I A TEXT- BOOK OF INORGANIC CHEMISTRY BY GEORGE SENTER, D.Sc. (LOND.), PH.D. \ * .ECTURER ON CHEMISTRY AT ST. MARY'S HOSPITAL, UNIVERSITY OF LONDON LECTURER ON PHYSICAL CHEMISTRY, CASS TECHNICAL INSTITUTE, ALDGATE, B.C. J FORMERLY EXAMINER IN CHEMISTRY TO THE ROYAL COLLEGE OF PHYSICIANS OF LONDON AND THE ROYAL COLLEGE OF SURGEONS OF ENGLAND NEW YORK D. VAN NOSTRAND COMPANY 23 MURRAY AND 27 WARREN STS. 1911 PREFACE UP till about twenty-five years ago, Inorganic Chemistry was based mainly on certain fundamental laws and theories, notably the Laws of the Conservation of Mass and of Energy, the Laws of Chemical Combination, the Atomic Theory, Avo- gadro's Hypothesis, and the Periodic System of Classifying the Elements. The recent development of the subject is char- acterized by the discovery and applications of certain laws and theories usually associated with Physical Chemistry, more par- ticularly the principles of Chemical Equilibrium (including the Law of Mass Action), the conception of Osmotic Pressure and its application to the determination of Molecular Weights in Solution, and the Electrolytic Dissociation Theory. The newer laws and theories are a necessary supplement to, and extension of the older principles established by the labours of Boyle, Lavoisier. Richter, Dalton, Avogadro and Mendeleeff, but their application to the problems of Inorganic Chemistry has to a large extent revolutionized the study of that subject. It is now generally recognized that the newer views have contributed enormously to the development of chemistry, but in spite of this fact their general adoption into courses of elementary instruction in this country has been very slow. The present book is written throughout from the modern standpoint, and it is hoped that it may contribute in some degree to the wider use of the newer principles at a relatively early stage in chemical courses. Chemistry is an experimental science, and no adequate know- ledge of it can be gained without practical work. The study of 254496 vi A TEXT-BOOK OF INORGANIC CHEMISTRY a text-book should proceed simultaneously with the performance of a number of experiments illustrating the general principles of the subject. From considerations of space, it has been found impossible to give full practical details for experimental work in this book, but simple experiments for illustrative purposes can readily be devised. The order of treatment of the different parts of the subject presented considerable difficulty, and has been decided by my own teaching experience. One of the guiding principles has been that a study of the facts on which laws and theories are based should precede the statement of the laws and theories themselves. For this reason the study of oxygen, hydrogen and water is taken up at a very early stage, and in connexion with these subjects the general properties of gases, liquids and solu- tions are discussed, and a brief account of combustion is given. Chlorine and hydrogen chloride are then considered, and in this way a sufficient number of facts are available to illustrate the Laws of Chemical Combination, the Atomic Theory, the deduc- tion of Molecular Formulae and the writing of Equations, which are dealt with in succeeding chapters. The fundamental distinc- tion between facts and theories, so often inadequately appreciated by the student, is emphasized by this arrangement of the subject- matter. It follows from the above statement that the theoretical part of the subject is distributed throughout the book. The different topics are dealt with as opportunity offers, but the treatment of any particular subject is postponed till facts illustrating it have been mentioned. The subject in question is then considered as fully as space permits a mode of procedure which appears preferable to dealing with it in a fragmentary manner in different parts of the book. The latter method may usefully be employed for lectures, but the former method greatly increases the value of the book for purposes of reference and revision. As regards the order, I have also been guided to some extent by the historical development of the science. PREFACE vii Subject to the above considerations, the more important branches of chemical theory, including the Principles of Chemical Equilibrium, the determination of Molecular Weights in Solu- tion, Electrolysis and Electrolytic Dissociation, are introduced at a fairly early stage of the work, and their applications are repeatedly illustrated in the latter part of the book. The Periodic System is described after the non-metals, and serves as the basis for the discussion of the metals and their com- pounds. One of the main tasks of a teacher is so to familiarize the student with a general principle that he learns to apply it with confidence. The mere statement of a principle, in the expectation that the student will straightway apply it in his later work, is useless. It will be found that the subjects usually included in an Elementary Course, such as Stage I. of the Board of Education, are all dealt with in the early part of the book, and the more advanced work in the later chapters. Further division of the subject into stages leads inevitably to the result that information regarding any important element has to be sought for in different parts of the book, thus rendering it much less useful for purposes of reference and revision. The book is designed for use in University, Technical Institute and other general classes on the subject, and contains all that is usually included in a B.Sc. Course. In the preparation of the book, which is based on my lectures, a number of larger works, more particularly Abegg's Anorganische Chemie and Ostwald's Anorganische Chemie, have been frequently consulted ; the Abstracts published by the Chemical Society have also proved very serviceable. The references to Physical Chemistry in the text are to the Outlines of Physical Chemistry (2nd edition, 1911), where further information on certain branches of the subject may be found. In conclusion, I desire to express my sincere thanks to Dr. J. T. Hewitt, F.R.S., the General Editor of the Series, who viii A TEXT-BOOK OF INORGANIC CHEMISTRY read the whole of the manuscript and made many valuable suggestions. I must also express my obligations to Prof. A. W. Porter, F.R.S., and Dr. H. Burrows for useful criticisms, and to Mr. T. J. Ward, Mr. R. W. Davies, and Dr. J. M'Donald for assistance in correcting the proofs. G. S. LUMSDEN, ABKRDEENSHIRE, September 1911. CONTENTS CHAP. I' AGE I. INTRODUCTORY ILLUSTRATIONS OF CHEMICAL CHANGE I II. CONSERVATION OF MASS AND OF ENERGY CHEMICAL ATTRACTION ... ... III. THE CHEMICAL ELEMENTS IJ IV. OXYGEN COMBUSTION ...... 2O V. HYDROGEN GENERAL PROPERTIES OF GASES . . 31 VI. WATKR GENERAL PROPERTIES OF LIQUIDS . . 51 VII. SOLUTION 76 VIII. CHLORINE AND HYDROCHLORIC ACID ... 86 IX. LAWS OF CHEMICAL COMBINATION THE ATOMIC THEORY ...... . 101 X. DETERMINATION OF ATOMIC WEIGHTS COMBINING WEIGHTS AND CHEMICAL EQUIVALENTS FORMULAE AND EQUATIONS VALENCY . . . . Ir 5 XI. OZONE AND HYDROGEN PEROXIDE THERMO- CHEMISTRY 133 XII. THE HALOGENS AND HALOGEN ACIDS . . . 149 XIII. CHEMICAL EQUILIBRIUM THERMAL DISSOCIATION . 164 XIV. OXIDES AND OXYGEN ACIDS OF THE HALOGENS . 176 XV. OSMOTIC PRESSURE AND MOLECULAR WEIGHT IN SOLUTION . . . . . . . .192 x A TEXT-BOOK OF INORGANIC CHEMISTRY CHAP. PAGE XVI. NITROGEN, THE ATMOSPHERE AND THE ELEMENTS OF THE HELIUM GROUP ..... 2OO XVII. COMPOUNDS OF NITROGEN WITH HYDROGEN AND WITH THE HALOGENS 213 XVIII. OXIDES AND OXYACIDS OF NITROGEN . . .222 XI*. PHOSPHORUS 238 XX. ELECTROLYSIS AND ELECTROLYTIC DISSOCIATION . 257 XXI. SULPHUR, SELENIUM AND TELLURIUM . . .273 XXII. CARBON 308 XXIII. COMBUSTION AND FLAME . . . . -337 XXIV. SILICON AND BORON 347 XXV. CLASSIFICATION OF THE ELEMENTS THE PERIODIC SYSTEM GENERAL PROPERTIES OF THE METALS AND THEIR COMPOUNDS ..... 362 XXVI. THE ALKALI METALS . . .. . . . 378 XXVII. METALS OF THE COPPER GROUP ... . . 407 XXVIII. METALS OF THE ALKALINE EARTHS . . -432 XXIX. METALS OF THE ZINC GROUP .... 444 XXX. METALS OF THE ALUMINIUM GROUP . . . 463 XXXI. METALS OF THE TIN GROUP .... 476 XXXII. METALS OF THE ARSENIC GROUP .... 494 XXXIII. METALS OF THE CHROMIUM GROUP . . '5*7 XXXIV. METALS OF THE MANGANESE GROUP . . . 527 XXXV. METALS OF THE IRON GROUP .... 535 XXXVI. METALS OF THE PLATINUM GROUP . . '55^ XXXVJI. RADIO-ACTIVITY . . . . . . .564 INDEX . . . 569 A TEXT-BOOK OF INORGANIC CHEMISTRY CHAPTER I INTRODUCTORY ILLUSTRATIONS OF CHEMICAL CHANGE OF the different branches of human knowledge, the study of natural objects is the most complex and comprehensive ; it constitutes the domain of natural science. For the sake of con- venience, it has been found desirable to distinguish between the biological sciences, which are concerned with living things, and the physical sciences, such as astronomy, geology, physics and chemistry, which are primarily concerned with the behaviour and properties of non-living things. The different branches of natural science are, however, by no means sharply marked off from each other. The geologist, for instance, deals with such questions as the origin of a deposit of chalk and its relative age with regard to other deposits, but questions as to the composition of chalk, the nature of the changes produced by heat- ing, etc., belong primarily to the provinces of physics and chemistry. In many of the problems with which the geologist has to deal, how- ever, a knowledge of the behaviour of his materials from a chemical and physical point of view is indispensable, so that these branches of knowledge overlap in many respects. The same is true to a greater or less extent of all the natural sciences. At this stage of our work it is not possible to give a satisfactory definition of the province of chemistry, but the following illustrations will give a preliminary idea of the nature of chemical changes, more particularly with reference to the broad distinctions between physical and chemical changes. i 2 A TF,XT-BOpK OF INORGANIC CHEMISTRY Physical and Chemical Changes If a stick of sulphur is briskly rubbed with a dry cloth it acquires the power of attracting light objects, such as small pieces of paper, and is said to be electrified. If olaced in hot water it acquires a new property, that of being able to give up heat to other bodies at the ordinary temperature. Further, i placed in a test-tube and carefully heated over a Bunsen flame, the sulphur melts to a yellowish or brownish liquid (according to the degree of heating), but on removing the tube from the source of heat the contents soon become solid again, and then show all the pro- perties of ordinary sulphur. It should be clearly realized that substances can only be recognized, that is, distinguished from other substances, by their properties. Thus the stick of material which we know as sulphur is characterized by its colour, by its shape, by its apparent weight, the temperature at which it melts, etc. Every substance has an almost infinite number of properties, but for purposes of recognition some pro- perties are much more important and characteristic than others. Further experience will show us that it is advantageous to classify the properties of a substance as (i) characteristic, those which pertain to it under all circumstances, (2) non- characteristic, those which may alter without the substance losing its identity. For instance, the shape of a substance, such as our sulphur, is not a characteristic property ; its colour, on the other hand, is a char- acteristic property. We shall learn later that the most important characteristic property of a substance is its composition. In examining the properties of a substance we do not confine our attention to those which can be observed directly, but extend it to the effects produced by altering the conditions. Much information, for instance, is gained by noticing the effect of heat upon substances. Regarding the experiments with sulphur described above in the light of these considerations, we see that only one or two of the properties non-characteristic properties of the sulphur have been altered, and further, the changes are merely temporary. If left to itself, the electrified sulphur soon loses the property of attracting light objects, and the temperature of the heated sulphur soon falls to that of its surroundings. Such changes, which are temporary and affect only a few of the properties of a substance, are termed physical changes. When a little sulphur is placed on platinum foil and brought in contact with a flame, it catches fire and burns with a bluish flame, giving rise to a characteristic sharp odour ; in a short time it com- ILLUSTRATIONS OF CHEMICAL CHANGE 3 pletely disappears. In this case a much more fundamental change has taken place, and the change is permanent ; no substance having any of the original characteristic properties of sulphur remains. On the other hand a new substance, characterized by a sharp, choking smell, is formed and escapes into the atmosphere. The sulphur has in this case undergone a chemical change. Another instructive experiment is to place successively in the same flame a piece of platinum wire and a piece of magnesium ribbon. The platinum wire becomes white-hot, but on removal from the flame regains all its original properties, so that the change is a physical one. The magnesium ribbon, on the other hand, burns with an extremely bright flame, and when the change is complete there remains a new substance in the form of a white powder, which differs in ill its properties from the original metal. Hence a chemical change has occurred. On the basis of these experiments, a broad distinction can be drawn between physical and chemical changes. When the change is more or less temporary and concerns only a few of the properties of the substance it is physical; when, on the other hand, new sub- stances, characterized by entirely different properties, are formed, the change is a chemical one. It should be carefully noted that the chemical change of one substance (or number of substances) to others is not gradual, but is perfectly sharp and definite. This is well illustrated by the formation of the new substance on burning magnesium ribbon. This change is usually incomplete, especially it performed in a crucible, and particles of the white powder may be observed lying side by side with particles of magnesium. The pro- perties of the two are entirely distinct ; each particle has either changed completely, or not at all. In the extreme cases discussed above, there is no difficulty in distinguishing between physical and chemical changes, but, as we shall learn later, the matter is by no means always so simple. It is evident from the foregoing that the essential point is to be able to detect the formation of new substances, and, for the sake of clear- ness, one or two definitions will now be given. Everything around us is said to be composed of matter. A satisfactory definition of such a fundamental conception as matter cannot be given, but for our present purpose matter may be regarded as anything which occupies space or has weight. A definite limited portion of matter, isuch as a piece of sulphur, a piece of granite, or a knife, is termed a body or thing. Bodies differ, however, in 4 A TEXT-BOOK OF INORGANIC CHEMISTRY complexity. The material of which the piece of sulphur is made up is homogeneous to the naked eye, and even under the microscope, whereas in granite three constituents can readily be distinguished, a white crystalline part termed quartz, a gray portion, felspar, and nearly colourless, lustrous scales called mica. Each of these three components is in itself uniform or homogeneous to the naked eye. Such homogeneous materials are termed substances. Whilst sulphur is made up of a single substance, granite is made up of three sub- stances, each of which is characterized by its special properties. The important distinction between bodies and substances should be carefully noted. Needles, chains, hammers, and nails are different bodies, but all may be composed of the same substance, steel. Further Illustrations of Chemical Change. Mixtures and Chemical Compounds When a piece of sulphur is ground in a mortar and mixed with about its own weight of iron filings, an apparently homogeneous, grayish powder is obtained. On ex- amination with a microscope, however, the separate particles of iron and sulphur can readily be detected. Moreover, if a magnet is held just above the mixture, the iron can be removed, leaving the sulphur, and a separation can also be effected by shaking up the mixture with a liquid called carbon disulphide, which takes up (dissolves) the sulphur, leaving the iron. It is evident that we are dealing with a mechanical mixture of iron and sulphur, in which both substances retain all their properties unimpaired. If now the mixture is placed in a test-tube and heated, the contents of the tube soon begin to glow, and this glowing increases and spreads through the whole mass, even if the tube is removed from the flame. After the tube has cooled to room temperature, a dark mass remains, unlike either the iron or sulphur. From this mass the iron cannot be abstracted by means of a magnet, nor can the sulphur be dissolved out by means of carbon disulphide. Further, if a few drops of dilute hydro- chloric acid are added to a small portion of the dark mass in a test-tube, a gas with a very disagreeable odour is given off, quite different from that obtained by adding a little of the acid to the mechanical mixture of iron and sulphur. It is evident, therefore, that on heating the mixture a chemical change has taken place between the iron and the sulphur, resulting in the formation of a substance whose pro- perties are entirely different from those of the original substances. The new substance is a definite chemical compound which, in allusion to its formation from iron and sulphur, is called iron sulphide. There can be no doubt that iron sulphide contains both iron and ILLUSTRATIONS OF CHEMICAL CHANGE 5 sulphur, as it is possible by special methods to obtain these sub- stances from the sulphide. The fact that the latter shows none of the properties of the component substances illustrates the pro- found nature of the change which we call chemical combination. In striking contrast to the behaviour of a chemical compound, the properties of a mechanical mixture of two (or more) substances is the mean of the properties of the components. Another instructive chemical change will now be described. A small amount of the red powder called mercuric oxide is cautiously heated in a. test-tube until it turns black. If at this point the tube is removed from the flame, the powder regains its original colour and other properties on cooling, and the change in question is, therefore, a physical one. If, however, the heating is continued, it will be noticed that a mirror soon begins to form on the cooler upper part of the tube, and, further, if a glowing splinter is inserted in the tube it will burst into flame. The tube is now removed from the source of heat, and the mirror rubbed with a glass rod, when small shining globules will be formed, characteristic of mercury. The property of causing a glowing splinter to burst into flame indicates the presence of a colour- less gas, called oxygen. In this case, therefore, under the influence of heat, the red powder has given rise to two new substances, the liquid metal mercury and the colourless gas oxygen. In order still further to illustrate the variety of chemical changes, an experiment of a different type will now be described. When a few crystals of ordinary sodium carbonate (washing soda) are shaken up with water, they soon disappear and form a homogeneous mixture with the water. Such a homogeneous mixture is called a solution, and the sodium carbonate is said to have dissolved in the water. A solution in water of the substance called calcium chloride is prepared in the same way. On mixing these clear solutions a white substance is producec , which, in course of time, partially settles to the bottom of the mixture. The appearance of this substance with new pro- perties indicates that a chemical change has taken place between the sodium carbonate and the calcium chloride. The solid substance may be separated from the remainder of the mixture by pouring the contents of the beaker on a filter paper supported in a glass funnel, as shown in Fig. i. Filter- paper is a porous form of paper which readily allows liquids to pass through, but retains solid sub- stances. After most of the liquid has passed through, the solid residue on the filter-paper is treated with a little water to remove adherent solution, and on subsequent examination it may be shown 6 A TEXT-BOOK OF INORGANIC CHEMISTRY to have all the properties of ordinary chalk (calcium carbonate). It is still possible, however, that another substance may have been formed by interaction of sodium carbonate and calcium chloride in solution, but has remained dissolved in the water. The simplest way of testing this possibility is to heat the solution which has passed through the filter-paper known as the filtrate until all the water is driven off. Finally a residue is left which is salt to the taste, and has all the properties of common salt (sodium chloride). The change in FIG. i. this case is, therefore, a somewhat complicated one, two substances entering into chemical combination and giving rise to two new sub- stances. The foregoing experiments, which should be performed by the student, illustrate the three most important types of chemical change, which are as follow : (i) Chemical Combination, when two or more substances unite to form a single substance. Example, iron and sulphur. If the symbols A and B represent the two substances, this type of chemical change may be represented thus A + B-AB, ILLUSTRATIONS OF CHEMICAL CHANGE 7 where the approximation of the letters on the right of the arrow indicate that the substances are chemically.combined. (2) Chemical Decomposition, when a chemical compound splits up into two or more substances. The decomposition of mercuric oxide by heat is a good illustration of this type of change. In symbols it is represented thus AB-^A + B. (3) Double Decomposition, when two (or more) substances interact to form two (or more) new substances. Example, sodium carbonate and calcium chloride, as just described. A simple case of double decomposition is represented symbolically thus A + B-^C + D, where A and B stand for the reacting substances and C and D for the products. One important point in connexion with the chemical changes described is that the reacting substances must be in actual contact before any reaction occurs. This is an important characteristic of all chemical changes. Elements and Chemical Compounds The foregoing ex- amples show that under certain conditions a chemical compound can be split up into simpler substances, as in the case of mercuric oxide. It is natural to inquire whether, by further heating or otherwise, these substances can be split up into anything still simpler. So far, this has not been done ; both mercury and oxygen have up to the present resisted all attempts at further simplification. Substances of this type are termed elements. The accepted view of the constitution of the uni- verse is that it is made up of a large number of different kinds of matter the elements, and no method is known by which any element can be' further simplified or can be converted into another element at will. 1 Chemical compounds are made up of two or more elements in chemical combination. A binary compound is one which contains two elements only. Many of the common metals, such as iron, lead, tin, and zinc are elements, as are sulphur, phosphorus, and charcoal (carbon). Up to the present about eighty elements have been discovered. They are enumerated, and some of their more important properties briefly considered, in Chapter III. , ! As will be pointed out in detail later (chap, xxxvii.) it has been found quite recently that certain elements, more particularly radium, have the power of de- composing spontaneously with ultimate production of other elements, but up to the present no method of initiating or controlling such changes has been discovered. CHAPTER II CONSERVATION OF MASS AND OF ENERGY- CHEMICAL ATTRACTION Conservation of Weight in Chemical Changes So far, ^-^ we have considered chemical changes from the qualitative point of view only, but it is also necessary to consider them from the point of view of the relative amounts of the reacting substances. As a preliminary to this inquiry, it is first necessary to ascertain whether chemical changes are associated with changes in weight of the reacting substances. The indispensable instrument in such in- vestigations is the balance, by means of which the mass, or quantity of matter, is determined. The balance (Fig. 2) consists essentially of a long beam, supported on a knife-edge at its middle point, and provided with scale-pans supported on knife-edges at either extremity of the beam. The substance of unknown weight is put on one scale-pan and known weights on the other until the beam is exactly horizontal, as indicated by a pointer in front of a scale. The balance does not indicate directly the mass or quantity of matter in a body, but its weight, that is, the force with which it is attracted toward^ the earth. The weight of a substance is the product of its mass and the force of gravity at the place of observation. Hence, as at any one place the weight and the mass bear a constant ratio to one another, 1 two bodies of the same weight have equal masses or contain equal quantities of matter. We are now in a position to investigate the question as to whether there is any alteration in mass (or in weight) when substances enter into chemical combination. Some well-known experiments appear at first sight to show that such changes in weight actually occur. Into an ordinary crucible provided with a lid some pieces of magnesium ribbon are placed, and the crucible and contents weighed. The covered crucible is then strongly heated over a Bunsen flame, 1 The ratio between weight and mass, in other words the force of gravity, is different at different parts of the earth's surface. CONSERVATION OF MASS AND OF ENERGY 9 FIG. 2. the lid being raised occasionally so as to admit air. When the glowing of the metal has practically ceased,-the crucible and contents are allowed to cool and again weighed. If the experiment has been carefully performed it will be found that there is a. gain m weight. This, how- ever, does not show conclu- sively that the products of the chemical change weigh more than the substances which entered into reaction, as in the course of the heat- ing something may have been taken up from the air. When a candle burns in the air it slowly disappears, and in this case it seems as if the chemical change is attended by a loss of weight. There is, however, the possibility that something may in this case be passing into the air, thus escaping being weighed, and this may be shown to be the case by the following experiment. A piece of candle is fixed on a piece of cork cut so as to fit the bottom of a glass cylinder (lamp glass), the cork being provided with a large number of holes for the admission of air. In the upper part of the cylinder a piece of wire gauze is supported, and the space above it is filled with sticks of a substance known as caustic soda, which, as we shall see later, has the property of taking up the substances formed when a candle burns in the air (Fig. 3). The whole arrangement is then placed on one pan of a balance and weights just sufficient to bring the pointer to the middle of the scale placed on the other. As the candle burns, it will be observed that that side of the balance is slowly depressed, showing a gain in weight. The same remark applies here, how- ever, as in the magnesium experiment j the increase in weight may be accounted for by the absorption of some- thing from the air. It is evident that in order to ensure that all the reacting substances and all the products are taken into account, the reaction must be carried out in a closed space. A convenient chemical change FIG. 3. io A TEXT-BOOK OF INORGANIC CHEMISTRY for our purpose is the action of sodium carbonate on calcium chloride in aqueous solution, which has already been described in another con- nexion (p. 5). A wide-mouthed glass bottle (Fig. 4) contains a small amount of an aqueous solution of sodium carbon- ate, and a small tube, containing an aqueous solution of calcium chloride, is also placed in the bottle in such a way that no mixing of the solutions takes place ; the bottle is then well corked and the whole weighed. The bottle is then inclined in such a way that the solutions mix and double decomposition (p. 7) takes place. On weighing again, it will be found that the weight is unaltered. Experiments with other substances lead to the same conclusion ; provided that all the reacting sub- stances and all the products of reaction are taken into FIG. 4. account, there is no change in weight as the result of a chemical change. The short statement of these results is called the LAW OF THE CONSERVATION OF MASS, and should be remembered in the following form : When a chemical change occurs, the total weight (or mass] of the reacting substances is equal to the total weight (or mass) of the products. The law holds for physical as well as for chemical changes. The above statement is sometimes called the law of the conservation of matter, and is regarded as indicating that the quantity of matter in the universe cannot be altered in consequence of chemical (or any other) changes ; in other words, matter is indestructible. The defini- tion in italics is, however, to be preferred, as it is purely experi- mental, whereas the latter statement introduces difficulties with reference to the proper definition of matter. The law of the conservation of weight, which was firmly established by the investigations of Lavoisier towards the end of the eighteenth century, is of the most fundamental importance for chemistry. Al- though it may seem at first sight that the law is self-evident, it must be clearly realized that the law is a purely experimental one, and can only be regarded as established within the limits of the unavoidable errors of the experiments which have been made to test it. Landolt's Experiments on Conservation of Weight The most comprehensive series of experiments undertaken with the object of testing the validity of the law in question are due to Landolt. In the majority of Landolt's experiments glass tubes with two limbs were used. The reacting substances were placed separately CONSERVATION OF MASS AND OF ENERGY n one in each limb, the tube carefully sealed and weighed. The tube was then inclined so as to mix the two solutions, and when the chemical action was complete and the temperature had fallen to that of the atmosphere, was again weighed. A special form of balance was used, capable of detecting extremely small differences of weight. The final result of Landolt's experiments, which extended over more than twenty years, is that the law of the conservation of weight has been verified to a very high degree of approximation ; in no case were the very slight deviations observed greater than the possible experimental error. The Atmosphere and Burning It now remains to reconcile the fact 1 hat an increase of weight is observed when magnesium and a candle burn in the air with the law of the conservation of weight. The observed results could be satisfactorily accounted for if some- thing is taken up from the atmosphere during the process of burning, so that the respective increases in weight observed with the magnesium and with the candle are exactly balanced by the loss of weight suffered by the atmosphere. The detailed proof of the validity of this suggestion can only be given at a later stage (p. 28), but it can readily be shown that the air has weight. A large glass globe (Fig. 5) provided with a stopcock (a brass stopcock is suitable) is placed on one pan of the balance and weights added to the other pan until it is in equilibrium. Part p IG ,- of the air is then pumped out, the stopcock is closed, and after replacing on the pan the globe and contents prove to be considerably lighter than before. As will be shown in detail later, the constituent which is taken up from th(; atmosphere during burning is oxygen, a gas obtained, as already mentioned, by the action of heat on mercuric oxide. There is now no difficulty in understanding the burning of magnesium and of a candle in the atmosphere. In the former case the magnesium enters into chemical combination with the oxygen of the atmosphere, forming a white powder which necessarily weighs more than the magnesium first taken. Similarly, the constituents of the candle combine with oxygen during burning, giving rise to gaseous products which are taken up by the sticks of caustic soda. The arrangement increases in weight during the burning, to an extent determined by the weight of the oxygen absorbed. The Conservation of Energy. Chemical Energy It was shown in describing the influence of heat on a mixture of iron 12 A TEXT-BOOK OF INORGANIC CHEMISTRY and sulphur (p. 4) that when chemical combination has properly started the reaction proceeds of itself, and is accompanied by a considerable evolution of heat and light. Further investigation has shown that chemical changes are invariably accompanied by heat changes ; in some cases heat is given out, in other cases it is absorbed. Heat is a form of energy. Energy may be defined as that property of a body which diminishes when work is done by the body, and its diminution is measured by the amount of work done by the body. Other forms of energy, besides heat, are potential energy, kinetic energy, electrical energy, and radiant energy. For a full discussion of this subject a text-book on physics should be consulted. One kind of energy can be transformed into another. Thus if the energy of a falling weight be used to drive a stirrer in water, the water becomes hot, and the potential energy which the weight possessed before it began to fall has been converted into heat. Further, if a Bunsen flame is placed under a hot-air engine, the latter is set in motion, so that the heat in this instance is partially transformed to mechanical (kinetic) energy. The reader will be able to supply many other illustrations of transformation of energy from his own experience. Now it has been found that when a certain amount of one form of energy disappears, an equivalent amount of another form of energy makes its appearance ; in other words, energy can neither be created nor destroyed. This result, which is purely experimental, is termed the LAW OF THE CONSERVATION OF ENERGY, and may be expressed as follows : The energy of an isolated system is constant, that is, it cannot be altered in amount by interactions between the parts of the system. By an isolated system we mean one which is neither receiv- ing energy from outside nor giving up energy to its surroundings. If, for example, two substances capable of entering into chemical combination are contained in a closed vessel, through the walls of which no energy enters or passes out, the total energy inside is the same before and after chemical combination. We may assume that the universe is made made up of two things, and of two things only, matter and energy, neither of which can be altered as regards total quantity, although they may be altered in form. During the combination of iron and sulphur a considerable- amount of heat is given out, as already mentioned, and therefore, according to the law of the conservation of energy, an equivalent amount of another CONSERVATION OF MASS AND OF ENERGY 13 form of energy must have disappeared. The latter form is con- veniently called chemical energy, and we state that iron and sulphur, in the un combined condition, possess a considerable store of chemical energy, part of which is transformed into heat and light when they enter into chemical combination. It must not be assumed that iron sulphide has no chemical energy we know, as a matter of fact, that it has bat it possesses less energy than the free elements before combination. Chemical energy may, however, readily be transformed into other kinds of energy than heat. This may be shown very readily by means of the action of a mixture of sulphuric acid and water on metallic zinc. When the zinc and dilute acid are brought together in a test-tube a very vigor- ous evolu:ion of gas takes place, and the tempera- ture of the solution in- creases considerably, as shown by a thermometer placed in the tube. The same change may now be brought about in another way. A plate of zinc and a plate of copper are joined by means of wires to an instrument (Fig. 6) known as a galvanometer, the needle of which moves along the scale when an electric current passes through. The metals, without being allowed to touch, are dipped into a mixture of sulphuric acid and water, when it will be found that the zinc dissolves and simultaneously an electric current passes through the galvanometer. In this case the chemical energy which disappears when zinc and sulphuric acid enter into chemical combination appears, in part at least, as electrical energy. The conversion of chemical into other forms of energy is of the utmost commercial importance. Our most important source of energy i.s the burning of coal, a process in which the materials of the coal combine with the oxygen of the atmosphere, the enormous store of chemical energy which the substances contain before com- bination thus becoming available (p. 330. Conversion of Heat and of Electrical Energy into Chemical Energy The converse of the transformations of FlG i 4 A TEXT-BOOK OF INORGANIC CHEMISTRY energy described above can also be readily performed. Many compounds can be split up by heat into simpler substances which possess more chemical energy than the original compounds ; in such cases heat has been partially transformed into chemical energy. The action of heat on mercuric oxide is a process of this nature, as free mercury and free oxygen possess more energy than the corre- sponding amount of mercuric oxide. In a similar way, an electric current may be used to split up complex substances into simpler ones possessing more chemical energy ; in such a case, electrical energy has been transformed into chemical energy. On account of the great im- portance of this principle, we will illustrate it by an account of the decomposition of water by the agency of an electric current. Up till about 130 years ago water was regarded as a chemical element, and the experiment now to be described is one of the most convenient for the demonstra- tion of its compound nature. The apparatus used for the purpose, known as a voltameter, is represented in Fig. 7. It con- sists essentially of a tube bent in the form of the letter U (so-called U-tube) joined at the lower part to a longer tube ending in a bulb B. In the lower part of the two limbs of the U-tube are two platinum plates, a, a, connected with platinum wires which are sealed through the glass and can be joined (usually by copper wires) to the positive and negative poles respectively of a battery. In order to perform an experiment, the stopcocks at the upper ends of the two limbs are opened and dilute sulphuric acid * is poured into B till it rises to the level of the stopcocks, which are then closed. As soon as con- nexion is made with the battery, bubbles of gas are given off from each platinum plate, and rise to the upper part of the two tubes. After the current has passed for some time, it will be observed that the volume of gas which has collected above the platinum plate connected with the negative pole of the battery is about double that in the other limb. The gases are colourless, and can easily be shown to be both odourless and tasteless. The gas present in larger proportion can 1 Water containing a little sulphuric acid. FIG. 7. CONSERVATION OF MASS AND OF ENERGY 15 be collected in a small test-tube by inverting the latter over the top of the vo tameter tube and cautiously opening the tap. When a light is put to the mouth of the tube the gas takes fire and burns with an almost colourless flame. By means of these and other tests, it can be recognized as hydrogen, a gas which can be made in quantity by more convenient methods (p. 31). It can further be shown, from its property of igniting a glowing splinter and by other tests, that the gas in the other tube is oxygen, the preparation of which from mercuric oxide has already been described (p. 5). The effect of passing an electric current through acidulated water is therefore very remarkable, inasmuch as at one plate the negative plate only hydrogen is given off, at the other plate only oxygen is given off, and the volume of the hydrogen is about double that of the oxygen. We shall see later on that a mixture of hydrogen and oxygen can be caused to combine with formation of water, and in this process a large amount of heat is given out (p. 37). The process just described represents the reverse of this, inasmuch as water has been split up into two component elements, hydrogen and oxygen, and a large amount of electrical energy has been absorbed in the process. At a later stage evidence will be given for the statement that it is really the water which is split up, but a little sulphuric acid must also be added, as pure water does not conduct the electric current. The platinum plate connected with the positive pole of the battery is termed the positive pole, positive electrode or anode j the plate connected with the negative pole of the battery is called the nega- tive pole, negative electrode or cathode. The process of splitting up a chemical compound by means of the electric current is called electrolysis (that is, splitting up by means of electricity). The nature of the chemical changes occurring in the electrolysis of water will be dealt with later, under the heading Electrolysis (p. 263). The Cause of Chemical Change It has already been pointed out that when platinum wire is heated in the air it suffers no chemical change, whilst magnesium ribbon under the same circumstances gives rise to a new compound, and it is natural to inquire into the reasons for this difference of behaviour. Chemists are accustomed to state that there is a certain attraction the so-called chemical affinity between magnesium and one of the constituents of the atmosphere (the oxygen) which comes into play when the magnesium is heated, and leads to chemical combination, whereas there is little or no chemical affinity or chemical attraction between platinum and 16 A TEXT-BOOK OF INORGANIC CHEMISTRY oxygen (p. 558). Similarly, the mercury and oxygen in mercuric oxide are held together by chemical attraction, and the substance has to he raised to a high temperature in order to overcome this attraction and liberate the elements. Very little is known with certainty as to the nature of this so-called affinity or attraction. The term "affinity" denotes likeness, and was introduced into chemistry at a time when it was thought that chemical combination took place most readily between substances of like properties. It will be shown later, however, that this view is erroneous ; the most stable compounds are formed by combination of substances of unlike properties. Summary. Characteristics of Chemical Change The first two chapters have been mainly devoted to a consideration of the nature of chemical change. It will be useful to summarize here the more important characteristics of chemical change. (1) As the result of a chemical change new substances make their- appearance. These new substances are recognized by their pro- perties, more particularly by their composition. The composition of a substance, that is, the elements of which it is built up and the proportion in which they are present, is its most characteristic property. (2) When a chemical change takes place, the total weight of the products is equal to the total weight of the reacting substances (Law of Conservation of Mass). (3) A chemical change is always attended by the evolution or absorption of heat, or more generally, the total chemical energy of the final products always differs from that of the reacting substances. (4) Chemical changes occur only between substances which are in actual contact. (5) A fifth characteristic will be mentioned here for the sake of completeness, and will be fully considered in later chapters. It is that chemical changes always occur between definite weights (or volumes, in the case of gases) of the reacting substances. From this it follows that the composition of a definite chemical compound is constant, no matter how it is prepared, or what is its source. These characteristics of chemical change are important, and should be carefully remembered. CHAPTER III THE CHEMICAL ELEMENTS WE have learnt in the previous chapters that many of the sub- stances with which we are acquainted can be split up into simpler substances by various methods, for example by the action of heat or of electrical energy. Thus from a definite quantity of mercuric oxide by the action of heat, two substances, mercury and oxygen, are obtained, each of which weighs less than the original oxide, whilst the sum of their weights is equal to that of the original oxide. So far, no method has been discovered by means of which mercury and oxygen can be split up into "simpler" substances; in other words, it has not been found possible to separate a definite amount of either of those substances into two or more other sub- stances, each weighing less than the original substance. For this reason mercury and oxygen are classed as elements. It must be care- fully noted that elements are substances which so far have resisted all attempts at decomposition, but it does not in the least follow that they are really undecomposable ; it is, in fact, highly probable that in the near future methods of simplifying at least some of the elements will be discovered. The above definition of an element is due to Boyle (1661), and was later adopted by Lavoisier. It has proved extremely serviceable, as many compounds which in Lavoisier's day were classified as elements have since been shown to be chemical compounds. Bearing this in mind, it might be supposed that some substances accepted as elements at the present day are comparatively simple compounds of the same type as water, which could be split up by a more energetic use of the means now at our disposal. We shall find later, however, that this suggestion is highly improbable. While it may be accepted that in one sense the elements are complex, their complexity is of a different order from that of ordinary chemical compounds (p. 366). At present about eighty elements are known. The exact number cannot be definitely stated, as some of them are present only in very small proportion on the earth, and considerable difficulties are met 2 *7 1 8 A TEXT-BOOK OF INORGANIC CHEMISTRY with in obtaining them in a pure condition. The names of the elements now recognized by the International Committee on Atomic Weights are given in alphabetical order in the following table. The more important elements are printed in ordinary type, the less im- portant in italics. Aluminium Fluorine Molybdenum Silver Antimony Gadolinum Neodymium Sodium Argon Gallium Neon Strontium Arsenic Germanium Nickel Sulphur Barium Glucinum (beryllium] Nitrogen Tantalum Bismuth Gold Osmium Tellurium Boron Bromine Helium Hydrogen Oxygen Palladium Terbium Thallium Cadmium Indium Phosphorus Thorium Caesium Iodine Platinum Thulium Calcium Iridium Potassium Tin Carbon Iron Praseodymi u m Titanium Cerium Krypton Radium Tungsten Chlorine Lanthanum Rhodium Uranium Chromium Lead Rubidium Vanadium, Cobalt Lithium. Ruthenium Xenon Columbium Lutecium Samarium . Ytterbium (Neoytterbium] Copper Magnesium Scandium Yttrium Dysprosium Manganese Selenium Zinc Erbium Mercury Silicon Zirconium An inspection of this table shows that many elements are familiar to us in everyday life. Thus copper, silver, gold, iron, lead, zinc and tin are elements, as are sulphur, phosphorus, carbon and oxygen. Copper, silver, gold, lead and a number of other elements are called metals; they show metallic lustre and conduct heat and electricity. The elements which do not possess these properties are called non- metals. This class includes oxygen, hydrogen, sulphur, phosphorus, carbon, and many other elements, the properties of which are con- sidered in detail in the earlier part of the book. The division of the elements into metals and non-metals is not in all respects a satis- factory one, but the lines on which the elements can be classified can only be adequately discussed after we have become familiar with their more important physical and chemical properties. As regards the physical state of the elements, a few of them, such as hydrogen and oxygen, are gases at the ordinary temperature, two only, mercury and bromine, are liquid, and the great majority are solid under ordinary conditions. A few of the elements, such as copper, silver, gold, oxygen and sulphur occur free (that is as elements) in nature, but most of them are found naturally only in the form of chemical compounds. The THE CHEMICAL ELEMENTS methods employed in obtaining elements from their compounds are considered in detail in connexion with the mdividual elements. We have already seen that heat and electrical energy can be successfully employed for this purpose in some cases. The relative proportions in which the elements occur in the part of the earth accessible to us are very unequal. Oxygen, which occurs in the air, in water, and in the earth's crust, constitutes about 50 per cent, by weight of the part of the earth (including the sea) known to us. The nine elements, oxygen, silicon, aluminium, iron, calcium, magnesium, sodium, potassium, and hydrogen (free or in combination) together make up more than 99 per cent, of the earth's crust. The following table, compiled by Clarke, contains an estimate of the relative amounts of the elements occurring in the part of the earth known to us : Earth's Crust. Ocean. Atmosphere. Entire Earth. Per cent. Per cent. Per cent. Per cent. Oxygen . . oihcon . 47.29 27.21 85-79 23.00 49.98 2 5-3o Aluminiun 7.81 7.26 Iron .... 5-46 5.08 Calcium 3-77 0.05 3-5i Magnesium . 2.68 0.14 2.50 Sodium . 2.36 0.14 2.28 Potassium . . 1 2.40 0.04 2.23 Hydrogen , . 0.21 10.67 0.94 Titanium . 1 -33 0.30 Carbon ... 0.22 0.002 0.21 Chlorine . . o.oi 2.07 0.15 Bromine O.OO8 Phosphorus O.IO 0.09 Sulphur. . . j 0.03 O.Og O.O4 Nitrogen . 77.00 0.02 CHAPTER IV OXYGEN COMBUSTION WITH this chapter we commence the systematic study of the elements and their- more important compounds. In general, the order of treatment will be as follows : History, Occurrence, Methods of Preparation, Physical Properties, and Chemical Properties of the particular element. The more important compounds which the element forms with other elements will then be considered. History The discovery of oxygen was made and publicly an- nounced by the English chemist Priestley in 1774. He obtained it by enclosing red oxide of mercury in a tube over mercury, and concentrat- ing the rays of the sun on it by means of a powerful lens. Priestley observed that substances which burned in air burned still more rapidly in the new gas, and, for reasons which will be mentioned later (p. 29), he termed the gas dephlogisticated air. Oxygen was independently discovered by the Swedish chemist Scheele, who in 1775 described a number of methods of preparing it. A recent study of Scheele's original papers has shown, however, that this chemist discovered oxygen in 1773, about a year before Priestley. In the publication of his discovery, Scheele was, however, anticipated by Priestley, as already indicated. Lavoisier observed that many of the products obtained by burning substances in Priestley's dephlogisticated air readily dissolved in water, and the resulting solutions possessed a sour taste and showed certain other properties generally regarded as characteristic of acids (p. 98). For this reason he called the new gas oxygen, or acid producer (ogvs, acid, and yei/i/aco, I produce), and expressed the view that it is the acidifying principle, to which acids owe their characteristic properties. Later investigation has, however, shown that this view is erroneous, as some of the best known acids, for example hydrochloric acid, contain no oxygen. Occurrence Free oxygen is a very important constituent of the atmosphere, in which it occurs mixed with about four times its volume of another colourless gas, called nitrogen. OXYGEN COMBUSTION 21 In chemical combination it constitutes about eight-ninths of the weight of water, the remaining ninth being hydrogen. It also forms a very important constituent of nearly all rocks and earthy substances in fact, 44 to 48 per cent, of that part of the crust of the earth known to us consists of combined oxygen (p. 19). Oxygen is also one of the essential constituents of practically all animal and vegetable substances. Chemical compounds in which oxygen is combined with only one other element are termed oxides. Preparation Laboratory Methods (i) Oxygen may be pre- pared in small amount by heating mercuric oxide, as already de- scribed (p. 5). For this purpose a small quantity of the oxide is placed in a hard-glass tube, closed at one end, and provided with a cork and delivery-tube, as shown in Fig. 8. For the collection and manipulation of gases over water the arrangement shown in the figure is very convenient. The hollow cylindrical vessel A, which has a small central opening on the upper surface, and is open on the lower side, is placed in the vessel B, and completely covered by water. The end of the delivery tube is then placed below the lower edge of the vessel A, which is grooved for the purpose, and the hard-glass tube carefully heated with a Bunsen burner. The gas is allowed to bubble through the water for some time in order to 22 A TEXT-BOOK OF INORGANIC CHEMISTRY drive out the air. A gas-collecting jar is then inverted over A, as shown, when the oxygen displaces the water in the jar. If, as is necessary in some cases, a gas has to be collected over mercury, a vessel known as a pneumatic trough, provided with a permanent shelf in the interior, is used for the purpose. (2) Oxygen can be obtained most readily for laboratory purposes by heating potassium^chlorate, a substance which occurs in colourless crystals. When the chlorate is heated in a test-tube it melts, and at a rather higher temperature is slowly split up into oxygen and another colourless substance termed potassium chloride. The reaction is one of simple decomposition (p. 7), and is repre- sented by the equation l Potassium Chlorate = Potassium Chloride -I- Oxygen. If the potassium chlorate is mixed with about one-fourth of its weight of a black powder called manganese dioxide, oxygen is given off rapidly at a temperature below the melting-point of the chlorate. This is the most convenient method for the laboratory preparation of oxygen, although the gas obtained in this way is not quite pure. The gas may be collected over water, as described on the previous page, and several jars should be filled with it, in order to demonstrate its properties (p. 25). The most remarkable fact about the change just described is that although the manganese dioxide greatly accelerates the liberation of oxygen from potassium chlorate, and thus enables it to take place at a much lower temperature than when the chlorate is heated alone, yet the manganese dioxide can be recovered unaltered in amount at the end of the process, and therefore plays no apparent part in the reaction. This is only one of many instances in which a substance greatly accelerates a chemical change, while itself remaining unaltered at the end of the reaction. In such cases the substance is said to exert a catalytic effect, and is termed a catalyst for the change in question. It should be clearly understood that the use of these terms does not suggest any explanation for the effect in question, but merely classes together a number of phenomena having certain important features in common. We shall meet with many examples of catalytic action in the course of our subsequent work. (3) Very pure oxygen may be obtained by heating potassium permanganate, a substance which occurs in reddish-black crystals. 1 The term equation denotes that, in accordance with the law of conservation of weight, the sum of the weights of the substances on one side of the = sign is equal to that on the other (cf. p. 10). OXYGEN COMBUSTION 23 The experiment is carried out as described under (i). The chemical change in this case is rather complex, and will be considered under potassium permanganate (p. 532). (4) Oxygen may be obtained by the effect of heat on many substances other than those mentioned, for example, manganese peroxide, but higher temperatures are usually required, and the reactions are in other respects less suitable for laboratory purposes than those already described. (5) Oxygen may also be obtained from substances such as potassium permanganate and potassium bichromate by heating with sulphuric acid. These methods are dealt with in detail at a later stage. Commercial Preparation of Oxygen (6) Brirts Oxygen Process Oxygen gas, compressed into steel cylinders, is now an article of considerable commercial importance. For preparing a small quantity of oxygen in the laboratory, the convenience of the method is the chief consideration, and the cost a matter of secondary importance. In preparing a substance on the commercial scale, however, the cost is of prime importance, and it is therefore necessary to find a plentiful source of the raw material from which the substance required may readily be obtained. In the case of oxygen, it is natural to think of the atmosphere as a source of supply, since it is mainly composed of free oxygen and nitrogen. A method suitable for the present purpose would be to use a substance which enters into chemical combination with oxygen, but not with nitrogen, and the resulting compound must readily give up its oxygen under suitable conditions. A substance which answers these requirements is barium oxide. When this substance, in the form of a white powder, is heated at low red heat (about 500) in the air, it takes up oxygen forming a compound called barium peroxide, which, for the same proportion of barium, contains twice as much oxygen as the oxide. When the peroxide is heated to a bright red heat (about 1000), it gives up half of its oxygen, and barium oxide is reformed. The same quantity of barium oxide may be used over and over again, being alternately heated while air is passed over it, and the product then raised to a higher temperature to drive off part of the oxygen. The reaction may conveniently be represented by the following equation Barium peroxide:Barium oxide + oxygen in which the change in the direction of the upper arrow takes place at a bright red heat ; that in the direction of the lower arrow at a 24 A TEXT-BOOK OF INORGANIC CHEMISTRY low red heat. A change of this type, which may proceed in either direction depending on the conditions, is termed a reversible reaction. In actual practice, the decomposition of the barium peroxide is effected more economically by reducing the pressure instead of by raising the temperature. The process thus consists in passing air over heated barium oxide at atmospheric pressure, and then removing the oxygen at the same temperature under reduced pressure. Before being admitted to the barium oxide, the air is passed over lime to remove a gas called carbon dioxide, which otherwise would combine with the barium oxide and render it useless for absorbing oxygen. The process just described is termed Brin's oxygen process. The compressed oxygen is not pure, but contains a few per cent, of nitrogen. (7) On the commercial scale, oxygen is now obtained almost entirely by the evaporation of liquid air. The latter is a mixture of liquid oxygen and nitrogen. As liquid nitrogen goes into vapour more readily than does liquid oxygen, a mixture rich in oxygen can be obtained when liquid air is allowed to evaporate slowly. As a number of factors are concerned in the process, which at this early stage have not yet been met with, the consideration of this method of preparing oxygen is postponed to the section dealing with the lique- faction of gases (p. 73). Physical Properties Oxygen is a colourless, odourless, taste- less gas. It is 1.105 times heavier than air. The usual standard to which the density of gases is referred is that of hydrogen (the lightest gas) taken as unity; on this basis the density of oxygen is 15.90. The weight of a litre of oxygen at o and 760 mm. pressure is, according to Morley, 1.4290 grams, according to Rayleigh 1.4295 grams. Oxygen may be liquefied at and below - 119, which is its critical temperature (p. 71). At its critical temperature about 58 atmos- pheres pressure are required to liquefy it (the so-called critical pressure), and the further the temperature is lowered below the critical temperature the less is the pressure required to convert the gas to a liquid. At its boiling-point the specific gravity of the liquid is 1.131. Dewar has succeeded in obtaining oxygen as a bluish, snow-like solid by cooling with liquid hydrogen ; the solid melts at -227. Liquid oxygen is attracted by a magnet. Oxygen is slightly soluble in water ; i c.c. of water dissolves at o 0.0489 c.c., at 20 0.031 c.c, and at 30 0.026 c.c. of the gas, measured in each case under i atmosphere pressure (760 mm.). According to OXYGEN COMBUSTION Winkler, the solubility, C, of oxygen in water diminishes as the temperature rises in accordance with the formula C = 0.0489 - o.oo 1 34 1 3/4- 0.00002 8 3^ - 0.000000295 34/ 3 . It is a remarkable fact that oxygen is fairly soluble in fused silver, but escapes almost entirely when the silver solidifies. Chemical Properties In describing the chemical properties of an element, we are mainly concerned with its power of combining directly with other elements, the conditions under which the chemical changes occur, and the nature of the compounds formed. When an element has the power of combining readily with a large number of other elements it is said to be chemically active. Many substances, such as sulphur, phosphorus, carbon, and iron, combine rapidly with oxygen when the reaction has been started by heating. This may readily be shown as follows. A number of gas jars are filled with oxygen by heating a mixture of potassium chlorate and man- ganese dioxide (p. 22), and are closed by glass covers (lubricated with vaseline or grease) until required. The sulphur (phosphorus or carbon) is placed in a spoon of the form shown in Fig. 9, heated till it just begins to burn, and then plunged into the oxygen, when it continues to burn, but much more rapidly than in air. In order to show the burning of iron in oxygen, a bundle of thin iron wires is bent round a small piece of sulphur at one end, and this end is then heated in a flame. The combination of the iron and sulphur gives out a considerable amount of heat, which starts the combination of the iron and oxygen of the air. If the wires, held in a tongs, are quickly immersed in a jar of oxygen, they continue to burn with an extremely brilliant flame, giving rise to a shower of sparks, while particles of the fused product of the reaction fall on the bottom of the gas jar and solidify. If, after action has ceased, the contents of the jars are shaken up with water, and a few drops of litmus solution added, it will be noticed that the litmus is turned red in the jars in which sulphur, phosphorus, and carbon have been burned, whilst no change of colour occurs in the jar in which iron has been burned. The change in the colour of the litmus to red indicates the presence of an acid (p. 98), from which it follows that the oxides obtained by burning sulphur, phosphorus, and carbon in the air dissolve in water to form acids. FIG. 9. 26 A TEXT-BOOK OF INORGANIC CHEMISTRY The oxide of iron is insoluble in water, and therefore does not affect the colour of the litmus. The products formed when substances burn in air are the same as those formed when they burn in oxygen, the only difference being that in the latter case the reactions are more vigorous. This is easily understood when it is remembered that the atmosphere is essentially a mixture of oxygen and another gas, nitrogen ; as the latter is a comparatively indifferent gas, and combines directly with very few substances, air behaves towards the great majority of elements simply as diluted oxygen. Oxygen combines with all the other elements except fluorine and the rare elements helium, neon, argon, krypton, and xenon. In the case of a few elements, such as silver and gold, the oxides can only be obtained indirectly. When an element or compound enters into chemical combination with oxygen, the process is termed oxidation, and the element or compound is said to be oxidized. Combustion The rapid combination of certain elements with oxygen forms a good illustration of combustion, which may be defined as a chemical change which proceeds with the evolution of light and heat. From the definition it follows that the term combustion is not confined to combinations with oxygen, although these are by far the most familiar. When iron filings and sulphur are heated together in a test-tube, they combine with evolution of light and heat, and therefore, according to the above definition, this is a process of combustion. The combination of many elements with oxygen, which proceeds very rapidly at high temperatures, may also proceed at a low tem- perature, though much more slowly. The familiar glowing of phos- phorus in the dark is an accompaniment of the combination of oxygen and phosphorus, and leads to the formation of the same compound as when phosphorus burns brightly. Ftirther, the amount of energy given out when a definite amount of an element combines with oxygen is the same whether the action takes place quickly or slowly. When combustion is rapid, the heat is given out quickly and can thus raise the products of combustion to a high temperature. When, however, the change is slow, the heat of reaction escapes into the surroundings and the temperature may not rise much above the ordinary temperature. Importance of the Study of Combustion for the De- velopment of Chemistry Many of the chemical phenomena OXYGEN COMBUSTION 2 7 with which we are familiar in everyday life, such as the burning of coal and \vood in air, the rusting of iron, fhe glowing of phosphorus, the burning of sulphur in air, are processes of combustion. A point of fundamental importance in all these processes is that the product or products of combustion weigh more than the original substance. At first sight there appears to be a loss of weight in some of these inactions, for example, when a candle burns in air; but it has already been pointed out that when the experiment is performed in such a way that the products of combustion are not allowed to escape into the atmosphere, there is an evident increase of weight in this case also. When it is remembered that oxygen is one of the constituents of air, these observations are readily understood. The burning of the candle is essentially a process of chemical combination between its constituents and oxygen, and it may be anticipated that, in accordance with the law of conservation of mass, the excess in weight of the products is just balanced by the loss in weight of the atmosphere. Although this explanation of the phenomena of combustion is simple and easily understood, it was only arrived at after much experiment and discussion. Its general acceptance towards the end of the eighteenth century was due to Lavoisier. A brief account of his classical experiment on the oxidation of mercury in a confined volume of air, and of the conclusions he drew from his observations, will now be given. Four ounces of mercury were placed in the glass retort A (Fig. 10), the neck of which was bent as shown and dipped into the glass vessel B, which contained air confined over mercury. The retort containing the mercury was then heated for twelve days at a temperature near the boiling-point of mercury, and it was noticed that the surface of the metal gradually became covered with red scales (mercuric oxide). After the apparatus had cooled, it was noticed that the volume of the confined air had diminished ; out of 50 cubic inches originally taken only 42 remained. The red scales were collected, and on heating gave off 7-8 cubic inches of a gas (oxygen) which supported com- bustion much more energetically than ordinary air. Lavoisier also found that the gas remaining in the vessel, which had not combined with the mercury, was no longer capable of oxidizing metals, and a lighted taper was extinguished in the gas as if it had been plunged into water. It follows from this experiment that air is a mixture of two gases, one of which (oxygen) combines with metals such as mercury, the other, present in larger proportion (nitrogen), does not 28 A TEXT-BOOK OF INORGANIC CHEMISTRY combine with metals and does not support the combustion of a lighted taper. The further point, that the gain in weight of the products is balanced by a loss in weight of the atmosphere, was proved more conclusively by Lavoisier in a further experiment on the combustion of tin. Some pieces of tin were sealed up in a glass vessel along with air, and the vessel heated for some time. It was then found that, although the tin had altered in appearance and, as subsequent investigation showed, had gained in weight, there was no difference in the weight of the closed vessel before and after heating. When, however, the vessel was cautiously opened air rushed in, and it was found that FIG. 10. the weight of the air which entered was equal to the gain in weight of the tin during the heating. The simplest laboratory method of showing that a gas is taken up from the atmosphere in the process of combustion is to burn a piece of phosphorus in a confined volume of air. The phosphorus is placed on an inverted crucible lid floating on water, and the whole is covered by a bell-jar provided with a well-fitting stopper (Fig. 11). The stopper is removed for a moment and the water brought to the same level inside and outside the jar, the phosphorus ignited by touching with a heated rod, the stopper immediately replaced, and the bell-jar pressed down against some blotting-paper or other soft material in the bottom of the dish while the combustion OXYGEN COMBUSTION 29 lasts. 1 The bell-jar becomes filled with white fumes (of phosphorus pentoxide), and when practically all the oxygen is used up the phos- phorus becomes extinguished. The jar is then cautiously raised without allowing the lower edge to rise above the surface of the water, when it will be found that water enters and fills about one-fifth FIG. ii. of the space previously occupied by air. The gas remaining in the jar is incapable of supporting the combustion of a burning taper. The conclusion which might be drawn from this experiment, that about one-fifth of the atmosphere by volume is oxygen, is fully confirmed by further investigations (p. 203) ; the remaining four-fifths is mainly nitrogen. The Phlogiston Theory Before Lavoisier's time another view of the nature of combust ion , termed the phlogiston theory, was held almost universally through- out the chemical world. According to this view combustible substances, such as carbon and sulphur, contain a large amount of phlogiston, which escapes during combustion. The residue, if any, was said to be dephlogisticated (i.e. deprived of phlogiston) in this process, as in the case of the so-called "calx" of lead or tin, 2 but it could be restored to its original condition by heating it 1 The a,r expands at first owing to the heat given out in the combustion, and part of it may escape below the edge of the bell-jar unless the precaution above mentioned is taken. 2 The products obtained by burning metals, such as lead and tin, in air were termed ca!xes ; thus the substance now known as lead oxide was called the calx of lead. 30 A TEXT-BOOK OF INORGANIC CHEMISTRY with a substance such as charcoal, the latter being supposed to be rich in phlogiston. The adherents of the phlogiston theory to some extent lost sight of the fact that there is a gain in weight as the result of combustion, whilst if phlogiston is a material substance there ought to be a loss in weight. The overthrow of the phlogiston theory and the establishment of our present views on the subject were due to Lavoisier. As a matter of fact, however, the phlogiston theory, although unsuitable in many respects, contributed materially to the development of chemistry. It will be clear from what has been mentioned above with regard to the conversion of lead to its calx and the reconversion of the latter to metallic lead by charcoal, that the escape of phlogiston denotes what we now term oxidation, and the addition of phlogiston constitutes reduction. On this basis the chemists of that period were enabled to classify together many changes previously regarded as quite different, which greatly facilitated experimental investigation. The function of the air in combustion, according to the adherents of the phlogiston theory, was to take up phlogiston, and when it was no longer able to support combustion it was said to be saturated with phlogiston. As oxygen was a much better supporter of combustion than ordinary" air, it is now easy to see why it was termed by Priestley dephlogisticated air. On the same principle, nitrogen was termed by Priestley phlogisticated air. CHAPTER V HYDROGEN GENERAL PROPERTIES OF GASES History Paracelsus, in the sixteenth century, was familiar with the fact that an inflammable gas is produced by the action of dilute acids on certain metals, but the gas was then looked upon as a form of ordinary air. Cavendish, in 1766, was the first to prepare pure hydrogen and to recognize it as a definite substance ; he obtained it, like Paracelsus, by the action of acids on metals. As has already been pointed out, hydrogen is a constituent of water, hence its name, which is derived from the Greek words vdwp, water, and yevi>aa>, I produce. Occurrence Free hydrogen, mixed with other gases, is given off from volcanoes, 1 and is also found in small amount enclosed in meteorites. It also occurs free in traces in the atmosphere ; accord- ing to Rayleigh the average proportion does not exceed i in 30,000 by volume, whilst according to Ramsay the amount does not exceed i in a million. Spectroscopic observations indicate that free hydrogen is present in very large proportion in the sun, and in many fixed stars and nebuLe. In the combined form, hydrogen occurs very largely on the earth. It forms over n per cent, of water, and is one of the essential ele- ments in plants and animals. It also occurs, combined with carbon, in natural J[as, marh__ga.Sj and petroleum. It is an essential con- stituent of acids. Preparation (A) From Water. As water is a chemical com- pound of hydrp_gen and oxygen, and is always available, it is natural to have recourse to it as a source of hydrogen. The elements may be considered as being held together by chemical attraction, and this attraction must be overcome in some way in order to obtain free hydrogen. Another and preferable way of regarding the matter is that hydrogen and oxygen give out a large amount of energy, mainly in the form of heat, when they combine to form water, and 1 The ga^ es given off from Mount Pelee, Martinique, during the eruption of 1902, contained over 20 per cent, by volume of free hydrogen (Moissan). 31 32 A TEXT-BOOK OF INORGANIC CHEMISTRY in order to obtain the free elements from water this energy must be supplied in some way (p. 14). From these considerations two principal methods might be suggested for obtaining hydrogen from water : (a) to supply a large amount of energy to water, for example, in the form of heat or of electrical energy ; (<) to bring water in contact with a substance which has a greater attraction for oxygen than the oxygen has for hydrogen. The more important methods for the preparation of hydrogen from water, and incidentally for the de- composition of water, will now be briefly considered. (1) When steam is passed through a platinum tube, heated to a temperature exceeding 1000, and the gases leaving the tube are rapidly cooled, very small amounts of hydrogen and oxygen are obtained. This method is not, of course, a practical one for pre- paring hydrogen in quantity, but is of considerable theoretical interest. It is further referred to in connexion with water (p. 52). (2) Hydrogen, and also oxygen, can readily be obtained from water by the employment of electrical energy. The apparatus used for this purpose, known as a voltameter, and the method of pro- cedure, have been described in a previous chapter. Hydrogen obtained in this way is very pure. (3) The second general method of obtaining hydrogen from water can be illustrated by means of the soft metals sodium and potassium, which are acted on by water at the ordinary temperature with liberation of hydrogen. With the former metal the experiment is best performed by enclosing a small piece in a so-called sodium spoon (a small chamber of wire-gauze at the end of a long rod), and holding it below the surface of water beneath an inverted test- tube filled with water. The sodium acts rapidly on water, liberat- ing hydrogen, which displaces the water in the test-tube. The experiment with potassium is best performed by throwing a small piece of the metal on the surface of water in a glass dish. Potas- sium acts on water more vigorously than sodium does, and the heat given out is so great that the hydrogen catches fire, and burns with a violet flame, the colour being due to the vapour of potassium (p. 401). The water in which sodium (or potassium) has been dissolved makes the fingers slippery and turns red litmus paper blue. This indicates the presence <-,f a new substance, which can be obtained as a white residue by evaporating the water. It is called sodium hydroxide or caustic soda. A hydroxide, as its name indicates, contains both hydrogen and oxygen. The chemical change which HYDROGEN GENERAL PROPERTIES OF GASES 33 occurs when sodium is thrown into water is, therefore, represented by the equation Sodium + water = sodium hydroxide + hydrogen. Substances which in solution have a soapy feel, turn litmus blue and show certain other characteristic properties are termed bases, and The solutions are said to have an alkaline reaction. Potassium hydroxide and sodium hydroxide are typical bases. Certain other metals, including lithium, barium, and calcium, also FIG. 12. act on water at the ordinary temperature, liberating hydrogen and forming the corresponding hydroxides. (4) Metals such as magnesium and zinc, although practically without action on water at room temperature, slowly decompose boiling water, and act still more energetically when heated in a current of steam. The arrangement in the case of magnesium is illustrated in Fig. 12. The metal is strongly heated in a bulb tube A and steam from a boiler B passed over it. When nearly red-hot, the metal catches fire in the steam, forming magnesium oxide and hydrogen. The latter can be ignited at the end of the tube as it escapes. At the end of the experiment the magnesium oxide, in the form of a white powder, is found in the bulb. 3 34 A TEXT-BOOK OF INORGANIC CHEMISTRY (5) Iron at a red heat also decomposes steam, forming hydrogen and an oxide of iron. The arrangement of the apparatus for this purpose will readily be understood from the illustration (Fig. 13). FIG. 14. The iron tube, which contains iron in the form of nails, is raised to a red heat by means of the burners, and steam is passed through the tube till all the air is expelled. The hydrogen escaping at the end of HYDROGEN GENERAL PROPERTIES OF GASES 35 the delivery tube can then be collected by displacement of water and tested in the usual way. (6) The most convenient laboratory method for the preparation of hydrogen is by the action of zinc on dilute sulphuric acid. The zinc is placed in a two-necked bottle (the so-called Woulf's bottle) ; one opening is fitted with a thistle funnel, the other with a gas delivery tube (Fig. 14). Dilute sulphuric acid is poured into the funnel, and the hydrogen is collected in the usual way. It is a remarkable fact, which has not yet been adequately ex- plained, that pure dilute sulphuric acid has practically no action on perfectly pure zinc. If, however, a few drops of a solution of copper sulphate or of platinum chloride are added, the action at once becomes vigorous. Commercial granulated zinc, which contains other metals as impurities, is readily acted on by sulphuric acid. The other product formed when sulphuric acid acts on zinc is termed zinc sulphate, and may be obtained in crystals on evaporating the solution. The equation expressing the action of sulphuric acid on zinc is, therefore, as follows : Zinc + sulphuric acid = zinc sulphate + hydrogen. Hydrogen may also be obtained by the action of other acids on metals otl er than zinc, as will be shown at a later stage. In order to obtain hydrogen by the action of . sulphuric acid on zinc as re- quired, the apparatus devised by Kirjrj and represented in Fig. 15 is very con- venient and economical. It consists essentially of two glass bulbs B and C, joined by a narrow glass tube ; into the top of the bulb B a third bulb A, extended at its lower part into a long tube, is fitted airtight in such a way that the tube extends nearly to the bottom of the lower bulb C. The zinc is placed in B, and is prevented from falling down into C by the narrow neck, which is nearly, but not entirely, filled by the prolongation of the bulb A. p ]G x - With the stop-cock, D, open, dilute sul- phuric acid is poured through A into C till it reaches the zinc in B, when hydrogen is given off. On closing the stopcock, gas continues 36 A TEXT-BOOK OF INORGANIC CHEMISTRY to be given off till all the liquid is forced out of B (part of it rising into A through the long inner tube), when the action ceases automati- cally. If desired, more dilute acid can be poured into A when the stop-cock D is closed, and on opening D the gas will be given off under increased pressure, owing to the higher level of the liquid in A. This apparatus may also be used for preparing hydrogen sulphide, carbon dioxide, and other gases. (7) Although granulated zinc has very little action on water even at boiling-point, hydrogen is readily given off when water is heated with zinc coated with copper, the so-called zinc-copper couple. The couple is prepared by immersing granulated zinc for a few minutes in a dilute solution of copper sulphate ; the excess of the sulphate solution is then poured off, and the couple washed two or three times with cold water. The chemical change taking place when the couple is heated with water is represented by the equation Zinc + water = zinc oxide + hydrogen, so that the copper apparently plays no direct part in the change. The action is therefore catalytic, and is no doubt analogous to that of the same metal in accelerating the reaction between sulphuric acid and pure zinc. (8) Very pure hydrogen may be obtained by heating metallic zinc with a solution of sodium hydroxide. The reaction in this case is represented by the equation Zinc + sodium hydroxide = sodium zincate + hydrogen. Aluminium may be used in place of zinc. Physical Properties Hydrogen, like oxygen, is a colourless, odourless, tasteless gas. It is the lightest gas known, its density being about one-sixteenth that of oxygen. One litre of hydrogen, at o and 760 mm. pressure, weighs 0.089873 grams (Morley). The lightness of hydrogen may be shown very strikingly by suspending a beaker, mouth downwards, from one arm of a balance and placing weights in the other pan till the pointer is at zero on the scale. If then the contents of a jar of hydrogen are poured upwards into the beaker, the movement of the pointer will indicate that the beaker and contents weigh less than before. On account of its lightness hydrogen is employed for inflating balloons. Hydrogen was first obtained as a coherent, colourless, transparent liquid by Dewar in 1898, and somewhat later by Travers. Its boiling- HYDROGEN GENERAL PROPERTIES OF GASES 37 point on the helium scale is -252.5, and its melting-point -259 at a pressure of 49-50 mm. (Travers). Trie critical temperature of liquid hydrogen is about -241 and its -critical pressure 14 atmos- pheres (O;szewski). The density of liquid hydrogen is 0.07 : in other words its density is only one-fourteenth that of water. Hydrogen is only very slightly soluble in water. I c.c. of water dissolves at o 0.0215 c.c., at 10 0.0198 c.c., and at 20 0.0184 c.c. of hydrogen under i atmosphere pressure (Timofejeff). Chemical Properties If a jar is filled with pure dry hydrogen, and a lighted taper is applied to the mouth of the jar, the gas catches fire and burns with an almost colourless flame. After the flame has gone out, moisture will be observed on the sides of the jar, indicating the formation of water. The burning of hydrogen in air is, in fact, a chemical change in which the oxygen of the air unites with hydro- gen to form water. If care has not been taken to expel all the air from the Woulf's bottle before collecting the hydrogen, the mixture of hydrogen and air will explode when a light is brought to the mouth of the jar. The explosive character of mixtures of hydrogen and oxygen may be more strikingly demonstrated by filling a soda-water bottle by displacement with a mixture of one volume of oxygen and approxi- mately two volumes of hydrogen. The bottle is wrapped in a cloth to protect the hand in case of accident, and on applying a light to the mouth a violent explosion occurs. The gases combine under ordinary conditions only when the temperature is sufficiently high. The application of the taper to the mouth of a bottle containing the mixed gases leads to the combination of a small portion of the mixture, and the heat given out raises the rest of the mixture above the temperature of combination, so that the change is practically instantaneous. The water formed by the combination occupies a volume which is negligible in comparison with that of the gases, so that there is a partial vacuum in the bottle after combination. The noise of the explosion is due to the inrush of air to fill the bottle. The fact that water is formed when hydrogen burns in air may be illustrated by causing a jet of burning hydrogen to impinge against a cold surface, as shown in Fig. 16. The hydrogen prepared in the Woulf's bottle A is dried by passing through the U -tubes B and C, which contain anhydrous calcium chloride, and the burning jet is in contact with the flask D, which contains a large quantity ofcoldw^ter. In a few minutes sufficient water will have collected to allow of its recognition. 38 A TEXT-BOOK OF INORGANIC CHEMISTRY Hydrogen has so great an affinity for oxygen that it not only combines with the free gas, but even removes it from combination with other elements. This may conveniently be shown by heating FIG. 16. black copper oxide in a bulb tube A (Fig. 17), and passing a stream of dry hydrogen over it. In a few minutes the black powder will be observed to turn red, and finally only metallic copper remains in FIG. 17. the bulb. The chemical change in this case is represented by the equation Copper oxide + hydrogen = copper + water. The water may be collected in the tube B placed behind the bulb tube. HYDROGEN GENERAL PROPERTIES OF GASES 39 Oxidation and Reduction It has already been pointed out that the addition of oxygen to another element or to a chemical com- pound is termed oxidation, and the same term is applied when hydro- gen is removed from a chemical compound. The removal of oxygen from copper oxide by hydrogen, just described, is called reduction, and the same term is applied to the adding on of hydrogen to another ele- ment or compound. Reduction is therefore the converse of oxidation. Occlusion of Hydrogen Certain metals, more particularly platinum, palladium, and iron, possess the remarkable property of absorbing or occluding many times their own volume of hydrogen when heated, and retaining it at the ordinary temperature. The following table gives the maximum volume of hydrogen, referred to o and 760 mm., which can be retained by one volume of the finely divided metals : Palladium black Platinum sponge Gold precipitated Iron reduced 873 vols. Nickel reduced 493 Cobalt 46 Copper 19.2 Silver powder 18 vols. 153 4.8 0-95 The occluded hydrogen does not appear to enter into chemical combination with the metals, but the phenomenon is by no means well understood. No other gas is occluded by metals in general to the same extent as hydrogen. Collection of Gases by Dis- placement of Air A very light gas, such as hydrogen, may be collected by downward displace- ment of air. The delivery tube is passed r early to the bottom of the inverted jar (Fig. 18 d), and after some time it will be found that the jar is full of hydrogen. If the jar is kept inverted, the hydro- gen will be retained for some time, bu: when it is placed mouth & a upwards the gas soon escapes. FIG. 18. Similarly, a gas heavier than air (e.g. chlorine) can be collected by upward displacement of air (Fig. 18 J). 46 A TEXT-BOOK OF INORGANIC CHEMISTRY GENERAL PROPERTIES OF GASES General It is well known that matter can exist in three forms or states of aggregation, the gaseous, liquid, and solid states. Further, the particular form in which a definite substance occurs depends on the external conditions ; thus when the temperature is raised solid water or ice changes to liquid water, and at a still higher temperature water changes to steam, which is a gas. Solids have definite form, and the volume of a solid does not alter greatly when the external conditions, such as temperature and pressure, are altered. Liquids differ from solids in that they readily take the shape of the vessel containing them ; like solids they have a definite volume, which is not greatly affected by changes of temperature and pressure. Gases, on the other hand, have no definite volume; they are characterized by their tendency to fill completely, and to a uniform density, any available space. This fact cannot be illustrated very conveniently with colourless, odour- less gases, such as hydrogen and oxygen, but may readily be shown with bromine, which is red in colour. If a few drops of liquid bromine are placed, by means of a pipette, at the bottom of a tall gas jar containing air, and the jar is covered and set aside, it will be found in course of time that the gaseous bromine, which at first was only to be found at the bottom of the jar, has become uniformly distributed throughout the confined space, as shown by the colour. In giving the weights of a litre of hydrogen and of oxygen respec- tively, the temperature and pressure under which the gases were measured have been given. This implies that the volume of a definite quantity of a gas depends on the conditions under which it is measured, and such is the case. The laws expressing the behaviour of gases under varying conditions will now be briefly considered. The most remarkable fact about these laws is that they are to a great extent independent of the nature of the gas ; the volume of all gases is affected by changes of temperature and pressure to much the same extent. Relation between Volume and Pressure for Gases When the pressure upon a confined volume of gas is increased, the volume diminishes. The exact relationship between pressure and volume was discovered by Boyle (1661), and is known as Boyle's law ; the law may be formulated as follows : At constant tempera- ture the volume of a given mass of a gas is inversely proportional to the pressure to which it is subjected. This means that if the pressure HYDROGEN GENERAL PROPERTIES OF GASES 41 on a gas is doubled, the volume is halved, if the pressure is made four times as great, the volume is reduced' to one-fourth, and so on. Boyle's law may therefore be put in the alternative form, that the product of the pressure/ and volume v of a definite mass of a gas is constant at constant temperature. If v is the volume at pressure^ and ^ the volume at pressure^ then/ v = A ^i = constant. The validity of Boyle's law may be tested by means of the arrange- ment represented in Fig. 19. A straight tube, A, which must be ABC FIG. 19. more than 80 cm. long, and closed at one end, is filled with mercury and inverted in a mercury trough in such a way that no air enters. It will be found that the mercury only fills part of the tube, the upper unshaded part is a vacuum. The mercury is supported in the tube by the pressure of the atmosphere, and the difference of level a b be- tween the surface of the mercury in the trough and that in the tube measures :he pressure of the atmosphere. The average height of the column of mercury is about 760 mm. (30 inches.) The apparatus is 42 A TEXT-BOOK OF INORGANIC CHEMISTRY called a barometer. Another tube, B, has a long limb which is open to the air and a short limb which is closed. A little mercury is poured into the open end, and so adjusted that a column of air is confined in the shorter limb, while the mercury stands at the same level in both limbs. Under these circumstances the confined air is necessarily under atmospheric pressure, since the atmosphere is pressing upon the surface of the mercury in the open tube,- and the enclosed air upon that in the closed tube ; the fact that the mercury is standing at the same level in both tubes indicates the equality of the two pressures. The length of the confined column of air is then measured, and mercury poured into the open end of the tube till the difference in level between the mercury surfaces in the two limbs corresponds with one atmosphere pressure, as read off on the barometer. The enclosed air is now under two atmospheres pressure, and it will be found by measurement that its volume has been reduced to one-half, in accordance with Boyle's law. If a longer tube is taken, the validity of the law can be tested by applying still higher pressures. The apparatus with which Boyle established the law was very similar to that just described. Boyle's law is not strictly true for any actual gas, but is very nearly so for gases such as hydrogen, oxygen, and nitrogeji, which are very difficult to liquefy. Under ordinary conditions hydrogen is less com- pressible, and all the other gases more compressible, than the law indicates. Relation between Volume and Temperature for Gases When a gas is heated at constant pressure it expands. Careful experiment has revealed the remarkable fact that the increase in the volume of a gas for a given rise of temperature is a constant, independent of the nature of the gas. If a definite volume of a gas is raised i in temperature at constant pressure, it expands by 1/273 of its volume at o ; if it is raised 10 in temperature it expands by 10/273 of its volume at o. In the same way, if a definite volume of a gas is cooled i, the volume diminishes by 1/273 f its value at o. These statements are summarized in Charles's law, which may be expressed as follows : At constant pressure a gas expands or con- tracts by 1/273 of its 'volume at o for every change in temperature of i C. This law may be put ir a more concise form on the basis of the following considerations. Imagine a gas confined in a tube graduated in c.c. (Fig. 20) by the air-tight weightless piston (shaded in the diagram) which moves without friction in the tube, and that the HYDROGEN GENERAL PROPERTIES OF GASES 43 volume of the gas occupies 273 graduations of the tube at o C. If now the te nperature of the gas is raised i a , the volume increases by 1/273 of that at o, that is, by one graduation, and the new volume is 274 c.c. Similarly if the temperature is raised to / \ 100, the new volume is 273 i 273/ ; 37C 310 "A ^~ ) = 373c.c., and if the temperature is lowered to 100, the new volume is 273 100 273 173 c.c. It is evident, therefore, that the arrangement in question might be used as a thermometer for measuring the tem- perature, as every change of temperature of x 260CC degrees corresponds with a change of .r graduations on the scale. Further, if the same rule continues to hold, the volume of the gas will theoretically be zero when the temperature has fallen ^-273 C. and no lower temperature can be registered by toocc our gas thermometer. These considerations have led to the establishment of a new scale of tempera- faocc ture, the so-called absolute scale, the zero on the absolute scale being -273 C. The temperatures /60CC on the absolute scale, the so-called absolute tem- peratures, are shown on the right hand of the scale, and it is evident that absolute temperatures are /20CC obtained by adding 273 to the corresponding tem- peratures on the Centigrade scale. It is now /O Q CC clear that our constant pressure gas thermometer measures absolute temperatures, or to put the socc matter in another way: At constant pressure, the volume of a gas is proportional to the absolute eocc temperature. This is the form of Charles's law which is most useful in calculations dealing with gaseous volumes. The method of employing it will now be illustrated by two examples. (i) If the temperature of a quantity of gas which measures 100 c.c. at 10 is changed to 130 at con- stant pressure, what is the new volume ? As the volume is proportional to the absolute temperature, 273 c.c. at o C. will measure 283 c.c. at 10 and 273 + 130 = 403 c.c. at 130. Hence 100 c.c. at 10 will measure 100 x ~cr = 142.4 c.c. at 130. 20CC 10 CC occ 17 C Z90A 7" C MO /? 0C 273 A -I3C 2GOA -23C 250A - -33C 240 /? -53 C 220* -7JC 200A -93C I80 A ~I33C 140 A I73C /00A -/93C 80A 2/3C 60* -233C 40A -ZS3C 1QA -263C /0A -Z73C 0A FlG. 20. 44 A TEXT-BOOK OF INORGANIC CHEMISTRY (2) What will be the volume of the same quantity of gas if measured at - 80 C. ? 283 c.c. at 10 C. will measure (273 - 80) = 193 c.c. at - 80. Hence the volume at 8o ( must be 100 x -^ = 68.2 c.c. 283 The important point to be observed in making the above correc- tions is the proper placing of numerator and denominator. If the volume at a lower temperature is required, the smaller absolute temperature is of course used as numerator, as the final volume is smaller than the initial volume ; if the volume at a higher tem- perature is required, the greater absolute temperature is used as numerator. The expansion which unit volume of a gas shows when raised i c in temperature at constant pressure may be termed its coefficient of ex- pansion, and it has been pointed out in this section that the coefficient of expansion of all gases is approximately the same, and is equal to 1/273 or 0.003665 of its volume at o. The coefficient of expansion is often indicated by the letter a, and if v t is the volume of the gas at / C. and VQ its volume at o C. then v t = v (i + /) = ^ +0.003665 /) according to the law of Charles. This equation may be used to correct gases for changes of temperature, but is not so convenient as the absolute temperature method already described. So far, it has been assumed that the coefficient of expansion of all gases at constant pressure is exactly the same. This, however, is by no means the case, the law of Charles, like that of Boyle, being only an approximate one. The deviations from the simple law depend both on the nature of the gas and on the temperature at which the observations are made, and in all cases the coefficient approaches the value 0.003665 the more nearly the lower the pressure, and the further the gas is removed from its temperature of liquefaction. This is shown by the following table Coefficients of Expansion of Gases at Constant Pressure Gas. Pressure at o in cm. Mercury. Coefficient of Expansion. Gas. Pressure at o in cm. Mercury. Coefficient of Expansion. Hydro- 76 cm. 0.0036613 Carbon 76 cm. 0.003710 gen diox- 252 0.003845 , , 254 ,. 0.0036616 ide. Air 76 0.003671 " 257 ,, 0.003695 HYDROGEN GENERAL PROPERTIES OF GASES 45 General Equation for Gases So far, we have investigated the relationship between the volume and pressure of a gas at con- stant temperature (Boyle's law), and between the volume and temperature of a gas at constant pressure (Charles's law). It still remains to discuss the relationship between pressure and tempera- ture at constant volume. A little consideration shows, however, that the law in question can readily be deduced when Boyle's law and Charles's law have been established. Suppose, for example, we have a definite quantity of a gas contained in a vessel of fixed capacity at 273 abs., and raise its temperature to 546 abs, ; if free to expand, its volume would be doubledl at constant pressure. As, however, its volume is kept con- stant, its pressure, according to Boyle's law, must be doubled. From these considerations we deduce the third of the simple gas laws : At constant volume, the pressure of a gas is proportional to the absolute temperature. Like the other laws the above law is only approximately true. It can readily be shown that the three gas laws can be expressed in the simple formula _ = constant, 1 o 1 1 where p Q and v are the pressure and volume of a definite quantity of gas at the absolute temperature 7" , and p l and v^ the pressure and volume at the absolute temperature T r The use of this formula in finding the volume of a gas when both temperature and pressure vary may be illustrated by the following example. A definite quantity of a gas measures 500 c.c. at 20 and 500 mm. pressure, what is its volume at - 30 and 900 mm. pressure ? Substituting in the general equation, we have 500 x 500 _ 900 x v\ 273 + 20 273-30 Whence ^goox 500x243^ cc . 900 x 293 The same result may also be obtained by correcting first for the change of pressure by Boyle's law, the new volume being then corrected for change of temperature by Charles's law. The general 46 A TEXT-BOOK OF INORGANIC CHEMISTRY formula should not be used until the student is also familiar with the latter method of calculation. The general formula, like the laws of which it is a summary, is only approximately valid, but is the more nearly true the higher the temperature, the lower the pressure, and the further the gas is removed from its temperature of liquefaction. The theoretical bearing of this statement is discussed later. Diffusion of Gases It has already been stated that when two gases are brought together they ultimately become uniformly mixed, even against the force of gravity. This may be shown by a modification of the experiment described on p. 40. If a few drops of bromine are carefully placed, by means of a f^. pipette, at the bottom of a jar of hydrogen, and the jar is covered, it will be found after a short time that the colour is equal throughout the space, showing that the bromine gas or vapour, although eighty times heavier than hydrogen, has become uniformly distributed through it. The same fact can be illustrated by placing a covered jar of hydrogen mouth to mouth over a covered jar of the heavy greenish-yellow gas chlorine, and care- fully withdrawing the covers. 1 After a time it will be noticed that the chlorine has become distributed through the hydrogen in the upper jar, and on bringing a light to each jar an explosion will occur, showing that both contain the two gases in con- siderable proportion. This process of mixing is termed the diffusion of gases. The same pheno- menon is observed when two gases are separated by means of a porous plate ; in this case it may be assumed that the gas particles readily pass through the pores, so that diffusion is not hindered. Experiment shows that gases diffuse at very different rates, and that a gas diffuses the more rapidly the lower its density. This may be shown very instructively by means of the apparatus re- presented in Fig. 21. A long glass tube, the upper end of which is fixed, by means of a cork, into an inverted porous pot A (a battery jar), is fitted with a cork into one of the openings of a Woulf's bottle C, which is partly filled with water. Into 1 As the gases combine explosively in sunlight, this experiment must be per- formed in diffused daylight. } 1 1 1 UJ FIG. 21. HYDROGEN GENERAL PROPERTIES OF GASES 47 the other opening of the bottle is fitted a cork carrying a short glass tube, the lower end of which dips in the liquid in the bottle, the upper end being drawn out to a point. If now a beaker, B, is placed over the porous pot and hydrogen is passed up into B, diffusion of the hydrogen into the pot will take place much more rapidly than the air diffuses outwards, the pressure inside the pot is considerably increased, and water is forced out at the narrow end of the short tube in the form of a jet. If now the beaker with the hydrogen is removed, the jar is again surrounded with air, the mixture of gases in the pot (mainly hydrogen) diffuses out more rapidly than the air enters, the pressure inside the pot diminishes, the water falls in the long tube, and bubbles of air are drawn into the bottle. The Law of Gaseous Diffusion The law expressing the relative rates at which gases diffuse was first established by Thomas Graham, and may be formulated as follows : The relative rates of diffusion of two gases are inversely proportional to the square roots of their densities. If the velocity of diffusion is represented by v, and the density of the gas by d, Graham's law can be expressed by the formula In order ':o illustrate the law, we may consider the relative rates at which hydrogen (density i) and chlorine (density 35.5) diffuse through a porous partition into air. The ratio of the square roots of the densities is V^v/SS-S* approximately 1:6, hence the rela- tive rates of diffusion, being inversely as the square roots of the densities, are as 6:1. The res alts obtained by Graham, and given in the following table, indicate that the experimental results are represented very satis- factorily indeed by the above formula. Gas. Density, d, compared with Air as Unity. w, Observed Rate of Diffusion compared with Air as Unity. Hydrogen . . Nitrogen . Oxygen . Carbon dioxide Sulphur dioxide 0.0695 0.9713 1.1056 1.5290 2.247 3-794 1.014 0.951 0.809 0.667 3-83 1.014 0.949 0.812 0.68 48 A TEXT-BOOK OF INORGANIC CHEMISTRY Not only may the rate of diffusion of a gas be calculated from its density, but the density may be determined indirectly by ob- serving the rate of diffusion. This principle has been found useful on more than one occasion in chemical investigations. The differences in the rates of diffusion of gases have been taken advantage of to effect a partial separation of the constituents of a gaseous mixture a process known as atmolysis. If, for instance, a mixture of hydrogen and oxygen is passed through a long tube of porous material, such as a series of tobacco pipe stems, and the issuing gas is collected over water, it will be found to be much richer in oxygen than the original mixture. The rate of passage of gases through a very small hole in a plate (preferably a platinum plate) was also investigated by Graham, and was found to follow the same law as gaseous diffusion. The pheno- menon is termed gaseous effusion. The Kinetic Theory of Gases It is natural to consider whether any mental picture of the nature of gases can be sug- gested which may serve to account for the simple laws which have been found to represent their behaviour, and also to account for the deviations from these laws. Such a mechanical representation was brought forward by Bernoulli as far back as 1738, and has been developed by later workers into the kinetic theory of gases. According to the theory, gases are made up of small, perfectly elastic particles (the chemical molecules, p. 106), which are in continual rapid motion, colliding with each other and with the walls of the containing vessel. The particles of any one gas are supposed to be identical, but differ from the particles of other gases in respect to mass, speed, and other properties. The space actually filled by the gas particles is supposed to be smaller than that which they inhabit under ordinary conditions ; they have, therefore, a comparatively large free space in which to move, and are practically free from each other's influence except during a collision. The average distance over which a particle moves before colliding with another particle is termed the mean free path of the particle. According to this theory, the pressure exerted by a gas on the walls of the containing vessel is due to bombardment by the moving particles. It io evident therefore that the magnitude of the pressure must depend on the mass and the velocity of the particles. It can be shown that -the pressure exerted by a single particle is proportional to its mass and to the square of its velocity, HYDROGEN GENERAL PROPERTIES OF GASES 49 and the total pressure is the sum of the. pressures exerted by each particle. It should, however, be remembered that, owing to collisions and for other reasons, the speed of the particles in any one gas is probably by no means uniform, but varies considerably about a mean value. It will now be shown that the simple gas laws are in full accord with the view as to the constitution of gases just stated. If at constant temperature the volume in which a definite mass of gas is confined is halved, the number of impacts on the walls in a given time is doubled ; in other words, the pressure of the gas is doubled. This is Boyle's law, that the product of the pressure and volume of a given mass of gas is constant at constant temperature. Further, we have seen that at constant volume the pressure of a given mass of gas is proportional to the absolute temperature. As increase of temperature cannot alter the number of the particles, the observed increase of pressure must, according to the kinetic theory, be due to an increase in the speed of the particles, resulting in a greater number of impacts on the walls of the vessel in a given time. We have already learnt that the pressure of a gas is proportional to the square of the rectilineal velocity of the particles, hence it follows that the square of the velocity of the particles is proportional to the absolute temperature. These considerations throw an interesting light on the physical meaning of the absolute zero. As the temperature falls, the velocity of the panicles steadily diminishes and finally, at the absolute zero, they must theoretically come to rest. The absolute zero is therefore the lowest temperature theoretically attainable. In actual practice it has never been reached, but in his recent investigations on the liquefaction of helium (p. 210), Kammerlingh Onnes has got within 2 of the absolute zero. On the other hand, there is no theoretical upper limit to the speed of the particles, and therefore no upper limit of temperature. Not only does the kinetic theory of gases afford a satisfactory interpretation of the gas laws, but it also accounts for the more important deviations from these laws. In deducing Boyle's law, we have tacitly assumed that when the volume is halved the space in which the particles move has been halved ; in other words, we have neglected the space filled by the particles in comparison with that which they inhabit. Further, we have made no allowance for a possible attraction between the particles. If such an attraction exists, it mist be the greater the nearer the particles approach each 4 50 A TEXT-BOOK OF INORGANIC CHEMISTRY other. Hence, when the pressure is doubled, the effect of the attraction, which must lead to a diminution of volume, is super- imposed on the regular contraction according to Boyle's law, so that the value of ^, and hence the product, pv, is less than the calculated value. On the other hand, a little consideration shows that the effect of the finite size of the particles is such that the value of pv tends to increase with increasing pressure. Let us assume that a particle is moving backwards and forwards between the walls a and b and that the distance between the walls is 20 times I 1 , the diameter of the particle. The dis- | tance traversed between each impact is clearly 19 diameters. If now the distance between the walls is halved, the distance traversed between each impact is only 9 diameters, and therefore the number of impacts and the pressure is more than doubled by halving the volume. It is evident from the foregoing that the two causes which bring about the deviations from the gas laws act in opposite directions. For most gases, the effect due to the attraction of the particles is greater at ordinary pressures than that due to the finite size of the particles, and therefore pv diminishes at first as the pressure is increased (p. 42), but at high pressures increases with the pressure. For hydrogen, however, the effect of the volume correction counter- balances from the first the attraction correction, and pv increases continuously with the pressure. The kinetic theory also accounts satisfactorily for the fact that the deviations from the gas laws are the smaller the lower the pressure. Under these circumstances, the attraction between the particles becomes negligible owing to their distance apart, and the volume of the particles is negligible in comparison with the total volume. As already mentioned, no actual gas behaves exactly according to the simple gas laws, but the deviations in the case of hydrogen, helium, nitrogen, and other gases which are difficult to liquefy are very slight at ordinary temperatures and pressures. A gas which would follow the gas laws accurately is termed a perfect or ideal gas. In the foregoing it has been implicitly assumed that the particles themselves are incompressible ; the diminution in volume on com- pressing a gas is ascribed entirely to a diminution in the free space between the particles. CHAPTER VI WATER PHYSICAL PROPERTIES OF LIQUIDS T Tistory Water was for a long time regarded as an element. The discovery that it is in fact a chemical compound of "inflammable air" (hydrogen) and of " dephlogisticated air" (oxygen) was made by Cavendish in 1781, and he showed at the same time that the leases combine in the proportion of two volumes of hydrogen to one volume of oxygen. The method employed by Cavendish was briefly as follows. A large graduated vessel was filled with a gaseous mixture containing approximately one volume of oxygen to two volumes of hydrogen, and was connected by a bent tube to a glass globe provided with a brass stopcock and with two sealed-in platinum wires, between which an electric spark could be passed. The glass globe, which had previously been exhausted by the air-pump, was filled with the gaseous mixture, and the latter then exploded by a spark. It was then noticed that the walls of the previously dry globe were covered with moisture, showing the formation of water. The globe was again filled from the graduated vessel and the mix- ture exploded, the process being repeated until sufficient water was collected to put its identity beyond doubt. As practically no gas remained in the apparatus after the experiment, the evidence was conclusive that two volumes of hydrogen and one volume of oxygen combine to form water. Cavendish's discovery as to the composition of water was confirmed by Lavoisier (1783), who passed steam through a heated iron tube and collected and measured the hydrogen. Decomposition of Water into its Elements The methods by which it may be proved that water is a compound substance have already been considered when dealing with hydrogen. It has been shown that hydrogen is readily obtained from water when the latter is brought into contact with a substance having a greater attraction for oxygen than the latter has for hydrogen. On the same principle, free oxygen may be obtained from water by bringing it in contact with a substance which readily enters into chemical combination with 51 52 A TEXT-BOOK OF INORGANIC CHEMISTRY hydrogen. A suitable substance for the purpose is gaseous chlorine. When steam and chlorine are passed through a red-hot tube, free oxygen and a chemical compound of hydrogen and chlorine are formed, and the oxygen may be collected over water in the usual way. This important action is further referred to in connexion with chlorine (P- 87). The direct decomposition of water into its elements by the action of heat or of electrical energy has also been referred to. The amount of decomposition which steam undergoes on heating increases with the temperature, and quite recently the variation of the decomposition with the temperature has been determined with considerable accuracy. Some of the results are given in the accompanying table Temperature . . 1027 Amount of decomposi- > Q tion per cent. ) 1207 1288 i882 c 1.18 1984 1.77 The reaction is a reversible one, being represented by the equation water^hydrogen + oxygen, and the above results show that increase of temperature favours the reaction represented by the upper arrow. The Composition of Water by Synthesis The exact ratio by volume in which hydrogen and oxygen unite together to form water can readily be determined by a slight modification of Cavendish's original method. The apparatus for this purpose is illus- trated in Fig. 22. It consists essentially of a graduated tube, called a eudiometer ; provided at its upper end with two platinum wires sealed through the glass, the ends inside the tube being so placed that an electric spark may be passed between them. The tube is first filled with mercury, inverted in the mercury trough, and a few c.cs. of pure oxygen, ob- tained by heating potassium permanganate (p. 22), is passed up into it. The volume of the oxygen is carefully read off on the graduated scale. In order that the volume of the oxygen may be ascertained under conditions comparable with that of the hydro- gen, its temperature and pressure must be known ; the temperature is read off on the thermometer in the neighbourhood of the eudiometer, and the pressure is obtained by subtracting the FIG. 22. WATER PHYSICAL PROPERTIES OF LIQUIDS 53 height of the column of mercury in the eudiometer tube from that of the barometer at the time of the experiment. Since the pressure of the atmosphere is balanced by the pressure of the oxygen + that of the mercury column in the tube, it is evident that the pressure of the oxygen is obtained as just stated. The volume of the oxygen is then corrected to normal temperature and pressure by the method already describee. A quantity of hydrogen, the corrected volume of which is four to five times that of the oxygen, is then introduced, the volume of the mixed gases ascertained, and corrected to normal temperature and pressure as before. The tube is then firmly pressed down below the mercury in the trough upon a plate of caoutchouc, and the gases ignited by a spark from an induction coil. A bright flame is seen to pass down the tube, and on cautiously raising the end of the tube from the pad, a con- siderable amount of mercury will enter. After the temperature has fallen to that of the atmosphere, the volume of the residual gas (un- combined hydrogen) is read off and reduced to normal conditions. The mode of calculating the combining volumes of oxygen and hydrogen from these data is best illustrated by an example. Corrected volume of oxygen .... 31.82 c.c. mixed gases . . . 185.75 c.c. residual hydrogen . . 90.35 c.c. Hence Total volume of hydrogen used 185.75-31.82=153.93 c.c. Volume of hydrogen which has ) 153.93-90.35 = 63.58 c.c. combined with oxygen ) Therefore Volume of Oxygen : Volume of Hydrogen as 31.82 : 63.58 or i : 1.998. In accurate experiments a number of corrections must be applied which have not been referred to in the above brief sketch of the process. For example, the volume of the water formed, although very small in comparison with that of the gases, is not entirely negligible. According to Morley, who employed all conceivable precautions, the true combining ratio is i : 2.00269 at o. Volumetric Composition of Water by Analysis The ratio in which hydrogen and oxygen are present in water can also be determined by electrolyzing water between platinum electrodes 54 A TEXT-BOOK OF INORGANIC CHEMISTRY and collecting the gases separately (p. 14). The volume of the oxygen is always rather less than half that of the hydrogen for two reasons: (i) oxygen is more soluble in water than hydrogen; (2) a denser modification of oxygen, termed ozone, is formed in small amount at the anode. Volumetric Composition of Steam If the apparatus in which the mixture of hydrogen and oxygen is exploded be kept at such a temperature that the water produced remains in the form of FIG. 23. steam, it is found that two volumes of hydrogen and one volume of oxygen give rise to two volumes of steam when all the gases are measured under the same conditions. This may readily be shown by means of the apparatus represented in Fig. 23. It consists essentially of a U-tube, one limb of which acts as a eudiometer tube, the other is open to the air. A mixture of hydrogen and oxygen in the propor- tions to form water, prepared by electrolysis, is introduced into the eudiometer tube, which is surrounded by a wide tube through which the vapour of a high-boiling liquid may be passed. When the vapour from the boiling liquid in the flask has been passed round the eudiometer tube sufficiently long to ensure that the temperature is constant, the pressure is adjusted to that of the atmosphere by altering WATER PHYSICAL PROPERTIES OF LIQUIDS 55 the quantity of mercury in the open limb, the volume of the mixed gases is carefully noted, and the mixture then exploded by an electric spark. When the pressure is again adjusted to that of the atmosphere by pouring mercury into the open limb until the level is the same in the two cubes, it will be found that the volume of the gas in the eudiometer has been reduced by one-third. Amyl alcohol, boiling-point 132, may conveniently be used as jacketing vapour ; it is condensed on leaving the tube, as shown in the figure. Steam under increased pressure may also be used for the same purpose, although less advantageously. Gravimetric Composition of Water The composition of water by weight may be determined (a) from the densities and com- bining volumes of the gases (&) directly. (a) According to Morley, the relative densities of hydrogen and oxygen are as I : 15.9 and the combining volumes 2.0027: i. Hence the relative weights of hydrogen and oxygen which combine to form water are Hydrogen 2.00269 x i = 2.0027. Oxygen 1x15.900 = 15.900. 17.9027. Hence 17.9027 parts of water contain 2.0027 parts of hydrogen and 15.900 parts of oxygen ; otherwise expressed, water contains 11.186 per cent, of hydrogen and 88.814 P er cent, of oxygen. (b} Direct Method The method which has been most largely employed for the direct determination of the composition of water by weight depends on the reduction of copper oxide (p. 38). A known weight of the oxide is heated and pure dry hydrogen passed over it ; the water formed is collected in suitable vessels which are weighed before and after the experiment. The bulb containing the copper oxide is also weighed after the experiment in order to find the weight of oxygen which is contained in the amount of water formed. The weight of the hydrogen is determined by difference. An elaborate investigation of the composition of water by this method was carried out by Dumas and Stas in 1843. The arrange- ment used is illustrated in Fig. 24. Hydrogen, generated from zinc and sulphuric acid in the bottle A, and carefully purified and dried by means of the reagents in the U -tubes, is passed over the thoroughly dried copper oxide heated in B, the water formed being collected in the vessel C and the three succeeding drying tubes. From the weight of the vessel and drying tubes before and after the experiment, the 56 A TEXT-BOOK OF INORGANIC CHEMISTRY WATER PHYSICAL PROPERTIES OF LIQUIDS 57 weight of the water formed was obtained, and from the loss in weight of the copper oxide tube the quantity of oxygen which went to its formation. As a result of their experiments, Dumas and Stas found that i part by weight of hydrogen combines with 15.96 parts by weight of oxygen, a result which does not agree very well with modern results. In the hands of W. A. Noyes, however (1889), this method yielded results very close to the value now generally accepted. An improve- ment subsequently introduced was to weigh the hydrogen as well as the oxygen and the water ; this was effected by absorbing a consider- able weight of hydrogen in palladium, heating to drive off the gas, which was then led over the copper oxide, the tube containing the palladium being weighed before and after the experiment. Morley (1895) weighed the hydrogen occluded in palladium, and oxygen in ihe gaseous form, burned the hydrogen in the oxygen, and weighed the water formed. The result obtained by Morley, which is practically the mean of the most trustworthy results of other observers, is that water contains 2 parts by weight of hydrogen to 15.879 parts by weight of oxygen, or 11.186 per cent, of hydrogen to 88.814 per cent, of oxygen, in exact agreement with the result obtained from the densities and combining volumes. The fact established in the foregoing sections, that pure water, whatever its source or mode of preparation, has invariably the same composition, is a matter of the utmost importance. We shall see later that the same is true of all other definite chemical compounds. This important result is known as the law of constant composition, and may be stated as follows : A definite chemical compound always contains th? same elements in the same proportions by weight. Physical Properties of Water Water at ordinary tempera- tures is an odourless, tasteless liquid, colourless in thin layers, but in thick layers showing a slight bluish-green colour. The colour is most readily seen by looking at a white object through a column of pure water several yards in length, contained in a tube with blackened sides. The striking blue colour of certain Swiss lakes, fed by glacier streams, is probably due to the intrinsic colour of the water. It should, however, be remembered that a very finely-divided solid suspended in water also gives rise to a blue colour under certain conditions. As is well known, the freezing-point of water is taken as the zero point, o, and the boiling-point under 760 mm. pressure as 100 on the Centigrade scale. 58 A TEXT-BOOK OF INORGANIC CHEMISTRY The compressibility of water is very small. 20,000 volumes are reduced to 19,999 volumes by increasing the pressure by one atmos- phere. The effect of change of temperature on the volume of water is also comparatively small, as is shown in the accompanying table, in which the volume and relative density (specific gravity) of water at temperatures from o to 20 are given, referred to water at 4 as unit. Tempera- ture. Volume. Rel. Density. Tempera- ture. Volume. Rel. Density. O .OOOI22 0.999878 5 . 000008 0.999992 I .000067 0-999933 6 .000031 0.999969 2 .000028 0.999972 8 .000118 0.999882 3 . 000007 0.999993 10 .000261 0.999739 4 .000000 1. 000000 20 .001730 0.998270 The figures in the table show the remarkable fact that water attains its greatest density at a temperature in the neighbourhood of 4. In other words, when water is warmed from o, it contracts till the temperature reaches 4 (accurately 3.945) and beyond that point expands with rising temperature. Water expands on freezing, one volume of water at o giving 1.09082 volumes of ice at the same temperature. This fact is of great import- ance in effecting the disintegration of rocks, the water penetrating into cracks and exerting enormous pressure on solidification. The burst- ing of water-pipes in winter is due to the same cause. Water is a very bad conductor of heat. If a piece of ice be held in the lower part of a test-tube by a piece of wire gauze, the water in the upper part of the tube may be boiled without melting the ice. The three properties of water just mentioned are of enormous importance in Nature. When a large surface of water, such as a lake, is subjected to a low temperature the upper layers are first cooled ; they become denser and sink to the bottom. A continuous circulation is thus set up, the cooler layers sinking and the warmer rising, until the whole mass has fallen to 4. On further cooling, the upper layers become lighter and remain on the top till they solidify. As, however, ice is less dense than water, it remains on the surface, so that only the upper layers solidify. The lower layers are further pro- tected against cooling by the very small conductivity of the upper layers. If it were not for the factors just enumerated, more particu- larly the existence of a point of maximum density for water above zero, the mass of water would solidify as a whole, the heat of summer would be quite insufficient to melt the accumulations of ice formed in WATER -PHYSICAL PROPERTIES OF LIQUIDS 59 winter, and the climate of a large part of Europe would approach that of the Polar regions. The latent heat of fusion of ice is about 80 calories 1 per gram, in other words, 80 calories must be supplied in order to change one gram of ice at o to water at the same temperature. The latent heat of vaporization of water is 537 calories, that is, 537 calories must be supplied in order to change one gram of water at 100 to steam at the same temperature. Natural "Waters Owing to the very great solvent power of water, it is never found pure upon the earth, but always contains dissolved substances in larger or smaller amount. According to the amount of dissolved substances they contain, natural waters are roughly classified into (i) Fresh waters, in which the proportion of substances in solution is relatively small ; (2) Mineral waters, in which the dissolved impurities are perceptible to the taste. Fresh waters are sometimes divided, according to their origin, into rain, river and spring waters, and each of these classes will now be briefly considered. Rain Water is the purest form of natural water. The only impurities it contains in appreciable amount when collected in the country are ammonium salts, nitrates, and traces of organic matter, which it takes up from the atmosphere. Rain water collected in towns is much less pure, owing to contamination of the atmosphere; it often contains traces of free sulphuric acid, resulting from the slow oxidation of sulphur compounds. The amount of solid matter in rain- water collected in the country is very variable, being greater at the beginning than at the end of a shower owing to the gradual removal of the impurities from the atmosphere. The average amount of dissolved matter does not exceed 0.03 to 0.04 parts per 1000 of water (0.03 to 0.04 grams per litre). Spring Waters are always less pure than rain water, owing to their solvent action on the strata through which they pass. The nature of the dissolved substances naturally depends upon the chemical composition of the strata, but most fresh spring waters contain the sulphates, carbonates, chlorides and silicates of mag- nesium, calcium, iron, potassium and sodium in very varying propor- tions, and also dissolved gases, more particularly carbon dioxide. The proportion of salts in spring waters varies from 0.05 to 3 grams per litre. As already indicated, spring waters containing an 1 A calorie is the amount of heat required to raise i gram of water i in temperature. 60 A TEXT-BOOK OF INORGANIC CHEMISTRY exceptionally large proportion of dissolved substances are termed mineral waters. River Waters usually contain less dissolved matter than spring waters ; but, owing to the fact that rivers are mainly fed by surface drainage, they contain more organic matter, suspended and dissolved, than spring waters. The mineral substances in river waters consist largely of calcium salts, more particularly the carbonate and sulphate. The proportion of salts varies from 0.05 to 1.6 grams per litre. The average amount of total solids per litre in some well-known rivers is as follows : Neva, 0.055 5 Dee a * Aberdeen, 0.057 ; Thames (upper part), 0.307 (lower part) 1.617 ; Nile, 1.580. The composition of the water of lakes varies enormously. The water of Loch Katrine, from which the city of Glasgow is supplied with water, contains about 0.03 grams of solid matter per litre, that of the Dead Sea 240 grams per litre. Mineral Waters Sea water, though not usually classed as a mineral water, may appropriately be mentioned here. When collected far from land, the composition of sea water is very constant ; it con- tains about 36 grams of solid matter per litre. The chief salt present is sodium chloride, but magnesium, potassium, and calcium salts are also present, mainly as chlorides, bromides, and sulphates. In the water of the Irish Sea the proportion of the more important salts in grams per litre is approximately as follows (Thorpe, 1870) : sodium, as chloride, 27 grams ; magnesium, as chloride, 3.2 grams, as sulphate, 2 grams, as bromide, 0.07 grams ; potassium, as chloride, 0.75 grams ; calcium, as sulphate, 1.4 grams. The typical mineral waters are classified according to their most important constituents. The chief types are as follows : carbonated waters, which contain a large proportion of dissolved carbon dioxide examples, Seltzer, Apollinaris ; (2) alkaline waters contain much sodium bicarbonate example, Vichy ; (3) saline waters contain other salts than sodium bicarbonate. Chalybeate waters contain iron salts in solution; sulphur waters contain hydrogen sulphide and alkali sulphides, as at Harrogate ; the wells at Epsom contain chiefly magnesium sulphate, the .well-known Epsom salts, etc. Certain of the mineral waters are cold as they escape from the ground, others, as at Carlsbad, are hot. Potable Waters Water used for drinking purposes should be colourless, odourless, and free from materials injurious to health. The salts usually found in river and spring waters are not detri- mental ; in fact the presence of them in small proportion is advan- WATER -PHYSICAL PROPERTIES OF LIQUIDS 61 tageous. The contamination of drinking water most to be feared is the presence of the germs of various diseases such as cholera and typhoid fever, which often reach water in sewage ; and the organic substances present in sewage may also yield dangerous products. It is evident therefore that bacteriological examination of water used for drinking purposes is essential. A chemical examination is also of importance, not so much on account of danger from salts that may be present, but because the presence of certain constituents in large proportion indicates comparatively remote contamination with sewage. The presence of appreciable amounts of nitrogen, com- bined in complex organic compounds (albuminoid nitrogen), indicates recent contamination. In course of time, however, this form of nitrogen becomes oxidized to nitrates and nitrites, and the pre- sence of these salts, of sodium chloride, and of free ammonia, above a certain small proportion in a sample of water is usually a sign of previous contamination. It often happens that river and spring waters not entirely free from pollution have to be used for drinking purposes. Such waters may be rendered safe by boiling, which destroys bacteria, but the same object is generally attained on the large scale by filtration. This process consists in passing the water through beds of sand or gravel witr free exposure to air, by which means the bacteria are retained and the organic matter oxidized to comparatively harmless substances. The filter-beds must of course be renewed from time to time. Ozone is now sometimes employed for the purification of water, as it rapidly oxidizes organic matter and can afterwards be readily removed. On the small scale, a filter of porcelain, the so-called Pasteur- Chamberland Filter, may be used ; in this case the water has to be forced through under pressure. Filters of powdered charcoal are also in use. Distillation of Water Water can be freed from most of the ordinary impurities by converting it into steam, which is then passed through a cooling arrangement and condensed again to water. This process is known as distillation. An apparatus which can be used for this purpose is shown in Fig. 25. The water is boiled in the flask A, and the steam passes along the inner tube of the " Liebig's condenser'' B, where it is cooled by a stream of cold water passing through the outer tube of the condenser. The condensed water collects in the flask C, which is termed the receiver. The first 62 A TEXT-BOOK OF INORGANIC CHEMISTRY portion of the distillate, which contains most of the dissolved gases, is rejected, and the last portion of the water, which contains the non-volatile impurities, is left in the distilling flask, so that the dis- tilled water is much purer than the original sample. Steam has, however, a slight solvent action on the glass of the condenser, and in order to avoid this source of contamination, condensers made of materials not acted on by steam, such as tin or platinum, are largely used. Water may also be purified by partial freezing, the impurities FIG. 25. remaining in the fluid portion. The ice is separated, allowed to melt, and the process repeated if necessary (Nernst). The best method for determining the purity of a sample of water is to measure its electrical conductivity (p. 260). Distilled water has a flat taste, owing to the absence of the dissolved gases (chiefly air and carbon dioxide) which impart to ordinary water a refreshing taste. SOME GENERAL PROPERTIES OF LIQUIDS As water is in many respects a typical liquid, it will be con- venient to deal here with certain general properties of liquids, WATER PHYSICAL PROPERTIES OF LIQUIDS 63 illustrated mainly by reference to the properties of water. The great majority of pure substances which exist in the liquid form can also be obtained in the solid and gaseous forms, and we are chiefly concerned with the conditions under which the change of one form into the other takes place, the phenomena accompanying these transformations, and the conditions under which two or more forms of a substance can exist together. The Change of Liquid to Vapour. Equilibrium It is a well-known fact that water and all other pure liquids, under suitable conditions, tend to give off vapour, and if the space into which the vapour escapes is sufficiently large, a given amount of a liquid may be changed completely to vapour. The process of transformation of a liquid into a vapour is known as evaporation or vaporization. The phenomenon may be studied most advantageously when vaporization takes place into a confined space. Three glass tubes, A, B, and C, closed at one end, and about a metre long, are filled with mercury and inverted in a bath of the same metal (Fig. 26). The mercury stands at the same level in each tube. If now, by means of a bent pipette, a little water is introduced into the tubes B and C, the mercury will be depressed. Part of the water at once evaporates, and the vapour exerts a pressure which is measured by the difference of level of the mercury in these tubes and in the com- parison tube A. If one of the tubes is lowered a little in the bath so as to diminish the space above the mercury some vapour at once condenses ; if the tube is raised a little so as to increase the space over the mercury more vapour is formed, and in both cases the vapour pressure regains the original value. When a liquid is in contact with its own vapour under such conditions, the space is said to be saturated with the vapour. It follows from the experi- ments just described that at constant temperature water exerts a definite vapour pressure, which can be reached both from a more saturated and from a less saturated state, and which is independent of the amount of liquid present. The above results are best expressed by the statement that in a confined space water rapidly attains a state of equilibrium with its vapour ; at constant temperature the space above the liquid contains a definite amount of vapour per unit volume, which therefore exerts a definite pressure termed the vapour pressure of the liquid. The vapour pressure of all liquids increases as the temperature rises. If ihe tube C is jacketed and heated by means of steam from a boiler, t will be noticed that the mercury is more and more 64 A TEXT-BOOK OF INORGANIC CHEMISTRY depressed as the temperature rises, and finally, when the whole is heated to 100, the mercury stands at the same level outside and inside the tube. This indicates that at 100, the temperature at which water boils, its vapour pressure is equal to the vapour pressure FIG. 26. of the atmosphere. The same is true for all other liquids, and hence the boiling-point of a liquid is that temperature at which its vapour Pressure is equal to atmospheric pressure. In the following table the vapour pressure of water is given in mms. mercury for a few tempera- tures between 10 and 150 : WATER PHYSICAL PROPERTIES OF LIQUIDS 65 TempenUure. Vapour Pressure. Temperature. Vapour Pressure. -IC 2.16 mm. + 4 54.97 mm. + c 4-58 ,, + 50 92.17 ,. + 10 9-i7 .. + 70 233-8 ,, + 20 17-41 .. + 100 760.0 ,, + 3 31-56 ,, + 150 3S 8 o-o As is \vell known, the boiling-points of liquids, that is, the tem- peratures at which they exert a vapour pressure equal to that of the atmosphere, are very different. Thus, of the pure substances already mentioned, liquid hydrogen boils at -252.5, liquid oxygen at 182.5, sulphur dioxide at 10, water at 100, and mercury at 356. " At the boiling-point of a liquid bubbles of vapour form and rapidly escape into the atmosphere, which gives rise to the characteristic bubbling ;md agitation of the liquid at this point. It is evident that a liquid will boil at a lower temperature when the pressure above it is reduced in other words, when the pressure of the air is reduced below atmospheric the vapour pressure of the liquid will be able to overcome it at a lower temperature. This may readily be shown by placing a flask containing water at 40 or 50 in connexion with an air-pump and rapidly exhausting, when the water will boil vigorously. The equilibrium between liquid and vapour may also be considered from the standpoint of the kinetic theory. Liquids, like gases, may be regarded as being made up of small particles in more or less rapid movement. When water is placed on the top of mercury in a vacuum, the particles in most rapid motion find their way into the space above the liquid. As the number of particles in the vapour space increases, more and more of them find their way back into the liquid, and finally a condition of equilibrium is attained in which as many particles enter as leave the liquid in a given time. Accord- ing to this view, the equilibrium between a liquid and its vapour is of a kinetic and not of a static character. It will be shown later that this view can be extended to chemical as well as to physical equilibrium. The amount of a substance in unit volume is conveniently termed the concentration of the substance. Thus we may say that at a definite temperature water is in equilibrium with a definite concentration of vapour, and the higher the temperature the greater is the concentra- 5 66 A TEXT-BOOK OF INORGANIC CHEMISTRY tion of vapour. This definition of concentration should be carefully noted, as it is of great importance. The equilibrium is not deter- mined by the absolute amount of vapour, but by the amount per unit volume the concentration. In the foregoing, we have considered the equilibrium between liquid and vapour in the absence of other substances. It is important to remember, however, that the presence of another vapour or gas does not affect the magnitude of the vapour pressure of a liquid. If, for example, one of the tubes represented in Fig. 26 contains air, the final concentration of water vapour in the space above the liquid is the same as when water vaporizes into a vacuum. The only effect of the presence of another gas or vapour is that the equilibrium is not reached so rapidly. Heat of Vaporization If a little ether is poured on the hand it rapidly evaporates, and a powerful cooling effect is felt. This is an illustration of the fact that when a liquid changes to a vapour heat is absorbed. The heat thus taken up in bringing about a change of state is termed latent heat ; when the change is from liquid to vapour it is called latent heat of vaporization. It has already been mentioned that the latent heat of vaporization of water at its boiling- point is 537 calories ; in other words, it requires the expenditure of 537 calories to convert i gram of water at 100 to water vapour at the same temperature. The same amount of heat is of course given up when the vapour is condensed. We can now understand why the temperature of boiling water remains constant although heat is being continuously supplied ; the heat is used up in bringing about a change of state. An alternative statement of the facts just mentioned is that steam has much more energy than an equal weight of water. According to the kinetic theory, the heat supplied is mainly transformed into kinetic energy ; it is used up in overcoming the attraction between the particles (p. 49), and in overcoming the pressure of the atmosphere (p. 146). The degree of cooling which may be attained by the rapid evapora- tion of volatile substances is very considerable. Thus the tem- perature of liquid ethylene, which boils at - 103, can be reduced to - 120, and in the same way, by the rapid evaporation of liquid oxygen, which boils at -182.5, a temperature of -210 may be reached. The Change of Liquid to Solid When a pure liquid is pro- gressively cooled it finally changes to the solid form, and conversely, WATER PHYSICAL PROPERTIES OF LIQUIDS 67 when the solid modification is heated it is again changed to the liquid fom ; in other words, it melts or fuses at a definite temperature. This property of solidifying or melting at a constant temperature is one of the most important characteristics of pure substances. It may be mentioned that some liquids do not solidify when the melting- point is reached, but may require to be supercooled several degrees before the solid form appears ; when solidification once begins, how- ever, it proceeds at constant temperature, which is the same as that at which the solid fuses (compare next section). In the case of water, and of all other liquids which expand on solidification, the temperature at which solid and liquid are in equi- librium is lowered by increase of pressure, but the effect is invariably small. In the case of water between i and 336 atmospheres an increase of pressure of n atmospheres lowers the melting-point of ice by 0.0074 n degrees, so that ice under a pressure of 336 atmospheres melts at -2.5. At higher pressures the coefficient is somewhat greater. Thus under a pressure of 1155 atmospheres ice melts at - 10, and under a pressure of 2200 atmospheres at - 22 (Tammann). The melting-point of substances which contract on solidification is raised by increase of pressure. Latent Heat of Fusion When heat is supplied to a solid its temperature rises till the melting-point is reached, and then remains constant till all the solid is melted. The phenomenon is exactly analogous to that occurring when a liquid is vaporized, the heat being used up in bringing about a change of state. Similarly, when a liquid is progressively cooled, the temperature falls till it begins to solidify, and then remains constant till solidification is complete, during which process the latent heat is given out. In order to convert I gram of ice at o into water at the same temperat ire 80 calories must be supplied ; in other words, the latent heat of fusion of ice is 80 calories. The values for most other sub- stances are lower than for water. Thus the latent heat of fusion of tin is i4-::5 calories per gram, and of sulphur 9.37 calories per gram. Equilibrium of the three Modifications, Ice, Water, and Vapour The proper understanding of the questions discussed in the previous paragraphs is greatly facilitated by a graphical repre- sentation of the variation of physical properties with temperature and pressure. The accompanying diagram (Fig. 27) affords such a representation for the different forms of water. Two lines are drawn at right angles, the so-called rectangular co-ordinates and tem- peratures are measured along the horizontal axis, pressures along the 68 A TEXT-BOOK OF INORGANIC CHEMISTRY B Liquid Solid vertical axis. The line OA represents the variation of the vapour pressure of water with temperature ; it is obtained by drawing a curve through the points at which lines drawn parallel to the pressure axis from points on the horizontal axis representing particular temperatures meet the lines drawn parallel to the temperature axis from points on the vertical axis representing the corre- sponding pressures. The point O represents the temperature at which a mixture of water and ice is in equili- brium under the pressure of their own vapour. This temperature is very near to, but is not exactly o. The latter temperature has already been defined as that temperature at which ice and water are in equilibrium under atmospheric pressure. At present, Vapour Temperature FIG. 27. however, we are concerned with the equilibrium between ice and water under the pressure of their own vapour. As this amounts at o to about 4.6 mm., which is practically an atmosphere below atmospheric pressure, and as diminution of pressure raises the melting-point of ice, the temperature corresponding with the point O is about +0.0074. Like water, ice exerts a definite vapour pressure, as is evident from the fact that it slowly evaporates at temperatures below zero. The line or curve OC represents the variation of the vapour pressure of ice with temperature, and it should be observed that it is not a direct continuation of the curve AO. The diagram indicates that at o (strictly speaking at +0.0074) water and ice have the same vapour pressure, and this has been proved both experimentally and theoreti- cally. Suppose that water and ice at o are contained separately in the limbs of the bent tube shown in Fig. 28, and let us assume for a moment that the vapour pressure of one of the modifications, say the water, is greater than that of the other. Under these circum- stances the water would continuously pass into vapour, which would solidify to ice in the other limb, until all the water has disappeared. It is therefore evident that the two modifications can only be in equilibrium when they exert the same vapour pressure. As water can be supercooled considerably, its vapour pressure can be measured for some degrees below zero. The vapour pressure of WATER PHYSICAL PROPERTIES OF LIQUIDS 69 supercooled water is represented by the line OA' on the diagram, which lies above OC ; it follows, therefore, that the vapotir pres- sure of supercooled water is greater than that of ice at the same temperature. This at once explains why supercooled water cannot exist in contact with ice ; if they are contained in the bent tube (Fig. 28) at a temperature below o, the water will completely evaporate into the other limb in virtue of its higher vapour pressure and solidify. Supercooled water is sometimes said to be unstable, as it solidifies at once in contact with ice. It is preferable to use the term meta- stable in this connexion in order to indicate that supercooled water has little or no tendency to crystallize in the entire absence of the solid form. These statements apply to other substances as well : the metastable form of a substance has invariably a higher vapour pressure than the stable form. It will, of course, be understood that relative stability or instability is entirely a question of conditions. Water is stable above o, metastable below o, ice stable below o and would be metastable above o if it could be superheated. The equilibrium diagram for water may be completed for our present purpose by drawing the line OB, which represents the effect of pressure on the melting-point of ice. As increased pressure lowers the melting-point, the line is slightly inclined towards the pressure axis, as shown. It is evident that at points on the curves only two forms are in equilibrium, solid and liquid along OB, liquid and vapour along OA, and ice and vapour along OC. At only one point, the point O, are three modifications in equilibrium, and the point O is therefore termed a triple point. The temperature (0.0074) an d pressure (4.6 mm.) at the triple point are the only values at which the three forms can exist together. If the pressure is increased, vapour disappears ; if it is diminished, water disappears. If the temperature is increased, ice disappears ; if it is lowered, water disappears. As the diagram shows, at the values of temperature and pressure represented by the regions between the curves only one form is capable of existence. Delay in Appearance of New Forms. Phases Water, ice, and water vapour are often termed different phases of the same substance. Each phase is homogeneous throughout and is separated by a definite surface from other phases. The term is chiefly used when one is dealing with so-called heterogeneous systems, which are made up of two or more phases. Water in contact with water vapour is such a heterogeneous system, made up of two phases, a liquid and 70 A TEXT-BOOK OF INORGANIC CHEMISTRY a vapour phase. This term will be found very useful in our later work. We have already learnt that ice does not necessarily appear when water is carefully cooled below o. This and similar facts may be generalized in the statement that when the conditions are such that a new phase may appear, it does not necessarily form unless a small amount of it is already present in the system. We shall see later that the amount of a new phase necessary to ensure the appearance of the latter in quantity under favourable conditions is exceedingly small (cf. p. 85). A further illustration of retardation in the appearance of a new phase is that water may be heated in a clean vessel several degrees above its boiling-point without the formation of vapour. At a certain point, however, a considerable quantity of vapour suddenly escapes, thus giving rise to that irregular form of boiling known as " bumping." In the absence of liquid water (and of dust), water vapour can be obtained under pressures greater than its vapour pressure under the experimental conditions without the appearance of water. On the other hand, as already mentioned, ice cannot be superheated ; the new phase appears as soon as the temperature exceeds o. Liquefaction of Gases We have seen in Chapter V. that a liquid is in equilibrium with its own vapour at a definite temperature and pressure. If at constant pressure the temperature is raised and kept at the new value the whole of the liquid will vaporize ; if under the same circumstances the temperature is kept below the equilibrium value, the whole of the vapour will liquefy. Similarly, at constant temperature the formation of liquid is favoured by increasing, the formation of vapour by lowering, the pressure. From this it would appear to follow that any gas or vapour can be liquefied by sufficiently lowering the temperature and increasing the pressure, and, as a matter of fact, every known gas has now been liquefied by applying this principle. One important point must. however, be noted in this connexion. If carbon dioxide is confined in a tube at room temperature and the pressure gradually increased, at a certain point liquid will appear in the tube, and by further increasing the pressure the whole of the gas liquefies. Under the same circumstances, however, no pressure, however great, brings about the liquefaction of oxygen. These facts puzzled the older chemists, but the problem was finally solved by Andrews, who showed that for every gas or vapour there is a temperature (which differs for each gas) below which it can be liquefied by pressure, but WATER PHYSICAL PROPERTIES OF LIQUIDS 71 above which no amount of pressure can bring about transformation to the liquid state. This temperature is -called the critical tempera- ture. It is now easy to understand the different behaviour of carbon dioxide and of oxygen under pressure. The critical temperature of carbon d : oxide is 31, so that at room temperature it is below its critical temperature ; on the other hand, the corresponding value for oxygen is 1 19, so that at room temperature it is far above its critical temperature. The pressure just sufficient to liquefy a gas at the critical tempera- ture is termed the critical pressure. For carbon dioxide at 31 the critical pressure is 72 atmospheres. For oxygen at -119 it is 58 atmospheres. It will, of course, be readily understood that the further a gas or vapour is cooled below the critical temperature the smaller is the pressure required to liquefy it. The physical constants, including the critical temperatures and pressures, of some of the commoner gases is given in the accom- panying table. Boiling- point. Melting- point. Critical Tempera- ture. Critical Pressure. Density at Boiling- point. Atmos. Helium . -268.5 -268? 2-75 0.15 Hydrogen Nitrogen . -252.5 -195.6 -259 -213 -241 -149 J 4 27-5 0.07 0.791 Carbon monoxide -190 -207 -136 33-4 Oxygen - 182.5 -223 -119 58.0 1.131 Ethylene . - i3- 5 -169 + 9 58.0 0.571 Nitrous oxide . - 89.8 - 102.7 + 37 73- 1.226 Carbon cioxide - 80 + 3 T -3S 72-3 Ammonia . - 33.5 - 75-5 + 131 "3 Chlorine . - 33-4 -102 + 141 84 i57 Sulphur dioxide - 10 - 76 + I5S-4 79 Methods used in Liquefying Gases The first example of the conversion of a substance which is a gas at the ordinary temperature into a liquid by pressure was chlorine (Northmore, 1806). Later, Faraday succeeded in liquefying a number of the commoner gases, such as sulphur dioxide, nitrous oxide, and ammonia. For this purpose he used bent tubes of strong glass (Fig. 28) ; substances for generating the gas were FIG. 28. 72 A TEXT-BOOK OF INORGANIC CHEMISTRY placed in one limb a, the end b of the tube was then sealed up, and on heating the other end the gas was given off, and under its own pressure part of it condensed in the limb b. If necessary, this end could be placed in a freezing mixture. Faraday did not succeed in liquefying oxygen or air ; the credit of having first obtained oxygen as a coherent liquid is due to Pictet (1877). The gas was obtained under very great pressure by heating potassium chlorate in a confined space, while the copper tube con- taining the gas was surrounded by liquid carbon dioxide kept at- 120 to- 140 by rapid evaporation under reduced pressure (p. 66). The liquefied carbon dioxide used in this experiment was obtained by compressing the gas in a tube surrounded by liquefied sulphur dioxide evaporating under reduced pressure (temperature - 60) ; the liquid sulphur dioxide in its turn being obtained by compressing the gas at room temperature. Prior to the working out by Linde and Hampson (independently) of the modern method of liquefying air and other so- called "permanent" gases, Pictet's method was exclusively used for this purpose. Simultaneously with Pictet, Cailletet obtained liquid oxygen, though only in the form of a mist and in small drops, by subjecting the gas to great pressure, which was then suddenly released. The intense cooling thus obtained caused the momentary liquefaction of part of the compressed gas. The older methods have now been completely displaced, as regards the less condensible gases, by the Linde-Hampson method, which has led to the liquefaction of hydrogen and helium. The principle of the method is that when a gas is allowed to pass from a high to a low pressure through a small opening (jet) or series of small openings (porous plug) it becomes cooled (Joule-Thomson effect). The cooling effect is due to the performance of "internal" work in overcoming the mutual attraction of the particles, and is therefore observed only for "imperfect" gases (p. 50). The effect is the greater the lower the temperature at which the expansion takes place, and the greater the difference of pressure on the two sides of the valve. The cooling effects thus obtained are "summed up in a very ingenious way by the principle of "contrary currents," the same quantity of gas being caused to circulate through the apparatus several times ; after expand- ing through the jet it is caused to flow over and cool the tube through which a further quantity of gas is passing on its way to the jet. The apparatus employed is represented diagramatically in Fig. 29. By means of the pump A the gas is compressed in B to (say) 100 atmospheres, the heat given out in the process being absorbed by WATER PHYSICAL PROPERTIES OF LIQUIDS 73 surrounding B with a vessel through which a current of cold water is passed. The cooled, compressed gas then passes down the central tube G tovards the jet E, being further cooled on the way by the gas passing up the wide tube D, which has just expanded through the jet. After passing through E, and thus falling to its origina; pressure, the gas passes up- wards over the central tube G and again reaches A by the tube C and the left- hand valve at the bottom of A. The direction of the circulating stream of gas is indicated by the arrows. In course of time the temperature becomes so low that part of the gas is liquefied, and collects in the vessel F. More air is drawn into the apparatus as required, and the process is continuous. By means of an apparatus constructed on this principle, Dewar, and, some- what later, Travers, succeeded in ob- taining liquid hydrogen in quantity. The hydrogen before expanding through the jet was cooled to about - 200 by means of liquid air boiling under reduced pressure. In the same way Kammerlingh Onnes of Leyden has just succeeded in liquefying helium, the preliminary cooling in this case being effected by means of liquid hydrogen. All known gases have now been liquefied. The Linde - Hampson machine is chiefly used for making liquid air, which is now a relatively cheap com- mercial article. As will be shown in detail later, air is essentially a mixture of the two " permanent " gases oxygen and nitrogen. The boiling- point of ]iquid air lies between those of the two components, oxygen -182.5 an d nitrogen 195.6, and when the liquid is fresh is about 190. It contains about 50 per cent, of oxygen. As nitrogen boils at a lower temperature than oxygen, the former passes off in much FIG. 29. 74 A TEXT-BOOK OF INORGANIC CHEMISTRY greater proportion as the liquid boils, and finally a mixture very rich in oxygen is obtained. As already mentioned, this method is used for the commercial preparation of oxygen (p. 24). The first fractions, which consist of almost pure nitrogen, are also of commercial value (p. 236). This process of separating a liquid into its components by taking advantage of a difference in their boiling-points, is known as fractional distillation. A much more efficient separation of the constituents of air is secured by methods devised by Claude and by Linde, which are based on a process which has long been in use for extracting spirit from a weak solution of alcohol and water. The weak alcohol is made to trickle down a tower containing zig-zag shelves or baffle plates, and a current of steam is admitted at the bottom and passed up through the shelves. At each stage some of the alcohol (which boils at a lower temperature than water) is vaporized, and some of the steam condensed, and finally a vapour very rich in alcohol escapes at the top and almost pure water trickles out below. In the application of this method to liquid air, the nitrogen, as the more volatile substance, corresponds with the alcohol and the oxygen with the water. Liquid air trickles down through a rectifying column up which nearly pure oxygen (obtained by evaporation of liquid air) is passed. As the gas passes up the column there is a continuous exchange of substance ; at each stage some of the rising oxygen is condensed, and some of the nitrogen in the downcoming liquid is evaporated ; finally almost pure oxygen is obtained at the bottom, and nitrogen containing a comparatively small proportion of oxygen passes off at the top. Quite recently, Claude has still further im- proved the process and obtained a complete separation of the oxygen and nitrogen. Claude's Method of Preparing Liquid Air 1 It has been pointed out that in the Linde method of preparing liquefied air no external work is done ; the cooling is done by internal work. Many attempts have been made to liquefy air by allowing it to expand in a working cylinder with performance of external work (p. 146), but at first the question of lubrication at such low temperatures presented a serious difficulty. Suitable lubricants have now been found by Claude in petroleum ether, which becomes viscous, but does not solidify even at -160; at still lower temperatures liquid air itself is used. In the first stage 1 An excellent account of modern methods of liquefying air and other gases is given by Ewing The Mechanical Production of Cold (Cambridge University Press). WATER PHYSICAL PROPERTIES OF LIQUIDS 75 of Claude's process, air is cooled by expansion in a working cylinder to temperature rather below 140 (the critical tem- perature of air) ; this cooled air is then used to cool the remainder of the air, with the result that the latter is liquefied at the pressure (50 atmospheres) at which it is supplied to the apparatus. It is doubtful whether the preparation of liquid air by external expansion a^one is more advantageous than the Linde method, but the use of expansion with external work for the first stage of the cooling, the cooled air being then liquefied by expansion through a throttle valve (without external work) appears to present un- doubted advantages. Investigations at low temperatures have been greatly facilitated by the use of so-called Dewar flasks glasi vessels with double walls, the space between the walls being completely exhausted of air (Fig. 30). In such vessels the conduction of heat from the outside is reduced to such an extent that liquid air can be preserved for hours with very little loss. A still better result is ob- tained by coating the walls with silver ; by this means radiant heat is reflected. FIG. 30. CHAPTER VII SOLUTION IN the preceding chapters, frequent reference has been made to solutions, more particularly those of gases and of solids in liquids, and it will be convenient to summarize at this stage some of the more important phenomena of solution, illustrated mainly by the examples already quoted. When a little sugar is shaken up with water, the former quickly disappears and a perfectly homogeneous mixture is obtained, which has a sweet taste, but in which no particles of sugar can be detected, even under the highest power of the microscope. The sugar does not settle out of the water, no matter how long the mixture may be kept, but can be obtained in its original form by evaporating off the water. If, on the other hand, a little chalk is shaken up with water, a milky mixture is obtained in which the particles of chalk can readily be detected under the microscope, and on standing the chalk sinks to the bottom, leaving the water clear. In the former case the sugar is taken up or dissolved by the water, forming a true solution, in the latter case the chalk is merely suspended in the water. In these extreme cases there is no difficulty in distinguishing between a suspension and a true solution, but we shall see later that some- times the distinction is by no means easy to draw. A solution may be defined as a homogeneous mixture of two or more substances, and to this we may add that the composition of a solution can be varied continuously within certain limits. The latter part of this definition serves to distinguish between a solution and a chemical compound. The latter, as we have already learnt, are of definite and invariable composition, whereas the composition of a solution of sugar in water may vary within wide limits. The component of a solution which is present in largest propor- tion is often called the solvent, and the substance taken up by the solvent is termed the dissolved substance or solute. In some cases, however, as in a mixture of alcohol and water in equal volumes, it cannot be said that one of the compounds has more right than the 7 6 SOLUTION 77 other to be regarded as the solvent. Strictly speaking, therefore, no sharp distinction can be drawn between the terms solvent and solute, but as a matter of convenience they ae often employed in the sense already indicated. The extent to which one substance can take up another to form a solution depends largely on the nature of the substances. Thus water and alcohol are miscible in all proportions, whilst chalk is scarcely taken up at all by water. When the miscibility is limited, and a solvent has taken up as much of a solute as it can retain in contact with the undissolved solute (compare p. 84), the solution is said to be saturated, A saturated solution represents a state of equilibrium between two phases, the solution and the undissolved substance. Solutions are usually classified according to the physical state of the components. The more important classes are as follows : (1) Solutions of gases in gases, (2) Solutions of gases in liquids, (3) Solutions of liquids in liquids, (4) Solutions of solids in liquids, and each of these will be briefly considered. Solutions of Gases in Gases Gases are miscible in all proportions, in other words, the mutual solubility of gases is unlimited. An important law which has been established as the result of the investigation of mixtures of gases is that when two gases, which do not act chemically on each other, are mixed the total pressure is the sum of the pressures of the two gases taken separately. Otherwise expressed, the pressure exerted by each of the gases is the same as if the othe- gas were not present, and the same mass of any one of the gases occupied the total space. The pressure exerted by any one gas ir a mixture of gases is termed its partial pressure. The fact that the vapour pressure of water is the same in air as in a vacuum under equivalent conditions (p. 66) is an excellent illustra- tion of the law. The above law, which was discovered by Dalton, is a particular case of a more general law, which states that when the components of a gaseous mixture exert no mutual influence, the properties of the mixture are the sum of the properties of the con- stituents. This law is of the same order of validity as the gas laws already enumerated (p. 40), and is the more nearly true the smaller the concentrations of the gaseous components. Solutions of Gases in Liquids Unlike the solubility of gases in gases, the solubility of gases in liquids is limited. The 78 A TEXT-BOOK OF INORGANIC CHEMISTRY amount of a gas which a liquid can take up depends upon four factors : (a] the nature of the gas, (b) the nature of the liquid, (c) the temperature, (d) the pressure. (a) The amounts of different gases which the same solvent can take up vary enormously with the nature of the gas. Thus I c.c. of water at o absorbs 0.0203 c.c. of hydrogen, 1.713 c.c. of carbon dioxide, and 1148 c.c. of ammonia gas under 76 cm. pressure. (b) The amounts of the same gas taken up by different solvents under comparable conditions vary greatly with the nature of the solvent. The volumes of hydrogen taken up by i c.c. of each of the following solvents at o under a gas pressure of 76 cm. is as follows : Solvent Solubility" . Water. Aniline. 0.0203 0.0303 Acetic Acid. 0.0617 Ethyl Ethyl Acetate. Alcohol. 0.0764 0.0851 0.0862 Acetone. (c] The solubility of gases diminishes regularly as the temperature rises. This rule is illustrated by the following table, which gives the volumes of hydrogen, nitrogen, and carbon dioxide taken up by i c.c. of water at temperatures between o and 50 under a gas pressure of 76 cms. : Solute. " Absorption Coefficient " in Water at Hydrogen . Nitrogen . Carbon dioxide . 0.0203 0.0239 T -7i34 10 20 0.0190 0.0177 0.0196 0.0164 1.194 0.878 30 0.0163 0.0138 0.665 50 0.0146 0.0106 0.436 The .majority of gases can be completely removed from solution by raising the temperature. This, however, is not invariably the case ; hydrogen chloride, for instance, cannot be completely removed from aqueous solution by boiling (cf. p. 95). (d) The amount of a gas absorbed by a given volume of liquid increases with increased pressure. The law connecting the amount of absorption and pressure, usually known as Henry's law, is as follows : At constant temperatiire, the amount of gas absorbed by a given volume of liqtiid is proportional to the pressure of the gas. Another way of stating Henry's law is that the 'volume of gas raken up by a given volume of liquid is independent of the pressure. This is clearly equivalent to the first statement, because when the pressure SOLUTION 79 is doubled the quantity of gas absorbed is doubled, but since its volume, according to Boyle's law, is halved, the original and final volumes absorbed are the same. A third instructive method of stating Henry's law is that the ratio of the concentration of the dissolved gas to that in tl.e free space above the liquid is independent of the pressure. In the case of hydrogen at o, for example, the ratio of the concen- tration in water to that in the gas space is 0.0203 : i, or approximately i : 50. When the pressure is doubled, the concentration both in the gas space and in water is doubled, but the ratio remains unaltered. Henry's law is approximately valid for such gases as hydrogen and oxygen, but does not hold for very soluble gases nor for those gases which enter into chemical combination with the solvent (p. 95). For comparative purposes, the solvent power of a liquid for a gas is best expressed in terms of the "solubility" or "coefficient of solubility," which is the volume of the gas taken up by unit volume of the liquid at a definite temperature (Ostwald). As is clear from the foregoing, the "solubility" is simply the ratio in which the gas distributes itself in the liquid and in the gas space, and has the great advantage of being independent of the pressure in so far as the gas follows Henry's law. Solubility measurements are still sometimes expressed in terms of the so-called "absorption coefficient" of Bunsen, which is the volume of gas, reduced to o and 76 cm. pressure, absorbed by i c.c. of a liquid at a definite temperature under a gas pressure equal to 76 cm. of mercury. A little consideration shows that the only difference between the "solubility" s and the "absorption coefficient" a, is that in the former case the volume of the gas is taken at the tem- perature of the experiment and in the latter case is reduced to o. For measurements at o, therefore, the two factors coincide, and at other temperatures the relationship is expressed by the equation s\a= (273 + /)/273, where / is the temperature of observation on the Centigrade scale. A convenient form of apparatus for determining the solubility of gases in liquids, used in Ostwald's laboratory, is represented in Fig. 31. The graduated measuring tube A is connected by a rubber tube with :he tube B, and by altering the height of the latter till the liquid (mercury or water) stands at the same height in the tubes, a quantity of gas confined in the upper portion of A can readily be measured at atmospheric pressure. The tube A is connected with the absorp:ion vessel C by a flexible metal tube; a and b are three- way stopcocks and c is a simple stopcock. In making an experiment 8o A TEXT-BOOK OF INORGANIC CHEMISTRY the vessel C, the volume of which must be known, is filled with the air-free solvent and a quantity of the gas is brought into A. The volume of the gas at atmospheric pressure is then noted, A and C are put in connexion by means of the stopcocks a and b, and a measured volume of liquid run out of C, its place being taken by the gas. The tap b is then closed, the gas and liquid in C are vigorously shaken in order to secure equilibrium, the tap b again opened and the diminution in the volume of the gas read off in A after adjustment to atmospheric pressure. As the volume of liquid in C and the volume of gas absorbed are known, the solubility of the gas at the temperature of the experiment can readily be calculated. In order to secure constancy of temperature, A and C are immersed in a bath of water or are surrounded by vapours of known temperature. Solubility of Mixed Gases- It has already been pointed out that each component of a gaseous mixture exerts its effect quite independently of the other components. The solubility of mixed gases is a particular case of this law ; the amount of any one gas taken up by a given volume of liquid is proportional to its partial pressure in the mixture. This law was established by Dalton (1807), and is known as Dalton's law of partial pressures. In calculating the amounts of each gas taken up from a mixture, the solubility of the gas, as well as its partial pressure, must of course be taken into account. Thus if a mixture of two volumes of hydrogen and one volume of nitrogen, which have about the same solubility in water at o, is shaken up with water, the volume of hydrogen dissolved will be about double that of the nitrogen, but if a mixture of two volumes of hydrogen and one volume of oxygen are used, approxi- mately equal volumes of the gases will be dissolved, as the coefficient of solubility of oxygen is about double that of hydrogen. Solutions of Liquids in Liquids As regards the mutual FIG. 31. SOLUTION 81 solubility of liquids three cases may be distinguished: (i) The liquids mix in all proportions, e.g. alcohol and water; (2) the liquids are practically immiscible, e.g. benzene and water ; (3) the liquids are partially miscible, e.g. ether and water. In some respects the partially miscible liquids are the most in- teresting. The mutual solubility of ether and water may be shown by shaking approximately equal volumes of the two liquids in a separating funnel. On allowing to stand, a separation into two layers takes place, the upper layer consisting of a solution of water in ether, and the lower of a solution of ether in water. The lower layer may be separated from the upper by carefully opening the tap, and the presence of ether may be shown by warming the solution in a flask provided with a cork and glass tube ; the escaping ether can be ignited at the end of the tube. The presence of water in the lighter ethereal layer may be shown by adding a small amount of anhydrous copper sulphate (cf. p. 414), the white colour of which changes to blue in the presence of moisture. In most cases the mutual solubility of two partially miscible liquids increases with the temperature, and it may therefore be anticipated that liquics which in certain proportions form two layers at the ordinary temperature may become completely miscible at higher temperatures. This is the case, for instance, with ordinary phenol and wr ;r, which mix in all proportions at temperatures above 68.4. Solubility of Solids in Liquids The solubility of solids in liquids depends upon the nature of the solid, the nature of the solvent, and upon the temperature, but is only very slightly affected by change of pressure. In all cases the solubility is limited but varies within very wide limits ; thus chalk is practically insoluble in water, but the latter solvent can take up its own weight of cane sugar at the ordinary temperature. Two general methods are used for deter- mining the solubility of a solid in a liquid. According to the first method, the finely divided solid is shaken with a definite volume of the liquid at constant temperature till no more will dissolve; part of the solution is then removed and the amount of the dissolved substance it contains determined. The second method depends upon the fact that in the great majority of cases the solubility in- creases as the temperature is raised. The solvent is heated with excess of ;he solid to a temperature higher than that at which the solubility is to be determined, and is then cooled to the required temperature in contact with the solid, when the excess above that required to form a saturated solution separates out. 6 82 A TEXT-BOOK OF INORGANIC CHEMISTRY As already indicated, a saturated solution represents a state of equilibrium between two phases, and in the present case a liquid and a solid phase are in equilibrium. It is instructive to compare an equilibrium of the type now under discussion with that between water and water vapour. In the latter case, there is a definite con- centration of vapour in equilibrium with water at a definite tem- perature, and the vapour concentration is independent of the amount of vapour (that is of the size of the space accessible to it) and of the amount of water. In the same way, the concentration of a saturated solution in contact with undissolved solid is independent of the volume of the solution and of the amount of undissolved solid. The extent of the solubility of a solid in a liquid under definite conditions may be expressed in a number of ways : (i) As the number of parts by weight (grams) of the solute present in (a) 100 parts by weight (grams) or (b] 100 parts by volume (c.c.) of the saturated solution ; (2) as the number of parts by weight (grams) of the solute taken up by (a} 100 parts by weight (grams) or (b) 100 parts by volume (c.c.) of the solvent. The second method appears to possess certain advantages, and will be mainly used in the present book. Effect of Temperature on the Solubility of Solids in Liquids The solubilities of a number of salts in the same solvent, water, and the varia- tion of the solubilities with temperature, is represented graphically in Fig. 32. The ordinates represent the num- ber of grams of salt taken up by loo grams of water and the abscissas represent tempera- tures. The curve for potas- sium nitrate is very steep, indicating that the solubility of the salt increases very rapidly as the temperature rises. Thus it can be gathered from the curve that at 10 100 grams of water dissolve about 20 grams of the salt, and at 50 about 85 grams, so that a rise of temperature of 40 has more than quadrupled the solubility. ' ^ I ftl l . -NO A I3 1 5 (^ ^ ' 120 < 1 \^ i-, '"' ic 1 >v, cS 110 ^ J / / / / ^ 80 / \ 8 / / r\ ^ . e ^ > j" ^ f i j .^ '7 / s* "> 50 j j s* ?* ,LA ^ ** **** ^ f 40 2x * - "i = ^: = 40 i s - <* ^ -= "vfc jj- *d f r J< j> * \* o / s ^-* ""rrd S V ^ " / j B g.* ! 0" '10 20 30 40 50 60 70 60 90 IOO*C Temperatures > FIG. 32. SOLUTION On the other hand the solubility of sodium chloride at o is about 36 grams and at 100 nearly 40 grams in 100 grams of water, so that the change of solubility with temperature is very slight. The temperature coefficients of solubility of the great majority of salts in water are greater than that of sodium chloride but less than that of potassium nitrate, as :he diagram indicates. The solubility of sodium sulphate in water is remarkable inasmuch as it increases fairly rapidly with the temperature up to 33 and then slowly diminishes as the temperature is further increased (Fig. 33). An examination of the solid in contact with the solution will show that at below 33 it consists of a compound of the salt and water, a hydrate (p. 91), above 33 the solid in equilibrium with the solution consists of the anhydrous salt. We may therefore say that the part BC of the curve represents the solubility cf a hydrate, the branch CD that of the anhydrous salt. In a few exceptional cases the solubility of a solid in a liquid diminishes $ as the temperature rises, "o When fresh lime is shaken g up with water for some Js time at the ordinary tem- perature and the excess of solid is allowed to settle, a clear saturated solution of calcium hydroxide in water, familiarly known as lime water, is obtained. On warming the solution it becomes turbid from separation of solid, showing that the solubility diminishes with increase of temperature. Some further examples of negative temperature-coefficients of solubility, as they are called, will be mentioned in connexion with the respective salts. The effect of temperature on the solubility is closely connected with the question whether heat is given out or taken up when solution takes place. Certain substances, such as potassium nitrate, dissolve in water with absorption of heat, as is shown by the fact that the temperature falls while the salt is being dissolved ; other substances give out heat during solution and the temperature rises (example, bromine in water). It has now been found that for substances which _c Z 30 20 30 40 50 60 70 Temperatures > FIG. 33. 84 A TEXT-BOOK OF INORGANIC CHEMISTRY dissolve with absorption of heat the solubility increases as the tem- perature is raised, whilst with substances which dissolve with evolution of heat the solubility becomes less as the temperature is raised. A good illustration of this rule is the solubility of sodium sulphate (Figr-33) ; corresponding with the fact that the hydrate dissolves with absorption of heat, the solubility increases with rise of temperature (branch BC), the anhydrous salt, on the other hand, dissolves with evolution of heat, and its solubility diminishes with rise of temperature (branch CD). The usefulness of this rule is considerably lessened by the fact that it refers to the heat effect of dissolving a substance in its own saturated (or practically saturated) solution and not to whether heat is given out or absorbed when a substance is dissolved in the pure solvent. It sometimes happens that although heat is absorbed when a definite amount of salt is dissolved in a large volume of water, heat is given out when the same amount of salt is dissolved in a solution containing a large proportion of the same salt. The rule in question should therefore be used with considerable caution (cf. p. 172). Supersaturated Solutions If water is saturated at a definite temperature, say 20, by shaking with excess of a salt the solubility of which increases with rise of temperature, and a portion of the solution, free from undissolved solid, is heated to, say, 40, it will then be capable of dissolving more salt, and is therefore termed im- saturated. If, on the other hand, a concentrated solution of the same salt is prepared at, say, 40, and then allowed to cool slowly in the complete absence of undissolved solid, solid salt does not necessarily separate when the temperature has fallen below that at which the solution is saturated in contact with solid. Such a solution is said to be supersaturated. When a minute fragment of the solid is added to such a solution the excess of dissolved solid at once separates, and the concentration falls to that of a saturated solution at the tempera- ture of the experiment. The above statements may be illustrated as follows. If crystallized sodium sulphate is heated with its own weight of water till a perfectly clear solution is obtained, and the neck of the flask is closed with a plug of cotton-wool, the solution will remain perfectly clear when the temperature has fallen to that of the room, although it is then highly supersaturated. If, however, a minute crystal of the solid is added, crystallization at once takes place. The amount of a solid sufficient to start crystallization in a super- saturated solution is extremely small. According to Ostwald, the SOLUTION 85 almost inconceivably minute amount of icr 10 gram l (a ten-millionth part of a milligram) brings a supersaturated solution of sodium sulphate to crystallization. As small fragments of many salts are present in the floating dust in the air of laboratories, it is usually necessary ro plug with cotton-wool the neck of the flask in which a supersati rated solution is being prepared. Supersaturation is by no means confined to solutions of solids in liquids. Under certain conditions, highly supersaturated solutions of gases in liquids may be obtained. It is evident that there is a close analogy between supersaturation and the phenomenon of supercooling. Both are illustrations of the general ru e that when the conditions are favourable for the appear- of a new phase, it does not necessarily appear unless a small amount of it is already present in the system. If a supersaturated solution of sodium sulphate is cooled sufficiently, the solid salt begins to separate without the necessity of adding a small crystal of the solid phase. A distinction is sometimes drawn between the inetastable region, in which crystallization does not occur in the absence of a fragment of the solid phase, and the labile region, in which crystallization of a supersaturated solution occurs spon- taneously (cf. p. 277). Solution and the Kinetic Theory The experimental fact that a salt can remain in equilibrium with a solution at a definite temperature, under which circumstances the solution is said to be saturated, is instructive when considered from the standpoint of the kinetic theory. We may suppose that when a solid is in contact with a liquid it tends to send out particles into the liquid. These particles move about in the liquid, and some of them will presumably return and redeposit on the solid. The number thus returning to the solid state will be the greater the greater the concentration of the dissolved substance, and ultimately a stage will be reached when the number sent out is just balanced by the number returning to the solid this is the equilibrium condition. The tendency of a solid to send out particles into a liquid in contact with it is known as its solution pressure; at equilibrium the solution pressure is balanced by the pressure of the dissolved substance. 1 Very small amounts are most conveniently expressed by means of negative indices. Thus lo" 1 gram is T ^ gram, io~ 3 gram is j^Vs gram, and so on. CHAPTER VIII CHLORINE AND HYDROCHLORIC ACID "O EFERENCE has already been made to both these important -tv substances in the chapter on water. It has been pointed out that chlorine is an element which can be made to combine with the hydrogen of water, forming hydrogen chloride and setting free oxygen. The chemical equation representing the change is as follows : Chlorine + water = hydrogen chloride + oxygen. Hydrogen chloride is sometimes called hydrochloric acid gas ; its solution in water is hydrochloric acid. For several reasons it is convenient to deal with these two substances at the present stage. CHLORINE History Chlorine was discovered in 1774 by Scheele, who obtained it by heating hydrochloric acid with a naturally occurring oxide of manganese. It was for a number of years called oxymuriatic acid) being supposed to contain oxygen (p. 92). In 1810 Sir Humphry Davy proved it to be an element, and gave it its present name in allusion to its greenish-yellow colour (x\oip6s = greenish- yellow). Occurrence Chlorine does not occur free in nature, on account of its great chemical activity, but is found in large amount combined with sodium, potassium, magnesium, and other metals. Its com- pound with sodium, known as sodium chloride or common salt, is, as already mentioned, the salt present in the largest proportion in sea- water, and is also found in mines in Galicia and elsewhere. Potassium and magnesium chlorides are found along with sodium chloride in salt deposits at Stassfurt in Germany. Both chlorine and hydro- chloric acid are, however, obtained almost exclusively from sodium chloride. Preparation (i) As hydrochloric acid is a chemical compound of chlorine and hydrogen which is easily obtained in quantity, it is 86 CHLORINE AND HYDROCHLORIC ACID 87 natural to use it as a source of chlorine. For this purpose the hydrogen has to be removed, and this is most conveniently done by causing it to combine with oxygen. The equation representing this action is as follows hydrogen chloride + oxygen^water + chlorine and is the reverse of the equation on the previous page, which represents the decomposition of water by free chlorine. The reversi- bility of the reaction is conveniently represented by oppositely directed arrows (cf. p. 164). The apparatus used is represented in Fig. 34. A stream of air is passed through the bottle, which contains concentrated hydrochloric acid, and the mixture of air and hydrogen chloride thus obtained is FIG. 34. led throigTi the bulb tube, which is strongly heated. The tube contains pieces of pumice stone soaked with a solution of copper chloride (or sulphate). The gas escaping at the end of the tube contains chlorine, which can be recognized by its smell and by other characteristic tests mentioned below. It is, however, mixed with excess of air and of hydrogen chloride, and cannot be obtained even approximately pure by this method. In the absence of copper chloride practically no chlorine is formed when a mixture of air and hydrogen chloride is passed through a red- hot tube. It is evident, therefore, that the copper salt has accelerated the change represented by the upper arrow, whilst the salt can be recovered unchanged at the end of the reaction. This is a further illustration of what we have already termed catalytic actions. The method just described is used for preparing chlorine on the 88 A TEXT-BOOK OF INORGANIC CHEMISTRY commercial scale, and is called the Deacon process. The salt of copper chiefly used as a catalyst for the Deacon process is cuprous chloride (cf. p. 412). (2) Instead of using free oxygen to decompose hydrogen chloride, it is much more convenient to use combined oxygen in the form of certain peroxides and other compounds containing a large proportion of oxygen. The substance generally used for this purpose is the black powder known as manganese peroxide or pyrolusite, which occurs largely in nature. The apparatus used for preparing chlorine by this method is repre- sented in Fig. 35. The mixture of pyrolusite and a concentrated aqueous solution of hydrochloric acid is heated gently in a large flask A, and the issuing chlorine is first passed through water in a two-necked bottle B (the so-called Woulf's bottle) to free it from hydrogen chloride, and then, if required perfectly dry, through a second bottle containing concen- trated sulphuric acid. As chlorine is fairly soluble in water, and acts chemically on mercury, it cannot be collected over either of these liquids. As it is considerably heavier than air, it can readily be collected by allowing it to issue from the delivery tube at the bottom of a gas jar C, as shown in FIG. 35. the figure. This method of collecting heavy gases by upward dis- placement of air has already been mentioned (p. 39). The equation representing the chemical change just described is as follows : Manganese peroxide + hydrochloric acid = chlorine + manganese chloride + water. The manganese chloride which, as its name indicates, is a chemical compound of manganese and chlorine, remains behind in the flask. It is evident that the change in this case is one of double decomposi- tion (p. 7). CHLORINE AND HYDROCHLORIC ACID 89 (3) Inszead of manganese peroxide, chlorinated lime, lead peroxide, potassium bichromate or potassium permanganate, all compounds Containing oxygen, may be used to effect the decomposition of hydrochloric acid. The reactions are described in connexion with the compounds themselves. Chlorinated lime, otherwise called bleaching powder, is particularly suitable, as the chlorine is rapidly given off in the cold. (4) By electrolysis. Chlorine can also be prepared from hydro- chloric acid by electrolysis in concentrated aqueous solution. An apparatus similar to that shown on p. 14 may be used ; owing, how- ever, to the fact that chlorine attacks platinum, carbon electrodes must be used. The gas which comes off at the negative pole is not pure chlorine, but always contains oxygen, and the proportion of the latter gas is the greater the more dilute the hydrochloric acid solution. The oxygen probably results from the action of the freshly liberated chlorine on the water. Compounds of chlorine with metals, the so-called metallic chlorides, also yielc chlorine on electrolysis. Chlorine is now prepared on the large scale by the electrolysis of sodium chloride in aqueous solution, as described on p. 382. Physical Properties Chlorine is a yellowish-green gas with a powerful, disagreeable odour. It attacks the mucous membranes strongly, and if inhaled in considerable amount causes death by suffocation. Its density, referred to air as unity* is about 2.45, referred to hydrogen about 35.5. One litre of chlorine, at normal temperature and pressure, weighs 3.22 grains, Chlorine is very easily liquefied, at o six atmospheres pressure is required, and at -34 one atmosphere pressure is sufficient. The last statement indicates that the boiling-point of liquid chlorine is 34. At 102 it solidifies ; both the liquid and the solid have a yellow colour. Liquid chlorine is now a commercial article, being conveyed in steel cylinders lined inside with lead. The critical temperature of chlorine is 146, and its critical pressure 94 atmospheres. Chlorine is fairly soluble in water; at room temperature (18) I volume of water takes up about 2.2 volumes of the gas. The variation of the absorption coefficient, , with temperature, is represented by the formula # = 3.0361 -0.046 1 96/4- o.ooo 1 1 07/2. The aqueous solution has the odour and colour of the gas, and is known as chlorine water. It is very unstable, especially in sunlight ; 9 o A TEXT-BOOK OF INORGANIC CHEMISTRY owing to the affinity of the chlorine for the hydrogen of the water, hydrogen chloride and oxygen are formed (see below). This may be shown by filling a small retort with chlorine water and exposing it to bright sunlight while supporting it in the position shown (Fig. 36). In a short time a gas, which proves to be oxygen, collects in the upper part of the retort. Chemical Properties Chlorine is characterized by great chemical activity, combining at room temperature with a number of other elements with evolution of light and heat. Finely powdered arsenic and antimony, when shaken into the gas, at once catch fire, forming white fumes of the respective chlorides. Finely divided copper, lead, tin, and " Dutch metal," which consists of a mixture of copper and zinc hammered out into thin sheets like gold leaf, also catch fire in chlorine, and phosphorus at first melts and then burns feebly in the gas. Chlorine does not unite directly with carbon, with nitrogen, or with oxygen. Reference has already been made to the great affinity between chlorine and hydrogen. A jet of hydrogen burning in the air will continue to burn when lowered into a jar of chlorine, and in the same way a gas jet or lighted candle continues to burn in chlorine, hydrogen chloride being formed, and the carbon set free in the form of soot. The same facts may be illustrated very strikingly by soaking a strip of filter-paper with turpentine, which is a chemical compound of carbon and hydro- gen, and suspending it in a jar of chlorine. In a few moments the turpentine catches fire and dense clouds of soot are formed. It is a remarkable fact, which has not yet been adequately ex- plained, that many of the reactions just described do not take place if the chlorine is perfectly dry. Thus Dutch metal is not affected when put into a jar of chlorine which has been dried by means of concentrated sulphuric acid. We shall meet later with many examples of the influence of traces of moisture on chemical reactions. The effects belong to the class of catalytic phenomena. Bleaching Action of Chlorine If chlorine gas is passed into a solution of a colouring matter, such as indigo or litmus, the colour rapidly disappears and is said to be bleached. Further, if a piece of Turkey red cloth is damped and put into a jar of chlorine it is rapidly bleached, but if it is previously carefully dried by keeping it for some time in a vessel over concentrated sulphuric acid, and is CHLORINE AND HYDROCHLORIC ACID 91 then suspended in a jar of chlorine gas .which has stood for an hour or two in a carefully closed gas jar containing a layer of sulphuric acid at the bottom, no bleaching occurs. As the presence of moisture is thus shown to be essential for bleaching to occur, it is assumed that the actual bleaching effect, the conversion of a coloured to a colour- less substance, is due to oxygen liberated from water by the action of chlorine. It has been found, however, that free oxygen has very little bleaching action under ordinary conditions, and it is therefore often assumed that oxygen just liberated as the result of a chemical change, the so-called "nascent" oxygen, has much more powerful bleaching and other properties than has ordinary oxygen. In all probability, however, the matter is more complicated, and will be referred to again in connexion with the oxygen compounds of chlorine (p. 1 80). Chlorine bleaches ordinary ink, but has little or no action on printer's ink. The colour of the latter is due to particles of carbon, which at the ordinary temperature is not acted on either by oxygen or chlorine. Chlorine has also a powerful destructive action on low organisms, such as disease germs, and is therefore largely used as a disinfectant. This accion is probably also due, in part at least, to its property of liberating oxygen from water. As the oxygen set free from water by the agency of chlorine can readily effect oxidations, chlorine is regarded as a powerful oxidizing agent. Many examples of this property will be met with in the course of our subsequent work. Chlorine is also used commercially in the extraction of gold from its ores, as it forms a readily soluble chloride with this metal (p. 428). Chlorine Hydrate When chlorine gas is passed for some time into water cooled to o, a greenish crystalline substance, consisting of chlorine and water in chemical combination, separates. The crystal- line substance is called chlorine hydrate. It is fairly stable at low temperatures, but above 9.6 splits up into water and chlorine. The composition of chlorine hydrate will be referred to later (p. 112). The name hydrate is applied to a substance formed by the com- bination of water as a whole with other compounds, or, as in the present case, with elements. An enormous number of well-known compounds, such as common salt, copper sulphate, sodium carbonate, etc., form hydrates with water, but very few hydrates are known which have ari element as one of the components. 92 A TEXT-BOOK OF INORGANIC CHEMISTRY HYDROCHLORIC ACID (HYDROGEN CHLORIDE) History The aqueous solution of hydrogen chloride known as hydrochloric acid has been known since the middle of the seventeenth century. It was obtained by heating sodium chloride (common salt) with sulphuric acid, the method which is still employed on the com- mercial scale. It was formerly known as spirit of salt, and also as __j muriatic acid. As chlorine is usually prepared by oxidizing hydrochloric acid, and was formerly supposed to contain oxygen, it is now easy to see why it used to be called oxy- muriatic acid (p. 86). The mistake was at the time a very natural one, but we know now that oxidizing agents do not add on oxygen to hydrogen chloride, but remove hydrogen from it. The gas itself, hydrogen chloride, was discovered by Priestley, who was the first to collect gases over mercury. Preparation (i) Hydrogen chloride can be prepared by direct combination of its elements, or by the action of chlorine on many compounds containing combined hydrogen, as already mentioned. The combination of the elements can be shown very instruc- tively in a thick-walled tube (Fig. 37) provided with a stopcock at each end, and divided into two equal parts by a stopcock in the middle. The two halves of the tube are filled in red light with dry hydrogen and chlorine respectively by dis- placement of mercury, the central stopcock is opened to allow the gases to mix, and the tube is then exposed to diffused daylight. The greenish colour of the chlorine gradually dis- appears, and after a time it will be found that the gases have combined completely to form hydrogen chloride. If then one of the end stopcocks is opened carefully under mercury, no gas will escape and no mercury will enter. This experiment ' 37> illustrates two very important points : (i) hydrogen and chlorine combine in equal volumes to form hydrogen chloride ; (2) hydrogen and chlorine combine without change of volume to form hydrogen chloride. If equal volumes of the mixed gases are exposed to bright sunlight or to a magnesium flash-light, combination takes place explosively. This experiment should not be tried in the tube just described, but may be done quite safely by supporting a sealed thin-walled bulb, filled with the mixed gases, behind a glass screen and exposing it to a magnesium flash-light. Combination is practically instantaneous, CHLORINE AND HYDROCHLORIC ACID 93 and the bulb is reduced to fragments. The preparation of the bulbs is effected by blowing a few of them from "a single piece of glass tube in such a way that they remain connected by thin capillary tubes. A stream of mixed hydrogen and chlorine, prepared by electrolysis, is passed through the series of bulbs for a considerable time in dark- ness, and then the connexions between them are sealed off with a small flame. They can also be bought filled and ready for use. All kinds of light are not equally effective in causing combination of hydrogen and chlorine. Blue and violet rays are most active in this respect, whereas, as already indicated, red light is practically without effect. The combination of the gases can also be brought about by passing an electric spark or by applying a lighted taper to the mouth of a jar containing them, as in the case of a mixture of hydrogen and oxygen. The fact that hydrogen and chlorine combine slowly under the in- fluence of light is taken advantage of in the construction of one form of actinometer, an instrument for measuring the chemical activity of light. The gases are confined over water and kept at constant pressure, and from the rate of diminution of volume (the hydrogen chloride which is formed at once dissolves in the water) the chemical activity of a source of light can be estimated. It is a curious fact that when the gases are exposed to light, practical: y no combination takes place for a short time ; after- wards the rate of combination is quite regular. This initial period of no reaction is termed the "period of induction." In spite of a great deal of work on the subject, the phenomenon is by no means understood. (2) Hydrogen chloride is almost invariably prepared, both in the laboratory and on the commercial scale, by the action of sulphuric acid on common salt. When the acid is not used in too great excess, the chemical change is represented by the equation sodium chloride + sulphuric acid = sodium sulphates- hydrogen chloride. The apparatus used is similar to that represented in Fig. 35. Dry sodium chloride is placed in a flask provided with a thistle funnel, and concentrated sulphuric acid is added through the funnel. On gently warming, hydrogen chloride is given off in a steady stream ; it can be dried by passing through a bottle containing pumice stone soaked with concentrated sulphuric acid, and then collected over mercury or by upward displacement of air (p. 39). As the gas escapes into the air 94 A TEXT-BOOK OF INORGANIC CHEMISTRY white fumes will be noticed ; these are due to the combination of the gas with the moisture of the air. As already mentioned, hydrogen chloride is usually employed in the form of its solution in water, known as hydrochloric acid. A method of preparing such a solution is shown in Fig. 44. The glass tube dips just below the surface of the water, and in order to prevent the water from being sucked back into the flask, the glass tube has a bulb in the middle. In such a case, when the entrance of water into the generating flask might cause an accident, it is well to interpose an empty bottle between the flask and the absorption vessel. Hydrogen chloride may be obtained very conveniently for labora- tory purposes by allowing concentrated sulphuric acid to drop from a funnel provided with a stopcock into commercial concentrated hydrochloric acid. An apparatus similar to that represented in Fig. 44 may be used. Little or no heating is required. The commercial preparation of hydrochloric acid is again referred to in connexion with sodium carbonate (p. 388). Physical Properties Hydrogen chloride is a colourless gas with a suffocating odour. Its density referred to hydrogen is 18.3, so that it is about 1.26 times heavier than air. By cooling under pressure it can readily be converted into a colourless liquid which boils at - 83 ; in other words the vapour pressure of liquid hydrogen chloride at -83 is one atmosphere. At o, the vapour pressure is equal to 26 atmospheres. Its critical temperature is 52, so that it can be liquefied by pressure alone at the ordinary temperature ; its critical pressure is 86 atmospheres. One of the most remarkable properties of hydrogen chloride is its great solubility in water. This may be illustrated in a striking way by passing a few drops of water, by means of a bent pipette, into a quantity of the gas confined over mercury. The mercury will be observed to rise rapidly in the tube, showing that water can absorb many times its own volume of the gas. A still more striking method of illustrating the same fact is indicated in Fig. 38. The flask is completely filled with hydrogen chloride through the long glass tube, which is open at both ends ; the outer end of the tube is then placed in a vessel of water, as shown, and the liquid sucked up the central tube by means of the side tube. When the water just reaches the top of the inner tube the side tube is closed. The first few drops of water dissolve nearly all the gas and produce a partial vacuum, into which the water is forced by the pressure of the atmosphere. At o i volume of water absorbs 503 volumes of hydrogen chloride, CHLORINE AND HYDROCHLORIC ACID 95 but the solubility diminishes rapidly with rise of temperature. The effect of pressure on the solubility of hydrogen chloride in water is not represented by Henry's law ; as a matter of fact, the amount dissolved increases only slightly with considerable increase of pressure. This is illustrated by the following table, in which the upper line represents the pressure in mms. of mercury and the lower line the weight of acid dissolved by 100 grams of water at o (Roscoe). P = Grams of hydro- gen chloride. 100 65.7 200 70.7 300 73-3 500 78.2 700 1000 81.7 85.6 It has already been stated (p. 79) that Henry's law does not hold for very soluble gases. At 15 water takes up about 43 per cent, by weight of hydrogen chloride under a pressure of the gas equal to 760 mm. of mercury, and this is, of course, the most concentrated acid which can be obtained under ordinary con- ditions ; i;s density is 1.21. The concentrated acid of commerce contains about 36.5 per cent, by weight of the acid, the density of the solu- tion is i. i 6. When the concentrated aqueous solution is heated, hydrogen chloride alone is given off at first, and the concentration of the solution gradually diminishes till it reaches 20.2 per cent., when the remainder of the mixture distils un- changed in composition at 110. If, on the FIG. 38. other hand, a dilute solution is boiled, only water is given off at first, and the concentration of the solution in the distilling flask gradually increases (with simultaneous rise of temperature) till it again reaches 20.2 per cent., when the residue distils at constant temperature as before. If we happen to start with a 20.2 per cent, solution, the whole of it distils at constant temperature, like a pure liquid (p. 64), and the distillate is throughout of the same composition as the solution in the distilling flask. The behaviour of mixtures of hydrogen chloride and water on dis- tillation is different from that of the great majority of binary mixtures of liquids. We have already seen that a liquid boils when its vapour pressure s equal to that of the atmosphere, and in the same way 9 6 A TEXT-BOOK OF INORGANIC CHEMISTRY a binary mixture boils when the sum of the partial pressures of its components is equal to atmospheric pressure. In a mixture of two liquids of different boiling-points, one of the components has in the great majority of cases, not only a lower vapour pressure than the other, but a lower vapour pressure than a mixture of the components in any proportions. When such a, mixture is heated, the less volatile component will tend to remain behind, and a partial separation can thus be effected, and this will take place the more readily the greater the difference in the boiling-points of the components. If, however, a mixture of the components in a certain definite proportion happens to have a lower vapour pressure than that of either of the components or of a mixture of the two in any other proportion, it is evident that on distillation a mixture of this composition will tend to remain behind. We have an excellent illustration of the last case in the mixture of hydrogen chloride and water containing 20.2 per cent, of the acid, and the reason why this mixture tends to remain behind on distillation, and finally distils at constant temperature, will now be understood. It was long thought that this constant-boiling mixture of hydrogen chloride and water is a definite chemical compound, but Roscoe and Dittmar (1860) showed that the composition of the mixture depends on the pressure. Thus at a pressure of 250 cm. of mercury the constant-boiling mixture contains 18 per cent, at 76 cm. (atmos- pheric pressure) 20.2 per cent, (boiling-point 110), and at 10 cm. 22.9 per cent, of hydrogen chloride (boiling-point 62). Chemical Properties The most striking feature with refer- ence to the chemical properties of hydrogen chloride is the remark- able difference in activity between the liquefied gas and the aqueous solution. Whilst the solution of the gas in water turns litmus red, at once acts on lime, and dissolves such metals as zinc and magnesium, the liquefied gas, in the entire absence of moisture, shows none of these properties. It is evident, therefore, that only the aqueous solution acts as an acid ; anhydrous hydrogen chloride has no acid properties. The explanation of this remarkable alteration of pro- perties in contact with water will be given later. The aqueous solution is a typical acid ; it turns red litmus blue, has a sour taste, and, like sulphuric acid (p. 35), acts on metals such as magnesium and zinc, hydrogen being set free. The general properties of acids are further considered below (p. 98). Composition of Hydrogen Chloride The composition by volume of hydrogen chloride, like that of water, may be determined CHLORINE AND HYDROCHLORIC ACID 97 by analytical or by synthetical methods. It has already been shown by a synthetical method (the double tube experiment, p. 92) that the hydrogen and chlorine combine in equal volumes to form hydrogen chloride, and that by the combination of one volume of hydrogen and one volume of chlorine two volumes of hydrogen chloride are formed. That combination is complete when equal volumes of the gases are taken can be shown by cautiously opening the lower stopcock under water, when the latter will enter and completely fill the tube. The conclusions drawn from this experiment can be confirmed by an analytical method described by Roscoe. The apparatus used is represented in Fig. 39, and consists essentially of a U-tube, one limb of which is open, and the other is provided with a stopcock. Both limbs are at first filled with mercury, and then dry hydrogen chloride is drawn into the left-hand limb until it is about two- thirds full at atmospheric pressure. Some liquid sodium amalgam (a solution of sodium in mercury, p. 381) is then poured into the open limb of the U-tube, the end of the latter firmly closed with the thumb, the hydrogen chloride transferred to the right-hand limb by inclining the tube and kept for a little time in contact with the amalgam. Under these circumstances the hydrogen chloride is com- pletely decomposed, the chlorine combining with the sodium to form sodium chloride, and the hydrogen is set free. The latter is transferred to the left-hand limb, the pressure again adjusted to that of the atmosphere, and it will then be found that the hydrogen occupies exactly half the volume of the original hydrogen chloride. The residual gas may be shown to be hydrogen by driving it out at the stopcock and igniting. It still remains to find the volume of chlorine which has dis- appeared. This is done by taking a long narrow glass tube drawn out at both ends, filling it with the mixed gases obtained by electro- lysis of concentrated hydrochloric acid (p. 89) and sealing both ends. One end is then dipped into a solution of potassium iodide and the end broken, when it will be found that the solution enters the tube and occupies exactly half the volume. Free chlorine acts upon potassium iodide, forming potassium chloride and setting free iodine, which dissolves in the solution, producing a brown colour. The residual gas is hydrogen. It follows that the chlorine which has 7 98 A TEXT-BOOK OF INORGANIC CHEMISTRY disappeared occupied a volume equal to that of the hydrogen, and by combining this result>with that of the previous experiment it is evident (i) that hydrogen chloride is entirely made up of equal volumes of hydrogen and chlorine ; (2) that the gases combine without change of volume to form hydrogen chloride. From the densities and combining volumes of the gases it can be calculated that hydrogen chloride contains 35.46 parts of chlorine to 1.0078 parts of hydrogen by weight. Direct determinations of the composition by weight lead to the same result. We have here a further illustration of the law that definite chemical compounds are of constant composition. The composition of water and of hydrogen chloride have been considered in some detail on account of their fundamental importance for the theory of chemistry. Acids, Bases and Salts Reference has already been made on several occasions to acids, and under this heading we have grouped substances which have certain properties in common. The more important of these properties are (i) the power of turning litmus red ; (2) a sour taste ; (3) the liberation of hydrogen when they are brought into contact with certain metals, such as zinc or magnesium. It follows from the statement of the third property that, as already pointed out, all acids contain hydrogen. It must be carefully remembered, however, that a substance con- taining hydrogen is not necessarily an acid. Thus liquefied hydrogen chloride, although it contains hydrogen, has no acid properties, nor has water itself. The full discussion of this important subject is postponed to a later stage. We have also learned that when metallic sodium is added to water, a solution is obtained which has a soapy feel and turns red litmus blue. These properties indicate the presence of a base in the solution. The particular base in this case is sodium hydroxide ; it contains sodium, hydrogen, and oxygen, and can be obtained as a white solid on evaporating the solution. A solution which turns litmus blue is said to have an alkaline reaction. A solution which turns litmus red has an acid reaction. If now to a solution of sodium hydroxide, containing a little litmus, a dilute solution of hydrochloric acid is added a little at a time, the mixture being stirred after each addition, the blue colour will remain for some time, but after a definite amount of hydrochloric acid has been added, it will suddenly turn red. Experiment will show that a drop or two of acid or alkali is sufficient to change the colour from CHLORINE AND HYDROCHLORIC ACID 99 blue to red or vice versa, but a point can be reached at which the solution is neither definitely blue or red, but violet. Such a solution has none of the properties of an acid or a base, it is said to be neutral. On evaporating off the water, crystals of sodium chloride (common salt) are obtained. The formation of a substance with new properties isj of course, evidence cf a chemical action. We therefore obtain the very im- portant result that when a typical acid, such as hydrochloric acid, is brought in contact with a typical base, such as sodium hydroxide, the chemical change results in the complete disappearance of the acid and basic properties, and a new substance, with entirely different properties, results. The process just considered is termed neutraliza- tion, and a substance resulting from the neutralization of an acid by a base is termed a salt. The statements just made are quite general ; when any acid is brought in contact with any base a salt is formed. In the present example we have seen that sodium chloride is one of the products of the neutralization. As it contains only sodium and chlorine, there remains the hydrogen of the hydrochloric acid and the oxygen and hydrogen of the sodium hydroxide to be accounted for. These are, however, the components of water, and, as a matter of fact, water is the second product of the reaction. The complete equation is therefore Hydrochloric acid + sodium hydroxide = sodium chloride + water. When an acid is neutralized by a base, water is always one of the products. The formation of the salt may be regarded from a rather different point of view, as resulting from the displacement of the hydrogen of the acid by a metal, in this case sodium. It should, however, be mentioned that if the hydrogen of an acid is only partially replaced by a metal, the resulting substance is nevertheless termed a salt. A salt may, therefore, be defined as a substance formed by the complete or partial displacement of the hydrogen of an acid by a metal. It does not, of course, matter whether the displacement is effected directly by the metal itself or, as in the present case, by a compound of the metal. We have already met with a number of instances of the direct displacement of hydrogen from an acid by a metal with formation of a salt. It has been pointed out that sodium chloride in aqueous solution is neutral to litmus. The same is true of many other salts, such as calcium chloride and sodium sulphate, but is by no means generally ioo A TEXT-BOOK OF INORGANIC CHEMISTRY true. The aqueous solutions of some salts have an acid reaction, those of others an alkaline reaction. Oxides of Chlorine Although chlorine cannot be made to combine directly with free oxygen, yet two oxides of chlorine can be obtained by indirect methods. These oxides are briefly referred to here in connexion with the fundamental laws of chemistry discussed in the next chapter, and will be dealt with more fully in Chapter XIV. in connexion with other compounds of chlorine. Chlorine monoxide is obtained by passing dry chlorine over red oxide of mercury in the cold, and is a pale yellow, very explosive gas. Analysis shows that it contains 35.46 parts of chlorine to 8 parts of oxygen by weight. Chlorine peroxide is formed when concentrated sulphuric acid is added to a small amount of dry potassium chlorate in a test-tube, and the mixture is very cautiously warmed. It is a deep yellow, extremely explosive gas. It contains 35.46 parts of chlorine to 32 parts of oxygen by weight. A third oxide of chlorine, containing a higher proportion of oxygen than the peroxide, has also been described (p. 179). CHAPTER IX LAWS OF CHEMICAL COMBINATION THE ATOMIC THEORY T aw of Constant Composition In the previous chapters -* ' we have seen that the composition of pure water is constant ; no matter what its source or how it has been prepared, it always contains 83. 184- per cent, of oxygen and 11.816 per cent, of hydrogen (approximately 8 parts by weight of oxygen to i part of hydrogen). In the same way we have shown that hydrogen chloride is of con- stant composition ; it contains about 35.5 parts of chlorine to i part of hydrogen by weight. Further, it may be shown that mercuric oxide always contains 92.6 per cent, by weight of mercury and 7.4 per cent, by weight of oxygen. If oxygen is used in large excess, nevertheless the elements combine in the above proportions, and the excess of oxygen remains uncombined. The composition of an enormous number of other pure substances has been deter- mined with very great accuracy, and in all cases it has been found that they are of constant composition. We are, therefore, justified in assuming that this result is a law of nature. It is usually called the Law of Constant Composition, and is formulated as follows : A definite chemical compound always contains the same elements in the sam* proportions by weight. Law of Multiple Proportions The determination of the com- position of a number of chemical compounds has shown that the same elements may unite in more than one proportion to form chemical compounds. At the end of the last chapter it was stated that there are two well-defined oxides of chlorine, which for a fixed pro- portion 35.5 parts by weight of chlorine contain 8 and 32 parts of oxygen respectively. Both are definite chemical compounds of constant composition, but for a fixed proportion of chlorine one contains four times as much oxygen as ( the other. A further illustration is to be found in the two iodides of mercury. 102 A. TEXT-BOOK OF INORGANIC CHEMISTRY If 20 grams of mercury are rubbed in a mortar with 12.7 grams of iodine, a few drops of alcohol being added occasionally to facilitate combination, the globules of mercury gradually disappear, and finally a yellowish-green compound is obtained, containing the mercury and iodine in the above proportions. If now a further 12.7 grams of iodine are added, and the rubbing continued, a new chemical compound is obtained in the form of a red powder. Both are definite chemical compounds of constant composition, but for the same amount of mercury one contains twice as much iodine as the other. The examination of a large number of compounds shows that the above result, that for a fixed amount of one element there is a simple relationship between the amounts of the other element present, is a general rule, and we are therefore entitled to express it in the form of a law. The Law of Multiple Proportions, which summarizes the facts, may be stated as follows : When two elements unite in more than one proportion, for a fixed amount of one element there is a simple relationship between the amounts of the other element present. Later results will show that the ratio in question need not be i : 2 or i : 4 ; it may be 2 : 3, 3 : 4, or even more complex. For simplicity, the law has been deduced from the composition of binary compounds, but it applies to chemical compounds in general, whether simple or complex. When the composition of a compound containing two elements is expressed in percentages, it is not evident at first sight that the law of multiple proportions holds. Thus of the two iodides of mercury, the yellow compound contains 61.16 per cent, ot mercury and 38.84 per cent, of iodine, and the red compound 44.05 per cent, of mercury and 55.95 per cent, of iodine. When, however, the compositions are referred to a fixed amount, say i part, of mercury, the green compound is found to contain 0.635 parts, and the red compound 1.27 parts of iodine to i part of mercury, the former numbers being in the ratio of 1:2. Chlorine monoxide contains 81.59 per cent, of chlorine and 18.41 per cent, of oxygen ; the dioxide contains 52.56 per cent, of chlorine and 47.44 per cent of oxygen. It may readily be shown from these figures that the law of multiple proportions holds. The Law of Combining Weights The study of the com- position of chemical compounds has led to the establishment of a still more comprehensive law, of which the two laws already deduced are special cases. As a preliminary to the deduction of LAWS OF CHEMICAL COMBINATION 103 the law, the percentage composition of a number of familiar chemical compounds is given in the following table. (1) Mercuric Oxide. (3) Water. (5) Mercuric Sulphide. Mercury 92.59 Hydrogen 11.18 Mercury 86.18 Oxygen 7.41 Oxygen 88.81 Sulphur 13.82 (2) Mercuric Iodide. (4) Hydrogen Iodide. (6) Hydrogen Sulphide. Mercury 44.05 Hydrogen 0.78 Hydrogen 5.92 Iodine 55.95 Iodine 99.22 Sulphur 94.08 The first step in comparing the combining proportions of the elements in these six compounds is to fix on one element as a standard, to which the weights of the other elements may be referred. As hydrogen appears to be the element present in smallest proportion in the compounds in question, we may con- veniently take the quantity of hydrogen present in a compound as unity, to which the amounts of the other elements present are to be referred. In our list there are three compounds containing hydrogen, and it may easily be calculated that the combining weight of oxygen in water is approximately 8, that of sulphur 16, and that of iodine 127, when referred to unit quantity of hydrogen. We will now consider the composition of a compound containing one of the above three elements, say iodine, and any other element. Mercuric iodide is a convenient substance for our purpose. It may easily be shown (p. 102) that in this compound 127 parts of iodine are combined with 100 parts of mercury. We now take a further step and find the combining weight of, say, sulphur by finding the amount of that element combined with 100 parts of mercury (the com- bining weight of mercury referred to iodine) in mercuric sulphide ; the amount in question is 16. It has, however, already been found by an entirely independent method the analysis of hydrogen sulphide that the combining weight of sulphur is 16. From this the pro- visional conclusion may be drawn that the proportions in which the elements enter into chemical combination their so-called combining proportions or combining weights are independent of the nature of the elements with which they are combined. Further investigation fully confirms this conclusion with a slight amplification. It is possible to find a combining weight for each element, which re- presents the proportion, referred to the combining weight of hydro- gen as unit (more accurately, to the combining weight of oxygen taken as 8), in which it enters into chemical combination, but the amount of the element entering into combination may also be an io 4 A TEXT-BOOK OF INORGANIC CHEMISTRY integral multiple of the combining weight. The latter part of this statement, that an element may enter into combination as a simple integral multiple of its combining weight, is illustrated by the composition of mercurous iodide, quoted on p. 102. This com- pound contains 200 parts of mercury to 127 parts of iodine (equivalent to i part of hydrogen), the former number being twice the combining weight. The law of combining weights is formulated as follows : For each element a fixed number, termed its combining weight ', can be found, such that the clement enters into chemical combination in the ratio of its combining weight, or in simple integral multiples of this ratio. For reasons which will be given later, the combining weights of the different elements are referred to that of oxygen taken as 8, and not to that of hydrogen taken as unit. The accurate combining weight of hydrogen, referred to the oxygen standard, is 1.008. The combining weights of the elements so far dealt with are as follows : Oxygen. Sulphur. Mercury. Iodine. Hydrogen. 8 1 6 100 127 1.008 From the composition of hydrogen chloride the combining weight of chlorine (which is 35.5) can be deduced, and hence the composition of the compounds of oxygen and chlorine can be deduced. The expected values should be compared with the experimental values already given (p. 102). It is clear from the foregoing that the experimental determination of the combining weight of an element is a comparatively simple matter ; we have only to determine the proportion in which it enters into combination with the combining weight of another element. As already indicated, hydrogen and oxygen are chiefly used as reference elements for this purpose. We may therefore, from an experimental standpoint, define the combining weight of an element as the smallest amount of it which can combine with (or take the place of) 1 .008 parts by weight of hydrogen or 8 parts by weight of oxygen. Gay-Lussac's Law of Volumes Having discussed the laws of combination by weight, we now proceed to consider the laws of combination by volume ; and in this connexion we confine ourselves to the volumes of gases and vapours, as offering the simplest relation- ships. It has been shown that hydrogen and oxygen combine in the ratio of two volumes of the former to one volume of the latter, giving two volumes of steam, when all the gases are measured under the THE ATOMIC THEORY 105 same conditions. Further, one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrogen chloride. These remarkably simple relationships, and others of a similar kind, were discovered by the French chemist Gay-Lussac, and led him to the enunciation of the Law of Gaseous Volumes, which may be stated as follows : Gases combine in simple ratios by volume, and the volume of the gaseous product bears a simple ratio to the volumes of the reacting gases. Subsequent investigation has shown that the Law of Volumes, like the other gas laws, is very nearly, but not exactly true. The Atomic Theory The four laws of chemical combination just considered are purely experimental, and independent of any view we may lake as to the constitution of matter. In connexion with the properties of gases, however, we have already discussed the theory that matter is not continuous, but is made up of extremely small, practically incompressible particles, which are far apart in gases but much closer together in solids and liquids, and it is natural to inquire if this theory throws any light on the laws of chemical combination. The view of the constitution of matter just indicated originated with the Greek philosophers, but was first developed to a consistent theory b> Dalton (1808) ; it is termed the atomic theory. According to this tneory, matter is made up of small particles called atoms, which cannot be further divided by any means at our disposal. The atoms of any one element are identical in all respects, but differ, at least in weight, from those of other elements. By the union of the atoms of different elements in simple numerical proportions, chemical compounds are formed. As was first pointed out by Dalton, the laws of chemical combina- tion by weight find a ready explanation on the atomic theory. For simplicity we will consider only binary compounds, that is, compounds made up of two elements ; but the reasoning is the same for more complicated compounds. As the ultimate particles of a binary com- pound are made up of two kinds of atoms, and a fixed number of each kind, the compound must be of constant composition, since it is made up of an enormous number of such ultimate particles. If, for example, we assume that the ultimate particles of hydrogen chloride are made up of one atom of hydrogen, weight 1.008, and one atom of chlorine, weight 3546, the compound must be of invariable composition, con- taining hydrogen and chlorine in the ratio 1.008 : 35.46. Similarly, the law of multiple proportions follows at once from the 106 A TEXT-BOOK OF INORGANIC CHEMISTRY atomic theory. If an ultimate particle of a compound contains an atom of one element, of weight x, and an atom of another element, of weight y, the compound will contain the two elements in the ratio x :y. If the same two elements unite to form a second compound, the ultimate particles of which contain one atom of the first element and two of the second, the ratio of the two elements in the second compound will be x : 2y. From this it follows that for a fixed amount x of one element, the other element is present in the two compounds in the exact ratio 1:2, in accordance with the law of multiple proportions. Finally, the law of combining weights is also seen to be a logical consequence of the theory, the experimentally found combining weights bearing a simple relation to the relative weights of the atoms. The fixing of the relative weights of the atoms is clearly a matter of the utmost importance, and we shall see in the following paragraphs how the problem was satisfactorily solved. Avogadro's Hypothesis We have seen that, according to Gay-Lussac's law, gases combine in simple ratios by volume. Further, according to the atomic theory, chemical combination takes place between one or two (or some small number of) atoms of one kind and one or two atoms of another kind. By combining these two state- ments it is evident that there is a simple relationship between the number of particles in equal volumes of different gases under the same conditions. Strong support is lent to this deduction by the facts already stated with reference to the similar behaviour of all gases on altering the temperature or pressure. The most obvious suggestion in this connexion is that equal volumes of all gases contain the same number of atoms, and Gay-Lussac him- self was at first inclined to adopt this view. It was soon found, however, to be untenable ; and the view held at the present day was shortly afterwards put forward by Avogadro (1811). Avogadro drew a distinction between atoms, the smallest particles of matter which can take part in chemical changes, and molecules, the smallest par- ticles of matter capable of independent existence, and formulated his hypothesis as follows : Equal volumes of all gases, under the same conditions of tempera- ture and pressure, contain the same number of molecules. There should be no difficulty in drawing a clear distinction between molecules and atoms. The molecules of a chemical compound, the smallest particles capable of independent existence, are made up of atoms of different kinds. On the other hand, the molecules of which THE ATOMIC THEORY 107 a quantity of an elementary substance is made up, are usually formed by the urion of two or more atoms of an 'element ; in a few cases by a single atom of the element. We shall see later that the atoms of elements already considered hydrogen, oxygen, chlorine have not been obtained independently ; the molecule of hydrogen, for example, is formed by the union of two atoms of hydrogen. On the other hand, the molecule of mercury in the gaseous state is made up of a single atom. It will be shown later that Avogadro's hypothesis leads to the con- clusion that the molecule of hydrogen chloride is made up of one atom of hydrogen and one atom of chlorine. Assuming this result for the o @@ @ @ @ @ @ @ @ @ @ @ @ @ moment, the accompanying diagram gives a graphic representation of Avogadro's hypothesis applied to hydrogen chloride, hydrogen, and mercury. The symbol @ represents an atom of hydrogen, (Cl) represents an atom of chlorine, and (Hg) represents an atom of V S mercury. The spaces a, b, and c represent equal volumes, and the diagram illustrates the following important points: (i) Equal volumes of all gases and vapours contain the same number of mole- cules under the same conditions, no matter whether the molecules are those of elements or of chemical compounds. (2) The volumes of the molecules themselves (the " particles " of the kinetic theory, p. 48) are sim-ll compared with the total volume occupied. (3) The mole- cules of chemical compounds are made up of atoms of different kinds. io8 A TEXT-BOOK OF INORGANIC CHEMISTRY (4) The molecules of elements are made up of atoms of the same kind, and the number of atoms in the molecule varies from one upwards. These four statements contain the essential features of the mole- cular theory of matter. The second statement at once enables us to understand how the same number of molecules of different substances, although they may be very different in size, may still occupy equal volumes in the gaseous form under equivalent conditions. Complexity of the Molecules of Hydrogen and of Oxygen We have seen that one volume of hydrogen and one volume of chlorine unite to form two volumes of hydrogen chloride. Suppose in these two volumes there are 2000 molecules of hydrogen chloride. Each of these molecules contains some hydrogen and chlorine, and must contain at least one atom of each. Therefore, 2000 atoms of hydrogen and 2000 atoms of chlorine must have been concerned in the formation of the 2000 molecules of hydrogen chloride. But, according to Avogadro's hypothesis, if two volumes of the hydro- gen chloride contain 2000 molecules, the one volume of hydrogen must contain 1000 molecules. It follows, therefore, that the 2000 atoms of hydrogen have been obtained from 1000 molecules of the gas, so that each molecule of the gas must contain (at least) two atoms. By exactly similar reasoning it may be shown that each molecule of chlorine is made up of two atoms of that element. It is evident that the above reasoning only fixes a lower limit to the complexity of the hydrogen molecule ; if the molecule of hydrogen chloride contains more than one atom of hydrogen, the hydrogen molecule must be still more complex. At a later stage, however (p. 116), evidence will be given that hydrogen chloride contains only one atom of each element, and therefore the molecules of hydrogen and of chlorine contain only two atoms of the respective elements. From the experimental fact that two volumes of hydrogen and one volume of oxygen give two volumes of steam, it may be shown in exactly the same way that the oxygen molecule is formed by the union of two atoms. Avogadro's Hypothesis and Molecular Weights On the basis of Avogadro's hypothesis, the determination of the relative weights of the molecules of gases is very simple. As equal volumes contain the same number of molecules, it is evident that the relative weights of equal volumes of gases give us the relative weights of the molecules contained in them ; in other words, the ratio of the densities of two gases is the ratio of the weights of their molecules. It only remains to choose some substance as standard, to which the weights THE ATOMIC THEORY 109 of the molecules of other substances are to be referred. As hydrogen is the lightest gas, and has already been used as standard in density determinations, it is natural to use it also as standard for molecular weights. On this basis the weight of the molecule of hydrogen might be taken as unity. It has been pointed out, however, that the hydrogen molecule contains two atoms, and there are certain advan- tages in taking the weight of the atom of hydrogen as unity, to which all molecular weights are referred. On this basis, the molecular weight of hydrogen is 2. In previous chapters it has been pointed out that the density of oxygen in round numbers is 16, referred to hydrogen as standard, and this, according to Avogadro's hypothesis, indicates that the oxygen molecule is sixteen times the weight of the hydrogen molecule, or thirty-two times the weight of the hydrogen atom. The molecular weight of oxygen is therefore found by doubling its vapour density. This method is clearly of general application, and may be stated as follows : The molecular weight of a gas is double its vapour density referred to hydrogen as tmit, In order that the basis of this important rule may be clearly realized, we may think of a certain distance which, when expressed in feet, measures 15, and in yards measures 5. It is clear that, since in the first case the distance is referred to a unit one-third of that used in the second case, the number expressing the distance in the first case must be three times that in the second case. In exactly the same way, molecular weights, being referred to a standard half that to which the densities are referred, must be represented by numbers which are double those representing the densities. A slightly different way of stating Avogadro's hypothesis is that the same volume is occupied x by the molecules of all gases under the same conditions ; in other words, the molecular volumes of all gases are equal. It has been found convenient to represent each molecule as occupying two unit volumes, and this statement will be largely used later on. No assumption is made as to the actual volume occupied by a molecule, the definition being only used relatively. 1 The term occupied is here used in the sense of inhabited, and does not mean that the space is filled with the particles. The true meaning is rendered evident by the diagram on p. 107. In equal volumes of different gases there are the same number of particles, therefore single molecules of different gases inhabit the same average space under corresponding conditions. For purposes of com- parison wi take this average space as two unit volumes. no A TEXT-BOOK OF INORGANIC CHEMISTRY It can readily be calculated from the density given on p. 36, that the molecular weight of hydrogen in grams (that is, 2.016 grams of hydrogen) occupies about 22.40 litres at o and 760 mm. pressure. As, however, the molecules of all gases occupy the same volume under the same conditions, the molecular weight of any other gas in grams must also occupy 22.40 litres under normal conditions. This is the definition of molecular weight which should be remembered by the student, and it will be repeated for the sake of definiteness. The molecular weight of a gas is that weight of it, expressed in grams, which occupies 22.4 litres at o and 760 mm. pressure. The validity of the above statement may perhaps be shown still more clearly by taking it in the converse way. In 22.4 litres of different gases there are the same number of molecules, according to Avogadro's hypothesis, and the relative weights of these equal volumes are clearly in the ratio of the molecular weights. That particular volume, how- ever, contains the molecular weight of hydrogen in grams, and therefore the molecular weights of the other gases in grams. The enormous importance of the statements just made is obvious. If a substance can be vaporized without decomposition, its molecular weight can be determined without difficulty. We shall see later that other methods of determining molecular weights are used for substances which cannot be converted into vapour without decom- position (cf. Chapter XV.). Symbols and Formulae Before proceeding to discuss the methods used in determining atomic weights, it is desirable to con- sider briefly the methods used in chemistry for representing the com- position of elements and compounds. Instead of writing out in full the name of an element each time it is mentioned, it is conveniently and shortly indicated by using the first letter, or, in certain cases, two letters of its Latin name. Thus, O represents oxygen, N stands for nitrogen, C for carbon, S for sulphur, and so on. The symbol for chlorine is Cl ; for tin Sn (from stannum) ; for lead Pb (plumbum) ; for iron Fe (contracted from ferrum) ; and antimony is represented by Sb (stibium). When the names of more than one element begin with the same letter, the initial letter is generally used for the better known element, and as symbols for the other elements two letters are used. A full list of the symbols for the elements is given on the back page of the cover, and the student should gradually familiarize himself with them. Chemical compounds are formed by the combinaton of two or more elements, and it is natural to indicate the composTjaon of a com- THE ATOMIC THEORY in pound by \\ riting side by side the symbols of the elements composing it. It will, however, obviously be an enormous advantage if the symbols can be used to indicate, not only the qualitative but also the quantitative composition of a chemical compound. The first step towards this end is to use the symbol of an element, not merely to indicate the presence of the element, but to represent the atomic weight of the element. Thus the symbol H, when used to represent a component of a chemical compound, stands for I part by weight of hydrogen, and Cl stands for 35.46 parts of chlorine. As already indicated, the molecule of hydrogen chloride is formed by the union of one atom of hydrogen and one atom of chlorine. This is indicated by HC1, which shows not only that hydrogen chloride contains hydrogen and chlorine, but that it contains I part by weight of hydrogen to 35.46 parts by weight of chlorine. HC1 is termed the formula for hydrogen chloride. In the same way, the formula for common salt, NaCl, indicates that it contains sodium and chlorine, and that these elements are present in the ratio of 23 parts by weight of sodium to 35.46 parts by weight of chlorine. It sometimes happens that the molecule of a chemical compound contains more than one atom of a particular element. The number of atoms present is indicated by the appropriate figure written at the lower righi -hand side of the symbol. Thus H 2 O, which is the formula for water, indicates that the molecule contains two atoms of hydrogen to one atom of oxygen, and as the atomic weights of these elements are in the approximate ratio of I : 16, the formula indicates that water is made up of 2 parts by weight of hydrogen to 16 parts by weight of oxygen, as has already been proved experimentally. The remark- able simplicity and convenience of this system of formulation will be clear from the example just given. From the formula of ordinary chalk, CaCO 3 , it may easily be calcu- lated that it contains 40 per cent, of calcium, 12 per cent, of carbon, and 48 per cent, of oxygen. The composition of potassium chlorate, KC1O 3 , and of sulphuric acid, H 2 SO 4 , may also be calculated from the formulae. 1 In some cases a number of chemical compounds contain the same group of e^ments. We have just seen that the formula for potassium chlorate is KC1O 3 , and it will be shown later that all chlorates contain the group C1O 3 ; that is, they contain chlorine and oxygen in the proportion of one atom of the former to three atoms of the latter. 1 Approximate atomic weights: H = i; = 12; O = i6 ( '; 8=132; 01=35.5; ii2 A TEXT-BOOK OF INORGANIC CHEMISTRY Sometimes, however, a molecule contains more than one such group of elements ; for example, barium chlorate contains two C1O 3 groups associated with one atom of barium. This might, of course, be indicated by writing its formula BaCl 2 O 6 , but this method has the disadvantage that we could not tell at a glance that we are dealing with a chlorate. It is therefore preferable to write the formula of barium chlorate either as Ba(ClO 3 ) 2 or as Ba2ClO 3 . An integer on the line, as in the last example, is to be understood as multiplying all that follows, up to the next comma. Still another type of formula may be mentioned here for the sake of completeness. We have already seen that some salts form chemical compounds with water, generally called hydrates. Thus when sodium chloride separates from solution at low temperatures, one molecule of it is found to be combined with two molecules of water. When the temperature is raised, the two molecules of water are driven off and sodium chloride is left. The compound in question is repre- sented by the formula NaCl,2H 2 O. When two parts of a formula are thus separated by a comma, it usually indicates that the compound is comparatively easily broken up at the point in question. The water associated with such compounds is sometimes called water of crystallization. The meaning of the formula C1 2 ,8H 2 O, ascribed to chlorine hydrate, will now be readily understood. On the basis of Avogadro's hypothesis, we have decided to use the convention that the molecule of any substance in the state of vapour occupies two unit volumes. The formulas of the volatile substances quoted in the present section, since they represent the formulae of molecules, are termed molecular formula, and stand for the amounts of the substances which occupy two unit volumes in the form of vapour. A formula such as HC1 is therefore remarkably expressive, inasmuch as (1) It shows what elements are present in the compound. (2) It shows the relative proportions by weight of the components. (3) It represents two unit volumes of the substance in the state of vapour. We have now met with two sets of numbers, the combining weights and the atomic weights, both of which represent the propor- tions in which elements enter into chemical combination. The exact relationship between these two sets of numbers, and the methods employed in fixing atomic weights, form the subject matter of the next chapter. THE ATOMIC THEORY 113 Fact- Generalization or Natural Law Hypothesis- Theory Chemistry, like all the experimental sciences, is based on facts, established by experiment. A large number of such facts have already been mentioned ; for example, that certain chemical com- pounds, v r hich have been investigated with the greatest care, always contain the same elements in the same proportions. A mere collec- tion of facts, however, does not constitute a science. When a certain number of facts have been established, the chemist proceeds to reason from analogy as to the behaviour of systems under conditions which have not yet been investigated. For example, Proust showed that there are two well-defined oxides of tin and that the composition of each is invariable. From the results of these and a few other investigations he concluded from analogy that the composition of all pure chemical compounds is invariable, although, of course, very few of them had then been investigated from that point of view. To proceed in this way is to generalize, and the short statement of the conclusion arrived at is termed a generalization or law. It will be evident that a law is not in the nature of an absolute certainty ; it comprises the facts experimentally established, but also enables us to foretell a great many things which have not been, but which if necessary could be investigated experimentally. The greater the number of cases in which a law has been found to hold, the greater is the confidence in its validity, until finally a law may attain practically the same standing as a statement of fact. We may confidently antici- pate that, however greatly our views regarding natural phenomena may change, such generalizations as the law of constant proportions will remain eternally true. Natural laws can be discovered in two ways : (i) by correlating a number of experimental facts, as just indicated ; (2) by a speculative method, on the basis of certain hypotheses as to the nature of the phenomena in question. The meaning to be attached to the term hypothesis is best illustrated by an example. In the present chapter, we have seen that the laws of chemical combination are accounted for satisfactorily on the view that matter is made up of extremely small, discrete particles, the atoms. Such a mechanical representation, which is more or less inaccessible to experimental proof, is termed a hypo- thesis. A hypothesis may, therefore, be defined as a mental picture of an unknown, or largely unknown, state of affairs, in terms of something which is better known. Thus the state of affairs in gases, which is and will remain unknown to us, is represented, according to the kinetic theory, in terms of an enormous number of rapidly H4 A TEXT-BOOK OF INORGANIC CHEMISTRY moving, perfectly elastic particles, and on this basis it is possible to deduce certain of the laws which are actually followed by gases (p. 49). There does not appear to be any fundamental distinction in the use of the terms "hypothesis" and "theory." A theory may be defined as a hypothesis, many of the deductions from which have been confirmed by experiment and which admits of the convenient representation of a large number of experimental facts. CHAPTER X DETERMINATION OF ATOMIC WEIGHTS COMBIN- ING WEIGHTS AND CHEMICAL EQUIVALENTS- FORMULAE AND EQUATIONS VALENCY IN Chapter IV. we have been led to assign definite combining weights to each element, and a few of these combining weights are given in the accompanying table. The combining weight of an element has already been defined as the smallest quantity of it which combines with, or takes the place of, I part by weight of hydro- gen, or 8 parts by weight of oxygen. On the other hand, we have made use in the last chapter of numbers called atomic weights ; these are supposed to represent the relative weights of the atoms of other elements, referred to hydrogen as unity, more accurately, to oxygen = 16, as in the case of the combining weights. A comparison of the experimentally found combining weights, and of the atomic weights, for a number of elements is shown in the accompanying table : Hydrogen. Oxygen. Sulphur. Chlorine. Mercury. Combining Weights . 1.008 8.0 16.0 3S-4 6 IO Atomic Weights . . 1.008 16.0 32.0 35-4-6 200 It is evident from these data that the atomic weights are either the same as, or are simple integral multiples of, the combining weignts. There must clearly be some good reason why atomic weights are preferred to combining weights in representing chemical formulas, and these will be fully set forth in the following sections. Determination of Atomic Weights by Volumetric Methods The molecular weight of a volatile substance the weight of the molecule referred to the atom of hydrogen as unit can be determined on the basis of Avogadro's hypothesis ; it is the amount of the substance in grains which, at o and 760 mm. pressure, occupies 22.4 litres. The molecule is made up of atoms, and the molecular weight is the sum of the weights of the atoms. The amount contributed to the molecular weight by each of the "5 n6 A TEXT-BOOK OF INORGANIC CHEMISTRY atoms can of course readily be determined by analysis. Experiment shows that the molecular weight of hydrogen chloride is 36.46 approximately, and an analysis of the compound shows that about i part is hydrogen and the remainder, 35.46, is chlorine. If we know further how many atoms of hydrogen and of chlorine are present in the molecule, we at once obtain the atomic weight. If, for example, only one atom of chlorine is present, the atomic weight of this element must be 35.46, if two atoms are present it must be 17.73. The problem is solved if we know how many atoms of an element are present, but this is just where the difficulty comes in. It may safely be assumed, however, that if we investigate a sufficient number of volatile compounds, we shall meet with some which con- tain not more than one atom in the molecule. The relative amount of the element in those molecules in which it occurs in the smallest proportion is therefore the atomic weight required. The atomic weight of an element is therefore the smallest amount of it which occurs in a molecule of one of its compounds , referred to the atom of hydrogen as unity (strictly 1.008). As an illustration of the method, we will use it to find the atomic weights of hydrogen, oxygen, and chlorine, elements which have been considered in the previous chapters, and also of carbon and mercury. Weights of Constituents in Molecular Weights. Compound. Molecular Weights. Hydrogen. Oxygen. Chlorine. Carbon. Mercury. Hydrogen chloride 36.5 I 35-5 Chlorine monoxide 87 16 7i ... * L - Chlorine dioxide 67.5 32 35-5 ... Water . . . 18 2 16 Methane . . . 16 4 12 Ethylene . . . 28 4 24 Acetylene 26 2 24 Carbon monoxide 28 l6 12 Carbon dioxide 44 32 12 Mercurous chloride 23S-5 35-5 200 Mercuric chloride . 271 7i 2OO From an examination of the table it is an easy matter to pick out the smallest weight of a particular element occurring in the mole- cule of a compound, and this by definition is the atomic weight of the element. Thus, among the compounds of carbon examined DETERMINATION OF ATOMIC WEIGHTS 117 none contain less than 12 parts of that element in the molecule, and therefore 12 is the atomic weight of carbon. In the same way, no known compound contains less than 16 parts of oxygen in the mole- cule, and therefore 16 is the atomic weight of oxygen. In the same way, the table shows that the atomic weight of hydrogen is i (strictly 1.008), that of chlorine 35.5, and that of mercury 200. In some cases a molecule contains a simple multiple of the atomic weight of an element, for example, carbon dioxide contains 32 parts of oxygen. In this case we make the obvious assumption that two atoms of oxygen are present in the molecule. It is, of course, evident that this method gives only an upper limit for the atomic weights, but if sufficient volatile compounds are examined, some are certain to contain only one atom in the molecule. Although it is now clear that atomic weights may be deduced from the results of vapour density determinations alone, by application of Avogadro's hypothesis, yet far greater confidence will be felt in the results if they can be corroborated by other methods. Fortu- nately, there are at least four other methods which may be used for this purpose, and they confirm in the most satisfactory way the conclusions drawn from volumetric data. The more important methods for fixing the relative values of the atomic weights are as follows : 1. Volumetric methods (already dealt with). 2. Purely chemical methods. 3. Methods based on specific heat determinations. 4. Methods based on isomorphism. 5. Methods based on the periodic system of the elements. Methods (2), (3) and (4) will now be briefly considered ; the fifth method will be referred to at a later stage (p. 369). (2) Chemical Methods of Fixing Atomic Weights The first step in using this method is exactly the same as for the volumetric method, that is, an analysis is made of the substance in order to determine the relative amounts of the components. The further procedure may be illustrated by reference to water. Water contains in round numbers i part by weight of hydrogen to 8 parts by weight of oxygen. If the formula for water is HO, it is clear that the atomic weight of oxygen is 8, referred to hydrogen as unit. If, on the other hand, the formula is H 2 O, the ratio of the amounts of the two elements is, as before, i : 8 or 2 : 16, and therefore the atomic weight of oxygen is 16. If, on the other hand, the formula is H 3 O, n8 A TEXT-BOOK OF INORGANIC CHEMISTRY the atomic weight of oxygen must be 24 in order to maintain the experimentally found ratio I : 8. Now it has been found that the hydrogen can be displaced from water in two stages (by means of metallic sodium, for instance), and in two stages only. As not less than one atom can be displaced, it follows that there are at least two (and probably only two) atoms of hydrogen in the molecule of water, the formula is therefore H 2 O, and the atomic weight of oxygen is 16. The last conclusion is confirmed by the experimental fact that the oxygen cannot be displaced in stages, and therefore only one atom is present. The method will now be further illustrated by application to two compounds of carbon with hydrogen, called methane and ethylene respectively. This case is of historical interest, as it led Dalton to the discovery of the law of multiple proportions. Methane contains 3 parts of carbon to i of hydrogen (in other words, the equiva- lent of carbon is 3), and ethylene contains 6 parts of carbon to i of hydrogen. Dalton, who based his atomic weights to some extent on assumed simplicity of composition, was of opinion that ethylene contained one atom of each element, and hence that the atomic weight of carbon was 6. We now know, however, that one- fourth of the hydrogen in methane can be displaced by chlorine ; hence methane contains four' hydrogen atoms, and to retain the experimental ratio of 3 : i the atomic weight of carbon must be 12, provided only one atom is present. The formula for methane is therefore CH 4 , and it can be shown by similar reasoning that the formula of ethylene is C 2 H 4 . The same conclusion can be reached more readily by the volumetric method, as shown in the table, p. 116. (3) Fixing of Atomic Weights from Specific Heats Dulong and Petit, in 1818, made the remarkable observation that when the specific heats of the elements in the solid form were multiplied by the respective atomic weights, the products were in all cases approximately the same, the average value being about 6.2. This is illustrated in the table on the following page, the values given for the specific heats being the mean between 17 and 100. The mean value of the product for the metals quoted in the table is about 6.1, and it will be seen that only in the case of silicon is there a serious deviation from the average value, the product being much smaller. The same applies to three other light elements, the names and products of which are as follows : beryllium, (9.1) 3.7 ; boron, (n) 2.6; carbon, (12) 2-2.8. It has been shown, however, that the specific heats of these elements increase very rapidly with the tern- DETERMINATION OF ATOMIC WEIGHTS 119 perature, so that the deviation from the normal value of the product is much smaller at high temperatures. Element. Atomic Weight. Specific Heat. Product. Lithium . . . 7 0.94 6.6 Magnesium 24.4 0.247 6.0 Aluminium 27.1 0.217 5-9 Silicon . . . 28.4 0-175 5-o Iron . . . 56 O.I 10 6.2 Copper . . . 63-4 0.093 8 Silver . . . 108 0.056 6.1 Tin . . . H9 . 0.0556 6.6 Antimony . . . 120 0.0503 6.0 Platinum . . . . *95 0.0310 JU- Gold . . . . 197 0.0310 6.1 Mercury . . . 200 -0335 6.7 Bismuth . . . . 208.5 0.0303 6-3 Another way of regarding Dulong and Petit's law is that quantities of different substances in the ratio of their atomic weights require the same expenditure of heat to raise them through an equal number of degrees. The product of specific heat and atomic weight may be called the atomic heat, and the purport of Dulong and Petit's law is that the atomic heats of all solid elements are approximately equal (6.2 calories being required to raise the atomic weight in grams i in temperature). It is certainly a most remarkable fact that the amounts of heat required to raise 7 grams of lithium and 208 grams of bismuth through the same number of degrees are equal. The method of using Dulong and Petit's law to fix the atomic weights will be obvious from the above. Once the law gained the confidence of chemists, it proved very valuable in deciding whether a number or some multiple or sub-multiple represented the true atomic weight. (4) Fixing of Atomic Weights by Isomorphism It was ob- served by Mitscherlich (1822) that the salts known as sodium hydrogen phosphate and sodium hydrogen arsenate separate from solution with the same number of molecules of water, i2H 2 O, are identical or nearly so in crystalline form, and when a solution containing the mixed salts is evaporated crystals containing both substances separate. It is clear that there is a very great similarity between the crystalline forms of the two salts, which is expressed by saying that the salts are isomorphous, or have the same crystalline form. On the basis of 120 A TEXT-BOOK OF INORGANIC CHEMISTRY these and similar observations Mitscherlich brought forward his Law of Isomorphism^ according to which compounds of the same crystal- line form are of analogous constitution. Thus when one element replaces another in a compound without alteration of the crystalline form, it is assumed that they displace each other atom for atom. A little consideration shows that a method of fixing atomic weights may be based on the above statements ; if an element of known atomic weight is displaced by another of unknown atomic weight, the amounts are in the ratio of the atomic weights. To take a simple case, the two salts, potassium chloride and potassium iodide, are isomor- phous, and we assume that the atomic weight of chlorine is known. Analysis shows that potassium chloride contains 39 parts of potassium to 35-5 parts of chlorine, and that for the same amount of potassium the other salt contains 127 parts of iodine. As iodine is assumed to have displaced chlorine atom for atom, it follows that 127 is the atomic weight of iodine. Further reference to crystal form and isomorphism is made at a later stage. The accepted atomic weights for the elements are given in the table published by an International Committee (1911), which for convenience of reference is printed on the back page of the cover of the book. Combining Weights, Chemical Equivalents, Atomic "Weights From the foregoing sections the advantage of using atomic weights in place of combining weights will be evident. If, for example, we use the combining weight for oxygen, 8, the least number of combining weights in a molecule will be 2, and this intro- duces a needless complication. In the same way, each molecule of a carbon compound would contain 4 or some multiple of four combin- ing weights. Still another advantage of the atomic weights over the combining weights is seen in Dulong and Petit's law. There is no such con- nexion between the specific heats and the combining weights as has been shown to hold for specific heats and atomic weights. Finally, it will be shown that the remarkable relationships between the atomic weights according to the periodic system are entirely absent when combining weights are used instead. The combining weights may be regarded from a slightly different point of view. The following represents the quantities of a number of elements which unite with 8 parts by weight of oxygen : Hydrogen. Sodium. Zinc. Silver. Copper. Mercury. I 23 32.5 108 31.7 100 COMBINING WEIGHTS 121 and these are the numbers hitherto termed combining weights. As, however, they are the amounts of different elements which displace the same amount of hydrogen when acted on by acids, it is evident that these amounts are equivalent to i part by weight of hydrogen or to 8 parts by weight of oxygen, and therefore they are often termed chemical equivalents. The chemical equivalent of an element is defined as that amount of it which combines with or takes the place of i part (strictly, 1.008 parts) by weight of hydrogen (or 8 parts by weight of oxygen). It is clear that the chemical equivalents are practically identical with what have hitherto been termed combining weights. There is, however, a slight difference, due to the fact, which will be fully considered later, that certain elements have more than one chemical equivalent. When an element has only one chemical equivalent, the combining weight and the chemical equivalent are identical. When, however, the element has more than one equi- valent, there is invariably a simple integral ratio between them, and the combining weight corresponds with the smallest of the equivalents. Establishment of Molecular Formulae When the atomic weights are known, it is a relatively simple matter to establish the molecular formula of a chemical compound. In the first instance, the percentage composition must be known. This can be determined by direct analysis or by synthesis, that is, by finding the relative amounts of the components which unite to form the compound. The next step is to divide the relative proportions of the different elements present by the respective atomic weights ; this, reduced to its simplest terms, gives us the ratio between the number of atoms in the molecule. The formula thus obtained is termed the empirical formula, and the molecular formula may be the same as, or a simple multiple of, the empirical formula. In order to settle which multiple is to be used, the molecular weight has finally to be determined. The method will be fully understood from an example. It was found by analysis that hydrogen peroxide (p. 138) contains 5.93 per cent, of hydrogen and 94.07 per cent, of oxygen. Dividing these numbers by the respective atomic weights in order to obtain the relative number of atoms, we have 5.88 : 5.88 or i : i, that is, the same number of atoms of each element are present, and the empirical formula is HO. The molecular formula is therefore (H.O)* where x is an integral which can be found from the molecular weight. The latter determined directly is 34, which is satisfied by making .r=2, so that the formula of hydrogen peroxide is H 2 O 2 . 122 A TEXT-BOOK OF INORGANIC CHEMISTRY When the components and the substance itself can be obtained in gaseous form, the establishment of the molecular formula is still simpler. It has been found, for instance, that two volumes of am- monia yield on decomposition one volume of nitrogen and three volumes of hydrogen when all the gases are measured under the same con- ditions. Applying Avogadro's hypothesis, that one molecule occupies two unit volumes, we have 2 vols nitrogen + 6 vols hydrogen->4 vols ammonia Since the four volumes of ammonia must also contain two molecules, it follows at once that the formula for ammonia is NH 3 . In this case it is necessary to know the number of atoms in the molecules of hydrogen and of nitrogen, and it must be remembered that the mole- cules of all substances do not contain two atoms. These methods of establishing molecular formulas, and modifica- tions of them, will be frequently used throughout the book. Equations and Calculations As we are now familiar with the methods of representing chemical compounds by their formulae, we are in a position to represent the chemical changes considered in the previous chapters much more concisely. The combination of hydrogen and chlorine to form hydrogen chloride, for example, is represented by the equation and the combination of hydrogen and oxygen to form water as follows : As chemical reactions take place between molecules, it is preferable to write the equation representing the formation of water as above, instead of the method occasionally used, H 2 + O = H 2 O. In representing a chemical change by an equation, all the reacting substances and all the products must be known, as well as their respective molecular formulas. The first step is to put on the left- hand side the formulae for the reacting substances, and on the right- hand side the formulae for the products. The next step is to satisfy the law of the conservation of mass, that the sum of the weights of the products is equal to the sum of the weights of the reacting sub- stances. The simplest way of doing this is to ensure that the FORMULA AND EQUATIONS 123 same atoms and the same number of each appear on the two sides of the equation. These statements will now be illustrated by writing the equation representing the action of sulphuric acid, H 2 SO 4 , on zinc, Zn, giving zinc sulphate, ZnSO 4 , and hydrogen, H 2 , as follows : Zn + H 2 SO 4 = ZnSO 4 + H 2 . It is clear that each atom occurs the same number of times on the two sides of the equation, so that the latter is complete as it stands. A slightly more complicated case is the effect of heat on potassium chlorate, KC1O 3 , giving rise to potassium chloride,, KC1, and oxygen, O 2 . The first stage is as follows : but it is evident that in this form the atoms do not occur the same number of times on each side ; according to the usual expression the equation is not " balanced." For this purpose, sufficient potassium chlorate must be taken to yield an even number of oxygen atoms. The completed equation is as follows : Proceeding in a similar way, the reversible reaction between water and chlorine at high temperatures is represented by the equation and the action of steam on iron, which is also reversible, by the equation The student should now try to write a number of equations for himself, bearing in mind that all the reacting substances and products, as well as their formulae must first be known. Calculations It has already been pointed out that the symbols in a chemical formula have a quantitative significance, inasmuch as they stand for quantities of the different substances proportional to their atomic weights. This consideration at once enables us to calculate from a chemical equation the relative amounts of the react- ing substances and the products. This may be illustrated by the equation representing the action of sulphuric acid on zinc. Zn + H 2 SO 4 = ZnSO 4 + H 2 2 2 i2 4 A TEXT-BOOK OF INORGANIC CHEMISTRY The sum of the weights of the atoms in each molecule is found, and the equation shows that 65 parts of zinc react with 98 parts of pure sul- phuric acid to give 161 parts of zinc sulphate and 2 parts of hydrogen. When gases take part in a chemical change, we can further calculate the relative volumes concerned, bearing in mind th,e result already given (p. 1 10) that the molecular weight of any gas in grams occupies 22.4 litres at normal temperature and pressure. If, in the above equation, the amounts are expressed in grams, it is evident that 65 parts of zinc liberate 2 grams or 22.4 litres of hydrogen measured under normal conditions. The statements will now be illustrated by means of further examples, (i) What weight and what volume of oxygen, measured under normal conditions, will be obtained by completely decomposing 10 grams of mercuric oxide ? The equation is as follows : 400 + 32 32 and shows that 432 grams of mercuric oxide yield 32 grams of oxygen, so that 10 grams yield ^x 32 = 0.74 grams of oxygen. Further, since 432 grams of mercuric oxide yield 32 grams = 22. 4 litres of oxygen at N.T.P., 10 grams will yield - x 22.4 = 0.518 litre or 518 c.c. of oxygen. (2) What weight of chlorine will be obtained by heating 100 c.c. of a 10 per cent, solution of hydrochloric acid with excess of manganese dioxide, and what weight of sodium chloride can be obtained by complete combination of the chlorine thus formed with sodium ? The first reaction (compare p. 88) is represented by the equation 4x36.5 2x35.5, As only the relationship between the weights of hydrochloric acid and of chlorine is required, the molecular weights of the other substances need not be calculated. As 146 grams of pure hydrogen chloride yield 71 grams of free chlorine under the conditions described, 10 grams of hydrogen chloride (the amount present in 100 c.c. of a 10 per cent, solution) yield- -^x 10 = 4.86 grams of chlorine. FORMULA AND EQUATIONS 125 The second equation concerned is 46 71 117 And as 71 grams of chlorine yield 117 grams of sodium chloride, 4.86 grams of chlorine yield 4.86 x = 8.0 grams of sodium chloride. The relationships between the volumes of gases concerned in chemical changes can be deduced very simply on the basis of the deduction from Avogadro's hypothesis already mentioned (p. 109), that the molecule of any substance in the form of vapour occupies two unit volumes. As already indicated, this has no reference to actual numerical values, but is merely used for purposes of compari- son. On this basis the volumes of the gases concerned in the re- versible reaction between steam and chlorine (p. 86) are as follows : 2C1 2 + 2H 2 O$4HC1 + O 2 4 vols. 4 vols. 8 vols. 2 vols. This equation shows that if the reaction proceeds completely from left to right there is an expansion of 8 volumes to 10 ; if from right to left there is a contraction from 10 volumes to 8. In the foregoing we have assumed that the formulae are known, and have deduced the corresponding volume changes. It is evident, however, that the rule may be employed in the converse way to deduce chemical formulas on the basis of the observed changes of volume accompanying chemical changes. It is on this principle that the formula for ammonia has already been deduced (p. 122), and numerous other illustrations will be met with later (cf. nitric oxide, p. 232, and sulphur dioxide, p. 285). In the present section we have on several occasions made use of the molecular weight of a substance in grams. As these amounts are used very largely in chemistry, it is convenient to have a shorter name, and for this purpose the term mol, introduced by Ostwald, is suitable. According to the molecular theory, mols of different substances contain the same number of molecules in each case, and are naturally often employed for comparative purposes. Practical Determination of Chemical Equivalents The chemical equivalent of an element has already been defined as that amount of it which combines with or displaces 8 parts by weight of oxygen, or 1.008 parts by weight of hydrogen. It is desirable 126 A TEXT-BOOK OF INORGANIC CHEMISTRY that the student should determine some equivalents, and in order to illustrate the methods employed, a few examples will be given. (1) Determination of the Equivalent of Zinc The equivalent of zinc (and of other metals which readily dissolve in acids, giving off hydrogen) may be determined by dissolving a known weight of zinc in excess of hydrochloric or sulphuric acid and measuring the hydrogen evolved. A convenient form of apparatus for this purpose is shown in Fig. 40. The wide graduated tube A, open at the lower end, is provided with *a side tube connected by rubber tubing to a glass tube pass- ing through the rubber cork closing the wide-mouthed bottle B. At the commencement of the experiment a weighed amount of zinc is placed in B, as is also a small bottle containing more than sufficient acid to dissolve the zinc. The cork is then inserted, and the water in the wide graduated tube brought to a convenient height by means of the short rubber tube and clip at the top of A ; the latter is then closed by the clip. The bottle B is now inclined, so that the acid comes in contact with the zinc ; the hydrogen given off displaces the water in A. When all action has ceased, the graduated tube is arranged so that the water stands at the same level outside and inside, the volume of the hydro- gen is then read off, and the tempera- ture of the water in the jar C noted. The volume of the gas is then corrected to normal temperature and pressure (p. 45), and the weight of the gas calculated in the usual way (p. 124). The weight of zinc which would give I part by weight of hydrogen is by definition the chemical equivalent. (2) Determination of the Equivalent of Copper by Analysis of the Oxide The equivalent of copper can readily be determined by reducing a known weight of the oxide to copper by means of hydrogen and weighing the resulting copper. The arrange- ment of the apparatus is shown in Fig. 17. The bulb tube is weighed empty, and then again after a small quantity of dried copper FIG. 40. FORMULA AND EQUATIONS 127 oxide (say, i gram) has been placed in it. The bulb is heated and dry hydrogen passed over it till reduction is complete, as shown by the colour ; it is then allowed to cool in the stream of hydrogen and subsequently weighed. The loss of weight represents the oxygen with which the copper was combined ; the amount of the latter is the difference between the final weight of the bulb and the weight of the latter when empty. From the results the chemical equivalent the amount of copper combined with 8 parts by weight of oxygen can readily be calculated (p. 55). The same experiment has already been referred to in connexion with the composition of water (p. 58). (3) Equivalent of Copper by Displacement with Mag- nesiumA weighed amount of magnesium ribbon, say, 0.25 gram, is put into excess of a warm solution of copper sulphate. After some time it will be found that the magnesium has completely disappeared, and a red precipitate of copper is obtained. The precipitate is separated by filtration, washed, dried, and weighed. It will be found that 0.25 gram of magnesium displaces about 0.65 gram of copper. The chemical equivalent of magnesium is about 12.2, so that i equivalent of magnesium displaces 12.2 x = 31.7 grams or i equivalent of copper. Practical Determination of Vapour Densities In the present chapter we have had many illustrations of the great part which a knowledge of vapour densities has played in the development of chemical theory. The principle of the methods employed in vapour density determinations is extremely simple. The volume occupied by a known weight of the vapour under known conditions of temperature and pressure is found, and the ratio of this weight to the weight of an equal volume of hydrogen under the same conditions is the required density. Space only admits of a brief account of two of the methods in actual use ; for a full account of vapour density deter- minations a text-book of physics should be consulted. (i) Victor Meyers Method The method most largely used was introduced by Victor Meyer in 1878. The most remarkable feature of the method is that it is not the volume of the vapour itself which is measured, but that of an equal volume of air displaced by the vapour. The apparatus consists of a cylindrical vessel, A, of about 200 c.c. capacity, ending in a long neck provided with two side tubes, as shown in Fig. 41. One of these side tubes, /, from which the displaced air issues during an experiment, is bent in such a way that 128 A TEXT-BOOK OF INORGANIC CHEMISTRY t, its free end can be brought under the surface of water in a suitable vessel. On the other tube, / 2 , is fitted a rubber tube carrying a glass rod, which can be moved outwards and inwards, and, at the commencement of the experiment, serves to retain in place the small glass bulb shown in the figure, containing a weighed quantity of the liquid the vapour density of which is to be determined. The top of the main tube is closed by a cork, which is kept in place throughout an experiment, and a little asbestos or mercury is placed in the bottom to guard against fracture of the glass when the bulb drops. The apparatus is kept at a constant temperature throughout the greater part of its length by means of the vapour of a liquid boiling in the outer bulb- tube B ; the temperature should be at least 20 above the boiling-point of the liquid to be vaporized. At the commencement of an experiment the bulb and rod are placed in position and the cork inserted ; the jacketing liquid is then boiled till air ceases to issue from the end of the tube t and bubble through the water, showing that the tempera- ture inside the bulb A is constant. A graduated measur- ing tube, C, full of water, is then inverted over the end of the delivery tube, and the small bulb allowed to drop by drawing back the glass rod. When air ceases to issue from the end of the delivery tube, the graduated tube is closed by the thumb, removed to a deep vessel containing water, allowed to stand till the temperature is constant, and the volume of air read off when the water outside and inside is at the same level. The temperature inside the tube A is the same before and after the experiment ; the only difference in the conditions is that a certain volume of air is displaced by an equal volume of vapour. The observed volume of air is therefore that which the vapour would occupy after reduction to the temperature and pressure at which the air is measured (provided that the vapour and air are equally affected by changes of temperature and pressure, which is approximately the case under suitable conditions). The temperature is that of the water, and the pressure that of the atmosphere less the vapour pressure of water at the temperature of observation. It is evident that it is not necessary to know the temperature at which the sub- stance is vaporized, and this is one of the advantages of the method. B FIG. 41. VALENCY 129 The mode of calculating vapour densities and molecular weights from the observed data may be illustrated b*y the following example : 0.220 grams of chloroform when vaporized displaced 45.0 c.c. of air, measured at 20 and 755 mm. pressure. As the vapour pressure of water at 20 is 17.4 mm. (p. 65), the actual pressure exerted by the gas is 737.0 mm. Therefore, as 0.220 grams of vapour, at 20 and 737.6 mm. pressure, occupy 45.0 c.c., we have to find what weight in grams will occupy 22,400 c.c. at 273 abs. and 760 mm. pressure, and this will be the required molecular weight. Now 45.0 c.c. at 20 and 737.6mm. pressure measured 45 x ^rr x ~^6 c 7 == 4-7 c.c. at N.T.P., and weigh 0.220 grams. Therefore, 22,400 c.c. at N.T.P. weigh 0.220 x ~'^ = 121, which is the required molecular weight, as compared with the value (119.5) calculated from the formula of chloroform. CHC1 3 . The vapour density can be calculated from the experimental data as follows. We have seen that 40.7 c.c. of the vapour weigh 0.220 grams at N.T.P. As 2 grams of hydrogen measure 22,400 c.c. at N.T.P. ,40. 7 c.c. of hydrogen weigh 2 x -7 = 0.00363 grams. Hence Wt. of substance 0.220 Vapour density = ^777 r -- i -- i FTT ~~ ^ = 60.6. 3 Wt. of equal vol. of H 2 0.00363 The molecular weight, being by definition double the vapour density, is 121, as has just been shown by the alternative method. (2) Hofinanrts Method A graduated glass tube is filled with mercury, and inverted in a bath of the same metal. A small glass tube, containing a weighed quantity of the substance whose vapour density ha^ to be determined, is inserted under the lower edge of the graduated :ube ; it rises to the top, and the liquid vaporizes, displacing part of the mercury. The volume of the vapour, its temperature and pressure, are then read off, and from these data and the weight of the substance the vapour density and molecular weight are calculated in the usual way. The advantage of Hofmann's method is that it admits of the determination of densities under reduced pressure. Valency and Structural Formulae We have learnt in previous chapters that the formula of hydrogen chloride is HC1, and of water H 2 O. One atom of chlorine always combines with one, and not more than one, atom of hydrogen ; similarly, one atom of oxygen combines with two atoms of hydrogen, neither more nor less. 9 130 A TEXT-BOOK OF INORGANIC CHEMISTRY It is evident that chlorine and oxygen have a different combining capacity for hydrogen. The combining capacity or chemical value of an element is termed its valency, and is usually measured with regard to hydrogen. An element such as chlorine, one atom of which com- bines with one atom of hydrogen, is said to be univalent, whilst oxygen, one atom of which combines with two atoms of hydrogen, is said to be bivalent. Elements of valency three are trivalent or tervalent, of valency four quadrivalent^ and so on. Many elements, however, such as sodium and potassium, do not form well-defined compounds with hydrogen. When we compare the formula of sodium chloride (common salt), NaCl, with that of hydro- chloric acid, HC1, it is evident that sodium and hydrogen displace each other atom for atom. Sodium and hydrogen are, therefore, of equal chemical value, and sodium is also univalent. Based on the foregoing considerations, the valency of an element is measured by the number of hydrogen atoms which it can combine with or displace. The question now arises as to whether the valency of an element is constant or variable. This question was widely debated by chemists for many years, but it is now agreed that the same element may have different valencies in different compounds. The highest valency shown by any element is eight. Hydrogen is invariably univalent, as are sodium and potassium. Oxygen is almost always divalent. Chlorine is univalent with regard to hydrogen, but in binary compounds with oxygen and in ternary compounds with oxygen and hydrogen, appears to have several different valencies. Barium and calcium appear to be always bivalent. The valencies of the elements will be discussed in con- nexion with their detailed consideration. A knowledge of the common valencies of the elements is of great value in writing chemical formulae. We have seen that salts arc derived from acids by displacement of the acidic hydrogen by metals. Suppose we wish to know the formula of barium chloride. It is derived from hydrochloric acid by displacing the hydrogen by barium. As the latter is a divalent element, and therefore displaces two atoms of hydrogen, the chloride is derived from 2HC1, and its formula is consequently BaCl 2 . Groups of elements may also be regarded as having a definite valency. Thus since the formula of sulphuric acid is H 2 SO 4 , the group SO 4 is bivalent, as it is associated with two atoms of hydrogen. Similarly, the group C1O 3 is univalent, as it is associated with one VALENCY 131 atom of potassium, a univalent element, in potassium chlorate, KC10 3 . The valencies in a compound may be shown by means of bars or links between the constituent atoms, each bar representing a single valency. The formula for water is written thus : H O H, showing that the two valencies of oxygen are satisfied by the two atoms of hydrogen. Such formulae are termed graphic, structural or constitu- tional formulae. The graphic formula of hydrogen chloride is H - Cl. The formulae of the oxides of chlorine, C1 2 O and C1O 2 , may serve as further illustrations. The first oxide is written thus : Cl - O - Cl, showing that it is of the same type as water. The other oxide, C1O 2 , presents more difficulty from this point of view. It may be written O = C1 = O. showing the chlorine as quadrivalent ; but other methods of formulation may be used. We shall see in detail later that the graphic formula of a compound is an attempt to represent in a brief way its characteristic behaviour. As the be mviour of chemical compounds is many-sided, it is not surprising that opinions often differ with regard to the most suitable graphic formulae of particular compounds. Chemical Equivalent, Atomic Weight, and Valency The chemical equivalent of an element has been defined as that quantity of it which combines with or displaces one part by weight of hydrogen. The atomic weight of an element, on the other hand, may displace one or more parts by weight (one or more atoms) of hydrogen, according as the element is univalent or polyvalent. If univalent, the atomic weight of an element displaces one part by weight of hydrogen : Na + HOH->NaOH + H if divalent, two parts by weight of hydrogen : Zn + H 2 SO 4 ->ZnS0 4 +H 2 and so on. From the foregoing it is evident that Atomic weight Valency : Chemical equivalent. This important relationship should be carefully remembered. If the atomic weight and the chemical equivalent of an element have been determined independently by the methods described in the present chapter, the valency can at once be deduced. If, on the other hand, 132 A TEXT-BOOK OF INORGANIC CHEMISTRY the atomic weight and valency are known, the chemical equivalent can at once be calculated. In the light of this statement, the chemical equivalents on p. 115 should be compared with the corresponding atomic weights. Summary The last two chapters deal with many of the funda- mental principles of chemistry, and it will be useful briefly to summarize some important points. The basis of chemical theory is Avogadro's hypothesis, which states that equal volumes of all gases, under the same conditions of temperature and pressure, con- tain the same number of molecules. This principle enables us at once to fix the molecular weights of volatile substances, which are referred to the atom of hydrogen as unit. From the data on molecular weights, the atomic weights of the elements can be determined, the atomic weight being by definition the smallest quantity of an element which occurs in a molecule, referred to the atom of hydrogen as unit. When the atomic weights are known, the molecular formula of a compound can readily be established as follows. The percentage composition is determined by analysis, the relative proportion by weight of each element is then divided by the corresponding atomic weight, and the result, reduced to its simplest terms, gives us the ratio between the number of atoms in the molecule, the so-called empirical formula. The molecular formula is the same as, or an integral multiple of, the empirical formula (p. 121). It must be emphasized that the methods of determining molecular weights only give approximate values, which are sufficiently accurate to show which multiple of the empirical formula is to be taken. Molecular formulae can also be established by applying the deduction from Avogadro's hypothesis that the molecule of any substance in the state of vapour occupies two unit volumes. The atomic weights deduced on the basis of Avogadro's hypothesis are in full agreement with those determined by the alternative methods. Similarly the molecular formulas obtained as above are in agreement with the conception that chemical reactions take place between molecules. It is, in fact, possible to determine the molecular weight of a substance from its chemical behaviour, on the assumption that the molecule is the smallest quantity of a substance which can take part in a chemical change. The exact connexion between atomic weights and chemical equivalents, and the advantages of representing composition in terms of atomic weights instead of in terms of equivalents, have been fully explained in the course of the chapter. CHAPTER XI OZONE AND HYDROGEN PEROXIDE- THERMOCHEMISTRY IN the present chapter we are concerned with two interesting sub- stances closely allied to two elements, hydrogen and oxygen, already dealt with in detail. We shall learn that ozone is simply a form of oxygen, whereas hydrogen peroxide, as its name indicates, is, like water, a chemical compound of hydrogen and oxygen. Both substances are energetic oxidizing agents. OZONE Formula, O 3 . Molecular weight =48. Density=;24. History It has long been known that when an electrical machine is in operation a peculiar smell, somewhat like that of dilute chlorine, is noticeable in its vicinity, and in 1785 Van Marum showed that it was due to the action of the electrical discharge on the oxygen of the atmosphere. The same smell is sometimes noticed after a thunder- storm. In 1840, Schonbein showed that the smell was due to the formation of a definite substance, which he named ozone (oeit>, to smell), and described other methods by which the new substance could be obtained. Occurrence It is generally assumed that traces of ozone are normally present in the atmosphere, but recent investigations render this extremely doubtful (see below). Preparation (i) Ozone, mixed with a large excess of oxygen, can be obtained by passing a silent electric discharge through oxygen. A suitable arrangement for demonstration purposes is represented in Fig. 42. It consists of a wide glass tube AA, into which a cork carrying a narrower tube B is fitted as shown. The outer tube is covered with tinfoil, and connected to one pole of an induction coil ; the inner cube contains some conducting material which is connected to the other pole of the coil. Under these circumstances the discharge 133 i 3 4 A TEXT-BOOK OF INORGANIC CHEMISTRY takes place between the two tubes, not in sparks, but as a luminous glow ; this is termed a silent discharge. Oxygen or air is then passed through the space between the tubes, and it can be shown by the characteristic smell, and by application of the tests described below, that the issuing gas contains ozone. The oxygen is always present in large excess, but under favourable conditions a mixture containing 5 to 6 per cent, of ozone can be obtained by means of the silent^discharge. The yield of ozone is increased by keeping the temperature low. (2) The oxygen obtained at the anode when dilute sulphuric acid is electrolyzed contains more or less ozone, depending on the condi- tions. It has recently been shown by Fischer and Massenez that under favourable conditions a mixture containing no less than 28 per cent, by weight of ozone can be obtained in this way. The best yields were obtained with solutions containing 10 to 15 per cent, by weight of the acid and a high current density. The anode was of platinum, the exposed surface being very small in order to minimize sub- sequent decomposition of the ozone, and ar- rangements were made jr ]G 42 for keeping the tem- perature low. (3) The oxygen formed by a number of the chemical methods already mentioned, for example, by the action of concentrated sul- phuric acid on potassium permanganate or manganese peroxide, contains sufficient ozone to answer the ordinary tests. (4) Ozone is formed during the slow oxidation of phosphorus at the ordinary temperature. If a freshly scraped stick of phosphorus is exposed to the air for a short time in a covered gas jar, the smell of ozone can readily be detected, and a strip of moist filter-paper impregnated with potassium iodide and starch (see tests for ozone) suspended in the jar will turn blue. Physical Properties At ordinary temperatures ozone is a gas which, in thick layers, is bluish in colour. Under atmospheric pres- sure it boils at 1 19, so that it can be obtained in the liquid form by passing the gas, mixed with oxygen, through a U-tube immersed in liquid oxygen (boiling-point - 182.5). In this way a liquid is obtained OZONE AND HYDROGEN PEROXIDE 135 containing only a small proportion of liquid oxygen, and the latter can be almost completely removed by evaporation, being much more volatile than ozone. Liquid ozone is deep indigo-blue in colour, and is extremely explosive. Ozone is much more soluble in water than is oxygen. Chemical Properties Pure gaseous ozone is so unstable as to be practically unknown, but ozone mixed with excess of oxygen is fairly stable at ordinary temperatures. When the mixture is heated to 250-300, however, it is rapidly reconverted to oxygen. This may readily be demonstrated by passing the gas from the ozone apparatus through a glass tube heated by a Bunsen flame, when the issuing gas will no longer give the tests for ozone. The facts just mentioned show that the equation representing the formation of ozone from oxygen is reversible ; the action represented by the upper arrow is favoured by the silent electric discharge (which supplies the neces- sary energy), but when a fairly large proportion of ozone is present the reaction proceeds rapidly, and almost completely, in the direction indicated by the lower arrow at 250. The interesting relationships between oxygen and ozone are considered more fully in the chapter on chemical equilibrium (p. 174). When ozone splits up into oxygen, a large amount of heat is given out. The reaction 2O 3 ->3O 2 liberates about 68,200 cal. (Jahn) when gram-molecular quantities are used, that is, when 96 grams of ozone are transformed to oxygen. It follows at once by Le Chatelier's theorem (p. 171) that raising the temperature must favour the reaction which proceeds with absorption of heat ; that is, increase of temperature must favour the formation of ozone from oxygen, and this conclusion is fully borne out by experiment. The apparent contradiction between this result and the rapid decomposition of ozone at 250 will be dealt with larer (p. 174). The most characteristic chemical property of ozone is that it is a powerful oxidizing agent, being much more active than free oxygen in this respect. It oxidizes both mercury and silver at room tempera- ture, the latter becoming blackened owing to formation of an oxide, Ag 2 O 2 . Ozone readily bleaches indigo, litmus, and some other colouring matters, and destroys india-rubber connexions. When passed into a solution of potassium iodide containing a little mucilage 136 A TEXT-BOOK OF INORGANIC CHEMISTRY of starch, the solution turns blue. The reaction in this case is represented by the equation that is, potassium hydroxide, oxygen and free iodine are formed, and the latter gives rise to the deep blue colour with starch. The reaction just described is often made use of as a test for ozone as follows. A little starch is boiled with a few c.c. of water for some minutes, and a few crystals of potassium iodide dissolved in the thick liquid. Strips of filter-paper are then smeared with the solution and when dried constitute the so-called iodide of potassium starch paper, which at once turns blue when moistened and exposed to ozone. We shall find, however, that this test paper is turned blue by many oxidizing agents besides ozone. The last equation serves to illustrate the important fact that when ozone acts as an oxidizing agent oxygen is generally set free. Further, since both O 3 and ,O 2> being single molecules, occupy two unit volumes, the volume of the resulting oxygen is the same as that of the ozone. It is therefore evident that the oxidation is effected by the extra atom of oxygen in the ozone molecule. The oxidizing power of ozone is closely connected with the fact that it has much more energy than an equal amount of oxygen. On the commercial scale ozone is employed for freeing water from micro-organisms (p. 61), for bleaching flour and for other purposes. Formula of Ozone The facts that ozone can be obtained from pure oxygen alone, and that on heating oxygen and oxygen only is obtained, prove that it consists simply of oxygen. In the foregoing we have assumed that whereas the molecule of oxygen contains two atoms, O 2 , that of ozone contains three atoms of oxygen, and is therefore represented by O 3 . It remains to indicate the evidence on which this formula is based. (i) The question could, of course, at once be settled by determining the density of the pure gas, but, as has been pointed out, the latter cannot be prepared. The difficulty has, however, been got over by determining the density of a mixture of oxygen and ozone con- taining a known proportion of ozone (Ladenburg). Suppose, for example, the weight of a litre of the mixture at N.T.P. containing 100 c.c. of ozone is found to be 1.500 grams. The weight of i litre of oxygen at N.T.P. is 1.429 grams, and of 900 c.c. 1.286 grams. The remainder, 1.500-1.286 = 0.214 grains, represents the weight of 100 c.c. of ozone at N.T.P. The weight of ozone which would OZONE AND HYDROGEN PEROXIDE 137 occupy 22.4 litres at N.T.P., that is, the molecular weight of ozone (p. ii o), is therefore 0.2 14X^^5 = 48 grams approximately, whence 100 it follows that the formula for ozone is O 3 . (2) Another proof depends upon the fact that ozone is completely absorbed from its mixture with oxygen by shaking with ordinary turpentine. Now, it has been found that when ozone is formed from oxygen there is a definite diminution of volume, and when the ozonised oxygen is shaken with turpentine a further diminution of volume double the former one is observed. Suppose, for example, 100 c.c. cf oxygen become 95 c.c. when partly ozonized ; the volume is reduced to 85 c.c. by shaking with turpentine. It is evident that the 10 c.c. of ozone absorbed by the turpentine were formed from 100-85=^15 c.c. of oxygen, so that three volumes of oxygen yield two of ozone. This observation can only be satisfied if the change of oxygen to ozone is represented by the formula 3O 2 ->2O 3 , according to which 6 unit volumes of oxygen yield 4 unit volumes of ozone (p. 109). (3) It has been pointed out (diffusion of gases, p. 47) that the relative rate of diffusion of gases is inversely as the square roots of their densities. Soret found that ozone diffused more slowly than oxygen, and that the relative rates were in accordance with the conclusion that the density of ozone is 24 when oxygen is taken as 1 6. TestH for Ozone The tests for ozone are rather important in connexion with the disputed question as to the presence of this gas in the atmosphere. The oxidizing properties already mentioned, such as the liberation of iodine from potassium iodide solution, the bleaching of indigo, etc., are not characteristic, as one or other of them is also given by other substances such as hydrogen peroxide and oxides of nitrogen, although not by oxygen. The following are useful. (a) \\ hen a strip of potassium iodide starch paper, coloured violet with litmus, is exposed to an atmosphere containing ozone, potassium hydroxide, as well as iodine, is formed (see above), and the former turns the litmus paper blue. Hydrogen peroxide also answers this test. (b) When very pure mercury is exposed to ozone, the metal becomes less mobile, and adheres to the glass. This effect is doubtless due to the formation of traces of oxide. (c) A bright silver surface is blackened on exposure to ozone, 138 A TEXT-BOOK OF INORGANIC CHEMISTRY owing to the formation of silver peroxide. This test is quite characteristic, but is not very sensitive. (d) An alkaline solution of an organic compound known as meta- phenylene diamine becomes red in ozone ; neither hydrogen peroxide nor nitrous acid answer this test. As already mentioned, it is very doubtful whether ozone is a normal constituent of the atmosphere. It is now known that the positive results of the earlier tests were due mainly to the presence of nitrous fumes and of hydrogen peroxide. HYDROGEN PEROXIDE Formula, H 2 O 2 . History Hydrogen peroxide was discovered by Thenard (1818), who obtained it by the action of acids on barium peroxide. Occurrence Minute traces of this compound appear to be present under normal conditions in the atmosphere. Preparation (i) Hydrogen peroxide is formed in small amount when hydrogen burns in the air. This is best shown by directing the flame on the surface of cooled water or on to a piece of ice ; the rapid cooling thus secured prevents the splitting up into water and oxygen which takes place in the flame under ordinary conditions. (2) Hydrogen peroxide is obtained on the large scale by the action of dilute sulphuric acid on barium peroxide, represented by the equation BaO 2 + H 2 SO 4 = BaSO 4 + H 2 O 2 . The best results are obtained by slowly adding powdered hydrated barium peroxide, BaO 2 ,8H 2 O, to a mixture of I part of concentrated sulphuric acid and 5 parts of water till the acid is nearly neutralized (see p. 99), the temperature being kept low by immersing the vessel in cold water or in a mixture of water and ice. Barium sulphate is insoluble in water, and can be removed by filtration, a dilute aqueous solution of hydrogen peroxide being thus obtained. A more concentrated solution of the peroxide may be obtained by distilling the dilute solution under reduced pressure (p. 65) ; at first almost pure water passes over, and then a fairly concentrated solution of the peroxide. By repeating this process of fractional distillation several times, hydrogen peroxide can be obtained quite free from water. Instead of sulphuric acid, hydrochloric acid may be used to OZONE AND HYDROGEN PEROXIDE 139 decompose the barium peroxide, but in this case barium chloride, BaCl 2 , is the other product, and as it is soluble in water it could not be separated from the hydrogen peroxide by filtration. In the course of our work we shall meet with many other illustrations of the importance of so choosing the reacting substances that the products can readily be separated. Phosphoric acid acts on barium peroxide forming hydrogen peroxide and barium phosphate, and as the Litter substance is insoluble in water, it can be filtered off. This method is sometimes 'dsed in the commercial preparation of hydrogen peroxide. (3) Hydrogen peroxide is formed in small amount when zinc is shaken with water and air or oxygen, and also when lead, copper, and some other metals are shaken with air and dilute sulphuric acid. These remarkable reactions are by no means well understood. Very delicate methods for detecting the traces of peroxide formed are describe d below. Hydrogen peroxide can be detected in many other cases of oxida- tion by free oxygen. (4) Hydrogen peroxide is formed in fair amount when oxygen is bubbled through the solution surrounding a negative pole at which hydrogen is being liberated by electrolysis. The change can be represented by the equation H 2 + O 2 =H 2 O 2 , but is not fully under- stood. Under ordinary conditions hydrogen and oxygen do not unite to form hydrogen peroxide, and we must assume that the hydrogen in the act of liberation at the anode is endowed with exceptional activity. Physical Properties When quite free from water, hydrogen peroxide is a colourless syrupy liquid of specific gravity 1.458 at o. Under 29 mm. pressure it boils at 65 and under 65 mm. at 85. It can be obtained in colourless crystals by cooling the pure liquid in a mixture of ether and solid carbon dioxide (p. 321); the crystals melt at - 2. The heat of solution in water is 460 cal. per mol. The heat given C'Ut when hydrogen peroxide splits up into water and oxygen (see below) is, according to Berthelot, 21,700 cal. per mol (34 grams). Chemical Properties The most remarkable chemical property of hydrogen peroxide is its tendency to break up into water and oxygen, as represented by the equation and for this reason it is a fairly energetic oxidizing agent. The rate at which the peroxide splits up, however, depends very much on its 140 A TEXT-BOOK OF INORGANIC CHEMISTRY purity and on the substances with which it is brought into contact. As already mentioned, a solution of the pure peroxide in water can be concentrated by fractional distillation without very serious decom- position taking place, but an impure peroxide cannot be concentrated in this way. Many finely-divided substances, such as silver, gold, platinum in the form of platinum black, manganese dioxide, and even powdered glass, bring about a very rapid decomposition of the per- oxide. As the substances remain unchanged at the end of the re- 1 action, we have to do with catalytic phenomena. It must be assumed | that under all circumstances hydrogen peroxide is slowly decomposing 1 and that the rate of this decomposition is greatly increased by the j catalysts mentioned. Even the rough surfaces in ordinary glass bottles facilitate the decomposition of the peroxide, but this effect can be greatly minimized by coating the interior of the bottles with solid paraffin. The German firm Merck supply a perfectly pure 30 per cent, aqueous solution of hydrogen peroxide in paraffin-coated bottles, which keeps its strength remarkably well. The peroxide is very unstable in alkaline solution, but traces of acid retard decom- position very markedly. The rate of decomposition is increased by raising the temperature and also by exposure to light. In the cases just mentioned free oxygen is liberated, but many substances take up an atom of oxygen from the peroxide and become oxidized, whilst water remains. Thus black lead sulphide, PbS, is changed by hydrogen peroxide to lead sulphate, PbSO 4 , which is white. From an aqueous solution of potassium iodide free iodine is liberated and at once turns starch emulsion blue (compare ozone, p. 136), the equation being as follows : When hydrogen peroxide is added to a solution of barium or calcium hydroxide the corresponding peroxide is formed, thus : A comparison of the respective formulae shows that barium peroxide may be regarded as being formed from hydrogen peroxide by putting in an atom of barium for two atoms of hydrogen, just as barium chloride, BaCl 2 , is derived from two molecules of hydrochloric acid. From this point of view barium peroxide may be regarded as being a salt derived from hydrogen peroxide. Other illustrations of the oxidizing action of hydrogen peroxide are mentioned under Tests (below). OZONE AND HYDROGEN PEROXIDE 141 As hydrogen peroxide is an oxidizing agent, it can, of course, be used for bleaching purposes. It can be safely employed for bleaching very delicate materials, such as silk, hair, and ivory, as the only products are water and oxygen. For the same reasons it is a valuable antiseptic. It is a remarkable fact that hydrogen peroxide sometimes acts as a reducing agent, that is, it removes oxygen from other substances. It does not itself become more highly oxidized, however, but gives up an atom of oxygen which with a further atom from the other sub- stance forms a molecule of oxygen, which is liberated. Thus it reduces silver oxide, Ag 2 O, to metallic silver according to the equation Ag 2 + H 2 2 = Ag 2 + H 2 + 2 , and it reduces an acidified solution of potassium permanganate, KMnO 4 . to colourless salts. The latter reaction can readily be performed by adding hydrogen peroxide to a solution of potassium permanganate acidified with sulphuric acid ; a vigorous evolution of oxygen will be noticed and the original deep purple colour of the solution will be discharged. The equation expressing this change is rather complicated for the present stage of our work, but will be better understood if taken in two stages. (1) 2 (2) 5 We mny suppose that potassium permanganate and sulphuric acid tend to react according to the first equation, giving rise to potassium sulphate, K 2 SO 4 , manganous sulphate, MnSO 4 (p. 530), water and oxygen, and that this tendency becomes effective in the presence of hydrogen peroxide, which is able to remove the oxygen, as repre- sented in (2). When ozone and hydrogen peroxide are brought together in not too dilute solution, water and oxygen are formed according to the equation These reactions are usually explained as being due to the tendency of the loosely-held atoms in the peroxide and the other substance to combine with formation of molecular oxygen, as indicated above. It is perhaps equally likely, however, that the oxidizing agent oxidizes 142 A TEXT-BOOK OF INORGANIC CHEMISTRY the hydrogen of the peroxide to water, as represented by the equation so that all the oxygen comes from the peroxide. Estimation of Hydrogen Peroxide Solutions The reaction between potas- sium permanganate and hydrogen peroxide in acid solution, just referred to, may be applied to the estimation of solutions of hydrogen peroxide, the liberated oxygen being collected and measured over water or mercury. The apparatus used for this purpose is the same as that already described in connexion with equivalents (Fig. 40). A measured (or weighed) quantity of the solution con- tained in the small tube is placed as shown in the bottle B, which contains excess of potassium permanganate and dilute sulphuric acid and is connected as shown with the graduated measuring tube, which contains water (or mercury). The position of the liquid in the measuring tube is adjusted to zero on the scale at atmospheric pressure by means of the levelling tube, the pinchcock being momentarily opened for this purpose. The levelling tube is then lowered, and the apparatus tilted so that the contents of the small tube become thoroughly mixed with the permanganate solution. When the reaction is over, the liquid is brought to the same level in both tubes, and the volume of gas read off. As the original air in the apparatus is at atmospheric pressure both before and after the experiment, the extra volume is that of the oxygen. From these data the volume of the oxygen at o, and hence the weight of oxygen (or of hydrogen peroxide) in the quantity of solution taken, may readily be calculated. 1 The concentration of a solution of hydrogen peroxide is often expressed in terms of the volume of oxygen given off under these conditions ; thus a 2O-volume solution is one of which i c.c. yields 20 c.c. of oxygen at N.T. P. Since i mol (34 grams) of hydrogen peroxide gives with permanganate 32 grams = 22, 400 c.c. of oxygen at N.T. P. , a 2o-volume solution contains x 34=0.03 grams in i c.c. of solution, i.e. the solution contains a little over 3 per cent, of hydrogen peroxide. If the same quantity of hydrogen peroxide were decomposed by heat alone, it would only yield half the volume of oxygen obtained by the use of perman- ganate. It follows that a 2O-volume solution expressed on the former basis would be double the strength of a 2o-volume solution measured by means of permanganate. Tests (i) A delicate test for hydrogen peroxide depends upon the fact that when a little of it is added to an acidified dilute solution of potassium dichromate a deep azure-blue solution is obtained (see p. 523). A dilute solution of the peroxide is shaken up with ether, a small quantity of an acidified solution of potassium dichromate added, 1 If considerable accuracy is required, account must be taken of the fact that the oxygen is saturated with aqueous vapour. The actual pressure of the oxygen is that of the atmosphere less the pressure of aqueous vapour at the temperature of the experiment. THERMOCHEMISTRY 143 and after further shaking the upper layer of ether, which separates on standing for a short time, is coloured deep blue. If no ether is used, the blue compound rapidly decomposes in aqueous solution. (2) The most delicate test for hydrogen peroxide depends upon the production of a coloured solution (orange-red in moderately concen- trated solution, yellow in very dilute solution) when the peroxide is added to a colourless solution of titanium dioxide, TiO 2 , in sulphuric acid. The coloured substance is titanium trioxide, TiO 3 . One part of hydrogen peroxide in ten million parts of water can be detected by this test. (3) The tests just described are characteristic for the peroxide. Other reactions, such as the liberation of iodine from an acidified solution of potassium iodide, already referred to, are also brought about by other oxidizing agents, such as ozone. Hydrogen peroxide does no: affect an emulsion of potassium iodide and starch in the absence of acid, but at once turns it blue when a little ferrous sulphate is added. Graphic Formulae of Ozone and of Hydrogen Per- oxide The graphic formula of ozone may be written thus O - O in which all the oxygen atoms are divalent, but the alternative formula O = O = O, in which one of the oxygen atoms is quadri- valent, appears to represent better the readiness with which an atom of oxygen is split off, leaving a molecule of oxygen O = O. That oxygen can act under certain conditions as a tetrad has been estab- lished by Collie and others from the behaviour of certain organic compounds. The formula for hydrogen peroxide may be Represented thus H-O--OH, in which the oxygen atoms are divalent, or thus TT Ti>O==O, one of the oxygens being divalent and the other quadri- valent. The latter formula is perhaps more in harmony with the great tendency of the peroxide to split up into water and oxygen. Briihl, on the basis of optical measurements, has suggested the formula H-O = O-H, both oxygens acting as tetrads. THERMOCHEMISTRY General It has been pointed out at a very early stage of our work that chemical changes are invariably associated with energy 144 A TEXT-BOOK OF INORGANIC CHEMISTRY changes in the system. It very often happens that heat and other forms of energy are given out as the result of chemical changes, and examples have been met with in the combination of hydrogen and oxygen to form water, and in the combination of hydrogen and chlorine to form hydrogen chloride. On the basis of the law of the conservation of energy, we state in such cases that the free elements possess more chemical energy than the resulting compounds, and that the chemical changes are accompanied by a simultaneous trans- formation of chemical energy to an equivalent quantity of other forms of energy, mainly heat. The department of chemistry which is concerned with the heat equivalent of chemical changes is called thermochemistry. In representing the results of thermochemical measurements, it is convenient to deal, not with equal weights of substances, but with quantities which are chemically comparable, that is, with molecular or molar amounts. Further, when heat is given out in chemical change, this is conveniently represented by writing the number of calories, preceded by a + sign, after the equation representing the chemical change. Thus the equation 2H 2 O 2 ->2H 2 O + O 2 + 2 x 23,000 cal. indicates that when 68 grams (twice the molar weight) of hydrogen peroxide split up into water and oxygen gas, 2 x 23,000 calories are given out ; in other words, the chemical energy associated with 68 grams of hydrogen peroxide is greater by 46,000 cal. than that associated with its products of decomposition. When heat is absorbed in a chemical change, the amount so absorbed is preceded by the - sign. Thus the equation 2 H C1->H 2 + C1 2 - 44,000 cal. indicates that in splitting up 73 grains of hydrochloric acid into 2 grams of hydrogen and 71 grams of chlorine, 44,000 calories are taken up ; in other words, the chemical energy associated with 73 grams of hydrochloric acid is 44,000 calories less than that associated with its products of decomposition. When a definite amount of heat is given out in a chemical change, according to the law of the conservation of energy, exactly the same amount must be supplied in order to regain the original substances in the same amounts. Thus the thermochemical equation represent- THERMOCHEMISTRY 145 ing the combination of hydrogen and chlorine to hydrochloric acid is as follows : H 2 +C1 2 ->2HC1 + 44,000 cal. As the heat is given out when the change proceeds in the direction of the arrow, the + sign is used. The heat given out or absorbed when a substance is formed is termed its heat of formation; the heat given out or absorbed when a substance is split up into its components is termed its heat of decom- positio?i. Most substances, like hydrogen chloride and water, are formed with evolution of heat, and their heat of formation is said to be positive, a few, such as ozone (p. 135) and the oxides of chlorine (p. 177), are formed with absorption of heat, so that their heat of for- mation is negative. The thermochemical equation for the formation of ozone is as follows : 3O 2 ->2O 3 - 2 x 34, 100 cal. Chemical compounds formed with evolution of heat are said to be exothermic, whilst substances formed with absorption of heat, like ozone, are endothermic. It is, of course, evident that heat must be supplied to split up exothermic substances, whereas endothermic substances decompose with evolution of heat. A little consideration will show that, when the other conditions are kept constant, rise of temperature favours the decomposition of exothe rmic compounds, but favours the formation of endothermic compounds from their components (Le Chatelier's theorem, p. 171). Two o .her terms employed in thermochemical work may be men- tioned. The heat of combustion of a substance is that amount of heat given out when a mol of it is completely burned, and the term is generally applied to combustion in oxygen. Heat of solution is the heat given out when a mol of a substance is dissolved in a large excess of the solvent. Energy Content of Different Forms of Substances- It is very important in writing thermochemical equations to see that the particular forms of the substances taking part in the reactions are cleaily stated, as these may differ considerably in their energy content. When hydrogen and oxygen unite to form liquid water, 68,400 cal. per mol are given out. Part of this heat, however, is due to the change of gaseous to liquid water, and as the heat of vaporization is 537 cal. per gram at 100 (p. 66), the heat per mol due to the change of state is 537 x 18 = 9666 calories. It follows that 10 146 A TEXT-BOOK OF INORGANIC CHEMISTRY the heat given out when hydrogen and oxygen unite to form a mol of steam at 100 is about 68,400-9670 = 58,730 cal. The difference in the energy content of water and ice at o is, on the same basis, 80 x 18= 1440 cal. per mol. A further important correction has to be taken into consideration when gases are formed or disappear as the result of a chemical change. When a gas is generated under atmospheric pressure, it does work and heat is taken up ; when, on the other hand, a gas disappears heat is given out. The relationship between the volume of gas which is formed or disappears and the heat change cannot be worked out here, but may be stated in this form. When a mol of any gas is formed against external pressure at the absolute temperature T, the heat absorbed is 2T, and when a mol of any gas disappears 2T calories are given out. It must be carefully noted that these heat changes are in addition to those associated with the chemical changes. It follows from the above that in the reaction 2H 2 + O 2 = 2H 2 O (liquid) +2x68,400 cal., the observed heat of combustion includes 3 x 2T calories = 1640 cal. (taking the temperature of the experiment as o = 273 abs.) due to the disappearance of three mols of gas. It is evident from the case of ozone and oxygen that the different allotropic modifications of a substance may differ considerably in energy content, and many illustrations of this will be met with later. Hess's Law A very important law of thermochemistry, first established experimentally by Hess, may be expressed as follows : When a chemical change takes place between definite amounts of different substances, the amount of heat given out is always the same provided the initial and final products are the same in each case. As an illustration of this law we will consider the combination of carbon and oxygen to form carbon dioxide. This reaction may take place slowly or quickly, and may take place in one or more stages, but in all cases, if one starts with 12 grams of carbon and 32 grams of oxygen, and finishes up with carbon dioxide alone, 94,300 cal. are always liberated. It is, of course, true that the temperature attained may be very different according to the speed of combination. When the combustion is slow, the heat evolved is communicated to the surroundings and conducted away, and the temperature attained is much lower than when the combustion is rapid. The total amount of heat given out is, however, the same in both cases. THERMOCHEMISTRY 147 The chief importance of Hess's law is that it enables us to deter- mine the thermal equivalent of many reactions which cannot readily be carried out directly. Suppose, for example, we wish to determine the heat ;^iven out or absorbed when 12 grams of carbon combines with 4 grams of hydrogen to form marsh gas, CH 4 (p. 329). We can assume that the reaction is carried out in one stage CH 4 + 2O 2 = CO 2 + 2H 2 O + 2ii.9oocal. (a) or in the following two stages ;r cal. () = CO 2 + 94,300 cal. As, by Hess's law, the heat given out in the reaction (a) is neces- sarily the same as that given out when the same reaction takes place in two stages (3) and (c), we have 2 1 1, 900 = # + 94,300 +136,800, when x- 19,200 cal. As 19,200 cal. are absorbed when methane is split up into its elements, the same amount is given out when it is formed from its elements ; in other words, methane is an exothermic compound. Relationship between Chemical Reactivity and Heat Of Reaction It will be evident, on consideration of the reactions already discussed, that chemical changes may be divided into two classes: (i) those which under the conditions of experiment are spontaneous or proceed of themselves once they are started, (2) those which only proceed when forced by some external energy, for example, when heat is continuously supplied. As examples of the first class we have the combination of hydrogen and oxygen to form water and of hydrogen and chlorine to form hydrogen chloride ; as examples of the second class the splitting up of mercuric oxide and of potassium chlorate by heat. The former reactions proceed rapidly to completion with explosive violence, the latter stop at once when heating is stopped. The energy relations of the systems throw a great deal of light on these changes. We have already seen that the formation of water and of hydrogen chloride from the respective elements are highly exothermic reactions (p. 145). On the other hand, the thermochemical equation for the decomposition of mercuric oxide is as follows : 2HgO->2Hg + O 2 - 2 x 20,600 cal. 148 A TEXT-BOOK OF INORGANIC CHEMISTRY So that a large amount of heat has to be supplied. From these results the conclusion might be drawn that a chemical reaction pro- ceeds of itself in the direction in which heat is given out, that is, that spontaneous chemical changes are exothermic. Experience has shown that this is actually the case for a very large number of chemical changes, more particularly those with large heats of reaction, and the above statement may be used as an approximate rule by the student. It does not apply in every case, however, as there are many changes, for example, the dissolution of ammonium chloride and certain other salts in water, which proceed of themselves and produce a considerable cooling effect. It would lead too far to discuss these relationships fully, 1 and it is quite sufficient for general purposes to make use of the approximate rule already mentioned. A reaction which is endothermic or only feebly exothermic, and which scarcely proceeds at all under ordinary conditions, may often be brought about by associating it with some other change, so that the total change is now strongly exothermic. The reaction usually adduced in illustration of this statement is that between hydrogen sulphide and iodine (p. 161). When dry hydrogen sulphide is passed over dry iodine at room temperature no appreciable reaction occurs. As a matter of fact, the reaction is endothermic : H 2 S + I 2 ->2HI + S - 7300 cal. When, however, hydrogen sulphide is passed into iodine suspended in excess of water at room temperature the reaction occurs readily, hydriodic acid and sulphur being formed. Taking into account the fact that the hydrogen iodide is finally present in aqueous solution 2 the thermochemical equation now becomes H 2 S,Aq+I 2 ->2HI,Aq + S + 17,000 cal., that is, the total change is now strongly exothermic and the reaction proceeds spontaneously. The alteration in the thermal character of the reaction is due in this case to the high heat of solution of hydriodic acid (p. 163). Other illustrations of this important principle will be met with in the course of our work. 1 Cf. Physical Chemistry, p. 148. 2 When a substance in dilute aqueous solution takes part in a chemical change, this is sometimes indicated by appending Aq to the formula of the substance, as above. CHAPTER XII THE HALOGENS AND HALOGEN ACIDS IN Chapter VIII. chlorine and its compound with hydrogen, hydrogen chloride, have been considered. In the present chapter we shall deal with three elements fluorine, bromine, and iodine which have so many analogies with chlorine, both as regards the elements themselves and their more important compounds, that all four elements are said to belong to the same family. They are called halogens (from a\s salt, and yei/j/do>, I produce) in allusion to common salt, the chief source of the chlorine compounds. Like chlorine, the three other elements each form a compound with hydrogen of the same type as hydrogen chloride ; the formulas and names are as follows : Hydrogen fluoride, HF ; hydrogen bromide, HBr ; hydrogen iodide, HI. All these compounds are gases at the ordinary temperature (hydrogen fluoride boils at 19.5), and dissolve in water to form strong acids. These hydrogen compounds are also described in this chapter. FLUORINE Symbol, F. Atomic weight, 19. Molecular weight, 38. Occurrence The chief source of fluorine compounds is calcium fluoride, CaF 2 , which occurs naturally as fluor-spar or Derbyshire spar. The pure mineral forms colourless cubical crystals, but many samples are brilliantly coloured by traces of impurities. Another naturally occurring compound of fluorine is cryolite, AlF 3 ,3NaF, a double fluoride of sodium and aluminium. Fluorine occurs to a small extent in bones and in the enamel of teeth. Preparation It is an interesting fact that fluorine was not obtained in the free condition till 1886. We have learnt that when hydrogen chloride is electrolyzed, part of the chlorine acts on the water and sets free oxygen. Fluorine has a still greater affinity for hydrogen than chlorine has, and when an aqueous solution of hydrogen fluoride is electrolyzed the fluorine set free at the anode acts 150 A TEXT-BOOK OF INORGANIC CHEMISTRY immediately on the water, with liberation of oxygen and reformation of hydrogen fluoride. It appears at first sight as if this difficulty could be overcome by using anhydrous hydrogen fluoride, but Gore found that the latter does not conduct the electric current, so that no chemical change occurs. Moreover, fluorine acts energetically on the great majority of elements and of chemical compounds, so that it is not FIG. 43. surprising that the attempts made to isolate this element were at first unsuccessful. The problem was solved by Moissan, who subjected to electrolysis anhydrous hydrogen fluoride in which potassium fluoride had been dissolved to make a conducting solution. One form of apparatus used is represented in Fig. 43. The U-tube (made of an alloy of platinum and iridium), which contains the mixture of hydrogen fluoride and potassium fluoride (about 4 parts to i by weight) is provided with two side tubes, A and B, and the open ends are closed with stoppers, C, made of fluor-spar and wrapped in thin sheet THE HALOGENS AND HALOGEN ACIDS 151 platinum. The electrodes, which are also made of a mixture of platinum and iridium (one of the few materials which #re not much affected by fluorine), pass through the stoppers, and are kept in place by screws. During the electrolysis the apparatus is kept at a temperature of - 23 by means of boiling methyl chloride. The fluorine, which is given off as a gas at the positive pole, is passed through a platinum spiral kept at a low temperature in order to remove any hydrogen fluoride, and may then be collected and examined in a platinum tube, the ends of which are closed with plates of fluor-spar. In his earlier experiments Moissan used an apparatus of platinum- iridium, but found subsequently that an apparatus of copper was equally convenient and much less costly. The chemical changes may be assumed to take place mainly in accordance with the equations (i) (2) The potassium liberated at the cathode as the result of the primary change represented by equation (i) immediately reacts with hydrogen fluoride, liberating hydrogen and reforming potassium fluoride. Physical Properties At ordinary temperatures fluorine is a greenish-yellow gas, much lighter in colour than chlorine ; it has a very pungent odour. The density of the gas is approximately 19, corresponding with the molecular formula F 2 . Fluorine has been obtained by Moissan and Dewar as a bright yellow liquid boiling at -187; the liquid is almost without action on glass. On cooling with liquid hydrogen, fluorine forms a pale yellow solid melting at -223, and on further cooling to -252 the solid becomes perfectly white. Chemical Properties Fluorine is the most chemically active element known ; it combines directly with all other elements except oxygen and the elements of the helium family, and in many cases with explosive violence. It combines with hydrogen explosively in the dark at ordinary temperatures, and even at -210 the gases unite immediately, producing a flame. Sulphur, phosphorus and iodine melt and then inflame in fluorine ; arsenic, antimony and boron become incandescent, and crystalline silicon burns in it with great brilliancy. Fluorine does not combine with oxygen under any conditions so far realized. Most of the metals combine readily with fluorine at the ordinary 152 A TEXT-BOOK OF INORGANIC CHEMISTRY temperature, but gold and platinum are only slightly affected, much less than when acted on by chlorine. With copper a thin coating of the fluoride is formed, which protects the metal against further action. Glass, whether moist or dry, is at once attacked by fluorine. Owing to its great affinity for hydrogen, fluorine immediately decomposes water according to the equation 2F + 2HO->4HF + iO. The liberated oxygen contains a large proportion (up to 12 per cent.) of ozone* HYDROFLUORIC ACID (HYDROGEN FLUORIDE), HF Preparation (i) Hydrogen fluoride is most conveniently ob- tained by heating calcium fluoride (fluor-spar), CaF 2 , with sulphuric acid in a vessel of lead or platinum, as glass is readily acted on by the fluoride. The equation is as follows : CaF 2 + H 2 SO 4 ->CaSO 4 + 2HF. The volatile fluoride is absorbed in water, and the aqueous solution, which is termed hydrofluoric acid, is kept in lead, paraffin, or rubber vessels. (2) Anhydrous hydrogen fluoride is best prepared by heating dry hydrogen potassium fluoride, HF,KF, in an apparatus of platinum, the condenser and receiver being surrounded by ice. The double compound is decomposed on heating, according to the equation and the hydrogen fluoride is collected in the cooled receiver as a colourless liquid. Hydrogen fluoride may also be prepared by other methods analogous to those described under hydrogen chloride (g.v.}. Physical Properties Hydrogen fluoride is a colourless, fuming liquid, boiling at 19.5. The vapour is very pungent, and when inhaled has a very injurious effect on the mucous membrane. It is usually kept in platinum vessels, but when perfectly free from water it has no action on glass (Gore). It mixes with water in all proportions, and when the aqueous solution is distilled a mixture of constant boiling-point (120 at 760 mm. pressure) containing 36 per cent, of hydrogen fluoride is obtained. At 88 the vapour density of hydrogen fluoride is approximately 10, THE HALOGENS AND HALOGEN ACIDS 153 corresponding with the formula HF (molecular weight = 20), but as the temperature is lowered the density progressively increases, and at 26.4 is 25.6, corresponding with a molecular weight of 51.2. The simplest explanation of this remarkable fact would appear to be that as the temperature falls the simple HF molecules unite among them- selves to form compounds of the type (HF)^ where x is a whole number. If the whole of the fluoride was present as H 2 F 2 molecules, the molecular weight would be 40, but this number is already exceeded at 26, so that some of the molecules must be still more complex, In such a case the substance is said to associate or poly- merize as the temperature falls. Many other cases of association are known, and they will be fully considered in the next chapter under the heading of Chemical Equilibrium. Chemical Properties As in the case of hydrogen chloride, a clear distinction must be drawn between hydrogen fluoride, which has no acid properties, and a mixture of the fluoride with water, which is a typical acid, hydrofluoric acid. The latter acid attacks many metals, liberating hydrogen and forming fluorides. The for- mulas of the fluorides correspond with those of the chlorides. Most of them are soluble in water, but the fluorides of calcium and barium are insoluble. The most striking property of hydrofluoric acid is that it acts on glass, and it is therefore largely used for etching purposes. The object to be etched is covered with wax, and the figures or other marks cut through the wax with a pointed instrument. The prepared surface is then exposed to the fumes of hydrogen fluoride for a time or dipped into an aqueous solution of the acid. Only the parts of the glass where the coating is removed are affected by the acid, and on removing the rest of the coating the design will be found etched on the glass. The effect just described depends upon the action of the hydro- fluoric acid on silicates (of which glass is composed, p. 438), according to the equation the silicon fluoride, SiF 4 , which is the main product of the action, escaping as a gas (p. 349). Acids such as hydrochloric acid and hydrofluoric acid, which con- tain only one hydrogen atom replaceable by a metal, are termed monobasic; acids containing more than one replaceable hydrogen atom are termed polybasic. 154 A TEXT-BOOK OF INORGANIC CHEMISTRY BROMINE Atomic weight, 79.92, Molecular weight, 159.84. History Bromine was discovered in 1826 by Balard in the mother liquor obtained by removing most of the sodium chloride from sea-water. The name is derived from /Spvpos, a stench, in allusion to its irritating odour. Occurrence On account of its great affinity for other elements bromine is never found free in nature. It occurs in small amount in sea-water in combination chiefly with sodium, potassium, and magnesium, and is also found in certain mineral springs. It also occurs, mainly as magnesium bromide, associated with carnallite, MgCl 2 ,KCl,6H 2 O, a double chloride of magnesium and potassium, in the salt deposits at Stassfurt, and this is now the chief commercial source. Preparation (i) Bromine can readily be obtained by a method analogous to that already mentioned under chlorine (p. 88) by the action of manganese dioxide and sulphuric acid on a bromide. The reaction may be regarded as taking place in two stages : (1) 4KBr + 2H 2 SO 4 -2K 2 SO 4 + 4HBr. (2) 4HBr+MnO 2 ->MnBr 2 + 2H 2 O + Br 2 . If, however, excess of sulphuric acid is used, all the bromine is liber- ated and manganese sulphate remains behind : 2KBr+2H 2 SO 4 +MnO 2 ->MnSO 4 + K 2 S0 4 + 2H 2 O + Br 2 . On gently warming the reaction mixture, the bromine is given off as a vapour, and is condensed in a cooled receiver. (2) Another method of obtaining bromine from bromides depends upon the fact that chlorine is a more active element than bromine and displaces the latter from combination : 2NaBr+Cl 2 ->2NaCl + Br 2 . The bromine may be obtained by warming gently and condensing the vapour as before. In order to remove traces of chlorine, some potas- sium bromide is added to the bromine, and the latter then redistilled. (3) When an electric current is passed through a solution of a bromide, bromine is set free at the anode. This method is now used commercially for obtaining bromine directly from mixtures of bro- THE HALOGENS AND HALOGEN ACIDS 155 mides and chlorides, as under suitable conditions all the bromine can be liberated before the chlorine begins to separate. Commercial Preparation All three methods just referred to are in use for the commercial preparation of bromine, but most of it is prepared from crude carnallite by the second method. The car- nallite is treated with water and the less soluble salts (chiefly potas- sium chloride) separated by crystallization. The mother liquor, which contains a considerable proportion of sodium and magne- sium bromides, is allowed to trickle down a tower filled with round stones, and chlorine is passed up the tower from below. The liberated bromine is carried over in the form of vapour and condensed in the usual way. It can be freed from chlorine by distilling over powderec potassium bromide, as already mentioned. Physical Properties Bromine is a dark-red liquid at ordinary temperatures. Its density at o is 3.188. It boils at 59, and solid bromine melts at -7.3. It has a high vapour pressure at room temperature, giving off brown fumes which have a very irritating effect on the nose and throat and also on the eyes. Bromine dissolves fairly readily in water, forming a red solution known as bromine water; 100 grams of water at o dissolve 3.6 grams, and at 20 3.2 grams of bromine. On cooling the aqueous solution a crystalline hydrate of bromine, Br 2 ,8H 2 O, is obtained (cf. chlorine, p. 91). The vapour density of bromine up to 750 is about 80, hence the molecular weight is about 160 ; and as the atomic weight of bromine, the smallest quantity of it which occurs in a molecule referred to the atom of oxygen as 16 (p. 116) is 80, the molecular formula is Br 2 . It was found by Victor Meyer, however, that at higher temperatures the density of bromine diminishes, and at 1500 is considerably less than 80. The simplest explanation of this observation is that at high temperatures bromine is partly split up or dissociated with single atoms, according to the equation This important point will be further referred to in the present chapter (p. 159). Chemical Properties The chemical properties of bromine are summarized in the statement that it behaves like chlorine, but acts less energetically. It unites directly with hydrogen, but the combination does not proceed explosively in sunlight, as in the case ofchloiMie. It combines directly with non-metals like phosphorus 156 A TEXT-B.OOK OF INORGANIC CHEMISTRY and arsenic, and also with many metals. It has a slight bleaching action, which may be accounted for as described under chlorine (p. 91). The great chemical activity of chlorine as compared with bromine is further shown by the fact already mentioned, that the former displaces the latter from combination. HYDROGEN BROMIDE (HYDROBROMIC ACID), HBr Preparation (i) Hydrogen bromide can be obtained by direct combination of its elements. Hydrogen and bromine do not combine explosively even when a lighted taper is applied to the mixture, but when the mixed gases are passed through a red-hot tube containing finely divided platinum (catalytic agent, p. 22) fumes of hydrogen bromide are produced. Further, hydrogen burns in bromine vapour, giving rise to hydrogen bromide. (2) As hydrogen bromide is readily volatile, 1 it can be obtained by the action of a suitable acid ; for example, phosphoric acid, H 3 PO 4 , on a bromide : (3) It might appear that in the above method of preparation phos- phoric acid could advantageously be replaced by sulphuric acid ; but if the latter acid is used, it will be observed that brown fumes of bromine are also liberated. The reaction in this case takes place in two stages : (1) 2KBr + H 2 SO 4 (2) 2 HBr+H 2 SO 4 ->2H 2 the bromine resulting from a secondary reaction, in which the sul- phuric acid is reduced to sulphur dioxide and water, and the hydrogen bromide oxidized to bromine and water. (4) The most satisfactory method for the preparation of hydrogen bromide is to decompose phosphorus tribromide, PBr 3 , with water: PBr + HOH->HBr + HPO 3 . As the other product of the reaction, phosphorous acid, H 3 PO 3 , is non-volatile, the substances can readily be separated. In practice, the phosphorus bromide is formed and decomposed in 1 The formation of a volatile product is conveniently shown by an arrow directed upwards, the formation of a precipitate by a downwardly directed arrow. THE HALOGENS AND HALOGEN ACIDS 157 one operation. Red phosphorus, mixed with a little water, is placed in the flask A (Fig. 44), the latter being closed by a cork carrying a dropping-funnel containing bromine. The bromine is added drop by drop to the mixture, and the issuing gas passed through a LJ-tube B containing glass beads mixed with red phosphorus, in order to remove any free bromine which may be carried over. The hydrogen bromide may be collected in a gas jar by upward displacement of air (p. 39), or an aqueous solution may be prepared by supporting the end of the delivery tube just over the surface of water in a bottle C. (5) An aqueous solution of hydrogen bromide is readily obtained by passing hydrogen sulphide, H 2 S, through bromine water till the colour of the latter is discharged, and then removing the precipitated sulphur by filtration. The main reaction is represented by the equation Br 2 +H 2 S->2HBr + Sj. Physical Properties Hydrogen bromide, like hydrogen chloride, is a colourless gas with a sharp odour, and fumes in con- tact with moist air. It can be condensed to a liquid, which boils at -68.7, and on further cooling solidifies; the crystals melt at -80. One volume of water dissolves about 600 volumes of the gas at 10. When distilled, hydrobromic acid, like hydrochloric acid, forms a mixture of constant boiling-point (126 at 760 mm. pressure) contain- ing 48 pc r cent, of hydrogen bromide. The thermochemical equation representing the heat of formation of gaseojs hydrogen bromide from its components in the gaseous form is as follows : H 2 + Br 2 =2HBr + 2 x 12,100 cal. Further, the heat of solution (p. 145) of hydrogen bromide is 20,000 calories, so that 12,100 + 20,000 = 32,100 cal. are given out when I gram of hydrogen and 80 grams of bromine combine and the product is dissolved in excess of water. Chemical Properties Liquefied hydrogen bromide has no acidic properties, but the solution of the gas in water, known as hydrobromic acid, is a typical acid. It acts on many metals, forming bromides and liberating hydrogen. Oxidizing agents set free bromine from hydrobromic acid, a change which takes place much more readily than the corresponding one with hydrochloric acid. This is a furthe- illustration of the fact already mentioned, that there is much less affinity between hydrogen and bromine than between hydrogen and chlorine. 158 A TEXT-BOOK OF INORGANIC CHEMISTRY IODINE Atomic weight, 126.7. Molecular weight (at 600), 253.4. History Iodine was discovered by Courtois (1812) in the course of experiments designed to prepare potassium nitrate from solutions obtained by lixiviating the ashes of sea-weeds. The name is derived from t'oet5i72HI + O. As, however, the affinity between hydrogen and iodine is com- paratively small, the reaction does not take place at all unless a reducing agent (e.g. sulphurous acid, p. 284), is present to remove the oxygen. This may be taken as a further illustration of the effect of the heat of reaction on the direction of a chemical change ; in virtue of the additional heat given out by the oxidation of sulphurous acid, the total change becomes strongly exothermic and proceeds spontaneously. THE HALOGENS AND HALOGEN ACIDS 161 The most characteristic test for iodine is the deep blue colour produced when even traces of it are added to starch solution. The blue compound is termed "iodide of starch," but its nature is not well understood and there is no evidence that it is a true chemical compound. On warming to about 80 the colour is discharged, but it returns as the solution cools. It is a remark- able fact that the colour is not given by iodine alone in water, but only when a soluble iodide is also present. HYDROGEN IODIDE (HYDRIODIC ACID), HI Preparation The methods used in preparing hydrogen iodide are almost identical with those used for preparing hydrogen bromide, and may therefore be very briefly dealt with. (1) A certain amount of hydrogen iodide is obtained when a mixture of hydrogen and iodine vapour is passed over heated, finely divided platinum, but under all circumstances the combina- tion is only partial. (2) As in the case of hydrogen bromide, hydrogen iodide can be obtained by heating an alkali iodide with phosphoric acid ; the hydrogen iodide, being volatile, readily passes off. Sulphuric acid cannot be used in place of phosphoric acid because the hydriodic acid first formed immediately reduces the sulphuric acid to sulphur dioxide, thus As hydrogen iodide is a much more powerful reducing agent than hydrogen bromide, very little hydrogen iodide is given off in the above process. (3) An aqueous solution of hydrogen iodide is readily obtained by pass ng hydrogen sulphide through iodine suspended in water. The reaction is a reversible one, as represented by the equation and is incomplete if a fairly high concentration of hydriodic acid is reached. In dilute solution, however, the reaction proceeds practically to completion in the direction of the upper arrow (cf. p. 148). As already mentioned, the total change is exothermic in virtue of the great heat of solution of hydriodic acid. (4) The most satisfactory method of preparing hydrogen iodide ii 1 62 A TEXT-BOOK OF INORGANIC CHEMISTRY is to mix red phosphorus and iodine in a dry flask and allow water to drop slowly on the mixture (Fig. 44). The escaping gas is freed from traces of free iodine by passing it over red phosphorus in a U-tube, B, and may b*e collected by upward dis- placement of air or over mercury (p. 39). If an aque- ous solution is required, the gas is delivered through a tube of the form shown, which dips just below the surface of the water. Should the water be sucked back during the operation, it is prevented from entering the U-tube by the wide bulb on the delivery tube. Physical Properties Hydrogen iodide is a colourless gas with a sharp odour ; it fumes in contact with moist air (p. 163). It can readily be condensed to a colourless liquid which boils at -35. 7 and melts at - 50.8. One volume of water dissolves 425 volumes of the gas at 10, forming a strongly acid solution (hydriodic acid). When distilled, hydriodic acid, like the other two hydrogen acids already referred to, FIG. 44. forms a mixture of constant boiling-point (127 at 760 mm. pressure) containing 57.7 per cent, of hydrogen iodide. Hydrogen iodide is an endothermic compound ; the thermochemical equation for its formation in the gaseous form is H 2 +I 2 ->2HI-2 x6,ioo cal. THE HALOGENS AND HALOGEN ACIDS 163 The heat of solution of hydrogen iodide in water is high, amount- ing to 19.200 cal. per mol. Chemical Properties The aqueous solution of hydrogen iodide is a typical acid, which dissolves many metals to form salts, the iodides. The latter resemble the chlorides and bromides in chemical properties, most of them are soluble in water. The chief differences in behaviour between hydriodic acid and the other halogen acids are connected with the comparatively small affinity between hydrogen and iodine, already referred to. For this reason hydriodic acid is a powerful reducing agent, hydrogen being given up in the process and iodine set free. An illustration of this behaviour has already been met with in the reaction between sulphuric acid and potassium iodide (p. 161). Hydriodic acid is largely used in organic chemistry for reducing purposes. In the same connexion is the fact that solutions of hydriodic acid rapidly turn brown in the air owing to liberation of iodine : Although iodine is practically insoluble in water, it readily dissolves in the excess of hydriodic acid. It has already been mentioned that hydrogen iodide splits up partially on heating according to the reversible reaction The amount of decomposition is the greater the higher the temperalure. It has been mentioned in the course of the chapter that hydrogen chloride, bromide, and iodide give rise to fumes when they escape into moist aij. This is due to the condensation of the moisture to minute drops, forming a kind of fog. CHAPTER XIII CHEMICAL EQUILIBRIUM THERMAL DISSOCIATION IN the present section the general principles underlying certain changes already referred to will be more fully considered. We have seen that when a definite amount of hydrogen iodide is sealed up in a glass tube and heated in the vapour of boiling sulphur it begins to split up into hydrogen and iodine, but the change stops when about 21 per cent, of it is decomposed. Under these circum- stances, therefore, the mixture contains 79 per cent, of hydrogen iodide and 21 per cent, of hydrogen and iodine, and the proportion remains unaltered at 445, no matter how long the mixture is heated. On the other hand, if equivalent amounts of hydrogen and iodine, contained in a sealed tube, are heated at 445 till no further change occurs, it is found that the resulting mixture of gases contains 79 per cent, of hydrogen iodide and 21 per cent, of hydrogen and iodine. It is evident, therefore, that just as water is in equilibrium with a certain concentration of water vapour at a definite temperature, so hydrogen, iodine and hydrogen iodide are in equilibrium at a definite temperature when a certain concentration of each is present. The equilibrium between water and water vapour is often termed a physical equilibrium, whilst that between hydrogen, iodine and hydrogen iodide, being reached as the result of a chemical change, is called a chemical equilibrium. All the facts are conveniently represented by the equation for a reversible reaction which has already been fully explained. The kinetic theory throws a great deal of light on the nature of equilibria of this kind. As the mere presence of hydrogen and iodine does not retard the decomposition of hydrogen iodide, it may appear at first sight as if complete decomposition of the latter ought to occur if sufficient time be allowed. As, however, hydrogen and iodine can combine to form hydrogen iodide under the same conditions, this reaction sets in as soon as any of the elements are present, and the 164 CHEMICAL EQUILIBRIUM 165 apparent static equilibrium is really a kinetic one, being the point at which the amount of hydrogen iodide decomposed in a given time is just balanced by the amount formed by recombination of its com- ponent elements. We are now in a position to consider rather more fully the influence of the conditions upon the equilibrium in a chemical system. The first point to notice is that an equilibrium can only be established when all the reacting substances remain in the system. This follows at once from the kinetic interpretation of a chemical equilibrium if one of the products of the reaction is continually removed from the system the back reaction is no longer possible. When excess of sulphuric acid is added to sodium chloride at the ordinary tempera- ture, the equilibrium in the solution is represented by the equation NaCl + H 2 SO 4 ^NaHSO 4 + HC1 f , the amount of the products formed in a given time, as represented by the upper arrow, being just balanced by the amount of the original substances reformed by interaction of the products. If, however, the mixture is heated, the hydrogen chloride, being readily volatile, escapes from the system as fast as it is formed, the reaction repre- sented by the lower arrow cannot take place, and therefore the chemical change proceeds practically to completion in the direction of the upper arrow. It should be carefully noted that sulphuric acid does not displace hydrochloric acid from combination with sodium because the former is the stronger acid ; in fact, as we shall learn later, hydrochloric acid is stronger than sulphuric acid under equiva- lent conditions. Similar considerations account for the fact that the action of sulphuric acid on barium peroxide, represented by the equation (p. 138) BaO 2 + H 2 SO 4 ^BaSO 4 | + H 2 O 2 , proceeds almost to completion in the direction of the upper arrow. As barium sulphate is practically insoluble in water, it is removed from the system in the solid form as fast as it is produced, so that the back reaction, represented by the lower arrow, does not take place to any appreciable extent. Influence of Concentration of Reacting Substances on Equilibrium. Law of Mass Action The next point to be considered is the effect on the equilibrium of varying the relative amounts of the reacting substances. Suppose, for instance, we repeat the experiment of heating together hydrogen and iodine, with the 1 66 A TEXT-BOOK OF INORGANIC CHEMISTRY single difference that double the quantity of iodine is taken for the same volume. It is then found that, instead of 79 per cent., about 93 per cent, of the total quantity of hydrogen is converted into hydrogen iodide at equilibrium ; in other words, with regard to the reaction an increase in the proportion of iodine displaces the equilibrium in the direction of the upper arrow. If instead of adding more iodine the proportion of hydrogen is increased, the effect on the equilibrium is in the same direction. From the kinetic standpoint, this means that an increase in the proportion of hydrogen or of iodine in the system causes a relative increase in the rate of the reaction represented by the upper arrow. According to the molecular-kinetic theory, the molecules of hydrogen and iodine must come into contact in order that hydrogen iodide may be formed, and it is plausible to suppose that the rate of formation of the iodide is proportional to the number of collisions per unit time between the reacting molecules. If in a definite volume of the gases the amount of iodine is doubled, the number of collisions per unit time will also be approximately doubled, and the rate of the reaction represented by the upper arrow correspondingly increased. A new state of equilibrium will finally be reached when the amounts of hydrogen iodide formed and decomposed in unit time again balance, and this is only possible by a falling off in the amounts of hydrogen and of iodine in the given volume and a corresponding increase in the amount of hydrogen iodide. This means that the equilibrium is displaced in the direction of the upper arrow by the addition of hydrogen or of iodine, in accordance with the experimental result. From the experimental data of this and similar experiments, the exact relationship between the rate of a chemical reaction and the amounts of the reacting substances present can be calculated. It would lead too far to work the matter out in detail, 1 but the result may be briefly stated as follows : The rate of a chemical reaction is Proportional to the molecular concentration of each of the reacting substances. It is important to note that the rate is not proportional to the amount of each substance present, but to its concentration, this is, to the amount per unit volume. This result is also in entire accord with the above molecular-kinetic considerations. If two vessels of equal volume contain the same amount of iodine and quantities of 1 Cf. Physical Chemistry, p. 156. CHEMICAL EQUILIBRIUM 167 hydrogen in the ratio i : 2, the number of collisions and therefore the speed of the reaction in the second vessel will be double that in the first vessel, corresponding with the fact that the concentration of hydrogen in the former is double that in the latter. If, however, the contents of the first tube are mixed with an equivalent amount of hydrogen, so that the resulting volume is greater than at first, the quantity of hydrogen is doubled, but not its concentration, and it is an experimental fact that the speed of the reaction is not doubled. It is evident that owing to the greater free space which the molecules occupy in the latter case the number of collisions is not doubled by doubling; the amount of hydrogen. The convenience of expressing concentrations in mols per litre instead of in parts by weight will be obvious from what has been stated in previous chapters. As chemical reactions between different substances always take place in molar or molecular proportions, there is a great gain in simplicity. The important result just stated, that the rate of a chemical change (in other words, the amount of a chemical change taking place in a given time) is proportional to the molecular concentration of each of the reacting substances, is usually called the Law of Mass Action, Instead of " molecular concentration" the term "active mass" is often used in stating the law ; but the former method is preferable, as the latter statement appears to imply that the rate of a chemical change is doubled by doubling the mass of one of the reacting sub- stances, which is not the case unless the total/volume is kept constant. The substance of the present section may be partially summarized by a further consideration of the reaction represented by the equation NaCl + H 2 SO 4 ^NaHSO 4 +HCl. Assuming that under the conditions of the experiment all the reacting substances remain in the system, the equilibrium is determined by two distinct kinds of factor: (1) The specific chemical affinity between the sulphuric acid and sodium chloride and that between the hydrochloric acid and sodium hydrogen sulphate. These specific affinities are assumed to be inde- pendent of the concentration, but depend on other factors, such as temperature, etc. (2) The relative concentrations of the reacting substances. The relationship between the molar concentration and the amount of chemical action is expressed in the law of mass action, already fully discussed. 1 68 A TEXT-BOOK OF INORGANIC CHEMISTRY This method of representing the facts, although not in all respects satisfactory, will be found suitable for our present purpose. These considerations can be put much more clearly in mathematical form. 1 Assume that two substances A and B react to form two new substances C and D, according to the equation and that the respective molar concentrations at equilibrium are a, b, c, and d. It is clear from the foregoing that the rate of the direct reaction is proportional both to a and to b, and is therefore proportional to their product. We may put this in the following form where k is a constant depending on the affinity between A and B. In the same way Rate re verse =k'cd, where k' is a second constant, which depends on the affinity between C and D. At equilibrium the rates of the direct and reverse actions are equal by definition, that is, kab=k'cd. The above equation may be written ab k' where K, being the ratio of the two constants k' and k, is also constant, and like k and K ', is independent of the concentrations. K is often termed the equilibrium constant. The above result may be put in the following form : At equilibrium the product of the concentrations on one side, divided by the product of the concentrations on the other side, is constant at constant temperature. This is the most general statement of the Law of Mass Action. When a molecule each of A and B react to form a molecule of a third substance C, according to the equation the equilibrium equation ab=Kcd simplifies to ab^Kc. The law of mass action can be experimentally illustrated by means of the reversible reaction between ferric chloride, FeCl 3 , and ammo- nium thiocyanate, NH 4 CNS, which form ferric thiocyanate, Fe(CNS) 3 , and ammonium chloride, NH 4 C1, according to the equation Solutions of the salts are first prepared ; the thiocyanate solution con- tains 3.7 grams of the salt to 100 c.c. of water and the ferric chloride 1 For details see Physical Chemistry, chap. vii. THERMAL DISSOCIATION 169 solution 3 grams of the commercial salt and 12.5 c.c. of concen- trated hydrochloric acid to 100 c.c. of water. Five c.c. of each of the solutions are added to 2 litres of water, and the solution divided between four beakers. Of the four salts present in solution, the ammonium salts are colourless, ferric chloride is very pale red in dilute solution, and ferric thiocyanate deep blood-red. As the con- tents of the four beakers prepared as above are pale-red, it is evident that the equilibrium is considerably displaced in the direction of the lower arrow. To the contents of two of the beakers are added 5 c.c. of the ferric chloride and thiocyanate solutions respectively, and it will be observed that the solutions become blood-red, owing to the displacement of the equilibrium in the direction of the upper arrow, in accordance with the law of mass action. On the other hand, the addition of 50 c.c. of a concentrated solution of ammonium chloride to the contents of the third beaker makes it practically colourless ; the equilibrium is thus displaced in the direction of the lower arrow, again in accordance with the law of mass action. The solution in the fourth beaker is kept for comparison. Thermal Dissociation When a chemical compound splits up into simpler substances on heating, and the products are capable of recombining to form the original compound on cooling, the latter is said to undergo dissociation. It is clear that the term only applies to reversible reactions ; the decomposition of potassium chlorate by heat into potassium chloride and oxygen, an irreversible reaction, is not a dissociation. A number of instances of dissociation by heat have already been met with, such as the splitting up of hydrogen iodide into its elements (p. 164), the splitting up of steam, according to the equation 2H 2 O^2H 2 + O 2 ; of iodine, according to the equation I 2 ^t2l, and so on. Many other illustrations of this process will be met with later. Owing to its importance a further example will be mentioned here, namely, the dissociation of phosphorus pentachloride, PC1 5 , into the trichloride, PC1 3 , and free chlorine ; The methods employed in detecting and measuring thermal dis- sociation depend on the nature of the reaction. The dissociation of hydrogen iodide at high temperatures can be measured by cooling the mixture rapidly to room temperature, which stops the re- action, and determining the relative concentrations of the reacting substances at leisure. A modification of this method, which may be used to show the occurrence of dissociation in water vapour at high 170 A TEXT-BOOK OF INORGANIC CHEMISTRY temperatures, has already been described (p. 32). In the case of phosphorus pentachloride, the latter and the bichloride are nearly colourless in the form of vapour, whilst chlorine is light green, and the progress of dissociation as the temperature is raised is indicated by the gradual increase in the depth of colour. The phenomenon can be illustrated still more satisfactorily by using phosphorus penta- bromide, PBr 6 , which dissociates in an analogous way, owing to the pronounced colour of the bromine. From a quantitative point of view the progress of dissociation can often be followed by measurements of density. When iodine exists as I 2 molecules its molecular weight is 254, and its density, referred to hydrogen at the same temperature, is 127. If, however, it is com- pletely split up into I atoms its molecular weight (which under these circumstances is the same as the atomic weight) is 127, and its density 63.5. Victor Meyer found that its density at 1250 is about 81.5, a number which lies between that for I 2 molecules and I atoms. We therefore assume that the element has undergone partial dissociation according to the equation and it can readily be calculated that the density at 1250 corresponds with a degree of dissociation of about 72 per cent. The above remarks on the experimental investigation of dissociation apply also to chemical equilibria, of which dissociation is simply a special type. An important question in connexion with the present subject is the effect of excess of one of the products of dissociation of a substance on its degree of dissociation. When a definite weight of phosphorus pentabromide is vaporized in a closed vessel at say 200, the vapour is deep red, indicating the presence of a considerable proportion of free bromine : If, however, the vessel previously contains excess of the tribromide, it will be found that the vapour is only very slightly coloured, indicating that under conditions otherwise the same the degree of dissociation is much less in the presence of excess of one of the products of dis- sociation. The same fact can also be shown quantitatively by determining the density of the vapour of the pentabromide alone and in the presence of excess of the tribromide. It is considerably higher in the latter case, indicating less complete dissociation. THERMAL DISSOCIATION 171 The above conclusion also follows at once from the kinetic inter- pretation of the law of mass action. The addition of more tribromide increases the reaction velocity in the direction of the lower arrow, but does not affect the specific rate at which the pentabromide decomposes (represented by the upper arrow), so that the equilibrium is displaced towards the left ; in other words, the degree of dissocia- tion is diminished. This result may be stated as follows: The degree of dissociation of a compound is diminished by addition of one of the Products of dissociation, provided that the 'volume is kept constant. Homogeneous and Heterogeneous Equilibria So far attention has been confined almost entirely to equilibrium in homo- geneous systems systems which consist of a single form of matter and are of the same composition at all points. For the sake of com- pleteness, reference 'must also be made to heterogeneous systems, which consist of two or more distinct portions so-called phases which are themselves of uniform composition but are separated by definite surfaces from other phases (p. 69). Liquid water in equi- librium with its vapour is a heterogeneous system made up of two phases, Mater and vapour, but the equilibrium in this case is of a physical nature. A more complicated heterogeneous system is that in a saturated solution of a salt ; in this case there are three phases, namely solid salt, solution and vapour in equilibrium. A more detailed consideration of heterogeneous systems will be given at a later stage. Effect of Change of Temperature and Pressure on Chemical Equilibrium. Le Chatelier's Theorem When the temperature of a system in equilibrium is altered, two main types of change may occur (i) the equilibrium may be displaced, all the original substances remaining in the system ; (2) one or more of the phases may disappear, an equilibrium of an entirely different type resulting. To take an example of the latter case first, if a mixture of ice and water in equilibrium with water vapour (p. 68) is raised in temperature even a fraction of a degree, and kept at the new tempera- ture, the ice melts and a new equilibrium is established between water and water vapour (cf. curve OA, p. 68). In the same way, when certain systems in chemical equilibrium at a definite temperature are heated or cooled one of the phases may entirely disappear. For our present purpose equilibria of the first type, those in which alteration of temperature produces only a displacement of the equi- librium, are more important. It has been found that the direction in 172 A TEXT-BOOK OF INORGANIC CHEMISTRY which the equilibrium is displaced is closely connected with the heat given out or absorbed in the chemical change. In order to simplify matters as much as possible, we assume that the volume is kept con- stant throughout. The rule connecting heat of reaction and displace- ment of equilibrium may now be stated as follows: At constant volume, increase of temperature favours the reaction in which heat is absorbed, lowering of temperature favours the contrary reaction. This rule may first be illustrated by means of the ozone-oxygen equilibrium (p. 135) As heat is absorbed in the formation of ozone, a rise of temperature ought to displace the equilibrium towards the left ; in other words, the proportion of ozone in equilibrium with oxygen should be the greater the higher the temperature, and the experimental facts are in entire accord with this deduction. The corresponding deduction, that at low temperatures the equilibrium is very near the right-hand side, has also been confirmed experimentally. We shall meet later with many illustrations of this rule. One very important instance has already been mentioned, namely, the effect of change of temperature on the solubility (p. 83). It follows at once that elevation of temperature should increase the solubility of salts such as potassium nitrate, which dissolve with absorption of heat, and diminish the solubility of compounds such as calcium hydroxide, which dissolve with evolution of heat. It will also be evident why the rule only applies to the solubility of substances in solutions already practically saturated with them, since it applies only to systems which are approximately in equilibrium. If no heat is given out on displacement of equilibrium, the latter ought not to be affected by altering the temperature at constant volume. This deduction also is borne out by experiment. The position of equilibrium is also generally altered by change of pressure ; the rule is as follows : At constant temperature increase of pressure displaces the equilibrium in the direction in which the volume diminishes, whilst decrease of pressure has the contrary effect. As an illustration we will again take the ozone-oxygen equilibrium Increase of pressure should displace the equilibrium in the direction of diminution of volume, that is, towards the right, and, on the con- trary, a lowering of pressure should favour the production of ozone. As THERMAL DISSOCIATION 173 it happens, the rule cannot readily be tested directly in this case on account of experimental difficulties, but it has been proved valid for hundreds of other equilibria (some of which will be men- tioned subsequently), and may safely be assumed to hold for this case also. It follov/s at once from this rule that if the volume does not change on displacement of equilibrium at constant temperature, alteration of pressure should have no effect on the position of equilibrium. Such a system is that containing hydrogen, iodine and hydrogen iodide: H 2 + I 2 ^ 2 HI 2 vols. 2 vols. 4 vols. and in accordance with this deduction it has been found by Bodenstein that the equilibrium point is the same, within the limits of experi- mental error, at total pressures varying between wide limits. As the rule implies, the influence of change of pressure on equilibria in liquid and solid systems is very slight, corresponding with the small changes of volume in such systems. The above rules with regard to the influence of temperature and pressure on physical and chemical equilibria are special cases of a very important rule or theorem usually associated with the name of the French chemist Le Chatelier. The theorem may be stated as follows : If one cr more of the factors determining an equilibrium, namely, con- centration, pressure or temperature, is altered, the equilibrium becomes displaced in the direction which tends to neutralize the effect of the alteration. The student should satisfy himself that all the examples already mentioned in this section are in accordance with the rule, and other illustrations of it will readily occur to the mind. For example, the conclusions can at once be drawn that ice must be melted by raising the temperature at constant pressure, as heat is absorbed in the process, and further, increase of pressure at con- stant temperature should also cause ice to melt, as this change is attended with diminution of volume. As has already been pointed out, these deductions are confirmed by experiment. Velocity of Reaction We have already met with a number of illustrations of the fact that the speed of chemical changes depends greatly on the nature of the substances, and is also markedly influenced by the conditions. When hydrochloric acid and sodium hydroxide are mixed, the rate of reaction, as shown by the change of colour of the indicator, is practically instantaneous. On the other hand, a 174 A TEXT-BOOK OF INORGANIC CHEMISTRY mixture of hydrogen and oxygen may be kept for years at the ordinary temperature without any apparent change taking place (p. 37), but if the temperature is sufficiently high combination at once occurs. The remarkable influence of so-called catalysts in accelerating chemical changes has already been repeatedly referred to, and many other illustrations of catalytic action will be met with in the course of the book. Light has a remarkable effect in accelerating certain chemical changes, such as the combination of hydrogen and chlorine, but it appears to have little or no influence on the great majority of chemical reactions. On the other hand, practically all chemical changes are remarkably accelerated by rise of temperature. A careful study of this effect has shown that the speed of chemical changes is generally doubled or trebled for a rise of temperature of 10. Suppose we assume that the speed is doubled for this increment of temperature. It can readily be calculated that the rate of reaction is increased more than 1000 times by raising the temperature by 100, and more than 1,000,000 times by raising the temperature by 200. It is therefore easy to understand that a reaction which proceeds fairly rapidly at a high temperature may be so slow at the ordinary temperature that no change can be detected in weeks or even months. Instances of this which have already been met with are the combination of hydrogen and chlorine to form hydrogen chloride, and of hydrogen and oxygen to form water. We are now in a position to understand fully the phenomena associated with the oxygen-ozone equilibrium, and this example will serve to summarize some of the general principles just discussed. When a silent electric discharge is passed through oxygen a mixture of the latter with ozone is obtained, and the ozone concentration is much higher than corresponds with the thermal equilibrium at the ordinary temperature. We would therefore expect that as soon as the mixture is removed from the influence of the electric discharge the ozone will begin to decompose, and doubtless this is the case. Owing, however, to the very small velocity of this reaction at the ordinary temperature, the mixture appears to be practically stable. When, however, the temperature is raised to 250, the speed with which the system moves towards equilibrium is enormously in- creased and the ozone rapidly decomposes. It has been calculated that the concentration of ozone in true equilibrium with oxygen THERMAL DISSOCIATION 175 at atmospheric pressure and ordinary temperature is only about o.ooi per cent. The above considerations show that it is of the utmost importance to draw a clear distinction between the two effects of temperature (i) in displacing the equilibrium; (2) in accelerating chemical changes. Allotropic Modifications. Polymers. Isomers When the same element exists in different forms, as in the case of oxygen and ozone, the forms are termed allotropic modifications. Phosphorus, sulphur, carbon, and many other elements are met with in different allotropic modifications. The allotropes of the same element differ in energy content (cf. p. 241). When two or more molecules of the same kind combine to form complex molecules of the same empirical composition, the process is termed polymerization or association^ and the product is said to be a polymer or polymeride of the simpler substances. Hydrofluoric acid is a case in point; the simpler substance has the formula HF, the polymer s represented as H 2 F 2 . The term is applied both to elements and chemical compounds. Some allotropic modifications of elements are undoubtedly polymers of the simpler forms. When two substances have the same empirical formula (p. 121) and the same molecular weight, but differ in properties, the phenomenon is termed isomerism and the substances are said to be isomers, Urea and ammonium cyanate are isomers (p. 327), both having the empirical formula CoN 2 H 4 . CHAPTER XIV OXIDES AND OXYGEN ACIDS OF THE HALOGENS eneral It has already been mentioned (p. 100) that although chlorine does not combine directly with free oxygen, three oxides of chlorine, of the respective formulas C1 2 O, C1O 2 and C1 2 O 7 , can be prepared by indirect methods. The first and last of these oxides are acidic. When passed into water, chlorine monoxide forms an acid, hypochlorous acid, HC1O, according to the equation C1 2 O + H 2 O->2HC1O. Acids of this type are called oxygen acids or oxyacids, as in addition to hydrogen and another element, in this case a halogen, they contain oxygen. A number of oxygen acids of chlorine besides hypochlorous acid are known ; they are described below. From their property of forming acids with water, acidic oxides are sometimes called acid anhydrides. No oxides of fluorine have been obtained up to the present, either by direct or indirect methods, and oxygen acids of fluorine are also unknown. No oxides of bromine are known, but two oxyacids of this element, hypobromous acid, HBrO, and bromic acid, HBrO 3 , have been prepared. One stable oxide of iodine, iodine pentoxide. I 2 O 5 , and two oxygen acids are known. These compounds will now be described in detail. OXIDES AND OXYACIDS OF CHLORINE Three oxides of chlorine are known, viz. : Chlorine monoxide (hypochlorous anhydride) . CLO Chlorine peroxide ....... C1O 2 Chlorine heptoxide (perchloric anhydride) . . Cl 2 Oj. Four oxyacids have been obtained ; the names and formulae are as follows : Hypochlorous acid ....... HC1O Chlorous acid ....... HC1O 2 Chloric acid ........ HC1O 3 Perchloric acid ....... HC1O 4 176 OXIDES AND OXYGEN ACIDS OF HALOGENS 177 As it will be necessary, in describing the preparation and properties of the oxides, to mention certain salts derived from the oxyacids before their systematic description is reached, it will be well to note here that the salts corresponding with hypochlorous acid are known as HypochloriteS) those corresponding with chlorous acid as chlorites, those corresponding with chloric acid as chlorates, whilst the salts of perchloric acid are termed perchlorates. The principles on which this nomenclature is based are discussed later. CHLORINE MONOXIDE (HYPOCHLOROUS ANHYDRIDE), C1 2 O Preparation This compound is obtained by passing chlorine over dry precipitated mercuric oxide at a low temperature : 2HgO + 2C1 2 = C1 2 O + HgO-HgCl 2 . The oxide passes off as a gas, a brownish crystalline compound, HgO'HgCl 2 , remaining behind. Properties Chlorine monoxide is a brownish-yellow gas at ordinary temperatures. It can readily be condensed to a liquid, which boils at 5. The oxide is very unstable ; on gentle warming, or in contact with organic matter, it breaks up explosively into its elements. This behaviour accords with the fact that the monoxide is a highly endothermic compound (p. 145). On breaking down into its elements, according to the equation 2C1 2 O->2C1 2 +O 2 2 x 17,800 calories are given out. The monoxide is fairly soluble in water, with which it combines to form hypochlorous acid : Composition The method by which the composition of the monoxide was first established is instructive. The gas was collected in a tube over mercury and split up into its elements by gentle heating, when it was found that two volumes of the gas gave three volumes of the mixture. The latter was then treated with potassium hydroxide, which absorbed the chlorine, leaving one volume of oxygen ; the mixture therefore contained two volumes of chlorine. These results can be satisfied only by the equation 2 C1 2 O -> 2C1 2 + 2 4 unit vols. 4 unit vols. 2 unit vols. 12 178 A TEXT-BOOK OF INORGANIC CHEMISTRY which indicates that the gas contains its own volume of chlorine and half its volume of oxygen. The formula C1. 2 O is confirmed by the vapour-density, which is about 43.5, corresponding with the molecular weight 87 (2 x 35.5 + 16). CHLORINE PEROXIDE, C1O 2 Preparation (i) On gently heating a mixture of concentrated sulphuric acid and potassium chlorate, KC1O 3 , chlorine peroxide is given off as a gas Potassium perchlorate, KC1O 4 , and potassium acid sulphate, KHSO 4 , remain in the retort. If the temperature is allowedlo rise above a certain point, the peroxide splits up into its elements with a violent explosion. (2) Chlorine peroxide, mixed with chlorine, is obtained when potassium chlorate, KC1O 3 , is heated with hydrochloric acid: The mixture of chlorine dioxide and chlorine obtained in this way was regarded by Davy as a definite compound of chlorine and oxygen, and was called euchlorine. (3) A steady stream of chlorine dioxide, mixed with an equal volume of carbon dioxide, is obtained by heating a mixture of potassium chlorate (40 grams) and oxalic acid (150 grams) with a little water (20 c.c.) at 60, direct sunlight being excluded : 2KC1O 3 + 2H 2 C 2 O 4 ->K 2 C 2 O 4 + 2H 2 O + 2CO 2 + 2C1O 2 . Properties Chlorine dioxide is a deep yellow gas, with a peculiar odour, somewhat resembling chlorine. It can be condensed to a deep reddish-brown liquid, which boils at 10, and on further cooling forms yellow crystals, which melt at -79. The dioxide is extremely unstable, exploding violently on warming or in contact with organic matter ; it also decomposes into its elements on exposure to sunlight at the ordinary temperature. Owing to the readiness with which it gives up oxygen, it is a powerful oxidizing agent. This may be shown very instructively by adding to a mixture of sugar and potassium chlorate a drop of concentrated sulphuric acid. The mixture at once bursts into flame, the liberated chlorine dioxide OXIDES AND OXYGEN ACIDS OF HALOGENS 179 igniting the sugar, which then burns at the expense of the oxygen in the chlorate. It is also a powerful bleaching agent. Chlorine peroxide is readily soluble in water. When exposed to sunlight the dissolved peroxide decomposes into its elements, but in the dark more complicated changes take place. When passed into an aqueous solution of potassium hydroxide, a mixture of potassium chlorate and chlorite in equivalent proportions is formed : C1O 2 + 2KOH->KC1O 2 + KC1O 3 . The formula C1O 2 , ascribed to this compound, is confirmed by vapour density determinations. CHLORINE HEPTOXIDE, C1 2 O 7 This compound is obtained by slowly adding phosphoric pentoxide to per- chloric acid cooled below - 10, and then distilling the mixture on a water-bath. The heptoxide is a colourless liquid, which boils at 82, and is more stable, and therefore a less powerful oxidizing agent, than the other oxides of chlorine. HYPOCHLOROUS ACID, HC1O Preparation (i) As already mentioned, hypochlorous acid is obtained by passing chlorine monoxide into water. (2) It is also obtained by passing chlorine in excess through mercuric oxide suspended in water : 2Cl 2 +H 2 0-HgCl 2 As mercuric chloride, HgCl 2 , is soluble in water, the products cannot be separated by filtration, but on distilling the solution a mixture of hypochlorous acid and water passes over. (3) When chlorine is passed into a cold solution of a base, such as potassium hydroxide, KOH, or calcium hydroxide, Ca(OH) 2 , a mixture of the chloride and hypochlorite of the base is obtained 2Cl 2 ->CaCl 2 -fCa(OCl) 2 + 2H 2 0. If an amount of dilute nitric acid equivalent to the hypochlorite is added to a mixture of chloride and hypochlorite, the former is practical !y unaffected, whilst hypochlorous acid is set free from the hypochlc rite according to the equation Ca(OCl) 2 + 2HNO 3 ->Ca(NO 3 ) 2 i8o A TEXT-BOOK OF INORGANIC CHEMISTRY On distilling the mixture a dilute solution of hypochlorous acid is obtained in the receiver. Instead of a mixture of chloride and hypochlorite, a pure hypo- chlorite could of course be used for preparing the acid, but they cannot be obtained so readily. (4) When chlorine is dissolved in water an equilibrium is estab- lished, represented by the equation but this reaction cannot conveniently be used for preparing hypo- chlorous acid. Properties Hypochlorous acid has never been obtained free from water. The dilute aqueous solution is light yellow in colour, has a chlorous smell, and is moderately stable ; the concentrated solution is golden-yellow and is very unstable, splitting up into hydrochloric acid and oxygen : 2HC1O->2HC1 + O 2 . Owing to the readiness with which it yields oxygen, hypochlorous acid is a powerful oxidizing and bleaching agent. In this respect it is twice as energetic as an equivalent quantity of free chlorine, as is evident from a comparison of the two equations In moderate concentration hypochlorous and hydrochloric acids react to form free chlorine and water, represented by the upper arrow in the equation but, as already explained, this reaction is reversible, the position of equilibrium depending on the relative concentrations of the reacting substances and other factors. The removal of chlorine from the system (owing to its volatility and relatively slight solubility) naturally favours the reaction represented by the upper arrow. The hypochlorites are derived from hypochlorous acid by displace- ment of the hydrogen by metals, and may be obtained pure by neutralizing the free acid with the appropriate bases. They are largely employed for bleaching and disinfecting purposes, and in this respect a solution of sodium hypochlorite, NaCIO, known as eau de Javelle, and a mixture of calcium chloride and hypochlorite, known as bleaching powder, deserve special mention. The latter will be further referred to under calcium (p. 436). OXIDES AND OXYGEN ACIDS OF HALOGENS 181 CHLOROUS ACID, HC1O 2 Chlorous acid has up to the present been obtained only in dilute aqueous solution by adding dilute sulphuric acid to a mixture of potassium chlorite and chlorate (obtained by passing chlorine peroxide into a solution of potassium hydroxide, p. 179) and removing chlorine dioxide and free chlorine by means of arsenious acid. A solution containing free chlorous acid, mixed with chloric acid, is thus obtained ; it decomposes fairly rapidly with formation of chlorine dioxide. Chlorous acid is a powerful oxidizing agent. The chlorites, on the other hand, have only a weak oxidizing action. CHLORIC ACID, HC1O 3 Preparation Chloric acid is most readily obtained by the action of dilute sulphuric acid on barium chlorate in equivalent proportions : Ba(ClO 3 ) 2 +H 2 SO 4 ->BaSO 4 | + 2HC1O 3 . The acid is separated from the insoluble barium sulphate by filtra- tion and concentrated in a vacuum over sulphuric acid. In this way a solution containing about 40 per cent, of the acid is obtained. When the attempt is made to remove more of the water the acid decomposes into chlorine, oxygen, and perchloric acid, HC1O 4 . Properties The concentrated acid is an extremely powerful oxidizing and bleaching agent ; wood and paper at once catch fire when plunged into it. The salts of chloric acid are called chlorates. The most important salt is potassium chlorate, KC1O 3 , which is obtained, mixed with potassium chloride, by passing chlorine into a hot concentrated solution of potassium hydroxide : 6KOH + 3C1 2 ->KC1O 3 + 5KC1 + 3H 2 O. The two salts may be separated by taking advantage of the fact that the chlorate is much less soluble in water than the chloride. All chlorates are soluble in water, and have only weak oxidizing properties. On prolonged heating at a high temperature potassium chlorate splits up completely into potassium chloride and oxygen : If, however, the heating is stopped at an intermediate stage, the mixture is found to contain a salt with more oxygen than the chlorate. This sait is called potassium perchlorate, KC1O 4 . The progress of 182 A TEXT-BOOK OF INORGANIC CHEMISTRY the decomposition of the chlorate depends upon the temperature and other factors, but it is probable that at 400 the first stage is as follows : At higher temperatures, the perchlorate is completely decomposed into chloride and oxygen. We shall meet later with many instances in which, as in the present case, a compound of intermediate type splits up into one of a lower and another of a higher type. The effect of heating a mixture of potassium chloride and hypochlorite which leads to the same result as when chlorine is passed into hot potassium hydroxide, is clearly of the same nature. 1 PERCHLORIC ACID Preparation (i) A certain quantity of perchloric acid is formed when a concentrated aqueous solution of chloric acid is heated (see chloric acid). (2) The acid is usually prepared by distilling potassium perchlorate with excess of concentrated sulphuric acid in a vacuum : KC1O 4 + H 2 S0 4 ->KHSO 4 + HClO 4 f. Properties The pure acid is a colourless liquid at ordinary temperatures. It boils at 16 under 18 mm. pressure with slight decomposition; its density at 20 is 1.7676. With water it forms a constant boiling mixture (p. 95) which contains 72.4 per cent, of acid and' boils at 203 under atmospheric pressure. The acid forms a number of stable compounds with water, including a monohydrate, HC1O 4 ,H 2 O, and dihydrate, HC1O 4 ,2H 2 O. The concentrated acid is a powerful oxidizing agent, wood and paper immediately catching fire in it. The dilute acid is a much less powerful oxidizing agent than a solution of chloric acid of the same concentration. This is shown, for instance, by the fact that no chlorine is given off when a dilute solution is gently warmed with hydrochloric acid. Further, solutions of perchloric acid have no bleaching properties. 1 It is, of course, evident that only the hypochlorite undergoes change in this case: 3KC1O->KC1O 3 + 2KC1. OXIDES AND OXYGEN ACIDS OF HALOGENS 183 The salts of this acid, the perchlorates, are all soluble in water, but the solubility of potassium perchlorate is small, and advantage has been taken of this to effect a partial separation of potassium from the other alkali metals. OXYACIDS OF BROMINE As already mentioned, no compounds of bromine and oxygen are known, but two oxyacids, hypobromous acid, HBrO, and bromic acid, HIirO 3 , have been obtained. As their methods of preparation and properties are closely analogous to those of the corresponding chlorine compounds, they can be dealt with very briefly. HYPOBROMOUS ACID, HBrO Preparation The acid can readily be obtained by shaking up precipitated mercuric oxide with bromine water : 2HgO + 2Br 2 + H 2 O-> + HgOHgBr 2 + 2HBrO. On distilling under reduced pressure, a dilute solution of the acid is obtaii ed. Properties The concentrated aqueous solution of hypobromous acid is straw-yellow in colour, and has a powerful bleaching and oxidizing action. On heating, it readily splits up into hydrobromic acid and oxygen. The hypobromites are also oxidizing agents. Sodium hypobromite, which is the best-known of these salts, forms a yellow solution with water, and is employed as a mild oxidizing agent in organic chemistry. BROMIC ACID, HBrO 3 Preparation (i) By treating barium bromate, Ba(BrO 3 ) 2 , with the calculated quantity of sulphuric acid : Ba(Br0 3 ) 2 + H 2 SO 4 ->BaSO^ + 2H BrO 3 . (2) By passing excess of chlorine through bromine water : Br 2 + 5Cl 2 + 6H 2 O-32HBrO 3 +ioHCl. In this interesting reaction, bromine, though itself an oxidizing agent under certain conditions, is oxidized by chlorine, a more powerful oxidizing agent. 1 84 A TEXT-BOOK OF INORGANIC CHEMISTRY (3) By treating the sparingly soluble silver bromate, AgBrO 3 , with bromine : Properties Bromic acid is known only in aqueous solution. On heating it splits up into bromine, water and oxygen. It has oxidizing and bleaching properties. The corresponding salts, the bromates, are mostly soluble with difficulty in water. Those of the alkalis decompose into bromide and oxygen on heating, without the intermediate formation of per- bromate. OXIDE AND OXYACIDS OF IODINE Hypoiodous Acid When iodine is added to a cold solution of potassium hydroxide a colourless solution is obtained. When freshly prepared the solu- tion bleaches indigo, and on addition of weak acids, such as acetic acid, iodine is set free. If, however, the solution is kept for a considerable time, or is heated and allowed to cool, it no longer shows the above properties. The simplest explanation of these observations is that the reaction between iodine and cold potassium hydroxide is similar to that with chlorine : I 2 + 2KOH->KI + K.IO + H 2 O, and that the hypoiodite, KIO, has a bleaching action. Weak acids set free hydriodic and hypoiodous acid, which react immediately with liberation of iodine : HI + HIO->I 2 +H 2 O. The mixture of alkali iodide and hypoiodite is, however, unstable, and a mixture of iodide and iodate results on keeping or on heating the solution : The iodate, KIO 3 , yields oxygen much less readily than the hypoiodite, and the solution containing it has therefore no bleaching properties. IODINE DIOXIDE, IO 2 * The existence of a compound of this formula was first mentioned by Millon, and quite recently Pattison Muir has described its preparation by heating iodic acid with sulphuric acid. It occurs in small lemon-yellow crystals, which begin to decompose at 130, and on further heating decompose completely into iodine and oxygen. IODINE PENTOXIDE, I 2 O 5 Preparation Iodine pentoxide is the anhydride of iodic acid, and is obtained by heating this acid at 199 to 200 : 2HIO 3 ->I 2 6 +H 2 O. OXIDES AND OXYGEN ACIDS OF HALOGENS 185 Properties It is a white crystalline substance, which splits up completely into iodine and oxygen on heating at 300, and dissolves in water to form iodic acid. IODIC ACID, HIO 3 Preparation (i) By treating barium iodate with the calculated amount ol" dilute sulphuric acid, filtering, and evaporating in vacuo : Ba(IO 3 ) 2 +H 2 SO 4 ->BaSO 4 (2) Iodic acid is readily obtained by the oxidation of iodine in presence of water, for example, by boiling iodine with nitric acid till the red colour disappears : 3l 2 + ioHNO 3 ->6HIO 3 + ioNOf + 2H 2 O. Instead of nitric acid, chlorine may be used as oxidizing agent (cf. p. 90). Properties Iodic acid occurs in colourless crystals, which dissolve in water to form a strongly acid solution. Iodic acid is an oxidizing agent, but not so energetic as chloric or bromic acids ; when brought in contact with hydriodic acid, HI, the hydrogen of the latter is oxidized to water, and iodine and water are the sole products : When iodic acid is heated to 1 10, it loses part of its water, and yields the compound HI 3 O 8 : 3 HIO 3 -H 2 O->HI 3 O 8 . At 190 to 200 all the water is driven off and iodine pentoxide remains. The iodates, like the chlorates and bromates, are obtained by the action of the free halogen on a hot solution of the base, and can be separated from the iodides, which are simultaneously formed, by taking advantage of the smaller solubility of the iodates. PERIODIC ACID, H1O 4 ,2H 2 O or H 5 IO 6 Preparation (i) This compound is obtained by the action of sulphuric acid on barium periodate : Ba(I0 4 ) 2 +H 2 SO 4 + 2 H 2 0-BaSO 4 + 2 HIO 4 (H 2 O) 2 . (2) Alkali periodates are obtained by the electrolytic oxidation of 1 86 A TEXT-BOOK OF INORGANIC CHEMISTRY alkali iodates. By the same method periodic acid can be obtained from iodic acid. (3) A periodate of sodium is obtained by the action of chlorine on sodium iodate in hot alkaline solution. Properties The acid is obtained, on evaporation of its aqueous solution, as colourless crystals of the composition HIO 4 ,2H 2 O, which melt at 130, and on heating at a slightly higher temperature split up into iodine pentoxide, water, and oxygen. The anhydrous acid, HIO 4 , has not been obtained, but salts, such as AgIO 4 and NaIO 4 , derived from it, are known (see next section). Anhydrides, Acids, and Hydrates of Acids An acid has been provisionally defined as a substance containing hydrogen, the whole or part of which can be replaced by metals, with formation of salts. An anhydride or acidic oxide is a compound derivable from an oxyacid by abstraction of one or more molecules of water, and, conversely, the formula of an oxyacid is derived from that of the anhydride by addition of one or more molecules of water. When the anhydride corresponding with an acid is known, the latter cannot always be obtained directly from the former, or the former from the latter. In the halogen group, as we have seen, these changes can generally be accomplished, but we shall see later that the acid corresponding with nitrous oxide, N 2 O, cannot be obtained directly from the latter, nor can phosphorus pentoxide, P 2 O 5 , be obtained directly from the corresponding acid. In many cases the anhydrides corresponding with oxyacids have not so far been obtained. Thus, from analogy, the existence of an anhydride, I 2 O 7 , corre- sponding with periodic acid, HIO 4 , would be anticipated, but so far it is unknown. Many other examples of this have been met with in the present chapter. We shall, however, sometimes find it convenient, for purposes of formulation, to assume the existence of these hypothetical oxides. As acids are derived from acidic oxides by addition of water, it is conceivable that one oxide; by association with different amounts of water, may give rise to more than one oxyacid. That this is the case we have already learnt in con- nexion with iodic acid (p. 185). For purposes of comparison, the acid itself may be written, I 2 O 5 ,H 2 O ( = 2HIO 3 ), and the other acid, obtained by heating the ordinary acid to 110, as 3l 2 O 5 ,H 2 O ( = 2 HI 3 O 8 ). It has been found that salts of different types are obtained from solutions con- taining periodates, for example, AgIO 4 and Ag 4 I 2 O 9 , and these may be regarded as being derived from acids of the types HIO 4 , H 5 IO 6 , H 4 I 2 O 9 , etc. It is not at first sight evident that these acids are all periodic acids, but further consideration shows that they may be regarded as being formed by the association of different quantities of water with the hypothetical periodic anhydride, I 2 O 7 , thus, I 2 O 7 ,H 2 O =2HIO 4 ; I 2 O 7 ,sH 2 O = 2H 5 IO 6 , and I 2 O 7 ,2H 2 O=H 4 I 2 O 9 . It is interesting to note that, as already pointed out, the stable acid is not HIO 4 , but H 5 IO 6 (which may be written HIO 4 ,2H 2 O), and the majority of the salts of periodic acid are derived from the latter acid. On the other hand, the salts derived from perchloric acid, HC1O 4 , are formed by the displacement of one hydrogen only by metals, although a number of compounds of the acid and water are known (p. 182). We OXIDES AND OXYGEN ACIDS OF HALOGENS 187 must therefore suppose that the water in the dihydrate of perchloric acid, HClO^HoO, is less intimately associated with the rest of the molecule than in the case of the dihydrate of periodic acid, as in the former case none of the hydro- gens is disp aceable by metals. For this reason the water associated with per- chloric acid may appropriately be regarded as water of crystallization ; in the case of periodic acid as water of constitution. As already mentioned, acids of which only one hydrogen is displaceable by metals are termed monobasic, and nearly all the acids hitherto met with are of this type. Acids which contain more than one displaceable hydrogen are said to be polybasii , and some of the complex periodic acids just considered are of this type. This question will be fully considered later in connexion with the oxyacids of phosphorus. Strong and Weak Acids We have seen that hypochlorous acid can be obtained by adding nitric acid to a solution of a hypo- chlorite, for example, potassium hypochlorite, and distilling : One reason why the reaction proceeds practically to completion in the direction of the arrow is that hypochlorous acid is volatile and escapes from the system, but it is of interest to inquire what occurs when equivalent proportions of potassium hypochlorite and nitric acid are mixed in the cold, under which condition practically all the hypochlorous acid remains in the system. As would be anticipated, an equilibrium is then established, represented by the equation KC1O + HNO 3 ^tKN0 3 + HC1O, and it can be shown, by methods which it would lead too far to describe, that it lies very near the right-hand side ; in other words, nitric acid displaces hypochlorous acid almost completely from com- bination with a base. In such a case we are accustomed to state that nitric acid is a much stronger acid than hypochlorous acid. The above is one of die standard methods by which the relative strengths of two acids is compared. They are allowed to compete, under such condi- tions tha: all the reacting substances remain in the system, for an amount of base insufficient to saturate both of them, and the ratio in which the; distribution takes place is measured. The stronger acid is the one which takes possession of most of the base. Of the acids so far considered, hydrochloric, hydrobromic and hydriodic acids are among the strongest, whilst hydrofluoric acid is considerably weaker. Most of the oxyacids of the halogens are strong acids, but, as already mentioned, hypochlorous acid is an exception. 1 88 A TEXT-BOOK OF INORGANIC CHEMISTRY In connexion with the relative strengths of acids the effect of differences of volatility and of solubility in disturbing the equilibrium must be borne in mind. Thus sulphuric acid is actually a weaker acid than hydrochloric acid, but on account of its volatility the latter is displaced almost completely from combination by the former on heating. Nomenclature The system adopted in naming the halogen oxyacids is in general use for all oxyacids, and will now be briefly described. It has already been explained that the names of binary compounds those containing two elements only end in ide. The name of the most important oxyacid usually ends in ic example, chloric acid, phosphoric acid. Acids containing more oxygen have the prefix per, e.g. perchloric acid, whilst the names of acids containing less oxygen end in ous, e.g. chlorous acid, sul- phurous acid. Those containing still less oxygen than the ous acid have, in addition, the prefix hypo, e.g. hypochlorous acid. The names of the salts derived from the ic acids ends in ate, e.g. potassium chlorate, potassium sulphate, those derived from the ous acids end in tie, e.g. potassium chlorite, potassium sulphite. The salts of the per . . . ic acids are, of course, called per . . . ates, e.g. potassium perchlorate, and those of the hypo ... ous acid are called hypo . . . ites, e.g. potassium hypochlorite. The< same principles are used in naming the oxides. If there is only one, the name of the element followed by oxide is generally used, e.g. magnesium oxide. If there are two, the name of that con- taining less oxygen ends in ous, that containing the higher proportion of oxygen in ic, e.g. mercurous oxide, mercuric oxide. An oxide with more oxygen than the ic oxide has the prefix per, e.g. lead peroxide, one with less oxygen than the ous oxide has the prefix hypo, e.g. hypochlorous oxide. Some exceptions to these rules will be met with in the course of our work. Another, and in some respects preferable, system is to indicate the number of atoms of oxygen by prefixing the corresponding Latin or Greek number, for example, chlorine monoxide, chlorine dioxide, iodine pentoxide. Valency in the Halogen Group The conception of valency has already been explained, and we have seen that it is not usually constant for any one element, but depends on the nature of the other elements with -which it is combined. The variable nature of valency is well illustrated by the halogen compounds discussed in this chapter. OXIDES AND OXYGEN ACIDS OF HALOGENS 189 With regard to hydrogen, chlorine, bromine and iodine are invari- ably univalent, the only known compounds having the graphic for- mulas H-C1, H-Br, H-I. Hydrogen fluoride, however, appears to have the formula H 2 F 2 under certain conditions. In this case fluorine may perhaps be trivalent, the graphic formula being written H-F = FH. When the attempt is made to measure valencies by means of oxygen compounds, uncertainty arises from the possibility of direct linkings between two atoms of oxygen. The graphic formula of chlorine CK monoxide may be written thus pO and that of the corresponding CV H\ acid as yO, the chlorine being univalent and the oxygen CY divalent. When, however, we come to chlorine dioxide, C1O 2 , diffi- culties arise. The chlorine could be represented as divalent, /o ^o CIC I , or as quadrivalent, Cl^. , but as we have no means of >0 ^O deciding the question, there is nothing gained by discussing the matter further. The oxyacids of chlorine have the formulae HC1O, HC1O 2 , HC1O 3 , and HC1O 4 respectively. These can be represented graphically as follows : Cl OH; O = C1 OH; ^Cl OH ; )H the valency of the chlorine being one, three, five, and seven respec- tively. It is an interesting fact that the valencies of polyvalent elements, as in the above instance, often increase by pairs of units. There are, however, many exceptions to this rule (chlorine, for instance, is apparently divalent in certain compounds), and the formulation is often somewhat arbitrary. We have here a striking illustration of the statement already made that the /alency of an element is not constant, but depends on the elements with which it is associated. Compounds of the Halogens with each other Iodine forms compounds with each of the other halogens, and a compound of chlorine and bromine is also known. Iodine itionochloride, IC1, is obtained in the form of a dark reddish-brown liquid i go A TEXT-BOOK OF INORGANIC CHEMISTRY by passing dry chlorine over iodine. It has been obtained in two crystalline forms, which melt at 27.2 and 13.9 respectively. It is decomposed by water, with formation of hydrochloric and iodic acids and liberation of iodine : Iodine trichloride, IC1 3 , is obtained by the further action of chlorine on iodine monochloride or by the action of chlorine in excess on iodine. It occurs in orange needles, and dissociates on heating: Iodine trichloride is soluble in water, and the solution is used as an antiseptic. Iodine monobromide , IBr, is a crystalline compound which melts at 36 and is decomposed by water with liberation of iodine. Iodine pentafluoride , IF 5 , is obtained by direct combination of the elements (Moissan, 1902), and is a colourless liquid which boils at 97 and decomposes on heating to 400. Bromine triftuoride, BrF 3 , obtained by direct combination of its elements, (Prideaux, Lebeau, 1905), is a colourless liquid which reacts violently with water, liberating oxygen and forming chiefly hypobromous and hydrofluoric acids. Comparison of the Halogens and Summary As the halogens are the first family of elements we have considered in detail, a comparison of their behaviour is of great interest. The most important fact brought out by this comparison is that there is considerable resemblance between the properties of the elements themselves, and between those of their corresponding compounds, and, further, there is a gradual variation both of physical and chemical properties corresponding with increasing atomic weight. Thus, just as bromine is intermediate to chlorine and iodine as regards atomic weight, it is also intermediate as regards physical and chemical properties. This generalization is illustrated, as regards physical properties, in the accompanying table. Property. Fluorine. Chlorine. Bromine. Iodine. Atomic weight . 19 35.46 79.96 126.8=; Melting-point -223 IO2 -7 + 114 Boiling-point . . -187 -34 ' 59 184 Density . . . J -33 3- J 9 4.95 (solid) Colour .... /pale greenish 1 yellow greenish yellow reddish brown bluish black Without exception, the magnitude of the physical properties increase, or become more pronounced, as the atomic weight increases. OXIDES AND OXYGEN ACIDS OF HALOGENS 191 The same is true of the chemical behaviour. Thus the affinity for hydrogen is greatest in the case of fluorine, and gradually diminishes until, in the case of iodine, the rate of combination is very slow and is incomplete. This is perhaps best shown by a comparison of the heats of formation of the hydrogen compounds (in gaseous form) from their elements (p. 145). H 2 + F 2 = 2 H F + 2 x 38,600 cal. H 2 +C1 2 =2HC1 + 2X22,000 H 2 + Br 2 = 2HBr + 2 x 8,400 H 2 +I 2 =2HI -2x6,100 The order of affinity of the halogens for metals is the same as for hydrogen, as is shown by the fact that the element of smaller atomic weight displaces that with a larger atomic weight from combination. In contrast to the behaviour with hydrogen and the metals, the affinity for oxygen increases with the atomic weight of the halogen. Thus fluorine forms no compound with oxygen, the compounds of chlorine and oxygen are unstable, whilst iodine pentoxide is rela- tively stable. Similarly, the halogens of higher atomic weight, under certain circumstances, displace the others from combination in oxygen compounds, thus potassium chlorate is transformed almost com- pletely into potassium bromate by fusing with potassium bromide : KBr->KBrO 3 The more important points of resemblance in the halogen group is that all form compounds of the type HX with hydrogen (X = halo- gen) ; these compounds are colourless gases, which fume in the air, and form strong acids when dissolved in water. Further, the salts containing the same metals and different halogens are of the same type, and the compounds with the alkali metals occur in cubic crystals. It is of interest to note, however, that the behaviour of fluorine differs more from that of the other halogens than the latter do among themselves. Thus hydrogen fluoride tends to polymerize (P- ! 53)> double salts of the type KF'HF exist, and silver fluoride, unlike the other silver halides, is readily soluble in water. CHAPTER XV OSMOTIC PRESSURE AND MOLECULAR WEIGHT IN SOLUTION eneral When a compound can be converted into vapour with- VJT out decomposition, its molecular weight can be determined by the volumetric method already described (p. 115). Many compounds, however, cannot be vaporized without decomposition, and their molecular weights cannot therefore be determined by the standard method. An example of this already met with is hydrogen peroxide. Analysis shows that it contains hydrogen and oxygen in equal atomic proportions, so that its molecular formula is either HO or an integral multiple of HO, such as H 2 O 2 , H 3 O 3 , etc. A more or less definite decision between these formulae can be arrived at on the ground of its chemical behaviour. For example, the readiness with which it splits up into water and oxygen appears to be most in harmony with the generally accepted formula. It will, however, clearly be of great advantage in this and similar cases if the molecular weight can be determined, in solution, for instance, by a direct method. Within the last thirty years such methods have been devised, mainly owing to the labours of the Dutch chemist, van 't Hoff, and have contributed enormously to the advance of chemistry. These methods depend ultimately on the conception of osmotic pressure, which will now be considered. Osmosis. Osmotic Pressure A piece of parchment paper or animal membrane is tied over the mouth of a thistle funnel A, which is then inverted, filled up to a point on the stem with an aqueous solution of copper sulphate, and then supported in a vessel of water B, so that the liquid outside and inside is at the same height (Fig. 45). After a time it will be observed that the solution has risen to a higher point in the tube and the level of the water has slightly fallen, Further, the water in the vessel is slightly blue, showing that a little copper sulphate has passed through the membrane. The phenomenon is clearly analogous to the interdiffusion of two gases separated by a porous membrane (p. 46). Water passes into 192 OSMOTIC PRESSURE the tube more quickly than the copper sulphate diffuses out, with the result that the bulk of solution inside increases. The process is called osmosis. The inward movement is sometimes termed endosinosis, and the outward stream exosmosis. When a bladder is used as membrane one of the substances (water) diffuses through it much faster than the other (copper sulphate), and it might be suggested that with a suitable membrane the passage outwards of the dissolved substance might be stopped altogether without interfering with the entry of water. This has been acti ally realized with a membrane of copper ferrocyanide, when cane sugar is used as solute. A membrane which allows one substance to pass through and entirely prevents the passage of another is said to be semipermeable. The copper ferrocyanide membrane is usually pre- FlG ' 45 ' pared by interaction of solutions of copper sulphate and of potassium ferrocyanide. It is much too weak to use in place of the bladder in the experiment described above ; but this difficulty can be overcome by depositing it in the walls of a porous pot, as first suggested by the German botanist Pfeffer (1877). The pot (one of those commonly used for experiments in gas diffusion is suitable) is first carefully washed, soaked in water for some time, then nearly filled with a solution of copper sulphate (2.5 grams per litre) dipped nearly to the neck in a solution of potassium ferrocyanide (2.1 grams per litre), and allowed to stand for some hours. The salts diffuse through the walls of the pot, and at their junction form a membrane of copper ferrocyanide which, owing to its being supported by the walls of the vessel, is capable of withstanding fairly high pressures. The cell is then taken out, washed, rilled with a concentrated solution of sugar, and closed with a well-fitting cork through which passes a glass tube open at both ends (Fig. 46). When a cell thus prepared is immersed in water the latter passes in and the level of liquid in the narrow tube slowly rises, and finally reaches a point at which, if the cell has been properly 13 B i 9 4 A TEXT-BOOK OF INORGANIC CHEMISTRY prepared, it remains constant for days. Under these circumstances the pressure inside the cell is greater than that outside by an amount measured by the difference in height of the liquid inside and outside the cell. This excess of pressure, which must prevail inside the cell in order to prevent more water flowing in through the semi-permeable membrane, is termed the osmotic pressure of the solution. If the cell is dipped into a more concentrated solution of sugar instead of into water, water passes outwards from the more dilute to the more concentrated solution, so that the concentrations tend to become equal. In the same way, if solutions of different substances are separated by a semi-permeable membrane, the solvent always passes through the membrane from the solution with lower to that with higher osmotic pressure. In other words, the direction of flow is always such that the pressures on the two sides of the membrane tend to become equal. The exact mode in which the osmotic pressure arises is not well understood. 1 According to one view, the extra pressure which pre- vails inside the cell is due to the bombardment of the interior walls by the particles of solute ; just as, according to the kinetic theory (p. 48), gas pressure is due to impacts of the particles on the walls of the containing vessel. It is interesting to note that, just as a mem- brane of copper ferrocyanide allows water, but not dissolved sugar, to pass through, a membrane made of metallic palladium allows hydrogen to pass through, but is impervious to other gases. On this principle an apparatus can be constructed which admits of the measure- ment of the partial pressure of a gas such as nitrogen, when mixed with hydrogen. A vessel of palladium, which can be heated to any con- venient temperature, contains nitrogen at a pressure of say half an atmosphere, as observed on a manometer. When the vessel, heated to 600, is surrounded by hydrogen at atmospheric pressure, the latter passes inwards through the membrane and the pressure inside in- creases. When equilibrium is finally attained, the pressure inside the cell is about \\ atmospheres. The excess of pressure inside over that outside is clearly due to the nitrogen. It is evident that there is a close analogy between this experiment and that by means of which the osmotic pressure of dissolved cane sugar is determined. If a concentrated solution of copper sulphate is carefully placed, by means of a pipette, at the bottom of a beaker of water so that there is a definite boundary between the two layers, it will be found after a time that the copper sulphate is uniformly distributed through the 1 Ci'. Physical Chemistry, p. 106. OSMOTIC PRESSURE 195 water (diffusion of liquids). We may in this case assume that the cause of diffusion is the higher osmotic pressure of the copper sulphate in certain parts of the system (just as the cause of gaseous diffusion is the unequal partial pressures of the gases in different parts of the system), and equilibrium is attained when the concentration (and the osmotic pressure) of the solute is uniform throughout. Osmotic Pressure and Molecular Weight in Solution Regard ng osmotic pressure as analogous to gas pressure, the relationsh p between volume, osmotic pressure and temperature can now be investigated for solutions, as has already been done for gases. Firstly, it may be shown experimentally that the osmotic pressure is approximately proportional to the concentration of the solution. This is evident from the following results obtained by Pfeffer for solutions of cane sugar, in which the concentrations are expressed in grams per 100 c.c. of solution and the pressures in cm. of mercury. Concentration, C . . . . i 2 4 6 Osmotic pressure, P . . . 53.1 101.6 208.2 307.5 Ratio, P/C 53.5 50.8 52.1 51.3 As the concentration is inversely proportional to the Volume V in which a definite quantity of solute is dissolved, we obtain, by substi- tuting i/V for C, the relationship PV = constant, that is, the osmotic pressure exerted by a definite quantity of a solute x the volume of solution in which it is dissolved is constant at constant temperature a result exactly analogous to Boyle's law for gases. Further, it may be shown that the osmotic pressure, like the gas pressure, increases proportionately to the absolute temperature T at constant volume. Pfeffer found the osmotic pressure of a I per cent, solution of cane sugar to be 51.0 cm. at 14.2. If P is proportional to T, the osmotic pressure at 32 should be 51 x -|^- =54.1 cm., whilst the value actually found was 54.6 cm. Finally, it remains to find the relationship between the magnitudes of the gas pressure and osmotic pressure under comparable condi- tions. This can also be done on the basis of Pfeffer's experiments with cane sugar, for example, from the observation that at o a i per cent, solution of cane sugar at o exerts an osmotic pressure equal to 49.3 cm. of mercury. We have already seen that the molecular weight of a gas is that quantity of it which, when present in 22.4 litres at V, exerts' a pressure of i atmosphere (76 cm.). Now we know 196 A TEXT-BOOK OF INORGANIC CHEMISTRY from chemical considerations that the molecular weight of cane sugar is 342. As, according to Pfefifer, I gram in 100 c.c. exerts an osmotic pressure of 49.3 cm. at o, the molecular weight of cane sugar, 342 grams, when present in 22.4 litres or 22,400 c.c., must exert an osmotic pressure of 49.3 x ^~ x - =75 cm. approx. which is, within the limits of experimental error, the same as the gas pressure it would exert if present as single molecules in the same volume in the gaseous form in the absence of the solvent. That this result is not accidental can be shown by similar experiments with other substances of known molecular weight. The molecular weight of glucose (grape sugar) is 180, of ethyl alcohol 46, and it has been found that in order to obtain an osmotic pressure of I atmosphere at o, 1 80 grams of grape sugar, or 46 grams of ethyl alcohol, must be present in 22,400 c.c. of solution. From this and other considera- tions van 't Hoff drew the very important conclusion that the osmotic pressure exerted by any substance in solution is the same as it would exert if present as gas in the same volume as that occupied by the solu- tion, provided the solution is sufficiently dihtte. On this basis is founded a method of determining molecular weights in solution which exactly corresponds with that used for gases and vapours (p. 1 10). The molecular weight of a dissolved substance is that quantity of it which, when present in 22.4 litres at o, exerts an osmotic pressure of i atmosphere. The general result of molecular weight determinations in solution by this method is that the values obtained usually correspond with the simplest formula ascribed to a substance from its chemical behaviour. Certain important exceptions to this rule will be con- sidered later. Determination of Molecular Weights in Solution from Depression of the Freezing-point and Elevation of the Boiling-point of Solutions l The results described in the last section may be summed up in the statement that the osmotic pressure of solutions containing the same number of molecules of different solutes in equal volumes of the same solvent is the same ; in other words, the osmotic pressure depends on the number and not on the nature of the particles. The determination of molecular weights by this method is, however, too complicated for general use, and 1 For details see Physical Chemistry, pp. 109-138. OSMOTIC PRESSURE 197 it has been found much more convenient to measure certain other magnitudes which are proportional to the osmotic pressure. The two most important of these are the depression of the freezing-point and the elevation of the boiling-point produced by the addition of a known weight of a soluble substance to a known weight or volume of solvent. It has been known for a long time that a dilute solution freezes at a lower temperature than the pure solvent. For example, a solution containing 5 grams of sodium chloride in 100 grams of water freezes at 2.9, pure ice separating. For a fixed quantity of solvent the lowering of the freezing-point is proportional to the amount of soluble substance added, and for a definite quantity of dissolved substance the depression is inversely proportional to the amount of solvent ; in other words, the depression is proportional to the concentration of the solution. Further, like the osmotic pressure, the depression caused by the same number of molecules of different solutes in equal amounts of the same solvent is the same. The lowering produced by i mol of solute in 100 grams of solvent is called the molecular freezing-point depression ; it vaneswith the nature of the solvent. The molecular depression K for a solvent is found once for all by experiment with a number of substances of known molecular weight, and the molecular weight of any other substance is that quantity of it which, when dis- solved in 100 grams of the solvent, produces the depression K. The value of K for water is 18.5, for benzene 50, and for acetic acid 39. In determining molecular weights by the freezing-point method, it is not of course necessary to use solutions of the concentrations stated above. The depression is determined in relatively dilute solutions, and the weight in grams which would produce a depression K in 100 grams of solvent is calculated on the assumption that even in concen- trated solutions the depression is proportional to the concentration of the solution. The formula for calculating the results is readily ob- tained as follows : If g grains of solute, of unknown molecular weight M, dissolved in L grams of solvent (that is, ioo-/L grams in 100 grams of solvent) lowers the freezing-point 8 degrees, whereas, according to the above law, M grams of solute in 100 grams of solvent lowers the freezing-point by K degrees, we have In order to illustrate this formula, we may calculate the molecular weight of hydrogen peroxide from the observation that the lowering 198 A TEXT-BOOK OF INORGANIC CHEMISTRY of the freezing-point of a solution containing 0.0943 grams of the com- pound in 1 8.90 grams of water is 0.270. Substituting in the above formula 100 x 0.0943 x 18.5 M= .8.90x0.270 '=34.2 in good agreement with the value, 34.0, calculated for the formula H 2 2 . The effect of dissolved substances in elevating the boiling-point of a solvent exactly corresponds with the effect in depressing the freezing- point, and need not be considered in detail. In the former case also the elevation is proportional to the concentration of the solution, and the same number of molecules of different substances in equal amounts of the same solvent raise the boiling-point of the solvent to the same extent. The molecular elevation constant can be found by experiments with substances of known molecular weight in the usual way. The value for water is 5.2, for ethyl alcohol 11.5, and for benzene 26.7. It is important to note that the methods of determining molecular weights just described only apply when the pure solvent, unmixed with the solvent, separates (as solid or vapour). The magnitude of the two effects just discussed is much less than that of the osmotic pressure. Thus a mol of solute in a litre of water lowers the freezing-point by 1.85, raises the boiling-point by 0.52, whereas the osmotic pressure of the solution is about 22.4 atmos- pheres. Eutectic Mixtures. Solid Solutions As we have seen, the lowering of the freezing-point of a fixed quantity of solvent by a dissolved substance is proportional to the amount of the latter added. An alternative method of stating the facts is that the temperature at which solution and solid solvent are. in equilibrium is the lower the greater the concentration of the solution. The extent to which the freezing-point of the solvent can be lowered is limited, however, by the fact that the solubility of solids in liquids is limited. The effect of potassium iodide in lowering the freezing-point of water is shown in Fig. 47, the ordinates representing temperatures and the abscissse concentrations. The point A represents the freezing-point of water (o), and the curve AO the temperatures at which ice and solution are in equilibrium with gradually increasing concentrations of salt. O represents the point at which the solution is saturated with potassium iodide ; at this time, therefore, solid potassium iodide, as well as ice., OSMOTIC PRESSURE 199 must be in equilibrium with the solution. The curve BO represents the effect of temperature on the solubility of potassium iodide ; it is the curve along which solid potassium iodide is in equilibrium with the solution. Since the solubility diminishes as the temperature is lowered, the direction of the curve must -be as shown. If the tem- perature of a saturated solution is progressively lowered, salt will separate, and the concentration will fall till the temperature is reached at which ice also begins to separate. At the latter temperature salt and ice are in equilibrium with the solution, which is the case at the point O, and at no other point. The latter is therefore the point of intersection of the freezing-point and solubility curves. If a mixture corresponding with the composition at the point O is cooled till the freezing-point is reached, ice and salt separate in the proportions in which they are present in the mixture. 30 Hence the composition of the solution does not alter during solidification, and 10 therefore the solution freezes at constant tem- perature (p. 68) like a pure substance. The mixture of -20 the solid components in _^ Q equilibrium with the solu- tion is termed a eutectic mixture ; when one of the components is ice it is termed a cryohydrate. The latter term was employed because it was formerly supposed that "cryohydrates" were definite compounds of salt and water, but it can be shown by microscopic examination and in other ways that they are in fact mechanical mixtures of the two components. It sometimes happens that when a solution is cooled, homogeneous crystals containing both components separate. Mixtures of this kind are known as mixed crystals or solid solutions. They resemble liquid solutions in being homogeneous, and also inasmuch as the composi- tion varies continuously between certain limits. The freezing-point of a mixture whose components form solid solutions may be higher or lowe than those of the components, depending upon the propor- tion in which the latter separate from solution. Grams KI in 100 grams solution i i i . i i i i i "30 40 50 60 FIG. 47. 70 80 90 CHAPTER XVI NITROGEN, THE ATMOSPHERE AND THE ELEMENTS OF THE HELIUM GROUP IT has already been stated (p. 29) that the atmosphere is composed essentially of the two gaseous elements oxygen and nitrogen. Within the last twenty years it has been discovered that no less than five hitherto unknown elements are always present in the atmosphere, though in comparatively small amount. Besides these elements, the atmosphere also contains varying proportions of water vapour, carbon dioxide, ammonia, oxides of nitrogen, and various other substances In the present chapter the nitrogen will first be considered, then the atmosphere, and finally the preparation and properties of the rare elements just referred to. NITROGEN Symbol, N. Atomic weight, 14.01. Molecular weight, 28.02. History Nitrogen was first isolated by Rutherford, Professor of Botany at Edinburgh, in 1772. He kept animals in a confined volume of air for some time, and after removing the "fixed air" (carbon dioxide) with caustic potash, found that the residual air was incapable of supporting life or combustion. He termed the gas mephitic air. Lavoisier was the first to regard mephitic air as an element. He termed it azote (from a, privative, and o>??, life) in allusion to its inability to support life. Chaptal first suggested the name nitrogen, because it is contained in nitre or saltpetre. The opinion held up to 1894, that atmospheric nitrogen is a single substance, was based on the work of Cavendish (cf, p. 207). Occurrence Nitrogen occurs free in the atmosphere, of which it constitutes about 75.6 per cent, by weight or 78 per cent, by volume. Free nitrogen also occurs to a very small extent occluded in certain minerals. In combination with carbon and hydrogen it forms an essential constituent of plants and animals. In combination with hydrogen it forms ammonia, and in combination with oxygen and NITROGEN 201 other elements it occurs as nitrates. At present one of the chief sources oi combined nitrogen is natural sodium nitrate (Chili saltpetre). Preparation Nitrogen can be obtained by two principal methods : (A) by removing the oxygen from purified air by means of easily oxidizable substances ; (B) from compounds containing nitrogen by chemical methods. The nitrogen obtained from air is only about 99 per cent, pure, as it is mixed with argon and other gases of the helium gioup which cannot easily be removed. (A) The methods of obtaining nitrogen by removal of oxygen from atmospheric air must be so chosen that the products of oxidation can readily be separated from the nitrogen. The more important are as follows : (1) By means of phosphorus. A small piece of phosphorus is caused to burn in air confined over water under a bell-jar. Dense white fumes of an oxide of phosphorus (phosphorus pentoxide) are produced, which dissolve in the water to form an acid. The nitrogen obtained in this way is impure, because the phosphorus becomes extinguished before the oxygen is completely removed. (2) By red-hot copper. The air is first freed from moisture and carbon dioxide by passing through U -tubes containing suitable absorbing agents, and is then drawn through a long tube containing copper filings and heated in a furnace. The nitrogen thus prepared contains argon, but is otherwise fairly pure. (3) In the last method, the copper slowly becomes inactive through conversion to cupric oxide. If, however, a mixture of air and ammonia gas is drawn through the furnace the copper oxide is continually reduced to metallic copper by the ammonia, and the process can be continued indefinitely. 1 (4) By ilkaline pyrogallol. A very rapid and convenient absorption of oxygen is secured by shaking air with a solution containing pyro- gallol and sodium or potassium hydroxide. The oxygen is used up in oxidizing the pyrogallol to dark-coloured complex compounds. This method is employed for estimating oxygen in gaseous mixtures. (B) For the preparation of nitrogen by chemical methods,ammonium compounds or salts of nitrous and nitric acids are generally used. (5) When a concentrated aqueous solution of ammonium nitrite is heated, it splits up directly into nitrogen and water: 1 Equation: 3CuO + 2NH 3 -3Cu + N 2 +3H 2 O. The excess of ammonia can be removed by passing the gas through water. 202 A TEXT-BOOK OF INORGANIC CHEMISTRY As ammonium nitrite does not keep well, it is more usual to employ a mixture of sodium nitrite and ammonium chloride : (6) Nitrogen is also obtained by heating ammonium bichromate (in practice a mixture of ammonium chloride and potassium bi- chromate is used) : (NH 4 ) 3 Cr 2 7 ->N 2 + Cr 2 3 + 4H 2 0. (7) By passing an oxide of nitrogen, for example, nitric oxide, NO, over red-hot copper : 2NO + 2Cu->2CuO + N 2 . (8) By the action of chlorine on ammonia, the latter being kept in excess : Physical Properties Nitrogen is a colourless, odourless, taste- less gas. The density of pure nitrogen referred to air is 0.9673, whilst that of atmospheric nitrogen is 0.9721 (Rayleigh). It was this dis- crepancy in the densities of nitrogen prepared by chemical methods and that obtained from the atmosphere that led in 1894 to the discovery in "atmospheric" nitrogen of a hitherto unknown element, argon, the density of which is about 1.4 times that of nitrogen. Nitrogen has been obtained both in the liquid and solid form. Liquid nitrogen boils at 196.6 under atmospheric pressure; its critical temperature is 146 and its critical pressure 35 atmospheres (Olszewski). Solid nitrogen is a white crystalline substance melting at -210 (Fischer and Alt) ; at -252.5 the density of the solid is 1.0265 (Dewar). Nitrogen is only very slightly soluble in water. At o i c.c. of water dissolves 0.0239 c.c., at 20 0.0164 c.c., and at 40 0.0118 c.c. of nitrogen at 76 cm. pressure. Chemical Properties At ordinary temperatures nitrogen is a very inactive element. At higher temperatures it combines directly with a number of other elements, more particularly lithium, magnesium, barium, calcium, and boron, to form nitrides. In the case of lithium, combination takes place slowly at the ordinary temperature. The formula of lithium nitride is Li 3 N, of magnesium nitride, Mg 3 N 2 . Under the influence of the electric spark or the electric arc, nitrogen and hydrogen form small quantities of ammonia, NH 3 , and nitrogen and oxygen combine to form a brown oxide of nitrogen, NO 2 (p. 227). THE ATMOSPHERE 203 THE ATMOSPHERE As already mentioned, the atmosphere is mainly composed of oxygen and nitrogen, in the proportion' of I part of oxygen to 4 parts of nitrogen by volume. The question of the nature of the atmosphere has 'naturally engaged man's attention from the very earliest times, but the first substantial advances in knowledge were made in connexion with the phenomena of combustion. Mayow, as early as 1674, showed' that only a portion of the air was absorbed in combustion and respiration, and similar observations were made in 1692 by Boyle. The complete explanation of the phenomena of combustion, and the clear recognition of the fact that air is essen- tially composed of an elementary substance (oxygen) which combines with other substances during combustion, and of another element (nitrogen) which plays no part in combustion, is due to Lavoisier (P. 27). Composition of the Atmosphere The relative propor- tions of oxygen and nitrogen in the atmosphere by volume can be determined by absorbing the oxygen from a known volume of air by ore of the methods already described (p. 201), and measuring the residual nitrogen. A more convenient method is to add to a measured volume of air in a eudiometer excess of hydrogen, explode the mixture by means of an electric spark, and measure the residual gas. As one volume of oxygen combines with two volumes of hydrogen, and the volume of the water formed is negligible, ^ of the contraction represents the volume of oxygen in the original volume of air. The relative proportions of oxygen and nitrogen by weight can be determined by drawing air, freed from carbon dioxide and moisture, over heated copper, which combines with the oxygen ; the nitrogen passes en into a previously evacuated flask and is weighed. The oxygen is determined by finding the increase in weight of the copper. This method was used by Dumas and Boussingault (1841). The result of numerous analyses of air from the most varied sources shows that there are slight but distinct variations in its composiiion. According to Angus Smith, the proportion of oxygen in towns, especially during foggy weather, is as low as 20.82 per cent, by volume, whereas on open moors and mountains it is as high as 21.00. As a mean of all experiments, it has been found that the atmosphere contains 20.93 volumes of oxygen and 79-7 volumes 204 A TEXT-BOOK OF INORGANIC CHEMISTRY of nitrogen (including argon), or 23 per cent, of oxygen to 77 per cent, of nitrogen by weight. A litre of air (free from carbon dioxide and moisture), at o and 76 cm. pressure, weighs 1.2933 grams. The Atmosphere a Mixture Owing to the remarkable constancy of the ratio of oxygen and nitrogen in the atmosphere, it has sometimes been suggested that the gases are in chemical combination. This suggestion, however, is untenable for many reasons : 1 i ) There are distinct though small variations in the composition of the air, whereas the composition of a chemical compound is constant. (2) The relative amounts of the two gases do not bear any simple relation to their combining weights. (3) Heat is neither given out nor absorbed when the gases are mixed in the proportions in which they are present in air, and the resulting mixture has all the properties of air.- (4) The oxygen and nitrogen retain their separate properties in air, as shown (a) by the fact that a partial separation can be effected by taking advantage of their different rates of diffusion, (b) when air is shaken up with water, the gases dissolve in accordance with their respective solubilities and partial pressures (p. 77). The last point is rather interesting as it indicates a comparatively simple method of obtaining a mixture rich in oxygen. It can be calculated that the air expelled from water by boiling contains 35 volumes of oxygen and 65 volumes of nitrogen. Substances present in smaller Proportion in the Atmosphere Besides oxygen, nitrogen, and the inactive gases, the atmosphere contains varying proportions of aqueous vapour, carbon dioxide, ammonia, nitric acid, hydrogen, and a number of suspended impurities, including dust particles, bacteria and other organisms. In the neighbourhood of towns, sulphur dioxide and other substances, mainly resulting from the combustion of coal, are found. The presence of ozone in normal air is doubtful (p. 138). Each of these constituents wilT now be briefly considered.' Water Vapour At a given temperature a definite volume of air can take up only a certain 'amount of water vapour ; when this maximum is reached the air is said to be saturated. Air saturated with moisture contains at o 4.871 grams, at 10 9.362 grams, at 20 17.157 grams, and at 30 30.095 grams of aqueous vapour per cubic metre. On the average, the amount of moisture is about two-thirds of the saturation value, but considerable varia- THE ATMOSPHERE 205 tions occur. When the temperature falls below that at which the amount of moisture is just sufficient for saturation, the excess is deposited as dew, mist or rain. The amount of aqueous vapour in the atmosphere can be deter- mined by drawing a known volume of air through bulbs containing pumice stone soaked with concentrated sulphuric acid, and determin- ing the increase in weight. Hygrometers are also used for the same purpose. Carbon Dioxide The normal proportion of carbon dioxide in country air is about 3 volumes in 10,000, but in large towns, where much coal is burned, it may rise to 6 to 7 parts in 10,000. Air contain- ing more than 7 volumes of this gas in 10,000 is considered harmful for respiration, but this is due less to the gas itself than to the accompanying impurities. The proportion of carbon dioxide in the atmosphere is continually increasing owing to combustion of coal, etc., and to respiration (p. 335), but it is simultaneously being taken up in enormous quantities by plants in the process of assimilation. A more or less accurate compensation is thus attained, and it is not certain whether at the present time the amount of carbon dioxide is increasing or decreasing. The proportion of carbon dioxide in the air is estimated by drawing a measured volume of it through a solution of barium hydroxide of known strength : Ba(OH) 2 + CO 2 ->BaCO 3 -hH 2 0, and finding by volumetric analysis theamount of unchanged hydroxide. Ammonia This substance results from the decay of nitro- genous organic matter (p. 214). Angus Smith found 0.5 to i.o gram in 10,000 grams of air from different sources, but still greater varia- tions occur. It does not exist to any great extent free in the air, but forms salts with the carbon dioxide, nitrous and nitric acids which are also present. These are washed into the soil during rain, and form a very important source of the combined nitrogen required for the growth of vegetation. Nitrous and Nitric Acids Under the influence of electrical discharges in the atmosphere, the nitrogen and oxygen com- bine to form oxides of nitrogen, the latter then reacting with water vapour to form nitrous and nitric acids. As already mentioned, these compounds are washed into the soil by rain and form a valuable source of nitrogen for plants. 206 A TEXT-BOOK OF INORGANIC CHEMISTRY It was formerly supposed that plants could not utilize atmospheric nitrogen, but it has recently been shown that leguminous plants, such as beans, peas, and clovers, can make use of nitrogen by the agency of bacteria. These bacteria are to be found in nodules on the roots of the plants in question, and by means not yet under- stood they convert the atmospheric nitrogen (which reaches them through the pores of the soil) into compounds which are readily assimilated by plants. As already mentioned (p. 31), hydrogen is a normal constituent of the atmosphere, but the proportion is very small, probably not more than i part in a million. Suspended Impurities in the Atmosphere The air contains a large amount of suspended impurities, which are rendered visible when a ray of light enters a darkened room. These consist partly of inorganic substances, including particles of salts, etc., and partly of organic substances, including bacteria and other micro- organisms. Air contains under normal conditions only 4 to 5 micro-organisms per litre on the average, whereas a pure, unfiltered river water contains from 6,000 to 20,000 in i c.c., and un- disturbed soil about 100,000 in i c.c. The majority of the micro- organisms in the air consist of the spores of yeast and of moulds which produce fermentation. The bacteria found in the air, in- cluding the bacilli of various diseases, seldom float free in the air, but are almost always attached to dust particles, and the latter, with their adherent bacteria, tend to settle out of the air owing to their relatively great specific gravity. Bacteria thus reach the air chiefly from the soil, and their distribution in the atmosphere is favoured by dryness of the soil and air currents. Air can be almost entirely freed from suspended impurities by drawing it through a tube loosely packed with cotton wool. The dust particles in the air act as nuclei for the condensation of moisture, giving rise to fogs. This may be well illustrated as follows. A large flask containing a little water at the bottom is partially ex- hausted by means of a pump ; the temperature is thus considerably lowered, and a fog forms in the flask. If, however, the air, before being admitted to the flask, is filtered through cotton wool, no fog is produced on rapid exhaustion. THE ELEMENTS OF THE HELIUM GROUP 207 THE ELEMENTS OF THE HELIUM GROUP Helium, at. wt.=3.99. Neon, at. wt.- 20.2. Argon, at. wt. =39.88 Krypton, at. wt. = 82.9. Xenon, at. wt. = 130. 2. History In a paper published in 1785, Cavendish describes experiments designed to find whether "atmospheric nitrogen" is exclusively composed of one substance. He confined a mixture of atmosphere nitrogen and oxygen over potassium hydroxide solution and passed electric sparks through the mixture till no further absorp- tion took place. Only a small bubble remained, and Cavendish drew the conclusion that " if there is any part of the phlogisticated air (nitrogen) of the atmosphere which differs from the rest, and cannot be reduced to ' nitrous acid,' we may conclude that it is not more than T |Q part of the whole." In 1894 Lord Rayleigh, in the course of careful experiments on the densities of gases, found that whereas a litre of nitrogen obtained from chemical compounds weighs 1.2521 grams under normal con- ditions, a litre of nitrogen prepared from the atmosphere weighs 1.2572 grams. The experiments undertaken by Rayleigh and by Ramsay to account for this remarkable difference led to the discovery of argon, a gas about 1.4 times heavier than nitrogen. It is an interesting fact, in connexion with the statement of Cavendish just mentioned, that the atmosphere contains rather less than i per cent, by volume of argon. While searching for other sources of argon, Ramsay (1895) heated a uranium-containing mineral called cleveite, and obtained a gas the spectrum of which gave a line (the D 3 line) not previously observed with any terrestrial gas, but which had been detected in the solar spectrum. Lockyer ascribed this characteristic line in the solar spectrum to an element which he called helium, and this name was retained by Ramsay when the element was discovered on the earth. Argon and helium differ from all other elements in being remark- ably inactive ; so far it has not been found possible to bring them into chemical combination. For reasons which will be mentioned later, Ramsay was led to search for other inactive gases, and in 1898 three other gases, belonging to the same family as argon and helium, were discovered by Ramsay and Travers. The lightest of the three, neon, was obtained by fractionating a large quantity of argon, the other two, krypton (density 41.5) and xenon (density 65.1), were found in :he residues after the evaporation of a large quantity of liquid air. Each of these gases will now be briefly described. 208 A TEXT-BOOK OF INORGANIC CHEMISTRY ARGON Symbol, A. Density, 19.94. Atomic and molecular weight, 39.9. Occurrence Argon is present in the atmosphere to the extent of 0.94 per cent, by volume, or 1.3 per cent by weight. It is also found occluded in certain minerals, and is present in very small proportion in most natural waters, especially in those from certain springs. Methods of Preparation (i) Argon may be obtained from the atmosphere by passing an electric discharge through a mixture of air and oxygen in the presence of alkali till no further change of volume occurs (p. 223). The excess of oxygen is then removed by alkaline pyrogallate and the residue is argon. This method was used by Lord Rayleigh in the original preparation of argon, but is much less convenient than that now to be mentioned. (2) Argon may also be prepared from air by first removing the oxygen with red-hot copper and then passing the residual gas, freed from carbon dioxide and aqueous vapour, backwards and forwards over heated magnesium, or better, over heated calcium (a mixture of calcium oxide and magnesium powder is used) till no further diminu- tion of volume occurs. The magnesium or calcium combine with the nitrogen to form nitrides. The argon thus obtained is contaminated with the other inactive gases, but can be purified by the use of liquid hydrogen (Ramsay and Travers). When the crude argon is passed into a tube immersed in liquid hydrogen, helium passes on, but neon, argon, krypton and xenon are liquefied or solidified. On allowing the mixture to warm up, the neon, being very volatile, first distils off", then argon begins to distil, and if the temperature is kept at a suitable point, pure argon distils off, while the less volatile krypton and xenon remain behind in the solid form. Properties Argon, like all the other members of this group, is a colourless, odourless gas. Its density is 19.94, and therefore its molecular weight is 39.88. So far, all attempts to make argon enter into chemical combination have been fruitless, so that its atomic weight cannot be determined by the usual volumetric method (p. 1 15). It has, however, been determined by the following method that 'argon and the other gases of this group have only one atom in the molecule, so that the atomic and molecular weight are identical. It can be shown 1 that the molecular heat 2 of a monatomic gas at 1 Physical Chemistry, p. 46. 2 The product of specific heat and molecular weight. THE ELEMENTS OF THE HELIUM GROUP 209 constant volume, that is, the heat required to raise i mol of the gas i in temperature, is 3 calories, but more heat is required to raise a mol of a polyatomic gas i in temperature, as in this case heat is also absorbed in doing work within the molecules. Since it was found by experiment that the molecular heat of the inactive gases was 3 calories, it follows that they are monatomic. The molecular heat of mercury vapour is also 3 calories at constant volume, and the con- clusion that it is monatomic has been confirmed by other methods. Liquid argon boils at 186 and solid argon melts at - 188. Its critical temperature is -117.4 an d the critical pressure 53 atmos- pheres. Argon is considerably more soluble in water than nitrogen ; at 15 100 volumes of water dissolve about 4 volumes of argon. The spectrum of argon is very complex, and changes markedly when the nature of the discharge is altered. With the intermittent discharge the glow in the tube is red, and few blue lines appear in the spectrum ; with an oscillating discharge the glow is bright blue, the red lines in the spectrum disappear or become faint, and many new green and blue lines appear. HELIUM Symbol, He. Density, 2.0. Atomic and molecular weight, 3.99. Occurrence Helium occurs in certain rare minerals containing uranium or thorium, notably in cleveite, uraninite, monazite sand, and thorianite. Onnes obtained the 200 litres of helium used in the liquefaction of the gas by heating monazite sand (p. 477). Ramsay found that thorianite, a rare mineral from Ceylon containing 76 to 78 per cent, of thorium oxide, along with the oxides of uranium and the cerium metals, yielded 9 c.c. of helium per gram when heated to redness. Helium is also present in certain mineral waters, notably in those at Bath, which contain argon with 8-10 per cent, of helium. It is also present in very small proportion, according to Ramsay to the extent of 0.0004 P er cent, by volume, in the atmosphere. Preparation Helium is most conveniently obtained from the minerals containing it ; the finely-powdered material is heated alone or with dilute sulphuric acid. The other gases generally present are removed by the usual methods. Helium can also be obtained from the non-condensible gases obtained in the process of making liquid air. The gases escaping from the liquefier contain a considerable proportion of helium and neon. All the gases except helium are condensed and in fact 14 2io A TEXT-BOOK OF INORGANIC CHEMISTRY solidified at the temperature of liquid hydrogen, so that the helium can be pumped off practically pure. Properties Next to hydrogen, helium is the lightest gas known, its density being only 2. After many failures, it was liquefied for the first time by Kammerlingh Onnes of Leyden in 1908. 200 litres of the gas, cooled by the evaporation of liquid hydrogen under reduced pressure, were circulated round the liquefier (of the Linde type) for about three hours, and finally about 60 c.c. of liquid helium was obtained. The liquid boiled about 4.3 abs. (-268.7 C.) and its density was 0.15 ; the critical temperature is about 5 abs. By evaporation of liquid helium a temperature within 3 of the absolute zero was reached. Helium is only very slightly soluble in water ; at o 100 c.c. of water dissolve 0.015 c - c -> anc ^ a ^ 2O 0.0138 c.c. of the gas. With the intermittent discharge at 7 to 8 mm. pressure, helium gives a yellow spectrum, the characteristic so-called D 3 line (p. 207) attain- ing its greatest intensity ; when the pressure is reduced to I to 2 mm. the tube emits a brilliant green light. The yellow D 3 line, a red line, and two lines in the green are most characteristic. NEON Symbol, Ne. Density, 10.1. Atomic and molecular weight, 20.2. Neon was isolated from the less condensible portion of liquid air, as already mentioned under helium. The complete separation from argon and helium was rendered possible by the preparation of liquid hydrogen in quantity by Travers. The argon and neon were solidified in a bulb immersed in liquid hydrogen and the helium pumped off: the temperature was then allowed slowly to rise. As neon boils at a lower temperature than argon, the first fraction was rich in neon and contained very little argon. Properties The boiling-point and melting-point of neon have not been accurately determined. When an electric discharge is passed through it in a vacuum tube the colour is orange red. The spectrum contains a number of lines in the red, a prominent line, D 6 , in the yellow, and a number of strong green lines. KRYPTON AND XENON Krypton, Symbol Kr. Density, 41.5. Atomic weight, 83.0. Xenon, Symbol X. Density, 65.1. Atomic weight, 130.2. Preparation These gases were isolated from the residue obtained in the evaporation of a large quantity of liquid air. The THE ELEMENTS OF THE HELIUM GROUP 211 oxygen and nitrogen were removed by the usual methods, the residue, containing krypton and xenon, was again liquefied and the substances separated by fractional distillation, krypton being considerably more volatile than xenon. Properties Krypton boils about* -152 C., and the solid melts at - 169 ; xenon boils about - 109 C. and the solid melts about - 140. Like helium and neon, the spectrum of krypton is independent of the nature of the electric discharge. The spectrum contains a few distinct lines in the red, but a yellow and a green line are the most brilliant. The green line appears to be identical with a prominent line in the spectrum of the aurora borealis. The spectrum of xenon, like that of argon, varies with the nature of the discharge. With the intermittent discharge the glow is blue ; when a Leyden jar and spark gap are introduced in the circuit, the colour is green. Summary of Group The members of this group, like the halogens, form a family of elements, whose properties vary regularly with increasing atomic weight. The more important physical pro- perties of the elements are summarized in the table : Helium. Neon. Argon. Krypton. Xenon. Atomic weight 3-99 20.2 39.88 82.9 130.2 Boiling-point . -268.7 -233 -i 86 -151.7 -I0 9 Melting-point . ... -188 - 169 - 140 Critical temperature -268 -117.4 - 62.5 + 14-75 Atomic volume ... 32.9 38.5 37-0 The elements of this group resemble each other in being chemically inactive ; so far none of them have been brought into chemical reaction. We may therefore say that the valency of the group is zero. They are all monatomic, as shown by considerations based on their specific heat. It may be stated that a number of attempts have been made to obtain other elements belonging to this group, but hitherto without definite success. Ramsay and Moore examined the residues from the evaporation of over 100 tons of liquid air made by Claude's method (p. 74), but found no element denser than xenon. 212 A TEXT-BOOK OF INORGANIC CHEMISTRY The proportion of these gases present under normal circumstances in the atmosphere is given in the following table (Ramsay) : Parts by Weight Parts by Volume in i Part of Air. in I Part of Air. Helium 0.00000056 0.000004 Neon 0.0000086 0.0000123 Argon 0.0136 0.00937 Krypton 0.00028 0.00009 Xenon ..... 0.00005 0.0000 1 T CHAPTER XVII COMPOUNDS OF NITROGEN WITH HYDROGEN AND WITH THE HALOGENS HRKE compounds of nitrogen and hydrogen are known. Their names and formulas are as follows : Ammonia ...... NH 3 Hydrazine N 2 H 4 or NH 2 'NH 2 Hydrazoic acid or azoimide . . . N 3 H It will also be convenient to describe in this chapter a compound of nitrogen with hydrogen and oxygen. Hydroxylamine NH 2 'OH AMMONIA Formula, NH 3 . Molecular weight, 17-03. History Compounds of ammonia have been known from the earliest times. Ammonium chloride was formerly imported into Europe from Egypt, where it was prepared (by sublimation) from the soot obtained on burning camel's dung, which was used as fuel near the temple of Jupiter Ammon in Libya. From this circumstance the salt was called sal-ammoniac, whence the name ammonia is derived. At a later period compounds of ammonia were prepared by the dry distillation of horns, hoofs and other animal matter. From this moce of preparation is derived the name spirits of hartshorn, sometimes applied to ammonia. Ammonia gas was first obtained by Priestley (1774), who collected it over mercury. Occurrence As already mentioned, ammonia occurs in the atmosphere, partly in the free state, but chiefly as carbonate and nitrate. It is found as the chloride and sulphate near active vol- canoes. As it is a product of the decay of animal and vegetable matter, it is found in the combined state in soils and in natural waters. 213 2i 4 A TEXT-BOOK OF INORGANIC CHEMISTRY Preparation (i) Traces of ammonia are obtained by passing a silent electrical discharge through a mixture of nitrogen and hydrogen : As the reaction is a reversible one, and the equilibrium lies very near the left-hand side, only a small proportion of ammonia is formed, but if it is removed by an acid as fast as it is produced the reaction goes completely in the direction of the upper arrow. This is an excellent illustration of the influence of the concentration of the reacting substances on the equilibrium (p. 165). (2) By the dry distillation of organic nitrogenous compounds (animal and vegetable matter). When coal, which results from the slow decay of vegetation, is heated out of contact with air, ammonia is one of the products. It is absorbed in water, and the solution, the so-called "ammoniacal liquor of the gas-works," is the chief com- mercial source of ammonia compounds. The ammoniacal liquor is impure, but a fairly pure ammonium salt can be obtained by heating the liquor with slaked lime, and absorbing the ammonia in hydro- chloric or sulphuric acid (cf. p. 402). (3) When organic nitrogenous compounds are heated for some hours with concentrated sulphuric acid, the nitrogen is completely converted into ammonia, and remains in the solution as ammonium sulphate. Kjeldahl's method for estimating nitrogen in organic compounds is based on this reaction. The nitrogen of many organic compounds is also converted into ammonia on heating with a con- centrated solution of an alkali, especially in the presence of potassium permanganate. (4) Nitrates and nitrites are reduced to ammonia by "nascent" hydrogen, that is, by hydrogen generated in the solution : NaNO 3 This method is made use of in the estimation of nitrites and nitrates in drinking water. (5) Pure ammonia gas is best prepared by heating ammonium chloride with calcium hydroxide. Ammonium salts contain the univalent group NH 4 , which plays the part of a univalent metal (p. 402). The formula for ammonium chloride is therefore NH 4 C1, and the reaction just described may be written as follows : NITROGEN-HYDROGEN COMPOUNDS 215 The gas is dried by passing it through a long tube containing lumps of calcium oxide, and can be collected over mercury or by downward displacement of air. Physical Properties Ammonia is a colourless gas, with a pungent smell and caustic taste. Its density referred to air is 0.597. It can bo condensed to a colourless liquid, which boils at -32.5, and on further cooling yields a solid which melts at -77. Its critical temperature is 131 and critical pressure 113 atmospheres. Liquefied ammonia is an excellent solvent. It absorbs a large amount of heat on evaporation, a fact taken advantage of in machines for making artificial ice. Ammonia gas is very soluble in water. At o i c.c. of water absorbs 1299 c - c> ? a * Io 7^3 c - c -j an d at 28 595 c.c. of ammonia, measured at o and 760 mm. pressure (Raoult). The gas is com- pletely expelled from its aqueous solution by boiling. Quite recently the equilibrium 2NH 3 ^N 2 + 3H 2 has been thoroughly investigated. The equilibrium mixture at 700 and 30 atmospheres pressure contains less than i per cent, of ammonia ; at ordinary pressures the proportion is of course much smaller. As the heat of formation of ammonia from its elements is positive (11,890 cal. for r; grams) the equilibrium is displaced in favour of the ammonia as the temperature is lowered (p. 172). Chemical Properties The most striking chemical property of ammonia is that it combines directly with acids to form salts, thus NH 3 +HC1 = NH 4 C1 H 2 SO 4 = (NH 4 ) 2 SO 4 . As already mentioned, the salts obtained in this way are termed ammonium salts, the NH 4 , or ammonium group, playing the part of a metal. Ammonium salts can be vaporized by heating, and the densities of the vapours are considerably less than the values cal- culated from their formulae. In the case of ammonium chloride, for instance, the molecular weight calculated from its formula is 53.5, but the density, instead of being 26.7, is only about half that value. The simplest explanation is that in the form of vapour ammonium chloride is dissociated almost completely into ammonia and hydrogen chloride, according to the reversible equation NH 4 C1^NH 3 +HC1. 2i6 A TEXT-BOOK OF INORGANIC CHEMISTRY This suggestion has been fully confirmed by experiment, a partial separation of the products of dissociation having been effected by taking advantage of their different rates of diffusion. The conclusion which may be drawn from the above results, that ammonia is basic in character, is confirmed by the observation that its aqueous solution is alkaline to litmus. It has already been stated that all bases in solution contain the OH group, and we may therefore assume that am- monia reacts with water to form ammonium hydroxide, NH 4 OH, corresponding with the alkali hydroxides KOHandNaOH. Ammonia does not burn in air, but in oxygen it burns with a yellow flame. The chief products are water and nitrogen, but traces of oxides of nitrogen are also formed. This in- teresting reaction is best shown by means of the arrangement represented in Fig. 48. Ammonia obtained by boiling the aqueous solu- tion in the flask A passes through a glass tube en- closed in a wider tube. When oxygen is admitted to the cylindrical space between the two tubes, the ammonia can be burned at the end of the narrow tube ; FIG. 48. but if the supply of oxygen is cut off, the flame becomes smaller and finally goes out. Ammonia is completely decomposed by chlorine : NITROGEN-HYDROGEN COMPOUNDS 217 When excess of ammonia is present, the hydrochloric acid is con- verted into ammonium chloride. Many metals decompose ammonia into its elements at high tem- peratures, and in this way a number of metallic nitrides have been obtained. Nitrides may in fact be regarded as being derived from ammonia, by displacing the hydrogen by metals. They are decom- posed by vater, forming ammonia and metallic hydroxides : Just as water unites with a number of metallic salts as so-called water of crystallization, so many solid compounds of metallic salts with ammonia are known. Of these, the compounds with silver chloride, 2AgCl,3NH 3 and AgCl,3NH 3 , and the compounds with copper sul- phate, CuSO 4 ,4NH 3 ; CuSO 4 ,2NH 3 and CuSO 4 ,NH 3 , are most familiar. Ammonia can be liberated from its compounds by warming with an alkali, and is detected by its characteristic smell, and by the fact that it turns litmus paper blue. Traces of ammonia, or of ammonium compounds, are detected by Nessler's reagent (p. 459). Composition of Ammonia If the formula of ammonia is NH.,, the equation representing its formation from nitrogen and hydrogen must be as follows : N 2 + 3 H 2 = 2NH 3 2 vols. 6 vols. 4 vols. that is, one volume of nitrogen and three volumes of hydrogen unite to form two volumes of ammonia. Perhaps the simplest method of confirming the formula is to submit a measured volume of the gas to electric sparks in the apparatus represented on p. 223 till its volume no longer changes. On again adjusting to atmospheric pressure, it will be observed that the volume has practically doubled, in accord- ance with the above equation. Oxygen is then added in excess, and an electric spark passed through the mixture. From the contraction, due to the formation of water, the volume of hydrogen in the original mixture can be calculated, and it may be shown that one volume of nitrogen to three of hydrogen were present. The fact that ammonia is decomposed by chlorine may also be taken advantage of to show the composition of the former gas. A long glass tube, provided with a stopcock and a funnel above, is marked off externally into three equal sections by indiarubber rings (Fig. 49), and filled with chlorine by displacement of a strong solution 218 A TEXT-BOOK OF INORGANIC CHEMISTRY of brine. A concentrated solution of ammonia is poured into the funnel, and then admitted, a drop at a time, into the chlorine by cautiously turning the tap. At first a yellowish flame follows the admission of each drop, but after a time this effect ceases, as all the chlorine is decomposed. When a considerable excess (three or four c.c.) of ammonia solution has been added, so as to ensure complete conversion into ammonium chloride and nitrogen (p. 220), excess of dilute sulphuric acid is added through the tap to neutralize the ammonia, and water is then cautiously drawn into the tube in the same way, until no more will enter. If the experiment has been performed in such a way that no gas has escaped through the tap, the residual gas, which is practically under atmospheric pressure, fills one of the divisions of the tube, and may easily be shown to be nitrogen. As hydrogen and chlorine combine in equal volumes, the original three volumes of chlorine must have combined with three volumes of hydrogen, and one volume of nitrogen has been liberated. It follows that ammonia is formed by the FIG. 49. combination of one volume of nitrogen with three volumes of hydro- gen. Hence the formula is (NHg)*, and as the molecular weight is 17, it follows that the molecular formula for ammonia is NH 3 . The same fact may also be proved by electrolyzing a concentrated aqueous solution of ammonia, to which some ammonium sulphate has been added, in the apparatus represented on p. 14. Hydrogen is liberated at the negative and nitrogen at the positive pole, and the volume of the former gas is three times that of the latter. HYDRAZINE, N 2 H 4 Preparation Hydrazine, or diamide, H 2 N-NH 2 , was discovered by Curtius (1887). It may be obtained by the reduction of hyponitrous acid (p. 234) : RON - - NH 2 +2H 2 O, but is most easily prepared by adding sodium hypochlorite to excess of a con- centrated solution of ammonium hydroxide, which contains about 0.02 per cent. of glue. The excess of ammonia is removed by boiling, the solution is concen- trated and on addition of sulphuric acid hydrazine sulphate, N 2 H4,H 2 SO4, crystallizes out (Raschig). Hydrazine Hydrate, N 2 H 4 ,H 2 O, is obtained by distilling a salt of hydrazine, for instance, the sulphate, with a very concentrated solution of potassium hydroxide. It is a colourless liquid, which boils at 120, smells of ammonia, and has a powerful corrosive action. It can be manipulated only in vessels of silver or platinum. Hydrazine hydrate was known for some years before hydrazine NITROGEN-HYDROGEN COMPOUNDS 219 itself had bee i obtained. The latter is prepared from the hydrate by heating with barium oxide : N 2 H 4 'H 2 + BaO = N 2 H 4 + Ba(OH) 2 . Properties At ordinary temperatures hydrazine is a colourless liquid which boils at 113, and can easily be obtained in the fofm of a solid melting at i. The density of the liquid at 15 is 1.0114. Like ammonia, hydrazine is a base, combining with acids to form salts. One molecule is able to combine with one and with two molecules of hydrochloric acid, forming the compounds N 2 H 4 ,HC1 and N 2 H 4 ,2HC1. It combines very readily with water to form the hydrate, N 2 H 4 ,H 2 O. Hydrazine is also a powerful reducing agent, owing to the readiness with which it can be oxidized to water and nitrogen : HYDRAZOIC ACID OR AZOIMIDE, N 3 H Preparation This remarkable compound was first obtained by Curtius in 1890. It is nost conveniently prepared by passing nitrous oxide, N 2 O, over fused sodium amide, NaNH 2 , at 190, and distilling the resulting sodium hydra- zoate with dilute sulphuric acid : NaNH 2 + N 2 O = 2 NaN 3 + H 2 S0 4 =2N 3 Properties -The anhydrous acid is a colourless, mobile liquid, which boils at 37, has a penetrating odour, and readily decomposes explosively into its elements : Hydrazoic acid is monobasic and its salts closely resemble those of the halogen acids. The similarity in the behaviour of the N 3 group and that of the Cl and Br atoms is v >ry striking ; it is also observed in organic compounds. A comparison of the three compounds of nitrogen and hydrogen just con- sidered shows that the basic character diminishes as the proportion of nitrogen increases. Hydrazine is a weaker base than ammonia, and hydrazoic acid is an acid of medium strength. HYDROXYLAMINE, NH 2 OH Preparation Hydro xylamine, discovered by Lossen in 1865, can be obtained by the action of " nascent " hydrogen on nitric oxide or nitric acid : =2NH 2 OH The most convenient method is to pass nitric oxide through a mixture in which hydrogen is being generated by the action of hydrochloric acid on tin, hydroxyl- amine hydro :hloride, NH 2 OH'HC1, being thus obtained. The tin is removed by means of hydrogen sulphide, and the hydroxylamine hydrochloride dissolved out with abs< lute alcohol. Hydroxylamine can be obtained by reducing nitric acid by means of hydrogen 220 A TEXT-BOOK OF INORGANIC CHEMISTRY generated electrolytically. By using a cathode of amalgamated lead (lead coated with mercury), 80 per cent, of the nitric acid used can be converted into hydroxyl- amine. Anhydrous hydro xylamine can be obtained most readily by heating hydroxyl- amine phosphate in a vacuum, the free base passing over at 135 to 137 : (NH 2 OH) 3 -H 3 P0 4 - 3 HN 2 OHt + H 3 P0 4 . Properties Hydroxylamine occurs in colourless needles which melt at 33. The liquid boils without decomposition at 56 to 58 under 22 mm. pressure, but decomposes at higher temperatures, explosively above 100, the chief, but not the exclusive products, being ammonia, nitrogen, and water : It is readily soluble in water, forming an alkaline solution. It is a considerably weaker base than ammonia, from which it may be regarded as being derived by the replacement of a hydrogen atom by the OH group. In connexion with the weakly basic character of hydroxylamine, it is an interesting fact that one mole- cule of a monobasic acid can combine with more than one molecule of this base. For example, the hydrochloride (NH 2 OH) 2 'HC1 is known, as well as the normal compound, NH 2 OH'HC1. Hydroxylamine is a powerful reducing agent. It immediately precipitates silver and gold from solutions of their salts and reduces cupric salts in alkaline solution to red cuprous oxide : COMPOUNDS OF NITROGEN WITH THE HALOGENS Three definite compounds of nitrogen with the halogens, nitrogen trichloride, NC1 3 , nitrogen tribromide,*^^^ and an iodide, NH 3 ,NI 3 , are known. They may be regarded as ammonia in which the hydrogen is partially or completely dis- placed by halogen atoms. They are all extremely explosive. Nitrogen Chloride, NC1 3 . Preparation This compound is obtained by the action of chlorine upon ammonium chloride : NH 4 C1 + 3 C1 2 = 4 HC1 + NC1 3 . When a jar of chlorine is inverted in a solution of ammonium chloride kept at 30-40, oily drops of the trichloride settle out. Properties Nitrogen trichloride is a dark-yellow, oily liquid of density 1.653. I* explodes violently at the ordinary temperature in contact with traces of organic matter or of phosphorus, and sometimes spontaneously, without any apparent cause. It also explodes on sudden heating, or on exposure to sunlight. It boils at 71 and can be distilled, but the operation is a very dangerous one. The vapour strongly attacks the eyes and the mucous membrane. It is soluble in benzene, chloroform, and other liquids, and these solutions are fairly stable. The trichloride is decomposed by excess of ammonia : NC1 3 + 4 NH 3 = 3 NH 4 C1 + N 2 , and there is therefore no danger in preparing nitrogen from ammonia and chlorine (p. 202) if the former is in excess. NITROGEN-HALOGEN COMPOUNDS 221 As would be anticipated from its very explosive character, nitrogen chloride is highly endothermic. In the formation of one mol from its elements, about 40,000 cal. are absorbed. Nitrogen Tribromide is obtained as a red, extremely explosive, oily liquid by the action of potassium bromide on nitrogen trichloride. Its composition has not been conclusively established, but, from analogy with the chloride, is generally assumed to be represented by the formula NBr 3 . Nitrogen Iodide, NI 3 'NH 3 . Preparation Nitrogen iodide is obtained as a brown powder by the action of a concentrated solution of ammonia on powdered iodine, on icdine dissolved in alcohol, or on iodine chloride, IC1. Properties When pure, nitrogen iodide occurs in lustrous, copper-coloured needles of density 3.5. When moist, it has not much tendency to explode, but in the dry st.ite it is extremely unstable, even a touch with a feather or the tread of a fly being sufficient to bring about explosion. The constitution of nitrogen iodide has been the subject of a great deal of discussion since its discovery by Courtois in ;8i2, but within the last ten years it has been definitely established by Chattaway, Orton and Stevens, and by Silberrad, that the formula is as above. CHAPTER XVIII OXIDES AND OXYACIDS OF NITROGEN 1TJMVE oxides of nitrogen are known Nitrous oxide (hyponitrous anhydride) . N. 2 O Nitric oxide NO Nitrogen trioxide (nitrous anhydride) . N. 2 O 3 Nitrogen peroxide NO 2 or N 2 O 4 Nitrogen pentoxide (nitric anhydride) .' N 2 O 6 Three acids are known, corresponding with the three anhydrides referred to above Hyponitrous acid H 2 N 2 O 2 Nitrous acid HNO 2 Nitric acid HNO 3 As all the other oxides and oxyacids are derived directly or in- directly from nitric acid, it will be convenient to deal first with this important compound. NITRIC ACID, HNO 3 History Nitric acid was probably familiar to the ancient Egyptians. It was prepared in the ninth century by the alchemist Geber by heating together nitre (potassium nitrate), alum and copper sulphate. The method of preparation now in use action of sulphuric acid on an alkali nitrate appears to have been introduced by Glauber about 1660. In 1784 Cavendish obtained it by passing electric sparks through a mixture of nitrogen and oxygen, and bringing the product in contact with water; in this way the constitution of the acid was established. In allusion to its powerful solvent action on metals, i was formerly called aquafortis. The name now in use refers to its usual mode of preparation from nitre. Occurrence Nitric acid does not occur free in nature, except perhaps in traces in the atmosphere, but in the form of its salts, the OXIDES AND OXYACIDS OF NITROGEN 223 nitrates, it is widely distributed. Sodium nitrate, NaNO 3 , known as Chili saltpetre, occurs in enormous deposits in the rainless regions of South America. Further, nitrates represent the final oxidation pro- duct of nitrogenous animal and vegetable matter, and for this reason they are found in all soils. It is interesting to note that this oxidation is effected by atmospheric oxygen with the co-operation of certain bacteria (p. 399). The occurrence of nitrates, chiefly ammonium nitrate, in the atmosphere has already been mentioned. Preparation (i) The preparation of nitric acid by passing sparks through a mixture of oxygen and nitrogen-, and treating the product with water, has already been referred to. The reaction may be carried out by means of the arrangement represented in Fig. 50. Sparks frc-m the Ruhmkorff coil are passed through the mixture in the globe until the colour is cieep brown owing to the formation of nitrogen peroxide, NO 2 ; on shaking with water the solution will be found to have an acid reaction : Nitric acid and nitric oxide are formed, and the reaction is reversible. A modification of this process has been introduced in recent years for the purpose of preparing nitric acid on th-s large scale from the atmosphere ; it is briefly described on p. 236. (2) Nitric acid is usqally prepared in the laboratory by heating potassium or sodium nitrate with concentrated sulphuric acid in a glass retort, and collecting the acid in a cooled receiver (Fig. 51) : NaNO 3 + H 2 SO 4 ->NaHSO 4 + HNO 3 f . Sodium acid sulphate, NaHSO 4 , remains in the retort. The acid thus obtained contains water and nitrogen peroxide. The water is almost entirely removed by adding concentrated sulphuric acid and redistilling, and the nitrogen peroxide by drawing a current of carbon dioxide or air through it. On the commercial scale, sodium nitrate and sulphuric acid are mixed in i cast-iron still in the proportions required by the above equation, ihe acid is distilled off under reduced pressure in order to minimize decomposition (see below); and is condensed in a series of 224 A TEXT-BOOK OF INORGANIC CHEMISTRY earthenware pots. If anhydrous sodium nitrate is used, the distillate consists of acid almost free from water. Physical Properties According to recent investigations, loc per cent, nitric acid exists only at low temperatures in the form of colourless crystals, which melt at 41.3, forming a liquid which i< yellowish in colour owing to slight decomposition : A mixture containing 98.6 per cent, of the acid is colourless and moderately stable at the ordinary temperature. It fumes in the air. and can be distilled without decomposition under reduced pressure, but under atmospheric pressure it boils at 86 with partial decomposi FIG. 51. tion according to the above equation. The commercial acid contains 68 per cent, of real nitric acid, and is quite stable. This composition corresponds with a constant boiling mixture of the acid and water (cf. hydrochloric acid, p. 95), which boils at 120.5 under at- mospheric pressure and has a specific- gravity of 1.414 at 15'. Nitric acid is completely decomposed into water, nitrogen peroxide and oxygen, according to the above equation, when heated to 260. Chemical Properties The more important chemical proper- ties of nitric acid can be classified under three heads (i) it is a typical monobasic acid, combining with bases to form salts, the nitrates ; (2) on account of the readiness with which it yields oxygen it is a power- ful oxidizing agent ; (3) it acts on certain organic compounds so thai: 1 Under reduced pressure the acid distils at a much lower temperature. OXIDES AND OXYACIDS OF NITROGEN 225 one or more atoms of hydrogen are displaced by univalent NO 2 groups. Nitrates are obtained by the action of nitric acid on the oxides or carbonates of the metals or on the metals themselves. With very few except ons, they are readily soluble in water. . On heating strongly they are all decomposed, the majority giving off oxygen and nitrogen peroxide a id leaving an oxide of the metal. Nitric acid acts on all metals except gold, platinum, iridium and rhodium, nitrates or oxides of the metal being formed, whilst the nitric acid is reduced to lower oxides of nitrogen. The vigour of the action and also the stage to which the nitric acid is reduced depend on the nature of the metal, the concentration of the acid, the tempera- ture and other factors. In contrast to other acids, hydrogen is very seldom evolved when nitric acid acts on metals. 1 We may assume that it is used up in reducing part of the nitric acid, being itself oxidized to water. The different stages in the reduction of nitric acid are represented by nitrogen peroxide, NO 2 , nitrogen trioxide, N 2 O 3 , or nitrous acid, HNO 2 , nitric oxide, NO, nitrous oxide, N 2 O, nitrogen, N 2 , and finally ammonia, NH 3 . When copper is acted on by nitric acid, the main reaction is represented by the equation and the action on mercury and on lead is represented by correspond- ing equations. With zinc, on the other hand, reduction to nitrous oxide (and under certain circumstances even to ammonia) takes place : Although gold and platinum are not acted on by nitric acid, they are readily dissolved by a mixture of nitric and hydrochloric acids, the so- called aqita regret (p. 235). Nitric acid also exerts an oxidizing action on many non-metals, with formation of oxides and oxyacids. Thus iodine is oxidized to iodic acid, HIO 3 (p. 185), phosphorus to phosphoric acid, H 3 PO 4 , and sulphur to sulphuric acid, H 2 SO 4 . The action of nitric acid on metals and non-metals is further referred to in connexion with the elements themselves. When nitric acid acts on glycerine, nitroglycerine, the principal constituent of dynamite, is obtained, and similarly, by the action of nitric acid on cotton-wool, the highly explosive gun-cotton is 1 When nitric acid acts on magnesium, hydrogen is liberated. 15 226 A TEXT-BOOK OF INORGANIC CHEMISTRY obtained. As already mentioned, these compounds are formed by the displacement of hydrogen atoms in the organic compounds by NO 2 groups, and illustrate the third main property of nitric acid. A similar kind of change takes place when dilute nitric acid acts on the skin, yellow compounds being produced. Concentrated nitric acid, on the other hand, causes painful wounds. Fuming nitric acid is obtained by distilling concentrated nitric acid with concentrated sulphuric acid, or by passing nitrogen trioxide into nitric acid. It consists of concentrated nitric acid containing nitrogen peroxide in solution, is red in colour, fumes strongly in the air and has powerful oxidizing properties. Methods of writing Complicated Equations An equa- tion such as that representing the action of nitric acid on copper can readily be written in stages. In such a case it is perhaps simplest to assume that the oxygen of the nitric acid is used up in converting the metal to oxide, the latter then uniting with more nitric acid to form the nitrate. As a preliminary to writing the equation, the stage to which the nitric acid is reduced must be determined experi- mentally, and the formula of the nitrate must be known. The steps in the action of nitric acid on copper are as follows : (i) (2) (3) 3 CuO + 6HN0 3 =3Cu(N0 3 ) 2 Adding up A rather more complicated case is that in which iodine is oxidized to iodic acid. Assuming that the nitric acid in this case also is reduced to nitric oxide, and remembering that the oxide corresponding with iodic acid is I 2 O 6 , we have : (i) (2) (3) Equations (i) and (2) have to be multiplied by 5 and 3 respectively in order that all the available oxygen may be used up. Adding : 3l 2 +ioHNO 3 ->6HIO 3 +ioNO + 2H 2 O. The student should apply this method to other reactions, for example to the action of nitric acid on phosphorus, on sulphur, and on zinc. OXIDES AND OXYACIDS OF NITROGEN 227 An alternative method is to represent the first stage of the reaction as involving the liberation of hydrogen, which is then used up in reducing the nitric acid to a lower stage of oxidation. It is some- times assumed that this represents the actual stages through which the reactions pass, but in reality they are probably much more compli- cated. The remarkable fact that pure dilute nitric acid has little or no action on copper or mercury unless a little nitrous acid is added cannot easily be accounted for on this view as to the mode in which the reaction proceeds. NITROGEN PENTOXIDE (NITRIC ANHYDRIDE), N 2 O 6 Prepa:ration-i(i) By the abstraction of the elements of water from nitric acid by means of phosphorus pentoxide: Phosphorus pentoxide is added to pure concentrated nitric acid in the cold till a little remains undissolved ; the mixture is then cautiously distilled and the pentoxide collected in a cool receiver. (2) By the action of chlorine on silver nitrate : 4 AgN 3 + 2Cl 2 ->4AgCl + 2 + 2N 2 O 6 . Dry chlorine is passed over silver nitrate in a U-tube kept at 50-60 ; the pentoxide distils off and is collected in a cool receiver. Properties Nitrogen pentoxide occurs in colourless, lustrous, rhombic crystals, which melt at 29.5 with partial decomposition. The reddish liquid begins to boil at 45 and decomposes rapidly, giving off brown fumes of nitrogen peroxide : It dissolves in water to form nitric acj^, with evolution of much heat: NITROGEN PEROXIDE, NO 2 OR N 2 O 4 Preparation (i) By passing electric sparks through a mixture of nitrogen arid oxygen : / ( 2) By direct combination of nitric_oxide and oxygen: 228 A TEXT-BOOK OF INORGANIC CHEMISTRY (3) By heating the nitrate of a heavy metal, e.g. lead nitrate : The carefully dried salt is heated in a glass tube and the products of decomposition passed into a U-tube immersed in a mixture of ice and salt, the nitrogen peroxide condensing as a nearly colourless liquid. Physical Properties At ordinary temperatures nitrogen per- oxide is usually met with as a reddish-brown gas which can readily be condensed to a liquid. On further cooling it forms colourless crystals, which melt at 10. Even at its melting-point the liquid is pale yellow, and the colour gradually deepens to orange as the temperature rises. The liquid boils at 26, changing to a reddish- brown vapour, which becomes progressively deeper in colour as the temperature is further raised. These changes of colour take place in the reverse order as the temperature is lowered. Simultaneously with the deepening of colour, a diminution in density occurs as the temperature is raised, as the following figures show : Temperature. . 26.7 60.2 90.0 121.5 1 35 l $ Density. . . 38.3 30.1 24.8 23.5 23.1 23.0 These numbers indicate that at 26.7 the molecular weight is about 76 and at 150 46. The formula NO 2 corresponds with a molecular weight of 46 and the double formula N 2 O 4 with a molecular weight of 92. The simplest explanation of the above numbers is that at 150 the gas consists entirely of NO 2 molecules and at lower tem- peratures of a mixture of NO 2 and N 2 O 4 molecules. From the changes of colour we may assume that the compound N 2 O 4 is colour- less and NO 2 deep brown. We are therefore dealing with an equilibrium represented by the equation N 2 4 ^t 2NO 2 2 unit vols. 4 unit vols. and it can be calculated from the above figures that at 26.7 the mixture contains 33 per cent, of NO 2 molecules whilst at 150 dis- sociation is practically complete. On heating nitrogen peroxide to higher temperatures, it dissociates (p. 169) into nitric oxide and oxygen: OXIDES AND OXYACIDS OF NITROGEN 229 Under atmospheric pressure, dissociation is practically complete at 620. Nitrogen peroxide is a supporter of combustion, provided the tem- perature is sufficiently high to liberate oxygen. Thus a taper is extinguished when plunged into the gas, but phosphorus burning in air continues to burn brilliantly in nitrogen peroxide. It acts as a powerful oxidizing agent towards many substances. When water in moderate excess acts on nitrogen peroxide at low temperatures a mixture of nitrous and nitric acids is obtained: At higher temperatures nitric acid and nitric oxide are formed : With a cold aqueous solution of potassium hydroxide, a mixture of nitrite ci.nd nitrate is obtained: These facts show that nitrogen peroxide, like chlorine peroxide, is a mixed anhydride, since it forms two acids with water and two salts with bases. Nitrogen Trioxide, N 2 O 3 Preparation (i) When equal volumes of nitrogen peroxide and nitric oxide are passed through a tube cooled to - 20, nitrogen trioxide is obtained as a blue liquid. (2) By heating nitric acid (density 1.35) with arsenic trioxide, and passing the gas through a tube which is kept cool, the same product is obtained : As 2 O 3 + 2HNO 3 ->As 2 O 6 +H 2 O + (NO + NO 2 ). Properties Until quite recently nitrogen trioxide was only known in the liquid form. Below -21 the liquid is fairly stable/.but at higher temperatures it decomposes into NO and NO 2 , and in the form of vapour, as shown by density measurements, is decomposed almost completely into these two oxides : Baker ( 1907) has shown, however, that if the liquid trioxide is dried extremely carefully before vaporizing, it has the density 38, so that 1 Liquid nitrogen trioxide is deep indigo blue below -2, but is green (and doubtless partially decomposed) at the ordinary temperature. On cooling in liquid air it forms deep blue crystals, which melt at - 103. 2 3 o A TEXT-BOOK OF INORGANIC CHEMISTRY no dissociation had taken place. In fact, in many of Baker's experi- ments a density exceeding 38 was observed, which probably indicates partial polymerization to N 4 O 6 molecules. Nitrogen trioxide dissolves in cold water to form nitrous acid : N 2 O 3 +H 2 O = 2HNO 2 . Nitrous Acid, HNO 2 Preparation The acid itself is un- stable, and has never been obtained in the free condition, but is moderately stable in dilute aqueous solution below o. It is obtained by, dissolving nitrogen trioxide, N 2 O 3 , in cold water : N 2 O 3 + H 2 O->2HNO 2 . The salts of nitrous acid, the nitrites, are quite stable. The nitrites of the alkali metals, e.g. potassium nitrite, are obtained by heating the nitrates alone, or better, with metallic lead : 2KNO 3 ->2KNO 2 + O 2 Properties Even in dilute aqueous solution nitrous acid decom- poses slowly at room temperature, more rapidly on warming, into nitric acid, nitric oxide, and water : For this reason brown fumes are given off when dilute acids (e.g. dilute sulphuric acid) are added to solutions of alkali nitrites, the nitrous acid which is presumably first liberated decomposing accord- ing to the above equation (distinction from nitrates). Nitrous acid is an acid of medium strength (p. 187) ; it is much weaker than nitric acid. When a strong oxidizing agent, such as potassium permanganate, is added to nitrou^acid, the latter is oxidized to nitric acid : + O 2 =2HN0 3 , and can therefore act as a reducing agent. On the other hand, nitrous acid gives up oxygen to certain easily oxidized substances, such as hydrogen sulphide, hydriodic acid, or indigo solution, being itself reduced to a lower oxide of nitrogen, for example, nitrous oxide : According to the conditions, nitrous acid can therefore act as an oxidizing or as a reducing agent. In these respects it resembles OXIDES AND OXYACIDS OF NITROGEN 231 hydrogen peroxide, which is also able to reduce an acidified solution of potassium permanganate and to oxidize indigo and hydriodic acid (p. 141). NITRIC OXIDE, NO Nitric oxide was first recognized as a definite chemical compound by Priestley (1772). It is the substance first formed by the direct com- bination of nitrogen and oxygen at high temperatures or under the influence of the electric discharge. Preparation (i) By the action of nitric acid of density 1.2 (30 to 35 per cent, of the pure acid) on copper, the temperature being kept low : The gas obtained in this way, although sufficiently pure for some purposes, always contains other oxides of nitrogen in small amount, and may be purified by passing into a cold concentrated solution of ferrous sulphate, which absorbs only nitric oxide (see below) and then heating the solution. The gas is collected over water. (2) Pui e nitric oxide is obtained by heating nitric acid with a mixture of ferrous sulphate, FeSO 4 , and dilute sulphuric acid, or ferrous chlo- ride and hydrochloric acid : [2FeSO 4 + H 2 SO 4 + O->Fe 2 (S0 4 ) 3 + H 2 O] x 3. Adding Fe 2 (SO 4 ) 5 is ferric sulphate (p. 546). Physical Properties Nitric oxide is a colourless gas. Its density referred- to hydrogen is 15, corresponding with the formula NO ; unlike nitrogen peroxide, it shows no tendency to polymeriza- tion, even at -70. It can be obtained as a colourless liquid, which boils at -153.6 ; the critical temperature is -93.5 and the critical pressure 71.2 atmospheres (Olszewski). Solid nitric oxide occurs in colourless crystals, which melt at - 167. Nitric oxide is very slightly soluble in water. At o I volume of water absorbs 0.074 c.c., at 10 0.057 c.c., and at 20 0.047 c.c. of the gas. As nitric oxide is partially decomposed into its elements at high temperatures in the presence of catalytic agents, the reaction is reversible. Nernst has shown that at 1760 C. the equilibrium 232 A TEXT-BOOK OF INORGANIC CHEMISTRY mixture contains 0.64 per cent., at 2210 2.05 per cent., and at 3000 about 5 per cent, of nitric oxide. The displacement of the equilibrium in favour of nitric oxide with increasing temperature is in accordance with its endothermic character, the thermochemical equation being N 2 + O 2 = 2NO -2 x 21,500 cal. It follows that at ordinary temperatures nitric oxide is unstable, its apparent stability being due to the fact that the reaction 2NO-^N 2 + O 2 is exceedingly slow under these conditions. Chemical Properties Owing to the excessive slowness with which it splits up into its elements, nitric oxide is the most stable of all the oxides of nitrogen.. Even at a reil_heat it is practically unaffected. For the same reason it is not a supporter of combustion in the ordinary sense, as a lighted candle or phosphorus burning feebly are extinguished when plunged into it. If, however, the tem- perature is sufficiently high to decompose it, it supports combustion ; thus burning magnesium or brightly burning phosphorus continue to burn in it. As already mentioned, nitric oxide combines with oxygen to form brown fumes of nitrogen peroxide : The combination is practically complete at low temperatures, but above 1 50 dissociation of the peroxide becomes appreciable. Nitric oxide is readily absorbed by aqueous solutions of ferrous salts, e.g. ferrous sulphate, forming a dark-coloured compound in solution which is readily decomposed on heating (p. 231). The consti- tution of this compound has not been satisfactorily established, but is probably FeSO 4 'NO (Manchot and Zechentmayer, 1907). On the formation of this compound is based the so-called " brown ring " test for nitrates. To a solution of a nitrate a crystal of ferrous sulphate is added, and after shaking and pouring concentrated sulphuric acid down the inside of the tube, a brown ring of the above compound forms at the junction of acid and solution. Composition Nitric oxide is completely decomposed on heat- ing with metallic sodium or iron, and the residual nitrogen has just half the volume of the original gas : 4 vols. 2 vols. A molecule of nitric oxide, therefore, contains one atom of nitrogen, OKIDES AND OXYACIDS OF NITROGEN 233 weight 14, and as the density of nitric oxide is 15, and therefore its molecular weight 30, the molecule contains 30-14 = 16 parts by weight, or one atom of oxygen, and its formula is therefore NO. NITROUS OXIDE (NITROUS ANHYDRIDE, LAUGHING GAS), N 2 O This g;is was discovered by Priestley ; its composition and more important properties were established by Davy. Preparation On cautiously heating ammonium nitrate (the salt, not a. solution) it splits up directly into water and nitrous oxide : NH 4 NO 3 ->N 2 Instead of ammonium nitrate a mixture of ammonium sulphate and sodium nitrate in equivalent proportions may be used ; by double decomposition ammonium nitrate is formed, which then decomposes as above. Ammonium chloride should not be used instead of ammo- nium sulphate, as a very impure gas is then obtained. When required pure, e.g. for anaesthetic purposes, the gas is passed through a solution of ferrous sulphate to remove nitric oxide, through alkali to remove traces of chlorine, and is collected over hot water or over mercury. Nitric acid can be reduced to nitrous oxide by the action of metals such as zinc or copper under certain conditions, but a mixture of gases is always obtained in these reactions (p. 225). Physical Properties Nitrous oxide is a colourless gas with an agreeable sweetish "odour and taste. Its density is 22. It can readily be condensed to a colourless liquid, which boils at -89.5; its critical temperature is 35.4 and critical pressure 75 atmospheres. The liquid is obtainable commercially, compressed in steel cylinders. Solid nitrous oxide occurs in colourless crystals, which melt at - 102.3. It was first observed by Davy that when nitrous oxide is inhaled in small amount, it gives rise to excitement often accompanied by hysterical laughter ; hence the name laughing gas. On continued inhalation it induces insensibility, and is therefore used as an anaes- thetic, especially in dental operations. Nitrous oxide is fairly soluble in waten One c.c. of water at 5 absorbs 1.048 c.c., at 15 0.7377 c.c., at 25 0.5443 c.c. of the gas (Geffcken). Nitrous oxide is decomposed into its elements at a much lower temperature than is nitric oxide ; at 700 the speed of the reaction 234 A TEXT-BOOK OF INORGANIC CHEMISTRY can readily be measured. Like nitric oxide, nitrous oxide is an endothermic compound, about 20,000 calories being absorbed in the formation of 30 grams of the gas from its elements, and its apparent stability at ordinary temperatures is due to its slow rate of decomposition. The direct formation of nitrous oxide from its elements has not hitherto been observed. Chemical Properties Owing to its relatively slight stability, nitrous oxide is a supporter of combustion, a glowing splinter plunged into it bursts into flame, and brightly burning sulphur and phos- phorus (though not feebly burning sulphur) continue to burn in it. It can at once be distinguished from oxygen, however, by the fact that no red fumes are produced when it is mixed with nitric oxide. It does not combine directly with oxygen even at a red heat. Composition When a confined volume of nitrous oxide is heated with metallic sodium, sodium oxide and nitrogen are formed, and the volume of the nitrogen is equal to that of the nitrous oxide taken. It follows that the molecule of nitrous oxide contains a molecule or two atoms of nitrogen, and its formula must be NgO*, where x is a whole number. As, however, its density is 22 and mole- cular weight 44, it contains 44- 28= 16 parts or i atom of oxygen, and its formula is therefore N 2 O. The same conclusion is also reached by exploding nitrous oxide with an equal volume of hydrogen. Water is formed, and a volume of nitrogen equal to that of the original gas liberated : N 2 O + H 2 = N 2 + H 2 O (2vols.) (2vols.) (2vols.) (liquid). Hyponitrous Acid Preparation The alkali salts of this acid, the hyponitrites, are formed when nascent hydrogen (obtained by the action of water on sodium amalgam) acts on an alkali nitrate, or, better, on a nitrite : The acid itself cannot be obtained by addition of sulphuric acid to a hyponi- trite in aqueous solution, as it immediately decomposes into nitrous oxide and water :' but is obtained by adding silver hyponitrite to a solution of hydrogen chloride in ether, filtering and evaporating off the ether in a desiccator : Properties Hyponitrous acid occurs in colourless crystalline leaflets, and is highly explosive. It is readily soluble in water, forming a moderately stable solution ; but decomposes into nitrous oxide and water on warming. OXIDES AND OXYACIDS OF NITROGEN 235 Freezing-point determinations in aqueous solution prove that the acid has the double formula, H 2 N 2 O 2 . Nitrous oxide does not form hyponitrous acid with water, so that the reaction is not reversible. Nitramide, NH 2 NO 2 , an isomer of hyponitrous acid, is obtained by the action of cold sulphuric acid on potassium nitrocarbamate : N0 2 -NK-COOK + H 2 S0 4 ->NH 2 N0 2 +K 2 S0 4 + C0 2 . It forms colourless crystals, which melt at 72-75 and are readily soluble in water ; the solution is strongly acid. Under the influence of catalytic agents it splits up into nitrous oxide and water. Its graphic formula has not been definitely settled. COMPOUNDS OF OXYGEN, NITROGEN AND THE HALOGENS Nitrosyl Chloride, NOC1 Preparation (i) By direct combination of nitric oxide and chlorine : (2) A mixture of i part nitric acid (density 1.2) and and 3 parts hydrochloric acid (density 1.12) is known as aqua regia, because it possesses the property of dissolving the so-called "noble" metals gold and platinum. It owes its activity to free chlorine, but also contains nitrosyl chloride : From the Mixture nitrosyl chloride can be separated by fractional distillation. Properties Nitrosyl chloride at ordinary temperature is an orange-yellow gas, which is readily condensed to a reddish liquid which boils at -5.6. It begins to decompose into its components only when heated to 700. It is readily hydrolyzed by water, with formation of nitrous and hydrochloric acids : NOC1 + HOH->HNO 2 + HC1. Nitrosyl Bromide, NOBr, formed by direct combination of nitric oxide and bromine, is a brownish-black liquid at low temperatures. It boils at -3, and at room temperature the vapour is already partially decomposed into nitric oxide and bromine. Nitrosyl Fluoride, NOF, is formed when nitrosyl chloride is passed over silver fluoride heated to 200-250 in a platinum tube : NOC1 + AgF-^NOF + AgCl. Nitrosyl fluoride is a colourless gas, which can be condensed to a colourless liquid boiling at -56. It is immediately hydrolyzed by water with formation of nitrous an 1 hydrofluoric acids. A Nitryl Fluoride, NO 2 F, has also been described. 236 A TEXT-BOOK OF INORGANIC CHEMISTRY Utilization of Atmospheric Nitrogen 1 As nitrogen is one of the elements essential for the growth of plants and is con- tinually being removed in the course of agricultural operations, it becomes a matter of the utmost importance to discover means of keeping up the supply. Ammonium salts and niUates, especially sodium nitrate (Chili saltpetre), have long been and still are in use as nitrogenous manures, but quite recently methods have been devised for utilizing the almost unlimited stores of nitrogen in the atmosphere. The most promising of these methods is based upon the combination of nitrogen and oxygen in the electric arc to nitric oxide. According to the Birkeland-Eyde method, used in Norway, the electric arc is drawn out by means of powerful electro-magnets into a disk of flame several feet in diameter. Air is passed through the arc thus formed, and at the high temperature (about 3000) a certain proportion of nitric oxide, about i per cent of the gas leaving the tube, is formed (p. 231). The gases are cooled and passed into the oxidation chambers, in which the nitric oxide is completely converted to peroxide (p. 232), and are then passed up towers filled with quartz or coke, over which water trickles, the nitrogen peroxide being thereby almost completely transformed to nitric acid : The nitric oxide combines with more oxygen and is again passed into the absorption tower, till finally absorption is practically complete. The nitric acid is treated with calcium carbonate or hydroxide, forming basic calcium nitrate, which is sold as a manure, and is also used in the manufacture of nitric acid. Another method of utilizing atmospheric nitrogen depends upon the fact that when impure calcium carbide is heated in nitrogen, the latter is absorbed with formation of a mixture of calcium cyanamide and carbon : The reaction is exothermic in the direction of the upper arrow and is reversible. The crude mixture of calcium cyanamide and carbon is known as nitrolim. Its importance depends on the fact that the cyanamide reacts with water to give calcium carbonate and ammonia : 1 An excellent account of the utilization of atmospheric nitrogen, with illustra- tions, is given by Crossley, Pliarm. Journal, 1910, 84, 329. OXIDES AND OXYACIDS OF NITROGEN 237 and hence it may be used directly as a manure. In the soil the decomposition of nitrolim is doubtless more complicated. Within the last few years a number of factories have been established for the manufacture of nitrolim, the largest being in Norway, where water power is cheap. The carbide is made by heating calcium oxide and anthracite coal in the electric furnace (P- 3 1 5)>^ ie nitrogen by the liquefaction and subsequent fractionation of air according to the Linde process. Nitrolim is now used for the commercial production of cyanides ; for this purpose, as it contains the requisite substances, it is only necessary to heat it with a flux ' : CaNCN+C->Ca(CN) 2 . Molar and Normal Solutions It has been pointed out (Chapter XV.) that solutions of certain substances of the same molar concentration have the same osmotic pressure, and therefore such solutions are directly comparable for certain purposes. A molar solution contains one mol of solute in a litre of solution. The neutra- lization of hydrochloric acid by sodium hydroxide is represented by the equation HC1 + NaOH->NaCl + H 2 O, and in this case one mol of hydrochloric acid maybe regarded as equivalent to one mol of sodium hydroxide. The neutralization of sulphuric acid by sodium hydroxide is represented by the equation and in this case one mol of sulphuric acid is equivalent to two mols of sodium hydroxide. The term normal solution is applied to a solution which contains the equivalent of one hydrogen in grams per litre. Hence a normal solution of hydrochloric acid contains 36.5 grams of the acid per litre, whereas a normal solution of sulphuric acid contains half the molecular weight, that is, % s or 49 grams of acid per litre. Normal solutions of acids are therefore equivalent in strength, as they contain the same concentration of replaceable hydrogen. From the definition it is evident also how normal solutions of bases, and also of oxidizing and reducing agents, can be prepared. 1 A flux is a readily fusible material which is added to infusible substances in order to bring them more intimately into contact at high temperatures, chemical combination being thereby facilitated. CHAPTER XIX PHOSPHORUS Symbol, P. Atomic weight 31.0. Molecular weight = 124. /Chemical Relations The compounds of phosphorus are in ^-^ many cases of the same type as those of nitrogen, due to the fact that the former element, like the latter, is mainly trivalent and pentavalent. The best-known hydrogen compound of phosphorus is PH 3 , corresponding with NH 3 , and phosphorus forms two principal oxides, P 2 O 3 (or P 4 O 6 ) and P 2 O 6 , which, like the corresponding oxides of nitrogen, are acidic. From the first-mentioned oxide is derived phosphorous acid, H 3 PO 3 , and the corresponding salts, the phos- phites ; from P 2 O 5 is derived phosphoric acid, H 3 PO 4 , and the phosphates. The formulae just given show that the oxyacids and oxysalts of phosphorus are not quite of the same type as nitrous and nitric acids and the corresponding salts. History Phosphorus was discovered in 1669 by Brand of Hamburg, who obtained it by evaporating urine (which contains phosphates) to dryness and distilling the residue. The process was kept secret for some years, but was independently discovered by Kunkel and by Boyle (1680) ; the latter obtaining phosphorus by distilling urine with sand. The method now in use, which depends upon the use of bone-ash (impure calcium phosphate) was discovered by Scheele (1770). The name "phosphorus" was at that time used to designate any substance which glowed in the 'dark, and the element now under consideration was known as Brand's phosphorus, or phosphorus mirabilis. Occurrence Phosphorus is never found free in nature, but in the form of compounds, chiefly phosphates, is very widely distributed. Calcium phosphate, Ca 3 (PO 4 ) 2 ,/^^^3CaSiO 3 + SCO + 2P. The sand combines with the calcium to form calcium silicate, CaSiO 3 , a readily fusible compound, which can be drawn off at intervals through the opening b. The phosphorus passes off through the pipe a, and is condensed under water in the usual way. As the reaction proceeds, fresh charges of the reacting substances can be added, and 2 4 o A TEXT-BOOK OF INORGANIC CHEMISTRY the process is thus made continuous. It may be pointed out that the process is not one of electrolysis ; the function of the electric arc is simply to give a very high temperature. The crude phosphorus obtained by the above methods is very impure, being often red or almost black in colour. By melting and stirring under water containing potassium bichromate and sulphuric acid, part of the impurities are oxidized away and others rise to the surface of the water. Mechanical impurities are best removed by melting under water and forcing through chamois leather. The purified phosphorus is then cast into sticks. In order to prevent oxidation it is kept under water. Physical Properties The modification of phosphorus obtained as above is a transparent, almost colourless (slightly yellowish) wax-like solid, of density 1.82 at 20. At o it is brittle, at room temperature it can be cut like wax, under water it melts about 44, forming a liquid which can be very considerably super- cooled without solidifying (p. 70). In absence of air it boils at 290. Even at room temperature it is distinctly volatile, as is evident from the characteristic smell. Up to 1000 the vapour density of phosphorus is 62, correspond- ing with the formula P 4 ; in the neighbourhood of 1 500, however, it is considerably less, indicating partial dissociation, probably into P 2 molecules. Yellow phosphorus is insoluble in water, slightly soluble in alcohol, fairly soluble in ether and benzene, readily soluble in carbon di sul- phide. On evaporating the solution in the latter solvent, out of contact with air, phosphorus is obtained in well-formed crystals (rhombic dodecahedra). As shown both by the boiling-point and freezing-point methods (p. 196), phosphorus in solution has also the formula P 4 . Chemical Properties When exposed to the air, white phosphorus is slowly oxidized, chiefly to phosphorous oxide, P 4 O (; , which forms white fumes, the process being accompanied by a garlic odour and a slight luminosity visible in the dark. During the slow oxidation of phosphorus, ozone (p. 134), hydrogen peroxide, and other products are formed. Dry phosphorus does not combine with perfectly dry oxygen, and, what is still more remarkable, there is a definite pressure (which depends on the temperature) above which combination with moist oxygen does not occur, although it takes place readily when the oxygen pressure is reduced. Phosphorus catches fire in the air when heated to about 35, but PHOSPHORUS 241 in the finely-divided form it catches fire at room temperature. This can be strikingly shown by soaking a piece of filters-paper with a solution of phosphorus in carbon disulphide and exposing to the air; when the solvent has evaporated the phosphorus bursts into flame. Yellow phosphorus is a very poisonous substance, o.i of a gram being usi ally fatal for an adult. Exposure to the vapour of phos- phorus ultimately causes necrosis of the jawbones and teeth and other injurious effects, Red Phosphorus. Preparation This element also exists in a second modification, which can be obtained by heating yellow phosphorus in absence of air at 240-250. On the commercial scale, yellow phosphorus is heated in a closed iron vessel for some time at 240 and then at a rather higher temperature in order to complete the change. The product is then ground under water, boiled with sodium hydroxide to remove traces of yellow phosphorus, then washed with water and dried. The change of yellow to red phosphorus can also be brought about in solution ; for example, by boiling a solution of yellow phosphorus in phosphorus tribromide. The change is also produced under the influence of light, and the reddish appearance which sticks of ordinary phosphorus sometimes present is doubtless due to a coat- ing of red phosphorus. The reaction yellow phosphorus->red phosphorus is strongly exo- thermic, about 27,300 cal. being given out in the transformation of 31 grams, so that red phosphorus is the stable form at high temperatures, and, judgingfrom the relativesolubilities,alsoat ordinary temperatures. 1 Physical Properties Red or "amorphous" phosphorus is a dark to violet-red powder of density 2.15. It is not luminous, has no taste or srnell, is insoluble in carbon disulphide, is not poisonous, and does not ignite in the air below 200 in all these respects pre- senting the most striking contrast to yellow phosphorus. At fairly high tern peratu res red phosphorus is vaporized, and the vapour condenses as yellow phosphorus. In a vacuum at 100, however, red phosphorus can be sublimed unchanged. Red phosphorus is sometimes termed amorphous phosphorus, but it consists, at least in part, of minute crystals. It is probably not a uniform substance, and at present its exact nature is not understood. Cohen (1910) regards red phosphorus as a "solid solution" (p. 198) of yellow phosphorus in Hittorf's phosphorus. 1 The sol ibility of the unstable modification, like its vapour pressure (p. 69), is always greater than that of the stable modification. 16 242 A TEXT-BOOK OF INORGANIC CHEMISTRY Although chemically less active than the yellow modification, red phosphorus is readily attacked by the halogens and by nitric acid (see below). Hittorf' s Phosphorus Hittorf obtained a definite modifica- tion of this element by heating red phosphorus with fused metallic lead in a sealed tube. On cooling, the phosphorus separates in yellowish- red transparent crystals of density 2.34. This modification is known as metallic phosphorus or Hittorf's phosphorus. Matches The chief commercial use of phosphorus is in the preparation of matches. Ordinary matches are made by dipping the heads first into melted paraffin (or fused sulphur) and then into a paste consisting of yellow phosphorus, an oxidizing agent (potassium chlorate, red lead, or manganese dioxide), and glue. When rubbed on a rough surface (e.g. sand-paper) the phosphorus catches fire, and by the agency of the sulphur or paraffin the wood then becomes ignited. These matches are very poisonous on account of the presence of phosphorus, and for the same reason are very injurious to the health of the workmen. These disadvantages are not present in the so-called " safety " matches, which are now very largely used. The heads are composed of a mixture of potassium chlorate or dichro- mate, antimony trisulphide, a little powdered glass to give greater friction, and glue, and they are ignited by rubbing on a prepared surface containing red phosphorus, antimony trisulphide, and glue. At the high temperature produced by friction the phosphorus burns in the oxygen of the oxidizing agents, thus setting the match on fire. COMPOUNDS OF PHOSPHORUS AND HYDROGEN At least three compounds of phosphorus and hydrogen are known, namely, gaseous hydrogen phosphide, PH 3 , liquid hydrogen phos- phide, P 2 H 4 , and a solid phosphide (P 4 H 2 ) n , probably Pi 2 H 6 . None of them can be obtained by direct combination of the elements. Gaseous Hydrogen Phosphide or Phosphine, PH 3 . Preparation (i) By boiling phosphorus with sodium hydroxide in absence of air: The second product of the reaction, NaH 2 PO 2 , is sodium hypo- phosphite. The arrangement of the apparatus is shown in Fig. 53. Phosphorus and a concentrated solution of sodium hydroxide are placed in the flask, from which the air is then removed by a stream PHOSPHORUS 243 of coal-;jas passed through the tube. Heat is then applied, and as the gas escapes into the air each bubble immediately catches fire, forming vortex rings of metaphosphoric acid. As a matter of fact, the spontaneous inflammability is due to the presence of traces of P 2 H 4 , and if the latter compound is removed by passing through FIG. 53. alcohol or hydrochloric acid, the gas no longer catches fire at the ordinary temperature. (2) By the action of water or of dilute hydrochloric acid on calcium phosphide: Owing to secondary reactions traces of the other hydrides are also obtained, and the gas is spontaneously inflammable. (3) Pui-e phosphine is obtained by the action of water or potassium hydroxide on phosphonium iodide, PH 4 I (see below): PH 4 I + KOH->PH 3 + KI + H 2 O. 244 A TEXT-BOOK OF INORGANIC CHEMISTRY Physical Properties Phosphine is a colourless gas with a dis- agreeable odour, recalling that of decaying fish. The liquefied gas boils at -86. It is very slightly soluble in water, more soluble in alcohol. It readily decomposes into its elements on heating. Chemical Properties When heated to about 100 in the air phosphine burns to water and metaphosphoric acid : PH 3 + 2O 2 =HPO 3 + H 2 O. It does not combine with oxygen at atmospheric pressure, but com- bination takes place explosively when the pressure is reduced. In this respect it resembles phosphorus, which combines with oxygen only below a certain limit of pressure, which in both cases depends on the temperature and the proportion of moisture present. The aqueous solution of phosphine does not affect litmus, so that it is much less basic than ammonia. It has, however, a slightly basic character, as shown by its capacity to combine directly with the halogen acids, HX, forming so-called phosphonium compounds, PH 4 X (analogous to ammonium halogen compounds, NH 4 X), which are dealt with in the next section. Unlike ammonia, phosphine does not form salts with oxygen acids. Phosphonium Compounds (a) Phosphonium Iodide, the most stable of these compounds, is obtained by direct combination of its components, or, better, by adding iodine to phosphorus dissolved in carbon disulphide, distilling off the solvent, adding water cautiously, then subliming the phosphonium iodide: 5 I + 9 P + i6H 2 0-^ S PH 4 I + 4 H 3 P0 4 . Phosphonium iodine occurs in colourless, well-formed crystals. (b] Phosphonium bromide is obtained in colourless crystals by direct combination of phosphine and hydrogen bromide in a vessel immersed in a freezing-mixture. It readily dissociates into its components. (<;) The corresponding chloride, PH 4 C1, is also obtained in crystalline form by direct combination of its components in a freezing-mixture, but is still less stable than the bromide, and dissociation can only be avoided by keeping the tempera- ture low, or by subjecting it to considerable pressure. All three compounds are decomposed by water, forming the halogen acid and phosphine, the latter passing off as gas. Liquid Hydrogen Phosphide, P 2 H 4 , is obtained as a colourless liquid when the gas produced by the action of water on calcium phosphide, and which consists chiefly of the hydride PH 3 , is passed through a u-tube immersed in a freezing- mixture. The liquid boils at 57 to 58 ; on exposure to air it immediately ignites and burns with a brilliant flame. On exposure to light in absence of air it rapidly decomposes into the gaseous and solid hydrides : S P 2 H 4 ->6PH 3 +P 4 H 2 . PHOSPHORUS 245 Solid Hydrogen Phosphide, (P 4 H 2 ) 3 or P 12 H 6 , obtained as just described, is a yellow flocculent powder, insoluble in water and alcohol. It is soluble in melted yellow phosphorus, and from the lowering in the freezing-point of the solvent (p. 196) its formula has been proved to be Pi 2 H 6 . COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS Phosphorus combines with all the halogens, forming the following compounds PF 3 (gas). PC1 3 (liquid). PBr 3 (liquid). P 2 I 4 (solid). PF 5 (gas). PC1 5 (solid). PBr 6 (solid). PI 3 (solid). All of them can be obtained by direct combination of the elements, and all are decomposed by water, giving rise to a mixture of halogen acid and an oxyacid of phosphorus. This behaviour has already been utilized in the preparation of hydrobromic and hydriodic acids (pp. 156 and 161). Besides the halogen compounds themselves, certain oxyhalogen compounds are known. The most important is phosphorus oxy- chloride or phosphoryl chloride, POC1 3 . Phosphorus trifluoride, PF 3 , is obtained by dropping arsenic trifluoride into phosphorus trichloride : AsF 3 + PCl 3 ->AsCl 3 + PF 3 . It is a colourless gas, which at room temperature is decomposed only very slo\vly by water. Phosphorus pentafluoride, PF 5 , is obtained by direct com- bination of the trifluoride and fluorine (Moissan) or by the action of arsenic trifluoride on phosphorus pentachloride (Thorpe) : [t is a colourless gas, which fumes in moist air and is decomposed >y water to form ultimately hydrofluoric and phosphoric acids : Unlike the corresponding chlorine and bromine compounds (^.-z/.), it can be raised to a high temperature without dissociating into the trifluoride and fluorine. Phosphorus Trichloride, PC1 3 , is obtained by passing dry chlorine over red or yellow phosphorus gently heated in a retort ; 246 A TEXT-BOOK OF INORGANIC CHEMISTRY the trichloride distils off, and is collected in a cooled receiver. In order to remove pentachloride, a little phosphorus is added and the liquid redistilled. Properties Phosphorus trichloride is a colourless liquid, which boils at 76 ; its density is 1.58 at 20. It fumes in moist air, and is at once decomposed by water with formation of phosphorous and hydrochloric acids : Phosphorus Pentachloride, PC1 5 , is best prepared by passing dry chlorine into a large flask containing the trichloride. The flask is kept cool, and when the contents become quite dry it is known that the change is complete. Properties Phosphorus pentachloride forms light-yellow almost colourless crystals with a pungent odour. When heated it sublimes rapidly above 140 without melting, forming a vapour which is almost colourless at low temperatures, but becomes yellower as the temperature is raised. Parallel with this change in colour goes a diminution in density, both phenomena being accounted for by increasing dissociation into the trichloride and chlorine according to the equation as already described (p. 169). That this is the true explanation of the phenomena is further shown by the fact that the products of dissocia- tion can be separated by diffusion. From the densities it has been calculated that at 200 under atmospheric pressure the pentachloride is dissociated to the extent of 48.5 per cent., at 250" to 80 per cent., and at 300 dissociation is practically complete. The effect of excess of chlorine or of the trichloride in diminishing the degree of dissocia- tion has already been referred to (p. 171). With a small quantity of water the pentachloride is converted into phosphoryl chloride and hydrochloric acid : PC1 6 + H 2 O-POC1 3 With excess of water phosphoric acid is formed: (or H 3 PO 4 ). Phosphorus pentachloride is largely used, both in inorganic and organic chemistry, for replacing hydroxyl groups by chlorine. With PHOSPHORUS 247 sulphuric acid, for instance, which can be written SO 2 (OH) 2 , the following reaction takes place /OH /Cl SO 2 < + PC1 5 ->SO 2 +POCL + HC1. \OH \OH The point above referred to, that phosphorus pentachloride sub- limes before the melting-point is reached, is simply a question of the magnitude of the vapour pressure at the melting-point. It happens that before the melting-point is reached the vapour pressure exceeds i atmosphere, and therefore volatilization becomes very rapid. When heated in a closed vessel the pentachloride melts under the pressure of its own vapour at 148. Phosphorus Oxy chloride, POC1 3 , is obtained, as mentioned above, by the action of a small amount of water on the pentachloride,, It is a colourless, fuming liquid, which boils at 107, and when solid melts at 1.5. On treatment with excess of water it yields phosphoric and hydrochloric acids, as stated above. Phosphorus Tribromide, PBr 3 , is a colourless, fuming liquid, which boils at 175; the pentabromide\s a yellow solid, which dissociates into the tribromide and bromine even more readily than does the pentachloride. Both bromides are obtained directly from their elements, and they behave with water like the corre- sponding chlorides. Phosphorus Diiodide, P 2 l4, is obtained by adding the calculated amount of iodine to phosphorus dissolved in carbon disulphide and distilling off the solvent. It forms orange-yellow crystals, which melt at 110 to a reddish liquid. Phosphorus Triiodide, PI 3 , is obtained by the method described for the di- ioclide by using more iodine. It occurs in dark-red crystals, which melt at 55, and are decomposed by water with formation of hydriodic and phosphorous acids. OXIDES AND OXYACIDS OF PHOSPHORUS Three oxides of phosphorus are definitely known Phosphorous oxide (phosphorous anhydride), P 4 O 6 Phosphorus tetroxide, ^2^1 Phosphorus pentoxide (phosphoric anhydride), P 2 O 5 (P 4 O 10 ) A number of oxyacids of phosphorus are known, most of which are derived from the oxides just mentioned Hypophosphorousacid,H 3 PO 2 ,orPH 2 O(OH), ) no corresponding Hypophosphoric acid, H 2 PO 3 \ oxides 248 A TEXT-BOOK OF INORGANIC CHEMISTRY Phosphorous acid, H 3 PO 3 , or P(OH) 3 ? > , p Metaphosphorous acid, HPO 2 f Orthophosphoric acid, H 3 PO 4 , or PO(OH) 3 ) Pyrophosphoric acid, H 4 P 9 O 7 , or P 2 O 3 (OH) 4 [ from P 2 O 5 Metaphosphoric acid, HPO 3 , or PO 2 (OH) ) a .Phosphorous Oxide, P 4 O 6 , is formed, mixed with the pentoxide, when phosphorus is burned in a glass tube in a icurrent of dry air. In order to separate the oxides the products of combustion are passed into the inner tube of a Liebig's condenser, the outer tube of which contains water at 60. The inner tube contains a plug of glass wool, which retains the pentoxide ; whereas the lower oxide passes on, and is condensed in a cooled U^tube. Properties Phosphorous oxide occurs in colourless crystals, which melt at 22.5 to a colourless liquid ; the latter boils at 173.1. The oxide is very poisonous. It combines slowly with cold water to form phosphorous acid : but when treated with hot water a vigorous action takes place, red phosphorus, hydrogen phosphide and orthophosphoric acid being the chief products. When heated to 50-60 in the air it inflames and burns to the pent- oxide. When heated in a sealed tube to 440 it decomposes rapidly into the tetroxide and phosphorus : 2P 4 6 ->3P 2 4 +2P. Phosphorus Tetroxide, P 2 O 4 , is obtained by heating phos- phorous oxide in a sealed tube as just described. It occurs in colour- less, lustrous crystals, which react with water to form a mixture of phosphorous and phosphoric acids (cf. N 2 O 4 , p. 229) : P 2 4 + 3 H 2 0->H 3 P0 3 +H 3 P0 4 . Phosphorus Pentoxide, P 2 O 5 , is obtained by burning phos- phorus in excess of dry air or oxygen. For experimental purposes it is conveniently obtained by burning phosphorus under a bell jar; after a time the white fumes settle as a soft white powder. It may be obtained quite pure by distilling in a current of dry oxygen. Properties Phosphorus pentoxide is a colourless, snowlike amor- phous substance, which when pure is odourless. At a tempera- ture a little below white-heat it is converted into vapour, which condenses in the crystalline form. PHOSPHORUS 249 It has a very great affinity for water, and is the most efficient drying agent known ; its application for the complete removal of moisture from gases has already been referred to. The first action of water is to form metaphosphoric acid : P 2 O 5 +H 2 O->2HP0 3 . On prolonged standing, much more rapidly on heating, another mole- cule of water is taken up and orthophosphoric acid, H 3 PO 4 , is obtained. An illustration of the action of phosphorus pentoxide in abstracting the elements of water from other compounds has already been met with in the preparation of nitrogen pentoxide : Vapour density determinations have shown that the formula for this oxide at 1400 is P 4 O 10 , but in writing equations representing its behaviour it is more convenient to use the simpler formula. Hypophosphorous Acid, HPH 2 O 2 The acid is most readily obtained by the action of sulphuric acid on barium hypophosphite in aqueous solution : The barium sulphate is removed by filtration, and the solution con- centrated, first in porcelain and finally in platinum vessels, the tempe- rature not being allowed to rise above 105. On cooling the acid separates n colourless crystals. Barium hypophosphite is obtained in solution by boiling phos- phorus with barium hydroxide (cf. p. 242) : Properties Hypophosphorous acid occurs in colourless leaflets which melt at 17.4. On heating it yields phosphine and orthophos- phoric acid : 2H 3 P0 2 ->H 3 P0 4 Owing to ts tendency to pass into phosphorous and phosphoric acids it is a strong reducing agent. Gold, silver, and mercury are precipi- tated from solutions of their salts : Ag 2 SO 4 +H 3 PO 2 +H 2 O->2Ag + H 3 PO 3 +H 2 SO 4 . From solutions of copper salts, copper hydride, Cu 2 H 2 , is precipitated 2 5 o A TEXT-BOOK OF INORGANIC CHEMISTRY on warming ; a reaction which is characteristic for hypophosphorous acid and its salts : ->6H 3 PO 3 + 2NaHSO 4 Hypophosphorous acid is a monobasic acid, that is, only one of the hydrogen atoms can be displaced by metals. The hypophosphites decompose into phosphine and the corresponding phosphate on heating. They are all soluble in water ; a number of them are used in medicine. The monobasic character of hypophosphorous acid may be expressed by writing its formula thus : /H O = P-OH \H the assumption being made that the acidic properties pertain only to the hydrogen attached to oxygen, and not at all to those attached directly to phosphorus. Phosphorous Acid, H 3 PO 3 Preparation (i) By dissolving phosphorous oxide in cold water (p. 248) : P 4 O 6 + 6H 2 O->4H 3 PO 3 . (2) By the gradual addition of water to phosphorus trichloride, with simultaneous cooling : The product is then evaporated, the temperature being finally raised to 1 80 to drive off all the hydrogen chloride; on cooling phosphorous acid crystallizes out. Properties Phosphorous acid occurs in colourless crystals, which melt at 71.1. On heating it decomposes into orthophosphoric acid and phosphine : 4H 3 PO 3 ->3H 3 PO 4 It is a strong reducing agent, precipitating silver from solutions of silver salts and copper (not copper hydride, difference from hypo- phosphorous acid) from solutions of copper salts. Although it contains three hydrogen atoms, only two of them are replaceable by metals, and it is therefore a dibasic acid. In accord- ance with this, it forms two series of salts. The two phosphites of sodium, for instance, have the formulas NaH 2 PO 3 and Na 2 HPO 3 . PHOSPHORUS 251 The behaviour of phosphorous acid may be expressed, in accordance with the principle discussed under hypophosphorous acid, by the graphic formula /H O = P^-OH \OH Metaphosphorous Acid, HPO 2 , is obtained in lustrous crystals by interaction of phosphine and oxygen under reduced pressure: PH 3 + O a ->HPO 2 + H 2 . It dissolves in water tp form phosphorous acid : HPO 2 + H 2 O-H 3 PO 3 . Hypophosphoric Acid, H 2 PO 3 , is formed, along with phos- phorous and phosphoric acids, by the slow oxidation of phosphorus in moist air. The pure acid occurs in small colourless crystals, melting at 53. It is a less powerful reducing agent than phosphorous acid. From the fact that it forms two series of salts, and for other reasons, it is regarded as a dibasic acid. Phosphoric Acid (Orthophosphoric Acid), H 3 PO 4 Pre- paration^} On the large scale, by the action of sulphuric acid on calcium phosphate (bone-ash or natural phosphate, p. 239) : The calcium sulphate is removed by filtration or by decantation and the phosphoric acid concentrated by evaporation. (2) A purer acid is obtained by boiling yellow or red phosphorus with nitric acid of density 1.2 : 6P+ioHN0 3 + 4H 2 O->6H 3 PO 4 +ioNO. Properties Phosphoric acid is usually met with as a colourless, syrupy liquid, which, however, still contains water. When the solu- tion is concentrated at 150 and then allowed to cool, the pure acid separates as colourless crystals, melting at 40.7. Phosphoric acid is a tribasic acid, forming three series of salts, primary or dihydrogen phosphates, e.g. NaH 2 PO 4 , secondary or monohydi-ogen phosphates, e.g. Na 2 HPO 4 , and tertiary or normal phosphates, e.g. Na 3 PO 4 . Salts are also formed from phosphoric acid by displacing the hydrogen by different atoms or groups, thus 252 A TEXT-BOOK OF INORGANIC CHEMISTRY sodium ammonium phosphate, NaNH 4 HPO 4 , is well known as mien- cosmic salt, and ammonium magnesium phosphate, Mg(NH 4 )PO 4 , is met with in analysis (p. 448). The phosphates of the heavy meta]s are mostly normal salts ; thus the only known silver phosphate is Ag 3 P0 4 . Sodium dihydrogen phosphate is slightly acid in aqueous solu- tion. The monohydrogen phosphate, Na 2 HPO 4 , is, however, slightly alkaline in solution, owing to its being partially decomposed by water : Na 2 HPO 4 + H 2 O^NaH 2 PO 4 + NaOH. The alkaline reaction is due to the sodium hydroxide formed. The normal salt is still more strongly split up by water : and, in fact, can only be isolated from aqueous solution when a large excess of sodium hydroxide is present. The hydroxide acts, according to the law of mass action, by displacing the equilibrium in the direction of the lower arrow. This decomposing action of water on salts is termed hydrolysis and will be fully considered at a later stage (p. 267). The normal phosphates of the alkalis are not affected by heat, but the secondary lose water and form pyrophosphates : O, and the primary under similar conditions give metaphosphates : NaH 2 PO 4 ^NaPO 3 + H 2 0. Both types of reaction are reversible. Pyrophosphoric Acid, H 4 P 2 O 7 Preparation (i) By heating orthophosphoric acid for some time at 210 to 220. (2) In a purer form by decomposing the difficultly soluble lead pyrophosphate (prepared by double decomposition from sodium pyro phosphate) with hydrogen sulphide : Pb 2 P 2 O 7 + 2 H 2 S->H 4 P 2 O 7 + 2PbSj,. Properties The pure acid occurs in minute microscopic needles which melt about 60. The aqueous solution is fairly stable, but on heating or on the addition of traces of other acids (catalytic action). it takes up water and forms orthophosphoric acid. Pyrophosphoric acid is a tetrabasic acid, but, curiously enough., forms only two series of salts, of the types Na 4 P 2 O 7 and Na 2 H 2 P 2 O 7 . PHOSPHORUS 253 As already mentioned, the pyrophosphates are obtained by heating the secondary phosphates. The aqueous solutions are not affected even on boiling, but at still higher temperatures the salts take up water and form orthophosphates. Metaph osphoric Acid, HPO 3 Preparation (i) By allow- ing phosphorus pentoxide to deliquesce in moist air, or by dissolving the pentoxide in a small amount of water : H 2 O->2HPO 3 . (2) By heating orthophosphoric acid in a gold crucible at about 400 till a molecule of water is driven off: H 3 PO 4 ->HPO 3 +H 2 O. Properties Metaphosphoric acid is a semi-transparent glassy sub- stance, which is usually sold in the form of sticks, and is known as "glacial phosphoric acid." It is readily soluble in water, and the solution is fairly stable at the ordinary temperature, but on boiling, or on the addition of a little mineral acid, it rapidly takes up water and forms orthophosphoric acid. This change can be detected by taking advantage of the fact that metaphosphoric acid coagulates a solution of albumen (white of egg) whereas the ortho acid does not [see below). Metaphosphoric acid is a fairly strong monobasic acid. The corre- sponding salts, the metaphosphates, can be prepared directly from the acid or by heating the dihydrogen phosphate (p. 252). Many metaphosphates, however, are complex in constitution, being most readily derived from acids of the general formula (HPO 3 ) M , where n is a whole number varying from I to 6 and probably higher. The compound Li. 2 K 4 (PO 3 ) 6 ,4H 2 O, for instance, is obviously derived from the hexabasic acid H 6 (PO 3 ) C ; that is, n = 6. At high temperatures the acid can be volatilized, and its vapour density corresponds with the formula (HPO 3 ) 2 . At ordinary temperatures, however, it is doubt- ess much more complex. Tests The three phosphoric acids can be distinguished by their action on a solution of albumen (white of egg) and on silver nitrate. With silver nitrate, orthophosphoric acid or any soluble orthophos- phate gives a yellow precipitate of silver orthophosphate, Ag 3 PO 4 , pyrophosphoric acid and pyrophosphates a white precipitate of silver pyrophosphate, Ag 4 P 2 O 7 , and metaphosphoric acid a white precipitate of silver metaphosphate, AgPO 3 . Of the three acids, only meta- phosphoric coagulates albumen. Further, orthophosphates give a 254 A TEXT-BOOK OF INORGANIC CHEMISTRY yellow precipitate when ammonium molybdate and excess of nitric acid are added and the solution heated. hospkoricAcid,H s PO 5 , and Pcrphosphoric y4pf-OH -> O = pf-OH + H 2 O. HO/ \OH \OH 1 Schmidlin pnd Massini, Berichte, 1910, 43, 1162. PHOSPHORUS 255 Similarly, from the highest nitrogen hydroxide, N(OH) 5 , which is unknown, we have N(OH) 5 - and from chlorine, the maximum valency of which is 7, is derived per- chloric ac d, thus C1(OH) 7 -> O==C1- It has been suggested that the hydroxyl compound corresponding with the maximum valency of the element should be called the orthocom- pound, thus N(OH) 5 would be orthonitric acid, and Cl(OH) r , ortho- chloric acid. As, however, many of these compounds are unknown, the convention is not generally observed. Thus orthophosphoric acid, PO(OH) 3 , is the first anhydride of the true ortho acid. Exactly the same considerations apply to bases. The behaviour of the hydroxide is determined by the nature of the group E ; when this is a metal, the hydroxide has basic properties. Compounds are often met with derived not from the normal hydroxide but from one of its anhydrides ; thus many bismuth compounds are derived, not from Bi(OH) 3 , but from BiO(OH), which is Bi(OH) 3 - H 2 O. The connexion between acidic oxides and acids has already been fully explained (p. 186). The graphic formulas of the phosphoric acids, which best represent their chemical behaviour, are as follows : /OH OH OH OH OH /OH = PfOH; \/ ; O = P<^ \OH 0=P-0-P=0 ^O Orthophosphoric acid Pyrophosphoric acid Metaphosphoric acid the phosphorus being quinquevalent throughout, and the basicity represented by the number of hydrogen atoms attached to oxygen. The graphic formula of phosphorous acid presents rather more difficulty. As it is readily formed from phosphorus trichloride, PC1 3 , /OH in which the phosphorus is trivalent, the formula P^-OH might be sug- gested. A s, however, only two of the hydrogen atoms can be displaced /H by metals, the alternative formula O = P^-OH is preferred, for the 256 A TEXT-BOOK OF INORGANIC CHEMISTRY reasons already stated, We shall meet with many illustrations of the fact that hydrogen atoms attached to oxygen are more reactive than those attached directly to the central atom. The graphic formulas of the two remaining oxyacids may be repre- sented as follows : /H O = P^-H hypophosphorous acid ; \OH /OH and O = P\ ? hypophosphoric acid. \OH Compounds of Phosphorus and Sulphur We shall find at a later stage that there is considerable analogy in the formulae of compounds containing oxygen and sulphur respectively ; due in part to the fact that both can function as bivalent and quadrivalent elements. This is shown to some extent in the formula? of the sulphides of phosphorus, which are obtained by heating the elements together in different proportions, the following three being definitely known, P 4 S 3 , P 2 S 6 (corresponding with P 2 O 6 ) and P 4 S 7 . It is shown more clearly in the existence of compounds of the formulas PSC1 3 , Na 3 PSO 3 (corresponding with Na 3 PO 4 ) and Na 4 P 2 S 7 (corresponding with Na 4 P 2 O 7 ), which cannot be described here. Phosphorus Pentasulphide, P 2 S 5 , the best-known sulphide of phosphorus, is obtained by heating the components, in the calculated proportions, in an atmosphere of carbon dioxide. The pure compound occurs in light-yellow crystals which melt at 274-276, and is decomposed by water with formation of phosphoric acid and hydrogen sulphide : CHAPTER XX ELECTROLYSIS 'AND ELECTROLYTIC DISSOCIATION "plectrical Conductivity Substances which conduct elec- t-' tricky may be divided into two classes : (i) conductors of the first class, in which the passage of electricity is not accompanied by a chemical change ; (2) conductors of the second class, in which the passage of electricity is accompanied by chemical decomposition. The more important conductors of the first class are the metals and alloys, but a few non-metals, such as graphite (p. 310), also convey the electric current without decomposition. For our present purpose, however, conductors of the second class, the so-called electrolytes, are more interesting. To this class belong aqueous solutions of salts, and of so-called "strong" acids and bases, all of which are good con- ductors. It must be carefully remembered, however, that water itself is not a conductor, nor are the pure salts, acids, or bases themselves under ordinary conditions. We have already seen that liquefied hydrogen chloride does not conduct the electric current nor does it show any acid properties, but its solution in water is a good conductor and also a strong acid. Although salts as a class under ordinary conditions, do not conduct, it should be mentioned that all of them readily conduct electricity in the fused condition, with simultaneous decomposition. Fused salts form practically the only exceptions to the general rule that pure substances (as dis- tinguished from mixtures) are non-conductors. Organic acids, bases, and salts are electrolytes in aqueous solution, but practically all other organic compounds, whether fused or in solution, are non- electrolytes. Finally, it should be mentioned that the property of electrical conductivity is not confined to aqueous solutions, but is also shown in a pronounced manner by solutions in certain other solvents. It has already been pointed out that the decomposition of an electrolyte when an electric current passes is termed electrolysis, and the terms employed in this branch of the subject have already been mentioned. The plates by which the current enters and leaves 17 2 57 258 A TEXT-BOOK OF INORGANIC CHEMISTRY the electrolyte are called electrodes. The electrode from which the positive current enters the electrolyte (that connected to the positive pole of the battery (Fig. 54)) is called the anode, the electrode at which the positive current leaves the electrolyte is the cathode. The further development of the subject is best illustrated by means of an example. When a solution of sodium sulphate, coloured with neutral litmus solution, is placed in a voltameter, and the electrodes connected with a battery, gases are immediately given off at the anode and cathode, and as electrolysis proceeds the solution in the neigh- bourhood of the anode turns red (showing the formation of an acid) and the solution round the cathode turns blue (indicating the formation of an alkali). Further investigation shows that oxygen only is given off at the anode and hydrogen only at the cathode. The important point to notice is that the chemical changes occur only at the electrodes ; there is no apparent change in the bulk of solution between the poles. It could, however, be shown by special experiments that as electrolysis proceeds, the amount of sodium sulphate in the solution regularly diminishes. The simplest way of accounting for these facts is to assume that the dissolved sub- stance is continually moving towards the electrodes, and that chemical changes occur only when the materials are actually in contact with the electrodes. This was the view taken by Faraday, who termed the moving particles ions (that is, travellers) ; the ions moving towards the anode are called anions, those moving towards the negative pole or cathode cations. The next point to be considered is the nature of these carriers in any particular case. As regards the sodium sulphate solution, we assume, for reasons given more fully later, that the cations are sodium atoms charged with positive electricity, while the anions are SO 4 groups charged with negative electricity. When the sodium ions reach the cathode they give up their electrical charges and become ordinary metallic sodium, which at once acts on the water, forming sodium hydroxide and liberating hydrogen according to the equation 2Na + 2H 2 O = 2NaOH + H 2 t (at cathode). Similarly, when SO 4 ions reach the anode, they give up their charges and at once react with the water : = 2H 2 SO 4 + O 2 f (anode). ELECTROLYSIS 2 59 A little consideration shows that this explanation accounts satis- factorily for the phenomena just described. The electrolysis of sodium sulphate is rather complicated by the reaction of the primary products of electrolysis with the solvent water. A simpler case would be the electrolysis of fused silver chloride, AgCl, with an anode of carbon and a cathode of silver. In this case we may assume that the ions are silver and chlorine, and during electrolysis silver is deposited on the cathode and chlorine liberated at the anode. Faraday's Laws We will now consider the relationship between Me amount of chemical action and the quantity of elec- tricity passed through a solution. The amount of chemical action might be estimated by measuring the volume of gas liberated at one of the: poles, or by weighing the metal deposited on an elec- trode. This question was investigated by Faraday, who established a law which bears his name and which may be stated as follows : For the same electrolyte, the amount of chemical action is propor- tional to the quantity of electricity which passes. Further, Faraday measured :he relative quantities of substances liberated from different electrolytes by the same quantity of electricity, and was thus led to the discovery of his so-called second law : The quantities of sub- stances liberated at the electrodes when the same quantity of electricity is passed turough different solutions are proportional to their chemical equivalents. In a previous chapter it has been shown that Atomic weight = chemical equivalent> Valency The second law therefore states that when the same quantity of electricity is passed through solutions of such electrolytes as sodium sulphate, silver nitrate, AgNO 3 , copper sulphate, CuSO 4 , and auric chloride, AuCl 3 , the relative amounts of hydrogen, oxygen, and the metals liberated are as follows : Electrolyte Na 2 SO 4 AgNO 3 CuSO 4 AuCl 3 Electrochemical) R O== i6 A 108 Cu = 6 3 .6. Au= i 97 Equivalent } 2 I 2 3 The above law may also be stated rather differently as follows : The electrochemical equivalents (the proportions of different elements set free by the same quantity of electricity) are proportional to the chemical equivalents. 260 A TEXT-BOOK OF INORGANIC CHEMISTRY To the different methods of determining chemical equivalents already fully discussed (p. 126), the electrical method just described has to be added. It is evident that that quantity of electricity which passes through an electrolyte when the chemical equivalent of an element in grams is liberated at each electrode must be an important quantity in electrochemistry. Since I ampere in i second (i coulomb) liberates 0.00001036 grams of hydrogen, it follows that when the chemical equivalent of hydrogen (or any other element) is liberated 1/0.00001036 = 96,540 coulombs must pass through the electrolyte. Further, 96,540 coulombs will liberate 35.45x0*00001036 = 0.000368 grams of chlorine, 108x0.00001036 = 0.001118 grams of silver, and 31.7x0.00001036 = 0.000328 grams of copper (from the solution of a cupric salt). Theories of Electrical Conductivity We have now to consider the mechanism of the conduction of electricity in an electro- lyte. A very ingenious theory to account for the experimental facts was brought forward by Grotthus in 1805. He assumed that each molecule is built up of positively and negatively charged particles the molecule of sodium chloride, for instance, of a positively charged sodium atom and a negatively charged chlorine atom and that these molecules are irregularly distributed throughout the electrolyte. When, however, electrodes in such a solution are connected with a battery they exert a directive effect on the molecules, the positive part of all the molecules being turned towards the negative electrode and the negative part towards the positive electrode. The sodium atom next the cathode then gives up its charge to the latter and becomes metallic sodium, simultaneously the chlorine half of another molecule gives up its charge to the anode and becomes free chlorine. The chlorine and sodium atoms thus left free in the solution unite with the sodium and chlorine halves of the neighbouring molecules, so that an exchange of partners takes place all along the line. Under the influence of the charged electrodes the molecules then swing round and the process just described is repeated. About 1857, however, Clausius pointed out that the work done in decomposing an electrolyte is entirely expended in overcoming the resistance of the solution, and that therefore no work is done in pulling apart the molecules of the solute, which would necessarily be the case if the theory of Grotthus were valid. To get over this difficulty, Clausius made the very important assumption that under ordinary conditions a small proportion of the molecules of the solute in ELECTROLYSIS 261 an electrolyte are split up into their positive and negative components, During their free moments the positive and negative ions would travel towards the oppositely charged electrodes, and on reaching these would lose their electrical charges, as Grotthus assumed. The theory of Clausius accounts for the fact* that no work is done in splitting up the molecules, as the ions are already free before electrolysis is started. The Theory of Arrhenius. Electrolytic Dissociation The view of Clausius was that at any moment only a minute fraction of the molecules were decomposed into their ions. In 1887 a great step forward was made by Arrhenius, who succeeded in cor- relating the conductance of solutions with certain other of their properties. We have seen in an earlier chapter that aqueous solutions of salts and so-called strong acids and bases have abnormally high osmotic pressures, their freezing-points are much Ipwer, and their boiling- po nts much higher, than those calculated according to van 't HoiPs theory of solution. As the magnitude of the osmotic pressure, ihe extent to which the freezing-point is depressed, etc., depend only on the number and not on the nature of the particles, the above result may be stated in the form that such solutions behave as if they contained more solute particles than correspond with the ordinary formula. In the case of sodium chloride, for instance, the ratio of the observed osmotic pressure and that calculated on the assumption that only NaCl molecules are present, is about 1.7 : I. Arrhenius now showed that there is a close connexion between abnormally high osmotic pressures and electrical conductivity, only those solutions which, according to van '/ Hojfs theory, have abnorm- ally high osmotic pressures, conduct the electric current. He ascribed the high osmotic pressures to a partial dissociation of the solute into charged ions, the ions acting as independent particles as far as their effect on the osmotic pressure and allied properties is concerned. In solutions of sodium chloride, for instance, we have the equilibrium NaCl^Na + Cl. and in normal solution the salt is dissociated to about 70 per cent, into its ions. Similar equilibria occur in other salts of the same type, for example 262 A TEXT-BOOK OF INORGANIC CHEMISTRY Some salts may, however, decompose into more than two ions. For sodium sulphate, two stages of dissociation are possible : (i) Na 2 SO 4 ^Na + NaSO 4 . (2) and similarly for calcium chloride : (i) CaCl 2 ^CaCl + Cl. (2) CaCl^tCa + Cl. In the case of acids the negative ion differs in each case, but all give rise to hydrogen ions : in the case of bases the positive ion differs in each case, but the negative ion is always the OH group : KOH^K + OH (i.) Ca(OH) 2 ^CaOH + OH. (ii.) CaOH^Ca + OH. It will be noticed that in the above equations the total number of positively charged ions is always equal to the number of negatively charged ions. This is necessarily the case, as the solutions are electrically neutral before as well as during electrolysis. As no atom has a smaller charge than hydrogen and other univalent elements. we assume that these elements have unit charge. Further, when sodium sulphate solution is electrolyzed two sodium atoms are dis- charged for every SO 4 group, and as the solution remains electrically neutral, the SO 4 ion must have had two negative charges to cor- respond with the single charges of the two sodium atoms. On the same principle are derived the other equations for ionic equilibria given on the previous page. The application of the theory of Arrhenius leads to the conclusion that in all cases the degree of dissociation increases regularly with dilution, and is only complete when the dilution is infinitely great. The great majority of salts are dissociated to more than 50 per cent. in normal solution, and those salts which split up into two univalent ions are almost completely dissociated in solutions containing I mol ELECTROLYSIS 263 of salt in 10,000 litres of water. When different acids are compared, however, there is found to be a much greater variation in the degree of dissociation. " Strong " acids are those which in ordinary dilutions are considerably ionised ; that is, the concentration of hydrogen ions under those circumstances is fairly high, whereas "weak" acids, such as hypophosphorous acid (p. 249), are only slightly ionised under ordinary conditions, and their hydrogen ion concentration is therefore small. Similar remarks apply to bases. In aqueous solutions of " strong" bases, such as potassium hydroxide, the OH' ion 1 concen- FIG. 54. tration is high ; in solutions of weak bases, such as ammonium hydroxide, the OH' ion concentration is comparatively small. The mechanism of electrical conductivity on the Arrhenius theory will be readily understood from the foregoing. The fundamental assumptio n is that the current is conveyed through the solution by the ions alone, the non-ionised molecules and the solvent playing no part in the process. When connexion is made with the battery the ions, which were previously moving irregularly through the solution, are attracted by the oppositely charged electrodes and move towards them (Fig. 54). When they reach the electrodes they give up their charges, which neutralize a corresponding amount of electricity on the electrodes, and are liberated as the ordinary uncharged substances a Instead of the + and - signs, the positive charge is conveniently indicated by a dot, and the negative charge by a dash, as shown. 264 A TEXT-BOOK OF INORGANIC CHEMISTRY we are familiar with. In some cases the liberated substances appear as such (e.g. silver and chlorine when fused silver chloride is electrolyzed), in other cases they attack the solvent (e.g. Na and SO 4 when sodium sulphate solution is electrolyzed), in still other cases they attack the electrodes. The objection is sometimes raised to the ionisation theory that solutions of sodium chloride, for instance, cannot contain free sodium and free chlorine, because these would at once attack the solvent. The theory, however, does not postulate the existence of free sodium and free chlorine in the solution, but sodium atoms charged with positive electricity and chlorine atoms charged with negative elec- tricity a very different matter. It is only when the respective charges are given up during electrolysis that the elements themselves are liberated. Relationships between Electrical Conductivity, Osmotic Pressure, and other Properties of Solutions Perhaps the most important feature of the theory of Arrhenius is the establishment of a quantitative relationship between electrical conductivity and the osmotic pressure and allied properties of electro- lytes. The degree of dissociation of an electrolyte can be deduced quite independently from electrical conductivity measurements, and from direct or indirect measurements of osmotic pressure (most con- veniently, perhaps, by freezing-point measurements) and the fact that the results obtained by these methods are in good agreement is one of the strongest arguments in favour of the theory. The deduction of the degree of dissociation from osmotic measure- ments is comparatively simple. Suppose that of 100 molecules in a solution the fraction a is dissociated, each molecule giving rise to n ions, the number of undissociated molecules is loo(i-a) and, the number of dissociated molecules being iooa, the number of ions is 100 na. The total number of particles is therefore 100(1 a + na)= 100(1 +(n- i)a}. The ratio of the number of particles actually present to that calcu- lated according to Avogadro's hypothesis van 't HofFs factor i is therefore z=loo{i+(- i)a}/ioo=i+(7Z- i)a, ora=izl. n i We have seen that in normal solution i for sodium chloride is about ELECTROLYSIS 265 1.7, and as n in this case is 2, we obtain 00 = 0.7 ; in other words, sodium chloride in normal solution is dissociated to the extent of 70 per cent, into its ions. As another illustration, we take a 0.18 molar solution of calcium chloride, for which z from freezing-point measure- ments = 2. 67. As in this case the molecule dissociates into three ions : = 3, and, substituting in the above formula a= 1.67/2 = 0.83. According to the theory / for a solution of sodium chloride cannot exceed 2, no matter how dilute the solution may be, and, similarly, the maximum value of z' for calcium chloride is 3. These deductions are fully confirmed by experiment. It would lead too far to describe fully the method of deducing the value of a from conductivity measurements, 1 and an indication of the principle of the method must suffice. Other things being equal, it is evident that the electrical conductivity of a definite amount of solute must, according to the theory, be proportional to the extent to which it is ionised. If then we determine the conductivity of a mol of salt at continually increasing dilutions, we must finally reach a value for the conductivity beyond which it no longer increases on dilution. Under these circumstances ionisation is complete, in other words, the who e of the salt is active in conveying the current. In more concentrated solutions the conductivity under equivalent conditions must be less, as only a part of the solute is active, and it is evident that, other things being equal, the ratio of the conductivity at a particular dilution -v to the maximum conductivity is a measure of the extent to which the solute is ionised at the dilution v. As has already been mentioned, the values of a deduced by this method are in satisfactory agreement with those obtained by osmotic methods. Acids and Bases On account of their great importance, acids and bases require rather fuller treatment from the present point of view. According to the ionisation theory an acid is a sub- stance A\hich contains hydrogen ions; the characteristic properties of acids, including sour taste, action on litmus, etc., are due to these ions, anc an acid is the more active the greater its ionisation. Thus normal solutions of nitric acid and of nitrous acid contain the same concentration of total hydrogen, and we ascribe the much greater activity of the former acid to the fact that nearly all its hydrogen is ionised, whereas the hydrogen in nitrous acid is nearly all combined. 1 Cf. Physical Chemistry, p. 261. 266 A TEXT-BOOK OF INORGANIC CHEMISTRY The relative activities of acids is shown very clearly by finding the ratio in which an amount of base insufficient to saturate both of them distributes itself between them. If, for instance, an equivalent quantity of nitric acid is added to a normal solution of sodium chloride, an equilibrium is established : NaCl + HNO 3 $NaNO 3 + HC1, and it can be shown that about half the sodium is present as the chloride, the rest as nitrate, so that the acids are approximately of equal strength. If, however, hydrochloric acid in equivalent amount is added to a normal solution of sodium nitrate, it can be shown that in the equilibrium the nitrous acid is almost completely displaced from combination with the sodium, which is present almost entirely (97 per cent.) as sodium chloride. Detailed investigation of such equilibria has shown that a base distributes itself between two acids in the ratio of their hydrogen ion concentrations as determined by other methods ; a further justification for our assumption that the activity of acids as such is to be measured, not by the total displaceable hydrogen, but by that portion of it which is present in the ionic condition. It should not, however, be assumed that all the properties of an acid are determined by its H' ion concentration ; in other words, that the non-ionised molecules are completely inactive. The above re- marks only apply to the properties common to all acids. The oxidiz- ing properties of nitric acid, for instance, are doubtless due, in part at least, to the non-ionised HNO 3 molecules. The relative activities of bases can also be compared by a distribu- tion method ; that is, by finding in what proportions a quantity of acid insufficient to saturate both of them distributes itself between them. In this way it may be shown that the ratio of the activities of potassium and ammonium hydroxides in normal solution is about 200: i, which is the ratio of the OH' ion concentrations in the two solutions. In order to illustrate the above remarks, the degree of dissociation, a, of a few acids, bases and salts at 18 is given in the accompanying table. Unless otherwise stated, the numbers refer to equivalent normal solutions. ELECTROLYSIS 267 Acids Nitric acid .... 0.82 Phosphoric acid, H'H 2 PO 4 . 0.20 Nitric acid (cone. 62 %) . . 0.097 Sulphurous acid, H ^803(25) 0.13 Hydrochloric acid . . . 0.80 Nitrous acid H'NO 2 (25) , 0.02 Hydrochloric acid (cone. 62 %) 0.138 Carbcfnic acid, H'HCO 3 (25) 0.0006 Sulphuric acid , H'HSO 4 . . 0.51 Hydrogen sulphide ,H ^8(25) 0.00024 Bases Potassium hydroxide . . 0.77 Calcium hydroxide, N/64 at 25 0.90 Sodium hydroxide . . . 0.73 Silver hydroxide, ^4630 at 25 0.64 Lithium hydroxide . . .0.64 Ammonium hydroxide . . 0.004 Salts Potassium chloride . . .0.76 Potassium sulphide . .0.54 Ammonium chloride . . 0.75 Potassium carbonate . . 0.53 Potassium nitrate . . . 0.65 Barium nitrate . . .0.38 Silver nitrate . . . .0.59 Magnesium sulphate . .0.26 Sodium acetate. . . .0.53 Copper sulphate . . .0.23 These numbers illustrate the important fact that salts, even those of weak acids, are ionised to a considerable extent in aqueous solution. Hydrolysis The decomposing action of water on certain salts, a process known as hydrolysis, has already been referred to, more particularly in connexion with the phosphates of sodium (p. 252). A much deeper insight into the phenomena of hydrolysis is obtained on the basis of the ionisation theory. So far. water has been regarded merely as a solvent, and the possi- bility th;.t it is electrolytically dissociated has not been taken into account. There is, however, a considerable amount of evidence to the effect that water itself is ionised, though to a very slight extent, according to the equation so that it contains hydrogen and hydroxyl ions. The available evidence appears to show that at room temperature the concentration of hydrogen and of hydroxyl ions in water is about o.ooooooi mol per litre ; in other words, there is I mol each of hydrogen ions and of hydroxyl ions (i gram of H' ions and 17 grams of OH 7 ions) in ten million litres of water. From the foregoing it is evident that water is an acid, as it contains H' ions ; and also a base, as it contains OH' ions. These facts are of the utmost importance in connexion with hydrolysis. 268 A TEXT-BOOK OF INORGANIC CHEMISTRY In the previous section it has been pointed out that when two acids are allowed to compete for the same base the latter distributes itself between the acids in proportion to their activities, and it has also been shown that the ratio of the activities of two acids is the ratio of the extent to which they are ionised. The same considerations apply for a salt in aqueous solution ; water, in virtue of its hydrogen ion concentration, being regarded as one of the competing acids. In the case of the salt of a strong acid, such as sodium chloride, it would not be anticipated that such a weak acid as water would take an appre- ciable amount of the base, and the available experimental evidence quite bears out this expectation. In other words, an aqueous solution of sodium chloride contains only Na" and Cl' ions and non-ionised sodium chloride in appreciable amount, and is therefore neutral. The case is quite different for a salt formed by a strong base and a weak acid, such as potassium cyanide. Here water as an acid is com- parable in strength to hydrocyanic acid (p. 328), and there is therefore a distribution of the base between the acid and the water according to the equation the proportions of potassium cyanide and potassium hydroxide de- pending upon the relative strengths of water and hydrocyanic acid under the conditions of the experiment. From the equation it is evident that: potassium hydroxide and hydrocyanic acid must be present in equivalent proportions, and since the hydroxide is much more highly ionised than hydrocyanic acid the solution contains an excess of OH' ions, and must therefore be alkaline, as is actually the case. Analogous considerations apply to the hydrolysis of salts formed from weak bases and strong acids, such as ammonium chloride. Although water is much weaker than ammonia as a base, yet the salt is slightly hydrolyzed : NH 4 C1 + H 2 O$NH 4 OH + HC1, and as hydrochloric acid is much more highly ionised than the hydroxide, the solution contains an excess of H' ions, and is therefore acid. It should be remembered that on hydrolysis of the salt of a strong base and a weak acid the solution becomes alkaline, whilst the solution of a salt of a weak base and a strong acid becomes acid on hydrolysis. As would be anticipated, salts of weak bases and weak acids are ELECTROLYSIS 269 hydroly?ed to a still greater extent under equivalent conditions than those belonging to the types just considered. As the H' and OH' ion concentration of water remains practically constant on dilution, while the H' ion concentration of the com- peting ;.cid or the OH' ion concentration of the competing base diminishes with dilution, the hydrolysis of salts of the first two types considered increases on dilution. Further, as the ionisation of water increases rapidly as the temperature rises, and the strength of other acids and bases is only slightly affected by increase of temperature, the hydrolysis of salts is markedly increased by raising the tempera- ture. The theory of hydrolysis will now be illustrated by application to the salts of crthophosphoric acid, already considered. It has been pointed out that in acids containing more than one atom of displaceable hydrogen ionisation takes place in stages, which in the, present case are as follows : (i) (2) (3) Further, in dissociation by stages, it is almost invariably the case that the ionisation in the first stage is much greater than in the second, and very much greater than in the third stage. In other words, phosphoric acid is a fairly strong monobasic acid, a much weaker dibasic acid, and as a tribasic acid is very weak indeed. It follows that the salt NaH 2 PO 4 is quite stable, the salt Na 2 HPO 4 is slightly hydrolyzed, whilst under ordinary conditions the salt Na 3 PO 4 is almost: completely hydrolyzed according to the equation Na 3 PO 4 +HOH^Na 2 HPO 4 + NaOH. As already explained, however, the hydrolysis can be greatly dimi- nished, according to the law of mass action, by addition of excess of sodium hydroxide. Further Applications of the Ionisation Theory The theory of electrolytic dissociation accounts satisfactorily for two rather remarkable facts in the domain of thermochemistry. It has been known for a long time that when equivalent amounts of strong acids are neutralized by strong bases in dilute solution the heat developc-d is the same in each case. For gram-equivalent quantities it amounts to about 1 3,700 calories at room temperature. Taking as an example the neutralization of sodium hydroxide by hydrochloric acid, 270 A TEXT-BOOK OF INORGANIC CHEMISTRY and assuming that ionisation of the acid, base and salt is complete, the equation may be written Since Na* and Cl' occur in equivalent amounts on each side, they may be neglected, and the equation reduces to OH'+H- = H 2 0, that is, the combination of H' and OH' ions to form water. The same equation applies to the neutralization of any other strong base by a strong acid ; provided that the solutions are so dilute that ionisa- tion is practically complete, the process in all cases consists simply of the combination of H* and OH 7 ions to form non-ionised water. We may therefore anticipate that the heat development will be the same for equivalent quantities, as is actually the case. The next point to consider is the so-called law of ther mo-neutrality , that no heat change occurs on mixing dilute solutions of two highly dissociated salts ; for example, sodium chloride and potassium nitrate. According to the ionisation theory, assuming that dissociation is prac- tically complete, the reaction may be represented as follows : The right and left hand sides of the equation are identical, so that the salts exert no mutual influence, which accounts for the fact that no thermal change occurs on mixing. The law of mass action applies to electrolytes (with some ex- ceptions) just as to equilibria in which undissociated chemical com- pounds are concerned. For instance, the addition of excess of one of the products of ionisation to an electrolyte diminishes the ionisation, just as excess of one of the products drives back ordinary dissociation (p. 171). Thus the strength of a weak acid, such as nitrous acid, is greatly diminished by adding excess of a highly dissociated nitrite : the equilibrium being driven towards the left. Also the strength of ammonia is greatly diminished by addition of an ammonium salt such as the chloride, and advantage is often taken of these facts in ana- lytical chemistry. In the foregoing we have confined our attention almost entirely to electrolytic dissociation in aqueous solutions, and this is justified by ELECTROLYSIS 271 the predominant use of water as a solvent in chemical operations. It should be mentioned, however, that other solvents have also been extensively investigated from this point of view, and it has been found that many of them are practically incapable of forming conducting solutions ; whilst others, such as liquefied ammonia (p. 215), form highly conducting solutions with many solutes. Electrolytic Dissociation and Valency The electro- lytic dissociation theory throws some light on the valency of electroly:es. According to the views just developed, the metallic components of salts in solution are looked upon as being positively charged, the number of charges corresponding with the ordinary valencies of the metals. Some important cations are K', Na*, Ag", NH 4 ', Ca", Hg", Fe", Al"', etc. The remainder of the salt mole- cule constitutes the negative ion (anion) which, like the positive ion, may have one, two or more (negative) electric charges corresponding with its ordinary valency. Among the more important anions, the majority of which have already been mentioned, are Cl', Br', I', NO 3 ', C1O3', SO/, CO 3 ", PO 4 "', etc. The nature of the ions is of course determined by examining the products of electrolysis at the anode and cathode respectively. It appears from the above that certain elements, more particularly hydrogen and the metals, have an affinity for positive electricity, and never appear alone in solution otherwise than as positive ions, whilst certain other atoms and groups have an affinity for negative electricity and occur in solutions only as negative ions. It should, however, be remembered that though metals never occur alone in the ionised condition otherwise than as cations, they may form part of complex anions, and, similarly, elements which alone appear only as negative ions may form part of com- plex anions. For example, the first stage in the dissociation of calcium chloride is CaCl^CaCl' + Cl', one of the chlorine atoms being present in the positive ion. It should be emphasized that the ordinary tests for the elements apply to them only when occurring as simple ions, thus KC1O 3 or KC1O 4 give none of the tests for chlorides, which are given only by the CT ion itself. With regard to the non-ionised salts it is natural to assume, although there is no direct evidence on the point, that the atoms are held together, at least partly, by the attraction of the contrary 272 A TEXT-BOOK OF INORGANIC CHEMISTRY electrical charges. The sodium chloride molecule, for instance, + - may be formulated thus, NaCl. It is clear, then, that salts (including acids and bases) contain atoms or groups with contrary electrical charges, but there is no evidence that this is true of the enormous number of chemical compounds which do not undergo electrolytic dissociation. We have seen, for instance, that chlorine and iodine, which are re- garded as electro-negative elements, unite atom for atom to form the well-defined compound IC1. Of the nature of the chemica. attraction in this case, whether electrical or otherwise, we kno nothing. The same remark applies to the molecules of the elemen: themselves. In some instances (e.g. hydrogen, oxygen) these ai so stable that even at the highest temperatures at our dispos; ' there is no evidence even of a partial decomposition into the coir-- ponent atoms. Finally, while on the subject of valency, a few words may f devoted to compounds formed by the association of two or mo' complete molecules, e.g. HF'KF, and NaCl,2H 2 O. Substances A? this type, of which the commonest are salts associated with "wart of crystallization," are often called "molecular compounds,' cannot be represented satisfactorily on the simple theory of valen If the affinities of the atoms of sodium and chlorine were complete- satisfied in forming sodium chloride, it is scarcely to be anticipate . that the molecule could bind, though weakly, two molecules of water. We may assume, then, that there remains over a little affinity, the so-called " free affinity," which enables the molecule to bind one or more molecules of another compound which also possesses free affinity. Few, if any, inorganic compounds appear to be entirely devoid of residual affinity or entirely saturated ; in other words, all of them appear to be capable of forming more or less stable "mole- cular compounds" with other compounds. Some substances belong- ing to organic chemistry, however (p. 331), appear to be almost completely saturated. At present, the subject of valency is in a rather unsatisfactory condition, and the views of chemists differ widely on many points of the subject. The matter is further referred to on p. 562. CHAPTER XXI SULPHUR, SELENIUM AND TELLURIUM SULPHUR Symbol, S. Atomic weight =32. Molecular weight =64 (at 1000). Relations The compounds of sulphur are in many ' ^ cases of the same type as those of oxygen, mainly because both iments are divalent. Sulphur, however, is also quadrivalent (as Jeed oxygen is in some compounds) and sexavalent. The best own hydrogen compound of sulphur is H 2 S (corresponding with i^ 2 O), which is a weak acid. The best known oxides have the respec- e formulae SO 2 , sulphur dioxide, and SO 3 , sulphur trioxide ; the 1 corresponding with the former is H 2 SO 3 , sulphurous acid, whilst most important acid derived from SO 3 is sulphuric acid, H 2 SO 4 . Occurrence Sulphur has been known from the earliest times, it occurs free in nature, more particularly in volcanic districts in Italy, Sicily, Iceland, China, India and in the Yellowstone Park in America. It is probably formed by interaction of hydrogen sulphide and sulphur dioxide : 2H 2 S + SO 2 ->2H 2 O + 38 | both of these gases being given off from active volcanoes. It also occurs as sedimentary deposits, not formed by volcanic action but partly at least by the agency of reducing bacteria. These deposits are particularly extensive in Texas and Louisiana, and are now the source of a considerable proportion of the world's sulphur supply. In combination with hydrogen, sulphur occurs as hydrogen sulphide in certain mineral springs. It also occurs very largely in combina- tion with metals, forming sulphides. Some important sulphides are galena, PbS, iron pyrites, FeS 2 , copper pyrites, CuFeS 2 , zinc blende, ZnS, stibnite or antimony sulphide, Sb 2 S 3 , and cinnabar, which is mercuric sulphide, HgS. It also occurs as sulphates, the more important being gypsum or 1 8 2 73 274 A TEXT-BOOK OF INORGANIC CHEMISTRY calcium sulphate, CaSO 4 , heavy spar or barium sulphate, BaSO 4 , and magnesium sulphate in the form of kieserite and Epsom salts. Sulphur is one of the essential constituents of albumen, and is there- fore found in all plant and animal tissues. Preparation (i) On the commercial scale, the sulphur is separated from the gypsum, limestone and other substances with which it is usually mixed, by heating out of contact with air, when it melts and flows away from the impurities. In Sicily the fusion FIG. 55. is effected by the wasteful method of burning a portion of the sulphur. The sulphur thus obtained is further purified by distillation, as shown in Fig. 55. It is first melted in the iron vessel A, and drawn off as required into B, in which it is heated to boiling, and the vapour con- veyed into the large bricklined chamber C. The first portions of vapour entering the chamber condense on the walls in the form of a fine powder, known as flowers of sulphur. As the walls become hot, however, the sulphur melts and collects on the floor of the chamber as an amber-coloured liquid, which is drawn off from time to time at SULPHUR, SELENIUM AND TELLURIUM 275 D and cast into sticks in wooden moulds. In this form it is known as roll sulphur. (2) By heating iron pyrites out of contact with air and condensing the vapour : (3) Sulphur is also obtained commercially from the "alkali waste" of the Leblanc soda process (p. 387), which contains a large propor- tion of calcium sulphide, CaS ; and also from the oxide of iron used in removing sulphur compounds from coal gas, and which finally con- tains a considerable proportion of sulphur. These methods are briefly described at a later stage (pp. 333 and 389). Physical Properties Sulphur usually occurs as a pale yellow, brittle, crystalline solid, insoluble in water, readily soluble in carbon disulphide, and more or less soluble in certain organic liquids such as turpentine, benzene and chloroform. It is a very bad conductor of heat and electricity. This ordinary form melts at 114.5 to a pale yellow mobile liquid. When further heated to 160 the liquid darkens, and is so viscous that the vessel containing it may be inverted without causing it to run out. At 260 the liquid is distinctly less viscous, at 400 it is quite mobile, at 445 it boils, giving off a yellowish-brown vapour. Just above its boiling-point and under fairly high pressure the density of the vapour approximates to 256, corresponding with the formula S 8 , from 860 to at least 1600 it is 64, corresponding with the formula S 2 , and it has quite recently been shown that in the neigh- bourhood of 2000 it falls to 48 (Nernst), indicating that the diatomic molecules are then split up to some extent into single atoms. In solu- tion in carbon disulphide the formula of sulphur is S 8 . Sulphur exists in at least three allotropic modifications, two of which, the " rhombic " and " prismatic " forms, are crystalline and the other amorphous. Allotropic Modifications (a) Rhombic Sulphur This is the form which is stable at the ordinary temperature, and therefore native suiphur and ordinary roll sulphur occur in this form. Good crystals are obtained by the slow evaporation of a solution in carbon disulphide at room temperature : they are orthorhombic pyramids. The density of rhombic sulphur is 2.06 ; it melts, as already mentioned, at 114.5. (b) "Mmodinic" or "Prismatic" Sulphur When sulphur is melted in a crucible, the mass allowed to cool till a crust forms, 276 A TEXT-BOOK OF INORGANIC CHEMISTRY and the still liquid portion poured off through a hole pierced in the crust, the inside of the crucible is found to be lined with long needle- shaped crystals. The crystals are almost colourless, the density is 1.96, and they belong to the monoclinic system. Monoclinic sulphur is stable above 96, but at lower temperatures gradually changes to rhombic sulphur. Just as water is stable above o and ice below o, while the two remain in equilibrium at that temperature, so 96 is the transition temperature, at which rhombic and monoclinic sulphur are in equilibrium, whilst at higher tem- peratures monoclinic, at lower temperatures rhombic, is the stable form. Neither of the changes is very rapid, so that it is possible to determine the melting-point of rhombic sulphur, which lies nearly 20 above the transition point, before appreciable change into the other form takes place. Substances such as sulphur, which are known in two crystalline forms, are said to be dimorphous; in general, when a substance is known in more than one crystalline form, it is said to be polymorphous. (c) Amorphous Insoluble Sulphur. Plastic Sulphur When sul- phur, heated nearly to boiling, is suddenly cooled by pouring into cold water, an elastic mass, resembling indiarubber, is obtained. On standing for a few days it hardens, and may then be shown to be a mixture of rhombic sulphur, which can be dissolved out by carbon disulphide, and a third modification, insoluble in carbon disulphide and devoid of crystalline form, which is usually termed amorphous sulphur. Amorphous sulphur is now usually regarded as supercooled viscous sulphur. At the ordinary temperature amorphous sulphur changes extremely slowly to soluble sulphur, at 100 the change is very rapid. For purposes of distinction, mobile liquid sulphur is often termed S\ and viscous sulphur (as well as the amor- phous modification) S^. Molten sulphur is regarded as an equilibrium mixture of S\ and S^, in proportions depending on the temperature. Amorphous sulphur is also contained in flowers of sulphur (prepared by condensing sulphur vapour on a cold surface), and is obtained by dissolving away the soluble sulphur with carbon disulphide. It can be obtained in the same way from the sulphur prepared by the action of acids on thiosulphates. Chemical Properties Sulphur combines directly with many other elements, both metals and non-metals. It burns in air or oxygen to sulphur dioxide, SO 2 . It combines directly with hydrogen when heated, forming hydrogen sulphide, H 2 S. If some sulphur is placed in the bottom of a test tube, which is loosely filled up with copper filings, SULPHUR, SELENIUM AND TELLURIUM 277 and the sulphur then boiled, the copper catches fire in the vapour and burns to the sulphide, CuS. Finely-divided iron and sulphur also combine directly (p. 4), giving out heat and light. Sulphur is used in the manufacture of^sulphuric acid and of sulphur dioxide, and forms one of the constituents of ordinary gunpowder. It is also used in medicine. Velocity of Crystallization. Supercooling The re- lationsl ip between supercooling and velocity of crystallization, as illustrated in the behaviour of S^, is of considerable general interest. When c liquid is supercooled to a certain extent in the absence of particles of the solid phase (p. 85) nuclei of the solid appear spon- taneously in the liquid and crystallization proceeds at a definite rate. When observations are made at temperatures more and more removed from the melting-point, it is found that the velocity of formation of the solid phase increases at first with the degree of cooling, attains a maximum, and at still lower temperatures is extremely slow. In the latter case the lowering of temperature diminishes both the tendency to the spontaneous formation of nuclei of the solid phase and also the velocity of crystallization. We can now understand that if the temperature be rapidly lowered beyond the point at which the change is rapid, the rate of crystallization or (as in the case of sulphur) the rate of transition may be extremely slow. COMPOUNDS OF SULPHUR AND HYDROGEN By fai the most important compound of hydrogen and sulphur is hydrogen sulphide, H 2 S. A number of other compounds of these elements, represented by the general formula H 2 S, where n stands for the integral numbers from 2 to 7, also appear to exist, and the following are definitely known, viz., H 2 S 2 ,H 2 S 3 , and H 2 S 5 . HYDROGEN SULPHIDE, H 2 S Occurrence This compound occurs in certain mineral waters, in springs at Harrogate, for example. It is formed in the decay of organic matter containing sulphur in the absence of air ; and is mainly responsible for the characteristic odour of rotten eggs. Preparation (i) By direct combination of the components on heating. The reaction is comparatively slow, but is finally practically complete at 310; in higher temperatures it is more rapid, but less complete (see below). 278 A TEXT-BOOK OF INORGANIC CHEMISTRY (2) By the action of dilute hydrochloric or sulphuric acid on ferrous sulphide at room temperature : FeS + 2HCl->FeCl 2 + H 2 Sf. This is the usual laboratory method for preparing the gas, which may conveniently be done in a Kipp's apparatus. It is, however, always contaminated with free hydrogen, owing to the presence of uncombined iron in commercial ferrous sulphide. The reaction pro- ceeds practically completely in the direction of the arrow, mainly because the hydrogen sulphide readily leaves the system as a gas, but partly also because it is a very weak acid (p. 187). (3) The pure compound is obtained by heating antimony sulphide with concentrated hydrochloric acid : Physical Properties Hydrogen sulphide is a colourless gas with a characteristic disagreeable odour. The liquefied gas boils at 60 and the solid melts at 83. It is moderately soluble in water ; the "solubility" being 4.37 at o, 3.58 at 10, and 2.9 at 20. It is extremely poisonous ; one part in 800 of air proved fatal to a dog, and one in 1500 is said to cause death to birds. The best antidote is the inhalation of very dilute chlorine. Hydrogen sulphide dissociates on heating, according to the equation As the compound is strongly exothermic, the dissociation increases with rise of temperature. It amounts to 2.3 per cent, at 627, 31.7 per cent, at 1137, and 76.1 per cent, at 1727. Chemical Properties Hydrogen sulphide burns in air with a bluish flame, forming water and sulphur dioxide : In a limited supply of air combustion is incomplete, and free sulphur is formed : and it is not unlikely that part of the free sulphur found in nature is produced in this way. Aqueous solutions of hydrogen sulphide are very unstable, owing to gradual oxidation by the oxygen of the air, water being formed and sulphur separating as a precipitate. SULPHUR, SELENIUM AND TELLURIUM 279 In harmony with the comparatively small affinity between its com- ponents (as illustrated by its dissociation on heating) hydrogen sulphide is a powerful reducing agent. The halogens are reduced to the corresponding halhydrogen acids, e.g. and concentrated sulphuric acid is reduced to sulphur dioxide and water On this account sulphuric acid cannot be used to dry hydrogen sulphide, but phosphorus pentoxide is suitable for this purpose. Hydrogen sulphide in aqueous solution acts as a feeble acid towards litmus, cind is, in fact, a dibasic acid. As with other dibasic acids, ionisation takes place in two stages : but the dissociation is small even in the first stage (p. 267), and in the second stage is very minute indeed. With certain monacidic bases, such as the alkalis, hydrogen sulphide gives both acid and normal salts, e.g. KHS and K 2 S. The former are prepared by passing excess of hydrogen sulphide into a solution of the base : KOH + H 2 S->KHS + H 2 0. Since H 2 S as a monobasic acid, although weak, is much stronger than water, the compound KHS is only slightly hydrolyzed. The normal sulphide, K 2 S, is prepared by adding to the acid sulphide an equivalent of KOH, and removing the water by evapora- tion : When, however, the normal sulphide K 2 S is dissolved in water, it is practically completely hydrolyzed according to the equation K 2 S + H 2 O^KHS + KOH, because the acid HS' is much weaker than the water, and the solution is, therefore, strongly alkaline. While the sulphides of the alkali metals, and of certain other metals, are soluble in water, those of a number of metals, such as zinc and iron, are insoluble in water but soluble in acids, and still another 2 8o A TEXT-BOOK OF INORGANIC CHEMISTRY group of metals, including copper, mercury, and lead, yield sulphides insoluble in dilute acids. The action of hydrogen sulphide on a solution of zinc chloride, for instance, may be represented as follows : but the equilibrium lies very near to the left, and no precipitation of zinc sulphide occurs. If, on the other hand, an acidified solution of copper chloride is used, the reaction is as follows : CuCl 2 + H 2 S->CuS j + 2HC1, and owing to the slight solubility of copper sulphide is practically com- plete in the direction of the arrow. On these facts is based a method of detecting and separating many of the metals fully described in books on qualitative analysis. The characteristic colours of many sulphides are also useful for purposes of identification. Composition When a piece of tin is heated in hydrogen sulphide over mercury, the sulphide and free hydrogen are formed, and on cooling it will be observed that no change of volume has occurred. It follows that the molecule of hydrogen sulphide contains two atoms of hydrogen, and its formula is H 2 S A -. As the molecular weight from the results of density determinations is 34, and the atomic weight of sulphur is 32, the molecule of hydrogen sulphide must contain one atom of sulphur, and its formula is therefore H 2 S. HYDROGEN POLYSULPHIDES When sodium sulphide is heated for some hours with varying proportions of sulphur in an atmosphere of hydrogen, and the products dissolved in water, solu- tions are obtained containing the following polysulphides of sodium : Na 2 S 2 , NagSg, Na 2 S 4 , and Na^Sg. When these solutions are allowed to flow into separate quantities of dilute hydrochloric acid, cooled in a freezing mixture, a yellow, oily liquid separates out. This liquid is a mixture of persulphides of hydrogen, and its composition varies with that of the sodium persulphide used in obtaining it. From the crude persulphide, by fractional distillation under reduced pressure, Bloch and Holm l have recently obtained the disulphide, H 2 S 2 , and the trisulphide, H 2 S ;? , in a pure condition. Hydrogen Bisulphide, H 2 S 2 , is a pale yellow, oily liquid, of density 1.376 ; it boils at 74 to 75. It decomposes slowly at the ordinary temperature, rapidly on waiming, giving off hydrogen sulphide and depositing rhombic sulphur. It is slowly decomposed by acids, very rapidly by alkalis, and is fairly stable in benzene solution. It is the analogue of hydrogen peroxide, H 2 O 2 . 1 Berichte, 1908, 41, 1961-1985. SULPHUR, SELENIUM AND TELLURIUM 281 Hydrogen Trisulphide, H 2 S 3 , is also a pale yellow, oily liquid, of density 1.496 at 15. In chemical behaviour it resembles the disulphide, but is less volatile and more stable towards alkalis. Hydrogen Pentasulphide, H 2 S 5 , also doubtless exists, but its properties have not been thoroughly established. The investigation of these sulphides is rendered very difficult owing to their tendency to dissolve sulphur. COMPOUNDS OF SULPHUR AND CHLORINE The following three well-defined compounds of sulphur and chlorine are known : Sulphur monochloride, S 2 C1 2 . Sulphur dichloride, SC1 2 . Sulphur tetrachloride, SC^. Sulphur Monochloride, S 2 C1 2 , is prepared by passing dry chlorine oyer fused sulphur in a retort. It is a yellow liquid which boils at 137 to 138; the solid melts at - 75 to - 76. It is the most stable of the chlorides of sulphur, being only slightly dissociated at its boiling-point. It is readily decomposed by water : It is an excellent solvent for sulphur, dissolving over 60 per cent, of the latter at room te nperature, and the solution is used in the vulcanization of rubber. Sulphur Dichloride, SC1 2 , is prepared by passing chlorine into the mono- chloride, ooled to - 15, until the theoretical gain in weight is attained. It is a dark redd sh-brown liquid of density 1.622 at 15 and boils at 59 under at- mospheric pressure. At] its boiling-point it is partially decomposed into the monochloride and chlorine. It is readily decomposed by water : 2 SC1 2 +2H 2 0-> 4 HC1+S0 2 + S. Sulphur Tetrachloride, SC1 4 , is obtained when the dichloride and chlorine are brough t together at temperatures below - 75. The pure compound is obtained by freezing out, the excess of chlorine being removed by centrifugal action. The crystals melt between -30 and -20. Even at low temperatures it is partially dissociatec and decomposition is practically complete at room temperature. It is readily decomposed by water : With b -online, sulphur forms the compound S 2 Br 2 , a brownish-red liquid, which boih with partial decomposition at 200. Compounds of sulphur and iodine have been described, but their existence is doubtful. OXIDES AND OXYACIDS OF SULPHUR The following four oxides of sulphur are known : Disvlphur trioxide (hyposulphurous anhydride) . . S 2 O 3 Sulphur dioxide (sulphurous anhydride) . . . >O 2 Sulphur trioxide (sulphuric anhydride) . . . SO 3 Sulphur heptoxide (persulphuric anhydride) . . S 2 O 7 . 282 A TEXT-BOOK OF INORGANIC CHEMISTRY The following oxyacids, the first five of which are derived from the oxides just mentioned, are known : Hyposulphurous acid . . H 2 S 2 O 4 O== ?' O = S'OH /OH Sulphurous acid . . . H 2 SO 3 OS\ \OH /OH Sulphuric acid . . . H 2 SO 4 O 2 S\ \OH /OH OH\ Pyrosulphuric acid . . . H 2 S 2 O 7 O 2 S^- O -^SO 2 (Nordhausen sulphuric aid) /OH Permonosulphuric acid . . H 2 SO 5 O 2 S/ (Caro's acid) \OOH /OH OH\ Persulphuric acid . . . H 2 S 2 O 8 O 2 S<( /SO 2 \ O O / /OH Thiosulphuric acid . . . H 2 S 2 O 3 O 2 S<^ \SH. Besides these, four less important acids, dithionic acid, H 2 S 2 O 6 , trithionic acid, H 2 S 3 O 6 , tetrathionic acid, H 2 S 4 O 6 , and pentathionic acid, H 2 S 5 O 6 , are known. SULPHUR DIOXIDE, SO 2 Occurrence As already mentioned, sulphur dioxide is given off from active volcanoes. It also occurs in traces in the atmosphere of towns, being formed from the combustion of the sulphur always present in coal. Preparation (i) By burning sulphur in air or oxygen : S0 2 , A little sulphur trioxide (in air up to 7 per cent.) is formed at the same time. (2) By burning sulphides, e.g. iron pyrites, FeS 2 , in air : SULPHUR, SELENIUM AND TELLURIUM 283 This method is used on the large scale in the manufacture of sulphuric acid. (3) By heating copper turnings with concentrated sulphuric acid : This is the most convenient laboratory method for preparing the gas, which can be collected over mercury or by upward displacement of air. (4) The reduction of concentrated sulphuric acid to sulphur dioxide can also bo effected by heating it with carbon or sulphur : These and the equations representing the reduction of sulphuric acid to sulphur dioxide can readily be obtained when the acid is written in the form SO 3 ,H 2 O (cf. p. 226). (5) By acting on sulphites with dilute sulphuric or hydrochloric acid : Physical Properties Sulphur dioxide is a colourless gas, with the penetrating and characteristic odour associated with burning sulphur. It can easily be obtained as a colourless mobile liquid, most conveniemly by passing the gas (obtained by the action of sulphuric acid on copper and dried by bubbling through concentrated sulphuric acid) through a tube immersed in a freezing mixture of ice and salt (Fig. 56). Liquid sulphur dioxide boils at - 8 ; at o its vapour pressure is 1.87 atmospheres, at 20 it is 3.24 atmospheres. It is a good solvent, especially for inorganic salts, and many of the solu- tions are *ood electrolytes. On account of its ready volatility and large heat of vaporization, liquid sulphur dioxide is used for obtain- ing low temperatures by its rapid vaporization (p. 66). The liquid is now obtainable commercially in syphons. Solid sulphur dioxide melts at -- 76. Sulphur dioxide is readily soluble in water, at o the "solubility" is 79.79, at 23 39.37, and at 40 18.766. It is completely expelled from solution by boiling. Chemical Properties The aqueous solution of sulphur dioxide has an acid reaction, owing to the fact that it combines with water to form sulphurous acid, H 2 SO 3 (see below). Sulphur dioxide is readily oxidized to the trioxide, SO 3 , and is therefore a useful reducing agent. It combines directly with oxygen 284 A TEXT-BOOK OF INORGANIC CHEMISTRY at high temperatures (p. 286), but its reducing character is more pro- nounced in aqueous solution, owing to the readiness with which sulphurous acid is converted to sulphuric acid. The free halogens, FIG. 56. iodates (p. 185), potassium dichromate (p. 521), and other oxidizing agents are readily reduced at the ordinary temperature : Papers soaked with potassium iodate and starch are sometimes used SULPHUR, SELENIUM AND TELLURIUM 285 for the detection of sulphur dioxide, as they turn blue when iodine is liberated. Sulphur dioxide in the presence of water is a useful bleaching agent. In most cases this is connected with its reducing properties, the colouring natters being reduced by the hydrogen liberated according to the equation SO 2 f2H 2 O->H 2 SO 4 +H 2 . In other cases (e.g. the bleaching of flowers) the change appears to depend upon direc: combination of the dioxide with the colouring matters, as the colour is restored on warming or on treatment with a dilute alkali. It is used for bleaching straw, wool, silk and other materials that would be damaged by chlorine. Sulphur dioxide, both in the gaseous form and in solution, is a useful dis- infectant. Composition The composition of sulphur dioxide can be determined with the apparatus represented in Fig. 57. A small quantity of sulphur is placed in the cup (connected to a copper wire) which is then placed in position in the bulb, the latter being filled with oxygen at atmospheric pressure. The position of the mercury in the right-hand tube is noted, the pressure reduced by running out mer- cury at the stopcock, and the sulphur ignited by passing an electric current, which heats to redness the thin platinum wire which has been arranged to dip in the sulphur. After combustion is complete and the apparatus has cooled, the pressure is again adjusted to that of the atmosphere by pouring mercury into the left-hand limb, and it will then be found that the mercury in the other limb stands at the same level as before. It follows that the volume of sulphur dioxide produced is equal to that of the oxygen used up in its formation ; FIG. 57. 286 A TEXT-BOOK OF INORGANIC CHEMISTRY in other words, sulphur dioxide contains its own volume of oxygen. By Avogadro's hypothesis the molecule of sulphur dioxide therefore contains a molecule or two atoms of oxygen and its formula is ?> X O. 2 , where x is a whole number. The molecular weight of the gas, deter- mined from its density, is 64, and as it contains 2 atoms or 32 parts of oxygen, there remains 32 parts of sulphur. Now no sulphur com- pound is known the molecule of which contains less than 32 parts of that element, so that 32 is the atomic weight of sulphur, and the formula of sulphur dioxide is SO 2 . Sulphurous Acid and Sulphites Sulphurous acid, H 2 SO 3 , has never been obtained otherwise than in aqueous solution. As a dibasic acid it forms two series of salts with alkalis, the normal sulphites, for example Na. 2 SO 3 , and the acid sulphites, for example NaHSO 3 , which are soluble in water. The normal sulphites in aqueous solution are considerably hydrolyzed, as sulphurous acid, though a fairly strong monobasic acid (H 2 SO 3 :H'+HSO 3 ') is a very weak dibasic acid (HSO/^tH 1 + SO 3 "), though not so weak as hydrogen sulphide. The sulphites of the other metals are mostly insoluble in water. SULPHUR TRIOXIDE Preparation (i) By direct combination of sulphur dioxide and oxygen : The combination is very slow, even at high temperatures, but is fairly rapid when the mixture is passed over finely-divided platinum, platinum asbestos or ferric oxide, heated to 400. On passing the vapours into a cooled receiver, the trioxide is obtained in colourless crystals. At 400 combination is practically complete under equilibrium conditions, but as the temperature is raised the trioxide undergoes dissociation, and the change is practi- cally complete in the direction of the upper arrow at 1000. As this reaction is now employed in the commercial preparation of sulphuric acid it is fully described at a later stage. (2) By heating pyrosulphuric acid (fuming sulphuric acid) (q.v.} : H 2 S 2 O 7 ->H 2 SO 4 (3) By strongly heating ferric sulphate : Fe 2 (SO 4 ) 3 ->Fe 2 O 3 + SULPHUR, SELENIUM AND TELLURIUM 287 (4) By dehydrating concentrated sulphuric acid by means of phos- phorus pentoxide : Physical Properties Sulphur trioxide is a liquid at the ordinary temperature ; it boils at 46, and gives dense white fumes when exposed to the air, owing to the combination of its vapour with moisture, f orming sulphuric acid. The crystals obtained by cooling the liquid melt at 14.8. A second, more stable form of the solid occurs in long, needle-shaped crystals resembling asbestos. This form vaporizes on heating without previously melting, and is pre- sumably -i polymer (p. 175) of the ordinary form, probably (SO 3 ) 2 . The vapour at low temperatures contains the two forms in equilibrium, but dissoc iates completely into the simpler form as the temperature rises, The dissociation of the trioxide into sulphur dioxide and oxygen at high temperatures has already been referred to. Chemical Properties Sulphur trioxide unites very vigorously with water to form sulphuric acid : SO 3 + H 2 O->H 2 S0 4 . It also unites directly with certain metallic (basic) oxides with forma- tion of salts : CaO + SO 3 ->CaSO 4 . The combination of basic and acidic oxides is one of the general methods for preparing salts (p. 374). SULPHURIC ACID History Sulphuric acid was a familiar substance to the al- chemists, being prepared by distilling ferrous sulphate, FeSO 4 ,7H 2 O (p. 544). The latter compound was khown as green vitriol, whence the name oil of vitriol, still sometimes applied to sulphuric acid. Preparation On the commercial scale sulphuric acid is pre- pared by bringing about the combination of sulphur dioxide and oxygen, and dissolving the resulting sulphur trioxide in water. The difficulty arising from the extreme slowness with which the gases combine directly (p. 286) is overcome by using catalysts. In the 288 A TEXT-BOOK OF INORGANIC CHEMISTRY "contact process," which will be first described, the catalyst is platinized asbestos ; in the " lead chamber process " oxides of nitrogen are used for this purpose. (1) The Contact Process Although this method is apparently a very simple one, serious difficulties were encountered in practice owing to the fact that the platinum soon lost its catalytic power. This was ultimately traced to the presence of impurities mainly arsenic and dust in the mixture of sulphur dioxide and air used in the process. This mixture is usually prepared by roasting iron pyrites (p. 273), which contain arsenic, hence the presence of this impurity in the reaction mixture. Methods for removing the arsenic and dust have been worked out by Knietsch (1899), to whom the commercial success of this method is mainly due. 1 In outline the contact process is as follows. The purified mixture of sulphur dioxide and air, of such composition that the ratio SO 2 : O 2 is approximately 2 : 3, is passed through tubes containing the platinized asbestos and kept at 400 to 450. As a large amount of heat is given out in the process of combination, and a higher temperature is disadvantageous, the cooling is effected by passing fresh portions of the reaction mixture round the outside of the tubes, the mixture being thus warmed to the most favourable temperature before enter- ing the reaction tubes. By increasing or diminishing the rate of flow of gas, the temperature can be regulated satisfactorily. The issuing trioxide is not passed into water, as part of it is apt to escape condensation, but is absorbed in 97 to 98 per cent, sulphuric acid, which is kept at that strength by the simultaneous addition of water. (2) The Lead Chamber Process As already indicated, this method consists essentially in the oxidation of sulphur dioxide to sulphuric acid by free oxygen in the presence of aqueous vapour, oxides of nitrogen being used to accelerate the reaction. A mixture of sulphur dioxide, air, aqueous vapour and nitrous fumes is allowed to interact in large leaden chambers at room temperature, and ultimately a dilute solution of sulphuric acid collects on the floor of the chambers. The simplest method of representing the reactions is as follows: (1) SO 2 + NO 2 ->SO 3 (2) SO 3 +H 2 O->H 2 SO 4 (3) 1 Other inventors, notably Winkler in Freiburg and Messel in London, appear also to have solved the difficulties of the process before the publication of Knietsch's patent in 1899, but their methods were kept secret. SULPHUR, SELENIUM AND TELLURIUM 289 As sulphur dioxide and nitrogen dioxide react much more rapidly than the former gas does with oxygen, we can easily understand why sulphuric acid is obtained so readily by an indirect method. The nitric oxide formed in equation (i) is rapidly reconverted to the peroxide, according to equation (3), so that theoretically a small amount of peroxide could convert an unlimited amount of sulphur trioxide to sulphuric acid. In practice, however, secondary reactions also occur to some extent (for example, reduction to nitrous oxide, N 2 O, which is no longer able to combine with oxygen, p. 234), result- ing in the loss of activity of part of the oxides of nitrogen, so that they have 10 be continuously supplied as the action proceeds. The reactions taking place in the lead chambers are probably more complicated than the above simple scheme would indicate, but they are not yet thoroughly understood. According to Lunge, reaction proceeds mainly in two steps ; the first leads to the formation of /OH nitrosyl-sulphuric acid, SO 2 \ , according to the equation \O-NO /OH ?,SO 2 + NO + NO 2 + O 2 +H 2 O->2SO 2 < \O-NO and the second stage consists in the rapid breaking up of nitrosyl- sulphuric acid by water, with formation of sulphuric acid, nitric oxide, and nitrogen peroxide : /OH /OH 2SO 2 < +H 2 O->2SO 2 < + NO + NO 2 . \0-NO M>H As the formula shows, nitrosyl-sulphuric acid is derived from sul- phuric acid by the displacement of one of the hydrogen atoms by the univalent - N = O group. This substance does not appear when the process is working properly, but if the supply of water is in- sufficient it separates on the sides of the chambers in colourless crystals, the so-called "chamber crystals." In a later paper Lunge modified this theory to some extent. 1 Raschig 2 on the other hand, considers t at nitrous acid, HNO 2 or NOOH, is /OH the active catalyst, and that two unstable intermediate compounds, SO 2 \^ Cf. Encyc. Britannica, nth Edition, vol. xxvi. p. 67. " Journ. Soc. Chem. Ind., 1910. 19 290 A TEXT-BOOK OF INORGANIC CHEMISTRY /OH nitroso-sulphuric acid, and the compound SO 2 \' />O have a transient exist- XOH ence during the change. Raschig's view is summarized in the following equations : (i) NO- (2) (3) NO(OH)S0 2 'OH_>NO + H 2 S0 4 (4) The production of sulphuric acid by this process can be shown FIG. 58. on the small scale in the laboratory. A large wide-mouthed flask (Fig. 58) is provided with a cork carrying five tubes, three of which are used for leading nitrous fumes, sulphur dioxide and air into the flask. In order to ensure the formation of the chamber crystals it is preferable to dry these gases by passing them through concentrated sulphuric acid. They are then admitted into the flask, and in a short time colourless crystals of nitrosyl-sulphuric acid will be observed on the sides and bottom of the flask. Steam is then admitted by a fourth tube, when the crystals soon disappear, oxides of nitrogen are regenerated, and sulphuric acid collects on the bottom of the flask. The arrangement of the apparatus used on the commercial scale is SULPHUR, SELENIUM AND TELLURIUM 291 shown in F ig. 59. The sulphur dioxide, obtained by burning sulphur, or more usually by roasting iron pyrites or other sulphides in a current of air (the pyrites burners are shown at P P), is first led into a long flue 292 A TEXT-BOOK OF INORGANIC CHEMISTRY in which the dust particles mechanically carried along settle out, and the proper proportion of air is added. The mixed gases, which are at a temperature of about 300, are then passed up the Glover tower, G, which is a high tower lined inside with lead and filled with stones, over which dilute sulphuric acid containing dissolved oxides of nitro- gen (nitrosyl-sulphuric acid) is trickling. 1 The nitrosyl-sulphuric acid is decomposed by the hot furnace gases, with formation of sulphuric acid, which is collected at the bottom of the tower (at A), and nitric oxide, which is carried along by the stream of gas, now at a much lower temperature, into the first of the chambers : OH 2S0 2 -NO /O 2 < \O- The chambers, which are constructed completely of sheet lead, often have a capacity of 150,000 to 200,000 cubic feet each, and are arranged in sets of three or four connected together (two only are shown in the figure). As the gases are slowly drawn through these chambers, water in the form of steam is injected, and the chemical reactions already described (p. 289) take place, resulting in the formation of sulphuric acid, which collects on the floors of the chambers. The gases escaping from the last of the chambers, which consist mainly of oxides of nitrogen and a large proportion of nitrogen from the air originally drawn into the chamber, are passed up a second tower, the so-called Gay-Lussac tower, H, which is filled with coke, over which 80 per cent, sulphuric acid continually trickles from a reservoir at the top of the tower. The acid used for this purpose is part of that drawn off at the bottom of the Glover tower, whence it is pumped up to the top of the Gay-Lussac tower. The object of the Gay-Lussac tower is to prevent the escape of the oxides of nitrogen ; these are absorbed almost completely by the concentrated acid with formation of nitrosyl- sulphuric acid. This acid is then pumped to the top of the Glover tower, where the oxides of nitrogen are removed from it by the hot furnace gases, and swept back into the chambers, as already fully explained. Owing to secondary reactions and unavoidable losses, however, the oxides of nitrogen are gradually used up, and the loss must be made good. This is effected by heating sodium nitrate with sulphuric acid in earthenware pots, which are placed in the path of the furnace gases 1 A mixture of acid from the chambers and from the Gay-Lussac tower is used for this purpose. SULPHUR, SELENIUM AND TELLURIUM 293 on their way to the chamber at N N. The nitric acid is rapidly reduced to nitrogen peroxide by the sulphur dioxide : When the sulphuric acid in the chamber has reached a density of i. 60 (68 per cent, of acid) it is withdrawn (more concentrated acid attacks the material of the chambers and dissolves oxides of nitrogen) and further concentrated to a density of 1.70 (77 per cent, of acid) by evaporation in leaden pans. 1 Above this strength lead would be dissolved fairly readily, so the final concentration is effected by heating in glass, platinum, or cast-iron vessels. The commercial acid thus obtained has a density of 1.83 to 1.84, and contains 93 to 98 per cent, of the pure acid. The commercial acid contains a number of impurities, more particularly arsenic from the pyrites, lead sulphate, and oxides of nitrogen. Most of the impurities, except the arsenic, are got rid of by adding a little ammonium sulphate and distilling from platinum stills ; the arsenic is removed by special methods. Physical Properties The purest sulphuric acid which can be obtained by distillation contains 1.5 per cent, of water, and its density is 1.842 at 15. The anhydrous acid is most conveniently prepared by adding the theoretical amount of sulphur trioxide to the 98.5 per cent, acid ; it is a colourless, odourless, oily liquid of density 1.838 at 15. When the pure acid is heated, it begins to decompose about 150, giving off white fumes of sulphur tri- oxide. At 338, when the concentration has fallen to 98.5 per cent, of acid through loss of sulphur trioxide, the liquid boils. The 98.5 per cent, acid is therefore a constant-boiling mixture, similar to that formed by hydrochloric acid and water. Even at ts boiling-point, the vapour of sulphuric acid is dissociated to a considerable extent into the trioxide and water vapour : Chemical Properties When sulphuric acid is added to water, a great quantity of heat is evolved. When i mol of water is added to i mol (98 grams) of acid, 6700 cal. are given out. More neat is given out, but in diminishing amount, as further quantities of water are successively added. The total heat evolved when a mol of acid is added to a very large excess of water is about 1 Chamber icid can also be concentrated up to 78-80 per cent, by passing it through the Glover tower. 294 A TEXT-BOOK OF INORGANIC CHEMISTRY 20,000 calories. This remarkable thermal effect is doubtless due mainly to chemical combination between the acid and water, and partly, in all probability, to progressive electrolytic dissociation. One well-defined hydrate of sulphuric acid, H 2 SO 4 .,H 2 O, has been obtained in prismatic crystals, melting at 8.5. There is some evi- dence of the existence of a second hydrate, H 2 SO 4 ,2H 2 O, but it has not been definitely isolated. Sulphuric acid itself may of course be regarded as the monohydrate of sulphur trioxide. It is interesting to note that whilst in the reaction SO 3 + H 2 O->H 2 SO 4 21,300 calories are given out, the combination with a further molecule of water, to form SO 3 ,2H 2 O or H 2 SO 4 ,H 2 O, liberates only 6700 calories. Further evidence of the great affinity between concentrated sul- phuric acid and water is to be found in the fact that it abstracts the elements of water from many compounds. In the case of organic compounds such as sugar and paper, which are composed of carbon, hydrogen, and oxygen, the removal of the last two elements leads to the liberation of free carbon, and the substances are said to be charred. The use of sulphuric acid for drying gases has already been repeatedly referred to. Sulphuric acid is also largely used for facilitating actions in which water is one of the products (e.g. the preparation of nitro-glycerine) ; it takes up the water and prevents a reverse reaction by which the product required would be decomposed. In aqueous solution sulphuric acid acts as a dibasic acid. As with other dibasic acids, it ionises in two stages : (i) H 2 SO 4 ^H' + HS0 4 ' (2) the second stage lagging considerably behind the first. In equivalent normal solution (49 grams per litre) about 50 per cent, of the total hydrogen is present in the ionic condition at 18, so that sulphuric acid, though a very strong acid, is not quite so strong as hydrochloric or nitric acids (80 to 82 per cent, ionisation) when solutions of equal concentration in total hydrogen are compared. With univalent metals, sulphuric acid forms salts of the two types, M'HSO 4 and M' 2 SO 4 ; with bivalent metals, the familiar salts are of the type M"SO 4 (e.g. CuSO 4 ,BaSO 4 ). As the ion HSO 4 ' gives a considerable concentration of H* ions, even in moderate dilution, the acid salts, e.g. NaHSO 4 , have an acid reaction in solution. The use of sulphuric acid in preparing other acids from their salts has been repeatedly referred to. Its value in preparing volatile acidi; (e.g. hydrochloric acid, nitric acid), depends upon two factors, its strength and its slight volatility. In virtue of its strength it displaces SULPHUR, SELENIUM AND TELLURIUM 295 a considerable proportion of the other acid from combination, and the mixture can be raised to a high temperature to drive off the volatile acid without much danger of the sulphuric acid volatilizing. Its application in the preparation of certain other acids (e.g. chloric acid, p. 181), depends upon an entirely different property, viz.: the insolubility of its calcium, barium, and lead salts in water. Thus, if a calcium salt of the acid in question is treated with the theoretical amount of sulphuric acid, the acid can be obtained practically pure by filtering off the calcium sulphate. Under certain circumstances sulphuric acid acts as an oxidizing agent, being reduced to sulphur dioxide. The oxidation of carbon, of sulphur, and of copper by sulphuric acid have already been mentioned in connexion with sulphur dioxide. The action of sulphuric acid on metals depends on the concentra- tion of the acid as well as on the nature of the metal. Dilute acid readily dissolves iron, zinc, and magnesium, with liberation of hydro- gen (p. 350): Zn + H 2 SO 4 ->ZnSO 4 + H 2 f, but has no action on copper or lead. Copper is, however, attacked by the concentrated acid on heating, with formation of copper sulphate and sulphur dioxide (p. 283) : It is probable, though not definitely proved, that the first stage in this reaction leads to the formation of hydrogen : Cu + H 2 SO 4 ->CuSO 4 +H 2 , which at the high temperature of the experiment reduces part of the sulphuric acid : H 2 SO 4 + H 2 ->2H 2 O + SO 2 1 . The action of concentrated sulphuric acid on mercury and on silver is similar to that on copper, as is its action on iron : Fe + 2H 2 S0 4 -FeSO 4 +SO 2 +2H 2 O. In the case of iron, therefore (and certain other metals), the products differ according as dilute or concentrated acid is used. Sulphuric acid is a substance of enormous importance in chemical industry. It is used, for instance, in the manufacture of sodium carbonate and of the more important mineral acids, in the manu- facture cf manures, and in the preparation of organic dyes. 296 A TEXT-BOOK OF INORGANIC CHEMISTRY Sulphates The more important points in regard to sulphates have already been mentioned in connexion with the chemical properties of sulphuric acid. All sulphates are soluble in water with the ex- ception of those of calcium, slrontium, barium and lead, which are practically insoluble. The normal sulphates of the alkalis are stable on heating ; the acid sulphates lose water and form pyrosulphates (see below) : 2NaHSO 4 ->Na 2 S 2 O 7 The sulphates of the heavy metals are decomposed by heat, the corresponding oxide and sulphur trioxide being formed : PbSO 4 ->PbO + SO 3 . The insolubility of barium sulphate in water is taken advantage of as a test for sulphates. Any solution containing SO 4 " ions gives, with a soluble barium salt, a white precipitate of barium sulphate, insoluble in acids. PYROSULPHURIC ACID (FUMING SULPHURIC ACID) H 2 S 2 7 (H 2 0,2S0 3 ) Preparation (i) By passing sulphur trioxide into concentrated sulphuric acid H 2 SO 4 + SO 3 ->H 2 S 2 O 7 . The trioxide for this purpose is made by the contact process. (2) The acid was formerly prepared by roasting ferrous sulphate in the air and then distilling from clay retorts (p. 544), the resulting sulphur trioxide being collected in water or sulphuric acid. Properties The pure acid is solid at the ordinary tempera- ture, the crystals melting at 36, and it fumes strongly in the air. The commercial "fuming sulphuric acid" consists of sulphuric acid con- taining varying proportions of dissolved sulphur trioxide. Fuming sulphuric acid is sometimes known as Nordhausen sulphuric acid, because it was formerly prepared at Nordhausen in the Hartz Mountains by distilling roasted ferrous sulphate. Fuming sulphuric acid is used in enormous quantities in the dye industry. The salts of pyrosulphuric acid, the pyrosulphates, are best pre- pared by heating acid sulphates : They are not known in aqueous solution, as on treatment with SULPHUR, SELENIUM AND TELLURIUM 297 water they immediately revert to the acid sulphates. The above reaction is therefore reversible. Persulphuric Anhydride, S 2 O 7 , is obtained in small drops by the prolonged action of a silent electric discharge on a mixture of oxygen and sulphur dioxide or tri oxide. It is unstable, breaking up slowly .even at room temperature into sulphur tri oxide and oxygen it dissolves in water to form persulphuric acid H 2 S 2 8 . Persulphuric Acid, H 2 S 2 O 8 Preparation (i) When concentrated sulphuric acid (50 pe - cent. ) is cooled and subjected to electrolysis with a platinum wire for anode no oxygen is evolved, but the solution round the anode is found to contain a new compound, persulphuric acid, H 2 S 2 O 8 . Its formation is readily understood when it is remembered that in concentrated solution sulphuric acid is ionised aln ost exclusively according to the equation When discharged at the anode, two HSO 4 groups unite to form the acid in question : (2) Persulphuric acid is also obtained under certain conditions when a con- centrated solution of hydrogen peroxide is added to sulphuric acid 2 H 2 S0 4 + H 2 2 ^H 2 S 2 8 + 2 H 2 0. The reaction is reversible. (3) The pure acid can be obtained in colourless crystals (m.p. 60) by the action of the calculated amount of 100 per cent, hydrogen peroxide on chlorosulphonic acid by the method described in connexion with Caro's acid (see below) /OH 2SO./ + H 2 2 ->H 2 S 2 8 + 2 HC1. V-^l 4. Persulbhates Persulphuric acid is very unstable (see below), but the corresponding alkali salts, the persulphates, are fairly stable, and are prepared by electrohsis of the acid sulphates. When a saturated solution of potassium hydrogen s ilphate is subjected to prolonged electrolysis with a large current, the solution being kept cool, potassium persulphate, K 2 S 2 O 8 , separates in colour- less crystals. The persulphuric acid first formed enters into double decom- position wii h the excess of acid sulphate and the slightly soluble potassium persulphate separates. Other persulphates are prepared from this salt or the sodium salt by double decomposition. Propertic: Persulphuric acid and its salts in aqueous solution readily give up oxygen, forming sulphuric acid and sulphates respectively, and are therefore powerful o>. idizing agents : 2H 2 S 2 O 8 + 2H 2 O->4H 2 SO 4 + O 2 . Solid persulphates are moderately stable. Persulphates can be separated from sulphates by taking advantage of the solubility of barium persulphate in water. 298 A TEXT-BOOK OF INORGANIC CHEMISTRY Permonosulphuric Acid (Caro's acid), H 2 SO 5 , is obtained under certain con- ditions by the action of hydrogen peroxide on sulphuric acid : H 2 S0 4 + H 2 2 ^H 2 S0 5 + H 2 , or, better, by adding 100 per cent, hydrogen peroxide to well-cooled chloro- sulphonic acid (p. 300). When the evolution of hydrogen chloride has ceased, the reaction mixture is allowed to warm slowly, and the remainder of the hydro- gen chloride removed by suction : /OH /OH S0 2 < + H 2 2 ->S0 2 < +HC1. xn \O-OH Caro's acid, obtained as above, is a crystalline mass, which melts at 45. In aqueous solution it is a strong oxidizing agent. Disulphur Trioxide, S 2 O 3 , is obtained by adding powdered sulphur to liquid sulphur trioxide. It occurs in bluish-green crystals, readily breaks up into its components, and is at once decomposed by water, the chief products being sulphuric acid and sulphur. Hype-sulphurous Acid, H 2 S 2 O 4 Preparation (i) When zinc acts on sul- phurous acid in aqueous solution, the hydrogen which is presumably the first product of the reaction reduces the sulphurous acid to hyposulphurous acid : 2H 2 SO 3 + H 2 -2H 2 O + H 2 S 2 O 4 , and the solution is found to have powerful reducing properties. (2) On the commercial scale, a solution containing sodium hyposulphite is prepared by the action of zinc on a concentrated solution of sodium hydrogen sulphite : 4NaHSO 3 + Zn->ZnSO 3 + Na 2 SO 3 + Na^C^ + 2H 2 O, and from this solution pure sodium hyposulphite can be prepared if necessary. Properties Hyposulphurous acid in aqueous solution is a yellow, unstable liquid, and, corresponding with the fact that it is the oxyacid of sulphur with least oxygen, is a powerful reducing agent. Both the acid and its salts rapidly absorb oxygen from the air. The mixture described in method of preparation (2) is used in the dyeing industry for reducing indigo to soluble indigo-white. THIOSULPHURIC ACID, H 2 S 2 O 3 The acid itself is extremely unstable, and decomposes immediately it is liberated from its salts. The salts, the thiosulphates, are quite stable, and the normal sodium salt, Na 2 S 2 O 3 , is largely used for photographic purposes under the name sodium hyposulphite, or " hypo." Preparation (i) The thiosulphates are readily obtained by boiling solutions of the corresponding sulphites with flowers of sulphur : Na 2 SO 3 + S->Na 2 S 2 O 3 . SULPHUR, SELENIUM AND TELLURIUM 299 Properties Sodium thiosulphate forms large colourless crystals, containing 5H 2 O, which are readily soluble in water. When a free acid is ;idded to a solution of a thiosulphate, the thiosulphuric acid, which may be assumed as the first product of the reaction, decom- poses nipidly into sulphur dioxide, watet, and free sulphur : Na 2 S 2 O 3 + 2HCl->2NaCl + H 2 S 2 O 3 H 2 S 2 O 3 ->H 2 O + S0 2 + S|. Thiosulphates in aqueous solution act on the insoluble silver halides forming soluble double salts : AgCl + Na 2 S 2 O 3 ->NaAgS 2 O 3 + NaCl. insoluble soluble The us? of sodium thiosulphate in photography depends upon this fact. POLYTHIONIC ACIDS As hae already been mentioned, four oxyacids of sulphur, each containing two atoms o" hydrogen and six atoms of oxygen, are known. The names and formula are Dithionic acid, H 2 S 2 O 6 ; trithionic acid, H 2 S 3 O 6 ; tetrathionic acid, H 2 S 4 O 6 ; and pentathionic acid, H 2 S 5 O 6 . The salts of these acids are quite stable compounds. The respective acids are usually prepared by the action of sulphuri : acid on the corresponding barium salts, They are not known in the free state, and are unstable even in aqueous solution. Dithionic Acid, H 2 S 2 O 6 The manganese salt is prepared by passing sulphur dioxide into water in which manganese dioxide is suspended : 2SO 2 + MnO 2 ->MnS 2 O 6 . From this compound the barium salt, and then the free acid, can be obtained. Trithionic Acid, H 2 S 3 O 6 The potassium salt is obtained by saturating a solution of potassium thiosulphate with sulphur dioxide : 2 K 2 S 2 3 + 3S0 2 ->2 K 2 S 3 6 + S. Tetrathionic Acid, H 2 S 4 O 6 As already explained, the sodium salt is obtained by the a:tion of iodine on sodium thiosulphate : 2Na 2 S 2 O 3 + I 2 ->Na 2 S 4 O 6 + 2NaI. Pentathionic Acid, H 2 S 5 O 6 A solution which contains chiefly pentathionic acid, mixed with other polythionic acids in small proportion, is obtained by passing a limited amount of hydrogen sulphide into a solution of sulphur dioxide : The ultimate products, when excess of hydrogen sulphide is used, are water and sulphur, as already stated. 300 A TEXT-BOOK OF INORGANIC CHEMISTRY OXYCHLORIDES OF SULPHUR Thionyl Chloride, SOC1 2 , is obtained by passing dry sulphur dioxide over phosphorus pentachloride : S0 2 + PC1 5 ->POC1 3 +SOC1 2 . It is a colourless liquid which boils at 78, fumes in the air, and is immediately decomposed by water : SOC1 2 +H 2 0-S0 2 + 2HC1. It may be looked upon as being derived from sulphur dioxide by replacing one oxygen by two chlorine atoms, or from sulphurous acid by displacement of two hydroxyl groups by two chlorine atoms. Sulphuryl Chloride, SO 2 C1 2 , is obtained by direct combination of equal volumes of sulphur dioxide and chlorine under the influence of sunlight, or with the addition of camphor as catalytic agent : SO 2 +C1 2 -SO 2 C1 2 . It is a colourless liquid, which boils at 70, fumes in contact with moist air, and is immediately decomposed by excess of water into a mixture of sulphuric and hydrochloric acids (cf. phosphorus oxychloride, p. 247) : SO 2 C1 2 +2HOH-SO 2 (OH) 2 +2HC1. When very little water is used only one of the chlorine atoms is displaced by hydroxyl, and chlorosulphonic acid is formed : 'OH + HC1. It is evident that chlorosulphonic acid is intermediate in composition to sulphuryl chloride and sulphuric acid. With excess of water it forms sulphuric acid. Graphic Formulas of Sulphur Compounds In order to economize space, the most important of these have already been given in connexion with the enumeration of oxyacids, and they are 0<^ /OH largely self-explanatory. The formula >, C , which has been W? \OH used for sulphuric acid, indicating that the two hydrogen atoms are joined to the sulphur atom through oxygen, is in accordance with the fact that in moderate dilution it is highly ionised as regards both hydrogen atoms. If, on the other hand, one of the hydrogen atoms were joined directly to sulphur, we would anticipate, from analogy with hydrogen sulphide, H - S H, that it would be very slighly ionised. The relationship between the acid and sulphuryl chloride, which can scarcely be written otherwise than as \S\ , explained O^ NCI SULPHUR, SELENIUM AND TELLURIUM 301 in the last paragraph, is also in good accord with the above formula. It should, however, be remembered that such formulas are at best very insufficient representations of the many-sided behaviour of most chemical compounds. The fact that the pure acid so readily splits up into sulphur trioxide and water is best shown by the simple formula, SO 3 ,H 2 (). As regards sulphurous acid, although generally represented thus /OH CKx /H O = S\ , the alternative formula \S\ has also been sue- X)H O^ \OH gested, and the fact that in the latter formula one hydrogen atom is joined directly to sulphur is quite in accord with the very slight ionisation of the second hydrogen atom (p. 286). Space does not admit of a fuller discussion of this interesting question. Crystallography In the foregoing chapters it has been fre- quently mentioned that many substances can be obtained in definite geometrical forms, so-called crystals. Crystals may be obtained on allowing a solution of the substance to cool (e.g. sulphur in carbon disulphide), by allowing a fused substance to solidify (e.g. crystals of prismatic sulphur), by allowing the vapour of a substance to condense (e.g. iodine), and in other ways. When a substance occurs in a non- crystalline form it is said to be amorphous. Sometimes the same substance may exist in a crystalline and in an amorphous form (e.g. sulphur) ; on the other hand, some substances occur invariably in the amorphous form. The most marked feature of a crystal is its regularity of form. The great majority of well-developed crystals show some kind of symmetry. From this point of view the possible type of crystals are considered with reference to imaginary points, planes, and axes, termed respec- tively points of symmetry, planes of symmetry, and axes or lines oj symmetry. A crystal has a point or centre of symmetry when for each face on the crystal there is another parallel to it on the other side of the crystal. A crystal has a plane of symmetry when an imaginary plane can be drawn through it so that the faces are in pairs symmetrically placed with reference to the plane; in other words, the plane cuts the crystal into two parts, one of which is the mirror image of the other. Crystals may have from o to 9 planes of symmetry. A crystal has an axis of symmetry when it can be rotated about an imaginary line (drawn through the centre of the crystal) in such a 3 o2 A TEXT-BOOK OF INORGANIC CHEMISTRY way that each face occupies the position previously taken by another, at least once in a rotation through 360. During the rotation through 360 this periodic recurrence of the same aspect may take place either two, three, four, or six times, and the axes are termed binary, trigonal, tetragonal, and hexagonal respectively. All known crystals may be classified in six systems, based upon the relations of their crystallographic axes. (1) The cubic (regular or isometric} system Crystals belonging to this system can be referred to three axes of equal length, all at right angles to each other. The regular cube (examples : common salt, fluor-spar (Fig. 60)), and the octahedron (example, the alums (Fig. 61)) belong to this class. (2) The hexagonal system Crystals belonging to this class are referred to four axes, three of which, in one plane, are equal in length FIG. 60. FIG. 61. FIG. 62. FIG. 63. and intersect at angles of 120 ; the fourth, known as the principal axis, is at right angles to the plane of the other three, and differs from them in length. Ice, quartz, calc-spar (Fig. 62), and many other important substances belong to this class. (3) The tetragonal system The crystals are referred to three axes, of which two only are of equal length. Example : stannic oxide (Fig. 63). (4) The orthorhombic {rhombic} system The crystals are referred to three axes at right angles, unequal in length. Examples : rhombic sulphur (Fig. 64), potassium nitrate, magnesium sulphate. (5) The monoclinic (monosymmetric) system The crystals belong- ing to this system are referred to three axes of unequal length ; two of these are not at right angles ; the third is at right angles to the plane of the other two. Examples : monoclinic sulphur (Fig. 65), ferrous sulphate heptahydrate, gypsum (Fig. 66). (6) The triclinic (asymmetric) system The crystals belonging to SULPHUR, SELENIUM AND TELLURIUM 303 this system are referred to three axes of unequal length intersecting one another at a point at angles which are unequal and not right angles. There are two divisions of this system ; in one the crystals have a centre of symmetry only, in the other division there is no symmetry. Examples : copper sulphate pentahydrate (Fig. 67), potassium bi- chromate. The more important terms relating to crystals have already been Fm. 64. FIG. 65. FIG. 66. FIG. 67. mentioned. Where the same substance occurs in different crystalline forms it is said to be polymorphous j if only in two forms, dimorphous. Substances of the same crystalline form are said to be isomorphous (p. 120). It sometimes happens that a dimorphous substance is iso- morphous with another dimorphous substance in both its forms ; this is termed isodimorphism. SELENIUM AND TELLURIUM In the group of which oxygen and sulphur are the first elements there are two other elements, selenium (atomic weight =79.1) and tellurium (atomic weight = 127.6), which show many analogies with sulphur, and are most conveniently considered here. Both form compounds with hydrogen, H 2 Se and H 2 Te, of the same type as hydrogen sulphide. Selenium gives an oxide, SeO 2 , and tellurium two oxides TeO 2 and TeO 3 , of the same type as the familiar oxides of sulphur. It will be shown, however, that, unlike the oxides of sulphur, those of tellurium have basic as well as acidic properties. Two oxyacids, H 2 SeO 3 , tellurous acid t and H 2 TeO 4 , telluric acid, are known. SELENIUM Symbol, Se. Atomic weight =79.1. Molecular weight = 158. 2. Occurrence Selenium, though a fairly widely-distributed element, is met with only in snu.ll quantities. It occurs in combination with certain metals, especially lead, mercury , and copper, in the Hartz Mountains ; but the chief sources are certain pyrites (FeS 2 ) in which the sulphur is partly displaced by selenium. When these pyrites are used in the manufacture of sulphuric acid the selenium is converted into the dioxide, and is partly deposited in the flues of the pyrite-burners and partly carried forward into the chambers, where it is reduced to selenium by 3 o 4 A TEXT-BOOK OF INORGANIC CHEMISTRY sulphur dioxide, and collects in the mud on the floor of the chamber. In this " chamber mud " selenium was discovered in 1817 by Berzelius. The name (from r), the moon) was given on account of its analogy with tellurium (from tellus, the earth), discovered not long before. Preparation Selenium is obtained from the flue dust or from the chamber mud by heating first with nitric acid, in order to convert it completely into selenic acid. The latter is then boiled with hydrochloric acid, whereby it is reduced to selenious acid, H 2 SeO 3 , with liberation of chlorine ; the selenious acid is then reduced by means of sulphur dioxide to selenium, which separates in red flakes : H 2 SeO 3 + 2SO 2 + H 2 O->Se + 2H 2 SO 4 . Properties Selenium exists in at least three allotropic modifications: (i)An amorphous red form, obtained, for example, by reducing selenious acid with sulphur dioxide. It is soluble in carbon disulphide, and separates from solution as a second modification (2) Red crystalline selenium, melting at 170 to 180. (3) A third modification, termed metallic selenium, is obtained by heating amorphous selenium to 97, or by quickly cooling melted selenium to 210 and keeping for some time at that temperature. It occurs in grey crystals, is insoluble in carbon disulphide, and conducts electricity. It is a remarkable fact that the electrical conductivity of metallic selenium is greatly increased on exposure to light. According to recent investigations, metallic selenium consists of two forms of very different conducting power in equilibrium, and the equilibrium is displaced by light. Selenium melts at 217 and boils at 680. Above 1400 its vapour density corre- sponds with the formula Se 2 ; at lower temperatures the molecule is more complex. Hydrogen Selenide, H 2 Se Preparation (i) By passing hydrogen over sele- nium heated to 400 : H 2 + Se->H 2 Se. (2) By acting on ferrous selenide with hydrochloric acid : Properties Hydrogen selenide is a colourless gas with an odour like horse- radish, and is more poisonous than hydrogen sulphide. It is moderately soluble in water, and the aqueous solution deposits selenium on exposure to the air owing to oxidation. When passed through solutions of salts of the heavy metals the selenides, being insoluble, are precipitated. Selenium Monochloride, Se 2 Cl 2 , and selenium tetrachloride, SeCl 4> are formed by direct combination of the elements. The former is a brownish-yellow oily liquid ; the latter a white crystalline solid which can be sublimed without decomposition, but undergoes partial dissociation into its elements when heated above 200. OXIDES AND OXYACIDS OF SELENIUM Selenium Dioxide, SeO 2 , the only known oxide of selenium, is obtained by burning selenium in the air. It occurs in long white needles, which sublime (without previously melting) at 310. It dissolves in water, forming selenious acid, H 2 SeO 3 . SULPHUR, SELENIUM AND TELLURIUM 305 Selenious acid, H 2 SeC>3, is formed when selenium dioxide is dissolved in water, and unlike sulphurous acid- can be isolated by evaporating the solution, when it separates in colourless prismatic crystals. It is a dibasic acid, forming normal and acid salts. Reducing agents, for example, sulphur dioxide or stannous chloride, reduce it to selenium : H 2 SeO 3 + 2SO 2 + H 2 O->2HSO 4 + Se. Selenic Acid, HaSeO^ is obtained by the action of chlorine (or bromine) on selenious acid : g + C1 2 + H 2 Cqt H^eC^ + 2 HC1. The actior is reversible, as selenic acid can be reduced to selenious acid by boiling with hydrochloric acid. Selenic acid is, therefore, a more powerful oxidizing agent than sulphuric acid, which oxidizes hydrobromic acid (p. 156), but not hydrochloric acid. The pure acid forms colourless crystals, which melt at 58. The )5 per cent, solution is an oily liquid, similar to sulphuric acid ; its density is 2.6. Barium selenate is as insoluble in water as barium sulphate. TELLURIUM Symbol, Te. Atomic Weight, 127.6 Occurrence Tellurium is a rare ^element. It occurs naturally to a small extent in tire free condition, but more commonly in combination with silver, gold, bismuth, and other metals in Transylvania, Hungary, California, Brazil, and Bolivia. The preparation of tellurium from its ores is a complicated process. Properties Tellurium, like the other elements of this group, occurs in different allotropic modifications. When precipitated from solution, a black, amorphous form is obtained, but on heating it fuses and solidifies as a silvery- white substance with metallic lustre ; its density is 6.24 and it melts at 452. Metallic tellurium conducts heat and electricity, but is not a good conductor. Hydrogon telluride, H 2 Te, is obtained by the action of hydrochloric acid on zinc telluride : ZnTe + 2HCl->ZnCl 2 + H 2 Tef. Properties Hydrogen telluride is a colourless gas with a disagreeable odour, and is very poisonous. It is fairly soluble in water, and the solution deposits tellurium o;i exposure to air. When passed into solutions of salts of the heavy metals, the corresponding tellurides are precipitated. Hydrogen telluride is more easily decomposed into its elements by heat than the corresponding sulphur and seleniu n compounds. Tellurium dichloride, TeCl 2 , and Tellurium tetrachloride, TeCl 4 , are ob- tained by d rect combination of the elements. The former occurs in small, nearly black crystals, the latter forms colourless crystals. Both can be volatilized with- out decomposition at high temperatures. They are decomposed by water, the latter according to the equation but the reaction is reversed when the hydrochloric acid concentration is con- siderable. 20 306 A TEXT-BOOK OF INORGANIC CHEMISTRY OXIDES AND OXYACIDS OF TELLURIUM Tellurium dioxide, TeO 2 , is obtained by burning tellurium in the air or by heating tellurous acid. It is a white crystalline powder, slightly soluble in water, and is at the same time an acidic and a basic oxide (see below). Tellurous acid, H 2 TeO 3 , is obtained by oxidizing tellurium with nitric acid. It occurs as a white powder, slightly soluble in water. It is a weak dibasic acid, forming normal and acid salts with the alkalis which are soluble in water. On the other hand, it forms salts with strong acids, which may be looked upon as being derived from the diacidic base, TeO(OH) 2 , or from the tetracidic base, Te(OH) 4 that is, H 2 TeO 3( H 2 O. Both types of salt are considerably hydrolyzed in solution, in accordance with the rule that when the same substance has both basic and acidic properties it is invariably weak, both as base and as acid. Tellurium trioxide, TeO 3> is obtained by heating telluric acid above 160 : H 2 TeO 4 ->H 2 O + TeO 3 . It is a yellow powder, practically insoluble in water, and its acidic properties are extremely weak. It splits up into the dioxide and oxygen on heating strongly. Telluric acid, H 2 TeO 4 , is obtained by the action of powerful oxidizing agents, for example, chromic acid, on tellurous acid. It is most conveniently obtained by fusing tellurium or tellurium dioxide with potassium carbonate and nitrate : TeO 2 + K 2 C0 3 + i0 2 -K 2 Te0 4 + CO 2 . Barium tellurate is then obtained from the potassium salt by double decomposi- tion with barium chloride and treated with the calculated amount of dilute sulphuric acid : BaTe0 4 + H 2 SO 4 ->BaSO 4 + H 2 TeO 4 . On filtering and evaporating the solution, telluric acid separates in the form of crystals of the composition H 2 TeO 4 ,2H 2 O. Telluric acid dihydrate, H 2 TeO 4 ,2H 2 O, or Te(OH) 6 , is a crystalline powder, very slightly soluble in water; its acidic properties are extremely weak, and it therefore shows practically no analogy to sulphuric acid. Like tellurous acid, it has weak basic properties. Summary of the Oxygen Group Sulphur, selenium, and tellurium, along with oxygen, constitute a family or group of elements which are very similar in chemical behaviour, and, like the halogens, show a gradual variation in physical and chemical properties with increase in atomic weight. The latter statement is illustrated as regards a few physical properties, in the accompanying table (cf. halogens, p. 190): As in the case of the halogens, however, there is a much closer resemblance between the last three elements among themselves than between these elements and oxygen. Sulphur and oxygen resemble each other in being bivalent, so that their compounds with hydrogen SULPHUR, SELENIUM AND TELLURIUM 307 and with the metals are of the same type, and are, in many cases, similar IE behaviour. Property. Oxygen. Sulphur. Selenium. Tellurium. Atomic Weight . 16 32 79-2 127.5 Melting-point -227 114.5 2I 7 452 Boiling-point -181.5 445 680 1400 Density 1.13 (atb.-pt.) 1.96 2.05 4.8 6.2 4 The last three elements are divalent, quadrivalent and sexavalent in their compounds, and therefore the compounds are of similar type, as shown, for example, in the hydrides H 2 S, H 2 Se and H 2 Te ; the oxides SO 2 and SO 3 , SeO 2 , TeO 2 , and TeO 3 , and"the acids of the two types H 2 KO 3 and H 2 EO 4 (E = element). What is much more striking, however, is that the similarity extends to the chemical behaviour of the compounds, as shown in detail in the foregoing pages. The affinity for hydrogen diminishes with increase of atomic weight, as shown by the relative stability of the hydrides. The affinity for oxygen also diminishes, but the affinity for chlorine increases, with increasing atomic weight. Perhaps the most important point, how- ever, is ihat the acidic character of the oxides diminishes with increasing atomic weight, and the oxides of tellurium show also weak basic properties. In this respect, as also in its appearance and in the relative stability of its chloride towards water, tellurium approaches the metals. CHAPTER XXII CARBON CARBON Symbol, C. Atomic weight = 12. Molecular weight unknown. /Recurrence Carbon occurs free in nature in two crystalline ^~-J modifications, diamond and graphite, and also in the amorphous form as charcoal. In combination with hydrogen, it occurs in marsh gas and in petroleum. In combination with oxygen, it occurs as carbon dioxide, which is an invariable constituent of the atmosphere (p. 205), and is found in many natural waters (p. 59) ; and also as car- bonates. Calcium carbonate is met with as chalk, marble and limestone in enormous quantities. The carbonates of magnesium, zinc, barium and of other elements also occur naturally. Carbon is an essential constituent of plants and animals. Coal is formed as the result of the slow decay of vegetable matter, and consists mainly of free carbon, but other substances are always present in greater or less amount (p. 314). ALLOTROPIC MODIFICATIONS As already indicated, carbon occurs in two crystalline forms, diamond and graphite, and also as amorphous carbon. Under the latter name are included all the numerous forms of carbon which have no definite crystalline form. The proof that all these substances are composed of the same element is that carbon dioxide is the sole pro- duct of their combustion in air. DIAMOND Occurrence Diamonds were originally obtained solely from India, being found in alluvial deposits. They have since been discovered in Brazil (about 1727), Australia, South Africa (1867), and also in the United States. They have also been found in meteorites. 308 CARBON 309 Diamonds vary much in size, and are usually comparatively small. A remarkable exception is the Cullinan diamond, discovered in the Premier mine at Kimberley in 1905 ; it originally weighed over 600 grams. Preparation of Artificial Diamonds After it was recog- nized that diamonds are simply crystallized carbon, many attempts were made to prepare them artificially. Great difficulties were, however, met with, but the problem was to some extent solved by Moissan. When carbon is dissolved in melted iron, and the fused mass allowed to cool, the carbon separates in the form of graphite. Moissan modified the method by heating the mixture of carbon and iron to 3000 in the electric furnace (see below), and then suddenly immersing in cold water the carbon crucible containing the fluid mass. A solid crust forms first on the surface, and as iron saturated with carbon expands on solidification, an enormous pressure is thus exerted on the interior partially liquid portion. The iron was finally dis- solved away by acid, and in the residue, which consisted mainly of graphite, minute crystalline particles, having all the properties of diamonds, were found. Some of the diamonds were transparent, others were black. The largest were only about 0.5 mm. in diameter. 1 The mode of formation of diamonds in nature is not thoroughly understood, but they are probably formed by crystallization of carbon in the interior of the earth at high temperatures under enormous pressure. Properties The diamond occurs in crystals belonging to the regular system, the octahedral and cubic forms being most common. When pure it is colourless and transparent, but is very frequently coloured by traces of impurities. When the colour is agreeable, such as blue, red or green, the stones may be as valuable for gems as the colourless variety. The black variety, known as car- bonado or bort, contains up to 2 per cent, of impurities, and is useless as a gem ; but on account of its extreme hardness is used for boring rocks, for cutting glass, and, in the form of a powder, for cutting and polishing precious stones, including diamonds. Diamond is the hardest substance known. It has a very high refractive index, and it is this property of scattering light possessed by the diamond to such a pre-eminent extent which renders it so valuable as a jewel. The average density is about 3.5, but that of the black variety is distinctly smaller, from 3.0 to 3.4. 1 There is reason to suppose that Moissan's artificial diamonds were not entirely free from silicon. 310 A TEXT-BOOK OF INORGANIC CHEMISTRY The diamond does not conduct electricity. It is not attacked by acids. When heated in absence of air to temperatures above 1000, it changes to graphite. When heated in air or oxygen, it readily burns to carbon dioxide, as was first shown by Lavoisier (1772). GRAPHITE (BLACKLEAD) Occurrence This allotropic modification of carbon is widely distributed in nature, being found in Bohemia, Ceylon, Spain, Siberia, California and Borrowdale in Cumberland. Modes of Formation (i) As already stated, graphite is ob- tained when diamond is heated to 1000 in absence of air. Further, when carbon is dissolved in melted iron, it separates in the form of graphite on cooling. (2) As natural graphite is generally very impure, and the substance is of considerable commercial importance, the discovery by Acheson of a cheap method of obtaining it from charcoal (coal or coke) proved of great value. The coke or finely divided coal is heated with a mixture of oxides (including those of iron and calcium) in the electric furnace. Carbides are first formed ; these at the high temperature decompose into graphite, which is deposited in a very pure condition, and the respective metals, which distil over into the colder parts of the furnace. Properties Graphite is a soft, shiny, grayish-black substance, with metallic lustre. The crystals are six-sided leaflets belonging to the monoclinic system. Its density varies from 2.17 to 2.32. Graphite, unlike diamond, is a good conductor of heat and electricity. Most natural graphite is very impure, containing from 40 to 70 per cent- of carbon, and leaving a large proportion of ash when burned. Graphite has the property of breaking off smooth, thin scales on rubbing or pressing, and on this is based its use for making lead pencils, and for lubricating purposes. In making pencils, the finely powdered natural or artificial graphite is mixed with a little clay as binding material, and the semisolid mass formed into narrow threads by squeezing it through a small opening. It was formerly supposed to contain lead, hence the common names plumbago and blacklead often applied to it. Graphite is not readily attacked by air or even by oxygen at high temperatures, and is not affected by the great majority of chemical compounds. For this reason, it is used, along with fireclay, in making the so-called plumbago crucibles. Further, on account of its refractory CARBON 311 nature and its high conducting power, it is largely used as an electrode material for batteries and in electrotyping. When finely divided graphite is treated with potassium chlorate and either nitric acid or a mixture of nitric and sulphuric acids a gray, apparently crystalline substance, the so-called "graphitic acid," is formed. The nature of " graphitic acid " is not understood, and it is doubtless a complicated mixture. On heating, it decomposes with explosion, and a black substance, " pyrographitic acid," remains behind. AMORPHOUS CARBON This term includes all the varieties of carbon which are non* crystalline. The more important are charcoal^ including among other kinds wood charcoal and animal charcoal, also lamp-black or soot, coke and gas or retort carbon. The majority of these sub- stances are prepared by heating carbon compounds in the absence of air. They are all more or less impure. The proportion of carbon varies from about 10 per cent, in animal charcoal to nearly 100 per cent, in carefully purified sugar charcoal. Charcoal is the general name applied to amorphous carbo obtained by heating substances (other than coal) containing carbon in the absence of air, that is, by destructive distillation (p. 214). An important variety is wood charcoal. Wood consists essentially of chemical compounds containing carbon, hydrogen and oxygen, and the process of carbonization consists in the more or less complete removal of the hydrogen and oxygen, along with part of the carbon, as water, compounds of carbon and hydrogen and a great variety of other products. This may be effected by destructive distillation, and also takes place when vegetable matter is allowed to decay in absence of air (see coal, p. 314). The method of carbonizing wood differs according as the other materials, including acetic acid, tar, etc., formed in the process are required or not. In the former case, the wood is placed in retorts, which are heated externally, and no air is admitted. The volatile products which are gases at the ordinary temperature are allowed to escape, and the others are condensed. Wood charcoal remains behind in the retort. The more wasteful process in which the products are allowed to escape is carried on in the forests by charcoal-burners. The wood is piled in heaps, covered with sods and earth, and ignited. The entry of air is carefully regulated so that the wood smoulders away and is finally completely carbonized. This process is also mainly one of destructive distillation, 3 I2 A TEXT-BOOK OF INORGANIC CHEMISTRY the requisite heat being obtained by combustion of part of the material. Wood charcoal retains the shape of the substance from which it is formed. The chief impurities are the mineral substances always present in wood, and also gases. Wood charcoal has a high absorp- tive power for gases, a property shared by some of the other forms of carbon. A much purer charcoal is obtained by carbonizing substances which contain no mineral matter, such as cane sugar. The charcoal ob- tained from the latter substance, after igniting in a stream of chlorine to remove absorbed hydrogen, is practically pure amorphous carbon. Animal Charcoal is obtained by the carbonization of bones in closed iron retorts. It contains only about 8 to 12 per cent, of carbon, up to 80 per cent, of calcium phosphate, and some calcium carbonate and sulphate, buf owing to its very porous character has great absorbing power for gases, colouring matters, etc. By treating it with acids, nearly pure charcoal can be obtained. Lamp black or Soot is obtained by burning substances rich in carbon, such as petroleum, turpentine and naphthalene in a limited supply of air. When freed from hydrogen and other compounds by heating in a stream of chlorine, it is a very pure form of carbon. It is used in making paints and Indian ink, etc. Coke When coal is subjected to destructive distillation, gases, tar and other products volatilize, and coke is left behind in the retort. Coke is very impure, containing most of the mineral matter which forms the ash of coal, as well as sulphur ; the proportion of carbon in coke may reach 80 per cent. It is largely used for metallurgical purposes. Gas or Retort Carbon is found as a hard deposit lining the walls of retorts used in the destructive distillation of coal. It results from the decomposition of volatile products of distillation on the hot walls of the retorts and is a fairly pure product, containing only i to 3 per cent, of ash. Its density is about 2.0, it is a good conductor of electricity, and is largely used for making carbon electrodes. Physical Properties of Ckarcal. AfL**rpti*n Carbon may be regarded as infusible. At the temperature of the electric arc, 3500, it is partly volatilized, and condenses in the form of graphite. The properties of charcoal depend greatly upon its method of preparation and on its degree of purity. For example, the density CARBON 313 of purified lamp-black under ordinary conditions is about 1.78, but can be raised to 1.87 by prolonged heating in absence of air. The fact that wood-charcoal floats in water is due to the presence of absorbed air ; when the latter is pumped out the charcoal sinks. The average density of wood charcoal is about 1.5 ; of sugar charcoal about 1.8. It has already been pointed out that whereas gas carbon and graphite are good conductors of electricity, diamond and ordinary charcoal are practically non-conductors. It is pro- bable that the conductivity depends on the presence of impurities. A remarkable property of charcoal is its power of taking up gases and vapours, and also, to a less extent, colouring matters and other substances from solution. This phenomenon is now called adsorption. It is a surface effect (a physical, not a chemical action), and is very pronounced for charcoal owing to the very great extent of surface for a com- parative small mass. It is very well shown by heating a piece of wood char- coal to drive out the condensed air, and then passing it up into ammonia gas confined in a tube over mercury (Fig. 68) ; the mercury will be observed to rise fairly rapidly ivi the tube. The amount of adsorption is greatest for the most readily liquefiable gases, it A~ \f 117 increases with increase of pressure, / although Henry's law is not followed, and p , is increased by lowering the temperature. Titoff found that i gram of cocoa-nut charcoal adsorbed, at o and 76 cm. pressure in each case, about 135 c.c. of ammonia, 66 c.c. of carbon dioxide, 13 c.c. of nitrogen, and 1.6 c.c. of hydrogen. Dewar found th.it i c.c. of charcoal absorbed at o and 76 cm. pressure 4 c.c., at- 185 and 76 cm. pressure, 35 c.c. of hydrogen corrected to N.T.P. Dewar has taken advantage of this strong adsorptive power of charcoal for gases at low temperatures to remove the last traces of gas from evacuated flasks. For the same reason charcoal is largely employed as an absorbent of deleterious gases in cisterns, etc., and as a trap in sewers. The adsorption from solution of substances of high molecular weight by charcoal has already been referred to. It can be shown 314 A TEXT-BOOK OF INORGANIC CHEMISTRY by shaking up water deeply coloured by caramel with animal char- coal and then filtering, when the water will pass through almost colourless. On account of its highly porous character, animal charcoal is most largely used in this connexion, for example, to remove colouring matters in the refining of sugar. Chemical Properties of Carbon When heated in the air, carbon combines with oxygen to form carbon dioxide, CO 2 , but the temperature at which combination begins varies enormously with the nature of the carbon, and appears to be the higher the greater the density. Wood charcoal catches fire about 300, sugar charcoal, previously ignited in chlorine, at about 450, whilst gas carbon, as already mentioned, can be raised to an extremely high temperature before igniting. When the supply of air is insufficient, carbon mon- oxide, CO, is formed to some extent. Carbon combines directly with hydrogen at 1100-1200 to form marsh gas or methane^ CH 4 . At the temperature of the electric arc another compound, acetylene, C 2 H 2 (q.v.), is formed by direct com- bination of the elements. Carbon combines with most of the metals, in some cases directly, to form carbides. The most important of these compounds, calcium carbide, CaC 2 , will be referred to at a later stage (p. 438). Carbon also combines directly with certain non-metals. With sulphur it forms carbon disulphide, CS 2 (p. 325). With silicon it combines in the electric furnace to form an extremely hard com- pound, carborundtim, CSi (p. 351). Coal As already indicated, coal is formed by the slow decay of vegetable matter in the absence of air. The process of coal formation consists essentially in the removal of the hydrogen and oxygen along with part of the carbon as water, carbon dioxide, marsh gas or methane and other products. The escape of methane from marshes is an illustration of this change at work. Materials representing all stages of the process are familiar to us, the order from the youngest to the most completely changed material being as follows : peat, brown coal or lignite, soft or bituminous coal, and hard or anthracite coal. The average composition of these materials, excluding ash and moisture, as regards carbon, hydrogen and oxygen, and the percentage of ash, is shown in the accompany- ing table. Anthracite coal burns with scarcely any flame or smoke, and is largely used as steam coal. Bituminous coal burns with a smoky flame, yields volatile hydrocarbons (compounds of carbon CARBON and hydrogen) on distillation, and is largely used for preparing coal gas and for ordinary household purposes. Carbon. Hydrogen. Oxygen. Ash. Wood (cellulose) . Peat .... 5 DO 6 6 44 ?A [1.5] c_- ->o Brown coal .... Bituminous coal . . Anthracite .... 67 81 94 5 5 3 27 8 3 ,3 'VJ LS-so i-io] [l-2] The Electric Furnace The electric furnace as used by Moissan is represented in cross-section in Fig. 69. It consists of two blocks of quicklime hollowed out so as to make a reaction chamber and fitting together as shown, the carbon electrodes rest- ing on the lower block. The heat is reflected on the crucible from the rounded upper surface of the chamber. The tempera- ture attainable in the electric furnace naturally depends upon the FIG. 69. voltage and strength of current used, but is limited by the degree of resistance to temperature of the furnace material. At 3000 lime melts, and it can stand a temperature of 2500 only for a limited period. The electric furnace is coming increasingly into use for \arious purposes. In most, if not in all cases the effect is purely thermal ; the furnace is simply a convenient apparatus for obtaining a high temperature. CARBON COMPOUNDS General Carbon is one of the elements invariably present in substances of animal and vegetable origin, and it was formerly supposed that such substances could only be formed as the result of vital activity. These substances were therefore classed as organic, 3 i 6 A TEXT-BOOK OF INORGANIC CHEMISTRY in contrast to substances of mineral origin, termed inorganic. Although we now know that vital activity is by no means in- dispensable to the production of carbon compounds, as many of them have been prepared in the laboratory from their elements, it has nevertheless been found desirable to retain the distinction between organic chemistry and inorganic chemistry, mainly on account of the enormous number and great importance of the carbon compounds. The chief reason for the existence of so many compounds containing this element is to be found in the capacity of carbon atoms to unite among themselves, forming, along with other elements (chiefly hydrogen, oxygen and nitrogen) chains and rings. Although the study of carbon compounds in general belongs to organic chemistry, it is usual to deal with the oxides of carbon and the carbonates in inorganic chemistry ; in fact, the com- pounds just mentioned are not regarded as organic compounds. This term is confined to compounds containing hydrogen as well as carbon. Compounds containing carbon and hydrogen only are termed hydro- carbons. On account of their importance in connexion with com- bustion, and for other reasons, we will deal very briefly with three simple hydrocarbons, methane, CH 4 , ethylene, C 2 H 4 , and acetylene, C 2 H 2 . A compound of carbon and nitrogen, cyanogen, C 2 N 2 , and some of its simple derivatives will also be shortly referred to. COMPOUNDS OF CARBON AND OXYGEN Three oxides of carbon are known : Carbon suboxide C 3 O 2 Carbon monoxide ....... CO Carbon dioxide CO 2 . Carbon dioxide forms a corresponding acid, carbonic acid, H 2 CO 3 . Carbon monoxide may also be regarded as the anhydride of an acid known as formic acid, HCOOH. Carbon suboxide was discovered so recently as 1906. Carbon Suboxide, C 3 O 2 , is obtained by heating malonic acid with phosphorus pentoxide : CH 2 (COOH) 2 +2P2O 5 ->C 3 O2 + 4HPO 3 . It is a colourless, poisonous gas with an acrid odour ; it can be obtained as a colourless liquid which boils at 7 and as a solid melting at - 107. It decomposes slowly at room temperature. Chemically it behaves like malonic anhydride, combining with water to form malonic acid. CARBON 317 CARBON MONOXIDE, CO Preparation (i) Carbon monoxide is formed when carbon dioxide is passed over red-hot carbon : This reaction takes place in a glowing coal fire. The air entering below combines with the carbon to form the dioxide, which on its passage upwards through the glowing fuel is reduced to the monoxide in accordance with the above equation. When the monoxide reaches the surface of the coals and comes into contact with the air, it burns to the dioxide with a characteristic bluish flame. The above reaction is reversible. At 450 the equilibrium mixture contains very little (less than 2 per cent.) of carbon monoxide, at 1000, on :he other hand, the reaction is practically complete in the direction of the upper arrow. (^0 Carbon monoxide is also obtained by reducing metallic oxides, e.g. zinc oZnO + CO. (4f When steam is passed over carbon raised to a white heat (over 1000) a mixture of carbon monoxide and hydrogen is obtained: C + H 2 O->CO + H 2 . The product, known as water gas, is employed commercially as a gaseous fuel. With the addition of certain hydrocarbons rich in carbon, it burns with a luminous flame and is used for lighting purposes. 1 (5) Carbon monoxide is readily obtained from formic acid or a formate by heating with concentrated sulphuric acid, water being abstracted : HCOOH->CO + H 2 0. (6) When under the same circumstances oxalic acid or an oxalate is used, a mixture of carbon monoxide and dioxide is obtained: (COOH) 2 ->CO + CO 2 +H 2 O. The carbon dioxide can be removed by passing the mixed gas through a solution of sodium or potassium hydroxide: C0 2 + 2NaOH->Na 2 CO 3 +H 2 O. 1 Producer gas, obtained by the partial combustion of coke in air, is a mixture of carbon monoxide and nitrogen. 3 i8 A TEXT-BOOK OF INORGANIC CHEMISTRY (7) When the gas is required in considerable amount it is most conveniently obtained by heating potassium ferrocyanide (p. 547) with five times its weight of concentrated sulphuric acid in a large flask: K 4 Fe(CN) 6 +6H 2 SO 4 + 6 + 3(NH 4 ) 2 SO 4 As the equation shows, the reaction is very complicated. Physical Properties Carbon monoxide is a colourless, odour- less, tasteless gas. Under great pressure at low temperatures it is obtained as a colourless liquid which boils at 191 ; on evaporation under reduced pressure it is obtained as a snow-like solid which melts at 208. The critical temperature is 13.9.5 and the critical pressure 35.5 atmospheres (Olszewski). Carbon monoxide is very slightly soluble in water ; the coefficient of absorption is 0.0354 at o and 0.0258 at 20. Carbon monoxide is a very poisonous gas owing to the fact that it combines with the haemoglobin of the blood to form a bright red compound, known as carboxy-haemoglobin. As this compound is much more stable than the compound of haemoglobin with oxygen, the latter is unable to displace carbon monoxide from combination, and the blood thus poisoned is unable to fulfil its function of supplying oxygen to the different parts of the body. Carbon monoxide is formed when combustion is incomplete, as in insufficiently ventilated stoves in which coke or coal is being burned, and as it does not betray its presence by any characteristic smell, many deaths have occurred owing to the escape of the gas into the room. A small proportion of the gas mixed with air causes headache. Chemical Properties The chemical properties of carbon monoxide are readily understood by regarding it as an unsaturated compound (p. 331). Its graphic formula may be written C = O or Ci=O, according as we regard the oxygen as divalent or quadrivalent. In any case, carbon tends to function as a quadrivalent and oxygen as a bivalent element, and therefore carbon monoxide readily adds on one atom of a bivalent or two atoms of a univalent element to form a saturated compound. Carbon monoxide burns in air with a bluish flame to form carbon dioxide : If the gases are perfectly dry, combination does not occur. The equation indicates that two volumes of the monoxide combine CARBON 3 i 9 with one volume of oxygen to form two volumes of the dioxide, and this can be confirmed by exploding a mixture in the above proportions. The gas unites directly with chlorine in sunlight to form carbonyl chloride or phosgene, COC1 2 , and with sulphur to form carbon oxy- sulphide, COS. It also unites directly with certain metals, especially nickel and iron, forming remarkable compounds known as carbonyls. The best-known is nickel carbonyl, Ni(CO) 4 , a colourless, mobile liquid, which boils at 43 under 750 mm. pressure, and is split up into its components on passing through a heated tube. Three iron car- bonyls are known. The pentacarbonyl, Fe(CO) 5 , is a yellow liquid which boils at 102.5 under atmospheric pressure ; the tetracarbonyl occurs in short, dark-green, lustrous, prismatic crystals, and diferronona- carbonyl, Fe 2 (CO) 9 , in orange-red crystals (Dewar and Jones). Carbon monoxide is readily absorbed by an ammoniacal or hydro- chloric acid solution of cuprous chloride at room temperature ; from the solutions a crystalline compound of the formula (CuCl) 2 ,CO,2H. 2 O has been obtained. Carbon monoxide does not combine directly with water to produce formic acid (p. 334), but combines with solid alkalis on heating to form the corresponding formates : KOH + CO->HCOOK. On account of the readiness with which it combines with oxygen, carbon monoxide is a powerful reducing agent. It reduces the oxides of certain metals, e.g. ferric oxide, to the metal (p. 543) : Fe 2 O 3 + 3CO->2Fe + 3CO 2 . The composition and thermochemical behaviour of carbon mon- oxide are referred to in connexion with carbon dioxide (p. 323). CARBON DIOXIDE, CO 2 History Carbon dioxide is the first gas which was definitely recognized as being different from ordinary air. This discovery was made in i;he beginning of the seventeenth century by van Helmont, who obtained the gas by burning wood and by the action of acids on chalk, and showed that it differed from ordinary air in not being a supporter of combustion. Joseph Black of Edinburgh discovered in 1757 that the gas was absorbed by caustic alkalis (alkali hydroxides), and for this reason called \\. fixed air. Finally, in 1781, Lavoisier proved that " fixed air " is a compound of carbon and oxygen. Occurrence Carbon dioxide is a regular constituent of the atmosphere ; on an average 10,000 volumes of air contain 3 volume 320 A TEXT-BOOK OF INORGANIC CHEMISTRY of the gas. It is present in all natural waters (p. 59). Certain mineral waters, for example those at Vichy, Selters, and Carlsbad, contain the gas under considerable pressure, and it escapes with effervescence when the pressure is reduced. It escapes from cracks in the earth, particularly in volcanic regions, as in the Grotto del Cane at Naples, and is given off from active volcanoes. Preparation (i) Carbon dioxide is obtained by burning carbon with free access of air : (2) The carbonates of all metals except the normal carbonates of potassium and sodium decompose on heating, giving the correspond- ing oxide and carbon dioxide. In practice calcium carbonate is gener- ally used : CaCO 3 ->CaO + CO 2 . With proper precautions a very pure gas is obtained by this method. (3) Carbon dioxide is also given off when a carbonate is decom- posed by a strong acid, such as hydrochloric acid : CaCO 3 + 2HCl->CaCl 2 + H 2 O + CO 2 . The nature of this reaction will be readily understood in the light of the considerations on chemical equilibrium already advanced. As carbonic acid is very slightly ionised (p. 267) it is almost completely displaced from combination by hydrochloric acid : Further, carbonic acid has a great tendency to dissociate into carbon dioxide and water : H 2 CO 3 ^H 2 O + CO 2 (ii.). The equilibrium lies very near the left-hand side, and as carbon dioxide is only slightly soluble in water it escapes from the system, and the reaction proceeds according to equation (i.) in the direction of the upper arrow. The reader should himself trace out the mechanism of other re- actions in the same way. This method of obtaining carbon dioxide is the most convenient for laboratory purposes, lumps of marble or limestone and dilute hydrochloric acid being used in a Kipp's apparatus in the ordinary way (p. 35)- (4) When a sugar, such as glucose, C 6 H 12 O 6 , undergoes fermentation CARBON 321 m the presence of yeast, ordinary alcohol, C 2 H 5 OH, and carbon dioxide are the main products : C 6 H 12 O G ->2C 2 H 6 OH + 2CO 2 . (5) Carbon dioxide, along with water and other products, is formed when organic substances (sugar, petroleum, oils, the materials of candles, e:c.) are completely burned in air or oxygen ; or when they are heatec with compounds which readily yield oxygen, for example, copper ox;de. Physical Properties Carbon dioxide is a colourless gas, with a distinct slightly pungent smell. As its density referred to hydrogen is about 22, it can be collected by upward displacement of air. It can be condensed to a colourless liquid, which boils at -79. At -30 the vapour pressure of the liquid is 14 atmospheres, at o 35 atmos- pheres, and at 20 58 atmospheres. The gas is sold commercially in strong iron cylinders, and is used in the preparation of aerated waters and for ot-ier purposes. As the above numbers show, it exerts con- siderable pressure on the walls of the containing vessels. A certain proportion of the commercial product is obtained by collecting the gas escaping from fermentation vats and compressing it into cylinders. The critical temperature of carbon dioxide is 31.4; its critical pressure 73 atmospheres. Liquid carbon dioxide is miscible with alcohol and with ether, but does not mix with water. When the liquid is allowed to escape from a small orifice (by partially opening the valve of the cylinder) into a canvas bag, the evaporation of part of the liquid takes up so much heat that the remainder solidifies and is retained by the bag. Solid carbon dioxide is a white, snow-like substance of density 1.5, which under atmos- pheric pressure passes directly into vapour without melting. This is owing to the fact that the temperature at which it exerts a vapour pressure of I atmosphere, -78, lies below its melting-point under these conditions. Under a pressure of 5.1 atmospheres it melts at - 56.4. Mixtures of solid carbon dioxide with ether or with alcohol are very useful cooling agents ; in both cases a constant tempera- ture of -So is obtained. The solid dioxide is now obtainable in London a; is. per Ib. At room temperature water dissolves about its own volume ot carbon dioxide. The "absorption coefficient" for different temperatures is as follows : 1.713 at o, 1.194 at 10, 1.019 at 15, 0.878 at 20, 0.665 at 30, and o. 530 at 40- Up to 3 or 4 atmospheres the solubility follows Henry's law, but at higher pressures the gas is less soluble than the 21 322 A TEXT-BOOK OF INORGANIC CHEMISTRY law indicates. This is shown by the following results, obtained by Wroblewski at 1 2.4 : Pressure (atmos.) i 5 10 20 30 Solubility 1.086 5.15 9.65 17.11 23.25 A solution of this gas in water, prepared under a pressure of 3 to 4 atmospheres, is known as soda-water. When a bottle is opened the excess of gas above that corresponding with the solubility at atmospheric pressure does not all escape; in other words, the solution remains supersaturated. The gas is completely expelled by boiling. The heat of combustion of amorphous carbon (12 grams) in oxygen amounts to 97,650 cal. ; for diamond the value is about 94,300 cal. The heat of combustion of carbon monoxide to the dioxide (at con- stant pressure) is 68,300 cal. It follows, according to Hess's law, that in the reaction 97,650-68,300 = 29,350 cal. are given out, so that much less heat is given out when carbon monoxide is formed from carbon than when the latter unites with a second atom of oxygen to form the dioxide. This may perhaps be explained by supposing that the molecule of carbon is very complex, and that part of the heat is taken up in simplifying the molecule. Chemical Properties Carbon dioxide is an extremely stable substance. Even at 1300 its dissociation, represented by the equa- tion 2CO 2 ^2CO + O 2 , is only 0.06 per cent. It is decomposed to a considerable extent by passing electric sparks through it, the degree of decomposition depending upon the pressure and other factors. It does not support combustion under ordinary conditions, but certain substances having a great affinity for oxygen, such as the alkali metals and magnesium, burn in the gas. If a piece of ignited magnesium ribbon is immersed in carbon dioxide, it continues to burn and carbon is set free. With metals having less affinity for oxygen, such as aluminium, zinc, and iron, the reduction proceeds only to carbon monoxide. Carbon dioxide is used as a fire extin- guisher, as it prevents access of oxygen, and does not itself under these conditions support combustion. The solution of carbon dioxide in water has a slightly acid taste and turns litmus wine-red, indicating the presence of an acid. The acid has not been isolated, but from analogy with its salts, the car- CARBON 323 bonates, we assume that its formula is H 2 CO 3 . It is a very weak dibasic ac^d, being ionised mainly according to the equation The second stage of the ionisation, HCO/^tH' + COg", is so slight as to be pracdcally negligible. Whether carbon dioxide in aqueous solution is partly present as such and partly in the hydrated form, or whether it is all present in the hydrated form, is not definitely known ; but the former alternative is the more probable. Animals brought into an atmosphere of carbon dioxide die on account of the absence of oxygen. Apart from this, however, the gas has a direct anaesthetic action ; and air containing a proportion of the gas much above the average amount should not be inhaled. Composition of Carbon Dioxide and Carbon Mon- oxide Atomic Weight of Carbon It can be shown, by means of the apparatus described in connexion with sulphur dioxide, that carbcn dioxide contains its own volume of oxygen. It follows that the molecule of the dioxide contains one molecule or two atoms of oxygen, and that its formula is C^O 2 . Direct determination of its composition shows that it contains 12 parts of carbon to 32 parts of oxygen, and as no gaseous compound is known whose molecule con- tains less than 12 parts of carbon, the atomic weight of this element must be 12 and the formula for carbon dioxide CO 2 . These considera- tions are confirmed by the observation that the density of the gas is 44. As above indicated, the accurate atomic weight of carbon can be found by determining the exact ratio in which it combines with oxygen to form carbon dioxide. This was done in a classical investigation by Dumas and Stas (1844). A platinum boat containing either diamond or graphite was carefully weighed and placed in a porcelain tube, which was raised to a red heat while carefully purified oxygen was passed through it. The carbon dioxide formed was completely absorbed in bulbs containing potassium hydroxide, which were weighed before and after the experiment, while the amount of carbon burned w&.s obtained by weighing the boat and its contents after the experimen:. The mean of a number of experiments was as nearly as possible : Carbon : oxygen = 3 :8, whence it follows that (for oxygen = 32) the atomic weight of carbon is 12.00. The most trustworthy later deter ninations of this constant have fully confirmed the accuracy of this figure. 324 A TEXT-BOOK OF INORGANIC CHEMISTRY As carbon monoxide combines with half its volume of oxygen to form carbon dioxide, whereas the latter, as shown above, contains its own volume of oxygen, it follows that carbon monoxide contains half its volume of oxygen. Hence, according to Avogadro's hypothesis, each molecule of carbon monoxide contains one atom of oxygen. As the molecular weight of the compound, determined from its vapour density, is 28, it follows at once that its formula is CO. Carbonates Although, as we have seen, carbonic acid, H 2 CO 3 , is unstable, the corresponding salts, the carbonates, are in most cases stable. As carbonic acid is dibasic, it forms with univalent metals two series of salts acid salts, e.g. NaHCO 3 , and normal salts, e.g. Na 2 CO 3 . When carbon dioxide is passed in excess through a solution of sodium hydroxide, sodium hydrogen carbonate (sodium bicarbon- ate) is formed : NaOH + CO 2 ->NaHCO 3 . The salt is, of course, ionised in solution according to the equation NaHCO 3 ^Na- + HCO 3 ', and as the HCO 3 ' ion splits off practically no H' ions, the solution is very nearly neutral. When sodium bicarbonate is mixed with an equivalent of sodium hydroxide, sodium carbonate, Na 2 CO 3 , is formed : The solution of normal sodium carbonate in water is alkaline, because it is partially hydrolyzed, as represented by the lower arrow in the above equation. This hydrolysis is connected with the exceeding weakness of carbonic acid as a dibasic acid (cf. sulphites, p. 286). When carbon dioxide is passed into a solution of calcium hy- droxide, calcium carbonate is at first precipitated according to equation (i) : (i) Ca(OH) 2 + CO 2 ->CaCO 3 +H 2 O (2) but on continuing the passage of the gas a clear solution is finally obtained. Further investigation shows that the solution contains an acid calcium carbonate, Ca(HCO 3 ) 2 , derived from two molecules of carbonic acid by displacing two hydrogen atoms by an atom of calcium. When the solution is boiled, the reaction is completely reversed in the direction of the lower arrow, the calcium carbonate CARBON 325 being rep-ecipitated. The calcium carbonate present in solution in natural waters occurs as the soluble acid carbonate, and is partly responsible for the hardness of such waters. The insoluble carbonates of certain other metals, e.g. magnesium, are also held in solution by excess of carbon dioxide. The carbonates of the alkali metals are readily soluble in water, those of the other metals are nearly all insoluble, and are therefore prepared by double decomposition. The normal alkali carbonates are stable even at high temperatures ; the bicarbonates lose carbon dioxide on heating, and the normal carbonates are left : 2NaHCO 3 ->Na 2 CO 3 + H 2 O + CO 2 . The carbonates of most other metals readily yield the oxide and carbon dioxide on heating. Percarbonie Acid, HoC-p,, (probably HO-CO'OO-CO-OH) is obtained as a salt by elect rolysis of a concentrated solution of an alkali carbonate or bicarbonate. The free acid is Unstable, decomposing into carbon dioxide and hydrogen peroxide. CARBON BISULPHIDE, CS 2 Preparation Carbon disulphide is obtained when the vapour of sulphur is passed over heated carbon (coke or charcoal). The method now most largely used on the commercial scale is due to Taylor. The heating is effected electrically. Coke is fed in between the electrodes, and sulphur is introduced just below. The heat volatilizes the sulphur, which passes over the glowing coke forming vapours of the disulphide, which are led off and condensed. The process is a continuous one. The disulphide can be purified by shaking with metallic mercury and mercuric sulphate and then distilling. Properties Carbon disulphide is a colourless liquid with an agreeable aromatic odour when pure, but the commercial article has a very disagreeable odour, owing to impurities. Its density is 1.264 at 20 ; it boils at 46.3. It burns in air or oxygen with a blue flame to carbon dioxide and sulphur dioxide : CS 2 + 3O 2 ->CO 2 -f-2SO 2 , and its vapour forms a highly explosive mixture with air. The equilibrium between carbon and sulphur at ordinary temperatures lies very near the right side of the equation, so that the disulphide is an unstable substance. The shock due to 326 A TEXT-BOOK OF INORGANIC CHEMISTRY the explosion of fulminating mercury causes it to split up explosively into its components. The rate of the reaction represented by the upper arrow is, of course, increased by raising the temperature. The disulphide is an endothermic compound, the thermochemical equation representing its formation being as follows : C + S 2 ->CS 2 (liquid) -22,300 cal., so that its tendency to decompose into its elements is readily under- stood (p. 147). Carbon disulphide is an excellent solvent for fats, oils, bromine, iodine, sulphur, phosphorus, and other substances, and is employed commercially in vulcanizing rubber, for extracting essential oils from plants, etc. When potassium sulphide is heated with carbon disulphide, potassium thiocarbonate, K 2 CS 3 , is obtained. From the latter sub- stance by the action of dilute acids, thiocarbonic acid, H 2 CS 3 , is liberated as a yellow, oily liquid, which decomposes readily into carbon disulphide and hydrogen sulphide. The close analogy of these compounds, CS 2 , K 2 CS 3 , and H 2 CS 3 , with the oxygen com- pounds, CO 2 , K 2 CO 3 , and H 2 CO 3 , is evident, and form a further illustration of the statement (p. 306) that oxygen and sulphur belong to the same family of elements. Carbon Oxysulphide, COS, prepared by passing carbon monoxide and sulphur vapour through a tube at a moderate tem- perature, is a colourless, odourless gas, which can be converted into a liquid boiling at -47. It burns in the air with a blue flame to carbon dioxide and sulphur dioxide. It is readily soluble in water; but is slowly decomposed in solution with formation of carbon dioxide and hydrogen sulphide. Carbonyl Chloride and Urea. Carbonyl chloride or phos- gene, COC1 2 , is obtained by mixing equal volumes of carbon monoxide and chlorine and exposing to sunlight (p. 319), or by passing the mixed gases over heated charcoal : COC1 2 . At the ordinary temperature carbonyl chloride is a colourless gas with a suffocating odour ; it can readily be condensed to a liquid which boils at 8. When treated with water it is decomposed into carbon dioxide and hydrochloric acid : CARBON 327 When treated with ammonia it forms urea : /Cl /NH 2 CO< + 4NH 3 ->CO< +2NH 4 C1. \C1 \NH 2 The urea can be separated from the ammonium chloride simul- taneously formed by taking advantage of its solubility in alcohol, from which it separates in colourless crystals. It is clear from its formula that urea may be regarded as being derived from carbonic acid, CO(OH) 2 , by displacement of the two hydroxyl groups by NH 2 or amido groups. Substances derived from acids in this way are called amides. Urea is therefore the amide of carbonic acid, and for this reason is also called carbamide. It can easily be hydrolyzed to carbon dioxide and ammonia: /NH 2 CO<( -fH 2 O->CO 2 + 2NH 3 . \NH 2 Urea is formed in the animal body as the result of the decomposition of nitrogenous compounds, and is excreted in the urine, from which it can be obtained by evaporation. It was formerly considered that urea and other " organic " compounds could only be obtained as the result of vital activity (p. 315). In 1828, however, Wohler made the important discovery that urea can be obtained by heating ammonium cyanate (see below) : NH 4 CNO->CO(NH 2 ) 2 , a compound which can be prepared in the laboratory from its elements. As a result of this and other discoveries it came to be recognized that there is no essential difference between the so-called " inorganic" and " organic" compounds. Cyanogen and allied Compounds Cyanogen, C 2 N 2 , appears to be formed by direct combination of its elements when the electric arc is passed between carbon poles in an atmosphere of nitrogen, but it cannot be detected in the gases drawn from the arc chamber. It is readily obtained by heating dry mercuric cyanide : Hg(CN) 2 ->Hg + C 2 N 2 , and by the action of potassium cyanide on a solution of copper sulphate : 328 A TEXTBOOK OF INORGANIC CHEMISTRY Cyanogen is a colourless gas, which is fairly soluble in water, and burns in air with a lavender-coloured flame to carbon dioxide and nitrogen : Like carbon disulphide and acetylene (g^v.} it is an endothermic compound. In its chemical properties cyanogen shows a remarkable resemblance to the halogens. It combines with hydrogen at high temperatures to form a monobasic acid, hydrocyanic acid, HCN. When passed into an alkali hydroxide, a mixture of alkali cyanide and cyanate is obtained (cf. chlorine, p. 179) : C 2 N 2 + 2KOH->KCN + KCNO + H 2 O. Hydrocyanic Acid, HCN, is prepared by direct combination of cyanogen and hydrogen, or carbon, nitrogen and hydrogen, in the electric arc, but most readily by heating potassium cyanide with sulphuric acid and collecting the distillate in a cooled receiver. Pure hydrocyanic acid is a colourless liquid, with an odour like that of bitter almonds ; it boils at 26.5. It acts as an extremely weak monobasic acid, so that the corresponding salts, the cyanides, are hydrolyzed to a considerable extent in aqueous solution. CyaJicites When potassium cyanide is fused with lead oxide, the latter is reduced to metallic lead, and potassium cyanate, KCNO, is formed : As has already been mentioned, the cyanate, mixed with cyanide, is obtained by the action of cyanogen on potassium hydroxide. The corresponding ammonium salt, NH 4 CNO, a colourless crystal- line substance, changes on heating into urea, as already mentioned. Thiocyanates As the name indicates, the salts are derived from the cyanates by displacement of the oxygen by sulphur. When potassium cyanide is evaporated down with ammonium sulphide, the free sulphur present in the latter (p. 405) converts the cyanide into thiocyanate. The same change can be effected by boiling an aqueous solution of potassium cyanide with sulphur : + S-KCNS. As already mentioned (p. 168, cf. also p. 548), thiocyanates give a deep red colour with ferric salts, owing to the formation of ferric thiocyanate. CARBON 329 SOME SIMPLE ORGANIC COMPOUNDS For the proper understanding of certain phenomena such as the nature of ilame and the carbon cycle in nature, it will be necessary to refer to certain compounds which belong to the organic division of the subject. These compounds are therefore very briefly con- sidered here ; for a full discussion of their properties books on organic chemistry must be consulted. It has already been stated that compounds which contain carbon and hydrogen only are called hydrocarbons. Four hydrocarbons, methane, ethane, ethylene, and acetylene are of importance for our present purpose. METHANE OR MARSH GAS, CH 4 Occurrence The gas escaping in bubbles from marshes, which results from the slow decay of organic matter, contains a large pro- portion (usually more than 80 per cent.) of methane, hence the name marsh gas. It is also produced in large amount during the forma- tion of the coal measures, and is enclosed, often under considerable pressure, in the coal. It escapes during the working of the coal seams, and is the chief constituent of the fire-damp which forms a dangerously explosive mixture with air. It is also the main con- stituent of the " natural gas " which escapes from the earth in coal and petroleum districts. Modes of Formation (i) Methane is obtained by direct combination of its elements when hydrogen is passed over sugar charcoal heated in a porcelain tube to 1100-1150. (2) By passing carbon disulphide vapour and hydrogen sulphide over heated copper : (3) By heating dry sodium acetate with sodium hydroxide : Properties Methane is a colourless, odourless, tasteless gas. It burns i.i excess of air or oxygen with a slightly luminous flame to carbon dioxide and water : and forms a highly explosive mixture with air. On account of its con- nexion wi.:h explosions in coal mines the properties of this mixture 330 A TEXT-BOOK OF INORGANIC CHEMISTRY have been frequently investigated from this point of view. It appears that a mixture of air and methane containing as little as 2 per cent. of the latter is dangerous. Chemically methane is a very inactive substance. It is, however, acted on by chlorine and bromine. In diffused daylight the hydrogen atoms are slowly displaced in successive stages by the former gas, according to the equations CH 4 + C1 2 ->CH 3 C1 + HC1 1 2 + HC1, and so on till finally all the hydrogen is displaced. Reactions of this type are called substitution reactions, so that the compound CH 3 C1 is a substitution product of methane. Such reactions are very frequent in organic chemistry. ETHANE, C 2 H 6 Preparation Ethane can readily be obtained from the sub- stitution product of methane, CH 3 Br, by treatment with metallic sodium : 2CH 3 Br + 2Na->2NaBr + C 2 H C . Properties Ethane is a colourless gas, which resembles methane very closely in all its physical and chemical properties. From its mode of formation it may be assumed that its formula is i.H 3 C-CH 3 , the carbon atoms being quadrivalent and the hydrogen -ii atoms univalent. The univalent group CH 3 , which occurs very frequently in organic compounds, is called the methyl group, the group CH 3 'CH 2 or C 2 H 5 is called the ethyl group. The compound C 2 H 6 C1, which is a substitution product of ethane, is called ethyl chloride, as it contains the C 2 H 6 or ethyl group. When ethyl chloride is heated with alkali hydroxide, the chlorine is displaced by the univalent OH group, forming the compound C 2 H 5 OH, which is ordinary alcohol : C 3 H 6 C1 + KOH->C 2 H 5 OH + KC1, The term alcohol is a general one, applied to compounds which are derived from hydrocarbons by the displacement of one or more atoms of hydrogen by a corresponding number of hydroxyl groups, and to distinguish it from other alcohols the compound C 2 H 5 OH is known as ethyl alcohol. CARBON 331 ETHYLENE, C 2 H 4 Preparation (i) This hydrocarbon is obtained by heating ethyl alcohol with concentrated sulphuric acid to 150 : The gas is purified by bubbling it through a solution of sodium hydroxide, and may be collected over water. From the above equa- tion it would appear that the process is one of simple dehydration, but in reality it is probably more complicated. (2) A much purer product is obtained by heating alcohol with con- centrated phosphoric acid at 200 to 220. The final result is dehydra- tion of alcohol, as represented by the above equation. Properties Ethylene is a colourless gas with an agreeable character stic odour. It can be condensed to a colourless liquid which boils at -- 103.4. It burns in air with a luminous flame to carbon dioxide and water, and is one of the constituents to which ordinary coal gas ewes its luminosity. Ethylene combines directly with two atoms of hydrogen, when the gases are passed through a heated tube, to form ethane, and it also combines directly with the halogens to form halogen derivatives of ethane, e.g. C 2 H 4 Br 2 . In these respects ethylene behaves as an unsaturaicd compound. As regards its formula, only two hydrogen atoms arc available for each atom of carbon, and in order to show carbon a:; quadrivalent, we may assume that the carbon atoms are attached by two valencies instead of by one, as in ethane, thus : H\ /H >C = C< . Such compounds, however, in which carbon atoms H/ \H are joined by more than one bond, have a great tendency to combine with elements or compounds with formation of substances which con- tain carbon atoms joined by single bonds only. This is the explana- tion of the tendency of ethylene to add on hydrogen, the halogens, etc., as ir.entioned above. Organic compounds which contain only carbon atoms joined by single valencies are said to be saturated; they have no additive charactei. Methane and ethane belong to this class. Compounds containing carbon atoms united by two or three bonds are said to be unsaturated; they are characterized by their additive character. Ethylene is an excellent example of an unsaturated compound. 332 A TEXT-BOOK OF INORGANIC CHEMISTRY ACETYLENE, C 2 H 2 Preparation (i) Acetylene is formed by direct combination of its elements in the electric arc : C 2 +H 2 ->C 2 H 2 . (2) It is obtained by the action of water or dilute acids on calcium carbide (p. 438) : CaC 2 +2H 2 O->C 2 H 2 + Ca(OH) 2 . (3) Acetylene is formed when coal gas burns in an insufficient supply of air (cf. p. 343). Properties Acetylene is a colourless gas with a disagreeable odour, and when inhaled in considerable amount is poisonous. It is soluble in its own volume of water at room temperature. It is a highly endothermic compound, the thermochemical equation ex- pressing its formation being as follows : 2 -53,2oo cal. We may therefore expect it to be an unstable substance, and as a matter of fact, when under increased pressure, a shock such as that produced by exploding mercury fulminate decomposes it into its elements with a violent explosion. It burns under ordinary circum- stances with a very luminous but somewhat smoky flame. With a special form of jet, however, so arranged as to secure free access of oxygen, it gives a very white luminous flame, and is therefore largely used for illuminating purposes. It is one of the illuminating con- stituents of coal gas (see below). The graphic formula for acetylene is HC^CH. It is still more unsaturated than ethylene, and combines directly with hydrogen, the halogens, etc. Its most characteristic property is the formation of a red precipitate of copper acetylide, C 2 Cu 2 , when the gas is bubbled through an ammoniacal solution of a cuprous salt. Copper acetylide is endothermic, like acetylene itself, and highly explosive. Yellowish-white silver acetylide, C 2 Ag 2 , is precipitated when acetylene is passed through an aqueous solution of a silver salt. When acetylene is passed through a red-hot tube, it is partially polymerized to benzene, C 6 H C , a hydrocarbon also present in coal gas. Coal Gas As the chief constituents of coal gas have just been described, it will be convenient to deal here with its preparation and composition. When coal is subjected to dry distillation in closed iron retorts, the chief products are (i) coal gas; (2) coal tar and (3) CARBON 333 ammoniacal liquor, which are condensed in the tubes leading from the retorts ; (4) coke, which remains behind in the retort. The coal gas is freed from a number of impurities, sulphur compounds, which are particularly deleterious, being removed by passing it over a mixture cf chalk and ferric oxide, and is C9llected over water. It is a mixture of gases, which may be divided into three groups : Illuminat- ing comt>'tuents, heating co?istituents and impurities. The more im- portant members of the three classes are as follows : Illuminating I Ethylene, acetylene, benzene and other unsaturated constituents / hydrocarbons. Heating ^ n constituents / H y dr S en > methane, carbon monoxide. Impurities Carbon dioxide, nitrogen, (hydrogen sulphide). The composition of coal gas depends greatly on the nature of the coal used and on the temperature to which it is heated. The average composition in volumes percent, is shown in the accompanying table. Hycrogen . . 49 Illuminants . . 4 Methane ... 34 Nitrogen ... 4 Carbon monoxide . 8 Carbon dioxide . . I Good gas should be practically free from sulphur compounds. As coal gas is largely used as a fuel as well as for lighting pur- poses, its value for heating purposes, so-called calorific power, is in many respects as important as its illuminating power. The calorific power is determined by burning a definite quantity in a calorimeter, and passing the products of combustion through tubes immersed in water, so that all its heat is given up to the water. Ethyl Alcohol Preparation (i) As already mentioned, ethyl alcohol is obtained by the action of potassium hydroxide on the ethyl halides : C 2 H 6 Br + KOH->C 2 H 5 O H + KBr. (2) It i c , however, always prepared on the large scale by fermenta- tion of sugar. To an aqueous solution of sugar (obtained from various sources) yeast is added, and the mixture kept about 25. The stu^ar is slowly decomposed into alcohol and carbon dioxide : C C H 12 C ->2C 2 H 5 OH +2C0 2 . As a result of recent investigations, the action of the yeast in this process is now fairly well understood. It gives rise to an organic 334 A TEXT-BOOK OF INORGANIC CHEMISTRY catalytic agent, zymase, under the influence of which the reaction represented by the above equation takes place. It is of interest to note that at present no other method of effecting the above change is known. Properties Ethyl alcohol is a colourless liquid with a characteristic odour ; it boils at 78. It is the active constituent of beer, wines, whisky, brandy and other beverages prepared by "alcoholic" fermentation. When acted upon by oxidizing agents, ethyl alcohol loses two atoms of hydrogen, and a compound CH 3 CHO, belonging to the class of aldehydes, is obtained. On further oxidation, an atom of oxygen is taken up and acetic acid, CH 3 COOH, is formed. Aldehydes and Acids The steps in the oxidation of ethyl alcohol, referred to in the previous paragraph, may be represented as follows : CH 3 'CH 2 OH -> CH 3 'CHO -> CH 3 'COOH Ethyl alcohol Acetic aldehyde Acetic acid. It is shown in books on organic chemistry that all aldehydes contain the CHO group of elements, and the general formula for alde- hydes may therefore be written R-C H where R is a univalent atom or group, such as the CH 3 or methyl group. Aldehydes have the property of taking up an atom of oxygen and forming compounds of the general formula R'COOH, where R has the same meaning as above. Substances of this type are acids, and just as all nitrates contain the NO 3 group, and all aldehydes the CHO group, so all organic acids contain the uni- ^O valent-C O-H or Carboxyl group. An acid containing one such group is a monobasic acid, e.g. CH 3 COOH, and the hydro- gen of the carboxyl group can be displaced by metals with formation of salts, e.g. CH 3 COONa, sodium acetate. From considerations of space, the properties of alcohols, aldehydes and organic acids cannot be described more fully here. It is of interest to mention, however, that when the univalent group R, referred to above, is a hydrogen atom instead of the methyl group, three very important compounds are represented : H'CH 2 OH orCH 3 OH H'CHO or CH 2 O H'COOH Methyl alcohol Formic aldehyde Formic acid. Formic acid has already been referred to in connexion with carbon CARBON 335 monoxide (p. 319). It will be necessary to refer to formaldehyde in the next section. Methyl alcohol and acetic acid are two of the chief volatile products formed by the destructive distillation of wood (P- 3ii)- The Carbon Cycle in Nature It is a familiar fact that the vital activity of animals and plants is attended by the oxidation of complex substances by the oxygen of the air, which is taken in by the lungs and conveyed to the different parts of the body by the red colouring matter of the blood. The ultimate oxidation products, as far as carbon and hydrogen are concerned, are carbon dioxide and water, the former of which is contained in expired air. At the same time a process almost exactly the converse of the above takes place in the green parts of plants under the influence of sunlight, as a result of which water and carbon dioxide, the latter taken in from the atmosphere, are converted into complex substances relatively rich in carbon and hydrogen, and an amount of oxygen equivalent to that of the carbon dioxide taken in is set free. If we assume for simplicity that the substance first formed in the plants is a simple sugar, for example glucose, the above changes may be represented by the following equation : 6CO 2 + 6H 2 O->C 6 H 12 O 6 + 6O 2 (i). These compounds, relatively poor in oxygen, are built up into starch, and further, along with nitrogen and other elements, into the living substance of plants, and finally into that of animals as well, since the latter ultimately depend for their nourishment ex- clusively on plants. As a result of the progressive simplification and ultimate oxidation of organic materials, man and other animals obtain the energy necessary for maintaining the body temperature and for o:her purposes. It appears therefore that the carbon in nature is passing through a continuous cycle, being alternately reduced in the green parts of plants from carbon dioxide to compounds relatively rich in carbon, which are again oxidized l by means of the oxygen of the air with reformation of carbon dioxide and water. The point of fundamental importance, however, is the energy relations of these reactions Again assuming for simplicity that the compound rich in carbon is glucose, the thermochemical equation expressing its oxidation is as follows : C G H 12 O + 6O 2 -^6CO 2 + 6H 2 O 4- 667,200 cal., 1 This oxidation takes place both in plants and animals. 336 A TEXT-BOOK OF INORGANIC CHEMISTRY that is, 667,200 cal. are given out when a mol of glucose is com- pletely burned to carbon dioxide and water. The reaction therefore proceeds spontaneously in the direction of the arrow (p. 147), and to reverse it this amount of energy must be supplied. This energy comes from that of the sun's rays, which, by a method not yet under- stood, is absorbed in the green parts of plants, and stored up in the form of compounds relatively rich in carbon and hydrogen. A little consideration will show us that ultimately we depend almost entirely on the energy thus absorbed by the green parts of plants from the sun's rays. It has already been pointed out that our nourishment comes directly or indirectly from plants, and by the combustion of these materials animals obtain the energy necessary for life. Further, the energy for industrial pro- cesses is obtained mainly by the combustion of wood and of coal. The slow change of wood to coal, being a spontaneous process, is attended by a loss of energy, but coal still represents a great amount of energy in a compact and convenient form. It is not correct to say that the energy resides in the coal exclusively ; it would be equally correct to say that it resides in the oxygen, as the combination of the two is essential before it can be obtained. As, however, oxygen is everywhere available, the only expense incurred is in obtaining the coal. The nature of the chemical process taking place in the green parts of plants has been the subject of numerous investigations, but has not been definitely elucidated. One view is that carbon dioxide and water react to produce formaldehyde and oxygen : CO 2 +H 2 O->CH 2 O + O 2 , the former then undergoing polymerization to form a simple sugar : 6CH 2 O->C 6 H 12 . As a matter of fact, traces of formaldehyde can be detected in the green parts of plants, but the proof that the first steps in assimila- tion are as above is by no means complete. CHAPTER XXIII COMBUSTION AND FLAME /"^ombustion Combustion has already been defined as the ^-^ chemical combination of two substances taking place with sufficient vigour to develop light and heat. The same chemical action ma}- or may not take place with combustion, depending on the conditions. Thus hydrogen and oxygen, in contact with finely divided platinum, combine slowly without sensible rise of tempera- ture or emission of light, whilst under other conditions (p. 37) the same reaction is accompanied by a flame and a great elevation of temperature. It must be remembered, however, that, in accordance with the law of the conservation of energy, the heat given out in a chemical change is independent of the way in which the change is carried out, provided that the final products are the same in each case. As in the processes of combustion with which we are most familiar one of the reacting substances is the oxygen of the air, the substance capable of combining vigorously with oxygen is said to be combust- ible, and the atmosphere is said to be a supporter of combustion. In the same way other gases, such as nitrous oxide and chlorine, which behave towards combustible substances more or less as air does, are said to bo supporters of combustion. Familiar combustible sub- stances are coal, coal gas, hydrogen, sulphur, hydrogen sulphide, etc. A third class of substances, including nitrogen, carbon dioxide, and sulphur dioxide, are neither combustible nor supporters of combustion under ordinary conditions. A little consideration shows us, however, that there is no essential difference between a combustible substance and a supporter of com- bustion. That this must be so is clear when it is realized that the essential feature of the process is chemical combination between two substances, and both are equally concerned in the change. This is well shown by the arrangement represented in Fig. 70, in which air is caused to burn in an atmosphere of coal gas. An ordinary lamp- glass is fi ted at the lower end with a cork carrying a central glass 22 337 338 A TEXT-BOOK OF INORGANIC CHEMISTRY tube and two side tubes, as shown, and on the upper opening is laid a cover with a round hole in the centre, which at the beginning of the experiment is closed by a small lid. The glass is filled with coal gas by the side tubes, and the issuing gas ignited at the lower end of the central tube. The lid is then cautiously removed from the top and the issuing gas lighted. The draught thus caused draws the flame to the upper end of the central tube, and we have now a jet of air burning in coal gas in the interior, and coal gas burning in air at the top. The interchangeability of the terms com- bustible and supporter of combustion is further illustrated by the fact that a jet of hydrogen burns in chlorine as well as a jet of chlorine in hydrogen (p. 90). This experiment also illustrates the fact, already repeatedly referred to, that the term combustion is not confined to reactions in which oxygen is concerned. The chemical aspects of combustion, and its importance for the development of chemistry, have already been fully discussed. Flame When the combination of gaseous substances takes place with the emission of light and a considerable amount of heat, the phenomenon is termed a flame. The restric- tion of the term to gaseous substances should be noted. When carbon combines with oxygen the substance glows and emits light, but there is no flame. The appearance of a flame when certain solid substances combine with oxygen is due to the intermediate produc- tion of gases or vapours. In the case of sulphur FIG. 70. and phosphorus, for instance, the heat of the reaction produces vapours of these elements, which then combine with oxygen, giving rise to a flame. The slight bluish flame which appears over burning coal is due to the intermediate formation and combustion of carbon monoxide. The same explanation applies to the burning of a candle, combustible gases being continuously produced by the heat of combustion. Flames differ very much in general appearance. The flame of hydrogen or of alcohol burning in air is practically colourless, that of carbon monoxide is blue, and that of cyanogen lavender-coloured. Some flames have a very high degree of luminosity. Thus coal gas COMBUSTION AND FLAME 339 under oroinary conditions burns with a bright yellowish flame, and acetylene with a brilliant, almost white flame. The important question of the luminosity of flames will now be considered. Luminosity Of Flames When a thin platinum wire is held in the almost colourless flame of hydrogen burning in air, the wire becomes luminous, and part of the energy of the flame is thus expended in the production of light. The luminosity of the flame of magnesiur i ribbon burning in air is due to the presence of particles of magnesium oxide, which become incandescent at the high tempera- ture of the flame. The luminosity of the ordinary coal-gas flame is also due to the presence of solid particles in this case particles of carbon heated to incandescence. This is easily shown by putting a cold object say a white porcelain dish in the flame, when the particles of carbon are deposited in the form of soot. The carbon particles exist as such of course only in the interior of the flame, where the oxygen supply is deficient ; in the outer part of the flame, where hot oxygen is avail- able, they are completely burned to carbon dioxide. It has already been stated that the luminosity of the coal-gas flame is connected with the presence in coal gas of unsaturated hydrocarbons, such as ethylene and acetylene. We now understand that in the process of combustion these gases must at some stage give rise to free carbon, upon the presence of which the luminosity depends. Methane, on the other hand, burns without the intermediate pro- duction of free carbon. Coal gas of small luminosity can be prepared more cheaply than a gas of high luminosity, and as the former is equally efficient for heating purposes, it is now largely used in commerce, being "enriched" by the addition of unsaturated hydro- carbons \\hen required for illuminating purposes. The tendency to the commercial production of a cheaper, slightly luminous gas has received further support by the extended use of incandescent mantles (see belov\ ). The presence of solid particles is not, however, the sole cause of luminosity in flames. Thus the vapour of carbon disulphide burning in oxygen or in nitric oxide gives a brilliant flp.me, although no solid matter is present. The luminosity in such cases is ascribed by some investigators to " luminescence." 1 The luminosity of flames may often be increased (i) by increasing i The emission of light arising from the temperature possessed by a body is termed incandescence; emission of light; arising from all other causes than tempera- ture is term-id luminescence. 340 A TEXT-BOOK OF INORGANIC CHEMISTRY the pressure; (2) by increasing the temperature. Thus Frankland showed that the flame of hydrogen burning in oxygen, which is practically non-luminous, becomes luminous when the pressure is increased. The effect of temperature on the luminosity is taken advantage of in the so-called recuperative burners, in which the gases are warmed before being introduced into the flame. It is a familiar fact that when lime is raised to a high temperature it glows and emits a very bright light. The same principle has been largely used in recent years in the construction of incandescent mantles, which consist simply of a skeleton of infusible oxides heated to incandescence in a non-luminous flame (Bunsen flame). For reasons as yet unexplained, a mixture of oxides gives much better results than any one oxide. The mixture most largely used consists of 99 per cent, thorium dioxide and I per cent, cerium dioxide. The luminosity of the ordinary gas flame is diminished (i) by cooling it, for example, by previously adding carbon dioxide or nitrogen to the gas ; (2) by preventing the liberation of carbon particles by previously mixing the coal gas with oxygen. The latter method is used in the Bunsen flame (g.v.}. Structure of Flame. The Bunsen Flame The ordinary coal-gas flame (or candle flame) consists of four distinct regions (Fig. 71). (i) The inner dark zone #, consisting of unburnt and practically FIG. 71. FIG. 72. unaltered gas. This may be proved by abstracting some of the gas by means of a glass tube inserted in the flame and lighting it (Fig. 72) COMBUSTION AND FLAME (2) Surrounding the inner zone is the luminous region , in which partial combination occurs. The un- saturated hydrocarbons (and perhaps other hydrocarbons as well) are decom- posed at the relatively high temperature of this region with liberation of carbon particles, the glowing of which is the source of the luminosity of the flame. (3) In the outer slightly bluish margin c, where excess of oxygen is available, complete combustion to carbon dioxide occurs. (4) The bluish region, d, at the base of the flame. In this region, which appears to correspond with the bluish- green sheath of the Bunsen flame, but does not extend over the whole flame, the partia' combination may reach the same stage as in the region of the Bunsen flame just mentioned (see below). The principle of the Bunsen burner, already referred to, is illustrated in Fig. 73. The gas issues from a small jet d at the bottom of the metal tube , and in passing upwards causes a reduction of pres- sure by means of which air is drawn in at the openings on either side of the jet. The air and gas mix on their way up the tube and burn at the top with a practically non-luminous (slightly bluish) flame. Under ordinary circumstances the amount of air taken in is considerably less than that re- quired for complete combustion of the coal gas. The flame maintains its position when the velocity with which the mixture of gas and air issues from the tube exceeds its velocity of inflammation ; if the velocity is less the flame " strikes back " and con- tinues to burn at the base of the tube . The Bunsen flame (Fig. 74) differs mainly from the flame described above in that the luminous region is absent. It consists of three regions (i) the inner cone where the gases from the inner cone finally burn completely to carbon dioxide and water. In the intermediate region there is practically no oxygen and a fairly high temperature, so that substances held in this region are readily reduced. In the outer mantle, on the contrary, there is excess of hot oxygen, and it therefore acts as an oxidizing flame. Investigation of Chemical Changes taking place in Flames It is evidently not quite an easy matter to investigate experimentally the changes taking place in the different regions of the Bun sen or any other flame. An arrangement for withdrawing the gases from different levels of the flame may itself cause disturbances and modify the changes. This difficulty is to some extent got over by means of an arrangement introduced independently by Teclu and by Smithells, known as aflame-separator. It consists of a wide glass tube which is fitted tightly over the upper part of a narrow tube, as shown in Fig. 75. The gas and air are admitted separately at the base of the inner tube, and by suitably regulating the supply of each the external bluish cone is made to burn at the upper end of the wide tube while the green inner cone burns at the top of the inner tube. By means of a side tube, a, on the wide glass tube, the interconal gases can be withdrawn and analyzed, and in this way the nature of the chemical action in the greenish cone has been definitely established. Owing to the experimental difficulties ot the investigation, the exact changes taking place in luminous flames, more FIG. 75. particularly the appearance of free carbon, are not well understood. It was once thought that the liberation of carbon was due to the preferential combustion of hydrogen, but this view is now abandoned. A more plausible explanation is, that at the high temperature of the COMBUSTION AND FLAME 343 flame ethylene dissociates into acetylene and hydrogen and the former into free carbon and hydrogen : CTT _w /" "LT I rj . f* TJ ^ /"* i TT 2 n^ ^v^ 2 n 2 ~r n 2 , ^ 2 ri 2 ^2v^-r-ti. 2 J in other words, dissociation precedes combustion. A modification of this view is that hydrocarbons yield free carbon by thermal decom- position without the intermediate production of acetylene. Bone, on the other hand, is of the opinion that the hydrocarbons first unite with oxygen forming " hydroxylated JJ compounds which, in the absence of excess of oxygen, break down into free carbon and other products. It would lead too far to discuss this interesting subject more fully here. Temperature of Flames The temperature of any particular flame is not quite constant, as it depends on the condition of the combustible substance, on the amount of oxygen supplied, etc. The temperature of the green inner cone of the Bunsen flame coal-gas-air is about 1500, that of the outer cone about 1800. The temperature of the outer cone when oxygen is supplied instead of air is 2200. The ordinary acetylene flame gives a temperature of about 1900, when a Bunsen burner is used it rises to 2500. The hydrogen- oxygen flame gives a temperature of about 2400, that of carbon monoxide burning in oxygen as high as 2800. The num- bers refer to the maximum temperatures observed. It might be anticipated that the temperature of a FIG ^ flame could be calculated from the heat of combustion of the products and their respective heat capacities (specific heats), provided that due allowance is made for the loss of heat by radiation. Taking as a simple illustration the hydrogen- oxygen ilame, the heat of formation of I mol (18 grams) of water i< about 68,400 calories, and as the specific heat of water vapour is about 9 calories per mol, we obtain for the calculated temperature 68,400/9 = 7 600 Even if allowance is made for the heat lost by radiation, this enormously exceeds the observed value. The explanation is that above a certain temperature the combination of hydrogen and oxgen is no Ion ^-er complete, and the higher the temperature the less comple the combination (p. 52). The further combustion is only rendered possible by the loss of heat owing to radiation, and the < temperature is a kind of equilibrium between these different fact. A suitable arrangement for obtaining a flame with two gases, e.g. hydrogen and oxygen, is illustrated in Fig. 76. One gas passes along 344 A TEXT-BOOK OF INORGANIC CHEMISTRY a moderately wide metal tube, the other is conveyed by a narrow metal tube in the axis of the wider tube, and the tubes are so arranged that the gases mix only just before reaching the flame. Ignition Temperature. The Davy Lamp The rate of combination of a mixture of hydrogen and oxygen, like that of other chemical reactions, depends greatly on the temperature. Below a certain temperature the change is comparatively slow, but above this temperature it takes place with explosion. Similarly phosphorus is slowly oxidized in the air at room temperature, but when the tem- perature is sufficiently high, the combination is so vigorous that heat and light are given out ; in other words combustion occurs. The lowest temperature which enables explosion or combustion to take place in a system is termed the ignition temperature. The occurrence of an ignition point in a system is closely connected with the heat given out in a reaction. Above the ignition point, combination of hydrogen and oxygen is so rapid that a large amount of heat is given out in a short time ; this raises a further quantity of the gas above the ignition temperature, this combines with emission of more heat, and so on, combination spreading quickly through the whole system. Below the ignition point, however, combination is slow ; the heat given out is dissipated over a wide area and the temperature never reaches that required for ignition. It does not follow, however, that explosion or combustion will necessarily occur when the temperature is raised at one point above that required for rapid combination ; this again depends upon the thermal behaviour of the system. If the heat given out in combustion is sufficient to raise the system above the ignition temperature, combustion will proceed once it has started without the further application of heat. This can of course only occur in exothermic reactions, and numerous examples have already been given. On the other hand, if the heat of combination is less than that required to raise the system above the temperature of ignition, combination cannot proceed unless energy is continuously supplied from without. The combination of nitrogen and oxygen under the influence of the electric discharge (p. 223) is an illustration of this case. On these considerations is based an alternative definition of the ignition temperature as the temperature at which the initial loss of heat due to conduction, etc.) is equal to the heat evolved in the same time by the chemical reaction (van 3 t HofF). The experimental difficulty of determining ignition temperatures with accuracy is very considerable. When solids are concerned, the state of division is of importance, and in the case of gaseous mixtures COMBUSTION AND FLAME 345 the results are influenced by the positive or negative action of the walls of the containing vessel. Thus the ignition temperature of phosphorus in air is usually given at 35, but we have seen that when deposited in a finely divided form from solution in carbon disulphide, it catches fire spontaneously at room temperature. The uncertainty of the results obtained with gases is well illustrated by the fact that the valuer given in the literature for the ignition temperature of electrolyti: gas (2H 2 : O 2 ) vary from 518 to 845. The difficulties in the case of gaseous mixtures have been largely overcome by an ingenious method due to Dixon, who has succeeded in igniting the gases out of contact with any solid material. One gas was passed through a wide porcelain tube, the other through a narrow tube fixed along the axis of the wide tube. The temperature was gradually raised till ignition occurred ; this took place at a point above the orifice, and special experiments showed, as was to be anticipated, that the ignition temperatures thus determined were independent of the materials of the tubes. The mean ignition temperatures for a few of the commoner gases at normal pressure are as follows : In Oxygen. 1 In Air. Hydrogen. 585 585 Carbon monoxide (moist) . 650 6 5 I Acetylene .... 428 429 Hydrogen sulphide . . | 227 3 6 4 As the table shows, the ignition temperatures of the first three gases are the same in air and in oxygen within the limits of experimental error ; that of hydrogen sulphide is much higher in air. The ignition point of a substance in air may be below room temperature ; it is then said to be spontaneously inflammable. Ex- amples of this have already been given. If a piece of wire gauze be held a little above a Bunsen burner and the gas lighted on the upper side (Fig. 77), the flame does not pass through to the gas on the lower side. This phenomenon is connected with the fact that wire gauze is an excellent conductor of FIG. 77- 346 A TEXT-BOOK OF INORGANIC CHEMISTRY heat, and the explanation usually given is that so much of the heat is conveyed away by the gauze, that the mixture of gas and air on the lower side of the gauze does not reach the ignition tem- perature. On this principle is based the well-known Davy safety-lamp used by miners. It con- sists of an ordinary oil lamp, the flame of which is enclosed in a glass cylinder sur- mounted by a cylinder of wire-gauze (Fig. 78). When this lamp is used in a mine the atmosphere of which contains methane, the flame does not ignite the explosive mixture outside, for the reason given above. The methane, however, passes through the gauze and burns inside, causing considerable alteration in the appearance of the flame. From the amount of alteration thus pro- duced, the proportion of methane in the atmosphere of the mine may be roughly estimated. If part of the flame is acci- dentally blown through the gauze, or even FlG - 7 8 - if the latter becomes strongly heated at any point, explosion may occur. CHAPTER XXIV SILICON AND BORON WITbl the exception of silicon and boron, which are discussed i a the present chapter, all the typical non-metals have now been considered. Silicon and boron do not belong to the same group of elements, but some of their compounds show considerable analogy. SILICON Symbol, Si. Atomic weight=28. 3. Molecular weight unknown. Silicon belongs to the carbon group of elements, and appears to function invariably as a quadrivalent element. As will appear later, there is a close analogy between carbon and silicon, and this applies to the elements themselves as well as to their compounds. Occurrence Silicon does not occur free in nature. In com- bination vvith oxygen, as silicon dioxide, SiO 2 , and with oxygen and different metals, as silicates^ it is very widely distributed. Quartz, flint, and white sand are nearly pure silicon dioxide. Silicates enter largely into the composition of rocks and the different varieties of clay, and are present in the soil, from which they are taken up by plants. Next to oxygen, silicon is the most abundant element in the earth's crust. Preparation Silicon, like carbon, occurs in different allotropic modifications, an amorphous and at least one crystalline form being known. (i) Amorphous silicon is obtained by heating potassium silico- fluoride (p. 350) with metallic potassium : (2) The same modification is obtained in still purer condition by heating silicon dioxide (powdered quartz, sand, or better, the pure precipita.ed oxide) with metallic magnesium : SiO 2 + 2Mg->2MgO + Si. 347 348 A TEXT-BOOK OF INORGANIC CHEMISTRY If excess of magnesium is used, as is advantageous, some magnesium silicide, Mg 2 Si, is formed, but it and the oxide are removed by treat- ment with hydrochloric acid, and fairly pure silicon remains. (3) The same form is obtained by passing the vapour of silicon tetrachloride over heated metallic sodium : The sodium chloride is removed by washing with water. (4) Silicon is obtained in octahedral crystals by heating in a crucible 3 parts of potassium silicofluoride, i part of sodium, and I part of zinc, and dissolving out the other products by treatment with acid. This method of preparation depends upon the fact that the silicon set free by the action of metallic sodium on the silicofluoride dissolves in fused zinc, and separates in crystalline form on cooling. (5) Crystalline silicon is now prepared on the commercial scale by heating quartz sand with coke in the electric furnace : SiO 2 +2C->Si + 2CO. Properties of Amorphous Silicon Amorphous silicon is a brown powder of density 2.35. It burns when heated with air or oxygen to the dioxide, but the latter compound, being non-volatile, coats the particles of the element and retards the action. It com- bines directly with fluorine at the ordinary temperature, with chlorine at 450, with bromine at 450, and with iodine at a red heat, forming compounds of the type SiX 4 . It combines directly with carbon (p. 351), with boron, and other elements when heated in the electric furnace. It is a powerful reducing agent at high temperatures ; for example, it reduces phosphorus pentoxide to phosphorus. It is readily dissolved by liquid hydrogen fluoride, with formation of hydrofluosilicic acid, H 2 SiF 6 , and evolution of hydrogen : The aqueous solutions of acids, even hydrofluoric, have very little action, but it is readily dissolved by a mixture of nitric and hydro- fluoric acids. It is readily dissolved on boiling with sodium or potassium hydroxide, alkali silicate and hydrogen being formed : Si+2NaOH + H 2 O->Na 2 SiO 3 +2H 2 . Properties of Crystalline Silicon This modification occurs in dark grey, opaque octahedral crystals somewhat resembling graphite; the density varies from 2.1 to 2.49. Crystalline silicon con- SILICON AND BORON 349 ducts electricity slightly ; the amorphous orm has no conducting power. The crystals are very hard, readily scratching glass. In chemical behaviour, the crystalline resembles the amorphous form, but being in a less finely divided condition is less active. The crystalline form is considerably more stable towards oxygen than the other modification. Hydrogen Silicides Two compounds of silicon and hydrogen, SiH 4 and Si 2 H , the analogues of methane and ethane respectively, are known. Hydrogen Silicide or Silicomethane, SiH 4 Prepara- tion When magnesium silicide, Mg 2 Si, prepared as already described (p. 348), is treated with concentrated hydrochloric acid, a spontaneously inflammable mixture containing the two hydrogen silicides and free hydrogen is obtained. In order to obtain the pure compounds, the gaseous mixture is passed through a tube dipping in liquid air. when the two silicides are condensed and the hydrogen passes on. The silicides are then separated by fractional distillation (cf. p. 208). Propenies Silicomethane is a colourless gas, which is not spontaneously inflammable at atmospheric pressure, but inflames when the pressure is reduced or when it is slightly warmed ; it then burns with a bright flame to silicon dioxide and water. When passed into the solution of an alkali, there is a vigorous evolution of hydrogen, and an alkali silicate is formed : SiH 4 + 2KOH + H 2 0->K 2 Si0 3 + 4H 2 . When heated above 400, it decomposes into its elements. Silico othane, Si 2 H The preparation of this compound has been described above. Properties Silicoethane is a colourless liquid, heavier than water ; it boils at 52, and the crystals melt at - 138. It is spontane- ously inflammable in air at the ordinary temperature, burning with a brilliant flame, silicon dioxide and amorphous silicon being deposited. It is completely decomposed into its elements on heating to 250. On passing into sodium hydroxide solution, hydrogen is given off and sodium si icate formed : COMPOUNDS OF SILICON WITH THE HALOGENS Silicon Fluoride, SiF 4 Preparation (i) By direct combina- tion of silicon and fluorine (Moissan). 350 A TEXT-BOOK OF INORGANIC CHEMISTRY (2) By heating silicon dioxide (sand) with hydrofluoric acid, or, more conveniently, by heating together calcium fluoride, sand and sulphuric acid. In the latter case, hydrofluoric acid is obtained on heating, and immediately attacks the silicon dioxide : + H 2 SO 4 ->CaSO 4 4HF + SiO 2 ->SiF 4 + 2H 2 O. The etching effect of hydrofluoric acid on glass (p. 153) depends on this reaction. (3) Pure silicon fluoride is obtained by heating dry barium silico- fluoride : BaSiF c ->BaF 2 +SiF 4 . The gas can be collected over mercury in the entire absence of moisture. Properties Silicon fluoride is a colourless gas, with a suffocating odour; it fumes in contact with moist air. It is stable even at high temperatures, corresponding with its great heat of formation, 23,980 cal, and is not decomposed even by electric sparks. When passed into water it is decomposed, with formation of hydrofluosilicic acid, which remains in solution, and silicic acid, which is precipi- tated : 3SiF 4 +3H 2 O->2H 2 SiF 6 + H 2 SiO 3 . As the insoluble silicic acid would soon stop up the tube conveying the fluoride, the experiment is best performed by pushing the end under mercury. As each bubble of gas escapes and comes into contact with water, it is decomposed with precipitation of silicic acid. Hydrofluosilicic Acid, H 2 SiF 6 When the precipitated silicic acid is removed by filtration, an aqueous solution of hydrofluosilicic acid, H 2 SiF G , is obtained. The acid is not known in the pure condi- tion ; when a concentrated aqueous solution is evaporated, silicon fluoride escapes most rapidly and hydrofluoric acid remains. The salts are decomposed in an analogous way on heating (see (3) above). Barium silicofluoride is practically insoluble in water, and this pro- perty is taken advantage of as a test both for barium and for fluosili- cates. Potassium fluosilicate is only slightly soluble in water (i in 833 at 17.5). Fluosilicates are decomposed in a rather remarkable way by alkalis in the cold, alkali fluoride being formed and silicic acid precipitated : SiO 3 + H 2 O. SILICON AND BORON 35 i Silicon Tetrachloride, SiCl 4 Preparation (i) By heating either amorphous or crystalline silicon in a stream of chlorine. (2) By the action of chlorine on a heated mixture of silicon dioxide and carbon: Neither carbon nor chlorine alone have any effect on silicon dioxide under the conditions of the above experiment. The prin- ciple of the method is that by using both substances much more energy is given out than would be the case if silicon or oxygen were products of the reaction (p. 147). The method is also used for other chlorides. Properties Silicon tetrachloride is a colourless liquid with a suffocating odour; it boils at 57, and its density at o is 1.52. It fumes in moist air, and is completely decomposed by water : When silicon is heated in a stream of hydrogen chloride, silicon chloroform, SiHCl s , the silicon analogue of chloroform, is obtained. It is a colourless liquid which boils at 33. Silicon Tetrabromide, SiBr 4 , and the tetraiodide, SiI 4 , can be obtained by methods analogous to those described in connexion with the chloride. The former is a colourless liquid which boils at 130, the iodide occurs in colourless, octahedral crystals, melting at 120.5. Carborundum, CSi, is prepared, according toAcheson, by heat- ing together in the electric furnace a mixture of silicon dioxide (sand), coke and common salt. Properties Carborundum occurs in hexagonal crystals, which are colourless when pure, the density is 3.2. It is not affected by any acid or mixture of acids, is scarcely affected by oxygen even at 1000, but is dec omposed by fused alkalis. Next to the diamond and boron carbide i: is the hardest substance known, and is largely used for grinding purposes. Silicon Dioxide, SiO 2 Occurrence As already indicated, silicon dioxide is very widely distributed in nature ; in fact the greater part of the crust of the earth is composed of this compound, and the salts of the corresponding acids, the silicates. Silicon dioxide occurs in two chief crystalline modifications () quartz with its numerous varieties, including rock crystal, amethyst, smoky quartz, etc. ; () tridymite; and also in the amorphous form as 352 A TEXT-BOOK OF INORGANIC CHEMISTRY flint, opal, jasper, kieselguhr, etc. It is also found in most plants, and is particularly abundant in the stems of grasses, equisetums (horse- tails), bamboo, etc. Modes of Formation (i) It is obtained by burning silicon in air. (2) Pure amorphous silicon dioxide is obtained by igniting silicic acid (q.v.) : H 2 SiO 3 ->H 2 O + SiO 2 . Properties The purest form of quartz, rock crystal, occurs in per- fectly colourless, six-sided crystals of density 2.6. Some of the crystals rotate polarized light to the right, others to the left, and this property is taken advantage of in the construction of polarimeters. Common quartz is white, but not transparent, the varieties already mentioned, including amethyst, smoky quartz, sand-stone, etc., owing their colours to traces of impurities. Owing to its great hardness, quartz (especially in the form of sandstone) is used as a grinding material. Quartz is only stable up to 900 ; above this temperature it changes to tridymite. The latter is found in small crystals in certain rocks, its density is only 2.3. It can be obtained both in hexagonal and in rhombic crystals. The amorphous forms of silicon dioxide, including flint, opal, agate and jasper, owe their colours to traces of impurities, often oxide of iron. Kieselguhr and the silicon dioxide obtained by ignition are practically colourless. Crystalline and amorphous silicon dioxide begin to soften when heated to 1 500, and are completely liquid at 1750. Pieces of chemical apparatus are now obtainable prepared by fusing quartz in the oxy- hydrogen flame. Quartz is remarkable for its extremely low co- efficient of expansion, and therefore a vessel made of silica may be made red-hot and dropped into cold water without risk of fracture. Silicon dioxide is not affected by water, or by acids other than hydrofluoric acid (p. 153). When fused with alkalis .or alkali car- bonates, silicates are formed : SiO 2 + 2Na 2 CO 3 ->Na 4 SiO 4 + 2CO 2 . Silicic Acids As silicon dioxide is an acidic oxide, the exist- ence of silicic acids of the types, H 2 O,SiO 2 or H 2 SiO 3 ; SiO 2 ,2H 2 O, or H 4 SiO 4 , and so on, might be anticipated. As a matter of fact, there appears to be a whole series of such acids of the general formula, ;rSiO 2 ,/H 2 O, where x andj are whole numbers. Up to the present none of them has been obtained in a pure condition ; but salts derived from them, the silicates, form the chief constituents of rocks. SILICON AND BORON 353 The alkali silicates, for example, sodium silicate, Na 2 SiO 3 , can readily be obtained by fusing silicon dioxide with an alkali carbonate, and are easily soluble in water. If hydrochloric acid is added to a solution of sodium silicate, silicic acid is thrown down as a gelatinous precipitate ; but if, on the other hand, a dilute solution of sodium silicate is cautiously added to concentrated hydrochloric acid, a clear solution containing silicic acid is obtained : Na 2 SiO 3 + 2HCl->H 2 SiO 3 4-2NaCl. The silicic acid is not, however, in true solution, but in what is termed colloidal solution. In contrast to substances present in true solution, for example, sodium chloride, colloids do not pass through certain animal and vegetable membranes, and advantage may be taken of this to effect the separation of the silicic acid and sodium chloride obtained in the above experi- ment. The arrangement for this purpose is shown in Fig. 79. Parchment paper or animal bladder is tied tightly over the end of a cylindrical tube ; the mix- ture containing silicic acid is then poured into the vessel, which is suspended in a p IG 7g . larger vessel containing water, as shown. The sodium chloride and excess of hydrochloric acid readily pass through the membrane, and by renewing the water occasionally can be ultimately almost completely removed from the solutior. The clear solution of silicic acid thus obtained is fairly stable, but if concentrated beyond a certain point the acid separates in a gelatinous form. The arrangement just described is termed a dialyser, and the process is called dialysis. Silicic acid is an extremely weak acid, as is shown by the fact that its solutions have no effect upon litmus and are without acid taste, and that the alkali silicates are hydrolyzed very considerably in solution. The acid H 2 SiO 3 is called metasilicic acid, and the compound 23 354 A TEXT-BOOK OF INORGANIC CHEMISTRY H 4 SiO 4 orthosilicic acid. As already mentioned, neither has been definitely isolated, and the silicic acid prepared as described above doubtless contains a mixture of acids. Silicates As already indicated, the silicates occurring as con- stituents of rocks, the majority of which are well crystallized, may be derived from hypothetical silicic acids of the general formula .rSiO 2l yH 2 O. From metasilicic acid, H 2 SiO 3 , are derived sodium sili- cate, Na 2 SiO 3 , and wollastonite, Ca 3 Si 3 O 9 ; from orthosilicic acid, H 4 SiO 4 , we obtain olivine, Mg 2 SiO 4 . Orthoclase, KAlSi 3 O 8 , is de- rived from the acid H 4 Si 3 O 8 , or 2H 2 O,3SiO 2 ; andalusite, Al 2 SiO 5 , from H 6 SiO 5 , or SiO 2 ,3H 2 O ; and meerschaum^ Mg 2 H 4 Si 3 O 10 , from the acid H 8 Si 3 O 10 , or 3SiO 2 ,4H 2 O. Zeolites are hydrated silicates, for example, thomsonite, CaAl 2 Si 2 O s ,3H 2 O. They have the remarkable property of losing water while still remaining a single phase, and the water is therefore not present as water of crystallization, but is simply absorbed. The silicates of which the original rocks of the earth's crust almost entirely consist are continually being decomposed by water and the carbon dioxide of the atmosphere. Perhaps the most powerful disinte- grating influence is due to alternate freezing and thawing. Moisture penetrates into small cracks, and on solidification exerts enormous pres- sure. Carbon dioxide, as it forms with water a stronger acid than silicic acid, displaces the latter from combination, silica being set free. Sand, which is found in large deposits, is formed in this way, the carbonates of the metals being washed into the soil, from which they are largely taken up by plants. Aluminium silicate, which along with alkali silicates occurs in many rocks, happens to be particularly stable, and when the minerals are decomposed under the combined influence of moisture and carbon dioxide, the alkali carbonates are washed away and the aluminium silicate, being deposited in a finely- divided form, constitutes clay. In the laboratory silicates are decomposed by fusing with excess o! alkali carbonate. Sodium silicate and carbonates of the bases are usually formed ; the former can be dissolved out with water, and the latter are usually soluble in hydrochloric acid. Colloidal Solutions We have already seen that a solution of silicic acid is unable to pass through an animal or vegetable mem- brane, whereas dissolved sodium chloride readily passes through. These two substances may be taken as types of two great groups O: compounds, crystalloids and colloids. The great majority of the com pounds so far considered (salts, acids, and bases) belong to the class SILICON AND BORON 355 of crystalloids. They diffuse relatively fast in solution, are able to pass through ani:nal and vegetable membranes, and can usually be obtained without difficulty in crystalline form. Colloids, on the other hand, diffuse very-j slowly in solution, are unable to pass through mem- branes, and when separated from solution have a gelatinous appear- ance, and cannot be obtained in crystalline form. Besides silicic acid, gelatine, gum, caramel or burnt sugar and arsenious sulphide are typical colloids. As diffusive power is proportional to osmotic pressure (p. 195), the above facts may also be stated in the form that the osmotic pressure of colloidal solutions is very small, from which it follows at once that colloids have very high molecular weights. The properties of colloidal solutions may be studied very con- veniently with a solution of arsenious sulphide, which is obtained by dissolving arsenious oxide in distilled water by boiling, and then passing hydrogen sulphide into the cooled solution. A yellow fluor- escent solution is thus obtained, which, when examined by transmitted light, appears homogeneous even under the microscope. When examined \vith the ultra-microscope, however, it is seen that the liquid is full of minute particles in extremely rapid motion. We have here an indication of the fundamental difference between solutions of crystalloids and of colloids ; in the former the particles of the solute are of molecular dimensions, and entirely beyond the range of any microscope In colloidal solutions, on the other hand, the particles are much larger, and in some cases can be seen with the ordinary microscope The precipitated colloid is termed a hydrogel; when in solution it is called a iiydrosol. Many colloids, e.g. arsenious sulphide, are at once precipitated from solution on addition of a small quantity of an electrolyte, while non- electrolytes are without action. Further, when two electrodes are immersed in a colloidal solution and a large difference of potential is established between them, the colloidal particles move either towards the positive or the negative pole. This phenomenon is best accounted for on the view that colloidal particles, like ions, are associated with electrical changes, either positive or negative. Arsenious sulphide and silicic acid are negatively charged ; ferric hydroxide is positively charged. Some colloids when separated in the gelatinous form (hydrogel) again go into colloidal solution (hydrosol) on treatment with water, while others, e.g. ferric hydroxide, are unaffected by further treatment 356 A TEXT-BOOK OF INORGANIC CHEMISTRY with water. The former are termed reversible, the latter irreversible colloids. It is not correct to assume that a " colloidal solution " is necessarily a solution of a substance which appears under all circumstances as a colloid. The colloidal form is also a condition which many substances usually appearing as crystalloids, e.g. arsenious sulphide, platinum. may be made to assume. The typical colloids, however, such as gum, gelatine and silicic acid, never appear in other than the col- loidal condition. BORON Symbol, B. Atomic weight =11.0. Molecular weight unknown. Chemical Relationships Boron acts exclusively as a trivalent element, and belongs to the same group as aluminium (p. 463). It shows comparatively little analogy to aluminium, and. apart from the types of the compounds, it shows much more analogy with silicon. One oxide of boron, B 2 O 3 , and a number of oxyacids derived from it are known. Occurrence Boron does not occur free in nature. In the form of boric acid, H 3 BO 3 , it is contained in the jets of steam so-called soffwni escaping from the earth in volcanic regions in Italy. An important source of boron compounds is the crude borax or tincal, Na 2 B 4 O 7 ,ioH 2 O, found in the United States (chiefly in Nevada and California) and in Tibet. Other minerals containing boron are colemanite, 2CaO,3B 2 O 3 ,5H 2 O, and boracite, 6MgO,8B 2 O3,MgCl 2 . Certain silicates, e.g. tourmaline, also contain boron. Modes of Formation The methods of obtaining boron are closely analogous to those used for obtaining silicon. ' (i) By heating boron trioxide (or fused boric acid) with metallic sodium, potassium or powdered magnesium in a covered crucible : The product is treated with dilute hydrochloric acid, separated from the solution by filtration, and finally washed with water. (2) By passing the vapours of boron trifluoride or trichloride over heated metallic sodium : (3) By decomposing an alkali borofluoride with an alkali metal or with magnesium : NaBF 4 + 3 Na->B + 4NaF. SILICON AND BORON 357 (4) The boron obtained by the methods already described is the amorphous variety. Crystalline boron can be prepared by heating together the amorphous modification and metallic aluminium at 1500 in absence of air for i to 2 hours, cooling, and dissolving out the aluminium with hydrochloric acid. Properties (a) Amorphous Boron This modification forms a brown to brownish-black powder of density 2.45 ; it does not fuse even at the temperature of the electric arc. It combines directly with chlorine at 410, with sulphur at 610, with oxygen at 700 and with nitrogen at 1000. In the latter case boron nitride, BN, is obtained, and therefore when boron burns in air a mixture of oxide and nitride is formed. Boron is not affected by water or by hydrochloric acid, but is oxidized on heating with nitric or with concentrated sulphuric acid : 2B + 6HNO 3 ->B 2 O 3 +6NO 2 +3H 2 O. On fusing with alkalis or alkali carbonates borates are obtained : 2B + 3Na 2 CO 3 ->2Na 3 BO 3 + 3CO. (b) " Crystalline oron"One substance described under this name occurs in black crystals, and appears to be a compound of aluminium and boron, A1B 12 . It is nearly as hard as the diamond. A second impure modification, occurring in yellow transparent crystals, has also been described. Up to the present crystalline boron has not been obtained pure. Boron Hydrides When boron trioxide is fused with excess of magnesium powder in a closed crucible, and the resulting mixture, which contains magnesium boride, Mg 3 B 2 , is treated with hydro- chloric acid, a mixture of gases, containing a large proportion of hydrogen and probably several hydrides of boron, are given off. Among these gases the compound BH 3 is presumably present : [as well as ?, second gas of the formula B 3 H 3 , which separates in the solid form when the original mixture of gases is passed through liquid air. The compound BH 3 has never been obtained free from hydrogen. The mixture of gases has a disagreeable odour, and burns with a green-edgeO flame. 358 A TEXT-BOOK OF INORGANIC CHEMISTRY Boron Trifluoride, BF 3 , is obtained by heating an intimate mixture of calcium fluoride and boron trioxide : + 2B 2 O 3 ->Ca 3 (BO 3 ) 2 or by heating together boron trioxide, sulphuric acid, and calcium fluoride : (i.) CaF 2 + H 2 SO 4 ->CaSO 4 (ii.) B 2 Properties Boron trifluoride is a colourless gas with a suffocat- ing odour. It fumes in the air, and is readily taken up by water. with formation of hydrofluoboric acid, HBF 4 , and boric acid ; the latter, on account of its limited solubility in water, partially crystal- lizes out : 4BF 3 + 3 H 2 0-> 3 HBF 4 +H 3 B0 3 . Hydrofluoboric acid, HBF 4 , may be regarded as being formed by the association of one molecule each of the trifluoride and hydrofluoric acid, BF 3 ,HF, and in fact, is partially decomposed into its com- ponents in aqueous solution : Hydrofluoboric acid is only known in aqueous solution. Its salts, the fluoborates, are stable. The whole behaviour of boron trifluoride strongly recalls that of silicon tetrafluoride. Boron Trichloride, ^^Preparation (\] By direct com- bination of the elements above 400. (2) By strongly heating a mixture of boron trioxide and charcoal and passing chlorine over it : Properties Boron trichloride is a colourless, mobile liquid, which boils at 1 8. It fumes in the air, and in conformity with the general behaviour of the chlorides of non-metals is immediately decomposed by water, hydrochloric acid and boric acid being formed : BC1 3 + 3HOH-B(OH) 3 Boron Nitride, BN, is obtained by direct combination of its components at a white heat, or by strongly heating amorphous boror; in ammonia gas or in nitric oxide. It forms a white amorphous powder, which is stable on heating in air, but is decomposed wher SILICON AND BORON 359 heated with steam at 200 under pressure, ammonia and boric acid being formed : It is also decomposed by fusing with potassium hydroxide : Boron Trisulphide, B 2 S 3 , is obtained by direct combination of its elements or by heating a mixture of boron trioxide and carbon in a current of carbon disulphide. It occurs in small, colourless needles, and is immediately decomposed by water, with formation of boric acid and hydrogen sulphide : B 2 S 3 + 6H 2 0->2HB0 OXIDES AND OXYACIDS OF BORON As already indicated, only one oxide of boron, B 2 O 3 , is known. Corresponding with this oxide are the well-defined oxyacids : Orthoboric acid ...... H 3 BO 3 Pyroboric acid ...... H 2 B 4 O 7 Metaboric acid ...... HBO 2 Boron. Trioxide, B 2 O 3 , is formed when boron is burned in air or oxyger, and is usually prepared by heating boric acid to redness in a crucible : Properties Boric acid is a transparent, glassy, hygroscopic solid, which melts about 580, and can be raised to a very high temperature without appreciable volatilization. On this account it displaces volatile acids from their salts at a high temperature ; carbonates and nitrates are completely, sulphates only partially, decomposed : It combines with many metallic oxides when fused, giving glasses of characteristic colours ; this fact is taken advantage of in testing for certain metals. Orthoboric Acid, H 3 BO 3 Preparation (i) On the com- mercial scale, boric acid is usually obtained from the "soffioni" in 360 A TEXT-BOOK OF INORGANIC CHEMISTRY Tuscany, in which it occurs to a very small extent (less than o.i per cent). The jets are caused to pass through water contained in brick reservoirs built round them, and after the water has been passed through several reservoirs, and has thus taken up a considerable proportion of the acid and become considerably concentrated by evaporation, it is further evaporated by the heat of the escaping vapours and the solution then set aside to crystallize. The source of the boric acid in the jets is not thoroughly under- stood. It probably results from the action of steam on boron nitride or on boron sulphide (q-v.} below the surface of the earth. (2) Boric acid is also prepared by adding sulphuric or hydrochloric acid to a concentrated solution of borax (sodium pyroborate, Na 2 B 4 O 7 ) : Na 2 B 4 O 7 + 2HC1 + 5H 2 O->2NaCl + 4H 3 BO 3 . Properties Boric acid occurs in colourless, lustrous, transparent six-sided leaflets, which are soapy to the touch. It is moderately soluble in water. At 13 100 c.c. of the saturated solution contain 3.84 grams, at 20 4.91 grams, and at 25 5.58 grams of the acid. When the aqueous solution is boiled a little boric acid escapes with the steam. The acid is readily soluble in alcohol, and when the solution is ignited it burns with a green-edged flame. This property is used as a test for the acid and its salts. Boric acid is an extremely weak acid. The aqueous solution colours litmus a sort of port wine colour, instead of the bright red produced by strong acids. Even in the first stage of the dissociation the ionisation is much less than for carbonic acid, and it is therefore impossible, owing to hydrolysis, to obtain corresponding salts orthoborates from aqueous solution. When boric acid is heated for some time at 100, it loses water, with formation of metaboric acid : H 3 BO 3 ->HBO 2 + H 2 O, and at 140 still more water is driven off, pyroboric acid being formed : 4HBO 2 ->H 2 B 4 O r +H 2 O. Metaboric Acid, HBO 2 , is obtained as already mentioned, by heating orthoboric acid to 100. It is a monobasic acid, and a SILICON AND BORON 361 number of corresponding salts, e.g. NaBO 2 , AgBO 2 , Ca(BO 2 ) 2 are known. Pyroboric Acid, H 2 B 4 O 7 , is obtained by heating orthoboric acid to i,.o. When pyroboric acid or metaboric acid is dissolved in water, orthoboric acid is immediately formed. Borates The best known borate is ordinary borax, sodium pyroborate, Na 2 B 4 O 7 ,ioH 2 O. It occurs naturally in tincal (p. 356), and is also obtained commercially by the action of sodium carbonate on colemanite: Ca 2 B 6 O n + 2Na 2 CO 3 ->Na 2 B 4 O 7 + 2 NaBO 2 + 2CaCO 3 . The solution is filtered to remove the calcium carbonate, and on evaporation the borax crystallizes out. The more soluble metaborate remaining in the mother liquor can also be converted into borax by passing carbon dioxide through the solution : On the laboratory scale, borax can be prepared by adding sodium carbonate to a boiling solution of boric acid, evaporating the solu- tion; and setting aside to crystallize : Borax occurs in large colourless prisms. On heating it loses water, swells up. and finally fuses to a clear bead, to which characteristic colours are imparted by certain metallic oxides, such as those of cobalt ana of manganese. If the formula is written thus, Na 2 O,2B 2 O 3 , we can understand that certain basic oxides can combine with part of the boron trioxide, forming definite chemical compounds. Borax is hydrolyzed to a considerable extent by water, and the aqueous solution has therefore an alkaline reaction. As already indicated, metaborates have also been obtained from aqueous solution ; thus silver metaborate, AgBO 2 , and calcium meta- borate, Ca(BO 2 ) 2 ,2H 2 O, have been prepared in this way. The weakness of the boric acids is strikingly shown by the fact that although borax contains a large excess of the acidic oxide, Na 2 O : 2L 2 O 3 , its aqueous solution is alkaline. Moreover, there is evidence that, as is to be anticipated, pyroboric acid is a consider- ably stronger acid than orthoboric acid. CHAPTER XXV CLASSIFICATION OF THE ELEMENTS THE PERIODIC SYSTEM GENERAL PROPERTIES OF THE METALS AND THEIR COMPOUNDS HAVING regard to the enormous number of facts already established with reference to the behaviour of the elements and their compounds, it is evident that some system of classifying the elements which will show their mutual relationships and facilitate the comprehension and utilization of the available data is in the highest degree desirable. Several methods of classification might be suggested. For ex- ample, the elements might be arranged according to their valencies, all univalent elements, for instance, being brought into one group. There are two serious drawbacks to this suggestion. In the first place, elements like potassium and iodine, which clearly have not the slightest analogy, would be brought together. In the second place, most elements have more than one valency, and might there- fore with equal justification be put in two or more groups. The division of the elements into metals and non-metals, already made use of, is much more promising, although by no means satisfactory. The main differences between the two groups will be. referred to in detail later, and be indicated here only very briefly. The chief characteristics of metals are (i) metallic lustre and ability to conduct heat and electricity ; (2) they combine with oxygen to form basic oxides ; (3) when forming constituents of salts in solution, they appear alone as positive ions only ; (4) their chlorides (and other binary halogen compounds) are fairly stable towards water. In contrast to the metals, the typical non-metals (i) have no metallic lustre and do not conduct heat and electricity ; (2) form acidic oxides ; (3) appear free in the ionised condition as negative ions only ; and finally, (4) their compounds with the halogens are readily hydrolyzed by water. It has already been pointed out, and further illustrations will be given later, that although these broad differences are quite sufficient to distinguish between typical metals and non- 362 CLASSIFICATION OF THE ELEMENTS 363 metals, a number of elements appear to stand on the border-line between .hese two great groups. By far the most successful system yet proposed for classifying the elements is based upon arranging them according to their atomic weights. In 1829, Dobereiner pointed out that there are a number of groups of three elements, so-called triads, the members of which have clo^ e chemical analogy, while the atomic weights differ by about 16 or a multiple of that number. Examples chlorine, bromine and iodine; calcium, strontium and barium. In 1864, Newlands showed that when the elements then known were arranged in the order of increasing atomic weights, although the successive elements showed no particular analogy, the eighth element was analogous to the first, the ninth to the second, and so on ; in other words, when the elements are arranged as above stated, there is a periodic recurrence of elements with similar chemical properties. As each such group or period con- tained seven elements, the discovery of Newlands was termed the law of octaves. In 1869, Mendeleeff, on the same basis, but without any knowledge of Newlands' work, developed the periodic system of the elements substantially in the form now accepted. Starting with helium = 4, which was unknown when Mendeleeff first brought forward his classi- fication, we have the arrangement shown in the first line of the accompanying small table He= 4 Li = 7 Gl = 9 B =11 C=i2 N = i4 O = i6 F =19 Ne^2o Na=2 3 Mg=2 4 Al = 2y Si= 2 8 P =31 8=32 1=35.5 in whicl the properties vary regularly from helium to fluorine. The next element, neon = 20, is an inactive gas and is placed below heliunv, sodium then falls into its proper place below lithium, phosphorus below nitrogen (p. 238), sulphur below oxygen (p. 306). It should be clearly realized that no elements are omitted in this arrangement ; the order is strictly that of increasing atomic weight, and the accuracy with which the elements fall into the places to which they would be assigned on purely chemical grounds is most striking. These groups of elements are termed periods, and the first two periods contain 8 elements each. The complete table, arranged on the above principle, is given on p. 364. A new period is started with argon, a third inactive gas, but in this case it is necessary to pass over 18 elements before another inactive gas (krypton), bearing a strong resemblance to argon, is reached . Such a period of 18 elements is termed a long period in o w 5 Bw xc * IN > > Sw a 3- 09 II CQ 1 * i w ... . ^ CT J c ^-3 c -0 00 00 Oo 'C J'C J'H J'C - CLASSIFICATION OF THE ELEMENTS 365 contrast to the two short periods of 8 elements each. The whole table is riade up of two short and five long- periods, but four of the long periods are incomplete, and the last one contains only 3 elements. A word must be said as to the significance of the blanks in the table. After molybdenum =96 the next element is ruthenium =101.7, which, strictly speaking, should come below manganese =55. The essential feature cf this arrangement, however, is the chemical similarity of elements in the same vertical rows. Now, ruthenium shows no analogy whatever with manganese, but shows certain analogies with iron =56. It is therefore placed under the latter element, and a blank is left below manganese, to indicate the position of a hitherto undis- covered clement. At first sight, this policy seems a somewhat arbitrary one, but it has been entirely justified by results (see below). The arrangement adopted in representing the long periods will be evident rom the table. The twelfth element shows some, but not a very close analogy to the second, the thirteenth is distantly related to the third, and so on. The former element of the pair is therefore placed below, but to the right of the latter, to indicate the absence of a close analogy. This system leaves 3 elements in the middle of the first, second and fourth long periods, the position of which presents some difficulty. Mendeleeff put these "transition" elements in a group bv themselves, the so-called eighth group (Group vill. in the table). A study of the elements, arranged as above, shows many striking regularities. Thus the valency with regard to hydrogen increases regularly up to the middle of a short period, and then falls regularly (ultimately to unity) in the latter half of the period, whilst the valency for oxygen increases regularly by units from the beginning up to the end of the period. The valencies in the long periods are not quite so regular, being complicated by the fact that most elements have several valencies, but, generally speaking, each of the halves of a long period behaves like a short period. It follows at once from this arrange- ment that the elements in the same vertical group should have the same valency, and a glance at the table shows that such is the case. The group under o comprises the inactive gases, which do not enter into chemical combination, and may therefore fittingly be regarded as having zero valency. The next group, under I., composed of univalert elements, comprises the alkali metals, which are typical univalent elements, and copper, silver and gold, all of which can act as univalent elements. Similar considerations apply to the succeed- ing groups. The position of the halogens in Group vn. is interesting, 366 A TEXT-BOOK OF INORGANIC CHEMISTRY and appears to be justified to some extent by the existence of com- pounds of the type HC1O4,KIO 4 , etc. (p. 182). The members of the eighth group would be expected to show a valency of 8. This is to some extent borne out by the existence of osmium tetroxide, OsO 4 , and ruthenium tetroxide, RuO 4 (p. 556), but the other elements do not appear to exert so high a valency. Many of the physical properties of the elements, such as the melting-point, the atomic volume, the density, conductivity of the metals for heat and electricity, heat of formation of oxides and chlorides, also vary regularly within such period. For example, the melting-points of the members of the first period gradually rise from helium to carbon and fall again to fluorine, and, similarly, the elements of highest melting-point (iron, cobalt, nickel, etc.) occur in the middle of the long periods. Not only the physical properties, but also the chemical properties of the elements vary regularly within each period. Thus the elements on the extreme left hand are inactive gases, those in the second group (Series A) decompose water and are strongly electro-positive, at the middle of the period the electrical character is much less pronounced (carbon, silicon), and towards the right hand strongly electro-negative elements are found. Besides this variation of properties in the horizontal series, there is a similar, but much less marked variation of both physical and chemical properties in the vertical groups. In the case of the alkali metals, for instance, the melting-points fall from lithium to caesium, and the electro-positive character increases. The statements in the last three paragraphs are summarized in the Periodic Law, due to Mendeleeff, which may be stated as follows : The properties (both physical and chemical) of- the elements and their compounds are periodic functions of the atomic weights. The Periodicity of Physical Properties It appears desirable to illustrate the periodicity of the physical -properties of the elements a little more in detail. (a) The Density of the Elements The density increases regularly from the beginning to the middle of a period, short or long, and then diminishes regularly towards the end of the period. Uncertainty often arises owing to the existence of a number of forms of some elements with different densities, and also as to the proper conditions for comparison. This rule is illustrated in the accompanying small table from the data for the second short period : CLASSIFICATION OF THE ELEMENTS 367 Element Na Mg Al Si P (red) S Cl (rhombic) (liquid) Density 0.97 1.75 2.67 2.49 2.14 2.06 1.33 The same periodic variation occurs in tlje density of the oxides of these elements : Oxide Na 2 O MgO A1 2 O 3 SiO 2 P 2 O 6 SO 3 C1 2 O 7 Density 2.8 3.7 4.0 2.6 2.7 1.9 (b) Tht Atomic Volumes of the Elements In the same way the atomic volumes of the elements that is, the relative volumes occupied by the atomic weight in grams of the elements in the solid state vary in a periodic manner, but in this instance the magnitude dimin- ishes towards the middle of a period and then increases regularly towards the end. The elements of smallest atomic volume therefore occur towards the middle of a period, whether short or long (examples ; boron, aluminium, iron, cobalt, nickel) and those of largest atomic volume .-it the ends of the period. These facts are illustrated graphically in the accompanying diagram (Fig. 80), in which the atomic volumes as ordinates are plotted against the corresponding atomic weights. The periodic variation of this property, resembling the successive swings of a pendulum of ever-increasing amplitude, is very striking. (c) The Melting-points of the Elements The melting-points of the elements rise towards the middle of a period, whether short or long, and fall < gain towards the end. This fact is illustrated by the data for the elements of the second short period : Element Ne Na Mg Al Si P S Cl Melting-point 97 <8oo 657 >i7oo 45 115 -102 The results in this instance are not quite so regular, but the general tendency of the numbers is quite definite. Uses of the Periodic System The periodic system is of use mainly in three ways : (1) As a system of classification which indicates in a fairly satis- factory way the physical and chemical relationships of the elements. (2) For predicting the existence and properties of elements hitherto undiscovered. (3) For enabling us to find the correct values of the atomic weights of elements which do not form volatile compounds. 368 A TEXT-BOOK OF INORGANIC CHEMISTRY CLASSIFICATION OF THE ELEMENTS 369 The first point has already been sufficiently illustrated in the fore- going paragraphs. In the remainder of the book, the periodic system w : ll be taken as the basis of discussion of the metals and their compounds, and its great value will then become apparent. As regards the second point, when the periodic system was first brought forward, there were more blanks 'in the table than there are at the present day, and Mendeleeff not only suggested that the positions of these blanks corresponded with hitherto undiscovered elements, but even foretold the properties of the missing elements from those of the known elements near them in the periodic table. At that time (1871) there was a blank in the table representing an unknown element with an atomic weight somewhat exceeding that of zinc, and Mendeleeff, from the properties of the surrounding known elements, zinc and arsenic in the horizontal and aluminium and indium in the vertical rows, foretold that the unknown element, which he termed eka-aluminium, would have an atomic weight of about 69, that it would be trivalent and form alums (p. 468) like aluminium, that its density would be about 5.9, that it would be more volatile than aluminium, and therefore its discovery might be expected by means of the spectroscope. Four years afterwards, Lecoq de Boisbaudran was led by spectroscopic observations, exactly as Mendeleeff had foretold, to the discovery of an element, which he called gallium, having all the properties predicted for eka-aluminium. Later, two other blanks were filled up by the discovery of scandium (Nilson and Cleve, 1879) and of germanium (Winkler, 1886), which also showed in all respects the properties predicted by Mendeleeff. The use of the periodic system for fixing atomic weights will be readily understood from the foregoing. When the equivalent of the element lias been determined, it is usually possible to decide which multiple of it is to be taken, as there will in general be only one position ',n the table into which the element can be satisfactorily fitted. The most familiar example is the controversy with regard to the atomic weight of beryllium (glucinum). From analysis of the chloride, \he equivalent of this element was found to be 4.5. If the metal is bivalent, the atomic weight must be 9.0 and the formula of the chloride is BeCl 2 , if, on the other hand, it is trivalent, the atomic weight is 13.5 and the formula of the chloride BeCl 3 . Mendeleeff pointed out that a trivalent element of atomic weight 13.5 cannot be fitted into the periodic table, but a bivalent element of atomic weight 9.0 fills the space between lithium = 7.0 and boron = ii.o. The state- ment gave rise to a controversy which lasted about ten years, and 24 370 A TEXT-BOOK OF INORGANIC CHEMISTRY which was finally settled in favour of Mendeleeff by Nilson and Petterson, two of the chief advocates of the trivalency of beryllium. These chemists determined the vapour density of beryllium chloride to be 41, hence the molecular weight of the chloride is 82, and its formula BeCl 2 . Deficiencies of the Periodic System Although the periodic system has proved of the highest value for the development of chemistry, it must be admitted that it is by no means satisfactory in all respects. We shall learn later that there is not that close analogy between the members of the copper group (copper, silver and gold) and the alkal imetals to be anticipated from their position in the first group, and it has been suggested that the first three elements belong more properly to the eighth group, coming after nickel, palladium and platinum respectively. There are several instances in which the order of the atomic weights is different from that required by the chemical behaviour of the elements. Thus argon must undoubtedly come below neon, and be followed by potassium, which fits into its proper place below sodium, yet the atomic weight of argon is greater than that of potassium. The question which has raised most discussion in this connexion, however, is the relative position of tellurium and iodine. Although from its chemical relationships (p. 305) the latter element must follow iodine, yet experiment shows that the atomic weight of tellurium is greater than that of iodine. It is at first natural to suppose that there must have been some error in determining the atomic weights, but recent investigations appear definitely to have disproved this sugges- tion. Finally, the chemical behaviour of cobalt indicates that it should precede nickel in the periodic table, whereas its atomic weight is somewhat greater than that of the latter element. There is no definite place for hydrogen in the table. Ramsay places it at the summit of the halogen group, but it must be admitted that the analogies are slight. Other chemists, with perhaps even less justification, place hydrogen at the summit of the alkali metals ; recent investigation shows that it has no metallic properties. The position of the metals of the rare earths in the periodic system is very uncertain (p. 474). These considerations show that the periodic system, although of the highest value, is only a first approximation to a satisfactory system. The difficulties of classification arise from the many-sided character of the different elements, as illustrated in their physical and more particularly in their chemical properties. GENERAL PROPERTIES OF THE METALS 371 GENERAL PROPERTIES OF THE METALS It will be of advantage, before taking up the study of the metals and their compounds in detail, to give a brief account of their general characters, more particularly as to the rhethods of preparation and general properties of the salts. It has already been stated that the metals are broadly distinguished from the non-metals (i) by their metallic lustre and conducting power for heat and electricity ; (2) by uniting with oxygen to form basic oxides, which latter unite with acids to form salts ; (3) the metal forming a constituent of a salt appears alone in the form of positive ions only ; (4) the chlorides are fairly stable towards water. A few general remarks on these characters will be of interest. The conducting power of metals for electricity is expressed in terms of the conductivity, in reciprocal ohms, of a cm. cube of the substance ; this magnitude is termed the specific conductivity, k. The metal of highest conductivity is silver, and copper comes very near to it in this respect. Some non-metals, in one or other of their allotropic modifications (e.g. carbon) show conducting power, but are much less efficient conductors than the metals. Solutions which show electrolytic conduction are all far inferior to the metals in this regard. The following small table, which illustrates these statements, is in- structive ; the numbers are valid for 18. Substance Silver Copper Mercury Gas 30 per cent. Carbon Sulphuric Acid Sp. conductivity 624,000 587,000 10,240 200 1.35. The points (2) and (3) are closely allied. The majority of basic oxides combine more or less readily with water to form hydroxides, which in solution are electrolytically dissociated into positive metallic ions and negative OH' ions. Hydroxides which behave in this way are termed bases, and are characterized by the properties already fully desc ribed. As regards those metallic oxides which are practi- cally insoluble in water, e.g. magnesium oxide, their basic character is recognized by the fact that with acids they form salts, whose positive ions in solution consist of charged atoms of the metal. The " strength " of a base is measured, as already stated, by the extent to which it is dissociated in solution, but the application of this criterion to metallic hydroxides is often complicated by the slight solubility of the latter. Another method of measuring the 372 A TEXT-BOOK OF INORGANIC CHEMISTRY strength of a base is to determine the extent to which a salt is hydrolyzed in aqueous solution. It is evident that if we compare in this respect salts formed by different metals with the same strong acid, for example hydrochloric acid, the salts of strong bases will undergo little or no hydrolysis, but as the basic character diminishes the hydrolysis will become increasingly pronounced (cf. p. 267). These considerations enable us to understand criterion (4) for the metals; when the basic character, in other words the tendency to form positive ions, is practically or entirely absent the chloride is readily decomposed by water (e.g. phosphorus trichloride). The hydroxides of non-metals, e.g. P(OH) 3 or H 3 PO 4 , have a tendency to form salts with bases ; in other words, the hydroxides are acidic. We shall learn later that the hydroxides of certain metals have both basic and acidic properties. Another way of regarding the matter is to consider the relative affinity for electricity shown by different elements in the atomic con- dition. Metals which have a great tendency to form positive ions may be regarded as having a great affinity for positive electricity ; in other words, such metals are strongly electro-posi ti ve (cf. p. 4 1 8). The non-metals have no affinity for positive electricity ; on the contrary, some of them, such as the halogens, have a great affinity for negative electricity, and are therefore said to be strongly electro-negative elements. They form negative ions in solution. Finally, certain elements, such as carbon and silicon, appear to have little or no tendency to associate with electricity ; they might be termed electro- neutral elements. It should, however, be stated that though there is a distinct parallelism, there is no direct proportionality between the strength of a base and the affinity of the metal for positive elec- tricity (cf. p. 419). The above considerations, which are of funda- mental importance for the proper understanding of our subject, will now be illustrated by brief references to the main groups of metals, arranged on the basis of the periodic system. The Principal Groups of Metals (i) The Alkali group. The closely allied elements, lithium, sodium, potassium, rubidium and caesium, which occur in the first group of the periodic table, are termed the alkali metals. They are strongly electro-positive, and this character increases with increasing atomic weight. Correspond- ing with this they are powerful bases and their salts are not hydro- lyzed in solution. They act exclusively as univalent elements. (2) The Copper group, comprising copper, silver, and gold also occur in the first vertical column of the periodic table. Corresponding GENERAL PROPERTIES OF THE METALS 373 with their position, they can all function as univalent elements, but copper and gold also show other valencies. The hydroxides of copper a id gold are relatively weak bases, and their salts are therefore partially hydrolyzed in solution. Silver hydroxide is a comparatively strong base. Their affinity for positive electricity is small. (3) Metals of the Alkaline Earths - These comprise calcium, strontium, and barium. They always act as bivalent elements ; the hydroxides are strong bases. They are strongly electro-positive elements. (4) The Zinc group. This group comprises beryllium, magnesium, zinc, cadmium, mercury. Their main valency is 2. They are weaker bases than the metals of the alkaline earths, and their salts with strong acids are partially hydrolyzed in solution. (5) The Aluminium group. Aluminium is the only important member of this family. In its compounds it is trivalent; the hydroxide has weak basic and also weak acidic properties. Never- theless, aluminium is a strongly electro-positive element (p. 465). (6) The Tin group. Tin and lead are the most important members of this jjroup. Corresponding with their position in the table, they are quadrivalent, but both can also function as bivalent elements. In the divalent condition they are weakly basic ; in the quadrivalent condition the acidic character predominates. (7) The Arsenic group. Arsenic, antimony and bismuth are the chief metals belonging to this group. They function mainly as triva- lent and pentavalent elements. The metallic character increases markedly with increase of atomic weight ; but even bismuth oxide is a weak base, and its salts with strong acids are partially hydrolyzed. (8) The Chromium group. Chromium, the chief metal in the sixth group, functions as a di-, tri-, quadri- and hexavalent element. The lowest oxide is weakly basic, the highest acidic. (9) The Manganese group. Manganese itself is the only metal in this group. It shows a number of valencies from 2 to 7. In its lowest state of valency it acts as a weak base, in its highest valency it has pronounced acidic properties. (10) The Iron group. Iron, cobalt and nickel are the members of this group. They act as divalent and trivalent elements, and are weak bases. ( 1 1 ) The Palladium and Platinum groups. The members of these groups show considerable diversity in valency. The lower oxides are weakly basic, the higher show slight acidic properties Fron: the above brief statement of the characteristics of the groups 374 A TEXT-BOOK OF INORGANIC CHEMISTRY certain important generalizations may be made. The electro-positive character (and also the electro-negative character) is most pronounced for univalent elements (e.g. the alkali metals, the halogens) ; but polyvalent elements may also show great affinity for electricity, e.g. aluminium. On the other hand, univalent elements are mostly strongly basic or acidic, and the oxides of polyvalent elements are often both weakly basic and weakly acidic (e.g. aluminium). When the same element forms both basic and acidic oxides, the higher oxides (that is, those containing the higher proportion of oxygen) are most acidic. Methods of Preparing Metals from their Compounds The methods will be described in connexion with the metals them- selves, but a brief outline of the processes will prove useful at the present stage. The compounds chiefly employed are the oxides and the halides, and the method , used in any particular case depends mainly on the volatility of the metal and on its chemical affinity for other elements. Provided the oxide is not too stable, it can be reduced by heating with carbon (examples : zinc, tin, iron, etc.) or with finely divided aluminium (examples : chromium, manganese, etc.). If the metal has so great an affinity for oxygen that this method cannot be applied, recourse is often had to the electrolysis of a fused salt, for example the fused chloride, the metal separating at the negative pole. Potassium, sodium, lithium, magnesium, alumi- nium and other metals are now obtained commercially in this way. In applying these general methods the naturally occurring com- pound of the metal has first to be converted into the oxide or chloride, and purification may also be necessary. The methods used depend largely on the nature of the ore. Sulphides, for example, are con- verted into oxides by roasting, the sulphur being burned away as sulphur dioxide. Methods of Preparing Salts The more important methods of preparing salts have already been described in connexion with the individual acids. If one or more of the reacting substances are used in solution, the method to be employed in any particular case depends greatly on whether the salt is or is not insoluble in water. If it is only slightly soluble the general method of preparation is by double decomposition between a solution containing the basic and one containing the acidic constituent, e.g. : ^ +2HC1 +2KNO 3 . If, on the other hand, the salt is soluble in water, it may usually be GENERAL PROPERTIES OF THE METALS 375 prepared by the action of the corresponding acid (if a moderately strong one) on the metal, oxide, hydroxide, or salt with a weak or readily volatile acid. Examples : Ba(OH) 2 + 2HCl->BaCl 2 4-2H 2 O Some sa ts (e.g. the halides or sulphides) are occasionally prepared by direct combination of the elements. It is, of course, evident that salts of a very weak base and a very weak acid cannot exist in contact with water, as they readily undergo hydrolysis. Such salts can, however, sometimes be prepared by reactions in the dry state. Chlorides may be obtained by direct combination of the elements (e.g. FeC! 3 ; SnCI 4 ) or by the action of hydrochloric acid on the metal, oxide or carbonate. The three chlorides which are practically in- soluble in water (silver, lead and mercurous chlorides) are prepared by double decomposition. The same remarks apply to bromides and iodides. Carbonates, being all insoluble in water with the exception of those of the alkali metals, are prepared by double decomposition. The alkali caibonates are prepared by special methods (q.v.}. Nitrates, being all soluble in water, are prepared by the methods used for obtaining soluble salts. Sulphates, which are mostly soluble in water, are prepared by the general methods. The sulphates of calcium, strontium, barium, and lead, which are only very slightly soluble in water, are prepared by double decomposition. Sulphides Most sulphides are insoluble in water, and are pre- pared by double decomposition. The sulphides of the alkalis, which are soluble, are obtained by the action of hydrogen sulphide on the corrc spending hydroxides. Phosp/'.atcs, with the exception of those of the alkali metals, are insoluble in water, and are obtained by double decomposition. The preparation of the phosphates of the alkalis has already been describee. Solubility of Salts The importance of a knowledge of tne solubility of salts in water has already been repeatedly emphasized. The accompanying table, a slightly modified form of one given 376 A TEXT-BOOK OF INORGANIC CHEMISTRY by Kohlrausch, gives the solubility, in grams per litre at 18, of most of the bases and salts in general use. The numbers in the squares represent the solubility of the salt whose positive ion stands at the top, and whose negative ion stands at the side of the column. Each square contains two numbers ; the top one represents the number of grams of salt (calculated as anhydrous) taken up by 1000 grams (i litre) of solvent at 18, while the lower one gives the number of grams of salt in i litre of the saturated solution at 18. ON 0' 4 LO ., CO CO d d do LO -tf- v^ O CO CO ^ H? H? H H 0*0 H M d d o^o 5 g g d d d O.OO2 0.002 88 d d N c o o Ti~ LO CO O 00 O tN M O O NJ- (N LO LO 00 O 9 9 o o S > H- i'* GO 01 u") to k 00 H M if) to o d if) to * 11 d d d d bjo oo o LOCO 10 Tf H O 0*00" H !! CO CO _ oo oo 2 S. o o <> d o p U i| H W O O d d LO ^t* CO O "* * d d HH ctf U CM O CO O tNVO (M O COVO ^ M 1 (N H MM o 2 S2 ^gN "7^ J^ 00 vO W CM K IN ci ci ^ ^ 00 00 8 8 d d CO CO d 6 ^ S R. LO O %,% d d coo $^88 ^^ tN tN M H (N) ( ^ M M W ON 00 vO co H H d d ^ ^ ^* ^ co co c5 ci IN IN O O M H d d d d CO CO coS 2* 8* VO VO Tl- 00 00 LO Tf 00 GO . co co O O tN tN co co 88 88 o'd o" d tN tN O O d d SCJ d d H co co d d -I- 1 d d d d O O ON O Cvj tN 00 ON ON ON ON or) co ^ co do * co r^- vo vO 00 O O Tt- ^ d d ON CO 44 H H LO d\ o rfiO * 88 d d d d d d LO CO CO M^M" IT n- N co oo sis s 99 H H O O d d vO vO LTJ (N vas drawn between sodium salts and potassium salts. The latter were chiefly obtained from vegetable sources, and were called kali or alkali. This distinction was first drawn by Duhamel du Monceau in 1736. Occurrence Sodium does not occur free in nature owing to its great affinity for oxygen. It occurs as silicstte in many rocks, by the disintegration of which it finds its way into the soil and thence into rivers and the sea. Sodium chloride occurs to the extent of 2.6 to 2.9 per cent, in sea- water, and the same salt is found in deposits of FIG. 81. jreat thickness in various parts of the world ; for example, in Russia, n German}-, and in Lancashire and Cheshire in this country. Sodium nitrate occurs in enormous amount in Chili and Peru, and is :ermed Chil saltpetre. Large deposits of cryolite, sodium aluminium Suoride, AlF 3 ,3NaF, occur in Greenland and Iceland. Preparation Metallic sodium was first obtained by Davy ^1807) by the electrolysis of sodium hydroxide. It is an interesting "act that at present the metal is prepared on the commercial scale ilmost entirely by this method. (i) Cast-tier's Electrolytic Process The method for this purpose ievised by Castner is illustrated in Fig. 81. D is a cylindrical steel 380 A TEXT-BOOK OF INORGANIC CHEMISTRY vessel with an opening at the bottom, through which the iron cathode C passes ; it is heated by gas-burners e e, so that all except the neck is kept at a temperature about 20 above the melting-point of the hydroxide. The anode A, which is conveniently made in the form oi a cylinder with vertical slits, surrounds the upper portion of the cathode. Within the anode is the collecting pot F, from which is suspended a cylinder of wire gauze, which surrounds the upper portion of the cathode. The products of electrolysis are sodium and hydrogen, which are liberated at the anode, and oxygen, which is liberated at the cathode, and escapes through a valve at the top of the vessel. The sodium rises to the surface of the fused electrolyte in F, and is removed from time to time by means of a perforated ladle. The hydrogen, also liberated at the cathode, serves to protect the sodium against oxida- tion ; it escapes at the loosely-fitting lid of the vessel F. (2) Other Electrolytic Methods Instead of using sodium hydroxide, it would obviously be advantageous in some ways if sodium chloride could be used as electrolyte : it is cheaper, and both products oi electrolysis, sodium and chlorine, are commercially valuable. Among the difficulties met with in using this method are the high melting- point of the chloride, the disintegrating effect of the fused electrolyte on the cell materials, and the combination of the liberated metal with the fused salt to form the " sub-chloride," Na 2 Cl. These difficulties do not appear to have been entirely overcome. (3) Chemical Methods The chemical methods for preparing metallic sodium, which have now been almost completely displaced by the electrolytic methods, depend upon the reduction of sodium hydroxide or carbonate with carbon or iron, or a mixture of both. In Castner's chemical process sodium hydroxide was mixed with iron and finely divided carbon (perhaps iron carbide, FeC 2 ), and the mixture distilled from iron retorts : The metal was condensed in flat iron receivers, with an outflow so arranged that the metal was collected under mineral oil. Physical Properties Sodium is a white, lustrous metal, which at room temperature can be moulded with the fingers, but is hard at -20. It melts at 97, and boils at 877; the vapour in thick layers has a purplish colour. The vapour density, determined in platinum vessels, is about 12, indicating that the metal is monatomic in this state. SODIUM 381 Chemical Properties Sodium is unattacked in perfectly dry air or oxygen, but in moist air the fresh surface becomes coated almost instantaneously with a film of oxide. It is vigorously acted upon by water at room temperature : 2Na + 2H 2 O->2NaOH + H 2 . The heat dven out in the reaction is not sufficient to ignite the hydrogen if the metal is allowed to move about on the surface of the water, but if the metal is confined to one point the hydrogen ignites and burns v/ith a yellow flame. Sodium dissolves in liquid ammonia to form a blue liquid. It com- bines with dry ammonia at 300-400 to form sodamide, NaNH 2 , a white solid which melts at 155. With mercury in certain proportions it forms a solid amalgam (p. 456), which contains definite compounds of the two elements. The action of water on sodium amalgam forms a convenient method of obtaining " nascent " hydrogen. With potas- sium, sodii'm forms a liquid alloy resembling mercury, which has been used for thermometric purposes. Sodium Hydride, NaH When sodium is heated with hydrogen in an iron vessel at 360, sodium hydride is formed, and condenses on a cooler part of the apparatus in the form of colourless crystals. The compound readily dissociates on heating above 430. It is stable in dry air, but is immediately decomposed by water. Sodium Oxides Two oxides are known, the normal oxide or monoxide, Na 2 O, and the peroxide, Na 2 O 2 . Sodium Monoxide is obtained by partial oxidation of sodium, the unchanged metal being removed by prolonged distillation in a vacuum. It is a white, amorphous, hygroscopic powder, which combines vigorously with water to form the hydroxide. Sodium Peroxide, Na 2 O 2 , is obtained by heating sodium in an iron tube at 300 in a current of dry air, free from carbon dioxide. It occurs as a light yellow powder, which dissolves in water under ordinary conditions with rise of temperature and considerable evolu- tion of oxygen. The primary reaction is probably as follows : but unless precautions are taken to keep the mixture cool, the hydrogen peroxide is largely decomposed into water and oxygen. The above reaction is reversible. When the peroxide is dissolved in dilute acids, the corresponding sodium salt and hydrogen peroxide are obtained. 382 A TEXT-BOOK OF INORGANIC CHEMISTRY An octahydrate, Na. 2 O 2 .8H 2 O, has been obtained in colourless crystals. Sodium peroxide is largely used for oxidizing and bleach- ing purposes. Sodium Hydroxide (Caustic Soda), NaOH Preparation (i) By the action of water on metallic sodium or sodium monoxide. The former method is used for the preparation of very pure hydroxide. (2) By boiling together sodium carbonate and calcium hydroxide in iron vessels : The reaction is a reversible one, and in practice the best yield is secured by using excess of calcium hydroxide and working at a definite dilution. The calcium carbonate is allowed to settle, the clear liquid decanted off, evaporated in iron vessels, and cast into sticks. (3) Caustic soda is now obtained commercially by electrolysis of sodium chloride. The methods differ according as the salt is em- ployed in the fused condition or in aqueous solution. As an illustra- tion of the former case the Acker process, in use at Niagara Falls, may be referred to. The salt is fused by the heat of the current itself, and is electrolyzed between a carbon anode and a lead cathode. The alloy of sodium and lead thus obtained is caused to flow by means of a steam jet into another compartment, where it is decomposed by water, the sodium hydroxide being run off and the lead returned to the electrolytic cell. The caustic soda solution thus obtained may be eva- porated if necessary, or may be used directly for soap-making and other purposes. Of the methods depending on the use of sodium chloride in solution space will only admit of a brief account of the Castner-Kellner process. A diagram of the apparatus used is given in Fig. 82. The cell is divided into three compartments by means of two vertical non- porous partitions, which do not reach right to the bottom but extend into two shallow grooves in the base of the cell. The bottom of the cell is covered with mercury, the liquids in the compartments being thus kept separate. The two side compartments contain brine, into which the carbon anodes extend ; the central compartment contains water and an iron grid which acts as cathode. The cell, which is closed, is supported on an eccentric at one end, by means of which a rocking motion can be imparted to it, which secures the circulation of the mercury. When the current passes, chlorine is evolved at the anodes and sodium is liberated at the surface SODIUM 383 of the mercury facing each anode, and dissolves in the mercury. In consequence of the rocking motion the amalgam passes into the cathode compartment, where it reacts with water, with formation of sodium hydroxide and hydrogen. The hydrogen and chlorine escape from the respective compartments by the pipes (not shown in the Fig.)- Properties Sodium hydroxide is a white, very hygroscopic substance, which melts at 318. It liquefies in the air owing to absorption of moisture, but finally becomes solid consequent on forma- nor FIG. 82. tion of the carbonate. Its aqueous solution has powerful caustic properties, and shows all the characteristics of a strong base. Sodium Chloride, NaCl As already mentioned, sodium chlo- ride is the chief saline constituent of sea-water, and it occurs in large proportions in certain lakes and also in deposits of great thickness in Galicia, Germany, and other parts of the world. These deposits were formed by the complete drying up of lakes or seas. As sodium chloride is less soluble than many of the other salts in sea-water, it separates out first, and above it are found layers of more soluble salts. The most famous of such deposits is at Stassfurt in Germany, where the salt layers are in some parts 1000 metres in thickness. Salt is prepared in warm countries from sea-water by evaporation, and from salt deposits either by mining the solid or by boring through the upper strata and then using water to form a strong brine, which is pumped through copper tubes and evaporated to crystallization. In order to economize fuel, the first stage of the evaporation is often accomplished by allowing the brine to trickle down large ricks of brushwood, so that free exposure to wind is secured. In order to 384 A TEXT-BOOK OF INORGANIC CHEMISTRY render it fit for household purposes, natural salt has to be purified by recrystallization from water. The deliquescent character of some samples of salt is due to the presence of magnesium chloride as impurity. Sodium chloride can be obtained in a perfectly pure condition by precipitating from aqueous solution by means of hydrogen chloride (cf. p. 423) and then heating strongly. Properties Sodium chloride is usually obtained in cubic crystals, which melt about 810. The solubility in grams per 100 grams of water is as follows : 35.63 at o, 35.82 at 20, 36.32 at 40, 37.06 at 60, and 39.12 at 100. The very slight change of solubility with tempera- ture is in accordance with the very small heat of solution. Only one hydrate of sodium chloride, NaCl,2H 2 O, is known ; it separates from solutions below o in monoclinic crystals, whilst at higher temperatures the anhydrous salt is obtained in cubic crystals. The fact that sodium chloride is an essential constituent of the food of animals is familiar to all. This salt is the source of all sodium compounds except the nitrate, and is the chief source of chlorine compounds. Sodium Bromide, NaBr, and Sodium Iodide, Nal, are prepared by the methods described under the corresponding potas- sium salts. They are isomorphous with sodium chloride. At room temperature the dihydrates separate from aqueous solution. The transition temperature, NaBr-NaBr,2H 2 O, is at 50.6 ; that of NaI-NaI,2H 2 O at 65. At low temperatures sodium iodide pen- tahydrate separates from solution ; the transition temperature, NaI,5H 2 O-NaI,2H 2 O, is at - 13.5. Composition Of Hydrates The general methods used in determining the composition of the hydrates formed by a salt or other substance and the limits within which they exist may conveniently be considered in connexion with sodium iodide. The diagram represent- ing the variation in the solubility of the salt with temperature is shown in Fig. 83 ; the ordinates represent temperatures and the abscissas concentrations in grams of salt per 100 grams of water. The curve AB represents the effect of the salt in lowering the freezing- point of water ; in other words, it is the curve along which ice and solution are in equilibrium (cf. the system, water-potassium iodide (p. 199)). When sufficient salt is added, however, a point is reached at which, on cooling the solution sufficiently, a hydrate, and not ice, separates. Thus, along the part BC of the curve the pentahydrate separates from solution ; in other words, BC represents the solubility curve of the compound NaI,5H 2 O. At the point B, - 32.5, ice and the SODIUM 385 pentahydrate are in equilibrium, and this is a eutectic point. As the quantity of salt is still further increased the point C is reached, at ~ I 3-5j wjien the pentahydrate and dihydrate are in equilibrium. With higher proportions of salt the solid in equilibrium with the solu- tion is the dihydrate, and its solubility curve is represented by CD. Finally, at D, 65, the dihydrate is in equilibrium with anhydrous salt, and the further part of the curve DE represents the solubility of g*0 I" I- ^ Concentration (grams in 100 grams water). FIG. 83. the anhydrous salt. At each of the points, B, C, and D, so-called breaks on the curve, a new phase appears ; and as the sections BC, CD, DE each represent the solubility curve of a separate substance, it is easy to understand why the whole curve is discontinuous. The number and composition of the hydrates can then be determined from the number and position of the breaks on the curve, controlled by analysis of the compounds in equilibrium with the solution at different calutions. 25 3 86 A TEXT-BOOK OF INORGANIC CHEMISTRY Formerly the number and nature of the hydrates formed by a salt was deter- mined by evaporating the solution at different temperatures, and analysing the solid compounds obtained. The modern method is to start with pure water, add successive portions of the salt, and determine the complete solubility curve, as shown in Fig. 83. If no hydrates exist under the experimental conditions, the curve representing the equilibrium between salt and solution will be con- tinuous, as in the examples in Fig. 32. The existence of hydrates will be indicated by breaks in the curve. The complete curve of the system includes the eutectic ice-salt and that part of the curve where the solution is in equilibrium with ice, but this part is generally omitted in the solubility curves. Other properties of the system, besides the solubility, may also be made use of to determine the existence and composition of hydrates. One of the most instructive of these properties is the vapour pressure of solid hydrates. If a crystal of copper sulphate pentahydrate is put in the vacuum of the barometric column (Fig. 19) it will be found to exert a definite pressure, due to water vapour, just as water itself does under the same conditions. At 50 the pressure in question is about 47 mm. If the copper sulphate is put in a desiccator over concentrated sulphuric acid, which absorbs the vapour, it will be found that the pressure remains constant at 47 mm. (if the observations are made at 50) till the salt has lost 2H 2 O, when it suddenly falls to 30 mm. , which is the vapour pressure of the trihydrate. As long as any trihydrate is present it remains constant at this value, but when two further molecules of water have been removed the pressure suddenly falls to 4.4 mm., which is the vapour pressure of the mono- hydrate. Finally, when the water is completely removed, the pressure of course falls to zero. It is evident from the above that the occurrence of intermediate hydrates can be deduced from observation of the pressure during dehydration. If the pressure fell at once from 47 mm. to zero, it would mean that no intermediate hydrates exist. Conversely, if water vapour at a pressure of say 5 mm. is brought into contact with anhydrous copper sulphate at 50, the monohydrate will be formed, but no higher hydrate. Only when the pressure is raised above 30 mm. is the trihydrate formed, and a pressure exceeding 47 mm. is required to form the pentahydrate. The reactions in question are therefore reversible, and the phenomenon is one of dissociation. These considerations enable us to understand the behaviour of a hydrated salt in the air. If the pressure of aqueous vapour exerted by the hydrate is greater than the partial pressure of aqueous vapour in the atmosphere, the salt will lose water and a lower hydrate will be formed. Under these circumstances the salt is said to effloresce in the air. At room temperature, the vapour pressure of copper sulphate pentahydrate is less than the average pressure of aqueous vapour in the atmosphere, and the salt does not effloresce ; the vapour pressure of sodium sulphate decahydrate, on the other hand, is greater than the average pressure of aqueous vapour in the atmosphere, and the salt effloresces. It should be pointed out that although we have spoken of 47 mm. as the dis- sociation pressure of copper sulphate pentahydrate at 50, a definite pressure is only attained when two solid phases, the hydrate in question and the next lower hydrate, are present. Thus, at 50, copper sulphate trihydrate is stable in the presence of pressures of aqueous vapour from a little above 30 mm. to a little SODIUM 387 below 47 mn . ; when it is kept at a value less than 30 mm. the trihydrate finally disappears a id monohydrate is formed exclusively ; when it is kept at a value above 47 mn . , the pentahydrate is finally obtained. Many hydrates thus remain unaltered in the air when the temperature and vapour pressure are altered within certain limits. When the vapour pressure of the saturated solution of a salt is less than the vapour pressure of the atmosphere, the salt takes up water, and a solution is finally obtained. Such a salt is said to be deliquescent example, calcium chloride. From the general principles of equilibrium, the dissociation pressure of a hydrated sal must evidently correspond with the vapour pressure of the aqueous solution witl which it is in equilibrium. Thus sodium iodide dihydrate separates from aqueous solution at all temperatures between 13 and +67, and we may anticipate, as is in fact the case, that it remains unaltered in pressures of aqueous vapour from about 2 mm. (the pressure at - 13) to 200 mm. (the pressure at 67), and therefore that it is stable in the air. Sodium Carbonate, Na 2 CO 3 Preparation This salt, which is extensively used in the manufacture of glass, soap, etc., is now prepared commercially from sodium chloride by three distinct methods (i) The Leblanc process, invented about 1790 ; (2) the ammonia-soda or Solvay process (1860) ; (3) the electrolytic process, introduced within the last few years. (i) The Leblanc Process There are three distinct stages in this process, in the first, sodium chloride is converted into the sulphate by heating; with sulphuric acid (salt-cake process) ; the sulphate is then reduced to the sulphide by heating with carbon, and finally from the sulphide and calcium carbonate (chalk) at a high tem- perature sodium carbonate and calcium sulphide are obtained : (a) 2NaCl + H 2 S0 4 ->Na 2 S0 (c) Na 2 S + CaCO 3 ->Na 2 CO 3 The last t ,vo stages are accomplished in one operation, the sodium sulphate teing mixed with chalk and small coal, and the mixture strongly heated. The product (black ash) is then treated with water, which dissolves out the carbonate, leaving the insoluble calcium sulphide. With reference to the details of the operation, the first stage of the salt-cake process is carried out by heating the salt and sulphuric acid in lar ; V e cast-iron pans (Fig. 84, d). The main reaction at this stage is as follows : H 2 S0 4 ->NaHS0 4 + HClt. 388 A TEXT-BOOK OF INORGANIC CHEMISTRY The practically dry product, containing sodium bisulphate and un- altered chloride, is raked out on the hearth (/;) of a reverberatory furnace 1 (Fig. 84) and heated still more strongly, with continual mixing, until the following reaction is complete : NaHSO 4 + NaCl->Na 2 SO 4 + HCl. The hydrogen chloride escaping in the two stages of the process is absorbed by passing it up towers filled with coke over which water is made to trickle. The production of the black ash is brought about by heating the reacting substances in a rotating cylinder (to secure thorough mixing) in a furnace, and the end of the reaction is recognized by the appear- FIG. 84. ance of jets of burning carbon monoxide, due to the action of carbon on the calcium carbonate : CaCO 3 + C->CaO 4- 2CO. The carbon monoxide formed towards the end of the operation has the effect of rendering the black ash porous, and therefore more readily extracted by water. 1 A reverberatory furnace is an arrangement whereby the material to be heated or fused does not ccme into contact with the solid fuel. A simple furnace of this type is shown in Fig. 85. The fuel is burned in the fire-box a, and the flame and heated gases are caused to act on the material spread out on the bed, c, of the furnace by "reverberation" from the low roof. The products of combustion escape by the flue d. The furnace represented in Fig. 84 has in addition the large cast-iron pan d, in which the mixture of sodium chloride and sulphuric acid is heated. SODIUM 389 The extraction (lixiviation) of the black ash is carried on in a series of vessels so arranged that the water passes in turn from one to the other until it is completely saturated. The process is carried on at 30 to 40, the temperature at which sodium carbonate is most soluble. The solution is evaporated in vessels heated in the flues of the black-ash furnace, when the monohydrate, Na 2 CO3,H 2 O, is pre- cipitated. When the monohydrate is strongly heated, all the water is driven off, and "calcined soda," largely used in commerce, is obtained. On dissolving the anhydrous salt in water and allowing it to crystallize, the readily soluble decahydrate, Na 2 CO 3 ,ioH 2 O, " washing soda," is obtained in large crystals. The insoluble residue from the black ash, which consists largely of calcium sulphide, is worked up in order to extract the sulphur. According to Chance's process for this purpose, the residue mixed with water is placed in a series of vessels through which carbon dioxide is passed. The final result of the reactions is that calcium carbonate and a gas rich in hydrogen sulphide are obtained. The latter is either burned completely to sulphur dioxide, which is used directly i i the sulphuric acid manufacture (p. 288), or is burned in an insufficient supply of air and the sulphur collected : (2) 77/6' Ammonia-Soda Process According to this simple and economical process, introduced by Solvay, sodium chloride and ammonium bicarbonate are brought together in fairly concentrated solution, v/hen sodium bicarbonate, NaHCO 3 , being relatively slightly soluble ir water, is precipitated, and is then converted into the normal carbonate by heating. In practice, ammonia and carbon dioxide (obtained by heating limestone) are led into a solution of sodium chloride. The reactions are represented by the following equations : (i) CO 2 + NH 3 + H 2 O->NH 4 HCO 8 . (2) (3) From the ammonium chloride formed in the course of the reaction the ammonia is regenerated and, along with the carbon dioxide formed ir (3), is used to decompose a fresh quantity of sodium chloride, s 3 that the only by-product is calcium chloride. The process is thus a very economical one, and it has the additional advantage 390 A TEXT-BOOK OF INORGANIC CHEMISTRY over the Leblanc process that the soda obtained is much purer. It must be clearly understood that the whole basis of the process is the relative insolubility of sodium bicarbonate as compared with ammonium bicarbonate. (3) Electrolytic Method Sodium hydroxide, obtained by an electro- lytic method such as that described on p. 382, is converted to the carbonate by means of carbon dioxide obtained by heating limestone. At present all these methods are in use. The great advantages of the Solvay over the Leblanc process are to some extent counter- balanced by the fact that the latter gives the valuable by-product hydrochloric acid. It is probable that in course of time both will be superseded by the electrolytic method. Properties of Sodium Carbonate The substance separating from solution at room temperature is the decahydrate, Na 2 CO 3 ,ioH 2 O. It has a high tension of aqueous vapour and therefore loses water (effloresces) on exposure to air, with formation of the monohydrate. An intermediate hydrate, Na 2 CO 3 ,7H 2 O, can be obtained by crystal- lization from warm solutions under certain conditions. The deca- hydrate and heptahydrate are in equilibrium with the saturated solution at 32, the heptahydrate and monohydrate at 35.4. The solubilities of the decahydrate and heptahydrate increase, that of the monohydrate diminishes, with rise of temperature, from which it follows that the salt must show a maximum solubility about 35.4 (cf. sodium sulphate). The anhydrous salt melts at 852. Above this temperature it begins to decompose, giving off carbon dioxide. The aqueous solution of sodium carbonate is slightly alkaline owing to hydrolysis. This has already been fully explained (p. 324). Sodium Bicarbonate, NaHCO 3 , is obtained by passing carbon dioxide into a solution of sodium carbonate : As already explained, this reaction is reversible ; even the aqueous solution of sodium bicarbonate gives off carbon dioxide when boiled. The aqueous solution is practically neutral owing to the very slight ionisation of the HCO/ ion. The solubility in grams per 100 grams of water is as follows : Temperature . o 10 20 30 40 50 60 Solubility . . 6.9 8.4 9.6 ii.i 12.7 14.5 16.4. The salt is used for making baking powder. SODIUM 39 i Sodium Sulphate This salt occurs naturally in the salt de- posits in the anhydrous form as thenardite (rhombic crystals). It is obtained on the large scale in the first stage of the Leblanc process (P- 387) and also in the preparation of nitric acid (p. 223). The salt which separates from solution at room temperature is the decahydrate, Na 2 SO 4 ,]oH 2 O, which has been known for centuries as Glauber's salt. Above 33, the anhydrous salt separates from solution. The solubility curve of sodium sulphate is shown in Fig. 33, where the ordinates represent solubilities and the abscissae temperatures. The curve BC represents the solubility curve of the decahydrate, which increases with rise of temperature, and CD that of the anhy- drous salt, which shows that the solubility diminishes with increasing temperature. The point C, at 32.4, represents the maximum solubility of the sr.lt ; it is that temperature at which decahydrate and anhy- drous salt are in equilibrium with the solution. It must be carefully remembered that the sharp change in direction at C corresponds with a change in the solid from decahydrate to anhydrous salt and has nothing to do with changes in the solution. The readiness with which sodium sulphate forms supersaturated solutions has already been referred to. If such a supersaturated solution is cooled to 5, crystals of a heptahydrate, Na 2 SO 4 ,7H 2 O, are obtained. This salt is unstable in contact with the saturated solution under all conditions. Sodi im Nitrate, NaNO 3 This salt occurs naturally as Chili saltpetre, and is purified by crystallization from water. It forms rhombic crystals, which, unlike those of the corresponding potassium salt, are deliquescent. It is largely used as a manure, 1 and also for the preparation of nitric acid and of potassium nitrate (g-v.). Sodium Nitrite, NaNO 2 , can be prepared by strongly heating sodium litrate, or, better, by heating the latter salt with metallic lead, and is purified by crystallization : NaNO 3 + Pb->NaNO 2 + PbO. It is used in the preparation of nitrous acid. Sodium Phosphates These salts have already been described in connexion with phosphoric acid (p. 251). The ordinary phosphate of commerce is the disodium salt, obtained by neutralizing phosphoric acid with sodium carbonate and crystallizing. Below 36 the dodeca- i In 19 38 over two million tons of nitrate, worth about 20,000,000, were exported from Chili. 392 A TEXT-BOOK OF INORGANIC CHEMISTRY hydrate, Na 2 HPO 4 ,i2H 2 O, separates from solution ; above this tem- perature the compound Na 2 HPO 4 ,7H 2 O is obtained. Sodium Silicates Silicates of sodium occur in many rocks. A solution of sodium silicate, known technically as water-glass, is prepared by fusing two parts by weight of sand or quartz with one pajt by weight of sodium carbonate, dissolving the fused mass in water and evaporating to a syrupy consistency. The pure silicate, Na 2 SiO 3 , is obtained in crystalline form by dissolving the calculated amount of freshly precipitated silicic acid (p. 353) in sodium hydroxide solution and precipitating by means of alcohol. As silicic acid is an extremely weak acid, solutions of sodium silicate are strongly alkaline owing to hydrolysis. Sodium Sulphides The acid sulphide, NaHS, is obtained by saturating sodium hydroxide solution with hydrogen sulphide and evaporating to dryness, when the salt separates in colourless crystals. The normal sulphide, Na 2 S, is obtained by adding to the acid salt an equivalent of sodium hydroxide and evaporating the solution, when it separates as the nonahydrate in colourless crystals : NaHS + NaOH->Na 2 S + H 2 O. The behaviour of aqueous solutions of these salts has already been explained. The acid sulphide is slightly, the normal salt very con- siderably hydrolyzed in aqueous solution. Tests for Sodium As sodium salts are all highly ionised in solution, the tests for the positive component of the salts are really tests for sodium ion. Since practically all sodium salts are soluble in water, it cannot usefully be tested for by precipitation reactions. Sodium salts are always recognized by placing a platinum wire, to which a little of the salt adheres, in a colourless Bunsen flame, when a characteristic intense yellow colour is obtained. Conclusive evidence as to the presence of sodium is obtained by examining the coloured flame with an instrument known as the spectroscope. The Spectroscope. Spectrum Analysis Not only sodium, but all the alkali metals, give characteristic colours to the Bunsen flame, and can thus readily be distinguished with the naked eye. When salts of more than one metal are present, however, this method cannot be used, as the more intense flame masks the others. By means of the spectroscope, however, the metals present in a mixture can be separately detected. Light of a definite colour, for example yellow, is made up of vibrations, whose wave-lengths vary within narrow limits. White SPECTRUM ANALYSIS 393 light, on the other hand, is made up of vibrations of all wave-lengths within a considerable range. When a beam of white light, which has passed through a slit, falls on a triangular prism (glass prisms are generally used), the rays are bent or refracted to a different extent on passing through the prism, those of short wave-length, e.g. violet, being most, and those of longer wave-length, e.g. red, being least refracted. Therefore, if a screen is placed behind the prism, a so-called " continuous " spectrum made up of the different colours of the rainbow is seen ; it is formed of a very large number of images of the slit placed side by side. The spectrum will also be seen by an observer properly placed behind the slit. The spectroscope, an instrument designed for the examination of spectra, consists essentially of a prism, a tube provided with a narrow slit at the end remote from the prism, and a movable telescope, placed behind the prism, by means of which the image of the slit can be seen. When a Bunsen flame coloured by sodium is observed through the spectroscope, only one yellow image of the slit is seen against a dark background. This is due to the fact that sodium light is all of one wave-length, it is therefore not scattered on passing through the prism, and only one image of the slit is seen. Light made up of vibrations of one wave-length only is termed mono- chromatic. If a flame coloured by potassium is examined in the same way, two images of the slit two lines, as they are called are to be seen, one in the region corresponding with violet in the ordinary spectrum, the other in the red region. Lithium is characterized by two lines, a well-marked one in red, and a more feeble one in the orange. It is only in very rare cases that lines belonging to two different metals overlap, and therefore the constituents in a mixture of lithium, sodium and potassium can readily be detected by means of the spectroscope. One of the great advantages of spectroscopic analysis is that only minute quantities of the different substances are required for the purpose ; ii. is stated that one ten-millionth of a gram of sodium can be detected in this way. As sodium is very widely distributed in dust, etc., it is difficult to avoid the appearance of the character- istic line in the spectrum. As each element has its own characteristic spectrum, it can readily be understood that many elements were first discovered by means of the spectroscope. Thus the two rare elements of the alkili sub-group, rubidium and caesium, were detected by Bunsen and Kirch hoff in the waters of a mineral spring at Durkheim 394 A TEXT-BOOK OF INORGANIC CHEMISTRY (p. 402), and helium, thallium, scandium, and many other elements were discovered in this way. Line spectra, such as we have been describing, are given only by substances in the state of vapour. The spectra of metals are obtained by volatilizing the corresponding salts in a colourless flame, as already indicated. For the same purpose, the permanent gases are sealed up in glass tubes, and rendered luminescent (p. 339) by means of an electric discharge. Incandescent solids, e.g. carbon particles in a luminous flame, give a continuous spectrum. POTASSIUM Symbol, K. Atomic weight- 39.1. Molecular weight=39.i (probably). Occurrence Potassium occurs in nature mainly in the form of silicates in rocks, especially in felspar and mica. When these rocks are broken up, the soluble potassium salts pass into the soil, from which they are taken up by plants. Potassium salts are essential for the growth of plants, and are taken up in much larger proportion than sodium salts. When plants are burned the potassium is found in the ash as carbonate, and this formerly represented the chief commercial source of potassium compounds. Part of the potassium salts from the disintegration of rocks finds its way into seas, lakes and mineral springs. When an inland sea evaporates, the less soluble sodium chloride separates out first, and above this, at a later stage, are deposited the more soluble salts, chiefly those of potassium and magnesium. The Stassfurt deposits, formed in this way, now constitute the chief source of potassium, the more important compounds being carnallite, MgCl 2 ,KCl,6H 2 O, kainite, MgSO 4 ,KCl,3H 2 O, and sylvine, KC1. Preparation of Metal Like sodium, potassium was first obtained by Davy (1807) by electrolysis of moist potassium hy- droxide. The commercial methods now in use are similar to those described under sodium. The chief chemical method consists in strongly heating a mixture of potassium carbonate and finely divided carbon : Potassium and carbon monoxide combine under certain conditions to form potassium carboxide, K (CO) 6 , a highly explosive compound, and in order to avoid this the metal must be rapidly condensed, POTASSIUM 395 best by using flat receivers surrounded by cold water. The metal is collected under mineral oil. The intimate mixture of carbonate and carbon is usually obtained by raising a mixture of the carbonate and tar to a low red heat. The chief electrolytic method used depends upon the electrolysis of the fused hydroxide (cf. metallic sodium). Properties In all its properties potassium closely resembles sodium. It is a soft, silvery-white metal, which can be obtained in cubic crystals by sublimation; it melts at 62, and boils at 757. The vapo ir a little above the melting-point is greenish, at higher temperatures it becomes violet. It is probably monatomic in the state of vapour. Potassii m oxidizes in moist air or oxygen, and combines with the halogens even more vigorously than sodium does. It reacts vigorously with water : 2H 2 O->2KOH + H 2 , and so much heat is given out in the process that the evolved hydrogen catches fire and burns with a lavender flame. Potassium Hydride, KH, resembles sodium hydride in its mode of preparation and properties. Potassium Oxides At least two oxides of potassium, the monoxide K 2 O, and a peroxide, K 2 O 4 , are definitely known. Potassium Monoxide, K 2 O, is obtained by incomplete com- bustion of the metal in dry oxygen, the excess of metal being removed by distillation in a vacuum. It occurs in yellowish-white crystals, and combines very vigorously with water. Potassium Peroxide, K 2 4 5 is tlle cm 'ef product obtained when potassium is burned with free access of air. It is a yellow powder, w hich is decomposed by water, with formation of potassium hydroxide, hydrogen peroxide, and oxygen : Potassium Hydroxide, KOH Preparation (i) By boiling potassiun carbonate with calcium hydroxide or by the action of the latter compound on potassium sulphate : +2KOH +2KOH. The details have already been given under sodium hydroxide. (2) As in the case of sodium hydroxide, by electrolysis of an aqueous solution c f potassium chloride. 396 A TEXT-BOOK OF INORGANIC CHEMISTRY Properties Potassium hydroxide is a white amorphous substance, which melts at 360. It is usually sold in sticks. Three hydrates, KOH,H 2 O ; KOH,2H 2 O and KOH,4H 2 O, are definitely known. In all its properties it closely resembles sodium hydroxide. When solu- tions of potassium or sodium hydroxide are evaporated by boiling the temperature gradually rises without the separation of solid at any stage, and finally the fused compound is obtained. This curious behaviour depends on the fact that at no stage during the evaporation is there a saturated solution, whose vapour pressure reaches one atmosphere. Potassium Chloride, KC1 This salt occurs in the Stassfurt deposits as sylvine and as a constituent of carnallite, MgCl 2 ,KCl, 6H 2 O. It can be prepared by the general methods for a soluble salt (p. 374), and commercially is chiefly obtained from carnallite. When the latter compound is dissolved in water, it splits up into its components, and on evaporating the solution the potassium salt, being least soluble, separates out first. Properties The salt crystallizes in the anhydrous form in cubic crystals. No hydrates of potassium chloride, bromide or iodide are known. Unlike sodium chloride, the solubility of potassium chloride in water increases considerably with rise of temperature. Potassium Bromide, KBr The salt can be prepared by the general methods, but is almost invariably obtained by adding bromine to a hot aqueous solution of potassium hydroxide as long as the colour disappears ; a mixture of bromide and bromate is thus obtained : If both bromide and bromate are required, the salts are separated by fractional crystallization, the bromate being least soluble. If only the bromide is wanted the solution is evaporated to dryness, the residue mixed with charcoal, strongly heated to reduce the bromate to bromide : and the potassium bromide purified by recrystallization. Properties The salt separates from solution in colourless, anhy- drous cubic crystals. It is used in medicine, and also in photography for preparing silver bromide (g.v.). Potassium Iodide, KI, is prepared by the method described under potassium bromide, iodine being substituted for bromine, and also as follows. Iron and iodine are brought together below water, when they combine, forming a solution of the compound Fe 3 I 8 ; the POTASSIUM 397 latter is then decomposed by potassium hydroxide with formation of the compound Fe 3 O 4 , which is precipitated, and potassium iodide, which remains in solution : Potassium iodide occurs in cubic crystals, which are very soluble in water. Its used in medicine and in photography. Potassium Chlorate, KC1O 3 Preparation (i) By the action of chlorine on a hot solution of potassium hydroxide (p. 181) : As potassium chlorate is not very soluble in cold water" the salts can readily be separated by crystallization. (2) By the above method only one-sixth of the potassium appears as chlorate, and it is much more economical first to prepare calcium chlorate by the action of chlorine on hot milk of lime : The proper amount of potassium chloride is then added and the solution e\aporated to a definite density and set aside, when potas- sium chlorate separates out. (3) Potassium chlorate is now largely prepared by electrolytic methods, for example, by electrolysis of a hot solution of potassium chloride, the chlorine and alkali formed at the anode and cathode respectively reacting to form the chlorate, which separates out on con- centrating the solution. In practice only dilute solutions of chlorate can be prepared by direct electrolysis, and several methods of over- coming th s difficulty have been devised. Properties Potassium chlorate occurs in anhydrous monoclinic leaflets, and melts at 370. The solubility in water (grams in 100 grams of v/ater) is given in the accompanying table : Temperature o 10 20 30 40 50 70 100 Solubili'.y 3.3 5.0 7.1 10.1 14.5 19.7 32.5 56.0 When heated a little above its melting-point, it rapidly gives off oxygen, and finally only the chloride remains : If the reaction is stopped at an intermediate point, the residue is found to contain potassium perchlorate, KC1O 4 , which can readily be 398 A TEXT-BOOK OF INORGANIC CHEMISTRY separated from the chloride and unaltered chlorate by means of its slight solubility in water. Potassium chlorate is used as a source of oxygen, as well as in the firework and match industry (p. 242), and is also made use of in medicine. Potassium Perchlorate, KC1O 4 , prepared from the chlorate as just described, forms colourless rhombic crystals. At o, 25 and 50 100 grams of water dissolve 0.71, 1.96 and 5.34 grains of the salt respectively. It is a less powerful oxidizing agent than the chlorate ; it begins to give off oxygen when the temperature exceeds 400. Potassium Carbonate, K 2 CO 3 This salt was formerly pre- pared by extracting the ash of plants with water and evaporating, but since the discovery of the Stassfurt deposits is prepared from the chloride by a method analogous to the Leblanc process ; or by electrolysis, and treatment of the hydroxide with carbon dioxide. On account of the considerable solubility of potassium bicarbonate in water, a method of preparation analogous to the ammonia-soda process is not workable. Potassium carbonate is a colourless salt, extremely soluble in water. When the anhydrous salt is exposed to the air it takes up moisture and becomes liquid (deliquesces), owing to the fact that the vapour pressure of its saturated solution is less than the average pressure of aqueous vapour in the atmosphere (p. 387). It appears to form a number of hydrates with water, among others a dihydrate, but the facts have not yet been clearly established. The aqueous solution of the salt is alkaline owing to hydrolysis (p. 267) : Owing to its great affinity for water, the anhydrous salt is used as a drying agent. Potassium Bicarbonate, KHCO 3 This salt can be pre- pared by passing carbon dioxide into a concentrated solution of the normal carbonate ; being less soluble than the latter, it separates from solution and can be purified by recrystallization : K 2 CO 3 + H 2 O + CO 2 $2KHCO 3 1 . Potassium bicarbonate is more soluble in water than the correspond- ing sodium salt. The aqueous solution of the pure salt is practically neutral owing to the very slight ionisation of the HCO 3 ' ion (p. 324). Potassium Sulphate, K 2 SO 4 This salt occurs in the Stass- POTASSIUM 399 furt deposits chiefly as kainite, MgSO 4 ,KCl, 3 H 2 O, and polyhalite MgSO 4 ,2CaSO 4 ,K 2 SO 4 ,2H 2 O. It is obtained from kainite by treating with cold water, when a separation into the difficultly soluble schonite, MgSO 4 ,K L ,SO 4 ,6H 2 O, and the readily soluble magnesium chloride occurs. The latter is completely removed, and the schonite treated with excess of potassium chloride, when potassium sulphate crystal- lizes out : Potassium sulphate is also prepared by heating the chloride with concentrated sulphuric acid. Potassium sulphate occurs in rhombic prisms, and, in contrast to sodium sulphate, forms no hydrates. About i part of the salt dis- solves in 10 parts of water at 15. It finds application as a manure. Potassium Hydrogen Sulphate (potassium bisul- phate), KHSO 4 This salt is prepared by heating potassium chloride with excess of sulphuric acid : KC1 + H 2 SO 4 -KHSO 4 + HC1. It forms monoclinic prisms which melt about 210, and on further heating break up into the pyrosulphate and water (p. 296). 2KHSO 4 ->K 2 S 2 O 7 + H 2 O. It is extremely soluble in water, and the solution is strongly acid owing to the fact that the anion, HSO 4 ', splits off H' ions (p. 294). Potassium Nitrate This salt has been known from the earliest times. It is formed when nitrogenous organic matter, ammoniacal compounds, dung, etc., decay in contact with potassium salts, con- tained, for example, in wood ashes, and therefore the conditions were favourable for its formation in Ihe neighbourhood of dwellings at a very early stage of civilization. The oxidation in this case is accom- plished by the oxygen of the air with the help of certain bacteria, and does not take place in the absence of the bacteria. This process is used commercially in the so-called " nitre plantations." Dung and other animal refuse are mixed with wood-ashes and exposed in heaps to the action of the air. After a year or two the potassium nitrate is washed out and purified by recrystallization. Potassium nitrate is now usually prepared by mixing hot saturated solutions of sodium nitrate and potassium chloride. In such a mixture the four ions K*, Na - , Cl' and NO./ are present, and the salt, whose solubility product (p. 423) is first reached under definite conditions, will deposit from solution. A glance at the solubility table shows 400 A TEXT-BOOK OF INORGANIC CHEMISTRY that at high temperatures sodium chloride is the least soluble salt which can be formed from the ions ; it therefore deposits in consider- able amount. When the solution is then allowed to cool, potassium nitrate becomes the least soluble salt, and it separates in a fairly pure condition. It is then purified by recrystallization. It is evident that the purification of a salt by crystallization is most readily accom- plished when, as in the present case, the solubility changes greatly with change of temperature. Properties Potassium nitrate is a dimorphous substance ; at low temperatures it separates from solution in rhombic, at high temperatures in rhombohedral crystals. It melts at 339, and at higher temperatures loses oxygen and forms potassium nitrite : 2KNO 3 ->2KNO 2 + O 2 . The fused nitrate has powerful oxidizing properties. Sulphur (or charcoal) thrown into a crucible containing it catches fire and burns vigorously. The salt is largely used in the preparation of gun- powder. Gunpowder Gunpowder is a mixture of potassium nitrate, sulphur and charcoal. The proportions per cent, of the ingredients are not quite constant, but are approximately as follows : Potassium Nitrate, 75; Sulphur, 10; Charcoal, 15. On this basis we might expect the decomposition to be represented by the equation As a matter of fact, the reactions are much more complex. Besides potassium sulphide, the solid products include potassium sulphate, potassium thiosulphate and other salts ; and among the gases, besides carbon dioxide and nitrogen, carbon monoxide and hydrogen sulphide are formed. About 42 to 45 per cent, goes off as gas, and 55 to 58 per cent, of solid matter remains. The efficiency of the powder depends upon the almost instantaneous production in a small space of a considerable quantity of gas, which therefore exerts a very high pressure, and this pressure is still further increased by the heat given out in the combustion of the powder. It has been found that the pressure of the gaseous products reaches 6400 atmospheres, and the temperature reaches 2000. Potassium Sulphides The normal and acid sulphides, K 2 S and KHS, are prepared as described in connexion with the corre- sponding sodium salts. The former separates from solution as the pentahydrate, K 2 S,5H 2 O, in reddish crystals. LITHIUM 401 Poly sulphides, e.g. K 2 S 2 , K 2 S 3 , K^ and K 2 S 5 are also known. The penta sulphide is obtained, along with potassium thiosulphate, by heating potassium carbonate with excess of sulphur to 250 for some time in a closed vessel : The mixture of these two compounds was formerly called liver oj sulphur. Tests for Potassium As already mentioned, compounds of potassium are recognized by the characteristic lavender colour im- parted to ;i colourless flame. Like the sodium salts, the potassium salts are, almost without exception, readily soluble in water. The exceptions are potassium platinic chloride, PtCl 4 ,2KCl, and potassium acid tartr.ite, which are soluble with difficulty in water. Potassium compounds alone, or in the presence of sodium salts, can therefore be detected by adding platinic chloride, concentrated hydrochloric acid and excess of alcohol, when the double salt separates in yellow crystals, or by adding to the solution excess of tartaric acid*, and shaking, when the acid tartrate separates out as a crystalline pre- cipitate. It may be mentioned that ammonium salts give similar precipitates under the same conditions. LITHIUM Symbol, Li. Atomic weight -7. 03. Occurrence In contrast to the compounds of sodium and potassium, lithium compounds are found very sparingly in nature. It occurs, usually in very small proportion, in a number of silicates. Those richest in the metal are: spodumene and petaliti, both lithium-aluminium silicates, the former containing about 3.8 per cent, and the latter about 2 per cent, of lithium, and lepidolite or lithium mica, an a kali aluminium fluorosilicate which contains 0.8 to 2.7 per cent, of lithium. It also occurs in small quantity in certain mineral springs, and in the ash of certain plants, such as tobacco. Preparation of Metal The metal is usually obtained by electrolysis of the fused chlor'.de ; the salt is heated to fusion in a porcelain crucible, and subjected to electrolysis, a carbon rod being used as anode and an iron wire as cathode. Instead of the chloride, a mixture of lithium bromide with 10 to 15 per cent, of the chloride, vhich fuses at a lower temperature than the chloride, may be used. The metal may also be prepared by electrolysis of a solution of the chloride in pyridine. Properties Lithium is a silvery-white metal, harder than sodium, but than lead. It is the lightest solid known, its density being only 0.534 ; it floats on mineral naphtha ; it melts at 186. Lithiun rapidly tarnishes in moist air ; when heated above 200 it burns in air with a while flame like that of magnesium, forming lithium oxide, Li 2 O. It acts 26 402 A TEXT-BOOK OF INORGANIC CHEMISTRY on water at the ordinary temperature, but much less vigorously than the other alkali metals, lithium hydroxide, LiOH, being formed and hydrogen set free. Lithium Oxide, Li 2 O, is prepared by burning lithium in air or by heating lithium nitrate. It dissolves in water, forming lithium hydroxide, LiOH. The latter compound can also be prepared by boiling lithium carbonate with calcium hydroxide (cf. sodium hydroxide, p. 382). Lithium hydroxide is a strong base. Lithium Chloride, LiCl, is very soluble in water. Between o and 20 the dihydrate, LiCl,2H 2 O, from 20 to 100 the monohydrate, and above 100 the anhydrous salt separate from solution. Lithium Fluoride, LiF, differs from the fluorides of the other alkali metals in being very slightly soluble in water (about i in 400 at 18). Lithium Carbonate, Li 2 CO 3 , is also much less soluble in water than the other alkali carbonates ; 100 grams of water dissolve at o 1.54 parts, and at 20 1.33 parts of the salt. Lithium Phosphate, Li 3 PO4, separates from solution as a white crystalline precipitate when ammonia and sodium phosphate are added to a solution of a lithium salt. It is very slightly soluble in water (i part in 2540 parts of water at 18), and this is taken advantage of as a test for lithium compounds and in sepa- rating it from mixtures. Teats for Lithium Apart from the behaviour of the phosphate, lithium is characterized by the crimson colour it imparts to flame, and by the presence of a ed and a yellow line in the spectrum (p. 393). RUBIDIUM AND CAESIUM These rare elements were discovered by Bunsen and Kirchhoff (1860-1861) in the mineral waters of Durkheim by means of the spectroscope. They also occur in certain lepidolites (p. 401), and rubidium is now obtained almost entirely from the Stassfurt deposits, in which it occurs in very small amount as rubidium car- nallite, MgCl 2 ,RbCl. It is separated from the other alkali metals by taking advantage of the comparative insolubility of rubidium alum, Al 2 (SO 4 ) 3 ,Rb.>SO4, 24H 2 O. The only naturally occurring substance rich in caesium is the rare mineral pollux, a silicate of aluminium and caesium, which contains up to 32 per cent, of the latter element. Both are soft, silvery-white metals, which ignite at the ordinary temperature in moist air and also in dry oxygen. Rubidium melts at 38, caesium at 26.5. The chemistry of rubidium and caesium is practically identical with that of potassium. The only outstanding difference is that the former elements give polyhalogen compounds containing three and five atoms respectively of the halo- gens, e.g. RbBr 3 ,CsBr 3 ,CsBr 5 , in addition to the normal compounds with one atom of halogen. AMMONIUM SALTS It has already been mentioned that the salts containing the NH 4 or ammonium group show a remarkable analogy with those of the alkali metals, especially potassium, and it is therefore convenient to deal with them here. The chief source of ammonia and its compounds is the ammoniacal AMMONIUM COMPOUNDS 403 liquor of the gas-works (p. 214). When steam is blown through the liquor the easily hydrolyzable salts are decomposed and the ammonia passes off. The stable salts remaining in the liquor are decomposed by boiling with milk of lime. The process can be made continuous by using a rectifying column on the top of a still, the easily hydro- lyzable salts being decomposed by steam in the column, the stable salts by li ne in the still. The ammonia is* absorbed in hydrochloric acid. The solution is evaporated, and the solid ammonium chloride thus obtained purified by sublimation. From the chloride all the other ami ionium compounds can readily be obtained. Ammonium Chloride (sal ammoniac), NH 4 C1, prepared as mentioned above, usually occurs in feathery groups of small octa- hedral (m ire rarely cubic) crystals, which have a salt taste. When heated it sublimes before the melting-point is reached, and the vapour appears (from vapour density measurements) to be completely dis- sociated into ammonia and hydrogen chloride : NH 4 C1^NH 3 +HC1. Baker has shown that neither of the above reversible reactions takes place when the substances are perfectly dry ; ammonium chloride does not dissociate, nor is the chloride formed when ammonia and hydrogen chloride are mixed. At o, 20, and 40 water dissolves 29.7, 37.2, and 45.8 grams respec- tively of the chloride. As ammonia is a weak base (see below) the chloride is slightly hydrolyzed in aqueous solution : NH 4 C1 + HOH^NH 4 OH + HC1, and this is the more pronounced the higher the temperature. Further, as the hydroxide in aqueous solution is partly dissociated into water and ammonia, the latter escapes when the solution is boiled, so that the aqueous solution of ammonium chloride (and of all other ammo- nium salts), which is very faintly acid under ordinary conditions, becomes much more acid on boiling. Ammcnium Hydroxide, NH 4 OH When ammonia gas is passed in;o water it is readily absorbed, and the resulting strongly alkaline sc lution undoubtedly contains ammonium hydroxide, N H 4 O H, which is largely ionised into NH 4 ' and OH' ions. There is evidence, however, that the combination with water is not complete ; in other words, the solution contains free ammonia. The state of affairs may therefore be represented by the equations 404 A TEXT-BOOK OF INORGANIC CHEMISTRY At present no quite satisfactory method is known by means of which the relative amounts of these substances can be determined. Ammonium Sulphate, (NH 4 ) 2 SO 4 The crude salt is obtained by absorbing the ammonia from gas liquors in sulphuric acid, and is purified by recrystallization. It separates from solution in colourless, anhydrous, rhombic crystals, isomorphous with potassium sulphate. The crude salt is largely used as a manure. Ammonium Carbonates Commercial ammonium carbonate, also known as ammonium sesquicarbonate, is obtained by strongly heating in a retort a mixture of ammonium chloride and calcium carbonate and condensing the vapours in a receiver. It is purified by sublimation. The substance is a loose combination of a molecule of ammonium acid carbonate and a molecule of ammonium carbamate NH 4 NH 2 CO 2 (that is, a molecule of the normal carbonate minus i H 2 O), and therefore has the formula, NH 4 HCO 3 ,NH 4 NH 2 CO 2 . The compound usually occurs in crystalline masses. It loses am- monia at the ordinary temperature, for which reason it smells strongly of the gas, and also carbon dioxide, and finally only the bicarbonate remains. On heating above 60 it decomposes completely into water, ammonia and carbon dioxide. When the sesquicarbonate is dissolved in water the carbamate combines with the latter to form the normal carbonate, (NH 4 ) 2 CO 3 , so that the solution contains a mixture of normal and acid carbonate. Finally, when sufficient ammonia is added to convert the acid car- bonate into the normal carbonate, a solution containing the latter salt only is obtained : (i) (2) The normal carbonate, (NH 4 ) 2 CO 3 , separates in the solid form when ammonia gas is passed into a concentrated solution of the commercial carbonate. It gives up ammonia at room temperature, forming the bicarbonate, and when heated to 60 decomposes completely into carbon dioxide, ammonia and water. The acid carbonate, NH 4 HCO 3 , is obtained in the solid form by decomposition of the normal carbonate or commercial carbonate, as stated above, and also by passing carbon dioxide into an aqueous solution of the commercial carbonate and evaporating : The dry salt gives up practically no ammonia at room temperature. AMMONIUM COMPOUNDS 405 At 12.5 100 grams of water dissolve 17.1 grams, at 21 21.6 grams of the salt. In aqueous solution the salts are all considerably hydrolyzed, as both base and acid are weak. Ammonium Sulphides When hydrogen sulphide is passed in excess into a solution of ammonia, ammonium hydrosulphide, NH 4 HS, is obtained ; and if an equivalent amount of ammonia is added it might be assumed that the nonnal sulphide, (NH 4 ) 2 S, is formed. As, however, both base and acid are weak, the latter com- pound is almost completely hydrolyzed in solution : Even the hydrosulphide is hydrolyzed to some extent under ordinary conditions. The hydrosulphide can be obtained in rhombic leaflets, stable at room temperature but dissociating into ammonia and hydrogen sulphide on heating. The normal sulphide cannot be obtained at room temperature in the solid form. Solutions which contain only the acid or normal sulphide are practically colourless. On standing, the hydrogen sulphide formed by hydrolysis in the solution of the normal sulphide is slowly oxidized by atmospheric oxygen to water and sulphur, and the latter then combines with the normal sulphide to form yellow polysulphid.es, e.g. (NH 4 ) 2 S t , (NH 4 ) 2 S 5 and (NH 4 ) 2 S 7 . The importance in analysis of these solutions of "yellow" ammonium sulphide is referred to later (pp. 503). Ammonium, NH 4 As the group NH 4 behaves in many respects like an alkali metal it is interesting to inquire whether it has been isolated. In spite of many attempts, this has so far not been done. When sodium amalgam is added to a strong solution of ammonium chloride a porous, bulky mass is obtained, which breaks up into mercury, ammonia, and hydrogen. It was formerly supposed that this bul'iy mass was "ammonium amalgam," a solution of NH 4 in mercury, but no conclusive evidence on this point has been obtained. The same amalgam is obtained when an aqueous solution of an ammonium salt is electrolyzed with a mercury cathode. A further attempt to obtain ammonium by the electrolysis at - 95 of ammonium iodide dissolved in liquid ammonia was also unsuccessful. Tests for Ammonium Ammonium compounds are easily distinguished by the characteristic odour of ammonia when they are warmed with sodium hydroxide. 406 A TEXT-BOOK OF INORGANIC CHEMISTRY COMPARISON OF THE ALKALI METALS AND SUMMARY The main characteristics of the members of this group are that all are soft metals, which readily oxidize in the air, and decompose water at room temperature. They are the most electro-positive of all the metals ; the hydroxides are very strong bases ; practically all the salts are readily soluble in water. They impart characteristic colours to the Bimsen flame. They are all univalent elements. As in other groups, the physical and chemical properties vary regularly with increasing atomic weight. This is shown, for the more important physical properties, in the accompanying table : Li. Na. K. Rb. Cs. Atomic weight Density (solid) Melting-point . Boiling-point . Atomic volume 6.94 0-534 1 86 red heat I3-I 23.00 0.9741 97 878 23-7 39.10 0.863 62.5 758 45-4 85.45 1.52 38-5 6 9 6 55-8 132.81 1.87 26.5 670 71 As regards chemical behaviour, we have seen that the activity increases with increasing atomic weight. This is shown in combining with oxygen, rubidium and caesium bursting into flame in dry oxygen at room temperature, in the behaviour towards water, and in the heat of formation of their compounds. There are also certain characteristic differences. Lithium differs in several respects from the other four, as is not uncommon in the case of the first member of a family. Thus lithium carbonate and phosphate (and also the fluoride) are much less soluble in water than the corresponding salts of the other alkali metals : properties which recall the alkaline earth metals. The salts of sodium are more generally soluble in water than those of the other metals. Many of the salts of lithium and of sodium form stable hydrates ; the salts of potassium, rubidium, and caesium, on the other hand, are nearly all anhydrous. CHAPTER XXVII ELEMENTS OF GROUP I. SUB-GROUP B PHIS sub-group includes the following three metals, generally J- known as the members of the copper group : Copper (Cu) 63.57 Silver (Ag) 107.88 Gold (Au) 197.2 These elements show only a comparatively distant resemblance to the alkalis, although both belong to the first group. They are all univalent elements, but whilst silver only functions with this valency, copper also acts as a divalent and gold as a trivalent element. The metals differ markedly from the alkalis in having small affinity for oxygen ; they also melt at high temperatures. COPPER Symbol, Cu. Atomic weight, 63.57. History On account of its small affinity for oxygen, and the consequent readiness with which it can be obtained from certain ores (e.g. the carbonate), copper has been known from the earliest times. According to Berthelot, a copper age followed the stone age and preceded the bronze age. Utensils consisting almost entirely of copper were in use in Egypt and Babylonia anterior to 4500 B.C. Occurrence Copper occurs free in various parts of the world ; in great quantity near Lake Superior and in New Mexico (United States). In the combined state it is very widely distributed ; the more important ores are as follows: Ruby ore (cuprite), Cu 2 O. Copper pyrites or chalcopyrite, Malachite, CuCO 3 ,Cu(OH) 2 . Cu 2 S,Fe 2 S 3l or CuFeS 2 . Azurite, 2CuCO 3 ,Cu(OH) 2 . Purple copper ore, 3Cu^S,Fe^5 3 , Copp<:r glance or chalcocite, Cu 2 S. or Cu 3 FeS 3 . Metallurgy of Copper (i) From Non-sulphur Ores The preparat on of copper from the oxide and carbonate is very simple, 407 408 A TEXT-BOOK OF INORGANIC CHEMISTRY the ores being mixed with coal or coke and smelted in a blast- furnace : Naturally occurring copper is separated from the accompanying impurities by grinding and washing. The metal obtained by these or other processes can be purified by the electrolytic method (see below). (2) From Sulphide Ores These ores usually contain iron as well as sulphur, and their removal presents considerable difficulty, mainly owing to the fact that sulphur has a greater affinity for copper than for iron, and, further, copper sulphide, unlike the sulphides of iron and certain other metals, is only converted with great difficulty into FIG. 85. the oxide by roasting in air. The first step consists in roasting the ore in a reverberatory furnace (Fig. 85), whereby volatile impurities (arsenic, antimony, etc.), are burned off, and the iron sulphide is partially converted to oxide. The roasted ore is then heated to fusion (smelted), whereby a mixture of copper and iron sulphides, known as "matte" or "coarse metal," is obtained. The matte is then fused with sand, whereby a readily fusible slag of iron silicate is formed, which floats on the copper and can be run off. This alternate roasting and fusing is repeated several times till all the iron is removed and a mixture of copper and copper sulphide remains. This is then heated in a reverberatory furnace, when interaction between sulphide and oxide occurs : Cu 2 S + 2CuoCM>6Cu + SO 9 . The impure copper thus obtained must now be refined. This may be COPPER 409 done in a furnace or electrolytically. According to the former method it is melted in a reverberatory furnace in a current of air, whereby the impurities are oxidized, and either volatilize or combine with the silica lining the hearth to form a slag. Finally, the oxide of copper, formed to some extent in this process, is reduced by stirring the melted mass with poles of fresh wood; the gases formed by the burning of the wood effect the reduction of most of the cuprous oxide. Instead of the lengthy process above indicated, sulphide ores are now worked up by the converter method. The ore is first roasted, then fused, and placed in a Bessemer converter (p. 539) lined with silica. A blast of air is then forced through the mass, whereby the sulphur is burned off as sulphur dioxide, the arsenic and antimony also escape as oxides, and the ferrous oxide combines with the silica to form a slag. The copper thus obtained is then refined by the furnace method, or, better, by electrolysis. (3) Wet Methods These methods of extraction are used for ores which cannot profitably be treated in the dry way. When the ore is free from sulphur it is extracted with sulphuric acid, and the solution, which coi tains the copper as cupric sulphate, treated with ferrous chloride and sulphur dioxide, whereby reduction to cuprous chloride takes place. From the latter the copper is obtained by displacement with iron ; 2CuCl + Fe->2Cu + FeCl 2 . A methoc. in use for sulphide ores is fb roast with common salt, whereby c upric chloride is formed ; the latter is then extracted with water, and the copper precipitated by addition of scrap iron : CuCl 2 + Fe->FeCl 2 + Cu. Electrolytic Refining of Copper Since impurities such as iron, cuprous oxide, etc., greatly diminish the value of copper for commercial purposes, it is important to have a convenient method for obtaining the pure metal, and the requirements are fully met by the electrolytic method. Thick plates of fairly pure copper, which form the anode, are suspended in an electrolyte containing 8 to 14 per cent, of copper sulphate and 4 to 10 per cent, of sulphuric acid. Thin plates of copper form the cathode. When the current passes, the coppc r is dissolved from the anode and deposited in a very pure state on the cathode, so that the net result of the process is the con- veyance of copper from one pole to the other. Some of the impurities, 4 io A TEXT-BOOK OF INORGANIC CHEMISTRY such as silver, antimony, and cuprous oxide, do not dissolve, and are found in the so-called anode mud ; other impurities, such as zinc, dis- solve at the anode and remain in solution. Finally, the pure copper is stripped from the cathodes, melted, and cast in blocks. As there is very little polarization in this process, a small E.M.F. (less than i volt) is sufficient. Properties Copper is a bright-red metal of density 8.94 to 8.96 ; it melts at 1084, and boils at 2310, in absence of air. It is fairly hard, but very tough and flexible ; it can be drawn out into thin wire and beaten into thin sheets. It is a very good conductor of electricity, being surpassed in this respect only by silver. Its conductivity and malleability are greatly diminished by traces of impurities. In dry air copper becomes coated with a very thin film of oxides, which scarcely affect the colour, but protect it against further oxida- tion. In moist air it becomes covered with a thin coating of basic carbonate (malachite). At a red heat it combines fairly rapidly with oxygen to form black cupric oxide. Copper is not acted upon by water at any temperature, nor is it affected by dilute acids (other than oxyacids) at room temperature in absence of air. Nitric acid rapidly dissolves copper at the ordinary temperature (p. 225), cupric nitrate being formed. When heated with concentrated sulphuric acid, cupric sulphate is formed and sul- phur dioxide given off. Finely-divided copper dissolves .slowly on boiling with concentrated hydrochloric acid, with formation of cuprous chloride : 2HCl-x>CuCl + H. With free access of air, copper dissolves slowly even in dilute acids, as a result of the simultaneous action of the latter and oxygen, e.g. Under the same circumstances, copper is rapidly dissolved by ammonia, with formation of a deep blue solution : Cu + O + H 2 O + 2NH 3 ->Cu(NH 3 ) 2 (OH) 2 . On account of its stability in the air and resistance to acids, copper is largely used commercially for coinage, for covering the hulls of ships, for conveying electricity, etc., and also in electrotyping. Some important alloys of copper are mentioned in the next section. COPPER 411 Alloys of Copper The more important are brass, which con- tains 16 to 35 per cent, of zinc ; the bronzes, which consist of copper, zinc and tin in varying proportions, and usually lead as well ; German silver, which contains 2 parts of copper, i part of nickel, and I part of zinc, and is nearly white ; gun-metal, copper with 10 per cent, tin ; bell-metal, copper with 20 to 25 per cent, tin ; aluminium bronze, copper with 5 to 10 per cent, of aluminium ; phosphor bronze, copper with 5 to 15 per cent, of tin and 0.25 to 2.5 per cent, of phosphorus ; manganese bronze, copper with 30 per cent, of manganese. The copper coin ige of this country contains 95 parts of copper, 4 parts of tin and i part of zinc. Compounds of Copper Copper is the first metal we have met with which forms more than one series of salts. In cuprous salts, which are of the type CuX (X = univalent anion), the copper is univalent, vhilst in cupric salts, of the type CuX 2 , the copper is divalent. Cuprous salts of oxygen acids are practically unknown, but the cupric salts of these acids are quite stable. The relative stability of cuprous and cupric salts depends greatly on the condi- tions. At high temperatures the cuprous halides are the more stable. Cuprous salts are usually prepared by reducing cupric salts, and they tend to absorb oxygen from the air, with reformation of derivatives of divalent copper. Pure cuprous salts with a colourless anion are colourless (the Cu* ion appears to be colourless) ; cupric salts in dilute solution arc greenish-blue (that is, the Cu" ion is blue). CUPROUS SALTS Cuprous Oxide, Cu 2 O This substance occurs naturally as ruby coppe ore. It is obtained, mixed with cupric oxide, by gently heating copper in the air, and is also formed by the action of alkalis, e.g. sodium hydroxide, on cuprous chloride. It is most conveniently obtained by heating a cupric salt in alkaline solution with a reducing agent such as grape sugar. Cuprous oxide is a bright red powder, insoluble in water. When treated with halogen acids, cuprous halides are obtained ; with oxy- acids, on the other hand, the corresponding cupric salts are formed and metall'c copper separates : Cu 2 O + H 2 SO 4 ->CuSO 4 + Cu + H 2 O. It is soluble in ammonia, and the solution, which is colourless in absence of air, probably contains the compound, Cu(NH 3 ) 2 OH. 4 i2 A TEXT-BOOK OF INORGANIC CHEMISTRY Cuprous Chloride, 1 CuCl This substance is obtained by dis- solving cuprous oxide in hydrochloric acid, or, more conveniently, by boiling cupric chloride with hydrochloric acid and copper turnings, and then pouring the solution into boiled water : , + Cu->2CuCl. Cuprous chloride settles out as a white precipitate. When rapidly washed with glacial acetic acid, then with absolute alcohol and quickly dried, it is stable in dry air. In moist air it turns green, due to the formation of cupric oxychloride, Cu(OH)Cl or CuCl 2 'Cu(OH) 2 - When treated with hot water, cuprous oxide is formed by hydrolysis. With cold water, the hydrolytic action is accompanied by partia 1 decomposition according to the equation 2CuCl-CuCl 2 + Cu. Cuprous chloride is soluble in hydrochloric acid with formation of the complex compound, HCuCl 2 , and in ammonia, with formation of an addition compound, Cu(NH 3 ) 2 Cl. Both solutions are colourless when fresh, but in contact with air they rapidly become blue owing to the formation of cupric salts by oxidation. The solutions have the power of rapidly absorbing carbon monoxide, and are used for this purpose in gas analysis. The absorption appears to depend mainly upon the formation of an unstable compound, 2CuCl,CO,2H 2 O, which has been isolated in the solid form. Cuprous Iodide, Cul, is obtained as a precipitate when potas- sium iodide is added to a solution of a cupric salt : We may assume that cupric iodide, CuI 2 , is first formed, but, being unstable, it decomposes immediately into cuprous iodide and iodine. Cuprous iodide occurs in small, colourless, octahedral crystals, practically insoluble in water. Cuprous Cyanide, CuCN, is obtained, analogous to the iodide, by warming a solution of cupric sulphate with potassium cyanide : Cyanogen escapes as a gas, and cuprous cyanide is obtained as a 1 Cuprous salts are sometimes written with the double formula, cuprous chloride, for example, as Cu 2 Cl 2 , but molecular weight determinations in solution and the way in which the compounds react (cf. p. 132) lend support to the simple formulae. COPPER 413 white precipitate, very slightly soluble in water, but readily soluble in a solution of potassium cyanide. The latter solution contains a compound KCu(CN) 2 which is partly dissociated into K* and Cu(CN) 2 ' ions. Cuprous Sulphide, Cu 2 S, occurs naturally as copper glance, and can be prepared by direct combination of the elements or by heating cupric sulphide in a current of hydrogen at 260. It is a dark brown crystalline powder. CUPRIC SALTS Cupric Oxide, CuO This compound is obtained by strongly heating copper in air or oxygen or by heating cupric hydroxide, carbonate or nitrate. It is a black powder, and has considerable power for absorbing gases, such as carbon dioxide and water vapour. When heated to 1000, it loses part of its oxygen, and a mixture of cuprous and cupric oxides is obtained. It gives up its oxygen fairly readily to reducing substances, such as free hydrogen, organic com- pounds, etc., and for this reason is largely used in organic analysis. Cupric; Hydroxide, Cu(OH) 2 , is formed as a blue flocculent precipitate: when sodium or potassium hydroxide is added to the solution of a cupric salt. When the liquid in which it is suspended is boiled i: loses water, black cupric oxide being formed. Cupric hydroxide readily dissolves in ammonia, and the resulting solution, vhich probably contains the compound Cu(NH 3 ) 4 (OH) 2 , has the remarkable property of dissolving cellulose. Cupric hydroxide is a very weak base, and therefore all the cupric salts have an acid reaction owing to hydrolysis. Cupric Chloride, CuCl 2 , is obtained by dissolving cupric oxide or carbonate in hydrochloric acid, and separates in blue needles of the dihydrate, CuCl 2 ,2H 2 O, on concentrating the solution. The anhy- drous salt can be prepared by heating the dihydrate in a current of dry hydrogen chloride, and is brownish-yellow. The concentrated aqueous solution of cupric chloride is green, but on progressive dilution the colour gradually changes and finally attains the deep blue colour characteristic of dilute solutions of all cupric salts. The simplest explanation of these observations is that con- centrated solutions contain almost exclusively the non-ionised salt, and that as the solution is diluted the colour approximates more and more to that of the divalent Cu" ion (or, more probably, of the hydratecl ion Cu(H 2 O) 4 "). Although the increased ionisation, accom- panied by increased hydration, is no doubt the main, it is not the 4 i4 A TEXT-BOOK OF INORGANIC CHEMISTRY only factor concerned in the changes of colour, as it has been shown that in concentrated solutions complex cations containing chlorine, e.g. CuCl', are also present. When ammonia is added in excess to a solution of cupric chloride a deep azure-blue solution is obtained, from which the compound Cu(NH 3 ) 4 Cl 2 ,H 2 O separates on concentration in dark blue needles. The compound Cu(NH 3 ) Cl 2 , a dark blue powder, is obtained by the action of ammonia gas on the anhydrous salt. When these com- pounds are heated to 150 ammonia is given off, and the compound CuCl 2 ,2NH 3 remains as a dark green powder. Compounds of Copper Salts with Ammonia When ammonium hydroxide is added to any cupric salt, azure-blue solutions are obtained which contain compounds of the salts with ammonia. There is evidence that the copper in these solutions (including that of cupric hydroxide) is present mainly if not entirely in the form of salts of the type Cu(NH 3 ) 4 X 2 , which on dissociation give rise to divalent Cu(NH 3 ) 4 " ions. Such salts can also be isolated in the solid form, but the composition in this case is very varied. Cuprous salts also dissolve in ammonia giving colourless solutions which very rapidly become blue in the air owing to the formation of cupric compounds by oxidation. In some cases at least these solutions contain complex salts of the type Cu(NH 3 ) 2 X, which yield cations of the type Cu(NH 3 ) 2 ' by ionisation. Cupric Sulphate (blue vitriol) CuSO 4 This salt is obtained by heating copper with concentrated sulphuric acid, or by dissolving the oxide or carbonate in dilute acid. On the commercial scale, it is prepared by roasting copper pyrites in a current of air, extracting with water and evaporating, when the salt separates in large blue triclinic crystals of the formula CuSO 4 ,5H 2 O. On heating for some time at 1 10 the salt loses 4H 2 O, and becomes anhydrous when heated for some time at 210. The anhydrous salt is grayish-white, has a great affinity for water, and is used as a dehydrating agent. On heating more Strongly, sulphur dioxide and oxygen are given off, and basic salts of the type ;rCuSO 4 ,jj/CuO are formed. -.At 15, 30, and 50, the solubility in 100 grams of water is 19.3, 25.5, and 33.6 grams respectively, calculated on the anhydrous salt. The behaviour of copper sulphate towards ammonia is similar to that of the chloride. From the azure-blue solution containing excess of ammonia the compound Cu(NH 3 ) 4 SO 4 ,H 2 O has been obtained COPPER 415 in blue crystals. A number of other solid compounds have been described, but their existence is somewhat doubtful. Cupiic Nitrate, Cu(NO 3 ) 2 , is obtained by dissolving copper, the dioxide or carbonate in dilute nitric acid and evaporating the solution. Above 24.5 a trihydrate separates from solution, at lower temperatures a hexahydrate, in blue tabular crystals, and at still lower temperatures the compound Cu(NO 3 )<>,9H 2 O. Cupric Carbonates The normal carbonate of copper is not known. When a solution of copper sulphate is heated with sodium carbonate, the bluish-green colloidal precipitate (of varying com- position) first obtained changes on standing in contact with the mother liquor to the crystalline basic carbonate 3CuCO 3 ,3Cu(OH)2,H 2 O. Important naturally occurring basic carbonates are malachite, CuCO 3 ,Ci (OH) 2 , and azurite, 2CuCO 3 ,Cu(OH) 2 , both of which have been prepared artificially. Cupric Sulphide, CuS, is obtained as a black precipitate by passing hydrogen sulphide through a solution of a cupric salt. It is practically insoluble in water and in hydrochloric acid, but is slightly soluble in yellow ammonium sulphide. Tests for Copper Copper salts (in the cupric form) are characterised by the blue colour of their solutions, by the azure-blue solution obtained on adding excess of ammonia, and by the formation of a black precipitate of cupric sulphide when hydrogen sulphide is passed through a solution. Oxidation and Reduction Up to the present we have re- garded the process of oxidation as being associated with the addition of oxygen to, or the removing of oxygen from, a compound ; con- versely, when oxygen is removed from, or hydrogen added to, a compound the latter is said to be reduced. Free oxygen, many oxides and peroxides, e.g. manganese peroxide, and many compounds rich in o>ygen, e.g. nitric acid, chloric acid, are familiar oxidizing agents. On the other hand, free hydrogen, "nascent" hydrogen (p.- 498) and hydriodic acid are reducing agents. Many substances which contain no oxygen act as oxidizing agents, e.g. chlorine, and similarly, substances containing no hydrogen may act as reducing agents. These indirect oxidizing and reducing agents usually act through the intervention of water. Thus chlorine combines with water to form hydrogen chloride, the oxygen becoming available for oxidizing purposes (p. 86) : 416 A TEXT-BOOK OF INORGANIC CHEMISTRY and similarly, sulphur dioxide, when it acts as a reducing agent, gives sulphuric acid and hydrogen : SO 2 + H 2 O->SO 3 +H 2 It has been repeatedly pointed out that oxidation is always accom- panied by reduction ; while one substance is oxidized, the oxidizing agent undergoes reduction. In connexion with metals such as copper, which have two valencies, it is necessary to extend somewhat the use of the terms oxidation and reduction. As cuprous salts are converted to cupric salts by the action of oxidizing agents, and conversely, cupric salts are converted to cuprous salts by reducing agents, we say that cupric chloride, CuCl 2 , for instance, represents a higher state of oxidation than cuprous chloride, CuCl, although neither oxygen nor hydrogen are directly concerned. When we consider that the oxide corresponding with cuprous chloride is Cu 2 O, and that corresponding with cupric chloride is CuO, it is evident that this way of representing the matter is justified. The same considerations apply to metals as compared with salts. As the oxide corresponding with cuprous chloride is Cu 2 O, it is usual to speak of the conversion of cuprous chloride (or any other salt of copper) to the metal as reduction, and the formation of the salt from the metal as oxidation. It follows that the combination of a metal with almost any element (except hydrogen) may be regarded as oxidation. The matter becomes clearer on the basis of the electrolytic dis- sociation theory. Comparing copper, cuprous chloride and cupric chloride, we may say that, as far as salts are concerned, an increase in the number of positive charges (or a diminution in the number of negative charges') denotes oxidation; decrease in the number of positive charges or an increase in the number of negative charges denotes reduction. Further illustrations of these statements will be given later. Electromotive Force 1 At a very early stage of our work, we saw that when zinc and copper are connected by a wire and dipped into dilute sulphuric acid, an electric current passes through the wire (p. 13). This process was regarded as a transformation of chemical into electrical energy, the source of the energy being the 1 For full details, see Physical Chemistry, chap. xiii. COPPER 417 oxidation of the zinc, which in the process is transformed into zinc sulphate. The hydrogen is liberated at the surface of the copper. An arrangement of this sort is termed a galvanic cell or battery. For our present purpose, it is rather more instructive to use a solution of copper sulphate as electrolyte in contact with the copper instead of the sulphuric acid previously used. Further, although a current is obtained when the electrodes are simply placed at some distance apart in the electrolyte, if is preferable, in order to minimize mixing of the solutions round the two poles, to separate them by some arrangement which at the same time allows the current to pass. An earthenware pot answers the purpose satis- factorily. The battery we are considering, then, is made by placing in the outer vessel a solution of copper sulphate in which a copper plate is partially immersed, and in the inner vessel a porous earthenware pot a rod of zinc dipping in a solution of sulphuric acid. When the poles are connected by a wire, an electric current passes round the circuit in the solution positive electricity flows from zinc to copper, and in the connecting wire from copper to zinc and simultaneously zinc dissolves at the zinc pole and an equivalent of copper is deposited on the copper pole. The direction of the current is shown in the accompanying diagram. In the external circuit. Zn | H,SO 4 | CuSO 4 | Cu In the solution. The pole where positive electricity leaves the solution in this case the copper plate is termed the positive pole, the electrode at which positive electricity enters the solution is called the negative pole. The chemical change accompanying the passage of the current may be written thus : Zn + CuSO 4 ->ZnSO 4 + Cu, or, from the ionic point of view : Zn + Cu"->Zn"+Cu ; that is, it is simply the displacement of copper from solution by zinc. This change, as we have already seen, takes place directly ; when zinc is put in a solution of copper sulphate, copper is deposited and an equivalent of zinc goes into solution. As the change takes 27 4 i 8 A TEXT-BOOK OF INORGANIC CHEMISTRY place spontaneously, we anticipate that there is a diminution of energy in the process, and this energy is obtained in the galvanic cell as electrical energy. A word is necessary regarding the factors of electrical energy. The work which an electric current can do depends not only on the quantity of electricity which passes, but on the potential or electromotive force in the circuit. The quantity of electricity is measured in coulombs, the electromotive force in volts. Electrical energy is proportional to the product of the electromotive force and the quantity of the electricity. Thus the same amount of light will be given out when 5 coulombs pass through an electric glow-lamp at a potential of 10 volts, as when 2 coulombs pass through at a potential of 25 volts. The Potential Series of the Elements It can be shown that the electromotive force of the zinc-copper cell, arranged as we have described, is rather more than one volt. If magnesium is used in place of zinc, the E.M.F. is about 1.8 volts. Similar cells can be built up with other metals in place of zinc and copper, and in every case, when one metal can displace the other from combina- tion an electric current is obtained. Now it can be shown that the E.M.F. of a cell built up with two -metals in this way is a measure of the tendency of one element to displace the other from combina- tion. Proceeding on these lines, the metals have been arranged in order in such a way that each metal can displace from combination all those following it in the series, and is displaced from combination by all those preceding it in the series. The so-called potential series of the metals is as follows, the potential of hydrogen being taken as zero : Na Mg Al Mn Zn Cd Fe Co Ni Pb + 2.58 1.482 1.276 1.075 0.770 0.420 0.334 0.232 0.228 0.151 H 2 Cu Hg Ag Pt Au o -0.329 -0.753 -0.771 -1.42? -1.65? These numbers represent the potential difference between a metal and the solution of one of its salts in normal solution, referred to the potential between a normal solution of an acid and a platinum electrode saturated with hydrogen, which is taken as zero. The actual zero chosen is a matter of indifference ; the important point is the difference of potential between one substance and another. The table shows that magnesium, aluminiym and zinc can displace SILVER 419 iron from its salts ; iron, on the other hand, can displace lead, copper and silver from combination. A very interesting point is the light thrown by this list on the behaviour of the metals with regard to the displacement of hydrogen from acids. All the metals preceding hydrogen in the list should be able to liberate it from acids, and we have seen in a number of instances (others are given later) that such is the case. On the other hand, copper, mercury and silver. cannot displace hydrogen from acids ; on the contrary, hydrogen should be able to displace these metals from their salts. This is not observed under ordinary circumstances, probably owing to the reactions being exceedingly slow, but hydrogen occluded in platinum can displace some of the metals from combination owing to the accelerating influence of platinum on the rate of reaction. Another method of stating the facts is that when the metals are arranged in the order of their potentials, they are arranged in the order of their solution pressures (p. 85). Metals such as sodium have so high a solution pressure that they decompose water at room temperature ; the solution pressures of copper and mercury, on the other hand, are extremely small. The metal, of course, goes into solution as positive ions, this being the only form in which it can dissolve in ordinary solvents. The metals near the beginning of the potential series, the strongly electro-positive metals, give rise to hydroxides which are strong bases ; whereas those near the end of the series give rise to hydroxides which are relatively weak bases. The magnitude of the electro-positive character and the basic character do not by any means run parallel, as will be evident in the sequel (cf. p. 465). SILVER Symbol, Ag. Atomic Weight 107.88. Chemical Relations In its compounds silver is almost invariably univalent. The hydroxide is a fairly strong base and therefore silver salts are not appreciably hydrolyzed. Like copper, it has a tendency to form complex ions. It has very little affinity for oxygen, and silver salts are therefore readily reduced to the metallic state. Occurrence Silver occurs free in nature, usually in small pro- portion distributed in quartz and other rocks, occasionally in large masses. In combination it occurs as sulphide mixed with lead 420 A TEXT-BOOK OF INORGANIC CHEMISTRY sulphide (galena), with which it is isomorphous. The more im- portant silver ores are : silver glance or argentite, Ag 2 S ; pyrargyrite or ruby silver ore, 3Ag 2 S,Sb 2 S 3 or Ag 3 SbS 3 ; proustite, sAg 2 S,As,S 3 or Ag 3 AsS 3 , and horn silver, AgCl. Silver is always present in native copper and (as sulphide) in copper pyrites (p. 407). The supply of silver comes mainly from the United States, Mexico, Bolivia, Germany and Australia. Metallurgy of Silver The processes used in obtaining silver depend upon the nature of the ores in which it occurs. When metallic lead is prepared from galena containing silver sulphide, the lead contains all the silver as metal, and its separa- tion is an important industry. Two processes are in use for this purpose : (i) the Pattinson process, (2) the Parkes process. (1) The Pattinson Process The alloy of silver and lead, which is relatively poor in silver (as little as 0.02 per cent, is commercially workable) is melted and allowed to cool, when lead separates out first in crystals, which are removed (cf. p. 197). This process is repeated till the eutectic point of lead and silver is nearly reached (about 2.25 per cent, silver). The alloy is then subjected to cupella- tion, which consists in heating it in a reverberatory furnace, the hearth of which is lined with bone-ash, while a blast of air is passed over the surface. The lead is thus converted into oxide, which at the high temperature melts and is driven by the blast of air over the edge of the "cupel." In course of time all the lead is thus removed, and its disappearance is indicated by the sudden appearance so-called "flashing" of the bright metallic surface of the silver. (2) The Parkes Process This process, which has now largely superseded the Pattinson process, depends upon the fact that when zinc is added to the alloy of lead and silver, and the metals are fused and thoroughly mixed, the zinc and lead separate almost completely into two liquid layers, and nearly all the silver passes into the upper zinc layer. The latter layer solidifies first, and is removed by means of sieves from the surface of the fused lead, heated gently to allow most of the lead to run off, distilled to remove the zinc, and finally the still remaining lead is removed by cupellation. Other Processes The so-called A malgamation Process, which has been in use in Mexico for centuries, is carried out as follows : The ores, chiefly sulphide, are first reduced to powder, copper sulphate and sodium chloride are then added, and thorough mixing is effected by the treading of mules. The silver sulphide reacts with the cuprir SILVER 421 chloride (formed by double decomposition) to give silver chloride and cupric sulphide : Ag. 2 S + CuCl 2 -x>Ag Cl + CuS. Mercury is then added and thoroughly incorporated by treading. The silver chloride is decomposed, with formation of mercurous chloride and liberation of silver, which dissolves in the excess of mercury : 2 AgCl + 2Hg-Hg 2 Cl 2 + 2 Ag. The amalgam is then separated and washed, and the mercury removed by distillation. According to an improved amalgamation process in use in Germany, the ores are roasted with common salt to form silver chloride, and the latter reduced to the metal by rotating the ore in casks containing iron plates ; the silver is then extracted by means of mercury, and the latter distilled off. A number of Wet Processes are also in use ; the following may be mentioned : (a) The Augustine Process. The ores are roasted with sodium chloride, and extracted with a concentrated solution of the same salt, in which silver chloride is soluble. (b) The Ziervogel Process. The ore is roasted in such a way that the silver sulphide is converted to sulphate, which is then extracted with water. In both (a) and (b) the silver is precipitated from solution by means of copper. (c) The Patera Process. The ores are roasted with salt, extracted with sodium thiosulphate, and the silver precipitated as sulphide by means of sodium sulphide. The latter is then converted to silver by roasting in a reverberatory furnace. Properties Silver is a white, lustrous metal which fuses at 964 and boils at 1950 ; it can be distilled in the oxyhydrogen flame. Its density is 10.49 to 10.50. It is the best conductor of heat and electricity among the metals. It is very malleable and ductile, and can be drawn out into very thin wires. Fused silver has the remark- able property of dissolving oxygen ; just above its melting-point one volume of silver takes up 20.3 volumes of the gas. As oxygen is prac- tically insoluble in solid silver, the gas escapes as the metal solidifies, which gives rise to the so-called "spitting" of silver, and causes curious excrescences on the solidified metal. Silver is obtained in the amorphous form by reducing silver sulphide with hyc rogen. The gray or black silver obtained by reducing the salts in solution, e.g. with ferrous sulphate, is ordinary silver in the form of minute needles. 422 A TEXT-BOOK OF INORGANIC CHEMISTRY Silver is obtained in " colloidal" solution by passing the electric arc between silver rods dipping under water, or by reducing silver salts in solution with certain reagents, e.g. with ferrous citrate. According to the nature of the reducing agent and the conditions, the solutions differ in colour (red, green, black, etc.), which is doubtless connected with the degree of fineness of division of the colloidal particles. Silver does not combine with oxygen even on heating. It is insoluble in hydrochloric or in dilute sulphuric acid, but on heating with concentrated sulphuric acid silver sulphate is formed and sulphur dioxide given off: It is readily dissolved by nitric acid, silver nitrate being formed : Alloys Pure silver is rather too soft to use for commercial pur- poses, and the alloy with copper is always employed. British silver coins contain 92.5 per cent, of the metal ; those of the United States and continental countries only 90 per cent. For jewellery, plate, etc., alloys containing 75 to 95 per cent, of silver are used. Oxides of Silver Two oxides of silver, the normal oxide, Ag 2 O, and a peroxide, AgO, appear to be definitely known. The existence of a third oxide, Ag 4 O, is doubtful. Silver Oxide, Ag 2 O, is obtained as a black, amorphous powder by adding potassium hydroxide to a solution of silver nitrate, or by boiling silver chloride with potassium hydroxide. By these methods we would expect the formation of silver hydroxide, AgOH, but this compound is unstable, and decomposes almost completely into silver oxide and water. The decomposition is not, however, quite complete, as the oxide is slightly soluble in water, and the solution is distinctly alkaline, and therefore contains silver hydroxide. Silver hydroxide, AgOH, although not so strong a base as the alkalis, is stronger than ammonium hydroxide, so that silver salts are not appreciably hydro- lyzed in solution. On heating to 250, silver oxide decomposes rapidly into silver and oxygen. It is reduced to metallic silver by heating in hydrogen at 100. It is readily soluble in ammonia, and the solution contains the complex compound Ag(NH 3 ) 2 OH, which is a strong base. Silver Peroxide, AgO (or Ag 2 O 2 ), is obtained as a black powder by the action of ozone on silver or by the electrolysis of a solution of SILVER 423 silver nkrate. As prepared by the latter method, it is always con- taminated with adhering silver nitrate. Silver Chloride, AgCl This salt is formed as a white curdy precipitate when hydrochloric acid or a soluble chloride is added to a solution of silver nitrate. It melts at 480-490, and can be vaporized without decomposition. On exposure to light, the originally white salt becomes violet, then brown, and the odour of chlorine can be detected. This change is due to the formation of a lower chloride according to the equation 4AgCl->2Ag 2 Cl + Cl 2 (see silver salts in photography, below). The solubility of silver chloride in water is very small, amounting to 1.56 :c io~ 5 mols per litre (i part in 430,000 parts of water) at 25. It is readily soluble in ammonia, in sodium thiosulphate, in potassium cyanide, in concentrated hydrochloric acid and in concentrated solu- tions of alkali chlorides with formation of complex ions. Solubility Product. Complex Ions containing Silver From the considerations advanced in previous chapters we know that in a saturated solution of a salt at a definite temperature there are two equilibria : (a) that between the solid salt and the non-ionised salt in the solution ; () that between the non-ionised salt and the ions. In the case of silver chloride, they may be represented as follows : Ag' + Cl'^AgCl (in solution) ^ u AgCl (solid). Further, since the solution is saturated, the concentration of the non- ionised salt in the solution must be constant at constant temperature, just as sugar has a constant solubility at a definite temperature. Applying the law of mass action to the equilibrium in solution (p. 166), we have [Cl'] = K[AgCl]=S. As the right-hand side of the above equation is constant at constant temper iture, the product of the concentration of the two ions the so- called solubility product, S, is also constant at constant temperature. The great importance of these considerations will be evident. If the solubility product in a saturated solution of silver chloride (which amounts to (1.25 x 10 B ) 2 =i.56 x io~ 10 mols per litre at 25) is exceeded in any way, for example, by adding Cl' ions, the 4 2 4 A TEXT-BOOK OF INORGANIC CHEMISTRY equilibrium is displaced towards the right, and the excess of non- ionised silver chloride falls out of solution ; this change proceeds till the original concentration of non-ionised silver chloride is re-established. If, on the other hand, silver ions are removed in some way, the product of the ions is no longer equal to the solubility product, silver chloride will dissolve, and this change will proceed until the solubility product is re-established. It is, of course, evident that the Ag' and Cl' ions need not be present in equivalent proportions. Equilibrium is established between ions and dissolved non-ionised salt when the solubility product is attained with any ratio between concentrations of the ions concerned. Two illustrations of the importance of the solubility product will be given. In quantitative analysis it is often necessary to precipitate a salt, e.g. silver chloride, as completely as possible from solution. The saturated solution of silver chloride contains about 0.0023 grams per litre of the salt at 25. If, however, sodium chloride is added till the solution is say i/ioo molar with reference to Cl' ions (that is, the Cl' concentration is increased 1000 times) silver chloride falls out of solution till the concentration of Ag' ions is about i/iooo of its former value (as is clear from the equation), and the amount of silver chloride remaining in solution is quite negligible. The converse change occurs when ammonia is added to a saturated solution of silver chloride in equilibrium with the solid salt. The Ag' ions combine with ammonia to form complex Ag(NH :j ) 2 ' ions, and the product of the ionic concentrations [Ag'] x [Cl'J falls far below the solubility product ; silver chloride dissolves to re-establish the equilibrium, and this proceeds till the solubility product is again reached. Similarly, the solubility of silver chloride in concentrated hydrochloric acid is accounted for by the formation of complex AgCl/ ions : which are only very slightly dissociated. into Ag 1 and Cl' ions. The behaviour of an unsaturated solution containing complex ions will clearly depend upon the stability of the complex ion. In a solu- tion of silver chloride in ammonia there are two equilibria : Ag(N H 3 ) 2 Cl$Ag(N H 3 ) 2 - + Cl' and but the Ag' ion concentration is extremely small. It follows that none of the ordinary reagents will produce a precipitate in an ammo- SILVER 425 niacal solution of silver chloride unless the solubility product of the substance which can be formed is excessively small. Silver sulphide is practically the only silver compound which answers these require- ments (solubility product, 3.9 x icr w mols per litre at 25), and, therefore, hydrogen sulphide causes a precipitate in an ammoniacal solution of silver chloride. Silver Bromide, AgBr, and Silver Iodide, Agl, are prepared by methods analogous to those described for the chloride. Both are yellow salts, which are still less soluble in water than the chloride. The solubility of the bromide amounts to 0.00012 grams, that of the iodide to 0.0000023 grams per litre at 25. The bromide is slightly soluble, the iodide practically insoluble in ammonia ; the explanation of this behaviour will be evident from the previous section. Both salts are soluble in sodium thiosulphate (p. 299). They are acted on by light in the same way as the chloride, and upon this depends their use in photography. Silver Fluoride, AgF, is obtained by dissolving silver oxide in hydrofluoric acid and evaporating, when, according to the conditions, the compound AgF,H 2 O, colourless quadratic crystals, or AgF, 2H 2 O, colourless prisms, is obtained. The anhydrous salt is amorphous and generally yellow, but the pure salt should presumably be colour- less. It differs from the other silver halides in being readily soluble in water. Silver Cyanide, AgCN This salt is obtained as an amorphous white precipitate when excess of silver nitrate is added to potassium cyanide. It is readily soluble in ammonia. When potassium cyanide is used in excess, the cyanide dissolves with formation of a complex compound, KAg(CN) 2 . This compound is ionised mainly according to the equation and the an ion is very slightly ionised thus : so that the Ag 1 ion concentration is very minute. The solution of potassium silver cyanide is used in electroplating. For this purpose the object to be plated is immersed in a solution of the cyanide, and forms the cathode, the anode being a silver plate. During electrolysis silver is dissolved from the anode and deposits on the object as a uniform adherent coating. The advantage of using the cyanide for this purpose appears to depend upon the small Ag' ion 426 A TEXT-BOOK OF INORGANIC CHEMISTRY concentration, as highly ionised silver salts, e.g. silver nitrate, give a much inferior coating. Use of Silver Halides in Photography Modern photo- graphy is based on the changes undergone by silver halides on exposure to light. The "dry" sensitive plates now almost exclusively used are glass plates coated with a thin film of gelatine containing silver bromide in suspension. When a plate of this kind is exposed in a camera, the silver bromide is affected in some parts more strongly than in others, according to the relative intensities of the light re- flected from the different parts of the object. At this stage no visible change has taken place on the plate, but when it is acted on by a " developer " for example, the solution of a reducing agent such a? ferrous oxalate silver is set free in the metallic form in those parts affected by the light, whilst the non-illuminated silver halide is un- affected. The unaltered halide is then removed by treatment with sodium thiosulphate, which does not affect the deposited silver this process is known as fixing and the result is the so-called " negative " on which the illuminated parts are dark (owing to the free silver), and the unilluminated parts clear (owing to the removal of the halide). Up to this stage, the process must be carried through in the absence of daylight. The negative is then laid on a paper coated with a sensitive film (e.g. silver chloride in gelatine) and exposed to direct sunlight. Those parts of the negative where the silver was deposited allow very little or no light to pass through, according to the amount of silver ; those parts free from silver allow all the light to pass. After " fixing," a positive which reproduces the illumination of the original object is obtained. Of the silver halides the bromide, being the most sensitive to light, is most largely used in photography. The chemical changes in the process are not thoroughly understood. The first effect of light is probably to form a "subhalide" or " photohalide," e.g. Ag 2 Br. Silver Sulphate, Ag 2 SO 4 , is prepared by the action of sulphuric acid on silver, silver oxide, or carbonate. It occurs in rhombic crystals, isomorphous with sodium sulphate. Silver Nitrate, AgNO 3 This salt is obtained by the action of nitric acid on silver. It occurs in colourless, rhombic crystals which are not hygroscopic ; it melts at 208. It is extremely soluble in water; at o 100 parts of water dissolve 115 parts, and at 20 215 parts of the salt. When heated above its melting-point it decomposes into silver oxide, oxygen and oxides of nitrogen ; above 300 complete decomposition into metallic silver occurs. GOLD 427 In contact with organic matter and other reducing agents it is reduced (most rapidly on exposure to light) to metallic silver, which is usually deposited in a black, finely-divided form. For this reason it is used in preparing marking inks ; it is also employed in medicine as a caustic. From solutions of silver nitrate in ammonia the compound AgNO 3 ,2NH 3 separates in rhombic crystals. Silver Carbonate, Ag 2 CO 3 , is obtained as a light-yellow powder by double decomposition between silver nitrate and potassium carbo- nate in solution. It is almost insoluble in water. Tests for Silver The formation of a white, curdy precipitate of silver chloride, soluble in ammonia but insoluble in nitric acid, when hydrochloric acid or a chloride is added to the solution of a silver salt, is characteristic. The precipitation of dark red silver chromate, Ag 2 CrO 4 , when potassium chromate is added to the solution of a silver salt, is also a distinctive test. GOLD Symbol, Au. Atomic weight, 197.2. General Characters Gold forms two series of salts, the aurous compounds, of the type AuX, in which it is univalent, and the auric compounds, AuX 3 , in which it is trivalent. Auric oxide, Au 2 O 3 , has, however, acidic as well as basic properties. All gold compounds are readily reduced to metallic gold by heat or by reducing agents. Occurrence Gold occurs in nature mainly in the free con- dition in veins of quartz and in alluvial deposits. It occurs free in small amount in many sulphide ores, e.g. iron, copper, and arsenical pyrites, and lead and zinc sulphides. In the combined condition it is always associated with tellurium, and also usually with silver. The more important ores are sylvanite t (Au,Ag)Te, 1 calaverite, (Au,Ag)Te 2 , with very little silver, and petzite, (Au,Ag)Te 2 , with a considerable proportion of silver. The chief gold-producing countries, in order of relative importance, are the Transvaal, the United States, and Australia. Metallurgy of Gold In the case of alluvial deposits gold is separated by washing, advantage being taken of its high specific gravity. When it occurs in quartz reefs, as in the Transvaal, the rock is crashed, reduced to a fine powder by stamping, and caused to 1 The formula (Au.Ag)Te denotes an isomorphous mixture of AuTe and AgTe in varying proportions. 428 A TEXT-BOOK OF INORGANIC CHEMISTRY flow, by means of a stream of water, over copper plates amalgamated with mercury, the latter retaining the gold. The amalgam is then scraped off and the mercury removed by distillation. Besides the above process, two methods are in use for the extraction of gold by dissolution, for which potassium cyanide and chlorine respectively are used. The cyanide method is carried out as follows : The material is first extracted with dilute alkali to remove certain impurities, and then with 0.35 per cent, potassium cyanide solution with free access of air. The gold is then precipitated by means of zinc or by electrolysis ; in the former case the zinc is removed by roasting. In order that gold may be dissolved by potassium cyanide, the presence of oxygen is necessary. The equation is as follows : In the electrolytic method of separation, iron anodes and lead cathodes are used. The latter are finally subjected to cupellation in order to obtain pure gold. The chlorine method is more troublesome, inasmuch as the ore has first to be roasted and then fused with common salt in order to convert all the other metals, except the gold, into chlorides. On treatment with chlorine, the gold passes into solution as auric chloride, AuCl 3 , and is precipitated as sulphide by hydrogen sulphide, or as metal by ferrous sulphate : Properties Gold is a bright yellow, rather soft, lustrous metal, which melts at 1063 : density, 19.3 to 19.5 at 18. It is extremely malleable and ductile, and can be beaten out into very thin sheets, which are green by transmitted light. It is a very good conductor of heat and electricity. Gold is not acted on by dry or moist air or oxygen, even on heating, and is not affected by any single acid (except selenic acid, p. 305). Chlorine readily dissolves it with formation of auric chloride, AuCl 3 , as does a mixture of nitric and hydrochloric acids (aqua regia), the action in this case depending upon the presence of free chlorine (p. 235). It is readily dissolved by potassium cyanide solution with access of air (see above). Fused alkalis also act on it. Compounds of gold are very readily reduced to the metallic condition; the precipitated gold varies in appearance, according to the nature of the reducing agent and the conditions. With ferrous sulphate or arsenious acid a brown precipitate is obtained. With GOLD 429 other reducing agents, such as formaldehyde and hydrazine, gold is obtained in colloidal solution showing brilliant colours, such as purple, red, and blue. Alloys 'Like silver, gold is too soft to use alone for coins and other articles, and is always alloyed with copper. The purity or fineness of gold is expressed in carats, pure gold being 24 carats. Gold jewellery, medals, etc., are 14 to 18 carats, that is, 24 parts of the alloy contains 14 to 18 parts of pure gold. The British gold coinage contains i of copper to 1 1 of gold (22 carats), that of the United States I of copper to 9 of gold. Aurous Compounds These compounds, which are of the same type as the silver salts, are characterized by the readiness with which, in presence of water, they yield a me compounds and the metal : 3 AuX->AuX 3 + 2Au. Aurous Oxide, Au 2 O, is obtained by adding an alkali hydroxide to the solution of an aurous salt. It is a dark violet powder, which decomposes rapidly into gold and oxygen when heated to 250. Aurous Chloride, AuCl, is obtained as a yellowish-white powder by heating auric chloride to c8o c : AuCl 3 _>AuCl + Cl 2 . It is decom :osed by water, with formation of auric chloride and metallic gold. With hydrochloric acid and the alkali chlorides it forms complex compounds of the type H AuCl 2 and KAuCl 2 . When heated to 230, it is completely decomposed into gold and chlorine. Aurous Oyanide, AuCN, like cuprous cyanide (cf. p. 412), is obtained under conditions ^uch that the formation of auric cyanide might be expected. It forms yellow microscopic plates, and dissolves in excess of potassium cyanide to form a complex salt, KAu(CN) 2 , potassium aurocyanide, which is used in gold-plating (cf. silver-p ating, p. 425). Auric salts dissolve in potassium cyanide to form potassium auricyanide, KAu(CN) 4 (= KCN,Au(CN) 3 ), also used in gold-plating. Auric cyanide, Au(CN) 3 , is known ; it occurs in large colourless plates with 3H 2 O. Auric Oxide, Au 2 O 3 , is obtained by heating auric hydroxide, Au(OH) 3 , to 140-150. At 155-165 it loses oxygen, and finally the lower oxide remains. The hydroxide, Au(OH) 3 (or perhaps AuO(OH)), is obtained in an impure state by boiling auric chloride with potassium hydroxide ; in pure condition by the action of magnesium carbonate on auric chloride, the magnesia being then removed by dilute nitric acid. It is a brown powder, soluble in excess of potassium hydroxide, and on evaporating the solution the compound KAuO 2 ,3H 2 O, potassium curate, separates in yellow needles: Au(OH) 3 + KOH-KAuO 2 + aH 2 O. This beha^ iour shows that the hydroxide has acidic as well as basic properties. 43 o A TEXT-BOOK OF INORGANIC CHEMISTRY Auric Chloride, AuCl 3 , is formed when gold is dissolved in aqua regia or in chlorine water, and is obtained as a dark red crystal- line mass on evaporating the solution and drying carefully at 1 50. When heated to 180 it decomposes into aurous chloride and chlorine ; in a closed space an equilibrium between the three substances is established : When hydrochloric acid is added to a neutral solution of gold chloride the colour changes to yellow, and on evaporating the com- pound HAuCl 4 ( = AuCl 3 ,HCl), chloranric arid, separates in long, light-yellow hygroscopic needles. Many salts are Derived from chlorauric acid, such as KAuCl 4 ,2H 2 O and NaAuCl 4 ,2H 2 O. In solu- tions of these salts, and in that of the acid itself, the gold forms part of a complex anion, AuCl/. In aqueous solution auric chloride is partially combined to form the compound HAuCl 3 (OH), which ionises as follows : and is therefore a weak acid. It is clear that there is considerable analogy between this compound and chlorauric acid ; the former is derived from the latter by the substitution of OH for Cl. Gold Sulphides Auric sulphide, Au 2 S 3 , cannot be obtained by interaction in aqueous solution, as it is immediately decomposed by water. It can, however, be obtained as a brown powder by the action of hydrogen sulphide on dry lithium chloraurate at -- 10, the lithium chloride being removed from the product by treatment with alcohol. \Vhen hydrogen sulphide is passed into a boiling solution of auric chloride, aitrotts sulphide, Au 2 S, mixed with sulphur, is obtained as a black precipitate. When the same reaction takes place at room temperature a higher sulphide, probably Au 2 S 2 , is obtained. The sulphides are soluble in solutions of alkali sulphides, forming thio- aurites, e.g. K 3 AuS 2 or 3K 2 S,Au 2 S, and thioaurales, KAuS 2 or K 2 S,Au 2 S 3 . (Cf. thioarsenites and thioarsenates, p. 503.) Purple of Cassius is a substance obtained under certain conditions by the action of stannous chloride on a solution of auric chloride. It forms a brownish-purple powder, and appears to be a mixture of finely-divided gold with stannic acid hydrogel (p. 355). It is used for colouring glass, enamels, etc. Tests for Gold All gold compounds finally yield the metal on heating. Gold is also obtained as a brownish precipitate by the GOLD 431 action of reducing agents on its salts. The formation of purple of Cassius is also characteristic. General Characters of the Sub-Group and Summary The resemblance between the members of this group among themselves is considerably less than in other groups so far considered. The metals melt at high temperatures, they are not affected by water, they have very little affinity for oxygen, and are among the least electro-positive of the medals. The small affinity is illustrated by the fact that the salts are readily reduced to the metal by heatini or by the action of reducing agents. They resemble each other fairly closely in their univalent - ous compounds ; the univalent halides are white, very slightly soluble in water, form complex com- pounds with hydrochloric acid of the type HMC1 2 , and the cuprous and silver halides are soluble in ammonia with formation of com- pounds of the type M(NH 3 ) 2 C1. A striking difference in behaviour is that silver forms only univalent compounds, copper forms stable divalent compounds, and gold trivalent compounds. The contrast in behaviour between the members of the copper and of the alkali sub-group is very marked. Almost the only resemblance is that all function as univalent elements, the crystals of cuprous, silver and sodium chloride belong to the regular system, and some corresponding salts of silver and sodium are isomorphous. The differences can readily be summarized from the above, and the summary of the behaviour of the alkalis (p. 406) already given. The difference in electro-positive character, the alkalis being very strong and the members of the copper group (except silver) very weak bases, is important. Further, the alkalis have no tendency, like the members of the copper group, to form complex ions. We have seen that the latter metals may enter both into complex cations, e.g. Cu(NH 3 ) 2 ', Cu(NH 3 ) t ", Ag(NH 3 ) 2 -, as well as into complex anions, e.g. Ag(CN)/, Cu(CN),", AuCl/, Au(CNY- In the readiness with which they are reduced to the metallic condition and their small affinity for oxygen, the members of this group resemble the platinum group. Copper in the divalent con- dition shows considerable resemblance to zinc, divalent nickel and divalent iron. T CHAPTER XXVIII ELEMENTS OF GROUP IL, SUB-GROUP A HIS sub-group comprises the following three metals, which are known as the metals of the alkaline earths : Atomic Weight. Calcium (Ca) 40-09 Strontium (Sr) . . . . . - 87.63 Barium (Ba) 137-37 The members of this family are invariably divalent in their compounds. The hydroxides are strong bases, almost as strong as the alkali hydroxides, so that the halides are stable towards water, and there is little or no tendency to the formation of basic salts. The metals themselves are soft, readily oxidize in the air, and are acted on by water at room temperature. CALCIUM Symbol, Ca. Atomic weight =40.09. Occurrence In the combined state calcium is very widely distri- buted in nature. As the carbonate it occurs in enormous amount in the form of chalk, marble, limestone, and coral. As sulphate it occurs in gypsum and selenite, CaSO 4 ,2H 2 O, and as anhydrite, CaSO 4 . As fluo- ride, CaF 2 , it constitutes^/^ rspar; and occurs as phosphate, Ca 3 (PO 4 ) 2 , in phosphorite. Calcium silicate, CaSiO 3 , is a constituent of many rocks. Calcium salts are found in all soils, from which they are taken up by plants; also in most natural waters. Bones are mainly com- posed of calcium phosphate. Preparation of Metal Calcium is obtained in a nearly pure state by heating calcium iodide with metallic sodium, the excess of sodium being finally removed by treatment with anhydrous alcohol (Moissan). A more convenient method, now used exclusively for commercial purposes, depends upon the electrolysis of the fused chloride. The salt is fused in a vessel of graphite, the walls of which form the anode ; 432 CALCIUM 433 and an iron rod, just touching the surface of the electrolyte, is used as cathode. As electrolysis proceeds the light metal collects at the end of the cathode, and the latter is slowly raised in order that the metal may solidify out of contact with the fused electro- lyte. In this way, by progressively raising the cathode, a long irregular rod of metal may be built up without interruption of the electrolysis. Properties Calcium is a silvery white metal somewhat harder than lead, its density is 1.52, and it melts at 800. It turns yellow on the surface on exposure to air, probably owing to the formation of nitride, Cu 3 N 2 . It burns vigorously when heated in air or oxygen, oxide and nitride being formed. It decomposes water slowly at room temperature, hydrogen being liberated. Calcium Hydride, CaH 2 , is obtained by passing hydrogen over calcium at a red heat. It is a white powder, which is vigorously acted on by water with liberation of hydrogen. Calcium Oxides Two oxides of calcium are known : the normal oxide, CaO, and the dioxide, CaO 2 . Calcium Oxide (lime, quicklime), CaO, is prepared by heat- ing calcium carbonate to a temperature exceeding 800 under such condition s that the carbon dioxide set free is continually removed : CaCO 3 ->CaO + CO 2 . On the commercial scale this process is carried out by heating lime- stone in brick kilns. The calcium oxide obtained by burning limestone is a white amor- phous powder which is infusible in the oxyhydrogen flame (lime- light), but readily fuses in the electiic furnace. It unites with water with the ('volution of much heat, calcium hydroxide, Ca(OH) 9 , being formed ; in this process li mps of quicklime crumble to powder. This is known as the slaking of lime, and the product is teimed slaked liriie. On account of its great affinity fcr water, lime is often used for drying gases (p. 215) and for removing the last traces of water frcm liquids. Calcium Hydroxide, Ca(OH) 2 , is a white, amorphous, hygro- scopic powder. When it dissolves in water he^.t is given out, and therefore the solubility diminishes with rise of temperature. At 20 100 grams of water dissolve 0.126 grams, at 50 0.098 grams, at ico 0.060 grams of the base. The solution is termed lime-water. Water containing a large excess of calcium hydroxide in suspension is teimed milk of lime. Milk of lime, lime-water, and calcium oxide itself readily 28 -434 A TEXT-BOOK OF INORGANIC CHEMISTRY absorb carbon dioxide from the air, with formation of calcium carbonate. Calcium Peroxide, CaO 2 , separates as the octahydrate, CaO 2 ,8H 2 O, when hydrogen or sodium peroxide is added to lime-watei. On heating to 130 it becomes anhydrous, and at a red heat decoir- poses into calcium oxide and oxygen. Mortar and Cement Mortar consists of lime and sand made into a paste with water. The setting is due to the escape ofwaten and the absorption of carbon dioxide to form calcium carbonate, which sets into a solid mass with the sand. The latter renders th mass porous, thus facilitating the entry of the carbon dioxide. N calcium silicate is formed till after the lapse of many years, 'so tha this substance plays no part in the process of hardening. Ordinary Cement is a mixture of lime (50 to 60 per cent.), silic (25 per cent.), and aluminium oxide (8 to 10 per cent.). It is mad by burning a mixture of limestone and clay. In some localities mb tures of limestone and aluminium silicates occur which yield cemer directly on burning. Cement has the advantage over mortar of se; ting to a hard mass even under water (hydraulic cement). Portla?~t, cement is of similar composition to ordinary cement. It is prepare; by burning at a high temperature a mixture in definite proportions limestone and clay rich in silica ; the hard mass or clinker is thei reduced to a fine powder. The hardening of cement is by no mean understood. According to Le Chatelier it depends mainly on tl change of a basic calcium silicate by absorption of water to the norm a hydrated silicate and calcium hydroxide : Calcium Carbonate, CaCO 3 , is perhaps the most familiar ex am pie of a dimorphous substance, being met with in the two cryste. 1 line forms : calcite, which generally occurs in rhombohedral crystal belonging to the hexagonal system ; and aragonite, in crystals belong ing to the orthorhombic system. Chalk, marble, limestone, Icelanc spar, etc., belong to the calcite modification ; aragonite is of somewh i rare occurrence in nature. When precipitated from solution by inter- action of a soluble calcium salt and sodium carbonate at roo;r temperature, calcium carbonate is amorphous ; but on standing ir contact with the mother liquor changes to minute crystals of calcite When, on the other hand, precipitation occurs from hot solution aragonite is obtained. Calcite appears to be the stable form abo/< CALCIUM 435 o under all conditions ; the apparent stability of aragonite is due to its slow rate of change (p. 174). Calcium carbonate is practically insoluble in water ; but is soluble in water containing carbon dioxide. In this case soluble calcium bicarbonate. CaH 2 (CO 3 ) 2 , is doubtless present in the solution : CaCO 3 + H 2 O + CO 2 ^CaH 2 (CO 3 ) 2 . The reaction is reversible ; on boiling carbon dioxide escapes and calcium carbonate is reprecipitated. The depositor "fur" in kettles and steam boilers is calcium carbonate which has fallen out of solution owing to the removal of the carbon dioxide on boiling. Hardness of Water As already indicated (p. 325), the hard- ness of water is due almost entirely to the presence of calcium and magnesium salts in solution, chiefly as carbonates and sulphates. A distinction is drawn between temporary hardness, due to the carbon- ates held in solution by carbon dioxide and which is removed by boiling, and permanent hardness, due to sulphates, which cannot be removed by boiling. The hardness of water is chiefly noticeable in its action on soap. A soft water, containing little dissolved salts, readily forms a lather with soap ; but hard waters use up a large quantity of soap before a lather is obtained. The hardness of water (including both temporary and permanent hardness) can, in fact, be estimated by adding to water a standard soap solution from a burette till the point is reached at which a lather is formed on shaking. Temporary hardness can be removed by boiling, as already indicated, and also by adding to the solution sufficient calcium oxide to combine with the carbon dioxide oresent, when all the carbonate is precipitated : + CaO->2CaCO 3 +H 2 O. Calcium Chloride, CaCl 2 , occurs *.$ tacliydrite (2MgCl 2 'CaCl 2 . 2H.,O) in the Stassfurt deposits. It is a bye-product in certain echnical operations, such as the preparation of ammonia from its alts (p. 214). It separates from solution at room temperature as he hexahydrate, CaCl 2 ,6H 2 O, in hexagonal crystals ; other hydrates re also known. The hexahydrate melts at 30.2. Above 260 all :ie hydrates change to the anhydrous salt ; the latter melts at 802. mhydrous calcium chloride has a great affinity for water, and is used )r drying gases and organic liquids. It forms a compound with mmonia, CaCl 2 ,8NH 3 , and cannot therefore be used for drying lis gas. 43 6 A TEXT-BOOK OF INORGANIC CHEMISTRY The hexahydrate and ice form a very useful freezing-mixture (p. 199), the temperature (cryohydric temperature) falling to -55'. Chlorinated Lime (bleaching powder) This substance is prepared by the prolonged action of chlorine on damp slaked lime at room temperature. The reaction may be represented by the follow- ing equation : /OC1 Ca(OH) 2 + Cl^Ca + H 2 O, but a certain amount of calcium hydroxide is always present. It was formerly supposed that bleaching powder was a mixture of equivalent amounts of calcium chloride and calcium hypochlorite, CaCl 2 ,Ca(OCl) 2 , "but conclusive evidence has been obtained that the solid compound contains no free calcium chloride. 1 The aqueous solution, however, appears to contain a mixture of chloride and hypochlorite : 2Ca(OCl)Cl->CaCl 2 + Ca(OCl) 2 . Dilute hydrochloric acid, as well as moist carbon dioxide, liberate practically all the chlorine from bleaching powder even in the cold. In the case of hydrochloric acid, we may assume that hypochlorous acid is first formed, and that it immediately reacts with hydrochloric acid to form water and chlorine : + 2HCl->CaCl 2 HC1O + HC1->H 2 O + C1 2 . Other acids presumably liberate both hydrochloric and hypochlorous acids, which then react as above. The use of this substance in bleaching depends upon the reactions just considered. The material to be bleached is dipped first in a solution of bleaching powder and then into a dilute acid ; the effect is due to the liberated chlorine. The "available" chlorine is the amount set free from bleaching powder by the action of excess of hydrochloric acid. It usually amounts to 30 to 36 per cent, but a product containing over 39 pe cent, can be obtained. Calcium Sulphate, CaSO 4 , occurs naturally in the anhydrous form as anhydrite (rhombic crystals), and as hydrate in gypsuiL, CaSO 4 ,2H 2 O (monoclinic crystals). Alabaster and selcnite are forms 1 The main evidence in favour of this view is that bleaching powder, unlil- o calcium chloride, is not deliquescent ; and further, alcohol, though agood;solve):t for calcium chloride, does not extract the latter from bleaching powder. CALCIUM 437 of gypsum. When calcium sulphate is prepared by double decom- position, the dihydrate is obtained. When gypsum is heated in an open vessel at 98, it loses water and the so-called half-hydrate, (CaSO 4 ) 2 ,H 2 O, is obtained ; on prolonged heating at 107 to 108, more readily at higher temperatures, all the water is driven off, and one form of the anhydrous salt, the so-called "soluble anhydrite" is obtained. When either the half-hydrate or soluble anhydrite is treated with water, rehydration takes place, and the mixture sets to a hard mass. Plaster of Paris consists mainly of the half-hydrate, and is made by heating gypsum in kilns at a temperature very little over 130. The setting is represented as follows : (CaSO 4 ),,H,O + 3H 2 O->2(CaSO 4 ,2H 2 O). When gypsum is heated for some time above 200, a second modifi- cation of anhydrite is formed, which sets very slowly with water. Gypsum thus treated is said to be dead burnt, and the product is use- less for making casts. Natural anhydrite does not set with water. Calcium sulphate is only slightly soluble in water. The solubility increases slowly with rise of temperature up to 35, and then diminishes (cf. p. 84). The amount dissolved depends somewhat on the size of the partides ; the finer the state of division the greater is the solu- bility. This rule is a quite general one, but the difference is only appreciable for substances of slight solubility. At o 100 grams of water take up about 0.195 grams, at 38 0.221 grams, and at 99 o.i 80 grains of the salt. Calcium Phosphates Tricalcium orthophosphate, Ca 3 (PO 4 ) 2 , Is the chief inorganic constituent of bones. It also occurs naturally in immense quantities as phosphorite and osteolite, and along with calcium fluoride in apatite ', 3Ca 3 (PO 4 ) 2 ,CaF 2 . As it is practically insoluble in water, it is at once precipitated when solutions contain- ing Ca" and PO/" ions respectively are mixed, e.g. by adding am- monia and sodium phosphate to a solution of calcium chloride. Although very slightly soluble in water, it is soluble in acids, even carbonic acid, and in solutions of salts, more particularly ammonium salts and :he chlorides and nitrates of the alkalis. These facts are of great importance in facilitating the taking up of calcium phosphate by plants fro n the soil, as plants can only take up soluble substances. When acted on by the requisite quantity of sulphuric acid, mono- calcium phosphate, CaH 4 (PO 4 ) 2 , and calcium sulphate are formed : 438 A TEXT-BOOK OF INORGANIC CHEMISTRY This mixture is known as superphosphate of lime, and is largely used as a manure. Monocalcium phosphate has the advantage of being readily soluble in water. The intermediate phosphate, dicalcium phosphate, Ca 2 H 2 (PO 4 ) 2 or CaHPO 4 , is obtained by decomposing the monocalcium salt with water : CaH 4 (PO 4 ) 2 ->CaHPO 4 j + H 3 PO 4 . It is more soluble in water than the triphosphate, but much less soluble than the monocalcium phosphate. Calcium Carbide, CaC 2 , is prepared by heating lime with carbon in the electric furnace (p. 315) : Pure calcium carbide is colourless ; the dark colour of the com- mercial article is due to impurities. It is used commercially in the preparation of acetylene. Calcium Cyanamide, CaNCN, has already been referred to in connexion with the utilization of atmospheric nitrogen (p. 236). Calcium Sulphide, CaS, is obtained by reducing calcium sulphate with carbon. When pure, it is a white, amorphous powder. It is very slightly soluble as such in water, but undergoes hydrolytic decomposition with formation of calcium hydrosulphide, Ca(SH).,, and calcium hydroxide, the former of which is readily soluble in water : 2H 2 O->Ca(SH) 2 + C For this reason no precipitate is obtained when ammonium sulphide is added to a calcium salt in solution ; owing to hydrolysis only the soluble hydrosulphide is formed. Ordinary calcium sulphide, as well as the sulphides of strontium and barium, after exposure to light, remain luminous in the dark. For this reason they are used in making luminous paint. The pure substances do not show this property, which is therefore connected with the presence of small amounts of impurities. Calcium Silicate. Glass The metasilicate, CaSiO 3 , occurs naturally as wollastonite, and can be prepared by fusing together calcium oxide and silica in the electric furnace. It occurs in crystals insoluble in water, and the temperature has to be raised to 1400 in order to fuse it. The different kinds of glass are mixtures of silicates of calcium or lead with alkali silicates, and contain excess of silica. Ordinary soft STRONTIUM 439 glass (window glass) is made by fusing together sodium carbonate, limestone, and sand. It is readily fusible. Crown or Bohemian glass contains potassium instead of sodium silicate, and is less readily ftsible than ordinary glass. Flint glass is a mixture of lead and potassium silicates. It is readily fusible, and has a high refrac- tive index. The colours of certain kinds of glass are due to traces of impurities. The green colour of the common glass used for making bottles and window-panes is due to silicate of iron. The appearance of milk-glass is due to the presence of finely-divided calcium phosphate. Glass has no definite crystalline properties, and when heated gradually softens without showing a definite melting-point ; it must therefore be regarded as an amorphous substance, and is in reality a highly supercooled liquid (p. 69). When kept at a high tempera- ture for a long time, some of the silicates may separate in crystalline form ; the glass is then said to be " devitrified," and is very brittle. As the alkali silicates are soluble in water (p. 354), it is not sur- prising that water (especially hot water) dissolves out small amounts of alkali from glass. This may be shown by grinding up some soft glass with water in a mortar and adding a few drops of phenol- phthaleir, when the solution turns pink. Tests for Calcium -Calcium compounds give a brick-red colour to the Bunsen flame. The most characteristic wet test is the precipita 1 ion of calcium oxalate by double decomposition ; this salt is insoluble in acetic, but readily soluble in hydrochloric acid. STRONTIUM Symbol, Sr. Atomic Weight=87.63. The compounds of strontium are very similar to those of calcium, and may therefore be discussed very briefly. Occurrence Strontium is much less widely distributed than the other members of the group. The chief naturally occurring com- pounds are celestine, SrSO 4 , and strontianite, SrCO 3 . Preparation of Metal Strontium, like calcium, is usually prepared by electrolysis of the fused chloride. Properties Strontium is a soft, silvery-white metal of density 2.55 ; it melts about 800. It is acted on by water and by dilute acids at room temperature more vigorously than calcium, and when heated in oxygen it burns vigorously to the oxide. When heated in hydrogen, strontium hydride, SrH 2 , is formed. 440 A TEXT-BOOK OF INORGANIC CHEMISTRY Preparation of Strontium Compounds Strontium com- pounds are prepared by dissolving the carbonate, strontianite, in the appropriate acid, or by reducing the sulphate to sulphide by means of carbon, and treating the latter compound with acids. Strontium Oxide, like the other oxides of this sub-group, is obtained by strongly heating the hydroxide, carbonate, or nitrate. It combines readily with water to form the hydroxide, Sr(OH) 2 ; the latter separates from aqueous solution in the hydrated form as Sr(OH) 2 ,8H 2 O in tetragonal crystals. Strontium hydroxide is more soluble in water than calcium hydroxide (see table, p. 376), and the solubility increases regularly with the temperature. It forms a com- pound with sugar which is insoluble in water, but is readily decom- posed by carbon dioxide, and for this reason is used in the sugar industry. Strontium Chloride, SrCl 2 , separates from solution as SrCl 2 ,6H 2 O in hexagonal crystals, isomorphous with calcium chloride hexahydrate. Unlike barium chloride, it is somewhat soluble in alcohol. Strontium Sulphate, SrSO 4 , obtained by double decomposi- tion, is less soluble in water than calcium sulphate, but more soluble than barium sulphate. At 20 100 grams of water take up 0.1479 grams, at 80 0.1688 grams of the salt. Strontium Nitrate, Sr(NO 3 ) 2 , separates from warm aqueous solutions as the anhydrous salt in octahedral crystals ; from cold solutions as Sr(NO 3 ) 2 ,4H 2 O in monoclinic crystals, It is practically insoluble in alcohol, and this property is taken advantage of in separating strontium from calcium, as calcium nitrate is soluble in alcohol. This and other strontium salts are used in pyrotechny, owing to the deep crimson colour they impart to flame. Tests- Strontium compounds are characterized by the crimson flame, very similar to that due to lithium. The spectra of the two metals are, however, quite distinct ; that of strontium has a number of bands in the red and a characteristic blue line. BARIUM Symbol, Ba. Atomic Weight = 137. 37. Occurrence The principal natural compounds of barium are wither ite } BaCO 3 , and heavy spar, BaSO 4 . Preparation of Metal The preparation of pure barium has proved a matter of considerable difficulty. The best results appear to have been obtained by warming a concentrated solution of barium BARIUM 44i chloride with sodium amalgam. From the resulting barium amalgam the mercury is removed by placing it in a tube and gradually increasing the temperature. Finally, at 1150, pure barium distils over. Properties Barium is a soft and, when pure, presumably a silvery-white metal, which melts below 1000 ; its density is 3.75. It becomes oxidized very readily in the air, and decomposes water vigorously at room temperature. Preparation of Compounds Barium compounds, like those of strontium, can be obtained by dissolving the carbonate in the appropriate acid, or by reducing the sulphate to the sulphide by heating with charcoal and then dissolving the sulphide in acids. A further method, starting with the very insoluble sulphate, is to fuse it with sodium carbonate : If a considerable excess of the carbonate is used, the action is almost complete in the direction of the upper arrow. The sodium sulphate and exce.ss of carbonate are removed by boiling the product with water. Barium Oxides Two oxides of barium are known, the monoxide, BaO, and the peroxide, BaO 2 . Barium Oxide, BaO, is usually prepared by heating the nitrate. It cannot conveniently be obtained by heating the carbonate alone, owing to the very high temperature at which the latter decomposes ; but when the carbonate is mixed with carbon, the reaction takes place at a much lower temperature : Barium oxide forms a white amorphous powder, which fuses in the electric furnace at a lower temperature than calcium oxide. When heated in air it combines with oxygen to form the peroxide. It com- bines very vigorously with water to form the hydroxide, Ba(OH) 2 , a large amount of heat being given out. Barium Hydroxide, Ba(OH) 2 , obtained as described above, separates from aqueous solution, at all temperatures between 20 and 109, in tetragonal crystals as the octahydrate, Ba(OH) 2 ,8H 2 O. The octahydrate effloresces in the air, with formation of the monohydrate. The anhydrous compound melts at a red heat without giving up water. Barium hydroxide is much more soluble in water than the other 442 A TEXT-BOOK OF INORGANIC CHEMISTRY bases of this group the solution is called baryta-water. At 10 100 grams of water take up 2.5 grams, at 20 4.3 grams, and at 40 8.2 grams of Ba(OH) 2 . Barium Peroxide, BaO 2 , is obtained by heating barium oxide in oxygen, or air free from carbon dioxide, at temperatures in the neighbourhood of 500. Raising the temperature favours the reverse decomposition. It is also precipitated in crystalline form as BaO 2 ,8H 2 O, when hydrogen peroxide is added to baryta water. Barium peroxide is a greyish powder, only very slightly soluble in water. By the action of dilute acids hydrogen peroxide is formed (p. 138). BaO 2 + H 2 SO 4 ->BaSO 4 + H 2 O 2 . Barium Chloride, BaCl 2 , prepared by the general methods (P- 374)i separates from solution in colourless, monoclinic crystals as BaCl 2 ,2H 2 O. Unlike calcium chloride hexahydrate, it is not deliquescent ; it is readily dehydrated on heating to 100. It is very soluble in water, and is precipitated from a concentrated solution by adding hydrochloric acid (cf. p. 423). Barium Sulphate, BaSO 4 , prepared by double decomposition, is a white powder, which can be obtained in the amorphous form and also as minute crystals. It is practically insoluble in water (0.235 milligrams in 100 grams of water at 18) and in dilute acids. It is readily soluble (to the extent of 10 to 12 per cent.) in concentrated sulphuric acid, doubtless owing to the formation of the acid sulphate, Ba(HSO 4 ) 2 , but is reprecipitated on the addition of water. Barium Nitrate, Ba(NO 3 ) 2 , obtained by the general methods, separates from aqueous solution at room temperature in the anhydrous form (octahedral crystals). It is readily soluble in water. It is used in pyrotechny for the production of green flame. Tests for Barium Barium salts impart a green colour to the Bunsen flame, and the spectrum is characterized by the presence of certain lines in the green and orange. In the wet way, the formation of the sulphate, which is always obtained, on account of its very small solubility product, when Ba" and SO 4 " ions are brought together, even in minute amount, is the most important test. It differs from the other two metals of the sub-group in giving a yellow precipitate of barium chromate with potassium chromate, insoluble in acetic acid. METALS OF THE ALKALINE EARTHS 443 COMPARISON OF THE ALKALINE EARTH METALS AND SUMMARY As is evident from the foregoing, there is a very close resemblance in the behaviour of the corresponding compounds of the members of this group, and, further, the properties vary regularly with increase of atomic weight. The metals themselves are almost as electro-positive as the alkali metals ; they quickly oxidize in the air and decompose water at room temperature. They impart characteristic colours to the Bunsen flame. They are all divalent elements. Their more important physical pro- perties are shown in the table Ca. Sr. Ba. Atomic weight 40.09 87.63 137-37 Density 1.52 2.55 3-75 Melting-point 800 about 800 about 8oo u Atomic volume 26.4 337 36.6 The chemical properties are, of course, determined by their strong electro-positive character, as illustrated by the fact that the chlorides are not hydrolyzed, and that there is very little tendency to form basic salts. Barium appears to be the most strongly basic, as is shown, for example, by the fact that the hydroxide can be fused without decom- position, whilst calcium hydroxide readily splits up into water and the oxide on heating ; the relatively great stability of barium carbonate also supports this conclusion. The alkaline earths differ markedly from the alkalis with regard to the small solubility of the hydroxides, carbonates, phosphates and sulphates, but in this respect, as we have seen, lithium approximates somewhat to the alkaline earth metals (p. 406). The solubility of the hydroxides and carbonates increases, that of the sulphates decreases with increasing atomic weight. CHAPTER XXIX ELEMENTS OF GROUP II SUB-GROUP B / ~PHIS sub-group comprises the following five metals : Beryllium (glucinum) (Be) . . . . 9- 1 Magnesium (Mg) 24.3 Zinc (Zn) 65.4 Cadmium (Cd) 112.4 Mercury (Hg) . . . . . 200.0 Corresponding with their position in the periodic table, they function as divalent elements only (regarding mercurous compounds cf. p. 456). They are considerably less electro-positive than the metals so far considered, and the electro-positive character diminishes from mag- nesium to mercury. From this it follows at once, since the hydroxides are relatively weak bases, that the halogen compounds undergo hydrolysis, and that there is considerable tendency to the forma- tion of basic salts ; further, that the hydroxides and carbonates are less stable than those of the alkalis and alkaline earth metals. BERYLLIUM (GLUCINUM) Symbol, Be. Atomic Weight = 9.1. Occurrence Beryllium is a comparatively rare element. The chief source is beryl, a silicate of aluminium and beryllium, 3BeO,Al 2 O3,6SiO 2 . Beryl coloured green by traces of impurities is called emerald ; when coloured bluish-green, aquamarine. Chrysoberyl, another mineral containing beryllium, has the formula BeO,Al 2 O 3 . Preparation of Metal Beryllium is most readily obtained by strongly heating beryllium potassium fluoride, BeF 2 'KF, with metallic sodium and excess of sodium chloride. After treatment with water, crystals of practically pure beryllium are found in the residue. Properties Beryllium is a silver- white metal of density 1.85; it melts below 1000. It is stable in the air at ordinary temperature, but on heating becomes coated with a layer of oxide, which retards further oxidation. It is not affected by water even at 100. It is readily dissolved by dilute hydrochloric and sulphuric acids, but nitric acid, whether dilute or concentrated, has very little 444 MAGNESIUM 445 action on it. It dissolves readily in potassium or sodium hydroxide, with evolution of hydrogen and formation of the compound Be(ONa) 9 . In this and other respects it shows a remarkable resemblance to aluminium. Compounds of Beryllium Beryllium oxide, BeO, is obtained by heating the hydroxide to 440; also by heating the carbonate or sulphate. It is a white powder ; which under ordinary circumstances is soluble in dilute acids, but after being strongly ignited is much less soluble. Beryllium Hydroxide, Be(OH) 2 , is obtained as a gelatinous precipitate when an alkali hydroxide is added to the solution of a beryllium salt. When heated at 440, it 'oses water and forms beryllium oxide." It dissolves in excess of alkali hydroxide, forming alkali beryllate, e.g. Be(ONa) 2 , but is reprecipitated on prolonged boiling. Beryllium Chloride, BeCl 2 , is obtained in the anhydrous condition by heating the oxide, mixed with charcoal, in a current of chlorine. It forms a colourless, minutely crystalline powder, which fumes in the air like phosphorus penta- chloride. It is readily soluble in water, and on evaporating the solution the tetrahydrate, BeCl2,4H 2 O, separates in monoclinic, deliquescent crystals. This compound cannot be completely dehydrated on heating, as under these circumstances the water reacts with the salt, hydrogen chloride escapes, and the oxide remains behind : BeCl 2 + H 2 O_>BeO + 2HC1. Beryllium Carbonate As Beryllium is a weak base and carbonic acid is a weak acid, it is not surprising that owing to hydrolysis basic carbonates are obtair.ed by double decomposition between a soluble beryllium salt and a soluble carbonate. Of these, the compound BeCO 3 ,2Be(OH) 2 is best known. The norn:al carbonate, BeCO 3 ,4H 2 O, is obtained by evaporating the aqueous solution n an atmosphere of carbon dioxide (cf. magnesium carbonates). The salt is dehydrated on heating to 100, and at a slightly higher temperature begins to lose carbon dioxide. Beryllium Sulphate, BeSO 4 This salt is obtained by dissolving beryllium oxide in dilute sulphuric acid ; on concentrating the solution, it separates as BeSO 4> 4U. 2 O in large octahedral crystals. It is very soluble in water. On heating the crystals practically all the water escapes ; at a slightly higher temperature sulphur trioxide is driven off and the oxide remains. MAGNESIUM Symbol, Mg. Atomic Weight = 24-32 Occurrence In the combined state, magnesium is one of the most abundant of the elements. In the Stassfurt deposits, it occurs as chloride in carnallite, MgCl 2 ,KCl,6H 2 O, and tachydrite, CaCl 2 ,2 MgCl 2 ,i2H 2 O, and as sulphate in kieserite, MgSO 4 ,H 2 O kainitc, KCl,MgSO 4 ,3H 2 O, and other compounds. As carbonate, it forms magnestte^ MgCO ;J , and is a constituent of dolomite, MgCC^CaCO.^ which forms whole mountain ranges. As silicate, it occurs in talc, Mg 3 H 2 (SiO 3 )4, serpentine, Mg 3 Si 2 O r ,2H 2 O, meerschaum, 446 A TEXT-BOOK OF INORGANIC CHEMISTRY Mg 2 Si 3 O 8 ,2H 2 O, asbestos and other compounds. As chloride and sulphate, it is found in many mineral springs. Preparation of Metal Magnesium is now usually pre- pared by the electrolysis of dehydrated carnallite, MgCl 2 ,KCl. The salt is kept at a temperature above the melting-point in an iron vessel, the walls of which serve as cathode, and the .carbon anode, surrounded by a porcelain tube to conduct away the chlorine, is introduced through a hole in the cover. Formerly magnesium was prepared by heating the double chloride of magnesium and sodium, MgCl 2 'NaCl(p. 447), with metallic sodium. A convenient modification of this method, still in use, is to treat dehydrated carnallite with metallic sodium. Properties Magnesium is a silvery-white, fairly hard metal of density 1.75 ; it melts about 600 and boils about 1100. It is fairly malleable ; on heating it becomes ductile and can be drawn into wire. It is stable in dry air, but in moist air becomes covered with a coating of oxide. When heated to redness in air, it burns with a brilliant white flame, the high luminosity of which is due to the great amount of heat given out and to the non-volatility of the oxide formed. For this reason it is used in signalling and in pyrotechny. Further, as the flame is very rich in chemically active rays, it is used as a flash-light in photography ; for this purpose magnesium powder is blown into a spirit flame. It acts on water very slowly even at 100, but when steam is passed over heated magnesium, the metal burns vigorously to the oxide and hydrogen is set free. It is readily dissolved by acids with evolu- tion of hydrogen. It is interesting to note that magnesium liberates hydrogen from nitric acid, a property possessed by no other metal. Unlike beryllium and zinc it is not affected by alkali hydroxides, even on boiling. Magnesium Oxide, MgO, is formed by burning the metal in air, or by heating the hydroxide or carbonate. It is a white, very light powder, which is used in medicine under the name of calcined magnesia. It is a very infusible substance (it can, how- ever, be fused in the electric furnace), and is therefore used in making fire-brick, crucibles, etc. It combines readily with water to form magnesium hydroxide, Mg(OH) 2 . Magnesium Hydroxide, Mg(OH) 2 , is obtained by the action of water on the oxide, or by adding sodium or potassium hydroxide to the solution of a magnesium salt. It is a white powder which loses water at a red heat, forming the oxide. It is very slightly MAGNESIUM 447 soluble in water, but sufficient is taken up to make the solution alkaline to litmus. It is fairly soluble in solutions of ammonium salts, e.g. ammonium chloride. The reaction in this case is re- presented by the equation and depends upon the fact that the OH 7 ion concentration in ammonium hydroxide solution is very- small, especially in the presence of ammonium salts. Magnesium hydroxide can only remain undissolved when the solubility product of Mg" and OH' ions is exceeded. In the above instance, OH' ions are taken up to form practically non-ionised ammonium hydroxide, as repre- sented bv the upper arrow, and if sufficient ammonium chloride is added the OH' ion concentration is reduced to such an extent that the solubility product is no longer reached ; in other words, magnesium hydroxide is dissolved. A paste of magnesium hydroxide and water absorbs carbon dioxide from the air, and s.ets into a hard mass of carbonate ; it is therefore a useful cement. Magnesium Chloride, MgCl 2 , is obtained by dissolving magnesium, the oxide or carbonate, in hydrochloric acid; on con- centrating the solution it separates in monoclinic crystals as MgCl 2 ,6H 2 O. Commercially it is obtained at Stassfurt from carnallite, MgCl 2 ,KCl,6H 2 O. As potassium chloride is the less soluble of the two salts, the greater part of it is first removed by crystallization, and from the mother liquor MgCl 2 ,6H 2 O is obtained on further con- centration. The crystals are deliquescent. The hcxahydrate cannot be completely dehydrated by heat, as the chloride undergoes partial hydrolysis, the oxide being formed and hydrogen chloride given off: The anhydrous salt is obtained by heating the metal in a current of chlorine, or from magnesium ammonium chloride, MgCl 2 ,NH 4 Cl,6H 2 O. When tl e latter compound is heated, it first loses water, then at a higher temperature ammonium chloride, and finally the anhydrous salt remains. Anhydrous magnesium chloride occurs in lustrous crystalline leaflets; it melts at 708 and can be distilled unchanged in dry air, but is decomposed on heating in moist air, in accordance with the above equation. 448 A TEXT-BOOK OF INORGANIC CHEMISTRY A number of basic chlorides of magnesium have been described, but their formulae have not been definitely established. Magnesium Sulphate (Epsom salts), MgSO 4 , occurs in Stassfurt deposits as kieserite, MgSO 4 ,H 2 O, and as a constituent of a number of double salts (p. 399). It is also found in many mineral springs, notably in that at Epsom, from which it was first obtained. It is usually obtained commercially by treating kieserite (a com- paratively insoluble salt) with hot water, and setting the solution aside to crystallize. Magnesium sulphate is usually met with as the heptahydrate, MgSO 4 ,7H 2 O, in rhombic prisms. Another unstable form of the heptahydrate (monoclinic prisms) sometimes separates from super- saturated solutions at room temperature. At 60 to 70 the hexa- hydrate separates from solution in monoclinic crystals, and above 68 kieserite, MgSO 4 ,H 2 O, is obtained. Other hydrates,e.g. MgSO 4 ,i2H 2 O, are also known. All the higher hydrates, when heated at 150, lose water and form the monohydrate ; above 200 the anhydrous salt is obtained. At o 100 grams of water dissolve 26.9 grams, at 15 33.8 grams, and at 30 40.9 grams of MgSO 4 . An acid sulphate, Mg(HSO 4 ) 2 , has been obtained in prismatic crystals from a hot solution of magnesium sulphate in sulphuric acid. Double salts of magnesium sulphate with the alkali sulphates, of the type MgSO 4 ,M 2 I SO 4 are known. The potassium and ammonium salts crystallize with 6H 2 O, the sodium salt with 4H 2 O. The potassium salt occurs naturally as schonite. Magnesium Ammonium Phosphate, MgNH 4 PO 4 ,6H 2 O, is always precipitated, on account of its very slight solubility in water, when Mg", NH 4 - and PO 4 "' ions are brought together, e.g. by adding ammonium chloride, ammonia and sodium phosphate to a solution of magnesium sulphate. Although only slightly soluble in water, it is still less soluble in solutions containing NH 4 - ions and free ammonia. It is used in the quantitative estimation of magnesium salts and of phosphates. On strongly heating, magnesium pyrophosphate, Mg 2 P 2 O 7 , is obtained: 2MgNH 4 PO 4 ->Mg 2 P 2 O 7 + 2NH 3 +H 2 O. Magnesium Carbonates The normal carbonate occurs naturally as magnesite, MgCO 3 , in rhombohedral crystals, and a basic carbonate, 3MgCO 3 ,Mg(OH) 2 ,3H 2 O, as hydromagnesile, in monoclinic crystals. As both base and acid are weak, only basic carbonates are obtained by double decomposition between mag- ZINC 449 nesium salts and carbonates in solution. The composition of these basic carbonates depends on the conditions of precipitation. When solutions of magnesium sulphate and sodium carbonate are mixed at room temperature, the precipitate has the approximate com- position 3MgCO 3 ,Mg(OH) 2 ,3H 2 O, the same as hydromagnesite. The carbonate thus obtained is known as light magnesium carbonate. A much denser precipitate, which may be rather more basic, is obtained by carrying out the precipitation in concentrated solution at the boiling-point ; it is known as heavy magnesium carbonate. There is no evidence that these basic carbonates are other than mixtures of carbonate and hydroxide, and in fact the composition of the com- mercial article varies within wide limits. When the basic carbonate is suspended in water and carbon dioxide passed through it, the carbonate dissolves, and the solution doubtless contains the bicarbonate, Mg(HCO 3 ) 2 . From this solution below 1 6 the normal carbonate separates as pentahydrate, and above 16 as trihydrate, MgCO 3 ,3H 2 O, in colourless crystals. The an- hydrous carbonate can be obtained by heating a solution of the carbonate in excess of carbon dioxide to 150 in a vessel provided with a porous stopper, so that the carbon dioxide can escape slowly. On heating at 200 both normal and basic carbonates are com- pletely decomposed, with formation of the oxide. Tests for Magnesium From solutions of magnesium salts, the alkali hydroxides precipitate magnesium hydroxide, insoluble in excess of alkali. For the reasons already stated, ammonia does not precipitate magnesium hydroxide from solutions containing ammonium chloride. From solutions of magnesium salts containing ammonium chloride and excess of ammonia, sodium phosphate throws down a white crystalline precipitate of magnesium ammonium phosphate, MgNH 4 PO 4 . ZINC Symbol, Zn. Atomic weight = 65. 37. Molecular weight =65.37. Occurrence Zinc is fairly abundant in nature as calamine or smithsonitc, ZnCO 3 , zinc blende, 2x&,franklinite t Zn(FeO 2 ) 9 , and red zinc ore, ZnO. The majority of the zinc ores contain a little cadmium. Preparation of Metal The first stage in obtaining the metal is to convert tljjpres to oxide. For this purpose the sulphide is roasted in air ; the carbonate only requires to be heated. The oxide is then mixed with coal and heated strongly in earthenware retorts, 29 450 A TEXT-BOOK OF INORGANIC CHEMISTRY when zinc distils over and is collected in iron receivers. The portion which first passes over is condensed in the cold receiver as a fine powder termed zinc dust. Besides the metal, zinc dust contains a few per cent, of oxide, and also some cadmium which, being more volatile than zinc, passes over first. As the retorts heat up, the zinc melts, and can be drawn off and cast into sticks. Zinc thus prepared is very impure, containing lead, iron, cadmium, arsenic and other elements. It can be partially purified by redis- tillation. Perfectly pure zinc is obtained by dissolving the metal in acid, and precipitating the carbonate, which is then heated and the oxide reduced with pure carbon. Properties Zinc is a bluish-white metal of density 6.9 to 7.2 ; it melts at 419 and boils at 920. It is usually obtained in hexagonal crystals. Ordinary zinc is moderately brittle at room temperature, at 150 it is malleable and ductile, at 200 it has again become so brittle that it can be powdered in a mortar. This behaviour is doubtless connected with changes in the nature of the crystals. The vapour density of zinc is about 32.7, from which it follows that it is monatomic in the state of vapour. In the air, zinc becomes coated with an extremely thin film of oxide (or perhaps basic carbonate) which protects it against further oxida- tion. On heating in air, it burns with a bluish flame, forming the oxide. It does not decompose water at 100. The commercial metal dissolves readily in dilute acids, hydrogen being given off, but very pure zinc is scarcely affected by acids. This difference in behaviour of pure and impure zinc is not entirely understood. In the latter case, differences of potential are probably set up between zinc and th< impurities (iron, lead, etc.) and the hydrogen is liberated at thes< metals as in an ordinary battery (cf. p. 416). With pure zinc, on th< other hand, the hydrogen probably accumulates on the surface and sets up a contrary E.M.F. which brings the reaction to a standstill. When boiled with solutions of alkali hydroxides, zinc dissolves wit] liberation of hydrogen and formation of alkali zincate : Zn + 2KOH->Zn(OK) 2 + H 2 f . Uses of Zinc Zinc is largely used in preparing galvanized iror , which is iron coated with zinc in order to protect it against oxidatior, Galvanized iron is not prepared by electrolytic deposition, as the name appears to imply, but by dipping clean iron sheets into molten zinc. Zinc in the form of sheets is also sometimes used for covering roofs, etc. ZINC 451 Zinc is aa important constituent of certain alloys, such as brass and German silver (p. 411), and is present in small proportion in copper coins. With some metals, e.g. cadmium, it is completely miscible in the fused state ; with others, such as lead, it is only slightly miscible (cf. p. 420). Zinc Oxide, ZnO, is obtained by burning zinc in the air or by strongly heating the basic carbonate. It forms a white amorphous powder, which becomes yellow on heating but regains its original colour on cooling. It does not combine with water to form the hydroxide, but dissolves readily in acids to form the corresponding salts. It is used as a pigment under the name of zinc-white. It has one r.dvantage over white lead in not being blackened by hydrogen sulphide. Like magnesium oxide, it becomes incandescent on heating, and even at a white heat does not fuse or decompose. Zinc Hydroxide, Zn(OH) 2 , is formed as a white gelatinous pre- cipitate wl en potassium, sodium or ammonium hydroxide is added to a solution of a zinc salt. It is soluble in excess of the alkalis as well as of ammonia. As regards the alkalis, this is due to the formation of alkali zincate, e.g. Na 2 ZnO 2 ; in the case of ammonia to the forma- tion of complex cations, Zn(NH3) 4 ", the solubility product of these and OH' ions being much larger than that of Zn" and OH' ions. The we.ikly basic character of zinc hydroxide is shown by the formation of complex ions with ammonia and by the hydrolysis of its salts ; its slightly acidic character by the existence of alkali zincates. Zinc Chloride, ZnCl 2 , is obtained in the anhydrous form by heating zinc in a current of chlorine, and in solution by dissolving the metal, oxide or carbonate in hydrochloric acid. On evaporating the solution partial hydrolysis takes place, and hydrogen chloride is given off; the product contains a certain proportion of oxychloride /OH Zn\ , and perhaps zinc oxide as well : >C1 ZnCl 2 +HOH->Zn(OH)Cl + HCl. The anhydrous salt is usually made into sticks. It is very deliquescent, and is sometimes used as a. dehydrating agent. When it is treated with water, the oxychloride generally present remains insoluble, but a clear solution can be obtained by adding hydrochloric acid. The pure anhydrous salt melts about 290, and boils at 730. When zinc oxide is added to a very concentrated solution of zinc chloride, the 452 A TEXT-BOOK OF INORGANIC CHEMISTRY mixture rapidly sets to a hard mass (mainly owing to the formation of oxychloride). Zinc chloride is used for preserving railway sleepers against decay ; also for cleaning metal surfaces in soldering, and as a caustic. Zinc Carbonates The normal carbonate occurs naturally as calamine in rhombohedral crystals isomorphous with calcspar. It is obtained by adding sodium bicarbonate to a solution of zinc sulphate. When the normal alkali carbonates are used, basic salts are ob- tained, the composition of which depends on the temperature, concen- tration and other conditions (cf. magnesium carbonates). Commercial zinc carbonate approximates to the composition ZnCO 3 ,2Zn(OH) 2 . Zinc Sulphate, ZnSO 4 , is obtained by dissolving zinc in sul- phuric acid. With the dilute acid the sulphate and hydrogen are formed ; with the concentrated acid the sulphate is formed and sulphur dioxide is given off: Zn + 2H 2 SO 4 ->ZnSO On the commercial scale, it is prepared by roasting the sulphide to sulphate, and dissolving out the latter with water. Zinc sulphate is usually met with as the heptahydrate, ZnSO 4 ,7H 2 O (hexagonal crystals, isomorphous with MgSO 4 , 7H 2 O), which separates from aqueous solution below 39. The crystals effloresce slightly in the air, above 100 they lose water and form the monohydrate. ZnSO 4 ,H 2 O, and above 240 the anhydrous salt is obtained. At o' loo grams of water dissolve 41.9 grams, at 15 50.8 grams, and at 35' 66.6 grams of the anhydrous salt. Above 39 the compound ZnSO 4 ,6H 2 O separates in monoclinie crystals, isomorphous with MgSO 4 ,6H 2 O. Zinc sulphate forms double salts with the alkali sulphates, e.g. ZnSO 4 ,K 2 SO 4 ,6H 2 O, isomorphous with the corresponding magnesium compounds (p. 448). Zinc Sulphide, ZnS, occurs naturally in regular crystals as zinc blende (coloured brown or black by traces of impurities) and ir. hexagonal crystals as wurtzite. It is obtained as a white precipitate when ammonium sulphide is added to the solution of a zinc salt, also when hydrogen sulphide is passed into an alkaline solution of a zinc salt, but not when the gas is passed into an acid solution. In alkaline solution, owing to the presence of a considerable concentration ofZn 1 and S" ions, the solubility product of ZnS is exceeded and precipita- tion occurs, but in acid solution, owing to the high H* ion concentra- tion, the S" ion concentration is reduced to such an extent that no CADMIUM 453 precipitation occurs. In neutral solution, owing to the accumulation of H- ions during the reaction, precipitation is only partial: but the addition of sodium acetate to the neutral solution prevents the accumulation of H* ions owing to the small ionisation of acetic acid : H' + CH 3 COO'+Na-->CH 3 COOH-|-Na-. The fact that zinc sulphide is insoluble in acetic, but soluble in hydro- chloric acid, can be accounted for on the same lines. Although zinc sulphide is practically insoluble in water, it has con- siderable lendency to pass into colloidal solution. This is often observed when the freshly precipitated compound is being washed with water. Tests for Zinc Zinc salts give no definite colour to the Bunsen flame. The more important tests are the precipitation, by means of ammonium sulphide, of white zinc sulphide, which, for the reasons explained above, is insoluble in acetic acid but soluble in hydrochloric acid. The formation of a white precipitate on addition of ammonium hydroxide, soluble in excess of the latter, is also a useful test. CADMIUM Symbol, Cd. Atomic weight = ii2.4. Molecular weight = 112. 4. Occurrence Cadmium is a comparatively rare element. It occurs in nature as greenockite, CdS, a very scarce mineral, and is present, generally not in excess of i per cent., in most zinc ores, from which it is exclusively obtained commercially. Preparation Of Metal Cadmium, being more volatile than zinc, passes over first when the oxides are heated with charcoal in the commercial preparation of zinc (p. 449). It can be separated from zinc by fractional distillation or by precipitation as sulphide, which, unlike zinc sulphide, is insoluble in dilute hydrochloric acid. The sulphide is then dissolved in concentrated hydrochloric acid, precipi- tated as c arbonate, and the latter converted by heating into the oxide, which is finally distilled with carbon. Properties Cadmium is a silvery-white crystalline metal of density 8.6 ; it melts at 321, and boils at 780. It is malleable and ductile at room temperature. In the air at room temperature it becomes coated with a thin film of oxide ; on heating it burns readily with formation of the oxide. It dissolves in dilute acids with forma- tion of the corresponding salts. 454 A TEXT-BOOK OF INORGANIC CHEMISTRY Cadmium Oxide, CdO, is obtained as a brown powder when cadmium is burned in air, and also by heating the carbonate. At a bright red heat in an atmosphere of oxygen it changes to dark-red cubic crystals. It is readily reduced by heating in a stream of hydrogen. Cadmium Hydroxide, Cd(OH) 2 , is obtained by the action of water on the oxide, or by double decomposition, as a white powder. It is very slightly soluble in water. Unlike zinc hydroxide, it is not soluble in sodium or potassium hydroxide, but is soluble in ammonia on account of the formation of complex ions. Cadmium Chloride, CdCl 2 , prepared by the general methods, separates from solution in colourless, monoclinic crystals as CdCl 2 ,2H 2 O. On heating it loses all its water without further decom- position. Cadmium iodide, CdI 2 , occurs in anhydrous crystals, very soluble in water. The behaviour of the cadmium halides in solution shows that the normal ionisation is small, and that complex anions containing the metal are present. Aqueous solutions of the iodide contain a con- siderable proportion of CdI 3 ' ions. The chloride has much less tendency to complex formation than the iodide. Cadmium Sulphate, CdSO 4 , obtained by the general methods for a soluble salt, usually separates from solution in colourless crystals as the hydrate, 3CdSO 4 ,8H 2 O. The composition of the hydrate is rather remarkable, but appears to be definitely established. The crystals effloresce on exposure to air. Cadmium Sulphide, CdS, is obtained as a bright yellow precipitate by passing hydrogen sulphide into the solution of any cadmium salt. It is soluble in concentrated hydrochloric acid, but insoluble in ammonium sulphide. It is used as a pigment in painting. MERCURY Symbol, Hg. Atomic weight = 200.0. Molecular weight = 200.0. Occurrence Mercury is sometimes found in the free condition as small globules in cavities of rocks, but the chief source is cinnabar, HgS, which forms dark red, often crystalline masses. This ore is found chiefly at Idria in Austria, Almaden in Spain, California, Peru. China, and Japan. Preparation of Metal As mercury has very little affinity for oxygen, it is readily obtained from its ores. According to one method, the sulphide is simply roasted in air and the metal condensed : MERCURY 455 FIG. 86. otherwise it is heated with lime, whereby the sulphur is retained as calcium sulphide and the vapours of mercury are condensed. At Idria the condensation is effected in large chambers connected in series ; in Almaden aludels are used (Fig. 86). These consist of pear-shaped earthen vessels arranged in series, the narrow neck of one being inserted in the base of the next, as shown. The mercury thus obtained is very impure, containing dissolved metals, such as lead, copper, zinc, and arsenic, and also mechanically mixed impurities. The latter are removed by filtering through chamois leather, the former by prolonged shaking with dilute nitric acid or with dilute sulphuric acid containing a little potassium bichromate. The agitation is conveniently effected by drawing a current of air through the liquid. These methods depend upon the fact that the usual impurities are much more easily oxidized and dissolved than mercury itself. The treatment with nitric acid is conveniently carried out in the apparatus represented in Fig. 87. The funnel A ends in a capillary, so that the mercury enters the acid (which is contained in a long glass tube, B) in a fine stream. It is then washed and dried. If required perfectly pure, it is finally distilled in a vacuum. Properties Mercury is the only metal which is liquid at room temperature. To this property it owes the names quicksilver and hydrargyrum (from tiSwp, water, and apyvpos, silver). The solid metal, which forms octahedral crystals, and is very malleable and ductile, melts at - 39, and the liquid boils at 357. The density at o is 13.596, at 20 13.546. Mercury is slightly volatile even at room temperature, as is shown by the fact that the surface of a gold leaf suspended over the liquid in a closed space becomes amalgamated. The vapour of mercury conducts the electric current at high temperatures and gives a peculiar greenish light. This is the principle of the " mercury vapour lamps " now widely used. Mercury is not affected by oxygen at room tem- perature, but on prolonged heating at its boiling-, point in contact with air red mercuric oxide is formed (p. 27). Hydrochloric acid and dilute sulphuric acid have FIG. 87. 456 A TEXT-BOOK OF INORGANIC CHEMISTRY no action on mercury, but when the metal is heated with concen- trated sulphuric acid mercuric sulphate is formed and sulphur dioxide given off: Hg + 2H 2 SO 4 -HgSO 4 Concentrated nitric acid acts on mercury with formation of mercuric nitrate, whilst with moderately dilute acid and excess of mercury mercurous nitrate, Hg 2 (NO 3 ) 2 , is the chief product. Amalgams Alloys of mercury with other metals are called amalgams. They are usually prepared by the direct addition of the other metal to mercury, sometimes also by electrolytic liberation of the metal at a mercury cathode. Amalgams may be liquid or solid, depending upon the nature and proportion of the second metal. Nearly all the metals form amalgams with mercury, platinum and iron being almost the only exceptions. In the majority of cases the amalgams contain definite chemical compounds, but in some instances the metals simply mix without combining. The alkali metals in certain proportions form solid amalgams with mercury, and a number of definite compounds have been isolated, e.g. NaHg 5 , NaHg G , LiHg 6 , KHg 12 , etc. The use of sodium amalgam in preparing hydrogen has already been referred to. The amalgams of zinc and of cadmium are of great importance in connexion with cells used as standards of E.M.F. The use of mercury in extracting gold from its ores has already been mentioned. Chemical com- pounds are probably formed in this case, but their formulae have not been established. Lead readily forms an amalgam with mercury, but in this case no chemical compounds appear to be formed. Compounds of Mercury Mercury forms two series of salts, in both of which it is bivalent. In the mercurous compounds the bivalent ion contains two atoms of mercury, e.g. Hg 2 Cl 2 , Hg 2 SO 4 , whereas the mercuric compounds contain a single bivalent atom, e.g. HgCl 2 , HgSO 4 . The mercurous compounds are sometimes written as if they contained univalent mercury, e.g. HgCl ; but for reasons which cannot be considered here the double formula is regarded as the correct one. Both mercurous and mercuric oxide are weak bases, so that many of the salts are hydrolyzed in solution. On the other hand, mercury differs from copper and other weak bases in that it does not combine with ammonia to form complex ions ; but reaction takes place in a different way. As the terms imply, mercuric salts can in general be obtained from MERCURY 457 mercurous salts by oxidation, and mercurous salts from mercuric salts by reduction. When excess of mercury is used, the mercurous salts are generally obtained. The soluble mercuric salts are very poisonous. MERCUROUS SALTS Mercurous Oxide, Hg 2 O, is obtained by the action of alkali on mercir'ous chloride : Hg,Cl 2 + 2 NaOH->Hg 2 O + 2NaCl + H 2 O. It is a da - k brown powder, which slowly decomposes at room tempe- rature, rapidly on exposure to light, into mercuric oxide and mercury. The corresponding hydroxide, Hg 2 (OH) 2 , has not been obtained ; it is presumably highly unstable. Mercurous Chloride (calomel), Hg 2 Cl 2 , is obtained by add- ing a chloride to a solution of mercurous nitrate, or more usually by heating a mixture of mercurous sulphate (or mercuric sulphate and mercury) with sodium chloride : If the vapour is condensed in a large chamber it is obtained as a fine powder ; if in a small chamber it forms a solid mass. Calomel is also formed by heating a mixture of mercuric chloride and mercury. Mercuious chloride usually occurs as a white powder, which darkens on exposure to light owing to the formation of mercury and mercuric chloride : Its vapour density above 400 is about 118, which appears to indicate that the formula is HgCl. It has been found, however, that the vapour rapidly amalgamates gold leaf, so that free mercury must be present, and this is readily accounted for if the calomel dissociates on heating into mercuric chloride and mercury, as in the above equation. Baker has found that in the perfectly dry condition the density corre- sponds with the formula Hg 2 Cl 2 , an observation which supports the view that the low density of the vapour in the ordinary condition is due to dissociation. Concentrated solutions of hydrochloric acid and other chlorides dissolve calomel to some extent, the above equilibrium being displaced towards the right. Calomel is largely used in medicine. Mercurous Iodide, Hg 2 I 2 , is obtained by direct combination of its elements. It is a yellow to yellowish-green powder, and is unstable, tending to break up into mercuric iodide and mercury. 458 A TEXT-BOOK OF INORGANIC CHEMISTRY Mercurous Nitrate, Hg 2 (NA) 2 , is obtained by the action of dilute nitric acid (the concentrated acid diluted with two or three parts of water) on excess of mercury at a gentle heat. It separates from solution as Hg 2 (NO 3 )2,2H 2 O in monoclinic crystals, which effloresce in the air. When treated with excess of water, basic salts are preci- pitated, e.g. Hg 2 (NO 3 ) 2 ,Hg 2 O,H 2 O, an orange-yellow powder ; but with excess of nitric acid a clear solution is obtained. Mercurous Sulphate, Hg 2 SO 4 , is obtained by adding a soluble sulphate to a solution of mercurous nitrate, or by rubbing together mercuric sulphate with the calculated quantity of mercury. It forms a heavy white crystalline powder, very slightly soluble in water. MERCURIC SALTS Mercuric Oxide, HgO, is obtained as a red crystalline powder by prolonged heating of mercury in contact with air at its boiling- point in a closed vessel (p. 27) ; more readily by strongly heating mercurous or mercuric nitrate. It is obtained as a yellow powder by double decomposition between mercuric chloride and sodium or potassium hydroxide. According to Ostwald, the only difference between -the two forms is the state of division, associated with a slight difference in solubility. By prolonged grinding in a mortar the red oxide becomes yellow, and its solubility in water slightly increases. The red oxide becomes black on heating, due to a new modification stable at high temperatures, but returns to its original colour on cooling. Mercuric Chloride (corrosive sublimate), HgCl 2 , is pre- pared on the large scale by subliming a mixture of mercuric sulphate and sodium chloride, a little manganese dioxide being added to prevent the formation of calomel. It forms rhombic crystals, which are anhydrous. At o 100 parts of water dissolve 5.73 grams, at 20 7.39 grams, and at 40 9.62 grams of the salt. It is still more soluble in alcohol and in ether than in water. It is remarkable that mercuric chloride is only very slightly ionised in aqueous solution ; in this respect it differs from all other metallic chlorides except that of cadmium. This perhaps accounts for its solubility in many organic solvents, and also for the fact that the salt is scarcely affected by concentrated sulphuric acid or nitric acid ; it can, in fact, be volatilized unchanged from boiling sulphuric acid. It is only very slightly hydrolyzed in aqueous solution. With hydrochloric acid mercuric chloride forms a number of MERCURY 459 crystalline double salts ; for example, HgCl 2 ,2HCl,7H 2 O and 2HgCl 2 ,HCl,6H 2 O ; and with the alkali chloride it forms compounds of similar type, e.g. HgCl 2 ,KCl,H 2 O and HgCl 2 .2KCl,H 2 O. In aqueous solutions of these compounds the mercury is mainly present as a constituent of complex ions, such as HgCl 3 ' and HgCl/'. Mercuric chloride is a powerful poison. It is also largely used as an antiseptic and for preserving anatomical specimens, etc. When used as 2Na 3 AlO 3 + 3CO 2 , the sodium aluminate, Na 3 AlO 3 , is extracted with water (the ferric oxide remaining undissolved), and carbon dioxide passed into the solution, when aluminium hydroxite is precipitated : The hydroxide is then washed, dried, and ignited. A simpler method of purification is to dissolve bauxite up to saturation in a solution of sodium hydroxide and then stir crystal- line aluminium oxide into the solution. As the latter is greatly super- saturated with reference to the crystalline hydroxide, the greater portion of the dissolved hydroxide separates out. The electrolysis is carried out in iron vessels lined in the interior with carbon ; the vessel itself forms the cathode, and carbon rods are used as anode. The electrolyte, usually a mixture of cryolite and aluminium fluoride, is kept fused by the current, and aluminium oxide is added from time to time as electrolysis proceeds. Properties Aluminium is a white metal of density 2.7 ; it melts at 657, and is not volatile at a white heat. It is very ductile and malleable ; it can readily be drawn into wire and beaten into sheets. It is very tenacious under ordinary conditions, but at high tempera- tures becomes brittle. It is an excellent conductor of heat and electricity. At room temperature, aluminium becomes coated with a thin film of oxide, which protects it against further oxidation. At high tem- peratures, however, it burns in air to the oxide with a brilliant flame ALUMINIUM 465 and the evolution of a large amount of heat. Pure aluminium is only very slightly acted on by water or steam, but the commercial metal is attacked by these reagents, doubtless owing to the formation of local currents set up between the metal and the impurities. The fact that Aluminium is a highly electro-positive metal is of importance in this connexion (cf. p. 418). Dilute sulphuric acid has very little action on aluminium, but the metal is dissolved by the concentrated acid with formation of aluminium sulphate and evolution of sulphur dioxide. Hydrochloric acid readily dissolves aluminium, but nitric acid, whether dilute or concentrated, is practically without action at room temperature, although at the boiling-point a vigorous reaction takes place. Phos- phoric acid, both in dilute and concentrated solution, readily dis- solves aluminium. Organic acids (e.g. acetic acid) are almost without action on aluminium at room temperature, but have a slight solvent action in the presence of sodium chloride. Aluminium is dissolved by alkali hydroxides with formation of aluminates and evolution of hydrogen : Uses of Aluminium Alloys On account of its lightness, relative cheapness and high electrical conductivity, aluminium is being used instead of copper for the conveyance of electricity. It is also used in making cooking utensils, pans, trays, etc., but the high expectations formerly held as to its usefulness have scarcely been realized in practice, owing to the fact, already mentioned, that it is attacked by alkalis, including soap solution, and also, unless Very pure, by v/ater and by acetic acid. The finely divided metal, mixed with oil, is used as a paint. Aluminium coated with mercury (aluminium amalgam), prepared by dipping sheet aluminium into a solution of mercuric chloride, decom- poses water at room temperature, with liberation of hydrogen. The greater activity of the amalgam,, as compared with aluminium, is partly at least due to the fact that the mercury prevents the formation of a protective coating of oxide on the surface of the metal. Certain other alloys of aluminium are in use. Aluminium bronze (copper v/ith 5 to 12 per cent, of aluminium) is golden-yellow in colour, very tough, takes an excellent polish, and gives good castings. Magnalium (aluminium with 10 to 25 per cent, of magnesium) is also used ; it i:> harder than aluminium and takes a good polish. The great affinity between aluminium and oxygen, already referred 30 466 A TEXT-BOOK OF INORGANIC CHEMISTRY to, is illustrated by the equation representing the heat of the forma- tion of the oxide : 4Al + 3O 2 ->2Al 2 O 3 + 2 x 380,000 cal. The use of aluminium for obtaining other metals from their oxides, according to the so-called " thermite " process (Goldschmidt), is based on this fact. The oxide in question (e.g. chromium, iron or man- ganese oxide) is mixed with finely divided aluminium, and the reaction initiated by a fuse of magnesium ribbon. It then proceeds with the evolution of much heat, the temperature, under favourable conditions, rising to 3000 ; the final products are the metal and aluminium oxide. The same process is used for welding pieces of metal, e.g. steel rails. The junction is surrounded by a mixture of ferric oxide and aluminium, which is ignited in the usual way, and under the influence of the high temperature the two pieces of metal are fused and thoroughly welded together. Aluminium Oxide (Alumina), A1 2 O 3 , occurs naturally in anhy- drous rhombohedral crystals as corundum, which is colourless ; as ruby, coloured by traces of chromium salts ; as sapphire, coloured blue by cobalt salts, and as amethyst. Emery is aluminium oxide contain- ing ferric oxide ; owing to its great hardness it is used for polishing. The hydrated oxide occurs naturally as bauxite, A1 2 O 3 ,3H 2 O, and as diaspore, A1 2 O 3 ,H 2 O. The oxide is obtained as a white amorphous powder by heating aluminium hydroxide, A1(OH) 3 ; in this form it dissolves readily in acids and in alkalis. When the oxide is strongly ignited it becomes practically insoluble in acids ; this is probably due, in part at least, to conversion into the insoluble crystalline form. The insoluble oxide is brought into solution by fusing with alkalis (p. 464). Aluminium Hydroxide, A1(OH) 3 , is obtained as a colourless gelatinous precipitate (hydrogel, p. 355) by adding an alkali hy- droxide (not in excess) or ammonium hydroxide to the solution of an aluminium salt : + 3NH 4 OH-A1(OH) 3 +3NH 4 C1. The precipitate thus obtained contains more or less adsorbed water (p. 313). On heating, the water is driven off and finally, on strong heating, the anhydrous oxide is obtained. The hydrates, A1 2 O 3 ,3H 2 O (=A1(OH) 3 ), A1 2 O 3 ,2H 2 O, and A1 2 O 3 ,H 2 O, are presumably successive stages in the dehydration of the gelatinous hydroxide. Aluminium hydroxide is at the same time a weak base and a weaV ALUMINIUM 467 acid. The basic character is indicated by the existence of salts such as aluminium chloride and aluminium sulphate ; and the fact that the hydroxide is a weak base is shown by the partial hydrolysis of these salts in solution, and by the fanct that salts with weak acids (e.g. carbonate, sulphide) cannot be obtained in the presence of water. The weakly acidic character of the hydroxide is indicated by the formation of salts, the aluminates^ with strong, but not with weak bases. This is illustrated by the solubility* of the hydroxide in excess of potassium hydroxide : and its insolubility in ammonium hydroxide solution. The acid HAlO 3 (cr HA1O 2 ) is so weak that potassium aluminate is hydrolyzed to a considerable extent in solution, and is decomposed even by carbonic acid, with precipitation of aluminium hydroxide (p. 464). The compound HA1O 2 , derived from the normal acid by abstraction of a molecule of water, may be termed meta-aluminic acid (cf. silicic acids, p. 353). The minerals spinelle, Mg(AlO 2 ) 2 , and chrysoberyl, Be(AlO 2 ) 2 , may be regarded as being derived from the latter acid. Aluminum hydroxide may be obtained in colloidal solution (hydrosol) by dissolving the hydrogel in a solution of aluminium chloride snd then removing the latter by dialysis (p. 353). When aluminium hydroxide is precipitated in a solution containing a colouring matter, the latter is carried down with the precipitate, leaving the liquid almost colourless. In such a case the hydroxide is said to have adsorbed the colouring matter. The precipitates are termed lakes. Advantage is taken of this fact in causing certain dyes to adhere firmly to cloth. The aluminium hydroxide is first precipitated in the fibres of the cloth, which is then immersed in the dye. The latter is removed from the solution by the aluminium hydroxide and "fixed" on the cloth. Substances used for the purpose of fixing dyes are known as mordants. Some dyes unite directly with the fibre, and in such cases mordants are unnecessary. When a substance is at the same time a weak acid and a weak base, both properties are necessarily weak. That this must be so is clear when it is remembered that the acidic and basic characters depend upon the presence of H' and OH" ions respectively, and that both cannot be present in considerable proportion in the same solution owing to their tendency to unite to form water. 468 A TEXT-BOOK OF INORGANIC CHEMISTRY Aluminium Chloride, A1C1 3 , is obtained in the anhydrous form by heating aluminium foil in dry chlorine or hydrogen chloride, or by strongly heating a mixture of aluminium oxide and charcoal in a current of chlorine ; the chloride is collected in a dry receiver. It is obtained in aqueous solution by dissolving aluminium oxide in hydrochloric acid; on evaporating the solution the hexahydrate, A1C1 3 ,6H 2 O, separates in colourless crystals. The anhydrous chloride occurs in colourless crystals, which fume in moist air; it vaporizes, without melting, at 183. At low tempera- tures the vapour density corresponds with the formula A1 2 C1 6 , but above 835 it is present entirely as A1C1 3 molecules. It is readily soluble in water, and undergoes considerable hydrolysis ; when the aqueous solution is evaporated to dryness the residue consists almost entirely of the hydroxide. Aluminium chloride is largely used as a catalytic agent in organic chemistry. Aluminium Sulphate, A1 2 (SO 4 ) 3 , is obtained on the large scale by dissolving bauxite in sulphuric acid (the product contains iron as impurity), and also by heating finely-divided clay (aluminium silicate) with concentrated sulphuric acid. The liquid is separated from the insoluble silica and other impurities and concentrated, when the salt separates as Al 2 (SO 4 ) 3 ,i8H 2 O in monoclinic crystals. The pure salt is obtained by dissolving the pure oxide or hydroxide in sulphuric acid. The salt with i8H 2 O becomes anhydrous on heating, and at a higher temperature is completely decomposed into aluminium oxide and sulphur trioxide. It is readily soluble in water, and the aqueous solution is acid owing to hydrolysis r A1 2 (S0 4 ) 3 + 2H 2 0$A1 2 (OH) 2 (S0 4 ) 2 + H 2 S0 4 . Aluminium sulphate is used as a mordant (p. 467). Alums Aluminium sulphate forms with the alkali sulphates a series of double salts known as alums, of which potassium alum, A1 2 (SO 4 ) 3 ,K 2 SO 4 ,24H 2 O (or A1K(SO 4 ) 2 ,I2H 2 O), is a type. The place of the potassium may be taken by sodium, rubidium, caesium, ammonium, silver, or thallium, and further, the isomorphous salts, which contain trivalent chromium or iron instead of aluminium, are also termed alums. The alums all form octahedral crystals with 24H 2 O, and are isomorphous. The general formula for any alum is from the above : M 2 III (S0 4 ) 3 ,M 2 I S0 4 ,2 4 H 2 or M III M I (SO) 2 ,i2H 2 O 2 . ALUMINIUM 469 When the alums are strongly heated, they lose water and sulphur trioxide, and finally yield a mixture of aluminium oxide and alkali sulphate. When ammonium alum is heated the final product is aluminium oxide. The alums are fairly soluble in water, and in dilute solution are decomposed into the component salts. The solutions have an acid reaction, owing to partial hydrolysis. They are used as mordants. Potassium alum, the most important of the alums, was known to the ancient Romans. It is prepared from the naturally occurring basic sulphate, alunite, A1 2 (SO 4 ) 3 ,K 2 SO 4 ,4A1(OH) 3 ,2H 2 O, by roasting, ex- tracting vvith water, whereby alumina remains undissolved, and then concentrating the solution. It may also be prepared by the action of sulphuric acid on bauxite or on clay ; to the aluminium sulphate thus obtained the requisite amount of potassium sulphate is added. Potass um alum loses all its water on heating to 100 in a current of air. The product dehydrated at a fairly high temperature is termed burnt alum; it is usually slightly basic owing to the loss of sulphur trioxide. The solubility of potassium alum in water increases rapidly with rise of temperature ; 100 grams of water take up at o 3.9 grams, at 20 15.1 grams, at 50 44.1 grams, at 100 357 grams of the hydrated salt. Ammonium alum is obtained in well-formed crystals, and is a familiar salt. Sodium alum, on the other hand, crystallizes with difficulty, and is very little used. Another class of double sulphates, sometimes termed pseudo-alums, are knov/n ; they differ from the alums in containing one atom of a divalent metal, e.g. manganese, copper, zinc, iron, or magnesium, instead of two atoms of a univalent element. The aluminium-iron pseudo-alum has the formula Al 2 (SO 4 ) 3 ,FeSO 4 ,24H 2 O. They are not isomorpbous with the alums. Aluminium Silicates Reference has already been made to the fact that double silicates containing aluminium are important constituents of rocks (granite, etc.). When these rocks are dis- integrated by moisture and carbon dioxide, aluminium silicate, being comparatively insoluble, remains, and the other constituents are dissolved and conveyed into the soil. The purest form of aluminium silicate is kaolin, Al 2 O 3 ,2SiO 2 ,2H 2 O, which is colourless and is used in making porcelain. Ordinary clay also consists largely of aluminium silicate, but, unlike kaolin, it is a sedimentary deposit, and contains ferric oxide (to which the brown colour is due), sand, and calcium carbonate as impurities. 470 A TEXT-BOOK OF INORGANIC CHEMISTRY When mixed with water, clays form a plastic mass, which can be moulded into any desired shape and sets hard on heating strongly (burning or " firing"). On this depends their use in the manufacture of bricks, earthenware and porcelain. Bricks are made from impure clay containing ferric oxide and calcium carbonate. Earthenware articles are prepared in the same way from rather purer clays ; the surface is usually covered by a glaze of fused silicate obtained by throwing sodium chloride into the vessel while in process of firing. Porcelain is made from pure clay (kaolin), and the pores are com- pletely filled with fused silicate, obtained by adding felspar and quartz before burning. The use of clay for making bricks, etc., was known to the ancient Babylonians and Egyptians at least 3000 years B.C., and our know- ledge of the early history of the Babylonians depends upon the fact that their records (letters, lists of laws, etc.) were impressed upon moist clay, which was afterwards fired. Ultramarine is a blue pigment obtained by heating- together kaolin, sodium carbonate, sodium sulphate, sulphur and carbon. It occurs in nature as lapis lazuli. It is usually regarded as a double silicate of aluminium and sodium associated with sodium poly- sulphides, but its constitution is not yet fully understood. It is decom- posed by acids with evolution of hydrogen sulphide and disappearance of the blue colour. It is largely used to neutralize the yellow tint in sugar and in linen and cotton goods. It is a constituent of laundry blue. Aluminium Sulphide, A1 2 S 3 , is obtained as a black amor- phous powder by adding sulphur to melted aluminium. It is com- pletely decomposed by water with formation of aluminium hydroxide and hydrogen sulphide : A1 2 S 3 + 6HOH->2A1(OH) 3 + 3H 2 S, and burns in air to the oxide and sulphur dioxide. As both base and acid are very weak, it cannot be obtained by interaction in aqueous solution owing to hydrolysis. Tests for Aluminium The formation of a white gelatinous precipitate when an alkali is added to the solution of an aluminium salt, the precipitate being soluble in excess of potassium or sodium hydroxide, but insoluble in ammonia. Further, when ammonium chloride is added to a solution of potassium aluminate, aluminium hydroxide is precipitated ; this test depends upon the formation and subsequent hydrolysis of ammonium aluminate. When aluminium THALLIUM 471 oxide is moistened with a few drops of cobalt nitrate solution and strongly heated on charcoal in the blowpipe flame, a blue mass is obtained. GALLIUM Symbol, Ga. Atomic weight =69. 9 Gallium was discovered in 1875 by Lecoq de Boisbaudran in zinc blende by means of the spectroscope. Its existence and properties had been foretold some years before by Mendeleeff on the basis of the Periodic System (p. 369). Gallium is a very rare element. The metal itself is white and melts at 30 ; its density is 5.9. In most of its properties it strongly resembles aluminium. Thus it has no action on water, it becomes coated with a film of oxide on heating in the air, it dissolves readily in hydrochloric acid and in alkali hydroxides, but is insoluble in dilute nitric acid at room temperature. In its most important compounds gallium is trivalent. The hydroxide, Ga(OH)o, is soluble in alkalis, the chloride, GaCl 3 , fumes in the air, and is hydrolyzed in aqueous solution, and the sulphate forms a well-defined alum, 4H 2 O, with ammonium sulphate. INDIUM Symbol, In. Atomic weight= 114.8. Indium is also a very rare element, and shows many analogies with aluminium. It was discovered in 1863 by Reich and Richter in a zinc ore from Freiberg by means of its spectrum, which is characterized by the presence of a bright blue line. Th.j metal itself is colourless and very soft; it melts about 155. It is stable in the air at room temperature, but on heating burns with a blue flame to the oxide, In 2 O 3 . It slowly decomposes water at room temperature with forma- tion of the hydroxide. In its most important series of compounds indium is trivalent. The hy- droxide, In(OH) 3 , is soluble in excess of potassium or sodium hydroxide, so that it has weakly acidic as well as basic properties. The chloride, InCl 3 , unlike those of gallium and aluminium, can be obtained without decomposition by evaporation of its aqueous solution at 100. The sulphate, In 2 (SO 4 ) 3 , forms an alum with ammonium sulphate. Besides acting as a trivalent element, indium forms two series of compounds in which it functions as a univalent and as a divalent element respectively. The halogen compounds of these series, e.g. IriCl and InCl 2 , are best known. THALLIUM Symbol, Tl. Atomic weight=204.o. Thallium was discovered by Crookes in 1861 in the mud of a sulphuric acid chamber at Tilkerorde in the Hartz mountains by means of the spectroscope. The spectrum is characterized by the 472 A TEXT-BOOK OF INORGANIC CHEMISTRY presence of a bright green line, hence the name (from 6a\\6s, a green twig). It is also found in the flue dust of the chambers. The only mineral which contains a considerable proportion of thallium is crookesite, a selenide of copper which contains 16 to 18 per cent, of thallium and 3 to 5 per cent, of silver. Preparation and Properties of Metal Thallium is best obtained from the dust in the flues of the sulphuric acid chamber by boiling with water, pouring off the solution, and adding to it sodium chloride, which precipitates the thallium as T1C1. The chloride is further purified, heated with sulphuric acid to obtain the sulphate, and the aqueous solution of the latter electrolyzed with a platinum anode and a cathode of copper foil. Thallium is a soft, heavy metal with a bluish tinge, like lead. Its density is n.8; it melts at 285 and boils at 1650. At room temperature it rapidly becomes coated with a mixture of oxides. It is stable in air-free water at ordinary temperature, but decomposes water at a red heat. The metal is slowly acted upon by water containing dissolved oxygen, with formation of thallous hydroxide, Tl(OH), which is soluble in water. When heated in air, it burns to the trioxide with a green flame. It is readily soluble in dilute nitric or sulphuric acids, but is only acted on slowly by hydrochloric acid, owing to the insolubility of the chloride, T1C1. Compounds of Thallium Thallium forms two well-defined series of compounds, thallous compounds, in which it is univalent, and thallic compounds, in which it is trivalent. The thallous salts resemble those of the alkalis inasmuch as the hydroxide and car- bonate are soluble in water, whilst the relatively slight solubility of the thallous halides recalls the silver halides. In its trivalent compounds, thallium resembles aluminium and gold. Thallous Compounds Thallous oxide, T1 2 O, is obtained as a black, hygroscopic powder by heating thallous hydroxide to 100. Thallous hydroxide, T1OH, is obtained by the simul- taneous action of water and oxygen on the metal or by double decomposition between barium hydroxide and thallous sulphate in solution. It is soluble in water and is a fairly strong base. Thallous chloride, T1C1, is obtained as a white curdy precipi- tate by adding a soluble chloride to the solution of a thallous salt. At 20 TOO grams of water take up 0.34 grams, at 100 2.4 grams of the salt. Thallous iodide, Til, is obtained by double decomposition as a yellow precipitate. At 20 100 grams of water dissolve 0.063 grams of the salt. Thallous carbonate, THALLIUM 473 T1 2 CO 3 , is obtained by saturating a solution of thallous hydroxide with carbon dioxide. It is fairly soluble in water (about 5 grams in loo grams of water at 18), and separates in long, prismatic needles o;i evaporating the solution. Thallic Compounds Thallic oxide, T1 2 O 3 , is obtained as a dark brown powder by burning thallium in the air, or by heating the hydroxide, T1(OH) 3 , or hydrated oxide, T1 2 O 3 ,H 2 O. It is in- soluble in water, is reduced by heating with hydrochloric acid with evolution of chlorine, and is also reduced by boiling with sulphuric acid, oxygen being given off and thallous sulphate formed. At a red heat thallic oxide is completely decomposed into thallous oxide ana oxygen. Thallic hydroxide, T1(OH) 3 ( = T1 2 O 3 ,3H 2 O), is said to be formed as a brown precipitate by the hydrolytic decom- position of thallic salts, but the compound thus obtained may be Tl 2 O 3 ,HvO. It is known that the voluminous reddish-brown pre- cipitate obtained on adding ammonia to a solution of a thallic salt has the composition T1 2 O 3 ,H 2 O. Thallic hydroxide is a much weaker base than thallous hydroxide. Thallic chloride, T1C1 3 , is obtained by passing chlorine into thallous chloride suspended in water till a clear solution is obtained ; on evaporation the tetra- hydrate, T1C1 3 ,4H 2 O, separates in colourless needles. By dehydrating in a vacuum over sulphuric acid, the anhydrous salt is finally obtained. The aqueous solution of the trichloride has an acid reaction owing to hydrolysis, and when it is considerably diluted the oxyhydrate, ~T1 2 O 3 ,H._,O, is precipitated. Thallic iodide, T1I 3 , is obtained in black, rhombic crystals by the action of iodine in alcoholic solution on thallous iodide. It is isomorphous with the tri-iodides of rubidium and caesium (p. 402). It very readily splits up into thallous iodide and iodine : Thallic .ndphate, T1 2 (SO 4 ) 3 , has been obtained in the anhydrous form by heating thallium hydrogen sulphate, T1 2 (SO 4 ) 3 ,H 2 SO 4 ,8H 2 O, at 220. The latter salt is obtained by dissolving thallic oxide in dilute sulphuric acid and evaporating the solution. Thallic nitrate, T1(NO 3 ) 3 , is obtained in colourless lustrous crystals with 3H 2 O by dissolving thallic oxide in concentrated nitric acid and evaporating the solution. It is hydrolyzed by water with precipita- tion of thallic oxide monohydrate. A number of double thallous-thallic salts, of the type . are kno\\ n. 474 A TEXT-BOOK OF INORGANIC CHEMISTRY Thallium salts are readily recognized by their characteristic spectrum. THE SCANDIUM SUB-GROUP As already mentioned, this sub-group consists of four rare elements, scandium, yttrium, lanthanum, ytterbium. They belong to the so-called rare earths, a group of elements which are so similar in behaviour that very great difficulty has been experienced in separating them. Fractional precipitation and fractional crystallization have been largely used for this purpose, and advantage has also been taken of the fact that the nitrates decompose at different temperatures on heating. Most of them have characteristic spectra. Scandium (Sc= 44.1) was discovered in 1879 by Nilson and by Cleve in certain rare minerals gadolinite and euxenite found in Scandinavia. Its properties were found to correspond closely with those of the till then unknown element ekaboron, as foretold by Mendeleeff. It forms only one oxide, Sc 2 O y , which occurs as a bulky white powder ; and a corresponding series of salts. The hydroxide, Sc(OH) 3> occurs as a white gelatinous mass, and is a very weak base. Yttrium was discovered by Mosander (1843) and, like scandium, occurs in the gadolinite earths. It always functions as a trivalent element. The oxide, Y 2 O 3 , is colourless, and dissolves in acids to form corresponding salts. Like the other rare earths, it gives a characteristic spectrum, consisting of numerous lines. Lanthanum, (La= 139.0) was discovered in 1839 in the cerite earths by Mosander. The metal is obtained by electrolysis of the fused chloride. It is malleable, becomes coated with oxide in dry air and burns vigorously to a mixture of oxide and nitride on heating. It decomposes water at room temperature. Lanthanum is invariably trivalent in its compounds. The oxide, La 2 O 3 , is colourless, and is the strongest base among the rare earths. The salts are also colourless, and are only very slightly hydrolyzed in aqueous solution. Ytterbium (Yb= 172.0) Quite recently (1907), Urbain and Auer von Welsbach showed almost simultaneously that the element hitherto regarded as ytterbium is a mixture of two elements, ytterbium proper (neoytterbium), and a new element termed by Urbain lutecium (At. wt. = 174.0). Comparison of the Members of the Aluminium Group and SummaryCorresponding with their position in -the periodic table, all the members of this group function as tri- valent elements. The first member of the boron sub-group boron itself is a typical non-metal, and the hydroxide, B(OH) 3 , is an acid, though a comparatively weak one. Aluminium in most of its pro- perties behaves as a metal, but, as already indicated, the hydroxide, A1(OH) 3 , is a weak acid as well as a weak base. Gallium, indium and thallium, the remaining members of the sub-group, are also weak bases in their trivalent compounds, as is indicated by the partial hydrolysis of the chlorides and by other properties. It is ELEMENTS OF THE ALUMINIUM GROUP 475 only in its trivalent character, however, that thallium behaves like other members of the group. The thallous compounds, of the type T1X, behave in some respects like the alkali metals, in other respects like silver; this is shown by the fact that thallous hydroxide and carbonate are soluble in water, whereas the thallous halides are comparatively insoluble. The fact that, in contrast to thallic hydroxide, thallous hydroxide is a comparatively strong base, is of considerable interest and importance. Metallic thallium strongly resembles lead. The variation of the properties of the boron sub-group with the atomic weight is shown in the accompanying table: B Al Ga In Tl Atomic weight . II. 27.1 69.9 II4.8 204.0 Density .... 2.45 2.7 5-9 7-4 1 1.8 Melting-point . . >2000 6 57 30.0 176 285 Atomic volume . 4-5 1 0.0 1 1.8 15-5 i7-3 CHAPTER XXXI ELEMENTS OF THE CARBON GROUP (GROUP IV) Sub-group A Sub-group B Titanium, Ti . . 48.1 Carbon, C 12.00 Zirconium, Zr . . . 90.6 Silicon, Si . . . 28.3 Cerium, Ce . . 140.25 Germanium, Ge . . 72.5 Thorium, Th . 232.4 Tin, Sn . . . . 119.0 Lead, Pb ... 207.1 THE sub-group A contains four rare elements ; and germanium, the middle member of sub-group B, is also a rare element. Of the remaining elements of sub-group B, carbon and silicon have already been dealt with among the non-metals. Tin and lead, on the other hand, are metals, but their oxides show also weak acidic proper- ties. In this, as in certain other families of elements already con- sidered, the electro-positive character increases with increase in atomic weight. In accordance with their position in the periodic table all the elements of this group are quadrivalent, but some of them, notably tin and lead, show other valencies as well. THE TITANIUM SUB-GROUP Titanium (Ti = 48. i) occurs naturally as the dioxide in the rare minerals rutile, anatase and brookite, and as ferrous titanate, FeTiO 3 , in certain iron ores. The element is obtained by reducing the dioxide with carbon in the electric furnace, but does not appear to have been obtained quite pure. It burns in oxygen to the dioxide, and combines readily with nitrogen at 800 to form the nitride, Ti 3 N 4 . In its compounds it shows considerable analogy with silicon (q.v.). Titanium dioxide, TiO 2 , has both acidic and basic properties. When fused with alkalis or alkali carbonates, alkali titanates, e.g. K 2 TiO 3 , are formed. The titanates are soluble in excess of hydrochloric or sulphuric acid, and the solutions presumably contain titanic chloride, TiCl 4 , or sulphate, Ti(SO 4 ) 2 . When alkali hydroxide or ammonia is added in excess to either of these solu- tions, titanic acid, Ti(OH) 4 ,.rH 2 O, is obtained as a voluminous white precipitate, Titanic chloride, TiCl 4 , is obtained as a colourless, fuming liquid by heating a mixture of titanium dioxide and carbon, and passing chlorine over it. The tetrachloride is completely hydrolyzed on addition of water : + 4HOH->Ti(OH) 4 + 4 HCl. 476 GERMANIUM 477 Titanium forms two other series of salts, in which it is divalent and trivalent respectively. The yellow colour obtained by the addition of hydrogen peroxide to a solution of titanium dioxide in concentrated sulphuric acid (p. 143) is due to the formation of titanium peroxide, TiO 3 . Zirconium occurs naturally chiefly as Zircon, ZrSiO 4 . The metal is obtained by reducing the dioxide, ZrO 2 , with carbon in the electric furnace under definite conditions. The hydroxide, Zr(OH) 4 , is obtained as a gelatinous precipitate by adding ammonia to a solution of a zirconium salt ; it is only very slightly soluble in aqueous alkalis, but when fused with the latter forms salts of the type Na 2 ZrO 3 , which are decomposed by water. The fact that zirconium sulphate, Zr(SO 4 ) 2 , can be recrysta lized from water shows that zirconium hydroxide is a base of moderate strength. Cerium occurs with other elements (p. 474) in the rare earths, more particularly in cerite, euxenite and monazite. The metal, which has been obtained by heat- ing the dioxide with magnesium powder, has a density of 6.7, melts at 623 and is permanent in dry air. Cerium forms two well-defined series of salts, the cerous salts, e.g. CeQ 3 , in which it is trivalent, and the eerie salts, e.g. CeCl 4 , in which it is quadrivalent. Cerous oxide, Ce 2 O3, is obtained'by heating the carbonate in a current of hydrogen. Ceric oxide, CeO 2 , is obtained as a pale yellow powder on heating the carbonate, nitrate or sulphate in air. Thorium, like cerium, is chiefly obtained from monazite sand (p. 209), it also occurs in tetragonal crystals as thorite, ThSiO 4 , isomorphous with zircon, rutile and cassiterite, and as thorianite, a mineral obtained from Ceylon, which is chiefly a nixture of thorium oxide, ThO 2 , and uranium oxide, UO 2 . Thorium sulphate, Th(SO 4 ) 2 , and the chloride, ThCl 4 , can be crystallized from aqueous solution, so that the hydroxide, which is insoluble in alkalis, is a fairly strong base. TIv; dioxide, ThO 2 , is used for the preparation of incandescent mantles, as already mentioned. Thorium compounds are radio-active (p. 567). THE TIN SUB-GROUP Germanium (06=72.5), was discovered by Winkler (1886) in an argentiferous mineral, argyrodite, 3Ag 2 S,GeS 2 (perhaps 4Ag 2 S,GeS 2 ), found at Freiberg in Saxony. It is also found in small proportion in some of the rare earths, for example in samarskite and in gadollnite. The metal itself is obtained from its salts by heating in hydrogen. It is a brittle, grayish-white metal of density 5.47 at 20. It melts below 900, and is only slowly vaporized at 1500. At room temperature it is not acted upon in the air ; at red-heat it burns to the dioxide, GeO 2 . Germanium is insoluble in hydrochloric acid, but nitric acid converts it into a hyd rated dioxide, GeO 2 ^H 2 O. Germanium, like tin, forms two series of salts, as it functions both as a divalent and tetra\alent element. Germanium dichloride, GeCl 2 , is obtained by passing hydrogen chloride over the heated sulphide, GeS. It has not been much investigated. Germanous sul- phide, GeS, is obtained by heating a mixture of germanic sulphide, GeS 2 , and germaniu n in a current of carbon dioxide. The sulphide occurs in grayish-black crystals, ;,nd is the best defined germanous compound. Germai'ium tetrachloride, GeCl 4 , is obtained by heating germanium in a current oi chlorine. It is a colourless liquid, which boils at 80, and is completely 478 A TEXT-BOOK OF INORGANIC CHEMISTRY decomposed by water, with formation of the hydroxide, Ge(OH) 4 and hydro- chloric acid. Germaniiim dioxide, GeO 2 , is obtained by burning germanium in air, or by heating the hydroxide. It is a white powder of density 4.7 at 78, and is not affected by heat. It is soluble in water, and the solution has an acid reaction, so that the dioxide has acidic properties, which is confirmed by its ready solubility in alkalis. It is also soluble in acids, and appears to possess slightly basic properties. Germanium hydroxide, Ge(OH) 4 , obtained as a gelatinous precipitate by hydrolysis of the tetrachloride, is a weak acid and also a very weak base. Germanium disulphide, GeS 2 , is obtained as a white precipitate by passing hydrogen sulphide through a solution of germanium dioxide in hydrochloric acid. It is slightly soluble in water, readily soluble in solutions of alkali hydroxides and of alkali sulphides (cf. stannic sulphide). TIN Symbol, Sn. Atomic weight, 119.0. Molecular weight uncertain. History As articles of bronze (an alloy of tin and copper) at least 4000 years old have been found in Egypt, it is clear that the metal must have been known at a very early period. The ores used by the ancient Egyptians as a source of tin appear to have come from northern Persia. In later times, tin was obtained, among other sources, from Cornwall and Devonshire, and it appears from Pliny that the earliest name applied to the British Islands was the Cassi- terides, from cassiterite, the chief ore of tin. The Latin name for tin is Stannutti) hence the symbol Sn. Occurrence Tin does not appear to occur naturally in the free condition. It occurs almost exclusively as the dioxide, SnO 2 , cassi- terite or tinstone, in tetragonal crystals, which are usually coloured brown or black by traces of iron. The ore usually contains (besides iron) silica, sulphur, arsenic, copper, lead and other impurities. Tinstone is found in Cornwall, Saxony, Bohemia, Bolivia, Mexico, but is now chiefly obtained (generally of a very high degree of purity) from Banca, Biliton and other islands in the Straits Settlements. Preparation of Metal The ore is first roasted, whereby the arsenic and sulphur usually present are removed, and is then washed to remove impurities such as copper sulphate and ferric oxide. The dioxide is then mixed with powdered coal and heated in a rever- beratory furnace : The tin thus obtained is melted at a low temperature and poured oft" from some of the impurities (the process is termed liquation), and is TIN 479 finally stirred with poles of green wood (cf. copper, p. 409) in order to reduce any oxide that may have been formed. Properties of Metal Tin is a silvery-white lustrous metal of density 7.29 at 15, it melts at 231.5, and boils at 2270. It is malleable and ductile at the ordinary temperature, and is so soft that it can be cut with a knife. When tin is bent a peculiar crackling noise, known as the cry of tin^ is noticed ; it is probably due to the friction of the crystalline particles. The crystalline character of tin is most clearly shown by etching the surface with warm hydrochloric acid or aqua regia. When heated to 200 tin becomes brittle, and can be powdered in a mortar. Tin exists in different allotropic modifications. At low temperatures, most rapidly at -48, in contact with an alcoholic solution of "pink salt," the ordinary white modification changes to a gray powder of density 5.8. E. Cohen has shown that the transition temperature for the two modifications is at 18 (cf. p. 276) ; below 18 ordinary tin in contact with gray tin changes to a powder of the latter form. This phenomenon is called the tin pest. It has sometimes been observed in organ-pipes which have been in use for many years. The brittle tin obtained by heating the metal to 200 is presumably a third modification. Tin is stable in the air at room temperature, but on heating strongly it burns to the dioxide, SnO 2 . It is dissolved by hot hydrochloric acid with formation of stannous chloride and hydrogen : Sn + 2HCl-SnCl, + H 2 , and is also attacked by hot concentrated sulphuric acid, stannic sulphate and sulphur dioxide being formed : 2H 2 SO 4 ->SnSO 4 + S0 2 + 2 H 2 0. The action of nitric acid on tin depends on the concentration of the acid and the temperature. With cold dilute nitric acid, stannous nitrate, Sn(NO 3 ) 2 , is the chief product, but a little stannic nitrate, Sn(NO 3 ) .;, may also be formed. The fairly concentrated acid vigor- ously attacks tin, metastannic acid (p. 482) being the chief product. The strongest nitric acid has practically no action on tm. Tin is dissolved by a- boiling solution of sodium or potassium hydroxide, with evolution of hydrogen and formation of alkali stannate : Sn + 2KOH + HoO->K 2 SnOo + 2H 9 . 4 8o A TEXT-BOOK OF INORGANIC CHEMISTRY Uses Of Tin. Alloys Tin is largely used as a protective coating for metals which are acted on by air. The ordinary " tin " utensils, widely used for household purposes, are constructed of tin plate, which is made by dipping carefully cleaned sheets of iron into melted tin. Many important alloys containing tin are in use. The alloys of tin and lead melt at a lower temperature than either metal (p. 199). Pewter contains 75 per cent, of tin and 25 per cent, of lead. Common solder contains one part of lead and one part of tin ; other solders contain the metals in different proportions. Tin is an important constituent of the bronzes, which have already been referred to under copper. Tin amalgam is used as a coating for mirrors. COMPOUNDS OF TIN Tin forms two series of compounds, stannous compounds, of the type SnX 2 , in which it is bivalent, and stannic compounds, of the type SnX 4 , in which the tin is quadrivalent. The hydroxides, Sn(OH) 2 and Sn(OH) 4 , are weak bases, and both have acidic properties. STANNOUS COMPOUNDS Stannous Oxide, SnO, and the Hydrate, 2SnO,H 2 O Stannous oxide hydrate, 2SnO,H 2 O, is obtained as a white precipitate on adding alkali carbonate to a stannous chloride solution : The hydrate is insoluble in ammonium hydroxide, but is dissolved by sodium or potassium hydroxide, with formation of an alkali stannite. Stannous oxide is obtained by heating the hydrated oxide in a current of carbon dioxide. It is a black powder, which dissolves in acids to form stannous salts. On heating in air, it catches fire and burns to the dioxide. Stannous Chloride, SnCl 2 , is obtained by dissolving tin in concentrated hydrochloric acid ; on evaporating the solution the dihydrate, SnCl,2H 2 O, separates in monoclinic crystals. The anhy- drous salt is obtained by heating pure tin, or stannous chloride dihydrate, in a current of dry hydrogen chloride. Stannous chloride melts at 250 and boils at 606. Vapour density determinations give values for the molecular weight which even at TIN 481 650 are not much larger than those corresponding with the formula SnCl 2 . When free from stannic salt, stannous chloride dissolves to a clear solution in a moderate amount of water, but with excess of water, a white precipitate of stannous oxychloride is formed by hydrolysis : + HOH->Sn(OH)Cl + HCl. Stannous chloride is a powerful reducing agent, especially in alkaline solution, being itself oxidized to stannic hydroxide or stannic chloride. It reduces mercuric chloride to calomel and finally to metallic mercury : l 2 + SnCl 2 ->Hg 2 Cl 2 +SnCl 4 Solutions of stannous chloride absorb oxygen from the air, and stannic hydroxide is precipitated : In presence of hydrochloric acid, stannic chloride is probably the first product, but it is ultimately more or less completely hydrolyzed to the hydroxide : The same equations express its behaviour with other oxidizing agents. Stannous chloride is used commercially as a mordant and as a reducing agent. Stannous Sulphide, SnS, is described in connexion with stannic sulphide (see below). STANNIC SALTS Stannic Oxide, SnO 2 , occurs naturally in crystalline form as cassiterite. It is prepared by burning tin in the air, but is usually obtained by strongly heating metastannic acid (q.v.). As thus pre- pared, it is a white amorphous powder which turns yellow on heating, but returns to its original colour on cooling. It is insoluble in acids or aqueous alkalis, but when fused with alkalis or alkali carbonates, alkali stannate is formed : Hydrates of Stannic Oxide At least two hydrates of stannic oxide, namely, stannic acid, H 2 SnO 3 (SnO 2 ,H 2 O), and metastannic 482 A TEXT-BOOK OF INORGANIC CHEMISTRY add, H 10 Sn 5 O 15 (5[SnO 2 ,H 2 O]), appear to be definitely known. A hydrate which corresponds in composition with orthostannic acid (or stannic hydroxide), Sn(OH) 4 , has also been obtained. Stannic Acid, H 2 SnO 3 , is obtained as a white gelatinous pre- cipitate by the action of ammonium hydroxide or of calcium carbonate on stannic chloride: When dried in the air, it has the composition H 4 SnO 4 ( = SnO 2 ,2H 2 O) ; when dried in vacuo or at 100, it has the composition H 2 SnO 3 . The precipitate, before drying, is soluble in sulphuric and in hydrochloric acids ; the former solution presumably contains stannic sulphate, Sn(SO 4 ) 2 , which shows the basic character of the hydroxide. The freshly precipitated acid is also' soluble in alkalis with forma- tion of stannates. The compounds Na 2 SnO 3 ,3H 2 O and K 2 SnO 3 ,3H 2 O are well-defined salts, readily soluble in water. They can also be prepared by fusing stannic oxide with the alkali hydroxides or car- bonates. The sodium compound is used as a mordant in dyeing under the name of " preparing salt." Metastannic Acid, H 10 Sn 5 O 15 , is obtained by the action of moderately concentrated nitric acid on metallic tin : 5 Sn + 2oH N O 3 ->H 10 Sn 6 O 15 + 2oN O 2 + 5 H 2 O. It is probable that in the first stage of the reaction stannic nitrate is formed and is rapidly hydrolyzed. Metastannic acid differs from stannic acid in being insoluble in acids. With aqueous alkalis it forms well-defined salts, termed metastan- nates, e.g. K 2 Sn 6 O n ,4H 2 O (or K 2 O,5SnO 2 ,4H 2 O) and Na 2 Sn G Oi 1 ,4H 2 O, showing that the acid is dibasic. When metastannic acid is fused with caustic alkalis, alkali stannates (not metastannates) are obtained. Conversely, when stannic acid is heated, it partially changes to metastannic acid. Stannic Chloride, SnCl 4 , is prepared by passing chlorine ovei finely divided tin heated nearly to its melting-point. It is a colourless liquid which boils at 1 14 and fumes in the air. When water is added to stannic chloride, heat is given out and a solid mass containing one or more hydrates is obtained : with more water a clear solution is formed. The solution is practically a non-conductor of electricity at first, showing that Sn"" ions can only be present in very minute amount, but the conductivity gradually increases, the solution mean- TIN 483 while remaining clear. It appears probable that the chloride under- goes slow hydrolysis according to the equation the stannic oxide remaining as a colloidal hydrosol (p. 355). On boiling the solution, stannic acid is precipitated. Hydrates of stannic chloride with 3, 4, 5, 8, and 9 molecules of water have been described. The pentahydrate is used as a mordant. Stannic chloride forms double salts with the alkali chlorides^.^-. SnCl 4 ,2HCl and SnCl 4 ,2NH 4 Cl. The latter, which is known as " pink salt," is used as a mordant. Tin Sulphides Stannotts sulphide, SnS, is made by heating tin foil in sulphur vapour or by passing hydrogen sulphide through an acidified solution of stannous chloride. It is a dark-brown powder which is insoluble in solutions of the monosulphides of the alkalis (K 2 S ; (NH 4 ) 2 S), but dissolves in solutions of alkali polysulphides (cf. p. 503) with formation of thiostannates : SnS + (NH 4 ) 2 S 2 -KNH 4 ) 2 SnS 3 . When excess of hydrochloric acid is added to a thiostannate solution, stannic sulphide, SnS 2 , is precipitated : Stannic Sulphide, SnS 2 , is obtained as above described and also by passing hydrogen sulphide into a solution of a stannic salt. Commercially it is prepared by heating in a retort a mixture of tin, mercury, sulphur and ammonium chloride till fumes are no longer given off. The sulphide which remains in the retort forms golden- yellow lustrous scales, and is used as a pigment under the name of mosaic gold. The thiDstannates, mentioned above, bear the same relation to stannic sulphide as the stannates do to stannic oxide. Tests for Tin All tin salts yield globules of metallic tin when mixed with sodium carbonate and reduced on charcoal. The re- actions of the sulphides of tin, as described above, are characteristic. The mercuric chloride [test for stannous chloride is often useful Stannic salts can be converted to stannous salts by nascent hydrogen and the test then applied. 484 A TEXT-BOOK OF INORGANIC CHEMISTRY LEAD Symbol, Pb. Atomic weight, 207.1. Molecular weight unknown. History Lead was known to the ancient Egyptians at least three thousand years ago. In ancient times, however, there was some confusion between tin and lead ; and Pliny was the first to distinguish clearly between plumbum nigrum, lead, and plumbum album or candidum, tin. The metal was used by the ancient Romans in making water-pipes. Occurrence The chief ore of lead is galena, PbS, which is very widely distributed. It also occurs in considerable quantities as cerussite^ PbCOo ; and in relatively small amount as anglcsite, PbSO 4 ; crocoisite, PbCrO 4 ; wulfenite^ PbMoO 4 ; and pyromorphite, PbCl 2 , 3 Pb 3 (P0 4 ) 2 . Preparation of Metal The metal is obtained almost exclu- sively from galena, which is generally a very pure material. The ore is first roasted in a reverberatory furnace (p. 388), whereby part of the sulphide is burned to oxide, part is oxidized to sulphate, and a con- siderable proportion remains unaltered : 2PbS + 3O 2 ->2PbO + 2SO 2 ; PbS + 2O 2 ->PbSO 4 . The temperature is then raised, when the sulphate and oxide react with the sulphide to form the metal ; 2PbO + PbS->3Pb + SO 2 . In North America and in Spain lead is sometimes obtained from galena by heating with iron ; the latter combines with the sulphur tc form ferrous sulphide, FeS. In practice, for economical reasons, ores which yield iron during the process are used instead of the metal they are mixed with galena and coke and strongly heated in a blast furnace. The ferrous sulphide impurities rise to the surface and the molten lead separates out below : As very pure lead is required for many purposes, e.g. for accumula tors, the metal obtained by either of the above processes must b( refined. Metals which are more easily oxidized than lead are con verted into oxides by heating the metal in an oxidizing atmosphere, and rise to the surface of the melted lead as a scum. Other impuri- ties are removed by "poling" (p. 409). LEAD 485 Properties Lead is a bluish-white, very soft metal, very lustrous on freshly -cut surfaces but rapidly becoming dull owing to surface oxidation. It melts at 327 and boils at 1525. Recent determina- tions of its vapour-density at 1870 show that it is monatomic. Its density at 20 is 11.34. Lead occurs in octahedral crystals belonging to the regular system, as can readily be shown by melting it in a crucible, allowing it partially to solidify, piercing a hole in the crust, and pouring off the still liquid portion. From the chemical point of view lead is a fairly active metal ; but the action of reagents on it is in many cases retarded by the formation of protective coatings and by the slight solubility of its compounds. The very thin black film which forms on the surface at room temperature is probably the suboxide, Pb/). When lead is heated strongly in air the oxide PbO is formed. The great influence of the state of division on the rate of reaction is well shown by the fact that very finely divided lead, obtained by reducing lead tartrate in a current of hydrogen at as low a temperature as possible, catches fire in the air at room temperature. Lead is not attacked by water free from oxygen, but water contain- ing oxygen dissolves it with formation of lead hydroxide : 2 0->Pb(OH) 2 . The solvent action of natural waters on lead is very important in con- nexion with the extended use of lead pipes for conveying water. Water saturated with air can retain in solution more than o. i mg. of lead per litre. Water containing carbonates or sulphates, however, dissolves only very minute quantities of lead owing to the very slight solubility of lead carbonate and sulphate ; and, further, these salts form a coat- ing on the metal which prevents further solvent action. Free carbon dioxide ur.der certain circumstances increases the solubility, and it is of advantage to add sufficient alkali to the water to convert it to the bicarbonate. It follows from the above that lead pipes can be used safely for most natural waters ; but very pure waters, e.g. rain water, may take up dangerous amounts of lead. The water of Loch Katrine, used by the city of Glasgow, is filtered through beds of chalk, and the carbonate thus taken up greatly lessens the solvent action of the water on the lead pipes. 1 Water containing ammonium salts, and 1 The mr in object in filtering the water through chalk was to supply the lime salts essent al for the proper nourishment of the human organism. It is now held, howe\er, that sufficient lime for this purpose is present in ordinary foods (especially farinaceous foods). 486 A TEXT-BOOK OF INORGANIC CHEMISTRY also water containing weak organic acids, e.g. acetic acid, have con- siderable solvent action on lead. Lead is rapidly dissolved by nitric acid. Hydrochloric acid is practically without action at room temperature ; but the hot concen- trated acid acts slowly on it, with formation of lead chloride and hydrogen. Hot concentrated sulphuric acid slowly dissolves lead as the sulphate, PbSO 4 ; but a less concentrated acid is practically with- out action, owing to the insolubility of lead sulphate. Lead, being a weakly electro-positive metal, is displaced from its salts in solution by zinc, magnesium, and other metals (p. 418). When a strip of zinc is suspended in a dilute solution of lead acetate, the lead is deposited as a branching crystalline structure known as the lead tree : Pb(C 2 H 3 Q 2 ) 2 +Zn->Zn(C 2 H 3 2 ) 2 + Pb. Uses of Lead. Alloys On account of the readiness with which it can be worked and its resistance to many reagents lead is largely used for commercial purposes. Lead pipes are made by heat- ing the metal till soft and squeezing it into shape by hydraulic pressure. Lead bullets, which contain a small proportion of arsenic, are made by forcing the metal into moulds. Alloys of lead with tin have already been referred to (p. 480) ; lead-antimony and lead-bismuth alloys will be mentioned at a later stage. COMPOUNDS OF LEAD Lead, like tin, forms two series of compounds, in which it functions as a divalent and as a quadrivalent element. In its divalent com- pounds, which are by far the more important, it is a base of moderate strength, and has very weak acidic properties ; in its tetravalent compounds it has scarcely any basic properties, the dioxide, PbO 2 , being distinctly acidic. The compounds containing divalent lead are usually called plumbic compounds; the compounds containing quadrivalent lead have no general designation. Oxides of Lead Five oxides of lead are known : Pb 2 O, PbO, Pb 2 3 , Pb 3 O 4 , and PbO 2 . Lead Suboxide, Pb 2 O, is obtained by heating lead oxalate in a current of carbon dioxide or nitrogen at the lowest temperature suffi- cient to effect decomposition (about 300) : It is also formed by the action of air on lead at temperatures below its melting-point. LEAD 487 Lead suboxide is a grayish-black powder, which is decomposed on heating into plumbic oxide, PbO, and lead. Acids act on it in an analogous way ; a plumbic salt goes into solution and lead remains : Plumbic Oxide (Litharge), PbO, is obtained as a yellowish-red amorphous powder by strongly heating lead in air, and is a by-pro- duct in the separation of lead from silver by cupellation (p. 420). It is also obta ned by heating the nitrate or carbonate, and also by heating any of the other oxides at a sufficiently high temperature in air. When lead oxide is heated above its melting-point, 835, and allowed to cool it solidifies to a crystalline mass known as litharge. There are at least two modifications of lead oxide, a yellow and a red, the latter being the stable form at room temperature. Lead oxide is slightly soluble in water, forming an alkaline solution which presumably contains lead hydroxide, Pb(OH) 2 . It dissolves in acids to form plumbic salts, and also dissolves on boiling with alkali hydroxides, owing to the formation of soluble plumbites, e.g. Na 2 PbO :! (see below). Lead Hydroxide, Pb(OH) 2 , is formed by the slow oxidation of lead in moist air, and also as a white precipitate, by adding an alkali hydroxide or ammonium hydroxide to the aqueous solution of a lead salt. The nature of the precipitate appears to depend on the condi- tions ; the hydrates 2PbO,H 2 O and 3PbO,H 2 O have been described. The basic character of lead hydroxide is shown by the fact that it dissolves to a small extent in water, forming an alkaline solution, and reacts with acids to form corresponding salts. Its acidic character is shown by its solubility in solutions of alkali hydroxides; the solutions contain plumbites, e.g. Na 2 PbO 2 : Pb(OH) 2 + 2 NaOH->Pb(ONa) 2 + 2H 2 O. Lead Sesquioxide, Pb 2 O 3 , is obtained as an orange-yellow powder by adding sodium hypochlorite to a solution of plumbic oxide in alkali. Acids decompose it into the monoxide and dioxide, the former of which dissolves to form plumbic salts ; and it may there- fore be regarded as a loose compound of the two oxides, PbO and PbO 2 (cf. red lead). Red Lead or Minium, Pb 3 O 4 , is obtained by prolonged heat- ing of lead monoxide in the air at 500. It is a bright scarlet powder, and is used as a pigment. When heated to 550 the dissocia- tion pressure of the oxygen is equal to the oxygen pressure in the 4 88 A TEXT-BOOK OF INORGANIC CHEMISTRY atmosphere, and the compound therefore changes to the monoxide. With acids the monoxide is dissolved out with formation of plumbic salts, and lead dioxide remains : so that the salt behaves as a mixture of the two oxides. It has, however, been shown that a definite compound of the formula Pb 3 O 4 (PbO 2 ,2PbO) exists, since the dissociation pressure of the oxygen in red lead is less than that of lead peroxide at the same temperature. The composition of the commercial article is not quite constant ; besides the compound, Pb 3 O 4 , it may contain excess of PbO or PbO 2 . Lead Peroxide, PbO 2 , is obtained, as just described, by the action of dilute nitric acid on red lead ; it is also obtained by the oxidation of lead monoxide, dissolved in alkali, with hypochlorite, and is deposited on the anode during the electrolysis of lead salts (cf. p. 491). It is an amorphous, dark-brown powder, which on heating to 350 gives up oxygen with formation of the lower oxides mentioned above ; when the temperature is sufficiently high the monoxide results. When heated with concentrated hydrochloric acid chlorine is given off: With concentrated sulphuric acid oxygen is given off : 2PbO 2 + 2H 2 SO 4 ->2PbSO 4 + 2H 2 O + O 2 . In both cases plumbic salts are formed. Lead peroxide shows both acidic and basjc characters. When boiled with a concentrated solution of potassium hydroxide it dis- solves, and from the solution, on cooling, potassium pluuibatc, K 2 PbO 3 ,3H 2 O, separates in colourless rhombohedral crystals, iso- morphous with potassium stannate trihydrate (p. 482). Salts corre- sponding with " orthoplumbic acid," Pb(OH) 4 , are also known, e.g. Ca 2 PbO 4 . This salt is obtained by heating in the air a mixture of calcium carbonate and lead oxide, oxygen being absorbed : 4CaCO 3 + 2PbO + O 2 ^2Ca 2 PbO 4 + 4CO 2 , The action is reversible, and upon this is based Kassnei j s method (no longer used) of obtaining oxygen from the air. The compound Pb 3 O 4 may be regarded as plumbic plumbate, Pb 2 PbO 4 . Lead Chloride, PbCl 2 , being only slightly soluble in water, separates when solutions containing Pb" and Cl' ions are brought LEAD 489 together. It forms colourless, lustrous rhombic crystals; at 20 100 grams of water dissolve about i.o gram, at 100 about 4 grams of the salt. The solubility is diminished by the addition of dilute hydro- chloric acid (and soluble chlorides), but is greater in the presence of concentrated hydrochloric acid. The lowering of solubility is due to the addition of a compound with a common ion (p. 424) ; the subse- quent increase to the formation of complex ions (p. 424). In the latter case, the solutions probably contain compounds of the type, HPbCl 3 , which ionise as follows: Lead Bromide, PbBr 2 , and Lead Iodide, PbI 2 , are prepared in a similar manner to the chloride; the former is colourless, the latter separates from aqueous solution in lustrous yellow crystals. At o 100 grams of water dissolves 0.044 grams, at 25 0.076 grams, and at 100 0.^36 grams of the iodide. All the lead halides form double salts with the alkali halides, ^>2PbCl 2 'KCl ; PbCl 2 '2KCl; PbI 2 'KI'2H 2 O. Solutions of these double salts contain the lead, in part at least, as a complex anion, e.g. PbCl 3 ' (see above). Lead Nitrate, Pb(NO 3 ) 2 , is obtained by dissolving lead, the oxide or carbonate in nitric acid ; on cooling it separates in anhydrous octahedral crystals. At o 100 grams of water dissolve 39 grams, at 20 56 grams of the salt. On heating it is decomposed into lead oxide, nitrogen peroxide and oxygen (p. 228). A number of basic nitrates of lead, e.g. PtyNOg^PbO and Pb(NO 3 ).'PbO, have been ' described. Lead Sulphate, PbSO 4 , occurs naturally in rhombic crystals as anglesite, and is formed when solutions containing Pb" and SO 4 " ions are mixed. It is a heavy white powder, practically insoluble in water, still less soluble in dilute sulphuric acid, but readily dissolves in concentrated sulphuric acid. The latter phenomenon is due to the formation of an acid sulphate, Pb(HSO 4 ) 2 ,H 2 O, which separates when the solution in'question is cautiously diluted. Several basic sulphates of lead have been described. Lead sulphate dissolves fairly readily in solutions of ammonium acetate and of the alkali acetates. In dilute solution this effect is due to the formation of slightly ionised lead acetate by double decomposi- tion and r consequent diminution in the Pb" ion concentration. In concentrated solution the formation of complex salts appears to come into account. Lead Carbonate, PbCO 3 , occurs naturally as cerussite. It is 490 A TEXT-BOOK OF INORGANIC CHEMISTRY obtained as a heavy white precipitate by adding a solution of lead nitrate or acetate to excess of a solution of commercial ammonium carbonate. When sodium or potassium carbonate is used, basic car- bonates of lead, _rPbCO 3x yPb(OH) 2 , are obtained, the composition of which depends upon the conditions of precipitation. The best known basic carbonate of lead, 2PbCO 3 ,Pb(OH) 2 , known as white lead, is largely used as a pigment. A number of methods for preparing it are in use, but the product of highest covering power is obtained by the so-called Dutch process, which has been used for centuries. The method depends upon the successive action of acetic acid and of carbon dioxide on the metal. The lead is cast in gratings, or sheet lead is twisted in spirals so as to expose a large surface, and the gratings or spirals are placed in earthware pots in such a way that they do not come into contact with the small quantity of vinegar (dilute acetic acid) which each pot contains. A row of pots is placed on dung or spent tan bark in a shed, and alternate rows of pots and layers of tan-bark or manure are piled up till the shed is nearly full. The whole arrangement is left for some weeks. The heat given out during the fermentation of the tan-bark vaporizes the acetic acid, which, along with oxygen, converts the lead to a basic acetate, Pb(C 2 H 3 O 2 ) 2 ,Pb(OH) 2 . The latter is then in turn acted upon by the carbon dioxide produced during the fermentation, the basic carbonate being formed. When the gratings are almost entirely converted to white lead, the latter is scraped off, ground up while wet, washed to remove lead acetate, and dried. White lead is also obtained by passing carbon dioxide into a solu- tion of basic lead acetate, or by rubbing lead oxide into a paste with lead acetate solution and then passing in carbon dioxide. Electrolytic methods have also been proposed. None of these methods gives a product equal to that obtained by the Dutch process. White lead has the disadvantage that it is poisonous, and it blackens when exposed to hydrogen sulphide, which is usually present in the atmosphere of towns, but these drawbacks are more than counter- balanced by its great covering power. In order to give good results it must be amorphous. Lead Sulphide, PbS, occurs naturally as galena, and is obtained as a black amorphous precipitate by passing hydrogen sulphide into a solution of a lead salt. As the sulphide is practically insoluble in water, hydrogen sulphide is used to detect minute amounts of lead in water. Lead sulphide is insoluble in dilute hydrochloric acid, but the concentrated acid converts it to lead chloride with evolution of LEAD 491 hydrogen sulphide. When lead sulphide is boiled with dilute nitric acid, lead nitrate is formed ; the concentrated acid converts it chiefly into the sulphate. When hydrogen sulphide is passed into a solution of lead chloride, a reddish precipitate is obtained which becomes black when excess of hydrogen sulphide is used. This phenomenon is due to the inter- mediate formation of one or more double compounds of lead sulphide and chloride. The only one which has been definitely isolated has the formula PbS'PbCl 2 ; it is a red powder. Lead Acetate, Pb(C 2 H 3 O 2 ) 2 ,3H 2 O, known as sugar of lead " on account of its sweetish taste, is obtained in monoclinic crystals by dissolving lead oxide in acetic acid and concentrating the solution. It is very soluble in water. With lead oxide it forms soluble basic salts ; two such compounds, Pb(C 2 H 3 O 2 ) 2 ,Pb(OH) 2 and Pb(C 2 H 3 O 2 ) 2 ,2Pb(OH) 2 , have been definitely isolated. Compounds of Quadrivalent Lead Mention has already been made of lead dioxide, P oO 2 , and of salts derived from the corresponding hydroxide, Pb(OH) 4 , functioning as orthoplumbic acid, and from its first anhydride, H^PbOg, meta- plumbic acid. The compounds derived from the hydroxide, Pb(OH) 4 , acting as a base will now be dealt with. Lead Totrachloride, PbCl 4> is formed when lead dioxide is dissolved in cold concentrated hydrochloric acid. It is most readily obtained by passing chlorine into water in which lead dichloride is suspended, and then adding ammonium chloride to the solution, when the compound, PbCl 4 ,2NH 4 Cl, separates in crystalline form. The latter compound is added to cooled concentrated sulphuric acid, whei the tetrachloride separates as a yellow, highly refractive liquid. Lead tetrachloride is a heavy yellow liquid of density 3.18 at o. With a small amount of water it forms a hydrate ; with excess of water it is completely hydro- lyzed to the dioxide and hydrochloric acid : PbCl 4 +4HOH->PbO 2 +4HCl + 2H 2 O. Lead Disulphate, Pb(SO 4 ) 2 , is obtained by the electrolysis of sulphuric acid of density 1.7 to 1.8 between electrodes of lead, the temperature not being allowed to rise above 30. Towards the end of the electrolysis the temperature is raised to 40 to 50, and on cooling the disulphate separates from the solution in the anode compartment in yellowish crystals. When pure it is colourless. Lead d sulphate is readily decomposed by water into the dioxide and sulphuric acid, a basic salt, PbOSO 4 ,H 2 O, being formed as intermediate product. It is a powerful oxidizing agent. Lead tetracetate, Pb(C 2 H 3 O 2 ) 4 , has also been obtained ; it occurs in colourless needles. The Lead Accumulator 1 The lead accumulator is a cell in which electrical energy is stored. It consists in its simplest form of two plate s, one of which when charged is coated with lead peroxide, the 1 Cf. Physical Chemistry, p. 377. 492 A TEXT-BOOK OF INORGANIC CHEMISTRY other with finely-divided metallic lead, and the plates dip into dilute sulphuric acid. When the two poles are connected by a wire, dis- charge takes place at a potential of two volts, and the potential remains practically constant for a long time. During the discharge, the lead dioxide becomes reduced to the oxide, which forms lead sulphate with the sulphuric acid ; at the other pole the lead is oxidized to lead sulphate. The chemical changes taking place when the cell is discharging are therefore represented by the equation : The essential feature of the accumulator is that it is readily rever- sible. When, after it is discharged, a current is sent through it in the opposite direction to the discharge current, the sulphate at one pole is oxidized to the dioxide, at the other pole it is reduced to metallic lead, the cell being thus restored to the same condition as before discharge. Tests for Lead Lead compounds are readily reduced on charcoal to the metal, which can be recognized by its softness and by its property of marking paper. The formation and behaviour of the sulphide, sulphate and iodide, already fully described, are characteristic. General Properties of the Carbon Sub-group and Summary Carbon, silicon, germanium, tin and lead constitute a natural family of elements and show the usual gradual change of physical and chemical properties with increase of atomic weight. The more important physical properties are summarized in the accompanying table : Carbon. Silicon. Germanium. Tin. Lead. Atomic weight 12.00 28.3 72.5 119.0 207.1 Density Melting-point BoHing-point Atomic volumes 2.25 to 3.6 very high 3000? 3-4 2.35 to 2.49 1600 II. 5-47 <9oo >iSoo 13.2 7.29 231-5 2270^ 16.6 11.34 334 1525 18.3 i As regards the chemical properties, all the elements are tetravalent, corresponding with their position in the periodic table, and the last three form well-defined divalent compounds. Carbon is also divalent in carbon monoxide. Further, the acidic character diminishes with increase of atomic weight. Carbon and silicon are typical non- ELEMENTS OF THE CARBON GROUP 493 metals ; no compounds are known containing carbon or silicon cations. The oxides of germanium, tin and lead are both basic and acidic, and this applies to the oxides in which the metals are divalent as well as to those in which they are quadrivalent. All the hydroxides of the type X(OH) 4 are, however, very weak bases (including H 4 PbO 4 ), and the halogen compounds of the type XCl 4 are immediately decom- posed by water. 1 In the case of the divalent compounds there is a definite increase of electro-positive character with increase of atomic weight. Stannous hydroxide, Sn(OH) 2 , is a fairly strong base, and stannous salts with strong acids are not greatly hydrolyzed in solution. Leadhyc roxide, Pb(OH) 2 , is a stronger base than stannous hydroxide, and plumbic salts with strong acids, e.g. lead nitrate, are scarcely hydrolyzed in aqueous solution at moderate dilution. In accordance with the gradation of properties in the group, only the first two elements form compounds with hydrogen. 1 Carbon tetrachloride, CC1 4 , is an apparent exception to this rule, as it is not affected by water in the cold. It is, however, completely decomposed by heating with water in a sealed tube, and the action is not reversible : CC1 4 + 2 H 2 O->CO 2 + 4 HC1. CHAPTER XXXII ELEMENTS OF THE NITROGEN GROUP (GROUP V) Sub-group A Sub-group B Vanadium, V. . . . 51.06 Nitrogen, N .... 14.01 Columbium, Cb (Niobium) . 93.5 Phosphorus, P 31.0 Tantalum, Ta 181.0 Arsenic, As .... 74.96 Antimony, Sb. . . . 120.2 Bismuth, Bi . . . . 208.0 THE three members of sub-group A are rare elements. Of the members of sub-group B, nitrogen and phosphorus have already been considered among the non-metals. This family presents a very striking illustration of the increase in electro-positive character with increase of atomic weight. Nitrogen and phosphorus are typical non-metals, arsenic behaves in practically all respects like a non- metal, antimony is intermediate in character, bismuth behaves in most respects like a metal. Throughout the group the principal valencies are three and five. THE VANADIUM SUB-GROUP Vanadium was discovered by del Rio in 1801 in vanadinite, sPbgfVO^'PbClg, isomorphous with pyromorphite (p. 484), which is found chiefly in Spain, Chili, and the Argentine. The metal was first obtained by Roscoe (1867) by heating the dichloride, VC1 2 , in a current of hydrogen, but was not pure. The pure metal was obtained for the first time (Weiss and Aichel, 1904) by the reduction of vanadium pentoxide by the thermite process (p. 466), a mixture of rare earth metals being used in place of aluminium. Methods for preparing vanadium by electrolysis have also been proposed. Vanadium is a white, lustrous, crystalline metal of density about 5.8 ; it is the hardest metal known, is fairly brittle, and melts about 1680. It is stable in the air at room temperature, but on heating burns to V 2 O 5 . Vanadium is remarkable for the numerous series of compounds derived from it. There are four well-defined oxides, VO (or V 2 O 2 ), V 2 O 3 , V 2 O 4 , and V 2 O 5 . These compounds are of analogous type to the oxides of nitrogen, but as regards the lower compounds there is very little resemblance in their chemical properties. VO is a strongly basic oxide, and well-defined salts corresponding with it are known. The sulphide, VSO 4 ,7H 2 O, forms reddish-violet crystals, isomorphous with ferrous sulphate heptahydrate, FeSO 4 ,7H 2 O. Vanadous chloride, VC1 2 , occurs in greenish hexagonal plates. The oxide V 2 O 3 is a base of moderate 494 ARSENIC 495 strength ; VC1 3 ,6H 2 O occurs in green crystals, V 2 (SO 4 ) 3 is a yellow powder (cf. compounds of trivalent iron). The oxide VO 2 is weakly basic and also weakly acidic. The tetrachloride, VC1 4 , is known, but most of the compounds containing quadrivalent vanadium are of the type VOX 2 ; that is, in solution divalent VO" ions are present. The pentoxide, V 2 O 5 , is acidic, and behaves in many respects like phosphorus pentoxide. The salts are derived from ortho- vanadic ac d, H 3 VO 4 , and metavanadic acid, HVO 3 ; the latter are the more stable. Arimonium metavanadate, NH 4 VO 3> is insoluble in ammonium chloride solution, ar.d advantage is taken of this in separating vanadium from its ore. Salts containing di- and trivalent vanadium ar# powerful reducing agents. Vanadium is now finding considerable application as a constituent of certain alloys. Columbian! and Tantalum occur together in the rare minerals columbite and tanta.ite, Fe[Cb(Ta)O 3 ] 2 . Both elements are always present ; when columbium is in excess it is termed columbite, with excess of tantalum tantalite. Niobium is a lustrous, fairly ductile metal which, when pure, is scarcely as hard as soft steel, but when traces of impurities are present, is nearly as hard as vanadium ; it melts about 1950. Metallic tantalum can be obtained by reduction of the dioxide, TaO 2 with carbon in the electric furnace ; the powder thus obtained is purified and obtained in coherent form by heating in the electric arc in a vacuum, when the oxide distils off, leaving the fused metal, which solidifies on cooling to a lustrous regulus. Pure tantalum is a lustrous metal of grayish colour, and is so ductile that it can be ber ten into foil and drawn out into wire ; it melts about 2275. At a white heat, it burns slowly in air to the pentoxide. As is well known, tantalum is now largely used as a filament for electric lamps. The pemoxides Cb 2 O 5 and Ta^O 5 are both basic and acidic. The salts derived from niobi : and tantalic acids are of several types ; their behaviour shows that the acids a -e weaker than vanadic acid. The remarkably close resemblance be- tween niobium and tantalum compounds is further shown by the fact that there is no definite difference in strength between niobic and tantalic acids. Niobiun and tantalum differ markedly from vanadium in showing very little tendency to form compounds of a low stage of oxidation. Only one compound of the type CbX 3 is known with certainty, and tantalum is exclusively quadrivalent and quinquevalent in its compounds. THE NITROGEN SUB-GROUP N= -14.01. P = 3i.o. As = 74.96. Sb= 120.2. Bi = 2o8.o. Of the members of this family only arsenic, antimony, and bismuth remain to be dealt with. ARSENIC Symbol, As. Atomic weight=75. Molecular weight 300. Occurrence Arsenic occurs free in nature, but is usually met with in combination. In association with oxygen it occurs as arseno- lite, As 2 O 3 , and with sulphur as realgar, As 2 S 2 , and orpiment, As 2 S 3 . 496 A TEXT-BOOK OF INORGANIC CHEMISTRY In combination with metals it occurs as arsenical iron, FeAs 2 , arsenical nickel or kupfernickel, NiAs, and cobalt speiss, CoAs 2 . In association with metals and sulphur it occurs in arsenical pyrites, FeSAs or FeS 2 ,FeAs 2 , in cobalt glance, CoSAs, and in other com- pounds. It occurs in most specimens of iron pyrites, and hence finds its way into commercial sulphuric acid and into products, such as phosphorus, into the preparation of which sulphuric acid enters. Preparation of Element (i) On the commercial scale it is obtained by sublimation of native arsenic or by heating arsenical pyrites in earthenware vessels in absence of air. The pyrites are decomposed into ferrous sulphide and arsenic, and the vapour of the latter is condensed in earthenware receivers : FeS 2 ,FeAs 2 ->2FeS + 2 As. It is purified by sublimation. (2) It is easily obtained by heating arsenious oxide with charcoal : Properties Arsenic exists in at least three allotropic modifica- tions. The best known form (metallic arsenic) occurs in rhombo- hedral crystals, which are steel-gray in colour, and have metallic lustre ; the density is 5.73. Metallic arsenic is brittle ; it is a good conductor of heat and electricity. It sublimes at 450, before the melting-point is reached, but can be melted under pressure. A second form is obtained as a black, apparently amorphous, but really minutely crystalline, powder by subliming ordinary arsenic ; its density is 4.72. At 300 it changes with evolution of heat into the stable metallic form. A third modification is obtained in light yellow crystals (density 2.026) by rapidly cooling arsenic vapour, e.g. by means of liquid air. This modification, like yellow phosphorus, is soluble in carbon disulphide, and smells strongly of garlic. Yellow arsenic changes to the metallic form at room temperature; the change is markedly accelerated by exposure to light. Other modifica- tions of arsenic have been described. 1 The vapour density of arsenic at 600 to 700 corresponds with the formula As 4 , but it gradually diminishes with increasing temperature, and at 1700 has fallen to half the original value; the formula is then As 2 . When heated in the air arsenic burns to the trioxide, As 2 O 3 . When 1 Cf. Erdmann and Reppert, Annalen, 1908, 361, T. ARSENIC 497 fragments of arsenic are thrown into chlorine or bromine at room tem- perature, a vigorous reaction occurs, and the trihalide (AsCl 3 ; AsBr 3 ) is formed. COMPOUNDS OF ARSENIC WITH HYDROGEN AND WITH THE HALOGENS Arsine (Arseniuretted Hydrogen), AsH 3 Preparation (i) Arsine is obtained, mixed with hydrogen, when a soluble compound of arsenic is added to a solution in which hydrogen is being generated, e.g. by the action of hydrochloric or sulphuric acid on zinc. It is also obtained when the hydrogen is generated electrolytically in a solution containing a compound of arsenic: (2) The pure compound is prepared by the action of dilute sulphuric acid on zinc or sodium arsenide : Zn 3 As 2 + 6HCl->3ZnCl 2 +-2AsH 3 . Sodium arsenide is easily obtained by heating to redness a mixture of metallic sodium and excess of arsenious oxide. Properties Arsine is a colourless gas with an odour of garlic, and is extremely poisonous. It can be condensed to a colourless liquid, which boils at 55. Arsine is an endothermic compound, and is readily decomposed into its elements by heat. This is readily shown by passing the mixture of arsine and hydrogen obtained in (i) through a hard glass tube and heating the latter at a definite point, when a mirror of arsenic will form in the tube just beyond the heated portion. Arsine burns in air with a vivid blue flame to the trioxide and water : 2 AsH 3 + 3O 2 ->As 2 O 3 + sH 2 0, but when the supply of air is insufficient it burns to arsenic and water : The reactions just mentioned form the basis of Marsh's test for the detection of arsenic. The apparatus used is shown in Fig. 88. Hydrogen is generated in the Woulf's bottle from zinc and dilute sulphuric acid (which must themselves be arsenic-free). When all the air has been displaced from the apparatus, the hydrogen is lighted 498 A TEXT-BOOK OF INORGANIC CHEMISTRY at the end of the tube, which is drawn out to a jet, and the arsenic solution added through the funnel. The hydrogen flame soon becomes livid blue, and when a white porcelain dish is held in the flame black spots of metallic arsenic are formed. Under the same circumstances antimony gives similar spots, but, unlike the arsenic spots, they are insoluble in a fresh solution of bleaching FIG. 88. powder. The decomposition of arsine by heat, already referred to, may be demonstrated by heating the tube as shown. Arsine is a powerful reducing agent ; it readily precipitates gold and silver from solution : From a solution of copper sulphate, copper arsenide, Cu 3 As 2 , is precipitated. The Nascent State In the previous section it has been pointed out that when hydrogen is generated in a solution of arsenious oxide, the latter is reduced to arseniuretted hydrogen. When, however. gaseous hydrogen (from a Kipp's apparatus, for example) is passed through the solution, no reduction of arsenious oxide occurs. Hydrogen generated in the solution the so-called " nascent " hydrogen there- fore appears to be a much more powerful reducing agent than ordinary hydrogen. The usual explanation of this phenomenon is that the hydrogen at the moment of liberation is in the atomic condition, and is therefore much more active chemically. The energy relations of the systems throw some light on this ques- ARSENIC 499 tion. In the reaction under consideration the total heat change is made up of two parts : (a) the heat of solution of zinc in acid, which "is positive and large in amount ; (b) the heat given out when arsenious oxide is reduced by hydrogen, which is negative although small. The whole reaction is therefore exothermic, whereas the reduction of arsenious oxide by free hydrogen is endothermic ; and it is there- fore to be anticipated that, quite apart from any question as to atomic hydrogen, the former reaction will proceed much more readily than the latter (p. 148). It is highly improbable, however, that the energy of one reaction can be transferred to another entirely independent reac- tion proceeding in the same system, and we must therefore assume that the reactions (a) and (fr) are connected in some way perhaps through the agency of an intermediate compound in order that the energy given out in reaction () may become available for the whole change. Apart from the question of the total amount of energy, the effect of catalytic agents on the speed of reaction in the present case, for instance, the effect of the zinc on the activity of the hydrogen is also of importance; but the matter cannot be further discussed at the present stage. Arsenic Halides The following compounds, AsF 3 , AsCl 3 ,AsBr 3 and As I, are definitely known. The only halide containing penta- valent arsenic which is definitely known is the pentafluoride AsF 5 . Arsenic Fluoride, AsF 3 , is formed by direct combination of its elements, but is most readily obtained by distilling a mixture of arsenious oxide, potassium fluoride, and excess of concentrated sulphuric acid from a lead retort: As 2 O 3 + 6H F->2 AsF 3 + sH 2 O. Properties Arsenic fluoride is a colourless, fuming liquid, which boils at 63, and is at once decomposed by water into the trioxide and hydrofluoric acid. On account of its great activity it is used in preparing other fluorides. It is best kept in platinum bottles. Arsenic Chloride, AsC] 3 , is formed by heating arsenic in a current of chlorine, but is most easily obtained by distilling a mixture of arsenious oxide, sodium chloride, and concentrated sulphuric acid : Properties Arsenic chloride is a colourless, fuming liquid, which boils at 130, and is very poisonous. It is decomposed reversibly by water, with formation of the oxide and hydrogen chloride: 500 A TEXT-BOOK OF INORGANIC CHEMISTRY With a small quantity of water an oxychloride, AsOCl,H 2 O, or As(OH) 2 Cl, is obtained as a white precipitate. Arsenic Bromide, AsBr 3 , and Arsenic Iodide, AsI 3 , are most easily obtained by adding powdered arsenic to a solution of the halogen in carbon disulphide. The bromide occurs in colourless prismatic crystals, which melt at 31 ; the liquid boils at 221. The iodide forms red, lustrous, hexagonal crystals, which melt at 146. OXIDES AND OXYACIDS OF ARSENIC Two oxides of arsenic are known, both of which are acidic : Arsenious oxide . . . As 2 O 3 (or As 4 O G ) Arsenic oxide .... As 2 O 5 No acid corresponding with arsenious oxide has been isolated, but salts derived from orthoarsenious acid, H 3 AsO 3 , and from metarsenious acid, HAsO 2 , are known. From arsenic oxide three acids, of the same type as the phosphoric acids, are derived : Orthoarsenic acid H 3 AsO 4 Pyroarsenic acid H 4 As 2 O 7 Metarsenic acid HAsO 3 Arsenious Oxide, As 2 O 3 This compound, also known as " white arsenic" or "arsenic," is the most important compound of arsenic, and was familiar to chemists in the Middle Ages. It occurs naturally as arsenolite, and is formed when the element burns in air or oxygen. It is obtained commercially in the roasting of arsenical pyrites (iron oxide remaining behind) and of other pyrites containing arsenic (often as secondary product), the vapours being condensed in long tubes. It is purified by sublimation from iron vessels connected with cylindrical receivers, in which it condenses in the vitreous form. Properties Arsenious oxide exists in three allotropic modifications, two of which are crystalline and one amorphous. (1) The octahedral modification (regular system). This modifica- tion, which is stable at room temperature, is obtained by rapidly- cooling the vapour of the oxide, or (in well-formed crystals) by allow- ing a solution of the oxide in hydrochloric acid to crystallize. The density is about 3.6. At 15 100 grams of water dissolve about 1.65 grams, at 25 about 2.04 grams of this form, but the rate of solution is extremely slow. (2) The prismatic form is obtained by allowing a solution of the ARSENIC 501 oxide in potassium hydroxide to evaporate slowly in the absence of nuclei of the octahedral form. This form is metastable at room temperature; its density is about 4.15. (3) The amorphous modification is obtained as a colourless, trans- parent, glassy mass by slow condensation of the vapour of the oxide at a relatively high temperature. It is unstable at room temperature, and slowly becomes opaque owing to its transformation to the octahedral form. As the change proceeds from without inwards, masses of oxide are often met with which consist of the octahedral form on the outside with a core of the vitreous form. In accordance with the general rule (p. 241), the vitreous form is the more soluble in water, but the exact solubility cannot be determined owing to the fact that the transformation is accelerated by water. The density of the vitreous form is about 3.71. Arsenious oxide passes directly into vapour on heating, but can be melted under pressure. The density of the vapour up to 1500 corre- sponds with the formula As 4 O G , at 1800 with the formula As 2 O 3 . The aqueous solution of arsenious oxide is slightly acid, indicating the presence of arsenious acid, H 3 AsO 3 (or perhaps HAsO 2 ). Arseric trioxide is a powerful poison. It is remarkable that when it is taken in gradually increasing doses, the human organism acquires increased resistance to its action, so that quantities may ultimately be taken without danger which would at once prove fatal to one unused to it. It is employed as a rat poison, in calico printing, etc., and is also used in medicine. Arsenious Acid and Arsenites Whether the acidity of a solutior of arsenious oxide in water is due to the presence of orthoarsenious acid, H 3 AsO 3 , or of meta-arsenious acid, HAsO 2 , has not been definitely established. Arsenious acid has not been obtained in the free condition ; on evaporating the solution the anhydride separates. It is a very weak acid, about the same strength as boric acid (p. 360). The salts of arsenious acid, the arsenites, are derived from the ortho acid, H 3 AsO 3 , the meta acid, HAsO 2 , the pyro acid, H 4 As 2 O 5 , and still more complex acids. The most important are silver arsenite, Ag 3 AsO 3 , which is yellow, and copper hydrogen arsenite or Scheele's green, CuHAsO 3 . The latter is used as a pigment. Arsenic Pentoxide, As a O 5 , is obtained by heating the tri- oxide with nitric acid, and dehydrating the arsenic acid thus obtained by heating to low redness : 502 A TEXT-BOOK OF INORGANIC CHEMISTRY Properties Arsenic pentoxide is a white, amorphous, deliquescent solid, which dissolves in water to form arsenic acid. On heating strongly, it decomposes into the trioxide and oxygen. Arsenic Acid, H 3 AsO 4 , and Arsenates Arsenic acid is most readily obtained by heating arsenious acid with nitric acid ; on evaporating the solution to dryness, and recrystallizing the resi- due from water, colourless crystals of the composition 2H 3 AsO 4 ,H 2 O are obtained. At 100, the water of crystallization is driven off and the anhydrous acid remains. At 140 to 1 80 two molecules of acid lose a molecule of water and pyroarsenic acid is obtained ; at 200 still more water is expelled and mcta-arsenic acid results ; finally, at a low red heat all the water is driven off and the pentoxide remains. The analogy with the phosphoric acids is obvious (cf. p. 251). However, pyro- and meta-arsenic acids, unlike the corresponding phosphorus compounds, change immediately to the ortho acid when dissolved in water, and, further, meta-arsenic acid, unlike meta- phosphoric acid, can be dehydrated by heat. The arsenates, derived from the three acids just mentioned, closely resemble the corresponding phosphates ; for example, phosphates are in many cases isomorphous with the corresponding arsenates. The rneta- and the pyro-arsenates change immediately to ortho- arsenates on dissolving in water. Ammonium magnesium arsenate, MgNH 4 AsO 4 , like the corresponding phosphate, is insoluble in water, and advantage is taken of this in estimating arsenic in solution. Silver arsenate, Ag 3 AsO 4 , is chocolate in colour, where- as silver phosphate is yellow, and this test serves to distinguish phosphates and arsenates. COMPOUNDS OF ARSENIC AND SULPHUR Three sulphides of arsenic are known : Arsenic disulphide (realgar) . . . As 2 S 2 Arsenic trisulphide (orpiment) . . . As 2 S 3 Arsenic pentasulphide .... As 2 S 5 Arsenic Disulphide, As 2 S 2 , occurs naturally as realgar, and is also obtained by fusing together arsenic and sulphur in the calculated proportions. On the large scale it is obtained (con- taminated with trioxide) by heating together iron pyrites and arsenical pyrites : ARSENIC 503 Properties Arsenic disulphide is a red crystalline solid of density 3.5 ; it was formerly used as a paint. A mixture of realgar, sulphur and potassium nitrate is used in pyrotechny under the name of Bengal fire or Greek white fire. It burns in air to the trioxide and sulphur dioxide. Arsenic Trisulphide, As 2 S 3 , occurs naturally as orpiment in yellow monoclinic crystals. It is obtained by fusing a mixture of its components in the calculated proportions/and is formed as a yellow precipitate when hydrogen sulphide is passed through an acidified solution of an arsenic compound, e.g. As 2 O 3 : Properties When freshly precipitated from solution, arsenic sulphide is amorphous, and dissolves to a considerable extent in cold water, forming a colloidal solution. When hydrogen sulphide is passed into an aqueous solution of the trioxide in the absence of acid, the solution becomes yellow but no precipitation occurs, the sulphide remaining in colloidal solution. It is at once pre- cipitated by the addition of acids or salts (cf. p. 355). Arsenic trisulphide is soluble in colourless ammonium sulphide solution, with formation of ammonium thioarsenite : As 2 S 3 + 3 (NH 4 ) 2 S->2(NH 4 ) 3 AsS 3 . It is also soluble in alkali hydroxide, a mixture of arsenite and thio- arsenite being formed : The thioarsenites, which may be regarded as analogues of the arsenites with the oxygen replaced by sulphur, are decomposed on addition of hydrochloric acid, the trisulphide being precipitated : If instead of colourless ammonium sulphide the yellow sulphide is used to dissolve arsenic trisulphide, the solution contains ammonium thioar senate, (NH 4 ) 3 AsS 4 , the analogue of ammonium arsenate : As 2 S 3 + 3(NH 4 ) 2 S + 2S->2(NH 4 ) 3 AsS 4 . On adding excess of hydrochloric acid to the solution, arsenic sulphidt is precipitated : So 4 A TEXT-BOOK OF INORGANIC CHEMISTRY Neither thioarsenious acid, H 3 AsS 3 , nor thioarsenic acid, H 3 AsS 4 , is known in the free condition. It is evident that the sulphides As 2 S 3 and As 2 S 5 bear the same relationship to these hypothetical acids as As 2 O 3 and As 2 O 5 do to arsenious and arsenic acids. Arsenic Pentasulphide, As 2 S 6 , is prepared as described above or by heating a mixture of the elements in the theoretical proportions. It is a light yellow solid which can be fused and sub- limed without apparent decomposition. It dissolves in ammonium sulphide to form ammonium thioarsenate, and in a solution of potassium hydroxide to form a mixture of potassium arsenate and thioarsenate : 4 As 2 S 6 + 24KO H-^3K 3 AsO 4 -H 5 K 3 As S 4 + 1 2 H 2 O. Tests for Arsenic The more important tests have been mentioned in connexion with the different compounds. They will be further referred to in connexion with antimony, which resembles arsenic very closely in many respects. ANTIMONY Symbol, Sb. Atomic weight, 120.2. Molecular weight, 120.2 (at 1800). History The chief ore of antimony, stibnite, Sb 2 S 3 , has been in use from very early times ; it is mentioned in the Old Testament as being employed by women for darkening their eyebrows. A vessel used by the ancient Chaldeans has been shown to be made of metallic antimony (Berthelot). Stibnite was called stibium by the Romans, hence the symbol for the element still in use. In the Middle Ages, compounds of antimony were largely used for medicinal purposes. Occurrence Rich deposits of metallic antimony of a fairly high degree of purity have recently been found in Australia. In combina- tion with oxygen, as Sb 2 O 3 , it constitutes the mineral senarmontite, and as Sb 2 O 4 it constitutes antimony ochre. The most important ore, however, and the one from which the metal is chiefly obtained, is stibnite or gray antimony ore, Sb 2 S 3 . A large number of ores contain- ing antimony, sulphur, and either lead, copper, silver or iron are also met with. Preparation of Metal The first step in preparing antimony from stibnite is to heat the ore in an earthenware pot, when the sulphide fuses and is drained away (through holes in the bottom of the pot) from the accompanying impurities. From this partially purified sulphide the metal is prepared by one of two processes. ANTIMONY 505 (1) The sulphide is heated with scrap iron, whereby ferrous sulphide is formed and floats on the top, the antimony settling out below. Sb 2 S 3 + 3Fe->3FeS + 2 Sb. (2) According to the second method the ore, mixed with charcoal to prevent it from caking, is roasted to convert it to the oxide (the tetroxide, Sb 2 O 4 , is the chief product) ; it is then mixed with a further quantity of charcoal and strongly heated, whereby the metal is liberated : Sb 2 S 3 + 5 O 2 ->Sb 2 O 4 + sSO 2 The antimony thus obtained contains copper, iron, arsenic and other impurities. It can be purified by fusing with a mixture of antimony sulphide and sodium carbonate. The sulphides of the other metals, formed by double decomposition, rise to the surface and can be removed. Properties Antimony exists in at least three different allotropic modifications ; (a) metallic antimony, the ordinary stable form, which occurs in hexagonal crystals ; (b) explosive antimony ; (c) yellow antimony. (a) Metallic antimony (the ordinary form) is a silver- white, lustrous, brittle metal of density 6.5 ; it melts at 630.6, and boils about 1300. It conducts heat and electricity, but is a much less efficient conductor than the typical metals. (b) Explosive antimony is obtained as a black shining substance on the cadiode when a solution of antimony trichloride in hydrochloric acid is electrolyzed with an antimony anode and platinum cathode. Its density is 5.78. It is unstable under ordinary conditions, and ex- plodes vigorously when scratched with a needle or rubbed in a mortar, with formation of ordinary antimony. It always contains a greater or less amount of adsorbed antimony trichloride. Whether explosive antimony is crystalline or amorphous has not been established. (c) Yellow antimony is obtained by the action of oxygen on liquid antimony hydride, SbH 3 , at 90 : Like the corresponding modifications of phosphorus and arsenic, it is soluble in carbon disulphide. It is very unstable, and slowly changes to the metallic modification even at 90. The vapour density of antimony at 1400 exceeds that required by 506 A TEXT-BOOK OF INORGANIC CHEMISTRY the formula Sb 2 , but at 1800 it is about 60, indicating that the metal under these conditions is monatomic. Antimony is stable in the air, but on heating burns to the trioxide. It combines directly with chlorine and bromine at room temperature. It is slowly dissolved on boiling with concentrated hydrochloric acid, with evolution of hydrogen. When heated with concentrated sul- phuric acid, antimony sulphate, Sb 2 (SO 4 ) 3 , is formed and sulphur dioxide given off. Dilute nitric acid oxidizes antimony chiefly to the trioxide, Sb 2 O 3 ; with concentrated nitric acid the pentoxide, Sb 2 O 5 (or one of the corresponding acids), is the chief product. Alloys The most important alloys of antimony are type metal (which contains antimony, tin and lead in varying proportions) and Britannia metal (which contains tin, antimony and copper). Antimony expands on solidification, and therefore the alloys containing it, when used for casting, give sharp impressions. Antimony Hydride (Stibine) SbH 3 The methods of prepara- tion and properties of this compound are very similar to those of arsine (p. 497). It is formed (i) when a soluble antimony compound is added to a mixture of zinc and dilute sulphuric acid ; (2) in much purer condition by the action of dilute hydrochloric acid on magnesium antimonide. Properties Antimony hydride is a colourless gas with a character- istic musty odour. The liquefied gas boils at - 18, the solid melts at -91. It is an endothermic compound, and slowly decomposes into its elements even at room temperature ; the decomposition is greatly accelerated by a mirror of antimony deposited on the tube. It burns in air with a livid-blue flame to the trioxide and water, but when the supply of air is insufficient, water and antimony are the first products. When therefore a porcelain dish is held in the flame, a black spot of antimony is obtained which, unlike the corresponding arsenic spot (p. 498), is insoluble in bleaching powder solution. When passed into a solution of silver nitrate, silver antimonide, SbAg 3 , is precipitated : Compounds of Antimony with the Halogens The following halides of antimony are known : SbF 3 SbCl 3 SbBr 3 SbI 3 SbF 5 SbCl 5 Antimony Trifluoride, SbF 3 , is obtained by dissolving the trioxide in hydrofluoric acid, and separates in colourless, transparent, ANTIMONY 507 rhombic crystals when the solution is evaporated. It does not fume in moist an - , and, unlike the other trihalides, forms a clear solution with water. Antimony PentafLuoride, SbF 5 , is best prepared by boiling antimony pentachloride with anhydrous hydrofluoric acid for some days ; the residue is then subjected to fractional distillation : The pentachloride is a colourless, thick, oily liquid which boils at 149-150, and forms a clear solution with water. Both the trifluoride and pentafluoride form complex compounds with the alkali fluorides, e.g. SbF 3 ,KF ; SbF 3 ,2KF ; SbF 5 ,KF ; SbF 6 ,2KF,:iH 2 O. Antimony Trichloride, SbCl 3 , is obtained when chlorine is passed over metallic antimony, but is most readily obtained by dis- solving antimony oxide or sulphide in concentrated hydrochloric acid : Properties Antimony trichloride forms a soft crystalline mass, which, in allusion to its consistency, is known as "butter of antimony." It melts at 73.2 to an oily liquid, which boils at 223. With a small quantity of water a clear solution is obtained, but when more water is added hydrolysis occurs, and one or both of the oxychlorides SbOCl and (SbOCl) 2 ,Sb 2 O 3 are obtained, depending upon the conditions : + 5H 2 O^(SbOCl) 2 ,Sb 2 O 3 +ioHCl. The first is the main reaction in cold, the second in hot solutions. By boiling with successive quantities of water complete hydrolysis can be effected : (SbOCl) 2 ,Sb 2 O 3 +H 2 O->2Sb 2 O 3 -f2HCl. AntimDiiy Pentachloride, SbCl 6 , is obtained when antimony burns in excess of chlorine, but is best prepared by passing chlorine into the trichloride : Properties Antimony pentachloride is a colourless (as commonly met with often yellowish) fuming liquid, which can be obtained in 5o8 A TEXT-BOOK OF INORGANIC CHEMISTRY crystals melting at - 6. When heated, it boils about 140, the vapour being partly dissociated into the trichloride and chlorine : Under reduced pressure it can be distilled without decomposition. With small quantities of water it forms two crystalline hydrates, SbCl 6 ,H 2 O and SbCl 6 ,4H 2 O. When more water is added, it undergoes hydrolysis, oxychlorides, probably SbOCl 3 and SbO 2 Cl, being precipi- tated. With excess of hot water it is completely decomposed, with formation of hydrochloric and antimonic acids (cf. PC1 5 , p. 1246) : SbCl 5 + 4H 2 O->SbO(OH) 3 + sHCl. Antimony pentachloride forms a large number of crystalline " addi- tion compounds" with the chlorides of other elements, and with organic compounds, e.g. SbCl 5 ,KCl,H 2 O ; SbCl;3,NH 4 Cl,H 2 O ; SbCl 6 ,FeCl 3 ,8H 2 O ; SbCl 6 '3HCN. Antimony Tribromide, SbBr 3 , and the Tri-iodide, SbI 3 , are prepared, like the corresponding arsenic compounds, by adding powdered antimony to bromine or iodine dissolved in carbon disul- phide. The bromide forms colourless, rhombic crystals, which melt at 90 to 94 ; the liquid boils at 275. The iodide occurs in three allotropic modifications, the most stable forming ruby-red crystals. Both halides are hydrolyzed by water. OXIDES AND OXYACIDS OF ANTIMONY Three oxides of antimony are known: Antimony trioxide (antimonious oxide) . Sb 2 O 3 (or Sb 4 O 6 ). Antimony tetroxide ..... Sb 2 O 4 . Antimony pentoxide .... Sb^Og. From the trioxide is derived Sb(OH) 3 , or H 3 SbO 3 , which has both acidic and basic properties. ' The acidic properties are shown in the existence of the compound, NaSbO 2 ,3H 2 O (octahedral crystals) which, however, is derived from metantimonious acid, HSbO 2 . The basic character of the hydroxide is referred to below. From the pentoxide three acids are derived : Orthoantimonic acid . . H 3 SbO 4 . Pyroantimonic acid . . H 4 Sb 2 O 7 . Metantimonic acid . . . HSbO 3 . ANTIMONY 509 Antimony Trioxide, Sb 2 O 3 , is obtained, mixed with the tetroxide, when antimony is burned in the air ; it is best obtained by boiling the trichloride with a dilute solution of sodium carbonate : Properties Antimony trioxide is a colourless substance which exists in two crystalline modifications (a) a rhombic form, density 5.6; (b] an octahedral form, density 5.3 (cf. arsenic trioxide). It is practically insoluble in water, and in dilute nitric or sulphuric acid, but dissolves in hydrochloric acid to form the trichloride, and also in a solution of potassium hydrogen tartrate, KHC 4 H 4 O G , to form potassium antimony tartrate (tartar emetic) : Sb 2 O 3 + 2KHC 4 H 4 O 6 ^2K(SbO)C 4 H 4 O 6 +H 2 O. It dissolves in a concentrated solution of alkali hydroxide to form an antimonite (see above). The univalent SbO group, which is met with in antimony oxychloride, SbOCl, and in tartar emetic, is further referred to under antimonious hydroxide (y.v.'). Antimony Tetroxide, Sb 2 O 4 , is obtained by heating the metal, the trioxide, or the pentoxide in air at a high temperature. Properties Antimony tetroxide is a white powder, insoluble in water. It cannot be fused or volatilized; when heated to 1000 it decomposes into the trioxide and oxygen. In many respects it behaves like a mixture of the trioxide and pentoxide, and may be regarded as antimonious antimonate, Sb 2 O 3 'Sb 2 O 5 . Antimony Pentoxide, Sb 2 O 5 , is obtained by the action of nitric acid on metallic antimony, and heating the antimonic acid thus obtained to a temperature not exceeding 300. Proper/us Antimony pentoxide is a yellow, infusible powder, insoluble in water. When heated to 440 it decomposes into the tetroxide and oxygen. It has weakly acidic properties, but is devoid of basic properties. Antimonic Acids Three antimonic acids, orthoantimonic acid, H ;J ^bO 4 , pyroantimonic acid, H 4 Sb 2 O 7 , and metantimonic acid, HSbO 3 , are known, corresponding with the phosphoric and arsenic acids. The ortho acid is obtained by acting on antimony with nitric acid, and heating the product at 100. ' When orthoantimonic acid is heated at 200 for some time, it loses a molecule of water, and the pyro acid is formed ; at a slightly higher temperature more water is 5io A TEXT-BOOK OF INORGANIC CHEMISTRY driven off and the me fa acid, HSbO 3 , results. The three acids are white powders, and show no differences in chemical behaviour. The antimonates are mainly derived from metantimonic acid, HSbO 3 , but some salts of pyroantimonic acid are also known. When antimony is fused with potassium nitrate, potassium met- antimonate, KSbO 3 , is formed as a white powder. When this salt is fused with potassium hydroxide, potassium pyroantimonate is formed : When boiled with water, potassium pyroantimonate is decomposed into free alkali and dipotassium pyroantimonate, K 2 H 2 Sb 2 O 7 : K 4 Sb 2 7 + 2H 2 0->K 2 H 2 Sb 2 7 + 2KOH. The latter salt is fairly soluble in water and is used as a test for sodium ; its use for this purpose is based upon the fact that the corresponding sodium salt, Na 2 H 2 Sb 2 O 7 ,6H 2 O, is only slightly soluble in water. Antimonious Hydroxide, Sb(OH) 3 , and its Salts Refer- ence has already been made to the acidic properties of the hydroxide, Sb(OH) 3 . Its basic character is shown by the existence of salts derived either from the acid itself, e.g. Sb(NO 3 ) 3 ; Sb 2 (SO 4 ) 3 , or from its first anhydride, 1 SbO(OH), e.g. SbOCl. Antimony sulphate ', Sb 2 (SO 4 ) 3 , is obtained by dissolving the trioxide in concentrated sulphuric acid, and separates from solution in long, colourless, lustrous needles. It is decomposed by water with formation of basic salts. The nitrate, Sb(NO 3 ) 3 , obtained by dissolving the trioxide in cold, fuming nitric acid, also occurs in colourless crystals. On heating gently, antimony pentoxide is formed. It is hydrolyzed by water. The existence of these two salts is conclusive proof of the basic character of antimony in its trivalent compounds. Antimony Sulphides Two sulphides of antimony are known : the trisulphide, Sb 2 S 3 , and the pcntasulphide, Sb 2 S 5 . Antimony trisulphide, Sb 2 S 3 , is obtained as a grayish-black mass by fusing together sulphur and antimony in absence of air, and as an orange- red precipitate on passing hydrogen sulphide into a solution of an antimony salt. It dissolves in concentrated hydrochloric acid to form the trichloride, is soluble in colourless ammonium sulphide solution to form ammonium thioantimonite, (NH 4 ) 3 SbS 3 , and dissolves in 1 The univalent group -Sb^O is sometimes termed the antimony I group. BISMUTH 511 yellow ammonium sulphide solution to form ammonium thioanti- monate, (NH 4 ) 3 SbS 4 (cf. arsenic sulphides, p. 503). Antimony Pentasulphide, Sb 2 S r> , is obtained on adding excess of an acid to a solution of a thioantimonate, e.g. (NH 4 ) 3 SbS 4 : It is a dark orange-red, amorphous powder ; on heating to 200 it decomposes into the trisulphide and sulphur. Thioant monious acid, H 3 SbS 3 , and thioantimonic acid, H 3 SbS 4 , are not known in the free condition, but a number of thioanti- monites aid thioantimonates are known. Sodium thioantimonate, Na 3 SbS 4 ,9H 2 O, which forms colourless tetrahedral crystals, is known as Schlippe's salt. Tests for Antimony As already mentioned, there is a resem- blance between the behaviour of arsenic and of antimony in their compounds. Both give Marsh's test, but the arsenic spot is soluble, the antimony spot insoluble in solution of bleaching powder. Anti- mony trisulphide is soluble, arsenic trisulphide practically insoluble in concentrated hydrochloric acid. When passed into solution of silver nitrate, both arsenic and antimony hydride produce black precipitates (p. 506), but whereas the arsenic remains in the solution as arsenious acid, the antimony is precipitated as silver antimonide, SbAg 3 . BISMUTH Symbol, Bi. Atomic weight =208.0. Molecular weight 208.0. Occurrence Bismuth is not a very abundant element, but is fairly widely distributed in nature. It occurs chiefly in the free condition, along with granite and with cobalt and silver ores. In the combi.ied state it is found as the sulphide, Bi 2 S 3 , bismuthite or bismuth glance, as the oxide, Bi 2 O 3 , bismuth ochre, and in many other forms. Preparation Formerly the metal was obtained by heating the ore containing it in inclined pipes when, owing to its low melting- point, the metal drained away, leaving" the impurities behind. Another method new largely used is to roast the ore and heat with coal and a flux. The impurities rise to the surface, and the fused bismuth can be separated. In order to purify the metal thus obtained, it is fused with potassium nitrate and sodium chloride, whereby the impurities are oxidized. A more thorough purification is effected by dissolving the pure metal 5i2 A TEXT-BOOK OF INORGANIC CHEMISTRY in nitric acid to form the nitrate, Bi(NO 3 ) 3 , and then adding water to the solution, whereby the oxynitrate, BiONO 3 , is precipitated. The oxynitrate is then heated to convert it to the oxide, and the latter reduced with carbon. Properties Bismuth is a white lustrous metal with a reddish tinge. Its density is 9.78 at 20 ; it melts at about 264 and boils at 1420. It is brittle and a rather inferior conductor of heat and electricity. It crystallizes in rhombohedral crystals belonging to the hexagonal system. Bismuth is stable in the air, but on heating burns with a bluish flame, forming the trioxide. It is scarcely affected by hydrochloric acid unless oxygen is present, but dissolves in hot sulphuric acid ; in nitric acid and in aqua regia it dissolves readily at room temperature. In each case compounds of trivalent bismuth are formed. Bismuth is used as a constituent of alloys, many of which fuse at very low temperatures. Rose's metal contains bismuth 2 parts, lead i part and tin i part, and melts at 94 ; Wood's metal, which consists of bismuth 4 parts, lead 2 parts, tin i part, cadmium i part, melts at 61. Bismuth, like antimony, expands on solidification, and therefore the alloys containing it give sharp castings. COMPOUNDS OF BISMUTH WITH THE HALOGENS Only the following compounds, containing trivalent bismuth, are definitely known : BiF 3 BiCl 3 BiBr 3 BiI 3 . Compounds of the type BiX and BiX 2 (for example BiCl and BiCl 2 ) have also been described, but the evidence regarding their existence is contradictory. Bismuth Trifluoride, BiF 3 , is obtained by dissolving bismuth trioxide in hydrofluoric acid. On evaporating the solution, it i& obtained as a grayish-white crystalline powder, which is practically insoluble in, and is not attacked by, water. Bismuth Trichloride, BiCl 3 , is obtained in- the anhydrous condition by passing dry chlorine over bismuth gently heated in a retort (cf. PC1 3 , p. 245) ; it is obtained in solution by dissolving the metal in aqua regia or the trioxide in hydrochloric acid. Properties Bismuth trichloride is generally met with in colourless deliquescent crystals which melt at 225-230 ; the liquid boils at BISMUTH 513 447. It is hydrolyzed by water, bismuth oxychloride, BiOCl, being precipitated : BiCl 3 +H 2 O^BiOCl + 2 Bismuth chloride forms a large number of complex compounds with alkali halides, e.g. NaBiCl 4 ,3H 2 O or BiCl 3 ,NaCl,3H 2 O ; K 2 BiCl 5 ,2H 2 O and (NH 4 ),BiCl 5 ,2H 2 O. Bismuth Tribromide, BiBr 3 , is prepared by slowly adding powdered bismuth to bromine, allowing to stand for several days and then distilling. It forms yellow, lustrous crystals and is decomposed by water with formation of the oxybromide BiOBr. Bismuth Triiodide, BiI 3 , can be obtained by heating the elements together in an atmosphere of carbon dioxide, but is best prepared by heating a mixture of bismuth trisulphide and iodine and then subliming the triiodide. The compound prepared by sublimation forms lustrous, black leaflets, and is decomposed by water with forma- tion of oxyiodide, BiOI. OXIDES OF BISMUTH Only the oxides BiO and Bi 2 O 3 are known with certainty. Mixtures containing a higher proportion of oxygen can also be obtained, but no definite compound has been isolated. Bismuthous Oxide, BiO, is obtained by heating the basic oxalate, (BiO) 2 C 2 O 4 , in a current of carbon dioxide: (BiO) 2 C 2 O 4 ->2Bi Properties Bismuthous oxide is a black powder, which when heated in air combines with oxygen to form the trioxide. When heated to 350 in a current of carbon dioxide it decomposes into bismuth trioxide and metallic bismuth. It acts as a reducing agent. Bismuth Trioxide, Bi 2 O 3 , is formed when the metal is burned in air, but is most conveniently prepared by heating the hydroxide, the basic nitrate or carbonate : 2BiONO 3 ->Bi 2 O 3 (BiO) 2 CO 3 -^Bi 2 O Properties Bismuth trioxide occurs as a yellow, amorphous powder, or in yelkm, lustrous crystals. It is stable in the air. It dissolves in excess of acids to form salts which contain Bi ions: 33 514 A TEXT-BOOK OF INORGANIC CHEMISTRY Bismuth Hydroxide, Bi(OH) 3 , is obtained as a white, amor- phous precipitate by pouring a solution of bismuth nitrate into excess of ammonium hydroxide : ) 3 + 3NH 4 OH->Bi(OH) 3 + 3NH 4 NO 3 . When dried at 100, it loses water and forms the first anhydride, BiO(OH) or Bi 2 O 3 ,H 2 O. Bismuth hydroxide has no acidic properties, but it acts as a base, combining with acids to form salts. The latter are derived either from the normal hydroxide, Bi(OH) 3 , or from the compound BiO'OH (see below). Higher Oxides of Bismuth (Peroxides of Bismuth} When bismuth trioxide is suspended in alkali and chlorine is passed into the mixture, oxides of bismuth are obtained which contain a higher proportion of oxygen than the trioxide. The product containing most oxygen is obtained by using a concentrated solution of potassium hydroxide. After saturating with chlorine, excess of concentrated nitric acid is added. The insoluble product is yellowish-red to scarlet-red in colour, and may be a mixture of Bi 2 O 4 and B5 2 O 5 , but no definite evidence on the matter has so far been obtained. It has further been stated that the production in question has acidic properties, and it has therefore been termed bismuthic acid (from analogy with antimonic acid). On this point also, however, the evidence is at present contradictory. Normal and Basic Salts As mentioned above, bismuth salts are derived exclusively from the hydroxide, Bi(OH) 3 , and its an- hydride, BiO(OH), acting as bases. The trichloride, BiCl 3 , and the oxychloride, BiOCl, have already been discussed. Bismuth tri- nitrate, Bi(NO 3 ) 3 , is obtained by dissolving the metal in dilute nitric acid, and separates on concentrating the solution in large triclinic colourless crystals, with 5H 2 O. The salt is decomposed by watei with formation of the basic nitrate, BiO'NO 3 , a white powder, which is used in medicine under the name subnitrate of bismuth. A numbei of basic nitrates are known. Bismuth Sulphate, Bi 2 (SO 4 ) 3 , is obtained by dissolving the trioxide in moderately concentrated sulphuric acid, and separates from solution in colourless, hygroscopic needles, which are decom- posed by water with formation of basic salts. Bismuth Trisulphide, Bi 2 S 3 , is the only sulphide of bismuth definitely known. It is obtained in crystalline form (grayish leaflets) by heating bismuth with excess of sulphur^ and as a brownish-black ELEMENTS OF THE NITROGEN GROUP 515 amorphous precipitate by passing hydrogen sulphide into a solution of a bismuth salt. Under great pressure the amorphous changes to the crystalline modification ; increase of temperature greatly accele- rates the change. The trisulphide is practically insoluble in solutions of alkali hydroxides and sulphides (cf. As 2 S 3 ; Sb 2 S 3 ), but dissolves in hot hydrochloric or nitric acid. Tests for Bismuth A valuable test for bismuth compounds is the formation of a white precipitate of basic salt when a solution (preferably of the chloride or nitrate) is poured into excess of water. The dark colour of the trisulphide, and its insolubility in alkali sulphides, at once distinguish bismuth compounds from those of arsenic and antimony. General Characters of the Elements of the Fifth Group and Summary The more important physical properties of the members of the nitrogen sub-group are shown in the accom- panying table, which also illustrates the variation of these properties with increasing atomic weight : Nitrogen. Phosphorus. Arsenic. Antimony. Bismuth. Atomic weight Density . . Melting-point Boiling-poii t 14.01 i. 026 at -252 210 -195.6 31.0 1.82 to 2.15 44 290 74.96 4.72 to 5.73 ca. 500 120.2 5.78 to 6.5 360.6 ca. 1300 208.0 9.78 264 1420 It will be observed that bismuth is out of accord with the remaining members of the series as regards its melting-point and boiling-point. The wide variation in the numbers quoted for the densities is due to the fact that some of the elements exist in different allotropic modifications. As regards the chemical characters, the first point to note is that the main valencies in the group are three and five. In the case of bismuth, however, the existence of quinquevalent compounds has not been definitely proved. The valencies are not exclusively three and five, however ; nitrogen, in particular, functions with other valencies as well. The second point to note is that the non-metallic character gradually weakens with increase of atomic weight ; whilst nitrogen is a typical non-metal, antimony and bismuth have distinct metallic characters. This change is illustrated by the appearance of the elements themselves, as also by the behaviour of their compounds. For instanc e, the oxides of nitrogen are exclusively acidic, antimony 516 A TEXT-BOOK OF INORGANIC CHEMISTRY trioxide is both basic and acidic, bismuth trioxide is exclusively basic. The chlorides of nitrogen and phosphorus are hydrolyzed irreversibly by water ; those of the remaining three elements are hydrolyzed par- tially and reversibly. Further, the first four elements form volatile hydrates of the type EH 3 , which is an indication of non-metallic character ; but they diminish in stability as the atomic weight in- creases, and so far no hydride of bismuth has been obtained. The existence of stable nitrates and sulphates of antimony and of bismuth is further proof of basic properties. Finally, the acidic properties of the trisulphides of arsenic and antimony is shown by the existence of thioantimonites, (NH 4 ) 3 As(Sb)S 3 ; but bismuth trisulphide is insoluble in alkali sulphide solution. CHAPTER XXXIII ELEMENTS OF THE CHROMIUM GROUP (GROUP VI, SUB-GROUP A) Sub-group A Sub-group B Chromium, Cr .... 52.0 Oxygen, O 16.00 Molybdenum, Mo . . . 96.0 Sulphur, S 32.07 Tungsten, W .... 184.0 Selenium, Se .... 79.2 Uranium, Ur . . . . 238.5 Tellurium .... 127.5 THE elements of sub-group B have already been considered in detail (pp. 20, 273) ; of the four members of sub-group A chromium is the only important element. All the members of the sub-group exhibit a number of valencies ; the lower oxides are basic, the higher acidic. Corresponding with the position of the elements in the sixth group of the periodic table, the typical oxides have the respective formulae CrO 3 ,MoO 3 ,WO 3 ,UrO3 ; they are almost exclusively acidic. CHROMIUM Symbol, Cr. Atomic weight=52.o. Occurrence Chromium does not occur free in nature. The ex- clusive source of chromium compounds is chrome ironstone or chromite^ FeO,Cr 2 O ; ; another natural compound of chromium is crocoisite, Preparation The element is now prepared commercially according to the Goldschmidt process (p. 466) by reducing chromium trioxide by means of powdered aluminium, the mixture being ignited by a fuse of magnesium ribbon. When a slight excess of the oxide is used the resulting metal is free from aluminium. Properties Chromium is a very hard, steel-gray metal, with high metallic lustre; it melts at 1489 (Burgess). Its density is 6.8. It is stable in the air at room temperature, but when strongly heated burns to Ihe oxide Cr 2 O 3 . It dissolves on warming with dilute hydrochloric or sulphuric acid, hydrogen being given off. Nitric acid does not dissolve it, but renders it " passive," and in this state it is 517 Si8 A TEXT-BOOK OF INORGANIC CHEMISTRY not attacked by hydrochloric or sulphuric acid. Other oxidizing agents have the same effect as nitric acid in rendering chromiurr inactive ; the nature of this " passivity " is not understood (cf. iron p. 54 2 )- It is remarkable that when dilute acids act on some specimens of chromium the evolution of hydrogen is periodic ; in other words, a period of rapid solution of the metal alternates regularly with a com- plete cessation of action. Chromium is now largely used in the preparation of chrome steel, which contains up to 3 per cent, of chromium. Compounds of Chromium Chromium forms three series oi compounds, which may be looked upon as being derived from the oxides, CrO, chromous oxide, Cr 2 O 3 , chromic oxide, and CrO 3 , chromic anhydride. Oxide. Character. Corresponding Compounds. (CrO) basic chromous salts ; example, CrCl 2 Cr 2 O 3 basic chromic salts CrCl 3 Cr0 3 acidic { chromates K 2 CrO 4 ( dichromates K 2 Cr 2 O 7 Chromous Compounds The compounds of divalent chromium are powerful reducing agents, having a great tendency to change into chromic salts by oxida- tion ; for this reason they are very difficult to prepare in pure condition. Chromous chloride, CrCl 2 , is obtained in solution by dissolving chromium in hydrochloric acid free from oxygen, and in the anhydrous condition as a white crystalline compound by heating chromic chloride in a current of hydrogen. The aqueous solution of chromous chloride is blue, but rapidly turns green in the air owing to oxidation : 4 CrCl 2 + 4 HC1 + O 2 -> 4 CrCl 3 + 2 H 2 O. Chromous sulphate, CrSO 4 ,7H 2 O, occurs in blue crystals isomorphous with ferrous sulphate. Chromous acetate, Cr(C 2 H 3 O 2 )2, a red crystalline powder obtained by adding a solution of chromous chloride to a saturated solution of sodium acetate, is the most stable chromous' salt. Chromous hydroxide, Cr(OH) 2 , obtained by double decomposition, is yellow and rapidly oxidizes in the air. The corresponding oxide, CrO, has not hitherto been obtained. Chromic Compounds Chromic salts are derived from chromic hydroxide, Cr(OH) 3 , acting as a base, and in practice can be obtained by dissolving the base in the appropriate acid. Chromic Hydroxide, Cr(OH) 3 ,.rH 2 O, is obtained as a bluish- green colloidal precipitate (hydrogel) on adding ammonium hydroxide to a solution of a chromic salt. When freshly precipitated it dissolves in a solution of sodium or potassium hydroxide to form a chromite (see below) ; but on boiling the solution it is reprecipitated in a less CHROMIUM 5r 9 highly hydrated, more insoluble form. When the hydroxide is heated in air it y ; elds chromic oxide, Cr 2 O 3 , a green insoluble powder. The latter is used as a pigment under the name of chrome-green. Chronic Chloride, CrCl 3 , is obtained in the anhydrous form as violet leaflets by heating a mixture of chromic oxide and charcoal in a streari of chlorine (cf. A1C1 3 , p. 468). The anhydrous chloride is practical!'/ insoluble in water (a little goes into solution on prolonged boiling), but when a trace of chromous chloride is added the trichlo- ride rapidly dissolves to form a green solution. This remarkable phenomenon is not yet understood. Chronvc chloride can also be obtained in solution by dissolving chromic hydroxide in hydrochloric acid, or by reducing chromates or dichromates in hydrochloric acid solution. The aqueous solution is usually green, and on evaporation green, deliquescent crystals of the formula CrCl 3 ,6H 2 O are obtained. When heated in air the hydrated chloride is decomposed, and green chromic oxide, Cr 2 O 3 , is formed ; when, however, the salt is heated in a current of dry hydrogen chlo- ride it yields the violet anhydrous chloride. Another chloride of chromium, CrCl 3 ,6H 2 O, is obtained in blue crystals by dissolving chrome alum in hydrochloric acid and passing hydrogen chloride into the solution. It dissolves in water to form a bluish-violet solu- tion, and it has been shown that the chromium is present as Cr'"ions and all the chlorine is ionised. In the green hydrate only one of the chlorine atoms is present in the ionic form ; its constitution, according to Werner, is represented by the formula [Cr(H 2 O) 4 Cl 2 ]Cl,2H 2 O (cf. p. 562). Chromic Sulphate, Cr 2 (SO 4 ) 3 ,is obtained by dissolving freshly precipitated chromic hydroxide in sulphuric acid in the cold ; on concentrating the violet solution in a vacuum violet crystals, of the formula Cr 2 (SO 4 ) 3 , 15 H 2 O, separate out. When the violet solution is heated it turns green owing to hydrolysis, but on standing for a long period nt room temperature it returns to the original violet colour. As in tie case of chromic chloride, only the violet solution contains the nornal salt. The nature of the green compound has not been conclusi/ely established; it may possibly be represented by the formula [Cr 4 O(SO 4 ) 4 ]SO 4 , as only about one-third of the SO 4 is ionised (p. 562). Chrome Alums When an alkali sulphate or ammonium sulphate is added in the requisite proportion to a solution of chromic sulphate and the solution is evaporated at a temperature not exceeding 30, a doublo salt of the type Cr 2 (SO 4 ) 3 ,M I 2 SO 4 ,24H 2 O, belonging to the 5$o A TEXT-BOOK OF INORGANIC CHEMISTRY class of alums, separates out in octahedral crystals, which appear dark violet by reflected, red by transmitted light. The potassium salt, Cr 2 (SO 4 ) 3 ,K 2 SO 4 ,24H 2 O, which is best known, is conveniently prepared by passing sulphur dioxide into a solution of potassium dichromate, acidified with sulphuric acid, the mixture being kept cool throughout : The aqueous solutions of the chrome alums, like that of chromic sulphate, turn green on heating, and slowly regain their original colour at room temperature. Chromites The slightly acidic character of chromic hydroxide is shown by the fact, already mentioned, that when freshly pre- cipitated it is soluble in potassium or sodium hydroxide solution. The solution presumably contains an alkali chromite, for example, Cr(OH) 2 OK or CrO'OK. The naturally occurring chrome iron- stone, FeO,Cr 2 O 3 , may be regarded as ferrous chromite, (CrO*O) 2 Fe. Chromium Trioxide, CrO 3 , and Chromates The ultimate source of all the chromium compounds is chrome ironstone, FeO,Cr 2 O 3 . The finely powdered ore is mixed with potassium carbonate and lime and roasted in a reverberatory furnace. The chief use of the lime is to keep the mass porous, and thus facilitate the oxidizing action of the atmospheric oxygen. The products of the reaction are potassium and calcium chromates, ferric oxide, and carbon dioxide : = 6K 2 CrO 4 +2CaCrO 4 The product is broken up and extracted with water, whereby the chromates of potassium and calcium are dissolved. Potassium sulphate is then added, and by double decomposition potassium chromate and calcium sulphate, the latter of which is almost insoluble in water, are formed. The solution is poured off and treated with sulphuric acid, whereby potassium dichromate, K 2 Cr 2 O 7 , is formed : The solution is then evaporated, and on cooling the dichromate separates in large red triclinic crystals. The object of preparing the dichromate instead of the chromate is that the former, being much less soluble in water, is more easily separated from solution and purified by recrystallization. CHROMIUM 521 Potassium Dichromate, K 2 Cr 2 O 7 , prepared as just described, is soluble about i in 10 of water at 15, and the solution has a strongly acid reaction. When the dry salt is heated it decomposes, giving off oxygen : When heated with concentrated sulphuric acid, chromium and potassium sulphates are formed and oxygen is given off: K 2 Cr 2 7 + 4H 2 S0 4 ->K 2 S0 4 +Cr 2 (S0 4 ) 3 + 4H 2 + 3 0. The same reaction takes place in the presence of a reducing agent, and therefore an acid solution of potassium dichromate is a powerful oxidizing agent. The equations expressing the oxidizing action of potassium dichromate can readily be written when the salt is represented as K 2 O,2CrO 3 , and it is noted that 2CrO 3 give Cr 2 O 3 and 30 available for oxidation : K 2 O,2CrO 3 ->K 2 O + Cr 2 O 3 + 30. The equations representing the oxidation of sulphur dioxide to sulphuric acid by potassium dichromate may therefore be written as follows : K 2 Cr 2 O 7 ->K 2 O + Cr 2 O 3 + 30 K 2 O + Cr 2 O 3 + 4H 2 SO 4 -K 2 SO 3[S0 2 + H 2 + 0]-> 3 [H 2 S0 4 ]. Adding On the same principle other equations representing the oxidizing action of potassium dichromate may readily be written. When it is heated with concentrated hydrochloric acid the latter is oxidized to chlorine and water : Adding K 2 Cr 2 O 7 + I4H Cl-^2CrCl 3 + 2KC1 + 7H 2 O + 3C1 2 . It also acts as an oxidizing agent in the absence of acid. When a film of gelatine containing potassium dichromate is exposed to light, reduction to chromic oxide takes place, and the latter forms with the gelatine a compound which, unlike gelatine itself, does not 522 A TEXT-BOOK OF INORGANIC CHEMISTRY dissolve or swell up in warm water. This property is taken advantage of for photographic purposes. Potassium dichromate is used as an oxidizing agent in batteries and in the preparation of organic dyes, etc. Sodium dichromate^ Na 2 Cr 2 O 7 ,2H 2 O, which is cheaper and much more soluble in water than the potassium salt, is often used in place of the latter. Ammonium dichromate, (NH 4 ) 2 Cr 2 O r , has already been mentioned (p. 202) ; it gives free nitrogen on heating : ( N H 4 ) 2 Cr 2 O 7 ->Cr 2 O 3 + 4H 2 O + N 2 . When chromates or dichromates are treated with more free acid, still higher chromates are obtained. Thus potassium trichromate, K 2 O,3CrO 3 or K 2 Cr 3 O 10 , and potassium tetrachromate, K 2 O,4CrO 3 or K 2 Cr 4 O 13 , have been isolated. Potassium Chromate, K 2 CrO 4 , is obtained by adding the requisite quantity of potassium hydroxide to a solution of potassium dichromate : K 2 Cr 2 O 7 + 2KOH->2K 2 CrO 4 + H 2 O. It occurs in yellow, anhydrous rhombic crystals, isomorphous with potassium sulphate. It is very soluble in water ; the solution has an alkaline reaction owing to hydrolysis. A number of chromates which are insoluble in water, and can therefore be prepared by double decomposition, are of interest. Lead chromate, PbCrO 4 , a bright yellow powder, is used as a pigment under the name of chrome yellow. Barium Chromate, BaCrO 4 , is also used as a pigment. Calcium chromate, CaCrO 4 ,2H 2 O, is isomorphous with gypsum, CaSO 4 ,2H 2 O. Silver chromate, Ag 2 CrO 4 , and mercurous chromate, Hg 2 OO 4 , are red. Chromates are sometimes prepared by double decomposition between a soluble salt and potassium dichromate. This is owing to the fact that in the solution of the latter salt CrO 4 " ions are also present to a small extent : and when the chromate is the less soluble salt it is precipitated. Chromium Trioxide, Chromic Anhydride, CrO 3 When concentrated sulphuric acid is added to a cold concentrated solution of potassium dichromate, after some time chromium trioxide, CrO 3 , separates in the form of red, needle-shaped crystals. The crystals CHROMIUM 523 are very deliquescent and extremely soluble in water, forming a red solution which probably contains dichromic acid, H 2 Cr 2 O 7 . On evaporating the solution, however, the trioxide crystallizes out. When heated in air chromium trioxide loses oxygen, and gives chromic oxide, Cr 2 O 3 . It is an extremely powerful oxidizing agent. Warm alcohol catches fire when dropped on it; it chars organic matter, paper, etc., at once. When a solution of chromic anhydride is neutralized with alkali, chromates and dichromates are formed. The former, as already mentioned, are of the same type as the sulphates, the latter correspond with the oyrosulphates, e.g. K 2 Cr 2 O 7 , K 2 S 2 O 7 . When to an acidified solution of potassium dichromate some hydrogen peroxide is added and the mixture is shaken up with ether, it will be observed that the layer of ether which separates out is coloured deep blue. If no ether is used, the blue colour observed when the hydrogen peroxide is added rapidly disappears, oxygen being given off, and the substance giving rise to it is therefore very unstable (p. 142). The blue compound has not been definitely isolated, but is believed to be a perchromic acid, of the formula HCrO 6 . Chromyl Chloride, CrO 2 Cl 2 , is prepared by heating a mixture of potassium dichromate, sodium chloride and concentrated sul- phuric acid ; the chromyl chloride distils over as a deep red, fuming liquid which boils at 118. It is readily decomposed by water into chromic oxide and hydrochloric acid : CrO 2 Cl 2 +H 2 O->CrO 3 Chromyl chloride corresponds with sulphuryl chloride, SO 2 C1 2 , and may be regarded as being derived from chromic acid, CrO 2 (OH) 2 , by substituting two Cl atoms for the two OH groups. The intermediate compound, CrO 2 (OH)Cl, chlorochromic acid, is not known, but its potassium salt, potassium chlorochromate, CrO 2 (OK)Cl, is obtained in yellovish-red crystals by crystallizing potassium dichromate from a concentrated solution in hydrochloric acid: K 2 Cr 2 O 7 + 2HCl$2CrO 2 (OK)Cl + H 2 O. Chromium Ammonia Compounds From chromic chloride, by the action of ammonia, a number of complex compounds have been derived, of the empirical formulae Cr(NH 3 ) ( iCl 3 , Cr(NH 3 ) 5 Cl 3 , Cr(NH 3 ) 4 Cl 3 , etc., which behave like the better known cobaltammines, of analogous constitution. It is therefore more convenient to deal with these compounds at a later stage (pp. 551, 561). 524 A TEXT-BOOK OF INORGANIC CHEMISTRY Tests for Chromium Compounds of chromium are readily recognized by their characteristic colours. Chromic salts give with alkali hydroxides a green precipitate of chromium hydroxide, soluble in excess of alkali, but reprecipitated on boiling. When fused with alkali and an oxidizing agent, chromic salts yield chromates (p. 520) which are characterized by a yellow colour, and formation of char- acteristic precipitates with soluble lead and silver salts. Further, both chromates and dichromates act as oxidizing agents, especially in acid solution, being thereby reduced to green chromic salts. MOLYBDENUM This comparatively rare element occurs chiefly as molybdenite, MoS 2 , which has a strong resemblance to graphite, and as wulfenite, PbMoO 4 . When molyb- denite is roasted, it is converted to the trioxide, MoO 3 , which is a white powder. Molybdenum can be prepared from the trioxide by heating in a current of hydro- gen, or by heating with carbon in the electric furnace. Molybdenum is a hard metal, of density 9.1, and strongly resembles iron; like the latter metal it takes up carbon and may be welded and tempered. The most important oxide of molybdenum is the trioxide, MoO 3 , which, like chromium trioxide, CrO 3 , is exclusively acidic. It readily dissolves in solutions of alkali hydroxide or ammonium hydroxide to form molybdates examples, sodium molybdate, Na 2 MoO4,ioH 2 O, and ammonium molybdate, (NH 4 ). 2 MoO 4 . When a strong acid is added to a molybdate solution, molybdic acid, H 2 MoO4,H 2 O, separates in colourless, lustrous crystals ; it is soluble in excess of acid. Ammonium molybdate, dissolved in excess of moderately concentrated nitric acid, is used as a test for phosphates. When the phosphate is added to the molybdate reagent and the mixture warmed, ammonium phosphomolybdate, of the approximate composition (NH 4 ) 3 PO4,iiMoO 3) 6H2O, comes down as a yellow precipitate. The composition of this precipitate indicates the tendency to the formation of polymolybdates, which is much more pronounced than the tendency to form polychromates (p. 522). Besides the trioxide, the compounds Mo 2 O 3 and MoO 2 are definitely known ; neither of them is definitely basic. Chlorides of the formulae (MoCl 2 ) 3 ; MoCU ; MoCl 4 and MoCl 3 have been described. TUNGSTEN OR WOLFRAM Tungsten is found in the relatively scarce minerals, wolfram, Fe(Mn)WO 4 , and scheelite, CaWO 4 . The element itself is made by reducing tungstic acid, H 2 WO 4 , with aluminium according to the Goldschmidt process. Tungsten is a very hard steel-gray metal of density about 19 ; it is stable in the air under ordinary con- ditions, but on heating forms the trioxide. When added to steel in the propor- tion of about 5 per cent. , a very hard alloy (tungsten steel) is formed. Tungsten is also used in making filaments for incandescent lamps. On fusing tungsten with sodium carbonate and extracting with water, a solution of sodium tungstate, Na 2 WO 4 , is obtained. By adding acid to the solution, tungstic acid, H 2 WO 4 ,H 2 O, is obtained; when the latter is ignited tungsten ELEMENTS OF THE CHROMIUM SUB-GROUP 525 trioxide, WO :J , remains. Like molybdenum trioxide, tungsten trioxide has a considerable tendency to unite with other compounds, forming complexes con- taining a large number of WO 3 groups. Phosphotiingstic acid, which is used as a reagent, has the formula H 3 PO 4 ,^WO 3 , where x is a large number, probably variable ace ording to the conditions of preparation. Four tungsten chlorides, WC1 2 , WC1 4 , WC1 5 , WC1 6> and two oxychlorides, WOoCl 2 ai d WOC1 4 , are known. Tungsten hexachoride, WC1 6 , prepared by the direct action of chlorine on tungsten, occurs in dark violet crystals, which melt at 275 . URANIUM The chie" source of uranium compounds is pitch-blende or uraninite, which is found in Aastria and in Cornwall; it consists mainly of the oxide, UO 2 '2UO 3 or U 3 O 8 . tn recent years pitch-blende has become very important owing to the fact that it contains radium (p. 564). Uranium is also of interest as being the element wiih the highest atomic weight (238.5). The met;il is obtained by heating the chloride with sodium or by reducing the oxide with carbon in the electric furnace. It is a hard silvery-white metal ; its density is 18.7. There are five oxides of uranium : UO 2 , which is exclusively basic ; U 2 O 3 ; U 3 O 8 , which is the most stable in the air ; UO 3 , which is both basic and acidic, and UO 4 , which is a peroxide. There are two chief series of salts. The uranous salts, in which the uranium is quadrivalent, are obtained by dissolving the dioxide in strong acids ; salts of this type are uranous chloride, UC1 4 , and uranous sulphate, L(SO 4 ) 2 ,8H 2 O. The second series of salts, which contain hexavalent uranium, are derived from the compound UO 2 (OH) 2 , which is usually called uranic acid. They thus con- tain the divalent group UO 2 , the so-called uranyl group. Among the salts belonging to this class are uranyl nitrate, UO 2 (NO 3 ) 2 ,6H 2 O ; uranyl sulphate, UO. 2 SO 4> 3H 2 O ; and ^lr any I chloride, UO 2 C1 2 . The salts are prepared by dis- solving uranic acid in the appropriate acid. As its name implies, the compound UO 2 ^OH) 2 has also acidic properties. The salts obtained by treating it with alkali hydroxides are, however, diuranates analogous in constitution to dichromates. Sodium diuranate, Na 2 U 2 O 7 ,6H 2 O, which is called uranium yellow, is used for tinting glass. Two uranium chlorides, UC1 4 and UC1 5 , are definitely known. Summary of Chromium Sub-group As chromium is the only important member of the sub-group, it is unnecessary to com- pare the properties of the elements in detail. They are characterized by forming several classes of compounds -with different valencies, but in the most important compounds they are hexavalent, corre- sponding with their position in the periodic table. The oxides CrO 3 , MoO 3 , \VO 3 and UO 3 are almost exclusively acidic ; the best known compounds derived from them are of the types K 2 Cr(Mo,W,Ur)O 4 and K 2 Cr 2 (Mo 2 ,W 2 ,Ur 2 )O 7 . There is a considerable tendency throughout to form complex salts containing a number of EO 3 groups. Sulphur 526 A TEXT-BOOK OF INORGANIC CHEMISTRY also shows this tendency to a slight extent (cf. pyrosulphuric acid, H 2 S 2 O 7 ). As regards analogies with members of sub-group A, the four elements now under consideration are most closely allied with sulphur. Many of their compounds are of the same type as those of sulphur, and in some cases are isomorphous with them. This is shown in K 2 CrO 4 and K 2 SO 4 , and in Na 2 CrO 4 ,ioH 2 O and Na 2 SO 4 ,ioH 2 O. The analogy of sulphuryl chloride, SO 2 C1 2 , with chromyl chloride, CrO 2 Cl 2 , has already been referred to. Chromium also 'shows analogies with metals belonging to other groups. The chromic compounds, for example, show similarity with those of aluminium and of trivalent iron, as is evident from the existence of the isomorphous alums. Further, there is some analogy between compounds of divalent chromium and divalent iron, e.g. CrSO 4 ,7H 2 O and FeSO 4 ,7H 2 O are isomorphous. Further, an oxide of iron of the formula FeO 3 is known (p. 544). CHAPTER XXXIV THE MANGANESE SUB-GRO.UP (GROUP VII, SUB-GROUP A) MANGANESE and the halogens are the only known members of the seventh group. The former element, like those belong- ing to the sixth group just considered, forms compounds belonging to a variety of stages of oxidation, and shows analogies with magnesium, iron, chromium and other elements, as well as with the halogens. MANGANESE Symbol, Mn. Atomic Weight =54. 93. Occurrence Manganese does not occur naturally in the free condition (except occasionally in meteorites), but in combination with oxygen ii occurs fairly abundantly as pyrolusite, MnO 2 ; also as hausmannite, Mn 3 O 4 , braunite, Mn 2 O 3 , manganite, Mn 2 O 3 ,H 2 O, and manganese spar, MnCO 3 . Preparation 'The metal is readily obtained pure by reducing the dioxide, MnO2, with aluminium according to the Goldschmidt process. Properties Manganese is a reddish-gray lustrous brittle metal, harder than iron. Its density is 8.0, and it melts at 1207. It becomes superficially oxidized in moist air and decomposes water at 100, with evolution of hydrogen. It dissolves readily in acids, even in weak acids such as acetic acid, hydrogen being given off and manganous salts formed. Manganese is largely used as an addition to iron (p. 542), and also in the production of manganese bronze (p. 411). OXIDES OF MANGANESE Six oxides of manganese are known : manganous oxide, MnO ; di- nmngaiuse trioxide, Mn 2 O 3 ; manganese dioxide, MnO 2 ; trimanganese tetr oxide, Mn 3 O 4 ; manganese trioxide, MnO 3 , and manganese hept- oxide, Mn 2 O;-. 527 528 A TEXT-BOOK OF INORGANIC CHEMISTRY Oxide. Character. Corresponding Compounds. MnO basic Manganous Salts ; example MnCl 2 Mn 3 O 4 basic (mixed oxide) Mn 2 O 3 basic Manganic Salts MnCl 3 MnO 2 acidic and peroxide Manganites CaMnO 3 MnO 3 acidic Manganates K 2 MnO 4 Mn 2 O 7 acidic Permanganates KMnO 4 Manganous Oxide, MnO, is prepared by heating manganous carbonate, MnCO 3 , in absence of air, or by heating any of the other oxides in a current of hydrogen. It is a green powder, which oxidizes to Mn 3 O 4 on heating strongly in air. Manganous Hydroxide, Mn(OH) 2 , is obtained as a white precipitate on mixing air-free solutions of an alkali hydroxide and a manganous salt. It rapidly absorbs oxygen from the air, forming the green hydroxide, Mn(OH) 3 . Trimanganese Tetroxide, Mn 3 O 4 , occurs naturally in red prismatic crystals as hausrnannite, and is obtained on heating any of the other oxides of manganese in the air at 1000. When it is heated with dilute sulphuric acid, manganous sulphate goes into solution and the dioxide remains (cf. Pb 3 O 4 , p. 487) : When it is treated with cold concentrated sulphuric acid, a mixture of manganous and manganic sulphates is probably formed. Manganic Oxide, Mn 2 O 3 , occurs naturally as braunite, and is obtained by heating any of the other oxides in oxygen at 650- 900. W T hen heated with dilute sulphuric acid, manganous sulphate dissolves and manganese dioxide remains insoluble. For this reason it is sometimes represented as MnO'MnO 2 ; on the other hand, corresponding salts, the manganic salts, are known. Manganese Dioxide, MnO 2 , occurs naturally in grayish- black crystals as pyrolusite. It is best obtained in a pure con- dition by cautiously heating manganous nitrate. When it is heated with hydrochloric acid, chlorine is given off and manganous chlorine, MnCl 2 , remains in solution : This method is used for the commercial preparation of chlorine, but is now to some extent replaced by electrolytic methods (p. 383). When the dioxide is treated with cold concentrated hydrochloric MANGANESE 529 acid, it dissolves to a dark liquid and very little chlorine is given off; on warming the solution chlorine is given off and manganous chloride remains in solution. The nature of the compound or compounds present in the cold solution is not settled. It may contain the tetra- chloride, MnCl 4 , the trichloride, MnCl 3 , or a mixture of the two. When manganese dioxide is heated with concentrated sulphuric acid, manganous sulphate is formed and oxygen is liberated : 2MnO 2 + 2H 2 SO 4 ->2MnSO 4 +2H 2 O + O 2 . Manganese dioxide is used in glass-making, and as an oxidizing agent in galvanic batteries. Manganous Acid (Hydrated Manganese Dioxide), Mn(OH) 4 , is obtained as a black precipitate by the action of a hypochlorite on a manganous salt in a neutral or alkaline medium. On partial dehydration, the compound MnO(OH) 2 or H 2 MnO 3 , a dibasic manganous acid, is obtained. The Wcldon process for the utilization of the manganous chloride formed in the preparation of chlorine from pyrolusite is based on the formation of salts, manganites, derived from the dibasic acid, Milk of lime is added to the chloride solution and air forced through the mixture, whereby the manganous hydroxide first precipitated is converted into calcium manganite, CaMnO 3 (or CaO.MnO 2 ) : MnCl 2 + 2CaO + O-CaMnO 3 + CaCl 2 . The manganite, being insoluble in water, slowly settles as a black mud, the calcium chloride solution is poured off, and the manganite treated with hydrochloric acid for the production of more chlorine : Manganese Trioxide, MnO 3 , prepared by adding potassium permanganate, dissolved in sulphuric acid, to dry sodium carbonate, occurs as a dark red mass. It is the anhydride of manganic acid (g.v.]. Dimanganese Heptoxide, Mn 2 O 7 , is obtained as a green, oily liquid by cautiously adding potassium permanganate to con- centrated sulphuric acid, the mixture being well cooled. The heptoxide is very unstable, and decomposes violently on warming, forming the dioxide and oxygen. It is the anhydride of perman- ganic acid, HMnO 4 (q.v.\ 34 530 A TEXT-BOOK OF INORGANIC CHEMISTRY MANGANOUS SALTS A 11 the oxides of manganese on heating with acids yield manganous salts, in which the metal is divalent. When the higher oxides are used, either oxygen is given off or the acid is oxidized. The manganous compounds are the only stable salts containing Mn ions. Manganous Chloride, MnCl 2 , obtained by dissolving any of the oxides in hydrochloric acid, separates from solution in pink crystals with 4H 2 O. It can be obtained in the anhydrous form by heating the tetrahydrate in a current of hydrogen chloride, or by heating manganous ammonium chloride, MnCl 2 ,2NH 4 Cl,H 2 O (cf. magnesium chloride, p. 447). Manganous Sulphate, MnSO 4 , prepared by the general method, separates from solution below 6 in pink, monoclinic crystals isomorphous with ZnSO 4 ,7H 2 O ; between 6 and 20 it is obtained as MnSO 4 ,5H 2 O in triclinic crystals isomorphous with CuSO 4 ,5H 2 O ; above 20, MnSO 4 ,4H 2 O separates in rhombic prisms. It forms double salts with alkali sulphates, of the type MnSO 4 ,K 2 SO 4 ,6H 2 O, isomorphous with the similarly constituted salts of magnesium, zinc and iron (p. 452). Manganous Sulphide, MnS, is obtained as a green powder by heating any of the oxides in a stream of hydrogen sulphide ; and in the hydrated form as a pink precipitate when ammonium sulphide is added to a solution of a manganous salt. It is soluble in dilute acids, even in acetic acid. MANGANIC SALTS Manganic Hydroxide, Mn(OH) 3 , from which the manganic salts are derived, has already been mentioned. It is an extremely weak base, and the salts are practically completely hydrolyzed on addition of water. Manganic Chloride, MnCl 3 , is probably formed by the action of cold hydrochloric acid on manganese dioxide (p. 528, but has not been isolated from the solution. It has, however, been obtained by treating manganese dioxide, suspended in carbon tetrachloride, with dry hydrogen chloride, and then extracting the trichloride by means of ether. It is a nearly black solid with a greenish tinge, and is immediately decomposed by water. MANGANESE 531 Manganic Sulphate, Mn 2 (SO 4 ) 3 , is obtained as a dark-green powder by gently heating hydrated manganese dioxide with con- centrated sulphuric acid. It is immediately hydrolyzed by water. With alkali sulphates it forms alums, e.g. Mn 2 (SO 4 ) 3 ,K 2 SO 4 ,24H 2 O, which a -e more stable than the salt itself. MANGANATES AND PERMANGANATES Manganates The nianganates are derived from manganic acid, H 2 MnO (or MnO 2 (OH) 2 ), which has not itself been isolated, but the oxide, MnO 3 , is known (p. 529). Potassium manganate is obtained as a green mass by fusing manganese dioxide with potassium hydroxide or carbonate with or without the addition of an oxidizing agent such as potassium nitrate or chlorate. When no oxidizing agent is added the oxidation is effected by oxygen taken up from the air : The fused mass is extracted with water and the solution evaporated, when potassium manganate is deposited in green rhombic crystals, isomorphous with potassium sulphate and potassium chromate. Potassium manganate is stable only in the presence of alkali ; when the solution is diluted and warmed, still more readily when a weak acid is added (for example, when carbon dioxide is passed into the solution;, the colour changes to deep purple, owing to the formation of potassium permanganate, manganese dioxide being precipitated : The remarkable reaction just considered is in the nature of a hydrolysis, and as the oxide corresponding with potassium perman- ganate is Mn 2 O 7 , and MnO 2 is also formed, it is clear that a simultaneous oxidation and reduction of the manganate (whose corresponding oxide is MnO ;; ) has occurred. 1 We may assume that in the first stage of the hydrolysis, the very unstable manganic acid is formed : 3 x[K 2 MnO 4 + 2H 2 O->H 2 MnO 4 + 2KOH] (i), and immediately decomposes into permanganic acid, HMnO 4 , and manganese dioxide : 3H 2 MnO 4 ->2HMnO 4 -f MnO 2 + 2H 2 O (2), ^- instances of the decomposition of a compound into two others, cne in a higher and one in a lower stage of oxidation than the original substance, have already been given (cf. potassium chlorate, p. 181). 532 A TEXT-BOOK OF INORGANIC CHEMISTRY the permanganic acid finally uniting with part of the potassium hydroxide to form potassium permanganate : + KOH->KMnO 4 +H 2 O]. Adding As Ostvvald has pointed out, it is simpler to write the reaction in terms of ions : This shows that H' ions are used up in the change, which therefore can only proceed in acid or nearly neutral solution. In alkaline solution, practically no H' ions are present, and therefore the manganate is stable. 1 Manganates act as oxidizing agents in alkaline solution, being thereby reduced to manganese dioxide (MnO 3 ->MnO 2 + O). Permanganic Acid and Permanganates The perman- ganates are derived from permanganic acid, HMnO 4 ; the solutions contain the MnO 4 ' ion, which is purple. Permanganic Acid, HMnO 4 , is obtained by adding to a solu- tion of barium permanganate the requisite quantity of sulphuric acid, and filtering from the precipitated barium sulphate ; on evaporating the solution, the acid is obtained in violet-blue crystals. The acid is also obtained in aqueous solution by dissolving manganese heptoxide, Mn 2 O 7 , in cold water. Permanganic acid is a strong monobasic acid, comparable in strength with hydrochloric acid, and is a powerful oxidizing agent. Potassium Permanganate, KMnO 4 , prepared as described above, occurs in purple rhombic prisms, isomorphous with potassium perchlorate. The aqueous solution of the salt is deep purple. At 20 100 grams of water dissolve 6.35 grams, at 50 16.9 grams of the salt. Potassium permanganate is a powerful oxidizing agent. When the dry salt is heated,, it decomposes thus : 2KMnO 4 ->K 2 MnO 4 + MnO 2 + O 2 . When it is heated in solution with an alkali, the reaction by which it 1 In the conversion of the MnO/' to the MnO 4 ' ion there is a diminution in the number of negative charges, and the reaction therefore conforms to the extended definition of oxidation already given (p. 415). MANGANESE 533 is formed is reversed, and a green solution of the manganate is obtained : When a reducing agent is present in alkaline solution, further reduc- tion to the dioxide occurs, so that from two molecules of permanganate in alkaline solution, three atoms of oxygen are available for oxidation (Mn 2 O 7 ->2MnO 8 +3O); In acid solution, in the presence of a reducing agent, reduction proceeds to manganous salt, so that from two molecules of per- manganate, five atoms of oxygen are available for oxidation In the light of the explanations given in previous cases, there should be no difficulty in writing the equations representing reactions of this type. Suppose, for instance, it is required to write the equa- tion representing the oxidation of ferrous sulphate, FeSO 4 , to ferric sulphate, Fe 2 (SO 4 ) 3 , in the presence of free sulphuric acid. As the oxide corresponding with ferrous sulphate is FeO, and that corre- sponding with ferric sulphate is Fe 2 O 3 , it is clear that one atom of oxygen will oxidize two molecules of ferrous to one of ferric salt. We have therefore (i) 5 x[2FeSO 4 +H 2 SO 4 + O-Fe 2 (SO 4 ) 3 + H 2 O]. (2) Adding 2MnSO 4 + 8H 2 O. Sulphur dioxide, hydrogen sulphide, hydrogen peroxide, nitrous acid and other compounds reduce potassium permanganate to man- ganous .salts in acid solution, and the equations can be written as in the above example. The partial equation (i) is the same in each case. Further, as manganous salts are almost colourless in solution, potassium permanganate can be used for the quantitative estimation of many readily oxidizable substances. If the permanganate is added to the reducing agent, the end of the reaction is marked by the appearance of a permanent pink colour due to free permanganate, and no other indicator is required. 534 A TEXT-BOOK OF INORGANIC CHEMISTRY On account of their powerful oxidizing action, permangates are largely used as disinfectants. The cheaper sodium permanganate, which does not readily crystallize, is sold in solution for this purpose under the name of Condfs disinfecting fluid. The action is identical with that of the potassium salt, being due to the MnO 4 ' ion. Tests for Manganese The formation of a flesh-coloured sulphide, MnS, when ammonium sulphide is added to the solution of a manganous salt is a characteristic test ; the sulphide is soluble in acetic acid. Further, the formation of a green mass of manganate when a salt of manganese is fused with alkali and an oxidizing agent is a useful test ; when the green mass is treated with water and the solution is boiled, it becomes pink. Manganese salts give an amethyst colour to a borax bead. Summary Owing to the isolated position of manganese in the periodic table, there are no elements which very closely resemble it ; but, on the other hand, it shows a number of points of analogy with elements belonging to other groups. The remarkable variety of valencies shown by manganese has already been mentioned ; it acts as a divalent, trivalent, hexavalent, septavalent, and possibly quadri- valent element. In the divalent manganous compounds it re- sembles magnesium and ferrous salts, the compounds MnSO 4 ,7H 2 O, MgSO 4 ,7H 2 O and FeSO 4 ,7H 2 O being isomorphous; in its trivalent compounds it resembles iron and aluminium, as is clearly shown by the existence of a manganese alum, Mn 2 (SO 4 ) 3 ,K 2 SO 4 ,24H 2 O ; in its sexavalent compounds it resembles sulphur and chromium, as is shown by the isomorphism of the compounds K 2 MnO 4 , K 2 SO 4 and K 2 CrO 4 ; finally, in its septavalent compounds it resembles the highest oxidation stage of chlorine, the compounds KMnO 4 and KC1O 4 being isomorphous. It follows that the formula of potassium manganate OA /OK is represented graphically thus : ^Mn^ , and potassium per- O<^ \OK manganate thus : \Mn 8g> oxide). As the air forced in at the base of the furnace is at a very high temperature (about 800) it combines with the carbon, forming carbon dioxide. The latter, on its way upwards through the hottest part of the furnace at f, is reduced by the red-hot carbon to 1 Some ores are calcined before being introduced into the furnace, the object being to expel water and carbon dioxide and render them more porous. IRON 537 carbon monoxide, which then reduces the ferric oxide to metallic iron according to the equation : The reduction of the oxide takes place mainly in the upper part of the furnace, where the temperature is not sufficiently high (600 to 900) to fuse the iron. On its way dowji the furnace, the finely divided iron reaches the regions of higher temperature at stable below 760. y-iron and /3-iron are non-magnetic, y-iron can retain a considerable proportion of carbon in solid solution ; (3 and a-iron dissolve little or no carbon. When y-iron containing dissolved carbon is very rapidly cooled it forms a hard, brittle alloy (hardened steel); it is highly supercooled, and the presence of carbon retards IRON 341 its transformation to the modification (a-iron) stable under these condi- tions. When, on the other hand, the iron containing carbon in solution is cooled so slowly that equilibrium is attained at every stage, the final product contains a-iron, a little carbide (known as cementite) and graphite. The moderate heating and subsequent cooling used in the tempering Df steel brings the alloy towards the equilibrium state to an exten; determined by the conditions. The properties of a steel depend upon the relative proportions of y-iroh, solid solution of carbon in iron, cementite, and a and /3-iron in the product, and on their state of distribution. Pearlite, a name often met with in the discus- sion of stees, denotes a eutectic of cementite and a-iron, and contains 0.9 per cem. of carbon. Properties of Pure Iron Pure iron can be obtained by heating ferric oxide in hydrogen, or by electrolysis of an iron salt, e.g. ferric chloride, in aqueous solution. It is a white lustrous metal which takes a high polish ; its density is 7.86. It is malleable and ductile, and is not very hard. It melts about 1500, but softens at much lower temperatures. It is magnetic, but, unlike steel, it rapidly loses this property when the magnetizing force is withdrawn. Iron is permanent in dry air, but in moist air it quickly becomes coated with a layer of ferrous and ferric oxides, and is said to rust. As the rust does not form a continuous coating, but peels off in sheets, the corrosion spreads to deeper layers of the metal. Rusting is greatly accelerated by the presence of carbon dioxide (from the air). The exact mechanism of the rusting of iron is at present a matter of dispute. Grace Calvert and, later, Crum Brown expressed the view that carbon dioxide is essential ; the first stage of the process is the formation of ferrous carbonate, FeCO 3 ; thus Fe + O + CO 2 ->FeCO 3 , which is then oxidized more or less completely to red ferric oxide. Moody {Trans. Chem. Soc., 1906, 89, 720) adopts this view, and con- tends that iron does not rust in air and water entirely freed from carbon dioxide. In the presence of carbon dioxide and water the iron dissolves to form ferrous bicarbonate, which then undergoes oxidation is stated above. Lambert and Thomson (Trans. C/iem. Soc.j 1910, 97, 2426), on the other hand, state that although rusting does not occur with perfectly pure iron in contact with pure water and pure oxygen, traces of impurities are sufficient to cause oxidation under the same conditions, even if the impurity be not of an acid nature or likely to produce an acid during the reaction. Walker, 542 A TEXJ-BOOK OF INORGANIC CHEMISTRY Cederholm and Bent (Trans. Amer. Chem. Soc., 1907, xxix. 1251) also state that carbon dioxide is not essential for rusting to occur, and advo- cate an electrolytic explanation of the phenomenon. In virtue of its solution pressure (p. 85) iron sends out Fe" ions even into pure water ; but in the absence of oxygen the reaction soon comes to an end, owing to the accumulation of gaseous hydrogen on the metal, which sets up an E.M.F. of polarization. In the presence of oxygen, however, the hydrogen is removed by oxidation, and the ferrous salt is removed from the solution by precipitation as oxide, so that the reaction is enabled to proceed. Iron readily dissolves in dilute hydrochloric and sulphuric acids, with evolution of hydrogen and formation of the corresponding salts. With dilute nitric acid, ferrous nitrate and ammonium nitrate are formed. Iron is not attacked by concentrated nitric acid, but is changed to the passive state, in which condition it no longer precipi- tates copper from its salts. This passive condition is also brought about by other oxidizing agents, e.g. potassium dichromate, hydrogen peroxide. The explanation of the passive state most in favour is that it is due to the presence of a thin film of oxide on the surface of the metal. There is a considerable difference of potential between a metal in the active and the same metal in the passive state. Iron Alloys The alloys with carbon have already been fully dealt with. Certain of the properties of steel can be enhanced by the addition of other elements. The addition of manganese (about 12 per cent.) confers increased hardness on steel, while the alloy possesses a high degree of ductility and elasticity. Nickel (up to 3 per cent.) makes steel more tough and elastic. Chrome steel (containing about 2 per cent, of chromium) combines intense hardness with a high elastic limit. Tungsten is also a useful addition under certain circumstances. OXIDES AND HYDROXIDES OF IRON Three oxides of iron are known : Ferrous oxide . . . . . . FeO Ferric oxide ....... Fe 2 O 3 Triferric tetroxide (magnetic oxide of iron) . Fe 3 O 4 The first two oxides are basic, giving rise to ferrous salts (type FeCl 2 ) and ferric salts (type FeCl 3 ) respectively. The third is a mixed oxide, yielding a mixture of ferrous and ferric oxides, FeO, Fe 2 O 3 . IRON 543 Salts derived from an acidic oxide, FeO 3 , are also known, e.g. potassium '"errate, K 2 FeO. A . Ferrous Oxide, FeO, is obtained by heating ferric oxide in hydrogen i t 300, or by heating ferrous oxalate out of contact with air. It is a black powder, which readily becomes oxidized on exposure to air, and dL c solves in acids to form ferrous salts. Ferrov.s Hydroxide, Fe(OH) 2 , is obtained as a white precipi- tate by mixing air-free solutions of a ferrous salt and an alkali hydroxide in the absence of air. It rapidly turns green in the air owing to oxidation, and finally becomes brown owing to the formation of ferric h) droxide. Ferric Oxide, Fe 2 O 3 , occurs in lustrous black, six-sided crystals, as specular iron ore, and in reddish masses as hcematite. It is obtained as a red anorphous powder by heating ferrous sulphate in the prepa- ration of Nordhausen sulphuric acid, and is also obtained when most iron salts are strongly heated in air. It is stable in air at a red heat, but above 1000 the compound Fe 3 O 4 is formed. It dissolves in acids {though very slowly, especially after strong ignition) to form ferric salts. The amorphous powder is used as jewellers' rouge for polish- ing purposes, and also as a pigment under the name of Venetian red. Ferric Hydroxide, Fe(OH) 3v *H 2 O, is obtained as a brown gelatinous precipitate by adding excess of ammonium hydroxide to a solution of a ferric salt. By careful dehydration a compound of the composition Fe(OH) 3 or Fe 2 O 3 ,3H 2 O appears to be obtained. The water in this, as in other cases is adsorbed by the hydroxide, and it is doubtful whether definite compounds are formed (cf. p. 466). The compound 2Fe 2 O 3 ,3H 2 O occurs in nature as limonite (brown hcematite}. Iron rust is mainly a hydrated ferric oxide, with ferrous oxide and carbonate in smaller proportion. Freshly precipitated and washed ferric hydroxide is soluble in a solution of ferric chloride. The latter can be almost completely removed by prolonged dialysis (p. 353), and a deep brown, practically tasteless colloidal solution of ferric hydroxide, the so-called dialysed iron, is obtained. Ferric hydroxide is a very weak base, and therefore the ferric salts are consicerably hydrolyzed in aqueous solution. Trifeiric Tetroxide {magnetic iron ore), Fe 3 O 4 , occurs natu- rally in nearly black octahedral crystals, as magnetite or lodestone, so called because it is magnetic. It is formed when iron is burned at a high temperature in oxygen or air (hammer scale), and also when steam is passed over heated iron (p. 34). When heated with acids it 544 A TEXT-BOOK OF INORGANIC CHEMISTRY yields a mixture of ferrous and ferric salts, and may therefore be represented as FeO 5 Fe 2 O 3 (cf. chrome iron ore, FeO,Cr 2 O 3 ). Ferrates When chlorine is passed into ferric hydroxide sus- pended in cold concentrated potassium hydroxide solution, a dark red solution is obtained, from which, on cautious evaporation, potassium ferrate, K 2 FeO 4 , can be obtained in dark red crystals, isomorphous with potassium sulphate and potassium chromate : Potassium ferrate is derived from ferric acid, H 2 FeO 4 , but neither the acid nor the corresponding oxide, FeO 3 (cf. SO 3 , CrO 3 ), has been iso- lated. When the solution of a ferrate is acidified, oxygen is given off and ferric and potassium salts remain in solution. FERROUS SALTS Most ferrous salts are light green in colour. They are obtained by dissolving the metal in the appropriate acid or by reducing ferric salts in acid solution by " nascent " hydrogen, hydrogen sulphide, etc. They are easily oxidized to ferric salts. Ferrous Chloride, FeCl 2 , is obtained in the anhydrous form in colourless crystals by heating iron in a current of dry hydrogen chloride. It is obtained as FeCl 2 ,4H 2 O in green, monoclinic crystals, by dissolving iron in hydrochloric acid, and evaporating the solution with exclusion of air. When it is heated in the air, a mixture of ferric oxide and chloride is obtained : e 2 O 3 . Ferrous Sulphate (Green Vitriol^, FeSO 4 ,7H 2 O, is obtained by dissolving iron in dilute sulphuric acid. Commercially it is prepared by exposing iron pyrites, FeS 2 , to atmospheric oxidation ; the result- ing ferrous sulphate is extracted with water and the solution evapo- rated. It separates from solution in green, monoclinic prisms, isomorphous with ZnSO 4 ,7H 2 O and MgSO 4 ,7H 2 O. When heated to 100 the crystals lose 6H 2 O ; at a higher temperature the monohy- drate, FeSO 4 ,H 2 O, is decomposed and ferric oxide is obtained. When exposed to air at room temperature, the heptahydrate efflor- esces slightly and becomes coated with a brown layer of basic ferric sulphate, Fe 2 O(SO 4 ) 2 (or perhaps Fe(OH)SO 4 ). The same product is formed when the heptahydrate is roasted in the air. The basic sulphate obtained by roasting was formerly used for preparing Nord- hausen sulphuric acid ; for this purpose it was distilled from clay IRON 545 retorts and the volatile products collected in water or sulphuric acid (p. 286) : Fe 2 O(SO 4 ) 2 ->Fe 2 O 3 + 2SO 3 . Ferrous sulphate forms double salts with the alkali sulphates. The most important is ferrous ammonium sulphate,FeSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O, which is loss easily oxidized in the air than ferrous sulphate, and therefore fnds application in volumetric analysis. Ferrous sulphate is used as a disinfectant, in the preparation of ink, etc. Ferrous Carbonate, FeCO 3 , occurs naturally as spathic iron ore, in rhombohedral crystals isomorphous with calc-spar. It is obtained by double decomposition when Fe" and CO 3 " ions are brought to- gether in solution, but is very rapidly oxidized in the air. It is used in medicine mixed with sugar, which to some extent protects it against oxidation. Ferric carbonate has not been obtained. When Fe" f and CO 3 " ions are brought together in solution, the ferric carbonate which may be momentarily formed is completely hydrolyzed by water and ferric hydroxide s precipitated. FERRIC SALTS Ferric Hydroxide, Fe(OH) 3 , is a very weak base, and therefore ferric salts with strong acids are considerably hydrolyzed in solution, and salts with weak acids (e.g. carbonic acid, hydrogen sulphide) cannot be obtained from aqueous solution. Ferric Chloride, FeCl 3 , is obtained in the anhydrous form, as lustrous, d;irk green crystals, by heating iron in a current of chlorine gas. It is obtained in solution by oxidizing ferrous chloride with chlorine, nitric acid or other oxidizing agent, and separates from solution as the hexahydrate, FeCl 3 ,6H 2 O, in yellow crystals. Lower hydrates of the salt are also known. It cannot be obtained in the anhydrous form by heating one of the hydrates in air, as at high temperatui es hydrochloric acid is given off. The anhydrous salt is fairly volatile. At 300 to 400 the vapour density corresponds approximately with the formula Fe 2 Cl c ; above 750 with the formula FeCl 3 . Ferric chloride is considerably hydrolyzed in solution, the hydroxide remaining dissolved in colloidal form (hydrosol) : 35 546 A TEXT-BOOK OF INORGANIC CHEMISTRY Ferric Sulphate, Fe 2 (SO 4 ) 3 , is obtained by oxidizing ferrous sulphate in sulphuric acid solution. Nitric acid may conveniently be used for this purpose : Adding The aqueous solution is brownish-red (owing to the presence of colloidal ferric hydroxide, formed by hydrolysis) ; on evaporation the salt separates in the anhydrous form as a grayish-white powder. When the calculated quantity of potassium or ammonium sulphate is added to ferric sulphate solution, and the solution is evaporated over sulphuric acid at room temperature, an iron alum, e.g. Fe 2 (SO 4 ) 3 ,K 2 SO 4 ,24H 2 O, separates in violet, octahedral crystals (p. 468). SULPHIDES OF IRON Ferrous Sulphide, FeS, is obtained as a dark metallic-looking mass by heating iron and sulphur together. The sulphide thus ob- tained is dissolved by acids, hydrogen sulphide being given off and a ferrous salt formed. Hydrogen sulphide is usually prepared by this method. It follows from the above that no precipitate is obtained when hydrogen sulphide is passed into the solution of a ferrous salt con- taining free acid, but the sulphide is obtained as a black amorphous precipitate when ammonium sulphide is added to a ferrous or ferric salt. When a ferric salt is used, free sulphur is also formed : + 3(NH 4 ) 2 S->6NH 4 The precipitation is preceded by the reduction of the iron from the ferric to the ferrous condition. The amorphous sulphide when moist is rapidly oxidized in the air to ferrous sulphate. Ferric Sulphide, Fe 2 S 3 , cannot be obtained by precipitation in presence of water (see above), but is said to be formed as a yellow mass when a mixture of the components in the calculated proportions is heated. Iron Disulphide, FeS 2 , occurs naturally as iron pyrites in yellow crystals belonging to the regular system (cubes, octahedra or other forms), and is also formed by gently heating a mixture of ferrous IRON 547 sulphide and sulphur. When heated in air it burns to sulphur dioxide and ferric oxide : 4FeS 2 +iiO 2 ->2Fe 2 O 3 + 8SO 2 , and is largely used in preparing" sulphuric acid (p. 291). The other sulphides of iron also give ferric oxide and sulphur dioxide when roasted in air. When it is heated in absence of air, part -of the sulphur is given off and ferros,i-ferri<: sulphide, Fe 3 S 4 , is formed: 3 FeS 2 ->Fe 3 S 4 + S 2 . COMPLEX IRON CYANOGEN COMPOUNDS The normal cyanides of iron, Fe(CN) 2 and Fe(CN) 3 , are not known, but two complex cyanides, potassitim ferrocyanide, K 4 Fe(CN) 6 , and potassium ferricyanide, K 3 Fe(CN) 6 , are known. Potassium ferrocyanide is obtained by fusing together potassium carbonate, iron filings and animal refuse such as charred blood, scraps of horns, hoofs, etc. ; the mass is extracted with water, and on evaporation potassium ferrocyanide separates in the form of light yellow crystals as K 4 Fe(CN) 6 ,3H 2 O. When potassium ferrocyanide is strongly heated, po'assium cyanide, nitrogen and iron carbide are formed : K 4 Fe(CN) 6 ->4KCN + FeC 2 + N 2 . Its employment in the preparation of carbon monoxide and of hydro- cyanic acid has already been mentioned. Potassium Ferricyanide, K 3 Fe(CN) 6 , is obtained bypassing chlorine into a solution of potassium ferrocyanide : The ferricyanide occurs in dark-red crystals which are readily soluble in water ; it acts as an oxidizing agent. It might appear that these salts can be regarded as compounds of potassium cyanide with ferrous and ferric cyanide respectively, thus |t 4 Fe(CN) 6 =4KCN,Fe(CN) 2 and K 3 Fe(CN) 6 = 3KCN,Fe(CN) 3 . As a matter of fact, however, the solutions contain the complex ions Fe(CN) 6 "' 7 and Fe(CN) G /// respectively ; in other words the group Fe(CN) 6 has four negative charges in the ferro- and three in the ferri- cyanides. The solutions give none of the reactions for iron salts (so that Fe" and Fe"* ions are absent) and are not poisonous (absence of CN' ior). 548 A TEXT-BOOK OF INORGANIC CHEMISTRY Potassium ferrocyanide and ferricyanide are used to distinguish between ferrous and ferric salts. Potassium ferrocyanide gives with a ferrous salt a light blue precipitate of ferrous ferrocyanide, Fe 2 "Fe(CN) G "" ; with a ferric salt a dark blue precipitate of ferric ferrocyanide, Fe 4 '"3[Fe(CN) c ]"", known as Prussian blue. Potassium ferricyanide gives with ferrous salts a dark blue precipitate of ferrous ferricyanide, Fe 3 "2[Fe(CN) 6 ]"'; with ferric salts a dark red colour, but no precipitate is obtained. The formulae for these compounds can readily be understood when it is noted that the ferrocyanides are derived from the tetrabasic hydroferrocyanic acid, H 4 Fe(CN) 6 , the ferricyanides from the tribasic hydroferricyanic acid, H 3 Fe(CN) 6 . The blue compounds are decomposed by alkalis into the correspond- ing hydroxide and alkali ferro- or ferricyanide, the colour being destroyed. Tests for Iron The majority of the tests for iron have been mentioned in the course of the chapter. Iron salts, whether ferrous or ferric, give with ammonium sulphide a black precipitate of ferrous sulphide, soluble in dilute acids. The behaviour of ferrous and ferric salts with ammonia and with potassium ferrocyanide and ferricyanide is characteristic. Ferric salts give with an alkali thiocyanate a deep red coloration, due to ferric thiocyanate, Fe(CNS) 3 . COBALT Symbol, Co. Atomic weight = 58. 97. Occurrence Cobalt is found free, along with iron, nickel, and other metals, in meteorites. In the combined state, it occurs mainly as smaltite, CoAs 2 , and cobalt glance, CoAsS. Preparation of Metal The metal is obtained by reducing one of the oxides in a current of hydrogen, or better, by reducing the oxide with aluminium powder (Goldschmidt's process). Properties Cobalt is a white, lustrous, tenacious metal capable of taking a high polish ; its density is 8.5. It melts at 1460, about 40 below iron. It is malleable, and when heated is ductile. It is magnetic, but not so strongly so as iron. It is practically unaffected by air at room temperature, but on heating strongly forms the oxide Co 3 O 4 . Hydrochloric and sulphuric acids dissolve it very slowly, but it is rapidly dissolved by nitric acid, with formation of cobaltous nitrate, Co(NO 3 ) 2 . COBALT 549 OXIDES AND HYDROXIDES OF COBALT Three oxides of cobalt are definitely known : (a) Cobaltous oxide, CoO ; (b) cobaltic oxide, Co 2 O 3 ; and (c) cobalto-cobaltic oxide, Co 3 O 4 . Cobalt dioxide, CoO 2 , has also been described, but does not appear to have been obtained in a pure condition. Like iron, cobalt forms two series of salts, cobaltous salts, in which it is divalent, and cobaltic salts, in which it is trivalent. The cobaltic salts are, however, extremely unstable. Cobaltous Oxide (cobalt monoxide], CoO, is obtained as a green powder by heating the hydroxide or carbonate in absence of air. It is stable in the air at room temperature, but at a red heat takes up oxygen and forms cobalto-cobaltic oxide, Co 3 O 4 . Cobaltous Hydroxide, Co(OH) 2 , is obtained as a red pre- cipitate by adding an alkali hydroxide to the solution of a cobaltous salt and then heating to decompose the basic salt first obtained. The hydroxide soon turns brown in the air owing to oxidation. It is insoluble in excess of alkali hydroxide, but dissolves in ammonia, with formation of complex compounds (see cobaltammine compounds, below). Cobaltic Oxide, Co 2 O 3 , is obtained as a black powder by cautiously heating cobaltous nitrate. At a red heat it loses a little oxygen, foiming the compound Co 3 O 4 . Cobaltic oxide dissolves in sulphuric or hydrochloric acid in the cold, and the solutions pre- sumably contain unstable cobaltic salts. On heating the sulphate solution oxygen is given off, whilst under the same circumstances the chloride solution gives off chlorine ; and in both cases cobaltous salts remain in solution (cf. manganese dioxide, p. 5 2 ^)- Cobaltic Hydroxide, Co(OH) 3 , is formed as a dark pre- cipitate by the action of sodium hypochlorite on cobaltous hydroxide. Cobalt o-Cobal tic Oxide, Co 3 O 4 , the oxide usually met with in commerce, occurs as a black powder. With acids it behaves like a mixture of cobaltous and cobaltic oxides. COBALTOUS SALTS The cobaltous salts are the only stable salts of cobalt. The solid hydrates a -e pink, and they form pink solutions at room temperature ; the anhydious salts are generally blue. The cause of the differences in colour are doubtless due to differences in constitution referred 'to later in connexion with the complex ammines (p. 551). 550 A TEXT-BOOK OF INORGANIC CHEMISTRY CobaltOUS Chloride, CoCl 2 , is obtained by dissolving any of the oxides of cobalt in hydrochloric acid, and separates on evapo- rating the solution in red monoclinic crystals as CoCl 2 ,6H 2 O. On heating to 120 all the water is driven off, and the resulting anhydrous salt is blue. The pink aqueous solution also turns blue on heating, but regains its original colour on cooling. Advantage is taken of this behaviour in preparing the so-called sympathetic ink. Writing done with an aqueous solution of cobalt chloride is practically invisible, but when the paper is warmed the blue anhydrous salt is formed, and the writing becomes visible. It fades away again on cooling, owing to rehydration of the salt. Cobaltous Sulphate, CoSO 4 ,7H 2 O, prepared by the usual methods, occurs in dark red monoclinic crystals, isomorphous with FeSO 4 ,7H 2 O. It can be dehydrated by heating, and the anhydrous salt is also red. It forms double salts with the alkali sulphates ; for example, CoSO 4 ,K 2 SO 4 ,6H 2 O, isomorphous with the correspond- ing iron salt. CobaltOUS Nitrate, Co(NO 3 ) 3 ,6H 2 O, forms red hygroscopic prisms, and is very soluble in water. It gives cobaltic oxide when cautiously heated. Cobaltous Sulphide, CoS, is obtained as a black amorphous precipitate by the addition of ammonium sulphide to a solution of a cobalt salt. When hydrogen sulphide is passed through a cobalt solution acidified with hydrochloric acid no precipitate is formed, but the sulphide obtained by precipitation in alkaline solution is practically insoluble in dilute hydrochloric acid. This curious be- haviour is not entirely understood. Sulphides of the formulas Co 2 S 3 , Co 3 S 4 and CoS 2 have also been described. COBALTIC SALTS As already mentioned, simple salts containing trivalent cobalt are highly unstable. Cobaltic Sulphate, Co 2 (SO 4 ) 3 , is formed in solution at the anode when a concentrated solution of cobaltous sulphate, CoSO 4 , is electrolyzed. It has been obtained in the solid form as Co 2 (SO 4 ) 3 ,i8H 2 O. The aqueous solution is dark green. Many complex compounds containing trivalent cobalt are known, and will now be briefly referred to. COBALT 551 COMPLEX SALTS CONTAINING COBALT Complex Cyanides The complex salts, K 4 Co(CN) 6 and K 3 Co(CN) 6 , corresponding with potassium ferro- and ferricyanide, are known. When excess of potassium cyanide is added to a cobaltous salt in the cold, a red solution is obtained, which, on evaporation at room temperature, yields potassium cobaltocyanide in violet crystals : + 6KCN-K 4 Co(CN) 6 When the red solution is boiled it becomes decolorized, hydrogen is given oft", and on evaporation the stable potassium cobalticyanide, K 3 Fe(CN) 6 , is obtained in colourless crystals. Complex Nitrite When potassium nitrite and acetic acid are added to the solution of a cobaltous salt, the red colour disappears, and in course of time the cobalt is completely precipitated in the form of a yellow crystalline substance, of the formula K 3 Co(NO 2 ) 6 , potassiiuh cobaltinitrite. The complex anion has the formula Co(NO 2 ) ( '". The salt can be classified from the point of view ot Werner's theory (p. 561). Cobaltammines When to a solution ot a cobaltous salt excess of ammonia is added, and the resulting solution is treated with an oxidizing agent, e.g. the oxygen of the air, a great variety of complex salts are obtained, the empirical composition of which corresponds with that of a cobaltic salt associated with 3, 4, 5 or 6 NH 3 molecules. The compounds, however, unlike the cobaltic salts, are quite stable, and therefore do not contain simple Co"' ions ; further, in some of them only part of the anion shows the reactions characteristic of it, for example, only part of the chlorine is precipitated by silver nitrate. In these and other respects the compounds resemble the chromium ammonia compounds, and they can be represented as follows : [Co(NH 3 ) ]-Cl 3 ; [Co(NH 3 ) 6 Cl]"Cl 2 ; [Co(NH 3 ) 4 Cl 2 ]'Cl ; [Co(NH 3 ) 3 Cl 3 ]. The complex enclosed in the square brackets is the cation. In the first compound it is trivalent, and all the chlorine atoms are ionised ; in the socond compound the cation is divalent, and in the third univalent. Finally, in the fourth compound it consists of a non- ionised complex, which gives none of the ordinary reactions for cobalt and chlorine ions. The behaviour of these compounds is satisfac- torily accounted for on Werner's theory (p. 561). 552 A TEXT-BOOK OF INORGANIC CHEMISTRY Tests for Cobalt All cobalt salts communicate an intense blue colour to a borax bead. The formation of a black precipitate with ammonium sulphide, insoluble in dilute hydrochloric acid, coupled with the fact that hydrogen sulphide in acid solution causes no precipitate, is characteristic of nickel and cobalt salts. The reactions whereby nickel and cobalt salts may be distinguished are referred to under nickel. NICKEL Symbol, Ni. Atomic weight, 58.68. Occurrence Nickel is found free in meteorites. In combina- tion it is met with in kupfernickel or niccolite, NiAs, in nickel glance, NiSAs, and as a silicate in garnierite, H 2 (Mg,Ni)SiO 4 ,2H 2 O. The ore last mentioned occurs in large quantity in New Caledonia, and is now the chief source of nickel. Preparation of Metal The metal may be obtained from the oxide by reducing with carbon or hydrogen, or by the Goldschmidt process with aluminium powder. Very pure nickel is now obtained commercially by Mond's process, which depends on the intermediate formation of nickel carbonyl, Ni(CO) 4 (p. 319). Nickel oxide (obtained by roasting nickel ores) is heated and carbon monoxide passed over it, whereby reduction to the metal first takes place, the latter then uniting with carbon monoxide to form the volatile nickel carbonyl. The vapour of the carbonyl is then passed through tubes at a higher temperature (about 200 C.), whereby it is decomposed, and nickel is deposited in coherent form. Properties Nickel is a silvery-white, lustrous metal of density 8.8 to 9. i ; it melts at 1435. It is malleable and ductile, and also hard and tenacious. It is magnetic, but loses this property on heating to 250. Like iron, it becomes passive on treatment with oxidizing agents. Nickel is practically unaffected by exposure to air. It is attacked very slowly by hydrochloric or sulphuric acid, but dissolves fairly readily in dilute nitric acid, nickel nitrate, Ni(NO 3 ) 2 , being formed. Alloys Of Nickel Nickel is a constituent of many useful alloys. German silver usually contains 50 per cent, of copper, 25 per cent, of nickel, and 25 per cent, of zinc, but the proportions vary to some extent. The nickel coins in use on the Continent and in the United States are made of an alloy of 75 per cent, of copper and 25 per cent, of nickel. It is noteworthy that in spite of the high NICKEL 553 proportion of copper this alloy is white. Nickel steel has already been mentioned (p. 542). Further, on account of its permanence in the air, rv'ckel is used as a coating for other metals. Nickel plating, as it is called, is done electrolytically, a solution of nickel sulphate or nickel ammonium sulphate being used as electrolyte and a nickel plate as anode. OXIDES AND HYDROXIDES OF NICKEL Three oxides of nickel are known : nickelous oxide, NiO, nickelic oxide, Ni.O 3 , and nickelo-nickelic oxide, Ni 3 O 4 . Only one series of nickel salts is known, corresponding with nickelous oxide, NiO. Nickelic oxide behaves towards acids as a peroxide. Nickelous Oxide, NiO, is obtained as a grayish-green powder by heating nickelous hydroxide out of contact with air. Nickelous Hydroxide, Ni(OH) 2 , is obtained as an apple- green precipitate when an alkali hydroxide is added to the solution of a nickel salt. The hydroxide is insoluble in excess of alkali hydroxide> but dissolves in ammonia to form a blue solution. Unlike the corre- sponding cobalt solution (p. 549), the ammoniacal solution of nickel hydroxide does not absorb oxygen from the air. It contains complex ions of the type Ni(NH 4 )^, in which the nickel is divalent. They are much less stable than the cobaltammines, which, however, are derived from tervalent cobalt. Nickelic Oxide, Ni 2 O 3 , is obtained as a black powder by decomposing nickelous nitrate by heat at as low a temperature as possible. When heated with hydrochloric acid it forms nickelous chloride and chlorine, whilst with sulphuric acid it gives nickelous sulphate and oxygen : It therefore behaves as a peroxide, and appears to have no basic properties. Nickelic Hydroxide, Ni(OH) 3 , is obtained as a black pre- cipitate "jy adding an alkali hypochlorite or a solution of bleaching powder ;o the solution of a nickel salt. Chemically it behaves like the com spending oxide. Nickelo-nickelic Oxide, Ni 3 O 4 , is said to be formed as a gray mass when moist oxygen is passed over nickel chloride 554 A TEXT-BOOK OF INORGANIC CHEMISTRY heated at 400. On heating strongly oxygen is given off and the monoxide, NiO, is formed. SALTS OF NICKEL As already mentioned, the salts of nickel are derived exclusively from the monoxide, NiO, and therefore contain the bivalent Ni" ion, which is green. The hydrated salts and solutions are green, the anhydrous salts yellow or brown. Nickel Chloride, NiCl 2 , is obtained by dissolving the oxide in hydrochloric acid, and separates from solution with 6H 2 O in green monoclinic prisms. When heated the anhydrous salt, which is yellow, is obtained. The latter absorbs gaseous ammonia, forming the compound Ni(NH 3 ) 6 Cl 2 , which separates from aqueous solution in blue, octahedral crystals. Nickel Sulphate, NiSO 4 , obtained by dissolving the metal or the oxide in dilute sulphuric acid, separates from solution below 20 with 7H 2 O in green, rhombic prisms, isomorphous with the corresponding ferrous and other sulphates. At a higher temperature (30 to 40) green, tetragonal crystals of the formula NiSO 4 ,6H 2 O are obtained. When heated at 100 it loses 6H 2 O, above 300 it becpmes anhydrous. With excess of ammonia a violet solution is obtained, from which on evaporation violet crystals of the formula NiSO 4 ,4NH 3 ,2H 2 O are obtained. This salt is presumably [Ni(NH 3 ) 4 (H 2 O) 2 ]SO 4 ; many other instances in which water displaces ammonia in such complexes are known. With alkali sulphates, nickel sulphate forms double salts, such as NiSO 4 ,(NH 4 ) 2 SO 4 ,6H 2 O, isomorphous with the corresponding iron salt. This salt is largely used in electro-plating. Nickel Sulphides Nickelous sulphide, NiS, is obtained as a black precipitate by adding ammonium sulphide to a solution of a nickel salt. Like the corresponding cobalt salt, it is practically insoluble in dilute hydrochloric acid, but is soluble to some extent in ammonium sulphide solution. The sulphides Ni 3 S 4 and NiS 2 have also been described, Nickel Cyanide, Ni(CN) 2 , is obtained as a green precipitate by adding potassium cyanide to a solution of a nickel salt. The pre- cipitate dissolves in excess of cyanide to form the compound K 2 Ni(CN) 4 , which can be obtained as the monohydrate, in yellow crystals, by evaporating the solution. Unlike the nearly analogous compound containing divalent cobalt, it is stable on boiling, but is much less stable than potassium cobalticyanide (see tests for nickel). THE IRON SUB-GROUPGROUP VIII 555 Tests for Nickel Nickel salts colour the borax bead yellowish brown in the oxidizing flame. The behaviour of the hydroxide and of the sulphide is made use of in detecting nickel. Nickel and cobalt compounds are distinguished by the facts (i) that only the latter give a yellow precipitate with acetic acid and potassium nitrite ; (2) when sodium hypochlorite or hypobromite is added to a solution of a nickel salt in excess of potassium cyanide a black precipitate of nickelic hydroxide is obtained, whereas a solution, of a cobalt salt in potas- sium cyanide after boiling gives no precipitate under these conditions. Summary of Metals of Iron Sub-group Corresponding with the fact that the members of this family differ only slightly in atomic \\eight, their physical constants do not differ very greatly. This is shown in the accompanying table : Iron. Cobalt. Nickel. Atomic weight 55.85 58.97 58.68 Densitv .... 7.86 8. 5 8.9 Melting-point . 1500 1460 1435 Atomic volume 7-i 6.9 6.6 Further, all three metals are grayish-white, magnetic, and readily become massive when treated with concentrated nitric acid. From the chemical point of view, all three metals form oxides of the type MO, which are strongly basic and give rise to stable series of salts. Further, all have oxides of type M 2 O 3 , but corresponding salts are only known for iron and cobalt ; nickel salts in which the metal is trivalent are unknown. Only iron forms salts corresponding with a (hypothetical) acidic oxide, FeO 3 , so that cobalt and nickel are more metallic than iron. The latter two elements resemble each other very closely in their compounds ; the chief difference is the greater ten- dency of cobalt to act as a trivalent element, as shown, for instance, in the cobaltammines. Iron has comparatively little tendency to enter into the formation of complex ions, and in this important respect differs from cobalt and nickel. As regards comparison with members of other groups, there is a close analogy, especially in the case of iron, with chromium, man- ganese .md copper. Reference to this has already been made in connexion with the latter metals. CHAPTER XXXVI THE PLATINUM SUB-GROUP (GROUP VIII) Ruthenium, Ru=ioi.7 Rhodium, Rh=io2.6 Palladium, Pd = io6.7 Osmium, Os =190.9 Iridium, Ir =193.1 Platinum, Pt =195.2 r T^HE remaining elements of the eighth group are naturally divided JL into two smaller groups containing three elements each. Just as iron, cobalt and nickel differ only slightly in atomic weight and in density (p. 555), so the atomic weights of the triad ruthenium, rhodium and palladium vary from 101.7 to 106.7, and their densities from 11.4 to 12.26; the atomic weights of the triad osmium, iridiuin and platinum vary from 190.9 to 195.2, and their densities from 21.5 to 22.5. As in the remainder of the periodic table, however, there is also considerable resemblance between the elements in the same vertical series ; thus the pairs ruthenium and osmium, rhodium and iridium, and palladium and platinum, show considerable chemical analogy. The six metals occur in the free state in nature, associated in alluvial deposits in the so-called platinum ore, which contains 70 to 80 per cent, of platinum, 5 to 8 per cent, of iridiuin, and a smaller proportion of the other metals. Gold and copper are also generally present. The platinum ores are found chiefly in the Urals and in Brazil. They are separated from the sand, etc., by washing, advan- tage being taken of their relatively great density. As platinum is the only important member of the group, the others will be described very briefly. Ruthenium (Ru= 101.7) is a steel-gray metal of density 12.26; it melts about 2000. When heated in the air it becomes covered with a brown film of the dioxide, and when heated in a current of oxygen the dioxide, RuO 2 , is formed. It is only slightly attacked by aqua regia, and is chiefly obtained from the alloy osmiridiitm, which remains when platinum ores are treated with aqua regia. Ruthenium forms the oxides, Ru 2 O 3 , RuO 2 and RuO 4 , and also salts corre- sponding with the acidic oxides, RuO 3 and Ru 2 O 7 . A salt corresponding with the trioxide is potassium ruthenate, K 2 RuO 4 ; when diluted it decomposes into potassium perruthenate, KRuO 4 , and an oxide of ruthenium (cf. manganates and 556 THE PLATINUM SUB-GROUPGROUP VIII 557 permanganates): The most stable ruthenium salts contain tervalent ruthenium, but compounds in which the metal is bivalent and quadrivalent are also known. Osmiuir (03 = 190.9) is a bluish-white crystalline metal ; its density is 22.5, and it melts about 2300. On account of its difficult fusibility, it is used as a filament in incandescent lamps. On heating in air, it burns readily to the easily volatile tetroxide, which is very poisonous, and has a highly injurious effect on the eyes. The metal is scarcely attacked by aqua regia. Osmiurr forms four oxides, OsO, Os 2 O 3 , OsO 2 andOsO 4 ; and salts, the osmates, are known derived from the unknown acidic oxide, OsO 3 . The best known is potassium :>smate, K 2 OsO4,2H 2 O, which occurs fn garnet red octahedral crystals. The most emarkable compound of osmium is the tetroxide, OsO 4 , usually known as osmic a Mel, although it does not appear to possess acidic properties. It occurs in colourless, lustrous needles, which melt about 40; the liquid boils about 100. It is used for staining specimens in biological work, the oxide being reduced, most readily by fats, to metallic osmium. It is in cresting to note that, corresponding with their position in the eighth group of the periodic table, ruthenium and osmium have a maximum valency of 8. Rhodium (Rh=io2.6) is a silver-white, lustrous, malleable metal of density 12.1, which fuses at a higher temperature than platinum. It is not attacked by aqua regia. Three oxides of rhodium are known, RhO, Rh 2 O 3 and RhO.>. The most stable salts are derived from the oxide Rh 2 O 3 . Unlike ruthenium and osmium, rhodium forms no compounds derived from an acidic oxide. The rhodium salts usually form red xqueous solutions, hence the name. The trichloride, RhCl 3 , forms stable compounds with the alkali chlorides, e.g. Na 3 RhCl 6 (or RhCl 3 ,3NaCl). Iridium (Ir= 193.1) is a white, lustrous, brittle metal of density 22.4. Next to osmium, it has the highest melting-point of the platinum metals. In the free condition it is not attacked by aqua regia, but when alloyed with much platinum it is dissolved to some extent. On account of their high melting-point and resist- ance to many reagents, alloys of platinum and iridium are used for various purposes. Iridium forms two basic oxides, Ir 2 O 3 and IrO 2> from which two series of salts are derived. The chlorides, IrCl 3 and IrC^, form complex compounds with the alkali chlorides of the types K 3 IrCl 6 or IrCl 3> 3KCl, and K 2 IrCl 6 or IrCl 4 ,2KCl, which conespond with the better known platinum compounds. Palladium (Pd = 106.7) i s found in platinum ore and also, alloyed with gold, in certain pa^ts of South America. It is a silvery-white, lustrous, malleable metal of density IT. 9, and melts at 1546. The most remarkable property of the metal is its power of absorbing hydrogen, already referred to (p. 39). The freshly ignited metal absorbs 600 times its own volume of hydrogen under ordinary conditions, and when made the cathode of an electrolytic cell at which hydrogen is being generated it takes up about TOOO volumes of the gas. The hydrogen is completely expelled on heating to redness. It does not appear to be chemically combined with the palladium, but is present in solid solution. The hydrogen absorbed in palladium has powerful reducing properties, clue partly to the catalytic effect of the palladium on the speed of reduction and partly to the very condensed condi- tion of the hydrogen. Two oxides of palladium, PdO and PdO 2 , are known. Well-defined palladium 558 A TEXT-BOOK OF INORGANIC CHEMISTRY salts are derived exclusively from the first-mentioned oxide. When palladium is dissolved in aqua regia, palladium tetrachloride is present in solution as H 2 PdCl 6 or PdCl 4 ,2HCl ; it has not been obtained in the free condition. When the solution is boiled, chlorine is given off and palladious chloride, PdCl 2 (or rather H 2 PdCl 4 ), is formed. PLATINUM Symbol, Pt. Atomic weight, 195.2. Preparation Of Metal The platinum ore is treated with dilute aqua regia, whereby platinum, palladium, rhodium and iridium are dissolved as the higher chlorides. The solution is evaporated to dryness and the fused mass heated at 125, at which temperature the palladium and rhodium salts are reduced to lower chlorides. The residue is extracted with water, acidified with hydrochloric acid and ammonium chloride added, when ammonium platinic chloride, PtCl 4 ,2NH 4 Cl, is precipitated in yellow crystals (see below). From the solution the corresponding iridium salt, IrCl 4 ,2NH 4 Cl, which is more soluble, is obtained by concentrating the solution. The double platinum salt, on ignition, yields the metal in spongy form. Platinum is obtained in lustrous, coherent form by fusing the spongy metal in a lime crucible by means of the oxyhydrogen flame. Properties Platinum is a silvery-white, malleable and ductile metal of density 21.42. It fuses at 1750, and therefore does not melt in the Bunsen flame, but readily fuses in the oxyhydrogen flame ; it can be welded at a white heat. Massive platinum does not oxidize in the air even at high temperatures. The metal is insoluble in any single acid, but is readily dissolved by aqua regia with formation of platinum tetrachloride. It forms alloys with lead and antimony, and certain other heavy metals, and combines with carbon, silicon and phosphorus to form brittle alloys. Further, it is attacked by fused alkalis, cyanides, nitrates and sulphides. These facts should be borne in mind in using platinum utensils (crucibles, wire, etc.), which, on account of the slight chemical activity of the metal, are largely employed in the laboratory. Platinum condenses gases, such as oxygen and hydrogen, on its surface, especially in the finely divided (spongy) form. Partly on this account, the metal has a remarkable effect in accelerating many chemical reactions, such as the combination of inflammable gases (hydrogen, marsh gas, benzene, etc.), with oxygen. A mixture of oxygen and hydrogen can be brought to explosion at room tempera- ture by bringing it in contact with finely divided platinum. In recent PLATINUM 559 years it has been found that the finely divided metal oxidizes much more readily than was formerly supposed, and it is not improbable that the intermediate formation of an active platinum oxide plays a part in these phenomena. Spongy platinum, as already explained, is obtained by heating ammonium platinic chloride ; platinum black, another finely divided form of tl e metal, is obtained by electrolysis of a solution of platinic chloride, or by adding" finely divided z;nc to the latter solution. These forms of platinum have very high absorptive power for hydrogen ; platinum black absorbs about 100 times its volume of the gas. Platinum black also absorbs about 100 times its volume of oxygen, but in this case partial chemical combination probably occurs. Platinum Compounds Two series of platinum compounds are kno\\ n ; platinous compounds, e.g. PtCl 2 , in which the metal is divalent, and platinic compounds, e.g. PtCl 4 , in which the platinum is quadrivalent. Oxides and Hydroxides of Platinum Platinous oxide, PtO, is obtained as a gray powder, and platinic oxide, PtO 2 , as a black powder, by cautiously heating the corresponding hydroxides. When strongly heated, they lose oxygen and yield the metal. Platinous Hydroxide, Pt(OH) 2 , is obtained as a dark pre- cipitate by the action of sodium hydroxide solution on platinous chloride. It is a typical base and dissolves in acids to form platinous salts. Platinic Hydroxide, Pt(OH) 4 , is obtained as the dihydrate, Pt(OH) 4 ,2H 2 O, or H 2 Pt(OH) , by boiling platinic chloride with sodium hydroxide solution, cooling and neutralizing with acetic acid. It is a yellow powder, which on heating turns brown and then black. It has weak basic and also weak acidic properties ; with potassium hydroxide it gives potassium platinate, K 2 PtO 3 . Platinous Chloride, PtCl 2 , is obtained by heating platinic :hloride, PtCl 4 , to 300. It forms a grayish-green powder, insoluble in water but soluble in hydrochloric acid. With the alkali chlorides it forms double salts, e.g. PtCl 2 ,2KCl and PtCl 2 ,2NH 4 Cl, both of which are red, well crystallized salts. On account of their behaviour, it is assumed that these salts contain the complex anion PtCl/', and they are therefore termed chloroplatinites, (NH 4 ) 2 PtCl 4 , and K 2 PtCl 4 , derived from chloroplatinous acid, H 2 PtCl 4 . The latter compound is presumably present in the solution of the dichloride in hydrochloric acid. 560 A TEXT-BOOK OF INORGANIC CHEMISTRY Platinum Trichloride, PtCl 3 , is obtained as a green powder by prolonged heating of platinic chloride in dry chlorine at 320 (Wohler and Martin, 1909). Platinic Chloride, PtCl 4 , is obtained by the direct action of chlorine on platinum at a high temperature. It is obtained more conveniently, however, by dissolving platinum in aqua regia, evaporat- ing to dryness, and heating the residue in a current of hydrogen chloride. It separates from aqueous solution in large red crystals as PtCl 4 ,5H 2 O. When platinum is dissolved in aqua regia and the solution is evaporated to dryness several times with hydro- chloric acid till all the nitric acid is expelled, the compound PtCl 4 ,2HCl or H 2 PtCl 6 , usually termed chloroplatinic acid, is obtained. It separates from aqueous solution in brownish-red prisms with 6H 2 O. Platinic chloride forms double compounds with the alkali chlorides, which may be regarded as being derived from chloroplatinic acid by displacing the hydrogen by metals. Owing to their different solu- bilities in water, they are used in analytical chemistry. Potassium Platinichloride, K 2 PtCl , obtained as a yellowish crystalline precipitate by adding a potassium salt to a solution of chloroplatinic acid, is only slightly soluble in water. At o 100 grams of water dissolve 0.70 grams, at 20 1.12 grams, and at 40 1.76 grams of the salt. : Potassium platinichloride is almost insoluble in alcohol. The corresponding ammonium salt is slightly, the sodium salt readily soluble in water, and on this is based a method of separating potassium and sodium in analysis. Platinocyanides and Platinicyanides Corresponding with the platinochlorides, platinocyanides, derived from platino- cyanic acid, H 2 Pt(CN) 4 , are known. The acid and the salts are characterized by remarkable differences in colour depending on the proportion of water with which they are associated, and further, the crystals show different colours in different directions. Platini- cyanides, derived from platinicyanic acid, H 2 Pt(CN) 6 , are also known. Platinum Sulphides Platinous sulphide, PtS, and platinic sulphide, PtS 2 , are obtained as black precipitates by passing- hydrogen sulphide through solutions of the respective chlorides. They are insoluble in hydrochloric acid, but dissolve in solutions of alkali sulphides, forming the salts. Platinum Ammonia Compounds Like salts of cobalt, chromium and other metals, platinum salts form compounds with ammonia which can be classified MODERN VIEWS ON VALENCY 561 on the b; sis of Werner's theory of valency (see below). Thus the following series of compounds containing divalent platinum is known : O c!H K '[ pta *]" K - All the compounds contain a special group, PtA 4 , the character of which varies with the n iture of the groups A. In the first two members it is basic, in the middle member it is neutral, and in the two last members it is acidic. The further behaviour of these compounds will be understood on the basis of the explana- tions give i under " molecular compounds " (below). The conpound Pt ^t? 3 ' 2 exists in two entirely different forms which differ in many of their reactions. This cannot be accounted for by formulae written on the plane of the paper, but is readily explained on the assumption of a different arrangement of the atoms in space. Another series of compounds, containing quadrivalent platinum is known : [Pt(NH ;s ),]--Cl 4 ; [Pt(NH 3 ) 5 ClJ--Cl 3 ; [Pt(NH 3 ) 4 Cl. 2 ]" C1 2 ; [Pt(NH 3 )Cl 3 ]' Cl ; [Pt(NH 3 ) 2 Cl 4 ]; [Pt(NH 3 )Cl 5 ]' K ; [PtCl 6 ]" K 2 . All these compounds contain a special group, PtA 6 , which in the first members is electro- positive, but gradually diminishes in valency as the number of Cl atoms in it increases, till it becomes electrically neutral (compound five), and finally elt ctro-negative. Werner's theory of valency, which affords a satis- factory n presentation of the behaviour of such compounds, will now be considered. "Molecular Compounds." Modern Views on Valency The study of the compoum s which ammonia forms with chromium, cobalt and platinum salts has thrown considerable light on the question of valency. In earlier chapters it has been poiired out that the valency of most of the elements is variable, depending on the nature of the other elements in the compound. Thus nitrogen is trivalent towards the halogens, but towards oxygen it shows several valencies. Further, one element may show more than one valency towards another ; thus iron behaves towards chlorine as a divalent and as a trivalent element. Although the old theory of valency gives moderately satisfactory results when applied to compounds of elements and groups which do not exist by themselves, serious difficulties arise when we come to deal with compounds between mole- cules whi^h are capable of independent existence. In some instances the difficulties have been overcome ; thus the compound formed by the combina- tion of sulphur trioxide, SO 3 , and water, H 2 O, is represented satisfactorily as H a rearrangement of valencies having occurred. C In othc; - cases, however, e.g. the compound Pt(NH 3 ) 6 Cl 3 , the representation according to the ordinary valency scheme presents serious, if not insuperable difficulties. A theory which we owe chiefly to Werner admits of a fairly satis- 36 562 A TEXT-BOOK OF INORGANIC CHEMISTRY factory representation of such compounds. It has already been stated that the following series of ammonia-platinic compounds is known : [Pt(NH 3 ) 6 ]--Cl 4 ; [Pt(NH 3 ) 5 Cl]-Cl 3 ; [Pt(NH 3 ) 4 Cl 2 ]"Cl 2 ; [Pt(NH 3 ) 3 Cl 3 ]' Cl ; [Pt(NH 3 ) 2 Cl 4 ]; [Pt(NH 3 )Cl 3 ]'K ; [PtCl 6 ]"K 2 . The first compound has a high electrical conductivity in solution, and all the chlorine atoms are readily precipitated by silver nitrate. In the third compound, on the other hand, the electrical conductivity is much smaller, and only two of the chlorine atoms react with silver nitrate ; finally, the fifth compound has practically no electrical conductivity in solution, and the chlorine does not react with silver nitrate. Werner accounts for this behaviour on the view that the groups within the square brackets are directly bound to the platinum "in the nner zone " ; chlorine bound in this way behaves like chlorine in an organic compound, and does not give the ordinary reactions ; in other words, it is non- ionised. The atoms in the outer zone, on the other hand, are less intimately bound to the central atom and are ionised. The number of groups attached directly to the central atom is known as the co-ordination number of the atom, and is determined from the behaviour of the compound, as in the example quoted. It is a remarkable fact that in a great many cases the co-ordination number is six. Examples from chromium, cobalt and platinum compounds have already been given. It will be noticed that the presence of ammonia molecules in the inner zone does not alter the valency of the central atom. Thus the valency of the complex [Co(NH 3 ) 6 ] is three, corresponding with the valency of the simple Co"' ion ; the only effect of the addition of the ammonia atoms is to increase the stability of the compounds containing trivalent cobalt (p. 551). The presence of water and of certain other groups is also without effect on the valency of the complex- cation. The NH 3 groups in the compound [Co(NH 3 ) 6 ]Cl 3 can be gradually displaced by H 2 O molecules, down to Co(H 2 O) 6 Cl 3 , the complex group contain- ing the central atom remaining trivalent throughout. On the other hand, for every NH 3 molecule displaced by a strongly negative univalent atom or group, the positive valency of the central group is diminished ty one unit. This is shown very clearly in the formulae of the platinum compounds already quoted. When four Cl atoms have entered the inner sphere, the four positive valencies of the platinum are neutralized and the compound is electrically neutral ; when a fifth chlorine atom enters, the group has become negative and univalent, with two Cl atoms it is negative and divalent. So far we have not discussed the nature of the valencies in such compounds, but this is a matter of fundamental importance. Werner distinguishes two kinds of valencies (i) the principal valencies, which are measured by the number of hydrogen atoms the element or group can combine with or displace, (2} the subsidiary valencies, which bind groups which are capable of existing as indepen- dent molecules. It is evident that the principal valencies correspond with those ordinarily accepted, which bind atoms or groups such as Cl, NO 2 , CH 3 . The subsidiary valencies, on the other hand, bind groups such as water and ammonia, which cannot be represented on the ordinary theory of valency. A further difference between the two kinds of valency is that some at least of the groups joined to the central atom by principal valencies can ionise, whereas groups joined by subsidiary valencies are incapable of ionisation. MODERN VIEWS ON VALENCY 563 It should be stated, however, that there is no perfectly sharp distinction between the two k nds of valency. It will be evident from the foregoing that the co-ordinakd groups may be joined either by principal or by subsidiary valencies ; in the con pound [Co(NH ;! ) 6 ]Cl 3 , for instance, they are all connected to the central atom by subsidiary valencies. It is a very remarkable fact that the co-ordination number is six for so many elements, largely independent of the nature of the co-ordinated groups, and Werner is of opinion that it is more a question of space than of affinity ; six is the max mum number of groups that can be^fitted round the central atom. The co-ordination number of carbon is 4, corresponding with its ordinary valency. The connexion between these views and the conception of affinity will be evident. n general, there will be a greater neutralization of affinity when atoms or r;roups are connected by principal valencies than when a group is connected by subsidiary valencies. The hydrogen and oxygen in water are connected by principal valencies, and there is a considerable neutralization of affinity. The neutralization is not, however, complete ; some free affinity remains, sufficient to bind the molecule of water (by a subsidiary valency bond) to another atom or molecule. Space does not admit of a fuller treat- ment of this interesting subject here ; for full details Werner's Neuere Anschaminjen anf deni Gebiete der anorganischen Chemie 1 should be consulted. 1 English Translation by Hedley (Longmans, Green & Co.), 1911. CHAPTER XXXVII RADIO-ACTIVITY ONE of the most striking properties of the Rontgen or X-rays is the fluorescence to which they give rise on screens of certain materials. In the course of an examination of a number of phosphorescent and fluorescent substances, under- taken in order to find out if they give rise to radiation of any kind, Becquerel, in 1896, discovered that uranium salts affected a photographic plate through several layers of black paper. This effect was clearly due to radiation from the uranium. The Becquerel rays, as this type of radiation came to be called, had the further property of rendering the air in their neighbourhood a conductor of electricity and therefore of discharging a gold leaf electroscope. 1 Substances giving out rays of this type are said to be radio-active. A little later, M. and Mme. Curie found that several minerals containing uranium, notably pitchblende from Austria, were found to be considerably more ra,dio-active than pure uranium compounds or even than uranium itself. It therefore appeared probable that the radio-activity was not due to uranium, but to some other constituent in uranium ores, and M. and Mme. Curie set themselves the task of isolating the substance. Ultimately, two very active substances were obtained ; one, closely allied in its chemical characters to bismuth, was named polonium (from Poland, Mme. Curie's native country), the other, closely allied chemically with barium, was termed radium. The amount of radium present in an average sample of pitchblende is extremely small, not exceeding i part in 10 million parts of the ore, and polonium occurs in still smaller proportion (not exceeding i part in 10,000 million parts of pitchblende). Pure radium salts are about 2 million times more radio-active than uranium compounds. The possibility of isolating elements present in such small proportion is due to the great readiness with which radio-activity can be detected. Up to that time the use of the spectroscope was the most sensitive means of detecting minute quantities of substances, but those substances which are radio-active can be detected in much more minute quantities. As little as io- 12 gram of radium can be detected in this way. At a later period, a third radio-active substance, called actinium, was found in pitchblende. Neither polonium nor actinium appear to have been obtained pure. Early in her investigations, Mme. Curie discovered that thorium com- pounds also show radio-activity. Properties of Radium Radium is a well-defined element, of atomic weight 226, and it occupies a position in the periodic table as the last member of the 1 A gold leaf electroscope consists essentially of two insulated gold leaves (sus- pended close together), which diverge when electrically charged, but fall together when the air in contact with them becomes conducting. 564 RADIO-ACTIVITY 565 barium group. All the work on radio-activity has hitherto been done with radium chloride, RaCl 2 , and radium bromide, RaBr 2 , but quite recently (1910) the element itself was isolated by Mme. Curie and Debierne. They subjected a solution of radium chloride to electrolysis with a mercury cathode, distilled off the mercury, c.nd obtained radium as a brilliant white metal, which melted sharply at 700 an J volatilized slightly at the same temperature. Like the other metals of the alkaline earths, radium is readily attacked by water, hydrogen being evolved, and it tan ishes rapidly in the air. Investigation of the radiation given off by radium (and by uranium ores) has shown thrt three distinct kinds of rays, known" respectively as a, /3, and y rays, can be distinguished. All of them affect a photographic plate, discharge a gold leaf electioscope, and render materials such as zinc sulphide and zinc silicate (willemitei fluorescent, but their relative activities are very different. The a rays have very little penetrating power, being completely stopped by three or lour layers of thin aluminium foil, by a sheet of paper, or by passing through s layer of air 2 to 3 cm. thick. Their electric effect is very marked. They are slightly deflected by a magnet, the direction of the deflection showing that they are positively charged. It is now generally accepted that when an a particle gives up its charges, an atom of helium remains. The 18 rays have at least 100 times the penetrating power of the a rays. Rutherford showed that on passing through 100 thin plates of aluminium their intensity vas only reduced by half. They are readily deflected by a magnet, the direction showing that they are negatively charged, and their mass is only about y-jjV, ,j of that of the hydrogen atom. It is generally agreed that the /3 par- ticles are negative electrons actual units of electricity. The y r lys have a penetrating power much greater than the /3 rays and affect a photographic plate strongly. They travel with the velocity of light. They appear to be identical with Rontgen or X-rays, and are therefore not particles, but a wave motion in the ether. Besides these rapidly-moving particles, radium compounds continuously give off a heavy vapour known as radium emanation. Thorium compounds give rise to an analogous emanation. The emanations are, however, unstable, and are only temp orarily radio-active. Radium emanation at first gives off only a rays. At a later stage it gives off/ 3 and y rays and other substances are also formed. It is usua to express the stability of a radio-active substance in terms of the time taken for its activity to fall to half its original value. The half-time period for radium emanation is about 3.85 days. Chemically, radium emanation behaves like a gas of the helium series (p. 209), and in a vacuum tube shows a characteristic spectrum of bright lines. It is unaffected by passing through a hot tube or through acids; it is condensed by passing tarough a tube immersed in liquid air, but vaporizes again when the temperature is allowed to rise. The liquid emanation boils at 71 C. It is colotirles.'- when first condensed ; when the temperature is lowered it freezes, and at fie temperature of liquid air glows with a bright rose colour. The density o ' the emanation in the form of vapour has been determined from its rate of diffusu n, and quite recently has been determined directly. The value obtained is about 120, which would bring it into the helium series, below xenon. The most surprising fact connected with the emanation is that one of the products of its decay is helium (Ramsay and Soddy). When a little of the 5 66 A TEXT-BOOK OF INORGANIC CHEMISTRY emanation is collected in a tube and its spectrum examined at intervals, no helium lines are noticeable at first, but after a time they appear and gradually become more intense. In the light of what has been mentioned as to the con- nexion between a particles .and helium, the appearance of the latter element will be understood. _ While the emanation is breaking up it gives off a particles, and when these lose their charges they become ordinary helium atoms. In the process of decay of the emanation, six further substances appear to be formed in succession, known respectively as radium A, B, C, D, E, and F. In some of these cases the change from one stage to another is attended by the emission of rays ; in other cases no rays are emitted. Disintegration Theory of Radio-activity We have now to consider whether any theory can be suggested to account for the remarkable facts stated in the pre- vious section. In this connexion two further important points should be mentioned : (i) radium compounds are continually giving out heat and maintain themselves at a temperature i to 2 above their surroundings ; (2) the radio-activity of a series of compounds of a radio-active element is directly proportional to the amount of the element present. As has already been pointed out, radium is undoubtedly an element, with a definite spectrum and a definite place in the periodic system. The simplest explanation of all the facts we have mentioned is that atoms of this element are continually disintegrating or exploding, throwing off charged particles and forming substances of smaller atomic weight. The first stage consists in the expulsion of an a particle (helium atom) and the forma- tion of radium emanation, which presumably has an atomic weight 4 units less than that of the radium atom. The atoms of the emanation in their turn explode, the first stage being the expulsion of an a particle with formation of radium A. Radium A in its turn expels an a particle and radium B results. The latter in its turn breaks down into radium C, but in this transformation no active particles are expelled, and so on. The time occupied in these different stages is very different. In the case of radium, the proportion of atoms exploding is very small, and it has been calculated that it requires about 2600 years for the activity of radium to fall to half its value. On the other hand, the half-time period for the emanation is about 3.85 days, and for radium A to radium B only three minutes. The successive stages in the disintegration of radium are given in the accom- panying table ; the rays given off (if any) at each stage are given in brackets, and the half-time period below : Radium () - Emanation(a) -> A(o) -> B -> C(a,/3/y) -> D -> E(/3,y,) -> F(a) 2600 years 3.85 days 3 min. 21 min. 28min. 40 yrs. 6 days 140 days. Radium F is polonium, discovered in pitchblende by Mme. Curie. It breaks down much more rapidly than radium, but the final product is unknown. It is doubtless a non-radio-active and stable element, and may possibly be lead (atomic weight 207), which is present in all uranium minerals. On the disintegration theory the source of the energy given out in a radio-active change is the internal energy of the atom ; the process of breaking up results in the formation of a succession of substances with less and less energy. Curie and Laborde found that i gram of radium evolves 100 calories per hour. It has been calculated that the heat given out in the complete disintegration of i gram of radium would raise 100 tons of water i in temperature ; or, in mechanical units, the available energy in i gram of radium would suffice to RADIO-ACTIVITY 567 raise 400 tons i mile high. An alternative illustration of the enormous amounts of energy concerned is that in the complete change of 1.3 cubic millimetres of emanatior 10,540 calories are evolved a quantity of energy four million times greater th in is evolved by the same volume of hydrogen and oxygen when they explode to form water. This gives us some idea of the immense stores of energy in the ato n. The chtmical activities of radium, for example the decomposition of water into hydrogen and oxygen, change of oxygen to ozone, etc. , are doubtless connected with the I beration of this large amount of energy. Radio-activity of Uranium and Thorium In concluding this brief account of radio-activifyi reference may be made to the radio-activity of uranium and thorium. As radium is continuously breaking down, it is evident that its presence in pitchblende can only be accounted for on the view that it is continuously being reformed from something else. As the quantity of radium in uranium ores is approxim itely proportional to the quantity of uranium present.it is natural to suppose Mat the latter (atomic weight 238.5) is the parent of radium. Whether this is so or not is not definitely settled. The first stage in the disintegration of uraniu.n is attended by the expulsion of an a particle and the formation of a new compound, uranium X, which in its turn breaks down into simpler sub- stances. There is some evidence that actinium (p. 564) is one of the disintegra- tion products of uranium, and radium may be a still later stage in the process. Thorium, like radium, gives off an emanation which gives rise to radiations, and in time loses its activity. There does not appear to be any connexion between the uranium and thorium products ; they belong to different series. It is suggestive that it is the three elements of highest atomic weight (radium, thorium, uranium) which undergo spontaneous disintegration ; there is no evidence that elements such as lead or bismuth are radio-active. Thus there appears to be a ccnnexion between the size of an atom and its instability. Radio-active change is entirely spontaneous, and no method of initiating it or influencing its rate is known. In particular, the widest range cf temperature, from that of liquid hydrogen to 2000, does not appear to affect the speed of atomic disintegration in the smallest degree. INDEX ABSOLUTE temperature, 43. - zero, 43 49. Absorption coefficient, 79. of gases by charcoal, 313. Accumulate', lead, 491. Acetic acid, 334. Acetylene, 332. Acid anhydrides, 176, 186. Acidic oxides, 186, 254. Acids, activity of, 187, 266. basicity of, 187, 254. general properties of, 98, 187, 265. strength of, 187, 263, 265. and salts, nomenclature, 188. Actinium, 564. Actinometei , 93. Active mass, 167. Adsorption, 313, 466. Affinity, chemical, 15. free, 272, 563. Agate, 352. Air, combustion in, 27. liquid, 72, 209. Alabaster, .(36. Alcohol, definition, 330. - ethyl, 333. Aldehydes, 334. Alkali meta s, comparison of, 406. - waste, 1:75, 387. Allotropic modifications, 175. Alloys, 381, 411, 465. Alum, burnt, 469. chrome, 519. Aluminates 467. Aluminium, 463. - alloys, 165. bronze, 411, 465. chlorid :, 468. hydrox de, 466. oxide, 466. silicates, 354, 469. sulphate, 468. sulphi< e, 470. Alums, 468. pseudo, 469. Alunite, 469. Amalgams, 456. Amethyst, 466. Amides, 327. Ammonia, 213. composition of, 217. soda process, 389 Ammoniacal liquor, 214, 402. Ammonium, 405. amalgam, 405. carbamate, 404. carbonates, 404. chloride, 403. dissociation of, 215, 403. dichromate, 202, 522. - hydroxide, 216, 403. ionic dissociation of, 263, 403. magnesium arsenate, 502. phosphate, 448. molybdate, 524. nitrate, 233. nitrite, 201. phosphomolybdate, 524. salts, 402. stannichloride, 483. sulphate, 404. sulphides, 405. thioantimonate, 511. thioantimonite, 510. thioarsenate, 503. thioarsenite, 503. - thiostannate, 483. Anatase, 476. Andalusite, 354. Anglesite, 484. Anhydride, 176. Anhydrite, 432, 436. soluble, 437. Anions, 258. Anode, 15, 258. Anthracite, 314. Antimonates, 510. I Antimonic acids, 509. Antimonious oxide, 510. Antimony, 504. 569 570 A TEXT-BOOK OF INORGANIC CHEMISTRY Antimony, explosive, 505. hydride, 506. nitrate, 510. ochre, 504. pentachloride, 507. pentafluoride, 507. ' pentasulphide, 511. pentoxide, 509. sulphate, 510. tetr oxide, 509. tribromide, 508. trichloride, 507. triiodide, 508. trioxide, 509. trisulphide, 510. Antimonyl potassium tartrate, 509. Apatite, 437. Apollinaris water, 60. Aqua regia, 225, 235. Aquafortis, 222. Aquamarine, 444. Aragonite, 434. Arc, electric, 236, 332. Argentite, 420. Argon, 208. discovery of, 207. Argyrodite, 477. Arsenates, 502. Arsenic, 495. acids, 502. disulphide, 502. oxychloride, 500. pentasulphide, 503, 504. - pentoxide, 501. tribromide, 500. trichloride, 499. trifluoride, 499. triiodide, 500. trioxide, 500. - trisulphide, 355, 503. white, 500. Arsenious oxide, 500. Arsenites, 501. Arseniuretted hydrogen, 497. Arsenolite, 495. Arsine, 497. Asbestos, 446. Association, 153. Atmolysis, 48. Atmosphere, 203. and combustion, n. Atmospheric nitrogen, utilization of, 206, 236. pressure, 41. Atom, definition of, 106. Atomic heat, 119. theory, 105. volumes, 367. Atomic weights, determination of, 115. and specific heats, 118. standard of, 115. Attraction, molecular, 49. Aurates, 429. Auric chloride, 430. salts, 429. Aurous salts, 429. Avogadro's hypothesis, 106. Azoimide, 219. Azurite, 407. BALANCE, 8. Barium, 440. chloride, 442. chromate, 522. hydroxide, 441. hypophosphite, 249. nitrate, 442. oxide, 23, 441. peroxide, 23, 442. sulphate, 442. Bases, 33, 98, 265. strength of, 263, 266, 371. Basic oxides, 371. salts, 448. - slag, 539. Basicity ot acids, 187, 254. Bauxite, 463. Bell metal, 411. Beryllium, 444. atomic weight of, 369. Bessemer process, 538. Bismuth, 511. glance, 511. - hydroxide, 514. ochre, 511. oxychloride, 513. oxynitrate, 514. peroxides of, 514. subnitrate, 514. sulphate, 514. tribromide, 513. trichloride, 512. -- trifluoride, 512. triiodide, 513. trinitrate, 514. trioxide, 513. trisulphide, 514. " Bismuthic acid," 514. Bismuthite, 511. Bismuthous oxide, 513. Black ash, 387. Blackband, 535. Blacklead, 310. Blast-furnace, 536. Bleaching, 90, 285. Bleaching powder, 436. INDEX Blue vitriol, 414. Boiling-point of a liquid, 64. elevi tion of, 196. Bone-ash, 2 \g, 437. Bones, composition of, 239 Boracite, 350. Borates, 361. Borax, 356, 361. Boric acids, 359-361. Boron, 356. - hydride,, 357. nitride, 358. sulphide, 359. trichlonde, 358. trifluoride, 358. trioxide, 359. Bort, 309. Boyle's law, 40. - accuracy of, 42, 50. and kinetic theory, 49. Brass, 411. Bniunite, 527. Bricks, 470. Erin's oxygen process, 23. Britannia metal, 506. Bromates, 184. Bromic acid, 183. Bromine, 154. trifluorile, 190. water, 55. Bronze, 411 Brookite, 476. Bunsen burner, 341. Burnt alum. 469. Butter of antimony, 507. Cadmium, 453. compounds of, 454. Caesium, 402. Calamine, 449. Calaverite, 427. Calc-spar, 434. Calcite, 434. Calcium, 432. bicarbonate, 324, 435. carbide, 332, 438. carbon; Lte, 324, 434. chlorate, 397. chloride- , 435. chrom ite, 522. cyanan ide, 236, 438. fluoride , 152. - hydride, 433. mangauite, 529. oxide, 433. - peroxi- e , 434. - phosphate, 437. silicate;;, 438. Calcium sulphate, 436. solubility in water, 437. sulphide, 387, 389, 438. Calculations, 123. Calomel, 457. Caloric, 59. Calorific power, 333. Calx, 29. Candle, burning of, 9. Carbamide, 327. Carbides, 314, 43 8 - Carbon, 308. allotropic modifications, 308. - cycle, 335. dioxide, 205, 319. solid, 321. disulpbide, 325. monoxide, 317. oxysulphide, 326. suboxide, 316. tetrachloride, 493. Carbonado, 309. Carbonates, 324. Carbonic acid, 322. Carbonyl chloride, 326. Carbonyls, 319. Carborundum, 314, 351. Carnallite, 394. Caro's acid, 298. Cassiterite, 478. Cast iron, 537. Catalysis, 22, 87, 140, 287. Cathode, 15, 258. Cations, 258. Caustic potash, 395. soda, 382. Celestine, 439. Cement, 434. Portland, 434. Cementation process, 539. Cementite, 541. Cerite, 477. Cerium, 477. Cerussite, 484. Chalcocite, 407. Chalcopyrite, 407. Chalk, 434. Chalybeate waters, 60. Chamber crystals, 289. Chance's process, 389. Charcoal, 311. Charles's law, 42. Chemical affinity, 15, 272. change, 3. characteristics of, 16. types of, 6. energy, n, 144, 336. equilibrium, 164. 572 A TEXT-BOOK OF INORGANIC CHEMISTRY Chemical equilibrium in electrolytes, 261, 423. Chili saltpetre, 158, 223, 391. Chlorates, 181. Chloric acid, 181. Chlorinated lime, 436. Chlorine, 86. bleaching action of, 90. chemical relations, 190. dioxide, 178. heptoxide, 179. hydrate, 91, 112. liquefaction of, 71. monoxide, 177. oxides of, 100, 178. preparation, 86. water, 89. Chlorites, 181. Chloroplatinates, 560. Chloroplatinites, 559. Chlorosulphonic acid, 298, 300. Chlorous acid, 181. Chromates, 522. Chrome alums, 519. iron ore, 517. Chromic acids, 523. anhydride, 522. chloride, 519. hydroxide, 518. oxide, 519. sulphate, 519. Chromites, 520. Chromium, 517. ammonia compounds, 523. trioxide, 520, 522. Chromous compounds, 518. Chromyl chloride, 523. Chrysoberyl, 444,- 467. Cinnabar, 454. Classification of the elements, 362. Clay, 354, 469. ironstone, 535. Coal, 314. Coal-gas, 332. Cobalt, 548. glance, 548. Cobaltammines, 551. Cobaltic hydroxide, 549. oxide, 549. sulphate, 550. Cobalticyanides, 551. Cobaltinitrites, 551. Cobalto-cobaltic oxide, 549. Cobaltocyanides, 551. Cobaltous salts, 549. Coke, 312. Colemanite, 356. Collection of gases, 21, 39. Colloidal solutions, 353, 354. Colloids, reversible, 356. Columbite, 495. Columbium, 495. Combination, chemical, 6. Combining weights, 103, 115. law of, 102. and atomic weights, 115, 120. chemical equivalents, 120. Combustion, 26, 337. definition of, 26. - heat of, 145. Complex ions, 414, 424, 430, 562. Compound, chemical, 4, 7. Concentration, definition of, 65. molecular, 166, 167. Condenser, Liebig's, 61. Conductivity, electrical, 257. Condy's fluid, 534. Conservation of energy, u. mass, 8. weight, 8. Constant composition, law of, 57, 101. Contact process, 288. Co-ordination numbers, 562. Copper, 407. alloys of, 411. ammonia compounds, 414. compounds. See under Cuprous and Cupric. equivalent of, 126. hydrogen arsenite, 501. properties of, 410. Corrosive sublimate, 458. Corundum, 463, 466. Critical data, 71. phenomena, 70. Crocoisite, 484, 517. Crookesite, 472. Cryohydrates, 199. Cryolite, 149, 379, 463. Crystallization, velocity of, 277. water of, 112, 384. Crystallography, 301. Crystalloids, 354. Crystals, 301. mixed, 199. symmetry of, 301. Cupellation, 420. Cupric carbonates, 415. chloride, 413. hydroxide, 413. nitrate, 415. oxide, 55, 413. sulphate, 414. sulphide, 415. Cuprite, 407. Cuprous chloride, 412. INDEX 573 Cuprous cyanide, 412. iodide, 412. Dxide, i-.-Li. sulphide, 413. Cyanates, 328. Cyanide process (gold), 428. Cyanides, commercial preparation, 237. Cyanogen, 327. D ALTON'S law, 77. Davy lamp 346. Deacon's process, 88. Decomposrion, double, 7. simple, 7. Degree of dissociation, 171, 264, 267. Dehydration, 386. Deliquescence, 387. Density of g-ases and vapours, 109, 127. Dewar flasks, 75. Dialysed iron, 543. Dialysis, 353. Diamond, 308. Dichromates, 521. Dichromic acid, 523. Diffusion of gases, 46. lav of, 47. of liquids, 194. Dimorphism, 276. Disintegration of atoms, 566. Dissociation, degree of, 171, 264. electrolytic, 261. anl valency, 271. therm il, 169. Distillation, 61. destructive (dry), 214, 311. under reduced pressure, 224. Distribution coefficient, 160. Disulphur trioxide, 298. Dithionic icid, 299. Dolomite, 445. Double decomposition, 7. Dulong and Petit's law, 118. Dutch metal, 90. Dynamite, 225. EARTHE>WARE, 470. Earth's ciust, composition of, 19. Eau de Juvelle, 180. Efflorescence, 387. Electric arc, 236, 332. battery, 417. furnace, 315. Electrical conductivity, 257. Electro-afinity, 372. Electrocl emical equivalents, 259. Electrodes, 15, 258. Electrolysis, 15, 257. of sodium chloride, 382. Electrolysis of water, 14. Electrolytes, 257. Electromotive force, 416. Electrons, 565. Elements, 7, 17. classification of, 362. list of, 18, and back of cover. natural occurrence, 19. potential series of, 418. Emerald, 444. Emery 1 463, 466. Empirical formula, 121. Endosmosis, 193. Endothermic actions, 145, 172. compounds, 145, 221, 232, 326. Energy, chemical, 13, 144, 336. conservation of, n. definition of, 12. electrical, 13, 418. transformations of, 12. Epsom salts, 448. Equations, 22. writing of, 122, 226, 521. Equilibria, displacement of, 171. Equilibrium, chemical, 164, diagram for water, 68. displacement of, 169, 171, 232. physical, 63 et seq. Equivalents, chemical, 121, 125, 131. electrochemical, 259. Ethane, 330. Ethyl alcohol, 333. Ethylene, 331. Exothermic actions, 145. compounds, 215, 322. Expansion of gases, 44. Eutectic mixture, 198. Euxenite, 477. Exosmosis, 193. FACT, 113. Faraday's laws, 259. Felspar, 4, 463. Ferrates, 544. Ferric acid, 544. alum, 546. chloride, 545. - hydroxide, 543, 545. - oxide, 543. sulphate, 546. sulphide, 546. Ferricyanides, 547. Ferrocyanides, 547. Ferrous ammonium sulphate, 545. carbonate, 545. chloride, 544. - hydroxide, 543. oxide, 543. 574 A TEXT-BOOK OF INORGANIC CHEMISTRY Ferrous sulphate, 544. sulphide, 278, 546. Filtration, 5. Fire-damp, 329. Flame, 338. oxy-hydrogen, 343. separator, 342. structure of, 340. Flames, luminosity of, 339. temperature of, 343.^ Flint, 352. Flowers of sulphur, 274. Fluoborates, 358. Fluorine, 149. Fluorspar, 149, 432. Flnosilicates, 350. Flux, 237. Formation, heat of, 145. Formulae, empirical, 121. graphic, 131. molecular, in, 121. significance of, 112. Franklinite, 449. Free affinity, 272. Freezing-mixture, 436. Freezing-point, depression of, 196. Fuming nitric acid, 226. sulphuric acid, 296. Furhace, blast-, 536. electric, 315. reverberatory, 388. Fusible metals, 512. Fusion, heat of, 67. GALENA, 484. Gallium, 471. Galvanic battery, 417. Galvanized iron, 450. Garnierite, 552. Gas carbon, 312. laws, 40. deviations from, 42, 44, 49. perfect, 50. Gaseous diffusion, 46. law of, 47. volumes, law of, 104. Gases, collection of, 21, 39. diffusion of, 46. effusion of, 48. expansion coefficient, 44. general equation for, 45. general properties of, 40. inactive, 207. kinetic theory of, 48. liquefaction of, 70. Gay-Lussac tower, 292. Gay-Lussac's law, 104. Generalization, 113. German silver, 411. Germanium, 477. Glass, 438. etching of, 153. Glauber's salt, 391. Glover tower, 292. Glucinum (see beryllium) Gold, 427. alloys, 429. properties of, 428. Goldschmidt's thermite process, 466. Graham's law. A?. Graham's law, 47. [-molecule, volume occupied by, i Gram no. Granite, 3, 463. Graphic formulae, 131. Graphite, 310. Graphitic acid, 311. Greenockite, 453. Guncotton, 225. Gun-metal, 411. Gun powderr40o. Gypsum, 432, 436. HEMATITE, 535. Halogen acids, 92, 149. heats of formation of, 191. Halogens, comparison, 190. - valency of, 325. Hampson's apparatus, 72. Hardness of water, 325, 435. H arrogate water, 60. Hartshorn, spirits of, 213. Hausmannite, 527. Heat, atomic, 119. molecular, of gases, 208. of combustion, 145. formation, 145. fusion, 67. neutralization, 269. solution, 145. vaporization, 66. Heavy spar, 440. Helium, 209. discovery of, 207. liquefaction of, 210. Henry's law, 78. Hess's law, 146. Heterogeneous equilibria, 171. Hofrnann's method for determining vapour densities, 129. Horn silver, 420. Hydrates, 112, 384. solubility of, 384. vapour pressure of, 386. Hydrazine, 218. Hydrazoic acid, 219. Hydriodic acid, 161. INDEX 575 Hydrobromir acid, 156. Hydrocarbons, 316, 329. Hydrochloric acid, 92. Hydrocyanic acid, 328. Hydrofluobo ic acid, 358. Hydrofluoric acid, 152. Hydrofluosilicic acid, 350. Hydrogel, 355. Hydrogen, 31. chemica' properties of, 37. combination with oxygen, 37. occlusion of, 39. physical properties of, 36. preparat on of, 31. antimonide, 506. arsenide, 497. bromide, 156. chloride, 92. composition of, 96. disulphide, 280. fluoride, 152. iodide, 161. thermal dissociation of, 163, 164, 166. pentasul')hide, 281. peroxide, 138. estimition of, 142. tests "or, 142. phosphu es, 242, 244. polysulphides, 280. selenitic, 304. silicide, 349. sulphide, 277. propc rties of, 278-280. trisulphiie, 281. telluride. 305. Hydrolysis, 252, 267. Hydromagncsite, 448. Hydrosol, 355. Hydroxides,' 254. Hydroxylam;ne, 219. sa.lts of, 220. Hypobromoiis acid, 183. Hypochloiites, 180. Hypochloroi:s acid, 179. anhydride, 177. Hypoiodous icid, 184. Hyponitrous acid, 234. Hypophosphoric acid, 251. Hypophosphorous acid, 249. Hyposulphuious acid, 298. Hypothesis, iefinition of, 113. ICK, melting-point of, 68. Iceland spar 434. Ignition temperature, 344. Incandescence, 339 Incandescent mantles, 340. Indium, ^j\. lodates, 185. lodic acid, 185. anhydride, 184. Iodine, 158. bromide, 190. chlorides of, 189. dioxide, 184. dissociation of, 159, 170. pentarluoride, 190. pentexide, 184. Ionic equilibrium, 261. lorn s.ation theory, 261. degree of, 264, 267. Ions, 258. complex, 424, 430, 562. Iridium, 557. Iron, 535. allotropic modifications, 540. alloys, 538, 542. alum, 546. carbonyls, 319. cast, 537. compounds of. See under Ferrous and Ferric. disulphide, 546. galvanized, 450. oxides c f, 543. passivity of, 542. pyrites, 535, 546. rusting of, 541. wrought, 538. Isodimorphism, 303. Isomerism, 175. Isomorphism, 119, 303. - law of, 120. JASPER, 352. Joule-Thomson effect, 72. KAINITE, 394, 445. Kaolin, 469. Kassner's process, 488. Kelp, 158. Kieselguhr, 382. Kieserite, 445. Kinetic theory of gases, 48. Kipp's apparatus, 35. Krypton, 210. Kupfernickel, 496, 552. LABILE state, 85. Lakes, 467. Lamp-black, 312. Landolt, 10. Lanthanum, 474. Lapis lazuli, 470. Laughing gas, 233. 576 A TEXT-BOOK OF INORGANIC CHEMISTRY Lavoisier, 10, 17, 27. Law, definition of, 113. - Boyle's, 40. Charles's, 42. Dalton's, 77. Dulong and Petit's, 118. Gay-Lussac's, 104. - Graham's, 47. Henry's, 78. Hess's, 146. Newlands', 363. Periodic, 366. of combining weights, 102. of constant composition, 101. mass action, 165. multiple proportions, 101 octaves, 363. thermoneutrality, 270. Laws, Faraday's, 259. Lead, 484. alloys, 486. - properties of, 485. accumulator, 491. acetate, 491. bromide, 489. carbonate, 489. chloride, 488. chromate, 522. disulphate, 491. hydroxide, 487. iodide, 489. monoxide, 487. nitrate, 228, 489. - peroxide, 488. - red, 487. sesquioxide, 487. suboxide, 486. sulphide, 490. sulphate, 489. tetracetate, 491. tetrachloride, 491. white, 490. Le Chatelier's thecrem, 173. Leblanc process, 387. Lepidolite, 401. Liebig's condenser, 61. Lime, 433. Lime-water, 433. Limestone, 434. Linde-Hampson apparatus, 72. Liquation, 478. Liquefaction of gases, 70. Liquid air, 73, 209, 211. fractionation of, 74. Liquids, diffusion of, 194. general properties, 62. miscibility of, 80. vapour pressures of, 63. Lithium, 401. mica, 401. salts of, 402. Lodestone, 543. Luminescence, 339. Luminosity of flames, 339. MAGNALIUM, 465. Magnesia, calcined, 446. Magnesite, 445. Magnesium, 445. ammonium arsenate, 502. phosphate, 448. bicarbonate, 449. carbonates, 448. chloride, 447. hydroxide, 446. nitride, 208. oxide, 446. pyrophosphate, 448. silicide, 348, 349. sulphate, 448. Magnetic iron ore, 535, 543. Magnetite, 543. Malachite, 407, 415. Manganates, 531. Manganese, 527. bronze, 411. dioxide, 88, 528. heptoxide, 529. trioxide, 529. Manganic acid, 531. chloride, 530. hydroxide, 530. oxide, 528. Manganite, 527. Manganous acid, 529 chloride, 530. hydroxide, 528. oxide, 528. sulphate, 530. sulphide, 530. Marble, 434. Marsh-gas, 329. Marsh's test, 497. Mass action, law of, 165. active, 167. conservation of, 8. Matches, 242. Matte, 408. Matter, 3, Mechanical mixture, 4. Meerschaum, 445. Melting-point, definition, 68. Mercuric ammonium chloride, 460. chloride, 458. cyanide, 459. diammonium chloride, 461. INDEX 577 Mercuric iodide, 459. nitrate, 459. oxide, 458. sulphide , 460. Mercurous chloride, 457. chromale, 522. iodide, 457. nitrate, 458. oxide, 457. sulphate, 458. Mercury, 4^.4. Metals and non-metals, 18, 371. general properties of, 371. preparation of, 374. relative displacing powers of, 418. Metantimonic acid, 509. Metaphospt oric acid, 253. Metaphospl orous acid, 251. Metarsenic icid, 502. Metasilicic ;icid, 353. Metastable .state, 69, 85. Metastannic. acid, 482. Meteoric iron, 535. Methane, 3:^9. Microcosmic salt, 252. Milk of lime, 433. Millon's base, 461. Mineral wa.ers, 60. Miscibility of liquids, 80. Mixed crystals, 199. Mixture, mechanical, 4. Molecular r.ttraction, 40. " Molecular compounds," 272, 561. Molecular concentration, 167. formule, in, 121. how to establish, 121. significance of, 112. heat oi gases, 208. weights, 108 in solution, 195. Molecule, definition of, 106. Molecules, complexity of, 108. Molybdenite, 524. Molybdenum, 524. Molybdic acid, 524. Monoperphosphoric acid, 254. Monopersi Iphuric acid, 298. Monazite, 209, 477. Mordants, 467. Mortar, 434. Multiple p;oportions, law of, 101 NASCENT hydrogen, 214, 219, 498. oxyge i, 91. state, 498. Neon, 210. Nessler's reagent, 459, 461. Neutralization, 99. 37 Neutralization, heat of, 269. Newlands' law of octaves, 363. Niccolite, 552. Nickel, 552. alloys of, 552. carbonyl, 319. cyanide, 554. sulphides of, 554. Nickelic oxide, 553. hydroxide, 553. Nickelo-jiickelic oxide, 553. Nickelous chloride, 554. hydroxide, 553. oxide, 553. sulphate, 554. Niobium, 495. Nitramide, 235. Nitrates, 222, 225. " brown ring test " for, 232. Nitre, 222, 378, 399. Nitric acid, 222. fuming, 226. oxide, 231.' Nitrides, 202, 217. Nitrites, 230, 391. Nitrogen, 200. atmospheric, utilization of, 206, 236. utilization by plants, 206. bromide, 221. chloride, 221. iodide, *zi. pentoxide, 227. peroxide, 227. dissociation of, 228. trioxide, 229. Nitroglycerine, 225. Nitrolim, 236. Nitrosyl bromide, 235. chloride, 235. fluoride, 235. sulphuric acid, 289. Nitrous acid, 230. oxide, 233. Nitryl fluoride, 235. Noble metals, 235. Nomenclature, acids and salts, 188. Non-electrolytes, 257. Non-metals, 18, 371. Nordhausen sulphuric acid, 296. Normal solutions, 237. OCCLUSION, 39, 313, 557. Octaves, law of, 363. Oil of vitriol, 165, 287, 300. Olivine, 354. Opal, 352. Organic acids, 334. compounds, 315. 578 A TEXT-BOOK OF INORGANIC CHEMISTRY Orpiment, 503. Orthoantimonic acid, 509. Orthoboric acid, 359. Orthoclase, 354. Ortho-phosphoric acid, 251. Orthoplumbic acid, 488. Orthosilicic acid, 354. Orthostannic acid, 482. Chmic acid, 557. Osmium, 557. Osmosis, 192. Osmotic pressure, 192. and gas pressure, 195. Oxidation, 39, 415. Oxygen, 20. chemical properties, 25. commercial preparation, 23, 74, 488. discovery of, 20. preparation, 21, 532. standard for atomic weights, 115. Oxy-hydrogen flame, 343. Ozone, 61, 133, 240. tests for, 137. Ozone-oxygen equilibrium, 135, 172, 174. Palladium, 557. occlusion of hydrogen by, 39, 557. Parkes process, 420. Partial pressures, law of, 77. Passivity of metals, 517, 542. Pattison process, 420. Pearlite, 541. Peat, 314. Pentathionic acid, 299. Percarbonic acid, 325. Perchloric acid, 182. anhydride, 179. Perchromic acid, 523. Perfect gas, 50. Periodic acid, 185. law, 366. system, 363. deficiencies of, 370. uses of, 367. Permanganates, 532. Permanganic acid, 532. anhydride, 529. Permonosulphuric acid, 298. Peroxides, 188. Perphosphoric acid, 254. Persulphates, 297. Persulphuric acid, 297. anhydride, 297. Petalite, 401. Petzite, 427. Pewter, 480. Phase, 69. Phlogiston theory, 29. Phosgene, 326. Phosphates, 251. Phosphine, 242. Phosphomolybdates, 524. Phosphonium compounds, 244. Phosphor bronze, 411. Phosphoric acid, 239, 251. Phosphorite, 238, 432, 437. Phosphorous acid, 250. oxide, 248. Phosphorus, 238. combustion of, 28, 240, 248. diiodide of, 247. halogen compounds of, 245. Hittorfs, 242. hydrides of, 242. metallic, 242. oxides of, 247. oxychloride, 247. oxyacids of, 247. pentabromide, 247. pentachloride, 246. dissociation of, 169, 246. pentafluoride, 245. pentoxide, 248. red, 241. sulphides of, 256. tetroxide, 248. tribromide, 247. trichloride, 245. trifluoride, 245. tri-iodide, 247. trioxide, 248. yellow, 240. Photography, 426. Physical change, 2, 3. Pig-iron, 537. Pitch-blende, 525, 564. " Pink salt," 483. Plaster of Paris, 437. Platinic chloride, 560. hydroxide, 559. oxide, 559. sulphide, 560. Platinicyanides, 560. Platinocyanides, 560. Platinous chloride, 559. hydroxide, 559. oxide, 559. sulphide, 560. Platinum, 558. ammonia compounds; 560. black, 559. ore, 556, 558. spongy, 559. Plumbago, 310. INDEX 579 Plumbates, 488. Plumbites, 487. Polonium, ',64. Polymeridei 153, 175. Polymerizat on, 153, 175. Polymorphism, 276. Polythionic acids, 299. Porcelain, 470. Potable waters, 60. Potash alum, 469. caustic, 395. Potassium, 394. aluminium sulphate, 469. antimo lates, 510. antimo'iyl tartrate, 509, auri chloride, 430. bicarbonate, 398. bromide, 396. carbonate, 398. carbox de, 394. chlorate, 181, 397. chlorieb, 396. chronu.te, 522, cobalt icy an ide, 551. cobalti litrite, 551. cobaltc cyanide, 551. cyanide, 328. dichroinate, 522. ferrate, 544. ferricyanide, 547. ferrocyanide, 547. fluorid ;, 150. fluosili^ate, 350. - hydride, 395. hydrogen fluoride, 152. sulphate, 399. hydro ide, 395. - hypochlprite, 179, 187. hypoiodite, 184. iodide 396. mang.'.nate, 531. mercu - ic iodide, 459. metamimonate, 510. mono ide, 395. nitrate, 399. nitrite 230. oxides of, 395. perchlorate, 182, 398. peroxide, 395. permanganate, 22, 532. persulphate, 297. platin chloride, 401, 560. polysvlphides, 401. polyehromates, 522. pyroa itimonate, 510. silver cyanide, 425. sulphate, 398. sulphides, 299, 400. Potential series of the elements, 418. Pressure, osmotic, 192. Properties, characteristic, 2. non-characteristic, 2. Proustite, 420. Prussian blue, 548. Puddling, 538. Purple of Cassius, 430. Pyrargyrite, 420. Pyrites, copper, 407. - iron, 535, 546. Pyroantimonic acid, 509. Pyroantimonates, 510. Pyroarsenic acid, 502. Pyrolusite, 527. Pyromorphite, 484. Pyrophosphoric acid, 252. Pyrosulphuric acid, 286, 296. Pyrosulphates, 296. QUARTZ, 351. Quicklime, 433. Quicksilver, 455. RADIO-ACTIVITY, 564. Radium, 564. emanation, 565. Rain water, 59. Rare earths, 474. Reaction, velocity of, 173. Reactions, reversible, 24, 87, 164. Reactivity, chemical, and heat of re- action, 148. Realgar, 502. Red lead, 487. Reduction, 39, 415. Residual affinity, 272, 563. Reverberatory furnace, 388. Reversible reactions, 24, 87, 164. Rhodium, 557. River water, 60. Rock crystal, 352. Rocks, constituents of, 354, 469. Rontgen rays, 564. Rose's metal, 512. Rouge, 543. Rubidium, 402. Ruby, 463, 466. Rusting of iron, 541. Ruthenium, 556. Rutile, 476. SAFETY lamp, 346. Sal ammoniac, 213, 403. Salts, basic, 448. definition of, 99. normal, 252. preparation of, 374. 580 A TEXT-BOOK OF INORGANIC CHEMISTRY Salts, solubilty of 82, 376. Sand, 347. Sapphire, 463, 466. Saturated compounds, 331. solutions, 77. Scandium, 474. Scheele's green, 501. Schonite, 448. Sea-water, 60. Selenic acid, 305, 428. Selenious acid, 305. Selenite, 432, 436. Selenium, 303. chlorides of, 304. dioxide, 304. Seltzer water, 60. Semi-permeable membranes, 193. Senarmontite, 504. Serpentine, 445. Siemens-Martin process, 539. Silent discharge, 133. Silica, 351. Silicates, 354. Silicic acids, 352. Silicides, 348, 349. Silicoethane, 349. Silicomethane, 349. Silicon, 347. carbide, 351. chloroform, 351. dioxide, 351. fluoride, 153, 349. tetrachloride, 351. tetrabromide, 351. tetraiodide, 351. Silver, 419. alloys, 422. antimonide, 506. arsenite, 501. arsenate, 502. bromide, 425. carbonate, 427. chloride, 423. chromate, 522. colloidal, 422. cyanide, 425. fluoride, 425. glance, 420. hydroxide, 422. hyponitrite, 234. iodide, 425. metaphosphate, 253. nitrate, 426. oxide, 422. periodate, 186. peroxide, 422. plating, 425. phosphate, 253. Silver, potassium cyanide, 425. sodium thicsulphate, 299. sub-halides of, 423, 426. sulphate, 426. Slag, 408, 537. Slaked lime, 433. Smaltite, 548. Smithsonite, 449. Soda, calcined, 389. caustic, 382. Sodamide, 381. Sodium, 378. acetate, 329. aluminate, 464. amalgam, 381, 456. aurichloride, 430. bicarbonate, 390. - bisulphate, 388. bromide, 384. carbonate, 387. hydrolysis of, 324. dichromate, 522. hydride, 381. - hydrogen carbonate, 390. hydrogen sulphate, 388. hydroxide, 381. hyposulphite, 298. metaphosphate, 252. metastannate, 482. monoxide, 381. nitrate, 223, 391. nitrite, 391. oxides of, 381. periodate, 186. peroxide, 381. phosphates, 251, 391. platinichloride, 560. pyroantimonate, 510. pyrophosphate, 252. silicates, 392. silver thiosulphate, 299. stannate, 482. stannite, 480. sulphate, 391. electrolysis of solution, 258. solubility of, 83, 391. sulphides, 392. sulphite, 286. tetrathionate, 299. thiosulphate, 299. Solder, 480. Solid solutions, 199. Solubility curves, 82. effect of temperature on, 81. of gases in liquids, 77. of liquids in liquids, 80. of solids in liquids, 81. product, 423. INDEX Solubility tables, 376. Solution, definition, 76. heat of, 145. pressure, 85. Solutions, boiling-points of, 198. freezing-points of, 197. general properties of, 76. saturated, 77. solid, 108. supersa urated, 84. unsatur ited, 84. Solvay process, 389. Soot, 312. Spathic iron ore, 535. Specific hea of gases, 208. of sc 'lids, 118. Spectroscope, 392. Spectrum analysis, 392. Specular ire n ore, 535. Speed of reaction, 173. Spiegeleisen, 539. Spinelle, 467. Spirits of hartshorn, 213. Spitting of silver, 421. Spodumene, 401. Spring water, 59. Stable state, elefinition, 69, 276. Standard temperature and pressure, no. Stannates, 482. Stannic acid, 482. chloride, 482. hydroxide, 482. oxide, 481. sulphide, 483. Stannous chloride, 480. oxide, 480. hydrate of, 480. sulphide, 483. Stassfurt deposits, 383, 394. Steam, action of, on iron, 34. on magnesium, 33. volumetric composition of, 54. Steel, 538. Sfibine, 506. Stibnite, 504. Strength of acids and bases, 187, 263, 265.. Strontianiie, 439. Strontium, 439. chloride, 440. hydroxide, 440. nitrate, 440. ' oxide, 440. sulphate, 440. Structural formulae, 131. Sublimation, 321, 404. Substance, 4. Substitution, 330. Sulphates, 296, 375. Sulphides, 279. Sulphites, 286. Sulphur, 273. allotropic modifications, 275. chlorides of, 281. dioxide, 282. dissociation of vapour, 275. flowers of, 274. heptoxide, 297. monoclinic, 275. plastic, 276. prismatic, 275. rhombic, 275. roll, 275. trioxide, 286, 296. Sulphuric acid, 287. chemical properties, 293. fuming, 296. Nordhausen, 296. physical properties, 293. preparation, 287. uses, 294, 295. Sulphurous acid, 286. anhydride, 282. Sulphuryl chloride, 300. Supercooling, 68, 70, 277. Superphosphate of lime, 438. Supersaturated solutions, 84. Sylvanite, 427. Sylvine, 394. Symbols, no. list of, back page of cover. Sympathetic ink, 550. TACHYDRITE, 435, 445. Talc, 445. Tantalite, 495. Tantalum, 495. Tartar emetic, 509. Telluric acid, 306. Tellurous acid, 306. Tellurium, 305. chlorides of, 305. oxides of, 306. Temperature, critical, 71. ignition, 344. of flames, 343. Tempering, 540. Tetrathionic acid, 299. Thallic salts, 473. Thallium, 471. Thallous salts, 472. Theory, definition, 114. Thermal dissociation, 169. Thermite process, 466. Thermochemical equations, 144. 582 A TEXT-BOOK OF INORGANIC CHEMISTRY Thermochemistry, 143. Thermoneutrality, law of, 270. Thioantimonatcs, 511. Thioantimonites, 511. Thioarsenates, 503. Thioarsenites, 503. Thiocarbonic acid, 326. Thiocyanates, 328. Thionic acids, 299. Thionyl chloride, 300. Thiostannates, 483. Thiosulphates, 298. Thiosulphuric acid, 298. Thomas and Gilchrist process, 539. Thomsonite, 354. Thorianite, 477. Thorite, 477. Thorium, 477. radio-activity of, 567. Tin, 478. alloys of, 480. amalgam, 480. combustion of, 28, 479. - pest, 479. Tinplate, 480. Tinstone, 478. Tincal, 356. Titanium, 476. Tourmalin^, 356. Transition temperature, 276. Tridymite, 351. Triple point, 67. Trithionic acid, 299. Tungsten, 524. Tungstic acid, 524. Type metal, 506. ULTRAMARINE, 470. Unit volume, 112. Unsaturated compounds, 331. Unstable state, definition, 69. 276. Uranium, 525. radio-activity of, 567. Urea, 327. VACUUM vessels, 75. Valency, 129, 271, 365, 561. and Faraday's laws, 271. and structural formulae, 129. Vanadinite, 494. Vanadium, 494. Vapour densities, determination'of, 12; pressure, 63. Vaporization, heat of, 66. Varec, 158. Velocity of crystallization , 277. of reaction, 173. and temperature, 174. Venetian red, 543. Vermilion, 460. Vichy water, 60. Vitriol, blue, 414. green, 544. oil of, 165, 287, 300. Voltameter, 14. Volume, atomic, 367. Volumetric composition of steam, 54. WASHING soda, 389. Water, 51. action of metals on, 32. as an acid, 268. as a base, 268. catalytic action of, 318, 403. composition by volume, 52. by weight, 55. decomposition into elements, 14, S 2 - distillation of, 61. hardness of, 325, 435. gas, 317. glass, 392. ionisation of, 267. of crystallization, 112, 384. physical properties of, 57. purification of, 61. softening of, 435. Waters, natural, 59. Weight, conservation of, 8. increase of, in chemical changes, 9. Weights, atomic, table of, back of cover. Weldon process, 529. Werner's theory of valency, 561. White lead, 490. precipitate, 460. Witherite, 440, Wolfram, 524. Wollastonite, 354. Wood, destructive distillation of 3"- Wood charcoal, 311. Wood's metal, 512. Wrought iron, 538. Wulfenite, 484. Wurtzite, 452. Xenon, 210. Ytterbium, 474. Yttrium, 474. ZEOLITES, 354. Zero, absolute, 43, 49. Zinc, 449. INDEX 583 inc, alloys of, 451. blende, 449. carbonates, 452. chloric!::, 451. equivalent of, 126. hydrox de, 451. Zinc oxide, 451. sulphate, 452, sulphide, 452. white, 451. Zircon, 477. Zirconium, 477. Printed by BALLANTYNE, HANSON &* Co. Edinburgh &* London REVIEWS OF THE FIRST EDITION OF " Outlines of Physical Chemistry ' by Dr. Senter W. OSTWALD in "Zeitschrift fur physikalische Chemie" " The book not only belongs to the class to which the descripti ' good ' can be applied, but takes high rank in that class. It is cleai and accurately written." NATURE " On the whole the book is one which can be recommended to all w wish to obtain a first acquaintance with the subject of Physical Chemistr In language it is clear and well expressed, and the practical illustratio which are appended to most of the chapters will be found very useful f laboratory work." AMERICAN CHEMICAL JOURNAL "... The various sections are closed by excellent practical illustratioi The presentation is exceptionally clear. The book is one of the b< elementary texts that have appeared, and deserves a cordial welcome." 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