BIOLOGY t\BRARY OUTLINES OF CHEMISTRY THE MACMILLAN COMPANY NEW YORK BOSTON . CHICAGO SAN FRANCISCO MACMILLAN & CO., LIMITED LONDON BOMBAY CALCUTTA MELBOURNE THE MACMILLAN CO. OF CANADA, Lm TORONTO OUTLINES OF CHEMISTRY A TEXT-BOOK FOR COLLEGE STUDENTS BY LOUIS KAHLENBERG, PH.D. PROFESSOR OF CHEMISTRY AND DIRECTOR OF THE COURSE IN CHEMISTRY IN THE UNIVERSITY OF WISCONSIN gorfc THE MACMILLAN COMPANY 1912 All rights reserved COPYRIGHT, 1909, BY THE MACMILLAN COMPANY. Set up and electrotyped. Published September, 1909. Reprinted January, September, October, 1910. Reprinted with corrections, January, September, 1911 ; August, 1912. October, 1912. Nmrfaoob J. 8. Gushing Co. Berwick & Smith Co. Norwood, Mass., U.S.A. PREFACE 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 views T 266991 yi 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 Bunsen 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 vil 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. CONTENTS CHAPTER I THE SCOPE OF 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 III 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 24 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 36 ix X CONTENTS CHAPTER V HYDROCHLORIC ACID AND CHLORINE PAOB 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 . . 49 V CHAPTER VI THE LAWS OF COMBINING WEIGHTS AND COMBINING VOLCTMES 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 58 CHAPTER VII OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE History, Occurrence, and Preparation of Ozone Relation between Ozone and Oxygen Allotropy Properties of Ozone History, Occurrence, and 'Preparation of Hydrogen Peroxide Properties of Hydrogen Peroxide Formula of Hydrogen Peroxide Uses of Hydrogen Peroxide Ozonic Acid 88 CHAPTER VIII 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 PAGE 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 96 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 Acidimetry 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 120 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 . . 139 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 150 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 . 161 CHAPTER XIII SULPHUR, SELENIUM, AND TELLURIUM Occurrence and Preparation of Sulphur Properties of Sulphur Uses of Sulphur Crystals and Crystal Systems Hydrogen xii CONTENTS 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 176 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 Sulphocyauates 210 CHAPTER XV HYDROCARBONS AND ADDITIONAL COMPOUNDS OP 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 . 232 CHAPTER XVI ILLUMINATING GAS AND FLAMES Illuminating Gas Flame Luminosity of Flame Structure of Flame Davy Safety Lamp 265 CHAPTER XVII THERMOCHEMISTRY General Remarks Calorimeters Laws of Thermochemistry Thermochemical Equations Thermochemical Data Tables Uses of Thermochemical Data 275 CONTENTS xiii CHAPTER XVIH SILICON AND BORON AND THEIR IMPORTANT COMPOUNDS PAGE 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 290 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 304 CHAPTER XX CLASSIFICATION OF THE ELEMENTS THE PERIODIC SYSTEM . . 337 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 . . . 343 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 . . 374 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 393 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 Determined in Solutions Osmosis and Osmotic Pressure Elec- trolysis Electrolytic Theories Electric Batteries Electro- chemical Series of the Metals 410 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 436 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 457 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 . . . 472 CHAPTER XXVIII CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM Occurrence, Preparation, arid Properties of Chromium Chromic Oxides and Hydroxides Chromous Compounds Chromic Salts xvi CONTENTS Chromates, Bichromates, and Chromium Trioxide Chromyl Chloride Analytical Tests for Chromium Molybdenum Tungsten Uranium 485 CHAPTER XXIX MANGANESE Occurrence, Preparation, and Properties Oxides Salts of Manga- nese Manganates and Permanganates Uses of Permanganates Analytical Tests for Manganese 495 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 502 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 522 INDEX . 529 LIST OF ILLUSTRATIONS 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 ....... 26 16. Burning of sulphur in oxygen 26 17. Oxyhydrogen blowpipe 31 18. Combustion of oxygen in hydrogen 32 19. Lavoisier's apparatus to show that mercury unites with oxygen when calcined .......... 34 20. Distillation 37 21. Demonstration of volumetric relations between oxygen, hydrogen, and steam 41 22. Desiccator 46 23. Composition of hydrochloric acid gas 50 24. Electrolysis of hydrochloric acid 51 25. Synthesis of hydrochloric acid gas by volume .... 52 26. Burning of arsenic in chlorine 54 27. Action of chlorine on water in sunlight 55 28. Ozone apparatus 88 29. Preparation of fluorine ......... 103 30. Preparation of hydrobromic acid 108 31. Sublimation of iodine in the laboratory 112 32. Sublimation of iodine on commercial scale ..... 113 33. Titration 129 34. Preparation of nitrogen from the air 140 35. Oxidation of nitrogen by means of the electric spark . . . 143 xvii xviil LIST OF ILLUSTRATIONS PAGB FIG. 36. Volumetric composition of ammonia gas 152 37. Decomposition of ammonia by the electric spark .... 153 38. Burning ammonia mixed with oxygen 154 39. Oxidation of ammonia by use of a platinum spiral . . . 154 40. Preparation of nitric acid 162 41. Heating sodium in nitric oxide ....... 167 42. Commercial distillation of sulphur ...... 177 43. Crystal of rhombic sulphur 178 44 to 54. Crystals of the isometric system 181 55 to 60. Crystals of the tetragonal system 182 61 to 67. Crystals of the hexagonal system 183 68 to 71. Crystals of the orthorhombic system 184 72 to 74. Crystals of the monoclinic system 184 75 to 76. Crystals of the triclinic system 185 77. When sulphur burns in oxygen the volume remains unchanged . 190 78. Bleaching of flowers by means of sulphur dioxide . . .191 79. Sulphuric acid by the contact process 193 80. Diagram of a sulphuric acid factory 197 81 and 82. Crystal forms of diamond 210 83. Acheson graphite furnace - . . 213 84. Typical arc furnace for experimental work 213 85. Absorption of ammonia gas by charcoal 214 86. Kipp apparatus 219 87. Siphoning carbon dioxide from one jar to another . . .221 88. Taylor's carbon bisulphide furnace 228 89. Yeast cells . 238 90. Acetic acid organisms 243 91. Lactic acid organisms 246 92. Formulae of dextro and laevo lactic acids 248 93. Polariscope 249 94. Crystals of dextro and laevo tartaric acid 251 95. Grains of potato starch ' . . 260 96. Grains of wheat starch 260 97. Grains of corn starch 260 98. Potato starch grains in polarized light 260 99. Manufacture of coal gas 266 100. Gases burn in the flame of a candle 268 101 to 103. Demonstration of the reverse flame 269 104. Burning oxygen in coal gas 270 105. Enriching carbon monoxide gas 270 106. Principle of the Bunsen burner 271 107. Zones of the flame of a candle . . . . . . 273 108 and 109. A flame will not pass through a wire gauze . . . 273 110. Davy safety lamp 274 111. Calorimetric apparatus 276 112. Combustion bomb and calorimeter 277 113. Right and left quartz crystals 292 LIST OF ILLUSTRATIONS xix FIG. PAGE 114. Crystal of tridymite 292 115. Dialyser 294 116. Making hydrofluosilicic acid 298 117. Retorts for making phosphorus 305 118. Electric furnace for making phosphorus 306 119. Making phosphine from phosphorus and caustic alkali . . 309 120. Phosphine from calcium phosphide 310 121. Marsh test for arsenic 319 122. Curve of atomic weights and atomic volumes. (L. Meyer.) . 341 123. Hopper-shaped crystal of sodium chloride ..... 356 124. Acker process of making caustic soda 357 125. Salt cake furnace, Le Blanc soda process 358 126. Revolving black ash furnace 359 127. Solubility curve of sodium sulphate ...... 362 128. Spectroscope 366 129. Spectra of some common elements 367 130. Tube for examining spectra of gases 368 131. Spectra of gases 369 132. Absorption of spectrum of blood 370 133. Making metallic calcium 375 134. Common limekiln 377 135. Glass pots, open and closed form 384 136. Solubility curve of magnesium chloride ..... 395 137. Iron flask for shipping mercury 402 138. Solubility curves of various salts 413 139. Making colloidal silver . 417 140 and 141. Explanation of osmosis 419 142. Demonstration of osmotic pressure ...... 421 143. Simple osmometer 421 144. Pfeft'er's osmotic apparatus 422 145. Grotthus's theory of electrolysis 427 146. Electrolysis according to the theory of electrolytic dissociation . 429 147. An electric battery 433 148. Measuring the voltage of a cell 433 149. Gravity battery 434 150. Electrolytic production of aluminum 458 151. Blastfurnace 503 152. Bessemer converter 507 153. Solubility curve of ferric chloride 511 154. Pyrite crystal 512 155. Dobereiner's lamp 526 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 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 tlie 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- aifine, or one of salt and one of paraifine, 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 C. 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 Caesium Indium Phosphorus Thorium Calcium Iodine Platinum Thulium Carbon Iridium Potassium Tin Cerium Iron Praseodymium Titanium Chlorine 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 8 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 Oxygen Silicon ...... 47.07 28.06 7.90 85.79 49.78 26.08 7.34 Iron 443 411 3.44 0.05 3.19 Magnesium Sodium Potassium 2.40 2.43 2.45 0.22 0.14 1.14 0.04 10.67 2.24 2.33 2.28 0.95 0.40 0.37 Carbon ...... Chlorine 0.20 0.07 0.002 2.07 0.008 0.19 0.21 0.11 0.11 Sulphur 0.11 0.09 0.09 0.11 0.09 0.07 0.07 0.03 0.03 Nitrogen ...... Fluorine ...... All other elements .... 0.02 0.50 0.02 0.02 0.48 100.00 100.00 100.00 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 percent Calcium 1.5 per cent Phosphorus 1.0 percent 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 FIG. l. 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 with 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. 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, hydrogen 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. When an electric current is passed through water acidulated with 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 + 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 fire and burns with a brilliant flame. Frequently the action is so violent as to result in explosions. Lithium, rubid- ium, caesium, 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, hydrogen and potassium zincate result. The latter is a salt which remains in solution; thus: Caustic Potash + Zinc = Potassium Zincate 4- 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 dilute 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 + 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. 18 OUTLINES OF CHEMISTRY the last-named temperature. Liquid hydrogen is clear and 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. 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 E 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 FIG. 7, 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 --- -^ is set on fire, but the flame of the candle is extinguished. At room temperatures and atmospheric pressure hydrogen is rather inert chemically, combining with vigor with but one element ; namely, fluorine. But in the sunlight, hydro- gen becomes more active, notably toward chlorine, with which it combines readily, forming hydrochloric acid. With nitrogen, hydrogen forms ammonia ; with sulphur, hydrogen sulphide; with carbon, marsh gas and other hydrocarbons. A compound of hydrogen with another element is called a hydride. Thus, water may be termed a hydride of oxygen; it may also be called an oxide of hydrogen. Since hydrogen readily unites with oxygen at elevated tem- peratures, it may be used to deprive some compounds of their FIG. FIG. 10. HYDROGEN oxygen content. So when copper is heated in the air it turns black because of union of the surface layers with oxygen of the air, and when this hot, black copper oxide is now brought into an atmosphere of hydrogen, the latter gas unites with the oxygen, forming water and bright, metallic copper. The process of union of the copper with oxygen is termed oxidation, whereas the process of depriving the copper oxide of its oxygen content is called reduction. Oxidation and reduction are thus opposite processes. At higher temperatures, hydrogen is similarly able to reduce quite a number of oxides, such as the oxides of lead, iron, mercury, zinc, and nickel. On account of its power to abstract oxygen from compounds, hydrogen is called a reducing agent. The process of oxidation and reduc- tion of copper is readily illustrated by means of the apparatus shown in Figs. 11 and 12. The bright copper FIG. 11. FIG. 12. crucible is heated by means of the burner as shown in Fig. 11. The black oxide of copper is thus formed on the surface of the crucible. The flame is now extinguished, and while the crucible is still quite hot it is brought into an atmosphere of hydrogen. This is accomplished by holding the large funnel, through which a strong current of hydrogen is being passed, down over the crucible so as to envelop it (Fig. 12); the black crucible thus quickly assumes a bright copper color. 22 OUTLINES OF CHEMISTRY Uses. Besides being employed as a reducing agent in various chemical operations in the laboratory and the indus- tries, hydrogen finds use on account of its light- ness and combustibility. Its lightness makes it specially suitable for filling balloons. Hydrogen prepared by electrolysis and placed, under pres- sures of 100 to 150 atmospheres, in steel cylinders (Fig. 13), is now an article of commerce. Mixed with carbon monoxide, hydrogen forms an im- portant part of water gas, which is used for heating purposes. Hydrogen is too expensive for use in ordinary heating. It is at times employed as fuel in operations where very high temperatures are required, as in working platinum and other metals having a high melting point. Hydrogen Equivalents of Metals. Whether a metal liberates hydrogen from water itself or from various dilute acids, the amount of hydro- gen which a given weight of a certain metal will set free is the same. The amount of hydrogen which a definite weight of a given metal is able to liberate may be ascer- tained by means of the apparatus shown in Fig. 14. Water is first placed in the beaker B. The graduated tube A is placed into the water as shown, and by means of suction at the upper end of A, while the cock is open, the water is drawn up the tube till it is filled to a little above the cock, which is then closed. The tube A is then raised slightly and its lower end placed into the little crucible D, 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 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 o-c FIG. 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, w T e 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 temperature, 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 the rate with which a chemical reaction proceeds is proportional to the chemical affinity that comes into play. 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. 24 OXYGEN 25 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 oxj^gen, 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 26 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 - -?~V:^* FIG. 16. OXYGEN 27 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. 28 OUTLINES OF CHEMISTRY 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 ; the large calorie is 1000 times as large, i.e. it is the amount of heat required to raise 1000 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 substance 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 : PEE CENT IRON PER CENT OXYGEN PARTS OXYGEN TO 77.75 PARTS IRON (1) 72.38 (2) 69.96 (3) 77.75 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 29 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 are as follows: (1) lead suboxide, a black substance formed when lead is heated at its melting point in the air ; (2) lith- arge, a yellow powder formed when lead is very strongly heated in air ; (3) lead sesquioxide, an orange-yellow powder formed when bleaching powder acts on litharge dissolved in caustic potash ; (4) red lead or minium, a bright red powder, which may be obtained by heating litharge in the air at a temperature not above 450; and (5) lead peroxide, a brown powder, which may be prepared by treating red lead with dilute nitric acid. The percentage composition of these oxides is as follows : NAME PER CENT LEAD PER CENT OXYGEN PARTS LEAD TO 3.72 PARTS OXYGEN (1) Lead suboxide . . .' (2) Litharge . 96.28 92 82 3.72 7.18 96.28 48.14 (3) Lead sesquioxide . . (4) Red lead . . 89.61 90 65 10.39 9.35 32.09 36.11 (5) Lead peroxide . . . 86.60 13.40 24.07 In the last column we have the amount of lead combined with 3.72 parts of oxygen in each of the oxides. Comparing the figures in the last column we note as follows : (1) and (2) (1) and (3) (1) and (4) (1) and (5) 96.28:48.14 = 2:1, 96.28:32.09 = 3:1, 96.28:36.11 = 8:3, 96.28:24.07 = 4:1. 30 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 31 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 oxy hydrogen 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 32 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- 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 notion dates back to antiquity. It was probably suggested by H OXYGEN 83 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 34 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 35 ers of the phlogistic theory. However, facts began to increase in favor of Lavoisier's explanations, and 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. 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 WATER 37 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 with 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 38 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. t 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 39 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 not 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- 40 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, and 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 41 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. 42 OUTLINES OF CHEMISTRY num terminals, across which an electric spark may be passed by connecting with the induction coil 0. 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 vapor 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 two 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 chemistry, 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 43 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 44 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 by 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 45 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 crystals 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 46 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 47 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 48 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. 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 49 50 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 FIG. 23. limbs A and B is kept constant by pouring HYDROCHLORIC ACID AND CHLORINE 51 mercury into B as required, it will be seen that wheu further addition of sodium no longer produces any action in A, the hydrogen obtained occupies just one half of the vocume 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, for 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, 52 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 ly 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 53 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. 54 OUTLINES OF CHEMISTRY Properties of Chlorine. Chlorine is a greenish yellow gas which is 2.5 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, FIG. 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 55 Chlorine is soluble in water. At 10 1 volume of water absorbs about 8 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 27 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 56 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- Iight 4 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 red 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 57 1 part of hydrogen to 8 parts of oxygen by weight. In hydro- chloric acid 1 part of h} r drogen 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 0, are able to unite to form chemical compounds with one another, the proportions by weight with which A. and unite to form the compound AB, and the proportions in which A and O unite, also determine the proportions in which B and O 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. CHAPTER VI THE LAWS OP 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 with 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 68 FUNDAMENTAL LAWS AND THEORIES 59 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, Lead : Oxygen : : 1 : 0.0387 () In the yellow oxide, Lead : Oxygen : : 1 : 0.0773 ( 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 + O 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 77 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 vapo-r 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 formulae 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 T8 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 Aluminum Mg Al 24.32 27.1 0.245 0.214 5.95 5.80 Phosphorus Sulphur Potassium P S K 31.0 32.07 39.10 0.202 0.203 0.166 6.26 6.50 6.49 Iron Fe 55.85 0.112 6.26 Copper Zinc Cu Zn 63.57 65.37 0.095 0.093 6.04 6.08 Silver Platinum Ag Pt 107.88 195.2 0.057 0.0325 6.15 6.34 Gold Au 197.2 0.0324 6.40 Mercury- Lead Hg Pb 200.6 207.1 0.0333 0.0315 6.68 6.52 Glucinum Gl 9.1 0.42 3.82 Boron B 11.0 0.24 2.64 Carbon (Graphite) Silicon C Si 12.0 28.3 0.200 0.177 2.40 5.01 FUNDAMENTAL LAWS AND THEORIES 79 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, 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 riot crystal- lize, that is, are non-crystalline, are called amorphous. All crystals may be classified into six crystal systems (which see). 80 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 formula 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 formulae 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 81 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 Ijave 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 in 1912 : INTERNATIONAL ATOMIC WEIGHTS, 1912 ELEMENT SYMUOL ATOMIC WEIGHT ELEMENT SYMBOL ATOMIC WEIGHT Aluminum Al 27.1 Neodymium Nd 144.3 Antimony Sb 120.2 Neon Ne 20.2 Argon A 89.88 Nickel Ni 58.68 Arsenic As 74.96 Niton (radium Barium Ba 137.37 emanation) Nt 222.4 Bismuth Bi 2080 Nitrogen N 14.01 Boron B 11.0 Osmium Os 190.9 Bromine Br 79.92 Oxygen 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 r>y 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 Sodium 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 Hydrogen H 1.008 Thallium Tl 204.0 Indium In 114.8 Thorium Th 232.4 Iodine I 126.92 Thulium Tm 168.5 Iridium Ir 193.1 Tin Sn 119.0 Iron Fe 55.84 Titanium Ti 48.1 " Krypton Kr 82.92 Tungsten W 184.0 Lanthanum La 139.0 Uranium U 238.5 Lead Pb 207.10 Vanadium V 51.0 Lithium Li 6.94 Xenon Xe 130.2 Lutecium Lu 174.0 Ytterbium Magnesium Manganese Mg Mn 24.32 54.93 (Neoytterbium) Yttrium Yb Yt 172.0 89.0 Mercury Hg 200.6 Zinc Zn 65.37 Molybdenum Mo 96.0 Zirconium Zr 90.6 82 OUTLINES OF CHEMISTRY Interpretation of a Chemical Formula. A chemical formula expresses (1) what elements occur in 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 83 is treated with zinc, as indicated by the equations (compare Chapter II) H-O-H+ Na = NaOH + H. Water + sodium = sodium hydroxide -h hydrogen. 2 NaOH + Zn = Na 2 O 2 Zn + H 2 . Sodium hydroxide + zinc = sodium zincate + 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 Cl ; 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 = Cl = 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 = C = 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 84 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 urn, like lithium, sodium, bariwm, 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 Tiydroxyl 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 + 2H 2 = Ca(OH) 2 +H 2 ; or when lime, calcium oxide CaO, is treated with water, i.e. is slaked, thus : CaO + H 2 O = 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 num- FUNDAMENTAL LAWS AND THEORIES 85 ber 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 PC1 6 is phosphorus pentachloride or phosphoric chloride. When more than two compounds are formed by two elements, the endings ous and ic are retained and the prefixes proto, hypo 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 swfoxide 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 86 OUTLINES OF CHEMISTRY H 2 SO 4 + Zn = ZnSO 4 + H 2 . Sulphuric acid -f zinc = zinc sulphate + hydrogen. 2KOH + Zn = K 2 2 Zn + H 2 . Potassium hydroxide + zinc = potassium zincate 4- hydrogen. 3KOH + Al = K 3 3 A1 + 3 H. Potassium hydroxide -f- aluminum = potassium aluminate-h 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 4. O. Mercuric oxide (on heating) = mercury -f 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 3 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 2 + 2 =2H 2 0. Mg + O = MgO. Cu + O = CuO. phosphorus pentoxide. C + 2 =C0 2 . 3 Fe + 2 O 2 = Fe 3 O 4 . 2Fe + 3O =Fe 2 O 3 , ferric oxide or sesquioxide of iron S + O 2 = SO a . FUNDAMENTAL LAWS AND THEORIES 87 (4) Reductions CuO + H 2 = Cu + H 2 0. Fe 8 O 4 + 4 H 2 = 3 Fe + 4 H 2 O. (5) Preparation of Chlorine MnCl 2 + manganous chloride. 4HCl + MnO 2 = MnCl 2 (6) Reactions of Chlorine H 2 0+ C1 2 =2HC1 + O. P + 3 Cl = PC1 8 , phosphorus trichloride. P + 5 Cl = PC1 6 , phosphorus pentachloride. Sb + 3 Cl = SbCl 3 , antimony trichloride. Cio H i8 + 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. CHAPTER VII OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE History, Occurrence, and Preparation of Ozone. When a fric- 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 89 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 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) 5* 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. 90 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 poiver. 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 91 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 : Ba0 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 nitration ; and thus a nitrate, 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. Insj^ad 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. 92 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 = PbS0 4 + 4 H 2 0. Potassium iodide in solution is oxidized thus : 2KI + H 2 2 =2KOH-i-I 2 . OZONE, ALLOTROPY, AND HYDROGEN PEROXIDE 93 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 limewater 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. While 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 + 2 . 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 + H 2 O + O 2 . Added to a potassium permanganate solution acidified with sul* 94 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 TT\ 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 95 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 : O 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. CHAPTER VIII THE HALOGENS The Halogen Family. The elements that belong to this group are fluorine, 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 bp 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- curic oxide : 2 HgO + 2 C1 2 = HgO HgCl 2 + C1 2 O. 96 THE HALOGENS 97 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 -f H 2 SO 4 = KHSO 4 -f HC1O 3 . chloric acid (2) 3 HC10 3 = HC10 4 + H 2 + 2 C1O 2 . 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 O + C1 2 7 . Chlorine heptoxide is a colorless oil which boils 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. 98 OUTLINES OF CHEMISTRY 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 2 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 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 134 OUTLINES OF CHEMISTRY 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 all the other factors very greatly, and so the action goes on in one ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 135 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 136 OUTLINES OF CHEMISTRY 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: H 2 O + C1 2 ^HOC1 + 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 ACIDS, BASES, SALTS, CHEMICAL EQUILIBRIUM 137 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 S0 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. 138 OUTLINES OF CHEMISTRY 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). CHAPTER X NITROGEN, THE ATMOSPHERE, AND THE ELEMENTS OP 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. 139 140 OUTLINES OF CHEMISTRY 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. 34), phosphorus pentoxide is formed, which is a solid that is readily absorbed by water, forming phosphoric acid : P 4 + 5 O 2 = 2 P 2 O 6 , and 3 H 2 + P 2 5 = 2 H 3 P0 4 . 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 NITROGEN, AIR, AND THE HELIUM GROUP 141 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 : 2 NH 3 + 3 C1 2 = 6 HC1 + 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 Cr 2 7 = Cr a O, + 4 H 2 O + N 2 ; or, by combining the two equations, K 2 Cr 2 7 + 2 NH 4 C1 = 2 KC1 + Cr a O 8 + 4 H 2 O + N 2 . 142 OUTLINES OF CHEMISTRY 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 2 . The potassium hypobromite solution as usually prepared con- tains an excess of caustic potash, which at once absorbs the carbon dioxide, forming potassium carbonate, which dissolves in water: 2 KOH + C0 2 = K 2 C0 3 + H 2 0. 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 than 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 8 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 . NITROGEN, AIR, AND THE HELIUM GROUP 143 FIG. 35. Due to electrical disturbances in the atmosphere, especially during thunder storms when lightning flashes from cloud to 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 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 144 OUTLINES OF CHEMISTRY 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 of 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 NITROGEN, AIR, AND THE HELIUM GROUP 145 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 + C0 2 =BaC0 3 + H 2 0. 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. L 146 OUTLINES OF CHEMISTRY Rain 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 and 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 ; arid 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 met in crowded, poorly ventilated rooms. NITROGEN, AIR, AND THE HELIUM GROUP 147 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 148 OUTLINES OF CHEMISTRY 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 yellow line in its spectrum. This line had previously been observed in the spectrum of the sun by Lockyer, who ascribed it to an element which he called helium, then unknown on the earth. Helium does not unite with any other element. Its molecular weight is 3.99 and its atomic weight is the same. About 1.4 cc. of helium are dissolved by 100 cc. of water at room temperature. Helium was liquefied in 1908 by Professor NITROGEN, AIR, AND THE HELIUM GROUP 149 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 . . = 0.0002T 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 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 150 AMMONIA AND OTHER NITROGEN COMPOUNDS 151 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 : KN0 3 + 8 H = KOH + 2 H 2 O + NH 3 . KN0 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 ammonia, 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- erystallization. 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 Kjeldahl's 152 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 irito the tube during the experiment. We thus see that 2 volumes of ammonia yield 1 volume FIG" 36 f 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 AMMONIA AND OTHER NITROGEN COMPOUNDS 153 Avogadro's hypothesis, equal volumes of gases contain an equal ftumber of molecules, hence : 3 molecules hydrogen + 1 molecule nitrogen = 2 molecules ammonia or 3 H 2 + N 2 = 2 NH 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. 154 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 * FIG. 38. 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 3 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 : NH 3 + HC1 = NPI 4 C1. 2 NH 3 + H 2 SO 4 = (NH 4 ) 2 S0 4 . NH 3 + HNO NH 4 NO 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 is called ammonium, the ending um being used to indicate that chemically it is analogous to a metal. When ammonia dis- AMMONIA AND OTHER NITROGEN COMPOUNDS 155 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 = 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 ammonium 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 4- 2 NH 3 = 2 NaNH 2 + H a . 2 K + 2 NH 3 = 2 KNH 2 + H 2 . 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, + H 2 = KOH + NH a . 156 OUTLINES OF CHEMISTRY The nitrides of lithium NLi 3 and of magnesium N 2 Mg 3 may 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 3 + 3 H 2 O = 3 LiOH + NH 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 157 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 8 (OH), which melts at - 40 and boils at 120. Its 158 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 Lessen 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 8 (OH)C1. 2 NH 2 OH + H 2 S0 4 = (NH 8 OH) 2 S0 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 + 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 159 Hydroxylamine readily decomposes into ammonia, nitrogen, and water, thus : 3 NH 2 OH = NH 8 + 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 O = H 2 O + NaN 3 . 2 NaN 8 + H 2 S0 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 : 2N 3 H = 3N 2 +H 2 . 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 : NH 3 + (N 3 )H = NH 4 (N 8 ). The empirical formula of NH 4 (N 3 ), 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 + 3 Cl a = 4 HC1 + NC1 8 . The compound may be prepared by the electrolysis of an aque- ous solution of ammonium chloride, the chlorine liberated acting 160 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 + N a . 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 T 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 8 with a solution of iodine in ether, triazoiodide IN 3 may be formed : It is a yellow powder of a very penetrating odor, and is ex- tremely explosive. CHAPTER XII OXY-ACIDS AND OXIDES OP 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. M 161 162 OUTLINES OF CHEMISTRY By treating a nitrate like potassium or sodium nitrate with concentrated sulphuric acid, nitric acid is liberated, thus : NaN0 3 + H 2 S0 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 HNO 3 = 2 H 2 O 4 NO O 2 . The nitrogen dioxide forms reddish brown fumes that dissolve in the nitric acid. This solution, which fumes strongly in the OXY-ACIDS AND OXIDES OF NITROGEN 163 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 : 8 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 + 2 ; 2 N0 2 . H 2 + 3 N0 2 5 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 164 OUTLINES OF CHEMISTRY 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 iiitro 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 OXY-ACIDS AND OXIDES OF NITROGEN 165 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 , hydroxylainine NH 2 OH, and ammonia NH 3 . 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 + HN0 3 = 2 H 2 + NOC1 + C1 2 . 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 6 = 2 HP0 8 + N 2 6 . 166 OUTLINES OF CHEMISTRY 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 8 + 2 C1 2 = 4 AgCl + 2 N 2 O 6 + O r 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 5 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 HNO 3 + 6 FeSO 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 OXY-ACIDS AND OXIDES OF NITROGEN 167 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 . . . . ^ .-n,. FlG. 41. 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 : 168 OUTLINES OF CHEMISTRY 2 Pb(N0 3 ) 2 = 2 PbO + 2 + 4 N0 2 . 2 Cu(N0 3 ) 2 = 2 CuO + 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 O 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.) OXY-ACIDS AND OXIDES OF NITROGEN 169 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 KNO 3 = 2 KNO 2 + O 2 . The nitrite may also be formed by heating saltpeter with lead or copper, thus : KN0 3 + Pb = PbO + KNO 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 3 , and water, thus : 2 HN0 2 = H 2 + N 2 3 . The latter then decomposes into nitric oxide NO and nitrogen N a 3 = NO + NO 2 . dioxide NO 2 , thus : 170 OUTLINES OF CHEMISTRY 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 8 . On the other hand, toward substances that will take up oxygen, nitrous acid plays the rdle of an oxidizing agent. So with hydriodic acid, the following reaction occurs : 2 HN0 2 + 2 HI = 2 NO + 2 H 2 O + I a . 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 OXY-ACIDS AND OXIDES OF NITROGEN 171 sodium amalgam on the aqueous solution, a salt of hyponitrous acid may be obtained thus : 2 KNO a + 4 H = K 2 N 2 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 2 + H 2 S0 4 = K 2 S0 4 + H 2 O + N 2 O. The reaction is not reversible, and so hyponifrous 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 O 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 NO 3 =2 H 2 O + N 2 O. 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. 172 OUTLINES OF CHEMISTRY Nitrous oxide is much more soluble in cold than in warm water; so at 0, 1.30 volumes are absorbed by 1 volume of 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 : OXY-ACIDS AND OXIDES OF NITROGEN 173 In ammonium salts it has a valence of five, thus : C1 In hydroxylamine, in hydrazine, in hydrazoic acid, and in nitrous oxide, nitrogen is trivalent, thus : ,H ft .H \ = N = N N^H ; NN-N/ ; X N/ ; \ o / . X OH W X 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=0 ^O V), and N=O 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 : >o This formula is not generally accepted, however. 174 OUTLINES OF CHEMISTRY In nitrogen trioxide and nitrous acid nitrogen is trivalent, thus : .O O 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 \o/' : ||| 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 wjiich 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 5 . 2N(OH) 4 minus4H 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 OXY-ACIDS AND OXIDES OF NITROGEN 175 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 O 2 C1 2 O . . . HC1O . . . NaClO. NO N 2 O 3 . . . HNO 2 NaNO 2 NaClO 2 . NO 2 . .:V ; ', . . . . C10 2 . N 2 5 . . . HN0 3 NaNOg I 2 O 6 . . . HC1O 3 . . . NaClO 3 . C1 2 7 ..HClO 4 ...NaC10 4 . 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 uneom- 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 algse 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. 176 SULPHUR, SELENIUM, AND TELLU-RIUM 177 Sulphur further occurs as hydrogen sulphide in the waters of 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 a . Fia. 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. The amount of gypsum produced in the United States alone in 1910 was 2,379,057 tons. Sulphur occurs in small quantities in com- 178 OUTLINES OF CHEMISTRY 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 is com- monly 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 orthorhombic 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 FlG - 43 - 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 sulphur 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, SULPHUR, SELENIUM, AND TELLURIUM 179 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 orthorliombic 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 two 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 -f 2 HC1 = 2 KC1 -h 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 180 OUTLINES OF CHEMISTRY explained by assuming that the molecules of the different modi- fications contain a different number of atoms, similar to the case of oxygen 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 3 , 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 2 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. We SULPHUR, SELENIUM, AND TELLURIUM 181 FIG. 44. FIG. 45. FIG. 46. FIG. 47. FIG. 48. FIG. 49. FIG. 50. FIG. 51. FIG. 52. ^x ^ PV ^ f\ / > t i \x / NJl V FIG. 53. FIG. 54. 182 OUTLINES OF CHEMISTRY 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. 56. FIG. 57. FIG. 58. FIG. 59. FIG. GO. 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 SULPHUR, SELENIUM, AND TELLURIUM 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 184 OUTLINES OF CHEMISTRY 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 Ca 3 (PO 4 ) 2 . The so-called rhombohedral FIG. FIG. 69. 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. 73. 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 SULPHUR, SELENIUM, AND TELLURIUM 185 FIG. 75. FIG. 76. 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 monoclinic 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 3 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 3 , 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- 186 OUTLINES OF CHEMISTRY sented 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 SO 4 = FeSO 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 hydro- 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 : 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 SULPHUR, SELENIUM, AND TELLURIUM 187 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 0. 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 : K 2 S + H 2 0:KSH + KOH. 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 HNO 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 188 OUTLINES OF CHEMISTRY 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 6 , 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 5 . 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 M OH, 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 5 and P 2 S 5 , Sb 2 O 3 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 other reagents as nitrogen. SULPHUR, SELENIUM, AND TELLURIUM 189 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 = 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 tetrachloride 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 T 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. 190 OUTLINES OF CHEMISTRY 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: PbO 2 -j-SO 2 =PbSO 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 8 + 4 SO 2 . In the laboratory, sulphur dioxide is commonly made by heating copper turnings with concentrated sulphuric acid : 2 H 2 S0 4 + Cu = 2 H 2 + CuS0 4 + SO 2 . It may also be formed by heating concentrated sulphuric acid with carbon or sulphur : 2 H 2 SO 4 + C = 2 H 2 O + CO 2 + 2 SO 2 . 2 H 2 SO 4 + S = 2 H 2 O + 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 + H 2 O. 2 MnO 2 + 4 $ = 2 MnS + 2 SO 2 . 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 SULPHUR, SELENIUM, AND TELLURIUM 191 does not depend upon the oxidation of the dyes, but rather 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 : S0 H 2 = H 2 S0 3 . With bases, sulphurous acid forms salts called sulphites, thus : H 2 SO 3 + NaOH = NaHSO 3 + H 2 O. HS0 + 2 NaOH = NaSO + 2 HO. H 2 S0 3 + Ca(OH) 2 = CaS0 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 192 OUTLINES OF CHEMISTRY oxygen from the air slowly converts sulphurous acid into sul- phuric acid in solution, thus : 2 H 2 S0 8 + 2 = 2 H 2 S0 4 . Chlorine, bromine, or iodine rapidly change sulphurous acid into sulphuric acid, thus : H 2 S0 3 + H 2 + C1 2 = H 2 S0 4 + 2 HC1. H 2 S0 3 + H 2 + I a = 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 : 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 : 2SO 2 +O 2 ^2SO 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 SULPHUR, SELENIUM, AND TELLURIUM 193 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 It; 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 194 OUTLINES OF CHEMISTRY 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 SO 4 + P 2 O 6 = SO 3 + 2 HPO 3 . K 2 S 2 O 7 = K 2 SO 4 + SO 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 bsy 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( \d and its decomposition products. SULPHUR, SELENIUM, AND TELLURIUM 195 In practice, the manufacture of sulphuric acid loy 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 S0 2 + 2 H 2 + 2 HN0 3 = 3 H 2 SO 4 + 2 NO. (3) 2 NO + H 2 + 30 = 2 HNO 3 , and (4) NO + O = N0 2 . (5) S0 2 + H 2 + N0 2 = H 2 SO 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 196 OUTLINES OF CHEMISTRY acid HO SO 2 - OH. The reactions involved in this explanation are : (1) SO 2 + HN0 3 = 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 Cr. 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, SULPHUR, SELENIUM, AND TELLURIUM 197 198 OUTLINES OF CHEMISTRY 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, SULPHUR, SELENIUM, AND TELLURIUM 199 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 200 OUTLINES OF CHEMISTRY 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. SULPHUR, SELENIUM, AND TELLURIUM 201 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 arid 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 T 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 ^l 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 : 202 OUTLINES OF CHEMISTRY 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 O 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. Persulphates. 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 was 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 Q , forms trithionates, like Na 2 S 3 O 6 . Tetrathionic Acid H 2 S 4 O 6 , forms tetrathionates, like Na 2 S 4 O 6 . Pentathionic Acid H 2 S 5 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 SULPHUR, SELENIUM, AND TELLURIUM 203 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 2 SO 3 = 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 + 2 HC1. 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 : S0 2 +C1 2 = S0 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 : S0 2 C1 2 + H 2 = S0 2 . Cl - OH + HCL 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 .C1.0H + H 2 = S0 2 (OH) 2 + HCL 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, 204 OUTLINES OF CHEMISTRY 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. SULPHUR, SELENIUM, AND TELLURIUM 205 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 Se0 3 + 2 S0 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 2 . 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 SeO 3 + H 2 O + C1 2 5* 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 M tiller von Reich- enstein, whose work was confirmed by Klaproth in 1798. The latter called the element tellurium, from tellus, earth. 206 OUTLINES OF CHEMISTRY 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 3 . 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 3 and K 2 TeO 3 . However, towards strong acids it behaves like a base. The salts thus formed SULPHUK, SELENIUM, AND TELLURIUM 207 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 2 = 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 + 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 tellurates 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 208 OUTLINES OF CHEMISTRY 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 Oxygen blue 16.0 1.124 (at - 181) -181.4 Sulphur yellow 32.07 1.96 to 2.0 114.5 450 Selenium red or 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 0, 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) - H /0-K /0-K 84=0 , Se^O , Te==0 \0 - H \0 - H X) - H In the compounds H 2 XO 4 , the elements are hexavalent, thus : 0-H O-H O-H O-H. )-H SULPHUR, SELENIUM, AND TELLURIUM 209 Toward the halogens, sulphur, selenium, and tellurium are bivalent and quadrivalent, while in some oxy -halogen deriva- tives they are hexavalent, thus : Cl Cl Cl /Cl s ^c! Se 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 the tip of the cone near B are the hottest parts of the flame. FIG. 108. 274 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 which 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 CH 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. CHAPTER XVII 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 va.por 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 gram 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 acid, 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 275 276 OUTLINES OF CHEMISTRY 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 THERMOCHEMISTRY 277 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- taneousty, 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. 278 OUTLINES OF CHEMISTRY 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 tlie 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 THERMOCHEMISTRY 279 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., 280 OUTLINES OF CHEMISTRY 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 grams of red phosphorus are converted into phosphorus trichloride, we have : [P] red + 3 Cl = (PC1 8 ) + 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 THERMOCHEMISTRY 281 chlorine is greater than that contained in mercuric chloride by 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) = (NaClaq) - 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 NaClaq, 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 terra 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. 282 OUTLINES OF CHEMISTRY 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 2 = 12 C0 2 + 11(H 2 0) + 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 THERMOCHEMISTRY 283 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: 284 OUTLINES OF CHEMISTRY 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 PI, Cl 22.0 39.3 HBr H, (Br) 86 28.6 HI H, [I] -6.1 13.1 H 2 O H 2 ,0 55.3 68.4 69.8 HA H 2 ,0 2 45.3 H 2 2 (H 2 0), -23.1 H,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 3 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 o, 2 -30.7 HC10 H, Cl, 31.65 HC1O 3 H, Cl, O 3 24.0 HC10 4 H, Cl, 4 18.3 38.6 HBrO 3 H, (Br) 3 12.3 HI0 3 H, [I], 3 57.9 55.7 HTO 4 H, [I], 4 - 47.6 SO 2 rhombic [S], O 2 71.0 78.8 S0 3 rhombic [S] , O 3 103.3 142.5 H 2 SO 4 H 2 , [S], 4 189.9 210.9 H 2 S 2 3 H 2 , [S], 3 141.7 N,0 N 2 ,0 -17.4 NO N, O -21.5 NA N 2 ,0 3 6.8 NO 2 N,0 2 -7.7 N 2 4 N 2 ,0 4 -2.6 N 2 5 N 2 ,0 6 . 13.1 29.8 HNO 8 H, N, 3 34.4 41.5 42.2 48.8 PA yellow [P 9 ],O 5 370.0 406.0 H 3 F0 4 H 3 , [P], 4 304.1 306.8 As 2 3 [As 2 ], 3 154.7 147.0 Sb 2 O 3 [SbJ, 3 166.9 Bi 2 3 [Bi 2 ], 3 139.2 BA amorph. [B 2 ], O 3 272.6 SiO 2 aq cryst. [Si], O 2 , 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 THERMOCHEMISTRY 285 TABLE 1 Continued COMPOUND FORMED FROM GASEOUS LIQUID SOLID DISSOLVED CO 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 s [Sb], Cl, 91.4 BiC] 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 SiiCl 4 [Sri], C1 4 122.2 129.8 (CN) 2 diamond [C 2 ], N, -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],O, H 103.2 116.5 NH 4 C1 N, H 4) Cl 75.8 71.9 NaOH [Na], 0, H 101.9 111.8 CaCl 2 [Ca], C1 2 169.8 187.2 LiOH [Li], O, H 112.3 118.1 ZnCl 2 [Zn], C1 2 97.2 112.8 NH 3 aq NH 3 , w(H 2 0) 20.3 A1C1 8 [Al], C1 8 161.0 237.8 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], O a , 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], 124.2 158.7 HgCl 2 (Hg), C1 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 MnO 2 [Mn], 2 125.3 AuCl 3 [Au], C1 3 22.8 27.3 FeO [Fe], 65.7 PtCl 4 [Pt], C1 4 59.8 79.4 Fe 3 4 [Fe 3 ],0 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], 3 269.9 275.4 Cu 2 O [Cu 2 ], 40.8 NaHC0 3 [Na],H,[C],0 8 227.0 223.7 PbO [Pb], 50.3 Na 2 S0 4 [Na 2 ], [S], 4 328.4 329.0 Pb0 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],[8],H 62.3 63.1 KC10 3 [K], Cl, 3 95.0 85.0 CaS [Ca], [S] 89.6 KBrO 3 [K], (Br),0 3 84.1 74.3 SrS [Sr], [S] 97.4 KI0 3 [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 CuS0 4 [Cu], [S], 4 1828 198.4 NaCl [Na], Cl 105.6 101.2 BaSO 4 [Ba], [S], 4 338.1 286 OUTLINES OF CHEMISTRY 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 FOEMULA 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 C 3 H 8 529.2 25.4 Benzene C 6 H 6 787.8 -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 9 2677.8 227.6 Starch c 6 H 10 o; 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 12 6 677.2 302.6 Lsevulose C C 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 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. THERMOCHEMISTRY 287 TABLE 5 HEATS OF NEUTRALIZATION VALUES GIVEN IN LARGE CALORIES. (THE SOLUTIONS CONTAINED 1 GRAM EQUIVALENT OF ACID OR BASE IN Two LITERS. SOME OF THE BASES USED AND SULPHATES FORMED ARE INSOLUBLE.) BASES HClaq HN0 3 aq CH 3 COOH aq |H 2 S0 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 1 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) 2 aq 13.8 13.8 15.6 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 JZn(OH) s 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, hydrfodic 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 2 288 OUTLINES OF CHEMISTRY 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 THERMOCHEMISTRY 289 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. 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 th#n 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 : 4Na = 4NaCl+Si. 3 SiCl 4 + 4 Al = 4 A1C1 3 + 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 = 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 290 SILICON AND BORON 291 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. In 1908, 500 tons of 90 per cent silicon were used 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 + 4F=SiF 4 . Hot solutions of caustic potash dissolve silicon : 2 KOH + H 2 + Si = K 2 Si0 3 + 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. 292 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 Fia. 113. , degree of rotation is propor- tional to the thickness of the layer traversed, and the deviation is either dextro or leevo 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 FIG. 114. SILICON AND BORON 293 of temperature. This is due to the 'fa.ct 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 : v 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 7 , H 4 Si 2 6 , H 2 Si 2 6 . 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 294 OUTLINES OF CHEMISTRY H 4 SiO 4 ; sodium silicate, or water glass, Na 2 SiO 3 is a salt of 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 CO 3 + SiO 2 = Na 2 SiO 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 295 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, 296 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 by 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 297 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 + 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 SO 4 = 2 HF + CaSO 4 , and SiO 2 + 4 HF = 2 H 2 O + SiF 4 ; or 2 CaF 2 + Si0 2 + 2 H 2 SO 4 = 2 CaSO 4 + 2 H 2 O + SiF 4 . Silicon tetrafluoride is a colorless gas of ve/y pungent odor. It boils at -65, and the solid melts at --77 . The tetra- fluoride is always formed when hydrofluoric jacid acts on sili- cates, and it is consequently produced when; that acid is used in etching glass. / Water decomposes silicon tetrafluoride : 3 H 2 O + 3 SiF 4 = H 2 Si0 3 + I H 2 SiF 6 . I The silicic acid formed separates out as i* gelatinous precipitate, while the hydrofluosilicic acid H 2 SiF 6 remains in solution. The latter may be concentrated to som extent by evaporation. The concentration must be carried on ima platinum dish, because hydrogen fluoride is formed during the process, and so glass or porcelain dishes would be attacked. Pure 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- 298 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 FIG. 116. 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 ar^d silica, thus : Si + 2 C1 2 = SiCl 4 , or SiOj + 2 C + C1 4 = 2 CO + SiCl 4 . The product is 1 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 t SiCl 4 -P 4 H 2 = 4 HC1 + H 4 Si0 4 . On heating silicon jn 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 ? vater. Bromine and iodine compounds of silicon, analogous to the chlorine compounds, have been prepared by similar methods. Silicon tetrabrpmide boils at 153 and melts at - 12, silicon tetraiodide SiI 4 forms octahedra that melt at 120 and boil at 290. SILICON AND BORON 299 Esters of Silicic Acid. Methyl silicate (CH 3 ) 4 SiO 4 , boiling at 121, and ethyl silicate (C 2 H 5 ) 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 : 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 300 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 301 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 7 + 5 H 2 + 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 4 O 7 . The latter on igni- 302 OUTLINES OF CHEMISTRY tion forms the trioxide or boric anhydride B 2 O 3 , 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 5 ) 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 4H 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 from the deposits in the United States in 1910 was 42,357 tons. Borax solutions have a slightly alkaline reaction toward indicators, which is explained by the fact that boric acid is weak arid 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 303 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 3 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 8 + 6 H 2 = 2 B(OH) 3 + 3 H a S. 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 -h Ca(ClF), though it is at times also found as wavellite 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 42 H 86 NPO 3 . 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 are formed by means of the following reaction : Ca 3 (P0 4 ) 2 + 3 H 2 S0 4 = 3 CaSO 4 + 2 H 3 PO 4 . The calcium sulphate is insoluble, while the phosphoric acid remains in solution and is drained off. This solution is evapo- 304 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 305 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 : Ca 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 (PO 4 ) 2 = 2 H 2 O + Ca(PO 3 ) 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 = 2 CaSiO 3 + 10 CO + 4 P. 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 : Ca 3 (P0 4 ) 2 + 3 Si0 2 + 50 = 3 CaSiO 3 + 5 CO + 2 P. 1 I 1 1 1 1 1 1 1 FIG. 117. 306 OUTLINES OF CHEMISTRY 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 corn- ea, us. 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 307 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- 308 OUTLINES OF CHEMISTRY teenth 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 matches 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 011 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 309 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 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 310 OUTLINES OF CHEMISTRY to prevent its clogging by phosphorus that mi ht disti11 over 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 3 P 2 + 6 H 2 = 3 Ca(OH) 2 + 2 PH 3 , or Ca 3 P 2 + 6 HC1 = 3 CaCl 2 + 2 PH 3 . 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 3 P0 3 = 3 H 3 P0 4 + PH 3 , or phosphorous acid phosphoric acid 2 H 8 P0 2 = H 3 P0 4 + PH 3 . hypophosphorous acid phosphoric acid When phosphonium iodide PH 4 I is treated with caustic alka- lies, phosphine is formed, thus : PH 4 I + NaOH = Nal + H 2 O + PH 3 . Water also decomposes phosphonium iodide : PH 4 I + H 2 O = HI + H 2 O + 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 + HI = PH 4 I ; which is analogous to PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 311 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 H 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 formulae 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 = 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 8 +C1 2 = PC1 6 , or = 4PC1 5 . The product is a light yellow, finely crystalline solid which can- not be 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 : 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: 312 OUTLINES OF CHEMISTRY With water, phosphorus pentachloride forms hydrochloric acid and phosphorus oxychloride : PC1 5 + H 2 O = 2 HC1 + POC1 8 . 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 3 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 6 forms yellow crystals, which on heating disso- ciate into bromine and phosphorus tribromide. Phosphorus triodide PI 3 forms dark red, prismatic crystals melting at 61. Phosphorus pentaiodide is not known, but a diphosphorus tetraiodide P 2 T 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 hydriodio 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 2 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 t^e 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 313 in a sealed tube to 440. Phosphorus pentoxide P 2 O 5 * s 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 : P 2 5 + H 2 = 2HP0 3 . 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 5 + H 2 O = 2 HPO 3 (metaphosphoric acid), P 2 O 5 -f- 2 H 2 O = H 4 P 2 O 7 (pyrophosphoric acid), P 2 Q 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 P0 2 ) 2 + H 2 S0 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) 6 by loss of one molecule of water. Pure phosphoric acid is prepared by action 314 OUTLINES OF CHEMISTRY 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. 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 + H 2 O. H 3 P0 4 + 2 KOH = K 2 HP0 4 + 2 H 2 O. H 3 P0 4 + 3 KOH = K 3 P0 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 decom- 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 : 2Na 2 HPO 4 :Na 4 P 2 7 + H 2 0; and sodium pyrophosphate NaH 2 P0 4 ^NaP0 3 + H 2 O. sodium metaphosphate PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 315 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 HP0 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 : 2 H 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 + HPO 3 ; or by treating phosphorus pentoxide with water; 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 316 OUTLINES OF CHEMISTRY 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 : 4H 8 P0 3 =3H 3 P0 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 : H orthophosphoric Q_H acid. \0-H O-H \ , , . pyrophosphoric metaphosphoric \)_ H acid ' O p>/H phosphorous \\0-H acid. \O-H /O-H P-O-H Nvri ^ hypophosphoric p//O-H acid. \ hypophosphorous acid. O-H PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 317 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 3 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 5 forms yellow crystals which melt at 275. The liquid boils at 518. With potassium sulphide it forms potassium sulphophosphate : With phosphorus pentachloride phosphorus sulphochloride PSC1 results: The latter compound is a colorless liquid of specific gravity 1.168 at 0. It boils at 125, and decomposes upon treatment with water : PSC1 8 + 4 H 2 O = 3 HC1 + H 2 S + H 3 PO 4 . Occurrence, Preparation, and Properties of Arsenic. Arsenic is very widely distributed in nature in minute quantities. It rarely occurs 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. 2 As 2 O 3 + 6 = As 4 + 6 CO. 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 318 OUTLINES OF CHEMISTRY 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 560 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 3 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. The reactions involved in these processes are typified by the follow- ing equations : (1) Zn 8 As 2 + 6 HC1 = 3 ZnCl 2 + 2 AsH 3 . AsNa 8 + 3 H 2 SO 4 = 3 NaHSO 4 + AsH 8 . (2) As 2 O 3 + 12 H = 3 H 2 O + 2 AsH 3 . PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 319 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 + 3 O 2 = 3 H 2 O + As 2 O 3 . On being heated, the gas readily dissociates into arsenic and hydrogen : 4AsH 3 = As 4 -f 6H 2 . 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 320 OUTLINES OF CHEMISTRY arsenic will deposit on the sides of the tube just beyond the 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 = As 2 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 3 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 321 and is formed when arsenic burns in the air or in oxygen. Arsenic trioxide is manufactured on a commercial scale by roasting arsenical pyrites in the air. In this process iron oxide remains as a non-volatile residue, sulphur dioxide escapes, and the arsenious oxide condenses as a white powder upon the brick waljs of the chambers. It is purified by resublimation. In the year 1910, 1497 tons of arsenious oxide were produced in the United States. About four 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 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 it 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- cine. 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 and 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 1775 by passing chlorine into arsenic trioxide suspended in 322 OUTLINES OF CHEMISTRY 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 HAs0 3 = H 2 + As 2 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 As 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 323 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 3 . 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 : As 2 S 3 + 3 (NH 4 ) 2 S + 28 = 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 5 + 6 NH 4 C1 + 3 H 2 S. Arsenic pentasulphide As 2 S 5 , 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 : 324 OUTLINES OF CHEMISTRY 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 O 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 fire 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 close 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 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 325 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 SbClg. 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 326 OUTLINES OF CHEMISTRY 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 = SbOCl + 2 HC1. (2) 4 SbCl 3 + 5 H 2 O = (SbOCl) 2 Sb 2 O 3 + 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 SbOl 5 -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 6 is an amorphous gummy mass. It readily enters into the formation of double salts. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 327 Antimony tribromide SbBr 3 forms white rhombic crystals that melt at 94. The salt boils at 275, and is decomposed by water. Antimony triiodide SbI 3 forms three different varieties of 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 5 . 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 4- 6 HC1 = 2 SbCl 3 + 3 H 2 O. Sb 2 O 3 + 2 KOH = 2 KSbO 2 + H 2 O. Sb 2 O 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 univalent 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- 328 OUTLINES OF CHEMISTRY 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 metantimonious 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 3 ; farther, 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 metantimonic acid, thus : Sb 2 4 + (C 4 H 4 6 )HK = (C 4 H 4 O 6 )SbO - K + HSbO 3 . The tetroxide is also obtained on igniting antimony pentoxide : Antimony tetroxide may be regarded as the antimonyl salt of metantimonic 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 8 , which on being heated with water passes into solution as potassium antimonate KH 2 SbO 4 . On fusing PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 329 potassium metantimonate with caustic potash, the pyroantimo- nate K 4 Sb 2 O 7 results : 2 KSb0 3 + 2 KOH = K 4 Sb 2 7 + H 2 O. Potassium pyroantimonate is decomposed by water : K 4 Sb 2 7 + 2 H 2 = 2 KOH + K a H 2 Sb a 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 (NH 4 ) 2 S = 2 (NH 4 ) 3 SbS 3 . 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 6 + 3 H 2 S. Antimony pentasulphide may also be obtained by treating anti- monic acid with hydrogen sulphide, thus : 2 H 3 Sb0 4 + 5 H 2 S = Sb 2 S 6 + 8 H 2 O. 330 OUTLINES OF CHEMISTRY 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 " 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 formulae 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 8 . In nitric acid, bismuth may readily be dissolved, forming the nitrate PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 33i Bi(NO 3 ) 3 ; likewise when the metal is treated with sulphuric acid, the sulphate, Bi 2 (SO 4 \ 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 = 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 rive have not been prepared, but a dichloride 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 332 OUTLINES OF CHEMISTRY 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 O 5 + 10 HC1 = 5 H 2 O + 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 = (BiO) 2 SO 4 4- 2 H 2 SO 4 + 2 H 2 O. With nitric acid, bismuth forms the nitrate Bi(NO 3 ) 3 , which crystallizes in triclinic forms with rive 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 -f 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. The 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. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 333 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, antimon}^, 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 WEIGHT COLOR SPECIFIC GRAVITY MELTING POINT BOILING POINT Nitrogen, N 14.01 colorless 0.885 (liquid) -210.5 -194.4 Phosphorus, P 31.0 yellow or 1.8-2.3 + 44.4 + 278.0 red Arsenic, As 75.0 gray, 5.7 500 450 lustrous (approx.) (approx.) Antimony, Sb 120.2 white, 6.8 625 1500 lustrous (approx.) Bismuth, Bi 208.0 reddish 9.8 268 1600 white (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 8 . So we have ammonia NH 3 , phosphine PH 3 , arsine AsH 3 , 334 OUTLINES OF CHEMISTRY and stibine SbH 3 . 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 5 PBr 8 PBr 5 P 2 I 4 PI 3 AsF 3 AsCl 3 AsBr 3 As 2 I 4 AsI 3 SbF s SbF 5 SbCl 8 SbCl 5 SbBr 3 SbI 3 Sbl, BiF 3 BiCl 3 BiBr 3 BiI 3 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 such 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 : PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 335 N 2 O NO NA (N0 2 ) ^A PA (P0 2 ) 2 PA As 2 O 3 As 2 5 Sb 2 O 3 (Sb0 2 ) 2 Sb 2 O 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 N 2 2 HNO 2 (HAsO 2 ) HSbO 2 H 3 P0 2 H 3 PO 3 H 3 As6 3 H 3 Sb0 3 H 3 P0 4 H 3 AsO 4 H 3 SbO 4 H 4 P 2 O r H 4 As 2 O r H 4 Sb 2 O 7 HN0 3 HP0 3 HAsO 3 HSbO 3 (HBi0 3 ) The sulphides are commonly of the general type V R 2 S 5 . They are given in the following table : or