the readily soluble ones are, of course, the most dangerous. Painters and others that work with lead or its compounds are liable to lead colic, which is produced by the gradual absorption of lead compounds, by breathing the latter in form of fine dust, or by constantly handling them. Analytical Tests for Lead. When mixed with soda and heated on charcoal before the blowpipe, lead compounds yield globules of metallic lead, which upon oxidation are transformed to plumbic oxide. Lead sulphide, sulphate, hydroxide, chlo- ride, and carbonate are all characteristic precipitates which may be obtained from solutions of lead salts as already stated. A lemon-yellow precipitate of lead chromate is produced by add- ing potassium chromate or bichromate to a solution of a lead salt. REVIEW QUESTIONS 1. Compare the chemical properties of lead and tin. How are these metals related to : (a) germanium, (6) carbon and silicon ? 2. How much carbon would be required to reduce one ton of cas- siterite ? Equation. 3. What are the characteristic physical properties of tin? Which prevents the formation of iron rust better, to coat iron with tin or with zinc ? 4. Why is it not advisable to cover roofs with tin in cold countries? 5. How do the following act on tin : (a) concentrated nitric acid, (6) hot concentrated sulphuric acid, (c) hydrochloric acid? Equations. 6. Mention three alloys of tin and state their uses. 7. How form the tetrachloride of tin: (a) from metallic tin, (6) from stannous chloride? TIN AND LEAD 511 8. What is a stannite? Give an example. What compounds of the following metals are analogous to the stannites : zinc, aluminum, antimony, lead. 9. Compare stannic acid with silicic acid and carbonic acid. Write the formulas of the compounds showing the relationship. 10. What two sulphides of tin are known? By equations show how they react with yellow ammonium sulphide. 11. Given a solution of stannous chloride, describe four different tests by means of which it could be shown that the solution contains tin. Equations. 12. What is the main ore of lead ? By equations indicate how lead is prepared from this ore. 13. Describe the uses of metallic lead. 14. What alloys of lead are in common use ? 15. What is litharge? How is it made and what is it used for? 16. What is a plumbate? Give an equation showing how sodium plumbate may be formed. Write the analogous equation for making sodium carbonate. 17. Write the formula expressing the composition of each of the following and state the use of each substance : sugar of lead, lead arsenate, white lead. 18. How test a paint to ascertain if it contains white lead? CHAPTER XXVIII CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM CHROMIUM (Cr 52.0), molybdenum (Mo 96.0), tung- sten (W 184.0), and uranium (U 238.5) are related to the members of the sulphur family about as titanium, zirconium, cerium, and thorium are to the carbon and silicon group, or as vanadium, columbium, and tantalum are to the phosphorus group. The members of the chromium family form the tri- oxides CrO 3 , MoO 3 , WO 3 , and UO 3 , which are analogous to SO 3 . Like the latter, they are acid anhydrides. With me- tallic oxides they form salts like K 2 CrO 4 , K 2 MoO 4 , K 2 WO 4 , which correspond to K 2 SO 4 . Again, the compounds MoO 2 , WO 2 , and UO 2 are analogous to SO 2 . On the other hand, chromium often acts as a base, forming compounds that are analogous to those of aluminum and iron, as in CrCl 3 , and Cr 2 (SO 4 ) 3 . The other members of the group do not form such salts, but they enter into other rather complicated com- pounds. This is particularly the case with uranium. Chro- mium is the most important member of the group. None of the elements of this family occur in nature in the free state. Occurrence, Preparation, and Properties of Chromium. Chro- mium is generally found in form of chromite, also called chrome iron ore, FeO Cr 2 O 3 or Fe(CrO 2 ) 2 , which crystallizes in octa- hedra and is isomorphous with spinel. Crocoisite PbCrO 4 , first found in Siberia, occurs more rarely, though it was in this mineral that chromium was discovered in 1797 by "Vauquelin. The name chromium was given the element because it forms colored compounds. Metallic chromium may readily be prepared by Goldschmidt's process, consisting of igniting a mixture of chromic oxide and finely divided aluminum by means of a fuse of magnesium rib- bon or a mixture of barium peroxide and aluminum powder. The reaction when thus started proceeds to completion. Chro- mium may also be obtained by reducing the oxide with carbon in the electric furnace. 512 CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 513 Chromium is a steel-gray, hard, brittle metal of high metallic luster. Its specific gravity is 6.8. It requires the electric furnace to melt chromium. The metal may be polished, and it remains unchanged in the air. At high temperatures it burns in oxygen or in the air, emitting a brilliant light and forming Cr 2 O 3 . Chromium is not attacked by nitric acid; but in warm dilute hydrochloric or sulphuric acid it dissolves with evolution of hydrogen. Chromium is commonly bivalent, trivalent, or hexavalent. Chromium is used in the steel industry for making chrome steel, in which process ferrochromium, an alloy of iron and chromium containing 60 per cent of the latter, is added to the steel. This ferrochromium is readily prepared by heating chromium ore with carbon in the electric furnace. Chromic Oxide and Hydroxides. With oxygen, chromium forms a sesquioxide, Cr 2 O 3 , which acts as a base, and a trioxide, CrO 3 , which acts as an acid. These are the only oxides of chromium that are known with certainty, though chromous hydf oxide Cr(OH) 2 , in which the metal is bivalent, has also been prepared, being formed as a yellow precipitate when caustic alkali is added to a solution of chromous chloride. Chromic oxide (chromium sesquioxide) Cr 2 O 3 is a grass-green powder formed by ignition of chromic hydroxide, chromium trioxide, or ammonium bichromate. The oxide may be obtained in the form of very dark green, lustrous, hexagonal crystals by passing the vapors of chromyl chloride CrO 2 Cl 2 through a red- hot tube, thus : 2 CrO 2 Cl 2 = Cr 2 O 3 + 2 C1 2 4- O. Amorphous chromic oxide dissolves readily in acids, but after strong ignition the latter scarcely attack it ; and like highly heated oxides of iron or aluminum, chromic oxide is then usually fused with bisulphate of potassium in order to effect its solu- tion. Chromic oxide is used as a pigment, chrome green, in paints. It also serves in coloring glass green. Chromic hydroxide is formed as a grayish blue precipitate of the composition Cr(OH) 3 2 H 2 O when ammonia is added to a solution of a chromium salt. On drying this substance over sulphuric acid in a vacuum, a residue of very nearly the com- position Cr(OH) 3 may be obtained. On heating the latter in 514 OUTLINES OF CHEMISTRY hydrogen to about 220, chromous hydroxide Cr(OH) 2 is formed, which upon strong ignition yields chromic oxide and water. Potassium or sodium hydroxide also precipitates hydrated chromic hydroxide from solutions of chromium salts, but the precipitate always contains some alkali. In excess of potassium or sodium hydroxide, chromic hydroxide is soluble, forming chromites, thus : - NaOH + Cr(OH) 3 = NaCrO 2 + 2 H 2 O. This behavior is similar to that of aluminum hydroxide. How- ever, on boiling solutions of chromites, they are decomposed, chromic hydroxide being precipitated, while aluminates are stable under similar treatment. Insoluble chromites are also known, chrome iron ore Fe(CrO 2 ) 2 being a compound of this character. With acids, chromic hydroxide forms chromic salts ; indeed, it generally acts as a base. Chromous Compounds. In these compounds chromium is bivalent. The hydroxide, Cr(OH) 2 , has already been men- tioned. Chromous chloride CrCl 2 and chromous sulphate CrSO 4 7 H 2 O have been prepared by dissolving the metal in hydrochloric or sulphuric acid. Chromous chloride has also been made by heating chromic chloride in hydrogen. Chro- mous salts have been studied but little. They are strong reducing agents, for they readily pass over into chromic com- pounds. Chromic Salts. These are made by action of acids on chro- mic hydroxide. Chromic chloride CrCl 3 may also be prepared by the action of chlorine on heated chromium, or by passing chlorine over a red-hot mixture of chromic oxide and carbon. Thus prepared, the compound sublimes in violet leaflets that are almost insoluble in water. By ignition in the air these pass over into chromic oxide. On long-continued boiling the salt slowly dissolves; but when chromous chloride is added, even in traces, the action progresses much more rapidly, a green solution being formed from which by evaporation green, deliquescent crystals, CrCl 8 6 H 2 O, may be obtained. The latter yield both water and hydrochloric acid on being heated in the air ; consequently, to prepare the anhydrous chloride from them, they are heated in a current of chlorine or hydro- chloric acid gas. CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 515 Chromic sulphate Cr 2 (SO 4 ) 3 15 H 2 O is deposited from cold solutions in the form of violet crystals. The salt forms violet solutions at ordinary temperatures. On boiling these, they become green because a hydrolysis takes place, the exact nature of which has not yet been definitely determined. From the green solutions no crystals are obtainable ; but on standing, these solutions slowly become violet again and deposit violet crystals. This peculiar behavior is also exhibited by other chromic salts. Chrome alums are double salts which chromic sulphate forms with sulphates of the alkalies. They have the same general formula as other alums, and are isomorphous with them. Potas- sium chrome alum KCr(SO 4 ) 2 12 H 2 O is the commonest of these compounds, and is generally called simply chrome alum. It forms octahedra which are of a dark violet color, but appear reddish by transmitted light. They effloresce on exposure to the air. Their solutions exhibit color changes similar to those of chromic sulphate. Chromates, Bichromates, and Chromium Trioxide. The chief source of chromium is chrome iron ore, as already stated. When this is pulverized, mixed with potash and calcium carbonate, and roasted in contact with air, ferric oxide, potassium chro- mate, calcium chromate, and carbon dioxide are formed, thus : 4 Fe(CrO 2 ) 2 + 6 K 2 CO 3 + 2 CaCO 3 + 7 O 2 = 6 K 2 Cr0 4 + 2 CaCr0 4 + 2 Fe a O 8 + 8 CO 2 . On lixiviating the mass with water, potassium and calcium chromates dissolve, and oh adding potassium sulphate to the solution, calcium sulphate is formed, thus : CaCr0 4 + K 2 S0 4 = CaSO 4 + K 2 CrO 4 . On evaporating the clear solution of potassium chromate thus obtained, the salt crystallizes out in rhombic pyramids of lemon color which are isomorphous with potassium sulphate. Potas- sium chromate is soluble in about 2 parts of water. The solu- tion has an alkaline reaction because of partial hydrolysis of the salt. On adding sulphuric acid to a solution of potassium chro- mate, the latter turns orange-red and readily deposits large, red, triclinic crystals of potassium bichromate K 2 O 2 O 7 . These are COLLE8I DWARWIACY 516 OUTLINES OF CHEMISTRY soluble in about 1 part of water at 100, whereas at room tem- peratures about 10 parts of water are necessary to effect their solution. The salt can consequently readily be purified by recrystallization. Potassium bichromate is analogous to po- tassium pyrosulphate K 2 S 2 O 7 . On heating potassium bichro- mate, it melts and then decomposes : 2 K 2 Cr 2 O 7 = 2 K a CrO 4 + Cr a O 8 + 3 O. When heated with concentrated sulphuric acid, chrome alum and oxygen are formed : K 2 Cr 2 O 7 -f- 4 H 2 SO 4 = 2 KCr(SO 4 ) 2 + 4 H 2 O + 3 O. With hot concentrated hydrochloric acid, chlorine is evolved : K 2 Cr 2 O 7 + 14 HC1 = 2 CrCl 3 + 2 KC1 + 7 H 2 O + 3 C1 2 . From these reactions it is clear that potassium bichromate is a strong oxidizing agent, each molecule yielding three oxygen atoms that are available to effect oxidations. If, in the roasting of chrome iron ore with calcium carbonate and potash, soda is substituted for the latter, sodium chromate is produced. It is a yellow salt that may be obtained in deli- quescent, prismatic crystals, Na 2 CrO 4 10 H 2 O, which are iso- morphous with Glauber's salt. On treatment with sulphuric acid, sodium chromate may be transformed to sodium bichro- mate Na 2 Cr 2 O 7 2 H 2 O, which forms red, triclinic crystals, soluble in about 1 part of water at ordinary temperatures. This salt is cheaper than potassium bichromate and hence is generally used in place of the latter in the industries. By adding potassium chloride to solutions of sodium bichromate, potassium bichromate may be obtained, for it is less soluble. Gelatine treated with potassium bichromate solution darkens on exposure to light and becomes insoluble, probably because of the formation of chromic oxide and partial oxidation of the gelatine. This fact is used in photographic processes, and also in making a glue that gradually becomes insoluble. Bichro- mates are used in dyeing and tanning, also as oxidizing agents in the laboratory and in the industries. Ammonium bichromate (NH 4 ) 2 Cr 2 O 7 forms readily soluble red crystals which on igni- tion continue to oxidize with brilliant scintillations : (NH 4 ) 2 Cr 2 7 = Cr.0. + 4 H 2 + N r CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 517 Lead chromate PbCrO 4 is a bright yellow precipitate formed by adding either a soluble chromate or bichromate to a solution of a lead salt. It serves as a pigment in paints, being called chrome-yellow. It may be fused without decomposition. On cooling, it then forms a brown crystalline solid, which when heated with carbon compounds readily oxidizes the latter, hence its use in organic combustion analyses. Barium chromate BaCrO 4 is a yellow precipitate prepared by adding a soluble chromate or bichromate to a solution of a barium salt. It is insoluble in acetic acid, but soluble in nitric or hydrochloric acid. It is also used as a pigment, under the name ultramarine yellow. Calcium chromate CaCrO 4 2 H 2 O is analogous to gypsum CaSO 4 2 H 2 O. Magnesium chromate MgCrO 4 7 H 2 O is analogous to Epsom salt MgSO 4 7 H 2 O. With the exception of barium chromate, the chromates of the alka- line earths all dissolve in acetic acid, which fact is used in analysis. Chromates of the heavy metals are insoluble in water and hence are readily prepared by precipitation. Silver chromate Ag 2 CrO 4 and mercurous chromate Hg 2 CrO 4 are characteristic red pre- cipitates. On adding concentrated sulphuric acid to a well-cooled con- centrated solution of sodium or potassium bichromate, chro- mium trioxide, or chromic acid anhydride, CrO 3 separates out in form of beautiful, dark red, deliquescent, rhombic needles. Its aqueous solutions probably contain bichromic acid H 2 Cr 2 O 7 , but the latter has not been isolated. On neutralization with alka- lies, such solutions of chromic acid yield bichromates and chro- mates. Chromium trioxide is a powerful oxidizing agent. It destroys organic tissues and oxidizes many compounds, thus : 2 CrO 3 + 3 H 2 S = Cr 2 O a + 3 S + 3 H 2 O. Cr0 3 + 6 HC1 = CrCl 8 + 3 H 2 O + 3 01. 2 Cr0 3 + 3 S0 2 = Cr 2 (S0 4 ) 3 . 2 Cr0 8 + 3 H 2 S0 4 + 3 C 2 H 5 . OH = Cr 2 (S0 4 ) 3 + 3 CH 3 CHO + 6 H 2 O. 2 Cr0 8 + 3 H 2 S0 4 + 3 (COOH) 2 = Cr 2 (SO 4 ) 3 + 6 H 2 O + 6 CO 2 . It should be borne in mind that compounds of chromium are poisonous, those that are readily soluble being most harmful. Chromyl Chloride Ci O. 2 C1 2 is analogous to sulphuryl chloride 518 OUTLINES OF CHEMISTRY SO 2 C1 2 . It is a dark red, fuming liquid commonly made by distilling a mixture of sodium chloride, sodium or potassium bichromate, and sulphuric acid, thus: 4 NaCl + Na 2 2 7 + 6 H 2 SO 4 = 6 NaHSO 4 + 3 H 2 O + 2 CrO 2 Cl 2 . Water decomposes the compound: CrO 2 - C1 2 + H 2 O = CrO 8 + 2 HC1 ; hence, in preparing it, a sufficient amount of sulphuric acid must be used to absorb the water formed during the reaction. Analytical Tests for Chromium Compounds. Because of their color, these compounds are readily detected. Chromous com- pounds readily pass over into chromic compounds. The latter yield a green coloration when heated in the borax or sodium metaphosphate bead. Caustic alkalies precipitate chromic hy- droxide, which is soluble in an excess of the precipitant, but is again precipitated on boiling. Alkaline carbonates precipitate chromic hydroxide. Hydrogen sulphide produces no precipi- tate in solutions of chromic salts, but ammonium sulphide pre- cipitates chromic hydroxide. The latter, like the hydroxide of aluminum, is scarcely affected by an excess of ammonium hydroxide. When fused with soda and saltpeter, chromic compounds yield chromates; these are characterized by their yellow color and by the characteristic insoluble precipitates which their solutions yield with silver nitrate, barium chloride, lead nitrate, and mercurous nitrate. Furthermore, when acidi- fied with sulphuric acid, chromate solutions readily oxidize oxalic acid, alcohol, ferrous salts, etc., yielding green solutions of chromic sulphate. Molybdenum. This metal is found in molybdenite MoS 2 , which looks like graphite, and in wulfenite PbMoO 4 , which forms yellowish tetragonal crystals. On roasting molybden- ite, it is converted to molybdenum trioxide MoO 3 , a white crys- talline powder which is the most stable oxide of molybdenum. This powder readily dissolves in caustic alkalies or ammonium hydroxide, forming molybdates. Of these ammonium molyb- date (NH 4 ) 2 MoO 4 is an important reagent in analytical chemistry, for with phosphoric acid or phosphates in dilute nitric acid solu- tion it yields a characteristic light yellow precipitate of ammo- nium phosphomolybdate (NH 4 ) 8 PO 4 - 11 MoO 8 6 H 2 O. The CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 519 composition of the latter varies somewhat according to the conditions of precipitation. In dilute acids the compound is insoluble, but in alkalies or an excess of phosphoric acid it dis- solves. On treating a solution of a molybdate with nitric acid, yellow crystals of molybdic acid H 2 MoO 4 H 2 O separate out, which upon drying yield H 2 MoO 4 . Similarly when ammo- nium phosphomolybdate is treated with aqua regia, crystals of phosphomolybdie acid H 3 PO 4 11 MoO 8 12 H 2 O are formed. This acid forms insoluble salts with salts of potassium, ammo- nium, and the alkaloids, and is consequently of value in ana- lytical tests for the latter. It is known as Sonnenschein's reagent. Molybdenum receives its name from the Greek word mean- ing lead, for molybdenite was regarded as graphite, which was called black lead. Scheele made molybdic acid in 1778 by treating molybdenite with nitric acid, thus showing that the mineral is not graphite. In 1790 Hjelm prepared the metal by heating the oxides or chlorides in a current of hydrogen, and Klaproth determined the true nature of wulfenite in 1797. Molybdenum may also be made by heating the oxide with car- bon in the electric furnace. It is a hard, lustrous metal of specific gravity 9.1. Its physical properties resemble those of iron. Like the latter, it may be welded and tempered. Besides the compounds already mentioned, molybdenum forms the oxides Mo 2 O 3 and MoO 2 ; the chlorides (MoCl 2 ) 3 , MoCl 3 , MoCl 4 , MoCl 5 ; and the sulphides MoS 3 and MoS 4 . It *also forms a large variety of additional, rather complicated, com- pounds. Tungsten. This metal is found in the minerals scheelite CaWO 4 , wolframite (FeMn)WO 4 , and stolzite PbWO 4 . By treating the pulverized ores with nitric acid, tungsten trioxide WO 3 is obtained as a yellow powder which is insoluble in acids. With caustic alkalies, it forms tungstates like Na 2 WO 4 and K 2 WO 4 , which are soluble, and from whose solutions tung- stic acid H 2 WO 4 -H 2 O is precipitated by means of acids. Sodium tungstate may also be made by fusing native tungstates with soda and lixiviating the mass with water. The salts Na 2 WO 4 2 H 2 O and Na 2 W 4 O 13 . 10 H 2 O are used as mordants. They are also employed in making fabrics fireproof. Tung- sten forms the oxides WO 2 and WO 8 ; the chlorides WC1 2 , 520 OUTLINES OF CHEMISTRY WC1 4 , WC1 6 , WC1 6 , WO 2 C1 2 , and WOC1 4 ; and the carbides W 2 C and WC. In addition, tungsten enters into a very large number of complex compounds. Phosphotungstic acid is known as Scheibler's reagent. It consists of a compound of tungstic and phosphoric acids, which is analogous to phosphomolybdic acid. Tungsten was discovered in 1781 by Scheele, and two years later the element was isolated by the d'Elhujar brothers. Tungsten may be made by igniting the oxides or chlorides in a current of hydrogen, by heating the oxides with carbon in the electric furnace, or by the Goldschmidt process. It is a hard, steel-gray, lustrous, brittle metal of specific gravity 19.13. Its melting point lies very high. On ignition, it burns to the trioxide. Tungsten is used in the steel industry, for its pres- ence to the extent of 5 to 8 per cent in steel makes the latter very hard and tough. Tungsten is now also being used in making filaments for incandescent electric lamps, which have a higher efficiency than the lamps using a carbon filament. Uranium. The mineral pitchblende or uraninite UO 2 -2 UO 3 or U 3 O 8 constitutes the chief source of uranium, which, like molybdenum and tungsten, belongs to the rarer elements. The true nature of pitchblende was first recognized by Klaproth in 1789; but metallic uranium was not prepared till 1841, when Peligot obtained it by heating metallic sodium and urarious chloride together. The element is named after the planet Uranus, discovered in 1787. Metallic uranium may also be obtained by reduction of its oxides with carbon in the electric furnace, by heating the oxides with aluminum, or by elec- trolysis of the molten double chloride UC1 4 2 NaCl. Uranium is a very hard, white, lustrous metal, not unlike iron in appear- ance. Its specific gravity is 18.7; and its melting point is about 1500. It has the highest atomic weight of all the ele- ments known, namely, 238.5. Uranium shows variable valence in its compounds, which are very numerous. The maximum valence exhibited by uranium is eight. On treating pitchblende with nitric acid, uranyl nitrate UO 2 (NO 3 ) 2 6 H 2 O may be obtained. This salt crystallizes in greenish yellow, rhombic prisms that show fluorescence, a characteristic of many uranium salts. By careful ignition, uranyl nitrate may be converted into uranium trioxide UO 3 , CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 521 a dark yellow powder, which is uranic acid anhydride. On digesting uranium trioxide with nitric acid, uranic acid UO 2 (OH) 2 or H 2 UO 4 may be obtained as an amorphous, yellow powder. Towards strong acids, this acts as a base, forming the uranyl salts containing the bivalent radical uranyl UO 2 . Thus we have uranyl chloride UO 2 - C1 2 ; uranyl sulphate UO 2 -SO 4 .3H 2 O ; uranyl nitrate UO 2 (N0 3 ) 2 .6H 2 O; uranyl acetate UO 2 (0 2 H 3 O 2 ) 2 ; uranyl ammonium carbon- ate UO 2 CO 3 - 2(NH 4 ) 2 CO 3 ; uranyl ammonium phosphate UO 2 -NH 4 -PO 4 . The latter salt is obtained as a precipitate, insoluble in acetic acid, when uranyl acetate is added to a solu- tion of a soluble phosphate containing ammonium chloride. This fact is sometimes used in the estimation of phosphoric acid. By the addition of caustic alkalies to solutions of uranyl compounds, uranates are precipitated. These are not deriva- tives of H 2 UO 4 , but of H 2 U 2 O 7 ; that is, they are diuranates, analogous to pyrosulphates. So there are potassium diuranate K 2 U 2 O 7 -3H 2 O and sodium diuranate Na 2 U 2 O 7 6 H 2 O. The latter is called uranium yellow, and is used in making a beautiful yellowish green, fluorescent uranium glass. Besides the uranyl salts, uranium forms a series of uranous compounds in which the element is quadrivalent. So there are uranous chloride UC1 4 , uranous sulphate U(SO 4 ) 2 8 H 2 O, uranous hydroxide U(OH) 4 , and uranous oxide UO 2 . The latter may be obtained as a steel-gray crystalline powder by ignition of the trioxide in hydrogen; it was regarded as the element itself, till Peligot showed that this was erroneous. Uranous oxide is used in making a fine jet-black glass. Uranium Carbide U 2 C 3 may be made by heating uranium oxides with carbon in the electric furnace. It is harder than quartz. The sulphides, US 2 and UO 2 S, are also known. The variable valence of uranium is readily apparent from the follow- ing list of its chlorides and oxides : UC1 3 , UC1 4 , UC1 6 ; UO 2 , U 8 8 , U0 3 , U0 4 . Compounds of uranium exhibit the phenomena of radio- activity which led to the discovery of radium as already stated. The radio-activity of uranium bearing minerals is proportional to their uranium content. 522 OUTLINES OF CHEMISTRY REVIEW QUESTIONS 1. Why is chromium so named? What valence has this element in its compounds? Illustrate. 2. What other elements are classed with chromium in the periodic system ? Why ? 3. Is chromium an acid or base forming element? Illustrate. 4. Describe a simple way of preparing metallic chromium. Equation. 5. What is ferrochromium and what use is made of it? 6. Write the formulas of the oxides and hydroxides of chromium. Compare these compounds with the corresponding ones of aluminum. 7. How much chromium is there in 1000 grams of pure chrome green ? What is the latter used for? 8. What is chrome iron ore chemically ? How may it be decomposed ? 9. Write the formulas of the following compounds : chromous chlo- ride, chromic sulphate, chrome alum, potassium chromate, potassium bichromate, silver chromate, chromic acid anhydride, barium chromate. 10. By formulas show the analogy between the chromates, molybdates, and sulphates, also between the bichromates and pyrosulphates. 11. By means of an equation express the change that occurs when potassium bichromate is treated with hot, concentrated hydrochloric acid. How many available oxygen atoms has each molecule of a bichromate? Explain. 12. Why is it that potassium bichromate is the commonest chromium compound in the market? What is the compound used for? 13. How demonstrate that there is chromium in a piece of lead chro- mate? Give at least four different tests, writing the equations. 14. What use is made of ammonium molybdate in analytical chem- istry? Discuss the importance of this reagent. 15. How may metallic tungsten be prepared? For what purpose is it used? 16. What is pitchblende? Why has this mineral attracted particular attention in recent years? 17. What can you say of the valence of the element uranium? How does its atomic weight compare with that of other elements ? 18. How much metallic chromium could be made from 300 grams of chromium sulphate ? 19. Given 700 grams of chromium hydroxide, how much chromic acid anhydride could be formed from it ? CHAPTER XXIX MANGANESE MANGANESE is a metal which forms quite a variety of com- pounds. The element is bivalent and basic in character in the manganous compounds which resemble those of the magnesium and iron groups. In manganic compounds manganese is trivalent, thus resembling aluminum, chromium, and trivalent iron. In its dioxide and the manganites, manganese is quadrivalent and shows analogies to tin and lead. Again, manganese exhibits similarities to sulphur and the chromium group by forming an acidic trioxide and manganates that are analogous to sulphates, chromates, molybdates, etc. Finally, the metal exhibits re- semblances to the halogens in forming a heptoxide and per- manganic acid and permanganates, which are analogous to chlorine heptoxide, perchloric acid, and the perchlorates. In the periodic system of the elements manganese occurs in group VII with the halogens, but it should be definitely stated that it is never univalent like the latter. Occurrence, Preparation, and Properties. Manganese is some- times found uncombined in meteoric iron, otherwise it occurs in chemical combination in pyrolusite MnO 2 , braunite Mn 2 O 3 , manganite Mn 2 O 3 -H 2 O, hausmannite Mn 3 O^, manganese blende MnS, and rhodochrosite, or manganese spar, MnCO 3 . The chief of these ores is pyrolusite. The main localities are Russia, Brazil, India, Germany, and the United States. In small amounts, manganese is very widely distributed in soils and rocks, also in traces in plants and animals. The annual output of manganese ores is about 800,000 tons, most of which comes from the Caucasus region. The United States produces but 5000 tons of manganese ores annually. Manganese is prepared by heating its oxides with carbon in an electric furnace, or more readily by the Goldschmidt process of igniting the oxides with aluminum. Manganese is a hard, steel-gray, brittle metal of specific gravity 8.0. In outward appearance it is not unlike cast iron. It melts at about 1300. On exposure to moist air, it gradually 523 524 OUTLINES OF CHEMISTRY assumes a reddish hue due to superficial oxidation. From boil ing water it evolves hydrogen. In dilute acids it readily dis- solves, forming hydrogen and manganous salts. Manganese is non-magnetic. Its atomic weight is 54.93. The alloys of manganese with iron are important in steel manufacture. The alloys usually employed are spiegeleisen, which contains from 10 to 20 per cent manganese, and ferro- manganese, which contains from 20 to 80 per cent. Alloys of manganese and copper contain about 30 per cent manganese; they are called manganese bronze, are very hard, and possess great tensile strength. Though pyrolusite MnO 2 was known for a long time, it was not until 1774 that its real nature, was discovered by Scheele. In 1807 Gahn reduced the oxide, and so isolated the metal. Oxides. Manganese forms the following oxides: the monox- ide MnO, the sesquioxide Mn 2 O 3> the protosesquioxide Mn 3 O 4 , the peroxide MnO 2 , the trioxide or manganic anhydride MnO 3 , and the heptoxide or permanganic anhydride Mn 2 O 7 . Of these the first three are basic in character. Manganese dioxide yields manganous salts on treatment with acids, and half of its oxygen becomes available to effect oxidations. With strong bases, it forms manganites. The oxides, MnO 3 and Mn 2 O 7 , are acidic. On ignition in the air, all oxides of manganese are finally changed to Mn 3 O 4 , which is probably Mn(MnO 2 ) 2 ; and hausmannite is therefore analogous to the spinels, though its crystals are tetragonal. Manganous oxide MnO is a green powder formed by heating higher oxides in hydrogen, or by ignition of manganous carbon- ate out of contact with the air. Manganous hydroxide is a white precipitate formed by the addition of caustic alkalies to solutions of manganous salts. It readily turns brown because of oxidation to manganic hydroxide Mn(OH) 3 ; the latter easily loses water and changes to MnO OH, which on careful ignition yields a brown powder, manganese sesquioxide Mn 2 O 3 . This forms manganous nitrate and manganese dioxide MnO 2 on digesting with nitric acid. By ignition of the nitrate, the di- oxide may also be obtained, though on being strongly heated it loses oxygen and passes over into Mn 3 O 4 . Manganese dioxide is also produced at the anode when manganous salts are electro- lyzed. It is a conductor of electricity, and like lead peroxide MANGANESE 525 it is frequently used as an anode in batteries. With hot hydro- chloric acid, manganese dioxide yields manganous chloride and chlorine; while in cold hydrochloric acid a dark brown solution is formed. This probably contains MnCl 4 , which decomposes into MnCl 2 and chlorine on warming. With lime, manganese dioxide forms manganites of the composition CaO-MnO 2 and CaO - 2 MnO 2 . These act like a mixture of CaO and MnO 2 , of which fact advantage is taken in the Weldon process of again using MnCl 2 liquors for making chlorine. Manganese trioxide is an unstable, dark red mass. Manganese heptoxide Mn 2 O 7 is a dark, reddish green, oily liquid obtained by treating potassium permanganate with concentrated sulphuric acid. The mixture must be carefully cooled with ice, for otherwise violent decom- position will occur. Salts of Manganese. The stable salts in which manganese acts as a base are the manganous compounds. In these the element is bivalent. They are obtained by dissolving any oxide or hy- droxide of manganese with an acid. In the latter process, the higher oxides or hydroxides yield oxygen that is available to effect oxidations. Manganous chloride MnCl 2 4 H 2 O forms pink, deliquescent, monoclinic crystals. The salt is obtained as a by-product when chlorine is prepared by the action of hydrochloric acid on man- ganese dioxide. Manganous chloride may be dehydrated by heating it in a current of hydrochloric acid gas. It is the only chloride of manganese that has been prepared in pure form. The double salt MnCl 2 - 2 NH 4 C1 H 2 O may readily be made. It forms crystals of the isometric system, which on heating yield anhydrous manganous chloride, just as by heating magnesium ammonium chloride, anhydrous magnesium chloride is obtained. Manganous sulphate MnSO 4 -7H 2 O separates from solutions at temperatures below 6 in form of pink monoclinic crystals that are isomorphous with other vitriols like FeSO 4 7 H 2 O, ZnSO 4 - 7 H 2 O, etc. Between 6 and 20 the salt crystallizes in triclinic crystals MnSO 4 5 H 2 O, that are isomorphous with blue vitriol CuSO 4 -5H 2 O; and above 25 orthorhombic crystals MnSO 4 -4 H 2 O are obtained. Manganous sulphate forms double salts with alkali sulphates, like K 2 SO 4 - MnSO 4 -6 H 2 O. These sulphates are isomorphous with similar salts of magnesium, zinc, nickel, cobalt, and iron. 526 OUTLINES OF CHEMISTRY Manganous nitrate Mn(NO 8 ) a 6 H 2 O, is a deliquescent, pink salt which melts in its water of crystallization at about 25. Manganous carbonate MnCO 8 occurs in nature in reddish, hexagonal crystals as manganese spar. It is also formed as a white, insoluble precipitate by adding soluble carbonates to so- lutions of manganous salts. In water charged with carbon dioxide it dissolves, which behavior is similar to that of calcium carbonate. Manganic chloride MnCl 3 is supposed to exist in the brown liquid obtained by dissolving manganese dioxide in cold hydro- chloric acid. It has never been isolated. Manganic sulphate Mn 2 (SO 4 ) 3 is a dark green powder formed by heating manganese dioxide with concentrated sulphuric acid. It is unstable and readily passes over into manganous sulphate, sulphur dioxide, and oxygen on heating. With alkali sul- phates, manganic sulphate forms double salts that are isomor- phous with alum, like K 2 SO 4 .Mn 2 (SO 4 ) 3 -24 H 2 O. Manganates and Permanganates. Manganates are salts of manganic acid H 2 MnO 4 , which has not been isolated. By fus- ing manganese dioxide with caustic potash, potassium manga- nate K a MnO 4 is formed, as a green mass, thus : 3 MnO 2 + 2 KOH = Mn 2 O 8 + K 2 MnO 4 + H 2 O. On treating the mass with water, a dark green solution is ob- tained, from which, on evaporation, greenish black, rhombic crystals of K 2 MnO 4 are deposited These are isomorplious with potassium sulphate and potassium chromate. Manganates of sodium or potassium may be obtained by fusing any oxide or salt of manganese with sodium or potassium carbonate or hydrox- ide, plus sodium or potassium nitrate or chlorate. The presence of the oxidizing agent, or even the action of the oxygen of the air, insures the more complete conversion of the manganese compound into manganate. Solutions of manganates are stable only when they contain an excess of caustic alkali. On dilution with water, manganates suffer decomposition, thus : 3 K 2 MnO 4 + 2 H 2 O = 2 KMnO 4 + 4 KOH + MnO 2 . This change is effected more readily by passing carbon dioxide through the solution of the manganate : 3 K a MnO 4 + 2 CO a = 2 KMnO 4 + 2 K 2 CO 8 + MriO a . MANGANESE 527 The change may also be produced by other dilute acids or by the addition of chlorine. The new salt thus formed is potassium permanganate KMnO 4 . Its solutions have a beautiful purple-red color, and so in the reactions just mentioned the green solution of the manganate gradually changes through blue and violet to the characteristic purple color of the permanganate. Because of this change of color, Scheele called potassium manganate chameleon mineral. Sometimes potassium permanganate solution is called chameleon solution, for on treating it with hot, concentrated, caustic alkali, it again turns green, because of the formation of the manganate, thus : 2 KMn0 4 + 2 KOH = 2 K 2 MnO 4 + H 2 O + O. Potassium permanganate forms very dark purple, lustrous crystals of the rhombic system. They are isomorphous with crystals of potassium perchlorate KC1O 4 . On heating potassium permanganate, it is decomposed : 2 KMn0 4 = Mn0 2 + K a MnO 4 + O 2 . When treated with concentrated sulphuric acid in the cold, per- manganic anhydride Mn 2 O 7 is formed, as already mentioned, thus : 2 KMn0 4 + H 2 S0 4 = Mn a O T + K 2 SO 4 + H 2 O. This heptoxide is unstable. It gradually decomposes : At somewhat- elevated temperatures, this reaction proceeds with explosive violence. The vapors of the heptoxide are violet. Paper, alcohol, ether, illuminating gas, and other combustible substances burst into flame when brought in contact with per- manganic anhydride, because of the tremendous oxidizing power of the ozone that is being liberated. By dissolving permanganic anhydride in water at 0, a purple red solution of permanganic acid HMnO 4 is obtained. This may also be made by treating a solution of barium permanganate with sulphuric acid, for thus barium is precipitated as barium sulphate. Though much more stable than manganic acid, nevertheless permanganic acid HMnO 4 -^H 2 O gradually decomposes, especially on being heated or exposed to light, thus : 2 HMnO 4 = 2 MnO 2 + H 2 O + ?. O. 528 OUTLINES OF CHEMISTRY Uses of Permanganates. On account of their great oxidizing power, permanganates are used as disinfectants, and as oxidiz- ing agents in many chemical processes. They also serve in the preparation and analysis of many substances in the laboratory. Potassium permanganate is commonly employed ; though the cheaper sodium permanganate NaMnO 4 , which does not crystal- lize, is also made and sold in solutions as Candy's disinfecting fluid. Wood alcohol is readily oxidized to formic aldehyde by potassium permanganate, which fact is used to produce formic aldehyde in fumigating infected houses, etc. In alkaline or neutral solutions, potassium permanganate yields oxygen that is available for oxidation, with concomitant formation of manganese dioxide : 2 KMnO 4 + H 2 O = 2 KOH + 2 MnO 2 + 3 O. Two molecules of permanganate thus yield three available oxy- gen atoms. The oxidation of wood alcohol to formic aldehyde would be expressed thus : 2 KMnO 4 + 3 CH 3 OH == 2 KOH + 2 MnO 2 + 3 HCHO + 2 H 2 O. In alkaline solutions potassium permanganate serves for the destruction of organic matter in the analysis of waters, ferti- lizers, etc., the nitrogen present in the substances being simul- taneously liberated as ammonia. When potassium permanganate is to be used in acid solution, sulphuric acid is commonly employed. Thus, when reducing substances are present, a colorless solution of manganous sul- phate and potassium sulphate is produced. In sulphuric acid solutions, two molecules of potassium permanganate yield five atoms of oxygen that are available for oxidation : 2 KMn0 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 3 H 2 O + 5 O. Since the oxidation of oxalic acid to carbon dioxide and water requires one atom of oxygen per molecule of oxalic acid, we have : (COOH) 2 + O = 2 C0 2 + H 2 0. Consequently if the oxidation is effected by means of potassium permanganate, we have : 2 KMn0 4 + 3 H 2 SO 4 + 5 (COOH) 2 = 10 CO 2 + K 2 SO 4 + 2 MnSO 4 + 8 H 2 O. MANGANESE 529 Ferrous sulphate is converted to ferric sulphate, thus : 2 FeSO 4 + H 2 SO 4 + O = Fe 2 (SO 4 ) 3 + H 2 O. Therefore, if the oxidation is carried on by means of potassium permanganate, we have : 2 KMnO 4 + 8 H 2 SO 4 + 10 FeSO 4 = 5 Fe 2 (SO 4 ) 3 + K 2 SO 4 + 2 MnSO 4 + 8 H 2 O. With nitrous acid the reaction is : 2 KMnO 4 + 3 H 2 SO 4 + 5 HNO 2 = K 2 SO 4 + 2 MnSO 4 + 5 HNO 3 + 3 H 2 O. Potassium permanganate and hydrogen peroxide mutually re- duce each other, thus : This reaction is commonly used to determine the content of a solution of hydrogen peroxide by means of potassium perman- ganate. Analytical Tests for Manganese. Manganous sulphide MnS is precipitated as a flesh-colored, hydrous, amorphous substance, when ammonium sulphide is added to an aqueous solution of any compound of manganese. This sulphide readily turns dark on exposure to the air, because oxidation takes place. In dilute acid, even in acetic acid, manganous sulphide is readily soluble. From solutions of manganous salts, soluble carbonates precipi- tate manganous carbonate, and caustic alkalies precipitate man- ganous hydroxide. The latter is soluble in solutions of ammonia or ammonium salts. Amethyst-colored beads are produced by manganese com- pounds with either borax or microcosmic salt. Furthermore, when fused with a mixture of soda and saltpeter, a green manganate is obtained which, when dissolved in water and treated with carbon dioxide, yields the characteristic purple color of permanganate solutions 530 OUTLINES OF CHEMISTRY REVIEW QUESTIONS 1. Justify the position which manganese occupies in the periodic system. 2. Compare the physical properties of manganese with those of chromium and iron. 3. Manganese may form manganous, and manganic compounds, also manganites, manganates, and permanganates. Give an example of each of these, writing the appropriate formula. 4. What is the main ore of manganese? How may the metal be obtained from it ? What practical use is made of the ore itself ? 5. On heating any oxide of manganese in the air, it finally passes over into Mn 3 04. How do you account for this ? What is the valence of manganese in Mn 3 04? 6. What salts of manganese are commonest? Describe them and write their formulas. 7. What practical use is made of potassium permanganate ? How may this compound be made from manganese dioxide? Equations. 8. How many available oxygen atoms are there in potassium per- manganate? Give a concrete illustration demonstrating your answer. 9. How many grams of potassium permanganate would be required to oxidize 25 grams of oxalic acid crystals to carbon dioxide and water? 10. How would you prove that potassium permanganate contains manganese? How show that manganous sulphate contains manganese? What is the action of bromine on a solution of a manganous salt? Ex- plain. 11. Explain how it is that the two oxidizing agents, potassium perman- ganate and hydrogen peroxide, will mutually decompose each other in sulphuric acid solutions, setting oxygen free. Write the equation. 12. How much available oxygen do 1000 grams of potassium perman- ganate contain when that salt is used for oxidation purposes in sulphuric acid solution ? 13. From 256 Ib. of manganese carbonate, how much manganese sul- phide could be formed ? 14. Discuss the occurrence of manganese compounds in nature. 15. How much oxalic acid will ten grams of potassium permanganate oxidize to carbon dioxide and water in sulphuric acid solution ? CHAPTER XXX IRON, NICKEL, AND COBALT THE atomic weights of the elements of this group are nearly alike, being Fe 55.84, Ni 58.68, and Co 58.97. More- . over, it will be recalled that the atomic weight of manganese, 54.93, is but little less than that of iron. The chemical simi- larities between manganese and iron have been mentioned. Iron forms ferrous and ferric compounds. In the former it is bivalent, and in the latter trivalent. Iron is also hexavalent, forming ferrates, which are analogous to manganates and chro- irates. The ferrous compounds are analogous to those of the magnesium group, also to cupric and manganous compounds,' while the ferric compounds are analogous to those of aluminum and chromium. Nickel forms but one series of salts. In these the metal is bivalent. Nevertheless, a sesquioxide and a cor- responding hydroxide of nickel are known, though salts of these are lacking. Cobalt, like iron, forms two series of com- pounds, the cobaltous, in which the metal is bivalent, and the cobaltic, in which it is trivalent. Occurrence of Iron. iron is very abundant and widely dis- tributed. In meteorites it is found uncombined, also in small grains in some of the crystalline rocks. Compounds of iron are found in plant and animal tissues, particularly in the chloro- phyll of plants and the hemoglobin of the blood of animals. Iron is necessary for plant and animal life, though its real func- tion in the vital processes is not understood. All soils and rocks contain compounds of iron. Ores of iron are found in enormous quantities. The most important of these are hema- tite Fe 2 O 3 , magnetite, or magnetic iron ore, Fe 3 O 4 , limonite 2 Fe 2 O 3 + 3 H 2 O, and siderite FeCO 3 . Iron is also found in combination with sulphur, as in pyrite FeS 2 , but these com- pounds are not used for making metallic iron. In rocks, iron is found in the form of oxides and silicates ; and so by the weathering of rocks iron gets into the soil and all natural waters, whence it enters plants and animals. 531 532 OUTLINES OF CHEMISTRY Metallurgy of Iron. Metallic iron has been known to man for thousands of years. The Assyrians used iron knives and saws, the ancient Egyptians reduced iron ores and made steel, both Homer and Hesiod mention the forging of iron for weapons. Through the influence of the Romans iron came into more general use. Still, the metal was costly, because the processes of preparing it from ores were relatively diffi- cult to carry out, imperfectly understood, and known to but few; and so iron was rather slow in replacing bronze for use in weapons and other implements. Iron is made by reducing its ores with carbon at high tem- peratures. The ores, which consist of the oxides, or car- bonate, do not require pre- liminary roasting. The re- duction is effected in blast furnaces, a cross section of a modern type of which is repre- sented in Fig. 151. These furnaces are from 80 to 100 feet high, and have a diameter of about 20 feet where they are widest. They are lined inside with fire brick. The lower end of the furnace is provided with openings of tubes, the tuyeres, through which hot air can be forced into the furnace. The latter is heated and then charged from the top with ore properly mixed with coke and limestone. Instead of coke anthracite coal or charcoal may be used. The purpose of the limestone is to form calcium silicates, or so-called slag, with the sand present in the ore. If the latter contains carbonates of calcium and magnesium, or other basic materials, FIG. 151. IRON, NICKEL, AND COBALT 533 sand is added instead of limestone to make the slag. In any case the material added to the ore to produce the fusible slag is termed the flux. The charge is carried to the top of the furnace by some form of mechanical conveyor, and is introduced through a bell trap, so arranged that while the material is put into the furnace practically no gases escape from the latter. As air heated to about 800 is blown up through the charge, the latter becomes very hot. In the lower part of the furnace carbon dioxide forms, due to the combustion of the coke. The carbon dioxide as it rises through the hot layers of coke is reduced to carbon monoxide, and the latter acts on the ferric oxide and reduces it to iron, thus : Fe 2 O 3 + 3 CO = 2 Fe + 3 CO 2 . The gases, still rich in carbon monoxide, pass out through the vent at the top of the furnace, and are used as fuel to heat the air before it is blown into the furnace. Of late they are also frequently used to run gas engines. The slag and iron settle to the lower part of the furnace, forming two layers, the liquid slag covering the heavier molten iron. The iron is tapped off from below and run into molds of sand about every eight hours. In this way rough bars of cast iron called pigs are produced. The slag runs continually from a lateral opening above the iron. Blast furnace slag is now frequently used for making Portland cement. As the material in the furnace melts down, the latter is fed from the top, so that the process once started is continuous, furnaces remaining in operation for years at a time. In the upper parts of the furnace the temperature is not high enough to melt the iron, which is formed in a porous or spongy condition, and carried down with the slag to the lower and much hotter parts. Here it takes up carbon, form- ing iron carbide in part, which with the iron yields a mixture that melts at a much lower temperature than iron free from carbon. It is this iron containing carbon and other impurities that is tapped from the furnace and made into pig iron. The latter is essentially the same as cast iron, besides which we also have wrought iron and steel, as the main varieties of iron. All kinds of iron used in practice contain carbon and various other impurities, pure iron being practically unknown. Cast Iron. The melting point of cast iron varies from about 584 OUTLINES OF CHEMISTRY 1050 to 1300, according to its content of carbon and othei impurities. Cast iron contains from 2.3 to 5 per cent carbon, besides smaller amounts of silicon, phosphorus, sulphur, and manganese. When cooled very slowly, most of the carbon in cast iron crystallizes out in the form of leaflets of graphite. This iron consequently appears dark gray on its fractured surface and is known as gray cast iron. It is used in making castings, for it has a low melting point, contracts uniformly on cooling, and can afterwards readily be worked with tools. On dissolv- ing gray cast iron in hydrochloric acid, the graphitic carbon remains behind, while the carbon which was combined with the iron in the form of carbides is given off with the hydrogen, being evolved as hydrocarbon gases. Good gray castings contain about 2 to 3 per cent of graphitic carbon, and 1 to 1.5 per cent combined carbon. On cooling cast iron rapidly, practically all the carbon re- mains combined with the iron, forming carbides. Such iron has a silver-white fracture and is known as white cast iron. It is very hard and brittle, and is consequently not employed for castings, but is converted into wrought iron. Cast iron con- taining from 5 to 20 per cent manganese takes up considerable amounts of carbon. Its fracture shows a coarsely crystalline structure, whence the iron is known as spiegeleisen. Ordi- nary cast iron contains from 0.5 to 4 per cent silicon, from 0.4 to 2 per cent phosphorus, and sulphur varying up to about 0.2 per cent. Sulphur and phosphorus make iron brittle and hence are quite objectionable. Wrought Iron. This is nearly pure iron, often containing less than 1 per cent of impurities. It is made by puddling cast iron. This process consists of heating the pig iron with iron oxide in a current of air on the hearth of a reverberatory furnace. Thus the impurities in the iron are oxidized. The carbon largely escapes as carbon monoxide. The silicon and phosphorus after oxidation unite with some of the iron, yielding a slag, which also contains the sulphur. As the heating proceeds, the molten iron becomes more and more viscous. It is continually stirred or puddled so that the air may gain access to it. Finally, the mass becomes so thick that it can be worked up into a ball, which is then taken from the furnace and rolled, or hammered by a steam hammer. IRON, NICKEL, AND COBALT 535 Thus the slag is removed from the iron, and a malleable, duc- tile product is obtained which often contains less than 0.2 pel cent carbon. Wrought iron melts at about 1600, and is plastic enough to be welded at from 900 to 1100. In welding, the two ends to be joined are brought to the welding temperature in a forge. Borax or sand is sprinkled over the parts, which are then again heated for a few moments till the borax or sand has formed a slag with the oxides of iron on the surface. This slag is a borate or silicate of iron. It protects the iron from oxidation. On now hammering the parts together, the slag flies off, and the iron coming into actual contact welds. Be- cause of its very low carbon content, wrought iron will not harden when rapidly chilled, as does cast iron or steel. By heating iron castings covered with pulverized iron ore for about 48 hours, and allowing them to cool slowly, they become sufficiently malleable for many purposes. The process abstracts some of the carbon from the iron, which then is called malleable iron. It is much cheaper than wrought iron, in place of which it is often used when possible. Steel. Steel contains more carbon than wrought iron, but much less than cast iron. It also contains practically no sul- phur or phosphorus, and ordinarily runs low in silicon. The amount of carbon in steel varies from about 0.2 to 1.6 per cent. That which contains the least carbon, 0.2 per cent, approaches wrought iron in quality and is called mild steel. Steel used for building purposes is termed structural steel. It contains from 0.2 to 0.8 per cent carbon, whereas tool steel has a higher carbon content, namely from about 0.8 to 1.5 per cent. Like cast iron, steel may be hardened by heating and then suddenly cooling it. On the other hand, when such hardened steel is again heated to redness and cooled slowly, the material is soft. As in the case of cast iron, the heating and sudden chilling leaves the carbon in the combined condition and' thus makes a hard steel; whereas on slow cooling, the carbon crystallizes out as graphite and so produces a soft, pliable product. By proper heating and cooling, steel of any desired hardness may be pro- duced. This process is called tempering. Pig iron is converted into steel by either the Bessemer process or the open hearth or Siemens- Martin process. In the Bessemer process the molten cast iron is poured into a converter (Fig. 152), 536 OUTLINES OF CHEMISTRY through whose perforated bottom compressed air is then blown into the metal ; thus carbon, silicon, and phosphorus are oxi- dized. Molten spiegeleisen of known carbon content is then added in proper quantity to produce steel of the exact carbon content desired. The converter is mounted so that it can be tilted to pour out or receive its contents. The slag, which con- sists of silicates of manganese, iron, calcium, and sulphide of iron, is first poured off, after which the molten steel is run into molds to form ingots. These are afterwards rolled into rails. FIG. 152. By means of a Bessemer converter about 20 tons of cast iron can be converted into steel in 20 to 30 minutes. The converter is made of iron and is usually lined with material similar to that of fire brick. In case the cast iron is rich in phosphorus, this siliceous lining is replaced by one of calcium and magnesium oxides obtained by calcining dolomite. This basic lining absorbs the phosphorus, yielding calcium and magnesium phosphates, which appear in the slag. This adaptation of the Bessemer process to the treatment of cast iron rich in phosphorus is know as the Thomas-Gilchrist process, and the slag produced by it is called the Thomas slag. The latter, being rich in phos- phates, is ground up and sold as a fertilizer. Thus, ores rich in IRON, NICKEL, AND COBALT 531 phosphorus, which were formerly discarded as unfit for Bes- semer steel, are very profitably turned to use. The open hearth or Siemens-Martin process consists essentially of heating the cast iron together with rusty scrap iron or other iron oxide, commonly hematite ore, on the hearth of a special type of reverberatory furnace, using gas as fuel. The oxida- tion of the carbon in the cast iron thus proceeds at the expense of the oxygen in the iron oxide. The process is continued till a sample taken from the material shows that the oxidation has proceeded far enough to produce the steel desired. The molten mass is then run off into molds. The process requires about 8 hours for its completion; but it yields an excellent, uniform steel, and utilizes the carbon in the cast iron for the reduction of ores. Hence it is more economical than the Bes- semer process, which it is rapidly displacing. Moreover, if the cast iron is rich in phosphorus, a hearth lined with the oxides of magnesium and calcium may be used to absorb the phosphorus in form of phosphates; with this modification the process is called the basic open hearth process. The United States produces about 35,000,000 tons of pig iron annually, a large portion of which is converted into steel. Nickel, manganese, chromium, silicon, tungsten, and molyb- denum are often added to steel in small proportions, thus form- ing alloys of mechanical properties that are desirable for certain special purposes. Properties of Iron. Pure iron maybe made by strongly ig- niting ferric oxide in a current of hydrogen. If heated thus to not higher than 450, the iron obtained contains hydrogen and is called pyrophoric iron, for it burns spontaneously on ex- posure to the air. Electrolytic iron, obtained by electrolyzing a ferrous sulphate solution, using a thick wrought iron anode and a thin iron foil as cathode, is also nearly pure; but it contains a few hundredths of one per cent of hydrogen, which renders it very hard and brittle. The melting point of pure iron has not been determined with certainty; it lies above that of wrought iron. The specific gravity of iron is 7.86. The metal is white, malleable, ductile, and fairly soft. It is attracted by a magnet, and becomes magnetic; but it loses its magnetism rapidly, while magnetized steel does not. Iron remains unchanged in dry air, while in moist air or in presence of air and salt solutions 538 OUTLINES OF CHEMISTRY it rusts, forming hydrated ferric oxide. This corrosion is hastened by local electrolytic action, and is generally guarded against by painting or tarring the exposed parts of the iron or steel. Dilute hydrochloric or sulphuric acid dissolves iron, forming ferrous chloride or sulphate and hydrogen, which in case of cast iron or steel is mixed with hydrocarbon gases and compounds of carbon, hydrogen, phosphorus, and sulphur that have a bad odor. When dipped into very concentrated nitric acid and then rinsed, iron no longer dissolves in nitric acid, nor does it pre- cipitate copper from solutions of its salts. The iron is said to be passive. It is thought by many that the phenomenon is due to a very thin, invisible coating of oxide on the iron, for when passive iron is scratched with a hard point, it at once dissolves rapidly in nitric acid. Other metals, like chromium and nickel, also exhibit the phenomena of the passive state. Oxides and Hydroxides of Iron. Ferrous oxide FeO is ob- tained as a black powder by heating ferric oxide in hydrogen or carbon monoxide to about 300. On exposure to the air it readily oxidizes further. Ferrous hydroxide Fe(OH) 2 is a white precipitate formed by adding caustic alkali to a solution of a ferrous salt. It readily oxidizes in the air, turning green and then brown. Ferric oxide Fe 2 O 3 is found in nature as hematite ore. It is found in large quantities in the Lake Superior region, and is the most important of the iron ores. It crystallizes in very dark red, hexagonal pyramids and prisms. When finely ground, ferric oxide is used as a pigment in paints under the name of red ocher or Venetian red. By ignition of ferrous sulphate or oxalate in the air, ferric oxide is obtained, which when finely ground is sold as a pigment or as rouge for polishing purposes. Ferric oxide is prepared in the laboratory by ignition of ferric hy- droxide, which is obtained by precipitating a solution of a ferric salt with caustic alkali. Ferric hydroxide Fe(OH) 3 is an amorphous, brown, flocculent precipitate to which alkali adheres very tenaciously. It dis- solves to some extent in a concentrated solution of ferric chloride, forming a very dark brown solution of a basic ferric chloride; from this the chlorine may be removed in the form of ferric chloride by dialysis, thus leaving a dark brown, tasteless, IRON, NICKEL, AND COBALT 539 colloidal solution of ferric "hydroxide or so-called dialyzed iron behind. Hydrated ferric oxides occur in nature. Thus we have li- monite 2 Fe 2 O 3 + 3 H 2 O or Fe 2 O 3 2 Fe(OH) 3 , also known as brown iron ore; pyrosiderite FeO'OH, and bog iron ore Fe 2 O(OH) 4 . These may all be considered as dehydration products of Fe(OH) 3 . Yellow ocher is an impure hydrated ferric oxide. It is used as a pigment in paints. Though the iron oxide pigments are not as bright in color as many others, yet they are valuable because they are permanent and cheap. Ferrous ferric oxide Fe 3 O 4 also known as magnetic iron oxide, is formed as the final product of continuous strong ignition of any oxide of iron in the air. In nature it occurs in black octahedra and dodecahedra as magnetite ore. It is often magnetic and is then called lodestone. The hammer black formed as a scale on iron when it is heated in the air is Fe 3 O 4 . Ferrous ferric oxide is isomorphous with spinel, and consequently probably is Fe(FeO 2 ) 2 . When iron turnings are fused with potassium nitrate, or when chlorine is conducted through a cold, concentrated solu- tion of caustic potash containing ferric hydroxide in suspen- sion, potassium ferrate K 2 FeO 4 is formed. It may be obtained in the form of dark red crystals, which are, however, unstable. The} r decompose, yielding oxygen, ferric hydroxide, and caustic potash. Barium ferrate BaFeO 4 is more stable. The ferrates are salts of ferric acid H 2 FeO 4 , which, like manganic acid, is not known in the free state. Ferrates are analogous to and isomorphous with sulphates and chromates. When caustic alkali solutions are electrolyzed, ferrates are formed at the anode if the latter consists of iron. Chlorides of Iron. Ferrous chloride FeCl 2 is obtained as a white mass by heating iron filings in a current of hydrochloric acid gas. It may also be obtained in the form of green, mono- clinic crystals FeCl 2 -4H 2 O, from aqueous solutions carefully kept from the oxygen of the air, for the salt is readily oxidized, thus : 6 FeCl 2 + 3 O = 4 FeCl 3 + Fe 2 O 3 . Ferric chloride FeCl 3 is obtained in the form of very dark green, lustrous, hexagonal crystals when iron is heated in a current of chlorine. The product may be sublimed. It is very 540 OUTLINES OF CHEMISTRY deliquescent. Ferric chloride may readily be formed by pass ing chlorine into a solution of ferrous chloride, by boiling the latter with aqua regia, or by dissolving ferric oxide or hydrox- ide in hydrochloric acid. On evaporation, a dark brown crys- talline mass, Fe 2 Cl 6 -12 H 2 O, is obtained. At higher tempera- tures different hydrates crystallize out and are in equilibrium with the saturated solution. This question has been carefully studied by H. W. Bakhuis Roozeboom, whose results are shown 500 400 300 200 1 100 Fe ,Cl e .7H Fe, C1 . F< 5HJL/ FlisCln 4H 2 q 60 40 20 20 Temperature in degrees C. FIG. 153. 40 60 80 in Fig. 153, which gives the solubility curves of ferric chlo- ride and indicates the ranges of temperature at wMch the dif- ferent hydrates are in equilibrium with the solutions (compare the case of magnesium chloride, Fig. 136). Ferric chloride is used in medicine. The salt is soluble in alcohol, ether, and many other liquids besides water. Ferrous bromide FeBr 2 and ferric bromide FeBr 8 are analo- gous to the corresponding chlorides. Ferrous iodide FeI 2 4 H 2 O consists of bluish green, mono- clinic crystals formed by heating iron filings and iodine to- gether under water. The salt is used in medicine in sirup of ferrous iodide. Ferric iodide is not known. IRON, NICKEL, AND COBALT 541 Sulphides and Sulphates of Iron. Ferrous sulphide FeS is formed by heating iron and sulphur together, or by adding ammonium sulphide to a solution of a salt of iron. If a ferric salt is used, a mixture of sulphur and ferrous sulphide is obtained, thus : 2 FeCl 3 + 3 (NH 4 ) 2 S = 6 NH 4 C1 -f 2 FeS + S, Ferrous sulphide is also formed by warming finely divided iron with sulphur in water. The fused sulphide forms a black, brittle, crystalline mass on cooling; whereas the precipitated sulphide is black and amorphous. Ferrous sulphide is soluble in dilute acids, hence it is not precipitated by hydrogen sulphide from acid solutions of salts of iron. Ferric sulphide Fe 2 S 3 is a greenish yellow mass, obtained by fusing iron and sulphur, or ferrous sulphide and sulphur, together in proper proportions. It is not formed by precipi- tating ferric salts with ammonium sulphide ; for the latter firsu reduces the ferric salt, and then precipitates ferrous sulphide. Iron disulphide FeS 2 occurs in nature as pyrite and marcasite. It may be made artificially by carefully heating iron and sul- phur together in proper proportions. Pyrite forms golden yellow crystals hav- ing a metallic luster. As pyrite, these crystals are cubes, octahedra, or pentago- nal dodecahedra (Fig. 154); and as mar- casite, they are orthorhombic. All sulphides of iron when roasted in the air finally yield sulphur dioxide and ferrous ferric oxide Fe 3 O 4 . On exposure to moist air the sulphides are gradually oxidized to sulphates. Ferrous sulphate FeSO 4 -7H 2 O, also called green vitriol or copperas, is formed by gently roasting pyrite so as to form ferrous sulphide, and then allowing the latter to oxidize to sulphate in moist air. From the mass the salt is readily extracted with water, and from the solutions green, monoclinic prisms, FeSO 4 -7H 2 O (Fig. 74), are obtained. These are isomorphous with other vitriols. The salt FeSO 4 5H 2 O is also known. It forms triclinic crystals that are isomorphous with CuSO 4 -5H 2 O (Fig. 75). Ferrous sulphate is used as a reducing agent, as a mordant, 542 OUTLINES OF CHEMISTRY as a disinfectant, and also in making ordinary writing-ink. The latter consists essentially of a solution of ferrous sulphate and extract of nutgalls, which contains tannin. Thus a ferrous tan- nate is formed, which on exposure to the air is oxidized to a ferric com'pound that is black and not readily soluble. Dextrine or gum arabic is generally added to ink to retard the precipita- tion of the ferric tannate as the ink stands in bottles. An antiseptic like carbolic acid or corrosive sublimate is often also introduced to prevent the growth of molds. Ferrous sulphate forms double salts with alkali sulphates. Of these ferrous ammonium sulphate (NH 4 ) 2 SO 4 -FeSO 4 -6H 2 O is of special importance. It is known as Mohr's salt and is used in chemical analysis. On exposure to the air this salt does not oxidize as readily as ferrous sulphate. Ferric sulphate Fe 2 (SO 4 ) 3 is formed by the oxidation of fer- rous sulphate, or by dissolving ferric hydroxide in sulphuric acid. In an anhydrous condition it is a white mass. This dissolves in water, slowly yielding a brown solution. With alkali sulphates it forms ferric alums; examples of these are, (N H 4 ) 2 SO 4 - Fe 2 (SO 4 ) 3 24 H 2 O and K 2 SO 4 . Fe 2 (SO 4 ) 3 24 H 2 O. Ferrous Carbonate FeCO 3 occurs in rhombohedral crystals in nature as siderite. It is formed as a white precipitate when alkali carbonates are added to solutions of ferrous salts. Fer- rous carbonate, like calcium carbonate, is soluble in water charged with carbon dioxide, thus forming ferrous bicarbonate Fe(HCO 3 ) 2 . On exposure to the air ferrous carbonate soon turns dark in color because of the formation of hydrated ferric oxide. Ferric carbonate is unknown. When treated with sodium carbonate, ferric salts yield ferric hydroxide. The reaction is similar to that of aluminum salts, thus : Fe 2 (S0 4 ) 3 + 3 Na 2 C0 3 + 3 H 2 O = 3 Na 2 SO 4 + 2 Fe(OH) 3 + 3 CO 2 . Al 2 (S0 4 ) 3 +3Na 2 C0 3 -|-3H 2 = 3Na 2 S0 4 +2Al(OH) 3 + 3C0 2 . Cyanides of Iron. The simple cyanides Fe(CN) 2 and Fe(CN) 3 are unknown, but many double cyanides of iron have been prepared. Potassium ferrocyanide K 4 Fe(CN) 6 , or 4KCN-Fe(CN) 2 , is made by fusing together potash, scrap iron, and animal refuse, like blood, hoofs, horns, scraps of hides, etc., as stated in Chapter XIV. On cooling and leach- ing out the cake with water, a yellow solution is obtained 2 L IRON, NICKEL, AND COBALT 543 which deposits beautiful lemon-yellow, monoclinic crystals K 4 Fe(CN) 6 - 3 H 2 O. These are also known as yellow prussiate of potash. They readily lose water, and on further heating they decompose, yielding potassium cyanide, as already stated. Potassium ferricyanide K 8 Fe(CN) e , or 3 KCN-Fe(CN) 3 , is formed by treating potassium ferrocyanide with chlorine, thus : 2 K 4 Fe(CN) 6 + Cl a = 2 KC1 + 2 K 3 Fe(CN) 6 . The salt consists of dark red, rhombic prisms which readily dissolve in about three parts of water, yielding a greenish brown solution. The compound is also known as red prussiate of potash. When treated with concentrated hydrochloric acid, a satu- rated solution of potassium ferrocyanide yields a white crystal- line precipitate which is H 4 Fe(CN) 6 , i.e. the free ferrocyanic acid. Similarly, ferricyanic acid H 3 Fe(GN) 6 may be obtained from K 3 Fe(CN) 6 . While the alkali salts of these acids are soluble, the salts they form with the heavy metals are insoluble, and may consequently be obtained by precipitation from solutions. Copper ferrocy- anide, for instance, is an insoluble ferrocyanide. It has already been described. When a ferric salt is added to a solution of potassium ferrocyanide, ferric ferrocyanide Fe'" 4 [Fe"(CN) 6 ] 3 is formed as an indigo-blue precipitate called Prussian blue: 4 FeCl 3 + 3 K 4 Fe(CN) 6 = 12 KC1 + Fe'" 4 [Fe"(CN) 6 ] 8 . If the ferric salt is added to a large excess of potassium ferro- cyanide, soluble Prussian blue K 2 Fe'" 2 [Fe' ' (CN) 6 ] a is formed. Caustic alkali decomposes Prussian blue : Fe'" 4 [Fe"(CN) 6 ] 8 + 12 KOH = 3 K 4 Fe(CN) 6 + 4 Fe(OH) 8 . Prussian blue is used as a pigment. When a ferrous salt is added to a solution of potassium fer- ricyanide, ferrous ferric cyanide Fe" 8 [Fe'"(CN) 6 ] a is precipi- tated: 3 FeS0 4 + 2 K 3 Fe(CN) 6 = 3 K 2 SO 4 + Fe" 3 [Fe'"(CN) 6 ] 2 . This precipitate is also indigo-blue in color. It is known aa Turnbull's blue. When treated with caustic alkali, it is decom- posed into ferrous hydroxide and potassium ferricyanide. 544 OUTLINES OF CHEMISTRY Blue Printing. When a ferric salt is added to a solution of potassium ferricyanide, a brown solution, but no precipitate, is formed. Paper treated with such a solution and dried in the dark is the sensitive paper used in blue printing. On exposing this paper to the light, the ferric salt is partially reduced to the ferrous state. When the paper is then washed with water, the insoluble Turnbull's blue formed remains, while the places pro- tected from the light appear white, because the original mix- ture on the paper is simply dissolved away. As a rule, ferric ammonium citrate and potassium ferricyanide are used in mak- ing blue print paper. Ammonium hydroxide or caustic alkalies dissolve both Turnbull's and Prussian blue, and these serve as inks to write white characters on blue prints. Other Compounds of Iron. Ferrous nitrate Fe(NO 3 ) 2 and fer- ric nitrate Fe(NO 3 ) 3 have also been prepared. The latter is deliquescent and yields basic nitrates on boiling with water. Ferrous phosphate Fe 3 (PO 4 ) 2 is an insoluble white precipi- tate, while ferric phosphate FePO 4 is a yellowish white powder, insoluble in water and acetic acid. Iron carbide Fe 3 C, also called cementite, occurs in cast iron and steel as so-called combined carbon. The silicides Fe 2 Si and FeSi have been obtained in crystalline form. They are hard and brittle. The phosphides Fe 3 P and Fe 2 P are known. Their presence in iron makes it very brittle. Ferric acetate Fe(C 2 H 3 O 2 ) 3 is unstable and readily hydro- lyzed, forming acetic acid and basic ferric acetates, which, being insoluble in acetic acid, enable the analytical chemist to pre- cipitate iron from an acid solution. This is of consequence in separating iron from manganese, for instance. Analytical Tests for Iron. = From solutions of both ferrous and ferric salts, ammonium sulphide precipitates black ferrous sulphide, which readily dissolves in acids. In solutions of ferrous salts, which are almost colorless, hydroxides, carbonates, and ferricyanides of the alkalies pro- duce characteristic precipitates that have already been de- scribed. In solutions of ferric salts, which are commonly brown, hydroxides, carbonates, and ferrocyanides of the alkalies also produce characteristic precipitates; 'while potassium sul- phocyanate KCNS forms a deep red coloration of soluble ferric IRON, NICKEL, AND COBALT 545 sulphocyanide Fe(CNS) 3 . In the borax bead, ferrous com- pounds yield a green, and ferric compounds a brown, coloration. The fact that ferric salts may readily be changed into ferrous salts by many reducing agents, like nascent hydrogen, stannous chloride, hydrogen sulphide, etc., is frequently used in analyt- ical chemistry. The change of ferrous salts to the ferric stat a by oxidizing agents like nitric acid, bichromates, or perman- ganates, is also often employed. Occurrence, Preparation, and Properties of Nickel. Nickel i found in meteoric iron. The chief ores of nickel are nicollite NiAs, gersdorffite, or nickel glance, NiAsS, and garnierite Mg 2 Ni 2 H 4 (SiO 4 ) 3 - 4 H 2 O. The chief localities are Ontario and New Caledonia. The commercial production of nickel is a complicated pro- cess which will not be described here. Pure nickel may be obtained by igniting the oxides or the oxalate in a current of hydrogen ; also by Goldschmidt's process, or by reducing the oxides with carbon. Nickel is a silver-white, lustrous metal, of specific gravity 8.9. It melts at about 1485. Nickel is malleable, ductile, and tenacious. It is but slightly altered on exposure to the air, and consequently it is frequently used in plating other metals. Nickel is but slowly attacked by sulphuric or hydro- chloric acid ; but nitric acid dissolves it readily. Like iron, nickel is magnetic, and exhibits the phenomena of passivity. Nickel is used in many alloys. German silver, or argentan, an alloy of brass with nickel, has already been mentioned. Man- ganine, an alloy of nickel with copper and manganese, is used for resistance wires. Nickel coins contain 75 per cent copper and only 25 per cent nickel, which fact demonstrates the great power of nickel to impart its color to alloys. About three per cent nickel added to steel produces a product of great strength; and much nickel is used in making nickel steel for armor plates. The Chinese have used nickel-copper alloys under the name paclcfong for many centuries. Through the work of Cronstedt and Bergmann, nickel was distinguished from other metals in 1751. Nickel coins came into use about fifty years ago. Nickel Oxides and Hydroxides. Nickelous hydroxide Ni(OH) 2 is an apple-green, amorphous precipitate, formed by adding 546 OUTLINES OF CHEMISTRY caustic alkali to a solution of a nickel salt. On ignition it yields a green powder, nickelous oxide NiO. Nickelic oxide Ni 2 O 3 is a black powder obtained by careful ignition of the nitrate. Nickelic hydroxide Ni(OH) 3 is formed when a solu- tion of a nickel salt is treated with an alkaline hypochlorite. It is a black precipitate. Salts of Nickel. In these nickel is always bivalent. The an- hydrous salts are yellow or brown, and the hydrous salts and the solutions are green. Nickelic oxide acts like a peroxide on treatment with acids ; there are no nickelic salts. Nickelous chloride NiCl 2 -6H 2 O forms green, monoclinic 3rystals that readily dissolve in water. Nickelous nitrate Ni(NO 3 ) 2 -6 H 2 O consists of green, deli- quescent monoclinic plates. Nickelous sulphate NiSO 4 7 H 2 O forms green, orthorhombic prisms that are isomorphous with the other vitriols. It may also be obtained as NiSO 4 6 H 2 O in form of tetragonal crystals. It is readily soluble in water. The crystals effloresce on ex- posure to the air. With ammonium sulphate, nickel sulphate forms nickel ammonium sulphate (NH 4 ) 2 SO 4 -NiSO 4 -6 H 2 O, which is isomorphous with Mohr's salt. Nickel ammonium sulphate is soluble in about 17 parts of water. It is used in nickel plating, in which process a thick nickel plate is used as anode, and the thoroughly cleaned object to be plated is the cathode. Nickelous cyanide Ni(CN) 2 is formed as a green precipitate by adding potassium cyanide to a solution of a nickel salt. In excess of potassium cyanide, the precipitate dissolves, forming Ni(CN) 2 2 KCN H 2 O, which may be obtained as reddish yellow monoclinic crystals. On boiling, the solution remains un- changed ; but on treatment with hypochlorites, black nickelic hydroxide is precipitated, which fact is used in separating nickel from cobalt. Nickelous sulphide NiS is formed as a black precipitate when alkali sulphides are added to solutions of nickel salts. The sulphide is slightly soluble in excess of alkali sulphides, yielding a dark brown solution from which, by addition of acetic acid, the sulphide is again precipitated. Nickelous sulphide dissolves but slightly in dilute hydrochloric acid. Nickel carbonyl Ni(CO) 4 is a colorless liquid boiling at 43 IRON, NICKEL, AND COBALT 547 and congealing at 25. It is formed by passing carbon mo- noxide over nickel obtained by reducing the oxide or oxalate by ignition in hydrogen. The vapors of nickel carbonyl are poisonous. The substance burns, depositing nickelous oxide and nickel. This compound is similar to iron carbonyl Fe(CO) 5 , which is formed by passing carbon monoxide over finely divided iron under a pressure of about six atmospheres at 40 to 80. An analogous compound of cobalt has not been made. Occurrence, Preparation, and Properties of Cobalt. Cobalt is found in meteorites and in smaltite CoAs 2 and cobaltite CoAsS, in which it is commonly associated with nickel, iron, and man- ganese. The principal localities are Canada, Sweden, Bohemia, Germany, and the Urals. The metal is obtained by igniting the oxides or oxalate in a current of hydrogen. Cobaltous oxide is also reduced by mix- ing it with starch or flour and making small cubes of the paste formed, which are then embedded in pulverized carbon and highly ignited in a crucible. Thus the oxide is reduced, and the metal is obtained in form of cubes. Metallic nickel may be obtained similarly. Cobaltous oxide may be reduced by the Goldschmidt process. Cobalt is a silver-white, malleable, tenacious metal of specific gravity 8.5. It melts at about 1500. On exposure to the air, it soon acquires a reddish hue. Red-hot cobalt decomposes steam. The metal is magnetic. It dissolves in nitric acid, but other dilute acids attack it but slowly. The atomic weight of cobalt is 58.97 and its valence is either two or three. The cobaltous compounds are the more common. The metal itself is not used in the arts, but its oxides are used to color glass and porcelain blue. The fact that silicates are thus colored by cobalt ores has been known since ancient times, but the metal was not obtained till 1735, when the Swedish chemist Brandt prepared it. The word cobalt means goblin. This name was given to the ores of cobalt by the miners ; for though these ores have a bright metallic luster suggesting a metallic content, it was not till relatively recent times that the metal was ob- tained from them. Oxides and Hydroxides of Cobalt. Cobaltous hydroxide Co(OH) 2 is a rose-red precipitate formed by adding caustic alkali to a solution of a cobaltous salt, and then boiling to 548 OUTLINES OF CHEMISTRY decompose the blue basic salt that is first precipitated. On igniting cobaltous hydroxide or carbonate out of contact with the air, cobaltous oxide CoO is obtained as a greenish powder. This is also formed by passing steam over red-hot cobalt. On heating cobaltous oxide in the air, cobaltous cobaltic oxide Co 3 O 4 , which is analogous to magnetite, is formed as a black powder. This is produced as the final product by igniting any oxide of cobalt in the air. It also results by calcining the nitrate. But if the latter is gently heated, cobalt sesquioxide Co 2 O 3 is first obtained as a dark brown powder, which yields oxygen and Co 3 O 4 on further ignition. Cobaltic hydroxide Co(OH) 3 is a dark powder formed by treating a solution of a cobaltous salt with hypochlorites. Cobalt dioxide CoO 2 is also known. Other Cobalt Compounds. The hydrous cobaltous salts are red, which is also true, as a rule, of their solutions at ordinary temperatures. The anhydrous cobaltous salts are generally blue, as are many of their solutions at higher temperatures. But few simple cobaltic salts are known. Cobaltous chloride CoCl 2 is a blue crystalline powder obtained by heating cobalt in^chlorine. From aqueous solutions, it may be obtained in form of deep red, monoclinic prisms of the com- position CoCl 2 6 H 2 O. The salt is also formed by dissolving cobaltous hydroxide or carbonate in hydrochloric acid. When the hydrous cobaltous chloride is heated, it turns blue, becom- ing red again as it absorbs water after cooling. So by writing on paper with a solution of cobaltous chloride (so-called sym- pathetic ink) the lines, which are at first invisible, become blue on gently heating the paper. On exposure to the air at ordi- nary temperatures, the salt again takes on water and the writ- ing fades. Cobaltous nitrate Co(NO 3 ) 2 forms red, deliquescent, mono- clinic prisms. Cobaltous sulphate CoSO 4 -7H 2 O is isomorphous with the other vitriols. It loses water readily on heating. The an- hydrous salt is red and is not easily decomposed, even on strong ignition. Cobaltic sulphate Co 2 (SO 4 ) 3 -18 H 2 O has been obtained at the anode by electrolysis of cobaltous sulphate. Cobalt silicates result when glass or silica and potash are fused with an oxide or other compound of cobalt. The blue IRON, NICKEL, AND COBALT 549 glass that results is pulverized and used as a pigment under the name smalt. It serves well in decorating porcelain and glass, where it is afterwards fired in ; but it is not satisfactory in oil or water paints. Cobaltous sulphide CoS is obtained as a black precipitate by adding alkali sulphides to solutions of cobalt salts. Like the sulphide of nickel, cobaltous sulphide is insoluble in dilute hydrochloric acid. The sulphides Co 2 S 3 , Co 3 S 4 , and CoS 2 occur in nature. Cobaltous carbonate CoCO 3 is a bright red powder. By add- ing alkali carbonates to solutions of cobaltous salts, red pre- cipitates of basic carbonates form, which decompose and turn blue on boiling. Cobalt silicides Co 2 Si and CoSi have been made in the electric furnace. They are very hard crystalline substances. Cobaltous cyanide Co(CN) 2 is formed as a reddish precipitate by adding potassium cyanide to a solution of a cobaltous salt. Cobaltous cyanide is soluble in an excess of potassium cyanide, forming potassium cobaltous cyanide K 4 Co(CN) 6 , which upon oxidation yields potassium cobaltic cyanide K 3 Co(CN) 6 . These compounds are similar to those of the corresponding double cyanides of iron. Nickel does not form analogous double cyanides. Potassium cobaltic nitrite 3 Co(NO 2 ) 3 -6KNO 2 -3 H 2 O is formed as a yellow precipitate when potassium nitrite is added to a solution of a cobaltous salt acidified with acetic acid. The precipitate separates out slowly. It is called cobalt-yellow. The reaction is used to separate cobalt from nickel. The lat- ter does not form an analogous salt. Cobalt amines are complex compounds which cobaltic salts form with ammonia. These compounds are made by treating solutions of cobaltous chloride with ammonia in sufficient amount to redissolve the cobaltous hydroxide that first forms. On standing in the air, these brown solutions absorb oxygen, thus becoming red. When hydrochloric acid is added to the red solution, brick-red crystals of roseo cobaltic chloride CoCl 3 -5 NH 3 - H 2 O are deposited. On boiling, the solution yields purpureo cobaltic chloride CoCl 3 (NH 3 ) 5 , which forms a bright red crystalline powder, or brownish luteo-cobaltic chlo- ride CoClg-GNH^ if much ammonium chloride is present in the strongly acid solution. Other cobaltous salts form similar 550 OUTLINES OF CHEMISTRY complicated compounds on treatment with ammonia. The three examples given must suffice here to indicate their general char- acter, though the compounds are very interesting and have in recent years received careful study in connection with the sub- ject of valence, particularly by A. Werner and S. M. Jorgensen. Analytical Tests for Nickel and Cobalt. When heated in the borax bead, cobalt compounds always give a blue coloration. Nickel compounds yield a brown bead in the oxidizing flame, and a milky bead in the reducing flame, for by the latter metallic nickel is formed. The reactions of salts of nickel and cobalt with alkali hydroxides, sulphides, hypochlorites, cyanides, and nitrites, as above explained, are used in the detection and separation of nickel and cobalt. REVIEW QUESTIONS 1. Why is it that iron is the most useful of all of the metals? 2. What other metals are closely related to iron and are grouped with it in the periodic system? Why are these metals in the last column of the periodic system? 3. Write the formulas of the chlorides, sulphates, nitrates, and oxides of iron. Compare these with the formulas of the corresponding com- pounds of cobalt, nickel, manganese, chromium, aluminum, zinc and magnesium, by grouping the compounds of analogous composition. 4. What are the most important ores of iron ? Write their formulas and the equations showing how iron is obtained from each of these ores. 5. Discuss the occurrence of iron: (1) in soils, (2) in plants and animals. 6. Why is limestone used in a blast furnace ? Equation. 7. What are the main chemical differences between cast iron, wrought iron, and steel? 8. Why are phosphorus and sulphur especially objectionable in iron? 9. What occurs in the process known as the tempering of steel ? 10. What is the chemistry of the use of borax or sand in welding? Equations. 11. Compare the essential features of the Bessemer process and the open hearth process of making steel. 12. Explain how phosphate fertilizer may be produced in connection with the production of steel. 13. Explain how blast furnace slag is employed in the Portland cement industry. 14. What is the chemical nature of the following: Venetian red, rouge, red ochre, yellow ochre, passive iron, Fool's gold, copperas, Mohr's salt, Prussian blue, cementite, electrolytic iron. IRON, NICKEL, AND COBALT 551 15. Explain the essentials of the process of blue printing. Equations. How write white characters on a blue print? Equation. 16. Mention four different ways of converting 10 grams of ferrous chloride into ferric chloride. How much of the latter would be formed? Write the equation in each case. 17. How demonstrate that a one per cent solution of ferrous sulphate contains iron? Equations. How show that a one per cent solution of ferric chloride contains iron? Equations. 18. What is the composition of a nickel coin? Of German silver? How prove that these contain nickel? Equations. 19. Why is nickel used : (1) in armor plates, (2) in plating iron, also copper and brass articles ? How is the process of nickel plating conducted ? 20. Compare the physical properties of nickel and cobalt. 21. What use is made: (1) of metallic cobalt, (2) of cobalt com- pounds ? 22. Describe the chemical characteristics of cobaltous chloride. 23. What is sympathetic ink? Upon what fact does its use depend? 24. How test a compound for the presence of cobalt ? 25. What is smalt, and what is it used for? 26. Given a solution containing a mixture of the chlorides of nickel and cobalt, how proceed to separate the nickel from the cobalt ? Reaction. CHAPTER XXXI THE METALS OF THE PLATINUM FAMILY THE metals of the platinum group and their symools and atomic weights are: ruthenium (Ru 101.7), rhodium (Rh 102.9), palladium (Pd 106.7), osmium (Os 190.9), iridium (Ir 193.1), and platinum (Pt 195.2). These metals always occur together in nature in the form of alloys con- sisting of small corroded metallic bits or nuggets found in alluvial sands, chiefly in the Urals, California, Brazil, Australia, Borneo, and Sumatra. These alloys contain from 50 to 85 per cent platinum, several per cent iridium, and generally less than two per cent of each of the other members of the group. Gold, copper, iron, and other metals are also usually present. By careful washing of the sands, the particles of platinum ore are obtained as in the case of placer mining of gold. Though platinum ore usually occurs in the form of mere grains, larger nuggets weighing up to about eight kilograms have been found in rare instances. It will be observed that ruthenium, rhodium, and palladium have approximately the same atomic weight; their specific gravities are also nearly alike, being 12.3, 12.1, and 11.5, re- spectively. Again, the second three metals, osmium, iridium, and platinum, have nearly the same atomic weight, which is about 90 units higher than that of the first three members. Osmium, iridium, and platinum also have about the same spe- cific gravities, namely, 22.5, 22.4, and 21.5, respectively. It will be recalled that the atomic weights of iron, nickel, and cobalt are also nearly the same, and that these metals have specific gravities that differ but slightly from one another. Extraction of Platinum from the Ores. The ores are freed from sand and other adhering impurities, and then treated with dilute aqua regia, which dissolves the gold, copper, and iron that may be present, leaving the platinum metals behind. On adding concentrated aqua regia to this residue, platinum, pal- 652 THE METALS OF THE PLATINUM FAMILY 553 ladium, ruthenium, rhodium, and some of the iridium are dissolved. The residue left consists of osmium and iridium together with some ruthenium. From the solution, the plati- num and iridium are precipitated by means of ammonium chlo- ride as (NH 4 ) 2 PtCl 6 and (NH 4 ) 2 Ir01 6 . There is usually but little iridium present in this precipitate. On ignition of the latter, metallic platinum, alloyed with a little iridium, is ob- tained as a spongy mass. As platinum containing minor amounts of iridium is excellent for crucibles, dishes, and other purposes, being stronger mechanically and more resistant to- wards reagents than pure platinum, this sponge is commonly used without further attempts at purification. From the filtrate containing palladium, rhodium, and ruthenium, these metals are precipitated by means of iron and then separated by special processes which will not be described here. Ruthenium (Ru 101.7) was discovered by Glaus in 1845. The element is named after the Ruthenians, a little Russian people in whose country the metal is found. Ruthenium occurs in the alloy of osmium and iridium, osmiridium, remaining as a residue when platinum ores are treated with concentrated aqua regia. Ruthenium is commonly prepared from this alloy. It is a brittle metal of steel-gray color. It melts above 2000. When heated in the air, ruthenium oxidizes. The oxides RuO, Ru 2 O 3 , RuO 2 , and RuO 4 are known. They are black powders. Ruthenium is not attacked by acids ; even aqua regia acts but slowly on it, forming the trichloride RuC1 3 . The dichloride RuCl 2 is known, but the tetrachloride RuCl 4 exists only in solutions. On fusing ruthenium with caustic potash and salt- peter, potassium ruthenate K 2 RuO 4 , readily soluble in water, is obtained, which fact is used in separating ruthenium from osmiridium. Rhodium (Rh 102.9) was discovered by Wollaston in 1803. It receives its name from the red color of its chloride. The metal is silver- white, ductile, and malleable. It melts at about 2000, and when pure it is not attacked by acids, not even by aqua regia, which, however, dissolves rhodium alloyed with platinum. By fusing rhodium with caustic potash and saltpeter and extracting the mass with nitric acid, oxides of rhodium are obtained. The oxides of rhodium are RhO, Rh 2 O 3 , and RhO 2 ; the hydroxides Rh(OH) 3 and Rh(OH) 4 are also known. The 554 OUTLINES OF CHEMISTRY red chloride RC1 3 , formed by the action of chlorine on the metal, is the only one isolated. It is insoluble in water and acids. It forms soluble double salts with alkali chlorides. Rhodium chloride forms RhCl 3 -(NH 3 ) 5 and other compounds analogous to cobalt amines. The double cyanide K 8 Rh(CN) 6 , which is analogous to K 3 Fe(CN) 6 , is also known. Palladium (Pd lOd.7) was also discovered by Wollaston in 1803. It is named from Pallas, an asteroid discovered in 1802. While always present in platinum ores, palladium also occurs in alloys with gold in Brazil. Silver generally contains traces of palladium. It is a silver-white, ductile, malleable metal, which melts at about 1535. Palladium absorbs over 300 times its volume of hydrogen under ordinary conditions, at 100 about twice as much hydrogen is absorbed. On electrolyzing water, using a palladium cathode, the latter absorbs nearly 1000 times its volume of hydrogen. The compound formed, palladium-hydrogen, is lighter than the metal ; its composition varies with the conditions under which it is prepared. It gives off its hydrogen completely when heated to redness, though evolution of the gas begins even at 100. The hydrogen liberated effects reductions like nascent hydrogen. Palladium is soluble in nitric acid, in aqua regia, and in hot sulphuric acid. The oxides are PdO and PdO 2 ; both are black. The dioxide yields the chloride PdCl 2 and chlorine, when treated with hydrochloric acid, PdCl 4 being unknown. The following typical salts are also well known : Pd(NO 3 ) 2 ; PdI 2 ; PdSO 4 ; K a PdCl 6 ; PdCl 2 (NH 8 ) 2 . Osmium (Os 190.9) was discovered in 1804 by Tennant. It is found in platinum ores, in osmiridium and iridosmium, two alloys of osmium and iridium. Osmium is a steel-gray, hard, brittle metal, which has the highest specific gravity of all known substances, 22.5. It melts in the electric furnace above 2500. Osmium is oxidized to OsO 4 by aqua regia, nitric acid, or by ignition in the air. Osmium tetroxide forms white, monoclinic crystals that read ily melt, and then boil at about 100, emitting a very pungent vapor, from whose unbearable odor osmium gets its name. Osmium tetroxide is used in hardening and staining histological specimens, being commonly called osmic acid or perosmic acid. THE METALS OF THE PLATINUM FAMILY 555 Its use depends upon the fact that in contact with organic tissues, especially with fats, it is reduced, forming finely divided metallic osmium. The readiness with which osmium tetroxide is formed and its volatility serve in separating os- mium from iridium and ruthenium, which also occur in osmi- ridium. The chlorides OsCl 2 , Os 2 Cl 6 , and OsCl 4 are known, as are also osmates like K 2 OsO 4 -2H 2 O. The oxides are OsO, Os 2 O 8 , OsO 2 and OsO 4 . Iridium (Ir 193.1) was also discovered by Tennant in 1804. It is a grayish white, rather brittle metal which melts at about 2200. It is found in ores of platinum and in osmiridium. Its separation from osmium has been mentioned in connection with the latter. From platinum it may be separated by using the fact that the compound (NH 4 ) 2 IrCl 6 is readily soluble in water. Aqua regia attacks pure iridium but slowly. The oxides are IrO, Ir 2 O 3 , and IrO 2 . The hydroxides Ir(OH) 8 and Ir(OH) 4 and the chlorides IrCl 2 , IrCl 3 , and Ird 4 are also known. Iridium gets its name from the fact that it forms compounds of various colors. Platinum (Pt 195.2) is a silver-white, tenacious, malleable, and ductile metal of specific gravity 21.47. The name plati- num comes from platina, which is a diminutive of the Spanish plata, meaning silver. Platinum melts in the oxyhydrogen flame at about 1777. At white heat it may be welded with- out the aid of a flux, for its surface is not covered with oxide. Even at high temperatures, platinum does not oxidize in the air. Aqua regia dissolves platinum, but the latter is not attacked by hydrofluoric, hydrochloric, sulphuric, or nitric acid. Molten caustic alkalies and alkali nitrates, cyanides, or sulphides attack platinum, and it readily forms alloys with most of the heavy metals and also with silicon, boron, phosphorus, ar- senic, and antimony. These facts are of importance, for the substances mentioned, and also such compounds as may yield these substances on ignition, must not be heated in platinum dishes. Platinum should also not be heated in a reducing flame, for thus the metal takes up carbon and becomes brittle. Platinum sponge is formed by ignition of ammonium platinic chloride; while platinum black is produced by electrolysis of platinic chloride solutions, or by adding finely divided magne- sium, iron, or zinc to the latter. In finely divided form, platinum 556 OUTLINES OF CHEMISTRY absorbs about 300 times its volume of hydrogen or about one third that volume of oxygen. The gases are said to be occluded in the metal. On ignition they escape. The catalytic oxidations and reductions effected by finely divided platinum depend upon the fact that the latter absorbs the gases mentioned, which then act far more vigorously than when in the ordinary state. Thus Dobereiner's lamp (Fig. 155) depends upon the fact that when a jet of hydrogen is directed against a platinum sponge the latter is heated to redness and so lights the. jet. The decomposition of hydrogen peroxide and the synthesis of SO 3 from SO 2 and O 2 are further typical instances of the catalytic action FlG 155 of platinum black. Besides being used for crucibles, dishes, and other utensils for chemical operations, much plati- num is required for making various electrical connections. Platinum salts are used in photography. There are two series of platinum compounds. In platinous compounds the metal is bivalent, while in platinic compounds it is quadrivalent. Platinic chloride PtCl 4 is formed by the action of chlorine on platinum at high temperatures, or by ignition of chlorplatinic acid H 2 PtCl 6 -6 H 2 O, which results when platinum is dissolved in aqua regia. With the exception of sodium chlorplatinate Na 2 PtCl 6 the alkali salts of chlor- platinic acid are insoluble in alcohol and sparingly soluble in water, hence the use of chlorplatinic acid in separating sodium from potassium in analysis. Potassium chlorplatinate K 2 PtCl 6 , or PtCl 4 .2KCl, forms a golden yellow, crystalline precipitate. The crystals are small octahedra ; they are isomorphous with ammonium chlorplatinate (NH 4 ) 2 PtCl 6 . When potassium chlor- platinate is ignited, potassium chloride and platinum remain ; whereas when ammonium chlorplatinate is heated, the residue consists simply of spongy platinum. Either of these salts, which are also called potassium platinic chloride and ammo- nium platinic chloride, is readily formed by adding the alkali chloride to a solution of chlorplatinic acid or platinic chloride. THE METALS OF THE PLATINUM FAMILY 557 On heating chlorplatinic acid to 300, or on passing chlorine over platinum sponge at about 245, platinous chloride PtCl 2 is obtained as a grayish green, insoluble powder, which on igni- tion leaves platinum as a residue. Potassium platinous chloride K 2 PtCl 4 is formed by reducing potassium platinic chloride with cuprous chloride. Platinous hydroxide Pt(OH) 2 is formed by treating platinous chloride with caustic alkali. On ignition of platinous hydroxide, platinous oxide PtO is formed, and finally platinum. Platinic hydroxide Pt(OH) 4 results when platinic chloride is treated with caustic alkali. By careful ignition of this hydroxide, platinic oxide may be obtained, which on strong ignition yields platinum. When platinic hydroxide is treated with an excess of caustic alkali, platinates are formed, thus : Pt(OH) 4 + 4 NaOH = Na 4 PtO 4 + 4 H 2 O. Platinates also result when platinum is placed in molten caustic alkalies. Platinous sulphide PtS and platinic sulphide PtS 2 are black precipitates, insoluble in acids, formed by adding hydrogen sulphide to solutions of platinous and platinic compounds, respectively. These sulphides are soluble in aqua regia. They also dissolve in alkali sulphides, with which they form sulpho-salts. Analytical Tests for Platinum. The sulphides described are characteristic. On heating any platinum compound with soda on charcoal, spongy platinum is obtained. This is soluble in aqua regia, and from the solution potassium chlorplatinate may be precipitated by adding potassium chloride. REVIEW QUESTIONS 1. Name the metals of the platinum family and state how they may be divided into two distinct groups. 2. Besides these metals, what other noble metals are there? Why are these not grouped with the platinum metals in the periodic system? 3. How do the platinum metals occur in nature, and where? 4. Outline how platinum is extracted from its ores. 5. What physical and chemical properties make platinum an espe- cially valuable metal: (1) in the chemical laboratory, (2) in the arts? 6. Write the formulas of the following compounds: platinic chlo- ride, platinous chloride, potassium platinic chloride, osmic acid. 558 OUTLINES OF CHEMISTRY 7. Explain what property of palladium makes this metal of special value in gas analysis. 8. What is so-called hard platinum? What use is made of it? 9. What use is made of osmic acid in histology? 10. Explain the action of platinum sponge: (1) in lighting a jet of hydrogen or illuminating gas, (2) in decomposing hydrogen peroxide, (3) in making sulphuric acid. 11. How is platinic chloride used in the detection and estimation of potassium? Equation. 12. Solutions of platinic chloride and ferric chloride are both yellow in color. In a solution containing a mixture of these salts how demon- strate the presence of both platinum and iron? Give three ways of doing this, writing the appropriate equations. INDEX AJbietic acid, 489 Absorption spectrum, 387, 388 of blood, 389 Acetates, 256 Acetic acid, 254, 255 glacial, 256 Acetone, 267 Acetylene, 247 Acheson graphite, 223 Acid, definition of, 128, 132, 452 Acidimetry, 135 Acids, 127 anhydrides of, 129 basicity of, 133 strength of, 144 Acker process, 376 Actinium, 412 Adsorption, 19, 225 Affinity, chemical, 11, 141, 142 Agate, 308 Air, 150 alkaline, 158 ammonia in, 152 bacteria and microbes in, 151, 153 carbon dioxide in, 152 composition of, 150, 151 fixed, 146, 234 humidity of, 38 liquid, 154 mephitic, 146 moisture in, 151 nature of, 150 phlogisticated, 146 Alabaster, 400 Albite, 310, 490 Albuminoids, 276 Albumins, 276 Alchemists, 203 Alcohols, 249, 250 Aldebaranium, 495 Aldehydes, 252, 253 Algarotus, 343 Algin, 118 Alkalies, 362 Alkali metals, 362, 384 Alkalimetry, 135 Alkaline earth metals, 394 detection of, 410 Alkaloids, 275 Allotropism (see Allotropy) Allotropy, 93, 217 Alloys of copper, 463 of manganese, 524 Alumina, 485 Aluminates, 486 Aluminum, 482 acetate,. 489 alloys of, 484 bromide, 488 bronze, 484 chloride, 488 fluoride, 488 hydroxide, 485 iodide, 488 oxide, 485 production of, 482 properties of, 482 silicates, 490 sulphate, 488 sulphide, 488 tests for, 492 uses of, 484 Alums, 489 Alunite, 489 Amalgamation process, 469, 476 Amalgams, 423 for filling teeth, 500 Amethyst, 292 Amines, 274 Ammonia, 158 action of, on metals, 163, 164 combustion of, in oxygen, 162 composition of, 160, 161 concentrated, 162 liquid, 164 of crystallization, 164 preparation of, 158 properties of, 160 soda process, 379 water, 162 Ammonium, 162 alum, 489 amalgam, 424 bichromate, 516 bisulphate, 390 bromide, 390 carbamate, 392 carbonate, 392 carbonate, acid, 392 559 560 INDEX Ammonium Continued chloride, 390 chlorplatinate, 556 hydrazoate, 167 hydrosulphide, 390 iodide, 390 molybdate, 518 nitrate, 391 nitrite, 391 persulphate, 390 phosphomolybdate, 518 platinic chloride, 392 salts, 163, 389 salts, detection of, 165, 391 sulphantimonate, 346 sulphantimonite, 346 sulpharsenate, 340 sulpharsenite, 340 sulphate, 390 sulphide, 390 tartrate, acid, 392 Amorphous substances, 47 Amyl acetate, 264 Anaxagoras, 70 Andalusite, 490 Anglesite, 504 Anhydrite, 400 Aniline, 274 dyes, 274 hydrochloride, 274 Anions, 447 Anode, 447 Anorthite, 490 Anthracite, 226 Antichlorine, 382 Antimonic acid, 345 Antimonious acid, 344 Antimonium, Triumphal chariot of, 345 Antimony, 340 cinnabar, 347 compounds of, with halogens, 342 compounds of, with sulphur, 346 nitrate, 345 oxides and oxy-acids, 344 oxychlorides, 343 pentachloride, 343 pentafluoride, 344 pentasulphide, 346 pentiodide, 344 pentoxide, 346 sulphate, 345 tetroxide, 345 tri bromide, 344 trichloride, 342 trifluoride, 343 triiodide, 344 trioxide, 344 trisulphide, 346 Antimonyl group, 344 nitrate, 345 sulphate, 345 Apatite, 321 Aqua fortis, 170 regia, 174 Arabite, 250 Aragonite, 396 Argentan, 545 Argentic nitrate, 473 Argentite, 468 Argentum, 471 Argol, 262 Argon, 154, 155 Argyrodite, 498 Aromatic series, 248 Arrhenius, 450, 451 Arrhenius's theory, 450, 453, 454 Arsenic, 334 acid, 338 compounds of, with halogens, 337 disulphide, 339 iodides of, 337 oxides and oxy-acids of, 337 pentasulphide, 340 pentoxide, 339 tribromide, 337 trichloride, 337 trifluoride, 337 trioxide, 337 trisulphide, 339 white, 337 Arsenious acid, 338 Arsenites, 338 Arseniureted hydrogen, 335 Arsenolite, 334 Arsine, 335 Asbestus, 415 Assaying, 477 Association in solution, 441 Asymmetric carbon atom, 259 Asymmetric system, 194 Atmosphere (see Air) Atomic theory of matter, 66, 69 volumes, relation to atomic weights, 359 weights, determination of, 76 weights, choosing of, 81, 82 weights, table of, 83 Atoms, 67, 79 Atropine, 275 Aurates, 479 Auric chloride, 478 compounds, 478 oxide, 479 sulphide, 479 Aurous chloride, 478 compounds, 478 cyanide, 479 oxide, 479 sulphide, 479 Avogadro, 72, 76, 79 Avogadro's hypothesis, 72, 74, 161 Azoimide, 167 Azote, 146 Azurite, 460, 468 INDEX 561 B Babbitt metal, 505 Baeyer, 99 Balard, 110 Band spectrum, 387 Barite, 407 Barium, 407 bromide, 408 carbonate, 409 chloride, 408 chromate, 517 compounds of, 408 dioxide, 408 ferrate, 539 fluoride, 408 hydroxide, 408 iodide, 408 nitrate, 408 oxide, 408 sulphate, 409 sulphide, 401, 409 tellurate, 216 Barley sugar, 269 Bases, 129, 132, 134, 144, 452 Basic open hearth process, 537 Bauxite, 482, 485 Becquerel, 410 Becquerel rays, 410 Beer, 250 Begasse, 270 Bengal lights, 407, 409 Benzene, 248 Benzine, 245 Benzoic acid, 256 Bergmann, 545 Berthelot, 294, 303 Berthollet, 55 Beryl, 414 Beryllium, 414 Berzelius, 42, 63, 64, 65, 72, 79, 82, 132, 213, 234, 306, 315, 446, 495 Bessemer converter, 461, 536 process, 536 Bicarbonates, 230 Bichromates, 515 uses of, 516 Binary compound, 78 Bioses, 269 Birkelund and Eyde process, 172 Bismuth, 347 bromide, 348 dichloride, 348 dioxide, 349 disulphide, 350 fluoride, 348 halogen compounds of, 348 iodide, 348 nitrate, 349 oxides of, 348 oxy bromide, 348 oxyfluoride, 348 oxy iodide, 348 oxynitrate, 349 pentoxide, 349 salts of oxy-acids, 349 subnitrate, 349 tetroxide, 349 trioxide, 348 trisulphide, 350 Bismuthyl sulphate, 349 Bitter almonds, oil of, 253 Black, Joseph, 234 Black ash, 378 Black-jack, 417, 419 Blast flame, 286 Blast furnace, 532 Blast furnace slag, 532 Bleach, 402 Bleaching powder, 103, 104, 402 Blood charcoal, 225 Blowpipe flame, 286 Blue cup battery, 456 Blue printing, 544 Boiling points of solutions, 439 Bonds, 85, 86, 247, 248 Bone black, 225 Borax, 318 Borax glass, 318 Bordeaux mixture, 467 Boric acid, 316 Boric anhydride, 318 Boron, 316 carbide, 319 chloride, 319 fluoride, 319 hydride, 318 nitride, 319 sulphide, 319 Bort, 221 Bottger, 325 Bouronite, 504 Boussingault, 151 Boyle, 321 Brandt, 321, 547 Brandy, 250 Brass, 419, 463 Bricks, 491 Erin's process, 27 Britannia metal, 341, 500 British thermal unit, 291 Bromates, 115 Bromic acid, 114, 115 Bromine, 110, 111 oxy-acids of, 114 uses of, 116 water, 111 Bromoform, 249 Bronzes, 464, 500 Brucine, 275 Bruyn, Lobry de, 166 Bullets, 505 Bunsen, 372, 385, 388, 394, 415 Bunsen burner, 285 562 INDEX Burnt alum, 489 Bussy, 415 Butane, 244 Butter, 266 fat, 266, 434 of antimony, 343 of tin, 501 Butyric acid, 256 C Cadaverine, 276 Cadmium, 421 bromide, 422 chloride, 421 compounds of, 421 hydroxide, 421 iodide, 422 nitrate, 422 oxide, 421 sulphate, 422 sulphide, 422 Caesium, 372 Calamine, 417 Calaverite, 476 Calcite, 396 Calcium, 395 aluminate, 487 bicarbonate, 230, 396 bromide, 402 carbide, 403 carbonate, 396 chloride, 401 chromate, 517 cyanamide, 403 fluoride, 401 hydride, 396 hydroxide, 397 iodide, 402 metaphosphate, 322 nitride, 396 oxalate, 257 oxide, 397 phosphate, 402 phosphide, 403 silicate, 403 silicide, 403 sucrate, 270 sulphate, 400 sulphide, 401 sulphite, 401 Calomel, 425 Calorie, 291 Calorimeters, 291 Cane sugar, 269 Cannizzaro, 79, 80, 81 Caoutchouc membranes, 445 Caramel, 269 Carats, 478 Carbohydrates, 267 Carbolic acid, 252 Carbon, 221 allotropic forms of, 221 amorphous, 224 atom, properties of, 228 atomic weight of, 77 chemical behavior of, 226 cycle of, 233 pyrophoric, 225 Carbonado, 221 Carbonates, 230 Carbon bisulphide, 238 Carbon bisulphide furnace, 239 Carbon dioxide, 127, 229 early work on, 234 formula of, 78, 227 physiological effects of, 232 properties of, 231 relations of, to life, 233 solid, 232 uses of, 232 Carbonic acid, 127, 229 Carbon monoxide, 234 absorption of, 237, 465 formula of, 78, 236 physiological effects of, 238 properties of, 236 Carbon oxysulphide, 239 Carbon tetrachloride, 249 Carbonyl chloride, 237 group, 267 iron, 237 nickel, 237 Carborundum, 315 Carboxyl group, 254 Carlisle, 446 Carnallite, 363, 415 Carnelian, 307 Casciorolo, 409 Cassiopeium, 495 Cassiterite, 498, 502 Cast iron, gray, white, 534 Catalytic action of platinum, 556 Catalytic agents, 201 Cathode, 447 Cations, 447 Caustic soda, 376 Cavendish, 13, 35, 37, 146, 170 Celestite, 406 Celluloid, 274 Cellulose, 273 nitrates of, 273 solvent for, 463 Cement, 398 Cementite, 544 Ceric hydroxide, 495 Cerite, 494, 495 Cerium, 315, 495 Cerussite, 504 Chalcedony, 307 Chalcocite, 460 Chalcopyrite, 460 Chalk, 396 Chamber crystals, 205 Chameleon mineral, 527 INDEX 563 Chameleon solution, 527 Chance, 378 Chancel, 326 Chaptal, 146 Charcoal, 225 Chemical change, 1, 2, 5 cause of, 11 factors affecting same, 12, 142 rate of, 12, 23 types of, 9 Chemical compound, distinguished from solution, 6 Chemical elements, 6 distribution of, 8, 9 Chemical formula, interpretation of, 84 Chemical reactions, in electrolytes, 453 in insulators, 453 Chemistry, branches of, 4, 5 organic, 228 scope of, 1 Chevreul, 265 Chili saltpeter, 118, 173, 380 Chloral, 253 Chloral hydrate, 253 Chlorates, 104 Chlorauric acid, 478 Chloric acid, 104 Chlorination process, 476 Chlorine, 54 action of, on water, 57, 143 bleaching with, 57, 104 compounds with oxygen, 58, 101 dioxide, 58, 101 elementary nature of, 128 heptoxide, 102, 106 monoxide, 101 peroxide, 58, 101 preparation of, 55 properties of, 56 reactions of, 89 uses of, 57 Chloroform, 249 Chlorous acid, 106 Chlorplatinic acid, 556 Chlorsul phonic acid, 212 Choke damp, 288 Chromates, 515 Chrome alums, 515 green, 513 iron ore, 514 steel, 513 yellow, 517 Chromic acid, 517 chloride, 514 hydroxide, 513 oxides, 513 sulphate, 515 Chromite, 514 Chromites, 514 Chromium, 512 analytical tests for, 518 properties of, 513 Chromium trioxide, 515, 517 Chromous chloride, 514 hydroxide, 513 sulphate, 514 Chromyl chloride, 513, 517 Chrysoberyl, 414, 487 Cinchonine, 275 Cinnabar, 422, 428 Citric acid, 263 Clarke, F. W., 8 Glaus, 553 Clausius, 449, 450, 451 Clausius's electrolytic theory, 449 Clays, 490 Cleve, 494 Coal, bituminous, hard, soft, 226 gas, 279, 280 tar, uses of, 279 Cobalt, 531, 547 amines, 549 analytical tests for, 550 compounds, 548 dioxide, 548 sesquioxide, 548 silicides of, 549 yellow, 549 Cobaltic hydroxide, 548 silicate, 548 sulphate, 548 Cobaltite, 334, 547 Cobaltous carbonate, 549 chloride, 548 cobaltic oxide, 548 cyanide, 549 hydroxide, 547 nitrate, 548 oxide, 548 sulphate, 548 sulphide, 549 Cocaine, 275 Codeine, 275 Colemanite, 318 Collodion, 274 Colloidal solution, 311 Colloids, 311 Columbites, 353 Columbium, 353 Combining weights, 63 table of, 65 unit of, 64 Combustion, 28, 283 earlier views of, 34 heat of, 30, 297 in air, 29 of oxygen in hydrogen, 34 temperature of, 29 Compounds, denned, 9, 10, 139 stability of, 303 Concrete, 400 Conductors of the first class, 445 of the second class, 445 Condy's disinfecting fluid, 528 564 INDEX Confieldite, 498 Congo red, 137 Consolute liquids, 433 Consolute pairs, 434 Copper, 460, 462 analytical tests for, 468 aresnite, 468 carbonates, 467 extraction of, from ores, 461 ferrocyanide, 466 ferrocyanide membrane, 443 glance, 460 in plants and animals, 460 matte, 461 nitrate, 467 pyrites, 460 refining, 462 salts of oxy-acids, 466 sulphate, 466 sulphate, basic, 467 sulphate, double salts with, 467 Copperas, 541 Corpuscles, 454 Corrosive sublimate, 426 Corundum, 485 Cotton-seed oil, 264 Courtois, 116 Crawford, 410 Cream of tartar, 262 Creosote, 252 Crocoisite, 504, 512 Cronstedt, 545 Crookes, 493 Crookesite, 493 Cryolite, 107, 482 Crystalline substances, 47 Crystalloids, 311 Crystals, 189 Crystal systems, 189 Cullinan diamond, 222 Cupellation, 470 Cup grease, 246 Cuprammonium sulphate, 467 Cupric ammonium chloride, 465 bromide, 466 chloride, 465 chloride, hydrolysis of, 139, 140 compounds, 460 cyanide, 466 fluoride, 466 hydroxide, 465 oxide, 464 sulphide, 468 Cuprous bromide, 465 chloride, 465 compounds, 460 cyanide, 466 fluoride, 466 iodide, 465 oxide, 464 sulphide, 468 Curie, M. and Mme., 411, 412 Curtius, 159, 167 Cyanates, 241 Cyanic acid, 242 Cyanide process, 476 Cyanides, 240 of copper, 466 of iron, 542 Cyanogen, 240 Cymogene, 245 D Dalton, 61, 63, 68, 69, 70, 71, 73, 76 Davy, 55, 116, 288, 363, 373, 446 Deacon process, 55 Debierne, 412 Decomposition, double, 9, 437, 452 hydrolytic, 138 Definite proportions, law of, 4, 60 Delafontaine, 496 d'Elhujar brothers, 520 Deliquescence, 48 Del Rio, 352 Democritus, 70 Desiccator, 48 Developers, 473 Deville, H. Sainte-Claire, 54, 144, 175 Dextrine, 271, 273 Dextrose, 268 Dewar, 108, 153 Dialysis, 310, 311 Diamond, 221 Diaspore, 485 Diastase, 270, 271 Didymium, 496 Dilution, heat of, 296 Dimorphism, 188, 194 Dioxogen, 99 Diphosphorus tetraiodide, 329 Dissociation, 114, 144 in solutions, 441 Disthen, 490 Distillation, destructive, 158, 226 Disulphuric acid, 210 Dobereiner, 355 Dobereiner's lamp, 556 Dolomite, 396, 417 Dulong, 42, 81, 168 Dulong and Petit, law of, 80 Dumas, 42, 151, 234 Dutch metal, 463 Dutch process, 509 Dyeing, of cotton, 487 of wool and silk, 488 Dynamite, 266 Dysprosium, 496 E Earth, infusorial, 309 Earthenware, 491 Earths, alkaline, 394 metals of, 482 Efflorescence, 48 INDEX 565 Eka-aluminum, 358, 492 Eka-boron, 358, 494 Eka-silicon, 358, 498 Elastine, 276 Electric batteries, 454, 456 Electric furnaces, 223 Electrochemical series of the metals, 457 Electrodes, 447 Electrolysis, 445, 447 Electrolytes, 445, 446 migration in, 451 Electrolytic dissociation, 450 process, 509 soda process, 380 solution tensions, 457 theories, 448 Electromotive force, 455 Electrons, 454 Electron theory, 454 Elements, acid-forming, 130, 132 base-forming, 130, 132 chemical, list of, 7 classification of, 355 groups of, 7, 357 transmutation of, 413 Emanations from radium, etc., 413 Emerald, 414 Emery, 485 Emulsin, 271 Emulsion, 434 Enantiomorphism, 263 Energy, 1 conservation of, 11 electrical, 455 free, 296 total, 296 transformation of, 11 Ending ide, 78 Endings and prefixes, 107 Endothermic changes, 290 Enzymes, 271 Epicurus, 70 Epsom salt, 48, 417 Equations, chemical, 78, 87 Equilibrium, chemical, 139, 141, 161 Equivalents, chemical, 22, 63, 69 Erbium, 496 Erythrite, 250 Esters, 264 Ethane, 244, 247 Ethers, 266 Ethyl amine, 274 borate, 317 chloride, 249 ether, 266 nitrate, 264 nitrite, 264 silicate, 315 sulphide, 267 Ethylene, 247 bromide, 247 Europium, 496 Euxenite, 494 Exothermic changes, 290 Facts relative to laws and theories, 70 Faraday, 446, 447, 448, 451 Faraday's law, 448 view on electrolysis, 449 Fats, 264 solubility of, 49, 434 Fatty series, 248 Fehling's solution, 268 Fermentation and ferments, 271 Ferric acetates, 544 acid, 539 ammonium alum, 490 bromide, 540 chloride, 540 . ferrocyanide, 543 hydroxide, 538 nitrate, 544 oxide, 538 oxides, hydrated, 539 phosphate, 544 sulphate, 542 sulphide, 541 Ferricyanic acid, 543 Ferrochromium, 513 Ferrocyanic acid, 543 Ferromanganese, 524 Ferrous ammonium sulphate, 542 bicarbonate, 542 bromide, 540 carbonate, 542 chloride, 539 ferric cyanide, 543 ferric oxide, 539 hydroxide, 538 iodide, 540 nitrate, 544 oxide, 538 phosphate, 544 sulphate, 541 sulphide, 541 Fertilizer from slag, 536 Fertilizers, 207 Fire brick, 416 Fire damp, 288 Fischer, Emil, 276 Flame, 282 hydrogen, 20 luminosity of, 284 oxidizing, 287 reducing, 287 reverse, 283 singing, 20 structure of, 286 Flash light powder, 415 Flash point, 245 Flint, 307 Flores zinci, 419 Flour, wheat, 273 566 INDEX Fluorine, 107, 108 Fluorspar, 107 Fluosilicates, 314 Fluosilicic acid, 314 Flux, 533 Food values, 303 Force, 2 Formaldehyde, 253 Formaline, 253 Formation, heat of, 296 Formic acid, 254 Formic aldehyde fumigation, 528 Franklin, E. C., 164 Franklinite, 417 Fraunhofer lines, 387 Freezing point, definition of, 46 change with pressure, 46 of solutions, 440 Fructose, 268 Fuel values, 303 Fulminating gold, 479 mercury, 428 silver, 475 Fumaroles, 317 Fumigating, 528 G Gadolinite, 494, 496 Gadolinium, 496 Gahn, 321, 524 Gahnite, 417, 487 Galactose, 270 Galenite (galena), 504 Gallium, 492 Galvanized iron, 419 Garnierite, 545 Gas, candle power of, 282 carbonum, 234 conductivity of, 446 detonating, 32 enriching of, 281 illuminating, 279 laws applied to solutions, 444 marsh, 244 natural, 244 sylvestre, 234 water, 235 Gases, absorption of, by liquids, 433 diffusion of, 18 law of combination of, 44, 71 spectra of, 388 Gasoline, 245 Gautier, 153 Gay-Lussac, 55, 72, 116 Gay-Lussac, law of, 42, 46, 71, 160 tower, 207 Geber, 170 Gelatine, 276 Gerhardt, 79, 81 Germanium, 498 German silver, 464, 545 Germicides in agriculture, 463 Gersdorffite, 545 lass, 403 Bohemian, 405 bottle, 406 colored, 405 composition of, 406 crown, 405 cut, 405 enamel, 406 flint, 405 hard, 405 history of, 406 ordinary, soda-lime, 404 plate, 404 potash-lime, 405 soft, 405 window, 404 Glauber, 51, 158, 170 Glauber's salt, 48, 381 Glazes, 491 Glover tower, 207 Glucinum, 414 compounds of, 414 Glucose, 268 Gluten, 273 Glycerine, 250 Glycol, 250 Gly colic acid, 257 Gold, 460, 476, 477 alloys of, 478 analytical tests for, 480 compounds of, 478 electroplating with, 477 metallurgy of, 476 production of, 478 Goldschmidt's process, 484, 513 Gram-molecule defined, 140 Granulose, 272 Grape, sugar, 268 Graphite, 222 Graphitic acid, 223 Greenockite, 421 Grotthus, 448, 449, 450 Grotthus's theory, 448 Guajacol, 252 Gun cotton, 274 Gun metal, 464 Gunpowder, black, 368 smokeless, 274 Gypsum, 186, 400 dead burned, 401 Hales, 234 Halides, of copper, 465 of lead, 507 of silver, 472 Halogens, 101 compounds of, with each other, 123 compounds of, with sulphur, 123 general relations of, to one another, 124 INDEX 567 Hammer black, 539 Hardness of water, 231, 396 Hartshorn, spirits of, 158 Hausmannite, 523 Heat of combustion, 297 tables of, 301 Heat of formation, 296 tables of, 299, 300 Heat of neutralization, 295, 452 table of, 302 Heat of solution, 51, 296 Heavy spar, 407 Helium, 155, 156 Helium group, discovery of, 146 Hematite, 531 Henry, 433 Heptane, 244 Hermann, 421 Hesiod, 532 Hess, 292 Hess, law of, 293 Hexagonal system, 192 Hexane, 244 Hippuric acid, 256 Hisinger, 446, 495 Hjelm, 519 Hjelmite, 494 Hoffman apparatus, 42 Homer, 532 Homologous series, 245 Hornblende, 415 Horn silver, 468 Human body, composition of, 9 Hydrargyrum, 422 Hydrates, 49 Hydrazine, 165 Hydrazoic acid, 167 Hydriodic acid, 119, 120 Hydrogel, 311 Hydrobromic acid, 112, 113, 114 Hydrocarbons, 244 general behavior of, 248 halogen substitution products of, 249 preparation of, 246 Hydrocarboxylic acids, 257 Hydrochloric acid, 51 chemical behavior of, 53 composition of, 52, 54 electrolysis of, 53 preparation of, 89 solutions of, 52 Hydrocyanic acid, 241 Hydrofluoric acid, 109 Hydro fluosilicic acid, 314 Hydrogen, 13 absorption of, by palladium, 554 adsorption of, by solids, 19 antimonide, 341 diffusion of, 18 dioxide, 95 history and occurrence of, 13 in the air, 153 occlusion of, by platinum, 556 preparation of, 14, 15, 16, 87, 88 properties of, 17 uses of, 22 Hydrogen peroxide, 95 bleaching with, 99 formula of, 98 properties of, 96 uses of, 99 Hydrogen persulphide, 197 selenide, 213 silicide, 312 sulphide, 194, 197 Hydrolysis, 137, 138 irreversible, 137 reversible, 139 Hydronitric acid, 167 Hydroquinone, 252 Hydrosol, 311 Hydroxides, 49 Hydroxylamine, 168 Hygroscopicity, 48 Hypo, 383 Hypobromites, 114 Hypobromous acid, 114 Hypochlorites, 102 Hypochlorous acid, 102 Hypoiodites, 121 Hypoiodous acid, 121 Hyponitrous acid, 179 Hypophosphoric acid, 330 Hypophosphorous acid, 330 Hyposulphite of soda, 211, 382 Hypothesis (see Theory) Ice, action of, on salt, 5 crystalline nature of, 47 machines, 164 Iceland spar, 396 Illuminants, 281 Incandescent lamp filaments, 520 Indicators, 137 neutrality toward, 138 Indium, 493 Ingots, 536 Ink, India, 226 ordinary, 542 sympathetic, 548 "white," 544 Insulators, 446 Inulin, 269 Inversion of sugar, 268 Invertase, 269, 271 lodates, 122 lodic acid, 121 Iodides, 121 Iodine, 116 monobromide, 123 monochloride, 123 oxide of, 121 oxy-acids of, 121 568 INDEX Iodine Continued pentafluoride, 123 preparation of, 116, 117, 118 properties of, 118 tincture of, 118 trichloride, 123 uses of, 119 lodocrol, 119 lodoform, 119, 249 Ion, 454 Ionic theory, 450 Ions, 447, 451, 452 Indium, 555 Iridosmium, 554 Iron, 531 action of, on sulphur, 3 alloys, 537 analytical tests for, 544 carbide, 544 cast, gray, white, 534 chlorides of, 539 dialyzed, 539 disulphide, 541 electrolytic, 537 metallurgy of, 532 occurrence of, 531 ore, 531 ore, bog, 539 ore, brown, 539 ore, chrome, 512 ore, titaniferous, 315 oxides of, 30, 538 passive, 538 phosphides of, 544 production of, 537 properties of, 537 silicides of, 544 wrought, 534 Isomers, 10, 260 Isometric system, 191 Isomorphism, law of, 82, 194 Isosmotic solutions, 445 Isotonic solutions, 445 Javelle water, 104 Jorgensen, 550 K Kahlenberg, 445, 448 Kainite, 363, 415 Kaolin, 490 Kelp, 116 Keratine, 276 Kermes mineral, 347 Kerosene, 245 Ketones, 267 Kieserite, 415 Kindling temperature, 29 Kipp apparatus, 230 Kirchhoff, 385, 388 Kjeldahl's method, 159 Klaproth, 214, 315, 495, 519, 520 Kreosol, 252 Krypton, 156 Kunkel, 321 Lactates, 258 Lactic acid, 257 bacteria, 258 Lactose, 270 Ladenburg, 94 Lsevulose, 268 Lakes, 487 Land plaster, 401 Lapis infernis, 474 lazuli, 492 Lampblack, 225 Lamy, 493 Lana philosophica, 419 Lanthanum, 494 Laplace, 292 Latent heat, of fusion, 290 of vaporization, 290 Laughing gas, 180 Laundry blue, 492 Laurent, 79 Lavoisier, 35, 36, 37, 60, 70, 146, 170, 234, 292 Law and theory, difference between, 70 Law denned, 4 Law of Henry, 433 Hess, 293 Lavoisier and Laplace, 292 maximum work, 294 octaves, 356 thermoneutrality, 452 Lead, 504 acetates, 508 alloys of, 505 analytical tests for, 510 arsenate, 508 bromide, 507 carbonate, 508 chamber process, 203 jchloride, 507 chromate, 517 hydroxide, 506 iodide, 507 nitrate, 508 oxides, 31, 506 oxybromide, 507 oxychloride, 507 peroxide, 506 pipes, 505 production of, 504 properties of, 504 sesquioxide, 506 sulphate, 508 sulphide, 508 tetrachloride, 507 INDEX 569 Lead Continued tree, 505 uses of, 505 water, 508 Le Bel, 259 Le Blanc soda process, 377 Le Chatelier, principle of, 47 Lecithine, 321 Lecoq de Boisbaudran, 492, 496 Lenssen, 355 Leucippus, 70 Lewes, 285 L'Hermite, 443 Liebig, 415 Ligroin, 245 Lime, 397 chloride of (see Bleaching powder) feldspar, 490 Limekiln, 397 Limestone, 396 Limewater, 397 Limonite, 531, 539 Liquid air, 154 Litharge, 506 Lithium, 383 compounds of, 383 r 384 nitride, 164 Lithosphere, composition of, 8 Litmus, 137 Liver of sulphur, 371 Lodestone, 539 Lessen, 166 Lubricating oils, 245 Lunar caustic, 474 Lunge, 204 Lutecium, 495 Luteocobaltic chloride, 549 M Magenta, 274 Magnalium, 484 Magnesia, 415 alba, 416 calcined, 415 usta, 415 Magnesite, 415, 416 Magnesium, 414 amalgam, 424 ammonium arsenate, 417 ammonium chloride, 416 ammonium phosphate, 331, 417 carbonates, 416 chloride, 416 chromate, 517 group, general remarks about, 430 group, metals of, 414 hydroxide, 415 metallic, 415 nitride, 164, 415 oxide, 415 phosphates, 417 pyroarsenate, 417 pyrophosphate, 417 silicide, 306 sulphate, 417 tests for, 417 Magnetic iron ore, 539 Magnetite, 531, 539 Malachite, 460, 467 Malic acid, 262 Malleable iron, 535 Malonic acid, 257 Maltose, 270 Manganates, 526 Manganese, 523 analytical tests for, 529 blende, 523 bronze, 524 carbonate, 526 chloride, 526 dioxide (peroxide), 524 heptoxide, 525 monoxide, 524 nitrate, 526 properties of, 523 protosesquioxide, 524 salts of, 525 sesquioxide, 524 spar, 523 sulphate, 525 trioxide, 525 Manganic acid, 526 chloride, 526 compounds, 523, 526 sulphate, 526 Manganine, 545 Manganite, 523 Manganites, 525 Manganous carbonate, 526 chloride, 525 compounds, 523, 524, 525, 526 hydroxide, 524 nitrate, 526 Mannite, 250 Marine acid air, 51 Marl, 396, 490 Marsh's test, 336 Mass action, 139 illustrations of, 141 law of, 139 Mass, conservation of, 10 Matches, 325 Matter, 1 McBride, 234 Meerschaum, 415 Membrane, semipermeable, 442, 445 Mendeleeff, 356, 358, 494, 498 Mendeleeff' s view of Arrhenius's theory, 454 Mercuric ammonium chloride, 429 bromide, 427 chloride, 426 cyanide, 240, 406 diammonium chloride, 429 570 INDEX Mercuric fulminate, 428 iodide, 427 nitrate, 427 oxide, 425 sulphate, 428 sulphide, black, red, 428 Mercurous ammonium chloride, 429 ammonium nitrate, 429 bromide, 426 chloride, 425 chromate, 517 iodide, 426 nitrate, 427 oxide, 425 sulphide, 428 Mercury, 422 compounds of, 424 compounds, physiological properties of, 429 halides of, 425 lamp, 423 oxides of, 425 salts, compounds of, with ammonia, 429 tests for, 429 Meta-aluminates, 487 Meta-aluminic acid, 486 Meta-arsenates, 338, 339 Meta-arsenic acid, 339 Metaborates, 318 Metaboric acid, 317 Metaformaldehyde, 260 Metalloids, 7 Metals of the earths, 482 Metamerism, 261 Metantimonious acid, 344 Metaphosphoric acid, 330, 332 Metaplumbic acid, 507 Metastannates, 502 Metastannic acid, 502 Metathesis, 9 Methane, 244 Methyl, amine, 274 butyrate, 264 chloride, 249 formate, 264 hydrogen sulphate, 264 iodide, 264 orange, 137 salicylate, 264 silicate, 315 sulphide, 267 Meyer, Lothar, 53, 356, 360 Meyer, Victor, 198, 335 Microcosmic salt, 331 Milk of lime, 397 Milk sugar, 270 Millon's reagent, 428 Minium, 506 Mirrors, 424 Mispickel, 334 Mitscherlich, 82, 194 Mixture, 432 Mohr's salt, 542 Moissan, 107, 108, 222, 315, 353 Molasses, 270 Molecular theory, basis of, 74 Molecular weight, determinations of, 74 determinations of, in solutions, 439 in solutions, discussion of, 441 Molecules, 72, 79 Molybdates, 518 Molybdenite, 518 Molybdenum, 512, 518 compounds of, 519 trioxide, 518 Molybdic acid, 519 Monazite, 316, 494, 495 Monocalcium phosphate, 322 Monoclinic system, 193 Monoses, 267 Monosymmetric system, 193 Mordants, 487 Morphine, 275 Mortar, 397 Mosaic gold, 503 Mosander, 494 Moth balls, 248 . Miiller von Reichenstein, 214 Multiple proportions, law of, 32, 60 Muntz metal, 463 Muriatic acid, 51 Mycoderma aceti, 254 N Naphtha, 245 Naphthalene, 248 Narcotine, 275 Nascent state, 89 Negatives, 473 Neodymium, 496 Neon, 156 Neoytterbium, 495 Nernst lamp, 496 Nessler's reagent, 165, 427 Neutralization, act of, 129 Arrhenius's view of, 452 heat of, 297, 302 Newlands, 356 Newton's metal, 348 Nicholson, 446 Nickel, 531, 545 alloys of, 545 ammonium sulphate, 546 analytical tests for, 550 carbonyl, 546 coins, 545 glance, 545 plating, 546 salts of, 546 steel, 545 Nickelic hydroxide, 546 oxide, 546 INDEX 571 Nickelous chloride, 546 cyanide, 546 hydroxide, 545 nitrate, 546 oxide, 546 sulphate, 546 sulphide, 546 Nicollite, 545 Nicotine, 275 Nilson, 494 Niobium, 353 Niter, 170 Nitrates, of mercury, 427 test for, 176 Nitric acid, 170 properties of, 173 red, fuming, 172 Nitric oxide, 175 Nitrides, 149 Nitrobenzene, 274 Nitrocellulose, 274 Nitrogen, 146 assimilation of, 152 compounds, general considerations, 181 compounds of, with halogens, 167 dioxide, 176 distribution of, 147 in rain water, 153 iodide, 168 molecular formula of, 150 oxides, composition of, 61 pentoxide, 174 preparation of, 147, 148 properties of, 149 tetroxide, 176 tribromide, 168 trichloride, 167, 168 trioxide, 179 valence of, 165 Nitroglycerine, 266 Nitrosyl chloride, 174 Nitrosyl sulphuric acid, 204, 205 Nitrous acid, 178 Nitrous oxide, 180 Nomenclature, 86, 106 Nonane, 244 Nordhausen sulphuric acid, 210 Nucleoproteins, 276 Nux vomica, 275 O Occlusion, 19 Ocean, composition of, 8 Octane, 244 Oil gas, 281 Oil of mirbane, 274 Oleic acid, 256 Olefiant gas, 247 Olivine, 309 Olive oil, 264 Onnes, Kammerlingh, 156 Opal, 307 Open hearth process, 537 Opium, 275 Optical activity, 258 Organic acids, 253 Organisms in water, 153 Orpiment, 334 Orthite, 494 Orthoclase, 310, 490 Orthophosphoric acid, 330 Orthorhombic system, 193 Osmates, 555 Osmic acid, 554 Osmiridium, 554 Osmium, 554 compounds of, 554 Osmosis, 441 Osmotic pressure, 441, 443 Oxalic acid, 256 Oxamide, 241 Oxidation, 21, 88 stages of, 30 Oxides, acid-forming, 127 Oxides of cobalt, 547 copper, 464 iron, 538 lead, 506 manganese, 524 nickel, 545 nitrogen, 170 silver, 471 tin, 501 Oximes, 167 Oxy-acids of nitrogen, 170 Oxygen, 26 atomic weight of, 73, 76 history, occurrence, 26 preparation, 88 properties, 27 Oxyhydrogen blowpipe, 33 Ozone, 92 in the air, 153 properties of, 94 relation of, to oxygen, 93 Ozonic acid, 99 Packfong, 545 Paint, 509 Painter's colic, 510 Palladium, 554 Palladium compounds, 554 Palladium-hydrogen, 554 Palmitic acid, 256 Paper, 273 sizing of, 489 Paracelsus, 13, 429 Paraffin, 245 Paraffin series, 244, 245 Paraformaldehyde, 260 Pararosaniline, 274 Paris green, 338 572 INDEX Parke's process, 469 Parting, 477 Passive state, 538, 545 Pasteur, 263 Pattinson's process, 470 Pattinson's white lead, 507 Pearlash, 369 Peligot, 520, 521 Pentane, 244 Pepsin, 271 Peptones, 276 Perchlorates, 105 Perchloric acid, 105 Percussion caps, 428 Periodates, 122 Periodic acid, 122 Periodic law, 360 Periodic system, 355, 357 Permanent white, 409 Permanganates, 526 uses of, 528 Permanganic acid, 527 Persulphates, 211 Persulphuric acid, 211 Petit, 81 Petroleum, 244 Petroleum ether, 245 Petzite, 476 Pewter, 500 Pfeffer, 444 Phenacite, 414 Phenolates, 252 Phenolphthalein, 137 Phenols, 252 Phenyl hydrazine, 166 Phlogistic theory, 35 Phlogiston, 35, 55 Phosgene, 237 Phosphate rock, 402 Phosphine, liquid, 328 solid, 328 Phosphines, 326 Phosphomolybdic acid, 519 Phosphonium compounds, 327 iodide, 327 Phosphor bronze, 464 Phosphoric acid, 127, 330 glacial, 315, 332 Phosphorous acid, 331 Phosphorus, 321 compounds of, with sulphur, 334 group, general considerations of, 350 in iron, 534 oxides and acids of, 329, 333 oxybromide, 329 oxychloride, 329 oxyfluoride, 329 pentabromide, 329 pentafluoride, 329 pentasulphide, 334 pentoxide, 127, 330 preparation of, 322 red, amorphous, 324 sulphochloride, 334 tribromide, 329 trichloride, 328 trifluoride, 329 triiodide, 329 uses of, 324 yellow or white, 323 Phosphotungstic acid, 520 Photography, 472 Physical change, 1, 2, 5 Physical mixture, 3 Pig iron, 533 Pineapple oil, 264 Pink salt, 501 Pintsch gas, 281 Pitchblende, 410, 520 Plants, nitrogen supply of, 152 Plaster of Paris, 400 Platinates, 557 Platinic chloride, 556 hydroxide, 557 oxide, 557 sulphide, 557 Platinous hydroxide, 557 oxide, 557 sulphide, 557 Platinum, 555 analytical tests for, 557 black, 555 care of, 556 catalytic action of, 556 extraction of, from ores, 552 family, 552 sponge, 555 use of, 556 Pleonast, 487 Pliny, 505 Plumbago, 222 Plumbates, 507 Plumbic acid, 506 oxide, 506 Plumbites, 506 Plumbum, 504 candidum, 498 nigrum, 498 Polariscope, 261 Pollux, 372 Polonium, 412 Polymerization, 260 Polymorphism, 188 Polysulphides, 197 Polythionic acids, 211 Porcelain, 491 Portland cement, 398, 399 Positives, 473 Potash, 369 Potash, caustic, 365 Potash, feldspar, 490 Potassium, 362 alum, 489 amalgam, 366 INDEX 573 Potassium Continued amide, 163 antimonate, 345 antimonyl tartrate, 344 arsenyl tartrate, 345 auric cyanide, 479 aurous cyanide, 479 bicarbonate, 370 bichromate, 515 bisulphate, 371 bisulphite, 371 boryl tartrate, 345 bromate, 367 bromide, 364 carbonate, 369 chlorate, 104, 366, chlorite, 107 chlorplatinate, 556 chromate, 515 chrome alum, 490, 515 citrate, 263 cobaltic cyanide, 549 cobaltic nitrite, 549 cobaltous cyanide, 549 compounds with halogens, 364 cyanate, 368 cyanide, 240, 241, 368 diuranate, 521 ferrate, 539 ferricyanide, 543 ferro cyanide, 241, 542 fluoride, 365 fluosilicate, 370 hydride, 363 hydroxide, 365 iodate, 367 iodide, 364 iodide, uses of, 119 manganate, 526 metantimonate, 345 metantimonite, 344 nitrate, 367 nitrite, 368 oxide, 366 perchlorate, 105, 367 permanganate, 527 peroxide, 366 persulphate, 211 phosphates, 370 platinic chloride, 372 platinous chloride, 557 polysulphides, 371 pyroantimonate, 345 pyrosulphate, 371 ruthenate, 553 silicate, 370 silver cyanide, 475 stannite, 501 sulphate, 370 sulphides of, 371 sulphite, 371 sulphocyanate, 242, 369 sulphophosphate, 334 sulphydrate, 371 tellurate, 216 tellurite, 216 tests for, 372 thiosulphate, 371 water glass, 370 zincate, 132 Pottery, 491 Powder of algaroth, 343 Praseodymium, 496 Precipitation, 437, 452 Prefixes and endings, 107 Preparing salt, 502 Preservation of railroad ties, 246 Priestley, 26, 35, 36, 51, 146, 158, 175, 234 Producer gas, 235 Propane, 244 Properties, chemical, 12 Propionic acid, 256 Propyl amine, 274 chloride, 249 iodide, 249 tartrate, 264 Proteins, or proteids, 275 Proust, 70 Proustite, 468 Prussian blue, 543 Prussiate of potash, red, 543 yellow, 241, 543 Prussic acid, 241 Ptomaines, 276 Puddling, 534 Purple of Cassius, 479 Purpureo-cobaltic chloride, 549 Putrescine, 276 Putty, 397 Pyrargyrite, 468 Pyridine, 275 use of, in osmosis, 445 Pyrite, 531, 541 Pyroantimonic acid, 345 Pyroarsenates, 339 Pyroarsenic acid, 339 Pyroboric acid, 317 Pyrogallol, or pyrogallic acid, 252 Pyrolusite, 523 Pyromorphite, 321, 504 Pyrophosphoric acid, 330, 332 Pyrosulphates, 210 Pyrosulphuric acid, 210 Q Quadratic system, 192 Quartation, 477 Quartz, 307 Quartz glass, 307 Quartzite, 307 Quicklime, 397 Quinine, 275 Quinoline, 275 674 INDEX Radicals, alkyl, 249 Radio-activity, 410 Radium, 410 rays, 412 salts of, 411 Ramsay, 146, 154, 155, 156 Rare-earth elements, 493 Rayleigh, 146, 153, 154 Reaction, heat of, 296 irreversible, 142 reversible, 142 Realgar, 334 Reciprocal proportions, law of, 58 61 Red lead, 506 Red ocher, 538 Red precipitate, 425 Red prussiate of potash, 543 Red zinc ore, 417 Reducing agent, 21 Reduction, 21, 89 Refrigeration, artificial, 164 Regnault, 18 Regular system, 191 Reich, 493 Resin soap, 489 Respiration of plants, 33 Respiration, role of oxygen in, 32 33 Rhigolene, 245 Rhodium, 553 compounds of, 553 Rhodochrosite, 523 Rhombic system, 193 Rhombohedral crystals, 193 Richter, J. B., 59, 62, 70 Richter, 493 Rinmann's green, 421 Rochelle salt, 262 Rock candy, 269 Rocks, disintegration of, 45 Roebuck, 203 Roozeboom, 540 Rosaniline, 274 Roscoe, 234 Roseo-cobaltic chloride, 549 Rose's metal, 348 Rosin, 489 Rotary power, specific, 262 Rouge, 538 Rubidium, 372 Ruby, 485 Ruby silver ore, 468 Rum, 250 Ruthenium, 553 compounds of, 553 Rutherford, 146 s- Saccharose, 269 Saccharum saturni, 508 Safety lamp, miner's, 288 Sal ammoniac, spirits of, 158 Saleratus, 158 Salivation, 429 Salt cake, 377 Salt, definition of, 130, 132 formation of, 130 formation, older view of, 131 spirit of, 51 Saltpeter, 367 Salts, 130 acid, 133 basic, 134 neutral or normal, 133, 135 Samarium, 496 Samarskite, 494, 496 Sandstones, 312 Saponification, 265 Sapphire, 485 Sarcolactic acid, 258 Saturation, degrees of, 436 Scandium, 494 Scheele, 26, 35, 36, 55, 65, 146, 263, 321 335, 409, 519, 520, 523 icheele's green, 338, 468 cheelite, 519 Schlippe's salt, 347 chonbein, 92 chonite, 370, 415 chrotter, 324 chweinfurt green, 338 chweitzer's reagent, 273, 463 efstrom, 353 elenic acid, 214 elenious acid, 214 elenite, 400 elenium, 212, 218 compounds of, 213, 214 erpentine, 310, 415 etterberg, 372 hot, 505 derite, 531 emens-Martin process, 537 lica, 307 licates, action of, on water, 312 decomposition of, 312 licic acids, 309, 312 licic acid, esters of, 315 iico-e thane, 313 lico-methane, 313 icon, 306 amorphous, 307 carbide, 315 chloroform, 314 compounds of, with halogens, 313 crystalline, 307 dioxide, 307 tetrabromide, 314 tetrachloride, 314 tetrafluoride, 313 tetraiodide, 314 ver, 460, 468 analytical tests for, 475 INDEX 575 Silver Continued bromide, 472 carbonate, 474 chromate, 475, 517 chloride, 472 cyanide, 474 extraction of, from ores, 469 fluoride, 472 fulminate, 475 glance, 468 iodide, 472 mirrors, 475 nitrate, 473 nitrite, 474 oxide, 472 peroxide, 472 phosphate, 332, 474 plating, 475 properties of, 471 pyrophosphate, 332 refining of, 470 solvent for, 471 sterling, 471 suboxide, 472 sulphate, 474 Smalt, 549 Smaltite, 547 Smithsonite, 417 Soap, 265 Soapstone, 415 Soda, 377 baking, 380 calcined, 378 caustic, 376 crystallized, 378 feldspar, 490 washing, 378 water, 232 Sodium, 373 alcoholate, 250 alum, 489 amalgam, 373, 424 amide, 163 bicarbonate, 380 bichromate, 516 bisulphate, 382 bisulphite, 382 benzoate, 256 borate, 383 bromide, 375 carbonate, 377 chloraurate, 479 chloride, 373 chlorplatinate, 556 chromate, 516 cyanide, 383 diuranate, 521 fluoride, 375 formate, 254 hydride, 373 hydroxide, 376 hydrosulphide, 383 iodide, 375 manganese alum, 490 metaphosphate, 331 nitrate, 380 nitrite, 380 oleate, 264 oxide, 375 perchlorate, 105 permanganate, 528 peroxide, 375 persulphate, 211 phosphates, 380, 381 pyroantimonate, 346 pyrophosphate, 331 silicate, 310, 383 stannite, 501 sulphate, 381 sulphides, 383 sulphite, 382 thiosulphate, 211, 382 tungstate, 419 water glass, 383 Soffioni, 316 Soil, nitrogen in, 158 organisms in, 153 Solder, 500 Solid solution, 432, 437 Solubility curve of, ferric chloride, 540 magnesium chloride, 416 sodium sulphate, 381 Solubility curves, 435 Soluble Prussian blue, 543 Solution, distinguished from chemical compound, 6, 432 heat of, 296 normal, 136 saturated, 436 standard, 136 supersaturated, 382, 437 unsaturated, 436 Solutions, 432 colloidal, 311, 438 nature and kinds of, 432 of liquids in liquids, 433 of solids in liquids, 434 solid, 432, 437 Solvay, 379 Solvay process, 379 Sonnenschein's reagent, 519 Soot, 225 Spark spectra, 388 Specific heats, table of, 80 Spectra of metals, 386 Spectroscope, 385 Spectrum analysis, 384 Spectrum, continuous, 387 reversed, 388 spark, 388 Spelter, 418 Spiegeleisen, 524, 536 Spinels, 487 Spiritus fumans Libavii, 501 576 INDEX Stahl, 35, 55 Stannates, 502 Stannic acid, 501, 502 chloride, 501 hydroxide, 502 oxide, 502 sulphide, 503 Stannous chloride, 500 hydroxide, 501 oxide, 501 sulphide, 503 Starch, 271 paste, 272 production of, 233 soluble, 272 Stas, 234 Stearic acid, 256 Steel, mild, structural, tool, 535 Stereochemistry, 260 Stereoisomerism, 260 Stibine, 341 Stibnite, 340 Stohmann, 304 Stoichiometry, laws of, 62 Stolzite, 519 Stoneware, 491 Storage battery, 456 Strohmeyer, 421 Strontianite, 406 Strontium, 406 carbonate, 407 chloride, 407 compounds of, 406 dioxide, 407 hydroxide, 407 nitrate, 407 oxide, 407 sucrate, 270 sulphate, 407 sulphide, 401 Structural formulae, 84 Strychnine, 275 Sublimate, corrosive, 426 Sublimation, 117 Substance, definition of, 2 Succinic acid, 257 Sucrose, 269 Sugar of lead, 508 Sugars, 267 Sulphates of, iron, 541 mercury, 428 Sulphides of, copper, 468 ; iron, 541 ; tin, 503 Sulphites, 200 Sulphocyanates, 241 Sulphostannates, 503 Sulphur, 185 allotropic forms of, 188 amorphous, 188 auratum, 347 bleaching with, 199 compounds with halogens, 197 dichloride, 198 dioxide, 198 ethers, 266 flowers of, 187 group, general considerations of, 216 hexafluoride, 197 hexaiodide, 198 in iron, 534 milk of, 188 monobromide, 198 monochloride, 198 monoclinic, 198 monoiodide, 198 occurrence of, 185 peroxide, 211 plastic, 188 precipitated, 188 preparation of, 187 properties of, 187 rhombic, 187 roll, 187 sesquioxide, 201 tetrachloride, 198 trioxide, 127, 201, 203 uses of, 189 Sulphuric acid, 127 action of, on salt, 143 contact process, 202 factory, 206 hydrates of, 210 lead chamber process, 203 monohydrate of, 209 properties of, 208 uses of, 208 Sulphuric ether, 266 Sulphurous acid, 127, 200 Sulphuryl chloride, 212 Superphosphate, 207, 322 Superphosphate of lime, 402 Sylvanite, 476 Sylvite, 363 Symbols, chemical, 65 Talc, 415 Tank waste, 378 Tanning, chrome, 516 Tantalites, 353 Tantalum, 353 Tartar emetic, 344 Tartaric acid, 262 Taylor, E. R., 238 Tellurium, 214, 218 compounds of, 215, 216 Tempering, 535 Tennant, 555 Terbium, 496 Ternary compound, 87 Tetraboric acid, 317 Tetragonal system, 192 Thallium, 493 compounds of, 493 INDEX 577 Thenard, 55, 95 Thenard's blue, 492 Thenard's method, 509 Theory and law, difference between, 70 Theory of electrolytic dissociation, 450 Theory, use of, 70 Thermite, 485 Thermochemical data, 298 uses of, 302 Thermochemical equations, 293, 294 Thermochemistry, 290 laws of, 292 Thionyl chloride, 212 Thiosulphates, 210 Thiosulphuric acid, 211 Thomas-Gilchrist process, 536 Thomas slag, 536 Thomsen, Julius, 294, 303, 454 Thorium, 315 Thorium X, 412 Thoulet's solution, 427 Thulium, 496 Tin, 498 amalgam, 424 analytical tests for, 503 Banca, 499 block, 499 chlorides of, 500 cry, 499 pest, 499 recovery of, from tin cans, 500 stone, 498 tetrabromide, 501 tetrachloride, conductivity of, 446, 447 tetraiodide, 501 uses of, 500 Tinned iron, 500 Titanium, 315 Titration, 137 Tombac, 463 Toning bath, 479 Triads, 355 Triazoic acid, 167 Triazoiodide, 168 Triclinic system, 194 Tridymite, 307 Trigonal system, 193 Trioleine, 264 Tripalmitine, 264 Tristearine, 264 Trithiocarbonates, 240 Trithiocarbonic acid, 240 Trypsin, 271 Tungsten, 512, 519, compounds of, 519, 520 Turmeric paper, 137 Turnbull's blue, 543 Turpeth mineral, 428 Type metal, 341 U Ultramarine, 491 Uranium, 512, 520 compounds of, 520, 521 glass, 521 rays, 410 yellow, 521 Uranium X, 412 Urbain, 495 Urea, estimation of, 149 Valence, 84, 85 Valentine, Basil, 203, 345, 347 Vanadium, 352 compounds of, 353 Van Helmont, 234 Van Marum, 92 van't Hoff, 259, 444 Varech, 116 Vaseline, 245 Vauquelin, 412, 512 Venetian red, 538 Verdigris, 462, 468 Vermilion, 429 Villiger, 99 Vinegar, 254 Vitriol, blue, 466 green, 541 oil of, 302 white, 420 Vivianite, 321 W Wanklyn, 116 Water, 38 action of, on rocks, 49 as solvent, 49 clarifying of, 487 comparison of, with hydrogen sulphide, 197 composition of, 42 compounds with, 48 contaminated, 41 distilled, 39 formula of, 73 gas, 235 germs in, 41 glass, 310 hard, 231 hardness of, 396 in animals, 38 in plants, 38 metastable condition of, 46 mineral, 41 natural, 39 of crystallization, 48, 164 pipes, bursting of, 45 potable, 40 properties of, 44 purification of, 41 2P 578 INDEX Water Continued supercooled, 46 thermal, 42 Wavellite, 535 Welding, 535 Weldon process, 525 Welsbach, Auer von, 495, 496 Welsbach light, 286, 495 Werner, 550 Wheels, grinding, 315 Whetstones, 315 Whisky, 250 White lead, 509 adulterants of, 509 White precipitate, 429 Whiting, 397 Wine, 250 spirit of, 250 Winkler, 498 Wintergreen, oil of, 264 Withering, 410 Witherite, 407 Wohler, 228 Wolframite, 519 Wollaston, 63, 553 Wood alcohol, 249 use of, in fumigating, 528 Wood's metal, 348 Wood, spirits of, 249 Work, 2 Wrought iron, 534 Wurtzite, 421 Wulfenite, 504, 518 Xenon, 156 Yeast, 250 Yellow ocher, 539 Yellow prussiate of potash, 543 Ytterbium, 495 Yttrium, 494 Yttrotantalite, 494 Zinc, 417, 418 amalgam, 424 blende, 417 bromide, 420 carbonate, 419 chloride, 419 dust, 418 fluoride, 420 iodide, 420 oxide, 419 oxychlorides, 420 spinel, 417 sulphate, 420 sulphide, 420 tests for, 420 white, 419 Zirconium, 315 Zymase, 271 Printed in the United States of America. JUL 7 QD31 Kahlenberg, L. 42130 v K12 Outlines of chemistry... 1920 Revised.