THE FOUNDATIONS OF CHEMICAL THEORY . BY R. M. CAVEN, D.Sc. F.I.C. Systematic Qualitative Analysis. FOR STUDENTS OF INORGANIC CHEMISTRY. A Short System of Qualitative Analysis. FOR STUDENTS OF INORGANIC CHEMISTRY. The Foundations of Chemical Theory. AN INTRODUCTORY TEXTBOOK. BY R. M. CAVEN, D.Sc. AND G. D. LANDER, D.Sc. Systematic Inorganic Chemistry, FROM THE STANDPOINT OF THE PERIODIC LAW. A TEXT- BOOK FOR ADVANCED STUDENTS. BLACKIE SON, LTD., LONDON, GLASGOW, BOMBAY THE FOUNDATIONS OF CHEMICAL THEORY The Elements of ^Physical and General Chemistry BY R. M. CAVEN D.Sc.(London), F.I.C. Professor of Inorganic and Analytical Chemistry in the Royal Technical College, Glasgow BLACKIE AND SON LIMITED 50 OLD BAILEY LONDON GLASGOW AND BOMBAY 1921 PREFACE Modern chemistry presents a multitudinous array of facts. It could not be otherwise, since the science undertakes to describe more than eighty separate elements and the compounds they form, one with another. The exhibition of these facts, as set forth in a comprehensive textbook, whether of organic or inorganic chemistry, may well dazzle and bewilder the student, who will scarcely be likely to appreciate the beauty of the science, or the glamour of the human achievement of which it is the monument, if he is required labori- ously to appropriate the facts without judicious selection and arrangement. Indeed, the complaint is often made by students that chemistry requires too much memory-work, and is therefore not so inspiring a science as physics, which deals all the while with fundamental principles, and properties of matter. Yet the facts of chemistry are the facts of nature, and nature is not chaotic. Consequently these accumulated facts present to the mind of man a powerful challenge. They must be systematized. Chemical theoiy, however, is worthy of the facts which form the subject-matter of the science, for it includes some of the greatest generalizations of any science. The atomic and molecular theories, the theories of molecular structure, the periodic law, the conceptions embraced in modern physical chemistry, which con- stitute, par excellence, the science on its intellectual side, are among the noblest achievements of the human mind. Every student of chemistry must become acquainted in some degree with chemical theory. The question therefore arises how this theory may best be communicated. Ought it to be inter- mingled dexterously with chemical facts with the descriptive part, so much easier to understand and more difficult to remember; V 458530 PREFACE or ought a special course of instruction in the theory of the science to be devised, apart from the descriptive lectures? The exigencies of teaching generally lead to the following solu- tion: that the student learns enough theory during his second year or intermediate course to serve him for the time being, and that in the later course for his degree he attends lectures in physical chemistry, which take for granted a solid foundation of elementary theory. A degree course in physical chemistry is a serious affair, and in the hands of a specialist is likely to make a severe strain on the mental resources of a student. To profit by such a course a student should be very sure of his foundations. The purport of this book may therefore now be stated. As its title indicates, it is an endeavour to disclose the foundations of the science, and make them plain and real to the average student. The author hopes that the general reader, who wishes to know what modern chemistry really means, will find within these pages the information he desires.; and that the student to whom chemical science offers an open field of glowing possibilities will find' the chapters of this book a not unwelcome guide in his earlier ex- cursions. Briefly, it is suggested that the book may be read by the student during or at the end of his second year's course, for the purpose of knitting together his chemical knowledge in view of the more advanced studies which lie before him later. The writer has utilized facts and methods of presentation con- tained in former books of which he is joint or sole author; and he desires to acknowledge his indebtedness to various other works, particularly to the introductory volume of A Textbook of Inorganic Chemistry, by Dr. J. Newton Friend. He also wishes to thank Dr. E. B. R. Prideaux for kindly reading the manuscript of the book. ROYAL TECHNICAL COLLEGE, GLASGOW, October, 1920. CONTENTS CHAP. Page I. THE ATOMIC AND MOLECULAR THEORIES - - - . . . 1 Composition of Matter The Elements Laws of Chemical Com- binationThe Atomic Theory The Molecular Theory Law of Volumes Avogadro's Theory. II. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 18 Equivalent and Atomic Weights Standard for Equivalent and Atomic Weights Determination of Equivalent Weights Methods of Determining Atomic Weights: (a) Method of Vapour Density and Avogadro's Theory Vapour Densities by the Methods of Dumas, Hoffmann, Victor Meyer; (6) Method of Chemical Displacement; (c) Method of Specific Heats (Dulong and Petit's Law); (d) Method of Isomorphism (Mitscherlich's Law); (e) Method of the Periodic Law Illustra- tion: The Atomic Weight of Carbon Molecular Weights in Solution: Cryoscopic and Ebulliscopic Methods Molecular Complexity Molecular Compositions of Compound Gases. III. VALENCY AND CHEMICAL CONSTITUTION 55 Historical Definition of Valency Bonds and Graphic Formulae Variability of Valency Double Bond in Carbon Compounds "Chemistry in Space" Criterion of Valency Nature of Valency. IV. CLASSIFICATION OF THE ELEMENTS THE PERIODIC LAW 67 Law of Octaves Development of the Periodic Law Periodicity of Physical Properties Periodicity of Chemical Properties- Periodicity of Valency Uses of the Periodic Law Correction of Atomic Weight Values Suggestiveness of the Periodic Law Objections to the Periodic Law. V. THE STATES OF MATTER, AND THE PROPERTIES OF GASES 88 States of Matter Gas Laws of Boyle and Charles The Gas Equation Diffusion of Gases; Effusion; Atmolysis Dalton's Law of Partial Pressures Deviations from Boyle's Law Pro- perties of Carbon dioxide Andrews's Experiment Critical State Liquefaction of Gases Methods of Liquefaction: I Simple Compression; II Cascade Method of Cooling; III Method of Adiabatic Expansion ; IV Method of Self-intensive Refrigeration Liquid Air Liquefaction of Hydrogen and Helium Practical Applications of Liquefied Gases i Ammonia, Chlorine, Carbon dioxide, Sulphur dioxide, Air. vii viii CONTENTS CHAP. Page VI. THE PROPERTIES OF LIQUIDS 112 Mobility Viscosity Density Specific and Molecular Volume Vapour Pressure and Boiling-point Distillation under Atmos- pheric and Reduced Pressure Relation between Boiling-points of Liquids in Homologous Series. VII. THE PROPERTIES OF SOLIDS ........ 120 Formation of Solids Solidification of Vapours; Sublimation Solidification of Liquids Melting-points of Solids Melting- points of the Elements Melting-points of Carbon Compounds Formation of Solids from Solution Solidification of Mixtures; Cryohydrates and Eutectics Crystals and Crystallography Crystallization Polymorphism and Allotropy. VIII. SOLUTIONS 143 Solvent, Solute, Solution Solutions of Gases in Liquids- Solubilities of Gaseous Mixtures Solutions of Liquids in Liquids Distillation of Mixed Liquids; Fractional Distillation Solu- tions of Solids in Liquids Solubility The Process of Solution Influence of Temperature on Solubility; Solubility Curves Relation between Chemical Composition and Solubility. IX. THE PROPERTIES OF DILUTE SOLUTIONS 162 Electrolytes Electrolytic Dissociation Laws and Process of Electrolysis Solution Pressure Chemical Reactions in Solu- tion Colours of Salt Solutions Complex ions Theory of Indicators. X. TYPES OF CHEMICAL COMPOUNDS 177 Hydrides Oxides and Hydroxides Bases, Acids, and Salts Halides Sulphides Oxy-salts Hydrated Salts Double and Complex Salts. XI. CHEMICAL CHANGE IN GENERAL 198 A Typical Reaction Reversible Reactions Chemical Equili- brium Thermochemistry Rate and Limits of Chemical Change Catalysis. XII. CHEMICAL CHANGES CLASSIFIED 215 Chemical Action of Heat on Compounds Thermal Dissociation Thermal Decomposition Chemical Interaction of Water with Elements and Compounds Hydrolysis Chemical Interaction of Acids with Metals Interaction of Nitric Acid and Metals Effect of Solubility and Volatility on Chemical Change Oxida- tion and Reduction. XIII. THE COLLOIDAL STATE 247 Solution and Suspension Crystalloids and Colloids Hydrosol and Hydrogel Typical Colloids The Ultra-microscope Size of Colloidal Particles Gradations between Suspension and Solution. XIV. EQUATION-BUILDING - - - - 253 The Meaning of a Chemical Equation Construction of Chemical Equations: Typical Examples. INDEX 263 THE FOUNDATIONS OF CHEMICAL THEORY CHAPTER I THE ATOMIC AND MOLECULAR THEORIES i. The Composition of Matter The term matter at first suggests to the unsophisticated mind ~ *s quickly r only to ERRATA sider, for Page 21. Last line but one. For 0'09 read 0-00009. itted that Page 89. 5th line from bottom of page. For oo read a . th pOSSGSS Page 94. 7th line from bottom of page. For 367'5 read 430. > n f ,1,0 Page 94. 6th line from bottom of page. For 430 read 367'4. Page 94. 4th line from bottom of page. For 367'5 read 367'4. Qd th6 ail> Page 111. 3rd line. For oo read oc . Apples On .* possesses permanent bulk, apparent colour, or At the conclusions suggested by these thoughts the Ancients arrived after their own fashion. Their fundamental classification of natural things included the three categories: earth, water, air, together with a fourth fire. These four were the elements, but their names really stood for qualities rather than separate species of matter. For instance, earth meant dryness and coldness, water wetness and coldness, and so on. Nevertheless, the first three terms at least suggest an outlook on the world which was essentially true, since they stand for the three fundamental forms of matter: solid, liquid, gas. ( D 60 ) 1 2 viii CONTENTS CHAP. Page VI. THE PROPERTIES OF LIQUIDS 112 Mobility Viscosity Density Specific and Molecular Volume Vapour Pressure and Boiling-point Distillation under Atmos- pheric and Reduced Pressure Relation between Boiling-points of Liquids in Homologous Series. VII. THE PROPERTIES OF SOLIDS . . . -*.-.. 120 Formation of Solids Solidification of Vapours; Sublimation Solidification of Liquids Melting-points of Solids Melting- points of the Elements Melting-points of Carbon Compounds Formation of Solids from Solution Solidification of Mixtures; Cryohydrates and Eutectics Crystals and Crystallography Crystallization Polymorphism and Allotropy. VIII. SOLUTIONS 143 Solvent, Solute, Solution Solutions of Gases in Liquids Solubilities of Gaseous Mixtures Solutions of Liquids in Liquids Distillation of Mixed Liquids; Fractional Distillation Solu- tions of Solids in Liquids Solubility The Process of Solution Influence of Temperature on Solubility ; Solubility Curves- Relation between Chemical Composition and Solubility. IX. THE PROPERTIES OF DILUTE SOLUTIONS 162 Electrolytes Electrolytic Dissociation Laws and Process of Electrolysis Solution Pressure Chemical Reactions in Solu- In X. TYPE* HI H* Co XI. CHEM A bri XII. CHEM Ch Elements and Compounds Hydrolysis Chemical Interaction of Acids with Metals Interaction of Nitric Acid and Metals Effect of Solubility and Volatility on Chemical Change Oxida- tion and Reduction. XIII. THE COLLOIDAL STATE 247 Solution and Suspension Crystalloids and Colloids Hydrosol and Hydrogel Typical Colloids The Ultra-microscope Size of Colloidal Particles Gradations between Suspension and Solution. XIV. EQUATION-BUILDING 253 The Meaning of a Chemical Equation Construction of Chemical Equations: Typical Examples. INDEX ....... 263 THE FOUNDATIONS OF CHEMICAL THEORY CHAPTER I THE ATOMIC AND MOLECULAR THEORIES i. The Composition of Matter The term matter at first suggests to the unsophisticated mind such qualities as bulk, shape, colour, hardness, weight. It is quickly recognized, however, that some of these qualities belong only to some kinds of matter, and are absent from others. Consider, for example, a log of wood floating on water. It will be admitted that the wood but not the water possesses hardness, though both possess weight; that the log but not the water has a permanent shape, though both have bulk. Above the water is the air, and the air is something, for it blows in a man's face, and raises ripples on water. By a proper instrument it can be shown that air possesses weight. So air is matter, though it is without form, permanent bulk, apparent colour, or hardness. At the conclusions suggested by these thoughts the Ancients arrived after their own fashion. Their fundamental classification of natural things included the three categories: earth, water, air, together with a fourth fire. These four were the elements, but their names really stood for qualities rather than separate species of matter. For instance, earth meant dryness and coldness, water wetness and coldness, and so on. Nevertheless, the first three terms at least suggest an outlook on the world which was essentially true, since they stand for the three fundamental forms of matter: solid, liquid, gas. (D60) 1 2 CHEMICAL THEORY : .*" : This teoncltlsion as to the threefold constitution of the world is reached by an extensive outlook upon nature; an intensive, an introspective view, such as the following illustration furnishes, leads to another conclusion. Sea-water is distinguished from fresh water by its saltness, that is, by its special taste. What proportion of salt water mixed with fresh water could be so distinguished depends upon the sensitive- ness of the human palate; but such a test would assuredly fail when the salt water was highly diluted. The addition of silver nitrate to the much diluted salt water would, however, serve to detect the presence of salt after the test of taste had failed, because of the turbidity or opalescence which the silver nitrate produces with even very small quantities of chlorides in solution. Would this test fail in its turn when the utmost delicacy was required, or would it detect the minutest quantity of salt? It might fail by reason of defective human vision, were it not that an instrument has been made on purpose to detect the slightest cloudiness in a liquid; but it must fail at last for quite another reason. For the test depends on the insolubility of silver chloride in water; but silver chloride is not quite insoluble in water, and on this account will not be precipitated when the salt solution is excessively dilute. So salt may perhaps be present in, and diffused through, water in quantity too minute to be detected by any test whatsoever. How far, then, may the dilution be carried; will salt still be present after infinite dilution? Or, more generally, is matter infinitely divisible? This is the question which arises directly out of an experiment which any novice in chemistry can perform. The same question presented itself to the alert minds of the ancient peoples of the East; not, it is true, by reason of experimental investigation, but because of meditation on the nature of the material world. Thus the question was: Is matter infinitely divis- ible or not? To believe the latter is the easier and more satisfying philosophy; this was the philosophy of the Ancients. So matter was supposed to consist ultimately of hard, indivisible, and in- destructible particles separated by vacuous interspaces; that is, of atoms. Of the two theories of the Ancients, to which the Greeks gave finished expression the theory of Elements and the theory of Atoms the former passed through strange vicissitudes, which THE ATOMIC AND MOLECULAR THEORIES 3 need not here be traced; but the latter persisted essentially un- changed till modern times, when it was found to be in accord with the conclusions derived by Dalton from experimental data. 2. The Elements A theory of the elements should precede a theory of atoms. So, dismissing the ancient theory of the elements, it may be said quite briefly that an element is an ultimate species of matter; or, to adopt the more usual and explicit definition: An element is a substance which hitherto has not been resolved into two or more dissimilar kinds of matter. We owe this idea of an element first of all to Boyle (16"78); it was Lavoisier (1789), however, who realized its provisional nature; and, indeed, some of Lavoisier's elements, such as lime and the alkalis, are now proved to be compounds. If this definition merely marked the present state of progressive human achievement it would not be a scientific definition. An assurance is necessary that some kind of finality has been or may be reached in the decom- position of substances; that at least the "elements" are equally elementary. This assurance may be given with every confidence, except, perhaps, as regards the metals of the rare earths. That the elements are absolutely undecomposable has, however, never been a settled belief of the chemist; on the contrary, he has held the opinion from time to time that they are derived from, and so are resolvable into, a common primordial substance. The phenomena of radioactivity now furnish evidence of the spontaneous and perpetual disintegration, into simpler forms of matter, of the atoms of certain of the elements. It is possible, however, as has lately been pointed out, that the elements are not truly homogeneous, but contain different parts which are indistinguishable and inseparable by physical or chemical means. These considerations, however, do not affect the chemist's work- ing theory of the elements. He knows that the eighty -five different kinds of matter into which he has resolved the many substances found in nature, and out of which he can elaborate a vast number of compounds to which nature has no counterpart, pass unchanged through the crucible of his operations. So he habitually regards the catalogue of the elements that hangs in his 4 CHEMICAL THEORY laboratory as a permanent record, not only of human skill, but of Nature's handiwork as well. 3. The Atomic Theory The atoms of Greek philosophy were indestructible; indeed, the indestructibility of matter has probably always been an axiom of science, notwithstanding the surprising and fantastic changes matter was supposed to undergo in the hands of the alchemists of the Middle Ages. This principle was first clearly illustrated, however, by Lavoisier in his application of quantitative methods to chemistry, and was subsequently demonstrated, within the limits of the most accurate experimental research, by Stas, Landolt, and others. 1 In 1770 Lavoisier gave an account of experiments he had performed to test the supposition that water is transformed into earth by boiling. A weighed quantity of water was boiled for 101 days in a weighed and sealed glass vessel; and at the end of that time it was found that while "earth" appeared in the vessel the water weighed the same as at first, and the weight of the "earth" was equal to the loss in weight which the glass vessel had incurred. Thus it was shown that the "earth" came from the glass and not from the water, and that water is not transformed into earth by boiling. In these experiments the use of the balance played an essential part; but this was a novelty in chemistry. The scientific achievements of such men as Boyle, Black, Cavendish, Priestley, Scheele, notwithstanding their great value, were chiefly of a qualitative nature. Henceforth, however, chemistry was concerned with weighing things, and a new era began. It was now but a step to the quantitative analysis of chemical substances. Lavoisier took this step in his investigation of mer- curic oxide, or the calx of mercury, which Priestley and Scheele had decomposed into mercury and oxygen. Soon there arose an important question, the answer to which could be found only by quantitative analysis. This was the question: Is a chemical com- pound necessarily constant in composition, or may its composition vary within certain limits according to the way in which it is prepared ? It may appear to be a truism that the same compound must *For an account of these researches see The Study of Chemical Composition, by I. Freund. THE ATOMIC AND MOLECULAR THEORIES 5 always have the same composition, so that it is better to state the problem in this way: Can the products of different chemical re- actions, designed to produce the same compound, really differ slightly in composition? Berthollet was of opinion that they could; that the composition of a compound might vary within certain limits according to the way in which it was prepared; indeed, that the conditions of its genesis are the overruling factors of its composition. Barium sulphate was cited as an example. All known specimens of this compound were found to be identical in composition, but this identity was due, not to any inherent pro- perty of the constituent elements of the compound, but to the fact that by uniting in such proportions these elements produced a compound of maximum insolubility in water. It was fair to suppose, therefore, that if the salt could be precipitated from some other medium than water it would have a different composition accom- modated to a new requirement of maximum insolubility. Such an idea was gravely erroneous, and was quite foreign to the principles on which the atomic theory was soon to be founded. Yet the idea appeared to have experimental support; and, indeed, it contained the germ of an important truth. In support of his belief, Berthollet showed that when nitric acid reacted with mercury or with tin the composition of the nitrate of mercury or oxide of tin produced varied within certain limits according to the concentration of the acid employed. Proust, on the other hand, maintained that "be- tween pole and pole compounds are identical in composition; their appearance may vary owing to their manner of aggregation, but their properties never". After a controversy carried on with Berthollet over a period of eight years (1800-8), Proust fully established his proposition, and showed that the variable products obtained by Berthollet were variable mixtures of invariable com- pounds. Thus was established the first law of chemical combina- tion the law of definite or fixed proportions: The same chemical compound always contains the same elements united together in the same proportions; or the proportions between the constituent elements of a chemical compound bear an unalter- able relation to each other, and to the proportion of compound formed. This was the first foundation of the atomic theory. Nevertheless, it was a pity that the truth in Berthollet's view was entirely overlooked in the victory of Proust: the truth that 6 CHEMICAL THEORY the proportions or concentrations in which reacting substances are present may determine the proportions subsisting between the products of a reaction, although the proportions in which elements or compounds actually react to form these products are quite be- yond the influence of external and variable conditions. Thus, in the case of the action of nitric acid on mercury, studied by Ber- thollet, the concentration of tiie acid determines whether mercurous or mercuric nitrate or a mixture of these two salts is produced, although it can have no influence on the unalterable chemical composition of either of the two salts. The existence of two nitrates of mercury is, however, a note- worthy fact, which appears the more striking when it is discovered that in one compound the proportion of mercury to nitrate is exactly twice what it is in the other. Further examples of this phenomenon were observed by Dalton, who showed that the pro- portion of hydrogen to a fixed quantity of carbon is twice as great in methane as in ethylene, and of oxygen to a fixed quan- tity of carbon, twice as great in carbonic acid gas as in carbonic oxide. Other examples of compounds showing analogous relations are two of the oxides of lead, one of which contains twice as much oxygen compared with lead as the other, and the five oxides of nitrogen, in which the quantities of oxygen combined with a fixed amount of nitrogen are as 1 : 2 : 3 : 4 : 5. Here was an important generalization, which was formulated by Dalton as the law of multiple proportions: When one element combines with another in more than one pro- portion, these proportions bear a simple relation to one another. The foregoing facts furnish material enough for the atomic theory. It is usual, however, to add to the laws of definite and multiple proportions a third law, the law of reciprocal proportions, which, however, follows logically from the other two laws. It was shown by Richter, about 1*780, that the ratio between the quantities of two acids which neutralize a fixed amount of alkali is the same whatever the alkali may be; and by Berzelius, in 1810-2, that 381 parts of lead combine separately with 58*73 parts of sulphur and 29-6 parts of oxygen, whilst 58*73 parts of sulphur combine in turn with 57 45 parts of oxygen. Now, 57 45 = 29 6 x 2 within the limits of the experimental error of the time; and these facts may be expressed diagrammatically thus: THE ATOMIC AND MOLECULAR THEORIES Sulphur -< >- Oxygen So is illustrated the law of reciprocal proportions: The proportions of two elements which separately combine with a fixed proportion of a third element are also the proportions of these elements which combine with each other, or else in accord- ance with the law of multiple proportions they bear a simple ratio to these proportions. This law has within it, especially in the way in which it was illustrated by Richter, the idea of chemical equivalents; and so it may be stated in this axiomatic way: Quantities of substances which are chemically equivalent to the same quantity of a third substance, are chemically equivalent to one another. Thus it appears that, granted the validity of the idea of chemical equivalents, which will be examined later, the law of reciprocal proportions requires no experimental justification. Although we owe the essence of the modern atomic theory to Dal ton alone, the precise way in which the theory took shape in the mind of its author has been rather problematical. At the close of a paper on the absorption of gases by water, Dalton wrote as follows: "An inquiry into the relative weights of the ultimate particles of bodies is a subject, as far as I know, entirely new. I have lately been prosecuting this inquiry with remarkable success." No hint is given in the context of the way in which the atomic values, which follow, were estimated, nor of the precise reason why such values were believed to exist. The idea that matter consists of discrete particles was, however, in the air. Apart from the ancient theory of atoms, a theory of particles had been held more or less firmly by F. Bacon, Boyle, Higgins, and others; whilst Newton made the following explicit statement: " It seems probable to me, that God in the beginning formed matter 8 CHEMICAL THEORY in solid, massy, hard, impenetrable, movable particles, of such sizes and figures, and with such other properties, and in such proportion to space, as most conduced to the end for which He formed them; and that these primitive particles, being solids, are incomparably harder than any porous body compounded of them, even so very hard as never to wear or break in pieces; no ordinary power being able to divide what God Himself made one in the first creation. . . . The changes of corporeal things are to be traced only in the various separations and new associations and motions of these permanent particles." It would almost appear from such a pronouncement tbat Newton and not Dalton was the author of the atomic theory. Yet this statement is not a chemical theory: it is a cosmic theory intimately related to Newton's great discovery of universal gravi- tation. Dalton, however, was greatly indebted to Newton and the idea of ubiquitous particles which the theory of gravita- tion involved; and it appears that he conveyed this idea into chemistry and employed it to explain the laws of chemical com- bination. To understand the atomic theory, therefore, is simply to under- stand how the theory of particles fits the chemical laws. This is quite easy. Let there be three elements, A, B, C, and let the areas of the squares: represent the combining weights of these elements on any arbitrary scale, being the quantity of B which is found to combine separately with I A I parts of A and proportions; whilst parts of C, according to the law of definite parts of C also combine with IT parts of A, according to the law of reciprocal proportions, so that the following compounds are formed: THE ATOMIC AND MOLECULAR THEORIES 9 Then, according to the law of multiple proportions, compounds such as these may be formed: A 1 C A HA~ c B C c C C B B &c. How else can these experimental facts be interpreted than by the idea of " permanent particles " ? The elements combine accord- ing to the laws of definite and multiple proportions because they combine atom by atom; 1 atom of A with 1 atom of B; 1 atom of A with 2 atoms of B; 2 atoms of A with 1 atom of B; and so on. That is Dalton's atomic theory, and the theory is expressed succinctly in the following statements: 1. All matter consists of discrete particles called atoms, which are indivisible by any known chemical process. 2. Atoms of the same element are ordinarily supposed to be similar in all respects. 3. Chemical compounds are formed by the union of the atoms of different elements in simple numerical proportions. 4. The proportions in which elements combine to form compounds are determined by the atomic weights of the elements. The transition from the laws to the theory is quickly made: it is taken, so to speak, in a stride; but the boundary line between them must not be obliterated. The laws of chemical combination are statements of experimental facts; the theory is an explanation of these facts which is very probably true, but it does not stand in the same category as the facts. In science, facts and theory must always be clearly distinguished. When the atomic theory is accepted it at once appears that the combining weights of the elements represent the combining weights of their atoms. The atomic theory involves the atomic weights. No atomic theory previous to that of Dalton involved atomic weights; these were a novelty, and their introduction constitutes Dalton's great contribution to chemical science. The following atomic weights are selected from a list published by Dalton, the atomic weight of hydrogen being 1. 10 CHEMICAL THEORY D ALTON'S ATOMIC WEIGHTS Hydrogen 1 An atom of water or steam, Azote 5 composed of 1 of oxygen Carbon or Charcoal 5 + 1 of hydrogen 8 Oxygen Phosphorus Sulphur Magnesia 7 9 13 20 An atom of ammonia, com- posed of 1 of azote + 1 of hydrogen 6 Lime Q 1 ... 23 Oft An atom of carbonic oxide, ooua Jso 00 composed of 1 of carbon A ron . . Potash oo 42 + 1 of oxygen 12 Zinc... 56 A 4- e u 'i Copper Silver 56 100 An atom or caroomc acici, 1 carbon -f- 2 oxygen 19 Gold 140 An atom of sulphuric acid, Mercury 167 1 sulphur -{- 3 oxygen 34 If the student compares these atomic weights with those in use at the present day, he will see that they differ widely from the modern figures. Inaccuracies in Dalton's values are to be expected, but it is not experimental error which attributes, for example, an atomic weight of 7 to oxygen, instead of 16. As a matter of fact these combining weights are not atomic weights at all, but are approximately what we now recognize as equivalent weights. For, in truth, Dalton had no means of determining atomic weights. The value 7 (or 8) for oxygen is derived from the analysis of water: 8 parts by weight of oxygen combine with 1 part by weight of hydrogen to form 9 parts by weight of water. Who shall say from this that the atomic weight of oxygen is 8? That depends on the number of atoms of each element which combine together to form a unit of water, a fact clearly recognized by Dalton. Thus, we have the ratio O : H = 8 : 1 or 16 : 2. or 24 : 3, &c., and if 1 atom of oxygen combines with 1 atom of hydrogen, then the atomic weight of oxygen is 8; if 1 atom of oxygen combines with 2 of hydrogen, the atomic weight of oxygen is 16; if 1 combines with 3, it is 24; if 2 combine with 1, it is 4; and so on. There was, however, no evidence on which to base a decision between these alternatives. Just at this point Dalton made a regrettable mistake. Instead of recognizing the limita- tions of his experimental knowledge, he made the assumption THE ATOMIC AND MOLEOULAE THEORIES 11 that since only one compound of hydrogen and oxygen was known, it necessarily had the simplest possible composition, and so was formed from 1 atom of each of its constituent elements. Consequently, the atomic weight of oxygen was thought to be 7 (or 8); and for a similar reason the atomic weight of nitrogen (azote) was supposed to be 5, and that of carbon also 5. It is worth while to notice, however, that Dalton applied the term atom to the ultimate particles of substances known to be compounds as well as to those of elements; it is noteworthy also that his numerical values furnish examples of the law of multiple proportions; for instance, the composition of the two oxides of carbon. Dal ton's system of atomic symbols was ingenious: stood for oxygen, (^) for hydrogen, ^p for carbon, &c.; whilst for compounds such formulae as /[\- . which stands for sulphuric acid (SO 3 ), had to be constructed. In these formulae, however, picturesqueness did not compensate for practical inconvenience; and the suggestion of Berzelius (1811), that initial letters should replace Dalton's hieroglyphics, found general acceptance. 4. The Molecular Theory Dalton made no further advance along the road that he had traversed. His assumption that the simplest formulae for a com- pound is the right one was a subterfuge which marked the end of the road. Advance must therefore be sought in another direction; and it is found in the study of gases; for gases are the simplest form of matter, and, if atoms exist, the properties of gases will best elucidate their existence. In 1805 Gay-Lussac and Humboldt studied the volume pro- portions in which oxygen and hydrogen combine to form water; and announced that " 100 volumes of oxygen required for com- plete saturation 199-89 volumes of hydrogen, for which 200 may be put without error". This is a single example of a law, Gay- Lussac' s law of volumes, which is thus expressed: The volumes in which gases combine are simply related to each other, and to the volume of the compound gas which is formed. 12 CHEMICAL THEORY For example: 2 volumes of hydrogen combine with 1 volume of oxygen to form 2 volumes of steam. 1 volume of hydrogen combines with 1 volume of chlorine to form 2 volumes of hydrogen chloride. 3 volumes of hydrogen combine with 1 volume of nitrogen to form 2 volumes of ammonia. A necessary corollary of this law is the statement that: the densities, i.e. the masses of unit volumes, of the * elementary gases are simply related to their combining weights. Thus, since 1 volume of hydrogen combines with 1 volume of chlorine, and also 1 grm. of hydrogen combines with about 35 '5 grm. of chlorine, the density of chlorine compared with that of hydrogen as unity is about 35-5. It further follows, if gases combine volume by volume, accord- ing to the law of Gay-Lussac, and also atom by atom, according to the theory of Dalton, that there is a simple connection between the volume and the atom; and, indeed, that equal volumes of hydrogen and chlorine, for example, contain equal numbers of atoms. This conclusion, which was quite valid so far as it went, was reached by Gay-Lussac, but was denied by Dalton, on account of a difficulty which arose when the volume of the product was considered. Now when two separate and different elementary atoms com- bine to form a compound atom, or whatever it may be called, it is one entity they form, not two. It is impossible, for instance, that 1 atom of hydrogen combining with 1 atom of chlorine can produce two compound atoms of hydrogen chloride. And yet 1 volume of hydrogen combining with 1 volume of chlorine forms 2 volumes of hydrogen chloride. That was a dilemma; and it was met by Dalton by a spirited denial of the law of Gay-Lussac. "The truth is", said Dalton, "that gases do not combine in simple proportions by volume; when they appear to do so, it is due to an error in our experiments" ! Now, Dalton was wrong; and yet what other solution can be found, unless indeetl the " atoms " are torn in pieces in the process of chemical synthesis, and the pieces are afterwards joined together again in a different way? That is precisely the solution of the difficulty suggested by THE ATOMIC AND MOLECULAR THEORIES 13 Avogadro, in his celebrated hypothesis. In this hypothesis, which will now be expounded, two orders of particles were distinguished, which we now call atoms and molecules. Atoms are indivisible in ordinary chemical changes; molecules are aggregates of atoms with a few exceptions which maintain their integrity in ordinary physical changes, but suffer disruption in the course of chemical change, so that their constituent atoms may be re- arranged to form fresh molecules. Now, when hydrogen chloride" is formed from its elements the volume of the product is twice the volume of the hydrogen or of the chlorine; therefore it is sufficient to assume that the molecules of hydrogen and chlorine consist of pairs of atoms, which break into single atoms, and recombine, thus: Cl Cl so that 1 volume of hydrogen plus 1 volume of chlorine gives 2 volumes of hydrogen chloride, instead of 1 volume, according to the scheme: It might be objected, however, that if the molecules of hydrogen chloride are intrinsically twice the size of the atoms of hydrogen and chlorine, out of which they are formed, the volume of the compound gas might be expected in any case to be twice that of either of the simple gases. Such an objection, however, is invalid, since the actual size of the molecules of a gas is very small com- pared with the molecular interspaces, and consequently the question of a molecule of hydrogen chloride being intrinsically larger than an atom of hydrogen or of chlorine does not arise. The formation of 2 molecules of steam from 2 molecules of hydrogen and 1 molecule of oxygen is thus represented: H H H H 14 CHEMICAL THEORY The above processes of combination may be set forth in terms of volumes, by using Dal ton's symbols, thus: 1 vol. hydrogen. 1 vol. chlorine. 2 vols. hydrogen chloride. 2 vols. hydrogen. 1 vol. oxygen. 2 vols. steam. Or by means of chemical equations: H 2 + C1 2 = 2HC1. 2H 2 + O 2 = 2H 2 O. Thus the molecular formula H 2 for water makes its appear- ance. The proof of this formula is contained in the preceding argument, which may be thus epitomized: Hydrogen and chlorine gases consist of diatomic molecules, since the volume of hydrogen chloride they produce is twice the volume of either single gas. Similarly, oxygen gas consists of diatomic molecules, since the volume of the steam is twice the volume of the oxygen it contains. The only formula for steam which agrees with the diatomicity of hydrogen and oxygen, as well as with the volumetric com- position of steam, is H 2 O. That the density of steam (H = 1) is 9 furnishes no additional evidence, since it is deducible from the densities of hydrogen and oxygen, and the volume of the steam. That the atomic weight of oxygen is 16 follows from the fact that its density is 16, and that, like hydrogen, it is diatomic. The argument would, however, be invalidated if it were shown that these gases are not diatomic, that in the mole- cules H x and O x , x is greater than 2. Underlying the whole of this argument is Avogadros hypothesis, which is stated thus: Equal volumes of all gases and vapours, under the same conditions of temperature and pressure, contain equal numbers of molecules. But why hypothesis? This statement is not a law, any more than Dalton's atomic theory is a law. When first put forward it was properly regarded as a hypothesis, which, indeed, suffered much at the hands of its friends. Now, however, it is firmly established, and is of fundamental importance. It ought, therefore, THE ATOMIC AND MOLECULAR THEORIES 15 to be dignified with the name of theory. Henceforth we shall speak of Avogadros theory. It will be seen that this theory is in accord with Gay-Lussac's law of volumes, and satisfactorily explains the phenomena of the combination of gases. Thus, 1 volume of hydrogen combines with 1 volume of chlorine to form 2 volumes of hydrogen chloride, because 1 molecule of hydrogen reacts with 1 molecule of chlorine to form 2 molecules of hydrogen chloride. The language of volumes may be exchanged for fche language of molecules; that is the significance of Avogadro's theory. That equal volumes of different gases contain equal numbers of atoms is true only when the molecules of these gases contain equal numbers of atoms. It is a statement of limited truth, and of no permanent importance. The same may be said of the state- ment that the densities of elementary gases are in the same ratio as their atomic weights. The important fact is that the densities of all gases are in the same ratio as their molecular weights', and further, that since the molecular weight of hydrogen is 2, and its density, which is taken as the standard, is 1, therefore the molecular weights of all gases are twice their densities. Thus, the molecular weight of a gas or vapour is revealed by its density, as the following approximate figures show: Elementary Gas or Vapour. Density. Molecular Weight. Atomic Weight. Molecular Formula. Hydrogen 1 2 1 H 2 Oxygen Nitrogen Chlorine 16 14 35-5 32 28 71 16 14 35-5 2 N 2 C1 2 Ozone 24 48 16 3 Phosphorus , 62 124 31 P 4 Mercury 100 200 200 Hg Sulphur 128 256 32 S 8 It may be remarked, incidentally, that the magnitude of the atomic weight of an element cannot be deduced from its gas or vapour density unless the number of atoms contained within the molecule of the element, i.e. its atomicity, is known indepen- dently. As a rule, however, the atomic weight of the element is known independently, and then the atomicity is deduced from the density. 16 CHEMICAL THEORY The breadth of Avogadro's generalization was not realized in the time of its originator; and, owing to the persistence of the volume-atom theory of Gay-Lussac, and its unwarrantable exten- sion by Berzelius, 1 there was much confusion on the subject until Cannizzaro, in 1858, reinstated Avogadro's theory on a permanent basis. It should be added that Avogadro's theory applies strictly only to an ideal gas. When a gas deviates from Boyle's law it deviates to the same extent from Avogadro's theory. A useful fact to remember in connection with gas densities is that a litre of hydrogen at C and 760 mm. pressure, i.e. normal temperature and pressure (N.T.P.), weighs almost exactly 0-09 grm., or that 1 grm. measures 11-12 litres. Thus a gram-molecule (i.e. the molecular weight in grams) of hydrogen at N.T.P. measures 22-25 litres; and from Avogadro's theory it follows that the volume of a gram-molecule of any gas or vapour, reduced to normal temperature and pressure, is 22-25 litres. To determine the weight in grains of 22-25 litres of any gas or vapour, reduced to and 760 mm., is therefore to discover its molecular weight. SUMMARY AN ELEMENT is a substance which hitherto has not been resolved into two or more dissimilar kinds of matter. LAWS OF CHEMICAL COMBINATION. 1. Law of definite or fixed proportions. The same chemical compound always contains the same elements united together in the same proportions; or, the proportions between the constituent elements of a chemical com- pound are always the same. 2. Law of Multiple Proportions. When one element combines with another in more than one proportion, these proportions bear a simple ratio to one another. 3. Law of Reciprocal Proportions. The proportions of two elements which separately combine with a fixed proportion of a third element are also the proportions of these elements which combine with each other, or else in accordance with the law of multiple proportions they bear a simple ratio to these proportions. x The practice of referring all gaseous molecules to 2 volumes, which was a pernicious outcome of the theorizing of Berzelius, appears now, fortunately, to be dying out. Why, indeed, should every molecule be regarded as a microcosm of 2 volumes, as if it could necessarily be dichotomized? THE ATOMIC AND MOLECULAR THEORIES 17 THE ATOMIC THEORY. 1. All matter consists of discrete particles called atoms, which are indivisible by any known chemical process. 2. Atoms of the same element are ordinarily supposed to be similar in all respects. 3. Chemical compounds are formed by the union of the atoms of different elements in simple numerical proportions. 4. The proportions in which elements combine to form com- pounds are determined by the atomic weights of the elements. GAY-LUSSAC'S LAW OF VOLUMES. The volumes in which gases combine are simply related to each other, and to the volume of the compound gas which is formed. Corollary. The densities of the elementary gases are simply related to their combining weights. 1 AVOGADRO'S THEORY. Equal volumes of all gases and vapours under the same conditions of temperature and pressure contain equal numbers of molecules. Corollary. Since the molecule of hydrogen contains 2 atoms, the molecular weight of any gas or vapour is twice its density compared with that of hydrogen as unity. A litre of hydrogen at N.T.P. weighs 0*09 grin., and 1 gram- molecule of hydrogen (2 grm.) measures 22-25 litres. It follows from Avogadro's theory that this is also the volume at N.T.P. of 1 gram-molecule of any gas or vapour. AN ATOM of an element is the smallest particle of matter which takes part in a chemical change; it is the unit of chemical exchange. A MOLECULE is the smallest particle of matter which exists independently; it is the physical unit. The molecule of an element contains similar, that of a compound dissimilar atoms. The number of atoms contained within the molecule of an element is called the atomicity of the element. ] The term "combining weight" has sometimes signified equivalent weight, and some- times atomic weight. Since the term is ambiguous, a use is found for it during the development of the molecular theory when non-committal language is employed. After- wards the term should be dropped. (D60) CHAPTER II EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS I. Equivalent and Atomic Weights It was shown in the last chapter that Dal ton's "atomic weights'* were really equivalent weights, and that the equivalent weight of an element, when not identical with its atomic weight, is a sub- multiple of the latter. Thus, whilst the equivalent weight of oxygen referred to that of hydrogen as unity is approximately 8, the atomic weight of this element, referred to the same standard, is approximately 16. In general Atomic weight = n x equivalent weight, where n is a small whole number, which indicates the valency of the element. Valency, or atomic value, is a new idea, necessary to connect together the ideas of atomic weight and equivalent weight. It will be more fully developed later. It will now be useful to define equivalent and atomic weights. EQUIVALENT WEIGHT. The equivalent weight of an element is that weight of it which combines with, or displaces from combina- tion, unit weight of a standard element. ATOMIC WEIGHT. The atomic weight of an element is the ratio between the weight of its atom and that of the atom of a standard element. When these definitions are considered, it appears that the equivalent weight of an element is an experimental value, inde- pendent of theory, whilst the atomic weight is connected with the atomic theory. It further appears that since equivalent and atomic weights are ratios, they are not really weights at all, nor masses, but pure numbers. That the atomic weight of an element is not the weight of one of its atoms appears plainly enough when it is considered that the standard of atomic weights has varied from time to time. 18 EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 19 Farther, since equivalent weights are values to be determined experimentally, their determination may well form the starting- point in the estimation of atomic weights. As a matter of fact the accuracy with which the atomic weight of an element is known depends as a rule on the accuracy with which the quantitative observation of some chemical transformation has been carried out, so as to determine its equivalent weight. In some cases, however, atomic weights have been estimated accurately by the determination of gas density. For determining equivalents, comparison between reacting quantities may be made by combination as well as by displace- ment, because an element combines with, as well as displaces, what is equivalent to itself. Thus, if there are two elements, A and B, the chemical equivalent of B referred to A as standard is found by estimating the amount of B which combines with a known weight of A, as well as by causing B to displace A, or A to displace B from combination with another element or group of elements. When the equivalent weight of an element is known, it is necessary to determine the value of n in the above equation before the atomic weight can be fixed. What multiple of the equivalent weight the atomic weight may be, has to be decided by reference to one or more of several distinct principles, which lie chiefly in the domain of physical chemistry, and will shortly be discussed in detail. Standard for Equivalent and Atomic Weights. The question of a standard needs first to be considered; and, since hydrogen has the least atomic weight of all the elements, and as small an atomic value (valency) as any element, it is natural to choose hydrogen as the standard both of atomic and equivalent weights, and so to make its equivalent and its atomic weight both equal to 1. Now, although hydrogen combines with non-metals, and a few metals, and is displaced from its combination in acids by some metals, its chemical activity is too limited to permit its use as a general standard of comparison. Oxygen, however, with very few exceptions, combines with all the elements, metals and non- metals alike; on this account it was called by Berzelius the "pole of chemistry". As a matter of practical experience, therefore, equivalent and atomic weights are more often estimated with 20 CHEMICAL THEORY reference to oxygen than to hydrogen; the hydrogen equivalent may then be calculated from the oxygen equivalent by multiplying the latter by the equivalent weight of oxygen, and thence the corresponding atomic weight may be found. Now, although Dalton (1803) chose hydrogen = 1 as the atomic weight standard, oxygen was soon adopted in preference, so that Wollaston (1814) used oxygen = 10, Thompson (1825) oxygen = 1, Berzelius (1830) oxygen = 100, and Stas (1860-5) oxygen = 16. Until recently the two standards H = 1 and O = 16 were in use, but the latter is now the standard adopted by the International Atomic Weights Committee. Although unity as the standard is sacrificed by this procedure, and O = 16 is really quite an arbi- trary standard, it has at least two advantages over the H = 1 standard. It was pointed out by Stas that the standard atomic weight should, as far as possible, be directly connected with the atomic weight to be determined, and this is the case when oxygen rather than hydrogen furnishes the standard. Otherwise the ratio H : O is involved in the calculation when the data are derived from the composition of an oxide; and whilst this ratio has been determined with great accuracy to be 1 : 1588, any future modification of the ratio would involve the recalculation of all atomic weights de- pendent upon it. If, however, the ratio is written 1008 : 16, the atomic weight of oxygen being fixed at 16, any future alteration will involve only the atomic weight of hydrogen. The advantage of this is plain. Another advantage of the modern system is the fact that when O = 16 several other important atomic weights approximate very closely to whole numbers; e.g. C = 12-00, N = 14 .01, Na = 23-00. It is to be hoped that no further modification of the standard will now be made, for an unfortunate confusion even now remains in the minds of those who have employed several standards. For example, the atomic weight of chlorine has been variously given as 35-37, 35-18, 35-46; and these differences are due not to different estimations of the atomic weight of this element but to the adop- tion of three different standards for oxygen, viz. = 15-96 (Dumas), 15-88, and 16-00. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 21 2. Determination of Equivalent Weights The following are the more important methods ordinarily employed in the laboratory to determine the equivalent weights of elements. i. The measurement of the volume of hydrogen displaced from dilute sulphuric or hydrochloric acid by a weighed amount of a metal. ii. The conversion of a weighed quantity of a metal into its oxide which is weighed, or the reduction of a weighed quantity of oxide to metal. iii. The displacement of a metal from a solution of one of its salts by a weighed quantity of a more chemically powerful metal. iv. The separation of elements at the electrodes during the passage of an electric current through a series of electrolytes. This method yields the electro-chemical equivalent of an element; but this value is numerically identical with the chemical equivalent. i. The chemical equivalent of magnesium, zinc, or aluminium may be easily determined by dissolving a weighed quantity of the metal in the dilute acid contained in a piece of apparatus designed for collecting the evolved hydrogen. The gas is measured over water at atmospheric temperature and pressure; it will consequently be moist, and the pressure of water vapour at the observed tempera- ture must be subtracted from the atmospheric pressure, before the volume of the gas is corrected to normal temperature and pressure. The weight of metal divided by the weight of the evolved hydrogen gives the hydrogen equivalent of the metal. This must be multiplied by 1008 if the equivalent on the modern atomic weight basis is desired; though in view of the likely experimental error such a correction is superfluous. ' The experiment may easily be carried out on the lecture-table or by students. The following result has been obtained by a student- Weight of magnesium taken = 0-033 grm. Volume of moist hydrogen measured at) _ 32 .g cu cm 12 C. and 756 mm. / Pressure of water vapour at 12 = 10 5 mm. Vo,u m e of d r h a rogen at N.T.P. = ^ = 30-6 cu. cm. Weight of hydrogen = 30-6X0*09 = 0-002754 Equivalent of magnesium = ' ' \= 12-0. 0-00275 I \ 22 CHEMICAL THEORY ii. Magnesium may be converted quantitatively into oxide by the ignition of the metal in the air under suitable conditions, or by dissolving it in dilute nitric acid, evaporating the solution, and igniting the nitrate until brown fumes cease to be evolved. These methods are not without sources of error, but it may be shown that 30 grrn. of magnesium yields almost exactly 50 grm. of oxide, so that the equivalent weight of magnesium is '3x8 .. 2 The method of conversion into oxide through the nitrate is applicable to such metals as zinc and copper, which dissolve in nitric acid and yield stable oxides by the decomposition of their nitrates. The equivalent of tin may be determined by the con- version of the metal into hydrated dioxide by means of nitric acid, since the ignition of the product yields the pure dioxide. It would be possible to determine the equivalent of carbon by burning a weighed quantity of the element in a stream of dry air or oxygen, and collecting and weighing the carbon dioxide formed; but the great difficulty of obtaining pure carbon free from hydrogen under ordinary conditions stands in the way of this determination. For the determination of an equivalent by the reduction of an oxide to metal, copper furnishes the usual example, since the reduction is easily carried out by passing a stream of hydrogen over oxide of copper contained in a boat in a heated tube. Thus I'OO grm. of black oxide of copper leaves a residue of 0-799 grm. of copper; whence the equivalent of copper in this oxide is 799 201 X 8 = 31-8. There is another oxide of copper, however, the red oxide, whose equivalent weight is 31-8 x 2 = 63-6. This fact is connected with the exhibition of a dual valency by copper, which again furnishes an example of the law of multiple proportions. This phenomenon will be further dealt with under the subject of valency. iii. A well-known example of the displacement of a metal from the solution of one of its salts by another metal is the action of zinc upon a solution of copper sulphate, when the zinc is sup- posed to displace from combination its equivalent of copper which may be collected and weighed. This takes place almost quantita- tively when a cold concentrated solution of copper sulphate is EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 23 employed; but the method is generally unreliable, because other reactions occur between the displacing metal and the solution simultaneously with the main reaction, and these vitiate the results. The method is not therefore to be recommended. iv. When a suitable electric current is passed through acidified water contained in a "voltameter", hydrogen and oxygen are evolved in the proportion of two volumes of hydrogen to one "volume of oxygen. Provided the densities of hydrogen and oxygen are known, and the conditions of temperature under which the gases were measured have been observed, the hydrogen equiva- lent of oxygen might be calcu- lated from the volume relations 0-08 g ^ s c Rj Hydrogen 0-OlOOHg /^ /3 : ~tLp r / M Copper 0-318 g i Silver 1-079 g Fig. 1 of the gases. The estimation would not, however, be very accurate, owing to several sources of experimental error. If, however, the same current passes in succession through several salt solutions for example, copper sulphate, silver nitrate, gold chloride solutions it will liberate at the cathodes or negative electrodes amounts of the metals chemically equivalent to the hydrogen which is liberated in the voltameter. Thus, whilst 0-01008 grm. of hydrogen gas is being evolved, and 0-08 grm. of oxygen, 0-318 grm. of copper, 1079 grm. of silver, and 0657 grm. 24 CHEMICAL THEORY of gold will be deposited in the successive electrolytic cells. The necessary arrangement is shown in fig. 1. Thus the equivalent weights of these metals are determined. 3. Determination of Atomic Weights It has been suggested in the previous pages that two distinct considerations have to be taken into account in the problem of atomic weight determination. These are: i. An exact estimation of the chemical equivalent of the element must be made, generally by carrying out some suitable chemical transformation, occasionally by other means. ii. A decision must be arrived at as to the order of magnitude of the atomic weight, so as to discover the small whole number by which the equivalent weight must be multiplied to give the atomic weight. The order in which the two parts of the problem are here placed is that which would naturally occur to the mind. Never- theless it is not the order of historic sequence in relation to modern atomic weights. The approximate magnitude of the atomic weights of all the elements has long since been settled and is not discussed in modern research upon atomic weights; but the determination of the exact values of all these atomic weights is a laborious task which is not yet completed. The methods for determining chemical equivalents which have been described above are suitable for demonstration purposes, but not all of them are equally useful in the actual determination of atomic weights. Illustrations of the methods that have been employed in accurate atomic-weight determinations will be given in the sequel. The principles which have led to decisions upon the order of magnitude of the atomic weights of the elements will now be dealt with. It has already been seen that Dalton was in need of some guiding principle to enable him to fix the magnitude of his atomic weights; and that such a principle came to light in the discovery by Gay-Lussac of the law of gaseous volumes, and the proper interpretation of this law by Avogadro. Thus, by means of Avogadro's theory it was shown that the atomic weight of oxygen is very probably 16 and not 8; but clearly this theory is limited in EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 25 its application, since it can only be of use in the case of gaseous or gasifiable substances. Here may be mentioned the method of chemical displacement, which is of some value in deciding the magnitude of atomic weights. In 1819 two other and quite distinct principles became avail- able in the law of specific heats of Dulong and Petit, and the law of isomorphism discovered by Mitscherlich. These laws are espe- cially valuable in furnishing guidance as to the magnitude of atomic weights, because they are applicable to solid elements and their solid compounds. The former of these two laws is the more important, and has the wider application. Finally, the periodic law, established by MendeleefF in 1869, has been of distinct value in several ways in fixing the approximate magnitude of atomic weights. So the five guiding principles that aid in settling the order of magnitude of atomic weights are: i. Avogadro's theory, ii. Chemical displacement. iii. Dulong and Petit's law of specific heats. iv. Mitscherlich's law of isomorphism, v. MendeleefFs periodic law. i. The Method of Avogadro's Theory. It will be remembered that according to Avogadro's theory the molecular weights, not the atomic weights, of gases and vapours are proportional to their densities. It follows, therefore, that the rela- tive magnitudes of molecular weights, and not of atomic weights, are directly deducible from Avogadro's theory. So the question arises how far a knowledge of the relative weights of molecules can be of use in fixing the relative weights of any of their con- stituent atoms. Such knowledge may be employed in two ways. Consider the following volatile hydrocarbons: Methane. Ethylene. Propane. Benzene. Naphthalene. Approximate Density) 8 14 22 39 64 (<> = 16) J Approximate Molecular) Weight / 10 28 44 78 128 Molecular Proportion of \ Carbon J 12 24 36 ' 72 120 26 CHEMICAL THEORY Approximate estimations of gas or vapour density yield ap- proximate molecular weights; whilst quantitative analysis shows the proportion of carbon within the molecular proportion of each compound. Now, it is evident that all these hydrocarbons, except the first, contain more than 1 atom of carbon in their molecules. The molecule of methane might indeed contain more than 1 atom, though the fact than no submultiple of 12 appears in the pro- portions of carbon in the other molecules is evidence, so far as it goes, that the figure 12 represents an indivisible unit, or in other words that 12 is approximately the atomic weight of carbon. And since by the examination of the very large number of hydrocarbons that exist, every molecular proportion has been found to contain 12, or a multiple of 12 parts of carbon, the probability that 12 is the atomic weight of carbon reaches a practical certainty. The principle thus illustrated may be put in the following words: The least proportion of an element found within the molecular propor- tion of any of its volatile compounds is likely to be the atomic weight of the element; and if the number of compounds which have been examined is large, the value indicated is very probably the atomic weight. The question may be asked, however, whether atomic weights can be determined exactly by the method of Avogadro's theory, i.e. by the determination of gas density, and the answer is in the affirmative, provided an ideal gas density can be determined and the molecular composition of the gas is known. The density of a gas is determined by weighing a large glass globe of about 10 litres capacity, first evacuated, and then filled with the gas, at known temperature and pressure, corrections being applied for the air displaced by the globe, and for the slight shrinkage which the glass undergoes when the globe is evacuated. Thus it has been found, as the mean result of the experiments of Rayleigh, Morley, and Leduc, that 1 litre of oxygen at and 760 mm. at the latitude of Paris weighs 1-42895 grm., whilst 1 litre of hydrogen, under similar conditions, accord- ing to the experiments of Morley and Leduc, weighs 0-08985 grin. To conclude, however, that the atomic weights of oxygen and hydrogen are in the ratio , although we know that the 0*08985 molecules of both gases are diatomic, would be erroneous, because it would be to assume that the gases are ideal gases which behave in perfect accord with the gas laws (q.v.), and so with Avogadro's EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 27 theory. Such, however, is not the case, and the deviation of these gases from the ideal must be discovered, and allowed for. This may best be done, in the present case, by determining the effect of the deviation upon the volume relations in which oxygen and hydrogen combine to form water. Now it has been estimated that 2-00268 litres of hydrogen combine with 1 litre of oxygen, at and 760 mm. at the latitude of Paris. This complex ratio is due, not to any discrepancy between the simple proportions in which the molecules of these two gases interact, but to the fact that equal volumes do not contain quite equal numbers of molecules because oxygen is a little more compressible than hydrogen. But since the densities relate to equal volumes it may be concluded that 2-00268 x 0-08985 grm. of hydrogen combine with 1-42895 grm. of oxygen, and therefore that the hydrogen equivalent of oxygen is 1.42895 2-00268 x 0-08985 = 7-9412; and its atomic weight 15-88 when H = 1; so that H = 1-0075 when O = 16. A similar method may be applied to determine the atomic weight of a constituent element of a compound gas. Thus, by the calculation of the ideal densities of carbon monoxide, carbon dioxide, methane, and acetylene, by applying a correction for compressibility to the estimated densities, several observers have accurately determined the molecular weights of these gases and thence the atomic weight of carbon. METHODS OF DETERMINING VAPOUR DENSITY. The determination of gas density always consists in weighing a certain volume of the gas; but for determining the vapour density of a volatile liquid or solid, an alternative procedure may be adopted : the volume of the vapour produced by a weighed quantity of the liquid or solid may be measured under known conditions. There are three well -recognized methods of vapour density determination: the methods of Dumas, Hoffmann, and Victor Meyer. In the first of these three methods the weight of a known volume of the vapour is ascertained; in the two latter the volume of a weighed quantity of the substance is measured. The method of Victor Meyer is the easiest and most often employed. 28 CHEMICAL THEORY (a) Dumas's Method of Vapour-density Determination A glass globe of the shape shown in fig. 2, and capable of holding from 50 to 100 cu. cm. or more, is weighed, and then filled with the vapour of the substance in the following manner. A few cubic centimetres of the liquid are introduced into the globe, which is then immersed in a bath of another liquid whose temperature is kept constant, and from 20 to 30 above the boiling-point of the liquid in the globe. As the latter liquid boils it displaces the air from the globe, and vapour issues from the neck as long as any liquid remains within the globe. When the stream of vapour ceases, the globe is filled with the vapour at atmospheric pressure, and at the tem- perature of the bath in which it is immersed. The neck is then sealed by means of a blowpipe; and the temperature of the bath, and the pressure of the atmosphere at the time of sealing are recorded. After being cleansed, the sealed globe is weighed, Fig o and the temperature and pressure of the air in the vicinity of the balance are also observed. Since the true weight of the sealed globe with its contents is equal to its apparent weight plus the weight of the air which it displaces whilst it is being weighed, the weight of this air must be calculated and added to the apparent weight. For this calcu- lation, as well as to ascertain the volume of the vapour at the time of sealing, the cubical capacity of the globe must be de- termined. This is done by breaking off the end of the neck of the globe under water, which should then enter and fill the globe. The quantity of water in the globe is determined by another weighing, the weight of the air displaced being in this case negligible, and the weight of the water in grams shows the volume of the globe in cubic centimetres with sufficient accuracy. From these data the weight of the known volume of the vapour contained by the globe at the temperature and pressure at which it was sealed is calculated. The volume is then reduced to N.T.P., and the weight of hydrogen or air corresponding to it is calculated. The ratio of the weight of the vapour to that of the hydrogen is the vapour density of the substance. EXAMPLE. Calculate the density of ether vapour from the following data: EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 29 Weight of open globe in air = 22 549 grm. Temperature of bath at time of sealing . = 60 C. Atmospheric pressure at times of sealing and weighing = 760 mm. Apparent weight of sealed globe + vapour in air = 22 662 grm. Temperature of air at time of weighing =15 Capacity of globe, indicated by weight of water it \ ._ 75 can contain / Weight of 1 cu. cm. of air at C. and 760 mm. = 0-001293 grm. Weight of 1 cu. cm. of hydrogen at C. and 760 mm. = 0-0000899 grm. Calculation Weight of air displaced when sealed globe is weighed x 273 = 0.092 grm. cu cm Weight of vapour in globe = 22 662 + 092 22-549 = 205 grm. Weight of an equal volume of hydrogen at 60 C. and 760 mm. = 0-0000690X76X878 = 0.00553 grm . So density of ether ((C 2 H 6 ) a O) vapour = *-, = 37-1. (b) Hoffmann's Method of Vapour-density Determination A weighed quantity (about 0-05 grm.) of the liquid con- tained in a small, drawn-out bulb or stoppered bottle is in- troduced into the Toricellian vacuum of a graduated barom- eter tube surrounded by tbe vapour of a liquid boiling at a suitable temperature, which may be below the boiling-point of the liquid whose vapour den- sity is being determined. As the liquid is vaporized it de- presses the mercury in the barometer tube; and when the volume has become constant it is read off, and the temperature of the vapour jacket is observed. The pressure of the vapour is equal to atmospheric pressure less the height of the mercury in the tube above its level in the vessel in which the tube stands. Strictly speaking, the height of 30 CHEMICAL THEORY the mercury column should be corrected for expansion by heat; but this need not be considered. From these data the vapour density of the liquid can be calculated; as the following example shows: Weight of stannic chloride (B.P.I 14) taken = 0-0445 grm. Volume of vapour Temperature of vapour jacket Barometric pressure Height of mercury column Whence pressure of vapour = 16-2 cu. cm. = 99 = 752 mm. = 512 mm. = 240 mm. _. 16-2 X 273 X 240 _( 3-75 372 X 760 I cu. cm. Volume of vapour reduced to N.T.P. Weight of 3-75 cu. cm. of hydrogen at j = 3 . ?5 x Q . 00009 grm< = 0-0003375 grm. Vapour density ot stannic chloride (SnCl 4 ) = 0-0445 0-0003375 = 131-8. Fig. 4 (c) Victor Meyers Method of Vapour-density Determination In this method a weighed quantity of the substance is made to evaporate into a space sur- rounded with the vapour of a boiling liquid whose boiling-point is at least 25 higher than that of the substance. The volume of the vapour is not directly measured, but the air displaced by it is collected and measured at atmospheric tempera- ture and pressure; while all the displacing vapour remains in the locality of its production. The weight of an equal volume of hydrogen is then calculated, and the weight of substance taken divided by this weight of hydrogen gives the vapour density of the substance, since the vapour of the substance, if it could be obtained at atmos- pheric temperature and pressure without conden- sation, would occupy the same volume as the air. The figure and description will make the process plain. The tube A is closed at the lower end, and is furnished with a bent delivery tube B which dips under water in the dish C. The upper end of A is closed by a rubber stopper. The lower part of the tube is heated by the vapour of a liquid, e.g. water, boiling in the outer jacket D, and, owing to expansion, air escapes by the side EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 31 tube through the water. When no more bubbles of air are seen, the graduated tube E is placed over the end of B, and a little tube cr loosely-stoppered bottle, containing a weighed quantity of the substance under investigation, is dropped to the bottom of A, being received on a pad of asbestos or glass wool, which prevents frac- ture. For the introduction of the little vessel containing the sub- stance the rubber stopper is removed, and quickly replaced, or the vessel may be held by a mechanical contrivance at the top of the tube A, where the air is nearly cbid, 1 and then allowed to fall at the right moment, without opening the tube. EXAMPLE. 0144 grm. of chloroform displaced 28 6 cu. cm. of moist air measured at 14 and 756 mm. pressure. Pressure of water vapour at 14 = 12 mm. Vol.ofairatN.T.P. = 28-6 X 273 X (756- 12) f 26-6 287 X 760 \ cu. cm. Weight of an equal vol. of hydrogen = 26-6 X 0-0000899 grm. = 0-00239 grm. Vapour density of chloroform } __ 0-144 _ fio ~ CHC1 3 J ~ 0-00239 ~ The method of Victor Meyer is more easily carried out than either of the other methods. It employs very little of the substance and is sufficiently accurate for most purposes. Consequently, it is the method usually employed. J It might be supposed that since the air is colder in the upper part of the tube, which extends beyond the vapour jacket, than in the lower part, which is within it, too much air will be displaced, and a high result obtained. This, however, is not the case, because of the contraction of the air that rises in the body of the tube to take the place of the air driven out. The following proof of this statement has been given by Dr. E. B. R. Prideaux. First, suppose temperature constant in the V. Meyer tube, so that the heated vapour immediately displaces its own volume of heated air, which is then cooled. Let T = abs. temperature of vapour and air when first expelled, and T abs. temperature of cooled air leaving the end of delivery tube under water; let V = vol. of vapour formed and therefore of air expelled, and V vol. of air collected. Then V = 5^. Second, let there be two temperatures T and Tj within the tube, with corresponding volumes of equal masses of air V and V x . T V Then let V be expelled into the T! region, and thereby become V lf so that V x = --. Vi, not V, will now displace its own volume of air, which will be cooled so as to become, T 1 "V say, V 2 at the end of the delivery tube. Then V 2 = -^ = =^^- = ^ = V ; and LI Ijl similarly with any number of temperature zones. Thus a temperature gradient within the V. Meyer tube docs not affect the volume of air displaced from the end of the delivery tube. 32 CHEMICAL THEORY ii. The Method of Chemical Displacement. Somewhat related to the above principle is another by which the molecular formula of a compound may be determined, and so the atomic weight of a constituent element. Consider methane. The hydrogen in this compound can be displaced by chlorine in four distinct stages, the following sub- stitution products being formed : methyl chloride, methylene chloride, chloroform, and carbon tetrachloride. The carbon, how- ever, cannot be displaced fractionally. From these facts the inference is drawn that the molecule of methane contains 4 atoms of hydrogen and only 1 atom of carbon; but if methane is CH 4 , Dalton's problem of the number of atoms in the molecule is solved, and the atomic weight of carbon is 12. A similar argument may be applied to water. The composition of sodium hydroxide proves that half the hydrogen of the water molecule has been displaced by sodium. By no means, however, can any fraction of the oxygen of the water molecule be displaced. Thence it is concluded that water is H 2 O, and that the atomic weight of oxygen is 16. The principle of this method of fixing the magnitude of atomic weights may be stated thus: When th of the proportion of a constituent element in a chemical n compound can be displaced by another element, a molecule of the compound contains at least n atoms of that element, iii. The Method of Dulong and Petit 's Law. The specific heat of a substance is the ratio of the amount of heat required to raise unit weight of it through one degree of temperature to the amount of heat required to raise unit weight of a standard substance through the same temperature interval. The standard substance is water. In 1819 Dulong and Petit published the specific heats of thirteen elements, and showed that the product of specific heat into atomic weight is approximately a constant quantity, the average of which, on our modern atomic weight basis, is 6-4. In the following table, containing the elements studied by Dulong and Petit, modern values are given throughout. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 33 Element Specific Heat. Atomic Weight. Atomic Weight x Specific Heat = Atomic Heat. Bismuth 0-0305 208-0 6-34 Lead ... 0-0315 207-1 6-52 Gold ... 0-03035 197-2 5-99 Platinum 0-03147 195-2 6-14 Tin ... 0-0559 119-0 6-65 Silver ... 0-0559 107-88 6-03 Zinc 0-0939 65-37 6-14 Tellurium 0-0475 127-5 6-06 Copper . . . 0-09232 63-57 5-81 Nickel... , 0-10842 58-68 6-16 Iron 0-10983 55-84 6-36 Cobalt ... 0-10303 58-97 6-08 Sulphur 0-1712 32-07 5-49 The law of Dulong and Petit may therefore be stated thus: The specific heats of the solid elements are in the inverse ratio of their atomic weights. The product of specific heat and atomic weight, which is ap- proximately a constant, is called the atomic heat because it is the heat capacity of a quantity of an element proportional to its atomic weight. Thus, for example, 55-84 parts by weight of iron require the same amount of heat to raise them through one degree of temperature as, say, 208 parts of bismuth. But these quantities of the elements contain equal numbers of atoms. So, in the words of Dulong and Petit, " the atoms of all substances have exactly the same capacity for heat". In order to reach this result, however, Dulong and Petit made some drastic changes in the accepted atomic weight values, which aroused the opposition of Berzelius, their author. Thus, taking the atomic weight of sulphur as a true magnitude, they halved the atomic weights of the metals in relation thereto. This pro- cedure was, however, justified, even in the opinion of Berzelius, after Mitscherlich, his pupil, had arrived at similar conclusions by an application of the law of isomorphism. Now, since or Specific heat X atomic weight = 6*4 (approx.) 6*4 atomic weight = ^ = - (approx.), specific heat tere is a valuable method for fixing the magnitude of the atomic ( D 60 ) 4 34 CHEMICAL THEORY weight of an element. All that it is necessary to do is to deter- mine the specific heat of the element, and divide 6 4 by this value. It must be clearly understood, however, that the value thus obtained is only approximate, for the atomic heat value, 6-4, is only approximate, since it is a mean value, even if the specific heat is accurately known. The method serves to indicate what multiple of an accurately determined equivalent weight is the atomic weight. To divide 6 4 by the given specific heat of an element, and report the quotient as its atomic weight, is a gross error. The following illustration will make plain the use of Dulong and Petit's law: Marignac 1 found that 100 grm. of lead yielded 134 201 grm. of the chloride. The specific heat of the metal is 0-0315; find its atomic weight; 01 = 35-46. The equivalent weight of lead is found from the proportion: Wt. of chlorine : wt. of lead : : equivalent Cl : equivalent Pb, so 34-201 : 100 :: 35-46 : 103-68. The approximate atomic weight of lead, as indicated by its specific heat, is: ^' 4 - = 203-2. *0ol5 Therefore the atomic weight of lead is twice its equivalent weight; so Pb = 103-68 X 2 = 207-36. Dulong and Petit's law applies strictly only to solid elements, generally metals, whose atomic weights exceed 30. The specific heats of other solid elements vary with temperature, but become approximately constant at high temperatures, when they give an atomic heat value of about 5 5. iv. The Method of the Law of Isomorphism. Isomorphism is similarity of crystalline form. It was supposed by the earlier mineralogists that identity of crystalline form generally indicated identity of chemical composition; but it was shown by Mitscherlich in 1819 that compounds of analogous as well as identical composition crystallize in similar forms belonging to the same crystal systems. Thus di-sodium hydrogen phosphate and di-sodium hydrogen arsenate, which are now represented by the formulae Na 2 HPO 4 -12H 2 O and Na 2 HAsO 4 -12H 2 O, were found to be isomorphous. Careful measurements of the crystal angles of isomorphous salts show that these angles are not quite equal, i Marignac, (Euvres Completes, 1846, I, 186. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 35 but the following criteria of isomorphism have been established: (i) great similarity of crystalline form, (ii) analogous composition, (iii) power to form mixed crystals by simultaneous crystallization, (iv) power of crystal overgrowth, so that a crystal of one com- pound may form the matrix on which the growth of the crystal may be continued by the deposition of another substance. With regard to the second criterion, it must be noted that isomorphism is sometimes observed in pairs of compounds which are not chemically analogous, but have the same numbers of atoms within their molecules. Thus calc-spar (CaCO 3 ) is isomorphous with Chili saltpetre (NaNO 3 ), and aragonite (CaCO 3 ) with nitre (KN0 3 ). Mitscherlich stated the law of isomorphism as follows: " The same number of atoms combined in the same manner produce the same crystalline form; the crystalline form is independent of the chemical nature of the atoms, and is determined solely by their number and mode of combination." Nevertheless isomorphism such as that illustrated by the case of sodium phosphate and arsenate is the rule; that is to say, not only do the molecules of isomorphous compounds contain the same number of atoms similarly combined, but these atoms themselves are analogous, as, for instance, are phosphorus and arsenic. Indeed, isomorphism is taken to be a sign of chemical analogy. Therefore, for practical purposes, the law of isomorphism may be stated more briefly: The molecules of isomorphous substances contain equal numbers of atoms, which when not of identical are of analogous elements. The consequence of this law, when applied to the case already mentioned, is that the atomic weights of phosphorus and arsenic can be directly compared, and if one atomic weight is known the other is derivable from the results of chemical analysis. A simple numerical example is furnished by the following results of the analysis of the isomorphous salts potassium sulphate and potassium selenate, carried out by Mitscherlich: K 2 SeO 4 100 parts contain K 44-83 O 36-78 S 18-39 100-00 100 parts contain K 35-29 O 28-96 Se 35-75 100-00 127-01 parts contain 44-83 36-78 45-40 127-01 36 CHEMICAL THEORY In the third column is shown the proportion of selenium in an amount of the selenate which contains the same amounts of potas- sium and oxygen as are shown in the percentage analysis of the sulphate. Whence it follows that 45-40 parts of selenium take the place of 18-39 parts of sulphur. Now, the law of isomorphism declares that the ratio between these quantities is the ratio between the atomic weights of the two elements. Therefore, if the atomic weight of sulphur is 32*0 that of selenium is 32 - |Q X f 4 = 79-0. Io-o9 '^^~ The phenomena of isomorphism are somewhat confused by those of dimorphism and polymorphism. Thus calcium carbonate, as shown above, is dimorphous in calc-spar and aragonite; ammonium nitrate, NH 4 NO 3 , is tetramorphous, crystallizing in four distinct forms; arsenious and antimonious oxides, As 4 O 6 and Sb 4 O 6 are isodimorphous, that is to say, they are both similarly dimorphous. Nevertheless, the phenomena of isomorphism have been of value, not only in confirming atomic -weight magnitudes derived from other considerations, but in correcting erroneous magnitudes. For example, previous to the recognition of isomorphism, Berzelius regarded various metallic monoxides MO as dioxides MO 2 ; similarly, Fe 2 O 3 was written FeO 3 , Cr 2 O 3 was CrO 3 , and CrO 3 was Cr0 6 . But when this chemist recognized the isomorphism of chromates with sulphates he altered Cr0 6 to CrO 3 to agree with SO 3 , the oxide known to be present in sulphates. Consequently, the former Cr0 3 became Cr 2 3 ; and since chromic and ferric alums were isomorphous, what was formerly Fe0 3 became Fe 2 3 , and so FeO 2 became FeO. But compounds of copper, nickel, cobalt, man- ganese, zinc, and magnesium are isomorphous with corresponding iron compounds, and so if FeO 2 should be FeO the corresponding dioxides of all these metals should really be monoxides. This sweeping change would involve the halving of a number of accepted atomic weights; nevertheless, the change was made by Berzelius in accordance with the principles of isomorphism; and it was at once ratified by the law of specific heat, which required the same atomic weight magnitudes for the elements concerned. The alums which conform to the general formula M' 2 SO 4 .X 2 (SO 4 ) 3 .24H 2 O are amongst the best -known isomorphous compounds; and the EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 37 atomic weight of the element X can be determined by an analysis of its alum. For this purpose it is best to ignite the ammonium alum, which leaves a residue of the oxide X 2 O 3 . Thus, if a grm. of the alum leaves b grin, of oxide, the value of X is calculated from the expression: a : b = (NH 4 ) 2 S0 4 .X 2 (S0 4 ) 3 .24H 2 : X 2 O 3 = [132 + 2X + 288 + 432] :[2X + 48]. By this means Lecoq de Boisbaudran, who discovered gallium, found that 3-1044 grm. of its ammonium alum left on ignition a residue of 0' 5885 grm. of the sesquioxide; whence Ga == 701. v. The Method of the Periodic Law. An adequate account of the periodic law is necessary to an appreciation of its value as a guide to the magnitudes of the atomic weights of the elements; but this must be deferred to a later chapter. It will be sufficient to state here that a natural connection exists between the properties of an element and its atomic weight, and therefore that the order of magnitude of the atomic weight of an element may be judged from a study of the properties of the element and its compounds. Examples of this use of the periodic law will be given later. The application of the foregoing methods of atomic - weight determination is well illustrated by the case of carbon. 4. The Atomic Weight of Carbon The determination of the atomic weight of carbon has consisted of two parts: i. The determination of the order of magnitude, ii. The determination of the exact value. i. Determination of the Order of Magnitude of the Atomic Weight. Dal ton and his contemporaries attributed the value 6 to the atomic weight of carbon, but this was really only an equivalent weight. The following is the evidence that the atomic weight is about 12: (a) Avogadro's Theory. Never fewer than 12 parts by weight of carbon are present in a molecular proportion of any of its gaseous or volatile compounds. (6) Chemical Displacement. Use might be made of the argu- 38 CHEMICAL THEORY ment that, for example, hydrogen is displaceable from methane in four equal fractions, but carbon not fractionally; whence it follows that the formula for methane is CH 4 and the atomic weight of carbon 12. (c) The Law of Specific Heat. Although Dulong and Petit's law does not apply strictly to an element whose atomic weight is less than 30, and the specific heats of diamond and graphite differ widely from each other at ordinary temperatures, at 600 the specific heats of these two allotropic forms of carbon, which vary with temperature, become almost constant and equal, and give* an atomic heat of 5 5, if C = 12, a value which is comparable with the atomic heats > of analogous elements. (d) The Law of Isomorphism. The iodides of carbon and silicon are isomorphous; therefore they are similarly composed, and the atomic weights of carbon and silicon are in the ratio 12 : 28. (e) The Periodic Law. With an atomic weight of 12, carbon is appropriately placed in the periodic table between boron (11*0) and nitrogen (14-01); and is thus the first or typical element of the fourth group. If carbon forfeited its place owing to an alteration in the magnitude of its atomic weight, there is no vacant place in the periodic table which this element could fill, nor is any element known which could occupy the place of carbon. ii. Determination of the Exact Value of the Atomic Weight. There are two ways in which the atomic weight of carbon has been exactly determined: (a) By estimating the densities of its gaseous compounds. (b) By the combustion of carbon or the analysis of its compounds. (a) It has already been pointed out that gas or vapour density is simply related to molecular weight only when Avogadro's theory is rigidly true. This, however, is never the case; but an "ideal" density can sometimes be calculated from carefully ascertained data. This has been done 1 for the three gases: carbon monoxide (CO), carbon dioxide (CO 2 ), and acetylene (C 2 H 2 ). CO CO 2 C 2 H 2 Experimental density (O 2 = 1) 0-87495 1-38324 0-8194 "Ideal" density 0-87516 1-37516 0-81331 Molecular weight 28-005 44-005 26-026 Atomic weight of carbon ... 12-005 12-005 12-005 'Leduc, Ann. Cliem. Phyt., 1898 [vii.], 15, 5; 1910 [viii.], 19, 441. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 39 It appears from these figures that the method leaves nothing to be desired from the point of view of accuracy. (6) The atomic weight of carbon has been determined by several chemists by burning diamond or carefully purified graphite, weighing the carbon dioxide produced, and then calculating the result from the proportion: Weight of CO 2 : weight of C : : 32 + atomic weight C : atomic weight C. The following are the results, as originally given, and as cor- rected by Scott: Uncorrected. Corrected by Scott. 6 Dumas and Stas l 11-9975 11-9938 Erdmann and Marchand 2 12-0093 12-0054 Roscoe 3 12-0029 11-9973 Friedel 4 12-0112 12-0056 VanderPlaats 5 12-0031 12-0017 The ignition of organic silver salts, such as the acetate and tartrate, which leave a residue of pure silver, serves as a means of estimating the atomic weight of carbon; or the silver may be estimated electrolytically, as was done by Hardin, 7 with the follow- ing results, obtained with silver acetate and benzoate respectively: (1) C 2 H 3 O 2 Ag : Ag = 100 : 64-637 whence atomic weight of carbon = 12-000. (2) C r H 6 O 2 Ag : Ag = 100 : 47-125 whence atomic weight of carbon = 12-001. [Ag = 107-880, H = 1-00762, O = 16-00.] The above results are selected from amongst others as typical; they serve to show the degree of accuracy which has been attained in the determination of the atomic weight of carbon. This value lies between 12-000 and 12-005, and may be taken to be 12-003. 5. Determination of Molecular Weights (in Solution) The establishment of molecular weights by the determination of gas and vapour densities has been fully considered in the pre- ceding pages. By the study, however, of the influence of dissolved 1 Dumas, Pogg. Annalen, 1838, 44, 110. 2 Erdmann and Marchand, J. prakt. Chem., 1841, 23, 159. 8 Koscoe, Com.pt. rend., 1882, 94, 1180. Friedel, Bull. Soc. Cliim., 1884 pi.], 41, 100. 6 Van'der Plaats, Compt. rend., 1885, 100, 52. 6 Scott, Trans. Chem. Soc., 1897, 71, 550. 'Hardin, J., Amer. Chem. Soc., 1896, 18, 990. 40 CHEMICAL THEORY substances on the solidifying- and boiling-points of liquids, the molecular weights of substances in solution in these liquids may be determined; and it will be appropriate to consider here these newer methods of molecular- weight determination. It is well known that salt water freezes at a lower temperature than fresh water, and that sea ice when melted yields fresh water. Thus, when a dilute solution of salt in water is cooled, crystals of pure ice begin to separate from the solution at a temperature a little below 0. Blagden, in 1788, showed that the depression of the freezing-point of water by a dissolved salt is directly proportional to the amount of salt present. The boiling-point of water, on the other hand, is raised by salt in solution, and the elevation of boiling- point is directly proportional to the amount of salt dissolved. In 1883-4 F. M. Raoult discovered that not only are the depression of freezing-point and rise of boiling-point of a solvent proportional to the number of molecules of a particular substance in solution, but that equimolecular proportions of different substances have the same influence on the freezing- and boiling-points. Raoult's law, which applies equally to freezing- and to boiling- points of solvents, may be stated thus: The depression of freezing-point and elevation of boiling-point of a solvent by any dissolved substance are directly proportional to the number of molecules of the substance in solution, and consequently inversely proportional to its molecular weight. Or, otherwise: Equimolecular solutions, with the same solvent, have the same freezing- and boiling-points. Evidently these facts provide a means of comparing molecular weights, or of determining them if a substance of known molecular weight is chosen as a standard of comparison. It should be added that the extent to which a freezing- or boiling-point is affected depends also upon the solvent; consequently the first procedure is to determine the freezing- or boiling-constant (K) for a particular solvent by the use of a substance of known molecular weight. This constant is the number of degrees the freezing-point is lowered or boiling-point raised by 1 grm.-molecule of the substance dissolved in 100 grm. of the solvent. For instance, 2 grm. of cane sugar dissolved in 100 grm. of water cause a depression of the freezing-point, A = 0-11. Since EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 41 the molecular weight of cane sugar (C 12 H 22 O n ) is 342, the freezing- constant, K, for water, sometimes called the molecular depression, is 0-11 X 342 _ , Q -jr The same quantity of sugar dissolved in the same amount of water raises the boiling-point of the water 0*030. Therefore the boiling-constant or molecular elevation, K, for water is: 0-030 X 342 _ - -2~ When the freezing- or boiling -constant K for a solvent is known, an unknown molecular weight is calculated from observed data as follows: Let K = depression or rise caused by 1 grm.-mol. of a substance in 100 grm. of solvent (known constant). S = weight of substance taken. L = weight of solvent taken. A = observed depression or rise. M = required molecular weight. Then, since the observed depression of freezing-point or rise of boiling-point is directly proportional to the amount of substance taken, and inversely proportional to the amount of solvent, A K X S X 100 M _ 100 KS MxL AL ' Practical Methods. The prime necessity ior the experimental determination of molecular weights of substances in solution is a thermometer which will indicate accurately hundredths of a degree. If this ther- mometer is to be used both for freezing- and boiling-points, it would appear necessary for it to have a long range in addition. Real temperatures, however, have not to be read; only temperature differences. Consequently a thermometer has been devised by Beckmann with a range of about six degrees, the scale being divided into hundredths of a degree, and furnished with a reser- voir of mercury from which mercury can be added if low tem- peratures are to be recorded, and into which mercury can be driven when the instrument is to be used for higher temperatures. By the use of this device the same thermometer can be employed for temperatures near the freezing- as well as the boiling-point of water or other solvent. CHEMICAL THEORY The Cryoscopic Method. The determination of molecular weights by the cryoscopic method, that is, by observing the depression of freezing-point, is carried out in the apparatus of Beckmann shown in the figure. The tube (A), furnished with a side limb for the introduction of the substance, is fitted with a cork through which the thermometer (T) and platinum stirrer (S) pass. The lower part of this tube is surrounded by a wider tube (B) which provides an air jacket between the tube (A) and the freezing-mixture contained in the outer vessel (C). This freezing - mixture, whose temperature should be about 5 below the freezing-point of the solvent employed, is also furnished with a stirrer (S 1 ). A weighed quantity of water, or other solvent, is placed in the tube (A) and then frozen. Owing to under-cooling the temperature indicated by the thermometer falls below the freezing- point, and then quickly rises again, and be- comes stationary at that point as soon as ice separates. When the freezing-point of the solvent has been indicated on the arbitrary scale of the thermometer, a weighed quantity of the substance is introduced and the freez- ing-point of the solution determined. The amount of substance added should produce a depression of about 0*5. The determination may be repeated after the addition of a further quantity of substance. The reading should, however, be taken when a minimum quantity of the pure solid solvent has separated, so that the concentration of the solution may not be appreciably increased. The following are important freezing-constants (K): Water 18 6; acetic acid 39; benzene &d; phenol 73. EXAMPLE. Successive quantities of 0-317, 0-394, and 0-5152 grm. of a substance were dissolved in 18-054 grm. of benzene, the depressions of freezing-point being 0-278, 0-348, and 0-452 respectively; what is the molecular weight of the substance? The Fig. 5 EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 43 molecular lowering of the freezing-point of benzene (K) is 50. (Institute of Chemistry, July, 1902.) 100 KS M = i. M = ii. M = iii. M = AL ' 100 X 50 X 0-317 0-278 X 18-054 100 X 50 X 0-394 0-348 X 18-054 100 X 50 X 0-5152 0-452 X" 18-054 = 315-8. = 313-6. = 315-7. The Ebulliseopic Method Beckmann's Apparatus. The tube (A) (fig. 6) em- ployed in the Beckmann appa- ratus for determining elevation of boiling-point resembles that in which freezing-point deter- minations are carried out; but, in addition to the side tube for the introduction of the sub- stance, it is provided with an- other tube (B) fitted with a reflux condenser for the con- densation of the vapour arising from the boiling liquid. In order to prevent super-heating of the liquid, and consequent irregular boiling, a short piece of stout platinum wire (C) is fused into the bottom of the tube, which also contains some small beads which surround the lower part of the thermometer bulb, and serve to break and distribute the bubbles of vapour as they rise. In addition to this, the boiling-tube is sur- rounded with a wider vessel (D) packed with some non-conduct- ing material to prevent loss of heat by radiation, or sometimes Fig. 6 44 CHEMICAL THEORY with a glass envelope containing the vapour of the boiling solvent The whole apparatus stands upon a sheet of asbestos (E), below which the burner for heating is placed. In carrying out an experiment a weighed quantity of the solvent is heated until it boils briskly, and its temperature has become constant. If the condenser is acting efficiently the solvent should not lose in weight; but about 0-3 grm. should be subtracted from its weight to allow for the quantity required to wet the internal walls of the tube and condenser. After the boiling-point of the solvent has been recorded, the weighed quantity of the substance is introduced, and a reading again taken when the temperature has become constant. As in the case of freezing-point determinations, successive quantities of substance may be added to the same quantity of solvent, and corresponding readings taken. If much time elapses between the observations of the boiling-points of solvent and solution, it is necessary to read the barometer, and make a correction for change of atmospheric pressure during the interval. The Modified Landsberger Apparatus. A method of determining elevation of boiling-point, introduced by Sakurai, 1 modified by Landsberger, 2 Walker and Lurnsden, 3 and others, and lately by Turner and Pollard, 4 consists in raising the solvent to its boiling-point by passing into it the vapour of the same liquid boiling in another vessel. The vapour condenses, and its latent heat eventually causes the solvent to boil, although boiling-point after the addition of the substance is above that of the pure solvent. As the amount of the solvent continuously in- creases by condensation of vapour, it is estimated by weighing or measuring after condensation has been arrested instead of before heating is begun. By this method all possibility of superheating is avoided, and accurate results are rapidly obtained. The construction of the apparatus is shown in fig. 7. The vessel (A), about 16 cm. high and 3 cm. in diameter, is fitted with a two- holed cork through which pass the thermometer (T) and the delivery tube (B) by which vapour is conveyed to the bottom of the vessel 1 Trans. Chem. Soc., 1892, 61 , 994. *Ber., 1898, 31 , 461. 3 Trans. Ckem. Soc., 1898, 73 , 502. 4 Trans. Chem. Soc., 1910, 97 , 1184. Proc. Chem. Soc., 1913, 29 , 349. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 45 from the liquid boiling in the flask (F). upper part of the vessel allows uncondensed va- pour to pass into the outer vessel (D),' where it forms a vapour jacket and is then condensed, either here or by sub- sequent cooling after escaping by the side tube (E). The boiling-constants (K) of important liquids are: Water 52; ethyl alcohol 11 V; chloroform 39; benzene 27. EXAMPLE. Turner found that 1*150 grm. of diphenylamine (C 6 H 5 ) 2 NH, dissolved in 42 -82 grm. of chloro- form caused the boiling- point of the solvent to rise 0-618 weight of diphenylamine. 100 KS = 100 X 39 X 1 AL A small hole (C) in the Q M = Fig. 7 K = 39. Find the molecular 150 0-618 X 42-82 Theory for (C 6 H 6 ) 2 NH = 169-1. 6. Molecular Complexity The methods and results of determining the molecular weights of gases and vapours and of substances in solution have been reviewed in the preceding pages; and it appears that the molecules of substances in solution are of the same order of magnitude as those of the same substances in the state of vapour. For example, ferric chloride in a state of vapour at about 750 consists of mole- cules represented by the formula FeCl 3 , and the elevation of the boiling-point of ether or alcohol by dissolved ferric chloride points to the same molecular formula. The reason for this identity of molecular state is to be found in the fact that the vaporous state 46 CHEMICAL THEORY and the state of solution are analogous to each other, and that the process of vaporization of a solid or liquid, with the consequent distribution of its molecules through space, resembles the process of solution of the same substance, and the distribution of its molecules throughout the solvent. What, however, is to be said about the molecules of liquids themselves, and of solids? The subject may be approached by the consideration of dense vapours near their point of liquefaction or solidification. Ferric chloride vapour, for instance, at as low a temperature as possible has a density corresponding nearly to the molecular formula Fe 2 Cl 6 ; and nitrogen peroxide gas, which at elevated temperatures consists of NO 2 molecules, is composed largely of N 2 4 molecules near its temperature of liquefaction, whilst solutions of this substance in acetic acid and other solvents consist almost entirely of N 2 O 4 molecules. These facts are illustrations of molecular association', and such association is known to take place in the case of some pure liquids themselves. For instance, there is evidence that liquid water and liquid acetic acid contain the double molecules (H 2 O) 2 and (CH 3 COOH) 2 ; whilst liquid benzene, C 6 H 6 , forms no complex molecules. Little is known about the molecular state of solids; but, al- though there is no reason to believe that it is complex as a rule, cases are known in which the molecules of solids appear to assume great complexity. Silica, Si0 2 , furnishes a striking example of molecular association in the solid state. As a rule, silicon com- pounds are even more volatile than the corresponding carbon compounds; for instance, carbon tetrachloride boils at 76 and silicon tetrachloride at 59-6, chloroform, CHC1 3 , boils at 61 and silico-chloroform, SiHCl 3 , at 34; but, whilst liquid carbon dioxide boils at 80, silica requires the oxyhydrogen blowpipe to melt it, and the electric furnace to turn it into vapour. What is the reason for this extraordinary difference in physical properties between carbon dioxide and silicon dioxide ? The answer is that whilst carbon dioxide is represented correctly by the mole- cular formula CO 2 , silica must undoubtedly have the composition (SiO 2 ) n , where n is very large; for silicic acid whence silica is derived by dehydration is a colloidal substance, whose molecule is probably at least (H 4 SiO 4 ) 500 , and the process of dehydration would tend rather to increase the molecular complexity than otherwise, EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 47 by causing two or more contiguous molecules to condense into one as they lose water between them. 7. The Molecular Compositions of Compound Gases The evidence on which the molecular formula H 2 O for water is based has already been considered. There are a number of com- pound gases whose molecular formulae may be established by the application of the principles set 'forth in this chapter; and these will now be dealt with. It has already been shown that the molecules of hydrogen, chlorine, and oxygen are diatomic. This follows, it will be re- membered, from the fact that the hydrogen chloride formed from equal volumes of hydrogen and chlorine occupies twice the volume of each separate gas; and that steam occupies twice the volume of its constituent oxygen at the same temperature and pressure. By an extension of the principle here employed the number of atoms of a gaseous element within the molecule of a compound gas may always be determined. Thus, since it can be shown that 2 volumes of ammonia gas yield when decomposed 3 volumes of hydrogen and 1 volume of nitrogen, it follows, provided the nitrogen molecule is diatomic, that ammonia must be represented by the formula NH 3 ; for the molecular change on the decomposition of ammonia is: 2 mols. ammonia yield 3 H 2 + N 2 , consequently 2NH 3 = 3H 2 + N 2 . The argument may be put in another way. Since the volume of the ammonia is to that of the hydrogen as 2 : 3, the atomic con- centration of hydrogen in ammonia is to that in free hydrogen as 3:2; and since the volume of ammonia is to that of nitrogen as 2 : 1, the atomic concentration of nitrogen in ammonia is to that in free nitrogen as 1:2; whence the formula NH 3 follows. In the case of a gas containing a solid element, such as sul- phurous anhydride, the additional estimation of the density of the gas suffices to show how many atoms of the solid element it contains, provided the atomic weight of this element is known. Thus, (a) the gas produced by burning sulphur in oxygen measures the same volume as the oxygen; therefore the molecule of this gas contains 2 atoms of oxygen; (6) the density of the gas is 32, and its molecular weight consequently 64, whilst the weight of oxygen 48 CHEMICAL THEORY within its molecule is 32, and the atomic weight of sulphur is 32; consequently it follows that the molecule of the gas contains 1 atom of sulphur, and that its molecular formula is SO 2 . The following -statements epitomize the evidence for the mole- cular formulae of a number of the best-known gases: Hydrogen Chloride. That 1 volume hydrogen -\- 1 volume chlorine give 2 volumes hydrogen chloride is fundamental to the molecular theory. The following facts suffice to prove this relation: (a) Electrolysis of an aqueous solution of hydrogen chloride under suitable conditions yields equal volumes of hydrogen and chlorine. (6) Sodium amalgam removes the chlorine from hydrogen chloride gas, and the remaining hydrogen occupies half the volume of the hydrogen chloride. Water and Steam. (a) Electrolysis of acidified water yields hydrogen and oxygen in the proportion of 2 volumes of the former to 1 of the latter. (6) When a volume of electrolytic gas, i.e. a mixture of 2 volumes of hydrogen with 1 volume of oxygen is exploded in a eudiometer kept at a temperature above the boiling-point of water, the volume of the resulting steam is two-thirds the volume of the mixed gases. Therefore 2 vol. hydrogen + 1 vol. oxygen yield 2 vol. steam. Carbonic Anhydride. When carbon is burnt in oxygen gas the volume of the gas remains unaltered. Therefore a molecule of the gaseous product contains 2 atoms of oxygen (O 2 = 32). The density of carbonic anhydride is 22; therefore its molecular weight is 44 Within this molecular proportion are 32 parts (O 2 ) of oxygen, and therefore 12 of carbon. But 12 is the atomic weight of carbon. Therefore carbonic anhydride is CO 2 , and is rightly called carbon dioxide. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 49 Sulphurous Anhydride. When sulphur is burnt in oxygen the volume of the gaseous product is the same as that of the oxygen. The density of sul- phurous anhydride is 32, and its molecular weight 64. The atomic weight of sulphur is 32; therefore, by the same argument as applies to carbon dioxide, sulphurous anhydride is sulphur dioxide, SO 2 . Hydrogen Sulphide. When hydrogen sulphide gas, confined over mercury, is decom- posed by electric sparks, or when its sulphur is removed by means of tin heated in the gas and so converted into sulphide, the volume of the remaining hydrogen is equal to the volume of the original hydrogen sulphide, whose formula is consequently H 2 S n . That 7i = 1 is proved by the fact that the gas density is 17 and molecular weight 34; for of this 32 parts must be sulphur, and 32 is the atomic weight of sulphur. Thus the formula for hydrogen sulphide is proved to be H 2 S. Nitrous Oxide. Potassium, sodium, copper, and other metals remove the oxygen from nitrous oxide when heated in the gas, leaving nitrogen. There is some risk of nitrite being produced if the two former metals are heated too strongly in the gas, but strongly heated copper removes only the oxygen, and leaves all the nitrogen in a pure state. By this means it may be shown that nitrous oxide contains its own volume of nitrogen, and therefore that its mole- cule contains 2 atoms of this element. The density of nitrous oxide is 22, and its molecular weight is 44, and this weight contains 28 parts (N 2 ) of nitrogen, and therefore 16 parts of oxygen. Since 16 is the atomic weight of oxygen the molecule of nitrous oxide contains 1 atom of this element, and therefore the molecular formula for the gas is N 2 O. The same conclusion is reached by mixing nitrous oxide with its own volume of hydrogen and exploding the mixture. After condensation of the steam pure nitrogen remains equal in volume to the nitrous oxide. Thus it is shown, not only that nitrous oxide contains its own volume of nitrogen, but that the oxygen it contains would occupy half that volume, since it combines with a volume of hydrogen equal to that of the nitrous oxide. These facts are sufficient to establish the formula N 2 O for nitrous oxide. (D60) 5 50 CHEMICAL THEORY Nitric Oxide. If potassium is heated in nitric oxide the vigorous combustion which takes place results in the formation of nitrite and nitrate; but a spiral of iron wire heated electrically removes all the oxygen from the gas without combining with the nitrogen, and the residual nitrogen then occupies half the volume of the original nitric oxide. This proves that a molecule of nitric oxide contains 1 atom of nitrogen (N = 14). The density of nitric oxide is 15, and, since its molecular weight is 30, the molecule contains 1 atom of oxygen (0 = 16), and the molecular formula is NO. Ammonia. When ammonia solution is dropped into chlorine gas, hydrogen chloride is formed, and nitrogen set free. The experiment may be carried out in a long graduated tube, sealed at one end and pro- vided at the other end with a cork furnished with a tap funnel. Ammonia solution is passed through the funnel into the chlorine, and the reaction is accompanied by a greenish flame and fumes of ammonium chloride. After the ammonia has been added in excess, dilute sulphuric acid is introduced to combine with the excess of ammonia, after which water is allowed to enter until the gas in the tube is at atmospheric pressure, when the flow of water ceases. Then it is found that the gas, which is nitrogen, fills one-third of the tube. Since hydrogen and chlorine combine in equal volumes to form hydrogen chloride, the hydrogen of the ammonia from which the nitrogen has been liberated would have occupied three times the volume of this nitrogen. This shows that ammonia, when decomposed, yields 1 volume of nitrogen to 3 of hydrogen; but since the volume of ammonia gas which is thus decomposed is unknown, all that this experiment reveals is that the molecule of ammonia is (NH 3 ) n . The relation between the volume of ammonia and the volumes of its decomposition products may be determined by confining a measured volume of the gas over mercury and passing electric sparks through it until expansion ceases. The gas will then have been decomposed into a mixture of hydrogen and nitrogen which will occupy twice the volume of the ammonia. That this mixture consists of 3 volumes of hydrogen and 1 volume of nitrogen may be shown by adding excess of oxygen and ex- ploding the mixture. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 51 Thus for example: Volume of ammonia = 10'Ocu. cm. Volume of nitrogen -}- hydrogen after sparking = 20-0 cu. cm. Volume after addition of oxygen = 75*0 cu. cm. Volume after explosion = 52 5 cu. cm. Thus 22 5 cu. cm. of gas have disappeared, of which 15-0 cu. cm. must have been hydrogen. So it follows that 10-0 cu. cm. of ammonia were decomposed by electric sparks into 15 cu. cm. of hydrogen and 5 cu. cm. of nitrogen; and, as shown before, this proves the molecular formula NH 3 for ammonia. Phosphine. The case of phosphine differs from that of ammonia because, when the gas is decomposed by electric sparks, the liberated phos- phorus remains as a solid whose volume is negligible. Conse- quently, the proportion of phosphorus in the molecule must be discovered by density determination as in the case of sulphur dioxide, hydrogen sulphide, &c. Two volumes of phosphine, decomposed by electric sparks, yield 3 volumes of hydrogen. Therefore a molecule of the gas contains 3 atoms of hydrogen. The density of phosphine is 17, and its molecular weight 34. Consequently, the proportion of phosphorus within the molecular proportion of phosphine is 31. But 31 is the atomic weight of phosphorus. Therefore the molecule of phosphine contains 1 atom of phosphorus, and so its formula is PH 3 . Carbon Monoxide. Carbon monoxide can be converted into carbon dioxide by exploding it with oxygen, when it is found that 2 volumes of the gas combine with 1 volume of oxygen to form 2 volumes of carbon dioxide; or, since the molecular formula CO 2 and O 2 are known, in the equation, 2 C x O y + O 2 = 2 CO 2 , x and y both = 1, so that the molecular formula CO is proved. This conclusion is confirmed by the density of the gas, which is 14, whence the molecular weight is 28; and C = 12, O = 16, so that CO = 28. 52- CHEMICAL THEORY Methane, Ethylene, and other Hydrocarbons, If a certain volume of a hydrocarbon is exploded with a known volume of oxygen used in excess, the resulting moist gas, measured at atmospheric temperature and pressure consists of carbon dioxide mixed with unused oxygen. The volume of carbon dioxide formed is estimated by absorbing this gas in potassium hydroxide solution, and the total volume of oxygen used, part of which has produced carbon dioxide, and part water, is shown by the difference between the original and the remaining volume of oxygen. These data are sufficient to establish the molecular formula of the hydrocarbon. For, consider the gaseous hydrocarbon C x H y . The result of its explosion with oxygen is represented by the equation C*H y + (x + |)0 2 = *C0 2 + |H 2 0. The volume of steam formed and condensed is not measured; but when the volume of carbon dioxide, referred to that of the hydrocarbon as unity, which is #, has been ascertained, the value of y is found by subtracting this from the total volume of oxygen used, referred to the same standard, and multiplying the remainder by 4. When x and y are found, the formula of the hydrocarbon is settled. Vapour density will confirm the formula, but is not necessary to establish it. Methane. When a mixture of 10 cu. cm. of methane with 30 cu. cm. of oxygen is exploded, the resulting gas, measured at the same temperature and pressure, is a mixture of 10 cu. cm. of carbon dioxide and 10 cu. cm. of oxygen. Thus 1 volume methane requires for combustion 2 volumes oxygen, and yields 1 volume carbon dioxide. So in the equation C x H y + (x + DO, = *C0 2 + |H 2 0, x = 1 and | = 1 ; therefore the formula for methane is CH 4 ; or otherwise, because the volume of the carbon dioxide produced is equal to the volume of the methane, a molecule of the latter contains 1 atom of carbon; EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 53 and, because the volume of the oxygen required to burn the hydrogen of methane is equal to the volume of the methane, the atomic concentration of hydrogen in the methane molecule is twice what it is in the free hydrogen molecule; i.e. there are 4 atoms of hydrogen in methane. Thus the molecular formula for methane is CH 4 . Ethylene. When a mixture of 10 cu. cm. of ethylene with 40 cu. cm. of oxygen is exploded, the resulting gas, measured at the same temperature and pressure, is a mixture of 20 cu. cm. of carbon dioxide and 10 cu. cm. of oxygen. Thus 1 volume ethylene requires for combustion 3 volumes oxygen and yields 2 volumes carbon dioxide. So in the equation C x H y + (x + |)0 S = #C0 2 + |H 2 0, x = 2 and | = 1|; consequently the formula for ethylene is C 2 H 4 . Or, to employ the alternative argument, since the volume of the carbon dioxide produced is twice the volume of the ethylene, a molecule of this hydrocarbon contains 2 atoms of carbon; and since the volume of oxygen required to burn the hydrogen of ethylene is equal to the volume of the ethylene, this hydrocarbon contains 4 atoms of hydrogen. Thus, again, the molecular formula for ethylene is C 2 H 4 . In a similar way the molecular formula of any gaseous hydro- carbon may be established. SUMMARY EQUIVALENT WEIGHT. The equivalent weight of an element is that weight of it which combines with, or displaces from com- bination, unit weight of a standard element. ATOMIC WEIGHT. The atomic weight of an element is the ratio between the weight of its atom and that of the atom of a standard element. The standard is: = 16-00. DETERMINATION OF ATOMIC WEIGHT: (a) Exact estimation of chemical equivalent. (b) Decision as to order of magnitude. -64 -CHEMICAL THEORY Guiding principles: i. Avogadro's theory. ii. Chemical displacement, iii. Law of specific heats, iv. Law of isomorphism. v. Periodic law. PRINCIPLE OF CHEMICAL DISPLACEMENT. When -th of the n proportion of a constituent element in a chemical compound can be- displaced by another element, a molecule of the compound contains at least n atoms of that element. LAW OF SPECIFIC HEATS: DULONG AND PETIT'S LAW. The specific heats of the solid elements are in the inverse ratio of their atomic weights; or the atoms of the solid elements have the same capacity for heat. Specific heat X atomic weight = 6-4 (approx.) /> A or atomic weight = ^~-, . specific heat LAW OF ISOMORPHISM: MITSCHERLICH'S LAW. The molecules of isomorphous substances contain equal numbers of atoms, which when not identical are analogous. DETERMINATION OF MOLECULAR WEIGHTS: RAOULT'S LAW. The depression of freezing-point, and elevation of boiling-point of a solvent caused by any dissolved substance are directly proportional to the number of molecules of the substance in solution, and, consequently, inversely proportional to its molecular weight; or equimolecular solutions, with the same solvent, have the same freezing- and boiling-points. CHAPTER III VALENCY AND CHEMICAL CONSTITUTION It has already been seen, with regard to an element, that Atomic weight = n X equivalent weight; and it lias been stated that n is the valency or atomic value of the element. Now it was shown by E. Frankland that whilst 1 atom of tin is capable of combining with 2 atoms of oxygen to form the dioxide Sn0 2 , a molecule of the compound tin diethyl, Sn(C 2 H 5 ) 2 , or SnEt 2 , can combine with only 1 oxygen atom forming the compound SnEt 2 O. Thus it appears that the tin atom has a certain saturation capacity, that it can combine with not more than 2 atoms of oxygen or their equivalent, and that its power of combining with oxygen is diminished by the extent to which it is already combined with other atoms or groups of atoms. The principle was further illustrated by Frankland by reference to such compounds as NH 3 , NI 3 ; PH 3 , PC1 3 , in which the atoms of nitrogen and phosphorus combine with 3 atoms of hydrogen or halogen; and by Kekule, who showed that the carbon atom can combine with four other atoms, as in the compounds CH 4 , CH 3 C1, CHC1 3 , CC1 4 . The phenomenon here illustrated is now called valency, about which the following statement may be made: The valency of an element indicates the number of other atoms with which one of its atoms can directly combine. An atom may be uni-, bi-, ter-, quadri-, quinque-, sex-, sept-, or even octa-valent; 1 equivalent terms are monad, dyad, triad, l The Greek prefixes mono-, di- t tri-, &c., which, when attached to valent, make hybrid words, are now being abandoned, 65 56 CHEMICAL THEORY tetrad, &c. In the compounds cited above the nitrogen and phos- phorus atoms are tervalent, and the carbon atom is quadrivalent; whilst the hydrogen, chlorine, and iodine atoms are univalent. Hydrogen is never more than univalent, and therefore its atom is chosen as the standard of valency; chlorine is univalent with regard to hydrogen and metals, and, indeed, probably to all ele- ments except oxygen; it may therefore replace hydrogen as a standard. The following hydrides exhibit the valency of a number of elements: Valency 1 2 3 4 FH OH 2 NH 3 CH 4 C1H SH 2 PH 3 SiH 4 BrH AsH 3 IH SbH. and the following halides 1 illustrate valency more extensively: Valency 1 23 4 5678 NaCl OC1 2 BC1 3 CC1 4 PF 6 SF B OsF 8 KI ZnCl 2 PC1 3 SiCl 4 AsF 5 TeF AgCl HgCl 2 AlClj SnCl 4 SbF 6 UF 6 Oxygen is here shown to be bivalent. It is seldom other than this; and if oxygen is bivalent, the large number of oxides that exist may be classified to show valency, on the assumption that the valency of an element is equal to twice the number of oxygen atoms with which one of its atoms combines. Valency 1 2 3 4 5 6 7 8 Na 2 O MgO B 2 3 C0 2 NA SO, CIA OsO 4 K 2 O CaO A1 2 3 Si0 2 P 2 6 Cr6 3 (IA) Ku0 4 Ag 2 ZnO Fe 2 3 PbO 2 BiA U0 3 Mn 2 O 7 These oxides are in a different category from that of the foregoing hydrides and halides. In those the molecular formulae have in all cases been established by vapour density or other measurements, and the valency of the element concerned is directly indicated by the number of attached hydrogen or halogen atoms. The formulse for these oxides, however, are not in all cases molecular. Silica, for instance, is (SiO 2 ) n , and phosphoric oxide is (P 2 O 5 ) 2 . Moreover, according to the above statement regarding valency, this property cannot strictly be judged from oxides at all, because these compounds do not exhibit, attached to a nuclear 1 Halide = fluoride, chloride, bromide, or iodide. VALENCY AND CHEMICAL CONSTITUTION 57 atom, a number of peripheral atoms corresponding to its valency. Nevertheless, there is good reason to regard the valencies indicated by oxides such as those in the table to be correct. The establishment of the idea of valency was soon followed by a device by which the facts of atomic union were represented graphically. Bonds were introduced by Couper to show the joining together of the atoms in the following way: ' H H Y Cl H, H O H, N, H C H. i H Thus graphic or constitutional formulae were constructed, with bonds to show units of valency, or units of affinity, which they might be called, if they are thought of as standing for the forces by which the atoms are united. Oxides containing bivalent oxygen are represented by graphic formulae, such as the following: /\ O O Mg=O, B-O-B, O=C=O, \N_O- N/ , O=S=O, \ _ / r\r N>n ii \ / O O I! or O=B-O-B=O, O O O O=C1-O-C1=O, 0=Os=O. A Formulae like these provoke several questions. For example, two graphic formulae are here given for B 2 3 . Does this mean that the substance may have one or other constitu- tion, according to circumstances, or does it mean that we do not know which is the correct graphic formula, and therefore that the student may make choice between them according to his fancy? It means the latter rather than the former in this case. We know that boron is tervalent and oxygen bivalent. Both formulae repre- sent these facts, and there 'are probably no other known facts to support one formula rather than the other. Is, then, the writing of graphic formulae merely an interesting geometrical exercise based on the facts of valency alone? Con- 58 CHEMICAL THEORY sider, for example, a substance with the formula C 3 H 6 O. Since carbon is quadrivalent, oxygen bivalent, and hydrogen univalent, two graphic formulas are possible for this substance: H O H H H H H-C C C H and H-C-C C=O. H (i) H H H (ii) Does it matter which formula is adopted? The answer is that two quite different substances are known, both of w^hich are C 3 H 6 O; and one of which, acetone, certainly has the constitution (i), whilst the other, propaldehyde, as certainly possesses the con- stitution represented by (ii). Thus graphic formulae are constitutional formulae, and only so far as they represent the ascertained constitution of compounds are they valid. Another question arises in connection, for instance, with the compound S0 3 . Why should not the oxygen atoms be considered univalent and the sulphur atom tervalent here; why should not the latter of the two formulae O=S=O, O-S O, be considered as good as the former? Now the constitution of a compound is to be judged by its chemical relationships and transformations; and closely related to sulphur trioxide, S0 3 , are sulphuryl chloride, SO 2 C1 2 , and sulphuric acid, H 2 S0 4 , which is formed by the action of water on either of the former compounds. The constitutions of these three substances are well represented by the graphic formulae whereas if the oxygen atom which is replaced by two other atoms or groups had been originally united by a single bond to sulphur, its replacement could not have been so satisfactorily accounted for. It might further be suggested, however, that the other two oxygen atoms in these molecules, even if themselves bivalent, need not be attached by double bonds to sulphur; that the formulae (X O Cl O O H VALENCY AND CHEMICAL CONSTITUTION 59 would equally well represent the facts. These formulse, however, introduce a chain of 2 oxygen atoms, and this mode of linkage suggests specific properties which are absent from the above com- pounds. Thus the representation of double bonds in the graphic formula for SO 3 appears warranted. But here, again, it is in the constitution of carbon compounds that the conception of the double bond becomes more perfectly justified. Indeed, the question of mode of attachment of two of the three oxygen atoms in sulphur trioxide is not of practical importance, since the properties of this substance and its derivatives are sufficiently represented by the formula (S0 2 )0, S0 2 C1 2 , S0 2 (OH) 2 , &c. This latter formula is often referred to as the constitutional rather than the graphic formula for sulphuric acid. Variability of Valency. Early in the development of the theory of valency the question arose whether valency is a fixed and inherent property of an atom, like its mass, or whether it can vary under varying circumstances. Kekul6, who showed the quadrivalency of carbon, believed valency to be unalterable; and the study of carbon compounds alone ap- peared to justify Kekule's opinion. The following compounds were cited by Kekule' to illustrate the constant quadrivalency of carbon: H H Cl Cl H H, H C Cl, Cl-C Cl, H-C Cl, O=C<( , i A i- A C1 0=C=0, S=C=S, H-C=N. Frankland, on the other hand, observed that nitrogen formed not only NH 3 , in which the element is evidently tervalent, but also NH 4 C1, in which it was apparently and is actually quinquevalent. Thus was expressed the idea of a maximum potential valency, and an actual valency, exercised in specific compounds, which might be less than this. And it was observed that the actual valency fre- quently fell short of the potential valency by two units, as, for example, in the pairs of compounds NH 3 , NH 4 C1; P 2 O 3 , P 2 O 6 ; SO 2 , SO 3 ; SnCl 2 , SnCl 4 ; so it was supposed that when valency decreased from the maximum it was always by two units, and that consequently the valency 60 CHEMICAL THEORY of an element remained either odd or even. Then it was thought that the two valencies that remained disengaged in the lower com- pounds satisfied each other, so that no valencies remained free. There are, however, some notable exceptions to this supposed rule, and it cannot be regarded as a natural law. Examples of these exceptions are shown in the sets of compounds NO, N 2 3 , N0 2 , N 2 6 ; C1O 2 , C1 2 O 7 ; IO 2 , I 2 O 6 ; FeCl 2l FeCl 3 ; InCl, InCl 2 , InCl 3 ; WC1 5 , WC1 6 . The lower compounds are unsaturated, and combine with oxygen, chlorine, &c., to form higher compounds. When rise of temperature causes dissociation it thereby causes the acting valency of the nuclear atom or atoms of a compound to diminish. Thus when ammonium chloride, NH 4 C1, in which the nitrogen atom is quinquevalent, dissociates into ammonia and hydrogen chloride, the nitrogen atom becomes in consequence tervalent. Similarly tungsten hexachloride, WC1 6 , dissociates at high temperature into the pentachloride WC1 5 and chlorine. Occasion- ally dissociation involves the halving of molecules, with consequent reduction of valency, as the following examples show: 0, ,0 - Ck /Cl Ck /Cl >Fe - Fe< -* CK \C1 d' C\\ /Cl Cl\ /Cl (_-'l~/Al - .f\..i\ v^i " \*>i /j CK X C1 CK It may be added that, unless there is reason to the contrary, association into double molecules, such as those represented above, is supposed to be effected by means of 1 unit of valency. Mole- cular association in liquids and solids must also be accounted for by the exercise of additional valencies. Thus liquid water contains double molecules, or molecules of even higher complexity; and the existence of these complex molecules is accounted for by assuming oxygen to be quadrivalent, thus: since oxygen is known to be quadrivalent in some other compounds. The existence of double salts and salts with water of crystal- VALENCY AND CHEMICAL CONSTITUTION 61 lization cannot be explained by the narrower conceptions of valency. Consider, for example, potassium alum, K 2 SQ 4 'A1 2 (SO 4 ) 3 24 H 2 O. The constitutional formula for potassium and aluminium sulphates are respectively K CK /O-$O 2 O\ ~X>SO 2 and Al^-O SO 2 O^Al; K ~^ \ -S0 2 -0/ but it is difficult to see how these formulae are to be united together, and 24 molecules of water to be incorporated in the scheme as well. At one time it was customary to describe such compounds as "molecular" rather than "atomic", but such a distinction is no longer regarded as valid, and several theories have been proposed to account for the constitution of these compounds on the ground that auxiliary or latent valencies come into play in their formation. These theories cannot, however, be considered here. The variation of valency with the kind of compound formed has been illustrated in the lists of hydrides, halides, and oxides already given. Thus it appears that whilst the valency of an element towards oxygen and the halogens 1 may rise as high as 8, valency towards hydrogen is never greater than 4; no single atom is known to combine with more than 4 hydrogen atoms. Now hydrogen and oxygen are reciprocally related, and it is a noteworthy fact that as the valency for hydrogen diminishes in a series of elements with increasing atomic weight, the valency for oxygen correspondingly increases, and the sum of the oxygen and hydrogen valencies remains equal to 8. This is shown in the following compounds, although fluorine and bromine fail to form oxides, and selenium and iodine are known to form not the oxides Se0 3 and I 2 O 7 , but only the oxyacids corresponding to them; moreover tin forms no hydride, but a corresponding methyl com- pound. CH 4 CO 2 NH N A FH SiH 4 SiO 2 PH 3 P 2 6 SH 9 SO 3 C1H C1 2 O 7 GeH 4 GeO 2 AsH 3 As 2 O 5 SeH 2 (SeO 3 ) BrH Sn(CH 3 ) 4 SnO 2 SbH 3 Sb 2 6 TeH 2 Te0 3 IH (IA) With regard to valency for the halogens, it must be noted that as a rule halides are not so stable as the corresponding oxides. For example, NC1 3 is so unstable as to be highly explosive, whilst N 2 O 3 1 The halogen elements are fluorine, chlorine, bromine, iodine. 62 CHEMICAL THEORY does not split off oxygen; PC1 6 dissociates into PC1 3 and C1 2 , whilst P 2 5 is stable; S0 2 may be united with oxygen to form S0 3 , whilst SC1 4 loses chlorine when removed from the freezing-mixture in which it has been made. Fluorides, however, are much more stable than the other halides: PF 6 and SF 6 are stable gases, and the existence of OsFg, 1 in addition to OsF 6 and OsF 4 , shows that a valency of 8 towards a halogen is realizable. The Double Bond in Carbon Compounds. Consider the two hydrocarbons ethane, C 2 H 6 , and ethylene, C 2 H 4 . The former is saturated, the latter is unsaturated; that is to say, it is capable of combining with 2 more hydrogen atoms or their equivalent. This state of unsaturation of ethylene is represented by a double bond, the graphic formula for the two compounds being Ethane. Ethylene. H H H H H Q C H and H-C=C-H. The question may be asked whether the double bond is simply employed to keep up the appearance of the quadrivalency of carbon, or whether it has any real meaning; whether, indeed, carbon is not really tervalent in ethylene, so that the formula might as well be H H H C C H. This question may be answered in the negative for several reasons. First, no such compound as CH 3 CH 2 is known, in which 1 atom of carbon is quadrivalent, whilst the other is tervalent; so that both atoms must be either saturated or unsaturated. Here, at least, the idea that the 2 unsaturated atoms in ethylene satisfy one another appears justified; and the double bond expresses their mutual dependence. Further, the double bond between carbon atoms, the ethylene linkage, as it is called, expresses something more than unsaturation; for the nature of this union differs from that represented by the single linkage. It is weaker than the single linkage, for when i Ruff and Tschirch, Ber., 1913, JJ6, 929. VALENCY AND CHEMICAL CONSTITUTION 63 a compound contains a chain of carbon atoms in which there is a double linkage, this is the point at which the chain breaks when the compound comes under disruptive influence. The fact that the double is weaker than the single linkage shows that no mechanical significance must be attached to bonds. There is still a third characteristic of the double bond, which, however, can only be made clear by the study of the stereo- chemistry of carbon compounds. " Chemistry in Space." How far, it may be asked, is the graphic formula H H-C-H supposed to represent the real configuration of the molecule of this simple hydrocarbon, methane? The answer may at once be given that it is probably an imperfect representation of the truth, because it is a flat formula, a formula in two dimensions, whereas matter exists in three dimensions; the formula has length and breadth, but the molecule of methane has thickness as well as length and breadth. Moreover, the adequacy of the formula may be tested in a very simple way. The formula suggests that there might be two methylene chlorides, CH 2 C1 2 , Cl I H C H and Cl H C-C1, in which the two chlorine atoms are respectively opposite and adjacent to each other. Two such chlorides do not, however, exist; therefore a method of formulation must be found which does not suggest their existence. Only when the valencies of the carbon atom are equally distributed in tridimensional space is this requirement met; that is to say, when they are directed from the centre to the angular points of a regular tetrahedron, thus: 61 CHEMICAL THEORY Since this figure is symmetrical, the positions of the .2 hydrogen and 2 chlorine atoms in methylene chloride shown upon it may be interchanged in any way without causing a difference in the relative positions of these 4 atoms. This conception of the disposition in space of the valencies of the carbon atom, which is due chiefly to van 't HofT, has been very fruitful in organic chemistry. The aspect of the science thus suggested has been called "chemistry in space", or stereochemistry. Space-formulae should, of course, be applied to all chemical compounds, and some progress has been made with elements other than carbon; but these formulae are mainly of use in elucidating the structure of carbon compounds, where the question of constitution is of such vital importance. It may be added that double and triple bonds are represented stereochemically by the joining of two tetrahedra along their edges and adjacent surfaces respectively. For example, ethylene, CH 2 CH 2 , and acetylene, CH~CH, are thus represented: Fig. 9 The Criterion of Valency. The facts recorded in the foregoing pages suggest that valency might furnish a means of chemical classification of the elements, were it not that the exercise of this property varies somewhat irregularly. On the other hand, an independent classification of the elements might be expected to furnish information regarding valency. Such information is supplied by the Periodic Classifica- tion, which will shortly be studied. It will be sufficient to state here that in this classification the elements fall into nine groups Groups O to VIII; and that the maximum valency of each element appears to be identical with the number of the group which contains it. Thus, the no-valency elements of the argon family are in Group O, the univalent metals of the alkalis in Group I, the bivalent metals of the alkaline earths in Group II, and so on. VALENCY AND CHEMICAL CONSTITUTION 65 Very seldom does the acting valency of an element exceed that indicated by the group to which it belongs; more often it falls below it. For example, the halogens belong to the seventh group, and should therefore have a maximum valency of seven. This is realized by chlorine in C1 2 O 7 , and by iodine in H 6 I0 6 ; but not by fluorine or bromine. Iron, nickel and cobalt, as well as osmium, ruthenium, &c., belong to the eighth group; but whilst the two latter metals realize octavalency an OsO 4 and RuO 4 , the three former metals appear never to be octavalent. Nevertheless the Periodic Law is the best criterion of the valency of an element. The Nature of Valency. A study of the operation of valency, however detailed, or the graphic representation of the union of elements in chemical compounds by the use of bonds or solid geometrical figures, leaves the nature of valency itself quite unexplained. It may be said that the force which binds the atoms together is chemical affinity; but this explains nothing, and, moreover, the term " chemical affinity " has received a meaning in physical chemistry which is not closely associated with the idea of units of valency acting in specific directions through space. More than a century ago H. Davy l expressed the opinion that " electrical effects are exhibited by the same bodies, when acting on masses, which produce chemical phenomena when acting by their particles". Berzelius extended this idea in his electro-chemical theory, whence is derived the method of classifying the elements as electro-positive and electro-negative. Faraday, later, showed that during electrolysis a definite quantity of matter is always associated with a definite quantity of electricity, a fact which suggests that electricity as well as matter is atomic. Quite recently atoms of electricity have been recognized in the electrons produced in the Crookes tube and during the disintegration of radium, and the idea has been conceived that electrons are the binding media in chemical compounds, the stuff of which " bonds " are made. Now electrons are atoms of negative electricity, and it is supposed that the atoms of the metals become electro-positive as they lose negative electricity in the form of electrons, whilst the atoms of non-metals become electro-negative as they gain electrons. i H. Davy, Phil. Trans., 1807, 1. ( D 60 ) 6 66 CHEMICAL THEORY Sodium chloride, represented by the graphic formula Na Cl, is really NaECl, E being an electron which has come from the atom of sodium and serves to bind it to the atom of chlorine. Sodium and chlorine are univalent in sodium chloride, for only 1 electron is concerned in the union of their atoms. The valency of an atom thus depends upon the number of mobile surface electrons it may lose or receive, and this number may rise as high as 8. This electronic theory of valency is due to J. J. Thomson and Ramsay. SUMMARY VALENCY. The valency of an element indicates the number of other atoms with which one of its atoms can directly combine. CHAPTER IV CLASSIFICATION OF THE ELEMENTS The Periodic Law When the elements are regarded collectively, and in view of their ascertained atomic weights and properties, two considera- tions present themselves: (i) How may the elements be classified? (ii) What is their origin? These considerations are related, for the classification of material species is likely to lead to questions regarding the origin of such species. Probably the first systematic classification of the elements was derived from the electro-chemical theory of Berzelius, to which reference has already been made. This theory grew out of the facts of electrolysis. Thus, if, for example, an electric current passes through an aqueous solution of sodium chloride, the sodium appears at the cathode or negative electrode, and the chlorine at the anode or positive electrode. Consequently sodium was regarded as electro-positive, being attracted to the electrode of opposite sign, while chlorine was, for a similar reason, electro-negative. Or, more generally, metals were considered to be electro-positive and non- metals electro-negative. Further, it was recognized that some metals are more electro-positive than others, power of metallic replacement being regarded as a criterion of electro-positiveness. For example, since zinc displaces copper from copper sulphate in solution, zinc is more electro-positive than copper; and, conversely, since chlorine displaces iodine from potassium iodide in solution, chlorine is more electro-negative than iodine. So this method of classification served not only for the distinc- tion of metals from non-metals, but also for the recognition of metallic and non-metallic intensity. When the atomic weights of a sufficient number of the elements had been established with some degree of accuracy, it was perhaps inevitable that numerical relationships should be sought for between 67 68 CHEMICAL THEORY them, and that attempts should be made to discover a connection between the properties of an element and its atomic weight. The first attempt of this kind was made by Dobereiner, in 1829, who showed that in various triads of related elements the central member of each group possesses properties and an atomic weight which are approximately the mean of the properties and atomic weight of the extreme members of the triad. These triads are: lithium, sodium, potassium; calcium, strontium, barium; phosphorus, arsenic, antimony; sulphur, selenium, tellurium; chlorine, bromine, iodine. It will be sufficient to give numerical details for the first and last of the triads. Element. Atomic Weight. Differences. Mean of Extreme Atomic Weights. Lithium 6-94 , 06 Sodium 23-00 16 ' >1Q 23-02 Potassium 39-10 Chlorine 35-46 44 46 Bromine 79-92 ' m 81-19 Iodine 126-92 It will be observed that the atomic weight of sodium is almost exactly the mean of the atomic weights of lithium and potassium, but that the atomic weight of bromine is considerably less than the mean of the atomic weights of chlorine and iodine. The relations suggested by these* triads are therefore approximate only. It has been objected, moreover, that triads should not be made up to the exclusion of other related elements; that, for example, there are four halogens, and that it is arbitrary to exclude fluorine by form- ing a triad with the other three. But when it is recognized that fluorine differs from the other halogens, not only in atomic weight relationship, but also widely in the properties of its compounds, this objection loses force. So that without doubt the relationships shown by Dobereiner's triads are remarkable; nevertheless their value is historic only, for they are now merged in the generalization known as the periodic law. Another kind of triad was, however, observed by Dobereiner, in which the three related elements have nearly identical atomic weights. These triads are: Iron 55-84 Cobalt 58-97 Nickel 58-68 Ruthenium 101-7 Rhodium 102-9 Palladium 106-7 Osmium 190-9 Iridium 193-1 Platinum 195-2 CLASSIFICATION OF THE ELEMENTS 69 They also find a place in the periodic classification. Dobereiner's observations were limited to the elements cited above. These observations could not give rise to a generalization, since they were concerned with only a minority of the elements; the majority did not form triads; and therefore it is difficult to see what significance could have been attached at the time to the existence of these triads. In 1863-6 J. A. R. Newlands arranged the elements in ascend- ing order of their atomic weights, commencing with hydrogen, thus : H Li Gl B C N O F Na Mg Al Si P S Cl K Ca Cr Ti Mn Fe, &c. Thus he discovered that the eighth element is " a kind of repe- tition of the first", the ninth a repetition of the second, and so on; Na, for example, is a repetition of Li, Si of C, Cl of F. This dis- covery he called the Law of Octaves. "Members of the same group of elements stand to each other in the same relation as the extremities of one or more octaves in music." This simple "law" did not apply to the elements of higher atomic weight; even in the above table manganese is wrongly classified with phosphorus; and it was suggested by a contemporary of Newlands that it would be as useful to arrange the elements in alphabetical order as in the order of their atomic weights! Never- theless, the law of octaves is valid as an introduction to the periodic law. About the year 1869 Mendeleeff arranged the elements in the order of their atomic weights, and discovered, quite independently of Newlands, a periodicity in the properties of the elements so arranged. The fact of this periodicity he enunciated as the Periodic Law, which may be expressed as follows: The physical and chemical properties of the elements and their compounds are periodic functions of the atomic weights; or If the elements are arranged in the order of increasing atomic weight, their properties vary definitely from member to member of the series, but return to a more or less similar value at fixed points in the series. The figure of a spiral, first employed by Chancourtois in 1862, gives graphic expression to this idea, which is essentially the same 70 CHEMICAL THEORY as that of the law of octaves. For if seven vertical lines are drawn at equal distances apart on the cylindrical surface of the spiral, and consecutive elements are placed where the spiral inter- sects these vertical lines, the eighth element will be found under the first, the ninth under the second, &c.; it is thus suggested that the eighth element has pro- perties analogous to but not identical with those of the first, that the ninth element is similarly related to the second, and so on. The periodic system may now be developed. Hydrogen, the element of lowest atomic weight, has no analogue amongst the elements; consequently it became the sole member of Series 1 in Mendele'eff's system. Series (2) and (3) are the same as in the octaves of Newlands, thus: Fig. 10 (2) Li Gl B (3) Na Mg Al The next two series are: (4) (5) K Cu Ca Zn Sc Ga C Si Ti Ge V As O F S 01. Cr Se Mn Br, Fe Co Ni and are linked together by the triad Fe, Co, Ni; since to place these three elements in the consecutive positions occupied by Cu, Zn, Ga, thus displacing all that follow them, would be not only to obliterate periodicity from the scheme, but also to ignore the peculiar relations these three elements bear to one other as members of a triad. Series (6) and (7) connected by another triad are: (6) (7) Rb Ag Sr Cd Y In Zr Sn Cb Sb Mo - Te I. Ru Rh Pd Now it will be seen that Series (4) and (6) begin with the alkali metals K and Rb, whilst Series (5) and (7) begin with the metals Cu and Ag, which, whilst allied to each other, differ widely from the alkali metals. Similar differences exist between subsequent members of odd and even series. Elements in vertical columns constitute groups; of which, according to Mendele'eff, there were eight: seven groups corresponding with Newlands's octaves, and an eighth group in which Dobereiner's triads of nearly equal atomic weight were placed. When the inert gases CLASSIFICATION OF THE ELEMENTS 71 were discovered, these were placed in a group by themselves: Group O, which preceded the other groups. Except with regard to the elements of Series (2) and (3) Groups I to VII were sub- divided into A and B Sub-groups, to show the above-mentioned differences between consecutive members of the same group. Thus the complete periodic system takes the following form. Groups I II III IV V VI VII VIII Sub- A B A B A B A B A B A B A B groups Series 1 H 2 He Li Gl B c N F 3 Ne Na Mg Al Si P s Cl 4 Ar K Ca Sc Ti V Cr Mn Fe Co Ni 5 Cu Zn Ga Ge As Se Br 6 Kr Kb Sr Y Zr Cb Mo Ru Rh Pd 7 Ag Cd In Sn Sb Te I 8 Xe Cs Ba La Ce ' Rare 9 __ Earth __ Metals 10 Ta W Os Ir Pt 11 Au Hg Tl Pb Bi 12 Ha Th u Oxides X 2 O XO X 2 3 XO 2 XA X0 3 X 2 7 XO 4 The above arrangement may be improved upon, and a clearer view obtained by recognizing the existence of short and long periods. Thus Series (2) and (3) constitute short periods; Series (4) and (5), with the linking elements of the eighth group, form one long period. Other long periods follow, and the whole scheme is shown in the chart on p. 72. The great advantage of this mode of presenting the Periodic System is that the A and B Sub-groups are separated, so that elements which have little resemblance to one another are not classified together. For instance, it may well be objected that Cu, Ag, and Au, being very unlike the alkali metals, should not be placed with them in Group I. This objection is sufficiently answered when it is shown that these metals occupy positions near the centres of the long periods, whilst the alkali metals are quite differently situated at the beginning of these periods. Similar remarks apply to the relation between manganese and the halogens. saoiaad ONOI 1 -t. 1 1 - t* w ff 1 1 CD . 5 C/7 1 3 m J CO Q O e 1 * o 0. H QC _. U C! 1-1 1 P ro | fc N TJ U 1 bo ffi cq O 3 If 1 p T 2 *4 A 1 J co Q. U (3 s 1 u i D ..- cc o CQ j < 3 B 1 w O ..... j 5 bo C s 1 1 1 - a S u 1 * * CO K z > g " g l [s 2 i 10 o: 1 ^' / p N u' 1 H t >. a 1 1 ro <3 c/5 ctf CQ 1 K OJ M S N U 1 1 - ^ M * I 1 CO Q. ID O tE cp CQ ^ CO CD O or cp QD 3 CO CLASSIFICATION OF THE ELEMENTS 73 The arrangement of elements in any group now takes this form, illustrated by Group I: (A) Jt 1 (B) Na K Cu Rb Ag Cs Au Attention may next be drawn to atomic -weight differences between analogous elements in consecutive short and long periods. These differences are shown for a number of the elements in the following tables: THE Two SHORT PERIODS He Li Gl B C N O F Ne Na Mg Al Si P S Cl Differences 16-21 16-06 15-22 16-1 16-3 17-03 16-07 16-46 THE FIRST Two LONG PERIODS Ar K Ca Sc Ti V Cr Mn Fe Kr Rb Sr Y Zr Cb Mo Ru Differences 43-02 46-35 47-56 44-9 42-5 42-44 44-0 45-86 Co Ni Cu Zn Ga Ge As Se Br Rh Pd Ag Cd In Sn Sb Te I Differences 43-93 48-02 44-31 47-1 44-9 46-5 45-24 48-3 47-0 It will be observed that the differences in the short periods are approximately 16, and in the long periods about 45; in the short periods 8 elements intervene before a recurrence of properties, and in the long periods 18 elements. The differences are by no means constant, for no mathematical relations exist between the atomic weights; but anomalies are seen in the differences between krypton and argon, palladium and nickel, tellurium and selenium, in accordance with the anomalies in the atomic weights of argon, nickel and tellurium, to which attention will be drawn. It must be confessed, however, that there are other anomalies which are not pronounced enough to affect the order of the atomic weights of the elements. The elements of the short, or so-called typical periods may be allied to those either of the A or the B Sub-groups. In the case of Group I, Li and Na are plainly related to the other alkali metals K, Rb, Cs, in the A Sub-group, rather than to Cu, Ag, and Au 74 CHEMICAL THEORY in the B Sub-group, but in Group VII, F and 01 are related to Br and I in the B Sub-group, rather than to Mn in the A Sub- group. This latter relationship obtains in all groups from II to VII. The periodic law states that the physical and chemical pro- perties of the elements and their compounds are periodic functions of their atomic weights. This statement must now be illustrated. Periodicity of Physical Properties. Perhaps the most obvious property of a solid element is its density. It was shown by Lothar Meyer, in 1870, that the densities of the elements vary periodically. Instead, however, of using the densities of the elements directly, L. Meyer calculated from them the atomic volumes, and plotted these values on a curve as ordinates, together with the atomic weights as abscissae. The atomic volume of an element is related to its density in the following manner. The reciprocal of the density is the specific volume, Specific volume = , 1 . = volume of unit mass ; density the atomic volume is this value multiplied by the atomic weight, thus: Atomic volume = atomic weight density For example, the atomic weight of copper is 63-6, and its density 89; consequently O Atomic volume Cu = ^ = 7*15. 8*9 This figure stands for the relative volume of a mass of copper proportional to the atomic weight of the element; it does not express the relative size of the copper atoms themselves; it could only do this if the atoms were packed without interspaces, or if the interspaces were constantly related in volume to the atomic material of the elements. What it does express is the relative volume of the atom plus its share of atomic interspace. The atomic volume curve shows a remarkable periodicity; for it is like a series of waves consisting of crests and hollows; moreover, the crests of successive waves increase in height with increasing atomic weight. The most important fact connected with the curve, however, is that related elements occupy analogous positions upon CLASSIFICATION OF THE ELEMENTS 75 76 CHEMICAL THEORY it. For example, the alkali metals, potassium, rubidium, and caesium, are at the apices of successive curves, the halogens, chlorine, bromine, and iodine, are on ascending, and the alkaline earth metals, calcium, strontium and barium, on descending parts of the curves. The following other physical properties of the elements and their compounds are periodic. Melting-point, malleability, coefficient of expansion, atomic refraction, conductivity for heat and electricity, colours of salts in solution. Consequently there are certain regions, which are similar on successive curves, where these properties are manifested, or reach their maxima. The student may test this statement with reference to the colours of salts in solution; remembering, however, whether the colour is due to the basic or the acidic radicle. Periodicity of Chemical Properties. The fundamental chemical division of the elements is into metals and non-metals; and, according to the classification of Berzelius, metals are electro - positive, and non-metals electro- negative. As the elements are traversed in the order of ascending atomic weights the variation of metallic and electro -chemical properties is periodic. Thus in the two short periods from lithium to fluorine, and from sodium to chlorine, there is continuous and regular transition from great metallic and electro - positive to extreme non-metallic and electro - negative characters. In the long periods which follow, for example the period from potassium to bromine, there are two phases; the first phase is from potassium through manganese to the eighth-group metals iron, cobalt, and nickel; the second phase is from copper to bromine. The transition from potassium to bromine is similar in degree to that from sodium to chlorine, but the period contains more than twice as many elements; and the stages of this transition present an interesting phenomenon. The elements of the first phase (K to Fe, Co, Ni) are all metals, but there is a continuous diminution of electro-positiveness throughout them; the elements of the second phase begin with the comparatively inert and electro-negative metal, copper, and there is actually a rise in metallic strength to zinc, followed by a regular fall to the non-metallic and electro-negative bromine. CLASSIFICATION OF THE ELEMENTS 77 Similar relations exist in the subsequent long periods; but the inertness of the central elements, i.e. those of the eighth group and of Group IB, becomes more pronounced with elements of higher atomic weight. When the transition of properties within the separate groups, i.e. the elements in vertical columns, is considered, an increase of metallic nature or decrease of non-metallic nature is found to be the rule. Thus, for example, the alkali metals increase in electro- positiveness with rise of atomic weight; and the halogens similarly show a diminution of electro-negativeness with rise of atomic weight. Within the region of chemical inertness and metallic electro-negativeness, i.e. the eighth group, Group IB, and to a less extent Group II B, an opposite state of things, however, exists; there 'is a diminution of electro-positiveness and chemical reactivity with rise of atomic weight. Thus the inert metals, platinum, gold, and mercury, occur consecutively as the last members of Groups VIII, IB, and II B. From all this it follows that the most powerful metals^ are to be found at the extreme left of the periodic diagram; caesium, the most electro-positive metal, being in the lower left-hand corner; whilst the non-metals occupy the upper right-hand portion of the diagram; fluorine, the most powerful non-metal, being in the upper right-hand corner. The dotted line in the diagram on p. 72 delimits the region of non-metals. Periodicity of Valency. The following statement is generally true. The maximum valency of an element corresponds with the number of the periodic group to which it belongs. The statement is illustrated by the formulae of the typical oxides appended to the table on p. 71. In Chapter III valency was illustrated by lists of hydrides, halides, and oxides, and in most of the formulae for the halides and oxides, but not the hydrides, the numerical value of the valency indicates the group to which the element belongs. A valency of seven is not always realized in the seventh group; less often is a valency 'of eight seen in the eighth group. On the other hand, copper and gold in Group IB show bi- and ter-valency respectively in CuCl 2 and AuCl 8 , but the elements of this group are in any case somewhat anomalous in their relationships. A more striking exception is shown in the case of 78 CHEMICAL THEORY boron, which forms the hydride B 2 H 6 , in which the element can hardly be less than quadrivalent. It was seen in the chapter on valency that the sum of the oxygen valencies and hydrogen valencies in volatile hydrides of an element is equal to eight; and this is true irrespective of the periodic group to which the element belongs. The elements of Groups I, II, and III, however, excepting boron, form no volatile hydrides, and exhibit only the lower valencies in the oxides. Thus in the first and second short periods the oxides and hydrides show valencies as follows: I II III IV V VI VII Li 2 O G1O B 2 O 3 CO 2 N 2 O 6 Oxides. BH 3 , &c. CH 4 NH 3 OH 2 FH Hydrides. Na 2 O MgO A1 2 O 3 SiO 2 P 2 O 5 SO 3 C1 2 O 7 Oxides. SiH 4 PH 3 SH 2 C1H Hydrides. Metals which form non- volatile hydrides exhibit the same valencies in these compounds as in the oxides; for example, K 2 0, KH; CaO, CaH 2 . The periodic law constitutes a valuable criterion of valency, because the periodic group to which an element belongs indicates almost invariably the valency of the element in the highest oxide which it can form. The existence of super-oxides, such as Na 2 O 2 and Ba0 2 , constitutes no real exception to this rule, because these compounds contain an oxygen chain, thus: Na-O O-Na, Ba/ \ , X) and so the valencies of the metals are the same as in the corre- sponding basic oxides. The elements of Groups VI, VII, and VIII often fail to realize their maximum valency; and indeed some, e.g. iron, never exhibit the group valency. Since the chemical character of a compound depends largely upon the active valency of its nuclear -element, the elements of these higher groups show great variety in the properties of their compounds, because they exhibit highly variable valency. The highest oxide is to be regarded as the typical oxide, provided it exhibits the valency of the group to which the element belongs; it is then found that lower oxides and their derivatives show relationships to oxides and their derivatives of similar type, but belonging to elements in other groups. CLASSIFICATION OF THE ELEMENTS 79 For example: Derivatives of Mn 2 O 7 in Group VII are isomorphous with those of Cl 2 O r in Group VII. Mn0 3 VII S0 3 VI. Mn 2 3 VII Fe 2 3 VIII, andA! 2 O 3 III. MnO VII FeO VIII, and ZnO II. Other examples might be given, all of which show that poly- valent elements, forming several classes of compounds, exhibit several relationships corresponding to these classes, and therefore that the type is the determining factor in chemical relationship. Consequently manganese, which can be septavalent, is not dis- qualified from appearing in Group VII by reason of relationships to metals in Groups VIII, VI, III, and II. Uses of the Periodic Law. Prediction of Unknown Elements. In the periodic scheme, as first formulated by Mendeleeff, there were some significant omissions. The positions now occupied by scandium, gallium, and germanium were left blank, since no elements were known qualified to fill them. If every available space had been filled with the known elements, placed in the order of their atomic weights, there would have been no periodic system, or but a distorted one, because every element which now follows a space that should have been left unoccupied would thereby have been moved one space forward, and the arrangement of analogous elements in groups would have been interfered with. On the other hand, deliberately to leave certain spaces blank, so as to preserve the desired periodicity, was to suggest that elements remained to be discovered to fill these spaces, and so to provoke a severe test of the truth of the periodic law. The latter alternative was chosen by Mendeleeff, and in par- ticular the existence of three elements was foretold, which were named provisionally eka-boron, 1 eka-aluminium, and eka-silicon. The first of these lay between calcium and titanium in the periodic table; the other two were placed consecutively to fill two blank spaces between zinc and arsenic. Moreover, by reference to the properties of neighbouring elements in series and in group, it 1 Eka is Sanskrit for one. 80 CHEMICAL THEORY was possible to foretell with considerable accuracy the properties of these undiscovered elements. This prophetic use of the periodic system by its discoverer has been rightly compared with the employment by Adams and Le Verrier of mathematical calculation to foretell the existence of the planet Neptune from observed irregularities in the movements of Uranus, and it has had an equally satisfactory vindication. For the elements scandium, gallium, and germanium, subsequently dis- covered, have been found to possess properties closely agreeing with those foretold by Mendeleeff. This is illustrated in the follow- ing comparison of eka-aluminium with gallium. EKA-ALUMINIUM. Atomic weight, dr. 68. Metal of density 5-9 and low melting-point; not volatile; un- affected by air; should decompose steam at a red heat and dissolve slowly in acids and alkalis. Oxide should have formula E1 2 O 3 , density 5-5, and dissolve in acids to form salts of the type E1X 3 . The hydroxide should dissolve in acids and alkalis. There should be a tendency towards the formation of basic salts. The sulphate should form alums. The sulphide should be precipitated by H 2 S or (NIL) a S. The anhydrous chloride should be more volatile than zinc chloride. The element will probably be dis- covered by spectrum analysis. GALLIUM Atomic weight, 69 9. Metal of density 5-94; melting at 30- 15; not volatile; unchanged in air; action on steam not known; dissolves slowly in acids and alkalis. Oxide, Ga 2 O 3 ; density not known; dissolves in acids, forming salts GaX 3 . The hydroxide dissolves in acids and alkalis. Salts readily hydrolyze and form basic salts. Alums are known. The sulphide can be precipitated by H 2 S or (NH 4 ) 2 S, but only under special circumstances. The anhy- drous chloride is more volatile than zinc chloride. Was discovered by spectrum analysis. There are other blank spaces in the periodic system which presumably correspond with hitherto undiscovered elements. The space under Mn and following Mo is noteworthy. Eighteen spaces are shown in the table between Ce and Ta, but it is unlikely that these represent undiscovered elements. It has been proposed to move up into these spaces the elements beneath them; but in any case there are a number of rare-earth metals whose atomic weights lie between those of Ce and Ta, but which ought not to fill these spaces. Probably an element remains to be discovered between W and Os, and possibly a halogen element to follow iodine, and an alkali metal more electro-positive than caesium. CLASSIFICATION OF THE ELEMENTS 81 Correction of Atomic Weight Values. Since the periodic law requires sequence of atomic weight values and sequence of properties to be in accord, grossly erroneous atomic weight value placed in sequence must disturb the sequence of properties; or, conversely, if sequence of properties is maintained it will necessitate a departure from atomic-weight sequence. In either case the erroneous value is revealed when the element in question is considered in the light of the periodic law. Indeed the erroneous atomic -weight value must cause a position to be claimed for the element, which, according to its properties, should be occupied by another element, and must consequently leave vacant a place suited to the element and in accord with its true atomic weight. Therefore the periodic law is of value, not only for detecting false atomic -weight values, but also for suggesting true ones. For example, the atomic weight of caesium was at first erroneously thought to be 123-4. This value would place caesium after antimony, and, of course, cause the displacement of tellurium, iodine, and other elements one place to the right. Such a condition cannot be thought of; therefore the value 123-4 is condemned. On the other hand, since caesium is an alkali metal it should follow rubidium in group, and consequently have an atomic weight of about 131-8, so that Cs - Rb = Kb - K = 45-35. The atomic weight of caesium is now known to be 132*81, and this metal occupies its proper place in the scheme. In the cases of glucinum, indium, and uranium the periodic law has furnished the means of deciding what multiple of the equivalent is the atomic weight. The equivalent weight of glucinum is 4-55, and the atomic weight of this element was at first thought to be 4-55 X 3 = 13-65. This value would place glucinum in an impossible posi- tion between carbon and nitrogen, whereas 4-55 x 2 = 9-1 would give it a place in harmony with the periodic law. Subsequent considerations have confirmed the value Gl = 9-1. Indium with the equivalent weight 38-27 was thought to have an oxide InO, and atomic weight 76-54, which would place this element between arsenic and selenium, where it cannot stand. An atomic weight of 38-27 x 3 = 114-8, with the corresponding oxide In 2 O 3 , would satisfy the periodic law; and this value has subsequently been accepted on the grounds of specific heat. The atomic weight of uranium was originally thought to be about 60, or else 120; but neither of these values enables the (D60) 7 82 CHEMICAL THEORY element to be placed suitably in the periodic scheme. The value 240 was required by Mendel^eff, so that the element might become the last member of the sixth group, following tungsten. This high value, or more accurately 238*5, has been supported by the vapour- density method applied to the halides, and by the fact that uranium is radio-active, since radio-activity is characteristic of the heaviest atoms. The criticism of the atomic weights ot the elements by means of the periodic law may be carried further. The accepted atomic weight of argon is greater than that of potassium, that of cobalt is greater than that of nickel, and that of tellurium is greater than that of iodine; yet the individual members of these three pairs of elements are placed in the reverse order of their atomic weights in the periodic scheme, because their properties do not permit of any other arrangement. Repeated attempts have been made in the case of tellurium to reduce the value of its atomic weight below that of iodine, but without avail, and it is at present recognized that the relationships of these three pairs of elements constitute exceptions to the periodic law. Such a conclusion is unsatisfactory; but recent researches in connection with radio-activity have shown that the mode of origin of an element has influence upon its atomic weight, and possibly a hint towards the solution of the difficulty may be found in this fact. The Suggestiveness of the Periodic Law, In spite of its apparent imperfections and the anomalies it contains, the periodic law is true; there is undoubtedly a connec- tion between sequence of atomic weights and sequence of proper- ties. This fact is a challenge to the scientific imagination; it must provoke numerous questionings in the mind of the trained student, questionings which may lead to research and discovery. For example, in a group of allied elements,; such as the alkali metals, Li, Na, K, Rb, Cs, there are series of compounds i- such as ; oxides, hydroxides, chlorides, sulphates, carbonates, and so forth, which may be expected to be related to one another somewhat as the metals themselves are related. The examination of the physical and chemical properties of these compounds may therefore be undertaken with a view to discovering the gradations which exist between them. Interesting relations will thus be established, and this fact will become apparent: that there is a break in the CLASSIFICATION OF THE ELEMENTS 83 gradation of properties between Na and K; in other words, that K, Rb, and Cs and their compounds are closely related, while Na and its compounds, as well as Li and its compounds, stand apart from them. The periodic classification affords an explanation of this phenomenon; it is that Na is situated in the second short period, whilst K occupies a different kind of position near the beginning of the first long period, and Kb and Cs follow K in quite analogous positions in subsequent long periods. Having observed this, the student may then remember that although caustic soda and caustic potash are thought of as very similar substances, sodium salts are after all not very similar to potassium salts, for they do not crystallize with the same amounts of water of crystallization as the latter, and frequently they are not iso- morphous with them, while their solubilities in water are so different from those of potassium compounds that solutions of sodium salts are used to precipitate potassium, and vice versa. At the other extremity of the periodic table the halogens pre- sent another interesting subject for study. The fact that the affinity for hydrogen diminishes from F to I in the hydrides HF, HC1, HBr, HI is well known, and is quite in accord with what occurs in other groups; e.g. in the hydrides OH 2 , SH 2 , SeH 2 , TeH 2 in Group VI, or NH 3 , PH 3 AsH 3 , SbH 3 in Group V; but fluorine is widely different from the other halogens. Why is this? This is a sort of question that perhaps cannot be answered; but one which on consideration elicits this remarkable fact: that all the elements of the first short period are unique, being widely separated in properties from those in the same groups which follow them. It suffices to draw attention to carbon, nitrogen, and oxygen, which cannot be properly classified with the elements succeeding them. Again, hydrogen fluoride differs remarkably in condensibility from the other halogen hydrides; is there any analogy to this phenomenon in neighbouring groups? Assuredly there is; if water were no more condensible than hydrogen sul- phide, the world would be a very different place to live in! In the region of the periodic chart where volatile hydrides occur the following compounds are found: BH 3 CH 4 NH 3 OH 2 FH SiH 4 PH 3 SH 2 C1H GeH 4 AsH 3 SeH 2 BrH SbH 3 TeH 2 IH. 84 CHEMICAL THEORY The periodic law suggests a comparison between them in series and in group; and thus the following gradations of properties are discovered. The hydrides diminish in stability with rise of atomic weight in every group. Thus, for example, in the fifth group ammonia is very stable, and is decomposed only slowly by the passage of electric sparks; phosphine, PH 3 , is less stable than ammonia, and is rapidly decomposed by the same agency; arsine, AsH 3 , is broken up into its elements when passed through a tube heated to 230, and stibine is similarly decomposed at 150. In series, i.e. in the hydrides standing in horizontal lines, there is an increase of stability with rise of atomic weight, corresponding with the increase of non-metallic characters, and also the diminu- tion of hydrogen valency, so that there is less hydrogen to be retained. Thus hydrogen fluoride is the most stable volatile hy- dride, and germanium hydride probably the least stable. It may be observed that Ge, As, and Sb are metalloids, that is, almost metals. No true metal forms a volatile hydride. The power to form alkyl compounds, i.e. compounds with radicles, such as methyl, CH 3 , and ethyl, 'C 2 H 5 , is more extensive than that to form hy- drides; so that some metals in the B sub-groups preceding in series the above non-metals form these so-called organo-metallic com- pounds. Perhaps the best known of these substances is zinc ethyl, Zn(C 2 H 5 ) 2 ; but, in addition to zinc, cadmium, mercury, tin, lead, and bismuth form them, and thus come into line with the above non-metals, all of which form alkyl compounds as well as volatile hydrides. Another interesting but rather difficult question is that of the relative acidic or base-producing power of these volatile hydrides. Consider the four hydrides: CH 4 , NH 3 , OH 2 , FH. Methane is inert; ammonia is base-producing, for its solution in water is alkaline owing to the formation of ammonium hy- droxide, NH 4 OH; water is neutral, and hydrogen fluoride is acid. Why is not methane, CH 4 , more base-producing than NH 3 ? the gradation of properties seems to require it to be. The answer is that in CH 4 carbon is already saturated with hydrogen, so that this substance cannot form an additive compound with water or CLASSIFICATION OF THE ELEMENTS 85 an acid as ammonia does; for the peculiar base-producing power of ammonia is an additive property. Consider again the hydrides: NH 3 OH 2 PH 3 SH 2 . There is a loss of base -producing power from NH 3 to PH 3 , and an apparently analogous increase, in acidity from OH 2 to SH 2 ; but it is difficult to generalize here, for ammonia is unique in base- producing power, just as nitrogen is unique as an element; and water, again, like oxygen, is unique in its properties. Moreover, it must not be concluded that increase in acidity of hydrides with rise of atomic weight in a group is general, for C1H, BrH, and IH are acids of about equal strength. The comparison of properties of the oxides of elements in the various groups of the periodic system is a simpler and more satis- factory exercise. For there is in general a loss of acidic and a corresponding gain of basic properties with rise of atomic weight in a group. This is shown, for example, in the oxides N 2 O 3 P 2 O 3 As 2 O 3 Sb 2 O 3 Bi 2 O 3 , and N 2 O 6 P 2 O 6 As 2 O 6 Sb 2 O 5 Bi 2 O 6 . In the trioxides there is a gradual transition from wholly acidic, through amphoteric 1 to purely basic properties, and in the pentoxides from powerfully to very feebly acidic properties. Again, the trioxides of Group VI A, Cr0 3 , Mo0 3 , W0 3 , U0 3 , form an interesting series; for, in accordance with the above generalization, basic properties actually appear, together with acidic properties, in the oxide UO 3 , which is basic with regard to one oxygen atom only, forming basic salts, such as U0 2 (NO 3 ) 2 , the uranyl salts. Objections to the Periodic Law. A consideration of the criticisms to which the periodic system has been submitted is valuable. If the criticisms are baseless, as some of them are, the process of their refutation will be illumi- nating; if they are valid, their consideration may exhibit the 1 Both basic arid acidic, anfarepos = both. 86 CHEMICAL THEORY relations of the elements from a new point of view, and so increase our knowledge concerning them. The most sweeping accusation which has been brought against the periodic system is that it places together dissimilar elements, whilst separating similar ones. It brings together the alkali metals and copper, silver, and gold in Group I, it is said a most unnatural alliance. This objection has already been met by a denial of the statement that these dissimilar metals are brought together. It is further objected that the periodic classification separates copper from mercury and barium from lead. But it may be maintained that such separation is proper; for the similarities between the metals in these several pairs are superficial rather than funda- mental, for copper and mercury are widely different in physical properties and in oxidizability; and, in spite of the fact that both metals form two series of salts, and that their lower chlorides are insoluble in water, there is little further resemblance between their corresponding salts. The differences between barium and lead are even more fundamental, so that to regard the elements as similar on account of the insolubilities of their sulphates, and the isomorphism of some other salts, is a grave error of judgment. The discovery of argon, and the determination of its atomic weight, furnished material for adverse criticism of the periodic law. For not only was it supposed that no room could be found in the scheme for an element with such extraordinary properties as argon possessed, but the atomic weight of argon was found to be greater than that of potassium; and it was manifestly impossible to place this element between potassium and calcium. Then other inert elements were discovered helium, neon, krypton, xenon, the companions of argon; and these have atomic weights less than those of the neighbouring alkali metals. Thus the atomic weight of argon is recognized as anomalous, like that of tellurium, and the inert gases therefore form a new group, which is like a buffer between the extremely different halogen elements and alkali metals; just as the metals of the eighth group intervene between manganese in Group VIlA and copper, silver, and gold in Group IB. So it is recognized that the elements of the argon family are properly placed as Group O, the periodic law is vindicated, and, in recog- tion of their analogy with the noble metals, the elements concerned are sometimes called the noble gases. CLASSIFICATION OF THE ELEMENTS 87 SUMMARY PERIODIC LAW. The physical and chemical properties of the elements and their compounds are periodic functions of the atomic weights; or If the elements are arranged in the order of increasing atomic weight, their properties vary definitely from member to member of the series, but return to a more or less similar value at fixed points in the series. USES OF THE PERIODIC LAW. Prediction of unknown elements. Correction of atomic weight values. Stimulation of thought and research regarding the elements. CHAPTER V THE STATES OF MATTER AND THE PROPERTIES OF GASES The states of matter which are commonly recognized are three solid, liquid, and vapour or gas. Matter in these three physical states is composed of molecules, the inter-relations of which deter- mine the state of the matter. A solid is characterized by a volume which is but slightly affected by changes in temperature and pressure, and by a shape of its own, which may be naturally assumed or artificially induced. By reason of the heat energy they possess, the molecules of a solid are in motion; but, inasmuch as the volume of the solid does not tend to change spontaneously, this motion does not affect the distance apart of the molecules; and, since the shape of the solid is permanent, the molecules do not change their positions relatively to each other. The motions of the molecules of a solid, therefore, are restricted to vibrations about mean fixed positions, and the molecules themselves are re- tained in these positions by the exercise of the force of cohesion, which in solids is at its maximum. Even in some solids, however, cohesion does not prevent vaporization or diffusion, as in the case of some solid metals. A liquid, like a solid, has a definite volume, little affected by changes of temperature and pressure; it has no definite shape of its own, however, but when placed in a containing vessel assumes for the time being the shape of the vessel. It is thus the property of a liquid to flow, to spread itself out, and adapt itself to external conditions, while maintaining its volume unchanged; consequently the molecules in their motions must keep the same mean distances apart while the liquid changes its shape. Thus the force of cohesion in a liquid, although weakened so that the molecules are not maintained in a fixed mean position as in a solid, is yet sufficient to confine these molecules within a definite volume so long as they remain part of the liquid. STATES OF MATTER AND PROPERTIES OF GASES 89 The characteristic of a gas is its power of indefinite expansion, so that it will distribute itself uniformly over whatever space is at its disposal. Thus a gas has neither definite shape nor definite volume, and it must be confined within an impervious envelope or it will be lost in space. The latter, of course, is true to a limited degree of most liquids and some solids; that is to say, they slowly disappear when left exposed in the air on account of evaporation. The reason that a gas tends thus to expand, and so exerts a pressure upon the envelope designed to keep it in bounds, is that its molecules are not held together by cohesion, but are indepen- dent of each other, and free to move in accordance with their inherent kinetic energy. Consequently they move in straight lines, according to Newton's first law of motion, until they encounter other molecules or the sides of the containing vessel, when the direction of their path is changed. Ideal gases are entirely devoid of cohesion between their mole- cules, whilst these molecules are wide enough apart to behave towards one another like points in space. In so far as gases conform to these conditions, they behave similarly, and indepen- dently of their chemical composition, under changes of temperature and pressure. In this respect they differ from liquids and solids, which show individuality of behaviour under such changes. Con- sequently the two fundamental gas laws connecting the volume of a gas with its temperature and pressure are of universal applica- tion; and, although in no case rigidly true, are only seriously departed from at high pressures and low temperatures, when the molecules of a gas are being brought into closer relations with one another, i. The Gas Laws The Law of Boyle (or Marriotte). The volume of a gas at constant temperature is inversely pro- portional to its pressure. p OQ. - or pv = constant, or, since density and volume are reciprocal, the density of a gas varies directly as its pressure. The Law of Charles (or Gay-Lussac), The volume of a gas, at constant pressure, increases by 5-*- 5 90 CHEMICAL THEORY (= 0-00367) part of its value at C. for every degree rise of temperature. Thus 273 vol. of gas at become 274 vol. at 1. 283 10. 263 ,,-10,&c, and if the law held for all ranges of temperature, the volume of a gas would vanish at 273. Although no such event takes place, because conditions are modified at extremely low tempera- tures, 273 is called the absolute zero of temperature, and (t + 273) is absolute temperature, or T. Consequently the law of Charles may be stated in another way the volume of a gas varies directly as its absolute temperature; and to correct the volume of a gas for change of temperature it is only necessary to add 273 to each temperature concerned, and multiply or divide by one or other absolute temperature, according to whether expan- sion or contraction is taking place. Thus, for example, a volume V at 10 becomes fffV at 15. It may be added that, if the volume of a gas is kept constant, while its temperature changes, the pressure of the gas varies directly as its absolute temperature. This follows from Boyle's law. The Gas Equation. Since pv is constant at constant temperature, and p or v varies directly as the absolute temperature, while the other remains con- stant, it follows that pv varies directly as the absolute tempera- ture, or pv = RT where R is a constant. Thus, if, with a given quantity of gas, p, v, and T undergo change, their values must be so related that R remains constant, or , This is a useful form of the gas equation which may be employed in correcting the volume of gas for simultaneous changes of temperature and pressure; for EXAMPLE. 100 cu. cm. of air at 15 at 750 mm. pressure are STATES OF MATTER AND PROPERTIES OF GASES 91 heated to 30, while the pressure becomes 770 mm. What is the new volume? = 100 X 750 X 303 770 X 288 The problem may, however, be solved quite easily without any formula by considering whether the changes in temperature and pressure, respectively, will cause increase or decrease of volume, and then arranging the data accordingly. 2. Diffusion of Gases Owing to their powers of indefinite expansion, gases mix when free to do so. Thus, if a jar of ammonia, hydrogen sulphide, or other strongly- smelling gas is opened in a room, the odour of the gas will soon be perceived at some distance from the jar. This process of mixing is called gaseous diffusion. It may be demon- strated by placing a jar of hydrogen closed with a glass plate above a similar jar of the brown gas nitrogen peroxide, and care- fully withdrawing the glass plates which separate the gases. The brown gas, although much heavier than hydrogen, will be seen to be rising in the upper jar, and after a time the contents of both jars will be uniformly brown. Or a jar of carbon dioxide may be opened beneath one of hydrogen, and very soon the presence of carbon dioxide in the upper jar may be proved by lime-water. These effects, taking place in spite of gravitation, which tends to keep the heavier gas in the lower cylinder, must be due to the intensely active motion of the molecules of the diffusing gases. Nevertheless the process takes a little time; it is not instantaneous, as the process of free expansion of a gas into a vacuum appears to be. A mental picture of the process would show the molecules of one gas rapidly threading their way through the obstructing molecules of the other gas, with consequent and frequent deflections and hindrances in their course, which are the cause of the time consumed before the mixing process is completed. The fact that mixed gases do not separate again on account of their different densities was known to Priestley; but it was Dal ton, in 1803, who proved that a lighter gas cannot rest per- manently upon a heavier, as oil upon water. Dobereiner, in 1823, observed that hydrogen escaped through a crack in a glass flask containing it, so that the gaseous pressure inside the flask 92 CHEMICAL THEORY diminished; and Graham, in 1832, examined systematically the phenomena of gaseous diffusion, and established the law regarding them. Graham employed a glass tube closed at one end with a plug of plaster of Paris. This material, as well as unglazed porce- lain, or a thin plate of artificial graphite, is porous, i.e. it will allow a gas to pass through it by diffusion whilst preventing its escape in bulk. Consequently an alteration of gaseous pressure may take place inside such a tube on account of diffusion; and this when measured will indicate the extent to which diffusion is taking place. Graham confined various gases in this tube over water, and observed sometimes a rise, sometimes a depression of the water, corresponding to a diffusion of the gas from the tube into the air at a faster or slower rate than that at which the air diffused into the tube. Then, by an analysis of the gas re- maining in the tube, data were obtained by means of which the following law of gaseous diffusion was established. The rates of diffusion of different gases are inversely proportional to the square roots of their densities. For example, oxygen is 16 times as heavy as hydrogen; so hydrogen diffuses ^/IG = 4 times as quickly as oxygen. The law is true whether diffusion takes place into another gas or into a vacuum. 1 The following results of Graham's experiments substantiate the law. /-^ _ Density : I Velocity of Dif- tras. Air = 1. V density fusion : Air = 1. Hydrogen 0-06949 3-7935 3-83 Methane 0-559 1-3375 1-344 Carbon monoxide 0-9678 1-0165 1-1149 Nitrogen 0-9713 1-0147 1-0143 Ethylene 0-978 1-0112 1-0191 Oxygen 1 - 1056 0-9510 0-9487 Hydrogen sulphide . 1-1912 0-9162 0-95 Nitrous oxide 1-527 0-8092 0-82 Carbon dioxide 1-52901 0-8087 0-812 Sulphur dioxide 2-247 0-6671 0-68 The statement of the law may be modified by saying that the times of molecular passage of gases through a porous septum are directly proportional to the square roots of the densities of the gases; 1 Probably the same law always applies to the rates at which different gases expand into a vacuum under comparable conditions ; but the speed is ordinarily too great for measurement. STATES OF MATTER AND PROPERTIES OF GASES 93 and this aspect of the law is illustrated by the following results of Graham. Gas. Time of Molecular Passage into Air at 100 mm. Pressure. V Density : = 1. Time of Mole- cular Passage into a Vacuum. Hydrogen Air Oxygen Carbon dioxide ... 0-2472 1-0000 1-1886- 0-2509 1-0000 1-1760 0-250,5 0-9501 1-0000 1-1860 The phenomena of diffusion may be illustrated in an interesting manner by employing a large glass U-tube partly filled with coloured water, and having a cylindrical porous pot attached to one end of it by means of a tightly-fitting rubber bung. When an inverted beaker of hydrogen is brought over the porous pot (fig. 11), the water in the adjacent limb of the U-tube is depressed, because hydrogen enters the pot quicker than air escapes from it, and conse- quently an excess of pressure is set up within the pot. When, however, the beaker is taken away, the hy- drogen now within the pot diffuses out quicker than air can enter, so that not only does the water rise again in the ad- jacent limb, but it goes beyond its original level, showing a tem- porary pressure within the pot less than that of the atmosphere. Finally, the water slowly falls till it attains the same level in both limbs. If the apparatus is now modified so that the pot is attached to the U-tube the other way up (fig. 12), a beaker of carbon dioxide may be brought round it so that the pot is immersed in this heavy gas. In this case the water will rise in the adjacent limb, because carbon dioxide diffuses into the pot more slowly than air diffuses Fig. 11 Fig. 12 94 CHEMICAL THEORY out of it; and when the beaker is removed the pressure will gradually be restored again, as the carbon dioxide escapes from the pot and air takes its place. The first part of the experiment may be modified so as to cause a fountain to play (fig. 13), or the displaced liquid may be caused to complete an electric circuit and so to ring a bell. The presence in mines of methane, which is lighter than air, may be indicated by an alarm given in this manner. When gases pass through a minute aperture (^-^ in. diameter) in a thin metallic plate, they obey the law of diffusion. This phenomenon was called by Graham the effusion of gases. Diffusion or effusion may be employed (a) to deter- mine the relative densities of gases, or (6) partially to separate the constituents of a mixture of gases. The latter process is called at- molysis. (a) The density of ozone relative to that of chlorine was determined by Soret, who showed that 227 volumes of chlorine dif- fused in the same time as 271 volumes of ozone; or that the rates of diffusion of the two gases were as 0-8376 : 1. If the density of chlorine is 35-46, then that of ozone is consequently 24-9; for 1 : 0-8376 : : V35-46 : Fig. 13 Similarly, Ladenburg found that a mixture of oxygen and ozone containing 86-16 per cent of ozone required 367-5 seconds to diffuse, while pure oxygen required 430 / seconds. Putting the density of oxygen = 1, then 367 $: 430:: VI :Vi- 3698; and if the density of this mixture of 86-16 per cent ozone with 13-84 per cent oxygen is 1-3698, that of pure ozone may be calculated to be 1-429 (O = 1). STATES OF MATTER AND PROPERTIES OF GASES 95 (6) Atmolysis takes place when electrolytic gas, i.e. a mixture of 2 volumes of hydrogen with 1 of oxygen, passes through an unglazed earthenware tube, such as the stem of a "church- warden" tobacco-pipe; the gas which is collected does not explode, but, owing to the escape of hydrogen by diffusion, contains a sufficient proportion of oxygen to ignite a glowing wood splint. Also, when ammonium chloride is vaporized in a glass tube through the centre of which an unglazed pipe-stem passes, the products of dissociation are partially separated by atmolysis; the lighter ammonia, passing through the stem, may be blown out against red litmus paper, which it turns blue, while the denser- hydrogen chloride, remaining in the glass tube, shows its presence by reddening blue litmus paper. Atmolysis has been employed partially to separate heavier argon from lighter nitrogen derived from air, but the process is imperfect and of little practical use. From the standpoint of dynamics, diffusion may be attributed to the velocity of the molecules of a gas, and the different rates of diffusion to differences in molecular velocities. Now the kinetic energy of a particle of mass m, moving with a velocity v, is J mi; 2 . Consider, therefore, two gases whose mole- cular weights are m x and m 2 ; when these two gases nare in thermal equilibrium, that is at the same temperature, the following relation will hold: m^v^ or d^ = d 2 or t>! : v 2 : : V This relation, however, expresses the law of diffusion if v 1 and V 2 stand for rates of diffusion. It may therefore be concluded from the kinetic theory of gases, to which the above argument belongs, that the rate of diffusion of a gas depends directly upon the velocity of its molecules. 3. Pressure of Gaseous Mixtures Barton's Law of Partial Pressures states that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases. Otherwise regarded, the law states that the different kinds of molecules present in a gaseous mixture do not interfere with 96 CHEMICAL THEORY each other, so that each gas exercises a pressure proportional to the concentration of its molecules, as if that gas alone occupied the whole space. A priori, there is no reason why this should not be true; and, indeed, it must be true if Boyle's law is true for all the gases in the mixture, and these gases have no chemical influence on each other. For, to use a well-worn analogy, gases occupy a space as the soldiers of an army might occupy a country, not in massed battalions but in widely separated units, so that the soldiers of an- other army might equally occupy the country if they were evenly distributed between the men of the first army. The law may be examined in this way. Suppose a rectangular box divided by removable partitions into several spaces, I, II, III, &c. ; let the volume of these spaces be respectively v v v 2) V 3 , &c., and let + &C. = V. Let different gases, A, B, C, &c., which have no chemical action on one another, be contained in these spaces, and let them all be at the same pressure, p. Now let the partitions be withdrawn, so that the gases mix by diffusion and their molecules become evenly distributed over the whole space. The partial pressure of gas A will then be, according to Boyle's law, -Sp, that of B -%>, and so on; the law then states that the sum of these partial pressures will equal the original pressure p: Again, it may be seen that the validity of this law depends on the validity of Boyle's law, for the above statement is reducible to this: v l + i> 2 4- ^3 4- &c. = V; and so it may be said that the volume occupied by a mixture of gases is equal to the sum of the volumes occupied by its constituents under the same conditions of temperature and pressure. The truth of this statement may be tested in the following way. From the densities of air, oxygen, and atmospheric nitrogen the percentage composition of air by weight may be calculated; it may also be determined from the volumetric composition, if the density of air is left out of account, by multiplying the STATES OF MATTER AND PROPERTIES OF GASES 97 relative volumes of the gases by their respective densities, and converting the resulting values into percentages. By calculation air has thus been found to contain 2321 per cent of oxygen and by estimation 2318 to 23*23 per cent. Dal ton's law underlies the correction which is made for the pressure of water vapour when a gas is measured over water, and is therefore saturated with water vapour. The pressure of water vapour at the temperature \of the gas, which is subtracted from the pressure under which the gas is measured, was estimated, originally, in vacuo; and yet the correction is applied to a gas at atmospheric pressure, under the justifiable assumption that the presence of the air makes no difference to the pressure the water vapour exerts. For example, if a quantity of gas is measured over water, and the barometric pressure is 765 mm., whilst the room temperature is 14 C., the pressure of the dry gas will be 765 12 = 753 mm., since the pressure of water vapour at 14 C. is 12 mm. 4. Deviations of Gases from Boyle's Law Boyle's law is true only for perfect gases, that is for gases whose molecules are independent units, behaving like points moving in space without attracting one another, and rebounding after impact with unabated energy. The law is most nearly true for gases at comparatively low pressures, and far from their liquefaction temperatures. It is therefore truer at ordinary temperatures for difficultly than for easily condensable gases. The following table shows the values of |^ at atmospheric tem- peratures, where p = 05 atmospheres and P = 1 atmosphere for a series of gases arranged in the order of their condensability. 1 For a perfect gas = 1. Gas. B.-P. Atmospheric Pressure. Temperature. pv PV Hydrogen (253 C.) 10-7 0-99974 Nitrogen (-196 C.) 14-9 1-00015 Air (about 190C.) 11-4 ! 00023 Carbon monoxide (-190 C.) 13-8 1-00026 Oxygen (-183 C.) 11-2 1-00038 Nitrous oxide ( 89-8 C.) 11-0 1-00327 Ammonia (-33- 5 C.) 9-7 1-00632 1 Lord Rayleigh, Proc. Roy. Soc., 1905, 74, *4 which seems to be unique, the hydrides of the elements fall into two classes: (a) non- volatile hydrides of powerful metals, (6) volatile, generally gaseous, hydrides of non-metals and metalloids. (a) Metallic Hydrides. The following metallic hydrides, in addition to (CuH) n> are known: NaH, KH, RbH, CsH, CaH 2 , SrH 2 , BaH 2 . The metals forming them are those of the alkalis and alkaline (D60) 177 13 178 CHEMICAL THEORY earths, the most intense of all the metals, which themselves are able to decompose water at atmospheric temperature. The hydrides are crystalline solids formed by the combination of the respective metals with hydrogen, and are decomposed by water with the evolution of twice the volume of hydrogen which is evolved when the metal alone reacts with water. Thus calcium hydride, known technically as " hydrolith ", reacts in the following way with water: CaH 2 + 2 H 2 O = Ca(OH) 2 + 2 H 2 . These hydrides stand in great contrast to the volatile hydrides of the non-metals next to be enumerated. (b) Non-metallic Hydrides. The following non-metallic hydrides are typical: BH 3 (?) CH 4 NH 3 OH 2 FH SiH 4 PH 3 SH 2 C1H GeH 4 AsH s SeH 2 BrH SbH s TeH 2 IH These compounds, and the relationships between them, have, how- ever, been fully studied in the chapter on the periodic system, to which the student is referred. Oxides and Hydroxides. The following classes of oxides can be distinguished: Neutral oxides, including suboxides. Basic oxides. Acidic oxides, including mixed anhydrides. Saline oxides. Peroxides, divided into poly- and super-oxides. Neutral oxides include the following: H 2 0, CO, N 2 0, NO, as well as the sub-oxides Cu 4 O, Ag 4 O, Pb 2 O. Water is a truly neutral oxide, or, more strictly, it is equally basic and acidic, since by the minute degree of ionization that takes place within it hydrogen and hydroxide ions are necessarily produced in equivalent quantities. Carbon monoxide is very slightly soluble in water, but the TYPES OF CHEMICAL COMPOUNDS 179 solution is neutral. Nevertheless this oxide is related to formic acid, H-COOH, in the following manner: CO + H 2 O +- H-C OH; that is to say, CO is produced from HCOOH by dehydration, but HCOOH is not produced from CO by hydration. The explanation appears to be that hydroxylation of CO thus: C=0 + H 2 does not take place; for even when another oxygen atom is present in CO 2 , to fortify the first oxygen atom, the acid produced, viz. carbonic acid, is unstable: /OH HO ^ O=C-OH. If CO is brought into contact with KOH at 100 C., however, the case is different; CO is slowly absorbed to produce formate, thus: C=O + H O K H C-OK. This reaction involves rearrangement of the atoms concerned, i.e. intra-molecular rearrangement, probably in the following way: C=0 + KOH -> C HC-OK; . but the reaction is possible owing to the basic power of the potash, and the stability of the salt produced. Nitrous oxide, N 2 O, whilst distinctly soluble in cold water, pro- duces no acid. It is nevertheless derived from hyponitrous acid, H 2 N 2 O 2 , by loss of water, thus: N=N N=N I I \/ + H 2 0. OH HO O Like carbon monoxide, nitrous oxide is an acidic anhydride from one side only; it is produced from an acid by loss of water, but does not combine with water to form that acid. A true acidic anhydride must produce an acid by combining with water; nitrous oxide is regarded as a neutral oxide because its aqueous solution is neutral. Nitric oxide, NO, is also classed as a neutral oxide, because so far as it dissolves in water it yields a neutral solution. There is, 180 CHEMICAL THEORY Jiowever, an uncommon acid, nitrohydroxylamic acid, from which this oxide is derived by loss of water: HO N-NO 2 H 2 O + 2NO. The suboxides Cu 4 O, Ag 4 O, Pb^, 1 contain less oxygen than the lowest basic oxides Cu 2 O, Ag 2 O, PbO, and they do not dissolve in water. Thus they form a small category by themselves, and since they are neither basic nor acidic they may be regarded as neutral. Basic oxides are numerous, and differ much in basic power. The oxides of the alkali metals slake very vigorously, producing hydroxides which are caustic alkalis. The oxides of the alkaline earth metals also slake, producing hydroxides which increase in solubility from calcium to barium. A few other metallic oxides, e.g. MgO and Ag 2 0, dissolve very slightly in water, giving faintly alkaline solutions. Metallic hydroxides may be produced either by the direct com- bination of the corresponding oxide with water, or by precipitation from a corresponding salt solution by an alkali. Occasionally both methods of preparation are possible; e.g. calcium hydroxide may be produced either by the slaking of lime thus: CaO + H 2 O = Ca(OH) 2 , or, on account of its slight solubility in water, by precipitation from a salt solution by means of alkali, thus: CaCl 2 + 2 NaOH = Ca(OH) 2 + 2 NaCl. Of these two methods of preparation the first is alone available for the alkali hydroxides on account of their solubility, the second can alone be employed for such hydroxides as those of copper, zinc, iron, and aluminium, on account of the inertness towards water of the corresponding oxides. Note on Bases, Acids, and Salts. A base is generally defined as a compound which neutralizes an acid to form a salt, water being eliminated in the process of neutralization. Originally the formation of water was overlooked, bases and acids being considered oxides which combined to form a salt, e.g.: SO 3 = CaO-S0 3 . 1 Pb 2 O appears to be basic, since halides corresponding to it are known. Denham, Trans. Ckem. Soc., 1917, 111, 29; 1918, 113, 249. TYPES OF CHEMICAL COMPOUNDS 181 Then it was discovered that hydrogen chloride is an acid which forms a salt with a base with elimination of water, thus: CaO + 2HCl = CaCl 2 + H 2 0, and so two kinds of acids were recognized: oxyacids such as H 2 S0 4 derived from acidic oxides such as SO 3 , and the so-called hydracids, which, containing no oxygen, are not derived from acidic oxides. And since acidic oxides were found* to combine with water to form acidic hydroxides or oxyacids, they were called acidic anhydrides, or simply anhydrides. Thus the nomenclature on the acidic side became definite; an acid was recognized as a compound containing hydrogen replaceable by a metal, and acidic oxides were no more called acids. The dualism which divides an oxysalt into basic and acidic oxides is sometimes retained for convenience, however, espe- cially in tabulating analytical data. A carbonate such as dolomite, for example, will be said to contain so much CaO, so much MgO, and so much CO 2 . Consider the following ways in which a salt such as calcium sulphate might be formed: CaO + H 2 SO 4 = CaSO 4 + H 2 O. Ca(OH) 2 + H 2 SO 4 = CaSO 4 + 2 H 2 O. Ca(OH) 2 + SO 3 = CaSO 4 + H 2 O. CaO + SO 3 = CaSO 4 . As a matter of fact, the last reaction does not take place; for in the entire absence of water basic and acidic oxides do not combine. It is open to question, too, whether CaO and SO 3 enter into combina- tion with acids and bases respectively without first combining with water. 1 Now, whilst it is definitely agreed that H 2 SO 4 or SO 2 (OH) 2 is an acid, while SO 3 is not, Ca(OH) 2 and CaO are indiscriminately described as bases, because apparently each will neutralize an acid with elimination of water. This practice, however, does not appear to the author to be scientifically justifiable; for CaO is a basic oxide and Ca(OH) 2 a basic hydroxide, just as SO 3 is an acidic oxide and H 2 SO 4 or SO 2 (OH) 2 an acidic hydroxide. And just as SO 2 (OH) 2 is an acid and SO 3 is not, so Ca(OH) 2 should be regarded as a base and CaO 1 It cannot be denied, however, that basic and acidic oxides appear sometimes to combine in absence of water, as, for example, in the formation of the saline constituents of igneous rocks, and of glasses and slags. 182 CHEMICAL THEORY not. Further, there is no valid reason why, etymologically speaking, the term " anhydride " should be connected with an acid only and not with a base. S0 3 is an acidic anhydride producing with water the acidic hydroxide or oxyacid S0 2 (OH) 2 ; similarly CaO might be properly regarded as a basic anhydride, yielding with water the basic hydroxide or base Ca(OH) 2 . The modern theory of neutraliz- ation supports this view. An acid yielding H* ions in solution is neutralized by a base which yields OH' ions. A base is thus a metallic hydroxide, rather than an oxide, whilst an acid need not be a hydroxide, with a corresponding anhydride, because it is required to yield only H*, not GET ions. The following table should make these ideas plain: Basic Anhydride. Base. Acidic Anhydride. Acid. Salt + Water from Base + Acid. CaO Na 2 O Ca(OH) 2 2NaOH SO 3 S0 2 H 2 S0 4 H 2 S0 3 CaSO 4 + 2H 2 O Na 2 SO 3 + 2 H 2 O MgO Fe 2 3 C0 2 H 2 C0 3 3HC1 MgC0 3 + 2H 2 FeCl 3 + 3H 2 Acidic Oxides. The oxides of non-metals are generally acidic oxides, and com- bine more or less readily with water to form oxyacids. Thus N 2 O 5 , P 2 O 6 , S0 3 , have a great affinity for water, and can with difficulty be separated from it; in this respect they resemble the oxides of the alkali metals, which combine with water with great vigour to form hydroxides. Other non-metallic oxides, such as CO 2 , SO 2 , B 2 O 3 , As 4 O 6 , do not so vigorously combine with water, and are more easily separated from it; Si0 2 , probably because it consists of polymerized molecules, does not directly combine with water. The strongest acidic oxides, like the strongest basic oxides, are those which combine most readily with water; and the feeblest are those which, like silica, have very little attraction for water. If the basic hydroxides are arranged roughly in order of diminishing basic strength, and the acidic hydroxides similarly in order of diminishing acidic strength, the interesting observation is made that the two series overlap in the centre thus: NaOH Ca(OH) 2 Mg(OH) 2 | A1(OH) 3 Diminishing basic strength. Al(OH) a PO(OH) 3 SO 2 (OH) 2 N0 2 OH Diminishing acidic strength. TYPES OF CHEMICAL COMPOUNDS 183 Thus A1(OH) 3 is both a feebly basic and a feebly acidic hy- droxide, and is said to be amphoteric. 1 This is shown by the fact that this hydroxide dissolves in acid to form an aluminium salt, and in alkali to form an aluminate. That it is both a base and an acid is shown by the following series of reactions which can easily be carried out: = Al(OH) 3 + 3NaCl (base) (salt) iv. AlCl 3 + 3NaOH (salt) (base) A1(OH) 3 + 3 NaOH = Al(ONa) 3 + 3 H 2 O (salt) (water) = Al(OH) 3 + 3NaCl (acid) (salt) = AlCl 3 -f 3H,O ii. (acid) (base) iii. Al(ONa) 3 + 3HCl (salt) A1(OH) 3 (base) (acid) 3HC1 (acid) (salt) (water). Thus A1(OH) 3 completes the cycle: base; acid; acid; base. The mineral known as spinel, MgO*Al 2 O 3 or MgAl 2 4 , contains alumina as aluminate; and analogous to this is chrome ironstone, which is ferrous chromite, FeO Cr 2 O 3 . Other amphoteric hydroxides are: Sn(OH) 2 , Sn(OH) 4 , Sb(OH) 3 , and perhaps Zn(OH) 2 ; though the existence of Zn(ONa) 2 is improbable, and of Zn(OH)(ONa) doubtful. It may here be recorded that certain of the less electro-positive metals form both basic and acidic oxides, the lower oxides being basic, the higher acidic. The best-known examples of such metals are chromium and manganese, to which may be added uranium on account of an interesting point in connection with the trioxide of this metal. The facts are set forth in the following table. Cr. Mn. U. CrO, wholly basic. CV 2 3 , basic, feebly acidic, e.g. in FeCr 2 O 4 . Cr0 3 , wholly acidic. (CrOgCLj is not a salt, but an acidic chloride). MnO, wholly basic. MnOp possibly basic, feebly acidic. MnO& wholly acidic. Mn 2 0 Pb Pb and Pb <8> Pb <0> Pb - It might be expected that tin, in the same group of the periodic system as lead, would form similar saline oxides. That it does not is probably due to the fact that SnO is not basic enough to combine with SnO 2 , for SnO and SnO 2 are undoubtedly more acidic than PbO and PbO 2 respectively. Peroxides. A peroxide is an oxide which readily yields some of its oxygen either as a gas or by behaving as an oxidizing agent. Thus PbO 2 , besides being an acidic oxide, is a peroxide which behaves as follows when heated alone (i) or with hydrochloric acid (ii): i. 2PbO 2 = PbO-f O 2 . ii. PbO 2 + 4 HC1 = PbCl 2 + 2 H 2 O + C1 2 . This description, however, is not precise enough; for BaO 2 be- haves like PbO 2 in regard to these two reactions, but is also capable of a third reaction, with dilute acid, viz.: Ba0 2 + 2 HC1 = BaCl 2 + H 2 O 2 . Now the difference between Ba0 2 and Pb0 2 thus revealed is fundamental, for lead belongs to the fourth periodic group, and is quadrivalent in Pb0 2 , whilst barium, belonging to the second 186 CHEMICAL THEORY group, is only bivalent. The constitutions of these two oxides are therefore represented thus: O=Pb=O; Ra,'/ and so BaO 2 yields H 2 O 2 with dilute acid thus: ,Cl H-0 a reaction of which PbO 2 is plainly incapable. Peroxides are divided into two classes, of which the above are examples. They are sometimes known as poly- and superoxides respectively (Mendeleeff) : POLYOXIDES PbO 2 . MnO 2 . C10 2 . N0 2 . SUPEROXIDES 2- BaO, Na 2 O 2 SA- TiO s . The term peroxide is applied to a particular oxide, irrespective of any other property it may possess. Thus, amongst the poly- oxides are C10 2 and NO 2 , which are also mixed anhydrides; and amongst the superoxides are those which are basic: BaO 2 and Na 2 O 2 ; and those which are acidic: S 2 O 7 and TiO 3 . The acidic, like the basic superoxides, contain a chain of 2 oxygen atoms in lieu of an increased valency of the nuclear atom. Thus the constitutions of the two latter oxides are: O O and 0=Ti Halides. The binary compounds of the halogens fluorine, chlorine, bromine, and iodine differ much in character, according to the diversity of the elements which form them. Attention may be confined to the chlorides, and those of silicon and sodium may be chosen first of all as typical of non-metallic and metallic chlorides respectively. Silicon tetrachloride, SiCl 4 , the chloride of a non-metal, may be TYPES OF CHEMICAL COMPOUNDS 187 prepared by the union of its elements, or more usually by passing chlorine over a heated mixture of silica and carbon, thus: SiO 2 + 2 C + 2 C1 2 = SiCl 4 + 2 CO. It is formed as a vapour, and may be condensed as a colourless liquid which is soluble in such a solvent as benzene. It cannot be prepared by the action of aqueous hydrochloric acid on silica, for it is instantly decomposed, i.e. hydrolyzed, by water, thus: SiCl 4 + 4H 2 O Si(OH) 4 + 4HCL It is not, however, decomposed by concentrated sulphuric acid, for silicon forms no sulphate. Sodium chloride, NaCl, presents an extreme contrast to silicon chloride in mode of preparation and properties. It is obtained by neutralizing base by acid in aqueous solution, thus: NaOH + HCl -* NaCl + H 2 0, and is consequently not decomposed by water. It is a crystalline solid, insoluble in a solvent like benzene, which dissolves silicon chloride; it is not volatile, except at high temperature, but it is decomposed by concentrated sulphuric acid, because sodium forms a sulphate, and HC1 is more volatile than H 2 SO 4 . Now consider aluminium chloride, A1C1 3 . Alumina, A1 2 O 3 , is amphoteric, and the chloride also is intermediate in character. Thus it is prepared anhydrous as a sublimate by passing either chlorine or hydrogen chloride over the heated metal, and it may also be obtained in aqueous solution by dissolving the hydroxide in excess of hydrochloric acid, and then crystallized as A1C1 3 '6H 2 O. If, however, the anhydrous or hydrated chloride is heated with water, hydrolysis takes place and hydrated alumina separates. So the following reaction is reversible: A1C1 3 + 3H 2 ^ A1(OH) S + 3HCI. Contrast with this: SiCl 4 + 4H 2 O Si(OH) 4 + 4HCl, and NaCl-t-H 2 O NaOH + HCl. Thus it is seen that, as with oxides, there are chlorides in which basic and acidic characters predominate respectively, and chlorides of intermediate or amphoteric character. 188 CHEMICAL THEO11Y Sulphides. In the course of qualitative analysis the student becomes acquainted with various types of metallic and metalloidal sul- phides. Non-metallic sulphides are not very important; carbon disulphide will serve as an example of these. Since oxygen and sulphur are in the same group of the periodic system, a comparatively close analogy between oxides and sulphides may be expected. Thus, for example, members of the following pairs of compounds may be expected to show somewhat close relationships to each other: C0 2 :CS 2 ; As 2 3 :As 2 S 3 ; K 2 O:K 2 S; and they do. Carbon dioxide and carbon disulphide differ in physical pro- perties no more than oxygen and sulphur differ; and they are distinctly analogous to each other in chemical properties. Thus, as carbon dioxide combines with basic oxides to form carbonates, so carbon disulphide combines with basic sulphides to form thio- or sulpho-carbonates: CO 2 + 2 NaOH = Na 2 CO 3 + H 2 O. CS 2 + 2 NaSH = Na 2 CS 3 + H 2 S. These oxy- and thio-salts are respectively decomposed by acids, thus: Na,COs + 2 HC1 = 2 NaCl + H 2 CO 3 . Na 2 CS 3 + 2HCl = 2NaCl + H 2 CS 3 . Thiocarbonic acid, H 2 CS 3 , separates in the latter case as a liquid, and is thus more stable than carbonic acid, H 2 C0 3 , which exists only in dilute solution. These acids, however, readily decompose into H 2 and CO 2 and H 2 S and CS 2 respectively. Practice in qualitative analysis teaches the student the relation between As. 2 O 3 and As 2 S 3 . Thus As 2 3 or more correctly As 4 6 is an acidic oxide, dissolving in alkali hydroxide to form arsenite, and similarly As 2 S 3 is an acidic sulphide dissolving in alkali hydro- sulphide to form thioarsenite; As 2 O 3 + 6 NaOH = 2 As(ONa) 3 -f 3 H 2 O. As 2 S 3 + 6NaSH = 2 As(SNa) 3 + 3 H 2 S. Frequently, however, As 2 S 3 is dissolved in NaOH, and then the reaction is: As 2 S 3 + 6 NaOH = As(SNa) 3 + As(ONa) 3 + 3 H 2 O. TYPES OF CHEMICAL COMPOUNDS 189 When, however, the solution formed in this way is acidified, As 2 S a is reprecipitated on account of its insolubility: As(SNa) 3 + As(ONa) 3 + 6 HC1 = As 2 S 3 + 6 NaCl + 3 H 2 O. Basic sulphides resemble basic oxides, thus: Na 2 O + H 2 O = 2NaOH Na 2 S + H 2 S = 2NaSH; hydrosulphide being analogous to hydroxide. When an alkali sulphide is dissolved in water, as might be expected, it reacts with water in this way: Na 2 S + H 2 = NaSH + NaOH, and is strongly alkaline on account of the hydroxide formed. Calcium sulphide, formed in the dry way as in the black-ash process, is scarcely soluble in water; nevertheless it is slowly hydro- lyzed by water, thus: 2 CaS + 2 H 2 O = Ca(SH) 2 + Ca(OH) 2 . Consequently it is not formed in presence of water; and calcium is not precipitated from solution as sulphide. A characteristic of many metallic sulphides is their exceeding insolubility in water; and in consequence they are unacted upon by this substance, and the corresponding metals are quantitatively pre- cipitated from aqueous solutions of their salts by hydrogen sulphide, as in the well-known analytical reactions. The sulphides of some metals, however, whose oxides are feebly basic or arnphoteric, cannot be precipitated in presence of water, so that the hydroxides appear in place of the sulphides. Thus when ammonium hydrosulphide is added to an aluminium salt solution, A1(OH) 3 is precipated, owing presumably to hydrolysis of the hydrosulphide: A1C1 3 -f 3 NH 4 SH = A1(SH) S + 3 NH 4 C1. A1(SH) S + 3 H 2 O = A1(OH) 3 + 3 H 2 S. A similar reaction occurs with chromium. Oxy-salts. The formation of oxy-salts, such as sulphates, is a criterion of a metal. A chloride may or may not be a salt, a sulphate necessarily 190 CHEMICAL THEORY is. Consider, for example, the following series of chlorides and sulphates : Chlorides: PC1 6 SiCl 4 A1C1 3 MgCl 2 Nad Sulphates: A1 2 (SO 4 ) 3 MgSO 4 Na 2 SO 4 . PC1 5 and SiCl 4 are certainly not salts, and phosphorus and silicon form no sulphates; A1C1 3 is a chloride of intermediate char- acter possessing some saline qualities, and A1 2 (SO 4 ) 3 is an imperfect salt somewhat hydrolyzed by water; MgCl 2 is also hydrolyzed when heated with water, MgSO 4 scarcely so; NaCl and Na 2 S0 4 are true salts yielding neutral solutions in which there is no hydrolysis. That the chlorides A1C1 3 and MgCl 2 appear more hydrolyzable than the corresponding sulphates is probably due to the volatility of hydrogen chloride, which escapes with the steam when solutions of these salts are evaporated. It appears that an element must possess a certain minimum metallic strength in order to form a sulphate; to form an acid sulphate a metal must be of the strongest character. Thus it is only the sulphates of the alkali metals which combine with sulphuric acid to form solid acid sulphates, e.g. NaHS0 4 and KHSO 4 ; though a few other metals appear to form such sulphates in solution. The study of carbonates is very instructive. The elements fall into five categories as regards power to form carbonates: i. No carbonates. ii. Basic carbonates only. iii. Normal as well as basic carbonates. iv. Normal carbonates only. v. Acid as well as normal carbonates. Non-metals form no carbonates, and power to form basic carbon- ates emerges amongst the metalloids. Thus, of the fifth-group elements, nitrogen, phosphorus, arsenic, antimony, bismuth, the last alone forms a carbonate, which however is basic, and so reveals itself as more metallic than any other of these five elements. The elements of the fourth group, carbon, silicon, germanium, tin, and lead, show similar relationships. Lead is the only member of these five that forms a carbonate, and this metal appears to be more metallic than bismuth, because in presence of excess of carbonic acid to prevent hydrolysis it yields the normal carbonate PbC0 3 . A number of other metals more readily form basic than normal carbonates; amongst these are copper, mercury, zinc, magnesium. TYPES OF CHEMICAL COMPOUNDS 191 Undoubtedly the failure to form a normal carbonate is due not only to the weakness of a metal but to the feebleness of carbonic acid. Thus magnesium, which is prone to form a basic instead of a normal carbonate by precipitation, shows no tendency to form a basic sul- phate, because sulphuric acid is sufficiently powerful to enable its salts to resist hydrolysis. Incidentally it may be supposed that a normal carbonate is first formed in the act of precipitation, and that the basic carbonate which actually appears is due to subsequent hydrolysis. For example, in the case of lead, the basic carbonate may be supposed formed from the normal carbonate, thus: 3PbCO 3 +2H 2 O 2PbCO 3 -Pb(OH) 2 + H 2 CO 3 . Sometimes this process of hydrolysis may actually be observed. For example, when sodium hydrogen carbonate is added to mercurous nitrate solution the precipitate first formed is almost white, and consists of the normal carbonate, the hydrolysis of which is pre- vented by the bicarbonate present. On dilution and warming, how- ever, the precipitate darkens, and basic carbonate, and finally oxide, is formed, because of the decomposition of the alkali bicarbonate in solution, and the consequent hydrolysis of the precipitated mercurous carbonate. The series of the metals must be traversed far in the upward direction before those metals are reached which form normal but not basic carbonates. Such are the metals of the alkaline earths and the alkalis. Calcium carbonate, for example, is normal when precipitated, and never becomes basic. This salt is soluble to a minute extent in pure water, to which it imparts a faintly alkaline reaction. It is true that hydrolysis of the dissolved salt takes place, but this results in the formation of hydroxide and bicarbonate, thus: 2CaCO 3 + H 2 O ^f 2Ca(HCO 3 ) 2 + Ca(OH) 2 , OH' ions showing alkalinity being due to the ionization of Ca(OH) 2 . The same phenomenon appears more markedly in the case of alkali carbonates, which show a strongly alkaline reaction. Thus, sodium carbonate reacts with water in this way: Na 2 CO 3 + H 2 O ^^ NaHCO 3 + NaOH, and so its strongly alkaline reaction is explained; for, according to the ionization theory, a salt producing basic and acidic ions only, e.g. 2Na- and C0 3 ", would be neutral in reaction. 192 CHEMICAL THEORY It has been seen above that only the most powerful metals form hydrogen or acid sulphates; the same is true in regard to hydrogen or acid carbonates. Thus it is only the alkali metals that form solid hydrogen carbonates, and these increase in stability from sodium to caesium in the series: NaHCO 3 , KHCO 3 , RbHCO 3 , CsHCO 3 . The alkaline earth metals, with ferrous iron and magnesium, form unstable hydrogen carbonates in solution however, e.g. Ca(HCO 3 ) 2 ; and there is this difference between these hydrogen carbonates in solution, and the solid hydrogen carbonates of the alkalis, viz. that the former are more soluble in water than the corresponding normal carbonates, whilst the latter are less soluble, and are precipitated, as, for example, in the case of NaHCO 3 , by passing carbon dioxide gas into cold saturated solutions of the normal carbonates. The superior solubility of calcium carbonate in water containing carbon dioxide in solution over its solubility in pure water is of profound importance in nature; for it is the cause, not only of the temporary hardness of water, but of the disintegration of calcareous rocks, as well as of their original formation through the agency of marine organisms, which form their shells from calcium carbonate held in aqueous solution by carbonic acid. These facts are repre- sented by the following reversible reaction: CaCO 3 + H 2 O + CO 2 ^= Ca(HCO 3 ) 2 . Hydrated Salts. Water of crystallization is of common occurrence in crystallized salts, and since its presence has a great influence on physical pro- perties, the student must on no account ignore it in formulat- ing a salt. The influence of temperature and atmospheric con- ditions on hydrated salts will be dealt with in another place; it may here be remarked, however, that when such salts are coloured the corresponding anhydrous compounds are invariably of a different colour. Thus, for example: CuSO 4 -5H 2 O is blue; CuSO 4 is white. FeSO 4 -7H 2 O green; FeSO 4 white. NiSO 4 -7H 2 O deep green; NiSO 4 yellow. CuCl 2 -2H 2 O bluish green; CuCl 2 brown. FeCl 3 -6H 2 O yellow; FeCl 3 iron-black. CoCl 2 '6H 2 O crimson; CoCl 2 blue. CoBr 2 -6H 2 O dark red; CoBr 2 green. CoI 2 -6H 2 O dark red; CoI 2 violet. TYPES OF CHEMICAL COMPOUNDS 193 The proportion of water varies much in different hydrated salts; ammonium oxalate, for example, has 1 molecule of water to 1 mole- cule of salt, ordinary sodium phosphate has 12, and the alums 24; whilst between these extremes there are salts containing 2, 5, 6, 7, and 10, and less frequently 3, 4, and 8 molecules of water. Occa- sionally, too, the same salt will crystallize with varying proportions of water according to the temperature of its formation. Thus, for example, manganous sulphate, MnSO 4 , forms crystallo-hydrates with 1, 4, 5, and 7 molecules of water at different temperatures. The way in which water is combined chemically in crystallo- hydrates constitutes a problem the discussion of which is beyond the scope of the present work; nevertheless it will be well to tabulate here the commonest hydrated salts according to the mole- cular proportions of water they contain: 2CaSO 4 -H 2 O. H 2 L 2 V 2H 2 3H 2 4H 2 O 5H,0 6H 2 O Na 2 C0 3 -H 2 0; (NH 4 ) 2 C 2 O 4 -H 2 O. BaCl 2 -2H 2 O; CuCl 2 -2H 2 O; CaSO 4 -2H 2 O. K 4 Fe(CN) 6 -3H 2 O. NaNH 4 HPO 4 -4H 2 0. Na 2 S 2 O 3 -5H 2 O; CuSO 4 -5H 2 O; Bi(NO 3 ) 3 .5H 2 O. CaCl 2 -6H 2 O; MgCl 2 -6H 2 O; CoCl 2 -6H 2 O; FeCl 3 6H 2 O; CrCl 3 -6H 2 O; FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 O, and similar double sulphates. 7H 2 O MgSO 4 -7H 2 O; ZnSO 4 -7H 2 O; FeSO 4 -7H 2 O; NiSO 4 -7H 2 O; CoSO 4 -7H 2 O. 8H 2 O . . Ba(OH) 2 -8H 2 O; BaO 2 -8H 2 O. 10H 2 O 12H 2 O 18H,O 24H 2 Na 2 CO 3 .10H 2 O; Na 2 SO 4 -10H 2 O; Na 2 B 4 O T -10H 2 O. Na 2 HPO 4 .12H 2 O; Na 2 HAsO 4 .12H 2 O. A1 2 (S0 4 ) 3 .18H 2 0. K 2 SO 4 -A1 2 (SO 4 ) 3 -24H 2 O, and other alums. Double and Complex Salts. DOUBLE SALTS are those which have a definite chemical indi- viduality in the solid state, but break up more or less completely in aqueous solution into their constituent single salts. Crystallized potassium alum, K 2 SO 4 A1 2 (S0 4 ) 3 24 H 2 O, for example, is un- doubtedly a chemical compound, and not a mixture of its two constituent salts; but when dissolved in water it gives the separate reactions of aluminium and potassium sulphates, so that its solution is a mixture of these two salts, the process of solution having been accompanied evidently by disintegration of the double salt. The formula for alum is sometimes halved, thus: KA1(SO 4 ) 2 -12H 2 O. (D60) 14 194 CHEMICAL THEORY Now it is always wise to accept the simplest available formula in default of evidence to the contrary; and there is no direct evidence regarding the molecular magnitude of a solid alum. It may be objected, however, that the above formula suggests a complex rather than a double salt, since it does not show complete molecules of the two constituent sulphates. This objection would perhaps have little weight were it not for a peculiar change which solid chromic alum undergoes when heated to 90 C. The violet crystals then turn green, with loss of water, changing into a salt which contains no free sulphate, since its solution gives no pre- cipitate with barium chloride. The change is thus formulated, ignoring water of crystallization: K 2 S0 4 .Cr 2 (S0 4 ) 3 -> K 2 [Cr 2 (S0 4 ) 4 ] or 2K[Cr(SO 4 ) 2 ]. So a double salt becomes a complex salt; potassium chromic sulphate becomes potassium chromisulphate, the potassium salt of chromisulphuric acid, HCr(SO 4 ) 2 , a compound which is actually formed when chromic sulphate is warmed with sulphuric acid. That K[Cr(SO 4 ) 2 ] is so different from K 2 S0 4 - Cr(S0 4 ) 3 24 H 2 O is a good reason for not writing the formula for any alum in a way to suggest relationship to the former of these compounds. Double salts are very numerous. Besides the double sulphates and isomorphous selenates there are double chlorides, bromides, and iodides, and less frequently double carbonates and nitrates. DOUBLE SULPHATES. The alums, and salts of which ferrous ammonium sulphate is a well-known example, may be mentioned. Alums are isomorphous salts of the type M 2 -SO 4 .X 2 -(SO 4 ) 3 -24 H 2 O, where M = Na, K, NH 4 , Kb, Cs, Tl, Ag, and X = Al, Fe, Cr, Mn, Ga, Ti, Rh. They are formed by mixing the constituent salts in aqueous solution, in proportions approximating to theoretical requirements, and crystallizing. The alums are less soluble than their con- stituent salts, and this is particularly the case with those of the extremely electropositive metals rubidium and caesium. Double sulphates of the ferrous ammonium sulphate type are the salt FeS0 4 .(NH 4 ) 2 SO 4 -6H 2 O and others in which Mg", Zn", Cu", Mn", Co", Ni" may take the place of Fe", and other alkali metals that of NH 4 . The relation between these double sulphates and the heptahydrated sulphates, e.g. FeS0 4 7H 2 O, is interesting. TYPES OF CHEMICAL COMPOUNDS 195 It was found by Graham that one of the seven molecules of water in this salt required a higher temperature for its expulsion than the other six. This seventh molecule Graham called con- stitutional water, because it appeared to enter into the constitution of the salt more intimately than the other six molecules. It is this molecule which seems to be displaced by ammonium or other alkali sulphate in the formation of the double salt. The relation- ship may be thus shown: FeSO 4 -H 2 O-6H 2 O : FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 O. It should be > remarked that the ammonium sulphate in the double salt exerts a protective influence over the ferrous sulphate, for ferrous ammonium sulphate is less oxidizable by the air than ferrous sulphate, and for this reason is preferred for the purpose of volumetric analysis. DOUBLE CHLORIDES. The mineral carnallite is KC1 MgCl 2 6 H 2 O, to which there corresponds the ammonium salt NH 4 Cl'MgCl 2 6H 2 O. The solubility of magnesium and manganous hydroxides in am- monium chloride solution, with the corresponding fact that the hydroxides of these metals are not precipitated by ammonia in presence of ammonium chloride, is sometimes attributed to the formation in solution of complex ions, such as (MgCl 3 )', derived from NH 4 Cl'MgCl 2 . The salts themselves, however, are usually regarded as double rather than complex salts. The double chloride NaCl-AlCl 8 is a volatile compound, the formation of which was a part of an early process for the preparation of metallic aluminium. Examples of double salts of another type are sodium potassium tartrate /Rochelle salt), NaKC 4 H 4 O 6 4H 2 O, microcosmic salt, NaNH 4 HP0 4 4H 2 O, and magnesium ammonium phosphate, MgNH 4 P0 4 -6H 2 O. These are formulated differently from the alums and other double salts, as containing two or more metallic radicles within the same molecule. Since, however, these salts show no complex ions in solution, and their molecular magnitudes are unknown, it may be that they should be put in the same category as other double salts. COMPLEX SALTS are those which, derived originally from single salts, are so stable as to maintain their individuality in solution, one of the metals appearing as a basic ion, whilst the other has become part of a complex acidic ion, so that its metallic nature is 196 CHEMICAL THEORY masked. Potassium ferrocyanide, K 4 Fe(CN) 6 , is a familiar example of a complex salt. It appears to be composed of 4KCN -f Fe(CN) 2 , and is indeed formed by adding potassium cyanide to ferrous sul- phate solution until the precipitated cyanide has been redissolved, and then boiling the solution. Thus a remarkable change takes place; the iron ceases to behave as a basic radicle and becomes part of an acidic complex, so that it gives no ferrous reactions in solution. No ferrous salt is present, only a potassium salt potas- sium ferrocyanide which ionizes in solution thus: K 4 Fe(CN) 6 ^= 4K- + [Fe(CN) 6 ]"". So profound is this change, and so stable the complex salt, that from its concentrated solution sulphuric acid separates hydroferro- cyanic acid, H 4 Fe(CN) 6 , as a white solid. Alum and potassium ferrocyanide, as representatives of double and complex salts respectively, present extremes, but there are gradations between them. The behaviour of nickel and cobalt salts with potassium cyanide furnishes a case in point. The fol- lowing reactions take place: NiS0 4 + 2KCN = Ni(CN) 2 + K 2 S0 4 ; Ni(CN) 2 + 2 KCN = K 2 Ni(CN) 4 . CoSO 4 + 2KCN = Co(CN) 2 + K 2 SO 4 ; Co(CN) 2 +4KCN = K 4 Fe(CN) 6 . Both K 2 Ni(CN) 4 and K 4 Co(CN) 6 are complex rather than double salts, for they do not contain nickelous and cobaltous ions; more- over, K 4 Co(CN) 6 is plainly analogous to K 4 Fe(CN) 6 . From each of these solutions, however, the simple cyanide Ni(CN) 2 or Co(CN) 2 is reprecipitated by dilute acid. These are examples of complex salts, therefore, which are less stable than ferrocyanide. When a solution of potassium cobaltocyanide is boiled in presence of air it undergoes oxidation to cobalticyanide thus: 2 K 4 Co(CN) 6 + H 2 O + O = 2 K 3 Co(CN) fl + 2 KOH, and this latter salt is much more stable than cobaltocyanide, in this respect resembling ferro- or ferricyanide. The fact that nickel forms no such stable complex salt, nickelic salts being unknown, underlies the well-known separation of these two metals. The student meets with other examples of complex acids and salts in the course of chemical analysis. Hydrofluosilicic acid, H 2 SiF 6 , is evidently composed of 2HF-fSiF 4 , but it contains the complex ion [SiF 6 ]". Potassium platinichloride, or chloroplatinate, TYPES OF CHEMICAL COMPOUNDS 197 K 2 PtCl 6) and the corresponding acid H 2 PtCl 6 , formed when platinum is dissolved in aqua regia, are of the same type, and so is the corresponding stannichloride, K 2 SnCl 6 . Potassium cobaltinitrite, K 3 Co(N0. 2 ) 6 , formed as a yellow crystalline precipitate when potas- sium nitrite is added to a cobaltous solution acidified with acetic acid, is of the same type as K 3 Co(CN) 6 and K 3 Fe(CN) 6 . Ammonium phospho-molybdate is a complex salt of a different kind, in which from 10 to 14 molecules of MoO 3 are combined with (NH 4 ) 3 PO 4 . It is formed in presence of nitric acid, and when dissolved by ammonia suffers hydrolysis into simple phosphate and molybdate. Potassium antimonyl tartrate, or tartar emetic, [KSbOC 4 H 4 O 6 ] 2 H 2 O, is a complex rather than a double salt, for it dissolves in water without hydrolysis, which antimonious salts will not do. It is therefore best regarded as the potassium salt of antimonyl-tartaric acid, [K(SbOC 4 H 4 O 6 )] 2 H 2 O. SUMMARY Types of Chemical Compounds HYDRIDES. Metallic and non-metallic. OXIDES AND HYDROXIDES. Neutral oxides, including suboxides. Basic oxides. Acidic oxides, including mixed anhydrides. Saline oxides. Peroxides divided into poly- and superoxides. HALIDES. Metallic and non-metallic. SULPHIDES. Metallic, metalloidal, and non-metallic. OXYSALTS. Sulphates, carbonates, &c. HYDRATED SALTS. DOUBLE AND COMPLEX SALTS. CHAPTER XI CHEMICAL CHANGE IN GENERAL A classical illustration of chemical change, at once simple and valuable, is furnished by the work of Priestley and Lavoisier on mercuric oxide. Priestley heated mercuric oxide by concentrating the sun's rays upon it with a lens, so as to decompose it into mercury and oxygen. The reaction is represented thus: 2HgO 2Hg + O 2 . This mercuric oxide could previously be obtained, as was shown by Geber, by gently heating mercury for a long time in the air, when atmospheric oxygen united with the metal, thus: 2 _* 2HgO. Lavoisier combined these two operations by first heating mercury at a moderate temperature in a confined space, and noting the volume of air absorbed, and then collecting the mercuric oxide formed and heating it more strongly; this resulted in the evolution of a volume of oxygen equal to that of the air which was previously absorbed. So the possibility of reversing a chemical reaction was established, a fact now represented thus: 2 :^ 2HgO. This simple illustration has been chosen because it gives rise to various questionings, the consideration of which leads far into the subject of chemical change in general. Thus it is a surprising thing that the compound of mercury and oxygen should be a red, crystalline powder, so different from its constituent elements, and the question at once occurs whether the whole of chemistry is full of surprises like that. The elementary student is rather led to suppose that it is. At least there are many such surprises which lend to chemical science a fantastic charm for the youthful mind. For example, the vapour of sulphur is led over 198 CHEMICAL CHANGE IN GENERAL 199 red-hot charcoal, and, instead of yellow crystals and black lumps, there appears a colourless liquid with a quite extraordinary smell; or ammonia and hydrogen chloride gases are brought together, and instead of a neutral gas resulting from the combination of an acid and an alkaline gas, there is -dense white smoke which settles down as solid sal ammoniac. " It is the unexpected that happens" might apparently be said of chemical change. Indeed, the difference between physical mixture and chemical combination often appears to be this the properties of a mixture are what might be expected, they are the mean of those of the constituent parts of the mixture, whilst the properties of a chemical compound often could not be expected, for they are unrelated to those of the constituent elements. Yet, if the impression gained from the facts above considered were true, and if the above epigram were a generalization of chemistry then there would be no chemical science; chemistry would be but a catalogue of curious material phenomena. The scientific thinker is thus met with the fundamental question of the relation between the properties of compounds and those of their constituent elements; he is led, indeed, to the threshold of a field of inquiry as broad as chemistry itself. This particular inquiry may, however, be carried a little further. The greatest differences in properties are seen between elements and their simplest compounds. When a compound is converted into another compound by the addition or substitution of other elements, the physical differences brought about are not so great. Consider, for example, the series of paraffin hydrocarbons C n H 2n+2 (see p. 118). An increment of CH 2 causes no surprising change in the properties of a hydrocarbon; on the contrary it causes an almost constant alteration of boiling-point and other physical properties. Or, consider the influence of the substi- tution of chlorine for hydrogen in the CH 3 group of acetic acid, CH 3 .COOH. The chloracetic acids CH 2 C1-COOH, CHC1 2 -COOH, CC1 3 'COOH, stand in order of increasing strength; thus the electro- negative element chlorine has had a specific influence in increasing the strength of the acid. The colour, also, of a complex chemical compound is definitely related to its constitution, and is modified by the substitution of one element or group for another within the molecule. The art of producing synthetic dyes depends among other things on a 200 CHEMICAL THEORY knowledge of the influence of certain substituents on the colour of the compound formed. The same is true regarding the thera- peutic properties of synthetic drugs. But, more generally, if, according to the periodic law, the properties of the elements and their compounds are periodic functions of the atomic weights, then this law should at least relate the properties of a particular compound to those of a similar compound of an analogous element. That it does this is shown by a systematic study of oxides, chlorides, and other simple com- pounds. Our present knowledge of the law, however, fails to account completely for the properties of a particular compound. It is known, for example, that the solid iodides of imperfect metals are brightly coloured, although the constituent ions are colourless; such iodides are: PbI 2 , Hg 2 I 2 , HgI 2 , SnI 4 , SbI 3 ; but why PbI 2 is yellow, for instance, and SnI 4 scarlet, is quite unknown. The same is true of the colours of some ions; thus there is no theory of the inherent property of manganese which causes the colour of the green manganate or crimson permanganate ion. These are examples of the questionings to which a consideration of the superficial properties of a simple chemical compound gives rise. The human mind desires an explanation of the unexpected. Why is mercuric oxide red; why is permanganate solution crimson; what hidden properties do the constituent elements of these com- pounds possess which are revealed in so striking a manner on combination? And these questions cannot at present be answered. Reversible Reactions. A second question for consideration suggested by the reaction between mercury and oxygen is that of the reversibility of a chemical change. Can all chemical reactions be reversed; if not, why not? Speaking generally, the possibility of reversing a chemical reaction depends on the realization of suitable conditions. Some chemical changes brought about by heat are so profound that their reversal in the narrower sense is not possible. Sugar, for example, is destroyed if heated strongly; it is commonly said to be burnt, and the final result of the burning of sugar in air is the conversion of its elements into carbon dioxide and water. Can such a change be reversed? Growing plants can reconvert the carbon of carbon dioxide CHEMICAL CHANGE IN GENERAL 201 and water into sugar, and the chemist can laboriously synthesize a kind of sugar; but that is not a true reversal of the chemical reaction of decomposition, because the synthetic changes do not follow the same route as the changes of decomposition. Now, reverting to the reaction between mercury and oxygen, it would appear from Lavoisier's experiment that there is a certain minimum temperature at which visible combination between the elements takes place, and a somewhat higher temperature at which visible decomposition of the compound formed sets in. These temperatures cannot be stated because they are conditioned, but it may be judged that below or above a limited range of temperature oxygen and mercury do not combine or remain in combination respectively. Such a statement, however, is not very satisfactory, because, whilst it recognizes the increased activity of mercury and oxygen molecules, due to rise of temperature, which first promotes com- bination between the two elements, and subsequently causes disruption of the compound formed, it takes no account of the physical state or concentration of the combining elements; or otherwise, since oxygen is a gas, that it may escape from the mercury altogether when evolved and so render a reversal of the reaction impossible. It is worth while to attempt to gain a clear mental picture of this reversible reaction, since it illustrates fundamental principles which underlie chemical reactions in general. Chemical .Equilibrium. Suppose the flask A in the figure con- tains mercury and mercuric oxide in con- tact with oxygen gas, the pressure of the latter being indicated by the manometer B, consisting of a U-tube containing mercury; and suppose that the flask is heated in a chamber, the outline of which is shown by the dotted line, to a temperature t C., within the limits between which a reaction be- tween mercury and oxygen is known to take place. Then the manometer will show in which direction the reaction is proceeding. If combination is taking place, the mercury will rise in the limb nearer the flask, owing to diminution of gas L Fig. 45 202 CHEMICAL THEORY pressure; if decomposition, the mercury in the nearer limb will be depressed because of increase of pressure. But in either case equilibrium will eventually result, and the mercury in the mano- meter will become and remain stationary, registering a certain gaUU 1500 1400 1300 1200 . 1100 1 100 ^ 900 1 800 2 700 1 60 & 500 400 300 200 ioo ( i i i / ( i / / ( /\ / / / A s / ^ ^ 360 380 400 420 440 460 480 Temperature. Fig. 46. Dissociation Pressure Curve of Mercuric Oxide If the student has followed the exposition of this simple chemical reaction thus far he will be prepared for the conclusion that the latter depends on three determining factors. First, there is the specific power of combination between mercury and oxygen. Com- pared with other elements, mercury has little intrinsic attraction for oxygen; if, for example, copper were substituted for mercury, the result would be very different. The quality of mercury thus mani- 204 CHEMICAL THEORY f ested may be called its chemical affinity, a convenient but originally rather vague term. Second, there is the effect of temperature. The power of mercury to combine with oxygen remains latent until by rise of temperature the molecules of metal and gas are made sufficiently active. Chemi- cal change is always confined within certain temperature limits. At very low temperatures matter is inert and incapable of chemical change ; at very high temperatures matter is too active, its atoms are too restless to enter into chemical union. Third, there is the effect of pressure, that is of the concentration of oxygen molecules around the mercury. If the surface of the mercury in the flask is thought of as in a state of flux, some mole- cules of the oxide being continuously produced whilst others are simultaneously undergoing decomposition, it will be understood that the chances of combination are increased and of decomposition simultaneously diminished by the crowding of oxygen molecules just above the surface. Thus increase of pressure promotes com- bination, and diminution of pressure decomposition, altogether apart from the question of temperature. In other words, the extent and apparent direction of a chemical reaction is dependent on the active masses of the reacting substances. Thus, apart from temperature, chemical action is determined by chemical affinity and active mass. By active mass is meant suf- ficient mass for the purpose. The mercury beneath the surface in the illustration has no action it need not be there ; neither need the oxygen far above the surface, except to transmit the pressure to the region of activity. What is of importance, therefore, is the degree of concentration of the reacting substances in the zone of reaction. It must of course be understood that the proportions of mercury and oxygen which are in contact do not in any way influence the proportions in which these elements combine. They influence the extent to which a compound is formed, not its quantitative com- position. That depends upon the relative masses '.of the combining atoms, a fact expressed by the law of definite proportions. Further examples of the influence of mass in chemical change will be given in the sequel. Thermochemistry. There is one other aspect of this simple chemical reaction between mercury and oxygen which must be considered before the study of CHEMICAL CHANGE IN GENERAL 205 it is completed. The effect of temperature on the chemical reaction has been noticed; but heat is applied in quantity, and some of this heat is transformed into other kinds of energy during the pro- gress of the reaction. In the decomposition of the oxide, for example, oxygen leaves the solid and becomes a gas; not only is heat used up in the actual dis- ruptive process, but also doubtless in the gasification of the oxygen, as well as in the vaporization of some of the mercury, though the latter heat is given up again if the mercury is condensed. Thus heat is absorbed in the change 2HgO and it has been estimated that when 2 gram-molecules (i.e. 433 2 grm.) of mercuric oxide are decomposed, producing 2 gram-atoms of liquid mercury (401 2 grm.) and 1 gram-molecule of oxygen gas (32 grm.) 44,000 calories are absorbed; and similarly, when these quantities of the separate elements unite to form mercuric oxide, the same quantity of heat is liberated. These facts may be expressed thus: 2 HgO = 2 Hg + O 2 44,000 cals. and 2 Hg + O 2 = 2 HgO + 44,000 cals. So mercury and oxygen in the system [2 Hg + O 2 ] contain between them an excess of energy over that in the system [2 HgO] amounting in terms of heat energy to 44,000 calories. It must, however, be clearly understood that nothing is here said or known concerning the total energy in either of these systems. The rise or fall of the tide may be measured against a rock in a sea the depth of which is unplumbed. Similarly, the rise or fall of energy content in a material system may be measured although the total energy in the system cannot be estimated. The energy of a system which is transferable by chemical change is called the free energy, the energy which is not transferable is the bound or latent energy] these two make together the total energy in the system. Now, every distribution of matter which we call a chemical change is accompanied by a corresponding distribution of energy. The branch of chemistry concerned with energy changes manifested by heat phenomena is called thermo-chemistry, the kind of equation which sets forth heat change accompanying changes in matter is called a thermo-chemical equation, and such an equation is based upon the 206 CHEMICAL THEORY principle of the conservation of energy, just as an ordinary chemical equation, which is a mass equation, is based upon the principle of the conservation of mass. It will be well to consider now some representative thermal equations chosen to illustrate the range of the subject. And first it may be remarked that thermal equations need not necessarily be thermo-chemical equations. They may express thermal phenomena accompanying physical rather than chemical change. For example, the transformations between ice, water, and steam, may be represented thus: H 2 O (water) = H 2 O (Ice) + 1440 cal. H 2 O (steam) = H 2 O (water) + 9670 cal. These equations express, in perhaps unusual guise, the latent heats of water and steam. Or the heat of transformation of one allotropic form of an ele- ment into another may be represented thermo-chemically thus: S (raonoclinic) = S (rhombic) + 64 cal. C (graphite) = C (diamond) + 500 cal. Thus when 32 grm. of monoclinic sulphur are converted into the same weight of rhombic sulphur, 64 calories are evolved; and simi- larly in the conversion of 12 grm. of graphite into diamond 500 calories would be evolved. These figures are not, however, deter- mined directly, but are derived from the differences between the heats of combustion of different forms of the element. For example, the heats of combustion of graphite and diamond are: Graphite, 94,810 cal.; diamond, 94,310 cal.; so the difference, 500 calories, must represent the heat of transfor- mation of graphite into diamond. It may next be observed that in the case of a reversible change, whether physical or chemical, if the evolution of a certain quantity of heat accompanies the change in one direction, the absorption of the same quantity of heat takes place when the change is reversed. This may be represented simply by transposing the terms of the thermal equation; thus in the case of ice and water H 2 O (water) = H 2 O (ice) + 1440 Cal.J but H 2 O (ice) = H 2 O (water) 1440 cal. Thus when 18 grm. of water are frozen, 1440 calories are removed CHEMICAL CHANGE IN GENERAL 207 from the water, and to melt the ice a similar quantity of heat must again be supplied. Changes in which heat is evolved are exothermic changes; those in which heat is absorbed are endothermic changes. These terms are of value in considering chemical changes in- volving much heat energy. Thus, reverting to the original example, the formation of mercuric oxide from its elements is an exothermic change; the decomposition of the* compound into its elements an endothermic change; the two changes, as regards energy, as well as matter, are equal and opposite. One, the heat of formation, is + 44,000 cal.; the other, the heat of dissociation, is 44,000 cal. Other kinds of thermo-chemical data of value are heats of solution of substances in water and other solvents; heats of evapo- ration of solvents containing specific dissolved substances; heats of neutralization of acids and alkalis. It may be pointed out that when chemical operations are being conducted on a large scale, thermal effects, often negligible on the small scale, assume much importance. For example, how to divert or usefully to dispose of the heat evolved in a chemical action, or how most economically to conduct the evaporation of a given solution, having regard to its latent heat of evaporation, often become problems of pressing importance to the chemical engineer. Here may be mentioned the law of Hess, which, briefly, states that the heat of formation of a compound is independent of its mode of formation. This law follows naturally from the principle of the conservation of energy, but it is capable of experimental proof. Lastly, the connection between heat evolved and chemical affinity may be briefly considered. Exothermic compounds are stable; endothermic compounds often unstable. Exothermic compounds may be compared with a boulder at the base of a mountain which has reached its position by rolling down the mountain's side, expending its potential energy the while, but having reached the lowest ground is likely to remain there. Endothermic compounds may be compared with a boulder raised and poised above the valley, and liable, after a slight disturbance, to roll down and expend energy in its fall. High-explosives, and detonators such as fulminate of mercury, are good examples of endothermic substances, as well as ozone, acetylene, and carbon disulphide. 208 CHEMICAL THEORY Ozone is formed by the absorption of energy from the electric sparks which generate it from oxygen; acetylene derives its in- herent energy from the electric furnace in which calcium carbide is produced; carbon disulphide is similarly formed from its elements at high temperature. These compounds are evidently highly artificial, and unlikely to be produced by nature in her quieter moods, though electric or volcanic action might produce them. Solar energy is, however, continuously being absorbed by growing plants, and the compounds produced by them, such as starch and sugar, may, in a wider sense, be regarded as endothermic. Similarly, wood, coal, and petroleum are in this sense endothermic, though it is unusual to extend the term to include these things. O In general, endothermic compounds are likely to be produced at high temperature, because the heat to be absorbed in their forma- tion is thus available; whilst exothermic compounds are likely to be formed at lower temperature, because the heat evolved in their formation can more easily be dispersed. The connection between the heat effect of a chemical change and the chemical affinities of the reacting substances may now be considered briefly. It is generally supposed that if the reaction between two chemical elements is vigorous and accompanied by the evolution of much heat, the chemical affinity between these elements is great; and that if the reaction is sluggish and accompanied by small heat evolution, the affinity between the elements concerned is small. This is illustrated by the heats of formation of the halogen hydracids, which are: HF 38,600 calories. HC1 22,000 HBr 8,400 HI -6,000 Thus the heats of formation fall from HF to HI, and that of the latter is actually negative; HI is an endothermic compound. The affinities of the halogens for hydrogen undoubtedly fall from fluorine to iodine, but that of iodine cannot be regarded as actually negative, as the thermal value might suggest. So that whilst the heat effect of chemical union is some indication of the strength of attraction between the elements, it cannot be regarded as a quan- titative measure of chemical affinity. CHEMICAL CHANGE IN GENERAL 209 There are other considerations bearing on thermo- chemistry, and the connection between heat effect and chemical affinity, which the student may postpone to a more advanced course. Rate and Limits of Chemical Change. The rate of a chemical change is much influenced by the condi- tions under which it takes place. Thus the physical state of a reacting substance or substances determines whether or how a chemical change occurs. So it is often the case that substances which interact in solution will remain in contact with one another in the solid state without chemical change. The familiar example of a mixture of sodium bicarbonate and tartaric acid, which effer- vesces when water is poured upon it, may be quoted. The modifying influence of temperature upon chemical change is noteworthy. Consider, for example, the formation of water from its elements or its dissociation into them. If hydrogen and oxygen gases are mixed together in the propor- tions necessary to form water, no change is perceived at atmospheric temperature, but if the mixed gases are gradually heated, combina- tion becomes perceptible at about 450 C., whilst .at 700 C. there is a rapid and complete combination, accompanied by explosion. It is possible, however, to raise the mixed gases to such a temperature as not to promote but actually to prevent their combination. Thus above 2000 C. steam is decomposed into its elements, and hydrogen and oxygen will not combine. So there are temperature limits within which the combination of hydrogen and oxygen can take place, but outside which no combination between these gases occurs. And what is true of this particular change is true of all chemical changes. Chemical reactions, vigorous under ordinary conditions, become sluggish with reduction of temperature, and cease altogether at very low temperatures. The velocity of a reaction is frequently halved by a fall of 5 C., and if this rate of diminution proceeds regularly, the velocity will become infini- tesimal, the reaction coming practically to a standstill before the temperature of liquid air is reached. On the other hand, chemical unions between the elements is impossible in the hotter stars; only in the coldest stars, for example, do hydrocarbons exist. The effect of heat in promoting change is too common to need specific illustration. The use of a Bunsen gas flame to further chemical -action soon becomes instinctive to the student. (D60) 15 210 CHEMICAL THEORY Catalysis. Further study of the interaction of hydrogen and oxygen reveals the fact that the degree of combination between these gases is much influenced by the presence or absence of water vapour, Thus, when the pure gases have been dried perfectly by long exposure to phosphoric oxide, they do not combine at a red heat, and silver can be melted in a mixture* of them at a temperature of 960 C. Minute quantities of other foreign substances also affect the rate of combination, and this is notably the case with certain metals, especially if they are finely divided. Thus, when the mixed gases are passed through a heated capillary tube containing a thread of platinized asbestos, gradual combination takes place, whilst the thread becomes red-hot, and would eventually cause the gases kept in contact with it to explode. The platinum, however, remains quite unaltered at the end of the process. These are examples of catalytic action, or catalysis, and the promoter of the chemical change is called a catalyst. Catalysis plays an important part in modern chemistry, and especially in chemical industry. A brief study of the subject may therefore be undertaken here. The student will already be familiar with several cases of catalysis. He knows that manganese dioxide causes potassium chlorate to yield its oxygen at a lower temperature than that at which the pure salt is decomposed; also that "nitrous fumes" promote the formation of sulphuric acid from sulphur dioxide, atmospheric oxygen, and water by a cycle of reactions which may be represented thus: 2NO + O 2 = 2NO 2 2H 2 O = 2H 2 SO 4 So nitric oxide carries oxygen from the air to sulphurous acid, being alternately oxidized and reduced. Platinized asbestos, however, causes sulphur dioxide and oxygen to combine thus: 2SO 2 -f 2 = 2S0 3 , without itself undergoing perceptible change. Cuprous chloride, employed in the Deacon process in the oxidation of hydrogen chloride to chlorine and water, behaves similarly to nitrous fumes in the chamber sulphuric acid process, as an oxygen carrier, thus: CHEMICAL CHANGE IN GENERAL 211 2CuCl 2 = Cu 2 Cl 2 Cu 2 Cl 2 + O = Cu 2 OCl 2 , Cu 2 OCl 2 + 2 HC1 = 2 CuCl 2 + H 2 O; the net result being seen by adding these reactions thus: 2HC1 + O = H 2 O + C1 2 . These well-known examples of catalysis illustrate the variety of the mechanism of the reactions .which may occur. Thus in the case of the older sulphuric acid process, as well as in the Deacon process, there is no doubt that a series of reactions takes place, involving the catalysts, whilst it is practically certain that finely divided platinum is not altered chemically when it acts as a catalyst. The case of manganese dioxide is perhaps uncertain. It is generally believed that this substance is oxidized by the chlorate to a higher oxide which is immediately decomposed again; but if this is the case it is hard to believe that a similar thing happens to ferric oxide which may take the place of manganese dioxide in promoting the decomposition of potassium chlorate. The following are well-known examples of catalytic agents and catalytic action. i. Influence of Water. It has already been seen that a trace of water-vapour greatly influences the combination of hydrogen and oxygen to form water. In the entire absence of water the following reactions will not take place: 2CO + O 2 = 2CO 2 . 2NO + O 2 = 2NO 2 . H 2 + C1 2 = 2HC1. CaO + SO a = CaSO 4 . CaO + 2NH 4 Cl = Also, mercuric chloride and ammonium chloride do not dissociate when quite dry, and sodium and potassium can be distilled in perfectly dry oxygen without combustion. There is thus a variety of reactions which depend for their progress on the influence of water as a catalyst. ii. Influence of Hydrogen Ions. An ester, such as ethyl acetate, may be formed or hydrolyzed according to the following reversible reaction: C 2 H 6 OH + H - C 2 H 3 O 2 ^= C 2 H 5 - C 2 H 3 O 2 + HOH. 212 CHEMICAL THEORY The progress of the reaction in either direction, however, is exceedingly slow in the absence of a catalyst; but hydrogen chloride, by reason of hydrogen ions, under suitable conditions accelerates both the formation and the hydrolysis of the ester. The hydrolysis of cane-sugar thus: C 12 H 22 O n + H 2 = C 6 H 12 6 + C 6 H 12 6 , (dextrose) (levulose) is also promoted by hydrogen ions, introduced as dilute acid, and the extent of hydrolysis, in this case as well as in that of an ester, is directly proportional to the concentration of the hydrogen ions; consequently, these reactions can be used to compare the strengths of acids. Another interesting example of the catalytic influence of hydrogen ions is the precipitation of a colloid. If hydrogen sulphide gas is passed into a dilute aqueous solution of tartar- emetic potassium antimonyl tartrate there is no precipitate, but a deep orange, faintly -opalescent liquid results. This liquid contains antimonious sulphide in the colloidal state. If then a few drops of dilute hydrochloric acid are added, an orange precipitate is formed, the colloidal state of quasi-solution having been dis- turbed by the hydrogen ions of the acid. iii. Influence of Finely Divided Metals. The catalytic influence of platinized asbestos has already been noticed. The contact process for the manufacture of sulphuric acid depends on the synthesis of sulphur trioxide in presence of platinized asbestos. The realization of this reaction on a commercial scale was found to depend on keeping the platinum free from dust, and especially from arsenious oxide, derived from the iron pyrites used as the source of sulphur. The arsenic has been said to "poison" the catalyst; it is indeed a negative catalyst, counter- acting the influence of the finely divided metal. Platinum black, precipitated from platinic solutions by reducing agents, is also very active in promoting gaseous synthesis. Finely divided osmium and uranium have been found valuable in promoting the synthetic formation of ammonia; platinum, as well as iron con- taining copper or bismuth, is employed in bringing about the oxidation of ammonia to nitric acid; finely divided nickel at 250 C. CHEMICAL CHANGE IN GENERAL filS causes hydrogen to be added to unsaturated hydrocarbons and their derivatives, a reaction employed for the hardening of fats. Another example of the catalytic action of a finely divided metal is that of silver on hydrogen peroxide. If aqueous hydrogen peroxide is added to silver oxide the oxide is reduced to metal with evolution of oxygen, and then if more hydrogen peroxide is poured upon the metal the evolution of PUS continues. Alternate oxidation and reduction of the silver O may here be postulated thus 1 : 2 Ag + H 2 O 2 = Ag 2 O + H 2 O. Ag 2 + H 2 2 = 2 Ag + H 2 + 2 . Colloidal metals are particularly active owing to their state of minute subdivision. Bredig has obtained platinum, silver, &c., in colloidal suspension in water by striking an electric arc between poles of the metal beneath the surface of the water. In this way the metal becomes diffused through the water in a colloidal state, and is very active catalytically. iv. Influence of Enzymes, Enzymes are unorganized ferments produced by living organisms, and they are very powerful catalysts. For example, diastase, produced in germinating seeds, converts starch into sugar, thus: 2C 6 H 10 6 + H 2 = C 12 H 22 O n , and so renders stored up food available for the growing plant; zymase, in yeast, also converts sugars of the type C 6 H 12 O 6 into alcohol and carbon dioxide, thus: C 6 H 12 6 = 2C 2 H + 2C0 2 . Similarly, there are unorganized ferments in the human alimentary canal, which aid the digestion of food. All these actions are catalytic. Theory of Catalysis. Consideration of the foregoing examples of catalysis show that 1 Hydrogen peroxide, moreover, is very sensitive to catalytic action ; precipitated man- ganese dioxide will decompose it, as well as finely divided platinum, and even powdered glass. Indeed, sharp points and edges promote its decomposition. If a concentrated solution of hydrogen peroxide is poured into a glass vessel the surface of which has been scratched, the evolution of gas bubbles along the line of the scratch may be observed. In this case catalysis is certainly not due to chemical action. 214 CHEMICAL THEORY in some cases intermediate reactions occur, while in others the action of the catalyst seems to be only physical. Thus finely divided platinum is probably physical in its action; the finely divided metal presents a very large surface on which the combining gases by a kind of condensation are brought into intimate contact. Some examples of catalysis must, however, still remain unexplained. A catalyst may be defined, in the words of Ostwald, as a " sub- stance which alters the velocity of a reaction, but does not appear in the end products ". A catalyst is generally present in relatively small amount, and its amount may be minute. Thus it serves over and over again, and may remain active for an indefinite length of time. The velocity of the reaction, however, is often proportional to the amount of catalyst present. Further, a catalyst is believed not to initiate, but only to hasten, or sometimes to retard, a chemical change. It is true that in the absence of the catalyst the change often appears not to occur, but probably, as in the synthesis of water, this is due to the reaction being exceedingly slow. Moreover, a catalyst does not effect the final state of equilibrium of the reacting substances; it only alters the time in which this equilibrium is reached. Lastly, where intermediate reactions are known to occur, they must, when taken together, necessarily pro- ceed faster than the original reaction in order to hasten it; there must be "a saving of time in the longer way round". SUMMARY THERMAL DISSOCIATION. Thermal dissociation is a reversible chemical change caused by heat. Exo- AND ENDO-THERMIC COMPOUNDS. An exothermic com- pound is a compound produced from its elements with the evolution of heat. An endothermic compound is a compound produced from its elements with absorption of heat. The heat of formation of a compound is the heat evolved in its formation from its elements. CATALYST. A catalyst is a substance which alters the velocity of a reaction, but does not appear in the end products. CHAPTER XII CHEMICAL CHANGES CLASSIFIED It is customary to classify chemical reactions in the following manner: I. Combination of elements or compounds, e.g.: 2 = 2HgO. S0 3 + H 2 = H 2 S0 4 . II. Decomposition of compounds into simpler compounds or ele- ments, e.g.: CaCO 3 = CaO + CO 2 . Ag 2 O = 2Ag + O 2 . III. Re-arrangement of the atoms within a molecule, e.g. the formation of urea from ammonium cyanate, thus: NH 4 CNO = CO(NH 2 ) 2 . IV. Condensation of two or more molecules into one molecule (polymerization), e.g.: 3 C 2 H 2 = C 6 H 6 . V. Single displacement, e.g.: CuSO 4 + Zn = ZnS0 4 + Cu. VI. Double decomposition, e.g.: BaCl 2 + H 2 SO 4 = BaSO 4 + 2 HC1. Such a classification is, however, rather formal, and not very illuminating, because it overlooks cause and effect in chemical change; and reactions which are similar in form may differ essen- tially in nature. Compare, for example, the above reaction of double decomposi- tion with the following: 2 NaCl + H 2 SO 4 = Na 2 SO 4 + 2 HC1. 215 216 CHEMICAL THEORY The two reactions are similar in form, but quite different in cause and effect. More chemistry may be learned by studying reactions according to the conditions under which they take place than according to a few set types accepted a priori. Thus, the study of changes effected by the action of heat on solid substances, either singly or mixed, or by bringing together solids and liquids such as non-metals and metals and their com- pounds on the one hand, and water, acids, and alkalis on the other, or of changes occurring in aqueous solution, will be found to cover a very wide field of chemistry, and yield much insight into the nature of chemical change. Or, again, the phenomena of oxidation and reduction well repay classification and thoughtful study. Chemical Action of Heat on Compounds. It has already been seen that the change in a substance produced by heat may be reversible or irreversible, and that in the former case the change is one of dissociation, in the latter one of decomposi- tion. A change which is really dissociation may, however, be described as decomposition if one or more products are gaseous and escape. Thus, limestone may be said to be decomposed by heat because the carbon dioxide is allowed to escape; similarly the action of heat on lead nitrate will generally be regarded as one of de- composition unless the reversibility of the reaction is recognized. Thermal Dissociation. The general principles of thermal dissociation have been studied in the previous chapter by reference to the case of mercuric oxide; the subject will now be further illustrated by a few chosen examples. Nitrogen Peroxide. Nitrogen peroxide exists at ordinary atmospheric temperature as a brown gas which becomes deeper in colour when heated, but paler when cooled, yielding at 10 C. a yellow liquid which at 10 C. forms colourless crystals. The alteration in colour of the gas is due to the following reversible change: N 2 O 4 ^ri 2NO* ^w^ N0 2 being deep brown, while N 2 O 4 is pale. Thus one molecule of N a O 4 is dissociated by heat into two similar molecules of NO 2 . The CHEMICAL CHANGES CLASSIFIED 217 progress of dissociation is indicated by an increase of pressure of the gas at constant volume, or by a reciprocal decrease in density at constant pressure -after a correction has been made in either case for change of temperature. Thus the density of nitrogen peroxide gas at 4 -2 C. is 2-588 (air = 1) to which a molecular weight of 74 -8 corresponds, whilst its density at 97-5 C. is 1-783, with a corresponding molecular weight of 51*5. These calculated mole- cular weights show that 62-6 and 11 9 per cent respectively of the mixed gases is N 2 O 4 . This is the simplest example of dissociation that can be given, for the whole system is gaseous, and there is only one kind of dis- sociation product. The formation of N 2 O 4 from two molecules of N0 2 may be regarded as -a case of polymerization. Phosphorus Pentachloride. The vapour of phosphorus penta- chloride readily dissociates thus: PC1 5 ^ PC1 3 + C1 2 , the dissociation at various temperatures being indicated by the fol- lowing densities, reduced to normal temperature and pressure: Temperature C 182 200 250 300 336 Density 73-3 70-0 57-6 52-4 52-5 The normal density of PC1 5 is 104 2; thus dissociation, considerable at 182 C. is complete at 300 C. Now, suppose that instead of vaporizing into empty space, PC1 5 is made to vaporize into a space containing much vapour of PC1 3 . Every C1 2 molecule derived from PC1 5 by incipient dissocia- tion would then encounter many molecules of PC1 3 , and its chance of remaining uncombined with one of these surrounding molecules would be very small. Thus PC1 3 molecules would hinder the dis- sociation of PC1 5 . This has been found to be the case. Wiirtz vaporized PC1 5 into PC1 3 vapour and found the PC1 5 vapour density at 160 to 175 C. to be nearly 104-2, that of the undissociated compound. This is a notable case of the influence of mass on chemical change. Vaporization of PC1 5 into chlorine gas would produce a similar result. Ammonium Chloride. This salt, when vaporized, dissociates into NH 3 and HC1, thus: NH 4 Cl ^=: NH 8 + HC1. 218 CHEMICAL THEORY The dissociation may be proved by the separation of ammonia from hydrogen chloride by the more rapid diffusion of the lighter am- monia through a porous pipe stem. Mercurous Chloride. The vapour density of mercurous chloride ordinarily corresponds to the molecular weight of HgCl, but it may be shown that the vapour is a mixture of Hg and HgCl 2 , since the mercury will diffuse through a porous tube and condense in globules outside. Consequently the formula for mercurous chloride is Hg. 2 Cl 2 , and the vapour dissociates thus: Hg 2 Cl 2 ^ Hg + HgCl 2 . It has been shown, however, by Baker that when water is very carefully excluded the vaporized salt does not dissociate, but shows a density corresponding to the formula Hg 2 Cl 2 . Calcium Carbonate. This compound dissociates thus: CaCO 3 ^^ CaO + (X) 2 , and the dissociation pressures of this salt at various temperatures are these: t C. Pressures. 547 27 mm. Hg. 610 46 740 255 810 678 865 1333 Dissociation ceases at any given temperature when the pressure reaches the corresponding value; consequently the accumulation of the gaseous product of dissociation hinders dissociation. This is a principle of mass action which is true in general. Its application in the case under consideration may be seen in this way. If pre- cipitated chalk is to be converted into quicklime by ignition in a crucible, whilst it is desirable to have as high a temperature as possible, the lid must not be allowed to cover the crucible too tightly or carbon dioxide will accumulate within it and hinder dissociation. This example of dissociation differs from the fore- going, because one of the dissociation products is permanently solid. CHEMICAL CHANGES CLASSIFIED 219 Barium Peroxide. Barium peroxide furnishes an important example of dissociation. When heated to a dull red heat it yields oxygen, thus: 2Ba0 2 ^= BaO + O 2 , and the resulting BaO will reabsorb oxygen at a lower temperature than that at which the peroxide dissociates. If, however, the material is kept at a constant temperature of 700 C., the direction of the reaction may be reversed by altering the pressure. In con- tact with air at 2 atmospheres pressure BaO at 700 C. absorbs oxygen forming BaO 2 ; but when the air is pumped away the re- action is reversed, and oxygen is evolved from the BaO 2 . This is the principle of Brin's process for obtaining oxygen from the air, and as the direction of the reaction depends upon the concen- tration of the reacting substances, it is a good example of the action of mass. Cry stallo- hydrates. The numerous salts and other ^substances containing water of crystallization afford material for illustrating the phenomena of dissociation. Water of crystallization is often but loosely retained by a salt, and it may sometimes be lost even at atmospheric temperature. The student will now appreciate the fact that the condition of a salt with reference to water of crystallization depends not only on temperature, but also on the concentration of water vapour in the vicinity of the salt. There is a certain range of temperature and external pressure of water vapour within which a given salt or its hydrate can exist permanently. Outside this range the com- pound will give rise to a product containing a different proportion of combined water. Lowering of temperature and increase of ex- ternal water- vapour pressure will promote the formation of a higher hydrate; raising of temperature and diminution of external vapour pressure will cause the salt to lose water, so as to form a lower hydrate or become anhydrous. Thus some hydrated salts are permanent in air because of favourable temperature and water -vapour pressure, while others tend to lose water or absorb it from the air. Crystals of blue vitriol, CuSO 4 5H 2 O, for example, neither lose nor gain water under ordinary atmospheric conditions, but crystals of washing-soda lose water according to the scheme: Na a C0 3 .10H 2 ^ Na 2 C0 3 -H 2 + 9H a O, 220 CHEMICAL THEORY because the vapour pressure of the system consisting of the deca- arid mono-hydrate is greater than the pressure of water vapour in the air at ordinary temperature.' So the salt effloresces, becoming opaque on the surface, owing to the formation of the powdery monohydrate, efflorescence being the phenomenon in which hydrated crystals lose water vapour to the air. Hydrated salts having small vapour pressures are frequently very soluble in water, and their saturated solutions may have vapour pressures less than the vapour pressure of water vapour in the air. Atmospheric water vapour will be absorbed by a salt of this kind until a liquid solution is produced of such a strength as to have a water-vapour pressure equal to that in the superincumbent atmosphere. Then the salt is said to , deliquesce. Crystallized calcium chloride, CaCl 2 6H 2 O, is a good example of a deliquescent salt. Thermal Decomposition. It has been seen in the previous chapter that whilst some reactions caused by heat are reversible, being cases of dissociation, others, such as the decomposition of complex organic substances, are undoubtedly irreversible. The question now arises as to the extent and boundaries of reversible and irreversible thermal change, and whether many changes which at present appear irreversible are really so. The investigation of this question leads to some interesting results, but not to a definite conclusion in every case. The thermal decomposition of oxides, hydroxides, and oxysalts furnishes a sufficiently wide field for investigation. Thermal Decomposition of Oxides, The student has already learned that the thermal decomposition of mercuric oxide is reversible; the action is one of dissociation, thus: 2HgO ^= 2Hg+0 2 . Silver and auric oxides are similarly decomposed by heat: 2Ag 2 O = 2Au 2 O 2 = 4Au + 3O 2 but these reactions are not reversible; gold and silver are not sus- ceptible of atmospheric oxidation; they are too electro-negative, too inert chemically, for that. Silver, however, is oxidized by ozone. CHEMICAL CHANGES CLASSIFIED 221 Cupric oxide is an interesting case; when heated to a white heat it loses half its oxygen, the following reaction being reversible: 4CuO ^= 2Cii 2 O + O 2 . No oxides of metals more electro-positive than mercury lose all their oxygen when heated, The loss of some of their oxygen by higher oxides under the influence of heat furnishes an instructive series of phenomena. Consider, for example, the oxides N 2 5 , P 2 O 5 , As 2 O 5 , Sb 2 O 5 , Bi 2 O 6 , which are decomposed by heat, thus: 2N 2 O 6 4NO 2 + O 2 . (P 2 O 6 stable. l2P 2 6 P 4 6 + 2 . 2As 2 O 5 As 4 O 6 + 2O 2 . 2Sb 2 O 6 - 2Sb 2 O 4 + O 2 . Bi 2 O 6 Bi 2 O 3 + O 2 . None of these reactions appears to be reversible. The behaviour of N 2 O 5 is unique, like that of the element nitrogen itself. P 2 O 5 is easily produced by the atmospheric oxidation of the element or the lower oxide; but -arsenic is less oxidizable than phosphorus, and the higher oxide must be formed by the chemical oxidation of the lower oxide, which alone is produced when the element burns. The oxide Bi 2 O 5 is a peroxide, formed only in the wet way; it has feebly acidic properties, but easily loses '-oxygen. Thus the oxides P 2 O 5 , As 2 O 5 , Sb 2 O 5 , Bi 2 O 5 , with the possible exception of Sb 2 O 5 , stand in the order of decreasing stability. The dioxides of the fourth group CO 2 , SiO 2 , GeO 2 , SnO 2 , PbO 2) are similarly related as regards stability; PbO 2 alone is decomposed by heat. Or consider the superoxides Na 2 O 2 , BaO 2 , SrO 2 , CaO 2 , MgO 2 , &c. Sodium peroxide does not lose oxygen when heated short of very high temperatures; BaO 2 ,and the oxides which follow it, are decomposed with increasing readiness as electro-positiveness diminishes. Only the more electro -positive metals form super- oxides, i.e. derivatives'- of hydrogen peroxide, at all. The reversible reaction, 6PbO + O 2 =F= 2Pb 3 O 4 , is instructive, and it is noteworthy that PbO 2 is not formed by the atmospheric oxidation of PbO; Pb 3 O 4 is a compound of 2 PbO and PbO 2 , that is, it is a salt, and the two oxides being mutually 222 CHEMICAL THEORY satisfied, the PbO in Pb 3 O 4 is not free to combine with more oxygen. If, however, another base is present with which PbO 2 can combine, PbO may be completely oxidized to Pb0 2 by atmospheric oxygen. Thus the following reaction is reversible: 2 PbO + O 2 + 2 Na 2 CO 3 ^ . 2 Na 2 PbO 3 + 2 CO 2 . This is also a good example -of mass action, since preponderance of oxygen or carbon dioxide determines the direction of the reaction. The case presented by the following reaction is similar: 4CrO 3 2Cr 2 O 3 + 3O 2 . The reverse reaction takes place only in the presence of alkali, when atmospheric oxygen may be absorbed at high temperature to form chromate, thus: 2Cr 2 O 3 + 3O 2 + 4Na 2 CO 3 4 Na 2 CrO 4 + 4 CO 2 . Amongst non-metallic oxides the case of sulphur trioxide is important. It is well known that the reaction 250 2 + O 2 -> 2SO 3 takes place to any considerable extent only in presence of a catalyst, whilst the reverse reaction, 250 3 -> 2SO 2 + O 2 , occurs when the vapour of the trioxide is passed through a tube heated to 1000 C. Nevertheless, when the dioxide is heated to 1200 C., or is submitted to the action of electric sparks, it undergoes the following reversible change: 3S0 2 ^= 2S0 3 Thermal Decomposition of Hydroxides. Hydroxides may be basic or acidic, and their stability depends on the basic or acidic intensity of the corresponding oxides. Consider, for example, the basic hydroxides: NaOH, Ba(OH) 2 , Ca(OH) 2 , Mg(OH) 2 , A1(OH) 3 , Cu(OH) 2 , AgOH, NH 4 OH. NaOH and Ba(OH) 2 , producible from the corresponding oxides by violent reactions with water, cannot be dehydrated by heat alone; Ca(OH) 2 , produced also from the oxide by slaking, is easily dehydrated by heating above 150 C.; Mg(OH) 2 is formed, not by CHEMICAL CHANGES CLASSIFIED 223 slaking, but by precipitation, and is stable at 100 C.; A1(OH) 3 , formed by precipitation in the cold, loses water, forming A1 2 O(OH) 4 or A1OOH, when heated at 100 C. or dried over sulphuric acid; Cu(OH) 2 is formed as a blue precipitate in the cold, which turns dark brown when warmed with water, forming Cu 4 3 (OH) 2 ; AgOH exists only in very dilute aqueous solution, for the precipitate formed by adding alkali to a silver salt is Ag 2 O, which dissolves very slightly in water, producing ^a faintly alkaline solution con- taining AgOH. Ammonium hydroxide. NH 4 OH, is unique. A hydrate, NH 3 H 2 O exists in crystals, melting at 78 C., but it cannot be asserted that this is ammonium hydroxide. A solution of ammonia in water contains much NH 3 as well as ionized NH 4 OH, so that the following reactions occur: NH 3 + H 2 O ^r: NH 4 OH. NH 4 OH :^= NH 4 '+ OH'. The presence of both NH 3 and hydroxide ions is proved by the use of ammonia solution to form an ammine, such as CuSO 4 -4NH 3 'H 2 0, as well as to precipitate a base such as Fe(OH) 3 . When an aqueous solution of ammonia is boiled, the above reactions are reversed, and all the ammonia escapes as gas. The following acidic hydroxides (oxy-acids) may be considered: S0 2 (OH) 2 , PO(OH) 3 , SO(OH) 2 , CO(OH) 2 . Sulphuric acid, S0 2 (OH) 2 , is fairly stable towards heat and may be distilled, for SO 3 is a powerfully acidic oxide, and as such holds its combined water tenaciously. At 440 C., however, dis- sociation according to the scheme H 2 SO 4 -> SO 3 + H 2 O is complete. Phosphoric acid, PO(OH) 3 , loses water when heated, thus: 2PO(OH) 3 -> P 2 O 3 (OH) 4 P 2 O 4 (OH) 2 Ortho- Pyro- M eta-phosphoric acid; but it is noteworthy that the meta-acid is the final dehydration product, from which water cannot be removed. The meta- and the pyro-acid both revert to the ortho-acid in contact with water, and there is evidence that the meta-acid passes through the pyro form on the way back to the ortho-acid. Sulphurous acid, SO(OH) 2 or SO 2 : H(OH), shares with carbonic 224 CHEMICAL THEORY and nitrous acids, which also have gaseous anhydrides, the property of existing only in aqueous solution; several crystallo-hydrates of this acid exist however. The following reaction is reversible: S0 2 + H 2 ^ SO(OH) 2 , though, like ammonia, possibly some sulphur dioxide exists in aqueous solution without being hydroxylated. It may be that sulphurous acid in aqueous solution passes into the unsymmetrical form, thus: SO(OH) 2 =^ at any rate, the sulphites appear to conform to the latter formula. Carbonic acid, CO(OH) 2 , is formed when carbon dioxide dis- solves in water, thus: C0 2 + H 2 ^ CO(OH) 2 ; the carbonic acid so formed being ionized to a considerable extent, thus: H 2 CO 3 ^i H* + HCO 3 '. When a solution of carbonic acid is boiled, both these reactions are reversed, and all the carbon dioxide escapes from the liquid. Thermal Decomposition of Oxy-Salts. The usual mode of thermal decomposition of oxy-salts is into basic and acidic oxides or their decomposition products; e.g.: CaCO 3 CaO + CO 2 . 2Cu(NO 3 ) 2 Whether such a reaction occurs depends upon the relative basic and acidic strengths of the respective oxides, and upon the volatility of the acidic oxide. Thus, when both basic and acidic oxides are powerful, as, for example, in the case of sodium sulphate, thermal decomposition does not take place, and even when the acidic oxide is feeble an alkali oxide is sufficiently powerful to retain it, as in the case of sodium carbonate, which is not decomposed at 1000 C. The sulphates of the feebler basic oxides lose sulphur trioxide when heated; for example, copper sulphate forms a basic salt at a dull red heat, and ferrous sulphate loses all its SO 3 , leaving a residue of ferric oxide. The behaviour of nitrates and chlorides when heated is instruc- CHEMICAL CHANGES CLASSIFIED 225 tive. Those of the alkali metals do not lose their nuclear nitrogen or chlorine; thus potassium nitrate forms nitrite, and chlorate chloride. Nitrates and chlorates of feebler metals, however, leave a residue of oxide; e.g. the lead salts decompose thus; 2Pb(NO 3 ) 2 2Pb(ClO 3 ) 2 An interesting question arises here as to reversibility. The reaction: 2KC1O 3 2KC1 + 3O 2 appears not to be reversible. Why is this? Hypochlorite easily undergoes self -oxidation and reduction thus: 3KOC1 2KC1 + KC1O 3 ; yet chloride seems incapable of oxidation to chlorate. The explana- tion may be that K Cl cannot become K O Cl by interpolation of an oxygen atom, but K O Cl can combine with other oxygen atoms to produce: K-0-< Besides relative basic and acidic strengths of oxides in an oxy- salt, the volatility of the acidic oxide determines the stability of the salt. Silica, for example, is not volatile like carbon dioxide, and so calcium silicate, unlike calcium carbonate, is not decomposed by heat. For the same reason silica at high temperature displaces carbon dioxide from chalk, thus: CaCO 3 + SiO 2 -* CaSiO 3 -f CO* a reaction common in metallurgy. Phosphates, moreover, are not decomposed by heat; for example, bone ash, Ca 3 (PO 4 ) 2 , is unchanged at a white heat; for, although phosphoric oxide itself is volatile, it cannot be separated from a combined basic oxide any more than it can be separated from water combined with it in metaphosphoric acid, H 2 P 2 O 6 . In general, the mode of decomposition of an oxy-compound when heated reveals the relative stabilities at high temperatures of the possible decomposition products. Thus the isomorphous salts KC1O 4 and KMnO 4 behave very differently when heated, because in spite of the superficial resem- blance between them and the similarity of their constitution, the elements concerned, namely chlorine and manganese, are widely (D60) 16 226 CHEMICAL THEORY different in chemical nature. The following reactions represent the manner of decomposition of these salts: KC10 4 = KC1 + 20 2 . 2 KMn0 4 = K 2 Mn0 4 + MnO 2 + O 2 . The chief ammonium oxy-salts are decomposed by heat thus: f(NH 4 ) 2 S0 4 \NH 4 HSO 4 NH 4 NO 3 NH 4 N0 2 (NH 4 ) 2 HP0 4 /(NH 4 ) 2 C0 3 \NH 4 HCO 3 = NH 4 HSO 4 4NH 3 . = H 2 O 4 SO 3 4 NH, = H 2 O 4 N 2 . = HPO 3 + H 2 O-f 21 = NH 4 HCO 3 4- NH 3 . = H 2 O 4 CO 2 4 NH, Chemical Interaction of Water with Elements and Compounds. The interaction of elements and pure water is limited almost completely to the behaviour of a few electro-positive metals. The metals 'of the alkalis and alkaline earths decompose water at atmospheric temperature, displacing hydrogen with formation of the hydroxide of the metal. The vigour of the reaction with water increases with the electro-positiveness of the metal from lithium to caesium, and from calcium to barium. The reaction is attributable firstly to the ionization of water, although this is so small, and secondly, to the solution pressure of the metal which is superior to that of hydrogen. The student has met with this latter idea before, under the subject of electrolysis, (q.v.). Thus metallic sodium, striving to assume the ionic state when brought into contact with, water, displaces the hydrogen ions of the latter, causing them to lose their electric charges and escape as gas. Besides the metals of the alkalis and alkaline earths, amalga- mated aluminium reacts with water at atmospheric temperature evolving hydrogen, and powdered aluminium as well as magnesium decomposes steam. The rusting of iron is believed to be due, first of all, to the action on water of the metal containing impurities which set up slight electro-potential differences throughout the mass; and it is well known that steam is decomposed by red-hot iron, whilst hydrogen can reduce heated iron oxide to metal. Copper, however, has no action on water or steam, for it is less electro-positive than hydrogen and does not displace this element from water or dilute acids. CHEMICAL CHANGES CLASSIFIED 227 So the following reactions are instructive: i. Mg + H 2 O -> MgO-fH 2 . ii. 3Fe + 4H 2 O ^ Fe 3 O 4 + 4H 2 . iii. Cu + H 2 O CuO + H 2 . Reactions (i) and (iii) proceed in single and opposite directions; reaction (ii) is reversible because of the intermediate character of iron. Of the non-metals the halogens alone react with water. Fluorine decomposes water in the dark, and chlorine in the day- light, both with evolution of oxygen. It is probable that the following is the first reaction between chlorine and water: C1 2 + H 2 O ^^ HC1 + HOC1; for chalk and chlorine water yield calcium chloride and hypo- chlorous acid, the latter not reacting with chalk. Hydrolysis. The chemical decomposition of compounds by water is hydro- lysis. This, again, is to be attributed to the hydrogen and hydroxide ions existing in water. Ideal salts, consisting of powerfully basic and acidic ions, are not hydrolyzed by water; their aqueous solutions are ionized but remain neutral. Normal salts are not necessarily neutral. Sodium chloride and trisodium phosphate, Na s P0 4 , are both normal salts, but whilst NaCl is also a neutral salt in the sense of giving a neutral solution with water, this is not the case with Na 3 PO 4 , the solution of which is alkaline on account of hydrolysis with production of free alkali, i.e. OH' ions in solution, thus: Na 3 P0 4 + HOH ^ Na 2 HP0 4 + NaOH. NaOH ^= Na' + OH'. Even the hydrogen salt Na 2 HPO 4 yields an alkaline solution owing to hydrolysis, thus: Na 2 HPO 4 + HOH ^= NaH 2 PO 4 + NaOH; for although the salt NaH 2 PO 4 is itself acid, it does not produce such a concentration of H* ions as the equivalent of NaOH pro- duces of OH' ions. 228 CHEMICAL THEORY The reactions of the three sodium phosphates are: Na 3 PO 4 , strongly alkaline; Na 2 HPO 4 , alkaline; NaH 2 PO 4 , acid; the point of neutrality thus lying between Na 2 HPO 4 and NaH 2 PO 4 . An alternative way of explaining the alkalinity of Na 3 PO 4 , applicable to the effect of mixing equivalents of NaOH and H 3 PO 4 , is that NaOH being stronger as a base than H 3 PO 4 is as an acid, the alkali provides a larger proportion of OH' ions than the acid of H* ions, and so an excess of OH' ions remains over after all the H* ions from the acid have been neutralized. These considerations have an important bearing on the volu- metric estimation of acids and alkalis. Only acids and alkalis which produce salts not appreciably hydrolyzed in aqueous solution can be titrated in the ordinary way; obviously phosphoric acid cannot be directly titrated with standard alkali. As regards salt hydrolysis in general, whilst the salts of strong bases with weak acids are alkaline ^in reaction, those of weak bases with strong acids are acid. An example of the latter is furnished by ferric chloride which is hydrolyzed in aqueous solution with formation of a basic salt and free acid, somewhat as follows: FeCl + HOH =^ The formation of the basic salt is shown by the darkening of the solution. The same thing is especially noticeable when ferric alum is dissolved in water. The pale-violet crystals yield a brown solution which becomes colourless when a little sulphuric acid is added to convert the basic salt into the normal salt. If a solution of ferric alum is poured into much boiling water a precipitate of ferric hydroxide separates owing to complete hydrolysis. Ferric acetate is similarly hydrolyzed by boiling water with the precipi- tation of a basic acetate. That salt hydrolysis is a frequent phenomenon is shown by the following facts: Soluble salts of the following acidic radicles have alkaline reactions in solution: Borate, carbonate, chromate, cyanide, hypochlorite, nitrite, phosphate, silicate, sulphide, sulphite. Soluble salts of the following metallic radicles possess an acid reaction: CHEMICAL CHANGES CLASSIFIED 229 Mercurous, mercuric, cupric, aluminium, chromic, ferrous, ferric, stannous, stannic, antimonious, bismuthous. Salts of weak bases with weak acids are generally insoluble in water, and on that account less amenable to hydrolysis. A fat may be hydrolyzed by superheated steam with the pro- duction of glycerol and fatty acid, thus: C 3 H 5 X 3 + 3HOH = C 3 H 6 (OH) 3 + 3HX. (Q,H 6 )'" = glyceryl. X = C 16 H 31 COO.(palmitate), orC^ The catalytic effect of H* ions in promoting the hydrolysis of an ester or of cane sugar has already been noticed (p. 211). Sometimes the basic salt produced by hydrolysis is insoluble and is precipitated. Thus, bismuthous and antimonious chlorides are not only hydrolyzed by water but yield precipitates of the basic chlorides: BiCl 3 + H 2 O :^= BiOCl + 2HCl. SbCl 3 + H 2 O ^ SbOCl + 2HCl. On this account hydrolysis is accentuated, for the reversal of the reaction is greatly hindered by the separation of the hydrolytic product in the solid state. Incidentally it may be observed that whilst BiOCl is stable towards water, SbOCl, on account of the feebler basic properties of antimonius oxide, loses all its chlorine when boiled with water. So far the hydrolysis of salts has been considered; but the process is not confined to salts. Consider the series of chlorides: NaCl, MgClfc A1C1 3 , SiCl 4 , PC1 5 . The series begins with a powerfully metallic chloride, and ends with a non-metallic chloride. The chloride of a powerful metal, being a true salt, is not hydrolyzed by water, whilst the chloride of a non-metal is at once and completely hydrolyzed. Between these two extremes there is a gradation of hydrolysis. MgCl 2 gives a neutral solution with water, but when this solution is evaporated nearly to dryness hydrogen chloride gas escapes with the steam, and the resulting pasty mass, containing the basic salt Mg(OH)Cl, reacts alkaline. By ignition the oxide MgO is eventually formed. AlCl z dissolves in water, reacting vigorously with it if anhydrous, 230 CHEMICAL THEORY and produces an acid solution by incipient hydrolysis; when this solution is evaporated to dryness and ignited, the oxide A1 2 O 3 remains. SiCli is instantly decomposed by water with separation of gelatinous silica H 2 SiO 3 : SiCl 4 + 3H 2 O -* H 2 SiO 3 + 4 HC1. To a certain extent the reaction is reversible, for H 2 Si0 8 is more soluble in hydrochloric acid than in water. PC1 5 is similarly decomposed, liquid POC1 3 being first formed, and then H 8 P0 4 thus: PC1 6 + H 2 POC1 3 + 2HC1. POC1 3 + 3H 2 O PO(OH) 3 + 3HC1. This is an extreme case of hydrolysis, and the reaction is not reversible. Chemical Interaction of Acids with Metals. It has been seen that there is an interaction between water and some metals, and that this is believed to be due to the slight ioniza- tion of water without which this substance would be inert. Now acids owe their essential nature to the presence of hydrogen ions, the strength of an acid in aqueous solution being measured by the concentration of these ions. Consequently, acids containing a much higher concentration of hydrogen ions behave much more vigorously toward metals than water. Consider the following metals arranged in order of decreasing electro-positi veness : Cs, Rb, K, Na, Ba, Sr, Ca, Mg, Al, Mn, Zn, Cd, Tl, Fe, Co, Ni, Sn, Pb, H, Sb, Bi, As, Cu, Hg, Ag, Pd, Pt, Au. Metals preceding hydrogen can displace it as a gas from dilute hydrochloric or sulphuric acid; those that follow H do not generate this gas in contact with dilute acid. Or, using an idea the student is already familiar with: metals preceding hydrogen have a greater, metals following it a less, solution pressure than this element. It has already been seen that metals as far as calcium, and under some circumstances aluminium, displace hydrogen from cold water, and that if steam is used and the metal is heated reactivity extends as far as iron. CHEMICAL CHANGES CLASSIFIED 231 The extent to which a metal displaces hydrogen from a dilute acid depends upon: i. The nature and purity of the metal. ii. The nature and state of dilution of the acid. i. The nature and purity of the metal. If dilute hydrochloric acid is poured upon fragments of magnesium, zinc, iron, and tin, in separate test-tubes, the reaction will be very vigorous with magnesium, less so with zinc, still less so with iron, and very slight with tin; indeed, it is necessary to heat tin with moderately con- centrated hydrochloric acid to dissolve the metal with evolution of hydrogen. All this is quite in accord with the order of electro- potential of the metals. The purity of a metal has a marked effect on its interaction with acid. If a stick of pure zinc is placed in dilute hydrochloric acid, very slow action takes place; but if a length of fine copper wire is coiled round the zinc, action becomes vigorous and the hydrogen is seen to be coming off from the surface of the copper, although this metal is found unchanged when the zinc has been dissolved. These different effects are due to the fact that in the first case the displaced hydrogen forms a protective film on the surface of the zinc, thus polarizing it, whilst in the second case electrical action is set up, hydrogen flows with the electric current from the zinc to the copper, and is evolved from the surface of the latter metal, whilst the zinc is left continuously exposed to the action of the acid. A similar effect is produced by adding a few drops of copper sulphate or platinic chloride to the acid in which the zinc is immersed. Copper or platinum is deposited on the zinc, and causes the production of electric circuits with the consequent promotion of chemical action. If the zinc is impure, and contains a small proportion of a less electro-positive metal, solution is hastened by this impurity without the help of an added metal. ii. The nature and state of dilution of the acid. The strength of an acid, i.e. its degree of ionization, determines the rate at which a metal displaces hydrogen from it. Thus, if parallel experiments are done with zinc in (a) dilute hydrochloric acid, (6) dilute acetic acid of equivalent strength, the rate of evolution of hydrogen from the acetic acid will be exceedingly slow, whilst that from the hydro- chloric acid may be vigorous. The explanation is simple. Hydro- chloric acid is a strong acid, almost completely ionized in dilute 232 CHEMICAL THEORY solution, whilst acetic acid is a weak acid which is but slightly ionized, and therefore approaches water in its behaviour towards metals. The state of dilution of a particular acid also determines the rate of evolution of hydrogen. If an acid is much diluted the concentra- tion of hydrogen ions is correspondingly reduced, together with the vigour of the reaction. If, however, too little water is present there may not be room enough in the solvent for extensive ionization, and consequently displacement of hydrogen by a metal will not be effected. . This is not the case with concentrated hydrochloric acid, which contains only about 32 per cent of hydrogen chloride; but it is the case with concentrated sulphuric acid, which consists of about 98 per cent of the absolute acid and 2 per cent of water. Thus zinc does not displace hydrogen from concentrated sulphuric acid; a reaction commences when this acid is heated with the metal; the gas evolved, howevejp, is not hydrogen but sulphur dioxide, thus: Zn + 2 H 2 SO 4 = ZnSO 4 + 2 H 2 O + SO 2 . Since a similar reaction takes place with copper, it is not necessary to assume that free hydrogen has anything to do with it. Interaction of Nitric Acid and Metals. Since nitric acid contains much oxygen, and is easily reducible, the question arises whether hydrogen can escape from the acid when displaced from its molecules by a metal, or whether it will neces- sarily reduce the acid instead. If the acid is reduced, since there are various reduction products of nitric acid, the possible reactions of the acid with metals are various. Hydrogen is among the gases evolved from dilute nitric acid by magnesium, the most electro-positive of the metals which do not react with cold water; but with this exception hydrogen is not obtained. The possible reduction products of nitric acid may be shown by separating water from the acid, and breaking up the resulting anhydride thus: , 2 HN0 3 = H 2 + N 2 6 , = H 2 O + 2 N0 2 + O t = H 2 + N 2 3 + 2 O, H 2 O + N 2 O + 4 H 2 O + N 2 + 5 O. CHEMICAL CHANGES CLASSIFIED 233 Besides the five reduction products, NO 2 , N 2 O 3 , NO, N 2 O, N 2 , hydroxylamine, NH 2 OH, and ammonia, NH 3 , are sometimes pro- duced. The question now arises whether displaced hydrogen Or the metal itself reduces the acid. Hydrogen cannot reduce nitric acid in the case of a metal which does not displace this element from an acid. It can scarcely be assumed that nascent hydrogen is respon- sible for the reduction of nitric acid by copper, since this metal does not displace hydrogen from any acid. As a matter of fact, metals may be divided into two categories as regards their behaviour to- wards nitric acid. The metals zinc, cadmium, iron, tin, and others more electro-positive than hydrogen may reduce nitric acid as far as ammonia; the metals bismuth, copper, mercury, silver, which are less electro-positive than hydrogen, do not reduce the acid beyond the stage of nitric oxide. Presumably the reduction to nitric oxide does not necessitate the intervention of hydrogen; reduction beyond this stage may. Dis- placed hydrogen must evidently play a part in the formation of NH 2 OH and NH 3 ; it may be indirectly the cause of the evolution of nitrogen, which may be derived from ammonium nitrite thus: NH 4 N0 2 = N 2 + 2H 2 0, whilst N 2 O may come from hyponitrous acid thus: HON : NOH = N 2 O + H 2 O. With regard to the reduction of nitric acid by metals of the copper series there is evidence that the pure, diluted acid, free from nitrous acid, does not react with these metals, but that nitrous acid is necessary to start the reaction, which proceeds thus: i. Cu + 4 HNO 2 = Cu(NO 2 ) 2 + 2 H 2 O + 2 NO; ii. Cu(N0 2 ) 2 + 2HNO 3 = Cu(NO 3 ) 2 + 2 HNO 2 ; iii. HNO 3 + 2 NO + H 2 O = 3 HNO 2 ; and adding iv. Cu + 3 HNO 3 = Cu(NO 3 ) 2 + HNO 2 + H 2 O. So, by the aid of a little HNO 2 more is produced, which, instead of accumulating, decomposes thus: v. 3 HNO 2 = HNO 3 + H 2 O + 2 NO. Thus nitric oxide gas is evolved, and the nitrate of the metal is formed in solution. 234 CHEMICAL THEORY This theory of the activity of nitrous acid does not invalidate the usual equation: 3 Cu + 8 HN0 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO, which may be obtained by multiplying equation (iv) by three, and adding the product to equation (v). The concentration of the nitric acid employed affects the nature of the reduction products. This may be seen by pouring very dilute nitric acid on zinc in a test-tube, and then gradually increasing the strength of the acid. At first no gas is evolved, then a colourless gas, and afterwards a brown gas. With the most dilute acid ammonia is produced, and combines with the acid, forming ammonium nitrate; then, with increasing strength of acid, nitrogen and nitrous oxide appear, and afterwards nitric oxide, and possibly still higher oxides of nitrogen with the strongest acid. In general, the more dilute the acid the more perfectly it is reduced. The following equations represent the formation of nitrous acid and ammonia respectively when zinc and nitric acid interact. 4 Zn + 10 KNTO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O. 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 -f 3 H 2 O + NH 4 NO 3 . That the quantities on the left side of the equation are the same in both cases is interesting. To argue that the same products should therefore result would be fallacious, for chemical equations do not represent the concentrations in which reacting substances are brought together. Effect of Solubility and Volatility on Chemical Change. There are two laws associated with the name of Berthollet which state (i) that if two substances by reacting in solution can yield a product less soluble than themselves, that product will be formed; and (ii) that if two substances when heated together can yield a product more volatile than themselves, that product will result. The facts thus expressed find ready explanation according to the principles of mass action. The direction of the reaction, AX + BY ^= AY + BX, will depend on the effective concentration of each component; if one component, from any cause, is placed partly or wholly outside CHEMICAL CHANGES CLASSIFIED 235 the sphere of action, its power of directing the course of the reaction will be so far diminished. Consider, for example, the following two reactions placed at the beginning of this chapter: BaCl 2 -fH 2 SO 4 BaSO 4 + 2HCl. 2NaCl + H 2 SO 4 Na 2 SO 4 + 2HCl. Superficially these reactions appear to be alike; actually they differ very much. The first is supposed to take place in dilute solution, BaSO 4 being precipitated and HC1 remaining in solution; the second takes place when solid NaCl and concentrated H 2 S0 4 are heated together, solid Na 2 SO 4 resulting, while HC1 escapes. If the conditions are reversed, the results are not so effective. Thus if solid BaCl 2 is heated with concentrated sulphuric acid the reaction will be only partial, because the BaCl 2 will be protected from the acid by a crust of BaSO 4 ; and if NaCl and H 2 SO 4 are mixed in dilute aqueous solution no HC1 gas escapes. The first reaction ordinarily proceeds to finality because of the insolubility of BaSO 4 . When BaCl 2 and H 2 SO 4 interact in aqueous solution, BaSO 4 and HC1 are first formed in solution, and if both remained in this state they would be effective in establishing an equilibrium with the original substance, and no change would be apparent. But BaSO 4 is at once precipitated because it is so very slightly soluble in water, and only the very minute quantity of this salt remaining in solution can have any effect in reversing the change. This effect is quite negligible, and, practically speaking, the reaction proceeds to completion from left to right. The second reaction proceeds to finality because of the volatility of hydrogen chloride. When NaCl is heated with concentrated H 2 SO 4 , torrents of HC1 gas are evolved, and the reversal of the reaction by means of this gaseous product which has escaped is out of the question. The principles here illustrated are very far-reaching, and the student should make sure he grasps them. Consider another example: The preparation of nitre by inter- action of sodium nitrate and potassium chloride in aqueous solution, thus: NaNO 3 + KCl ^^ NaCl + KNO 3 . All the four possible salts are soluble in water, but sodium chloride is the least soluble of the four. When, therefore, sodium nitrate 236 CHEMICAL THEORY and potassium chloride are mixed in equivalent proportions in solution and the latter is evaporated, sodium chloride is the first salt to crystallize; and thus potassium nitrate remains in solution, and may be crystallized after the sodium chloride has been re- moved. A salt such as sodium chloride may be eliminated more effectually if it is formed in contact with a non-aqueous solvent in which it is not soluble. An interesting example of this is furnished by the preparation of free hydroxylamine, NH 2 OH, from the hydro- chloride by causing the following reaction to take place in methyl alcoholic solution: NH 2 OH.HCl + NaOCH 3 = NaCl + HOCH 3 + NH 2 OH. Sodium methoxide, NaOCH 3 , reacts effectively with hydrogen chloride because of the insolubility of sodium chloride in methyl alcohol, CH 3 OH, and the methyl alcohol produced at the same time becomes part of the solvent, so that after the precipitated sodium chloride has been removed nothing remains in solution but free hydroxylamine. The reaction, Na 2 CO 3 + SiO 2 ^ Na 2 SiO 3 + CO 2 , which is of wide, terrestrial significance, serves further to illustrate the principle of volatility. If silica is heated with sodium carbonate, the reaction proceeds from left to right, because carbon dioxide is volatile and silica is not; but if carbon dioxide acts on sodium silicate in aqueous -solution, the reaction is reversed, because carbonic acid is soluble in water, whilst hydrated silica is not so soluble, and largely separates from solution. Thus, in the early ages of the earth's history much silica was probably combined with bases in the earth's crust, and much free carbon dioxide existed in the atmosphere: but "weathering" of siliceous rocks has now been going on for ages in presence of water and at moderate tempera- ture, with the fixation of carbon dioxide and consequent separation of silica. The phenomena of precipitation, so important in analysis, may now be seen in a clearer light. Some precipitates are amorphous, some crystalline. Suddenly - formed precipitates are frequently amorphous, because they have not had time to crystallize, and occasionally they become crystalline when kept in contact with the liquid from which they have been separated. CHEMICAL CHANGES CLASSIFIED 237 Thus magnesium ammonium phosphate, MgNH 4 PO 4 6H 2 O, when it is precipitated from concentrated solutions, appears amorphous, though when it separates gradually from dilute solutions it is distinctly crystalline. Similarly barium sulphate, precipitated from cold, dilute solu- tions, is so finely divided as to run through the pores of ordinary filter-paper, but when the precipitate is heated with water it slowly becomes granular, that is, micro-crystalline, and can then be easily filtered. The same condition is more quickly assumed if the pre- cipitation is carried out by mixing boiling solutions. Calcium carbonate is another example of an amorphous precipi- tate which may become crystalline; for when heated in the liquid from which it has been precipitated it gradually assumes the crys- talline form of aragonite. Sulphide precipitates are, however, colloidal, 1 and do not crystal- lize. They are exceedingly insoluble in water, though some of them may assume -the state of colloidal suspension. A curious case is presented by the sulphides of nickel and cobalt. These precipitates do not dissolve readily in dilute hydrochloric acid, though this acid prevents their formation. The explanation of this paradox is that on precipitation these sulphides polymerize, i.e. simple molecules combine to form more complex and therefore less soluble molecules. Thus dilute acid prevents the formation of the simple and more soluble molecules of the sulphides, but is unable to dissolve readily the more complex molecules formed by precipitation from alkaline solution. The student should now understand that crystallization and precipitation are linked phenomena. Indeed, crystallization from solution is slow precipitation; the slower the precipitation the larger and more perfect are the crystals; the more rapid the preci- pitation the smaller and less perfect are the crystals, until precipi- tation becomes too rapid, and the separated solid is amorphous. He will also understand that chemical affinity has little if any- thing to do with precipitation. For example, the reaction, NaCl + AgNO 3 = AgCl + NaNO 3 , takes place in aqueous solution, not because silver has a greater affinity for chlorine than sodium, for that is not the case, or because sodium has a greater affinity for nitrate than for chloride, but i Vide Chapter XIII, on the colloidal state. 238 CHEMICAL THEOKY because silver chloride happens to be practically insoluble in water, and a reaction which might be reversible fails to be so because the effective mass of the silver chloride is so small. Thus silver nitrate is a reagent for all chlorides, or, more properly, silver ions are a reagent for chloride ions from whatever source they are derived. Oxidation and Reduction. Berzelius called oxygen the pole of chemistry; and it is true that no other element occupies so significant a position among the rest. It has already been seen that the kinds of oxides an element forms go far to reveal the chemical nature of the element. Again, Lavoisier invented the name oxygen, thinking that this element was a constituent of all acids; and although this idea was false the name was not badly chosen after all, for the addition of oxygen to a substance generally enhances its acidic character. Consider, for example, the oxides of manganese: MnO, Mn 3 O 4 , Mn 2 O 3 , MnO 2 , MnO 3 , Mn 2 O 7 . The first oxide, MnO, is wholly basic, but with successive additions of oxygen the oxides lose their basic power, becoming more and more acidic, until Mn 2 O 7 , a strongly acidic anhydride, is reached. 1 Oxidation is the addition of oxygen to an element or compound', reduction is its removal. The removal of hydrogen is also regarded as oxidation, and sometimes also the addition of chlorine or an equivalent electro- negative atom or radicle. Thus, alcohol is oxidized to aldehyde by the removal of hydrogen in effect: C 2 H 6 + O = C 2 H 4 + H 2 0; and the conversion of a ferrous into a ferric compound is oxidation, however effected; e.g. = Fe 2 (SO 4 ) 3 + H 2 O, or 2FeCl 2 -hCl 2 = 2FeCl 3 . Oxidation generally involves an increase in the valency of the nuclear element of the compound oxidized; for example iron is bivalent in ferrous compounds and tervalent in ferric. This is not CHEMICAL CHANGES CLASSIFIED 239 always the case, however, for carbon is quadrivalent in both alcohol and aldehyde: H H CH 3 -C-OH : CH 3 -C=0, H and barium is bivalent in both BaO and Ba0 2 : / Ba=0 : Ba< | . X) Similarly, the addition of hydrogen to a molecule may be regarded as reduction, as for example in the formation of leuco- compounds from organic dyes through addition of hydrogen by means of sulphurous acid, thus: H 2 SO 3 -f OH 2 = H 2 SO 4 + 2 H. Oxidation and reduction are reciprocal processes; the oxidation of one substance often involves the reduction of another; as for instance the oxidation of ferrous sulphate by nitric acid: 6 FeS0 4 + 3 H 2 S0 4 + 2 HNO 3 = 3 Fe 2 (SO 4 ) 3 + 4 H 2 O + 2 NO. This reciprocal relationship necessarily exists when the reducing and oxidizing substances are both compounds; indeed the process is essentially the transference of oxygen from one compound to another. Common Oxidizing and Reducing Agents. The following are the principal oxidizing and reducing agents met with in inorganic chemistry. OXIDIZING AGENTS Free oxygen and ozone. Hydrogen peroxide and other per- oxides. Reducible basic oxides, e.g. Ag 2 O. The halogens, oxyacids and their salts, e.g. nitric and chloric acids and their salts. Higher oxides and oxyacids of metals, e.g. chromic and permanganic acids and their salts. REDUCING AGENTS Hydrogen, gaseous and "nascent". Unstable hydrides, e.g. H 2 S, HI. Oxidizable elements, e.g. carbon, the alkali metals, and magnesium. Lower oxides, e.g. CO. Lower oxyacids and their salts, e.g. sulphites, nitrites. Lower salts, e.g. ferrous, stannous. Cyanides, formates. 240 CHEMICAL THEORY Conditions of Action. Oxidation and reduction may take place either in presence or absence of water. These two conditions are well illustrated by the reactions of qualitative analysis carried out in the dry way and in solution respectively. A familiar example of oxidation in the dry way is the for- mation of a chromate by oxidizing fusion, thus: Cr 2 O 3 + 2 Na 2 CO 3 + 3 NaNO 3 = 2 Na 2 CrO 4 + 3 NaNO 2 + 2 CO 2 , or Cr 2 O 3 + 3Na 2 O 2 = 2 Na 2 CrO 4 + Na 2 O; and of reduction in the dry way the liberation of a metal from its oxide by heating it with charcoal in a reducing flame, e.g.: SnO 2 + 2C = Sn + 2CO. The dry reactions of qualitative analysis, which are frequently neglected, are at any rate of value because they illustrate in miniature important manufacturing processes. The oxidation of chromic oxide, for example, in presence of alkali, is employed to obtain chromium compounds from the natural source of chromium, chrome ironstone or ferrous chromite: FeO'Cr 2 3 . It is interesting to observe, moreover, that although Cr 2 O 3 is not oxidized when heated alone in the air, and that on the contrary CrO 3 loses oxygen under these conditions, nevertheless, in presence of a base with which CrO 3 can combine chromate is formed, thus: 4 FeO-Cr 2 O 3 + 7 O 2 + 8 CaCO 3 = 8 CaCrO 4 + 2 Fe 2 O 3 + 8 CO 2 . From this it may be inferred that the following is a reversible reaction : 2Cr 2 O 3 + 3O 2 ^= 4CrO 3 , the presence of a base with which it can combine promoting the formation of the acidic oxide CrO 3 , whilst the absence of a base, being the absence of the condition of stability of this oxide, permits its decomposition by heat. Further, it is easy to understand why CrO 3 oxidizes hydro- chloric acid according to the reaction 2 CrO 3 + 12 HC1 = 2 CrCl 3 + 6 H 2 O + 3 C1 2 , for the acid is oxidizable, and at the same time promotes the for- mation of the salt of the lower and basic oxide Cr 2 O 3 . Similar considerations apply to the oxides of manganese. The CHEMICAL CHANGES CLASSIFIED 241 oxide MnO 3 loses oxygen when heated alone, but manganates corresponding to it are formed when any other manganese oxide is heated in air with alkali; e.g.: 2 MnO 2 -f O 2 + 2 Na 2 CO 3 = 2 Na 2 MnO 4 + 2 CO 2 . Thus a deep-green mass of sodium manganate is produced when a trace of a manganese compound is heated strongly with fusion mixture. Without alkali, however, no manganate appears; so when manganese is tested for by means of a borax bead, the bead is not green, but amethyst in the oxidizing and colourless in the reducing flame, these colours being due to manganic and manganous borates, derived from Mn 2 O 3 and MnO respectively. The processes of metallurgy, i.e. the winning of metals from their ores, afford numerous examples of reduction in the dry way; and these are often imitated on a small scale by blowpipe and other laboratory reactions. The following reactions are typical of reduction in the dry way: K 2 CO 3 + 2C = 2K + 3CO. Fe 2 O 3 -f 3C = 2Fe + 3CO. 4KOH + 3Fe = 4K + Fe 3 O 4 + 2H,, Sb 2 S 3 + 3Fe = 2Sb + 3FeS. Na 2 SiF 6 + 4 Na = Si + 6 NaF. CoO + H 2 = Co + H 2 O. 2 AgCl 4- 2 Hg = 2 Ag -f Hg 2 Cl 2 . Cr a O 3 + 2Al = 2Cr + Al 2 O 3 . 2 PbS + 3 O 2 = 2 PbO + 2 SO 2 A 2PbO + PbS = 3Pb + SO 2 . PbSO 4 + PbS = 2 Pb + 2 SO 2 . J The last three reactions are remarkable. They concern the metallurgy of lead by what is paradoxically called the "air- reduction process". A similar reaction occurs in the metallurgy of copper, between cuprous oxide and sulphide, thus: 2Cu 2 O + Cu 2 S = 6Cu + SO 2 . Oxidation and reduction in solution are frequent laboratory operations. The reactions of nitric acid have already been studied. Sodium peroxide, or hydrogen peroxide in presence of alkali, is frequently employed, e.g. to oxidize precipitated chromic hydroxide to chromate: 2 Cr(OH) 3 + 3 Na 2 O 2 = 2 Na 2 CrO 4 + 2 NaOH + 2 H 2 O. (D60) 17 242 CHEMICAL THEORY Similarly sulphide is oxidized to sulphate: Na 2 S + 4 Na 2 O 2 + 4 H 2 O = Na 2 SO 4 + 8 NaOH. The oxidation of sodium sulphide, or rather hydrogen sulphide, by nitric acid, produces, however, not sulphuric acid in the first instance, but free sulphur, thus: 2HN0 3 = although by the action of hot, concentrated nitric acid sulphur is gradually converted into sulphuric acid. This difference between the manner of oxidation of a sulphide in acid and alkaline solution is a further illustration of the influence of conditions on the course of a reaction. Chloric acid, derived from a mixture of potassium chlorate and hydrochloric acid, is sometimes employed in qualitative analysis to oxidize finely-divided sulphur to sulphuric acid. When a chlorate is heated with hydrochloric acid a mixture of chlorine and chlorine dioxide, Davy's " euchlorine ", is evolved, the reaction being most simply represented, thus: 2 KC1O 3 + 4 HC1 = 2 KC1 + 2 C1O 2 + C1 2 + 2 H 2 O. It does not follow, however, that this proportion between chlorine dioxide and chlorine is maintained, for the former may in turn oxidize hydrochloric acid, thus: 2C1O 2 + 8HC1 = 4H 2 O + 5C1 2 , so that the proportion of chlorine is increased; but that is im- material, for the mixture of gases evolved represents quantitatively the oxygen content of the chlorate, so that an equivalent amount of iodine would always be liberated from hydriodic acid, thus: 2KC1O 3 + 12HI = 2KC1 + 6H 2 O + 6I 2 . A very sensitive reaction is that between iodate and iodide in presence of dilute acid, thus: KIO 3 + 5 KI + 3 H 2 SO 4 = 3 K 2 SO 4 + 3 H 2 O + 3 I 2 ; indeed the acidity of a solution may be estimated by the iodine liberated when it is mixed with excess of iodate and iodide, which do not interact in neutral solution. Among the most important oxidizing reactions in solution are CHEMICAL CHANGES CLASSIFIED 243 those of permanganate and dichromate, so commonly employed in volumetric analysis. The reactions of permanganate afford a most interesting example of the principle already noticed from time to time that differing conditions determine which of several possible reactions shall take place. The following scheme represents in terms of dxides the reduction of permanganate in stages, with corresponding colours of the reduc- tion products: Mn 2 O 7 -* 2MnO 3 + O -* 2MnO 2 + 3O - Crimson. Green. Brown. Pale Pink. Permanganate. Manganate. Hydrated manganese dioxide, Manganous or manganous acid. salt. These stages may be readily observed by the use of sulphite as reducing agent. Thus, if a little neutral or alkaline sulphite solution is added to a dilute solution of permanganate, made some- what alkaline, the colour becomes deep green, owing to the forma- tion of manganate. If more sulphite is added, and especially if the solution is not too alkaline, and is warmed, the green solution gives place to a brown and turbid liquid containing hydrated manganese dioxide. Finally, if sulphite is added to an acidified permanganate solution, the colour is instantly discharged, manganous salt being produced. Thus an alkaline solution with little reducing agent pro- motes the formation of manganate; a nearly neutral solution with more reducing agent causes the manganese dioxide stage to be reached, whilst the presence of acid is the best condition for complete reduction to manganous salt. All this is what might be expected; MnO 3 is an acidic oxide, most likely to be permanent as a salt in presence of alkali; MnO is a basic oxide, and its salts are likely to be formed in presence of acid; whilst MnO 2 is neither strongly basic nor acidic, and is there- fore likely to result when the solution is nearly neutral. It must not, however, be assumed that reduction cannot proceed beyond the stage of Mn0 2 in absence of acid, for this is not true. It is possible to create conditions of stability for compounds corresponding to MnO in presence of alkali. Thus, if a few drops of permanganate solution are added to a solution of alkali sulphide containing free alkali, the brown precipitate of hydrated MnO 2 first observed quickly becomes paler as it mixes with the solution, being converted by excess of reducing agent into the less-soluble MnS. 244 CHEMICAL THEORY Or, if the same permanganate solution is added to excess of alkali sulphite containing much ammonium chloride, the brown precipitate, when heated with the liquid, completely dissolves, yielding a colour- less solution, because, as the student should know, manganous solu- tion can remain unprecipitated by alkali in presence of much ammonium chloride. Occasionally a compound intermediate in oxygen content be- tween two other compounds may behave either as an oxidizing or a reducing agent, according to circumstances. Instances of such compounds are nitrous acid and the aldehydes. This double func- tion of nitrous acid may be expressed in terms of oxides, thus: -o +20 2 NO *- N 2 3 N 2 6 ; that is to say, nitrous acid may be oxidized to nitric acid or reduced to nitric oxide. If a solution of nitrite is carefully added to an acidified solution of permanganate the latter is decolourized, whilst the nitrite is oxidized to nitrate without evolution of gas, thus: 2 KMnO 4 + 5 HNO 2 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 3 H 2 O + 5 HNO 3 . If, however, nitrite solution is mixed with hydriodic acid or acidified potassium iodide, the following reaction occurs: 2HNO 2 + 2HI = 2H 2 O- the hydriodic acid being oxidized with liberation of iodine, whilst nitric oxide gas escapes as the reduction product of nitrous acid. Nitrous acid can also oxidize ammonia, nitrogen resulting, both as the reduction product of the former and the oxidation product of the latter; so ammonium nitrite decomposes, thus: NH 4 NO 2 = N 2 + 2H 2 O. Aldehydes are intermediate as regards oxygen content between alcohols and carboxylic acids, e.g. : CH 3 -CH 2 OH CH 3 COH CH 3 -COOH Ethyl alcohol. Acetaldehyde. Acetic acid. So an aldehyde may be reduced to an alcohol, thus behaving as an oxidizing agent, or oxidized to an acid, thus behaving as a reducing agent. Nascent hydrogen reduces an aldehyde thus: CH 3 -COH + 2H = CH 2 -CH 2 OH; CHEMICAL CHANGES CLASSIFIED 245 silver oxide oxidizes an aldehyde thus: CH 3 .COH + Ag 2 O = CH 3 COOH + 2Ag. Hydrogen peroxide may perform both oxidizing and reducing functions, though in both cases it is reduced to water, so that when it reduces a compound, oxygen gas is evolved. The following equations will make this plain: H 2 O 2 + X = XO + H 2 O, oxidizing function. H 2 O 2 + XO = X + H 2 O + O 2 , reducing function. The behaviour of hydrogen peroxide towards manganese com- pounds furnishes an admirable example of the influence of acidity or alkalinity on the course of a reaction. If H 2 O 2 is added to manganese hydroxide in presence of alkali, oxidation occurs, thus: Mn(OH) 2 + H 2 O 2 = MnO 2 + 2H 2 O; but if the liquid containing the precipitated MnO 2 is now acidified, reduction to manganous salt is brought about by H 2 O 2 , thus: MnO 2 + H 2 SO 4 + H 2 O 2 = MnSO 4 -f 2 H 2 O -f O 2 . Permanganate is similarly reduced in acid solution, thus: 2 KMnO 4 + 3 H 2 SO 4 + 5 H 2 O 2 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O -f 5 O 2 ; reduction of a compound by H 2 O 2 does not, however, necessarily take place in presence of acid, for chromic acid is oxidized in acid solution to the deep blue perchromic acid. But in general, since the lower oxides of a metal are basic, they are likely to be produced by reduction in presence of acids, whilst the higher oxides, if acidic, are more likely to result by oxidation in presence of alkali. The above reaction between permanganate and hydrogen per- oxide is a case of the mutual reduction of two oxidizing agents, each of which contributes one atom of oxygen towards a molecule of this gas. Similar examples are furnished by the reactions be- tween H 2 O 2 and silver oxide and ozone respectively: H 2 2 + Ag 2 = 2 Ag + H 2 + 2 . H 2 O 2 + O 3 = H 2 O + O 2 + O 2 . A reaction which may be regarded as the converse of the above is the partition of molecular oxygen between water and an oxidizable substance, as in the following reaction, which is believed to be part 246 CHEMICAL THEORY of what takes place in the process of solution of gold by potassium cyanide in presence of air: 2 Au + H 2 O + O 2 = Au 2 O + H-jO* There are a few compounds, intermediate in oxygen content, which, under suitable conditions, undergo self -oxidation and reduc- tion, producing compounds poorer or richer in oxygen respectively than themselves. Thus a hypochlorite solution, when boiled, passes into chloride and chlorate, thus: 3KOC1 = 2KC1 + KC1O 3 ; whilst chlorate, when heated suitably, yields chloride and per- chlorate, thus: 4KC1O 3 = KC1 + 3KC1O 4 . Phosphite and hypophosphite pass by heating into phosphate and phosphine, thus: 4 Na 2 HPO 3 = 2 Na 3 PO 4 + Na 2 HPO 4 + PH 3 . 2 NaH 2 PO 2 = Na 2 HPO 4 + PH 3 ; and sulphite and thiosulphate behave thus when heated: 4Na 2 SO 3 = 3Na 2 SO 4 + Na 2 S. 4Na 2 S 2 O 3 = In all these cases the compounds pass into others more stable under the given conditions. SUMMARY HYDROLYSIS. Hydrolysis is the chemical decomposition of a compound by water. OXIDATION AND REDUCTION. Oxidation is the addition of oxygen to an element or compound; reduction is its removal. CHAPTER XIII THE COLLOIDAL STATE When a finely -divided solid is mixed with water or other solvent, either it may dissolve completely or some or all of it may remain undissolved. These two conditions are easily distin- guished. If the solid dissolves completely the resulting liquid, whether coloured or not, is clear or transparent. If the solid does not dissolve completely, the liquid when shaken will appear turbid or opaque, and if the mixture is allowed to stand undisturbed the solid in suspension will in time settle, leaving the supernatant liquid clear. The distinction between a substance in solution and one in suspension appears fundamental; for a liquid containing suspended matter may be filtered to be made clear, but it is needless to filter a solution. Yet suspended solids differ in the fineness of their sub- division, and in the ease with which they are removed by subsidence or filtration. Sand, for example, will settle more quickly than pre- cipitated chalk, and both of these can be filtered more easily than precipitated barium sulphate or calcium oxalate. And the student is familiar with substances even more difficult to filter than the last named. The sulphur which separates when hydrogen sulphide gas is passed through an oxidizing solution, or when acid is added to a polysulphide, cannot be removed completely by means of ordinary filter-paper, the pores of which are evidently too large to retain the minute particles of which the precipitate consists. Indeed, although the solution may become practically trans- parent, yet the presence of suspended sulphur is revealed by a slight opalescence. Precipitated silver chloride presents a similar pheno- menon; whilst copper sulphide, imperfectly precipitated from cold solution, may yield a brown filtrate, which, although transparent, contains the sulphide in a very fine state of subdivision. It appears, therefore, that a liquid may contain suspended matter so finely 247 248 CHEMICAL THEORY divided as not to produce opacity or to be removable by subsidence or ordinary filtration. It is pertinent to ask, therefore, how such a suspension differs from a true solution. Meanwhile the subject may be approached from a different point of view. When an aqueous- solution of sodium silicate, or soluble glass, is acidified with dilute hydrochloric acid, silicic acid is liberated, and, if the solution is in a concentrated state, will separate from it in the form of "gelatinous silica". If, however, the solution is suffi- ciently dilute, there is no precipitate, the liquid remaining clear. It might be supposed that the difference between these two conditions depended simply on the amount of water present, there being enough water to hold the silicic acid in solution in the one case, but not in the other. If this were so, the gelatinous silica separated in the former case would be in equilibrium with 1 a saturated solution of the same substance. This, however, is not so; the phenomenon here exhibited is indeed quite different from an ordinary case of precipi- tation, as will appear in the sequel. This and kindred phenomena were first investigated by Graham in 1849 in connection with experiments on liquid diffusion. Graham found that the rates of diffusion into pure water of different sub- stances in aqueous solution were various, and that simple salts and acids passed rapidly through an animal membrane or parchment paper, whilst complex substances like gelatine or glue in aqueous solution did not penetrate these membranes. These latter substances Graham called colloids (/coXXa, glue), whilst acids and salts, being crystallizable or related to crystallizable substances, he called crystal- loids. So crystalloids and colloids can be separated from one another by aqueous diffusion through a parchment or other suitable membrane, fixed on a frame like a drum and floating on water, crystalloids passing through the membrane into the external water, whilst colloids remain behind in the drum. The process is called dialysis, because it involves separation of one substance from another by passing it through a semi-permeable membrane, through which crystalloids but not colloids in solution can pass. A diluted acidified solution of sodium silicate may be submitted to dialysis. The sodium chloride formed in the reaction, Na 2 SiO 3 + 2 HC1 = H 2 SiO 3 + 2 NaCl, and excess of hydrochloric acid pass through the membrane of the THE COLLOIDAL STATE 249 dialyser, leaving the silicic acid behind in pure aqueous solution. Such a solution may be concentrated by boiling to a strength of about 14 per cent. After this it changes to a jelly, similar to that obtained by acidifying a concentrated solution of soluble glass. Not only does concentration cause dialysed silicic acid to coagulate, but a trace of hydrochloric acid, or some simple salt, acting catalytically, produces the same effect. It thus appears that there are two forms of aqueous colloidal silicic acid; the clear form, which seems to be a solution, and the gelatinous form, which evidently is not. These two forms, in which colloids in general may occur, are called respectively the hydrosol and hydrogel, or simply sol and gel. Besides silicic acid and organic substances such as gums and resins, glue and gelatine, various inorganic substances occur or can be obtained in the colloidal state. Graham prepared the sols of ferric, chromic, and aluminium hydroxides by dialysis; the sols of arsenious and antimonious sulphides may be prepared by boiling arsenious oxide and tartar emetic respectively with water, and adding hydrogen sulphide to the solutions: the liquids become yellow and red respectively because of the formation of the sulphides in the sol condition. A drop of hydrochloric acid added to either solution precipitates the yellow arsenious or orange-red antimonious sulphide, the sulphides; thus assuming the gel condition. It will now be understood that the above-mentioned brown liquid obtained in precipitating copper sulphide contained the sol of this sulphide, whilst the opalescent liquid containing sulphur held this element also in the colloidal state. The sols of certain metals are interesting, and often display remarkable colours. Thus gold and silver may be separated from their salts by hydrazine, formaldehyde, &c. Faraday produced blue, violet, and rose-coloured liquids by reducing gold chloride by means of an ethereal solution of phosphorus floating on the surface of the solution; and Bredig obtained sols of gold, silver, platinum, &c., by an electric discharge through water between poles of the metal. It is easy to understand that colloids have been regarded by chemists with much interest from the time of their discovery to the present day. They are of practical importance because they embrace many common non-crystallizable organic substances such as gum, 250 CHEMICAL THEORY resin, glue, starch paste, egg-albumin, casein, and gelatine; but they are particularly interesting from the physico-chemical stand- point because they present a fresh phase of the great subject of molecular physics. Indeed, a transparent dialysed liquid, consisting of silicic acid and water for example, which on the addition of a suitable catalyst becomes a jelly that can be inverted without flow- ing, presents to the scientific mind a subject for investigation full of an interest that can scarcely be surpassed. What is this transparent liquid ? Is it a solution like a mixture of sodium chloride and water? If it is, why does the silicic acid remain in the dialyser whilst the sodium chloride passes through it? It has already been suggested that the process of dialysis is a kind of filtration, i.e. that the silicic acid molecules are too large to pass through the pores of the parchment paper. But filtration is applied to something in suspension. Is the silicic acid in this apparently clear liquid really in suspension? An answer to this question as regards colloids in general has been gained by the use of the ultra - microscope invented by Siedentopf and Zsigmondy in 1903. The principle underlying the use of this instrument is that illustrated by the vision of the " mote in the sunbeam ". It is well known that the moving dust of the air, which cannot ordinarily be seen, is made visible in a beam of sunlight entering a darkened room through a chink in a shutter. At the same time the track of the beam itself is clearly outlined; but if the air is free from dust the sunbeam dis- appears. This effect, studied by Professor Tyndall, is applicable also to liquids, and will reveal the presence of suspended particles within them in the same way that it shows aerial dust. Moreover, the lesson to be learned is that very intense and localized light, by increasing the intensity of reflection, greatly en- hances our powers of vision. And if the dust of the air, otherwise quite invisible, thus becomes apparent to the naked eye, particles too small for microscopic vision under ordinary illumination may be seen under illumination analogous to that of the sunbeam. This is the principle of the ultra-microscope, in which a beam of sunlight, or from the electric arc, passes through a slit horizontally, or is focused into a liquid which is examined by the microscope verti- cally. Any light which enters the microscope must then have been reflected from the surface of particles suspended in the liquid. THE COLLOIDAL STATE 251 Particles having a diameter only one -hundredth that of the smallest particles visible under ordinary illumination can then be seen as spots of light like planets in the darkness. And so colloidal liquids have been seen to be suspensions, and the size of the suspended particles has been estimated by counting the number of them in a volume of the liquid containing a known weight of material. Thus the particles of platinum, gold, and silver seen in colloidal suspensions of these metals have been discovered to have diameters ranging from 2xlO~ 4 to 6xlO~* mm. The smallest particles detectable by this method, when illuminated by bright sunlight, have a diameter of 4X10" 6 mm., whilst the individual molecules of substances like chloroform and alcohol have diameters of 0-4xlO~* to O-SxlO" 6 mm., and of hydrogen 0-1 xlO- 6 mm. Thus the particles of colloidal metals in aqueous suspension have diameters about a thousand times as great as those of mole- cules which form mixtures with water regarded as true solutions; whilst the smallest particles that can be rendered visible have only about ten times the diameter of gaseous and other simple molecules. These metallic suspended particles are not, however, molecules, but rather minute fragments of solid metal; since molecules of solid metals, consisting of definite aggregates of atoms, can scarcely be said to exist. Thus, they differ from silicic acid and complex organic substances which are known to consist of very large mole- cules. The molecule of egg-albumin is estimated to have a mole- cular weight of 17,000, and the molecules of the enzymes emulsin and invertin have molecular weights of about 45,000 and 54,000 respectively, with molecular diameters of about 6 x 10 ~ 6 mm. Such molecules can be seen by the ultra-microscope, but there seems no hope that the simpler inorganic molecules will ever be revealed to the eye of man, although they lie but a little way below his range of vision aided by this powerful instrument. Finally, although colloids cannot ordinarily be filtered, the fact that parchment paper retains them suggests that special methods of filtration might effect their separation as parchment paper does. Special filters have in fact been prepared by treating ordinary filter -paper with collodion or gelatine, which have pores varying in diameter between 930 x 10~ 6 mm. and 21 x 10~ 6 mm. By means of these filters various colloids have been differentiated and 252 CHEMICAL THEORY classified according to the sizes of their particles, with the following results: Suspensions of non-colloids. Colloidal platinum. Colloidal ferric hydroxide. Colloidal arsenious sulphide. Colloidal gold. 1 per cent gelatine. Colloidal silicic acid. Litmus. Dextrin. Solutions of crystalloids. Thus it is seen that colloids afford a gradation between what are commonly regarded as suspensions and solutions. SUMMARY DIALYSIS is the separation of substances in solution by the use of a semi-permeable membrarie, through which crystalloids in solution will pass but not colloids. COLLOIDS, e.g. silicic acid, can exist in two states, the hydrosol (or sol) state, and the hydrogel (or gel) state. Sols are converted into gels by catalysis. By means of the ultra-microscope colloidal liquids have been seen to be suspensions. Colloids have been separated by the use of special filters, and a gradation has been established between colloids in suspension and crystalloids in solution. CHAPTER XIV EQUATION-BUILDING The student of chemistry becomes familiar early with chemical equations. It may be that before he mastered the rudiments of the atomic and molecular theories he was taught to employ the equation as a brief and pointed way of stating what happens in chemistry; and that without fully understanding their significance he has committed to memory a number of equations representing the reactions that occur in the varied preparations of his elementary course. Consequently he has been in danger of magnifying the equation unduly and regarding it as a kind of talisman by means of which natural processes may be foretold or brought to pass. He may have supposed that by manipulating the formulae of certain reacting substances, and of other substances that might result from their interaction, a chemical change may be successfully expressed without any practical experience of that change. This, in fact, is how beginners often behave in answering questions. Their chief concern is to make the equation balance, in faithfulness to the principle of the indestructibility of matter, supposing that the exigencies of chemical science are completely satisfied if nothing is lost by the way. The student must learn that chemistry is not a branch of mathematics, that chemical equations are not to be solved like algebraic equations. To employ a chemical equation for a reaction until that reaction is understood qualitatively and quantitatively, so far as the distribution of matter it expresses is concerned, is unscientific and vain. What is the use of a chemical equation? it may be asked. Such a question is best answered by considering what chemistry would be without equations. The science might still exist, but it would be an exceedingly clumsy science, and would probably be in a much more rudimentary state than it is at present. Nature 253 254 CHEMICAL THEORY has not ordained the chemical equation; it is a human invention; and man might still be mixing or heating things together and watching for results if he had never invented a means of expressing briefly the results of his discoveries. What language is to thought, so roughly the equation is to chemistry; and just as the educated man chooses language to express his thoughts, and avoids the danger of allowing his language to outstrip his thoughts, so the chemist uses equations to express discovered chemical facts, and avoids allowing his equations to outstrip his facts. Chemical equations, then, are chemical notation, that is, a means of noting chemical facts in a convenient form. Facts first; equation afterwards. Let the student remember this, and he will at least be in the way of learning chemistry properly. Nevertheless, to express chemical reactions by equations, not learned by rote, but developed intelligently from well-understood principles, is an art which should be learned by every student of chemistry. It is the present purpose to show how even the most complex equations met with in inorganic chemistry may be built up when the underlying principles of the reactions they express are understood. Consider the simplest possible reaction; the preparation of hydrogen by the interaction of zinc and dilute sulphuric acid. If the student knows the symbol Zn and the formula H 2 SO 4 , and knows also that zinc displaces the hydrogen from the sulphuric acid, he may still go wrong with his equation. For how shall he decide whether: 2 Zn 4- H 2 SO 4 = Zn 2 SO 4 + 2H, or Zn + H 2 SO 4 = ZnSO 4 + 2H, or Zn + 2 H 2 SO 4 = Zn(SO 4 ) 2 + 4 H, or some other relation between the zinc and the displaced hydrogen is the right one? He may be told what the right relation is, and therefore what the proper equation should be, but he should also be told why. Now, whilst the equivalent weight of zinc is 32 65, the atomic weight of this metal, as shown by the method of specific heats and other methods is 65-3. Therefore, an atom of zinc displaces 2 atoms of hydrogen, and the equation becomes: Zn 4- H 2 SO 4 = ZnSO 4 4- 2 H. There remains another question, however; that is, whether it is EQUATION-BUILDING 255 quite proper to represent the atom of zinc or hydrogen as remaining single in the elementary state. The answer is that zinc is known so to exist, but that hydrogen gas consists of molecules H 2 . There- fore, amended, the equation finally becomes: Zn + H 2 SO 4 = ZnS0 4 -f H* When a student has his equation he is inclined to ask whether it explains everything. Then he has to be told that it explains nothing. A chemical equation does not explain a chemical reaction; it expresses it, with certain limitations. It is plain that the equation gives but a limited expression to the reaction between zinc and sulphuric acid; for it does not express the fact that the acid must be considerably diluted with water before any hydrogen can be obtained. The student will be aware that there are other circumstances of this reaction of which this simple equation gives no account. Nevertheless, it does express, qualitatively and quantitatively, a reaction between the metal and the acid. If concentrated instead of dilute acid is added to zinc, there is little if any evolution of hydrogen; but when the mixture is heated a vigorous reaction takes place, and much sulphur dioxide gas is evolved. If this quite different reaction is to be expressed by an equation, it must first be understood. Sulphur dioxide necessarily comes from the sulphuric acid, whence it is derived by reduction, which may be represented thus: H 2 SO 4 - O = H 2 O + SO 2 , or H 2 SO 4 = H 2 O + SO 2 + O. It is pertinent to ask how this reduction is effected. Now, since hydrogen is a reducing agent, a facile theory of the process is the following: Zn + H 2 SO 4 = ZnSO 4 + 2 H. 2 H + H 2 SO 4 = 2 H 2 O + SO 2 . The student may now add these two equations together, in this case treating them as if they were algebra; then, since hydrogen is eliminated, the final equation becomes: Zn + 2 H 2 SO 4 = ZnSO 4 + 2 H 2 O + SO 2 . Now this is the equation for the reaction in question whether 256 CHEMICAL THEORY hydrogen is liberated, and in the nascent state acts as a reducing agent, or not. But it is certainly open to question whether hydrogen is liberated at all in this reaction; and there is no need to assume that it is. Zinc itself is a reducing agent, so why should it not react directly? Now, whilst zinc oxide cannot appear in presence of excess of acid, it is symbolic of the state of oxidation assumed by the zinc; so that if the reaction is to be represented without assuming the agency of liberated hydrogen, the following device may be adopted: Zn + H 2 SO 4 = ZnO + H 2 O + SO 2 ZnO + H 2 SO 4 = ZnSQ 4 + H 2 O and adding Zn + 2 H 2 SO 4 = ZnSO 4 + 2 H 2 O + SO 2 . The same result is shown as above, and if it is not assumed that ZnO is actually separated, this method of arriving at the result is freer from assumption than the former one. Equations which often puzzle beginners are those representing the interaction of hydrogen sulphide with metallic salt solutions. The reaction CuSO 4 + H 2 S = CuS + H 2 SO 4 is simple enough, because the molecules are matched; but how is SbCl 3 to be represented as reacting with H 2 S? The answer is simple if it is remembered that the reaction is one of double decomposition; antimony chloride -f hydrogen sulphide give anti- mony sulphide + hydrogen chloride. That is, hydrogen and chlorine atoms must be in equal numbers on the left-hand side of the equa- tion to produce hydrogen chloride molecules on the right. This can result only if the equation is written: 2 SbCl 3 + 3 H 2 S = Sb 2 S 3 + 6 HC1; thus not only showing six complete molecules of hydrogen chloride, but also satisfying the tervalency of antimony. The reaction 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO may now be considered. Since copper does not displace hydrogen from dilute acid, the nascent hydrogen theory will not be assumed to account for the reduction of the nitric acid. The copper will be represented as being directly oxidized by the nitric acid, but as appearing as nitrate instead of oxide, because of the excess of acid present. It is EQUATION-BUILDING 257 important, however, to realize how the nitric acid is reduced to the state of nitric oxide, and this reduction may be represented thus: 2 HNO 3 = H 2 O + N 2 O 5 = H 2 O + 2 NO -f 3 O. To remove these 3 atoms of oxygen from 2 molecules of nitric acid 3 atoms of copper are needed, thus: = CuO; but nitrate is formed rather than oxide, thus: 3CuO + 6HNO 3 = 3Cu(NO 3 ) 2 + 3H 2 O. And now, adding together these three equations, the following equation is obtained as representing the actual reaction that takes place: 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO. Thus, by this method of construction it becomes apparent why 3 atoms of copper and 8 molecules of nitric acid are required; and the student should now recognize that the equation has been built by a sound method based on an understanding of the scientific principles involved. He may ask himself how otherwise he would produce the equation if he did not remember it. Only by hap- hazard and guesswork could he make the attempt. In the chapter on the Classification of Chemical Changes the following equations are found: i. 4Zn + 10HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O, and ii. 4Zn + 10HNO 3 = 4 Zn(NO 3 ) 2 + 3 H 2 O + NH 4 NO 3 . They may be built as follows: i. 2 HN0 3 = H 2 + N 2 O + 4O 4Zn + 4O == 4ZnO 4 ZnQ + 8 HNO 3 = 4 Zn(NO 3 ) 2 + 4 H 2 O adding 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O -f N 2 O; or, using the nascent hydrogen theory, which is perhaps preferable here: 2HNO 3 = H 2 O + N 2 C 4Zn + 8HNO 3 = 4Zn(NO 3 ) 2 8H-h4O = 4H 2 O adding Zn + 10HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O ii. HNO 3 + 8 H = 3 H 2 O + NH 3 4Zn + 8HNO 3 = 4Zn(NO 3 ) 2 + 8H . NH 3 + HNO 3 = NH 4 NO 3 Adding 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 3 H 2 O + NH 4 NO 3 (D60) 258 CHEMICAL THEOKY The student will see that the method of procedure is to fix on the essential reduction product, and consider how it can be formed. Now, in the above example, a molecule of ammonia is formed from one of nitric acid by the action of 8 atoms of hydrogen. These are therefore produced from 8 other molecules of the acid. So the final equation is easily and surely obtained. Allied to these reactions are those concerned in the use of permanganate and dichromate in volumetric analysis. Consider the following equations representing the oxidation of ferrous iron: 2KMnO 4 4-10FeSO 4 + 8H 2 SO 4 '== K 2 SO 4 +2MnSO 4 + 5Fe 2 (SO 4 ) 3 +8H 2 O; K 2 Cr 2 O r + 6 FeSO 4 + 7 H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 3 Fe 2 (S0 4 ) 3 + 7 H 2 O. It may be observed in passing that the numbers of molecules of sulphuric acid required and of water formed correspond with the number of oxygen atoms in the oxidizing agent. But this circumstance does not really furnish a key to the reactions. This is supplied by considering the available oxygen of the oxidizing agent and the oxygen essential for oxidizing the iron. Expressed in terms of oxides, these are shown thus in the case of permanganate: 2 KMnO 4 = K 2 O + Mn 2 O 7 = K 2 O -f 2 MnO 4- 5 O. 2FeO + O = Fe 2 O 3 . Thus, since 2 molecules of permanganate in presence of acid yield 5 atoms of oxygen for oxidizing purposes, together with potassium and manganous salts, to which the above oxides correspond; and 2 atoms of iron require 1 atom of oxygen to oxidize them from the ferrous to the ferric state, it appears that 10 molecules of ferrous sulphate are oxidized by 2 molecules of permanganate in presence of sufficient acid, the quantity of which participating in the reaction is easily arrived at by adding together the following equations: 2KMnO 4 = K 2 O 4- 2 MnO 4- 5 O 10 FeSO 4 4- 5 O 4- 5 H 2 SO 4 = 5 Fe 2 (SO 4 ) 3 -f 5 H 2 O K 2 O + H 2 SO 4 = K 2 SO 4 4- H 2 O 2H 2 SO 4 = 2MnSQ 4 _ Adding 2KMnO 4 +10FeSO 4 +8H 2 8O 4 = K 2 SO 4 +2MnSO 4 +5Fe 2 (SO 4 ) 3 +8H 2 O. The equation for the dichromate reaction is similarly obtained, the chromium being reduced from the state of CrO 3 to Cr 2 O 3 , thus: 2 CrO 3 = Cr 2 O 3 4- 3 O; EQUATION-BUILDING 259 and the iron oxidized by the available oxygen as before: K 2 O 2 O 7 = K 2 O + O 2 O 3 + 3 O 6 FeSO 4 + 3 O -f 3 H 2 SO 4 = 3 Fe 2 (SO 4 ) 3 + 3 H 2 O K 2 O -f H 2 SO 4 = K 2 SO 4 + H 2 O O 2 O 3 + 3 H 2 SO 4 = Cr 2 (SO 4 ) 3 + 3 H 2 O Adding K 2 Cr 2 OH-6FeSO 4 +7H 2 SO 4 = K 2 SO 4 +Cr 2 (SO 4 ) 3 +3Fe 2 (SO 4 ) 3 +7H 2 O. Other equations for other reactions of permanganate and dichro- rnate are built up similarly, e.g. the following: 2 KMnO 4 -f 5 H 2 C 2 O 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O + 10 CO 2 2 KMnO 4 + 5 H 2 O 2 + 3 H 2 SO 4 = K 2 SO 4 -f 2 MnSO 4 + 8 H 2 O + 5 O 2 . K 2 Cr 2 O 7 + 3 H 2 S + 4 H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 7 H 2 O + 3 S. K 2 Cr 2 O 7 + 6 HI + 4 H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 -f 7 H 2 O + 3 1 2 . The reactions of potassium chlorate with sulphuric and hydro- chloric acids are interesting, and the building of equations to repre- sent them is instructive. When sulphuric acid is added to potassium chlorate chloric acid is liberated thus: 2 KC10 3 + H 2 S0 4 = K 2 S0 4 + 2 HC1O 3 ; but this acid is dehydrated when it is heated with sulphuric acid. The anhydride C1 2 O 5 , however, does not exist; and in place of it 2C1O 2 + O are produced. The oxygen, which would otherwise accompany the chlorine dioxide, does not appear as gas, but oxidizes some of chloric to perchloric acid, which is stable in presence of sulphuric acid; or, otherwise expressed, chloric acid undergoes self- oxidation and reduction thus: 3 HC10 3 = 2 C10 2 + H 2 + HC10 4 . Consequently, the equation representing the reaction between potassium chlorate and sulphuric acid may be built up thus: 2KC10 3 + H 2 S0 4 = K 2 S0 4 + 2HC10 3 2 HC1O 3 = H 2 O + 2 C1O 2 + O KC1O 3 + O = KC1Q 4 Adding 3 KC10 3 + H 2 S0 4 = K 2 SO 4 + KC1O 4 -f H 2 O + 2 C1O 2 . If potassium chlorate is heated with hydrochloric instead of (D60) 18 a 2 260 CHEMICAL THEORY sulphuric acid, chlorine accompanies the chlorine dioxide, since hydrochloric acid is oxidized in preference to chloric acid: 2 KC1O 3 + 2 HC1 = 2 KC1 + 2 HC1O 3 2 HC1O 3 = H 2 O + 2 C1O 2 + O 2HC1 + = H 2 + C1 2 _ Adding 2 KC1O 3 + 4 HC1 = 2 KC1 + 2 H 2 O + 2 C1O 2 + C1 2 . The interaction of sodium hydrogen sulphite and sodium iodate in aqueous solution, by which iodine is liberated quantitatively from the iodate, is a reaction the equation for which may be built up from first principles, thus: 2NaI0 3 = Na 2 + I 2 5 I 2 O 6 + 5NaHSO 3 = 5NaHSO 4 + I 2 2NaHSO 4 + Na 2 O = 2Na 2 SO 4 + H 2 O _ Adding 2 NaIO 3 + 5 NaHSO 3 = 2 Na 2 SO 4 + 3 NaHSO 4 + H 2 O + I 2 equation for th built up as follows: 3 3 = 24 4 2 2 . The equation for the cyanide process for gold extraction may be built u fll 2 Au + O 2 + H 2 O = Au 2 O + H 2 O 2 2 Au + H 2 O 2 = Au 2 O + H 2 O 2 Au 2 O + 8 KCN + 2 H 2 O = 4 KAu(CN) 2 + 4 KOH Adding 4Au-f 8KCN + 2 H 2 O-f O 4 KAu(CN) 2 + KOH. The parts of this reaction are remarkable; for whilst the first two equations represent what cannot possibly take place alone, since gold is not susceptible of atmospheric oxidation, they likewise show the fact of the intermediate formation of hydrogen peroxide through the partition of the oxygen molecule between gold and water, which is not shown in the final equation. The preparation of phosphine by heating white phosphorus with sodium hydroxide solution is a reaction which is represented by the following equation: = 3 NaH 2 PO 2 + PH 3 , the product in solution being sodium hypophosphite. In endeavouring to build this equation the first question that occurs is: Where does the hydrogen come from to produce the phosphine ? The answer is that it comes from the sodium hydroxide, and the simplest representation of this fact is: NaOH + P = NaOP + H EQUATION-BUILDING 261 NaOP is not sodium hypophosphite, however; but if the elements of water are added the formula for this salt, NaH 2 P0 2 , is obtained. A further stage, therefore, is: NaOH + H 2 O + P = NaH 2 PO 2 + H. If now this equation is multiplied by three, and an atom of phos- phorus is added to give PH 3 , the above equation results thus: 3 NaOH + 3 H 2 O + P 4 = 3 NaH 2 PO 2 + PHs. Phosphonium iodide, PH 4 I, is prepared by dropping water on to an intimate mixture of finely-divided phosphorus and iodine in an inert atmosphere; and the following equation, representing the reaction, is one of the most difficult in inorganic chemistry: I + 16H 2 = 5PHJ + 4H 3 PO 4 . It presents two problems: First, to understand why these products result; and second, to discover the reason for the remarkable mole- cular proportions exhibited by the equation. There is apparently some connection between this reaction and the foregoing, though water alone seems to perform the function of the sodium hydroxide solution above. Then if hypophosphorous acid itself is formed instead of its sodium salt, it is unstable, and yields phosphoric acid and phosphine. Hydriodic acid is required, however, to combine with the phosphine, and this may be produced by the interaction of iodine and phosphorus to produce the iodide which is subsequently hydrolyzed, yielding hydriodic and phos- phoric acids. In this way the reaction may be accounted for, and the equation built as follows: 12H 2 O = 6 H 3 P0 2 + 2 PH 3 6 H 3 PO 2 = 3 H 3 PO 4 + 3 PH 3 P + 5I = PI 5 PI 6 + 4H 2 O = H 3 PO 4 5PH 3 + 5HI = 5PH 4 I Adding 9 P + 5 1 + 16 H 2 O = 4 H 3 PO 4 + 5 PH 4 I. This is a rather complicated theory to account for the formation of phosphonium iodide, but it may stand in default of a better; it has the merit of accounting for a complicated equation. 262 CHEMICAL THEORY These examples are sufficient to illustrate the art of equation- building. Many other examples will occur to the student; and he will discover by practice that any reaction of which he has an in- telligent knowledge can be expressed by an equation constructed by the application of the corresponding chemical principles. INDEX Absorptiometer, 145. Acidic hydroxides, 223. Acidic oxides, 178, 183. Acids and metals, interaction of, 230. Active mass, 204. Adams and Le Verrier, 80. Affinity, chemical, 204. Affinity, units of, 57. Air, liquefaction of, 110. Alcohols, melting-points of, 125. Allotropy, definition of, 137, 142. Allotropy, table of, 141. Ammonia, solubility of, 147. Ammonium chloride, dissociation of, 217. Ammonium hydroxide, 223. Amphoteric oxides, 183. Andrews, 100. Anhydrides, mixed, 178, 184. Anion, 163. Anode, 163. Arrhenius, 166. Atmolysis, 95. Atomic symbols, Dal ton's, 11, 14. Atomic theory, 4, 8, 9, 17. Atomic volume curve, 75. Atomic weight, 18. Atomic weights, Dal ton's, 10. Atomic weights, determination of, 24, 25, 53, 81. Atomicity, 15, 17. Atoms, 2, 13, 17. Avogadro, 14, 15, 16. Avogadro's theory, 14, 15, 16, 17. Avogadro's theory, method of, 25. Bacon, F., 7. Barium peroxide, dissociation of, 219. Bases, acids, and salts, note on, 180. Basic oxides, 178, 180, 197. Beckmann apparatus, 43. Berthottet, 5. Berzelius, 11, 16, 20, 33, 65. Black, 4. Boiling-points of liquids, 118. Bonds, 57. Bound energy, 205. Boyle, 3, 4, 7. Boyle's law, 89. Boyle's law, deviation from, 97, 98. Cailletet, 105. Calcium carbonate, dissociation of, 218. Cannizzaro, 16. Carbon, atomic weight of, 37. Carbon dioxide, liquid, 109. Carbon dioxide, solubility of, 147. Carbonates, 190. Carbonic acid, 224. Carre's freezing- machine, 108. Catalysis, 210. Catalysis, theory of, 213. Catalyst, 210, 214. Catalyst, negative, 212. Catalytic action, 210. Cathion, 163. Cathode, 163. Cavendish, 4. Chancourtois, 69. Charles's law, 89. Chemical affinity, 204. Chemical change, 198. Chemical change, rate and limits of, 209. Chemical changes classified, 215. Chemical combination, laws of, 16. Chemical composition and solubility, 159. Chemical compounds, types of, 178. Chemical displacement, 25, 32, 54. Chemical equilibrium, 201. Chemical properties, periodicity of, 76. Chemical reactions in solution, 172. Chemistry in space, 63. Chlorides, 190. Chlorides, double, 195. Chlorides, hydrolysis of, 229. Chlorine, liquid, 110. Claude process, 107. Cohesion, 88. Colloidal particles, size of, 251. Colloidal state, 247. Colloids, 252. Colloids and crystalloids, 248. Coloured ions, 172. Colours of salt solutions, 172. Combustion, heat of, 206. Complex salts, 193, 194, 195, 197. Composition of matter, 1. Constitutional formulas, 57, 58. Couper, 57. Critical phenomena, definitions of, 101. 203 264 INDEX Critical state, 100. Critical state, definitions of, 111. Crookes's tube, 65. Cryohydrates, 127, 128, 142. Cryoscopic method, 42. Crystal, definition of, 130, 142. Crystal systems, 133, 142. Crystalline and amorphous states, 126. Crystallization, 135. Crystallography, 142. Cry stallo- hydrates, dissociation of, 219. Crystals, 129. Crystals, study of classification of, 131. Cyanides, complex, 196. Dcdton, 6, 7, 8, 9, 10, 11, 12, 20, 91. Dalton and Henry, law of, 150, 161.- Dalton's law of partial pressures, 95. Davy, H., 65, 103. Definite proportions, law of, 5, 16. Deliquescence, 220. Densities of gases, 15. . Densities of liquids, 114. Dephlegmator, 154. Deviation of gases from Boyle's law, 97. Dewar vessel, 106. Dialysis, 248, 252. Diffusion experiments, 93, 94. Diffusion of gases, 91. Dilute solutions, properties of, 162. Dimorphism, 36. Dissociation, heat of, 207. Dissociation, thermal, 202, 203. Dissociation pressure, 203. Distillation, fractional, 153. Distillation under reduced pressure, 117. Distribution coefficient, 152. Dobereiner, 91. Dobereiner's triads, 68. Double bond, 62. Double salts, 193, 194, 197. Dulong and Petit, 25, 33. Dulong and Petit's law, 32, 33. Dumas, 20, 27. Dumas's vapour-density method, 28. Ebulliscopic method, 43. Efflorescence, 220. Effusion of gases, 94. Eka-aluminium, 80. Electrodes, 162. Electrolysis, 162, 163. Electrolysis, illustration of, 170. Electrolysis, laws of, 169, 176. Electrolyte, 162, 176. Electrolytic dissociation, 165, 176. Electrons, 65. Element, definition of, 3, 16. Elements, 2, 3. Elements, melting-points of, 124. Elements of the Ancients, 1. Enantiotropic change, 139. Endothermic change, 207. Endothermic compounds, 214. Energy, bound or latent, 205. Energy, free, 205. Energy, total, 205. Enzymes, influence of, 213. Equation building, 253. Equilibrium, chemical, 201. Equivalent and atomic weight standards, 19. Equivalent weight, 18, 53. Equivalent weights, determination of, 21. Equivalents, 7. Eutectic alloys, 129. Eutectic mixtures, 127, 142. Exothermic change, 207. Exothermic compounds, 214. Faraday, 103, 169. Fats, hydrolysis of, 229. Finely -divided metals, catalytic influence of, 212. Fixed proportions, law of, 5, 16. Formation, heat of, 207, 214. Fractional distillation, 153. Fractionating column, 154. Frankland, K, 55. Free energy, 205. Gallium, 80. Gas constants, 102. Gas equation, 90. Gas laws, 89. Gas laws, summary of, 111. Gaseous diffusion, law of, 92, Gaseous mixtures, 95. Gases, diffusion of, 91. Gases, liquefaction of, 101. Gases, molecular composition of, 47. Ammonia, 50. Carbon monoxide, 51. Carbonic anhydride, 48. Ethylene, 53. Hydrogen chloride, 48. Hydrogen sulphide, 49. Methane, 5*2. Nitric oxide, 50. Nitrous oxide, 49. Phosphine, 51. Sulphurous anhydride, 49. Water and steam, 48. Gases, solubility of, 145. Oay-Lussac, 11, 12, 16. Gay-Lussac's law, 11. Gel, 249. Graham, 92, 94. Graphic formulae, 57, 58. Halides, 186, 197. Heat, action of, on compounds, 216. Heat of combustion, 206. Heat of dissociation, 207. Heat of formation, 207, 214. Henry, law of, 149, 161. Hess, law of, 207. Higgins, 7. INDEX 265 Hoffmann, 27. Hoffmann's vapour-density method, 29. Holohedral and hemihedral forms, 133. Hydrated salts, 192, 197. Hydrated salts, melting-points of, 124. Hydrides, 178, 197. Hydrochloric acid, distillation of, 148. Hydrogel, 249, 252. Hydrogen and helium, liquefaction of, 107, 108. Hydrogen chloride, solubility of, 147. Hydrogen ions, catalytic influence of, 211. Hydrolysis, 227. Hydrolysis, definition of, 246. Hydrosol, 249, 252. Hydroxides, acidic, 223. Hydroxides, thermal decomposition of, 222. Ideal gas density, 26. Immiscible liquids, 152. Indicators, theory of, 174. Iodine, vapour pressures of, 121. lonization, 166. Ions, 163. Isodimorphism, 36. Isomorphism, illustration of, 36. Isomorphism, law of, 25, 35, 54. Isomorphism, method of, 34. Isothermals of carbon dioxide, 99. Joule-Thomson effect, 106. KekuU, 55. Ladenburg, 94. Landolt, 4. Landsberger apparatus, 44. Latent energy, 205. Lavoisier, 3, 4, 198, 201. Leduc, 26. Linde and Pfampson, 106. Lmde or Hampson's process, 107. Liquefaction of gases, 101. Liquefaction of gases, methods of, 111. Adiabatic expansion, 105. Cascade, 103. Self -intensive refrigeration, 105. Simple compression, 102. Liquefaction of gases, practical appli cations, 108. Liquid air, 107. Liquid density, 113. Liquids, 88. Liquids, mobile and viscous, 112. Liquids, molecular volumes of, 113. Liquids, properties of, 112. Liquids, solidification of, 122. Liquids, specific volumes of, 113. Liquids, viscosity of, 112. Litmus, 175. Marignac, 34. Mass, chemical action of, 204. Matter, composition of, 1. Membrane, semi-permeable, 252. Mendeldc/, 69, 80. Mercuric oxide, 198. Mercurous chloride, dissociation of, 218. Metal and acids, interaction of, 230. Meta-stable state, 123, 140, 155. Methyl orange, 175. Meyer, Lothar, 74. Meyer, Victor, 27. Meyer's, V., vapour-density method, 29. Mitscherlich, 25, 33, 34. Mixed anhydrides, 178, 184. Mixed liquids, distillation of, 153. Molecular association, 46. Molecular complexity, 45. Molecular compositions of gases, 47. Molecular compounds, 61. Molecular theory, 11. Molecular weights in solution, 39. Molecular weights of gases, 15. Molecules, 13, 17. Monge and Clouet, 102. Monotropic change, 140. Morley, 26. Multiple proportions, law of, 6, 16. Neutral oxides, 178, 197. Neutralization, 168. Newlands, J. A. K, 69* Newton, 7, 8. Nitric acid and metals, interaction of, 232. Nitrogen peroxide, dissociation of, 216. Northmore, 102. Octaves, law of, 69. Olszewski, 107. Onnes, K., 108. Osmotic pressure, 166. Oxidation, 238. Oxidation and reduction, 238. Oxidation and reduction, conditions of, 240. Oxidation and reduction, definition of, 246. Oxidation and reduction in solution, 241. Oxidation and reduction in the dry way, 240. Oxides, acidic, 178, 183. Oxides, basic, 178, 180, 197. Oxides, neutral, 178, 197. Oxides, saline, 178, 184, 197. Oxides, thermal decomposition of, 224. Oxides and hydroxides, 178. Oxidizing and reducing agents, 239. Oxy-salts, thermal decomposition of, 224. Ozone, density of, by diffusion, 94. Partial pressures, law of, 95. Partition coefficient, 152. Periodic law, 25, 67, 69, 87. Periodic law, method of, 37. Periodic law, objections to, 85. Periodic law, suggestiveness of, 82. Periodic law, uses of, 79, 87. Periodic series, 70. Periodic tables, 71, 72. INDEX Peroxides, 178, 185, 197. Phenol-phthalein, 175. Phosphoric acid, 223. Phosphorus pentachloride, dissociation of, 217. Physical properties, periodicity of, 74. Pictet, 105, 107. Polymerism, 138. Polymorphism, 36. Polymorphism and allotropy, 136. Poly-oxides, 178, 186. Precipitation, 236. Pressure on melting-point, effect of, 123. Prideaux, 31. Priestley, 4, 91, 198. Proust, 5. Raoult's law, 40, 54. Rayleigh, 26. Reactions, reversible, 200. Reciprocal proportions, law of, 7, 16. Recrystallization, 136. Regelation, 122. Reversible reactions, 200. Richter, 6, 7. Saline oxides, 178, 184, 197. Salts, reactions of, 228. Salts, slightly soluble, 160. Saturation, 155, 156. Scheele, 4. Scott, 39. Semi-permeable membrane, 166, 248, 252. Short and long periods, 73. Siedentopf, 250. Sodium carbonate, solubility of, 159. Sodium sulphate, solubility of, 159. Solids, 88. Solids, formation of, 120. Solids, melting-points of, 122. Sols, 249, 252. Solubilities, table of, 158. Solubility, 155. Solubility, coefficient of, 144. Solubility and chemical change, 234. Solubility curves, 158. Solubility of a gas, 144, 161. Solubility of a gas and temperature, 147. Solute, 143. Solution, process of, 157. Solution pressure, 171. Solutions, 143. Solutions of gases in liquids, 144. Solutions of Squids in liquids, 151. Solutions of solids in liquids, 154, 161. Solvent, 143. Specific heats, law of, 54. Stas, 20. Stereochemistry, 64. Sublimation, 121. Sulphates, 190. Sulphides, 188, 197. Sulphion, 163. Sulphur dioxide, liquid, 109. Sulphuric acid, 223. Sulphurous acid, 223. Superoxides, 178, 186. Supersaturation, 156. Suspensions and solutions, 252. Temperature and solubility, 157. Thermal decomposition, 220. Thermal decomposition of hydroxides, 222. Thermal decomposition of oxides, 220. Thermal decomposition of oxysalts, 224. Thermal dissociation, 202, 203, 214, 216. Thermochemical equations, 205. Thermochemistry, 204, 205. Thompson, 20. Total energy, 205. Transition temperature, 139. TyndaU, 250. Ultra-microscope, 250, 252. Unknown elements, prediction of, 79. Unsaturation, 156. Valency, criterion of, 64. Valency, definition of, 55, 66. Valency, nature of, 65. Valency, periodicity of, 77. Valency, units of, 57. Valency, variability of, 60. Valency or atomic value, 55. vant lloff, 64. Vapour-density methods, 27. Vapour pressure and boiling-point, 115. Vapour-pressure curve, 116. Volatility and chemical change, 234. Volume-atom theory, 16. Volumes, law of, 11, 17. Water, action of, on elements, and com- pounds, 226. Water, catalytic influence of, 211. WoUaston, 20. Wroblewski, 107. Z&iymondy, 250. v UNIVERSITY OF CALIFORNIA LIBRARY BERKELEY Return to desk from which borrowed. This book is DUE on the last date stamped below. 24Nov'50 LIBRARY USE JUN 1 4 1955 APR i 9 '65- LD 21-100m-9,'47(A5702sl6)476 17- 1958 11959 fha f oun chemlca Jul.2842 ations of 0ory Blol.Llb* THE UNIVERSITY OF CALIFORNIA LIBRARY