COLLEGE of MINING 
 
 DEPARTMENTAL 
 
 BEQUEST OF 
 
 SAMUELBENEDICTCHRISTY 
 
 PROFESSOR OF 
 
 MINING AND METALLURGY 
 1885 -1914 
 
INTRODUCTION TO 
 METALLURGICAL CHEMISTRY 
 
INTRODUCTION TO 
 
 METALLURGICAL CHEMISTRY 
 
 FOE TECHNICAL STUDENTS 
 
 BY 
 J. H. STANSBIE, B.Sc. (LOND.), F.I.C. 
 
 ASSOCIATE OF MASON UNIVERSITY COLLEGE 
 AND LECTURER IN THE BIRMINGHAM MUNICIPAL TECHNICAL SCHOOL 
 
 SECOND EDITION 
 
 NEW YOEK 
 
 LONGMANS, GKEEN & CO. 
 LONDON : EDWARD ARNOLD 
 
PREFACE 
 
 THIS little book is intended for the use of technical students 
 who are desirous of making a study of the metals employed 
 for industrial purposes. It must be regarded strictly as a 
 preparatory course, before the study of " metallurgy " proper. 
 Several years' experience in teaching large classes in metal- 
 lurgy has made it evident that many who are engaged in 
 business during the day, and have, therefore, but little time 
 for study, are unwilling to devote several years to a syste- 
 matic course of chemistry before commencing the subject in 
 which they are essentially interested. It has been found that 
 two, or in special cases three, evenings a week is the usual 
 limit to the attendance of such students. This permits of only 
 one subject being taken at a time. 
 
 The difficulty has been recognised for some years in this 
 school, and at first attempts were made to overcome it by 
 including chemical principles in the ordinary metallurgy 
 classes ; but the short time at the disposal of the teachers 
 made the information fragmentary, and of limited value. It 
 was then decided to establish a preparatory class in connection 
 with the metallurgical department, in which chemical principles 
 could be treated with special reference to the wwk of that depart- 
 ment ; and by careful correlation of this with the other classes, 
 it has been found that the students feel at once that the work 
 
vi PREFACE 
 
 they are doing is directly applicable to their particular branches 
 of industry. 
 
 It is becoming more evident every year that the aim in an 
 evening technical school should be rather to increase the general 
 capability of a large number of students than to turn out a few 
 highly-trained men. This end will be best attained by making 
 the subjects taught as self-contained as possible. 
 
 The plan of this book is simple. It assumes that those who 
 use it are practically interested in the common metals, and that 
 they have no further knowledge of their properties than has 
 been obtained by ordinary observation in a workshop or 
 foundry, so that it is necessary to commence their scientific 
 treatment at the very beginning. The physical properties of 
 the metals and their alloys being fully treated in standard 
 works on metallurgy, more than a passing notice is not devoted 
 to this very important part of the study of metals. Brief 
 references to the more useful properties will be found in different 
 parts of the text. The greater portions of the chapters are 
 taken up with a description of the chemical properties of the 
 common metals, and in the development of chemical principles 
 the metals are made to take as prominent a part as possible. 
 The non-metals are, for the purpose of this book, regarded 
 more as the servants of the metals than as their masters ; but 
 this is not from any want of appreciation of their importance, 
 and it is hoped that the common ones are sufficiently 
 described to make them quite familiar. Such of the non- 
 metals as are not of importance in dealing with the chemistry 
 of the metals are either entirely omitted or only briefly 
 noticed, further study being left for a more advanced stage of 
 the subject. 
 
 The practical work begins with the study of the effects of air 
 and water on metals. This is followed by a discussion of the 
 
PEE FACE vii 
 
 common properties of matter. In dealing with the relation 
 between metals and sulphur the opportunity is taken of intro- 
 ducing the preparation and properties of sulphuric acid. This 
 is followed by a description of the nature and occurrence of 
 common salt and saltpetre, from which, with the aid of sul- 
 phuric acid, the other common acids are prepared. By the 
 aid of these compounds common substances of importance in 
 the treatment of metals are prepared and described. Special 
 attention is paid to the reactions between metals and acids ; 
 and the chemical equivalents of some metals and non-metals, 
 together with their atomic weights, are considered at some 
 length. The great importance of carbon in the treatment of 
 metals is recognised, and its properties carefully described. 
 This chapter is followed by others in which the important 
 subjects of combustion and reduction are somewhat fully 
 considered. The pronounced part played by silicon and its 
 compounds in metallurgical work is also specially recog- 
 nised. 
 
 Even a superficial examination of the text will show that the 
 book is mainly practical in character, and that the theoretical 
 principles of the subject are developed as far as possible from 
 actual experiments, to be made either by the student in the 
 laboratory, or by the teacher as class demonstrations. Every 
 experiment has been carefully verified, and will give good 
 results, if the instructions are faithfully followed. 
 
 It may appear to the systematic chemist that some portions 
 of the subject have been brought into undue prominence, and 
 others either scantily treated or omitted entirely. This is due 
 to the fact that the book does not profess to be a work on 
 pure chemistry, but only an introduction to the chemistry 
 of the metals. It is hoped, however, that nothing of importance 
 has been left out in the development of first principles ; 
 
viii PREFACE 
 
 and that the book will be found useful both to teachers and 
 students. 
 
 The author's thanks are due to Mr. A. H. Hiorns for read- 
 ing the proofs, and to Mr. C. R. Clark for executing the 
 drawings. 
 
 METALLURGICAL LABORATORY, 
 
 BIRMINGHAM MUNICIPAL TECHNICAL SCHOOL. 
 
 November, 1903. 
 
CONTENTS 
 
 CHAPTER I 
 INTRODUCTORY 
 
 PAGES 
 
 Physical properties of bodies Physical states Chemical change 
 
 Energy 1-6 
 
 CHAPTER II 
 METALS AND OXYGEN 
 
 The atmosphere Oxidation of common metals when heated in air 
 Burning of phosphorus and charcoal Rusting of iron 
 Rough determination of the composition of air - 7-20 
 
 CHAPTER III 
 METALS AND WATER 
 
 Oxidation of common metals by liquid and gaseous water Decom- 
 position of water by the electric current and by heat Composi- 
 tion of water by volume - - 21-32 
 
 CHAPTER IV 
 PROPERTIES OF MATTER 
 
 Definition of matter Subdivision of matter Composition of red 
 oxide of mercury Elements and compounds Atoms and mole- 
 cules Chemical symbols Law of Avogadro Metals and non- 
 metals Classification of common elements Chemical formulae 
 and equations - - 33-54 
 
 ix 
 
x CONTENTS 
 
 CHAPTER V 
 METALS AND SULPHUR 
 
 PAOES 
 
 Sulphur Metals and sulphur Formation of sulphides Action of 
 heat and air on sulphides Preparation and properties of 
 sulphuric acid Determination of formulae of compounds - 55-68 
 
 CHAPTER VI 
 COMMON ELEMENTS AND COMPOUNDS 
 
 Common salt Hydrochloric acid Chlorine, bromine, and iodine 
 Potassium chlorate Saltpetre Nitric acid Ammonium 
 chloride and ammonia Caustic soda and potash - 69-84 
 
 CHAPTER VII 
 METALS AND ACIDS 
 
 Reactions of common metals and acids Volume of gas libera d 
 Metallic compounds formed Hydrogen Sulphur dioxide - 
 Oxides of nitrogen Iron and acids Alloys and acids GenciU- 
 tion of an electric current by dissolution of zinc in dilute 
 sulphuric acid Decomposition of metallic salts by an electric 
 current - - - - - - 85-115 
 
 CHAPTER VIII 
 
 CHEMICAL EQUIVALENTS AND ATOMIC WEIGHTS OF 
 COMMON METALS AND NON-METALS 
 
 Determination of equivalents of common metals and non-metals 
 Vapour density Atomic heat Isomorphism Atomic weights 
 Valency of elements - 116-134 
 
 CHAPTER IX 
 OXIDES, ACIDS, AND SALTS 
 
 Classification of oxides Formation of salts Crystallization of 
 salts Nomenclature of salts Chlorides, nitrates, sulphates, 
 carbonates, and other common salts Basicity of acids - 135-162 
 
CONTENTS 
 
 CHAPTER X 
 CARBON AND ITS COMPOUNDS 
 
 Occurrence of carbon Burning of carbon Composition of carbon 
 dioxide Native and prepared carbonates Action of heat on 
 carbonates Acids and carbonates Carbonic oxide : its forma- 
 tion and use Coal Combustion of coal in open and closed 
 grates "Air "gas Coal gas - 163-181 
 
 CHAPTER XI 
 REDUCTION 
 
 Action of heat and of heat and reducing agents on metallic oxides 
 and other compounds Carbon, hydrogen, carbonic oxide, 
 hydrocarbons, metals, and potassium cyanide as reducing 
 agents Reduction by reaction The electric current as a 
 reducing agent - - - 182-199 
 
 CHAPTER XII 
 COMBUSTION 
 
 Heat developed by the burning of common combustible bodies 
 Ignition-point Zone of combustion Flame Luminosity of 
 flame Transformation of chemical energy into heat during 
 combustion Measurement of heat - 200-220 
 
 CHAPTER XIII 
 PHOSPHORUS AND ITS COMPOUNDS 
 
 Common properties of phosphorus Extraction of phosphorus 
 Phosphoric oxide, phosphoric acid, and phosphates Metals 
 and phosphorus - - 221-225 
 
 CHAPTER XIV 
 
 SILICON AND ITS COMPOUNDS WITH OXYGEN AND 
 METALS 
 
 Silicon Silica Formation of silicates Natural and prepared 
 silicates Fusibility of silicates Nomenclature of silicates 
 Fireclay Glass Effects of complexity on the fusibility of 
 silicates General classification - - - 226-235 
 
xii CONTENTS 
 
 CHAPTER XV 
 WEIGHTS, MEASURES, AND APPARATUS 
 
 PAGES 
 
 The Metric System of weights and measures Measuring apparatus 
 
 The balance and its use - - 236-243 
 
 APPENDIX 
 
 Table of elements and atomic weights Principles used in the cor- 
 rection of the volume of a gas for variations in temperature, 
 pressure, and amount of water vapour - 244-247 
 
 INDEX 249-252 
 
INTRODUCTION TO 
 
 METALLURGICAL CHEMISTRY 
 
 CHAPTER I 
 INTRODUCTORY 
 
 THINGS which appeal to us from outside do so through the 
 agency of our senses. We are thus able to recognise the 
 different objects of the external world which come under our 
 notice, and to distinguish between them by means of those of 
 their properties which affect our sensory organs. 
 
 Natural objects are called substances when they are con- 
 sidered with regard to the quality of the material of which 
 they are made, without thinking about its quantity or the 
 position which the objects occupy in space. Thus we speak of 
 water, coal, and lead as substances ; but of a glass of water, a 
 lump of coal, and a piece of lead as bodies. In the first case 
 the reference is general, and the quality only of the material 
 is recognised ; while in the second case the quantity of the 
 material and its position, as well as its quality, are clearly in 
 mind. 
 
 The common and obvious characteristics of bodies by which 
 they are usually recognised belong for the most part to what 
 are called their physical properties. The general appearance 
 of a body and the various changes which may take place in 
 it, without altering the composition of its substance, are due 
 to its physical properties. The term "matter" is used to denote 
 
 1 
 
2 METALLUKGICAL CHEMISTKY 
 
 the substances of bodies generally, and any change in this 
 matter causes a change in the properties of the bodies which 
 contain it. 
 
 Physical States Of Matter. The substance of bodies 
 exists in three more or less well-defined physical states the 
 solid, liquid, and gaseous. A stone affords a good illustration 
 of a solid body, water is the typical liquid, while the atmo- 
 sphere is the most important gaseous body. Although 
 examples of solids, liquids, and gases are familiar to everyone, 
 it is necessary to consider their general properties sufficiently 
 to be able to distinguish clearly between them. 
 
 Solids. A solid has a definite external form, which it 
 retains as long as it is left to itself. If subjected to pressure 
 in any direction, it resists that pressure without requiring its 
 sides to be supported. A change in form may take place 
 under sufficient pressure, but when the resistance to this change 
 is equal to the pressure which is producing it, the new form is 
 retained or changes but very slowly, as long as the pressure is 
 constant. A hard solid, such as steel, offers great resistance 
 to deformation ; but a soft one, such as rubber, changes its 
 form very rapidly under pressure. Both bodies, however, 
 regain their original form on removal of the pressure, if it has 
 not exceeded their power of recovery. The common metals, 
 with the exception of mercury, are well-defined solids. 
 
 Liquids. A liquid must, on the other hand, have lateral 
 support if it is to retain any shape, for it has a tendency to 
 flow to a lower level. The sides of the vessel which contain 
 it supply this support, and thus prevent it from flowing. This 
 tendency to find a lower level causes a liquid to settle down in 
 the containing vessel, and present a definite limiting surface, 
 which is practically flat and horizontal, if the body of liquid 
 is small. With liquids which wet the sides of the vessel the 
 general surface is raised a little near the sides ; but with 
 liquids which do not wet them a slight depression is 
 
INTRODUCTORY 8 
 
 observed. These effects are noticed when water and mercury 
 are poured separately into dry test-tubes. 
 
 Some apparently solid bodies, such as cobbler's wax, show 
 a decided tendency to flow if left to themselves; and some 
 liquids, such as treacle, are so thick and viscous that they 
 flow very sluggishly. Most of the useful metals, though well- 
 defined solids if left to themselves, may be forced to flow 
 under the hammer, between rolls, or in the press, and so made 
 to take up new forms, which, however, they retain after the 
 operation. A large number of solids, including the metals, 
 may be converted into liquids if made hot enough, and this 
 without any change in their composition. 
 
 Gases. Gases possess the important and distinctive 
 property of indefinite expansion. Thus, if a small body of gas 
 be allowed to enter an empty vessel, it will expand and entirely 
 fill the vessel, irrespective of its size. The gas thins out in 
 the process, and if it expands through a very large space, 
 becomes highly attenuated ; its surface is limited only by the 
 sides of the vessel. This is equivalent to saying that if the 
 pressure of the gas were made very small, its volume would 
 become greater than that of any vessel into which it could be 
 put. A gas expands equally in all directions if unrestrained, 
 and so must completely fill any vessel which contains it. 
 
 Measurable volumes of gases fulfil, more or less accurately, 
 well-known laws when they are heated and compressed. 
 Also, gases mix together or penetrate each other very readily, 
 which is no doubt due to their property of expansion. Thus 
 a small quantity of an odorous gas liberated in a room will 
 rapidly expand through the air of the room, and may soon be 
 detected in any part of it. It is impossible to keep a gas in 
 an open vessel, for some of the gas escapes and air enters to 
 take its place. 
 
 Physical Change. Many bodies may be made to assume 
 all three physical states in succession, by heating or cooling 
 them. Thus, water cooled below its freezing-point becomes 
 
 12 
 
4 METALLURGICAL CHEMISTRY 
 
 ice, and, if heated continuously at its boiling-point, is con- 
 verted into steam. But the material constitution of the 
 body remains the same ; there is no alteration either in the 
 composition or the weight of its substance. All metals can 
 either be melted or converted into vapour at varying tempera- 
 tures, which in some cases are comparatively low, and in 
 others extremely high. 
 
 Chemical Change. There is, however, another kind of 
 change which bodies undergo when the necessary conditions 
 are observed. It is such as to bring about an alteration in 
 the material constitution of the body or bodies in or between 
 which it takes place. This is called chemical change, and is so 
 characteristic as to be easily recognised. It depends upon 
 the chemical properties of the interacting bodies, and requires 
 them to be in very close contact. Even then it is often 
 necessary to start the action by the application of heat. For 
 example, gunpowder may be stored for years without change ; 
 but if a light is applied to it the action commences, and the 
 change takes place very rapidly. The interaction of the 
 charcoal, sulphur, and saltpetre, of which the powder is made, 
 results in the formation of a very large volume of gas, which 
 in its expansion produces some of the observed effects. 
 
 Energy. Tt is such a common experience to see bodies in 
 motion, that we are apt to take the consequent changes in 
 their position as a matter of course, and scarcely give a 
 thought to the cause of the motion. When a man lifts a 
 heavy weight he experiences a sense of muscular effort, to 
 which he ascribes the movement of the weight. The cause of 
 the motion in this case is called muscular force. In the 
 same way any change in the position or motion of a body is 
 due to the action of some kind of force upon it. The study 
 of the visible motions of bodies forms a branch of mechanics 
 in which the term "force" has a definite meaning, which, 
 however, need not at present be discussed. On the other 
 hand, the parts of a body may be in motion, and this motion 
 
INTRODUCTORY 5 
 
 may be increasing or decreasing without any motion of the 
 body as a whole being visible. Thus, when a metal bar is 
 heated it expands, and although this expansion is invisible to 
 the unaided eye it is easily demonstrated by actual measure- 
 ment. Now, it is evident that the whole body could not 
 expand unless its smallest parts or particles either expand 
 themselves or move wider apart. If it is assumed that the 
 particles move wider apart, the heat which is imparted to the 
 body must be the cause of their increased motion. If then it 
 is necessary that force should be exerted upon a body of 
 sensible size in order to set it in motion, it must also be 
 necessary in the case of a very small body, or particle ; the 
 only difference in the two cases is in the relative magnitudes 
 of the forces required. 
 
 It follows from this that heat imparted to the metal must 
 exert force upon its particles, and so cause them to become 
 wider apart. There is every reason to believe that the ulti- 
 mate particles of the metal are already in motion, and that 
 the general effect of the additional heat is merely to increase 
 the rate of motion and range of these particles. The cause of 
 the increase may be called heat force, but a better name for 
 it is heat ensrgy. It must be understood that the motion 
 of the particles of a solid body due to heating it, is of a very 
 restricted character, and the spaces through which they move 
 very small. 
 
 When the charge of gunpowder in a cartridge is ignited, the 
 bullet in front of it is caused to move very rapidly, and through 
 a considerable distance. Now the energy which brings about 
 this motion must be stored up in the powder, for it is evident 
 that the mere blow of the hammer on the detonator will not 
 account for the effect produced. Thus, when gunpowder is 
 bought, the effective portion of the purchase is the energy 
 which is stored up in the material of the powder. Further, 
 this energy can be measured by the amount of work it will do 
 when set free. So that it is legitimate to speak about quantity 
 of energy. 
 
6 METALLUKGICAL CHEMISTRY 
 
 As the energy of the exploding gunpowder is developed, 
 the matter of the powder itself undergoes a change in its 
 properties, and new bodies are formed. This is a chemical 
 Change, and the energy which has been transformed into 
 heat, mechanical work, or other form of energy, was chemical 
 energy. It appears to be associated in some way with the 
 bodies between which the change takes place, but it is not 
 evident as long as they remain unchanged. It is, however, 
 possible to measure one form of energy indirectly, while it is 
 being transformed into another form of energy. 
 
 In dealing with bodies in general, and the processes going 
 on among them, two things have to be considered ; these are 
 the matter of which the bodies are composed, and the 
 energy invariably associated with it. It is most probable 
 that the association of the two makes the bodies what they 
 are, and it is certain that any alteration in this relation must 
 cause a change of some kind in the bodies themselves. If 
 during this change no new bodies are formed, it is physical in 
 character, but if new bodies are formed from the old ones, 
 then it is a chemical change. 
 
CHAPTER II 
 METALS AND OXYGEN 
 
 THE atmosphere which surrounds the earth is so essential to 
 life that we become familiar with it at an early age, and 
 unconsciously learn to recognise it by its general properties. 
 Although invisible, it appeals readily to the sense of touch, 
 especially in a high wind. The sensation thus produced is 
 that caused by a substance striking against the exposed parts 
 of the body, and compels us to admit that something material 
 is appealing to us through our senses. If, then, it is admitted 
 that the atmosphere, although invisible, is matter or substance, 
 an investigation of those of its properties which do not appeal 
 to us directly may be commenced. Such work, if methodi- 
 cally carried out, is sure to repay the time and trouble 
 expended upon it. 
 
 It will be assumed that the student knows very little about 
 the atmosphere, except that which he has acquired by every- 
 day experience. If the experiments described below, or as 
 many of them as time and the apparatus at command will 
 permit, are carried out, a considerable amount of information 
 about the properties of the bodies under manipulation, and 
 the changes they undergo in the presence of air, will be 
 obtained. As skill in manipulation is only acquired by 
 practice, it is well to bear in mind that an operation, although 
 only partially successful, is sufficiently so if it makes the 
 point it is intended to illustrate clear. 
 
 In the following experiments the atmosphere plays as 
 important a part as the other bodies, which are either liquid 
 
 7 
 
8 METALLURGICAL CHEMISTRY 
 
 or solid, and are therefore under direct observation. The 
 chief aim is to demonstrate the properties of the air, and the 
 behaviour of metals in contact with it. 
 
 ACTION OF HEAT AND AIR ON METALS. 
 
 Lead. This metal passes readily from the solid to the 
 liquid state when heated. It has the characteristic properties 
 of the common metals. The sheet lead used by plumbers has 
 a dull appearance, but if the surface is scraped with a knife, 
 it becomes bright, and shows the lustre common to metallic 
 bodies. It is soft, being readily cut with a knife, and can be 
 beaten or rolled into sheets: it possesses the property of 
 malleability in a marked degree. A metal is said to be 
 malleable when it can be hammered or rolled into sheets 
 without cracking at the edges. 
 
 EXP. 1. Test the softness and malleability of a piece of lead by 
 cutting it with a knife, and by hammering it. 
 
 The surface of molten lead when exposed to the atmosphere 
 undergoes a gradual change, which is now to be investigated. 
 
 EXP. 2. Cut two squares of glazed paper, about 3 by 3 in., put 
 one in each scale-pan of the balance (Chap. XV.), and make them 
 counterpoise. Roughly weigh about 25 grams * of sheet lead, add 
 to it a scrap of tin,t about 0*25 gram, and weigh the whole accu- 
 rately. Transfer the lead to an iron pan about 6 inches in diameter 
 
 (an ordinary sand-bath pan 
 will do), and place the pan on 
 a circular Bunsen burner, such 
 as that shown in Fig. 1 ; or on 
 a tripod-stand over an ordinary 
 Bunsen flarne. When the lead 
 has melted drop in the tin, and 
 
 I stir well with an iron scraper 
 
 p IG 1< made of a piece of narrow hoop 
 
 iron about 15 inches long, and 
 
 bent as shown in Fig. 1. Allow the metal to remain at rest for a 
 
 * Grain weights may be used instead of grams : in that case multiply 
 the number of grams by 15'5 to convert into grains. 
 
 t The presence of a small quantity of tin as an impurity in the lead 
 increases the rapidity of the action and renders the experiment less 
 tedious than with pure lead. Why it does so need not be considered now. 
 
METALS AND OXYGEN 
 
 time, and watch the rapid change in the appearance of its surface. 
 Resume the stirring, and keep the molten metal broken up into 
 globules, so as to expose as large a surface as possible to the air. 
 In from thirty to forty minutes the metal will be converted into a 
 coarse greenish powder, when the pan may be removed and allowed 
 to cool. When cold, transfer the powder with the aid of a small 
 brush to the counterpoised paper upon which the lead was weighed, 
 and weigh it. 
 
 EXAMPLE. Weight of powder ... 25*615 grams 
 
 metal 25-312 
 
 Increase O303 gram 
 
 This operation is technically known as dPOSSingf, and if a 
 longer time is taken in carrying it out, a larger increase in 
 weight is obtained. But the experiment may be considered 
 satisfactory if any actual increase in weight is shown, as it is 
 somewhat difficult to remove the whole of the powder from 
 the pan. This is especially the case when some of the dross 
 has softened, through getting too hot, and has stuck to the 
 pan. 
 
 EXP. 3 Transfer the dross obtained in the above experiment to 
 a clean porcelain mortar ; add about 10 c.c. of water, and grind up 
 the contents with the pestle ; stir well with the finger, and pour off 
 the muddy yellow liquid into a porcelain basin ; repeat the opera- 
 tion until the water runs away nearly clear. Allow the basin to 
 stand for a few minutes in order that the suspended solid matter 
 may settle. Then pour off the clear liquid, and carefully dry the 
 residue over gauze or on the sand-bath. When dry, remove the 
 residue from the basin, examine it, and reserve it for future use. 
 
 The light yellow powder thus obtained corresponds very 
 nearly to the ordinary massicot (or litharge) in the laboratory 
 bottle. It is usually lighter in colour than the commercial 
 substance, but the longer the dressing stage is continued, the 
 deeper the colour of the product. The washing operation is 
 called levigation, and if the residue from it is examined, it is 
 found to consist of small particles of unchanged lead, which 
 are readily recognised by flattening some of the larger ones 
 on an anvil and scraping them. It is thus evident that at 
 any intermediate stage in the drossing operation the dross 
 consists of the yellow powder and the unchanged metal. As 
 
10 METALLURGICAL CHEMISTRY 
 
 the action proceeds the proportion of the unchanged lead gets 
 smaller, until finally the whole of the metal is converted into 
 massicot. 
 
 EXP. 4. Put half the massicot obtained from several experiments 
 on one of the pans used for the dressing operation ; place the pan 
 on a piece of gauze over a moderate Bunsen flame ; stir the powder 
 occasionally, and continue the heating for twelve hours. The whole 
 operation need not be continuous, but the heating may be carried 
 out from time to time as opportunity occurs. 
 
 At the end of each heating the colour of the powder is 
 deeper than at the beginning. Finally, it is of a puce colour 
 when hot, and bright red when cold. It will bear comparison 
 with the red lead in the laboratory bottle. The operation, 
 which is technically known as colouring", should be shared 
 by those students who have contributed massicot to the 
 charge. The pan may be heated over a spare burner, and 
 needs very little attention. 
 
 The manufacture of massicot and red lead on the large 
 scale is carried on under similar conditions to the above ; but 
 reverberatory furnaces are used, and charges of a ton or more 
 dealt with in one operation. 
 
 Copper. This metal, which has a characteristic red colour, 
 is so very malleable and ductile that it is readily obtained in 
 the form of thin sheet or fine wire. If the end of a thin 
 copper wire is held in the hottest part of the Bunsen flame, it 
 melts. This indicates that the melting-point of the metal is 
 only moderately high. 
 
 EXP. 5. Cut a strip of thin sheet copper, about 1-5 inches wide 
 and 4 inches long, roll it into a loose coil, and weigh it accurately. 
 Put the coil on a clay roasting dish, and place it in a moderately 
 hot muffle. Let it remain there for about half an hour ; then take 
 out the dish with the coil on it, and, when cold, reweigh the coil. 
 Note the weight. Put the coil on a piece of smooth paper to prevent 
 loss, and squeeze it with the fingers. The copper has probably all 
 disappeared, and a brittle dark-red mass remains. Grind it to a 
 fine powder in a mortar, and transfer the powder to the dish. Put 
 the dish back into the muffle, and stir the powder from time to 
 time. Leave it in the muffle as long as convenient. When cold, 
 
METALS AND OXYGEN 11 
 
 brush the powder carefully from the dish to the piece of counter- 
 poised paper, and weigh it. The temperature of the muffle should 
 not be above a moderate red heat, or the powder will soften and 
 stick to the dish. 
 
 EXAMPLE. Weight of powder after 2nd heating 3'720 grams 
 coil 1st 3-502 
 
 . 3-000 
 
 Increase in weight ... ... 0'720 gram 
 
 A piece of the copper scale, if examined before being 
 powdered, will be found to be blue-black on the outside and 
 dark red inside. It is a mixture of two distinct bodies, one 
 black and the other red, in varying proportions. The pro- 
 portion of the black body increases with the time of exposure. 
 The powder obtained by crushing the scale is usually dark 
 red in colour, but after further heating it becomes black. On 
 prolonged heating in contact with air the weight of the 
 powder finally becomes constant. A muffle furnace gives the 
 best result but if one is not available the experiment may be 
 made as described below. 
 
 EXP. 6. Bend a strip of sheet copper into the form of a narrow 
 U ; weigh it, and suspend it on a piece of clay-pipe stem over a 
 Bunsen flame in such a position as to get the best heating effect 
 when the flame is inside the U. Place a sheet of paper under the 
 burner, and continue the heating for half an hour. Kemove the 
 strip, place it on a convenient support, and cover it at once with a 
 dry jar or beaker. When the metal is cold, collect any scale which 
 may have separated from the strip, and reweigh the whole. 
 EXAMPLE. Weight of strip after heating ... 10*580 grams 
 
 before ... 10'542 
 Increase 0'038 gram 
 
 As soon as the strip begins to cool, the thin scale formed by 
 the action of heat and air upon the metal commences to peel 
 off. This is caused by the rapid contraction of the cooling 
 metal. The change is due to surface action, so that the more 
 finely divided the metal is, the more rapidly it takes place. 
 With the finely-divided metal the maximum increase in 
 weight is obtained in a comparatively short time. 
 
 EXP. 7. Put a gram of finely-divided copper, either fine filings 
 or reduced copper (Chap. XI.), into a weighed porcelain crucible, 
 
12 METALLUEGICAL CHEMISTBY 
 
 and reweigh it. Support the crucible on a pipeclay triangle over a 
 good Bunsen flame. Observe the change which takes place, and 
 continue the heating for fifteen minutes. When the crucible is cold, 
 reweigh it. Grind the residue in a mortar, and examine it. 
 EXAMPLE. 
 
 Weight of crucible and copper after heating 11-900 grams 
 before ... 11-647 
 Increase 0*253 gram 
 
 The difference between the black, readily-powdered solid 
 and the metal from which it is formed is very marked. 
 
 Tin. This metal is soft, white, and very malleable. It 
 melts easily, but the change due to the combined action of 
 heat and air takes place slowly at temperatures below a red 
 heat. At a red heat, however, the change is fairly rapid, and 
 is easily effected in a muffle. It is more rapid when impure 
 metal is used. If exposed to too high a temperature, the 
 metal burns and white fumes escape. 
 
 EXP. 8. Weigh carefully 3 grams of tin, with which is included 
 0*05 gram of lead. Place the metal in a roasting dish, put the dish 
 into the muffle, and keep it at a moderate red heat for an hour. 
 Stir the metal occasionally with an iron scraper, and if it shows 
 signs of burning, draw the dish forward into a cooler part of the 
 muffle. Remove the dish, and, when cold, remove and weigh the 
 residue. Grind up the white powder in a mortar, put it back into 
 the dish, and heat it again for half an hour, with occasional stirring. 
 Weigh again, and note if any increase is obtained. 
 
 EXAMPLE. Weight of metal after heating ... 3-793 grams 
 before ... 3-000 
 
 Increase 0-793 gram 
 
 Compare the white powder with the metal, and with the "putty" 
 powder in the laboratory bottle. 
 
 Iron. This metal undergoes a somewhat rapid change 
 when heated in the air, and the black scale that collects round 
 an anvil on which red-hot iron is being forged is familiar 
 evidence of this change. 
 
 EXP. 9. Clean a few feet of thin iron wire by rubbing it with 
 emery-cloth, and hammer or roll it into a thin ribbon. Weigh 
 2 grams of the ribbon, and heat it to a bright red heat on a clay 
 dish in the muffle for half an hour. Allow the dish to cool, grind 
 up the product in a mortar, and, if it is not all reduced to powder, 
 
METALS AND OXYGEN 13 
 
 put the residue back into the muffle, and continue the heating. 
 When the whole has been reduced to powder, weigh it and note 
 the increase in weight. 
 
 EXAMPLE. 2 grams of iron increased to 2*766 grams of scale. 
 
 The general change which many metals undergo when 
 heated in the air may also take place at the ordinary atmo- 
 spheric temperature, though much more slowly. It is very 
 marked in the case of iron, especially if the air is moist and 
 contains acid vapours. On the other hand, tin will keep a 
 bright surface for a long time under ordinary circumstances, 
 and is used as a coating material for iron to protect it from the 
 rusting action of the atmosphere. 
 
 BURNING BODIES AND Am. 
 
 The visible change which takes place when a candle burns 
 is very familiar, for ordinary observation teaches that it dis- 
 appear s during the process, and that light and heat are given 
 out Further information is easily ob- 
 tained by simple experiments. 
 
 EXP 10. Fix a short piece of candle in a 
 deflagrating spoon, ignite it, and place it in 
 a large dry bottle, as shown in Fig. 2. At 
 first it burns as brightly in the bottle as in 
 the air outside, but the flame gradually 
 dwindles, and finally disappears. Also, the 
 sides of the bottle are covered with moisture. 
 Now pour a little clear lime-water into 
 another bottle, shake it up, and repeat the 
 burning of the candle in this bottle. When 
 the flame has disappeared, remove the candle 
 and shake the bottle again. The lime-water 
 is turned milky. 
 
 Three facts are demonstrated by this F 
 
 experiment : (1) The burning of the 
 candle is dependent upon the presence of air, for if the 
 quantity of air is limited the duration of the burning is 
 limited also; (2) moisture is formed during the burning; 
 (3) the air is so changed by the action that lime-water is 
 turned milky by it. 
 
14 
 
 METALLURGICAL CHEMISTRY 
 
 EXP. 11. Pour some lime-water into a wide-necked bottle. 
 Place a piece of charcoal on the bowl of the deflagrating spoon, 
 and hold the spoon in the Bunsen flame until the charcoal is red 
 hot. Then put it into the bottle, and allow it to remain there for a 
 minute or two. Remove the spoon, and shake the bottle. The 
 lime-water is turned milky, thus indicating that the burning of 
 charcoal in air produces one of the changes noticed during the 
 burning of a candle in the same body. 
 
 T he burning of phosphorus and the rusting of iron may be 
 made to give some important information about the nature of 
 the air in which these actions take place. 
 
 Phosphorus. This body is a yellow, waxy-looking solid, 
 which takes fire so readily that it must be kept in water. 
 
 EXP. 12. Fit up the apparatus shown in Fig. 3. A is a glass bell 
 jar of about 2 litres capacity, fitted with a well-greased stopper or a 
 rubber bung. The jar is first graduated by inverting it with the 
 
 stopper in position, and pour- 
 ing in 1,200 c.c. of water from 
 a c.c. measure (Chap. XV.). 
 The water-level is marked on 
 the side of the jar, 300 c.c. 
 more water added, and the 
 second level also marked. The 
 jar is then placed over water 
 in the glass trough B, and so 
 supported that the level of the 
 water coincides with the 1,500 
 c.c. mark when the stopper is 
 out. In this way 1,500 c.c. of 
 air is enclosed in the jar when 
 the stopper is in position, and 
 as the water in the jar forms 
 a movable bottom the air can 
 expand or contract without 
 
 gas escaping from or entering the enclosed space. A small por- 
 celain crucible containing a piece of dry phosphorus about the size 
 of a pea is floated on the water inside A. The stopper is removed, 
 the phosphorus touched with the hot end of a piece of wire, and the 
 stopper rapidly replaced. The phosphorus burns, and a dense white 
 vapour fills the jar. In a short time the flame disappears, and the 
 white vapour is dissolved by the water, which gradually rises up the 
 jar to the 1,200 c.c. mark. When the white fumes have disappeared, 
 it is seen that about one-fifth of the original volume of gas has been 
 removed by the burning phosphorus, and that the water has risen in 
 the jar to take its place, 
 
 B 
 
 FIG. 3. 
 
METALS AND OXYGEN 
 
 15 
 
 Water is now poured into B until the level of the liquid inside and 
 outside A is the same. This will allow the stopper to be removed 
 without any air entering the jar through the neck, as would be the 
 case if the water-level inside had to fall. The stopper is then 
 removed, and a burning candle lowered into the jar. The flame 
 disappears, thus proving that the residual gas is no longer able to 
 support the combustion of the candle. 
 
 As the phosphorus burns, a body is formed which is a 
 white solid at the ordinary temperature. This is easily shown 
 by burning a piece of phosphorus under a dry bell jar standing 
 on a glass plate. As the apparatus cools, a snow-white deposit 
 forms on the glass plate. With suitable apparatus it can be 
 proved that the white solid is heavier than the phosphorus 
 from which it is formed. This increase in weight comes from 
 the air, and Exp. 12 indicates clearly that as the burning 
 progresses the volume of the air concerned in it diminishes, 
 for as the new body is dissolved the water rises in the jar to 
 take the place of that portion of the air which took part in its 
 formation. 
 
 Iron. The rusting of iron takes place slowly under normal 
 circumstances, but if the metal 
 is finely divided and damp, the 
 action is more rapid. 
 
 EXP. 13. Fit up the apparatus 
 shown in Fig. 4. A is a glass 
 cylinder of about 500 c.c. capacity. 
 B is a deep glass or earthenware 
 dish in which A can be inverted 
 and supported above the bottom 
 in any convenient manner. A is 
 divided into five equal parts. To 
 do this, fill the cylinder with water, 
 pour it into a measuring cylinder, 
 and note the volume. Then pour the 
 water back one-fifth at a time, mark- J?IG. 4 
 
 ing the side of the cylinder at the 
 
 level of each fifth. Cut a strip of filter-paper a little narrower than 
 the internal diameter of A ; thoroughly wet the paper, and then rub 
 some fine iron filings on both sides of it. Double the prepared strip, 
 and hang it over a thin wire support placed on C. Invert A over it, 
 
 
 !\ 
 
 F 
 
 A 
 
 
 
 =-^F. 
 
 SI^~P 
 
16 METALLURGICAL CHEMISTRY 
 
 and pour water into B until it nearly reaches the first mark on A. 
 Set the whole aside, and, if possible, inspect it from time to time, 
 and note the gradual rise of the water in A. This will go on for a 
 day or two and then cease entirely. If the water rises one-fifth up 
 the cylinder, it shows conclusively that one-fifth of the original 
 volume of air has been removed. 
 
 When the action is finished, pour water into B until the level is 
 the same inside and outside A. Raise A a little, and carefully draw 
 out the paper and wire -stand, but be careful to keep the mouth of 
 the cylinder below the surface of the water in B, or air will enter 
 and spoil the experiment. Pass a greased plate under the mouth of 
 A, make it tight, lift out the cylinder with the four-fifths of gas 
 and one-fifth of water in it, and place it mouth upwards on the 
 table. Remove the plate, test the gas with a lighted taper, and 
 quickly replace the plate. Then put the cylinder back into its 
 original position in B, and remove the plate. 
 
 When the paper used in the above experiment is examined, 
 the particles of iron on it show distinct signs of rusting. It 
 appears, then, that while the iron is rusting gaseous matter is 
 being absorbed from the air concerned in the action. Also, that 
 the action ceases when about one-fifth of the total volume of 
 the air has been used up. This coincides with the result 
 obtained in Exp. 12, and points to the conclusion that only a 
 portion of the air is active in promoting the burning of phos- 
 phorus and the rusting of iron. 
 
 EXP. 14. Weigh roughly 10 grams of red lead, put the powder 
 into a piece of moderately thin- walled combustion- tube, closed at 
 one end and about 6 inches long. Tap the tube until the powder 
 forms a uniform layer about two-thirds up the tube when it is hori- 
 zontal. Fit the tube with a cork and delivery-tube, and support it 
 in a horizontal position, so that the end of the delivery-tube can be 
 passed under the mouth of A, Fig. 4. (See also Fig. 14.) Heat the 
 tube uniformly in the Bunsen flame by holding the burner in the 
 hand and moving it about. Allow the air driven out by the pre- 
 liminary heating to escape. Then pass the end of the delivery-tube 
 under A, and strongly heat the red lead until the jar is again full of 
 gas. Continue the heating, and collect a little of the gas in a test- 
 tube which has been filled with water and inverted over water in B. 
 
 Remove the test-tube, placing the thumb over its mouth to prevent 
 the escape of gas, raise the thumb, and insert a glowing splint. The 
 wood bursts into flame, and burns much more vigorously than in 
 the air. 
 
 Put the greased plate under A, take out the jar, and put it in an 
 
METALS AND OXYGEN 
 
 17 
 
 upright position. Place an empty jar by the side of it, and lower a 
 burning candle, first into one and then into the other. The candle 
 is found to burn equally well in the two jars. 
 
 Clearly, then, the rusting iron absorbs that part of the air 
 which gives to it the property of supporting the burning of a 
 candle, and the red lead, when heated, gives up a gas which 
 restores this property to the residual air. An examination of 
 the tube shows that the red lead has changed back to the 
 yellow substance from which it was formed. It is evident that 
 the red lead is formed from massicot by the absorption of gas 
 from the air. 
 
 Mercury. There is another experiment which brings out 
 this absorption of a definite gas from 
 the air by the changing metal. But 
 it can only be carried out under ex- 
 ceptional circumstances by the ordinary 
 student, as it requires a moderately 
 high and uniform temperature to be 
 maintained for many hours. The fol- 
 lowing description will, however, be 
 readily followed, and may be useful 
 for reference later. 
 
 EXP. 15. Fit a 12-ounce glass flask 
 with a good cork, through which is passed 
 a piece of glass tube about 3 feet long and 
 { inch bore. Place it securely in a deep 
 sand-bath arranged on a retort-stand, as 
 shown in Fig. 5. Eemove the cork, and 
 pour in mercury until the bottom of the 
 flask is well covered. Replace the cork 
 and tube, and put a Bunsen burner under 
 the sand-bath. Heat the bath until the 
 metal begins to boil; lower the flame 
 until the boiling ceases, and then raise 
 and lower it until a little change, either 
 
 one way or the other, makes the mercury boil or takes it off 
 the boil. Keep the flask at this temperature (about 350 C.) for 
 two or three days. A powder collects slowly on the surface of the 
 mercury, and when removed is found to consist of red scales. 
 
 FIG. 5. 
 
18 METALLUEGICAL CHEMISTKY 
 
 If some of this red powder were used in place of the red 
 lead in Exp. 14, a gas with exactly similar properties to that 
 collected from the red lead would be obtained, and liquid 
 mercury would collect on the cool parts of the tube. 
 
 Mercury vaporizes rapidly at temperatures near its boiling- 
 point, and the principal use of the long tube shown in Fig. 5 
 is to act as a condenser, and so prevent the escape of metallic 
 vapour ; it also serves for the interchange between the air 
 inside and outside the flask, which is necessary for the con- 
 tinuance of the change. 
 
 The Energy of Chemical Change. The changes 
 
 brought about during the experiments described in this 
 chapter are sufficiently characteristic to make it quite evident 
 that something else happens besides an alteration in the 
 properties of the reacting bodies. We must take notice of the 
 development of heat which accompanies the disappearance of 
 the phosphorus and carbon during their combustion. It may 
 be said that the heat is due to the energy of combustion, but 
 as combustion is only a particular case of chemical action, it 
 will be better to substitute the more general phrase, "energy of 
 chemical action." Clearly, this energy must be associated in 
 some way with the combustible bodies and with that part of 
 the air which assists in their burning ; but it is absent from 
 the products of the combustion. It has been transformed into 
 heat, which is simply another form of energy. The energy of 
 chemical action is stored up in the bodies which take part in 
 the change, and, as the action progresses, is converted largely 
 into heat in the ordinary processes of combustion. But other 
 forms of energy, such as light and electricity, may also appear. 
 There is also a very definite relation between the quantity of 
 heat and any other forms of energy developed, and the quantity 
 of chemical energy transformed. There is very little doubt 
 that no actual loss of energy can possibly take place. As 
 one form disappears, other forms appear in exactly the same 
 proportion. This subject is more fully dealt with in Chap. XII. 
 
METALS AND OXYGEN 19 
 
 SUMMARY. 
 
 Several important facts have been demonstrated in the fore- 
 going experiments. 
 
 The metals lead, copper, tin, and iron, undergo very char- 
 acteristic changes when heated in air, and new bodies are 
 formed, which differ in such a marked way from the metals 
 themselves as to suggest complete alteration in their funda- 
 mental properties This is to be recognised as chemical 
 change. It must be carefully thought about, and every effort 
 made to understand what it means, as far as the experiments 
 and observations will explain it. 
 
 The change in the metals is accompanied by an increase 
 in weight, and it is probable that in thinking about this in- 
 crease the mind will picture to itself the addition to the metals 
 of something which can be weighed. 
 
 The rusting of iron in a confined volume of air shows that 
 a definite part of the air is absorbed, which is also that part 
 which causes a candle to burn in air. For a candle will not 
 burn in the residual gas. The restoration of this property of 
 supporting combustion to the residual air by the addition of 
 the gas obtained by heating red lead, seems to indicate that 
 the change in the metals the burning of a candle, charcoal, 
 and phosphorus are caused by the same substance. 
 
 It is a gas, and is called Oxygen. 
 
 The increase in weight is due to the absorption of oxygen, 
 and the action itself is called oxidation. When the action 
 is very rapid, as in the case of burning bodies, it is known as 
 combustion. 
 
 The residual gas, after removal of oxygen from air, will not 
 support the combustion of ordinary combustible bodies. It is 
 called Nitrogen. 
 
 However varied the experiments may be, the volume of 
 oxygen absorbed is always found to be about one-fifth the 
 volume of the air acted upon. The composition of pure air 
 is, then, roughly : Oxygen i and nitrogen 4 by volume. 
 
 22 
 
20 METALLURGICAL CHEMISTRY 
 
 Oxygen. This is the active constituent of the atmosphere, 
 and as such will claim considerable attention later. 
 
 Nitrogen. This gas is inactive, and for our purpose need 
 not be considered further as an isolated body. It will be 
 remembered as a colourless, transparent, odourless gas, which 
 will not allow a candle or other combustible body to burn in 
 it. It reduces the activity of the oxygen of the air by being 
 simply mixed with and diluting it. What was formerly 
 known as pure atmospheric nitrogen is now found to contain 
 a small quantity of a heavier and even more inert gas called 
 argon. 
 
 The various bodies which have been formed with the co- 
 operation of the oxygen of the air in the foregoing experi- 
 ments are called oxides ; and, as will be proved later, some 
 of them are gases, which, as they form, mix with the residual 
 air. If the actions take place in the open air, as is often the 
 case, the gaseous oxides formed diffuse into the air, and form 
 part of it. As such they may be simply impurities, or they 
 may be essential constituents of a good working atmosphere. 
 This will be considered in future chapters. 
 
 QUESTIONS. 
 
 1. Descri.be the changes which take place when lead 01 
 copper is heated for some time in contact with air. 
 
 2. What do you understand by the term combustible 
 body 1 Describe an experiment to help your explanation. 
 
 3. How would you demonstrate the properties and pro- 
 portions of the two principal constituents of the air 1 
 
 4. What is meant by chemical change, arid how does it 
 affect bodies between which it takes place 1 
 
 5. Explain exactly what happens to phosphorus when it 
 burns in air. 
 
 6. How is the process of rusting explained 1 Give an 
 example. 
 
CHAPTER III 
 METALS AND WATEK 
 
 Water. This substance is the typical liquid, and is as im- 
 portant as it is common. It comes next to air as an essential 
 to animal and vegetable life. In the pure state it is a clear, 
 tasteless liquid, which is quite colourless when viewed in bulk. 
 It has the property of dissolving a large number of solids, 
 liquids, and gases more or less readily, and on that account is 
 never found perfectly pure in nature. The natural waters 
 taken in the order of their purity are rain, river, spring, and 
 sea water. The impurities may be removed ; therefore all 
 these varieties furnish samples of one well-defined invariable 
 body pure water. Suspended impurities, which make 
 water look dirty, are easily removed by filtering. Dissolved 
 gases are for the most part got rid of by boiling the water. 
 Solids in solution are left as a residue on the complete 
 evaporation of the liquid, and the condensed steam is practi- 
 cally pure water. Liquid impurities are not so readily 
 removed, but by repeated evaporation their complete removal 
 is also possible. The pure water used for particular purposes 
 is known as distilled water, but ordinary drinking water is 
 sufficiently pure for most purposes. It may be convenient to 
 use the substance as ice, water, or steam, but whichever is 
 used the matter is the same. Ice melts or water freezes at a 
 constant temperature, and water boils at another temperature, 
 which is just as constant, if the pressure on its surface is also 
 constant. These are the standard " points " for the ordinary 
 thermometer, and on the Centigrade scale, which is commonly 
 
 21 
 
22 
 
 METALLURGICAL CHEMISTRY 
 
 used for experimental purposes, they are marked C. and 
 1 00 C. Further information about the physical properties of 
 water will be found in most works on physics or chemistry. 
 
 ACTION OF METALS ON WATER. 
 
 Sodium. This is a soft metal which may be readily cut 
 with a knife. The freshly -cut surface has the ordinary 
 metallic lustre, and somewhat resembles freshly-scraped lead 
 in appearance, but it absorbs oxygen from the air so readily 
 that it tarnishes very rapidly. The metal is usually kept in 
 some liquid, such as mineral naphtha, which is not acted upon 
 by it, so as to exclude the air. 
 
 . EXP. 16. Fill a small gas cylinder of about 100 c.c. capacity with 
 water, and invert it over water, as shown in Fig. 6. Be sure that 
 no air is left in the cylinder, or an explosion may occur. Cut a 
 
 piece of sodium about the size 
 of a pea ; drop it into the water 
 in the trough, rapidly place 
 over it a gauze spoon, and 
 force it under the mouth of 
 the inverted cylinder. Bubbles 
 of gas escape through the 
 meshes of the gauze, rise in 
 the cylinder, and gradually dis- 
 place the water. Repeat with 
 similar pieces of the metal 
 until the cylinder is filled with 
 the gas. Sometimes the gas 
 does not get through the gauze 
 readily ; in that case tilt the 
 spoon a little, and let the metal 
 escape, when it will rise to the 
 surface of the water in the 
 
 cylinder, and there continue its action upon the liquid. Put a 
 greased plate over the mouth of the cylinder, and remove it from 
 the water. Hold the cylinder mouth downwards, remove the plate, 
 and bring a lighted taper near. The gas takes fire, and burns 
 slowly with a yellowish flame. 
 
 The visible result of this experiment is that the metal 
 dissolves in the water, and in doing so liberates a combustible 
 gas, which can be collected and burnt. 
 
 FIG. 6. 
 
METALS AND WATER 
 
 23 
 
 Magnesium. This metal is usually sold in the form of 
 ribbon or wire. It has the ordinary metallic lustre, but 
 tarnishes slowly in the air. It is a very light body for a metal, 
 and when heated in the Bunsen flame unites with oxygen so 
 readily as to take fire and burn with a brilliant light, leaving 
 a white residue, which contains magnesium and oxygen. 
 
 EXP. 17. Take a piece of magnesium ribbon, 15 inches long, and 
 make it into a compact coil by wrapping it round a pencil. Draw 
 down a piece of thin |-inch combustion-tube in the blowpipe flame, 
 and bend the drawn part as shown in Fig. 7. Push the coil of 
 magnesium into the tube and draw off the other end, bending it the 
 
 FIG. 7. 
 
 same as before, and leaving the body of the tube about 5 inches 
 long. Pass one end through a rubber bung fitted in a small flask 
 containing water, and attach a delivery- tube to the other by a piece 
 of rubber tubing. Arrange the delivery-tube so that it can be 
 directed readily under the mouth of a cylinder of about 200 c.c. 
 capacity, which has been filled, and inverted over water in the 
 pneumatic trough. Boil the water in the flask, and when steam is 
 coming off freely, dry the body of the combustion-tube by carefully 
 moving a Bunsen flame over it. When there is no tendency for 
 water to run back from the upper bend, place the Bunsen flame 
 directly under the metal coil, so as to heat it as strongly as possible. 
 Be ready to insert the open end of the delivery-tube under the 
 mouth of the cylinder directly the metal begins to burn. As the 
 tube is hot from the passage of steam through it, a rubber bung 
 slipped down to the middle eaves the fingers in handling it. When 
 
24 
 
 METALLUKGICAL CHEMISTRY 
 
 the action has ceased remove the delivery-tube, cover the mouth of 
 the cylinder with a glass plate, and lift it out of the trough. Hold 
 it upside down, remove the plate, and bring a lighted taper to the 
 mouth. The gas burns similarly to that obtained in the last experi- 
 ment, but the flame is not quite so yellow. 
 
 On examination the residue in the combustion-tube is 
 found to be a white powder. It will be useful to compare 
 this powder with that obtained by igniting a piece of mag- 
 nesium ribbon in the Bunsen flame, and allowing it to burn in 
 the air. 
 
 [The apparatus shown in Fig. 7 is all that is necessary, but 
 it will be safer to fix the combustion-tube in a clip to keep it 
 steady.] 
 
 Iron. The general appearance and common forms of this 
 metal are well known, and it is only necessary to point out 
 that the sample used should be clean, and show the usual 
 metallic lustre. 
 
 EXP. 18. Clean a strip of thin sheet iron by rubbing it with 
 emery-cloth, and cut it up into narrow strips about 1 inch by y ff inch. 
 
 Push a plug of asbestos to a convenient position in a thin iron tube 
 about 15 inches long and inch diameter, and put the clean iron 
 through the other end of the tube in sufficient quantity to fill about 
 6 inches of it. Fit the ends of the tube with rubber bungs and 
 delivery-tubes, and arrange it in a small table furnace, as shown in 
 
METALS AND WATER 25 
 
 Fig. 8. Wrap pads of blotting-paper round the bungs, and keep 
 them saturated with water. This prevents the rubber from becoming 
 overheated. Arrange a small flask for boiling water, so that the 
 steam may be passed directly into the tube. Prepare to collect any 
 gas which may be given off from the other delivery-tube. Heat the 
 iron tube, and boil the water. When the steam comes into contact 
 with the redhot iron, gas is liberated, and collects in the gas cylinder 
 placed to receive it. 
 
 Fig. 8 shows the arrangement of the apparatus, and when 
 about one-third of the tube is filled with iron at a good red 
 heat, the gas is liberated rapidly, so that several jars can be 
 easily collected. Closely adhering scale forms on the surface 
 of the iron. The material of the tube also helps the reaction. 
 The most suitable tube for the experiment is the thin, solid- 
 drawn material used for bicycle work. With this, if a furnace 
 is not available, the reaction can be brought about by heating 
 the tube in a good Bunsen flame, but in this case the gas is 
 liberated slowly, and it would be tedious to collect any con- 
 siderable quantity. By using the foot bellows and blowpipe 
 a better result is obtained. 
 
 Experiments with the Gas. The first jar of gas may 
 
 be rejected, as it contains the air which was present in the 
 apparatus at the commencement. 
 
 1. Hold a jar of the gas upside down ; remove the plate, 
 and bring a lighted taper near to the mouth. The gas takes 
 fire, and burns quietly with a pale-blue flame, and when the 
 burning taper is pushed through the flame into the gas above, 
 it is extinguished, but is rekindled on being pulled back 
 through the flame. The gas itself will burn (combustible), 
 but will not allow a taper to burn in it (non supporter of 
 combustion). 
 
 2. Place a jar with the mouth upwards, remove the plate, 
 and allow it to stand for twenty seconds. Test with a burn- 
 ing taper. The gas has escaped, and the jar is now full of 
 air. 
 
 3. Repeat the last experiment, but with the mouth of the 
 
26 METALLURGICAL CHEMISTRY 
 
 jar held downwards. Very little gas is found to have escaped 
 when a lighted taper is brought to the mouth. 
 
 4. Hold an empty jar with the mouth downwards, and 
 bring a jar of the gas into the position shown in Fig. 9. 
 Keep them in this position for about half 
 a minute; separate the jars and put 
 plates on them. On examination the 
 lower jar is found to be free of, and 
 the upper one to contain, the greater part 
 of the gas, which now burns with a slight 
 explosion when a lighted taper is put to 
 the mouth of the jar. The last three 
 experiments prove that the gas has a very 
 great tendency to escape in an upward 
 direction from vessels containing it. It 
 p q ^ is, therefore, a very light gas. 
 
 5. Place a jar of air and a jar of the 
 
 gas mouth to mouth, and allow them to stand, one on top of 
 the other, with the air cylinder at the bottom, for fifteen 
 minutes. See that the ground tops of the cylinders are well 
 greased, so as to prevent any escape of gas. Separate them, 
 and put plates, first on the bottom and then on the top 
 cylinder. Test both with a lighted taper. Equally loud 
 explosions in the two jars show that the mixture of the air 
 and the gas is perfect. 
 
 The name of the gas is Hydrogen. It burns with a pale 
 blue flame. It will not support the combustion of a taper, or 
 of similar bodies. It is a very light gas, and can be poured 
 upwards. It is colourless, transparent, and odourless when 
 pure ; but as collected in the ordinary way it has an odour, 
 due to impurities derived from the materials used in its pre- 
 paration. The property of gases which enables hydrogen and 
 air to mix together perfectly, although the vessel containing 
 the lighter gas is placed at the top, is known as diffusion, 
 and is a common property of all gases. Any number of gases 
 will diffuse into each other, and make a uniform mixture, 
 
METALS AND WATER 27 
 
 providing they do not act upon each other chemically. The 
 composition of a mixture of gases is the same throughout. 
 The heavier gases do not usually separate to the bottom 
 and the lighter to the top, however long the mixture may be 
 kept. 
 
 Action of the Electric Current on Water. The use 
 
 of the electric current for lighting purposes is now more or 
 less familiar to everyone, and the presence of the white-hot 
 filament in the incandescent lamp is readily noticeable. The 
 heat and light developed in the lamp are due to the resistance 
 offered by the carbon thread joining the terminals of the 
 lamp to the current passing through it. These terminals are 
 connected with copper wires, which are connected in turn 
 with the electric mains through which the principal current is 
 passing, and a small quantity is thus taken off to be used in 
 the lamp. It is an easy matter to cut one of the wires con- 
 nected directly with the lamp, and to connect the severed 
 ends with a vessel containing water, so that if the current is 
 to pass through the lamp it must also pass through the water. 
 Now it is found that perfectly pure water prevents the passage 
 of the current, and no visible effect upon the liquid is noticed; 
 but if a little acid is added to the water the current passes and 
 the lamp glows. The nature of electricity and the cause of 
 its flow through the wire and the water cannot be discussed 
 at this stage ; but it will be sufficient for our present purpose 
 to say that it can be drawn off and used to bring about heat 
 effects and chemical change, just as readily as coal-gas can be 
 passed into a Bunsen burner, and the heat developed by its 
 burning used for a similar purpose. 
 
 The apparatus used to contain the water is called a volta- 
 meter, and consists essentially of a glass vessel through the 
 bottom of which two platinum wires, with narrow strips of 
 platinum attached, are fused. These are shown in Fig. 10. 
 The cut ends of the lamp wire are joined to the ends of the 
 platinum wires outside the vessel, and acidulated water is 
 
28 
 
 METALLURGICAL CHEMISTRY 
 
 poured into the vessel to fill in the gap between the two 
 platinum plates. The lamp glows, and bubbles of gas collect 
 on the plates, rise to the surface of the 
 \ / water, and escape into the air. The gas 
 \nn/ thus liberated can be collected for ex- 
 it?! \T K ffi| animation by filling two tubes with water, 
 p -Q and inverting them over the plates. 
 
 EXP. 19. Fit up the apparatus shown in 
 
 Fig. 11. This voltameter is a very convenient form of the apparatus. 
 The platinum plates, P and P, are in the tubes A and B, and their 
 wires pass to the outside through the glass. The glass taps at the top 
 of A and B are opened, and water containing a little acid is poured 
 down C until it rises to the level of the taps, which are then closed. 
 The water is now at the same level in the three tubes, and on 
 making the connec- 
 tions with the lamp 
 wires the current 
 passes through the 
 water, and gases are 
 liberated from the plat- 
 inum plates. The bub- 
 bles of gas rise in A 
 and B, and force the 
 water back into C. It 
 is at once noticed that 
 the volumes of the 
 gases collecting in the 
 tubes are not equal, 
 and when sufficient 
 has collected it is seen 
 that the simple rela- 
 tion of 1 to 2 is main- 
 tained as long as the 
 action continues. When 
 the level of the water 
 in A nearly reaches the 
 
 MAINS 
 
 FIG. 11. 
 
 top of the platinum 
 plate, the connection 
 
 is broken, and the liberation of gas ceases. The larger volume of 
 gas is now tested by cautiously opening the tap, and at the same 
 time holding a lighted taper near. The issuing gas, driven out by 
 the fall of the water level in C, takes fire and burns with a pale blue 
 flame. When the gas in B is allowed to escape by opening the tap, 
 and a glowing splint is held near the jet, the wood bursts into flame, 
 
METALS AND WATEK 
 
 29 
 
 and burns vigorously. These tests are sufficient to indicate the 
 nature of the gases. The larger volume of gas is hydrogen, and 
 the smaller one oxygen. 
 
 If the taps are turned off and the connections reversed, the 
 hydrogen collects in B, and the oxygen in A. This shows 
 that the position of the gases in the apparatus depends upon 
 the direction of the current through the liquid. It is usual to 
 state that the oxygen is liberated on the plate from which 
 the current enters the liquid, and the hydrogen on the plate 
 through which it leaves the liquid. The acid put into the 
 water at the beginning of the experiment is there at the end, 
 unaltered either in quality or quantity. This shows that the 
 gases must come from the water. Thus we have additional 
 evidence that water is a combination of oxygen and hydrogen, 
 and, further, that they are present in the proportion of 
 1 volume of oxygen to 2 volumes of hydrogen. A simpler 
 form of the apparatus similar to that shown in Fig. 10, in 
 which two tubes, filled with water, are inverted over the 
 platinum plates in the open vessel, may be used. It can be 
 worked readily with two Bunsen cells if the platinum plates 
 are not too far apart. 
 
 The two gases can also be collected together, and a very 
 explosive oxy -hydrogen mix- 
 ture obtained. 
 
 EXP. 20. A is a small bottle 
 containing acidified water, and 
 fitted with a rubber bung, 
 through which pass the two 
 platinum wires tipped with 
 platinum plates, and a delivery- 
 tube leading under the inverted 
 cylinder B. When the current 
 is passed through the water in 
 A, the mixed gases are given off 
 and collect in B. If, when a small cylinderful has been collected, 
 it is removed and a lighted taper put to the mouth, a loud explosion 
 takes place. The gases combine together again, and water is 
 reformed. 
 
 If the lighting current is not available for these experi- 
 
 FIG. 12. 
 
METALLURGICAL CHEMISTRY 
 
 ments a voltaic battery of three or four cells may be used. 
 (See Chap. VII.) 
 
 Action Of Heat on Water. When steam is raised to a 
 high temperature, it is split up into its constituents, oxygen 
 and hydrogen, but as the temperature falls recombination 
 takes place. If, however, the gases as they are liberated are 
 allowed to mix with some neutral gas, the recombination can 
 be prevented, and evidence of the decomposition obtained. 
 
 EXP. 21. A small long-necked flask, A, is fitted with a bung and 
 delivery-tube. A spiral of thin platinum wire is stretched between 
 
 the lower ends of 
 two stout copper 
 wires which pass 
 through the bung, 
 and terminate in the 
 middle of the body 
 of the flask. The 
 water in A is kept 
 gently boiling, and 
 the current is sent 
 through the spiral, 
 which is thus raised 
 to a bright red heat 
 by the resistance of 
 the wire to the pas- 
 sage of the current. 
 FIG. 13. The steam in direct 
 
 contact with the 
 
 spiral is decomposed, and the liberated gases diffuse into the main 
 body of the steam. Bubbles of gas issue slowly from the delivery- 
 tube, and collect in B. On testing the evolved gas, it is found to 
 be explosive, and to resemble that obtained in Exp. 19. The intense 
 heat of the wire decomposes the water ; the current makes the wire 
 hot, but takes no further part in the action. 
 
 B 
 
 SUMMARY. 
 
 The experiments of this chapter all lead to the conclusion 
 that water is made up of two gases, oxygen and hydrogen. 
 In its decomposition by the electric current the information 
 obtained is very precise, for the water is proved to contain 
 the two gases in the proportion of 2 volumes of hydrogen 
 
METALS AND WATEK 31 
 
 to 1 volume of oxygen. Also, when the gases are mixed 
 together in these proportions, and a lighted taper is put 
 into the mixture, recombination takes place with great 
 violence. 
 
 Hydrogen has very characteristic properties. It burns with 
 a pale blue flame, and, as will be proved later, water is formed 
 during the combustion. It is a very light gas, being readily 
 poured upwards from one vessel to another. It also diffuses 
 into other gases very readily. The heavier a gas is the more 
 slowly it diffuses. The general law states that the rate of 
 diffusion of gases varies inversely as the square roots of their 
 densities. Hydrogen is the lightest body known, and has, 
 therefore, the highest rate of diffusion. 
 
 Ice when heated melts to water; water when boiled is 
 converted into steam ; steam when very strongly heated is 
 split up into oxygen and hydrogen. The first two changes 
 are purely physical, for the actual composition of the substance 
 is unchanged. The last is a change in composition, for other 
 forms of matter appear in proportion as the water disappears. 
 This is a chemical change. 
 
 The common metals, with a few exceptions, such as sodium 
 and potassium, do not attack pure water at ordinary atmo- 
 spheric temperatures, and may be left in contact with it for a 
 long time without appreciable change. But this is not the 
 case with natural water, which always contains impurities, 
 and these often act either by themselves, or help to bring about 
 an action between the metal and water. The dissolved oxygen 
 and carbonic acid gases, which are absorbed from the air by 
 water when exposed to it, are active agents in the case of 
 several metals. Thus iron rusts readily, and lead is acted 
 upon somewhat rapidly by such impure water. The presence 
 of solids in solution is also effective in promoting change. 
 This is noticed in the corrosion of those metal portions of 
 ships which are exposed to the action of sea-water. A large 
 proportion of the common metals, when at a red heat, attack 
 water vapour readily, and in all such cases oxides of the 
 
32 METALLUKGICAL CHEMISTKY 
 
 metals are formed. Copper, lead, and tin act very slightly, if 
 at all, even at a bright-red heat. 
 
 QUESTIONS. 
 
 1. How would you obtain a sample of pure water 1 ? 
 Describe its chief physical properties. 
 
 2. What kind of body is sodium? State generally what 
 takes place when it is thrown on to water. 
 
 3. How may hydrogen gas be obtained in sufficient quan- 
 tity to test its common properties ? Name these properties. 
 
 4. Make a sketch of a simple form of voltameter, and 
 explain how you would use it. 
 
 5. Can an analogy be drawn between the action of heat 
 and of an electric current upon water ? 
 
 6. Explain in a simple way the difference between physical 
 and chemical change. 
 
 7. Describe the property of diffusion of gases. 
 
 8. What is the nature of the bodies formed when metals 
 act upon water vapour ? 
 
CHAPTER IV 
 PEOPEETIES OF MATTEE 
 
 MATTER is the general name given to the stuff or substance of 
 which all bodies are made up. A little thought will make it 
 clear that the total number of different bodies is very large, 
 and it would seem an almost endless task to attempt their 
 classification. Very little difficulty is, however, experienced 
 in arranging the majority of bodies in three great divisions 
 the animal, vegetable, and mineral kingdoms. 
 
 Mass and Density. It may be stated that the "mass," 
 or quantity of matter in a body, is a property of the matter 
 which accompanies it without loss or gain through every 
 possible change which the body can undergo. Mass is most 
 easily estimated in practice by weighing, so that the weight 
 of a body may be taken as the measure of the amount of 
 matter in it ; but it must be remembered that the terms 
 weight and mass have somewhat different meanings. The 
 weight of a body, as measured on a spring-balance, varies 
 according to the position of the body with regard to the 
 earth, for the pull of the earth upon the body varies with the 
 distance of its centre from the centre of the earth. The force 
 of gravity, which is the cause of the weight of a body, 
 decreases as the distance of the body from the earth's centre 
 increases, so that a given body would stretch the spring of the 
 balance more at a place near one of the poles than it would at 
 a place nearer the equator, as at the latter the distance of the 
 earth's centre from it would be greater. Thus, it may be said 
 
 33 3 
 
34 METALLUKGICAL CHEMISTEY 
 
 that the mass of a body depends solely upon the amount of 
 matter in it ; and its weight upon the latter, together with the 
 proximity of other bodies to it. But in ordinary weighing on 
 a beam-balance, the body placed in one pan is counterpoised 
 by the weights placed in the other, so that any change in the 
 force of gravity affects the body and the weights equally, and 
 for practical purposes a sufficiently true estimate of the mass 
 of a body is obtained by determining its weight in the usual 
 way. 
 
 The density of a body is usually stated in terms of the 
 quantity of matter in a unit volume of the body as determined 
 by weighing. But to be accurate, it is necessary to state the 
 temperature of the body in the case of solids and liquids, and 
 both the temperature and the pressure in the case of gases or 
 vapours at the time of weighing. All bodies expand when 
 heated and contract when cooled ; and variations in the 
 pressure to which a body of gas is subjected cause changes in 
 its volume which cannot be neglected even in rough experi- 
 ments ; but changes in the volume of solids and liquids from 
 the same cause are so small as to be negligible in the ordinary 
 way. Thus the density of gold is 19-3 grams per cubic centi- 
 metre at 15 C. ; that of water 1 gram per cubic centimetre at 
 4 C. ; and that of oxygen OO0143 gram per cubic centimetre 
 at C. and 76 centimetres barometric pressure. (See 
 Chap. XV. and Appendix.) 
 
 The weighing of solids and liquids is a comparatively easy 
 operation ; but in the case of gases it is much more difficult 
 and tedious. In the method of weighing gases adopted by 
 Regnault, two equal glass globes fitted with stopcocks, and 
 having a capacity of several litres, are used. One of the 
 globes is exhausted by means of a good air-pump, and then 
 suspended from a hook fixed to the bottom of the weight pan 
 of a very sensitive balance. This globe serves as a counter- 
 poise for the other one, whi.h is also completely exhausted 
 and similarly suspended from the weighing pan of the balance. 
 The globes are now carefully counterpoised, and the one to 
 
PROPEETIES OF MATTEE 35 
 
 be used for the actual weighing is filled with the pure gas at 
 a pressure slightly gre tter than that of the atmosphere out- 
 side. It is again hung from the scale pan, and the stopcock 
 op ned and closed rapidly to bring the gas inside to the same 
 pressure as the air outside. The temperatut e of the balance 
 case in which the globes hang, and the barometric pressure of 
 the air, are the a noted, and the weighing completed. If, 
 then, the capacity of the globe, which determines the volume 
 of the gas, is known, and this volume is reduced by calculation 
 to that which the gas would occupy at the normal temperature 
 and pressure (N.T.P.), the weight of unit volume of the gas 
 under normal conditions is easily calculated. (See Appendix.) 
 Regnault found that : 
 
 1 c.c. of hydrogen weighs 0-0000896 gram 
 
 oxygen 0-0014304 
 
 air 0-0012936 
 
 In the determination of the density of the vapours of solids 
 and liquids which can be vaporized by heat at workable tem- 
 peratures, it is not necessary to actually weigh the vapours, 
 for the solid or liquid may be weighed and then converted 
 into vapour, and its volume, temperature, and pressure 
 observed. The volume of the known weighs of vapour at 
 N.T.P. is then easily calculated, and the density of the vapour 
 obtained. Sometimes it is preferred to weigh the liquid or 
 solid after it has been vaporized and condensed again. The 
 accuracy of these determinations depends upon the fact that no 
 change in the mass of the body takes place during the conversion 
 into vapour. 
 
 Conservation Of Matter. All practical experience leads 
 to the belief that, although bodies in general may be made 
 to undergo a variety of changes, the matter in them cannot 
 be destroyed. It preserves its quantity, and passes without 
 loss through every known process to which bodies can be 
 subjected. Nor can matter be created. The changes that 
 take place in bodies which are passing through a given process 
 
 32 
 
36 METALLURGICAL CHEMISTRY 
 
 may be very complicated, but there is no evidence of gain 
 in the quantity of matter they contain. These statements, 
 that matter can be neither created nor destroyed, refer 
 essentially to its quantity, and are supported by every ex- 
 periment in which the weights of the changing bodies are 
 considered. Any apparent increase in the weight of a body 
 which has undergone change can be traced to the transference 
 of matter to it from some other body that has taken part 
 in the change, and has suffered an equal decrease. This is 
 exemplified in the oxidation of metals in air. The general 
 principle is usually stated under the title of the conservation 
 of matter. 
 
 Divisibility of Matter. It may be taken as an admitted 
 fact that any mineral body can be divided into parts, and 
 that these parts, under exceptional circumstances, may be 
 exceedingly small. All ordinary experience is in favour of 
 this statement, for even the diamond, which is the hardest 
 known body, can be ground to a fine powder on the lapidary's 
 wheel ; but it is diamond cut diamond, for diamond dust 
 must be used in the process. But however finely divided 
 a body may be, its ultimate particles must have a definite 
 size, although they may be excessively small. Each particle, 
 however small, must have a definite volume ; that is, the 
 mass of matter in it must occupy a finite space to the 
 exclusion of all other matter frcm that space. This is some- 
 times expressed by saying that matter possesses volume. 
 Another general experience of all bodies that can be handled 
 is that they possess weight, and that this property is in 
 proportion to the amount of matter in them. If, then, these 
 general properties of matter are admitted, irrespective of the 
 size of the body, the smallest possible particles of a body must 
 occupy a definite volume, and contain a definite quantity of 
 the matter of the body. 
 
 Subdivision of Gold. This metal has many useful 
 properties, and one of the most impoitant is its great malle- 
 
PKOPEKTIES OF MATTEK 37 
 
 ability. It may be beaten out so thin that a sheet having an 
 area of 50 square inches weighs only 1 grain. A square inch 
 of this gold leaf is readily cemented to a sheet of glass, cut 
 into 100 strips by a sharp tool, and each of these divided into 
 100 parts by cutting across the square again at right angles 
 to the first direction. That is, -Jg- grain of gold may be 
 cut into 100x100 = 10,000 small pieces, or 1 grain into 
 500,000 small pieces ; and these can be seen with a good 
 pocket lens. But this is not by any means the limit to the 
 subdivision of the gold ; for with a special machine it is 
 possible to rule 10,000 lines across the square inch of leaf, 
 and, by ruling again at right angles to the first direction, to 
 cut the whole square into 10,000 x 10,000 = 100,000,000 small 
 pieces. Thus a single grain of gold can be mechanically 
 divided up into 100,000,000x50 = 5,000,000,000 pieces, and 
 these pieces can be- seen under a high-power microscope. 
 It may, then, be said that is possible to get a piece of gold 
 which is only ^oooWorr g rain in weight, and is still 
 visible to the aided eye. This division, extreme though 
 it may appear, is not the limit of possibility, for still finer 
 work can be done. 
 
 Dissolution of Common Salt. Compare the above 
 operation, which depends upon great mechanical skill, with 
 the simple process of dissolving common salt in water. The 
 solid disappears rapidly, and the most powerful microscope 
 fails to detect the particles into which it is divided. That it 
 is simply divided is readily proved. The liquid tastes of the 
 salt, and if the water is evaporated the particles run together 
 again, and the &olid salt is recovered. Very little is known 
 about the actual size of the salt particles in the solution, but 
 it is clear that they are smaller than a particle of gold weigh- 
 ing sinrWoinMny grain. Also, there is evidence in favour of 
 the statement that when a small quantity only of the salt 
 is present in the solution, its separate particles are as 
 small as they can possibly be. As will be explained later, 
 
38 METALLURGICAL CHEMISTRY 
 
 the term molecule, which means a small mass, is used to 
 designate the smallest possible particle of the salt which can 
 have a free existence. The molecules of salt are exessively 
 small, but none the less definite, and go to make up the mass 
 of a lump of salt of sensible size. 
 
 Gases. These bodies are for the most part invisible, but 
 it is certain that the body of any gas is made up of an enormous 
 number of very small particles, or molecules, which are in 
 constant motion. The readiness with which the molecules 
 of gases move about enables them to mix together. The 
 odour of musk, which is due to the escape of particles of the 
 body in the gaseous state, will pervade a whole room, but the 
 finest weighing would probably not detect any difference in 
 the weight of the substance from which the particles had 
 escaped to mingle with the air of the room. 
 
 Definite Bodies. The matter of all bodies is made up of 
 molecules. When the molecules of a body are all alike, it is 
 said to have a definite composition. Such bodies have well- 
 defined properties by which they can be recognised. They 
 may be formed in different ways, but they are always the 
 same in composition. It must not, however, be assumed that 
 every well-known body is of this description ; for in many 
 cases further examination brings out the fact that different 
 kinds of molecules are mixed together in the same body. 
 Such bodies are complex in character, and it is part of the 
 chemist's work to separate the molecules of the same kind, 
 and recognise the definite bodies to which they belong. Com- 
 mercial specimens of definite bodies are seldom perfectly pure, 
 but the adulterating substances are not usually present in 
 sufficient quantity to interfere with their ordinary properties, 
 or with experiments in which they are used. 
 
 Red Oxide Of Mercury. This well-known substance has 
 a definite composition. In its common form it is a heavy, 
 red, crystalline powder, which on close examination is found 
 
PEOPEKTIES OF MATTEK 
 
 39 
 
 to be quite homogeneous. Its smallest visible particles are 
 all alike. 
 
 EXP. 22. Weigh 2 grams of red oxide of mercury and introduce 
 it, by means of a gutter of glazed paper, into a piece of combustion- 
 tube closed at one end and about 5 inches long. Fit the open end 
 of the tube with a rubber bung and delivery-tube, and fix the tube 
 in a clip for heating in the Bunsen flame. Pass the end of the 
 deli very -tube under water in a shallow trough, and place over it 
 an inverted gas cylinder of about 100 c.c. capacity, previously filled 
 with water. Heat 
 the part of the 
 combustion - tube 
 which contains the 
 red oxide in the 
 Bunsen flame 
 until the whole of 
 the powder dis- 
 appears, or until a 
 very small residue 
 of impurity only 
 is left. The ar- 
 rangement of the FIG. 14. 
 apparatus is 
 
 shown in Fig. 14. When bubbles of gas cease to escape from the 
 end of the delivery-tube, remove it from the water by taking the 
 clip and tubes away bodily. When the combustion-tube is cold, 
 take out the bung and remove the whole of the mercury which has 
 collected on the sides near the open end. If necessary, use a little 
 water to bring away the last particles of metal. A glass rod tipped 
 with a short piece of rubber tube is useful. Dry the mercury by 
 rolling it about in a small tray, made by turning up the edges of a 
 piece of blotting-paper. Weigh the globule of metal. 
 
 The difference between the weight of the red oxide and 
 that of the metal obtained from it gives the weight of the gas 
 driven off during the heating. On testing the gas in the 
 cylinder with a glowing splint it is found to be oxygen. 
 
 EXAMPLE. Weight of oxide 
 
 ,, mercury 
 oxygen . 
 
 2-000 grams 
 1-846 
 0-147 gram 
 
 If a number of such experiments were made, using different 
 weights of the same red oxide, it would be proved on com- 
 parison that the relation between the weights of mercury and 
 
40 METALLURGICAL CHEMISTRY 
 
 oxygen obtained from the oxide is always the same. This is 
 one way of deciding whether a given body has a definite com- 
 position or not. It may then be taken for granted that red 
 oxide of mercury is a definite body, and that a weighed 
 quantity of it is made up of an enormous number of exactly 
 similar molecules. The observed effect of heat upon it is very 
 clear. It is split up into two other bodies, which are also as 
 definite in character as the red oxide itself, and each of these 
 bodies is made up of its own particular kind of molecules. 
 Where do these new molecules come from ? They must be 
 produced in some way from the disappearing molecules of the 
 oxide. Careful attention to the experiment will make it 
 evident that the action is a gradual one, and that the mercury 
 and oxygen appear in proportion as the oxide disappears. 
 Now, it should not be difficult to picture the individual mole- 
 cules of the oxide as being acted upon by the heat, and 
 divided up into particles of mercury and oxygen, which then 
 form the molecules of these new bodies. Further, these 
 bodies have both been made the subject of the most exhaus- 
 tive examination, but no one has yet been able to divide up 
 their molecules so as to obtain molecules of other new bodies 
 from them. Therefore, in the light of present knowledge, 
 the matter of either mercury or oxygen is as simple as it can 
 possibly be. 
 
 Clearly, then, definite bodies, of which red oxide of mercury, 
 mercury, and oxygen are examples, belong to two distinct 
 types : (1) Those from which only one kind of matter can 
 be obtained by any known process are called Elements ; 
 (2) those from which two or more distinct kinds of matter can 
 be obtained are called Compounds. 
 
 The number of elements is very small compared with the 
 number of compounds formed from them ; but the beginner 
 will easily understand the possibility of a very large number 
 of compounds being formed from a small number of elements, 
 if he compares it with that of the formation of the words of a 
 language from the letters of its alphabet. 
 
PROPERTIES OF MATTER 41 
 
 Matter, then, as we know it, is made up of elements and 
 compounds ; and as the result of a stupendous amount of work, 
 done for the most part during the last 150 years, very definite 
 statements can be made. Some of them are so far-reaching 
 that they are dignified by the name of laws. They will be 
 dealt with in due course. The following definitions will be 
 found useful : 
 
 An element is a body from which only one kind of 
 matter can be obtained. A compound contains two 
 or more elements in chemical union. Bodies which do 
 not come within these definitions may be regarded as more or 
 less intimate mixtures of elements or compounds, or of both 
 elements and compounds, as the case may be. They are 
 complex in composition, for their molecules are not all alike. 
 There are numerous examples of such bodies. 
 
 Now, it is quite justifiable to inquire whether the molecules 
 of elements are the smallest possible particles of these bodies. 
 Much information has been accumulated in answer to this 
 question. For example, are the particles of mercury and 
 oxygen which leave the red oxide during its decomposition 
 really molecules, or are they still smaller pieces of matter 
 which run together to form the molecules of these elements in 
 the free state 1 Some very concise statements can be made at 
 this point, and the beginner may accept them on the under- 
 standing that he will gradually become acquainted with the 
 experimental evidence upon which they are based. 
 
 Atoms. The molecule of oxygen is made up of two 
 smaller and exactly similar particles of oxygen called atoms. 
 These atoms are indivisible, for in them the extreme limit of 
 subdivision has been reached. 
 
 The molecule of mercury when in the gaseous state is not 
 divisible, but is the smallest possible particle of the metal. It 
 corresponds to the atom of oxygen in this respect, and is also 
 called the atom of mercury. So that when the smallest free 
 particle of an element is also the smallest possible particle of 
 
42 METALLUKGICAL CHEMISTRY 
 
 that element, it is called a one-atom or monatomic molecule. 
 The oxygen molecule is a more-than-one atom, or polyatomic 
 molecule. These remarks apply to the elements generally. 
 
 It is easy to imagine the oxygen (Exp. 22) escaping from 
 the red oxide as atoms, and these atoms running together in 
 pairs to form the molecules of the free gas ; and the mercury 
 escaping as atoms, which do not run together, and therefore 
 form molecules by themselves when in the form of vapour. 
 Also, in the oxidation experiments described in Chap. II., 
 the oxygen molecules in the air taking part in the change may 
 be pictured as breaking up into atoms, and joining with atoms 
 of the metals to form molecules of the metallic oxides, which 
 are the products of the change. 
 
 It must here be borne in mind that one of the intrinsic 
 properties of matter is its weight, and, however small the 
 particle of matter may be, it must possess this property. 
 Clearly, then, we may speak of the weight of an atom. 
 
 An atom may be defined as the smallest particle of an 
 element which can exist, or which can take part in the forma- 
 tion of a molecule of an element or compound. 
 
 A molecule is the smallest particle of an element or com- 
 pound which can have a free existence. The masses of all 
 bodies are made up of molecules. The molecules of the 
 vapours of some metals, such as mercury, zinc, and cadmium, 
 are monatomic. As soon as the possibility of the existence of 
 atoms became clear in the minds of the originators of the 
 theory, the want of some method of representing them by 
 signs or symbols was felt. At first these symbols were some- 
 what fanciful, but now the system has been reduced to a very 
 simple form, in which the symbol representing the atom of an 
 element consists of one or two letters taken either from its 
 common or from its scientific name. Thus the symbol for 
 oxygen is ; for hydrogen, H ; for copper, Cu (Cuprum) ; for 
 lead, Pb (Plumbum). But if represents the atom of oxygen 
 it also represents its weight ; for the weight of an atom is 
 
PKOPEKTIES OF MATTEB 43 
 
 intimately connected with the atom itself. Very little can be 
 said about the absolute weight of an atom, but it is quite 
 possible to compare the different elements with regard to the 
 relative weights of their atoms. Hydrogen, which is the 
 lightest body known, has been selected for the standard atomic 
 weight, and the other elements are compared with it. An 
 account of the experimental evidence upon which this rests 
 will be found in Chap. VIIL, and tables containing the 
 common elements, with their symbols and atomic weights, are 
 given on pp. 46-7. 
 
 The symbolic representation of atoms only requires a little 
 extension to adapt it to the representation of molecules. If 
 it is assumed that both the number and the nature of the 
 atoms in a given molecule are known, no difficulty presents 
 itself ; but if the nature only of the atoms in the molecule is 
 known, the best plan is to assume that their number is as 
 small as possible consistent with the composition of the mass 
 of the body. 
 
 The molecule of oxygen is said to contain two atoms of the 
 element ; it could, therefore, Ue r presented by OO, or by 
 20, but it is gene ally represented by O 2 , and this combination 
 of symbols is called the formula of the oxygen molecule. The 
 representation of molecules by definite formulae furnishes a 
 good example of the application of theory to the explanation 
 of the results of experiment, and a short descript'on of the 
 nature cf the experimental evidence upon which themole2ular 
 formula of oxygen is based will serve a useful purpose in making 
 it clear tha L formulae generally depend upon such evidence. 
 
 A general statement which is strongly supported by experi- 
 mental evidence can be made about gaseous bodies. It is 
 known as the Law of Avogadro, and is expressed thus : 
 
 Equal volumes of all gases under the same condi- 
 tions of temperature and pressure contain the same 
 number of molecules. 
 
 Thus equal volumes of oxygen and hydrogen contain the 
 
44 METALLURGICAL CHEMISTRY 
 
 same number of molecules, and double the volume of either 
 gas would contain double the number of molecules. When 
 two volumes of hydrogen are mixed with one volume of 
 oxygen at a temperature above the boiling-point of water, 
 simple mixture takes place, and three volumes of gas are 
 obtained. If the mixture is ignited, the gases combine and 
 s earn is formed, whijh on cooling to the original tempe-ature 
 of the mixture occupies exactly two volumes. Thus the 
 volume of the steam is double that of the oxygen which helps 
 to form it, so that the matter of which the oxygen molecules 
 are composed is distributed into twice as many molecules of 
 steam. That is, each oxygen molecule is divided into two 
 parts during the chemical change by which the hydrogen and 
 oxygen molecules are converted into molecules of water. If, 
 then, the law of Avogadro is true, oxygen molecules must 
 contain at least two atoms of the element, and as no change 
 is known in which they appear to be divided into more than 
 two pirts, the inference that oxygen molecules contain two 
 atoms is justified. 
 
 The weight of a given volume of any gas compared with 
 the weight of an equal volume of hydrogen, which is taken as 
 the unit weight, is called its density. If, then, it is assumed 
 that a molecule of hydrogen is made up of two atoms, each 
 of unit weight, the molecular weight of the gas is double its 
 atomic weight. By the application of Avogadro's law it 
 follows that the molecular weights of gases are proportional 
 to their densities. So that if the molecular weight of hydro- 
 gen is double its density, the molecular weights of other gases 
 are obtained by doubling their densities compared with that 
 of hydrogen. Thus if an element is a gas, or can be con- 
 verted into gas at a workable temperature, and its atomic 
 weight is known, the number of atoms in its molecule is 
 determined by comparing the weight of a given volume of the 
 gas with that of an equal volume of hydrogen under the same 
 conditions of temperature and pressure. This method of 
 determining molecular weights can be applied to all elements 
 
PROPERTIES OF MATTER 45 
 
 and compounds obtainable in the gaseous state. The assump- 
 tion that the molecule of hydrogen contains two atoms is 
 strongly supported by evidence similar to that given above in 
 the case of the oxygen molecule. 
 
 Symbols and formulae are used primarily to represent 
 atoms and molecules, but there is another way of using them 
 which is of great practical importance. It depends upon 
 reasoning such as that which follows : 
 
 A given volume of oxygen is made up of an enormous but 
 definite number of molecules, and if n represent this number, 
 the weight of the given volume is represented by 7i0 2 . The 
 atomic weight of oxygen is 0=16, and its molecular weight 
 is O 2 = 32. Now, it is easy to see that n may be such a number 
 that nQ% represents a definite weight of the gas, say 32 grams ; 
 so that n0. 2 might be used to represent a molecular weight of 
 oxygen in grams, or any other standard weights. This is 
 done, but it is usual to omit the n. Thus the formula 2 is 
 used in two senses: (1) to describe a single molecule of 
 oxygen ; (2) to denote a definite weight of the free gas, as 
 32 grams 32 pounds, etc. The number 32 always appears in 
 any statement which concerns the weight of a molecule of 
 oxygen, or any number of the same. The symbols and 
 formulae of bodies in general are used in the same way. .' /.-/ 
 
 Elements. These bodies have characteristic physical and 
 chemical properties, by which they can be recognised and 
 distinguished from each other, either by general inspection or 
 by chemical tests. They are divided into two main groups. 
 
 Metals. The larger group contains the metals, several of 
 which are so common, and have such well-defined properties, 
 that everyone is more or less acquainted with them. The 
 metals have two characteristic properties which serve to 
 distinguish them from other elements : (1) The smooth clean 
 surface of a metal has the peculiar appearance known as 
 metallic lustre ; (2) the metals are good conductors of heat 
 and electricity. Most of the common metals are very heavy 
 
46 
 
 METALLURGICAL CHEMISTRY 
 
 bodies, but they vary very much among themselves with 
 regard to this property. Platinum, the heaviest of the com- 
 mon metals, is twenty-five times heavier than potassium, one 
 of the lightest of them. 
 
 Non-metals. The smaller group includes a number of ele- 
 ments whose properties are so opposite to those of the metals 
 that they are known as non-metals. They have no common 
 physical characteristics, but their chemical properties are so 
 marked, that they are taken advantage of in classifying them. 
 This group includes such physically different bodies as sulphur, 
 carbon, and oxygen. A few of the elements occupy the border- 
 line between the well-defined metals and non-metals, and seem 
 to partake of the characters of both. The term metalloids is 
 applied to these, but they need not be considered at present. 
 
 Some elements occur in such enormous quantities as to be 
 excessively common ; others in much smaller quantities, but 
 so widely distributed as to be still common ; others again in 
 small quantities, and of limited distribution. These are called 
 rare elements, and, although interesting from the chemist's 
 point of view, are of little known practical value. A list of 
 them is given in the Appendix. 
 
 The following tables contain the names, symbols, and atomic 
 weights of the elements which are usually found, either in the 
 free state or as compounds, on the laboratory shelves. Full 
 lists will be found in the Appendix. 
 
 COMMON METALS. 
 
 
 2 
 
 ii 
 
 i 
 ? 
 
 11 
 
 
 1 
 
 J"* JS 
 
 Name. 
 
 fO 
 
 3 
 
 
 Name. 
 
 ^- 
 
 
 
 |S 
 
 Name. 
 
 #3 
 
 a 
 
 i 
 
 
 g 
 
 
 >. 
 
 
 
 > 
 
 
 
 e 
 
 p 
 
 
 OD 
 
 ^ 
 
 
 OQ 
 
 4* 
 
 Aluminium 
 Antimony 
 
 Al 
 Sb 
 
 27-0 
 120-0 
 
 Copper Cu 
 Gold Au 
 
 63-0 
 197-0 
 
 Nickel Ni 
 Potassium K 
 
 58-5 
 39-0 
 
 Barium Ba ! 137'0 
 
 Iron 
 
 Fe 
 
 56-0 
 
 Platinum Pt 
 
 194-0 
 
 Bismuth Bi 208 -0 
 
 Lead i Pb 
 
 207-0 
 
 Silver 1 Ag 
 
 108-0 
 
 Cadmium 
 
 Cd 
 
 112-0 
 
 Magnesium Mg 
 
 24-0 
 
 Sodium Na 
 
 23-0 
 
 Calcium 
 
 Ca 
 
 40-0 
 
 Manganese ; Mn 
 
 55-0 
 
 Strontium 
 
 Sr 
 
 87-0 
 
 Chromium 
 Cobalt 
 
 Cr 
 Co 
 
 52-5 
 
 58-5 
 
 Mercury 
 Molybdenum 
 
 Hg 
 Mo 
 
 200-0 
 96-0 
 
 Tin 
 Zinc 
 
 Sn 
 Zn 
 
 118-0 
 65-0 
 
PEOPEETIES OF MATTEE 
 
 47 
 
 COMMON NON-METALS. 
 
 Name. 
 
 i 
 
 14 
 
 g_bc ; 
 
 Name. 
 
 2 
 
 II 
 
 Name. 
 
 | 
 
 .2 *" 
 11 
 
 
 ^> 
 
 -** ^ 
 
 t 
 
 ijj jP 
 
 
 > 
 
 
 
 x 
 
 
 
 OQ 
 
 
 
 DO 
 
 ^^ 
 
 Arsenic 
 
 As 
 
 75-0 
 
 Fluorine 
 
 F 
 
 19-0 
 
 Oxygen 
 
 
 
 16-0 
 
 Boron 
 
 B 
 
 11-0 
 
 Hydrogen 
 
 H 
 
 1-0 
 
 Phosphorus 
 
 p 
 
 31-0 
 
 Bromine 
 
 Br 
 
 80-0 
 
 Iodine 
 
 i 
 
 127-0 
 
 Sulphur 
 
 s 
 
 32-0 
 
 Carbon 
 
 C 
 
 12-0 Nitrogen 
 
 N 
 
 14-0 
 
 Silicon 
 
 Si 
 
 28-0 
 
 Chlorine 
 
 Cl 
 
 35-5 
 
 
 
 
 
 
 
 Mercury is the only common metal which is liquid at atmo- 
 spheric temperatures. Arsenic is sometimes classed with the 
 metals, but its chemical properties are more like those of the 
 non-metals. The atomic weights given are only approxi- 
 mate, but they will be found sufficiently accurate for ordinary 
 work. 
 
 Compounds. These bodies are just as definite in their 
 chemical characters and in their constancy of composition as 
 the elements themselves. This cannot be too strongly urged. 
 In fact, the generalizations arrived at as the result of much 
 labour by many men are so all-embracing that they are put 
 forward with much confidence, and are looked upon as thor- 
 oughly reliable guides. They are known as the laws of 
 chemical combination. The first one may be stated here, the 
 others will be given later. 
 
 The first law of chemical combination is as follows : 
 
 A given compound always contains the same ele- 
 ments, and in the same proportions. 
 
 Verifications of this law will be constantly presenting 
 themselves, and the student should find no difficulty in 
 finally accepting it. Two examples may now be given. Red 
 oxide of mercury always contains mercury and oxygen in the 
 proportion of 12'5 parts of mercury to 1 part of oxygen by 
 weight. Water always contains oxygen and hydrogen in 
 the proportion of 1 part of oxygen to 2 parts of hydrogen by 
 
48 METALLURGICAL CHEMISTRY 
 
 volume (see Exp. 19). Or, since oxygen is sixteen times 
 heavier than hydrogen, water contains 8 parts of the former 
 to 1 part of the latter by weight. This will be confirmed 
 later. 
 
 Formulce of Compounds. Now, in Exp. 19 it is proved that 
 2 grams of red oxide of mercury contains 1-846 grams of 
 mercury and 0*147 gram of oxygen. Assuming that the weight 
 of each element is made up of a definite number of atoms of 
 the element, the ratio between these numbers can be found by 
 dividing each weight by the atomic weight of the element to 
 which it belongs. Thus : 
 
 = 0.0092 = 0-0092. 
 
 The ratio is 1 : 1 that is, the two weights contain the 
 same number of atoms. So that they must leave the oxide 
 molecules in pairs, and are, therefore, present in the mole- 
 cules in pairs. Clearly, then, the simplest way to represent 
 a molecule of the oxide is to put the symbols together thus 
 HgO. 
 
 In the decomposition of water by electrolysis it is proved 
 that the constituent gases are liberated in the proportion of 
 2 volumes of hydrogen to 1 volume of oxygen. Also, 
 by Avogadro's law, the double volume of hydrogen contains 
 twice as many molecules as the single volume of oxygen ; and, 
 since all these molecules are diatomic, it is probable that the 
 hydrogen and oxygen atoms leave the water molecules in the 
 proportion of 2 : 1. Therefore the simplest way of expressing 
 the water molecule is to write the symbols together thus 
 H 2 0. And there is much evidence in favour of this as the 
 actual composition of a molecule of water. These complex 
 symbols are called formulce. Thus, HgO and H 2 O are the 
 formulae for red oxide of mercury and water. Next consider 
 Exp. 2. It is not a case of the indifferent absorption of 
 oxygen into lead as a mass, but the individual atoms of 
 oxygen become associated with the individual atoms of lead, 
 
PEOPEETIES OF MATTER 49 
 
 and form molecules of oxide of lead. At any intermediate 
 stage the mass consists of two parts : (a) oxide of lead ; 
 (b) unchanged metal. It is either one or the other ; there is 
 no intermediate change. In this transformation of lead it is 
 found that 207 parts by weight of lead and 16 parts by weight 
 of oxygen combine to form 223 parts of the oxide. But as 
 these are the atomic weights of the two elements, it appears 
 that the atoms of lead and oxygen run together in pairs to 
 form the molecules of the oxide. Therefore, such a statement 
 as the following may be made : 
 
 n atoms of lead (nPlo) and n atoms of oxygen (nO) furnish 
 n molecules of oxide (?iPbO). Or, replacing and by + , and 
 furnish by =, and taking out the n, which is understood, it 
 may be written : 
 
 = PbO. 
 
 This statement is, as it stands, a chemical equation, and 
 expresses the chemical change under discussion. But, as the 
 oxygen of the air is in the form of molecules, it is usual to 
 write the equation thus : 
 
 Here 2PbO means 2[PbO], that is two molecules of the oxide, 
 and the same mode of expressing a number of molecules is 
 always used. 
 
 The decomposition of red oxide of mercury is expressed as 
 follows : 
 
 And the equation reads thus : 2 molecules of red oxide of 
 mercury furnish 2 molecules of mercury and 1 molecule of 
 oxygen. The beginner will probably have a little difficulty 
 in fully understanding the meaning of symbols, formula?, and 
 equations, but as he progresses with the practical work, and 
 becomes familiar with them by use, the difficulty will vanish. 
 
 Impurities. Ordinary specimens of elements and com- 
 
 4 
 
50 METALLURGICAL CHEMISTRY 
 
 pounds usually contain small quantities of other bodies which 
 are regarded as impurities, but these do not as a rule seriously 
 affect the characteristic properties of the bodies under con- 
 sideration. Almost pure specimens can, however, be obtained 
 if necessary. For example, gold containing less than 1 part 
 of impurity in 10,000 parts of the metal is to be obtained, but 
 it costs about three times as much as ordinary fine gold, 
 which contains about 3 parts of impurity in 1,000 parts of the 
 metal. Zinc and copper containing less than 1 part of im- 
 purity in 2,000 parts of the metals are easily obtained, and 
 many compounds are produced in a state of great purity. 
 
 Complex Bodies. These bodies are very numerous and 
 varied. Many of them are common, and have some well- 
 defined properties ; but they lack that absolute constancy of 
 composition which is characteristic of elements and compounds, 
 The atmosphere may be taken as an example. Pure mountain 
 air contains at least four substances of definite composition : 
 two elements, oxygen and nitrogen, and two compounds, 
 water vapour and carbon dioxide. The proportion of oxygen 
 to nitrogen is very nearly constant, but the evidence is all in 
 favour of a simple mixture of these gases in air. The pro- 
 portion of carbon dioxide varies slightly, and that of water 
 vapour considerably, from time to time. Granite is found to 
 contain at least three compounds, which on examination of 
 the polished surface of the stone are seen to be simply mixed 
 together. 
 
 Mixtures are often made for practical purposes, and gun- 
 powder, which contains saltpetre, sulphur, and charcoal, is a 
 good example. Such bodies are called mechanical mix- 
 tures, and when examined under the microscope the various 
 constituents are readily recognised. 
 
 The constituents of mixtures can usually be separated from 
 each other without resorting to chemical changes. Such 
 separation can be effected by mechanical means. The 
 removal of particles of iron from brass-dust, and of oxide of 
 
PROPERTIES OF MATTER 51 
 
 iron from magnetic sands by magnets, are examples. If 
 gunpowder is shaken up with hot water, the saltpetre in it 
 dissolves, and may be completely separated by filtration and 
 washing; the contained sulphur is just as easily dissolved 
 from the dried residue by digesting it with carbon bisulphide; 
 and the charcoal is left as a black residue. On evaporating 
 the solutions of the saltpetre and sulphur respectively, the 
 solids are obtained in the same condition as they were when 
 used to make the original mixture. Complete separation is 
 thus effected without any permanent change in the con 
 stituents of the mixture. 
 
 Theory of Atoms and Molecules. In the above state- 
 ments of facts about matter in general, and in the attempt to 
 explain them, the broad aspects of the atomic theory have 
 been used. The original idea of atoms is by no means a new 
 one, but it was introduced in its modern dress by Dalton, of 
 Manchester, in 1808, who used it most successfully in explain- 
 ing some facts of chemical combination which he had dis- 
 covered. Since his time other considerations have necessitated 
 an extension of the original ideas, so as to distinguish between 
 atoms and molecules. It is in this extended sense that 
 Dalton's atomic theory is used by the modern chemist. 
 
 When Dalton adopted the idea that matter is made up of 
 small particles (atoms), he assumed that they are all alike in 
 a given body of constant composition. If this is not so, then 
 it should be possible to produce two different specimens of 
 the same compound from its elements by a given chemical 
 change. For if the smallest particles of the elements differ in 
 size in the quantities used, the particles of the compound 
 should differ also. But no such variations in the properties 
 of different specimens of the same compound are known ; it is, 
 therefore, a fair inference that the atoms of an element are the 
 exact counterparts of each other. It also follows that if the 
 atoms of elements are alike, the molecules of compounds 
 formed from them must be alike. These assumptions enable 
 
 42 
 
52 METALLUKGICAL CHEMISTEY 
 
 us to account in a simple way for the constancy in the com- 
 position of elements and compounds. 
 
 It has already been noticed that there are two distinct 
 compounds of copper aud oxygen, and it is well known that 
 the weights of copper and oxygen in these compounds are in 
 the proportions of 63 : 8 and 63 : 8x2, and in no others. 
 It was evidence of this kind, obtained from experiments with 
 some of the compounds of carbon, that led Dalton to the 
 discovery of the law of chemical combination in multiple propor- 
 tions, which may be stated thus : 
 
 When one element unites with another element in 
 more than one proportion, the higher proportions 
 are simple multiples of the first proportion. 
 
 This law holds rigorously when two elements combine in 
 different proportions to form two or more distinct compounds; 
 and it seems very probable that the molecules of the com- 
 pounds containing the higher proportions are formed from the 
 molecules of those containing the lower proportion, by the 
 addition of like particles of the other element to them. This 
 was the view that Dalton took, and it enabled him to use the 
 atomic hypothesis in explaining the facts he had discovered. 
 
 It is necessary to distinguish clearly between a series of 
 facts and the efforts of the mind to explain them. A theory 
 is a mind picture, which keeps our work out of the mechanical 
 rut, and raises it from the region of mere groping to that of 
 intelligent effort. A good theory should not only be able to 
 explain what has happened, but also predict what may happen 
 under certain conditions. And the theory of atoms and 
 molecules has done so much in this direction, that it is 
 accepted with the utmost confidence, and is now regarded 
 almost as a proved truth. The student will, therefore, find 
 good mental training in looking for evidence in its support, 
 and in using it to explain observed facts. 
 
PROPERTIES OF MATTER 53 
 
 SUMMARY. 
 
 Matter is the substance of a body. The term " body " is used 
 to denote a limited quantity of matter which has a definite 
 volume, however great or small it may be. On the other hand, 
 substance is the general term for the matter of a body, with- 
 out reference to any particular quantity of it. Bodies can be 
 divided into a number of small parts, but there is a limit to 
 their division. Molecules are the smallest possible free par- 
 ticles of matter of any kind. Molecules under certain con- 
 ditions yield still smaller particles, called atoms, but these 
 cannot have a free existence in the popular sense. The 
 definitions of atom and molecule should be carefully noted, 
 and looseness in their use avoided. They are not interchange- 
 able. An atom is indivisible, and if the smallest free particles 
 of an element are indivisible, they are monatomic molecules 
 as well as atoms. The term "atom" is to be used only in con- 
 nection with elements, which are bodies containing but one 
 kind of matter. Atoms are the constituent parts of mole- 
 cules, and the latter are the smallest particles of compounds 
 and of free elements. 
 
 There is reason to believe that some compounds are, when 
 in dilute solution, split up or dissociated into their elements, 
 the atoms of which are more or less free. These are the 
 wandering atoms or ion-i of the electro-metallurgist. Also it 
 is possible to obtain some elements in a simpler state of 
 division than that in which they exist in the ordinary way, by 
 exposing them to a very high temperature. But in both 
 cases the atoms come together again when the separating 
 cause is removed. Atoms are represented by symbols, mole- 
 cules by formula). To indicate the number of similar molecules, 
 a figure is written in front of the formula of the molecule, as 
 2H 2 , 3PbO. 
 
 A chemical equation is the same as any other equation, 
 with respect to the equality of its two sides. But this equality 
 refers only to the number and quality of the atoms on the 
 
54 METALLURGICAL CHEMISTRY 
 
 two sides, which must be the same, however varied their 
 redistribution in the molecules may have been. Thus in the 
 equation 
 
 = 2PbO, 
 
 there are two atoms of lead and two atoms of oxygen on both 
 sides, but the atoms of the molecules of the elements on the 
 one side have been redistributed into molecules of oxide of 
 lead on the other. 
 
 A theory should always be accepted with an open mind, 
 and judged according to its ability to explain observed facts. 
 If it does this satisfactorily, as is the case with the atomic 
 theory, it may be used freely. 
 
 NOTE. The investigation of radio-active bodies, which received 
 such an impetus by the discovery of radium, seems to point to the 
 slow breaking down of atoms into still smaller particles. This 
 can, however, be only casually referred to in an elementary work 
 until the nature of these particles is better understood, and its 
 influence upon our present idea of atoms- more settled. 
 
 QUESTIONS. 
 
 1. Give a definition of matter. Give examples of the 
 division of bodies into small particles. 
 
 2. What is a refinite body ? Illustrate your answer by 
 reference to water. 
 
 3. Describe an experiment in which decomposition into 
 simpler forms of matter takes place. 
 
 4. Define the terms atom, molecule, element, compound. 
 
 5. How may symbols be used to represent the proportions 
 of the elements in a compound 1 
 
 6. What do you understand by the terms u hypothesis " 
 and " theory " ? 
 
 7. Find the formulae of the oxides of iron and copper, 
 having given that 4 grams of iron combine with 1*714 of 
 oxygen, and 6 grams of copper with 1-52 of oxygen. 
 
CHAPTER V 
 METALS AND SULPHUR 
 
 Sulphur. This well-known element is found associated with 
 earthy matter in volcanic districts. It is present in tbe 
 elementary state, and is called native sulphur. The crude 
 material is carefully heated, which causes the sulphur bo melt 
 and drain away, leaving the greater part of the earthy matter 
 behind. The substance thus obtained is still impure, and is 
 known as crude sulphur. To complete the purification the 
 crude material is put into a retort and strongly heated, when 
 the sulphur is converted into vapour. This passes through 
 an opening in the. retort into a cooling chamber, where it 
 condenses to the liquid state. The hot liquid collects in the 
 bottom of the chamber, from which it is allowed to run into 
 wooden moulds, and cast into cylindrical sticks. This variety 
 is called roll SUlphUF. If the chamber into which the vapour 
 passes is sufficiently cool the condensed material takes the 
 form of a fine powder, which is called flowers of sulphur. 
 These are the common forms of commercial sulphur, and also 
 of the purest. 
 
 A considerable quantity of sulphur is also obtained as a 
 by-product in certain chemical and metallurgical operations, 
 so that we are not entirely dependent upon the native form of 
 the element. 
 
 Action of Heat on Sulphur. Roll sulphur is a pale 
 yellow, brittle, crystalline solid, which, when heated, melts to 
 a clear amber-coloured liquid at 114'5C. On continuing the 
 heating, the liquid gets much darker in colour, and between 
 
56 METALLURGICAL CHEMISTRY 
 
 220 and 250 C. is thick or viscous like treacle ; above the 
 latter temperature it becomes thin again, but is still dark- 
 coloured; at 448 C. it boils, giving off a dark-red vapour. 
 On cooling, the vapour condenses to the liquid again, which 
 then passes through similar changes to the above, but in the 
 inverse order, and so returns to the solid state. If, when in 
 the thick condition, the liquid is poured into cold water, a 
 soft solid mass, called plastic sulphur, is obtained. 
 
 EXP. 23. Break up some roll sulphur into small pieces, and half 
 fill a test-tube with it. Heat the tube in the Bunsen flame, and 
 observe the changes which take place up to the boiling-point. Allow 
 the tube to cool, and when the liquid gets thick pour the greater 
 part of it into cold water. Note the change in colour of the remain- 
 ing liquid, and when it is amber-coloured pour some into water. 
 The residue solidifies on the side of the tube. Examine the water- 
 cooled portions : the one is a stringy, plastic mass ; the other 
 consists of hard, brittle lumps. 
 
 It may be inferred that to get sulphur into the plastic form 
 its temperature must be raised above 220 C. before it is 
 suddenly cooled. Also that some change in the molecules 
 themselves, or in their relation to each other, must take place 
 about that temperature, and that the sudden cooling prevents 
 the reversal of this change. The - plastic mass passes slowly 
 back to the hard yellow form at ordinary temperatures, arid 
 more rapidly at 100 C. 
 
 EXP. 24. Melt about ^ pound of sulphur in a large clay crucible, 
 and allow it to cool slowly until a crust forms on the surface. 
 Break the crust, and pour the remaining liquid into a dish of water. 
 The sides and bottom of the crucible are found to be covered with a 
 mass of long, transparent, needle-shaped crystals. 
 
 EXP. 25. Put about a gram of powdered sulphur into a test-tube, 
 and add to it about 5 c.c. of carbon bisulphide. Shake the tube 
 until most of the sulphur is dissolved ; then pour the clear liquid 
 into a small evaporating basin, nearly cover the basin with a clock 
 glass, and put it in a fume chamber to evaporate. When some well- 
 defined crystals have formed, pour off the remaining liquid, and 
 remove them for inspection. 
 
 On looking at the crystals with a strong magnifying-glass 
 or under a microscope of low power, some well-formed speci- 
 
METALS AND SULPHITE 57 
 
 mens will be observed. They are generally of two shapes, 
 similar to a, b (Fig. 15), but the first shape usually predominates. 
 They both belong to the rhombic 
 system. 
 
 The experiments described above 
 reveal three distinct forms of 
 sulphur. 
 
 1. Octahedral crystals belonging 
 to the rhombic system, which is 
 the stable form found in native 
 sulphur, roll sulphur, etc. 
 
 2. Needle- or prism-shaped crystals belonging to the mono- 
 clinic system, which gradually pass back to the stable form, 
 with evolution of heat. Melting-point 120 C. 
 
 3. The plastic form, which is insoluble in carbon bisulphide, 
 and gradually passes back to the stable form, with evolution 
 of heat. 
 
 These are allo tropic modifications of sulphur. This 
 subject will be referred to again in a later chapter, bub for 
 fuller details of the varieties of sulphur a larger work must be 
 consulted. 
 
 Action of Heat and Air on Sulphur. When any form 
 of sulphur is heated in air it melts, and then takes fire, burn- 
 ing with a pale blue flame. Oxygen from the air unites with 
 the sulphur to form a colourless transparent gas, having a 
 powerful odour, which is usually described as the smell of 
 burning sulphur. It dissolves readily in water, and the solu- 
 tion smells of the gas, which is a product of the oxidation of 
 sulphur, and is, therefore, an oxide. It is proved by experi- 
 ment to contain equal weights of sulphur and oxygen. The 
 atomic weights of its elements are S = 32 and = 16, so 
 that its simplest formula is S0 2 . This is also the true 
 molecular formula of the gas, as indicated by its vapour 
 density. 
 
 Metals and Sulphur. When the common metals are 
 
58 METALLUKGICAL CHEMISTRY 
 
 heated with sulphur a very marked change takes place, which 
 is as well defined as that brought about between the same 
 metals and the oxygen of the air under the influence of heat. 
 The compounds formed are metallic sulphides, and include a 
 number of very important metalliferous bodies. Accurate 
 quantitative results are easily obtained, so that experi- 
 ments on the formation of sulphides give much useful in- 
 formation. 
 
 Copper. EXP. 26. Weigh accurately 2 grams of finely- divided 
 copper, either in the form of filings or of the reduced metal. 
 Transfer it to a clean dry test-tube by means of a paper gutter. 
 Weigh roughly 1 gram of powdered sulphur, and add it to the 
 copper, well mixing the two by carefully shaking the tube. Heat 
 the end of the tube in the Bunsen flame until the action commences, 
 and when it has ceased return the tube to the flame, and heat it 
 strongly until the excess of sulphur is driven well up the tube. 
 Allow the tube to cool in a slanting position, so as to prevent any 
 liquid sulphur from running back on to the solid residue. Invert 
 the cold tube over a piece of glazed paper, and let the bluish black 
 residue fall out. Weigh this residue. Grind it up in a mortar, and 
 look at the powder under a good magnifying-glass. There is no 
 trace of either free copper or sulphur to be seen ; the body appears 
 to be perfectly homogeneous. Mix a little copper and sulphur 
 together ; look at the mixture under the glass ; put it into a test- 
 tube with a little water, and well shake the tube. The copper 
 settles to the bottom, and the sulphur is suspended in the water. 
 No such separation can be brought about with the residue after 
 the mixture has been heated. 
 
 EXAMPLE. Weight of residue 2-507 grams 
 
 ,, copper 2-000 
 
 ,, sulphur ... ... 0*507 gram 
 
 That is, 2 grams of copper combine with 0'507 gram of sulphur to 
 form 2-507 grams of copper sulphide. 
 
 The weight of the sulphur absorbed by the copper is 
 obtained by difference, but this is perfectly legitimate, for 
 practically no copper escapes under the conditions of the 
 experiment ; it is, therefore, quite clear that the increase in 
 weight is due to the absorption of sulphur by the metal. 
 
METALS AND SULPHUR 59 
 
 This definite body with its constant composition is copper 
 sulphide. 
 
 Now, bearing in mind the idea that the actual weight 
 of each element in the sulphide is made up of a large number 
 of atoms, and that the relative weights of these atoms are 
 Cu = 63, and S = 32, it is easy to understand that, by 
 dividing the actual weights by the relative weights, the 
 ratio between the numbers of the atoms of copper and sulphur 
 combining to form the molecules of sulphide is obtained. 
 Thus: 
 
 ^ = 0-032 and ^^ = 0-016. 
 oo oZi 
 
 The ratio is evidently 2:1, and the simplest formula for the 
 compound is Cu 2 S. 
 
 Iron. EXP. 27. Repeat the last experiment, using 2 grams of 
 fine iron filings and 1*5 grams of sulphur. After heating, the 
 residue sticks to the sides somewhat, and it is necessary to break 
 the tube and remove the pieces of glass from the residue before 
 weighing it. 
 
 EXAMPLE. Weight of residue 3-145 grams 
 
 ,, iron 2-000 ,, 
 
 ,, sulphur 1-145 ,, 
 
 That is, 2 grams of iron combine with 1*145 grams of sulphur to 
 form 3'145 grams of iron sulphide. 
 
 A magnet will remove the iron from the mixture of iron and 
 sulphur before it is heated ; but no such separation is possible 
 after the reaction has taken place. The definite compound iron 
 sulphide is formed. 
 
 It should be borne in mind that the metal and sulphur can 
 be mixed together in any proportions, and easily separated 
 again ; but neither variation in proportions nor easy separation 
 is possible when the elements are united in the sulphide. Thus 
 a clear distinction can be drawn between a mixture of two 
 elements, and a compound of the same. Read the statement of 
 the law on p. 47. 
 
 Taking the atomic weights Fe = 56 and S = 32, and using 
 
60 METALLUKGICAL CHEMISTEY 
 
 the principles employed in the last chapter, the formula for 
 iron sulphide is readily found with the data supplied by the 
 last experiment. 
 
 If the ratio is not evident from the numbers obtained by 
 the division with the atomic weights, the simple plan is to 
 divide each of these numbers by the lowest of them. This 
 always gives 1 as the lowest member of the ratio ; comparison 
 is then easy. 
 
 The chemical change may also be expressed by the 
 equation : 
 
 Fe + S = FeS 
 
 56 32 88 
 
 which is proved by the results of Exp. 27. 
 
 For 2 grams of iron require M45 grams of sulphur 
 * 1 ?) j > )> 
 
 .-. 56 = 32-032 grams of 
 
 sulphur. 
 
 A similar calculation made with the results of Exp. 26 proves 
 that 
 
 2Cu + S = Cu. 2 S 
 126 32 158 
 
 These examples furnish further evidence in favour of the broad 
 statement that atomic weights of elements expressed in grams, 
 or any other weight denomination, contain the same number 
 of atoms. They should be referred to when difficulties arise 
 as to the meaning of an equation. 
 
 Lead and Tin. Similar experiments may be made with 
 finely divided lead and tin, produced by vigorously shaking 
 the molten metals in a wooden box just as they are about to 
 solidify. But in this case some of the vapour condenses on 
 the sides of the tube along with the excess of sulphur, so 
 that a good quantitative result cannot be obtained. The 
 
METALS AND SULPHUE 61 
 
 compounds prepared as above are usually described as artificial 
 sulphides, to distinguish them from the naturally occurring 
 ones. 
 
 Natural Sulphides. The majority of the common metals 
 are found in the earth's crust as sulphides mixed with earthy 
 matter, and when the metals they contain can be profitably 
 extracted from them they form the useful ores of these metals. 
 
 Thus, Galena contains lead sulphide, PbS. Copper Pyrites 
 
 contains copper and iron sulphides, Cu 2 S, Fe 2 S 3 , and is used 
 as an ore of copper. Iron Pyrites is almost entirely a 
 sulphide of iron, FeS 2 . These bodies ofcen contain gold and 
 silver. 
 
 ACTION OF HEAT AND AIR ON SULPHIDES. 
 
 Most sulphides melt when strongly heated out of contact 
 with air, and some of them are converted into vapour; but no 
 other change is observed. If air is present, another change 
 takes place where the air comes into contact with the sulphide. 
 It is a surface change, and proceeds most rapidly when the 
 material is finely divided, and the temperature is kept below 
 the point at which the particles soften and clot together. If 
 the action is continued until no further change takes place the 
 sulphur disappears, and its place is largely taken up by 
 oxygen. The sulphide is thus converted into oxide. There 
 are some exceptions to this statement, but it is generally true. 
 It is noticed that as the action proceeds the material gets 
 more and more infusible, till finally it can be raised to a bright 
 red heat without partial fusion or clotting. This is because 
 oxides are generally more infusible than the corresponding 
 sulphides from which they are formed. The above operation 
 is technically known as roasting 1 . This kind of work requires 
 a muffle furnace to be properly carried out, and is fully dealt 
 with in works on metallurgy; but some useful information 
 can be obtained by simple experiments about general and 
 intermediate changes. 
 
62 
 
 METALLURGICAL CHEMISTRY 
 
 EXP. 28. Put a little finely-powdered iron pyrites into a small 
 test-tube, and carefully heat the tube in theBunsen flame. Sulphur 
 vapour is given off, which condenses to liquid on the sides of the 
 tube, and finally solidifies. A residue is left at the bottom of the 
 tube, which resembles the iron sulphide of Exp. 27. 
 
 EXP. 29. Put a little pyrites into the middle of a piece of glass 
 tube open at both ends ; hold the tube in a slanting position over 
 the Bunsen flame, so that the pyrites is heated. A pale blue flame 
 is seen over the heated powder. The position of the tube encourages 
 air to pass in at the bottom, through the tube, and out at the top. 
 The sulphur driven out of the pyrites by the heat burns in the 
 oxygen of the air passing through the tube, and the smell of burning 
 sulphur is readily detected in the gas issuing from the top. 
 
 It should be noted here that iron pyrites, FeS 2 , contains 
 twice as much sulphur as common iron sulphide, FeS, but 
 that half of it is more loosely held, and is driven off when 
 the compound is heated, leaving the simpler compound as a 
 residue. 
 
 FIG. 16. 
 
 EXP. 30. Fit up the apparatus shown in Fig. 16. A is a piece 
 of combustion-tube about 5 inches long, drawn off at one end, and 
 connected with the bottle B, which contains a little water. The 
 aspirator C, which is simply a large bottle with a tap at the bottom, 
 is nearly filled with water, and connected with B by a flexible tube. 
 When the tap at the bottom of C is opened water runs out, and air 
 passes through A and B into C to take its place. The rate at which 
 air flows through A can be regulated by the tap, and a steady 
 current thus kept up. If now some iron pyrites is put into A, 
 
METALS AND SULPHUR 63 
 
 and heated by a Bunsen flame underneath, while a current of air 
 is aspirated over it, a blue flame is seen in the tube, and white fumes 
 appear in B. The action may be continued for ten minutes, then 
 stopped, the bottle B removed, and the fumes blown out. The 
 water smells strongly of sulphur dioxide, which proves that one of 
 the changes resembles the burning of sulphur in air. 
 
 If the roasting is continued long enough the whole of the 
 sulphur is practically removed, and the residue, which is 
 bluish black in colour, is found to be an oxide of iron, having 
 the formula Fe 2 3 . Now part of the sulphur of the pyrites, 
 FeS 2 , is driven off by heat alone, leaving FeS. See Exp. 28. 
 The remainder is removed by the direct action of the oxygen 
 of the air. The various changes are expressed by the equa- 
 tions : 
 
 1. 2FeS 2 = 2FeS + 2S 
 
 Iron pyrites. Iron sulphide. Sulphur. 
 
 2. 2S + 20 2 = 2S0 2 
 
 3. 2FeS + 70= Fe^O. + 2S0 2 . 
 
 Ferric oxide. 
 
 Sulphuric Acid, H 2 S0 4 . By special means it is possible 
 to make" the sulphur dioxide take up a further proportion of 
 oxygen from the air. It is thus converted into another 
 definite compound, sulphur trioxide, S0 3 , a white volatile solid, 
 which unites readily with water. 
 
 EXP. 31. Fit up apparatus similar to that shown in Fig. 16, but 
 insert a loosely-fitting plug of asbestos, the fibres of which have 
 been coated with finely- 
 divided platinum. The 
 position of the plug is 
 shown in Fig. 17, Put 
 some coarsely - powdered 
 iron pyrites, or sulphur, in 
 the open end of A (Fig. 16), 
 heat it with the Bunsen 
 flame, and heat the asbes- 
 tos at the same time with 
 
 a separate burner. Aspi- FIG. 17. 
 
 rate a moderate current 
 
 of air through the apparatus. The sulphur dioxide formed by the 
 roasting of the pyrites, or the burning of the sulphur, is drawn, 
 
64 METALLURGICAL CHEMISTRY 
 
 together with the excess of air, through the hot platinized asbestos. 
 A should not be less than 1 inch internal diameter, and the quantity 
 of sulphide limited. A little added from time to time gives the best 
 results. As the air is only partially exhausted when it reaches the 
 asbestos, the sulphur dioxide takes up more oxygen, and is con- 
 verted into the trioxide. Dense white fumes now appear in B, and 
 readily dissolve into the water present. The last change is expressed 
 thus : 
 
 SO 3 + H 2 O=:H 2 S0 4 . 
 
 Remove the bottle B, move it about so as to make the liquid run 
 over the whole surface, drain it into a porcelain evaporating-dish. 
 and heat it carefully over the Bunsen flame until white fumes 
 appear. When cold the dish is found to contain a small quantity 
 of an oily liquid with very powerful acid properties, and resembling 
 in every way oil of vitriol, or sulphuric acid, H 2 S0 4 . 
 
 The pure acid is a heavy oily liquid of definite composition, 
 which mixes with water in all proportions. The commercial 
 acid usually contains a little water. The manufacture of the 
 commercial acid is based on the changes described above, but 
 oxygen is carried from the air to the sulphur dioxide by means 
 of oxides of nitrogen (Chap. VII.). The plant required for 
 its economical production is very extensive, and a detailed 
 account of the process will be found in any of the larger works 
 on chemistry. 
 
 It has been assumed that in roasting a sulphide the whole 
 of the sulphur is converted into sulphur dioxide, and its place 
 in the sulphide taken up by oxygen. In this case an oxide of 
 the metal is formed ; but if the temperature is kept down, 
 oxygen is actually absorbed into the mass, and a more complex 
 compound, containing oxygen in addition to the metal and 
 sulphur, is formed. 
 
 EXP. 32. Put about 5 grams of powdered copper sulphide, 
 Cu 2 S, into a fire-clay roasting-dish, and place the dish over the 
 Bunseu flame. Stir the powder constantly with an iron scraper, 
 and regulate the flame so that the mass just softens, but does not 
 clot sufficiently to stick to the dish. Continue the roasting for 
 twenty minutes, or longer, if time permits. Transfer the residue 
 to a porcelain mortar, and grind it to a powder again. Boil the 
 powder for a minute or two in 20 c.c. of water contained in a small 
 
METALS AND SULPHUK 65 
 
 beaker. Filter the solution into a clean porcelain dish, and, if it is 
 deep blue in colour, set it aside to crystallize ; if not, drive off some 
 of the water by boiling the solution over a small flame before setting 
 it aside. Boil up a little of the raw sulphide with water in a test- 
 tube, and notice that the solution remains colourless. 
 
 Blue crystals are obtained, which are readily seen to be 
 similar to the crystals of commercial blue vitriol. The 
 changes by which the new compound is formed are somewhat 
 complex, and cannot be considered at this stage ; but the 
 final result is equivalent to an absorption of oxygen by which 
 copper sulphate, CuS0 4 , is produced. 
 
 Percentage Composition. The usual way of expressing 
 the composition of a compound, after it has been experiment- 
 ally obtained, is to put it into the form of percentages. This 
 is known as percentage composition. 
 
 EXAMPLE. To find the percentage composition of copper sulphide 
 from the data given with Exp. 26. 
 
 2*507 grams of copper sulphide contain 2 grams of copper ; 
 
 o 
 .*. 1 gram of copper sulphide contains nT rams ^ c PP er 5 
 
 Therefore the percentage composition is | 
 
 .'.100 grams contain = 79*77 grams of 
 
 2 ' &u ' copper. 
 
 Similarly for sulphur =20*22 grams of sulphur. 
 
 79*77 
 = 20*22 
 99*99 
 
 To find the formula of the compound, the percentage of 
 each element is divided by its atomic weight, and each 
 quotient so obta : ned is divided by the lowest quotient. The 
 second division gives a series of numbers, from which it is 
 easy to derive the formula. Thus 
 
 7 |ir = l ' 25j and ^ ~-~~ ' 631 ; then o^ = 2 very nearly ' and 
 
 0*631 
 
 0*631 
 The required formula is Cu 2 S. 
 
 __ i 
 
66 METALLUKGICAL CHEMISTRY 
 
 The method of finding the formula of a compound depends 
 upon : (1) the invariable proportions of the elements in the 
 compound ; (2) the existence of atoms, which for the same 
 element are of invariable weight. If, then, it is found that 
 certain weights of two or more elements combine to form a 
 compound, and the idea of atoms is applied, the only inference 
 which can be drawn is, that the weight of any one of the 
 elements is made up of an enormous number of similar atoms. 
 And it is evident that if the atomic weight is known, the 
 relation between it and the known weight of the element in 
 the compound is proportional to the number of atoms in that 
 weight. 
 
 79*77 
 Thus the ratio = 1 -25 is proportional to the number of 
 
 DO 
 
 20-9O 
 atoms in 79-77 parts of copper, and the ratio ^ =0'631 is 
 
 proportional to the number of atoms in 20'22 parts of 
 sulphur. 
 
 Results obtained in this way lead to what are called 
 empirical formulae, for they only give the ratio between the 
 numbers of the different atoms in the molecule of the com- 
 pound. Thus Cu 2 S might be written Cu 4 S 2 , Cu 6 S 3 , etc. It 
 is, however, the only method of obtaining formulae which is 
 applicable to all compounds without exception. Other facts 
 have to be taken into consideration in deciding as to the 
 actual number of atoms in the molecule of a compound ; and, 
 as they are not applicable generally, there are many com- 
 pounds the actual molecular formula of which are not 
 known. 
 
 SUMMARY. 
 
 The matter of this chapter should receive careful attention, 
 especially that portion which refers to the determination of 
 formulae and the formation of equations. So much depends 
 upon getting clear ideas of the quantitative relations between 
 
METALS AND SULPHUR 67 
 
 the elements in a compound, that it is hardly possible to 
 devote too much attention to it at first. Some difficulty is 
 always experienced by those who are not acquainted with the 
 use of symbols in general, in their endeavour to grasp the 
 meaning of a formula or of an equation. The facts of a 
 particular reaction and the attempt to explain them are 
 sometimes confused ; so that it is a good plan to get the facts 
 well digested before attempting their explanation. 
 
 The formation of sulphides, and the method of finding their 
 formulae, should be carefully studied. The exact composition 
 of a compound can be found, either by building it up from its 
 elements (synthesis), or by separating it into its elements 
 (analysis). This is the first step, and whichever method is 
 adopted, both the nature of the elements in the compound 
 and their proportions by weight must be determined. Very 
 often it is possible to employ both methods for the same 
 compound, in which case they must confirm each other. Any 
 number of experiments, however diverse they may be, if they 
 are directed to the determination of the composition of a 
 given compound, should lead to the same result. 
 
 The distinction between the facts of a particular reaction 
 and the theoretical considerations used in their explanation 
 is well brought out by studying the equation : 
 
 Fe + S = FeS 
 56 32 88 
 
 The fact is that 56 grams of iron will combine with 32 grams 
 of sulphur to form 88 grains of iron sulphide. The theory 
 states, in scarcely less definite language, that for every atom 
 of iron in 56 grams of the metal there is a companion atom 
 of sulphur in 32 grams of that element ; that these atoms run 
 together in pairs to form molecules of iron sulphide ; and 
 that whatever takes place between a single pair of atoms is 
 simply repeated with every other pair in the reacting 
 mass. 
 
 This is equivalent to saying that whatever may be the 
 
 52 
 
68 METALLURGICAL CHEMISTRY 
 
 number of atoms in 56 grams of iron, there is the same 
 number of atoms in 32 grams of sulphur. 
 
 In practice, then, the symbols represent measurable pro- 
 portional weights of the elements. Theoretically they repre- 
 sent single atoms, and one side of an equation shows a group 
 containing the minimum number of atoms which can take part 
 in the chemical change represented by the equation, the other 
 the result of the change. Then, however great the reacting 
 mass may be, it is divided up into similar groups as the change 
 proceeds, and any excess of one kind of atoms over the others 
 takes no part in it. 
 
 QUESTIONS. 
 
 1. What changes take place when sulphur is heated out of 
 contact with air ? 
 
 2. How may copper be converted into its sulphide ? 
 Describe the experiment. 
 
 3. Write down a chemical equation, and say in your own 
 words what you think it means. 
 
 4. How may sulphuric acid be prepared from sulphur 1 ? 
 State the general properties of the acid. 
 
 5. What takes place when a sulphide is heated in air ? 
 
 6. What is understood by the percentage composition of a 
 compound 1 From the example given with Exp. 27 find the 
 weights of iron and sulphur in 100 grams of the sulphide, and 
 from the percentages so obtained deduce the formula of the 
 compound. 
 
 7. The empirical formula of a compound does not always 
 represent its molecule. Why is this 1 
 
CHAPTER VI 
 COMMON ELEMENTS AND COMPOUNDS 
 
 Common Salt. This well-known substance is found in the 
 earth in various parts of the world. The most extensive 
 deposits in this country are found in Cheshire, Worcester- 
 shire, and Yorkshire. It also forms the bulk of the dissolved 
 matter in sea water. It is a white solid which dissolves 
 readily in water, and crystallizes from its solutions in cubes. 
 The solid, or its solution, has a characteristic saline taste, 
 by which it is readily recognised. As it is easily purified, 
 the commercial salt is sufficiently pure for most purposes. 
 Common salt, or sodium chloride, contains the metal sodium 
 and the gas chlorine in the proportion of 23 parts of the metal 
 to 35-5 parts of the gas by weight. Its formula is, therefore, 
 NaCl. 
 
 Hydrochloric Acid. This very important compound, as 
 prepared for use, is a strongly-fuming liquid with very char- 
 acteristic properties. It is known also as muriatic acid and 
 smoking salts. The pure compound contains 1 part of 
 hydrogen to 35 -5 parts of chlorine by weight, and its formula 
 is HC1. It is prepared from common salt. 
 
 EXP. 33. Dry some common salt by heating it in an iron pan 
 over the Bunsen flame. Make a dilute solution of sulphuric acid 
 by pouring some of the strong acid slowly into its own volume of 
 water. Notice that the solution becomes very hot, and remember 
 that the acid must always be poured into the water, and not the 
 water into the acid. Neglect of this rule may cause an accident. 
 Put a little dry salt into a dry test-tube, add a little of the diluted 
 acid, and warm the mixture over the Bunsen flame. A gas is 
 
70 
 
 METALLURGICAL CHEMISTRY 
 
 given off which fumes strongly in the air, and has a strong acid 
 taste and smell. A strip of blue litmus-paper held in the fumes 
 is turned red. The gas is hydrochloric acid. 
 
 EXP. 34. Fit up the apparatus shown in Fig. 18. A is a 12-ounce 
 flask fitted with a bung and angle-tube. B is a drying bottle con- 
 taining a little strong sulphuric acid, and standing in a beaker of 
 water to keep it cool C is a gas cylinder with a delivery-tube. A 
 and C are connected with B by flexible rubber tubes. D is a long- 
 necked flask with a deli very- tube, and filled up to the mark with 
 water. Put about 10 grams of dry salt into A, and add 50 c.c. of 
 sulphuric acid (1 to 1) ; replace the bung, and heat the flask gently 
 on a piece of gauze over the Bunsen flame. Gas is given off, passes 
 
 FIG. 18. 
 
 through the strong acid in B, which dries it, and collects in C by 
 gradually displacing the air. White fumes will soon appear at the 
 mouth of the jar, but a considerable quantity of the gas must escape 
 before the whole of the air is expelled. When the jar is thought to 
 be full remove the delivery-tube, and put a greased plate over the 
 mouth of the jar. Fill another cylinder with the gas, and then 
 connect B with the deli very- tube in the bottle D, which contains 
 water up to the mark when the delivery-tube is in position. Allow 
 the gas to pass into D until the water shows signs of running back. 
 Disconnect from B, and note that the level of the water in D is con- 
 siderably above the mark, the delivery-tube still being in position; 
 also that the bottle feels warm to the hand. 
 
 The chemical change by which the gas is liberated is ex- 
 pressed by the equation : 
 
 NaCl + H 2 S0 4 = NaHS0 4 + HC1. 
 
COMMON ELEMENTS AND COMPOUNDS 71 
 
 The sodium hydrogen sulphate formed in the change remains 
 in the flask, and is easily obtained as a crystalline solid from 
 the residue. 
 
 Properties of Hydrochloric Acid. The dry gas col- 
 lected in the cylinders is colourless and transparent. On 
 removing the plate from a jar of the gas, fumes appear at the 
 mouth. The escaping gas, on coming into contact with mois- 
 ture in the air, dissolves in it, and causes its condensation as 
 an acid cloud. 
 
 A burning candle held in a deflagrating spoon and carefully 
 lowered into the gas ceases to burn, but just before the flame 
 disappears the blue portion of it turns green. The gas is non- 
 combustible, and does not support combustion. 
 
 A little blue litmus solution, when poured into the jar of 
 gas used in the last experiment, is turned red, showing that 
 the gas has acid properties. 
 
 When a jar of the gas is inverted over water and the plate 
 removed, the water rushes up and fills the jar. The gas is 
 very soluble in water, and as it is dissolved the water rises 
 in the jar to take its place. The dissolution of the gas in the 
 bottle D in Exp. 34 brings out two important facts : (1) that 
 the volume of the solution is greater than the original volume 
 of water ; (2) that heat is developed during the dissolution of 
 the gas. To be quite sure of the latter fact it is necessary to 
 put the drying bottle B into a vessel containing cold water to 
 ensure that cold gas passes into D. Also the weight of the 
 solution is greater than that of water bulk for bulk, and 
 increases with the quantity of gas absorbed, up to the point 
 of saturation. The strong solution at ordinary temperatures 
 has a density of 1*2, and contains 40 per cent, of HC1. To 
 produce this the water dissolves more than 400 times its own 
 volume of the gas. Thus the strong commercial acid contains, 
 roughly, 40 parts by weight of the gas dissolved in 60 parts 
 by weight of water. 
 
 Large quantities of the acid are obtained in the manufac- 
 
72 METALLUEGICAL CHEMISTEY 
 
 ture of salt cake from common salt and sulphuric acid. The 
 operation is conducted in a furnace, and at a higher tempera- 
 ture than that which can be used with a glass flask. In this 
 case double the quantity of salt is decomposed by the same 
 quantity of acid. The reaction takes place in two stages : 
 
 1. NaCl -f H 2 S0 4 = NaHS0 4 + HCl. 
 
 2. NaCl + NaHS0 4 = Na 2 S0 4 + HC1. 
 
 Sodium hydrogen Sodium 
 sulphate. sulphate. 
 
 The gas is absorbed while passing through towers filled 
 with coke over which water is allowed to trickle. The solu- 
 tion is collected in receivers at the bottom of the towers. The 
 crude acid is often impure, and must be purified for special 
 purposes, but the common acid is useful for many operations 
 in which purity is not essential. 
 
 CHLORINE. 
 
 This element is a yellowish -green gas which is present in 
 a number of compounds, of which common salt is the type. 
 It can be obtained either from common salt or from hydro- 
 chloric acid. Black oxide of manganese, which is used in its 
 preparation, is found in the earth as manganese ore or pyro- 
 lusite. Its formula is Mn0 2 . It is a black, easily-powdered 
 solid. 
 
 EXP. 35. Well mix a little dry salt with some powdered black 
 oxide of manganese. Put the mixture into a test-tube, add a little 
 dilute sulphuric acid, and gently warm the tube. A yellowish- 
 green gas is given off, which has a very irritating odour. The effect 
 of the small quantity which gets into the air is quite perceptible. A 
 piece of moistened litmus-paper held in the gas is quite decolourized. 
 This yellowish-green gas which is liberated from the salt by the 
 combined action of the oxide and the acid is chlorine, C1 2 . 
 
 EXP. 36. Use the apparatus shown in Fig. 18. The wash-bottle 
 need not be kept cool. Put about 20 grams of black oxide of man- 
 ganese into A, add 100 c.c. of strong hydrochloric acid, and heat the 
 flask gently on the gauze. Allow the gas to collect in C until the 
 jar seems full of a yellowish -green body. Cover the mouth with a 
 greased plate. Collect several jars of the gas in the same way. 
 
COMMON ELEMENTS AND COMPOUNDS 73 
 
 The experiment must be conducted in a fume-chamber, but 
 if one is not available the gas should be collected over water. 
 In that case a delivery-tube is passed directly from the flask A 
 to a pneumatic trough, and under the mouth of an inverted 
 cylinder filled with water. The gas is somewhat soluble in 
 water, but if a rapid stream of it is kept up no difficulty is 
 experienced in collecting a number of jars. Either the wet or 
 dry gas can be used for illustrating most of its properties. 
 
 The changes taking place in the above experiments are 
 expressed by the equations : 
 
 1. Mn0 2 + 2NaCl + 3H 2 S0 4 = MnS0 4 +2NaHS0 4 +2H 
 
 Manganese Manganous 
 
 dioxide. sulphate. 
 
 Manganous 
 chloride. 
 
 Properties Of Chlorine. It is a yellowish-green gas 
 when seen by daylight, or by the light of burning magnesium, 
 but the colour is too faint to be recognised by gaslight. It 
 has a very irritating odour, and care should be taken not to 
 inhale it in quantity. Fresh air is the best remedy. 
 
 Hydrogen or combustible bodies containing hydrogen burn 
 in this gas, and finely-divided metals take fire when brought 
 into contact with it. 
 
 Experiments with Chlorine. 1. Carefully lower a lighted 
 candle, supported in the bowl of a deflagrating spoon, into 
 a jar of the gas. It continues burning, but with a dull red, 
 smoky flame, and a thin deposit of soot is formed on the 
 inside of the jar. 
 
 2. Moisten a piece of tow with turpentine, hang it on 
 the bowl of a deflagrating spoon, and lower it into a jar 
 of the gas. The turpentine suddenly bursts into flame, 
 and a thick deposit of soot is formed inside the jar. The 
 candle and the turpentine contain carbon and hydrogen, and, 
 during the combustion the hydrogen only burns, the carbon 
 being set free and deposited as soot. When the spoon is 
 lifted out of the jar white fumes appear at the mouth. This 
 
74 METALLURGICAL CHEMISTEY 
 
 fuming is caused by the escape of hydrochloric acid gas. The 
 change is expressed by the equation : 
 
 H 2 + C1 2 = 2HC1. 
 
 3. Grind some antimony to a fine powder in a mortar ; put 
 it into a small bottle, and tie a strip of muslin over the 
 mouth. Sprinkle some of the metallic powder into a jar of 
 chlorine gas. A shower of brilliant sparks is seen. Repeat 
 the last experiment, using finely-divided copper in place of 
 antimony. 
 
 The changes are simple, and compounds resembling common 
 salt in general composition are formed. They are called 
 chlorides, and the formation of copper chloride may be taken 
 as typical : 
 
 Cu + Cl 2 = CuCl 2 . 
 
 Cupri<T 
 chloride. 
 
 The conversion of metals into their chlorides is sometimes 
 of considerable practical importance, and is effected on the 
 large scale both by free chlorine and by common salt. 
 
 Chlorine is somewhat soluble in water, and forms a solution 
 which smells strongly of the gas. When the solution stands 
 for some time slow decomposition of the water takes place, 
 with formation of hydrochloric acid and liberation of oxygen : 
 
 The change is more rapid in daylight, and very rapid when 
 a body which will take up oxygen readily is put into the 
 solution. On this account chlorine is sometimes called an 
 oxidizing agent. Colouring matter of vegetable origin is 
 readily decolourized by the gas in the presence of water. 
 This action is called bleaching. It is probable that the oxygen 
 liberated from the water combines with the colouring matter, 
 and forms colourless compounds with it. Mineral colouring 
 matter is not affected. 
 
 4. Moisten a strip of turkey-red cloth about half-way up 
 
COMMON ELEMENTS AND COMPOUNDS 
 
 75 
 
 and suspend it in a jar of dry chlorine. The colour is rapidly 
 discharged from the wet part, but very slowly from the dry 
 portion. 
 
 Chlorine gas is made on the large scale, and absorbed into 
 slaked lime to form bleaching powder, from which the chlorine 
 is readily liberated for use. Common salt and chlorine are 
 used in the metallurgy of copper, silver, and gold. 
 
 Potassium Chlorate, KC10 3 . This useful compound may 
 be prepared by passing chlorine gas into a strong solution of 
 caustic potash. A thistle funnel, 
 bent as shown in Fig. 19, is fitted 
 13 to a wash-bottle containing a little 
 water, and dips into the solution of 
 Y \- f " potash in water (1 to 2). Chlorine, 
 
 | ~] prepared as described above, is passed 
 
 through the wash - bottle into the 
 solution, by which it is rapidly 
 absorbed. A crystalline solid sepa- 
 rates from the solution. When the action is finished the 
 liquid is allowed to stand for further separation, poured off, 
 the solid dissolved in water, and recrystallized to purify it. 
 
 Action of Heat. When potassium chlorate is heated it 
 splits up into potassium chloride and oxygen, and this reaction 
 is made use of in the preparation of moderate quantities of 
 oxygen for experimental purposes. 
 
 EXP. 37. Put a little powdered chlorate into a dry test-tube, and 
 heat it strongly in the Bunsen flame. It melts and gives off bubbles 
 of gas, which, when tested by the introduction of a glowing splint, 
 is found to be oxygen. 
 
 FIG. 19. 
 
 The complete change is expressed thus : 
 
 2KC10, 
 
 30, 
 
 2KC1 + 
 
 Potassium 
 chloride. 
 
 2(39 + 35-5 + 48) 3x32 
 
 245 
 
 96 
 
76 METALLUKGICAL CHEMISTEY 
 
 By mixing the salt with about a quarter its weight of man- 
 ganese dioxide, and heating the mixture, the gas is liberated 
 much more freely and at a lower temperature. The oxide 
 undergoes very little change, and its action need not be con- 
 sidered now. 
 
 EXP. 38. Thoroughly dry some powdered potassium chlorate in 
 a porcelain basin on the sand-bath. Heat some powdered black 
 oxide of manganese in a porcelain crucible over the Bunsen flame. 
 Weigh accurately 2 grams of the dry chlorate, and put it into a dry 
 test-tube ; add about the same weight of the ignited black oxide, 
 and well mix the two by careful shaking. Now weigh the tube and 
 its contents accurately, and then heat it gently in the Bunsen flame 
 until no more gas is liberated. Two or three minutes' heating is 
 usually sufficient. Allow the tube to cool down, and reweigh it. 
 
 EXAMPLE. 
 
 Weight of tube and mixture before heating=15'641 grams 
 
 after = 14'863 
 
 Loss =00-778 gram 
 
 The loss is due to the escape of oxygen from the 2 grams of 
 chlorate, and may be compared \\ith the loss calculated by 
 using the weights given in the equation on p. 75. 
 
 For 245 grams of KC1O 3 lose 96 grams of 2 
 
 i ' ^ 
 
 245 
 1 ^- = 0-784 gram of 2 . 
 
 If the test-tube is fitted with a rubber bung and delivery- 
 tube, and is supported in a horizontal position by a clip, the 
 evolved gas may be readily collected over water. This method 
 is often used to obtain oxygen for experimental purposes, and 
 will be referred to later. 
 
 Bromine and Iodine. These elements resemble chlorine 
 in their chemical properties, and are of considerable use in 
 the laboratory. 
 
COMMON ELEMENTS AND COMPOUNDS 77 
 
 Bromine is a dark red, heavy liquid, 3-18 times heavier 
 than water. It gives off a red vapour at ordinary temperatures, 
 which is even more irritating than chlorine. It is somewhat 
 soluble in water, forming a red solution. It attacks water in 
 the presence of bodies which will take up oxygen, and on that 
 account is useful indirectly as an oxidizing agent. The com- 
 monest compound of bromine is potassium bromide, KBr. 
 
 Iodine is a greyish -black crystalline solid which melts 
 readily, and gives off a violet vapour with an irritating odour. 
 It volatilizes slowly at ordinary temperatures, so that even the 
 solid has an acrid smell. Its most important compound is 
 potassium iodide, KI, which is readily soluble in water. 
 Iodine is only slightly soluble in water, but is readily soluble 
 in water containing potassium iodide, and is often used in this 
 form. It also dissolves readily in alcohol, ether, and carbon 
 bisulphide. 
 
 Compounds of bromine and iodine are present in small 
 quantities in sea-water, and are absorbed into sea-plants during 
 their growth. When dry seaweed is heated to a low red heat 
 in a retort, combustible gas is driven off, and a black residue 
 left, which, on treatment with water, yields a solution con- 
 taining a considerable quantity of dissolved matter. On 
 careful evaporation in pans, potassium sulphate crystallizes 
 out, and is separated from the hot solution ; on cooling the hot 
 liquid a further deposit of potassium chloride is obtained. 
 The remaining solution, called the mother liquor, contains 
 compounds of bromine and iodine, and, after mixing with 
 sulphuric acid and standing for a time, is transferred to an 
 iron still, in which it is heated with black oxide of manganese. 
 Iodine distils over, and is condensed in receivers in the solid 
 state. When no more iodine passes over more black oxide is 
 added, and the heating continued. Now bromine distils over, 
 and is collected in a fresh set of receivers in the liquid state. 
 The bromine so obtained is usually sufficiently pure for ordinary 
 purposes, but the iodine needs purification, 
 
78 METALLUKGICAL CHEMISTRY 
 
 It will be noticed that the same materials are used to 
 liberate bromine and iodine as are used in Exp. 36 for the 
 liberation of the chlorine from common salt, and the chemical 
 changes taking place are also similar. Compare the equation 
 given below with No. 1 on p. 73 : 
 
 Mn0 2 + 2KBr + 3 H 2 S0 4 = MnS0 4 + 2K HS0 4 + 2H 2 + Br 2 . 
 
 In Germany bromine is largely prepared by the action of 
 chlorine on bromides in the mother-liquor after the separation 
 of potassium chloride from the mixed potassium salts which 
 occur in the Stassfurt salt-beds. The reaction is simple : 
 
 Cl 2 +2KBr = 2KCl + Br 2 . 
 
 Potassium Potassium 
 bromide. chloride. 
 
 Saltpetre, also called nitre or potassium nitrate, is obtained 
 in the form of colourless crystals, or as a white powder. It is 
 very soluble in water, and the solution has a saline taste. 
 An efflorescence which appears on the surface of the soil in 
 hot countries, notably in India, is found to consist largely of 
 potassium nitrate. If this surface layer is stripped off and 
 treated with water, the nitrate dissolves, and may be recovered 
 from the solution by crystallization. The formation of the 
 nitrate is brought about by changes in the organic matter of 
 the soil in the presence of decomposing rocks containing 
 potassium compounds. And it is only where these conditions 
 are present that the compound is formed. There is also a 
 method, based upon similar changes, for its artificial pro- 
 duction. The corresponding compound, Chili saltpetre or 
 sodium nitrate, occurs as large deposits in certain parts of 
 South America. This compound is very similar in its pro- 
 perties to the potassium salt, and, as it is cheaper, is used for 
 some purposes in place of the latter. 
 
 The formulse of the compounds are KN0 3 and NaN0 3 
 respectively. 
 
 Nitric Acid, HN0 3 . This compound is a strongly-fuming 
 liquid with very powerful properties. It is readily obtained 
 
COMMON ELEMENTS AND COMPOUNDS 
 
 79 
 
 B 
 
 from either of the nitrates described above by acting upon 
 them with strong sulphuric acid. 
 
 EXP. 39. Fit up the apparatus shown in Fig, 20. A is a glass- 
 stoppered retort of about 8 ounces capacity. B is a small flask 
 resting in a bowl of water, and covered with wet blotting-paper to 
 keep it cool. 
 Put about 20 
 grains of pow- 
 dered saltpetre 
 into the retort, 
 and pour in 
 strong'sulphuric 
 acid through a 
 funnel until the 
 solid is well 
 covered. Ee- 
 place the stop- 
 per, and gently 
 heat the mix- 
 ture. Eeddish 
 vapours appear 
 in the retort; FIG. 20. 
 
 drops of liquid 
 
 collect in the stem and run down into the receiver. Continue the 
 heating for twenty minutes. Then remove the receiver and 
 examine the liquid which has distilled over. 
 
 The acid formed by the reaction in the retort is gaseous at 
 the temperature of the experiment, and, passing in this form 
 into the cooler stem, condenses there, and runs into the receiver. 
 The liquid thus collected is usually coloured, which is due to 
 the presence of gaseous impurities formed by the decomposi- 
 tion of part of the acid by heat. If a stream of air is blown 
 through it for some time, the colour disappears, and a colour- 
 less, strongly-fuming liquid is obtained. 
 
 The change taking place in the retort is expressed by the 
 equation : 
 
 + H 2 S0 4 = KHS0 4 + HNO 8 
 
 Potassium 
 hydrogen sulphate. 
 
 39 + 14 + 48 1 + 14 + 48 
 
 101 
 
 63 
 
80 METALLUEGICAL CHEMISTEY 
 
 The potassium hydrogen sulphate is left as a residue in the 
 retort. The conversion of the nitrate into nitric acid is easily 
 seen to consist of an exchange of hydrogen for potassium. 
 Reference to the table of atomic weights will show that the 
 numbers under the formulae in the equation represent the 
 molecular weights of the compounds, and also the proportion 
 between the weight of nitrate and the weight of acid formed 
 from it. If sodium nitrate is used instead of saltpetre, the 
 reaction is expressed thus : 
 
 NaN0 3 + H 2 S0 4 = NaHS0 4 + HN0 3 
 
 Sodium hydrogen 
 sulphate. 
 
 23 + 14 + 48 1 + 14 + 48 
 
 85 63 
 
 On comparison it is seen that 85 parts of sodium nitrate 
 produce the same weight of nitric acid as 101 parts of 
 potassium nitrate. So that if the acid is the object of the 
 preparation, the sodium compound is the more economical 
 to use. 
 
 In the manufacture of the acid Chili saltpetre is used, 
 and the reaction is conducted in iron retorts lined with fire- 
 clay. A much higher temperature is employed and twice as 
 much acid obtained with a given weight of sulphuric acid. 
 This is explained by the equation : 
 
 2NaN0 8 + H 2 S0 4 = Na 2 S0 4 + 2HN0 3 . 
 
 Sodium 
 sulphate. 
 
 The acid vapour is condensed in a series of earthenware 
 receivers, and afterwards decolourized by blowing a current 
 of air through the liquid. 
 
 Properties Of Nitric Acid. The ordinary strong com- 
 mercial acid has a density of 1 -4, and contains about 65 per 
 cent, of the pure acid, together with 35 per cent, of water. 
 But by carefully redistilling this strong acid, mixed with half 
 its volume of strong sulphuric acid, a strongly-fuming liquid 
 
COMMON ELEMENTS AND COMPOUNDS 81 
 
 is obtained. This is the practically pure acid, containing 
 100 per cent. HN0 3 , and having a density of 1*53. So 
 that pure nitric acid is a colourless liquid, which is readily 
 vapourized and easily decomposed. It is much more stable 
 in the presence of water, with which it mixes in all pro- 
 portions. 
 
 A dilute solution of the acid turns blue litmus red, and has 
 the characteristic sour taste of acids in general. The strong 
 acid has a very caustic action upon the skin, which turns 
 yellow, then black, and finally peels off. It is a powerful 
 oxidizing agent, readily giving up part of its oxygen to bodies 
 which will absorb that element. Sodium and potassium 
 nitrates are similar to the acid in this respect. 
 
 EXP. 40. Make a piece of charcoal red-hot in the Bunsen flame, 
 place it on a tile, and carefully pour a few drops of strong nitric 
 acid upon the incandescent part. The glowing charcoal glows more 
 brightly, showing that increased oxidation is taking place, due to 
 the extra supply of oxygen obtained from the acid. 
 
 EXP. 41. Melt a little saltpetre in a test-tube supported by a 
 clip, and drop in a small piece of charcoal. When the charcoal 
 begins to burn, remove the flame, and the combustion will con- 
 tinue. The charcoal burns brightly in the oxygen derived from the 
 nitrate. 
 
 Repeat the last experiment, but drop into the tube a small 
 piece of sulphur instead of charcoal. The burner may be 
 removed before dropping in the solid, as the combustion 
 commences at once and continues vigorously. An iron tray 
 should be placed under the tube, as the heat developed is 
 sometimes sufficient to fuse the bottom and cause it to drop 
 out. 
 
 Ammonium Chloride. One of the by-products of the 
 manufacture of coal-gas is known as ammoniacal liquor. When 
 this liquid is mixed with lime and heated in a closed vessel, 
 ammonia gas is liberated, which is readily absorbed by strong 
 hydrochloric acid and converted into ammonium chloride. As 
 the acid liquid becomes saturated, the chloride crystallizes out, 
 
 6 
 
82 METALLURGICAL CHEMISTEY 
 
 is collected, purified, and put on the market as sal-ammoniac. 
 The reaction is : 
 
 NH 3 + HC1 = NH 4 C1. 
 
 Ammonia. Ammonium 
 
 chloride. 
 
 The pure compound is a white, crystalline, odourless solid, 
 readily soluble in water. It passes directly from the solid to 
 the gaseous state when heated. In the gaseous state its mole- 
 cules are largely split up into ammonia and hydrochloric acid 
 molecules : 
 
 NH 4 C1 = NH 3 + HC1. 
 
 But these recombine as the temperature falls, and the original 
 solid is reproduced : 
 
 NH 3 + HC1 = NH 4 C1. 
 
 This splitting up of the molecules at or above a definite tem- 
 perature, and their recombination below that temperature, is 
 known as dissociation. The last change is readily shown 
 by filling two gas-jars with the dry gases and bringing them 
 together mouth to mouth. The white chloride forms rapidly 
 and deposits on the sides of the jars. 
 
 Ammonia Gas, NH 3 . This compound has a very pungent 
 odour, and its presence is very readily detected, although it is 
 quite transparent and colourless. 
 
 EXP. 42. Mix together 5 grams of ammonium chloride and 
 4 grams of quicklime ; put the mixture into a test-tube fitted 
 with a bung and a straight delivery-tube about 8 inches long ; 
 fix the test-tube in a clip in the vertical position, and invert a wide 
 test-tube over the delivery-tube so as to rest on its end. Now 
 heat the mixture, and allow the gas to collect in the inverted tube. 
 When the tube is full, which is readily detected by the rapid escape 
 of the gas from the mouth, replace it by a fresh one, and make an 
 experiment with the full tube. 
 
 The heating of the mixture need not be continued during 
 the testing of the gas. Tests may be made to prove the 
 following statements : 
 
 1. If a piece of red litmus-paper is held in the gas, it is 
 turned blue. The gas has, therefore, alkaline properties. 
 
COMMON ELEMENTS AND COMPOUNDS 83 
 
 2. If a burning match or taper is put into an inverted tube 
 filled with the gas, the flame disappears, and the gas does not 
 take fire. It is non-combustible, and a non-supporter of com- 
 bustion under ordinary circumstances. If the gas is heated 
 strongly, or is mixed with oxygen, it burns with a somewhat 
 feeble, yellowish flame; water is formed, and nitrogen 
 liberated. 
 
 3. If a well-filled tube of the gas is placed mouth down- 
 wards in a dish of water, the liquid rises in the tube to take 
 the place of the dissolved gas, which is very soluble in water. 
 A piece of red litmus-paper dipped into the water before the 
 dissolution is unchanged ; but if dipped in after, is turned 
 blue. 
 
 4. If a piece of clay tobacco pipe is connected with the 
 delivery-tube by flexible rubber, and its end is strongly 
 heated in the Bunsen flame, while a current of gas is driven 
 through it by heating the mixture, the gas burns at the hot 
 end of the clay pipe. The change is : 
 
 The solubility of ammonia is its most remarkable property. 
 Cold water will absorb upwards of 1,000 times its own volume 
 of the gas. The ordinary ammonia liquor is simply a strong 
 solution of the gas, and has many useful applications. The 
 alkaline character of the solution, and the readiness with 
 which the gas escapes from it, has earned for the compound 
 the name of the volatile alkali. 
 
 The principal fixed alkalies are caustic potash and caustic 
 soda, and their solutions in water have properties very similar 
 to those of ammonia solution. But they are both solids, 
 which can be melted and raised to a high temperature without 
 decomposition. The result, then, of evaporating separate 
 solutions of these three alkalies would be a clean vessel for 
 the first, and solid residues for the other two. All three, 
 however, neutralize acids, turn litmus blue, or methyl orange 
 yellow, and have a caustic action on the skin. 
 
 62 
 
84 METALLURGICAL CHEMISTRY 
 
 The formulae for caustic soda and caustic potash are NaHO 
 and KHO, and, by analogy, the formula of the compound of 
 ammonia supposed to be present in its solution is written 
 (NH 4 )HO, orNH 4 HO. 
 
 Using the systematic names, the following comparison can 
 be made : 
 
 NaHO, sodium hydroxide. 
 KHO, potassium hydroxide. 
 (NH 4 )HO, ammonium hydroxide. 
 
 The group (NH 4 ), although it has no free existence, plays 
 an important part in ammonium compounds. It resembles an 
 atom of sodium or potassium in its chemical functions, and is 
 called a compound radicle on that account. 
 
 SUMMARY. 
 
 Common salt, saltpetre, and black oxide of manganese are 
 found ready to hand in Nature, and, in conjunction with 
 sulphuric acid, are used in the preparation of other equally 
 important compounds and elements. The common acids are 
 exceptionally important in the treatment of metals. Ammo- 
 niacal gas liquor is the usual source of ammonia and ammonium 
 compounds. 
 
 QUESTIONS. 
 
 1. Explain what takes place when common salt is heated 
 with moderately strong sulphuric acid. Sketch the apparatus 
 used. 
 
 2. How is chlorine gas prepared, and what are its most 
 important properties 1 Describe briefly the properties of its 
 companion elements. 
 
 3. What happens when potassium chlorate is heated 
 (a) alone, (b) mixed with black oxide of manganese 1 
 
 4. How is nitric acid prepared ? 
 
 5. Give a short description of the three common acids. 
 
 6. Write a short description of ammonia and its compounds. 
 
CHAPTER VII 
 METALS AND ACIDS 
 
 THE common forms of hydrochloric, nitric, and sulphuric acids 
 are now familiar, and their more evident properties have been 
 sufficiently studied to enable them to be handled intelligently. 
 The relation between metals and the common acids is an 
 important one, and the study of the reactions which take 
 place when they are mixed together can now be profitably 
 undertaken. In many cases a very marked change takes 
 place, which is readily noticed, and often easily investigated. 
 But in some cases no change is effected, and the metal and 
 acid may be left together for any length of time without 
 alteration. 
 
 EXP. 43. Put small pieces of zinc, tin, and platinum into three 
 test-tubes, and add to each 5 c.c. of dilute nitric acid. Warm the 
 tube, if necessary, to start the action, and note 'all the changes 
 which take place. 
 
 The acid which is added to the zinc boils up, and a rapid 
 escape of bubbles of gas is observed. The metal entirely 
 disappears. 
 
 The tin changes to a white powder, and a red gas with an 
 acrid odour is given off. 
 
 The platinum is apparently unchanged, and, if weighed 
 before and after the experiment, is found to be the same. 
 
 Two practical problems now present themselves : The first 
 is to learn as much as possible about the liberated gas ; the 
 second is to find out what happens to the metal. 
 
 85 
 
METALLUKGICAL CHEMISTKY 
 
 To FIND THE VOLUME OF GAS EVOLVED WHEN A WEIGHED 
 PIECE OF METAL is DISSOLVED IN AN EXCESS OF 
 DILUTE ACID. 
 
 The apparatus shown in Fig. 21 is to be fitted up. A is an 
 8-ounce flask fitted with a rubber bung and an angle-tube. 
 B is a 12-ounce bottle fitted with a rubber bung and two angle 
 tubes. C is a gas-cylinder of about 500 c.c. capacity. B is 
 connected with A and with the delivery-tube d by pieces 
 of sound, well-fitting rubber tube about 4 inches long. These 
 
 flexible connec- 
 tions allow the 
 bung to be re- 
 moved from or 
 inserted in the 
 neck of A, and 
 the delivery -tube 
 d to be raised 
 above B without 
 disturbing the 
 apparatus. Any form of pneumatic trough may be used, but 
 one with a beehive shelf is very convenient. The apparatus 
 must be perfectly gas-tight, so that care should be exercised 
 in the selection of the fittings and in the fitting up, or poor 
 results will be obtained. When ready, the apparatus may be 
 used for several experiments in succession, as a single experi- 
 ment requires a few minutes only for completion. 
 
 METALS AND HYDROCHLORIC ACID. 
 
 Zinc EXP. 44. Use the apparatus shown in Fig. 21. Add 
 a little water to the flask B, shake well so as to wet the sides, 
 allow it to stand for a minute, and then pour out the excess of 
 water. Remove the bung from the flask A, and add 25 c.c. of 
 dilute hydrochloric acid (1 to 1). Cut a strip of zinc foil about 
 1 gram in weight, and weigh it accurately. Fill the cylinder C, 
 invert it on the shelf of the trough, and arrange the end of the 
 delivery-tube under its mouth. Examine the connections to see 
 that they are well made, as the experiment is a failure if the whole 
 of the expelled gas is not collected in C. Drop the weighed metal 
 
 FIG. 21. 
 
METALS AND ACIDS 87 
 
 into A, and rapidly insert the bung. If, when the action has 
 ceased, the flask A is warm, due to the heat evolved during the 
 dissolution of the metal in the acid, place it in a shallow pan 
 containing water which has been standing in the laboratory for 
 some time. This will quickly bring the flask and its contents 
 back again to the original temperature, especially if the flask is 
 moved about in the water. When the level of the water in the 
 delivery-tube is constant, raise the tube and allow the liquid to 
 run back into B. Pour the water from B into a 50 c.c. measure, 
 and note its volume. This is the measure of the excess of gas 
 driven into C by the increase in the temperature in A. Now 
 moisten a narrow strip of filter-paper, and fix it on C a little above 
 the water-level in the cylinder ; depress the cylinder in the trough 
 until the water-level is the same inside and outside ; move the 
 paper till its lower edge marks this level. Kemove the cylinder 
 from the trough, and put a lighted taper to its mouth. The ^gas 
 is a mixture of air and hydrogen, and burns with a slight explosion. 
 Now fill the cylinder to the mark with water, rinse out the 500 c.c. 
 measure, pour in the water from the cylinder, and note its volume. 
 This is the measure of the volume of gas collected in C, and, there- 
 fore, of the volume of gas liberated from the acid by the weighed 
 quantity of zinc, which must displace that volume from A and B. 
 
 Further information about the experiment is obtained by 
 noting the height of the barometer column, which gives the 
 pressure of the gas, and by taking the temperature of the 
 water in the trough with a thermometer, which is also the 
 same as the temperature of the gas collected over it. For the 
 volume of a gas varies considerably with variations of its 
 temperature and pressure, and is increased by the presence of 
 water vapour in it. In accurate determinations it is necessary 
 to reduce this volume to that of the dry gas at the normal 
 temperature and pressure (N.T.P.). But the beginner may 
 leave this for a time. A series of experiments has shown that 
 an approximate correction for the volume of a gas collected 
 over water under ordinary laboratory conditions is obtained 
 by multiplying the observed volume by 0'916. 
 
 EXAMPLE. Weight of zinc taken 1 gram. 
 
 Volume of gas collected in C = 375 c.c. 
 ,, water in B = 3 ,, 
 
 ,, gas liberated = 372 ,, 
 
 Then 372 x 0'916 = 340'7 c.c. = the volume of dry hydrogen 
 at N.T J?. expelled from hydrochloric acid by 1 gram of zinc., 
 
88 METALLUEGICAL CHEMISTKY 
 
 The weight of a litre of hydrogen has been determined with 
 considerable accuracy, and may be used to find the weight of 
 any given volume of the dry gas at N.T.P. 
 
 1 litre = 1000 c.c. = 0-0896 gram. 
 
 Now, the volume of hydrogen liberated by 1 gram of zinc 
 having been found, it is easy to calculate the volume liberated 
 by 65 grams of the metal i.e., by an atomic weight of zinc in 
 grams. 
 
 340-7x65 = 22145-5 c.c. 
 
 And, as each 1,000 c.c. of this gas weighs 0*0896 gram, the 
 weight of the whole is 22145-5 x 0-0896 + 1000= 1-984. 
 
 The number thus obtained from a rough experiment is 
 sufficiently near the whole number 2 to suggest that an 
 atomic weight of zinc in grams liberates 2 atomic weights 
 of hydrogen in grams. From the formula HC1 it is seen that 
 2 molecular weights of hydrochloric acid in grams are 
 required to supply the 2 grams of hydrogen, and as the whole 
 of the chlorine combines with the zinc, the complete change is 
 expressed by the equation : 
 
 Zn + 2HC1 = ZnCl 2 + Ho 
 65 2(1 + 35-5) 65 + 71 2 
 
 Magnesium EXP. 45. Weigh 0-3 gram of magnesium ribbon 
 or wire. Add 50 c.c. of dilute hydrochloric acid (1 in 20). Make 
 the magnesium into a coil, drop it into the flask, and quickly 
 replace the bung. The gas comes off rapidly. Follow the instruc- 
 tions given for EXP. 44, and make the necessary measurements. 
 Prove that the liberated gas is hydrogen. 
 
 EXAMPLE. Weight of magnesium used = 0*3 gram. 
 Volume of gas in C = 312 c.c. 
 
 water in B = 7 ,, 
 
 gas liberated =305 
 
 NOTE. The principles upon which the above calculations are based will 
 be found in the Appendix. 
 
METALS AND ACIDS 89 
 
 Then 305x0-916 = 279-4. Compare this result with the 
 equation : 
 
 Mg + 2HClMgCl a +H 2 . 
 
 24 22320 c.c. 
 
 Iron. The dissolution of this metal is less rapid than that 
 of zinc or magnesium, and, when the evolution of the gas 
 slows down, it is necessary to heat the flask. This can be 
 done by placing it on gauze over a circular Bunsen burner. 
 The flexible connection between A and B, Fig. 21, enables 
 this to be readily done. The heating is continued over a small 
 flame until bubbles of gas cease to escape from the delivery- 
 tube, or escape very slowly. As commercial iron contains 
 impurities, the volume of gas liberated by a given weight of 
 the metal falls a little short of that which would be obtained 
 by using the pure metal. 
 
 EXP. 46. Weigh 0'7 gram of clean iron filings ; transfer it to 
 the flask A, Fig. 21 ; add 25 c.c. of strong hydrochloric acid, and 
 rapidly replace the bung. When the action slows down, heat the 
 flask gently until the dissolution is finished. Cool the flask and 
 make the necessary measurements. Prove that the gas is hydrogen. 
 During the heating a considerable volume of air will be expelled 
 from the flask, and a much larger quantity of water will run back 
 into B. When the bung is removed from A an offensive smell 
 will be noticed, which is due to the action of the acid on the 
 combined carbon present in the iron. 
 
 EXAMPLE. Weight of iron = 0'7 gram. 
 
 Volume of gas in C = 464 c.c. 
 
 water in B = 168 
 
 gas liberated = 296 
 Compare the result with the equation : 
 
 Tin. This metal dissolves even less rapidly than iron. 
 The most suitable form for general use is the thin sheet metal, 
 which is readily obtained by rolling down a piece of grain tin, 
 or by hammering it out on a clean anvil. 
 
90 METALLUEGICAL CHEMISTKY 
 
 EXP. 47. Weigh about 1-3 grams of tin ; add 20 c.c. of strong 
 hydrochloric acid to A, Fig. 21 ; put in the metal, and heat the 
 flask gently until the dissolution is complete. Measure and 
 examine the gas collected in C. 
 
 The volume of gas obtained is a little greater than that 
 given by the equation : 
 
 Sn + 2HCl=SnCl 2 + H 2 . 
 
 EXAMPLE. Weight of tin = 1-3 grams. 
 
 Volume of gas in C = 390 c.c. 
 
 ,, water in B = 117 ,, 
 
 liberated gas = 273 
 
 Aluminium. The thin sheet metal is readily obtained, 
 and dissolves in moderately strong hydrochloric acid without 
 the application of heat. The change is expressed by the 
 equation : 
 
 EXP. 48. Weigh 0'3 gram of aluminium ; add 15 c.c. of dilute 
 hydrochloric acid (1 to 1) to A, Fig. 21, and put in the weighed metal. 
 The action is sufficiently rapid to finish without heating the flask, 
 and very little water runs back into B. Measure and test the gas 
 collected in C. 
 
 EXAMPLE. 0'318 gram of aluminium gave 426 c.c. of hydrogen ; 
 or 426x0-916=390-2 c.c. of the dry gas at N.T.P. 
 
 Then, since 1,000 c.c. of hydrogen weigh 0*0896 gram, the weight 
 of hydrogen which would be liberated by the dissolution of an 
 atomic weight of the metal in grams is given by : 
 
 390-2x0-0896x27 
 ' 0-318 x 1000 2 '97 = 3 nearly. 
 
 So that 1 atomic weight of aluminium displaces 3 atomic 
 weights of hydrogen from the acid. 
 
 The metals used in the above experiments are those which 
 are most readily acted upon by hydrochloric acid. Some 
 common metals are dissolved very slowly, and others, again, 
 are not attacked by the acid, even on prolonged heating. A 
 metal which dissolves slowly may have its rate of dissolution 
 increased by passing a current of air through the acid liquid. 
 
METALS AND ACIDS 91 
 
 This is no doubt due to the formation of oxide of the metal by 
 the oxidizing action of the air, and its dissolution by the acid. 
 
 It may be stated generally that hydrogen gas is set free 
 when a metal dissolves rapidly in either hot or cold hydro- 
 chloric acid. 
 
 The compounds formed by the replacement of hydrogen 
 in the hydrochloric acid molecules by the respective metals 
 are called chlorides. They are readily separated from the 
 solutions by evaporating the water and excess of acid, or by 
 crystallization. (See Exps. 52 and 53.) 
 
 METALS AND SULPHURIC ACID. 
 
 When hydrochloric acid reacts with metals the general 
 change is a simple one, corresponding to the simple character 
 of the acid itself. But with sulphuric acid there is the possi- 
 bility of secondary changes, for, if its molecules are regarded 
 as consisting of the two groups H 2 and (S0 4 ), the second group 
 may undergo changes by which new compounds are formed. 
 The simplest kind of change is when the group (S0 4 ) moves 
 into or out of molecules in much the same way as an atom of 
 an element does. This is the simple explanation of the 
 ordinary dissolution of a metal in the dilute acid by which 
 hydrogen is liberated. But when the acid is hot and concen- 
 trated, another reaction becomes prominent, by which sulphur 
 dioxide gas is set free. The probable explanation of this is 
 that, at the temperature used, a considerable proportion of the 
 acid molecules are split up, or dissociated into (H 2 ) and 
 (S0 4 ) groups, and that the hydrogen set free by the action of 
 the metal on other acid molecules helps to attack the (S0 4 ) 
 groups, and abstracts oxygen from them. This is expressed 
 in the equations : 
 
 2. 
 
 The experiments with dilute sulphuric acid are made in 
 exactly the same way as with hydrochloric acid, so that a 
 
92 METALLUKGICAL CHEMISTRY 
 
 detailed description will not be required. The dilute solution 
 is made by pouring the measured quantity of strong acid into 
 the measured quantity of water. The mixture is then stirred 
 with a glass rod, and allowed to cool. The apparatus shown 
 in Fig. 21 is to be used. 
 
 Zinc EXP. 49. Use about 1 gram of the thin sheet metal 
 and 20 c.c. of dilute sulphuric acid (1 in 6). Measure and test the 
 liberated gas. 
 
 If the zinc is of good quality the action is very slow at first, 
 and it may be necessary to warm the flask a little. 
 
 EXAMPLE. 1 gram of zinc gave 377 x 0-916=345-3 c.c. of 
 hydrogen. 
 
 This result may be compared with that obtained by cal- 
 culation from equation No. 1 on p. 19. 
 
 Magnesium EXP. 50. Use about 0'3 gram of the metal ribbon 
 and 50 c.c. of dilute sulphuric acid (1 in 50). Measure and test the 
 liberated gas. 
 
 EXAMPLE. 0-3 gram of the metal gave 305xO'916=280-4 c.c. 
 of gas. 
 
 Iron EXP. 51. Use about 0-7 gram of the metal and 20 c.c. 
 of the dilute acid (1 in 6). Heat the flask gently, as described in 
 Exp. 46. 
 
 EXAMPLE. 0-7 gram of iron gave 296 x 0'916=27ri c.c. of gas. 
 Fe + H 2 S0 4 =FeS0 4 + H 2 . 
 
 On comparing the volume of gas obtained with that 
 determined by calculation from the equation, it is found to 
 be deficient. This is due to the impurities in the commercial 
 metal. 
 
 A comparison of the results obtained above shows that the 
 volume of gas liberated depends upon the weight of metal 
 used, and not upon the nature of the acid. 
 
 ACID. 
 
 ZN 1 GRAM. 
 
 Mo 0-3 GRAM. 
 
 FE 0'7 GRAM. 
 
 HC1 .. 
 H 2 S0 4 
 
 3407 c.c. 
 345-3 
 
 279-4 C.C. 
 2804 ,, 
 
 i 
 
 271 '2 c.c. 
 271-1 
 
METALS AND ACIDS 93 
 
 The other common metals are either insoluble or dissolve 
 with difficulty in dilute sulphuric acid. 
 
 The compounds formed by the replacement of the hydrogen 
 in the acid molecules by the metals are called sulphates, 
 and it will be useful at this stage to follow up the disappear- 
 ing metals and learn something about the bodies they pass 
 into. 
 
 EXP. 52. Dissolve 2 grams of zinc in 20 c.c. of dilute sulphuric 
 acid (1 to 7) contained in a small beaker. Pour the clear solution 
 into a porcelain basin ; raise it to boiling over the Bunsen flame, 
 and set the basin aside until the next lesson. Then drain away 
 the liquid, and dry the separated solid on a piece of filter-paper. 
 
 Repeat the experiment with 2 grams of clean iron filings. 
 
 The solids thus obtained consist of crystals of zinc and 
 iron sulphates, and contain the remnants of the sulphuric 
 acid molecules, together with water of crystallization. The 
 formulae of the salts are : ZnS0 4 , 7H 2 and FeS0 4 , 7H 2 0. 
 
 EXP. 53. Measure 8 c.c. of dilute sulphuric acid (1 to 12) into 
 a large porcelain crucible which has been carefully weighed ; cover 
 the crucible with a watch-glass, and add 0'3 gram of magnesium, a 
 small piece at a time. When the whole is dissolved, drain the 
 watch-glass into the crucible, and wash off any __ 
 adhering liquid with a fine jet of water from a 
 wash-bottle. Place the crucible on a pipe-clay 
 tripod, so that the bottom is just clear of the 
 sand in the sand-bath, and allow the liquid to 
 evaporate to complete dryness. Heat the 
 crucible over the Bunsen flame, at first with jr IGt 22. 
 
 the lid on in case of spurting, until white 
 fumes cease to be given off. Allow the crucible to cool, and weigh 
 it. The final heating must be done in a fume-chamber. 
 
 EXAMPLE. Weight of crucible + residue = 18'665 grams. 
 
 = 17-182 
 = 1'483 
 From the equation : 
 
 Mg + H 2 S0 4 = MgS0 4 + H 2 
 24 24 + 32 + 64 
 
 24 grams of magnesium give 120 grams of residue. Therefore 
 
 1 20 x 0'3 
 0'3 grain of magnesium should give z = 1*5 grams. 
 
94 
 
 METALLURGICAL CHEMISTRY 
 
 Thus the weight of sulphate formed is just as definite as the 
 volume of hydrogen liberated. Similar experiments may be 
 made with other metals and acids, with similar results. 
 
 PREPARATION OF HYDROGEN GAS. 
 
 The dissolution of zinc in dilute hydrochloric or sulphuric 
 acid may be used as a source of hydrogen for experiments 
 to demonstrate its properties. If the apparatus described 
 on p. 24 is not available, this method is to be employed for 
 obtaining a quantity of the gas. 
 
 EXP. 54. A is a 12-oz. bottle fitted with a rubber bung, through 
 
 which passes a thistle funnel, B, 
 and a delivery tube. About 10 
 grams of zinc is put into A, 
 covered with water, and strong 
 hydrochloric acid added through 
 B, a little at a time. The gas is 
 thus liberated steadily, and is 
 collected in the cylinder C. The 
 first jar of gas contains some air 
 displaced from the generating- 
 bottle ; the gas collected after this 
 is sufficiently pure for ordinary 
 experiments. 
 
 Experiments for demonstrating the properties of the gas are 
 described on p. 25. 
 
 B 
 
 PREPARATION OF SULPHUR DIOXIDE. 
 
 As the action of concentrated sulphuric acid on metals is 
 not susceptible of easy quantitative experiments, only one 
 example will be given, and that the one commonly used to 
 obtain sulphur dioxide in sufficient quantity to demonstrate 
 its properties. 
 
 EXP. 55. Fit up the apparatus shown in Fig. 24. A is a 12-ounce 
 conical flask ; B is a wash- bottle containing a little sulphuric acid 
 (2 parts of acid to 1 part of water) ; C is a dry gas-cylinder. Put 
 about 10 grams of sheet copper into A, and add 25 c.c. of strong 
 
METALS AND ACIDS 
 
 95 
 
 FIG. 24. 
 
 sulphuric acid. Heat the flask until the gas comes off freely, and 
 
 then regulate the flame so as to obtain a steady stream of bubbles 
 
 through B. Col- 
 
 lect two jars of 
 
 the gas by dis- 
 
 placement of air, 
 
 then transfer the 
 
 delivery - tube to 
 
 a bottle contain- 
 
 ing water, in 
 
 which the gas 
 
 will dissolve and 
 
 form a solution. 
 
 When the evolu- 
 
 tion of gas is 
 
 nearly finished, 
 
 disconnect the apparatus and add 50 c.c. of water to the flask A ' 
 
 raise the contents to boiling, and filter the liquid into a porcelain 
 
 dish. Concentrate the blue solution to about half its bulk by 
 
 evaporation, and set it aside to crystallize. Blue crystals of copper 
 
 sulphate, CuS0 4 , 5H 2 0, are obtained. 
 
 The principal change is expressed by the equation : 
 
 Cu 4- 2H 2 S0 4 = CuS0 4 + 2H 2 + S0 2 . 
 Properties of Sulphur Dioxide. The gas collected in 
 
 the jars is colourless and transparent, and has the odour of 
 burning sulphur. When a burning candle is lowered into 
 one of the jars, the flame disappears, and the gas itself does 
 not burn. A solution of blue litmus is immediately turned 
 red when poured into the jar. When the other jar is inverted 
 over water and the plate removed, the water rises rapidly 
 up the jar to take the place of the dissolved gas. The 
 solution, made as described above, smells strongly of the gas, 
 and is considered to contain sulphurous acid, H 2 S0 3 , formed 
 by the reaction between the gas and water, thus : 
 
 The pure acid has not been isolated ; it is only known in 
 solution, but it forms salts called sulphites. 
 
96 METALLURGICAL CHEMISTRY 
 
 METALS AND NITRIC ACID. 
 
 Many important changes depend upon the readiness with 
 which nitric acid parts with oxygen to bodies capable of com- 
 bining with that element, and an interesting series of reactions 
 may be brought about. It may be remarked at the outset 
 that metals are more readily soluble in nitric acid than in 
 either of the two acids already dealt with. The metals may 
 be roughly divided into three classes with respect to the action 
 of nitric acid upon them. 
 
 1 . The metal is converted directly into its oxide : Sn 
 and Sb. 
 
 2. The metal is converted into a soluble nitrate : Cu, Mg, 
 Hg, Pb, Fe, Ag, Zn. 
 
 3. The metal is not attacked by the acid : Au and Pt. 
 
 The solid non-metals also are more or less readily converted 
 into oxides, which then combine with water to form the 
 corresponding acids. 
 
 Tin EXP. 56. Heat a clean porcelain crucible over the Bunsen 
 flame. When cold, weigh it carefully, and weigh in it 0*5 gram of 
 tin. Add 5 c.c. of dilute nitric acid (1 to 1) from a pipette. Put a 
 watch-glass over the crucible, and heat it carefully on the sand-bath 
 or iron plate. When the action is finished, remove the watch-glass, 
 and if any white particles adhere to it, rinse them into the crucible 
 with a fine jet of water from the wash-bottle. Now place the 
 crucible on a pipe-clay tripod (see Exp. 53), and evaporate the 
 contents slowly to dryness. Then put the crucible in a pipe-clay 
 triangle on an iron tripod stand, and heat it carefully with the 
 Bunsen flame. Hold the burner in the hand, move the flame 
 about, and be ready to remove it if the residue shows signs of 
 spurting. When risk of loss is over, heat the crucible strongly 
 for ten minutes, and then for a minute or two over the foot blow- 
 pipe flame. If a muffle is available, it can be used for the final 
 heating. When cold, weigh the crucible and oxide. 
 
 EXAMPLE. Weight of crucible + oxide = 11*770 grams 
 
 = 0-635 gram 
 Then 0'635 -0*5 = 0135 = weight of oxygen absorbed. 
 
METALS AND ACIDS 97 
 
 0*5 
 To find the formula of the compound we have : vy^ = 0*0042 
 
 and -T* = 0*0084. Therefore the ratio of the number of atoms 
 ID 
 
 of tin to the number of atoms of oxygen in the molecule is 
 
 0-0042 
 
 0-0084 = * : 
 
 Antimony EXP. 57. Use 0-5 gram of powdered antimony, and 
 5 c.c. of dilute acid (1 in 3;. Follow the instructions given for the 
 last experiment. The final heating must be done with the crucible 
 open, for if a lid is used, reducing gases may get into the crucible 
 and cause loss by volatilization. 
 
 EXAMPLE. Weight of metal used = 0'5 gram; weight of oxygen 
 absorbed = 0-138 gram. Use the data obtained to find the fomiula 
 of the oxide. Sb = 120. 
 
 When metals dissolve in nitric acid their nitrates are 
 formed. They are, however, readily decomposed on being 
 heated, and, in most cases, are converted into the oxides of 
 the contained metals. There are two notable exceptions to 
 this : When nitrate of silver is heated it is completely 
 decomposed, and a residue of metallic silver is obtained ; 
 when nitrate of mercury is heated it is first reduced to the 
 oxide, which on further heating splits up into mercury and 
 oxygen. 
 
 During the dissolution of a metal by nitric acid in an open 
 vessel a red gas appears to be given off as long as any of the 
 metal remains. If, then, the solution is evaporated to dryness, 
 and the solid residue strongly heated, a further evolution of 
 red gas is noticed. The final residue is found to be the oxide 
 of the metal, and its weight compared with that of the metal 
 used indicates the weight of oxygen absorbed. 
 
 Iron EXP. 58. Put 0'5 gram of clean iron wire into a weighed 
 porcelain crucible, and add 5 c.c. of dilute nitric acid (1 to 1). 
 Dissolve, evaporate, and heat the residue as described in Exp. 56. 
 Allow the crucible to cool, and reweigh it. If the wire is thick, 
 it should be first flattened under the hammer and rubbed with 
 emery-cloth before being weighed and dissolved. 
 
 7 
 
98 METALLUKGICAL CHEMISTKY 
 
 EXAMPLE. Weight of metal used = 0'5 gram ; weight 'of oxygen 
 absorbed = 0-215 gram. 
 
 Then = 0-0089, and = 0-0134. Therefore 
 
 or, taking the nearest whole numbers, the ratio is 2 : 3, and the 
 formula of the oxide is Fe 2 O 3 . 
 
 Mercury EXP. 59. Use about 1 gram of mercury and 5 c.c. 
 of dilute acid (1 to 2). Proceed as in Exp. 58. After the evapora- 
 tion, to obtain the white residue of nitrate, the crucible must be 
 heated very carefully by holding the Bunsen burner in the hand 
 and moving it about. The white colour of the residue gradually 
 changes to red, and the action is finished when its colour is 
 uniformly red. It must not become dark-coloured during the 
 heating, or loss of metal and oxygen will take place. Eeweigh 
 the crucible when cold. 
 
 EXAMPLE. Weight of mercury 1'07 grams; increase 0'075 
 gram. 
 
 Find the formula of the oxide from the data thus obtained. 
 Hg = 200. 
 
 Silver EXP. 60. Use, 0*5 gram of silver and 5 c.c. of dilute 
 acid (1 to 1). A residue = 0*5 gram of silver is obtained. 
 
 Similar experiments may be made with zinc, magnesium, 
 lead, and bismuth. The result with the last-named metal 
 leads to the formula BLOo. 
 
 THE GASES EVOLVED WHEN METALS AND NITRIC ACID 
 REACT. 
 
 It will be useful to consider, as far as can be done in an 
 elementary way, the nature of the gases given off when metals 
 dissolve in nitric acid. It has already been noticed that a red 
 gas appears to be given off almost invariably when the dis- 
 solution takes place in an open vessel. But when the experi- 
 ment is made in a closed vessel, it is found that the gas 
 actually escaping from the liquid is not of constant composi- 
 tion for all metals. In fact, for the same metal even the 
 volume and composition may vary with (1) the concentration 
 of the acid ; (2) the temperature at which the reaction is pro- 
 
METALS AND ACIDS 99 
 
 ceeding; (3) the depth of the layer of acid liquid above the 
 dissolving metal, and through which the escaping gas has to 
 pass. It also varies with the nature of the metal. 
 
 During the changes which nitric acid undergoes several 
 oxides of nitrogen, and even nitrogen itself, may be liberated. 
 The hydrogen of the acid is almost invariably converted into 
 water by union with oxygen. Under exceptional circum- 
 stances it may be converted into ammonia, NH 3 , by union 
 with nitrogen, or even liberated in the free state. 
 
 There are three well-known and easily-obtainable oxides of 
 nitrogen : nitrous oxide, N 2 ; nitric oxide, NO ; and nitric 
 peroxide, N0 2 , or N 2 4 . 
 
 Nitrous Oxide, N 2 0, is a colourless gas with a faint odour 
 and sweet taste. It does not change when mixed with air, 
 and is somewhat soluble in water. If a glowing splint is put 
 into a jar of this gas, the splint bursts into flame ; but when 
 feebly-burning sulphur is brought into contact with it, the 
 flame disappears. This test distinguishes nitrous oxide from 
 oxygen. It is used by dentists to produce insensibility to 
 pain for a short time. Its popular name is " laughing gas." 
 
 EXP. 61. Fit up apparatus similar to that shown in Fig. 23, but 
 use an 8-ounce bottle, so as not to have a large volume of air for the 
 gas to displace, and arrange a small wash- bottle, nearly full of 
 strong nitric acid, in the path of the gas to absorb any nitric oxide 
 which may pass over. Put an excess of zinc into the bottle, and add 
 a small quantity of water through the thistle funnel to just cover 
 the bottom. Place the bottle in a dish of water to prevent a too 
 rapid rise in temperature, and add dilute nitric acid (1 to 1), a little 
 at a time, to keep up a steady stream of bubbles. Collect two jars 
 of the gas, over warm water preferably. Remove the plate from 
 one of the jars; no change in colour is observed. Plunge a glowing 
 splint into the jar ; the splint bursts into flame. Heat a little sulphur 
 in a deflagrating spoon until it shows a small flame, and plunge it 
 into a jar of the gas ; the flame disappears. Heat the sulphur until 
 it is burning vigorously, and put it into the jar again ; it continues 
 to burn, and with a brighter flame than when burning in the air. 
 
 Nitrous oxide supports combustion more vigorously than 
 air does, but the burning body must have a sufficient tempera. 
 
 72 
 
100 METALLUEGICAL CHEMISTRY 
 
 ture to decompose the gas and abstract its oxygen, if the 
 combustion is to continue. 
 
 Nitric Oxide, NO, is a colourless gas, which is converted 
 very rapidly into nitric peroxide when mixed with air or 
 oxygen. The gas thus formed has a red colour, so that the 
 change is easily observed, and it is commonly used to dis- 
 tinguish oxygen from nitrous oxide, which does not react with 
 nitric oxide. Thus, the addition of nitric oxide is a test for 
 oxygen, and the addition of oxygen is a test for nitric oxide 
 in mixtures of gases. Nitric oxide is only slightly soluble in 
 water, and does not support the combustion of ordinary burning 
 bodies. 
 
 EXP. 62. Eepeat the last experiment, using thin sheet copper in 
 the place of zinc. The bottle need not be cooled, and the wash- 
 bottle must be omitted. Collect two jars of the gas, and remove 
 them from the trough. Collect one tall cylinder about two-thirds 
 full of the gas, and leave it standing over water. Eemove the plate 
 from one of the jars ; a red gas forms in the mouth of the jar. 
 Plunge a burning candle into the jar : the flame disappears. Just 
 ignite a small piece of phosphorus in a deflagrating spoon and 
 plunge it rapidly into a jar of the gas : the flame disappears. Allow 
 the phosphorus to burn vigorously, and return it to the jar : it burns 
 more vigorously in the nitric oxide than in air. Decant half the 
 volume of oxygen into the jar standing over water : grip the side of 
 the jar in the hand while transferring the oxygen to it. A red gas 
 forms, and dissolves rapidly, the water rising in the jar after each 
 addition of oxygen, and the jar becomes quite warm in the hand. 
 
 Nitric oxide, although not a supporter of combustion in the 
 ordinary sense, will give up oxygen to bodies which have a 
 temperature sufficiently high to effect its decomposition. 
 
 The reaction between nitric oxide and oxygen is expressed 
 by the equation : 
 
 Heat is developed by the combination of the two gases. 
 
 Nitric Peroxide, N0 2 , is a red gas, very soluble in water, 
 with which it reacts, and forms an acid solution containing 
 nitrous and nitric acids. The peroxide itself is not an 
 
METALS AND ACIDS 101 
 
 acid-forming oxide. The change is expressed by the equa- 
 tion : 
 
 2N0 2 + H 2 = HN0 3 + HN0 2 . 
 
 Nitric Nitrous 
 acid. acid. 
 
 The formula N0 2 does not represent the composition of nitric 
 peroxide molecules under all conditions. At a low tempera- 
 ture the composition is more nearly represented by N 2 4 . If 
 the temperature is allowed to rise, the colour of the gas 
 deepens, and the complex molecules of N 2 4 are gradually 
 decomposed into the simpler molecules of N0 2 . At a 
 moderately high temperature the composition of the gas is 
 represented by N0 2 . At a still higher temperature the 
 whole of the gas is decomposed into nitric oxide and oxygen, 
 and the colour entirely disappears. As the temperature falls 
 the inverse changes take place. This is a good example of 
 " dissociation." 
 
 Nitric peroxide readily gives up part of its oxygen to 
 bodies eager to combine with that element. This is taken 
 advantage of in the manufacture of sulphuric acid, during 
 which nitric oxide is made to act as a carrier of oxygen from 
 the air to sulphur dioxide for its conversion into the trioxide, 
 thus : 
 
 NO + = N0 2 , and then N0 2 + S0 2 = S0 3 + NO. 
 
 These reactions are repeated as long as nitric oxide, sulphur- 
 dioxide, and oxygen are present in the acid chambers. 
 
 When silver dissolves into nitric acid, both nitric peroxide 
 and nitric oxide are liberated. 
 
 The principal reactions of zinc, copper, and silver with nitric 
 acid may be represented by equations : 
 
 1. 4Zn + 10HN0 3 = 4Zn(N0 3 ) 2 + 5H 2 + N 2 0. 
 
 2. 3Cu + 8HN0 3 = 3Cu(N0 3 ) 2 + 4H 2 O + 2NO. 
 3. 
 
 The nitrate of the metal is formed in each case. 
 
 It is usually stated that hydrogen gas cannot be obtained 
 
102 
 
 METALLURGICAL CHEMISTEY 
 
 by the action of- a metal on nitric acid ; but if magnesium is 
 substituted for zinc in Exp. 61, the bottle kept cool, and the 
 acid dilute, the greater part of the evolved gas is found to be 
 hydrogen. But if the acid is too strong, and the temperature 
 is allowed to rise, the proportion of hydrogen in the evolved 
 gas is considerably diminished. 
 
 EFFECTS OF DILUTION AND TEMPERATURE. 
 
 The effects of a variation in the strength of nitric acid, and 
 in the temperature at which the reaction takes place on the 
 volume of the escaping gas, are interesting, and instructive 
 
 experiments may be 
 made with the appa- 
 ratus shown in Fig. 25. 
 A is a test-tube fitted 
 with a bung and angle- 
 tube. B is a small tap 
 funnel. C is a fixed 
 delivery-tube con- 
 nected with the angle- 
 tube from A by a 
 piece of sound rubber 
 tube, and passing 
 under the mouth of the graduated cylinder D. If it is 
 desired to keep the test-tube A cool during an experiment, 
 it is immersed in a beaker containing cold water ; or, when 
 the beaker is away, the test-tube can be heated by a Bunseii 
 flame. 
 
 The conduct of an experiment is simple. When the parts 
 of the apparatus are in position, water is poured into B and 
 the tap opened. The water running from B into A displaces 
 the air in A and the delivery-tube connected with it. The 
 displaced air collects in D, and its volume can be measured. 
 Thus the volume of air in the apparatus at the beginning of 
 an experiment is known, if the volume of the added acid is 
 
 FIG. 25. 
 
METALS AND ACIDS 103 
 
 subtracted from the total volume obtained as above. The 
 water is then emptied out, and the necessary quantity of acid 
 added to A ; the coil of weighed metal is dropped in, and the 
 bung quickly replaced. If the tube is to be kept cool, the 
 beaker of water is placed in position. When the metal is 
 completely dissolved, water is poured into B, the tap opened, 
 and the whole of the gas driven from A to D. The volume 
 of gas in D is then measured by depressing the cylinder in 
 the trough until the level of the water is the same inside and 
 outside, and then taking the reading on the graduated scale. 
 Half the volume of oxygen is then passed into D by decant- 
 ing the gas from a graduated cylinder. When the red gas 
 formed by the reaction between the nitric oxide and oxygen 
 is all dissolved, and the residue is colourless, the cylinder is 
 depressed, and the volume of the residual gas measured. 
 Sufficient data are thus obtained for roughly finding the com- 
 position of the evolved gas. 
 
 When copper is the dissolving metal, nitric oxide, NO, is 
 the principal constituent of the liberated gas. Now, it has 
 been shown already that when this gas is mixed with oxygen 
 the red gas, nitrogen peroxide, N0 2 , is 'formed and it is easily 
 proved that nitric oxide combines with exactly half its own 
 volume of oxygen during the change. The reaction is shown 
 by the equation : 
 
 4 vols. 2 vols. 
 
 Iii the experiments made to obtain the following results, strips 
 of thin sheet copper of uniform size, weighing exactly 1 gram, 
 were used, together with 15 c.c. of the acid. In this way the 
 volume of air in the apparatus was kept the same in each 
 experiment. Now, it must be borne in mind that one-fifth of 
 the enclosed air is oxygen, which will combine with twice its 
 own volume of nitric oxide. 
 
 EXAMPLE. Volume of air in the apparatus after adding the 
 acid = 50 c.c. Total volume of gas in D after dissolution of the 
 metal, and displacement of gas in the apparatus by water run in 
 
104 
 
 METALLURGICAL CHEMISTRY 
 
 from B = 158 c.c. ; volume of oxygen added = 80 c.c. ; volume of 
 residual gas in D = 76 c.c. 
 
 Then 158 + 80 -76 -162 = the volume of NO and 2 , 
 which has disappeared, and, since two-thirds of this is NO 
 
 162 x 2 
 
 =108 = the volume of nitric oxide in the collected gas. 
 
 o 
 
 But 20 c.c. of NO are also accounted for by the 10 c.c. of 
 oxygen in the enclosed air at the beginning of the experiment. 
 Therefore, the total volume of NO liberated = 128 c.c. 
 
 Also the residual gas contains 40 c.c. of nitrogen from 
 the enclosed air, and 26 c.c. of added oxygen. Therefore, 
 76-40-26 = 10 c.c. of gas other than nitric oxide liberated 
 from the acid. 
 
 Therefore, the NO in the liberated gas = ^-^ = 
 
 92-7 per cent. 
 
 The following table shows the results of two experiments 
 conducted as described above. It demonstrates clearly that 
 the volume and nature of the gas liberated during the reaction 
 between nitric acid and a particular metal depends upon the 
 conditions under which the reaction takes place. One gram of 
 copper was used for each experiment. 
 
 It may be remarked that nitric oxide is soluble in the nitric 
 acid solution, and that its solubility decreases as the tempera- 
 ture increases. This would account largely for the difference 
 in the volumes of the gases liberated during the two experi- 
 ments. The changes taking place between nitric acid and 
 metals deserve more attention than they have yet received ; 
 but the subject is a difficult one, and the experiments described 
 above only give a rough indication of the method. 
 
 STRENGTH OF 
 ACID. 
 
 VOLUME 
 
 OF 
 
 VOLUME 
 
 OF 
 
 OXYGEN 
 
 VOLUME 
 
 OF 
 
 RESIDUAL 
 
 PERCENT- 
 AGE OF 
 NITRIC 
 
 TEMPERATURE. 
 
 ACID. 
 
 WATER. 
 
 GAS. 
 
 ADDED. 
 
 GAS. 
 
 OXIDE. 
 
 
 1 
 
 1 
 
 158 c.c. 
 
 80 C.C. 
 
 76 c.c. 
 
 927 
 
 Allowed to rise 
 
 1 
 
 1 
 
 106 
 
 52 
 
 84 
 
 77-5 
 
 Kept constant 
 
METALS AND ACIDS 105 
 
 ACTION OF ACIDS UPON COMMERCIAL IRON. 
 
 The different varieties of commercial iron known as cast 
 iron, steel, and wrought iron owe their variations in properties 
 very largely to the presence of carbon in the metal. This 
 element may be present, not only in varying proportions, but 
 also in different conditions, and thus cause a wide variation in 
 the properties of the metal. 
 
 For our present purpose carbon may be said to exist in 
 iron in two distinct forms. As Graphite it is mechanically 
 mixed in the form of crystalline scales with the particles of 
 the iron. These scales are readily seen on looking at the 
 fractured surface of a piece of grey pig-iron with a pocket 
 lens. As Carbide it is combined with part of the iron in the 
 form of iron carbide, Fe 3 C, which cannot be observed without 
 special preparation of the surface of the metal, and its examina- 
 tion under the microscope. 
 
 Cast iron contains on the average 3 '5 per cent, of carbon, 
 which exists largely as graphite in grey iroir, largely as carbide 
 in white iron, or in both forms in mottled iron. The amount 
 of carbon in steel varies considerably, and may be as high as 
 1*8 per cent. When the steel is in its normal or annealed 
 state, the carbon is in the form of carbide. Wrought iron 
 may contain up to 2 per cent, of carbon, which is present as 
 carbide. On account of these variations the common acids, 
 though able to dissolve the metal readily, give different results 
 with different varieties. 
 
 Hydrochloric and dilute sulphuric acids act similarly, but 
 the action varies with the condition of the carbon. The rate 
 of dissolution increases with the temperature. Graphite 
 separates as such, and being insoluble in the acid liquid, settles 
 out, and forms a black deposit on the bottom of the vessel in 
 which the dissolution is taking place. The carbide is decom- 
 posed by the acid, assisted by the newly liberated hydrogen, 
 and the freed carbon combining with hydrogen forms hydro- 
 carbons, which mix with the escaping gas, and give to it a 
 
106 METALLUEGICAL CHEMISTKY 
 
 very disagreeable odour. The presence of carbide carbon is 
 detected in this way. 
 
 Dilute nitric acid acts differently as regards the carbide 
 carbon, which in this case passes into the solution in combina- 
 tion with oxygen and hydrogen. The compound formed 
 imparts a distinct colour to the solution. Under similar con- 
 ditions the depth of the colour varies with the proportion of 
 carbon present. Graphite, however, is set free, and separates 
 out in the same way as with the other acids. 
 
 The other elements present in commercial iron are silicon, 
 phosphorus, sulphur, and manganese, and the purity of the 
 metal depends upon the proportions present, which may vary 
 from traces to several per cents. Each one has its own parti- 
 cular influence upon the properties of the metal, and has to be 
 reckoned with by the practical man. During the dissolution 
 of the metal in acids they undergo various changes. The 
 silicon is oxidized to silica and left as a solid residue, inter- 
 mixed with any graphite that may be present. Phosphorus 
 and sulphur are oxidized to phosphoric and sulphuric acids by 
 moderately strong nitric acid, and pass into solution. With 
 hydrochloric and sulphuric acids the sulphur escapes entirely 
 as sulphuretted hydrogen gas, and the phosphorus partially as 
 phosphoretted hydrogen. The iron and any manganese which 
 may be present dissolve in and form salts of the acids used. 
 
 ACTION OF ACIDS ON METALLIC ALLOYS. 
 
 Alloys are bodies containing two or more metals, which may 
 be present under different conditions. Sometimes the relation 
 between them is so intimate that the alloy may be regarded 
 as a chemical compound. Sometimes the body is simply a 
 solidified solution of its constituent metals. 
 
 As a rule, neither hydrochloric nor dilute sulphuric acid 
 will dissolve these bodies, and the more intimate the relation 
 between their constituents, the more difficult they are to dis- 
 solve. Dilute nitric acid is the best general solvent, and is 
 commonly used. If both metals in an alloy dissolve indi- 
 
METALS AND ACIDS 107 
 
 vidually in the acid, then the alloy is also soluble ; but if one 
 of the metals is insoluble, it is left as a residue of oxide or 
 metal, as the case may be. A copper-zinc alloy (brass) is 
 completely dissolved. A copper-tin alloy (bronze) is partially 
 dissolved, and a residue of white tin oxide is obtained. A 
 gold-silver alloy containing more than 2J parts of silver to 
 1 part of gold is partially dissolved, and a residue of metallic 
 gold left. 
 
 The last two reactions furnish the means of separating tin 
 from copper, and gold from silver. 
 
 EXP. 63. Put about 1 gram each of brass and bronze into 
 separate test-tubes ; add 10 c.c. of dilute nitric acid to each, and 
 heat the tube, if necessary, until all action ceases. Repeat the 
 experiment, using dilute sulphuric and hydrochloric acids in place 
 of the nitric acid. 
 
 When brass dissolves, copper and zinc nitrates are formed ; 
 and the partial dissolution of bronze results in the formation 
 of copper nitrate and hydrated tin oxide. 
 
 ELECTRO-CHEMICAL ACTION. 
 
 An important result of the dissolution of a metal in an acid 
 has now to be considered. It has already been noticed that 
 when the action is brisk, the vessel in which it is taking place 
 gets quite hot. It may, then, be said that heat is one of the 
 products of the chemical change ; and under ordinary circum- 
 stances, the amount of heat evolved is just as definite as the 
 volume of gas liberated, or the weight of salt formed. This 
 is easy to prove with suitable apparatus, but it is outside 
 ordinary elementary work, and must, therefore, be here taken 
 for granted. 
 
 Pure zinc is scarcely attacked by cold dilute sulphuric acid, 
 but if a piece of copper is placed in contact with the zinc in 
 the dilute acid, the action begins, and gas is liberated from the 
 surface of the copper. It would thus appear at first sight 
 as though the copper were dissolving, but it soon becomes 
 evident that this is not so. The solution remains colourless, 
 
108 METALLUEGICAL CHEMISTKY 
 
 which would not be the case if copper sulphate were being 
 formed. Also, the zinc gradually disappears, but the copper, 
 if weighed before being put in, and weighed again after being 
 taken out of the solution, is found to be unchanged. Clearly, 
 then, the contact of the copper with the zinc causes the latter 
 to dissolve more readily, but does not otherwise affect the 
 chemical change. And the solution gets hot, too, so that heat 
 is evolved as usual. 
 
 If ordinary sheet zinc is attacked by the dilute acid, it 
 should be cleaned by dipping it in moderately strong acid, 
 and then coated with mercury. Zinc so treated is said to be 
 amalgamated, and acts with the dilute acid in the same way 
 as the pure metal. Silver, platinum, or gas-carbon may be 
 used in the place of copper, to set up the action. 
 
 EXP. 64. Put a piece of pure zinc or the amalgamated metal 
 into a test-tube. Add a few cubic centimetres of dilute sulphuric 
 acid (1 to 20), and allow the tube to stand for a minute or two. 
 Note that no action is going on. Drop a piece of sheet copper on 
 to the zinc. The action commences at once, and soon bubbles of 
 gas are freely liberated from the copper. The solution does not 
 change colour, and the zinc gradually dissolves. 
 
 A silver coin, a piece of platinum foil, or a carbon rod may 
 be substituted for the copper with similar results. 
 
 If long strips of copper and zinc are only partly immersed 
 in the acid liquid, and the ends outside are made to touch, the 
 same result is obtained. Zinc dissolves, and hydrogen gas is 
 liberated from the copper, even though the immersed portions 
 are some distance apart. Also the ends of the plates may be 
 joined by a metal wire of considerable length with similar 
 results. But directly the connection is severed the action 
 ceases. 
 
 EXP. 65. Cut strips of thick sheet zinc and sheet copper, about 
 6 inches long by 1 inch wide ; immerse the zinc strip about three- 
 quarters the way up in moderately strong sulphuric acid for a few 
 seconds, and then rub a little mercury all over the cleansed surface 
 on both sides. Fix the plates, one on each side of a strip of wood 
 1 inch wide, so that their ends project a little above the upper 
 surface of the wood. This is easily done by driving short tacks 
 
METALS AND ACIDS 
 
 109 
 
 through the plates into the wood, but not in the same line, so that 
 there will be no danger of their meeting in the wood. Put the 
 plates into a beaker con- 
 taining dilute sulphuric 
 acid, which should 
 nearly cover the amal- 
 gamated part of the 
 zinc plate. Connect the 
 plates by forcing the 
 clean ends of a copper 
 wire between the plates 
 and the wood. Observe 
 that the action com- 
 mences as soon as the 
 connection is made. 
 Make a similar pair, 
 but with two copper A FIG. 26. B 
 
 plates instead of copper 
 
 and zinc. Put this pair into a solution of copper sulphate, and 
 connect the ends by copper wires, as shown in Fig. 26. Allow the 
 action to go on for a few minutes ; then take out the copper plates, 
 and examine them. 
 
 The copper sulphate solution makes the connection, and the 
 action goes on in A in much the same way as when the con- 
 nection is made direct by a single wire. The copper plate in 
 B, which is connected directly with the zinc plate in A, is 
 found to be coated with freshly deposited copper, and the 
 longer the action is allowed to proceed, the thicker the deposit 
 becomes. In fact, if the two plates are separately weighed 
 before and after immersion in B, one is found to increase and 
 the other to decrease in weight to exactly the same extent. 
 If the connections are reversed, copper is deposited on the 
 other plate. Evidently some kind of influence is transferred 
 from A to B, which carries the copper from one plate to the 
 other, and its direction depends upon the way in which the 
 connection is made. It seems to move from Cu in A to Cu l 
 in B, across the solution to Cu 2 , and back to Zn in A. This 
 is called an electric current. 
 
 If two or three cells similar to A have their alternate copper 
 and zinc plates connected by short thick copper wires, and 
 
110 METALLUEGICAL CHEMISTKY 
 
 the terminal copper and zinc are connected with carbon rods 
 or platinum plates in B, copper is still deposited on the plate 
 connected with the terminal Zn. But in this case nothing 
 dissolves from the other plate, and the copper sulphate is 
 gradually decomposed, oxygen being liberated from the plate 
 connected with the terminal Cu. Now, if the amount of heat 
 developed by the dissolution of zinc in the cell during the de- 
 composition of the copper sulphate is determined, it is found 
 to fall considerably short of the quantity which would be 
 developed by the dissolution of the same weight of zinc in the 
 ordinary way. The missing heat appears in the form of 
 electrical energy, which is carried round the system by the 
 current, and determines the decomposition of the sulphate. 
 
 If a very long thin iron wire is used to connect the two 
 plates in A (Fig. 26), nearly the whole of the energy developed 
 by the dissolution of the zinc is transferred to the wire. It 
 appears first as electrical energy, is then carried through the 
 wire by the current, and finally transformed into heat by the 
 resistance of the wire to the passage of the current. With a 
 short thick copper wire for connections, nearly the whole of 
 the heat appears in A. 
 
 The principle demonstrated in the last experiment is the 
 one upon which the electro-deposition of metals depends, and 
 it is put to a variety of uses. A large number of metals can 
 be so deposited, and it is only necessary that the surface upon 
 which the deposit takes place should be a conductor of 
 electricity. The common plan is to deposit a metal which has 
 specific properties, such as gold and silver, on account of their 
 lustre, and nickel in virtue of its hardness. It is even possible 
 to deposit two metals together, so as to form an alloy. Thus 
 the deposition of copper and zinc to form a deposit of brass is 
 common. The purification of metals by the use of electrical 
 energy is carried out on the large scale. Electrolytic copper 
 is a very pure metal, which is refined by this process, and is 
 largely used for electrical purposes, and for making high-class 
 copper alloys. 
 
METALS AND ACIDS HI 
 
 Electrolysis is the term used to denote the process through 
 which a metallic salt or oxide passes when it is decomposed 
 by an electric current. The body undergoing the change is 
 usually dissolved in water; but some solids which can be 
 melted without undergoing rapid decomposition are electro- 
 lized in the molten condition. The decomposing compound 
 is called the electrolyte, and the facts of electrolysis point 
 strongly to a decomposition of at least some of the molecules 
 of the compound by its dissolution in the solvent liquid alone. 
 These molecules, however, are only dissociated, for if the 
 solvent is removed by evaporation the original compound is 
 recovered intact. According to this, when copper sulphate is 
 the electrolyte it seems that an appreciable number of its 
 molecules are dissociated into Cu and S0 4 . The Cu re- 
 presents an elementary atom or simple radicle, and the S0 4 a 
 group of atoms or a compound radicle, which appears to have 
 properties similar to those of an elementary atom. It must 
 be clearly understood that these radicles are not supposed to 
 exist separately in the same way as an ordinary element or 
 compound. The existence of the one predicates the co-exist- 
 ence of the other in the solution of the salt. They are known 
 as ions, and the function of the current is supposed to be to 
 direct them towards the electrodes by which it enters and 
 leaves the solution. The first is the anode, to which the 
 anions move, and the second the cathode, to which the cathions 
 move. In the electrolysis of CuS0 4 , the S0 4 groups are the 
 anions and the Cu atoms the cathions. When copper elec- 
 trodes are used the Cu ions simply deposit on the cathode, 
 and lose their ionic condition, while the S0 4 ions deposit on 
 the anode and lose their ionic condition by abstracting copper 
 from the plate itself. But when platinum electrodes are used 
 the S0 4 ions attack water molecules in the proximity of the 
 anode, form H 2 S0 4 molecules, and set oxygen free, which 
 collects on the anode plate. The change with the copper 
 anode is in effect a transport of copper from one electrode to 
 the other ; with the platinum anode it is equivalent to the 
 
112 METALLURGICAL CHEMISTRY 
 
 decomposition of copper oxide into its elements. Much more 
 electrical energy is absorbed in the second case than in the 
 first; it is used up in effecting the chemical change in the 
 electrolyte. 
 
 This is shown by the equations : 
 
 The decomposition of water by the current (p. 27) when 
 regarded from this point of view, is really effected by the 
 electrolysis of a solution of sulphuric acid : 
 
 This is supported by the fact that pure water will not convey 
 the current. 
 
 The chief electrolytes are solutions of mineral acids and 
 metallic salts, molten salts and oxides, or solutions of oxides 
 in molten salts. 
 
 Voltaic Battery. The simple cell described in Exp. 65 
 falls off in strength very rapidly, owing to the accumulation of 
 a film of hydrogen on the copper plate. This offers consider- 
 able resistance to the passage of the current through the cell, 
 and also sets up a back pressure which diminishes the main 
 current. This is avoided in some constant cells by oxidizing 
 the hydrogen to water, and thus preventing its deposition on 
 the plate. The Bunsen cell is made up as follows : A circular 
 earthenware vessel containing dilute sulphuric acid, in which 
 is immersed a sheet of amalgamated zinc bent round into the 
 form of a hollow cylinder ; a porous earthenware pot contain- 
 ing strong nitric acid, in which is immersed a carbon rod. 
 The porous pot stands inside the zinc cylinder, and binding 
 screws are attached to the zinc and carbon plates. The liquids 
 soak into the wall of the porous pot from both sides, and a 
 conducting medium extends between the two plates. When 
 the binding screws are joined by a conductor, a much stronger 
 and more constant current than can be obtained by a simple 
 cell passes through the wire. The function of the nitric acid is 
 
METALS AND ACIDS 113 
 
 to prevent the deposition of hydrogen on the carbon plate, by 
 oxidizing it to water. By linking several of these cells to- 
 gether in series, a current sufficiently powerful for ordinary 
 experiments is obtained. 
 
 SUMMARY. 
 
 When a metal dissolves in hydrochloric or dilute sulphuric 
 acid, the reaction is simple, for hydrogen only is liberated, and 
 a salt of the metal and acid formed. The volume of hydrogen 
 set free depends upon the weight of the metal used, and is 
 different for the same weight of different metals. With hot 
 strong sulphuric acid the reaction is more complex, and 
 sulphur dioxide is the principal gas liberated. 
 
 Nitric acid dissolves most of the common metals, and gives 
 some very complicated reactions. The products of its decom- 
 position are varied, several oxides of nitrogen, hydrogen, 
 ammonia, and nitrogen itself being liberated under different 
 conditions. The nitrates of the metals are formed in most 
 cases, but with tin and antimony the oxides of the metals 
 are produced. The nitrates are readily decomposed by heat, 
 and, except in one or two cases, a residue of the oxides is 
 obtained. 
 
 The experiments on the oxidation of metals in air (Chap. II.) 
 and by nitric acid give interesting comparative results. If a 
 fixed weight of oxygen is taken, and the weights of different 
 metals which would combine with this fixed weight are cal- 
 culated from the results of experiments, an interesting table 
 can be constructed, as follows : 
 
 Combining weight of oxygen = 8. 
 
 METAL. 
 
 PARTS BY WEIGHT. 
 
 METAL. 
 
 PARTS BY WEIGHT. 
 
 Magnesium . . . 
 Iron ... 
 Tin 
 Copper 
 
 12 
 18-7 
 30-0 
 31-5 
 
 j Zinc ... 
 Bismuth 
 Mercury 
 | Lead ... 
 
 ... ! 32-5 
 ... i 69-0 
 lOO'O 
 103-5 
 
METALS AND ACIDS 115 
 
 The numbers are readily obtained, for in Exp. 69 : 
 
 0*251 gram of oxygen combines with 1 gram of copper 
 
 /. 1 , would combine with 
 
 .-. 8 grams 
 
 or copper 
 
 QUESTIONS. 
 
 1. Describe an experiment by which the volume of hydrogen 
 liberated by the dissolution of 1 gram of iron in hydrochloric 
 acid can be determined. 
 
 2. What changes take place (a) when zinc dissolves in 
 dilute sulphuric acid ; (b) when copper and the strong acid are 
 heated together ? 
 
 3. How would you find the weight of magnesium sulphate 
 formed when O5 gram of the metal dissolves in dilute sulphuric 
 acid? 
 
 4. Give a short description of the general reactions between 
 metals and nitric acid. 
 
 5. Name the three common oxides of nitrogen, and briefly 
 describe them. 
 
 6. What are the general effects of dilution and temperature 
 on the changes taking place between nitric acid and a given 
 metal 1 
 
 7. Which common acid is the best general solvent for 
 metals and alloys, and why ? 
 
 8. Describe a simple voltaic cell, and say how you would 
 obtain an electric current with it. 
 
 9. W T hat do you understand by the term "electrolysis"? 
 Give an example to illustrate your answer. 
 
 82 
 
CHAPTER VIII 
 
 CHEMICAL EQUIVALENTS AND ATOMIC WEIGHTS OF 
 COMMON METALS AND NON-METALS 
 
 THE student is now familiar with the fact that elements 
 combine together to form compounds, and is also aware of 
 the very definite character of these compounds, for he has 
 brought about various combinations, and has seen the extra- 
 ordinary changes effected in the properties of the elements 
 between which these combinations have taken place. Also, in 
 carrying out the experiments on the dissolution of metals in 
 acids, he has been directed to the conclusion that there is a 
 definite relation between the weight of metal dissolved and 
 the volume of gas liberated during its dissolution. It is also 
 fully recognised that the weight of a given body of gas is just 
 as definite as its volume. This is assumed in Chap. VIL, 
 together with the known weight of a litre of hydrogen, which 
 is used to calculate the weights of the volumes of gas evolved 
 in the experiments described ; but it is possible to use a more 
 direct method, if the pure gas only is allowed to escape from 
 the apparatus, and the loss in weight due to this escape is 
 noted. 
 
 The object of the following experiments is to make it clear 
 that certain proportional numbers or weights are inseparably 
 associated with the elements when they are entering into com- 
 bination with each other, or are being expelled from their 
 compounds ; and that these numbers, which are the results of 
 actual experiments, are connected in a simple way with the 
 atomic weights of the elements to which they refer. The 
 
 116 
 
EQUIVALENTS AND ATOMIC WEIGHTS 
 
 117 
 
 results we obtain may be only approximate ; but if they are 
 sufficiently accurate to give a clear idea of the facts they are 
 intended to illustrate, their object will be attained. Experi- 
 ments carried out by experienced workers using accurate 
 apparatus fully justify the deductions drawn from them, which 
 may therefore be accepted with confidence. 
 
 To FIND THE WEIGHT OF HYDROGEN EXPELLED FROM AN 
 
 ACID BY A KNOWN WEIGHT OF MAGNESIUM. 
 
 EXP. 66. Fit up the apparatus shown in Fig. 27. A is a 100 c.c. 
 flask fitted with a two-hole bung. A straight tube, B, fitted with 
 a cap to make it air tight, passes through one hole, and a bent tube, 
 C, through the other. The other end of C passes through a bung 
 in the U-tube, D. The bend of the U-tube 
 is filled with glass-wool, and each limb 
 with small pieces of fused calcium chloride, 
 and glass-wool on the top. The free limb 
 of the U-tube is fitted with a bung and 
 capped tube. The actual apparatus should 
 weigh about 70 grams. The caps are easily 
 made by pushing short pieces of glass rod 
 into pieces of rubber tube about twice their 
 length. 
 
 Add 10 c.c. of dilute H 2 S0 4 (1 in 12) to 
 the flask A, wipe the apparatus, and let it 
 stand for a time near the balance. Weigh 
 accurately about 0'2 gram of Mg ribbon, 
 and make it into a coil, which will pass 
 readily down the neck of A, and stand just 
 clear of the bung when dropped in. Put the apparatus and weighed 
 magnesium on the scale-pan, and weigh them together. Remove 
 the cap from the tube D. Take the bung out of A, drop in the 
 metal coil, and quickly replace the bung. As the metal dissolves 
 the coil gradually descends until complete dissolution is effected. 
 By this arrangement the evolution of gas is rendered uniform and 
 not too rapid. 
 
 Place the apparatus near the balance until it is nearly cold, and 
 then slowly aspirate about a litre of dry air through it. The air 
 is dried by passing through a U-shaped tube fitted in the same 
 manner as D, and attached to B by a rubber tube. Eeplace the 
 caps, and reweigh the apparatus. The loss in weight is due to the 
 escape of dry hydrogen. 
 
 The aspiration of air through the apparatus is of consider- 
 
 FIG. 27. 
 
118 
 
 METALLURGICAL CHEMISTRY 
 
 able importance, as a large portion of the liberated hydrogen 
 remains in the apparatus when the dissolution is finished. 
 Any form of aspirator may be used. In Fig. 28 the inlet of 
 the flask is connected with the drying-tube B, and the outlet 
 of the drying-tube, through which the gas escapes, with the 
 
 FIG. 28. 
 
 aspirator A. The tap C is opened carefully, and about 1 
 litre of water allowed to slowly escape. The hydrogen jis thus 
 replaced by air, and the atmosphere in the apparatus is now 
 similar to what it was at the beginning of the experiment. 
 
 EXAMPLE. Weight of magnesium taken = 0*201 gram. 
 
 ,, apparatus + magnesium =79-296 grams. 
 -hydrogen = 79279 
 
 ,, ,, hydrogen liberated = 0'017 gram. 
 
 Simplify this result by calculating the weight of magnesium 
 required to liberate 1 part by weight of hydrogen : 
 
 0-017 gram of hydrogen is replaced by 0*201 of magnesium. 
 
 o^rr 
 
 Or, 11 '82 parts of magnesium are equivalent to 1 part of 
 hydrogen. 
 
EQUIVALENTS AND ATOMIC WEIGHTS 119 
 
 To FIND THE WEIGHT OF HYDROGEN EXPELLED FROM AN 
 
 ACID BY A KNOWN WEIGHT OF ZlNC. 
 
 EXP. 67. Repeat the last experiment with about 0'75 gram of 
 thin zinc-foil and 10 c.c. of dilute sulphuric acid (1 in 6). The 
 stronger acid is recommended, as the action is not so violent as with 
 the magnesium. 
 
 EXAMPLE. Weight of zinc taken = 0'743 gram. 
 
 ,, ,, apparatus + zinc = 77*688 grams. 
 
 -hydrogen= 77'665 
 
 ,, hydrogen liberated = 0'023 gram. 
 Then, as before : 
 
 O023 gram of hydrogen is replaced by 0*743 of zinc. 
 
 0-743 00 o r 
 ' = 32-3 of zinc. 
 
 Or, 32 '3 parts of zinc are equivalent to 1 part of hydrogen. 
 
 These results may be used as starting-points in the deter- 
 mination of the equivalent weights of other elements. 
 
 To FIND THE WEIGHT OF OXYGEN WHICH WILL COMBINE 
 
 WITH A KNOWN WEIGHT OF MAGNESIUM. 
 
 EXP. 68. Heat a large porcelain crucible and lid (No. 2) over the 
 Bunsen flame, and set it aside to cool. Take a piece of magnesium 
 ribbon 18 inches long and weighing about 0'25 gram ; coil it round 
 a piece of thin glass rod, and cut the long coil into six short ones. 
 Weigh the crucible and lid, put in the coils, and weigh again. Put 
 the crucible over a good Bunsen flame on a pipeclay triangle, and 
 slope the lid a little so as to give a narrow air space on one 
 side. Heat the crucible strongly for fifteen minutes without raising 
 the lid. Then raise the lid carefully, and notice if there are any 
 signs of burning. If not, remove the flame, and break up the 
 residue with a clean glass rod, being careful to brush back into the 
 crucible any particles which may adhere to the rod. Continue the 
 heating without the lid for fifteen minutes longer. Allow the 
 crucible to cool, and weigh it. To be quite certain that the action is 
 finished, the crucible should be heated again for a few minutes, 
 cooled^ and reweighed to see if the weight is constant. 
 
120 METALLURGICAL CHEMISTKY 
 
 EXAMPLE. Weight of crucible + magnesium = 33-789 grams. 
 
 -33-544 
 
 = 0-245 gram. 
 . ,, after heating = 33-951 grams. 
 
 before 33'789 
 
 Increase = 0'162 gram. 
 
 The increase is due to oxygen absorbed from the air. 
 Take 12 the nearest whole number to the equivalent value 
 found for magnesium, and find the combining number for 
 oxygen : 
 
 0*245gramof magnesium combines with 0" 162 of oxygen. 
 
 0-162 
 
 
 ^245 
 
 ' 12 , ^^ = 7-93 grams. 
 
 Then 8 may be taken for the combining proportion of 
 oxygen. 
 
 To FIND THE WEIGHT OF COPPER WHICH WILL COMBINE 
 
 WITH A GIVEN WEIGHT OF OXYGEN. 
 
 EXP. 69. Heat a No. 1 porcelain crucible over the Bunsen flame. 
 Allow it to cool ; weigh it, and weigh in it 1 gram of finely-divided 
 copper reduced from the oxide. Put the crucible in a fire-clay 
 dish, and place the dish in a moderately hot munie. Continue the 
 heating for forty-five minutes ; remove the crucible ; cool, and 
 weigh it. Repeat the heating for ten minutes, and reweigh to see if 
 the weight is constant. 
 
 Grams. 
 
 EXAMPLE. Weight of crucible + copper after heating 12-851 
 
 before , 12-600 
 
 Increase = 0-251 
 Weight of copper taken = 1 gram. 
 
 Now, 0*251 gram of oxygen combines with 1 gram of copper ; 
 
 Q 
 
 .-. 8 grams of oxygen combine with - =31 -88 of copper. 
 
 0'251 
 
 31-5 may be taken as the combining proportion of copper. 
 
EQUIVALENTS AND ATOMIC WEIGHTS 121 
 
 NOTE. The reduced copper for the experiment is readily 
 prepared by reducing powdered copper scale by coal gas or 
 hydrogen. The common commercial oxide is too impure to 
 give a good result. 
 
 To FIND THE WEIGHT OF COPPER EXPELLED FROM ITS 
 COMPOUNDS BY A KNOWN WEIGHT OF ZINC. 
 
 EXP. 70. Dissolve about 3 grams of crystallized copper sulphate, 
 CuS0 4 , 5H 2 0, in 50 c.c. of water ; filter the solution, if necessary, 
 into a small beaker. Weigh accurately about 0*5 gram of thin zinc- 
 foil, and transfer it to the sulphate solution. Heat the solution 
 gently, and stir occasionally with a glass rod until the zinc is com- 
 pletely dissolved. This is determined by breaking up the precipitated 
 copper with the end of the glass rod. While the dissolution is 
 proceeding fold two filter-papers and 
 put them in a filter-funnel in the air- 
 oven, or over the sand-bath, to dry. 
 When dry place one in each pan of the 
 balance, and counterpoise them, cutting 
 the apex off the heavier one and then 
 snipping the one or the other at the top 
 until they counterpoise. Then put the 
 complete filter into the one which has 
 had the apex cut off, and open it out so 
 as to make a single filter with four folds 
 of paper on each side. Put it into the Fio. 29. 
 
 funnel, and when ready filter the sul- 
 phate solution. Transfer the whole of the precipitated copper to 
 the filter, and wash it thoroughly until every trace of the excess of 
 copper sulphate has been washed through. Use a wash-bottle in 
 which the water can be heated, as the washing is effected more 
 rapidly with hot than with cold water. Place the filter funnel and 
 filter in a cone over the sand-bath or in the air-oven. Be careful 
 not to overheat the filter, or the finely- divided copper will absorb 
 oxygen, and too high a result will be obtained. When the copper is 
 thoroughly dry, separate the two filters and put them on the scale- 
 pans in the same order as when they were counterpoised ; add 
 weights until they counterpoise again. The difference gives the 
 weight of the copper. Eepeat the drying for a few minutes to see if 
 the weight is constant. 
 
 By using two filters as described loss of copper is avoided, and 
 by putting them both in the funnel any loss of weight due to wash- 
 ing is the same in both. 
 
122 METALLURGICAL CHEMISTRY 
 
 EXAMPLE. Weight of zinc taken = 0-528 gram. 
 
 ,, ,, copper obtained = 0'509 
 
 Then using the equivalent for zinc found by Exp. 67, calculate 
 the equivalent for copper. 
 
 0"528 gram of zinc displaced 0-509 gram of copper ; 
 
 0-509 
 " ^528 
 
 0-509 x 32-3 
 .\32-3 -_ =31-47 of copper. 
 
 U DZO 
 
 Compare the results of the last two experiments. 
 
 To FIND THE WEIGHT OF SILVER EXPELLED FROM ITS 
 COMPOUNDS BY A KNOWN WEIGHT OF COPPER. 
 
 EXP. 71. Dissolve about a gram of thin sheet silver with 5 c.c. 
 of strong sulphuric acid in a small beaker. Put the beaker on a 
 sand-bath in the fume-chamber, and heat it carefully until the 
 metal is dissolved. Allow it to get cold, and then pour down the 
 side about 40 c.c. of distilled water. A white crystalline precipitate 
 of silver sulphate forms, which dissolves again on boiling the 
 solution. Weigh accurately about 0-25 gram of thin copper-foil, put 
 it into the hot silver solution, and allow to stand until the copper 
 is completely dissolved. The dissolution may be hastened by 
 keeping the solution hot, but it is better to let it stand in the cold. 
 Counterpoise two filter-papers as described in Exp. 70. Filter 
 off the precipitated silver, thoroughly wash and dry it. Weigh, 
 dry again, and reweigh until the weights are constant. The filtrate 
 contains silver, which can be recovered by immersing a strip of 
 copper in it. 
 
 EXAMPLE. Weight of copper taken = 0'251 
 ,, ,, silver obtained =0'860 
 
 Then, taking 31 "5 as the equivalent for copper : 
 
 0-251 gram of copper displaced 0-860 gram of silver ; 
 
 0-860 
 
 .-. 31-5 grams = 10>71 of silver. 
 
 108 may be taken as the equivalent of silver. 
 
EQUIVALENTS AND ATOMIC WEIGHTS 123 
 
 TO FIND THE WEIGHT OF CHLORINE WHICH WILL COMBINE 
 WITH A KNOWN WEIGHT OF SILVER. 
 
 EXP. 72. Weigh accurately about 0*5 gram of fine silver; 
 transfer it to a small beaker, and add 5 c.c. of dilute nitric acid 
 (1 to 1) ; heat the beaker gently on the hot plate until the metal is 
 dissolved ; make up the solution to about 100 c.c. with distilled 
 water ; raise it to boiling, and add a little dilute hydrochloric acid ; 
 stir well with a glass rod until the white curdy precipitate collects 
 together ; add a few drops more of the acid to see if the pre- 
 cipation is complete. Add the acid carefully, so as to have only 
 a slight excess in the solution when the whole of the silver is 
 precipitated. Dry and counterpoise two filter-papers as in Exp. 70. 
 Filter off the silver chloride, and wash it thoroughly with hot 
 water. Dry the chloride carefully, separate the filters, and weigh. 
 Kepeat the drying and weighing until a constant weight is 
 obtained. 
 
 EXAMPLE. Weight of silver chloride obtained = 0'67 5 gram. 
 
 taken=0-509 
 ,, chlorine found = 0-166 ,, 
 
 0*509 gram of silver combines with 0-166 gram of chlorine ; 
 
 0-166 
 " " " " 0-509 " 
 
 0-166x108 f 85-22 gra ms 
 )8 grams Q . 509 -| of ch i orine . 
 
 35-5 may be taken as the equivalent of chlorine. 
 
 If potassium bromide and potassium iodide solutions are 
 used to precipitate known weights of silver, in the same way 
 as described in the last experiment, the combining proportions 
 of bromine and iodine can be determined as readily as that of 
 chlorine. 
 
 If known weights of sodium chloride and potassium bromide 
 are precipitated with excess of silver nitrate, and the precipi- 
 tates collected and weighed, the combining proportions of 
 sodium and potassium, compared with chlorine and bromine 
 respectively, are obtained. 
 
124 METALLUKGICAL CHEMISTRY 
 
 SUMMARY OF RESULTS. 
 
 The experiments by which the results given in the previous 
 examples were obtained were made with special care, so as to 
 furnish reliable standards for the student to work to. There- 
 fore, the beginner must not expect too much from his own 
 initial experiments, and should be satisfied if his results are 
 sufficiently near the true ones to give him confidence in the 
 general principles which are under investigation, and in the 
 results which are stated as the outcome of many most accurate 
 investigations. 
 
 It should be borne in mind that if two or more distinct 
 methods are used to determine the chemical equivalent of a 
 particular element, the results should be practically the same. 
 This is illustrated in the determination of the equivalent of 
 copper : (a) by its combination with oxygen ; (b) by its pre- 
 cipitation by zinc. 
 
 The common starting-point is hydrogen, because it has the 
 lowest equivalent of any known element. It therefore fur- 
 nishes a very simple standard ; but any other element might 
 be used. Berzelius, who was the first to furnish a well- 
 defined list of the equivalents or combining proportions of 
 elements, adopted = 100 for his standard. But it is readily 
 seen that such a standard would be somewhat cumbersome in 
 its application. 
 
 The chemical equivalents or combining proportions of 
 elements are experimental numbers, and do not depend upon 
 any theory of the constitution of compounds, other than that 
 they contain elements in chemical combination. No difficulty 
 would, however, be experienced in assigning symbols to these 
 combining proportions, and in representing compounds with 
 them ; but formulae of compounds derived in this way would 
 not agree in many cases with those found by the application 
 of the theory of atoms. It is therefore necessary to consider 
 the question from another point of view, and endeavour to 
 trace a relation between chemical equivalents and atomic 
 weights. 
 
EQUIVALENTS AND ATOMIC WEIGHTS 
 
 125 
 
 EELATION OF CHEMICAL EQUIVALENTS TO ATOMIC WEIGHTS. 
 
 ELEMENT. 
 
 SYMBOL. 
 
 i 
 
 EQUIVALENT. 
 
 ii 
 
 ATOMIC WEIGHT. 
 
 RATIO OF i : ii 
 
 Hydrogen 
 
 H 
 
 1-0 
 
 i-o 
 
 1 
 
 Chlorine 
 
 Cl 
 
 35-22 
 
 35-5 
 
 1 
 
 Silver 
 
 Ag 
 
 10771 
 
 1077 
 
 1 
 
 Oxygen 
 
 
 
 7-93 
 
 16-0 
 
 2 
 
 Magnesium 
 
 Mg 
 
 11-82 
 
 24-0 
 
 2 
 
 Copper 
 
 Cu 
 
 31-47 
 
 63-0 
 
 2 
 
 Zinc ... 
 
 Zn 
 
 32-3 
 
 65-0 
 
 2 
 
 Bismuth 
 
 Bi 
 
 69-21 
 
 207-5 
 
 3 
 
 Tin* ... 
 
 Sn 
 
 29-4 
 
 117-4 
 
 4 
 
 Atomic Weights. Certain numbers, termed atomic 
 weights, have been used in connection with elements in pre- 
 ceding chapters, and a reference to the table given above 
 indicates that the combining proportion of an element as deter- 
 mined by experiment does not always agree with its atomic 
 weight as commonly used. Although the relation between the 
 two is a simple one, it requires some explanation if the beginner 
 is to use atomic weights with the confidence they deserve. It 
 may be stated that the determination of the combining pro- 
 portion of an element is always the first step in the fixing of 
 its atomic weight, and that the other considerations are also 
 based on experimental grounds. What these considerations 
 are depends upon the nature of the element and its compounds ; 
 so that a method which is very useful in dealing with some 
 elements may entirely fail with, others. Often two or more 
 methods may be applied to the same element, and if their 
 results agree, the deductions drawn from them are all the 
 more convincing. The method of deducing an atomic weight 
 may depend upon either the properties of the element itself 
 or on those of its compounds. The important methods used 
 for fixing the atomic weights of elements when their combining 
 proportions are known will now be briefly, but sufficiently, 
 described, to enable the student to get a general grasp of the 
 
 . * See Exp. 56. 
 
126 METALLUEGICAL CHEMISTRY 
 
 subject ; but for fuller details larger works on chemistry must 
 be consulted. 
 
 Vapour Density. A number of elements and compounds 
 are either gaseous at the temperature of the atmosphere, or 
 are converted into gas when raised to temperatures within the 
 range of accurate experimental work. The weights of 
 measured volumes of these gases or vapours can be determined, 
 and compared with the weights of equal volumes of hydrogen. 
 This gives the vapour density of the element or compound in 
 the gaseous state. 
 
 There are several methods employed for the determination 
 of vapour density ; but a brief outline of a comparatively 
 simple one will be sufficient for our present purpose. Suppose 
 the compound under consideration to be a liquid which is 
 easily converted into a combustible vapour. A glass flask 
 with its neck drawn off to a narrow tube is weighed, and an 
 excess of the liquid introduced. The flask is then immersed 
 in a bath, the temperature of which is well above the boiling- 
 point of the liquid, when a considerable volume of the vapour 
 escapes and carries with it the air which originally filled the 
 flask. When the air is all expelled, and the liquid is com- 
 pletely converted into vapour, the flask is full of it at the 
 temperature of the bath and the pressure of the outside atmo- 
 sphere. If the jet of vapour is allowed to burn as it issues 
 from the neck, the end of the vaporization is indicated by a 
 sudden drop in the flame. The end of the neck is now sealed 
 up by softening the glass in the blow-pipe flame, the flask 
 removed from the bath, and weighed when cold. The end of 
 the neck is then broken under water, the flask entirely filled, 
 and weighed again full of water. The weight of the water in 
 grams gives the volume of the flask, for a gram of water 
 measures a cubic centimetre. The weights obtained, together 
 with the corrections to reduce the volume of the vapour at 
 the temperature and pressure of the experiment to its volume 
 at N.T.P. (see Appendix), and for the expansion of the flask, 
 
EQUIVALENTS AND ATOMIC WEIGHTS 127 
 
 give the weight and volume of the vapour under examination. 
 It is then easy to compare this weight with the weight of an 
 equal volume of hydrogen under the same conditions of 
 temperature and pressure. It may, then, be stated that the 
 vapour density of an element or a compound which can be 
 converted into vapour without decomposition or dissociation 
 is an experimental number ; and full advantage may be taken 
 of the fact in helping to fix atomic weights. 
 
 If it is assumed that the bulk of 1 part by weight of 
 hydrogen represents the volume of an atom of the element, it 
 may be chosen as the unit of comparison for other atomic and 
 molecular volumes. Twice this bulk of hydrogen represents 
 a molecule of the gas, and, by Avogadro's law (see p. 43), 
 any other molecule in the gaseous state fills the same volume. 
 From the definition of the term atom, a molecule of any kind 
 cannot contain less than an atom of an element. 
 
 So that the smallest weight of an element ever 
 found in a molecular volume of the vapour of any of 
 its compounds may be taken as the atomic weight of 
 that element. 
 
 Further, it appears that vapour densities furnish a direct 
 method of determining molecular weights ; for since it is 
 assumed that a molecule of hydrogen contains two atoms, its 
 weight is twice that of a single atom, so that the weight of a 
 molecule of any other gas must be proportional to twice its 
 vapour density. If, then, the atomic weights of the elements 
 in a compound are known, and its vapour density can be 
 determined, a definite molecular weight is readily assigned to 
 it. For example, the smallest weight of oxygen ever found 
 in two volumes of any of its volatile compounds is 16; the 
 vapour density of water is 9; and, therefore, its molecular 
 weight is 18 ; so that a molecule of water contains an atomic 
 weight of oxygen, and 2 atomic weights of hydrogen and its 
 formula is therefore H 9 0. 
 
 The majority of the metals cannot be vapourized at workable 
 
128 METALLURGICAL CHEMISTRY 
 
 temperatures, nor do they furnish volatile compounds, so that 
 the application of the above principle to the fixing of their 
 atomic weights is limited. The problem must in their case be 
 solved by other means, as described below. 
 
 Isomorphism. Metals can be arranged more or less 
 definitely into groups, such that the members of any one 
 group can replace each other in solid compounds without alter- 
 ing the crystalline form of the compounds. These metals and 
 their crystallized compounds are said to be isomorphous, and 
 their atoms are regarded as fulfilling similar functions in the 
 molecules of the compounds. If, then, the atomic weight of 
 one metal in a group has been fixed by any other method, the 
 atomic weights of the other members of the same group can 
 be deduced. Thus the chemical equivalent of aluminium is 9, 
 and it forms only one oxide, in which the proportion Al : is 
 18 : 16 ; if it is inferred that the formula of this oxide is A10, 
 the atomic weight of the metal must be taken Al = 18. But 
 from other considerations the formula of the red oxide of iron 
 is given as Fe 2 O 3 , and the oxide of aluminium is isomorphous 
 with it, so that the more probable formula is A1 2 3 , and in 
 this case the atomic weight of the metal is Al = 27. (See also 
 "Double Salts and Isomorphism" in Chap. IX.) 
 
 Atomic Heat. The quantity of heat which a given weight 
 of a metal absorbs while its temperature rises through a given 
 range, is defined as the specific heat- of the metal. The 
 quantity of heat required to raise the same weight of water 
 through the same range of temperature is taken as the unit. 
 Now, it is found that the quantity of heat absorbed by a solid 
 metal is inversely proportional to its atomic weight. Or, if 
 weights of several metals proportional to their atomic weights 
 are taken, the same quantity of heat is required to raise each 
 of these weights through the same temperature. From this 
 it appears that the specific heat of a metal multiplied by its 
 atomic weight furnishes a quantity which is the same for all 
 metals. 6-2 is a mean value for this quantity, which may be 
 
EQUIVALENTS AND ATOMIC WEIGHTS 129 
 
 defined as the atomic heat of the metals, or as the specific 
 heat of their atomic weights. This is summarized in the 
 equation : 
 
 Specific heat x atomic weight = 6 -2 
 
 6*2 
 
 Therefore atomic weight = =. - 
 
 Specific heat 
 
 Thus, when the specific heat of a metal has been determined, 
 this principle, which is known as the law of Dulong and 
 Petit, must afford considerable help in fixing its atomic 
 weight. 
 
 The solid non-metals are usually included in the general 
 statement, but for carbon, boron, and silicon the atomic heat 
 is abnormal, arid varies for the different allotropic modifications 
 of these elements. If, however, their specific heats are deter- 
 mined at high temperatures, when their molecules probably 
 have a simpler structure, a much closer approximation to the 
 mean atomic value is obtained. So that the principle may 
 probably be said to include all solid elements. 
 
 Chemical Considerations. The formula for water was 
 at first written HO, and is correct in so far as it represents 
 the equivalent weights of the two elements in the compound 
 i.e., 1 : 8. But do these symbols represent atoms according 
 to the present view of the atomic hypothesis 1 One of the 
 first principles of the atomic theory is that when a chemical 
 change takes place in the properties of a substance, it is the 
 individual molecules which undergo the change, and this is 
 true for both elements and compounds. In Exp. 16 it is 
 proved that when sodium acts upon water, hydrogen gas 
 is liberated. Where does this hydrogen come from? From 
 the molecules of the water is the only satisfactory explanation. 
 Now, it is found that water disappears in the same proportion 
 as hydrogen appears, and, if a given weight of sodium is used, 
 an equivalent weight of hydrogen is obtained. Also, if the 
 residual liquid is evaporated to get rid of the excess of water, 
 a white solid residue is obtained. This is the common 
 
 9 
 
130 METALLURGICAL CHEMISTRY 
 
 compound known as caustic soda, which, since nothing but 
 hydrogen and excess of water have escaped, must contain the 
 remnants of the water molecules, together with the sodium 
 used to bring about the change. Now, if this solid is heated 
 with more sodium, a further quantity of hydrogen is obtained, 
 and this, too, must come from the remnants of the water 
 molecules associated with the first quantity of the metal. 
 The residue from this change is a solid when cold, and is 
 quite free from hydrogen. From this it appears that hydrogen 
 can be expelled from the water molecules in two distinct 
 portions. Other metals will displace this second proportion 
 of hydrogen from caustic soda, and their action can be more 
 readily controlled. 
 
 EXP. 73. Weigh 2 grams of caustic soda, and put it together 
 with 20 c.c. of water into the flask A (Fig. 21). Add 25 gram of 
 aluminium foil, and gently heat the mixture on the hot gauze. Test 
 the gas collected in C. The characteristic explosion of a mixture of 
 air and hydrogen is obtained. Hydrogen is displaced from the 
 caustic soda by the aluminium, and a double oxide of aluminium 
 and sodium is formed. 
 
 The inference to be drawn from these experiments is, that 
 since hydrogen can be expelled from the water molecules in 
 two distinct portions, there must be at least two atoms of 
 hydrogen in each molecule of water. No experiment has yet 
 been made in which the gas is displaced in more than two 
 portions. 
 
 Further, it is possible to displace oxygen from water, but 
 it all comes out at once. The simple explanation of this is 
 that there is only one atom of oxygen in the water molecule. 
 If, then, symbols are used to represent atoms, there is clearly 
 only one way to write the formula for water i.e., H 2 0. This 
 is also borne out by the results of Exp. 19, by which it is 
 proved that water contains hydrogen and oxygen in the 
 proportion of 2 to 1 by volume. For by Avogadro's law equal 
 volumes of gases under normal conditions contain the same 
 number of molecules. Therefore, the water molecules must 
 
EQUIVALENTS AND ATOMIC WEIGHTS 131 
 
 supply double the number of hydrogen atoms to make up 
 the double volume of that gas obtained from the water. 
 
 If, then, the formula H 2 is admitted, and the atomic 
 weights are based upon H = 1 for the standard, it is evident 
 that = 16, or the atomic weight of oxygen is double its 
 chemical equivalent. 
 
 The properties of water are so different from those of its 
 constituents that there must be a very intimate relation 
 between the atoms of oxygen and hydrogen in the molecules 
 of water. Some very powerful influence is at work holding 
 the atoms together, and this must reside in the atoms them- 
 selves, for the energy of combination cannot exist apart from 
 the constituents of the compound. It is not difficult, then, to 
 think of the atoms in a molecule as clinging together, merging 
 their individual properties into each other, and thus producing 
 a new particle, with well-defined properties of its own. If this 
 is so, an atom of oxygen must have twice the clinging power 
 of an atom of hydrogen, and this may be shown graphically 
 by writing the formula for water thus, H H. Several 
 names have been given to this important property of atoms, 
 and the best are given below. 
 
 Atom-combining 1 Power, Valency, and Atomicity. 
 
 When sodium pushes out part of the hydrogen from water, it 
 evidently displaces an atom from each molecule acted upon, 
 and an atom of sodium takes its place. This is shown by 
 writing the formula of caustic soda thus : Na H ; and it 
 is seen that the atom-combining power of sodium is the same 
 as that of hydrogen. This is shown again when the second 
 part of hydrogen is displaced by sodium, and the formula 
 becomes Na O Na. Again, magnesium or zinc pushes the 
 whole of the hydrogen out of water at one operation, and it 
 is not difficult to conceive that an atom of magnesium or zinc 
 enters each molecule to take the place of the two atoms of 
 hydrogen displaced. Hence the change is from H H to 
 Mg = or Zn = O, and the atom-combining power of these 
 
132 METALLURGICAL CHEMISTRY 
 
 metals is equal to that of oxygen. From this it seems that 
 the term atom-displacing power also expresses this property of 
 atoms. 
 
 When a mixture of hydrogen and chlorine is fired, rapid 
 combination takes place between the two gases, and they dis- 
 appear in exactly equal volumes to form hydrochloric acid 
 gas, the total volume of which is the same as that of the dis- 
 appearing gases. Now, it is practically certain that the 
 molecules of the elementary gases contain two atoms ; and by 
 Avogadro's law the equal volumes contain the same number 
 of molecules. Therefore, it may be considered that each pair 
 of hydrogen and chlorine molecules in the mixture divide up 
 and have their atoms rearranged. Thus : 
 
 | jy_H_| + j Cl |_ClJ = |jlJ_Cl I + j H | Cl | 
 
 2 vols. 2 vols. 2 vols. 2 vi Is. 
 
 This points to chlorine and hydrogen combining together 
 atom to atom. They have, therefore, the same combining 
 power. This information is useful, as it gives another standard 
 by which to measure the combining power of metals, since 
 they all unite with chlorine, and the relative number of atoms 
 in the molecules of their chlorides can be determined. 
 
 Bismuth unites with chlorine and forms a compound, the 
 formula of which may be written graphically, thus : 
 
 Cl Bi Cl, or BiCl 3 
 
 .Cl 
 Tin unites with chlorine to form the compound : 
 
 Cl 
 
 Cl Sn Cl, or SnCl 4 , stannic chloride. 
 
 It may be noticed here that tin also unites with two atoms 
 of chlorine to form the compound SnCl 2 , stannous chloride. 
 So that the total combining power of an elementary atom 
 
EQUIVALENTS AND ATOMIC WEIGHTS 
 
 133 
 
 need not be always used up in the formation of a molecule. 
 But the compound formed is " unsaturated," and can take up 
 a second proportion of the other element to form the higher 
 compound. 
 
 Reference to the table on p. 125 will now make it clear that 
 the number which expresses the ratio of the chemical equiva- 
 lent of an element to its atomic weight, also gives its atom- 
 combining power or valency. This part of the subject should 
 be carefully studied, and not put aside until the principles are 
 quite clear. 
 
 Atomic weight and valency are fundamental properties, and 
 are in constant use for determining the formulae by which 
 compounds are to be represented. Chap. IX. should be studied 
 in conjunction with this one, for a further development of the 
 subject of chemical formulae. The following table contains 
 the common elements arranged in groups according to their 
 valency, which is indicated by a number attached to the 
 typical element of each group. When an element appears in 
 more than one group it is an indication that its active valency 
 varies in different compounds. 
 
 TABLE OF ATOM-COMBINING POWER OR VALENCY OF 
 ELEMENTS. 
 
 MONADS. 
 
 DlADS. 
 
 TRIADS. 
 
 TETRADS. 
 
 PENTADS. 1 HEXADS. 
 
 H=l. 
 
 0-2. 
 
 O=2. 
 
 N = 3. 
 
 C = 4 
 
 P-ft. S-fi 
 
 Hydrogen 
 Chlorine 
 Bromine 
 Iodine 
 
 Oxygen 
 Carbon 
 Sulphur 
 Barium 
 
 Magnesium 
 Manganese 
 Mercury 
 Nickel 
 
 Nitrogen 
 Boron 
 Phosphorus 
 Antiinonv 
 
 Carbon Phosphorus 1 Sulphur 
 Sulphur Nitrogen Chromium 
 Silicon Antimony Cobalt 
 Aluminium Arsenic Iron 
 
 Fluorine 
 
 Calcium Platinum Arsenic 
 
 Chromium Bismuth 
 
 Platinum 
 
 Potassium 
 
 Strontium Iron 
 
 Bismuth 
 
 Iron 
 
 Tungsten 
 
 Sodium 
 Silver 
 
 Copper Lead 
 Cadmium i Tin 
 
 Gold 
 
 Lead 
 Platinum 
 
 
 
 
 Cobalt Zinc 
 
 
 Tin 
 
 
 
 
 Chromium 
 
 
 
 
 
 
 SUMMARY. 
 
 The chemical equivalents of the elements in general, or 
 their quantitative relations to each other, form the solid 
 
134 METALLUKGICAL CHEMISTKY 
 
 foundation upon which all systematic practical work is based. 
 Mere observation is useful for indicating the way, but it is 
 only when one begins to weigh and measure that the great 
 possibilities of the subject become apparent. The selection of 
 hydrogen as the standard weight renders comparison easy. 
 The chemical equivalent of an element is purely the result of 
 experiment, and is quite independent of any theory of the 
 constitution of matter. The atomic weight of an element, 
 on the other hand, assumes the existence of atoms, and is 
 dependent on the atomic theory for its definition ; but it is 
 none the less certain on that account, for it is simply the 
 chemical equivalent multiplied by a whole number. The 
 selection of this number for a particular element is also based 
 upon experimental methods, and the number itself indicates 
 the atom-combining power or valency of the element. The 
 great care which has been exercised in the determination of 
 atomic weights renders them perfectly reliable for use in the 
 most important calculations, and every confidence may be 
 placed in them. 
 
 QUESTIONS. 
 
 1. What do you understand by the chemical equivalent of 
 an element ? 
 
 2. Describe two methods by which the chemical equivalent 
 of copper can be determined. 
 
 3. What is meant by the valency of an element? Give 
 examples. 
 
 4. Describe experiments in which hydrogen is displaced from 
 water in two distinct operations. 
 
 5. What is the relation between the chemical equivalent of 
 an element and its atomic weight ? 
 
 6. Explain the difference between the atom-combining power 
 of an element and its combining weight. 
 
CHAPTER IX 
 OXIDES, ACIDS, AND SALTS 
 
 ALL the metals combine, either directly or indirectly, with 
 oxygen to form oxides; but those oxides which cannot be 
 formed by direct oxidation are unstable bodies of little practical 
 importance. The common metals which do not oxidize, when 
 exposed to air or oxygen under suitable conditions, are silver, 
 gold, and platinum. 
 
 A few of the common oxides unite directly with water to 
 
 form compounds, called hydroxides or hydrates. These 
 
 definite bodies contain the elements of the oxides and of 
 water. Thus, sodium oxide, Na 2 0, and potassium oxide, 
 K 2 0, form very stable compounds with water, which are not 
 decomposed at a bright red heat. This kind of change is 
 readily shown : 
 
 Na 2 + H 2 = 2NaHO. 
 
 Sodium 
 hydroxide 
 
 Calcium oxide (quicklime), GaO, and the corresponding oxides 
 of barium and strontium, combine readily with water, and a 
 considerable amount of heat is developed during the reaction. 
 Thus: 
 
 CaO + H 2 0=Ca(HO) 2 . 
 
 Calcium 
 hydroxide 
 
 These hydroxides, however, are decomposed at a red heat, and 
 the water escapes, leaving the original oxides as a residue. 
 The change shown in the last equation is then reversed. The 
 evolution of heat is readily observed on moistening a piece of 
 quicklime with water. 
 
 135 
 
136 METALLURGICAL CHEMISTRY 
 
 Hydroxides which can be formed by the direct union of 
 oxides with water are more or less soluble in the liquid, and 
 their solutions have very characteristic properties. They 
 produce a soapy feeling when rubbed between the thumb and 
 finger, and the very soluble ones have a caustic or burning 
 action upon the skin. They all turn reddened litmus blue. 
 This is called an alkaline reaction, and the compounds pro- 
 ducing it are alkaline hydroxides. 
 
 The soapy feeling mentioned above is really due to the 
 formation of a soap, by a reaction between the alkali and the 
 fatty matter in the skin. The common alkalies, caustic soda 
 and caustic potash, are largely used, on account of this property, 
 for the removal of grease from various articles. The soap 
 formed is readily washed away, and the grease thus removed. 
 Hydroxides of metals, the oxides of which do not combine 
 directly with water, can be formed in an indirect way. They 
 are insoluble in water, and when a solution of a soluble 
 hydroxide is added to a solution of a salt of one of these 
 metals, the hydroxide of the metal is generally formed, and 
 precipitated from the solution. 
 
 EXP. 74. Put a crystal of copper sulphate into a test-tube, and 
 dissolve it in a little water ; add a solution of caustic soda, a little 
 at a time, as long as a precipitate forms ; shake well, and allow the 
 precipitate to settle. Repeat the experiment with a solution of 
 'ferric chloride in place of the copper sulphate. 
 
 The green and brownish-red precipitates are hydroxides, 
 and the changes are shown by the equations : 
 
 CuS0 4 + 2NaHO = Cu(HO) 2 + Na,S0 4 . 
 Fe 2 Cl 6 + GNaHO Fe 2 (HO) 6 + GNaCl. 
 
 Basic Oxides. Metallic oxides which form hydroxides, 
 either soluble or insoluble, corresponding to the above, are 
 called basic Oxides, and it is very probable that, if the in- 
 soluble ones could be dissolved in water, their solutions would 
 give an alkaline reaction. The following are examples of basic 
 oxides : 
 
 K 2 0, CaO, MgO, HgO, PbO, ZnO, A1 2 3 , Fe 2 3 . 
 
OXIDES, ACIDS, AND SALTS 137 
 
 Peroxides. But some metals form more than one oxide, 
 and it is then necessary to enquire if all the oxides of a parti- 
 cular metal are basic in character. 
 
 Compare PbO and Pb0 2 ; BaO and Ba0 2 ; MnO and Mn0 2 . 
 The second oxide in each pair gives off oxygen gas when it is 
 heated alone, or with dilute sulphuric acid ; and when heated 
 with hydrochloric acid causes the liberation of chlorine from 
 the acid. For example : 
 
 2Ba0 2 = 2BaO + 2 . 
 Mn0 2 + 4HC1 = MnCl 2 + 2H 2 + C1 2 . 
 
 These higher oxides are not true basic compounds, and are 
 usually called peroxides. 
 
 Complex Oxides. Another class of oxides, of which red 
 lead, Pb 3 O 4 , and magnetic oxide of iron, Fe 3 4 , are examples, 
 are sometimes regarded as being a combination of two oxides 
 of the same metal. Compare Pb 3 4 and 2PbO.Pb0 2 ; Fe 3 O 4 
 and FeO.Fe Q . They are not true basic oxides. 
 
 Z o / 
 
 Acid-forming" Oxides. A few metallic oxides are even 
 more highly oxidized than the peroxides. A well-known 
 example of such compounds is the highest oxide of chromium, 
 Cr0 3 . This oxide when in solution has distinctly acid pro- 
 perties, and is altogether different from the basic oxides of the 
 same metal. It is an acid-forming oxide. 
 
 Metallic oxides may, then, be divided into four classes : 
 (1) basic oxides; (2) peroxides; (3) complex oxides; (4) acid- 
 forming oxides. 
 
 Oxides Of Non-Metals. The most powerful acid-forming 
 oxides are to be found among the oxides of the non-metals. 
 They unite with water to form hydroxides, which are mostly 
 soluble in water, and form solutions having very characteristic 
 properties. The solution of an acid hydroxide has a SOUP 
 taste, and turns litmus * red. This is the characteristic acid 
 
 * Litmus is a vegetable colouring matter. It dissolves readily in water, 
 and either the solution itself or porous paper stained with it may be used 
 for testing:. 
 
138 METALLUEGICAL CHEMISTKY 
 
 reaction, as distinguished from the alkaline or basic reaction, 
 in which reddened litmus is turned blue. The formation of 
 an acid hydroxide is represented thus : 
 
 24 . 
 
 Sulphur Sulphuric 
 
 trioxide acid 
 
 A few of the non-metals form peroxides, and some of the non- 
 metallic oxides are neutral in character. 
 
 Nomenclature Of Oxides. When there is only one oxide 
 of a particular element no difficulty is experienced in giving 
 it a name, as MgO, magnesium oxide ; ZnO, zinc oxide. But, 
 when there are two or more oxides of the same element, dis- 
 tinguishing names must be used. Thus, copper forms two 
 compounds, commonly called the red and black oxides, from 
 their respective colours. The systematic names are cuprOUS 
 oxide, Cu 2 O, and cuprie oxide, CuO. Similarly, we have 
 ferrous and ferric oxides, FeO and Fe 2 3 . The termination 
 -OUS indicates the oxide containing the smaller proportion, 
 and the termination -ic the larger proportion, of oxygen. The 
 prefix sesqui- is also used to denote the ratio 2 to 3. Sesqui- 
 oxide of iron is ferric oxide, Fe 2 3 . 
 
 The same terminations are used to distinguish hydroxides of 
 the same element from each other. There are ferrous and 
 ferric hydroxides; and the two important acid hydroxides 
 of sulphur are sulphurous and sulphuric acids, H 2 S0 3 and 
 H 2 S0 4 . 
 
 Another useful way of naming is to put a numerical prefix 
 before the word " oxide," and also before the name of the other 
 element, if necessary. The prefixes indicate the number of 
 atoms of oxygen or of the other element in the formula. 
 
 Thus sulphur dioxide, S0 2 ; sulphur trioxide, S0 3 ; and 
 triferric tetroxide, Fe 3 4 . Other examples are carbon mon- 
 oxide, CO ; carbon dioxide, C0 2 ; chromium trioxide, Cr0 3 ; 
 nitrogen tetroxide, N 2 O 4 ; phosphorus pentoxide, P 2 5 . 
 
 Acid-forming oxides are often called anhydrides, as 
 sulphuric anhydride, S0 3 . 
 
OXIDES, ACIDS, AND SALTS 139 
 
 Many compounds have common names, with which the 
 student will become familiar, and, with a little practice, he 
 will have no difficulty in giving the systematic name to any 
 compound, when he knows its formula. 
 
 Basic and acid-forming oxides are the most important, and 
 it may be stated definitely with regard to their distinctive 
 properties that, if there are two basic oxides of the same 
 metal, the lower oxide is the stronger base. Thus, FeO is 
 a stronger base than Fe 2 3 . On the other hand, of two acid- 
 forming oxides of the same element the higher oxide forms 
 the Stronger acid. Thus, sulphuric acid is more powerful 
 than sulphurous acid. With some of the metals, both acid- 
 forming and basic oxides are obtained. This is so with the 
 oxides of chromium : 
 
 Chromous oxide, CrO ; chromic oxide, Cr 2 3 ; 
 
 Strongly basic Basic 
 
 chromium trioxide, CrO 3 . 
 
 Acid -forming 
 
 REACTIONS BETWEEN BASIC OXIDES AND ACIDS. 
 
 The acid hydroxides are such well-defined compounds that 
 they were, for a long time, considered to be the typical acids, 
 and the oxygen they contain absolutely essential to the com- 
 position of an acid. But when it was conclusively proved 
 that hydrochloric acid contains no oxygen, this view had to 
 be modified. Now, clear proof has been furnished that, when 
 some of the metals are presented to the common acids, they 
 displace hydrogen and take its place, forming salts of the 
 acids. So that it would appear that the essential element in 
 an acid is the hydrogen which can be displaced from it by a 
 metal. Now, only a comparatively small number of metals 
 will displace hydrogen from acids in general, but a large number 
 of basic oxides and hydroxides will react with most acids and 
 form salts of the acids used. Thus copper does not dissolve 
 in either dilute sulphuric or hydrochloric acid, but cupric 
 
140 METALLURGICAL CHEMISTRY 
 
 oxide does dissolve, and forms the copper salts of the acids. 
 Thus : 
 
 1 . CuO + 2HC1 = CuCl 2 + H O. 
 
 2. 
 
 It is seen from the equations that the hydrogen of the acid is 
 replaced by the metal of the oxide, and that the displaced 
 hydrogen unites with the oxygen of the oxide to form water. 
 This is the general reaction. 
 
 The experiments which follow are intended to make the 
 beginner practically acquainted with the commoner reactions 
 between oxides and acids. 
 
 EXP. 75. Weigh roughly 3 grams of caustic potash, put it into 
 a porcelain basin with 10 c.c. of water, and stir the liquid with a 
 glass rod until the solid is dissolved. Test the solution with apiece 
 of red litmus-paper, and leave the paper sticking to the side of the 
 dish, so that a slight movement will bring the liquid into contact 
 with it. Now add dilute nitric acid (1 to 1) a little at a time, stirring 
 with a glass rod, until the blue colour of the litmus-paper is just 
 turned red when the solution is brought into contact with it. 
 Towards the end the acid must be added drop by drop. The first 
 portion of the acid may be run in from a test-tube, and the remainder 
 from a piece of glass tube pulled off to a jet. The dropping-tube is 
 put into the test-tube containing the acid, and when some of the 
 liquid has entered, the top of the tube is closed by the finger, and 
 the tube itself lifted out. By moving the finger the acid is allowed 
 to escape drop by drop. When the action is finished, remove the 
 litmus-paper, evaporate the solution down to about 15 c.c., and set 
 it aside to crystallize. Examine the crystals, and compare them 
 with the saltpetre crystals in the laboratory bottle. 
 
 In this experiment it is clear that the basic properties of 
 the alkali and the acid properties of the acid have entirely 
 disappeared, along with the bodies themselves, and a new body 
 has been formed. The reaction is expressed thus : 
 
 KHO + HN0 3 = KN0 3 + H 2 0. 
 
 Basic Acid Neutral Water 
 
 hydroxide hydroxide salt 
 
 This equation not only shows the general change, but also 
 expresses these quantitative relations between the reacting 
 
OXIDES, ACIDS, AND SALTS 141 
 
 compounds. These relations are just as definite as those which 
 exist between the elements in an oxide, or between the amount 
 of a dissolved metal and the gas liberated during its dissolution. 
 This is easily proved by using solutions of acids and alkalies 
 containing known weights of the reacting compounds, and 
 finding the volumes of the liquids required to be added together 
 to bring about neutralization. 
 
 Solutions containing 10 grams per litre of each of the 
 following pure compounds are easily prepared : sulphuric, 
 nitric, and hydrochloric acids, caustic soda and caustic 
 potash. Directions for the preparation of one of them is 
 given on p. 143, but they should be ready prepared for the 
 student's use. 
 
 The following exercises may be made successful with 
 moderate care, and will give a good insight into the quantita- 
 tive relations which exist during neutralization. Litmus has 
 been tised so far as an indicator of the acid or alkaline character 
 of a solution ; but, for the work now to be described, methyl 
 Orange* is to be preferred, as it is not affected by carbon 
 dioxide in solution. It is turned greenish-yellow by an alkali, 
 but changes to a bright red on the addition of a very slight 
 excess of an acid. 
 
 EXP. 76. Wash out the burette (Ch. xv.) with a little of the 
 standard acid solution, fix it in the clip, and fill it up to the mark 
 with the same solution. Transfer 10 c.c. of the caustic soda solution 
 to a conical flask, by means of a 10 c.c. pipette, add about 100 c.c. 
 of water to the flask, and one or two drops of methyl orange. Note 
 the level of the acid solution in the burette, and then run a little at 
 a time into the flask, shaking it between each addition, until the 
 yellow-green colour of the solution changes to red. Care must be 
 exercised towards the end of the reaction, as one drop of the acid 
 solution is sufficient to change the colour. Bead off the volume of 
 acid solution added. Eepeat the experiment with a second 10 c.c. 
 of the soda solution, and compare the two results. 
 
 An exactly similar experiment may then be made with 
 caustic potash solution. 
 
 * The compound known as Porrier's Orange No. 3 is the best ; 1 gram 
 per litre makes a good solution. 
 
142 METALLURGICAL CHEMISTRY 
 
 EXAMPLE. 10 c.c. of caustic soda solution required 12 '3 c.c. of 
 sulphuric acid solution, and 10 c.c. of caustic potash solution 
 required 8'7 c.c. of sulphuric acid solution. 
 
 Since each solution is known to contain 10 grams of the 
 compound in 1,000 c.c , O'Ol gram is contained in 1 c.c. 
 
 .'.01 gram of caustic soda neutralizes 123 gram of sul- 
 phuric acid, and 0-1 gram of caustic potash 0087 of the 
 acid. 
 
 These weights may be compared directly with the 
 equations : 
 
 2NaHO + H 2 S0 4 = Na 2 S0 4 + 2H 2 0. 
 
 Sodium 
 sulphate 
 
 2KHO + H 2 S0 4 = K 2 S0 4 + 2H 2 O. 
 
 Potassium 
 sulphate 
 
 Q.I 03 
 
 Also ,r^r- = 1*41 = the ratio of the weights of acid required 
 0*087 
 
 to neutralize the same weight of the alkalies. 
 
 Experiments with nitric acid and with hydrochloric acid 
 may be made in the same way, and the results compared. 
 
 EXAMPLE. 10 c.c. of caustic soda solution required 15'6 c.c. of 
 nitric acid solution, and 10 c.c. of caustic potash solution required 
 11 '2 c.c. of nitric acid solution. 
 
 NaHO + HN0 3 = NaNo 3 + H 2 0. 
 
 Sodium 
 nitrate 
 
 Also B =1-39 = ratio of weights of nitric acid required to 
 
 neutralize the same weight of the alkalies. 
 
 EXAMPLE. 10 c.c. of the caustic soda solution required 9'1 c.c. 
 of hydrochloric acid solution, and 10 c.c. of the caustic potash 
 solution required 6'5 c.c. of hydrochloric acid solution. 
 
 KHO + HC1 = KC1 + H 2 0. 
 
 Potassium 
 chloride 
 
 A.Q1 
 
 The ratio is 4-l-41. 
 
OXIDES, ACIDS, AND SALTS 143 
 
 It is evident from the above results that, if equal weights 
 of two bases are used to neutralize different acids, the ratio 
 of the weights of the acids neutralized is constant. This fact 
 was recognised as early as the latter part of the eighteenth 
 century, and work done in this direction by Kichter and others 
 helped to lay the foundation of the first law of chemical com- 
 bination, which is stated on p. 47. 
 
 Preparation of Solutions. In making an acid or an 
 alkaline solution of known strength, it is convenient to add 
 more than the required quantity of the compound, and after 
 determining the actual weight present in a given volume, to 
 bring the whole to the proper strength by the addition of a 
 calculated volume of water. The following proportions may 
 be used, and a large or small volume of the solution prepared 
 as required. 
 
 Concentrated sulphuric acid 7 c.c. 
 nitric acid 12 , 
 
 , . , , . . , __ , 800 c.c. of water 
 
 hydrochloric acid 25 ,, v , 
 
 in each case. 
 Caustic soda 10 grams 
 
 potash 10 J 
 
 The strength of an acid solution is usually determined by 
 taking advantage of its reaction with pure sodium carbonate. 
 With sulphuric acid this is : 
 
 Na 2 C0 3 + H 2 S0 4 
 46 + 12 + 48 2 + 32 + 64 
 
 106 98 
 
 The carbonate during its decomposition neutralizes the acid, 
 and by using the methyl orange indicator, the exact amount 
 of acid to be added is readily controlled. From the equation 
 it is clear that 106 parts by weight of Na 2 C0 3 neutralize 98 
 
 1 Ofi 
 
 parts by weight of H 2 S0 4 ; .*. =1-082 = the weight of 
 
 t/o 
 
 carbonate required to neutralize 1 part by weight of the acid ; 
 and 0'2164 gram is required for 0'2 gram of the acid. 
 
144 METALLUKGICAL CHEMISTEY 
 
 If, then, 0-2164 gram of pure sodium carbonate is dissolved 
 in about 100 c.c. of water, one drop of methyl orange added, 
 and the sulphuric acid solution run in from a burette until 
 the colour change takes place, the volume of the acid solution 
 used will contain exactly 0'2 gram of H 2 S0 4 . The whole 
 volume of the remaining solution may then be measured, and 
 the necessary volume of water added to bring it up to the 
 proper strength. 
 
 If V = the volume of the remaining acid solution, and v = 
 the volume used to neutralize 0'2164 gram of carbonate, 
 
 y 
 
 then (20 v) = volume of water to be added. 
 
 EXAMPLE. Volume of remaining solution = 750 c.c. ; volume 
 used to neutralize =17 c.c. 
 
 /. 7 1 y(20--17) = 75 !j^ = 132 c.c. of water to be added. 
 
 The strengths of the other acid solutions can be determined 
 in a similar manner by using the equations : 
 
 :Na 2 C0 3 + 2HN0 3 = 2NaN0 3 + C0 2 + H 2 0. 
 Na 2 C0 3 + 2HC1 = 2NaCl + C0 2 + H 2 0. 
 
 The atomic weights may be taken from the table in the 
 appendix. The strength of the alkaline solutions are then 
 readily determined with one of the acid solutions. Twenty c.c. 
 of the soda solution are transferred to a flask, diluted with 
 water, and one drop of methyl orange added. The standard 
 acid is then run in until the colour changes. For hydro- 
 chloric acid and caustic soda the reaction is expressed by the 
 equation : 
 
 NaHO + HC1 = NaCl + H 2 0. 
 23 + 1 + 16 1+35-5 
 
 ~~40~~ ^3(^T 
 1 gram of HC1= ^ = 1-096 gram of NaHO. 
 
OXIDES, ACIDS, AND SALTS 145 
 
 Then the number of cubic centimetres of acid required multi- 
 plied by 0-01096 gives the weight of NaHO in 20 c.c. of the 
 solution, and a simple calculation determines the volume of 
 water to be added to the alkaline solution to bring it down to 
 1 c.c. =0 - 01 gram of NaHO. Since each solution contains 
 10 grams of the pure compounds in 1,000 c.c., the bottles 
 may be labelled as follows : 
 
 Sulphuric acid, 1 c.c. = 0'01 gram H 2 S0 4 , etc. 
 
 NEUTRALIZATION BY INSOLUBLE OXIDES. 
 
 The majority of the oxides insoluble in water are dissolved 
 more or less readily by the common acids, and the correspond- 
 ing salts are formed. If, however, the salt to be formed is 
 insoluble in the acid solution, dissolution does not take place, 
 and the action between the oxide and the acid is somewhat 
 slow, but, as a rule, the salt is formed on continued digestion 
 with the acid. 
 
 Some oxides which will withstand a high temperature with- 
 out fusing dissolve very slowly after they have been strongly 
 ignited. This is probably due to a closer aggregation of the 
 molecules, by which they become more complex in character 
 and more difficult to decompose. This increase in molecular 
 complexity is known as polymerization. Strongly ignited ferric 
 and chromic oxides are examples. 
 
 Magnesia EXP. 77. Put 20 c.c. of dilute hydrochloric acid 
 (1 to 1) into a porcelain basin, and test it with blue litmus-paper. 
 Then add magnesia a little at a time ; stir well with a glass rod 
 between each addition, and continue the addition until some of the 
 powder is left undissolved. Test the solution with the litmus-paper, 
 and, if it turns red, boil the solution slowly for a minute or two to 
 see if the residue is dissolved. Filter the liquid to clear it, evaporate 
 it down to half the bulk, and set it aside to crystallize. 
 
 The crystals are usually small and not very distinct. If the 
 solution is evaporated carefully to dryness, and the dry mass 
 exposed to the air, it absorbs moisture, and becomes quite 
 moist. It is a deliquescent salt. The magnesium oxide, MgO, 
 
 10 
 
146 METALLURGICAL CHEMISTRY 
 
 dissolves in the acid liquid, and during its dissolution neutral- 
 izes the acid. It, therefore, establishes its claim to be a basic 
 
 oxide. 
 
 MgO + 2HC1 P. MgCl 2 + H 2 0. 
 
 Magnesium 
 chloride 
 
 A similar experiment may be made with sulphuric acid and 
 magnesia, and the magnesium sulphate formed is readily 
 crystallized. Its formula is MgS0 4 ,7H 2 0. 
 
 Copper Oxide EXP. 78. Put 2 grams of black oxide of copper 
 into a large test-tube, and add 15 c.c. of dilute sulphuric acid (1 to 6). 
 Boil the solution gently until the whole of the oxide has dissolved 
 or until the residue becomes distinctly red. Commercial black 
 oxide of copper usually contains a little of the red oxide, which is 
 not completely dissolved by sulphuric acid, and a residue of red 
 metallic copper is left. Filter the solution into a porcelain dish, 
 and, if it is deep blue in colour, set it; aside to crystallize ; but if 
 only light blue, evaporate it to about half the bulk before setting it 
 aside. 
 
 CuO + H 2 S0 4 = CuS0 4 + H 2 0. 
 
 The change with the red oxide is shown thus : 
 Cu 2 + H 2 S0 4 = CuS0 4 + H 2 + Cu. 
 
 The blue crystals (blue vitriol) have the composition CuS0 4 , 
 5H 2 0. 
 
 Zinc Oxide. EXP. 79. Repeat the last experiment, using zinc 
 white, ZnO, in place of black oxide of copper. Also, dissolve 2 grams 
 of zinc in dilute sulphuric acid, and set the two solutions aside to 
 crystallize. 
 
 On comparison the crystals are found to be exactly the same. 
 Zinc sulphate (white vitriol), ZnS0 4 ,7H 2 0, is formed in each 
 case. Compare the equations : 
 
 ZnO + H 2 S0 4 = ZnS0 4 + H 2 O. 
 
 Oxide of Lead. EXP. 80. Put 5 grams of lead oxide into a 
 test-tube, add 20 c.c. of dilute nitric acid (1 to 5), and heat the 
 tube gently until the oxide is dissolved. Filter the solution into 
 an evaporating basin, and set it aside to crystallize. Examine the 
 crystals. 
 
 PbO + 2HNO, = Pb(N0 3 ) 2 + H 2 0. 
 
 Lead nitrate 
 
OXIDES, ACIDS, AND SALTS 147 
 
 The lead nitrate crystals contain no water of crystallization, 
 and the formula of the crystalline salt is Pb(N0 3 ) 2 . Lead 
 sulphate, PbS0 4 , is insoluble in dilute sulphuric acid, but on 
 boiling the oxide for a few minutes with the moderately strong 
 acid, an appreciable quantity of the sulphate is formed. Hot 
 hydrochloric acid also dissolves lead oxide with formation of 
 the chloride, but the salt largely separates from its solution on 
 cooling. 
 
 EXP. 81. Boil a little lead oxide with moderately strong sulphuric 
 acid (1 to 3) for a few minutes ; allow the test-tube to stand until 
 the liquid becomes clear ; note the layer of white sulphate upon the 
 yellow residue of unchanged oxide. Repeat the experiment with 
 dilute hydrochloric acid (1 to 1). When the lead oxide is all dis- 
 solved set the solution aside to cool, and note the rapid separation 
 of lead chloride crystals. 
 
 Water of Crystallization. Potassium, sodium, and 
 
 lead nitrates crystallize from their solutions just as they are 
 formed ; but zinc, copper, and magnesium sulphates, when 
 they crystallize, carry out with them definite quantities of 
 water. This is called water of crystallization, and is associated 
 with the salts in a very definite manner. It may be taken 
 as a general statement that when a given salt crystallizes 
 under normal conditions it always takes out with it the same 
 proportion of water. The composition of a crystalline salt is 
 constant, but the relation between the salt and its water of 
 crystallization is not so intimate as that between the elements 
 of the salt itself. It is not so intimate even as that between 
 an oxide and the elements of water in its hydroxide. This is 
 shown in the formula by keeping the water separate. Thus : 
 crystallized copper sulphate is CuS0 4 ,5H 2 0. 
 
 In general, crystalline salts are freely soluble in water, but 
 there is a limit, differing with different salts, to the quantity 
 of the solid dissolved by a given weight of water. The solu- 
 bility of a salt, or other solid, is usually stated in terms of 
 parts by weight of cold water required to dissolve 1 part by 
 weight of the solid. Thus potassium nitrate dissolves in 
 
 102 
 
148 METALLURGICAL CHEMISTRY 
 
 7 parts, zinc sulphate in 2 -5 parts, and potassium chlorate 
 in 20 parts of cold water. 
 
 In most cases the salt is much more soluble in hot than in 
 cold water, and the cold saturation point is thus easily ex- 
 ceeded. As the solution cools the excess above that which 
 the cold solution will hold tends to separate in the solid 
 state i.e., to crystallize. But it is possible for the solutions 
 to become quite cold without crystallization setting in. Such 
 a solution is super-saturated, and will crystallize rapidly 
 when once the action is started. A small crystal of the same 
 substance dropped in is usually sufficient to set up the action, 
 and even a speck of dust may form a nucleus for the initial 
 crystals to form round. Generally, if the solution is above 
 the cold saturation point, and cools slowly, the crystallization 
 commences slowly, and the final crystals are larger and better 
 developed. If the solution is too strong the crystallization 
 will be rapid and the crystals very small, forming a powder 
 when dry. When once crystals are formed in the solution 
 they continue to grow, and the dissolved solid is reduced 
 somewhat below the point of cold saturation ; but it is im- 
 possible to separate the whole of the solid from the solution, 
 for there is a limit at which the solid dissolves from the 
 crystals already formed as fast as it is deposited upon them. 
 The residual solution from which the crystals have separated 
 is called the " mother liquor." 
 
 It will be readily understood that if two solids differ in 
 solubility, it is possible to separate a mixture of them by 
 repeated crystallization, as the more soluble one will not 
 crystallize as rapidly as the less soluble one under the same 
 conditions. Also, a small quantity of one soluble body can 
 be separated from a large quantity of another. These facts 
 are taken advantage of in the purification of salts by repeated 
 crystallization. Some salts dissolve almost as readily in cold 
 as in hot water. That is, their solubility increases very little 
 with a rise in the temperature at which the solution is 
 effected. Common salt is an example. Others, again, seem to 
 
OXIDES, ACIDS, AND SALTS 149 
 
 become less soluble after a certain temperature is passed. 
 This is so with sodium sulphate, which, below 33 C., crystallizes 
 with 10 molecules of water, thus : Na 2 S0 4 ,10H 2 0. But when a 
 solution which would barely crystallize at ordinary temperatures 
 is raised above 33 C., small crystals of the anhydrous (water- 
 less) salt, Na 2 SO 4 , are deposited. From this it would seem 
 that above 33 C. the ordinary crystallized salt does not exist, 
 and when the point of saturation for Na 2 S0 4 is passed crystals 
 are formed. 
 
 But this is really a case of equilibrium between the solid and 
 the solvent liquid, and comes within the scope of the phase 
 rule enunciated by Willard Gibbs. For information on this 
 subject the student is referred to recent works on physical 
 chemistry. 
 
 Crystals. A fully-developed crystal is a solid of definite 
 geometrical shape, bounded generally by plane surfaces, which 
 join each other and form solid angles. These surfaces are the 
 faces of the crystals, and an angle formed by the intersection 
 of two or more of these faces is invariable for the same 
 crystalline form. It is possible for a number of different 
 substances to crystallize in the same form, so that a classifica- 
 tion of crystals can be made. Thus common salt crystallizes 
 in the form of cubes, in which all the angles are right angles, 
 and therefore equal. Iceland spar crystallizes in rhombs, in 
 which the opposite angles only are equal. Quite a number 
 of other bodies crystallize in the same forms. All the known 
 crystalline forms have been classified into six systems, and 
 the expert crystallographer is able, from the measurement of 
 its angles and the position of its faces, to relegate a particular 
 crystal to its proper system. But the subject is a wide one, 
 and the student must be satisfied at this stage with such 
 information as he can get by an examination of the crystals 
 he produces. 
 
 When crystallization takes place en masse, as in ordinary 
 experiments, the individual crystals interfere with each other, 
 
150 METALLUEGICAL CHEMISTRY 
 
 and imperfect forms are obtained ; but such parts as are well 
 defined are perfect in themselves, and serve to establish the 
 form. To obtain a perfect crystal of considerable size it is 
 necessary to pick out a small perfect one, and then nurse it in 
 a saturated solution of the same salt. 
 
 Insoluble Salts. Some salts are practically insoluble in 
 water, and there is a very easy way of forming them by 
 taking advantage of this property. A solution of a soluble 
 salt of the metal is simply mixed with a solution of the acid, 
 or a solution of a soluble salt of -the acid and another metal. 
 The possibility of the formation of the insoluble salt is thus 
 assured, and it immediately commences to separate from the 
 solution as a precipitate, which may be crystalline, curdy, or 
 granular. 
 
 EXP. 82. Pour a few cubic centimetres of a solution of silver 
 nitrate into a clean test-tube and add to it dilute hydrochloric acid, 
 a little at a time, with vigorous shaking between each addition, as 
 long as a precipitate separates. Repeat the experiment, using a 
 solution of common salt in place of the hydrochloric acid. Filter, 
 wash and dry the solid salt. 
 
 The white curdy solid is silver chloride, and its formation 
 is expressed by the equations : 
 
 AgN0 3 + HC1 = AgCl + HN0 3 . 
 AgN0 3 + NaCl = AgCl + NaN0 3 . 
 
 Silver chloride is not soluble in nitric acid. 
 
 EXP. 83. Repeat the last experiment, using a solution of barium 
 chloride and dilute sulphuric acid, or a solution of potassium sulphate 
 in place of the acid. 
 
 The white, finely-divided solid is barium sulphate, and it is 
 insoluble in hydrochloric or nitric acid : 
 
 BaCl 2 + H 2 S0 4 = BaS0 4 + 2HC1. 
 BaCl 2 + K 2 S0 4 = BaS0 4 + 2KC1. 
 
 These reactions are used as tests for hydrochloric and sulphuric 
 acids and their soluble salts. 
 
OXIDES, ACIDS, AND SALTS 151 
 
 Direct Union of Elements to form Salts. Salts of 
 hydrochloric acid may be formed by the direct union of the 
 metals with chlorine. This has been already illustrated in the 
 experiments with chlorine. Thus, when chlorine gas is passed 
 over finely-divided silver, the chloride of the metal is formed : 
 
 2Ag + Cl 2 = 2AgCl. 
 
 Bromine and iodine act similarly, but not so readily, as 
 chlorine. 
 
 Direct Union of Oxides to form Salts. In the case of 
 some oxyacids it is possible to form their salts by the direct 
 union of the acid-forming and the basic oxides. 
 
 Thus: PbO + SO s = PbS0 4 . 
 
 Lead 
 sulphate. 
 
 = CaC0 3 . 
 
 Calcium 
 carbonate. 
 
 = Pb 2 Si0 4 . 
 
 Lead silicate. 
 
 These reactions indicate that there cannot be any strong 
 objection to writing the formulae of the compounds in such 
 a way as to show the combination of the two oxides. For 
 example: PbO.S0 3 ; CaO.C0 2 ; 2PbO.Si0 2 . Sometimes it 
 is convenient to use such formulae, especially in the case of 
 silicates. See Chap. XIV. 
 
 The decomposition of a carbonate by an acid, in which the 
 salt of the acid and of the metal in the carbonate is formed, 
 furnishes another general method of forming salts. Illustra- 
 tions of this method will be found in Chap. X. The following 
 is a summary of the chief modes of salt formation : (1) Direct 
 union of elements ; (2) direct union of acid-forming and basic 
 oxides ; (3) replacement of hydrogen in an acid by a metal ; 
 (4) reaction between acids and bases; (5) formation of in- 
 soluble salts by precipitation ; (6) decomposition of carbonates 
 by acids. 
 
 Nomenclature Of Salts. In the systematic naming of 
 
152 METALLUKGICAL CHEMISTRY 
 
 salts, terminations are used, as in the case of oxides. The 
 assumption that acids are also salts of the displaceable hydrogen 
 they contain, and are, therefore, of the same general character 
 as the compounds formed by the replacement of such hydrogen 
 by metals, is very useful. 
 
 According to this view hydrochloric acid, HC1, is also hydro- 
 gen chloride, and common salt, sodium chloride. Similarly 
 we have bromides, iodides, etc. Thus, when the acid and its 
 salts contain two elements only, the termination -ide is used 
 to designate the group to which the salt belongs, and the 
 name of the metal to particularize the salt. Potassium 
 bromide, KBr; zinc iodide, ZnI 2 ; bismuth chloride, BiCl 3 , 
 are examples. Nitric acid, HN0 3 , is also hydrogen nitrate, 
 and saltpetre is potassium nitrate. 
 
 But there is also another acid, HN0 2 , nitrous acid or 
 hydrogen nitrite, and its salts are nitrites e.g., potassium 
 nitrite, KNO 2 . 
 
 Thus when there are two acids containing the same elements, 
 but in different proportion, the terminations -OUS and -ite are 
 used for the lower acid and its salts, and -ic and -ate for the 
 higher compounds. Similarly sulphurous acid gives sulphites, 
 and sulphuric acid, sulphates. When there is only one oxy- 
 acid of a given element the higher terminations are used ; 
 e.g., carbonic acid gives carbonates. The student is advised 
 to keep the idea of acids as salts of hydrogen well in mind, as 
 it will assist him in getting a firm grip of the composition and 
 formulae of salts. 
 
 Constitution Of Acids. The molecule of an acid may be 
 looked upon as consisting of two parts, which are represented 
 graphically thus : (H)(C1), (H)(N0 8 ), (H 2 )(S0 4 ). Now, the 
 second part of the acid molecule is present in the molecule of 
 any salt of the acid, and is characteristic of it ; so that the 
 (Cl), (N0 3 ), and (S0 4 ) groups are characteristic of chlorides, 
 nitrates, and sulphates generally. The same may be said of 
 other acids and their salts. 
 
OXIDES, ACIDS, AND SALTS 153 
 
 There is good reason to believe that when an acid or its 
 salt is dissolved in a very large quantity of water, its molecules 
 split up or dissociate into the parts mentioned above. These 
 parts are called ions ; and even when the solution is moderately 
 strong, an average number of molecules are believed to be 
 constantly in this state of dissociation, or ionic condition. 
 This matter is closely studied in electro-chemistry, and is men- 
 tioned here in order to justify the statements made in the last 
 paragraph. It has already been stated that some metals form 
 two basic oxides, and when these dissolve in a given acid two 
 distinct salts are formed, Thus there are two sulphates of 
 iron : ferrous sulphate, FeS0 4 , and ferric sulphate, Fe 2 (S0 4 ) 3 ; 
 also two chlorides of mercury, mercurous chloride, Hg 2 Cl 2 , 
 and mercuric chloride, HgCl 2 . But it must not be inferred 
 that when there are two basic oxides of a metal there are also 
 two salts of any acid. For example, the two oxides, Cu 2 
 and CuO, both give copper sulphate, CuS0 4 , when heated with 
 dilute sulphuric acid, but the first is decomposed and leaves a 
 deposit of copper. Thus : 
 
 Cu 2 + H 2 S0 4 = CuS0 4 + H 2 + Cu. 
 
 There are, however, two chlorides of copper, Cu 2 Cl 2 and 
 CuCl 2 . 
 
 Some old-fashioned distinctive names which are still used 
 may be mentioned here. Thus ferrous sulphate is also 
 the protosulphate of iron, and ferric sulphate the per- 
 sulphate of iron. Also subchloride of mercury is the same 
 as mercurous chloride, and the perchloride is the mercuric 
 salt. 
 
 Basicity Of Acids. The valency of metals has already 
 been defined as their atom-replacing power, and is measured 
 by the number of atoms of hydrogen replaced by one atom 
 of the metal. Now, if acids are salts of hydrogen, it is easy 
 to see from their formulae that the number of atoms of the 
 replaceable hydrogen differs in different acids, and that the 
 
154 METALLURGICAL CHEMISTRY 
 
 property known as the basicity of the acid depends on this 
 number. 
 
 Nitric acid, HN0 3 is mono basic. 
 
 Sulphuric acid, H 2 S0 4 is di 1 
 
 Phosphoric acid, H 3 P0 4 is tri [-Poly basic. 
 
 Silicic acid, H 4 Si0 4 is tetr J 
 
 With a monobasic acid the hydrogen is all replaced in one 
 reaction, and the normal salt only is formed. But with a 
 polybasic acid the hydrogen may be replaced in two or more 
 distinct reactions, thus giving rise to two or more distinct 
 salts. Those in which hydrogen is present are called acid 
 salts. 
 
 Potassium hydrogen sulphate, KHS0 4 , acid salt. 
 
 Potassium sulphate K 2 S0 4 , normal salt. 
 
 Sodium dihydrogen phosphate, NaH 2 P0 4 , acid salt. 
 
 Disodium hydrogen phosphate, Na 2 HP0 4 , acid salt. 
 
 Trisodium phosphate, Na 3 P0 4 , normal salt. 
 
 The second salt of phosphoric acid above mentioned is the 
 common sodium phosphate. Sometimes salts are formed in 
 which an excess of the acid-forming oxide is present. Such 
 compounds are anhydro-add salts. The well-known example 
 is potassium bichromate, K 2 O 2 7 , which may also be written 
 K 2 Cr0 4 Cr0 3 . Potassium chromate, K 2 Cr0 4 , is the normal 
 salt. 
 
 Similarly some salts contain an abnormal quantity of the 
 basic oxide or hydroxide, and are then called basic salts. 
 The green mineral malachite is a basic carbonate of copper, 
 CuC0 3 .CuH 2 2 . 
 
 Only one acid of carbon has been mentioned, but a large 
 number of acids are found in organized matter, both animal 
 and vegetable, or are formed during organic changes. These 
 invariably contain carbon and hydrogen, but the whole of 
 the hydrogen is not as a rule replaceable in the same manner 
 as that of a mineral acid. This is well illustrated in the case 
 of acetic acid and its salts. The empirical formula of this 
 
OXIDES, ACIDS, AND SALTS 155 
 
 acid is H 4 C 2 2 , but to show its basicity it must be written 
 H.C 2 H 3 2 , for it contains only one atom of replaceable hydro- 
 gen which is written first. Thus the formula for lead acetate 
 (sugar of lead) is Pb(C 2 H 3 2 ) 2 . On the other hand, oxalic 
 acid, H 2 C 2 4 , is dibasic. The salts of these acids are termed 
 respectively acetates and oxalates. 
 
 Double Salts. It is found that when some soluble salts 
 are dissolved together in approximately molecular proportions 
 they also crystallize together in strictly molecular proportions 
 when the solution is sufficiently concentrated. In fact, so 
 constant is the composition of these double compounds, that 
 they are regarded as definite salts. The alums are the most 
 common and the best illustrations of these very interesting 
 bodies. Common alum is a double sulphate of potassium and 
 aluminium, and its empirical formula is KA1(S0 4 ) 2 ,12H 2 O. 
 The two sulphates are potassium sulphate, K 2 S0 4 , and 
 aluminic sulphate, A1_,(S0 4 ) 3 . So that the formula which 
 shows the composition best is obtained by doubling the above 
 thus, K 2 Al 2 (S0 4 ) 4 ,24HoO. Potassium sulphate is easily 
 crj^stallized alone, but the aluminium sulphate only with 
 difficulty, and the forms of the crystals are quite distinct. On 
 dissolving the salts together in molecular proportions, and 
 allowing the solution to crystallize, the double salt is obtained 
 in crystals of the regular octahedron form. 
 
 EXP. 84. Put 1 gram of alumina, Al. 2 3j into a small beaker, add 
 10 c.c. of dilute sulphuric acid (1 to 3), and boil slowly until the 
 solid disappears, and a somewhat milky-looking solution is obtained. 
 Dissolve 1 gram of caustic potash in 10 c.c. of water, add this solu- 
 tion to the first, raise the whole to boiling, and filter the solution 
 into a porcelain basin. Set it aside to crystallize. Examine the 
 crystals under a lens, and note their characteristic shape. Set aside 
 the best for reference. 
 
 The crystals are usually well defined, if the solution is not 
 too concentrated, and the crystallization takes place slowly. 
 If caustic soda is used in the above experiment in place of 
 potash, exactly similar crystals are obtained. The two salts 
 
156 METALLURGICAL CHEMISTRY 
 
 cannot be distinguished from each other by their crystalline 
 form. Also, if ferric oxide, Fe 2 3 , is substituted for alumina, 
 light green crystals are obtained, which cannot be distinguished 
 from the others by their crystalline form. 
 
 EXP. 85. Dissolve 5 grams of potassium bichromate in 25 c.c. of 
 water, and add 1 c.c. of strong sulphuric acid. Prepare sulphur 
 dioxide, and pass the washed gas (p. 95) through the solution until 
 it turns deep green in colour. Pour it into a large porcelain dish, 
 and set it aside to crystallize. Compare the crystals of chrome alum 
 with those of common alum from Exp. 84. 
 
 It will be noticed that the two metals in the sulphates are 
 monad and triad respectively ; any similar metal may be 
 substituted in the compound without altering its form. Thus 
 silver and the monad group ammonium (NH 4 ) may take the 
 place of potassium, and the triad metals, chromium and 
 manganese, that of aluminium. The general formula is 
 M'M'"(S0 4 ) 2 ,12H 2 where M' and M'" are monad and triad 
 atoms respectively. 
 
 If a crystal of common alum is put into a solution of iron 
 alum it increases in bulk ; the enlarged crystal may then be 
 transferred in turn to solutions of chrome alum and manganese 
 alum, in each of which there is a further growth, but the final 
 crystal keeps its fundamental form. Further, if a mixture of 
 various alums is crystallized, one crop of crystals containing 
 the different compounds is obtained. 
 
 Another class of double salts of practical importance contain 
 ammonium or potassium sulphate with iron, nickel, or zinc 
 sulphate. The well-known examples are ferrous ammonium 
 sulphate (NH 4 ) 2 Fe(S0 4 ) 2 ,6H 2 0, and nickel ammonium sul- 
 phate, (NH 4 ) 2 NiS0 4 ,6H 2 0. 
 
 Isomorphism. The similarity of form of different crystal- 
 lized compounds is looked upon as an indication of similarity 
 in the constitution of their molecules, and such compounds are 
 said to be isomorphous, which means, literally, of " equal 
 form." It is not confined to the double salts described above, 
 for two single salts crystallized separately may have the same 
 
OXIDES, ACIDS, AND SALTS 157 
 
 crystalline form. They are then judged to be made up of 
 similarly constituted molecules. 
 
 This supposed similarity of the molecules of isomorphous 
 compounds suggests similarity between the atoms of those 
 elements which are able to replace each other in them without 
 altering the crystalline form of the compounds ; and they are 
 also considered to be isomorphous. Thus, Na, K, Ag ; and Al'", 
 Fe'", Cr'" replace each other in the alums. They are regarded 
 as two groups of isomorphous metals. Similarly, Zn", Fe", Ni" 
 are isomorphous with regard to the class of isomorphous salts 
 of which (NH 4 ) 2 Ni(S0 4 ) 2 ,6H 2 is a well-known example. 
 This principle of isomorphism has been of considerable value 
 in helping to fix the atomic weights of a number of the 
 metals. It has also been used to explain some of the changes 
 which take place during the cooling and solidification of a 
 molten mass of mixed metals. A mixture of isomorphous 
 salts cannot be separated from each other by crystallization, 
 as they are deposited from the solution together ; but mixtures 
 of non-isomorphous compounds may be so separated if their 
 solubility in the mother liquor differs. Similar remarks 
 apply to isomorphous and non-isomorphous metals in a molten 
 alloy. 
 
 Action Of Heat upon Salts. Water of crystallization is 
 driven off, usually at temperatures very little above the boil- 
 ing-point, 100 C. But sometimes one or more molecules of 
 water hold more tenaciously to the salts, and require a higher 
 temperature for their separation. Further heating of the 
 anhydrous salt may cause it to melt and even to volatilize ; or 
 it may be decomposed into simpler compounds, and, in a few 
 cases, into its elements. 
 
 EXP. 86. Weigh exactly 2 grams of powdered crystallized copper 
 sulphate in a dry porcelain crucible, or on a clay roasting dish, and 
 heat it on a pipeclay triangle over a Bunsen flame until it turns 
 white. Allow the crucible to cool and reweigh it. Put the crucible 
 into a gas muffle, and keep it at a fair red heat for twenty minutes, 
 
158 METALLURGICAL CHEMISTRY 
 
 or until the residue is quite black. Allow the crucible to cool and 
 reweigh it. The residue is black oxide of copper. 
 
 EXAMPLE. Weight of crucible + salt =13140 < _ n 71ft 
 after first heating = 12 -422 j ~ 
 
 second = 11-776 = 0'646 
 of crucible =11-140 = 0'636 
 
 First loss = 0-718 gram H 2 0; second loss = 0'646 gram S0 3 ; 
 residue = 0'636 gram CuO, and these are the weights of the 
 three oxides in 2 grams of the compound. Now, divide each 
 weight by the molecular weight of the oxide to which it belongs. 
 CuO = 79; S0 3 = 80; H 2 = 18. 
 
 Then, - 0-0399; 0-008; = 0-008. 
 
 18 oU 79 
 
 These ratios show the relation between the molecules of the 
 oxides in the crystallized salt, and are readily reduced to 
 whole numbers by dividing each by the lowest. This gives 
 1 CuO : 1 S0 3 : 5 H^O, and the formula may be written, 
 CuO.S0 3 .5H 2 0. The complete reaction is expressed by the 
 equation : 
 
 CuS0 4 ,5H 2 = CuO + S0 8 + 5H 2 0. 
 
 Four molecules of water are separated at a temperature about 
 100 C. ; but the last molecule clings until the temperature 
 exceeds 200 C. A similar experiment may be made with 
 crystallized ferrous sulphate, FeS0 4 ,7H 2 0, which gives the 
 following reaction when heated : 
 
 2(FeS0 4 ,7H 2 0) = Fe 2 3 + S0 2 + S0 8 + 14H 2 O. 
 
 Most of the sulphates of the common heavy metals are decom- 
 posed in a similar manner, but a number of others, including 
 those of potassium, barium, lead, etc., may be heated very 
 strongly without undergoing decomposition. 
 
 The Nitrates are all more or less readily decomposed when 
 heated. Oxides of nitrogen and oxygen are set free, and a 
 residue of the oxide of the metal left. Nitrates of silver and 
 mercury are reduced to the metallic state. 
 
OXIDES, ACIDS, AND SALTS 
 
 159 
 
 The Chlorides fuse more or less readily, and volatilize when 
 strongly heated. Some are decomposed with liberation of 
 chlorine, and are thus reduced, either to a lower chloride or to 
 the metallic state. 
 
 The Carbonates of sodium and potassium will withstand 
 a high temperature without decomposition. The other common 
 carbonates are split up into metallic oxides and carbon dioxide. 
 Barium carbonate requires a bright red heat for its decom- 
 position. 
 
 A number of well-known acids are represented by only a 
 few common salts, but as these are of considerable importance 
 a list of them is given below. 
 
 ACID. 
 
 FORMULA. 
 
 SALT. 
 
 FORMULA. 
 
 COMMON NAME. 
 
 Sulphurous 
 Thio- 
 
 H 2 S0 3 
 H 2 SA 
 
 Sodium sulphite 
 Sodium thio- 
 
 Na 2 S0 3 ,10H 2 'Sulphite of soda 
 Na 2 S,0 3 .5H 2 Hyposulphite 
 
 sulphuric 
 
 
 sulphate 
 
 A 4 O & 1 
 
 of soda 
 
 Phosphoric 
 
 H 3 P0 4 
 
 Disodium hydro- 
 
 Na 2 HP0 4 ,12H 2 Phosphate of 
 
 
 
 gen phosphate 
 
 soda 
 
 Chloric 
 
 HC10 3 
 
 Potassium 
 
 KC10 3 Chlorate of 
 
 
 
 chlorate 
 
 potash 
 
 Chromic 
 
 H 2 Cr0 4 
 
 Potassium 
 
 K 2 Cr0 4 Chromate of 
 
 
 
 chromate 
 
 
 potash 
 
 
 
 Potassium 
 
 K 2 Cr 2 7 
 
 
 
 
 bichromate 
 
 
 
 Permanganic 
 
 H 2 Mn 2 8 
 
 Potassium per- 
 
 K 2 Mn 2 8 
 
 
 
 
 manganate 
 
 
 
 Metaboric 
 
 HB0 2 
 
 Acid sodium 
 
 Na 2 B 4 7 ,10H 2 \ Borax 
 
 
 
 borate 
 
 
 Nitrous 
 
 HN0 2 
 
 Potassium nitrite 
 
 KN0 2 
 
 Nitrite of 
 
 
 
 
 
 potash 
 
 Hydro- 
 
 HBr 
 
 Potassium 
 
 KBr 
 
 
 bromic 
 
 
 bromide 
 
 
 
 Hydriodic 
 
 HI 
 
 Potassium iodide 
 
 KI 
 
 
 Hydrofluoric 
 
 HF 
 
 Calcium fluoride 
 
 CaF 2 
 
 Fluor spar 
 
 Salts of these acids and the common metals are easily 
 obtained by one or other of the usual methods of salt 
 formation. 
 
 SUMMARY. 
 
 The basic oxides of the metals are a very important class of 
 compounds. They react with acids to form salts. In the 
 
160 
 
 METALLUEGICAL CHEMISTEY 
 
 o 
 
 88J3 
 
 S 
 
 cq 
 
 .2 6*d" 
 
 0,^00 
 
 HURIC 
 H 2 SO 
 
 DROCHLORIC A 
 HCI. 
 
 o 
 
 CO _5^ ^ -5- 
 
 T3 O O O O 
 
 OCOOQCOCC 
 
 r^^v^ ^ 
 
 0.00 
 
 o m qq, q, 
 
 5 - gg ff 5 
 
 402^5 ^ -5 s. 
 
 r~> "~ci^ O O 9 n 
 
 S^ 1 O?M 
 
 OJ^ S>R l M 
 
 
 
 c,- 
 
 .S3. 
 
 ^. * sM 
 
 I 
 
 . 
 
 
 
 
 p 
 
 3" 
 
 Sf 1' 
 
 ^ do 
 
 qp o 
 
 5 ob 
 
 ooo|o:o o 
 
 s * 
 
 
 
 ^J 
 
 &-S6S6SS SZ 
 
 s 
 
 
OXIDES, ACIDS, AND SALTS 
 
 161 
 
 _,SO 
 
 9o 
 
 O .1'' 
 
 W I 
 
 , 
 
 = Sw 
 
 
 3 J 
 
 QQOO CO QQ Q i"* 
 
 ll | |:: 
 
 1111 I.S.1 
 
 PH CH 02O2 OQ H N3 
 
 rST" PQ 
 
 . u-> j X 
 
 n ,"^j ^ 
 
 ^-3 5 I 
 
 r^T3 
 
 3 
 
 g S 
 
 -4-3 O 
 
 2 g 
 II 
 2 
 
 < 
 
 .a fe 
 
 2 ^ 
 S 2 
 
 ,-H ^ O 
 
 (D O3 
 
 3 5 
 
 ^ W^ 
 
 1^ 
 
 -2 .S 
 
 o S 
 61 
 
 g 
 
 -S co 
 
 I d 
 
 ,3 o-d 
 
 .S ^ 
 
 II 
 
 -j .2 o 
 1-3 o ^_ 
 S "'^l 
 
 -3 2 
 
 f- 1 _g .S C " ^ 
 
 2 g g'g g I 
 
 S W a3 
 
 i* 8 t- t 'T3 
 
 3 ^'3 
 
 I ss 
 
 g .S' 
 
 r-< 
 
 .S <D X> 
 
 a ^^ 
 
 S 1' 
 
 f! 
 
 O <^H r-3 
 
 ^,-r- v^ 
 
 oj s 
 
 
 I 
 
 I I -a 
 
 r -t 3 *3 
 
 S " 
 
 
 1 
 
 2 
 
 li 
 
 ^ g a 
 
 ^1 .s 
 
 ^'o a 
 
 S 
 
 ' 
 
 "Sg 
 
 s * 
 
 H 
 It 
 
 S p t 
 
 a> <3 
 
 2 S 
 
 S3 N 
 
 11 
 
162 METALLURGICAL CHEMISTRY 
 
 majority of cases the basic character of the oxide disappears 
 along with the acid character of the acid, and neutral bodies 
 are formed. Metallic hydroxides contain the elements of 
 water in addition to the oxides. They act in the same way 
 towards acids as the corresponding oxides. The peroxides 
 undergo partial decomposition when heated, and oxygen is 
 liberated. The acid-forming oxides unite with water to form 
 acids, and in this respect they resemble the acid-forming 
 oxides of non-metals. Salts are formed in a variety of ways, 
 but a given salt always has the same composition, however 
 it may be formed. The quantitative relations of neutraliza- 
 tion are just as definite as those of any other mode of chemical 
 action. 
 
 QUESTIONS. 
 
 1. How are oxides usually classified 1 ? Give examples. 
 
 2. What are the general effects of adding basic oxides to 
 acid solutions 1 
 
 3. Explain how an acid salt and a normal salt differ from 
 each other. 
 
 4. How are crystals usually formed 1 Why are some salts 
 said to be isomorphous 1 
 
 5. What is meant by the nomenclature of compounds 1 Give 
 examples to illustrate your answer. 
 
 6. Describe three ways of forming zinc sulphate. 
 
 7. Is there any similarity between the basicity of an acid and 
 the valency of a metal ? 
 
 8. Give a short account of the principle of isomorphism. 
 
CHAPTER X 
 CAEBON AND ITS COMPOUNDS 
 
 THE element carbon is a very common and useful substance. 
 It is a constituent of all animal and vegetable bodies; and 
 is present among the rocks and minerals of the earth's crust 
 in all natural carbonates, of which the calcium salt CaC0 3 
 is by far the most abundant. It is also the principal con- 
 stituent of coal, which is the product of the very slow decay 
 of vegetable matter out of contact with air. Such matter 
 contains carbon, hydrogen, oxygen, and nitrogen as essential 
 constituents, together with a small quantity of inorganic 
 or earthy matter. Air-dried wood contains about 50 per cent, 
 of carbon and 1*5 per cent, of inorganic compounds; the 
 remainder is made up of oxygen, hydrogen, and nitrogen. 
 During the extremely slow changes which the original wood 
 has undergone in its conversion to coal, the oxygen and 
 hydrogen, together with some of the carbon, have been largely 
 eliminated. Therefore the residue of coal contains a much 
 higher percentage of carbon than the vegetation from which 
 it was formed. In fact, the oldest variety of coal may be 
 described as impure carbon, containing as it does upwards of 
 90 per cent, of carbon. 
 
 Graphite or plumbago is also an impure form of carbon found 
 in the earth's crust ; and the diamond is the purest form of 
 natural carbon. 
 
 The various oils and fats of animal, vegetable, and mineral 
 origin are extremely rich in carbon. 
 
 163 112 
 
164 METALLURGICAL CHEMISTRY 
 
 Charcoal- When wood is heated in a vessel from which 
 the air is excluded, such as a porcelain crucible with the lid 
 on, a residue very rich in carbon is obtained; it is called 
 charcoal, and is a very useful fuel. 
 
 EXP. 87. Heat a piece of sound wood, about ^-inch cube, in a 
 porcelain crucible with the lid on until gaseous matter ceases to 
 escape. The crucible should be put on a pipeclay triangle over a 
 Bunsen flarne for a few minutes, and the heating finished over the 
 foot blow-pipe. When the crucible is nearly cold, remove the lid 
 and examine the residue. 
 
 Grind the piece of charcoal to a fine powder in a clean mortar, 
 put it into a small platinum crucible with a closely-fitting lid, and 
 heat it as strongly as possible for ten minutes in the blow-pipe 
 flame. 
 
 The experiment may be made quantitative by weighing the wood 
 before, and the charcoal after, heating. 
 
 The charcoal thus obtained contains about 92 per cent, of 
 carbon, 1*5 per cent, of earthy matter, which forms the ash 
 when the charcoal is completely burnt, and some oxygen and 
 hydrogen. 
 
 Some carbonaceous bodies are practically free from earthy 
 matter, and furnish very pure charcoal when properly treated. 
 Sugar is a very good illustration of these bodies. 
 
 EXP. 88. Melt some good lump-sugar in a large porcelain 
 crucible ; continue the heating until the combustible gas is all 
 expelled, and a black residue of " burnt " sugar is obtained. The 
 total quantity of sugar should be added to the crucible a little at a 
 time as it works down, or it will boil over and make the experiment 
 unsatisfactory. 
 
 Grind the black residue to a fine powder, and heat it very strongly 
 for ten minutes in a closed platinum crucible, as described above. 
 When cold, transfer the powder to a small stoppered bottle. 
 
 The carbon so obtained is not quite pure, but contains 
 about 0*1 per cent, of earthy matter and some oxygen and 
 hydrogen. The gaseous elements can be got rid of by strongly 
 heating the powder in a glass tube through which a stream of 
 dry chlorine gas is made to pass. The heating in contact with 
 chlorine is continued as long as fumes of hydrochloric acid 
 escape from the open end of the tube. A sample of carbon 
 
CARBON AND ITS COMPOUNDS 165 
 
 prepared as described in Exp. 88, and also subjected to the 
 action of chlorine, was found to contain 99*7 per cent, of 
 carbon and O'l per cent, of ash. 
 
 COMBUSTION OF CHARCOAL. 
 
 When charcoal is heated in a good current of air or in 
 oxygen, it burns, and the whole of the carbon in it is converted 
 into carbon dioxide, C0 2 . This compound is a colourless, 
 transparent gas, which dissolves in water, and forms a weak 
 acid solution. It is therefore an acid-forming oxide, and 
 combines with a number of basic oxides to form the salts 
 known as carbonates. It is rapidly absorbed by a solution 
 of caustic soda or potash, and forms a carbonate with the 
 alkali which remains in the solution. The reaction is shown 
 by the equation : 
 
 2KHO + C0 2 = K 2 C0 3 + H 2 0. 
 
 Potassium 
 carbonate. 
 
 This change is taken advantage of in estimating how much 
 carbon dioxide is formed by the combustion of a known weight 
 of carbon. The experiment requires considerable care, and 
 accurate weighings must be made ; but the following descrip- 
 tion will be readily followed. 
 
 EXP. 89. A small quantity (about O'l gram) of the purest char- 
 coal at hand is weighed on a bent strip of platinum foil, which is 
 then pushed into the middle of a glass combustion-tube, A (Fig. 30). 
 One end of A is then connected with the weighed potash bulbs, B. 
 which contain a strong solution of caustic potash, KHO, in water 
 (1 to 2). The other end of A is connected with the drying-tube C, 
 containing small lumps of caustic soda in the limb a, and calcium 
 chloride in the limb b. The part of the tube directly under the 
 platinum foil containing the charcoal is then heated in a Bunsen 
 flame, and a moderate stream of dry oxygen passed through the 
 apparatus. This is readily effected by using the apparatus shown in 
 Fig. 32. The bottle C is filled with oxygen, and the wash-bottle D, 
 containing a little strong sulphuric acid, is connected with the inlet 
 of the drying-tube C (Fig. 30). The charcoal burns in the oxygen ; 
 the carbon dioxide formed is carried forward by the excess of gas, 
 and is completely absorbed by the potash in B. The difference in 
 
166 
 
 METALLUKGICAL CHEMISTRY 
 
 the weight of the platinum foil containing the charcoal, before and 
 after the combustion, gives the weight of carbon burnt, and the 
 difference in the weight of the potash bulbs gives the weight of 
 carbon dioxide absorbed. A small quantity of ash is left on the 
 foil. 
 
 FIG. 30. 
 
 For very accurate work a drying-tube must be placed 
 between A and B. Powdered graphite and diamond dust 
 burai very slowly in a glass tube ; but a porcelain tube heated 
 in a table furnace (Fig. 8) gives very satisfactory results. 
 When pure carbon is used the ratio, carbon to carbon 
 dioxide, is found to be 12 : 44. That is, 12 parts of carbon 
 combine with 32 parts of oxygen to form 44 parts of carbon 
 dioxide. Therefore its formula is C0 2 ' 
 
 Experiments made with sugar carbon, purified graphite, 
 and diamond dust give practically the same result, and either 
 of them can be used to determine the composition of carbon 
 dioxide. 
 
 Experimental proof is thus furnished of the statement that 
 charcoal, graphite, and diamond are one and the same 
 Chemical element i.e., carbon. But there is a very marked 
 difference in their physical properties. These variations in 
 the properties of the same element are called allotropic 
 modifications (p. 57). The principal variations are in the 
 colour, hardness, and structure of the bodies. The diamond 
 is crystalline and very hard ; graphite is crystalline and soft ; 
 
CABBON AND ITS COMPOUNDS 167 
 
 charcoal is amorphous i.e., without crystalline form. Another 
 important difference is that, although the equal weights of the 
 three give the same weight of carbon dioxide, they do not 
 give out the same quantity of heat when burnt. The calorific 
 power of the diamond is less than that of graphite, and of 
 graphite less than that of charcoal. This is probably due to 
 the constitution and aggregation of their molecules. The 
 different forms of carbon can be changed from one to the 
 other. 1 hus, when a diamond is raised to the temperature of 
 the electric arc, it is converted into a black mass of coke -like 
 carbon. Amorphous carbon, after absorption by heated iron, 
 may be caused to separate from the molten mass in the form 
 of graphite. This form of artificial graphite is known as kisk. 
 Scales of crystalline graphite are easily recognised on the 
 fractured surface of a piece of grey pig-iron. The carbon 
 goes into the furnace in the form of coke. Amorphous carbon 
 has also been converted into diamond crystals ; but those 
 produced up to the present are too small to be of practical use. 
 
 CARBONATES. 
 
 Native Carbonates. Limestone, marble, and chalk, 
 which occur in enormous quantities in the earth's crust, are 
 varieties of calcium carbonate, CaC0 3 . The first two are dis- 
 tinctly crystalline in character. Calcite, or Iceland spar, a 
 purer form of the same compound, is found in the form of 
 rhombic crystals, which are sometimes as clear and trans- 
 parent as glass. Aragonite is another crystalline form of 
 calcium carbonate. 
 
 Magnesium carbonate, MgC0 3 , if isomorphous with the 
 calcium salt, and occurs along with it in dolomite, which is 
 usually described as a double carbonate of lime and magnesia. 
 Different specimens of dolomite contain the two carbonates in 
 varying proportions, which is consistent with the formation of 
 isomorphous mixtures (Chap. IX.). 
 
 Ferrous carbonate, FeC0 3 , which is the principal constituent. 
 
168 METALLUEGICAL CHEMISTEY 
 
 in spathic iron ore and clay ironstone, and manganous car- 
 bonate, MnC0 3 , are isomorphous with the calcium and 
 magnesium salts which often accompany them in their ores. 
 
 Zinc carbonate, ZnC0 3 , in calamine, and lead carbonate, 
 PbC0 3 , in cerussite, are also important natural carbonates. 
 
 Prepared Carbonates. Sodium carbonate, Na 2 C0 3 , is 
 the most important of these compounds. It is prepared both 
 in the anhydrous form, Na 2 CO 3 , and with water in soda 
 crystals, Na 2 C0 3 ,10H 2 0. Both forms are readily soluble in 
 water. Soda crystals lose water when exposed to the air, 
 and gradually break down to a white powder ; they are 
 efflorescent. When heated the crystals melt in their water 
 of crystallization, which is then removed by evaporation, 
 and the anhydrous salt is left as a white residue. This melts 
 to a clear liquid at a moderately high temperature. If mixed 
 with the corresponding potassium salt a more readily fusible 
 mass is obtained. Common fusion mixture contains Na 2 C0 3 + 
 K 2 C0 3 . 
 
 The majority of metallic carbonates are either insoluble or 
 only slightly soluble in water, and are precipitated as amorphous 
 solids when a solution of sodium or potassium carbonate is 
 added to solutions of the soluble compounds of the metals. 
 Very often, however, the oxide of the metal separates at the 
 same time, and basic carbonates are formed. 
 
 EXP. 90. Make a solution of sodium carbonate (1 in 10). Pour 
 about 10 c.c. of a solution of calcium chloride, CaCl 2 . into a large 
 test-tube, and add the sodium carbonate solution a little at a time, 
 with shaking, as long as a precipitate forms. Kaise the contents of 
 the tube to boiling, filter, and well wash the precipitate on the filter. 
 Dry the filter, and scrape off the white solid calcium carbonate. 
 Repeat the experiment with a solution of zinc sulphate in place of 
 the calcium chloride. Put some of the calcium carbonate in a test- 
 tube, add a little dilute hydrochloric acid, and pour the escaping gas 
 into a test-tube containing a little lime-water. Shake the lime-water 
 to bring it in contact with the gas, and note the milkiness caused by 
 the absorption of carbon dioxide. 
 
 Put a little of the dry zinc carbonate into a small test-tube, 
 heat it in the Bunsen flame, and prove that carbon dioxide is 
 evolved. 
 
CAEBON AND ITS COMPOUNDS 169 
 
 The reactions for the formation of the carbonates are as 
 follows : 
 
 CaCl 9 + Na 2 C0 3 = CaC0 3 + 2NaCl. 
 ZnSO* + Na 2 C0 3 = ZnC0 3 + Na 2 S0 4 . 
 
 Bicarbonates. These are acid salts, of which the best 
 example is bicarbonate of soda, NaHC0 3 . The normal soluble 
 carbonates have a decidedly alkaline reaction, for the weak, 
 acid-forming oxide, C0 9 , does not neutralize the strong soda 
 base, Na 2 0. The acid carbonates are milder because of the 
 excess of C0 2 present. They are readily decomposed at a 
 moderate heat into the normal salts, carbon dioxide, and 
 water. Normal carbonates are converted into the correspond- 
 ing acid salts when exposed to the action of carbon dioxide 
 and water, and sometimes an insoluble carbonate is thus con- 
 certed into a soluble one. When carbon dioxide is passed for 
 some time into a solution of sodium carbonate the bicarbonate 
 crystallizes out : 
 
 Na 2 C0 3 + C0 2 + H 2 = 2NaHC0 3 . 
 
 And when the gas is passed through water in which finely- 
 divided calcium carbonate is suspended, the solid passes into 
 solution. The acid carbonate is soluble : 
 
 CaC0 3 + C0 2 + H 2 = CaH 2 (C0 3 ) 2 . 
 
 Acid carbonate." 
 
 But on boiling the solution the C0 2 is expelled and the normal 
 carbonate precipitated. 
 
 ACTION or HEAT ON CARBONATES. 
 
 The sodium and potassium compounds are not decomposed 
 at a bright-red heat. Carbonates of Ba, Sr, Ca, and Mg are 
 decomposed at a good red heat into the basic oxides of the 
 metals and carbon dioxide; the barium carbonate requires 
 the highest temperature. Carbonates of the other common 
 metals are decomposed at a still lower temperature. Bicar- 
 bonates give off water in addition to carbon dioxide, but in 
 
170 METALLUBGICAL CHEMISTKY 
 
 the case of the alkaline bicarbonates only part of the CO 2 is 
 liberated : 
 
 2NaHC0 3 = Na 2 C0 3 + C0 2 + H 2 0. 
 
 EXP. 91. Weigh 1 gram of ground marble in a weighed porcelain 
 crucible ; heat the crucible, without the lid, in a gas muffle to a 
 bright-red heat for fifteen minutes. When the crucible is cold re- 
 weigh it. The loss in weight gives the carbon dioxide driven off. 
 If a muffle cannot be used, heat the crucible over a good circular 
 burner, and finish with five minutes over the foot blow-pipe. In this 
 case a platinum crucible is best. 
 
 EXAMPLE. One gram of marble lost 434 gram on being heated 
 for twenty minutes in a gas muffle. Compare this with the 
 equation : 
 
 CaC0 = 
 
 Sodium bicarbonate is readily decomposed in a porcelain 
 crucible heated over a Bunsen burner, and as the commercial 
 salt is comparatively pure, a good result is obtained. Zinc 
 and lead carbonates also furnish good examples. The residue 
 of zinc oxide is quite infusible ; but lead oxide is somewhat 
 fusible, and, unless the temperature is carefully regulated, it 
 will soften, stick to the crucible, and spoil it. 
 
 When a basic carbonate containing an excess of the an- 
 hydrous oxide is heated C0 2 only is driven off, but when the 
 hydroxide forms the excess of the base H 2 is also evolved. 
 In this case the loss includes both carbon dioxide and water. 
 In order to determine the proportions of the three oxides in 
 the salt it is necessary to determine the carbon dioxide by a 
 separate experiment, and then deduct its weight from the total 
 loss on heating. 
 
 EXP. 92. Weigh 1 gram of blue copper carbonate in a weighed 
 porcelain crucible, and heat it over the Bunsen flame until it turns 
 completely black. Allow the crucible to cool, and reweigh it. The 
 loss gives the weight of carbon dioxide and water driven off. 
 
 EXAMPLE. One gram of the blue carbonate lost 0-296 gram when. 
 heated for ten minutes over the Bunsen flame. 
 
CAEBON AND ITS COMPOUNDS 171 
 
 CARBONATES AND ACIDS. 
 
 All carbonates are decomposed by dilute acids with evolution 
 of carbon dioxide, and formation of the salts of the metals in 
 the carbonates and of the acids. If the salt formed is insoluble, 
 or only slightly soluble, in the acid solution, the action soon 
 ceases as the carbonate becomes coated with this insoluble 
 body, which protects it from the acid. Nitric acid is the best 
 general solvent. 
 
 Determination of the Weight of C0 2 in a Car- 
 bonate. The apparatus shown in Fig. 27 gives excellent 
 results, and is easy of manipulation. It may be used for any 
 carbonate. 
 
 EXP. 93. Add 10 c.c. of dilute nitric acid (1 to 8) to the flask A 
 (Fig. 27), and fit the parts of the apparatus together. Put the whole 
 on the scale-pan of the balance, and along with it a piece of filter- 
 paper which will conveniently hold 1 gram of the salt. Weigh the 
 whole accurately, then put a gram weight on the weight-pan, and 
 weigh 1 gram of blue copper carbonate on the filter-paper. Fold 
 up the paper containing the salt into a packet which can be readily 
 pushed through the neck of the flask. Take the bung out ; remove 
 the stopper from the outlet of the drying-tube ; push the packet 
 into the neck of the flask, and then into the body with the tube 
 passing through the bung, rapidly inserting the latter as the packet 
 falls in the acid. When the action is finished, connect the outlet of 
 the drying-tube with the aspirator (Fig. 28), and draw about a litre 
 of dry air through the flask to displace the remaining C0 2 still 
 present, and to restore the atmosphere in the flask to its original 
 condition. Eeplace the stoppers and reweigh the apparatus. The 
 loss gives the weight of C0 2 liberated from 1 gram of the salt, for 
 nothing else escapes from the apparatus. 
 
 EXAMPLE : Weight of 
 
 apparatus + filter-paper + 1 gram of the salt = 77v54 grams 
 ,, ,, ,, after experiment = 77*449 ,, 
 
 Loss due to escape of dry C0 2 = 0'205 gram 
 
 From the data supplied by Experiments 92 and 93 it is 
 found that 1 gram of blue copper carbonate contains O704 
 gram of CuO, O205 gram of C0 2 , and O091 gram of H 2 0. 
 Then by using the molecular weights, CuO -=79, C0 2 = 44, 
 
172 METALLUKGICAL CHEMISTKY 
 
 H 2 = 18, the ratio between the oxides in the carbonate can 
 be found and the formula deduced. Thus : 
 
 O^l =0 . 0891; 0^,0-00465; ^=0.005. 
 
 By reducing these quotients to the nearest whole numbers 
 the ratio 2 : 1 : 1 is obtained, and the empirical formula of 
 the compound is 2CuO.C0 2 .H 2 0, which is thus proved to be 
 a basic carbonate. The composition of the basic carbonates of 
 lead and zinc may be determined in the same way ; but as the 
 commercial compounds are not always pure, only approximate 
 results must be expected. 
 
 The determination of carbon dioxide in either prepared or 
 natural carbonates may be made with the apparatus used in 
 Exp. 93. Sodium carbonate and marble should be used 
 for preliminary experiments, as they are comparatively pure, 
 and thus furnish a check on the accuracy of the deter- 
 minations. 
 
 CARBON DIOXIDE (C0 2 ). 
 
 It will be useful at this point to examine this important 
 compound a little more closely than has already been done. 
 It is a constituent of the atmosphere, but the quantity present 
 varies somewhat under different conditions. Ten thousand 
 volumes of air contain about 5 volumes of carbon dioxide. 
 It is being constantly passed into the air as one of the 
 products of combustion, respiration, and putrefaction ; and as 
 constantly removed by growing plants, the leaves of which 
 absorb it readily. It is then decomposed under the influence 
 of sunlight, the carbon is assimilated by the growing tissue, 
 and the oxygen is returned to the atmosphere. In this way 
 a kind of rough balance is struck, and the average quantity 
 of carbon dioxide in the air kept about the same. 
 
 Preparation of Carbon Dioxide. If a considerable 
 volume of this gas is required to demonstrate its general 
 properties, it is most readily obtained by dissolving calcium 
 
CARBON AND ITS COMPOUNDS 
 
 173 
 
 carbonate in dilute hydrochloric acid. The reaction is as 
 follows : 
 
 2HC1 + CaC0 3 = CaCl 2 + H 2 + C0 2 . 
 
 EXP. 94. Put some small pieces of clean marble into the bottle 
 A (Fig. 31), and pour in a little water through the thistle funnel B, 
 so as to well cover the solid. Arrange the 
 delivery-tube in the cylinder C, and then 
 add strong hydrochloric acid, a little at a 
 time, through B. A steady stream of gas 
 is thus obtained, which, passing into C, 
 gradually displaces the air and fills the 
 cylinder. It is full when the flame of a 
 burning taper is immediately extinguished 
 on being brought into the mouth of the 
 jar. The gas may also be collected over 
 water. 
 
 FIG. 31. 
 
 Properties of Carbon Dioxide. 
 
 It is a colourless, transparent, odourless 
 gas, and is somewhat soluble in water, 
 which dissolves about 1*5 times its own volume of the gas at 
 ordinary temperatures. The solution has a sour taste, and 
 turns litmus red, thus showing its acid properties. It is 
 non-combustible, and a non-supporter of combustion in the 
 ordinary way; but burning magnesium continues to burn 
 in the gas with formation of magnesia and liberation of 
 carbon. Ked-hot charcoal also abstracts oxygen from it. 
 (See Exp. 96.) 
 
 Lime-water is turned milky when carbon dioxide is shaken 
 up with it. This is caused by the formation of insoluble 
 calcium carbonate. On further shaking to make the water 
 take up more gas the milkiness disappears, which is due to 
 the conversion of the normal into the soluble acid carbonate. 
 On boiling the solution the normal carbonate is thrown down 
 again. The reactions are shown by the equations : 
 
 CaH 2 2 + C0 2 = CaC0 3 + H 2 0. 
 CaC0 3 + C0 2 + H 2 = CaH 2 (C0 3 ) 2 . 
 
 EXP. 95. Shake up some lime-water in a jar of carbon dioxide 
 
174 METALLURGICAL CHEMISTRY 
 
 until the solution, which is first turned milky, becomes clear ; pour 
 the clear liquid into a test-tube, and boil it. 
 
 Carbon dioxide is heavy, and can be poured from one 
 vessel to another in the ordinary way. Its presence in the 
 second vessel is easily proved by putting in a lighted taper, 
 or by the addition of a little lime-water. It is 22 times 
 heavier than hydrogen, and 1*5 times heavier than air. 
 
 Carbonic Oxide, CO. The complete combustion of carbon 
 in air results in the formation of carbon dioxide ; but if the 
 supply of air is limited, and the temperature of the carbon is 
 sufficiently high, another oxide of carbon is produced. It is 
 probable that C0 2 is formed first, and then converted into CO 
 by abstraction of part of its oxygen. Thus : 
 
 C0 2 + C=2CO. 
 
 A minimum temperature of about 800 C., or a bright red 
 heat, is required for the change to go on rapidly. A good 
 table furnace is indispensable. The one shown in Fig. 8 
 has three Bunsen jets served from one gas tap, and the body 
 consists of a rectangular fire-brick chamber, surmounted by 
 a sheet iron tube to assist the draught. The reaction-tube 
 is a piece of solid-drawn steel tube, such as is used for cycle 
 work, and about 18 inches long. Rubber bungs carry the 
 inlet-tube A arid the outlet-tube B, and are kept cool by pads 
 of blotting-paper, which are wrapped round the ends of the 
 steel tube and kept well saturated with water. The charcoal 
 used should be previously heated in a closed clay crucible to 
 the highest temperature of a wind furnace to expel hydrogen 
 and gaseous hydrocarbons, which are not removed during the 
 ordinary process of manufacture. This preliminary heating 
 may be omitted, but then the first two or three jars of gas 
 collected will be impure. 
 
 EXP. 96. Insert a loose plug of asbestos in the steel tube about 
 j its length from one end, and fill the middle part with small pieces 
 of charcoal about the size of a pea. Connect the apparatus for 
 generating dry CO.,, A and B, Fig. 36, with A, Fig. 8, and by the 
 careful addition of dilute nitric acid keep up a moderate stream of 
 
AND ITS COMPOUNDS 175 
 
 hubbies through the wash-bottle, so as to insure a regular stream of 
 gas through the red-hot tube. If the carbonic oxide is to be quite 
 free from carbon dioxide, a wash-bottle containing a strong solution 
 of caustic soda must be inserted between the end of the steel tube 
 and the collecting cylinder. Collect several jars of the gas. 
 
 Properties Of CO. It is a colourless gas with a faint 
 oppressive odour, but must not be inhaled, as it is very 
 poisonous, and should not be allowed to escape unburnt. It 
 burns with a pale blue flame, and explodes when mixed with 
 air or oxygen. The pure gas does not turn lime-water milky, 
 but when it is burnt the product of the combustion does turn 
 the lime water milky, thus proving that carbon dioxide is 
 formed. 
 
 = 2C0 
 
 2 . 
 
 4 vols. 2 vols. 4 vols. 
 
 CO is only slightly soluble in water. Its density, compared 
 with air= 1, is 0'97. 
 
 Experiments With CO. (1) Add a little lime-water to a 
 jar of the pure gas, and shake the jar; burn the gas and 
 shake the jar again. (2) Mix 1 volume of the gas with 2'5 
 volumes of air in a tall cylinder, and ignite the mixture. The 
 flame runs rapidly down the jar and thus furnishes an example 
 of incipient explosion. (3) Mix 2 volumes of the gas with 
 1 volume of oxygen, and ignite the mixture. The combustion 
 is much more rapid, and a decided explosion is the result. 
 The violence of the explosion when an explosive mixture of 
 gases is ignited, depends upon the rapidity with which the 
 combustion extends through the mass. 
 
 Formation of CO when a Limited Supply of Air is passed 
 through Red-hot Carbon. EXP. 97. Disconnect the CO 2 apparatus 
 from the tube used in the last experiment, and replace it by any 
 convenient apparatus for sending a slow stream of dry air through 
 the red-hot charcoal. The form shown in Fig. 32 is very simple. B is 
 full of water ; C is full of air ; D contains some strong sulphuric acid. 
 By running the water slowly from B into C a slow stream of dry 
 air is forced through the red-hot charcoal, when the drying-bottle 
 D is connected with the tube containing the charcoal. If the 
 issuing gas is required to be free from CO 2 , the traces of this gas 
 
176 
 
 METALLURGICAL CHEMISTRY 
 
 B 
 
 can be removed by caustic soda, as described above. Collect two or 
 three jars of the gas. Add some lime-water to a jar of the gas 
 
 which has been puri- 
 fied by caustic soda. 
 If the liquid remains 
 clear the gas is free 
 from CO 2 . Put a 
 lighted taper to the 
 mouth of the jar, and 
 when the combustion 
 has ceased, shake up 
 the gas with the 
 lime - water, which 
 turns milky, thus 
 proving the presence 
 of the CO.,, formed 
 by the combustion of 
 the CO. Inflame 
 another jar of the 
 
 DJIIJII I'll | gas, and notice that 
 
 ll_ JPA it burns with the 
 
 characteristic blue 
 flame of CO. 
 
 FIG. 
 
 The carbonic 
 oxide formed oc- 
 cupies twice the volume of the oxygen absorbed, and is also 
 mixed with the nitrogen of the air from which the oxygen is 
 taken. Thus the mixture contains one-third CO and two- 
 thirds N 2 by volume. 
 
 Action of Red-hot Carbon on Water Vapour. The 
 
 principal reaction which takes place is one in which carbonic 
 oxide is formed and hydrogen gas liberated. 
 
 Another reaction which takes place in a much more limited 
 degree brings about the formation of marsh gas, CH 4 , and 
 probably other hydrocarbons of a similar character. 
 
 2 = CH 4 -f2CO. 
 
 If the temperature falls too low, hydrogen and carbon dioxide 
 are formed. 
 
CARBON AND ITS COMPOUNDS 177 
 
 So that when steam is passed through red-hot carbon at least 
 three combustible gases are formed. Both reactions are 
 endothermie i.e., more heat is absorbed in effecting the 
 decomposition of the water taking part in the reactions than 
 is evolved by the formation of the resulting compounds when 
 the change takes place ; and there is a definite limiting tem- 
 perature below which it ceases. It is therefore necessary to 
 keep the temperature above this limit by the application of 
 heat from outside, or by allowing some of the carbon to burn 
 to carbon dioxide. Similar remarks apply to the reaction 
 C + C0 2 = 2CO, which is also endothermie. On the other 
 hand, C + 2 = C0 2 is an exothermic change i.e., heat is 
 given out while it is taking place. 
 
 EXP. 98. Use the apparatus shown in Fig. 8, but replace the 
 iron by charcoal. Eegulate the boiling of the water in the flask so 
 as to give a steady stream of bubbles from the delivery-tube into 
 the collecting cylinder. Collect two or three jars of the gas. It is 
 a mixture of carbonic oxide, hydrogen, and marsh gas, and is called 
 " water " gas. Burn a jar of the gas, and note the faintly luminous 
 flame. Mix 1 volume of the gas with 2 volumes of air, and 
 explode the mixture. The explosion is louder than ..'th carbonic 
 oxide alone. 
 
 Ordinary air contains water vapour, and if it is used in 
 Exp. 97 instead of dry air, a small amount of hydrogen is 
 present in the gaseous mixture obtained. 
 
 Coal is a mineral of vegetable origin and of very complex 
 .composition. It contains C, H, 0, N, and ash-forming con- 
 stituents. Sulphur is also invariably present in small quantity. 
 The whole of the carbon and part of the hydrogen are avail, 
 able for the development of heat when the coal is burnt. In 
 practice it is usual to subtract from the total hydrogen one- 
 eighth the weight of the oxygen present, and to consider that 
 quantity of hydrogen as already combined in the form of 
 water, and therefore not available for heat production. The 
 remainder is known as available hydrogen. The older the 
 coal the more concentrated it is ; for the changes which the 
 vegetable matter undergoes during its extremely slow conver- 
 
 12 
 
178 
 
 METALLUKGICAL CHEMISTKY 
 
 sion into coal cause the percentage of carbon, and also the 
 ratio of the hydrogen to the oxygen, to increase. That is, the 
 percentages of oxygen and hydrogen are both decreased, but 
 the former more than the latter. 
 
 Coal containing much hydrogen undergoes very complicated 
 changes when it is strongly heated out of contact with air. 
 The greater part of the oxygen is liberated as water; the 
 nitrogen principally in the form of ammonia, and the residue 
 of the hydrogen partly in the free state, and partly in a very 
 numerous and complex series of carbon compounds. 
 
 EXP. 99. Draw off a piece of combustion-tube in the blow-pipe 
 flame ; put in it a few grams of powdered coal ; draw off and bend 
 the thinned-out portion, as shown in Fig. 33. A is the loaded tube 
 in position ; B is a bottle to condense liquid matter ; C is a cylinder 
 to collect the escaping gas. Strongly heat the tube as long as gas 
 
 FIG. 33. 
 
 collects in C. Remove the cylinder, and put a light to the mouth. 
 The gas burns with a bright white flame. B is found to contain a 
 watery liquid and a black oily liquid. On breaking A, a greyish- 
 black coherent mass of coke is obtained. 
 
 The coal is split up into: (1) a combustible gas; (2) a 
 watery liquid which contains ammonia; (3) tar, a liquid of 
 very complex character; (4) coke, a solid residue containing 
 carbon (about 90 per cent.), some hydrogen and oxygen, and 
 inorganic matter or ash. 
 
CAEBON AND ITS COMPOUNDS 179 
 
 APPROXIMATE COMPOSITION OF COAL GAS. 
 
 Hydrogen, H 2 ... ... ... ... 45 
 
 Marsh gas, CH 4 ... .... ... ... 40 
 
 Carbonic Oxide, CO ... ... ... 5 
 
 Hydrocarbons ... ... ... ... 5 
 
 Non-combustible gases ... ... ... 5 
 
 100 
 
 The experiment just described illustrates in a general way 
 the principles involved in the manufacture of illuminating 
 gas. If wood is substituted for coal in Exp. 99, a similar 
 result is obtained, and so striking is the analogy that one is 
 forced to admit the similarity in the composition of wood and 
 coal, and the probability that they have the same origin. 
 
 The Changes in an Open Coal Fire. When fresh coal 
 
 is put on an open fire the heat from the combustion of the 
 fuel already in the grate brings about the changes described 
 above. Water is expelled, and forms the greater part of the 
 first smoke. Then the volatile matter becomes combustible, 
 and burns with a white flame in the current of air passing 
 over the fire on its way to the chimney. The coal is gradually 
 coked, and the fire settles down to steady combustion. 
 Another current of air passing through the fire by way of 
 the grate-bars gives up its oxygen to the incandescent carbon 
 as soon as it enters, and carbon dioxide is formed. This gas 
 is carried into the heart of the fire, there to be converted into 
 carbonic oxide by the red-hot carbon. On emerging from the 
 upper surface of the fire this carbonic oxide burns with its 
 characteristic blue flame in the stream of air passing over the 
 top. The water vapour in the lower current of air also under- 
 goes similar changes to those already described, hydrogen and 
 carbonic oxide being added to the combustible gas. The 
 earthy matter of the coal is left in the grate as ash. It is 
 evident, from the accumulation of soot in the chimney, and 
 
 122 
 
180 METALLURGICAL CHEMISTRY 
 
 the contamination of the atmosphere, that combustion in an 
 open grate is far from perfect. 
 
 Closed Grate. If the top of the grate is closed so as to 
 cut off the upper air current, and the bar spaces are arranged 
 to admit a limited supply of air, the combustion is modified, 
 and the gas drawn off from the closed top contains upwards 
 of 30 per cent, of combustible matter. This is known as 
 * air ' gas or ' producer ' gas, and is made on the large scale for 
 heating purposes. By introducing more water vapour, up to 
 10 per cent., gas containing a larger percentage of hydrogen 
 is obtained. The most recent improvements are in this 
 direction. 
 
 AVERAGE COMPOSITION OF ' Am ' GAS. 
 
 Carbonic oxide, CO ... ... ... 25 
 
 Carbon dioxide, C0 2 ... ... ... 5 
 
 Hydrogen, H 2 ... ... ... ... 10 
 
 Nitrogen, N 2 ... ... ... ... 55 
 
 Hydrocarbons ... ... ... ... 5 
 
 100 
 
 The changes are for the most part endothermic, and the 
 complete combustion of a small percentage of the carbon must 
 take place to supply the necessary heat for the endothermic 
 changes. 
 
 SUMMARY. 
 
 Carbon is found in the free state in Nature. It is also a 
 constituent of many bodies of very diverse composition. It 
 exists in three distinct allotropic modifications. The car- 
 bonates are a very important class of metallic compounds, and 
 some of them are used for the extraction of the metals they 
 contain. They are readily decomposed by heat into metallic 
 oxides and carbon dioxide. Acids dissolve them with evolu- 
 tion of the same gas. There are only two oxides of carbon. 
 The dioxide is an acid-forming oxide, and the monoxide is a 
 
CARBON AND ITS COMPOUNDS 1S1 
 
 very useful combustible gas. Coal and its derivatives, coke 
 and combustible gas, are extensively used in the extraction 
 and after-treatment of the common metals. 
 
 QUESTIONS. 
 
 1. Give a short account of the different forms of carbon. 
 How are they proved to be one and the same element ? 
 
 2. What is the action of heat on carbonates generally ? 
 
 3. What are the properties of carbon dioxide, and how may 
 it be converted into carbon monoxide ? 
 
 4. What happens when coal or wood is heated in a closed 
 vessel ? 
 
 5. Explain fully what takes place when steam is brought 
 into contact with red-hot carbon. 
 
 6. Describe the changes taking place in an ordinary coal 
 fire. What is the result of closing the top of the grate ? 
 Give the equations expressing the various changes. 
 
CHAPTER XI 
 EEDUCTION 
 
 REDUCTION is a term used by metallurgists to express the 
 chemical change which takes place when a metal is set free 
 from its combinations with non-metals. As oxides and 
 sulphides are the most abundant natural compounds of the 
 metals, they furnish frequent examples of reducing action, 
 and are the chief sources of the prepared metals. But for 
 illustrating general principles it will be more convenient to 
 use prepared compounds, and so avoid the complications 
 which would be introduced by the presence of impurities in 
 the natural materials. 
 
 Reduction by Heat alone. The simplest change is that 
 which takes place when an oxide is decomposed by heat. 
 There are very few examples of this kind of change, for the 
 oxides generally are stable bodies, and not readily decomposed 
 by heat alone. Oxides of mercury, silver, and gold are the 
 only common examples. The decomposition of red oxide of 
 mercury (p. 39) may be taken as typical of the change. But 
 when the compounds are salts of these oxides the decomposi- 
 tion is more complex. Thus with silver nitrate, nitrogen 
 peroxide and oxygen are set free along with the metal. 
 
 2AgN0 3 = 2Ag + 2N0 2 + 2 . 
 
 EXP. 100. Put a crystal of silver nitrate into a small dry test- 
 tube, and heat it in the Bunsen flame. Note the red gas which is 
 given off, and test it with a glowing splint. Remove the reduced 
 silver from the tube, and rub it on a hard surface with the blade of 
 a knife. The particles of the reduced metal are pressed together, 
 
 182 
 
REDUCTION 183 
 
 and present the characteristic metallic appearance. The experiment 
 may be made quantitative by weighing the nitrate (about 0'5 gram) 
 before, and the residue of silver after, heating. 
 
 A partial reduction is often effected by heat alone, and a 
 simpler compound of the metal thus obtained. This is notably 
 so with the metallic carbonates, of which some examples have 
 been given (p. 169). The decomposition of lead carbonate is 
 easily brought about, but the residue of lead oxide is some- 
 what readily fused, and, if allowed to get too hot, would spoil 
 a porcelain crucible. Both the normal and the basic carbonates 
 are prepared, and a simple experiment will distinguish between 
 them. 
 
 EXP. 101. Weigh a convenient piece of a broken porcelain dish, 
 and weigh on it 1 gram of lead carbonate. Put the porcelain on a 
 pipeclay triangle and heat it until partial fusion takes place. Allow 
 the porcelain to cool, and re weigh it. 
 
 The loss is due to the escape of carbon dioxide, and the 
 result may be compared with the equation 
 
 PbC0 3 = PbO + C0 2 . 
 
 A partial reduction also takes place when some sulphates are 
 strongly heated, but the reaction is not quite so general as 
 with the carbonates. (See Exp. 86.) 
 
 EXP. 102. Weigh 2 grams of powdered green vitriol in a weighed 
 porcelain crucible, and heat it over a small Bunsen flame for fifteen 
 minutes. Allow the crucible to cool, and reweigh it. The loss is 
 due to the escape of the water of crystallization. Then put the 
 crucible into a gas muffle, and keep it at a moderate red heat for 
 thirty minutes. When the crucible is cold, reweigh it. The second 
 loss is due to the escape of oxides of sulphur. The residue, if 
 not overheated, is a red powder. Compare the results with the 
 equation : 
 
 2(FeS0 4 , 7H 2 0)=Fe 2 3 +S0 2 +S0 3 +14H 2 0. 
 
 Residue. Second loss. First loss. 
 
 This experiment illustrates the method used in the preparation 
 of rouge, and as the colour depends largely on the temperature 
 at which the reduction takes place, care must be exercised in 
 
184 METALLURGICAL CHEMISTRY 
 
 carrying it out. The value of the rouge also depends largely 
 on the grinding and levigation which the residue undergoes 
 after the reduction. (See Exp. 3.) 
 
 The partial reduction of metallic sulphides rich in sulphur 
 is demonstrated in Exp. 38. And it may be taken generally 
 that when there are two or more sulphides of the same metal, 
 the one containing the smallest proportion of sulphur is the 
 most stable, and will withstand a high temperature without 
 decomposition. The others are reduced to this stable form 
 when strongly heated. Iron and copper sulphides furnish 
 good examples : 
 
 1. 
 
 2. 
 
 The lower sulphide in each case may be fused at a bright red 
 heat out of contact with air. 
 
 Reduction by Electricity. The decomposition of water 
 (p. 27) is a simple reduction process as far as the observed 
 effects are concerned. For, though a little sulphuric acid is 
 necessary to the success of the experiment, the final result 
 is expressed thus : 
 
 If a solution of copper sulphate is substituted for the acid 
 solution in the voltameter, a deposit of copper is formed on 
 the plate from which the hydrogen escapes during the 
 decomposition of water, and oxygen gas escapes from the 
 other plate. The final result is therefore equivalent to the 
 reduction of copper oxide, whatever the intermediate changes 
 may be. 
 
 2CuO = 2Cu + 2 . 
 
 Free sulphuric acid appears in the solution. This is probably 
 due to secondary changes in which water takes part, and 
 furnishes the evolved oxygen. 
 
 1. CuS0 4 = Cu + (S0 4 ) 
 
 Cathion. Anion. 
 
 2. 
 
REDUCTION 185 
 
 Similar reactions take place when a current of electricity is 
 passed through a solution of any metallic salt, and if the 
 liberated metal does not act upon water, it is deposited upon 
 the cathode plate (see p. 107). But if the metal reacts with 
 water its hydroxide is formed, and hydrogen is liberated. 
 Common salt fuses to a clear liquid at a red heat, and if a 
 current of electricity is passed through the molten mass 
 between carbon electrodes, the metal sodium and the gas 
 chlorine are liberated and can be readily collected. 
 
 Sodium 
 chloride. 
 
 Similarly, the cheap production of aluminium depends upon 
 the passage of an electric current through fused cryolite, 
 AlF 3 ,3NaF, in which alumina is dissolved. The cryolite is 
 unchanged. 
 
 2A1 2 3 = 4 
 
 Reduction by Heat and Carbon. Every compound 
 
 tends to decompose when it is heated, and the higher the 
 temperature to which it is exposed^ the greater its proneness 
 to decomposition. Many compounds, however, resist a high 
 temperature, and if their reduction is to be effected the action 
 of heat must be assisted by that of a reducing agent, which 
 may be either an element or another compound. The reduc- 
 ing agent exerts a pull upon one or more constituents of the 
 compound, and thus assists the heat to effect its decomposition. 
 The action is most effective when the reducing agent forms a 
 compound which is volatile at the temperature of the reduc- 
 tion, and is rapidly removed from the sphere of action. But 
 this is not a necessary condition. Carbon, when it helps to 
 reduce a metallic oxide, forms either carbon dioxide or carbon 
 monoxide, according to the temperature necessary to bring 
 about the reduction. 
 
 EXP. 103. Well mix 0'2 gram of powdered charcoal with 8 grams 
 of lead oxide ; put the mixture into a small test-tube, and tit the 
 
186 
 
 METALLUEGICAL CHEMISTEY 
 
 tube with a bung and delivery-tube, as shown in Fig. 34. Heat 
 the mixture strongly for five minutes, and collect the escaping 
 
 gas in the jar. Test the gas 
 with a lighted taper ; it is a non- 
 supporter of combustion. Pour 
 the gas from the jar into a 
 beaker containing a little clear 
 lime - water, and shake the 
 beaker ; the gas is heavy, and 
 turns lime-water milky. It is 
 carbon dioxide, C0 2 . Examine 
 the residue in the tube for shots 
 of metallic lead. 
 
 A similar experiment may 
 be made with black oxide of 
 but the proportion 
 
 FIG. 34. 
 
 copper 
 
 between the charcoal and the oxide must be nearly three 
 times as great. 
 
 EXP. 104. Put about 2 grams of powdered wood charcoal into a 
 porcelain or platinum crucible ; close the crucible with a lid to 
 prevent the admission of air, and heat it strongly over the blow-pipe 
 flame for ten minutes. Allow the crucible to get cold before 
 removing the lid. Charcoal thus strongly heated contains about 
 92 per cent, of carbon. Weigh 054 gram of this charcoal, which 
 is roughly equal to 0'05 gram of carbon. Mix it thoroughly with 
 4 grams of dry litharge, and introduce the mixture, with the help 
 of a paper gutter, into a piece of combustion tube, ^ inch by 
 4 inches, closed at one end. Weigh the tube and mixture together ; 
 fix the tube in a clip, and heat it strongly in the Bunsen flame 
 until the mixture becomes pasty. Allow the tube to cool, and 
 reweigh it. 
 
 The carbon is converted into carbon dioxide by oxygen 
 obtained from the lead oxide, and the excess of the oxide 
 insures the combustion of the whole of the carbon. 
 
 EXAMPLE. Weight of tube + mixture = 18 '983 grams 
 
 after heating = 18*799 
 
 Loss = 0184 gram 
 
 The loss is due to the escape of carbon dioxide containing 
 0'05 gram of carbon. Therefore the weight of oxygen com- 
 
KEDUCTION 
 
 187 
 
 bined with this weight of carbon is (H34 gram. These weights 
 may be used to find the formula of the compound. 
 
 = 0-0083. .-. The ratio is 1:2, 
 
 For ^ = 0-0041; and ^| 
 
 and the formula is C0 2 , 
 
 Compare the results of the experiment with the equation 
 
 EXP. 105. Draw off a piece of -inch combustion-tube as shown 
 in Fig. 85. Weigh accurately 3 grams of lead oxide, and transfer 
 it to the tube with the 
 help of a paper gutter. 
 Weigh the tube and oxide 
 together. Fix it in a 
 clip, and connect the thin 
 end with a rubber tube to 
 the gas -tap. Turn on the 
 gas, and ignite it at the 
 open end of the tube, 
 allowing it to burn with 
 a moderate flame . Heat 
 the tube in the Bunsen 
 flame, and watch the re- 
 duction as it proceeds. 
 Note that moisture col- FIG. 35. 
 
 lects on the cool part of 
 
 the tube at first, but disappears as it gets warm. When the reduc- 
 tion is complete, which is indicated by the disappearance of the 
 oxide, remove the burner, but allow the gas to continue passing 
 through the tube until it has cooled somewhat ; then reweigh it. 
 
 Coal gas contains carbon and hydrogen, and the water 
 formed is the result of the oxidation of the latter (p. 189). 
 
 EXAMPLE. 3 grams of lead oxide lost 0*219 gram of oxygen. 
 
 Then 
 
 0-219 
 
 = 0-0137; and 
 
 = 0134, so that the ratio is 1:1. 
 
 16 207 
 
 Therefore the formula of the oxide used is PbO. 
 
 An exactly similar experiment can be made with copper 
 oxide. 
 
 The temperature required for the above reduction is com- 
 paratively low, and carbon dioxide is formed ; but when a 
 
188 METALLUEGICAL CHEMISTRY 
 
 high temperature is necessary for the reduction carbonic 
 oxide, CO, is the result. This is not easily demonstrated, 
 and must be left for more advanced work. The reduction of 
 tin and zinc oxides may be cited as examples, and the 
 simplest way of representing the reactions is shown in the 
 equations 
 
 Sn0 2 +2C = Sn + 2CO. 
 ZnO + C = Zn + CO. 
 
 But it is probable that carbon dioxide is first formed, and 
 then reduced by excess of carbon at the high temperature 
 required. Some oxides, such as alumina and lime, require the 
 excessively high temperature of the electric furnace to effect 
 their reduction. 
 
 Reduction by Hydrogen. The action of hydrogen gas 
 upon hot copper oxide has received more attention from 
 eminent chemists than any other experiment in the whole 
 range of the science. It was used by Dumas in his deter- 
 mination of the combining proportion of oxygen, and his 
 experiment has been repeated several times, with every possible 
 refinement of experimental research. For full details of this 
 work some standard book on the subject must be consulted ; 
 but a comparatively simple experiment may be described, and 
 even carried out with ordinary care. In this case approximate 
 values only must be expected. 
 
 EXPERIMENT TO PROVE THE COMPOSITION OF WATER BY 
 
 WEIGHT. 
 
 The apparatus shown in Fig. 36 is to be fitted up. A is a 
 bottle fitted with a tap funnel, and used for preparing the 
 hydrogen. B is a wash-bottle containing a little water to 
 wash the gas. C is a drying-tube filled with fused calcium 
 chloride and glass-wool. D is a combustion-tube containing 
 ignited copper oxide. E is the water-absorbing portion, which 
 consists^of a test-tube, the upper part of which is filled with 
 
EEDUCTION 
 
 189 
 
 glass-wool, and a calcium chloride drying-tube. The delivery- 
 tube from C dips into a small beaker containing a little strong 
 sulphuric acid. This serves to show the rate at which the 
 excess of hydrogen is escaping, and also prevents the absorption 
 of moisture from the air. 
 
 To make the experiment, D and E are weighed separately, 
 the rubber joints a, 6, and c being closed by the insertion of 
 bits of glass rod. The connections are then made, and a 
 moderate current of hydrogen passed through the apparatus, 
 escaping through the acid in the beaker. When the air has 
 been displaced, the copper oxide in D is heated ; the reduction 
 commences, and is continued until a quantity of water has 
 
 FIG 
 
 collected in the test-tube of E. The flame is then removed, 
 and the tube allowed to cool with the hydrogen still passing 
 through it. The parts are then separated and reweighed. If 
 the experiment is conducted with care, it proves approximately 
 that the loss of weight in D is to the gain of weight in E as 
 8 is to 9 ; or that 8 parts by weight of oxygen combine with 
 1 part by weight of hydrogen to form 9 parts by weight of 
 water. The reduction is expressed by the equation 
 
 CuO+H 2 = Cu+H 2 0. 
 
 Ferric oxide, Fe 2 3 , is also readily reduced by hydrogen, and 
 the reaction may be made the subject of a simple experiment. 
 
190 METALLURGICAL CHEMISTRY 
 
 EXP. 106. Fit up the apparatus shown in Fig. 36, with the 
 exception of the part E. Put 2 grams of ferric oxide into the tube D, 
 and pass a steady current of hydrogen through it. When the air 
 has been displaced, -heat the tube until the red powder turns com- 
 pletely black. Allow the tube to get cold while the gas is still 
 passing through it ; then transfer the reduced metal to a dry test- 
 tube and cork it up, to exclude the air from the metal. 
 
 The condensation of water in the cold neck of the com- 
 bustion-tube is readily observed. The equation is : 
 
 If ferric oxide is reduced at as low a temperature as possible, 
 a black pyrophoric powder is obtained, which takes fire spon- 
 taneously if allowed to come into contact with the air ; but if 
 reduced at a higher temperature the metal is much more 
 stable, and may be exposed to the air when cold without 
 alteration. If a gram of this reduced metal is heated in an 
 open crucible, it increases to 1*43 grams, which corresponds 
 to the formula Fe 2 O 3 . The iron is again converted completely 
 into its oxide. 
 
 A number of other metallic oxides are reduced by hydrogen 
 under similar conditions. 
 
 Reduction by Carbonic Oxide. The action of carbonic 
 
 oxide is almost as general as that of carbon and hydrogen. 
 It may therefore be described as a common reducing agent 
 for oxides. The reduction of oxide of iron is very interesting, 
 because the reduced metal reacts on the excess of the reducing 
 gas and absorbs carbon from it. The first change is definite, 
 and is expressed thus : 
 
 It is found that a considerable excess of carbonic oxide must 
 be used, or the reduction is not complete. The reaction is the 
 principal reducing one taking place in an iron-smelting blast 
 furnace, and it is well known that there must be at least twice 
 as much carbonic oxide as carbon dioxide in the waste gas if 
 the furnace is to work satisfactorily. The general reactions 
 
REDUCTION 191 
 
 by which the reducing gas is formed in the blast furnace are 
 described on p. 174; but by using a forced draught a very 
 much higher temperature is obtained in the lower part of the 
 furnace than in a closed grate. For the fusion of the slag and 
 of the reduced metal, as well as the formation of carbonic 
 oxide, has to be considered. A falling off in the proportion 
 of carbonic oxide would result in a partial reduction by which 
 ferrous oxide would be formed, and pass into the slag, with 
 consequent loss of iron. The reaction is : 
 
 Ferric Ferrous 
 
 oxide oxide 
 
 The second change, in which the reduced metal reacts with 
 carbonic oxide and abstracts carbon from it, is not so definite 
 as the actual reduction of the metal, and is usually considered 
 generally. Thus, if x is used to represent an unknown but 
 determinate quantity, the reaction may be represented thus : 
 
 xFe + 2CO = Fe^C + C0 2 . 
 
 The reactions expressed in the last three equations are readily 
 demonstrated, but the first and third only need be considered 
 here. The carbonic oxide required may be prepared in any 
 convenient way; but the decomposition of yellow prussiate 
 of potash by strong sulphuric acid furnishes a ready means of 
 obtaining the pure gas. 
 
 EXP. 107. Fit up a 2-litre bottle, as shown in Fig. 37, to serve 
 as a simple gas-holder. Put 6 grams of powdered yellow prussiate 
 of potash into an 8 ounce flask, cover it with strong sulphuric acid, 
 and fit the neck with a bung and delivery-tube. Fill the gas-holder 
 with water ; connect the flask with the tube A, and gently heat it 
 on gauze over a small flame. Carbonic oxide will be given off 
 steadily if the flask is not heated too rapidly, and will displace the 
 water through the tube B. The quantity of yellow prussiate used 
 should be arranged so that the liberated gas does not quite fill the 
 gas-holder. When the evolution of gas ceases close the tube A with 
 a pinch tap, and the apparatus is ready for use. Draw off a piece 
 of combustion-tube similar to that used in Exp. 105 ; put into it 
 2 grams of ferric oxide, and arrange it in a clip for heating over the 
 Bunsen burner. Connect A, Fig. 37, with a drying-bottle contain- 
 
192 
 
 METALLURGICAL CHEMISTRY 
 
 FIG. 37. 
 
 ing a little strong sulphuric acid, and B with the tap of an aspirator 
 
 filled with water. Place the aspirator on a stand so that its bottom 
 is on a level with the tube A. Connect the 
 outlet tube of the drying-bottle with the 
 drawn- off end of the reduction -tube ; heat 
 the part of the tube containing the ferric 
 oxide with a Bunsen flame, and drive a 
 steady stream of carbonic oxide over it by 
 opening the tap of the aspirator. The flow 
 of water through B forces the gas through 
 A, and the rate of flow can be easily regu- 
 lated. Place a Bunsen flame near the out- 
 let of the reduction-tube to burn the excess 
 of carbonic oxide as it escapes. This is 
 most important on account of the very 
 poisonous nature of the gas. When the 
 reduction is complete, fit a cork in the end 
 of the tube to prevent admission of air, 
 and turn off the tap of the aspirator at the 
 
 same time. When the reduced metal is cold transfer it to a dry 
 
 test-tube, and cork it up. 
 
 If a weighed quantity of this reduced metal is heated in an 
 open porcelain crucible, it is changed into the red oxide again, 
 but the increase in weight falls considerably short of that 
 obtained in Exp. 106 with the pure metal. This is due to the 
 presence of carbon, which is proved conclusively by the 
 following experiment : 
 
 EXP. 108. Transfer a little of the reduced metal to a combustion- 
 tube ; aspirate a current of air over it and through a wash-bottle 
 containing lime-water; heat the part of the tube containing the 
 metal, and note the milky precipitate formed in the lime-water. 
 This proves without doubt that carbon is associated with the iron 
 when it is reduced by carbonic oxide. 
 
 Coal gas contains about 5 per cent, of carbonic oxide, arid 
 if used for the reduction of ferric oxide the reduced metal 
 contains carbon, but in smaller quantity than when pure car- 
 bonic oxide is the reducing agent. 
 
 EXP. 109. Reduce about 2 grams of ferric oxide in the same 
 way i-hat lead oxide is reduced in Exp. 105, and compare the result 
 of heating 1 gram of the reduced metal in a porcelain crucible with 
 the similar results obtained as described above. 
 
REDUCTION 193 
 
 The complex character of coal gas makes the individual 
 reactions somewhat numerous when it is used as a reducing 
 agent, but the final result is the same as would be obtained 
 with carbon and hydrogen together. The reaction with the 
 principal hydrocarbon when lead oxide is reduced by coal gas 
 is shown by the equation : 
 
 4PbO + CH 4 = 4Pb + C0 2 + 2H 2 0. 
 
 Marsh gas. 
 
 Reduction by Metals. The great affinity which a par- 
 ticular metal has for an element, or group of elements, in a 
 compound often enables it to replace another metal already in 
 the compound, and so liberate it in the metallic state, when 
 the proper conditions are present. This is the case with 
 aluminium, which exerts a powerful reducing action on some 
 metallic oxides in consequence. The action is accompanied by 
 an extremely rapid evolution of heat, and the local tempera- 
 ture obtained is sufficient to melt the reduced metal, even in 
 the case of such difficultly fusible metals as iron and chromium. 
 Experiments illustrating this mode of reduction are easily 
 carried out. The aluminium used must be in a finely-divided 
 state. Filings, which are readily made from a piece of the 
 rod metal with a moderately coarse file, are suitable for opera- 
 tions on the small scale. The oxides in the form of fine 
 powder should be heated before use to insure their freedom 
 from moisture. 
 
 EXP. 110. Well mix together 10 grams of dry black oxide of 
 copper and 3 grams of aluminium filings ; put the mixture into a 
 small clay crucible ; make a small cavity in the centre of the 
 mixture with a rod, and into this pour a mixture of 2 grams of 
 barium peroxide and O'l gram of aluminium filings. Then insert 
 in this a piece of magnesium ribbon about 2 inches long in an 
 upright position. Place the pot on a brick, ignite the ribbon with 
 a lighted taper, and rapidly replace the lid. When the reaction is 
 finished and the crucible has cooled down, break it and extract the 
 button of metallic copper, which is usually buried in a mass of slag 
 at the bottom of the pot. Repeat the experiment with a mixture of 
 5 grams of dry ferric oxide and 2 grams of aluminium filings. The 
 
 13 
 
194 METALLURGICAL CHEMISTRY 
 
 same quantity of the barium peroxide mixture may be used. When 
 the pot has cooled, break it and extract the button of metal. 
 
 The change is a comparatively simple one, for it consists of 
 the oxidation of one metal at the expense of the oxygen in the 
 oxide of the other. 
 
 3CuO + 2 Al = A1 2 3 + 3Cu. 
 237 54 
 
 Fe 2 3 + 2 Al = A1 2 3 + 2Fe. 
 160 54 
 
 A slight excess of aluminium over the quantities found by 
 calculation from the equations is used in the experiments 
 described above, and part of this excess passes into the reduced 
 metal. If in the case of the ferric oxide an excess of that 
 compound is used, the reduction is only partial, and the tem- 
 perature obtained correspondingly lower. The mixture may 
 be arranged for reduction to ferrous oxide, thus : 
 
 3Fe 2 3 + 2 Al = 6FeO + A1 2 3 . 
 
 When this mixture is packed round the prepared joint of a 
 steel tube, which is to be hard soldered, and fired, the metal 
 is very rapidly raised to the temperature necessary to melt 
 the solder and make a sound joint. 
 
 This method of extracting metals from their oxides is useful 
 in the case of such oxides as those of chromium and manganese, 
 when comparatively small quantities of these metals are 
 required for special purposes. The aluminium used for their 
 reduction must be very finely divided, or the reduction is 
 incomplete. 
 
 The reducing action of metallic sodium when brought into 
 contact with fused chlorides of metals is very marked, and 
 was used for many years in the last stage of the extraction 
 of aluminium on the large scale. But the method has now 
 been superseded by the use of the electric current for the 
 extraction of the metal. The reaction depends upon the 
 
REDUCTION 195 
 
 readiness with which sodium combines with chlorine to form 
 common salt. The equation is : 
 
 A1C1 8 + 3Na = Al + 3NaCl. 
 
 Aluminium 
 chloride. 
 
 Potassium has similar reducing properties. 
 
 Some of the common metals will reduce others from thejr 
 combinations with sulphur when they are brought into contact 
 with the fused sulphides. The reaction with iron and lead 
 sulphide is a good example of this kind of change ; but the 
 mixture must be raised to a good red heat for effective 
 separation. 
 
 Reference should also be made to Exp. 70, and similar experi- 
 ments in Chap. VIII., as examples of reduction in which the 
 metal is displaced from its combination with a group of 
 elements. The reaction for the separation of copper from 
 copper sulphate by zinc is expressed thus : 
 
 CuS0 4 + Zn = ZnS0 4 + Cu. 
 
 Similar reactions are of great importance in a number of 
 metallurgical operations. In such cases the cheapest suitable 
 material is employed. Iron is much used. 
 
 Reduction by Potassium Cyanide. Simple reductions 
 
 can be brought about by heating some oxides and sulphides 
 with potassium cyanide, KCN, at a moderate temperature. 
 Great care must be taken in handling the cyanide, on account 
 of its highly poisonous nature. 
 
 EXP. 111. Put 3 grams of potassium cyanide into a porcelain 
 crucible, and melt it over a good Bunsen flame. When thoroughly 
 melted, add 1 gram of tin oxide, put on the lid, and continue the 
 heating for fifteen minutes. If the globules of reduced metal have 
 not run together, grip the side of the open crucible with a pair 
 of tongs, and gently shake it over the flame until a single globule 
 of metal is obtained. Pour the liquid contents of the crucible on to 
 a dry porcelain slab, and, when solid, detach the excess of cyanide 
 and weigh the button of metal. 
 
 132 
 
196 METALLUKGICAL CHEMISTRY 
 
 EXAMPLE. The button of metal from 1 gram of oxide weighed 
 0-782 gram. 
 
 /. 1 0*782=0-218 gram of oxygen removed. 
 
 Then = 0-0066; = 0-0136. 
 
 llo lo 
 
 What is the formula of the tin oxide used? (See p. 97) 
 The experiment by which this result was obtained was made 
 with practically pure tin oxide. The ordinary compound 
 (putty powder) is usually contaminated with oxide of lead, 
 and gives too high a result for the metal. The reaction is 
 thus expressed : 
 
 Sn0 2 + 2KCN = Sn + 2KCNO. 
 
 Potassium Potassium 
 
 cyanide. cyanate. 
 
 Sulphide of antimony (stibnite) may be reduced in a similar 
 manner ; but it is usually too impure to give a good quantita- 
 tive result. Oxide of lead is also easily reduced by the molten 
 cyanide. 
 
 Reduction by Reaction. This kind of reduction is 
 largely used in the separation of copper and lead from their 
 ores, and is very interesting on that account. It depends 
 upon the fact that if a mixture of compounds, containing the 
 elements of another compound, which can exist in the gaseous 
 state at a temperature below that at which the remaining 
 elements in the mixture are volatilized, is heated, the gas 
 forms and escapes from the mixture. The non-volatile element 
 is left behind in the free state. 
 
 EXP. 112. Mix together 1 gram of copper oxide, CuO, and 1 gram 
 of copper sulphide, Cu 2 S. Put the mixture into a small test-tube, 
 and heat it strongly in the Bunsen flame. Note that sulphur 
 dioxide is given off, and examine the residue when cold for particles 
 of finely divided copper. 
 
 The reaction is expressed by the equation : 
 
 EXP. 113. Mix together 1 gram of lead oxide and 1 gram of lead 
 sulphate, and heat as above. 
 
REDUCTION 197 
 
 The liberated sulphur dioxide is readily detected by its 
 odour, but the reduced metal is not so easily seen as the 
 copper in the residue from the last experiment. The equa- 
 tion is : 
 
 The value of these experiments is in the direct proof they 
 furnish of the formation of sulphur dioxide. In ordinary 
 metallurgical experiments the separation of the metals is 
 taken as an indirect proof of the way in which the reduction is 
 effected. 
 
 Reduction Step by Step. It is possible to reduce some 
 compounds in stages, and thus prove a definite relation between 
 the proportions of two elements in two or more distinct com- 
 pounds of the same. 
 
 EXP 114. Weigh 2 grains of lead peroxide, Pb0 2 ; transfer it to 
 the combustion-tube, Fig. 35, and weigh the whole. Heat the 
 powder until it turns yellow, but do not let it fuse. Allow the tube 
 to cool, and reweigh it. The loss is due to the escape of oxygen 
 expelled by heat. Replace the tube in the clip, connect the end 
 with the gas-tap, and reduce the yellow powder to metallic lea din a 
 stream of coal-gas. When the reduction is finished and the tube is 
 cold reweigh it. The loss is due to the removal of oxygen abstracted 
 by the constituents of the coal gas. 
 
 EXAMPLE. In a carefully worked experiment 2 grams of the 
 peroxide lost 0*133 gram on heating, and 0*129 gram on reduction 
 in coal-gas. Therefore 2 0'262 = 1*738 grams of lead are united 
 with 0133 and 0*129 gram of oxygen respectively. If these numbers 
 are divided by the atomic weights of the elements to which they 
 belong, the ratio 1 : 1 : 1 is obtained. 
 
 This proves that the oxygen is removed in two equal steps, 
 and makes it evident that the second compound may be 
 regarded as formed from the first by the addition of an amount 
 of oxygen strictly proportional to its atomic weight. This is 
 also true of similar series of compounds. Numerous examples 
 might be brought forward, and they would all bear testimony 
 in the same direction. The following are examples of such 
 series : CO, C0 2 ; Cu 2 0, CuO ; CrO, Cr 2 3 , Cr0 3 . 
 
198 METALLURGICAL CHEMISTRY 
 
 The generalization based upon these facts is known as 
 the law of chemical combination in multiple proportion (see 
 p. 52). 
 
 The series of compounds included in this generalization, 
 which, it must be remembered, is based entirely upon experi- 
 mental evidence, and is therefore a statement of facts, afford 
 very strong support to the atomic theory. For it is perfectly 
 clear that after a particular compound has been formed, a 
 second addition of one of its elements may take place, by 
 which another equally definite compound is produced. 
 Further, the amount of the added element is strictly propor- 
 tional to its atomic weight, for the second compound is 
 formed by a distinct jump from one proportion to the other. 
 It is, then, easy to imagine that the molecules of the higher 
 compound are formed by the addition of invariable particles 
 or atoms of the element to the molecules of the lower 
 compound. 
 
 SUMMARY. 
 
 Reduction is one of the most important operations in 
 metallurgical work. The selling price of a metal depends 
 largely upon the readiness with which it is reduced from its 
 ores. The reduction of a metallic compound may be complete, 
 and the metal set free ; or it may be only partial, and a 
 lower compound of the metal obtained. Only a few metals 
 are liberated from their compounds by heat alone. Assistance 
 has to be rendered by reducing agents, of which carbon, 
 carbonic oxide, and hydrogen are the most important for the 
 reduction of metallic oxides. The formation of carbon 
 dioxide and water is easily proved. The electric current 
 is coming more and more into prominence as a reducing 
 agent, and it is in use on the large scale for the extraction 
 of several metals. Metals themselves are not quite so 
 prominent in the list of reducing agents as they were, but 
 iron, copper, and zinc are still largely used in the metallurgy 
 of gold and silver. Potassium cyanide is only used on the 
 
KEDUCTION 199 
 
 small scale for the reduction of metals from oxides and 
 sulphides. Step-by-step reduction, and its bearing on the 
 second law of chemical combination, is of great theoretical 
 importance. 
 
 QUESTIONS. 
 
 1. What do you understand by the partial and complete 
 reduction of a compound ? Give examples. 
 
 2. How may a reduction experiment be used to determine 
 the formula of an oxide 1 
 
 3. Describe an experiment by which the composition of 
 water is determined. 
 
 4. What is the composition of coal gas T Write out equa- 
 tions showing the effect of its various constituents upon a 
 heated metallic oxide. 
 
 5. Write down the law of chemical combination in multiple 
 proportions. What is the nature of the support it lends to 
 the atomic theory ? 
 
 6. How is sulphur dioxide formed in a reduction process 
 
CHAPTER XII 
 COMBUSTION 
 
 IN Chap. XL some experiments are described in which metals 
 are reduced from their oxides by carbon and hydrogen re- 
 spectively ; and it is proved that carbon dioxide and water 
 are formed in the reactions which take place. From this it 
 is evident that metallic oxides can act as oxidizing agents, and 
 in doing so give up their oxygen to the bodies by which they 
 are reduced to the metallic state. But in order that these 
 changes may take place, it is necessary to heat the metallic 
 oxides in contact with the reducing agents above certain 
 limiting temperatures, below which no action is observed. 
 The reason for this is that the affinity of the oxygen for the 
 metal must be overcome before the carbon or hydrogen can 
 combine with it. The effect of the decomposing action of the 
 heat on the oxide, and the pull of the reducing agent upon 
 its oxygen, bring about the decomposition, when they are 
 together greater than the attraction of the elements in the 
 oxide for each other. Then the reducing agents burn in the 
 oxygen of the oxides. 
 
 When carbon and hydrogen oxidize freely, and the action 
 continues without the aid of external heat, the change is 
 commonly known as combustion : but the compounds formed 
 are the same as in the reduction of oxides. The heat developed 
 in their formation is, however, much greater than that used 
 up in rendering the changes possible, and so the action 
 continues as long as the necessary materials are supplied. 
 
 There is a certain limiting temperature for a given com- 
 
 200 
 
COMBUSTION 201 
 
 bustible body, below which it may be exposed to the air 
 without taking fire. This temperature is called the ignition 
 point, and varies for different bodies. If the body is exposed 
 to the air at this temperature, or a little above it, oxidation 
 commences, heat is developed, and a further rise in tempera- 
 ture takes place. This increases the action, and if the body 
 burns freely its maximum rate of combustion is soon reached, 
 and continues steadily as long as the conditions are constant. 
 The rate of combustion must depend upon the rate at which 
 the particles of oxygen are brought into contact with the 
 combustible body. Also, the local temperature depends 
 largely upon the rate of combustion, and this can be increased 
 by bringing more particles of oxygen into contact with the 
 burning body in a given time. This is commonly effected by 
 blowing, and the higher temperature obtained by burning the 
 combustible body in a rapid current of air is explained. A 
 further increase in the temperature takes place if the air is 
 heated before it comes into contact with the burning body ; 
 for the heat which would be used up in raising the products 
 of combustion and the residual nitrogen to the temperature 
 of the heated air, is now available for increasing the general 
 temperature. 
 
 It has been already stated that the quantity of heat liberated 
 during a chemical combination is just as definite as the 
 quantity of the compound formed. Thus the thermo-chemical 
 equation 
 
 12 32 44 
 
 means that 12 grams of carbon in burning completely to 
 carbon dioxide liberates sufficient heat to raise 97 kilograms 
 of water through 1 C. Similar thermal equations may be 
 written for other combustible bodies. 
 
 Further, it must be carefully noted that exactly the same 
 quantity of heat disappears if the compound formed during 
 combustion is decomposed into its elements again. And it is 
 also certain that there is a limiting temperature at which the 
 
202 METALLUKGICAL CHEMISTKY 
 
 compound is decomposed as fast as it is formed. This higher 
 limit is the temperature of dissociation of the compound, and 
 is the theoretical limit to the temperature obtained by the 
 combustion of the element forming the compound. The 
 practical limit, which is influenced by a variety of circum- 
 stances, seldom reaches this. Some assistance in understand- 
 ing this limit may be obtained by considering the experiment 
 with mercury, described on p. 17. Here it is evident that a 
 certain minimum temperature is necessary in order that the 
 mercury may oxidize ; but, if the oxide is exposed to a tem- 
 perature not much higher than that at which it is formed, 
 decomposition sets in, and the elements are reproduced. It is, 
 then, certain that there is a temperature at which the oxide is 
 decomposed as fast as it is formed, and that above this limit 
 no permanent oxidation is possible. Thus, however favourable 
 the conditions may be under which a body is burnt, there is a 
 limit to the temperature of the products of its combustion. 
 When air is used for burning a body, every molecule of 
 oxygen brought into contact with it has four molecules of 
 nitrogen in its train. Now, if these are replaced by oxygen 
 molecules, it is evident that a more rapid action must result. 
 This is what happens when pure oxygen is used in place of 
 air : it has the same general effect as a rapid current of air, 
 without the disadvantage of the presence of a considerable 
 body of nitrogen to absorb heat, and thus to lower the general 
 temperature. If, then, pure oxygen is used to surround the 
 burning body instead of air, the effect should be considerably 
 increased. This may be rendered perfectly clear by making a 
 few experiments with oxygen itself. 
 
 OXYGEN. 
 
 Preparation of Oxygen from Air (Erin's process). 
 This commercial process depends upon the fact that when 
 barium oxide, BaO, is heated to a certain temperature in a 
 current of air, it absorbs oxygen rapidly, and is converted into 
 
COMBUSTION 203 
 
 * 
 
 the peroxide, Ba0 2 , and that when this compound is heated 
 to a still higher temperature, it is decomposed into the lower 
 oxide again, oxygen being set free. It is another case of a 
 limiting temperature, above which the change is reversed. 
 But it is found that by reducing the pressure over the hot 
 peroxide the limiting temperature for decomposition may be 
 brought down to that at which absorption of oxygen takes 
 place under ordinary pressure. By taking advantage of this 
 the process is made an alternating one, by varying the 
 pressure instead of the temperature, which is much more 
 convenient. 
 
 Barium oxide is heated in iron tubes arranged in a furnace, 
 and a current of air, freed from carbon dioxide and water by 
 passing it through vessels containing caustic soda and lime, 
 is passed over it. Absorption takes place : 
 
 Barium Barium 
 
 oxide. peroxide. 
 
 The air current is then stopped, and the tubes carefully 
 exhausted by a pump, when, without interfering with the 
 temperature of the tubes, oxygen gas is liberated. 
 
 2Ba0 2 = 2BaO + 2 . 
 
 The first portion of the gas is allowed to escape, as it is con- 
 taminated with the nitrogen present in the tubes when the 
 air is cut off ; the remainder, which is pure oxygen, is passed 
 into a gas-holder as fast as it is liberated. The oxygen thus 
 obtained is forced in large quantities into strong portable iron 
 bottles, from which it can be drawn at will. The gas stored 
 in this way is very convenient for use at short intervals. 
 
 Preparation of Oxygen from Potassium Chlorate. 
 
 When small quantities of the gas are required at irregular 
 intervals, it is better to prepare it as required. 
 
 EXP. 115. Mix 10 to 20 grams of powdered potassium chlorate with 
 a quarter its weight of black oxide of manganese, and put the mixture 
 into a test-tube, 1 inch wide and 7 inches long, fitted with a bung and 
 
204 METALLUKGICAL CHEMISTKY 
 
 delivery-tube. Arrange the tube in a horizontal position, as shown on 
 p. 39, so that there is a clear channel between the mixture and the side 
 of the tube, and with the free end of the mixture not too near the 
 bung. Heat the tube carefully, beginning at the free end of the 
 mixture, and gradually carry the flame along as the escape of gas 
 slows down. Collect several jars of the gas, and one stoppered bell 
 jar of about 1,500 c.c. capacity. 
 
 One gram of the chlorate liberates about 270 c.c. of oxygen, 
 so that a rough calculation can be made for the quantity 
 required, and waste thus avoided. The change is given by 
 the equation : 
 
 2KC10 3 = 2KC1 + 30 2 . 
 
 Potassium Potassium 
 
 chlorate. chloride. 
 
 The black oxide undergoes very little change itself, but assists 
 the oxygen to escape at a lower temperature than when the 
 chlorate is used alone. 
 
 Properties Of Oxygen. It is a colourless, odourless, and 
 tasteless gas ; only slightly soluble in water, and the most 
 important supporter of combustion. The following experi- 
 ments may be made with the gas prepared as above : 
 
 1. When a glowing splint is put into the gas it is rekindled. 
 
 2. When a piece of good charcoal is made red hot in the 
 Bunsen flame, and then plunged into the gas, it burns vigor- 
 ously, emitting much heat and light, but without flame. The 
 residual gas turns lime-water milky ; it is carbon dioxide. 
 
 3. Sulphur when feebly ignited bursts into more vigorous 
 combustion, and burns with a violet flame. The residual gas 
 has the odour of burning sulphur, and turns litmus red. It 
 is the acid-forming sulphur dioxide. 
 
 4. When ignited phosphorus is put into oxygen it burns 
 with a very bright flame. The acid oxide P 2 5 is formed. 
 A common jar should be used for this experiment, as it is liable 
 to fracture from the great heat developed. 
 
 5. A piece of thin steel ribbon made into a spiral, and one 
 end tipped with a short piece of a wax match, burns very 
 vividly when the wax is ignited and the whole put into a 
 
COMBUSTION 
 
 205 
 
 bell -jar filled with oxygen, and standing over water. See 
 Fig. 38. 
 
 6. If two small candles supported in deflagrating spoons are 
 ignited and put into two cylinders, 
 one containing air and the other 
 oxygen, a very marked difference 
 is observed in both the rate of 
 combustion and the appearance of 
 the flames. In the air the candle 
 burns at the normal rate for a 
 short time, then the flame lan- 
 guishes, and finally disappears. In 
 the oxygen the flame shrinks, be- 
 comes much brighter, and the rate 
 of combustion is increased. An 
 
 increase in temperature is evident from the fact that the wax 
 melts, and runs down into the spoon. The flame finally dis- 
 appears. The formation of carbon dioxide during the com- 
 bustion is easily proved by pouring a little clear lime-water 
 
 into each jar; on shaking, the lime-water is turned 
 
 milky. 
 
 Combustion in Air and other Bodies. The 
 
 burning of a candle may be made the subject of an 
 experiment by which the increase in weight due to 
 the combination of oxygen with the constituents of 
 the combustible body may be proved. A lamp 
 chimney, Fig. 39, is divided into two parts by a piece 
 of wire gauze. The upper part is filled with a mix- 
 ture of caustic soda and calcium chloride in small 
 lumps, arid the upper end is fitted with a perforated 
 cork. Another perforated cork bearing a small candle 
 is inserted in the lower end. The whole is then hung 
 by a wire loop on to the end of the scale-beam of a balance 
 from which the scale-pan has been removed, and counter- 
 poised. The lower cork is then removed, the candle ignited, 
 
206 METALLURGICAL CHEMISTRY 
 
 and rapidly returned to the tube. Very soon the apparatus 
 begins to show an increase in weight, although the candle is 
 actually disappearing. The cause of this increase is readily 
 explained, if it is borne in mind that the carbon and hydrogen 
 of the candle combine with and thus abstract oxygen from the 
 air passing through the apparatus, and that the absorbing 
 material in the upper part prevents the escape of the pro- 
 ducts of the combustion. Thus the whole of the matter of the 
 candle is kept in the tube, together with the added oxygen. 
 This is evidently the cause of the increase in weight. 
 
 The formation of water during the combustion of hydrogen, 
 or combustible bodies containing hydrogen, is easily demon- 
 strated. The drying-tube of the part of the apparatus for 
 generating dry hydrogen is fitted with a piece of clay tobacco- 
 pipe, and the long tube passing into the wash-bottle is raised 
 just above the surface of the liquid (p. 189). The apparatus 
 shown in Fig. 19 is carefully dried, and the head of the thistle 
 funnel arranged over the clay jet, the angle-tube being con- 
 nected with an aspirator. The jet of hydrogen is then ignited, 
 and a current of air aspirated through the thistle funnel. 
 Very soon water appears in the bottle, and the only explana- 
 tion of its presence there is that it has been formed by the 
 burning of the hydrogen in air. A jet of coal gas or a burning 
 candle may be placed under the funnel with a similar result. 
 
 Combustion is not limited to air or oxygen, for hydrogen^ 
 combustible bodies containing hydrogen, and finely divided 
 metals burn readily in chlorine (p. 74). Nitrous oxide and 
 nitric peroxide are also supporters of combustion. Finely 
 divided metals will also burn in sulphur vapour. 
 
 EXP. 116. Put some sulphur into a wide test-tube ; fix the tube 
 in an upright position, and strongly heat it. When the sulphur is 
 boiling vigorously lower a spiral of thin copper ribbon into the 
 vapour. In a short time it becomes red hot, due to the rapid 
 sulphidation. The metal burns in the sulphur vapour. When the 
 spiral is withdrawn it is fouod to be changed into copper sulphide. 
 
 Explosive Combustion. In ordinary combustion the 
 
COMBUSTION 207 
 
 conditions are such that the burning body and the supporter 
 of combustion are supplied to each other steadily, and any 
 alteration in the rate of supply alters the rate at which the 
 action takes place. Explosion of a mixture of gases or solids 
 is simply rapid combustion, and takes place when the two 
 bodies are so intimately mixed that the entire action is spread 
 over a very short space of time. 
 
 The explosion of a mixture of coal gas and air is a familiar 
 example, and the energy of the change is unmistakable. 
 The working of an ordinary gas engine depends upon a series 
 of such explosions. Gunpowder is an example of an explosive 
 mixture of solids ; it is a case of the rapid oxidation of carbon 
 and sulphur by the oxygen obtained from the saltpetre. A 
 large quantity of gaseous matter is thus formed in a limited 
 space, and the enormous pressure which this gas 
 is able to exert easily accounts for the effects 
 produced. 
 
 Inversion of Combustion. The terms 
 
 combustible and supporter of combustion are 
 merely relative. In the ordinary way we say 
 that coal gas burns in air ; but it is easy to 
 invert this order of things, and make the air 
 burn in coal gas. A very simple experiment 
 will prove this. 
 
 EXP. 117. Fix a lamp chimney in a clip, insert a 
 cork, through which two tubes pass, as shown in 
 Fig. 40, in the bottom, and put a gauze cap on the 
 top. Connect the angle-tube a with the gas-tap, and 
 ignite the jet of gas escaping through the gauze. Tip 
 the end of a thin splint of wood with a fragment of 
 phosphorus, ignite it, pass the splint rapidly up the 
 straight tube b, and ignite the jet of air at the top. FlG - 40 - 
 
 The direct ignition of the jet of air as described, although 
 more convincing, is not always easy to bring about. Another 
 way is to close the top of the chimney with a cork, through 
 which a glass tube is passed, to turn on the gas freely, and 
 
208 METALLUEGICAL CHEMISTEY 
 
 to ignite it at both top and bottom tubes. If then the 
 supply is regulated carefully, the flame at the end of the 
 bottom tube will strike back up the tube, and appear at the 
 inner end as a jet of air burning in coal gas. 
 
 The current of coal gas passing in through a and up the 
 chimney draws in a current of air through b, so that a jet of 
 air is supplied to an atmosphere of coal gas. The air takes 
 fire and burns with a flame, which, though not so luminous as 
 the flame of coal gas burning in air, resembles it in general 
 appearance. 
 
 Similarly, oxygen may be burnt in hydrogen, and chlorine 
 in hydrogen. It can then be said that when two gases are 
 concerned in the act of burning either may be supplied in a 
 jet to the other as a supporting atmosphere. But the chemical 
 change is the same whichever way the action takes place. 
 Thus, for oxygen and hydrogen water is the invariable pro- 
 duct. If, however, one of the bodies is a non-volatile solid, 
 the inversion is not possible. 
 
 Flame. The burning of carbon and sulphur in oxygen 
 furnishes examples of combustion in which a marked difference 
 is noticed. The carbon becomes incandescent, and gives out 
 light and heat, but without flame. It is a non-volatile solid. 
 The sulphur, on the other hand, gets hot and vaporizes. This 
 vapour diffuses, penetrates the surrounding space for some 
 distance, and burns there. The space in which the vapour 
 burns is filled with flame, and in the whole of this region the 
 combustion is going on. In the first case the zone of combustion 
 is confined to the surface of the burning solid, and the heat 
 developed is concentrated in space. In the second the zone 
 of combustion is extended to the region about the burning 
 body, and the heat is distributed over a larger space. 
 
 The appearance of the space in which gaseous matter is 
 burning gives the general idea of flame, and the further the 
 vapour penetrates into the space before it is completely burnt 
 the larger will be the volume of the flame. Bearing this in 
 
COMBUSTION 209 
 
 mind it is easy to explain why a candle flame shrinks when it 
 is fed with pure oxygen. For the oxygen burns up the com- 
 bustible gases emanating from the wick of the candle before 
 they have time to penetrate very far into the surrounding 
 space, and a shrinking of the zone of combustion takes place. 
 The local temperature is increased by this, and also by the 
 more rapid burning. If it is required to localize the heat so 
 as to attain a high temperature in a small space, a combustible 
 body which preserves its solid state during the whole period 
 of its combustion should be used, and the gas in which it is 
 burnt should be supplied as fast as possible. On the other 
 hand, when it is necessary to spread out the heat, so to speak, 
 through a larger space, gaseous fuel must be used. Also, the 
 supply of air must be such as to allow of perfect combustion 
 in as large a region as possible. Charcoal, coke, and anthra- 
 cite are examples of fuel which, when burnt in a good supply 
 of air, give high local temperatures. But gas coal, or long 
 flaming coal, which contains a large percentage of volatile 
 combustible matter, is more suitable for the distribution of 
 heat. Gaseous fuel, however, may be used for obtaining a 
 high local temperature, if it is suitably burnt. The shrinking 
 of a candle flame in oxygen suggests the means to be adopted ; 
 and the oxy-hydrogen blow-pipe flame is an excellent example 
 of its application. The tendency to do away with the use of 
 solid fuel for all purposes to which combustible gas can be 
 applied is growing; and the use of gaseous fuel may be 
 regarded as a distinct scientific, economic, and hygienic 
 advance upon the dirty, wasteful methods of heat production 
 still so largely used. Producer gas, water gas, and Mond gas 
 are now quite familiar terms. 
 
 The Structure of Flame. The complexity of a given 
 flame varies with the constituents of the gas producing it, and 
 the conditions under which the gas is burnt. A general idea, 
 however, can be obtained by a few simple observations. If 
 the air- way of a Bunsen burner is stopped a luminous flame is 
 
 14 
 
210 METALLUEGICAL CHEMISTRY 
 
 obtained at the top ; and if the gas is turned down so as to 
 give a small flame without flickering, the flame is seen to 
 consist of three parts : (1) a blue portion, appearing to spring 
 from the top of the burner; (2) a luminous portion, which 
 occupies the upper part of the body of the flame ; (3) a faintly- 
 luminous mantle, which surrounds the whole body of the flame 
 and forms its limits. A dark portion, which does not belong 
 to the flame proper, is noticed in the lower part, and this 
 region is filled with unburnt gas. This is easily proved by 
 thrusting the head of a match rapidly through the envelope 
 into this space, and the fact that it does not ignite there is 
 sufficient evidence that combustion is not taking place in that 
 space. These observations are most successful in an other- 
 wise dark room. An ordinary candle flame has an exactly 
 similar appearance. The same conical space filled with un- 
 burnt gas is present in the ordinary Bunsen flame, and its 
 presence may be demonstrated by holding a piece of stiff" 
 white paper well down on the flame for a second or two, so 
 as to cut through this space. A black ring surrounding a 
 white centre shows where the paper has been in contact with 
 burning and non-burning gas. 
 
 The Luminosity of flame is of considerable importance 
 when it is a question of obtaining as much light as possible 
 from the burning gas. The white flame, which is so much 
 prized in illuminating gas, is not due to any single cause. 
 It is probably largely caused by the presence in that portion 
 of the flame of dense gases and solid particles, produced by 
 complex changes in the gas before its complete combustion. 
 An increase in the pressure under which a gas is burnt also 
 increases the luminosity of its flame, but this condition is not 
 commonly present to any extent. The hydrogen flame gives 
 out very little light, which is due to the extremely simple 
 character of its combustion; but if fine particles of solid 
 matter are passed into it its luminosity is at once increased. 
 This principle is adopted in converting the practically lightless 
 
COMBUSTION 211 
 
 flame of a Bunsen burner into a brilliant one by suspending 
 in it a mantle made of a very light infusible oxide, which is 
 rendered incandescent by the heat of the flame. The burning 
 of phosphorus, on the other hand, proves that it is possible to 
 have a highly luminous flame without the presence of solid 
 particles in it. For solid matter cannot possibly be present 
 in this flame. Also, the highly-luminous flame of burning 
 acetylene, C 2 H 2 , is probably largely due to the endothermic 
 character of the compound itself. Much heat is absorbed in 
 its formation, and this heat, which is liberated during its 
 decomposition, is added to that evolved by the combustion of 
 its constituents. A very high local temperature is thus 
 obtained with the other conditions for a luminous flame. 
 
 THE ENERGY OF COMBUSTION. 
 
 The Conservation of Energy. Since combustion is 
 the principal source of artificial heat, if it may be so called 
 to distinguish it from the natural heat derived from the sun, 
 it will be in place here to inquire a little into the nature of 
 heat. In the first place, heat is real, but it is not ponderable ; 
 that is, it lacks the intrinsic property of matter, which is best 
 denoted by substance or density. It has been defined as a 
 mode of motion ; but as we can form no idea of motion with- 
 out something to move, it does not appeal to us until it is 
 associated with bodies of some kind. Heat motion in bodies 
 is a movement of their particles, and not of the bodies as a 
 whole. A full discussion of the nature of heat will be found 
 in books on the subject ; but, for our present purpose, it is 
 sufficient to recognise its ability to cause an increase in mole- 
 cular motion when it enters bodies, and a decrease in that 
 motion when it leaves them. 
 
 When a body is moving it is said to possess energy, 
 for it is capable of doing work. This kind of energy is 
 kinetic or " moving " energy. Therefore, since we have no 
 experience of absolutely cold bodies, such bodies as we know 
 
 142 
 
212 METALLUKGICAL CHEMISTKY 
 
 possess a store of this moving energy, due to their molecular 
 motion. 
 
 But bodies may also possess energy due to their position 
 with respect to other bodies. Thus, they may be able to fall 
 and, in falling, to do work. This kind of energy is potential, 
 or " stored " energy. The potential energy of a body appeals 
 very strongly to the mind at times. If one is looking up at 
 a piece of overhanging rock on the top of a cliff, the feeling 
 will arise' as to what would happen if it were to fall. The 
 same piece of rock lying at one's feet does not create the same 
 impression. When a body is falling its potential energy, which 
 depends upon its position, is being rapidly converted into 
 kinetic energy, and just before it reaches the surface towards 
 which it is falling, all the potential energy due to its original 
 position with respect to that surface is converted into moving 
 energy. By impact with the surface the body is brought 
 to rest, and it would appear as if its energy of motion were 
 destroyed ; also, since it no longer possesses the potential 
 energy due to its former position, that a complete disappear- 
 ance of energy had taken place. But further examination 
 proves that the energy of motion of the mass is not destroyed ; 
 it is converted into increased motion of the particles of the 
 body and the surface struck by it. That is, it is converted 
 into heat. Further, the heat energy developed is exactly 
 equivalent to the kinetic energy lost. The same reasoning 
 holds for every modification of energy. This principle, which 
 is known as the conservation of energy, has been most 
 exhaustively studied in connection with the relation between 
 heat and mechanical work, and in the case of these two modifi- 
 cations of energy has been proved up to the hilt. Those who 
 are most qualified to give an opinion hold that the same is true 
 for every other modification of energy. 
 
 Every particle of matter has energy associated with it, and 
 when a redistribution of the particles of a body, or system of 
 bodies, takes place, there is also a redistribution of the energy 
 associated with them. The energy of chemical attraction is of 
 
COMBUSTION 213 
 
 the potential kind, and when the conditions are such that this 
 attraction is able to bring about chemical change in the matter 
 of the reacting bodies, a change in energy also usually results. 
 Take, for example, the case of a mixture of oxygen and 
 hydrogen : the two gases are ready to combine with each other, 
 and a considerable amount of energy is stored up in their 
 molecules. When the mixture is exploded a redistribution of 
 the atoms of the gases to form molecules of water takes place, 
 and this is accompanied by rapid conversion of the potential 
 energy of chemical attraction into the kinetic energy of the 
 heat evolved. The cause of the force exerted between 
 dissimilar bodies when chemical changes are affected, and the 
 nature of chemical energy, are quite unknown, but this does 
 not prevent us from using ideas suggested by the results. The 
 conversion of chemical energy into heat resembles very much 
 the conversion of the energy of a falling body into heat when 
 it is brought to rest by collision with the earth ; and it does 
 not appear at all improbable that the change in the motion of 
 the atoms of oxygen and hydrogen^ in their endeavour to form 
 new and more stable molecules is the cause of the conversion of 
 the stored chemical energy into heat. The action going on 
 when a piece of red-hot charcoal is surrounded by oxygen or 
 air may be pictured in the same way. The atoms of oxygen 
 become associated with the atoms of carbon, and form molecules 
 of carbon dioxide, and the transformation of energy goes on 
 vigorously. During the rush and turmoil of the action, 
 chemical energy disappears, and heat energy appears in the 
 same proportion. 
 
 No difficulty should now be experienced in explaining the 
 fact that different modifications of carbon appear to give out 
 different quantities of heat when they are completely burnt. 
 The different forms of carbon are admittedly due to a different 
 arrangement of the molecules or atoms of the element, and not 
 to a difference in the quality of the atoms themselves. It is, 
 therefore, quite feasible that the atoms of carbon in graphite 
 require more to bring them up to the condition necessary for 
 
214 METALLURGICAL CHEMISTEY 
 
 union with oxygen than do the similar atoms of carbon in 
 charcoal, and that the energy required to do this work of 
 preparation is subtracted from the total quantity of heat, 
 which would be developed by the burning of the free atoms of 
 carbon. In fact, it is analogous to the case of a body which 
 has to be lifted into a position from which it can fall. Here the 
 energy used up in lifting the body into position is to be 
 subtracted from the total heat developed by the impact of the 
 body when it is brought to rest. It may, then, be said that the 
 distribution of energy in the allotropic modifications of carbon 
 is not the same. This holds for sulphur too ; for the monoclinic 
 form of that element evolves heat when its crystals pass into 
 the more stable rhombohedrons (p. 57). Generally, when an 
 element passes from a more stable to a less stable form, heat is 
 absorbed, but heat is evolved in the opposite change. When a 
 group of molecules of different bodies are reacting among 
 themselves, and forming molecules of a new body or bodies, 
 evolution or absorption of heat takes place according as the 
 new molecules are more or less stable than the reacting mole- 
 cules. When heat is evolved the change is exothermic, 
 
 and when it is absorbed the change is endothermic. All 
 
 the common combustion processes are exothermic in character ; 
 and a number of changes will be met with in metallurgical 
 operations, which are distinctly endothermic in character. 
 They do not take place unless heat is supplied from without. 
 Energy then, as we know it, is a condition of matter, and like 
 matter can undergo a variety of changes, and, in any self- 
 contained system of bodies, the amount of energy, like the 
 amount of matter, is absolutely fixed. 
 
 MEASUREMENT OF HEAT. 
 
 When a given quantity of heat enters a body and distributes 
 itself uniformly through the substance of the body, a change 
 takes place in its thermal properties, and the temperature of 
 the body is said to rise. This change of temperature depends 
 
COMBUSTION 215 
 
 upon (a) the mass of matter in the body, and (b) its power of 
 absorbing heat. The latter property is termed its specific 
 heat, and is dependent upon the kind of matter concerned in 
 the absorption. Specific heat and temperature are regarded as 
 the two factors of heat energy, and are measured by reference 
 to arbitrary standards. The temperature of a body indicates 
 the intensity of its heat energy, and is the factor which 
 determines the possibility of heat entering or leaving the body 
 under given conditions. Thus if two bodies at different 
 temperatures are brought together, heat will pass from the 
 body at the higher to the body at the lower temperature, and 
 the exchange will go on until the intensity of the heat energy 
 is the same in both. The temperature of the one body rises 
 and that of the other falls, until it is the same for both. But 
 the quantity of heat which leaves the one body arid enters the 
 other also depends upon their specific heats i.e., upon the 
 other factor of heat energy. Thus if a number of bodies at 
 different temperatures are brought together and left to them- 
 selves, exchange of heat energy will take place between them, 
 and they will finally reach one dead level of temperature, the 
 value of which will depend upon the masses, specific heats, and 
 temperatures of all the bodies in the system. It is assumed in 
 this general statement that no process other than redistribution 
 of heat energy takes place in the system. 
 
 Measurement of Temperature. Advantage is taken of 
 
 the fact that bodies in general expand when heated and con- 
 tract when cooled, and that this expansion and contraction 
 may within certain limits be taken as a measure of the rise 
 and fall of their temperature, and, therefore, of that of other 
 bodies with which they are in contact. The instruments used 
 for registering moderate temperatures are called thermometers, 
 and those for high temperatures pyrometers. The mercurial 
 thermometer consists of a glass bulb and tube nearly filled 
 with mercury. The relation between the capacity of the bulb 
 and the diameter of the bore is such that the metal does not 
 
216 METALLURGICAL CHEMISTRY 
 
 quite leave the bore at the lowest, and does not quite fill it 
 at the highest temperature the instrument is intended to 
 register. The fixed points of the scale are obtained by first 
 immersing the instrument in melting ice, and then in the 
 steam from water boiling at normal pressure. The point on 
 the stem to which the mercury recedes when exposed to the 
 temperature of melting ice is marked as the freezing-point of 
 water, for ice melts and water freezes at the same constant 
 temperature ; the point on the stem to which the metal 
 expands when exposed to the steam is marked as the boiling- 
 point of water, for the steam from water boiling at a constant 
 pressure has a constant temperature. The space between these 
 points is divided into degrees, which are marked on the scale. 
 
 On the Centigrade scale there are 100 between the freezing 
 and boiling points. Thus the freezing-point is marked C., 
 and the boiling-point 100 C. The scale may, however, be 
 graduated above and below these points ; the lower limit is 
 the freezing-point ( - 39 C.), and the upper limit the boiling- 
 point (350 C.) of the mercury itself. But when the tempera- 
 ture of the metal approaches either of these limits the indica- 
 tions are uncertain, and it may be said generally that the rate 
 of expansion of a body is not strictly proportional to its 
 increase in temperature, and the hotter it gets the greater is 
 its deviation, so that in accurate instruments corrections have 
 to be made for this deviation. In the case of mercury the 
 deviation is small between and 100, and can be corrected 
 for at higher temperatures, so that some mercurial thermo- 
 meters are constructed to register upwards of 300 C. 
 
 The temperature scale commonly used in the British Isles 
 is the Fahrenheit, on which the freezing-point is marked 32, 
 and the boiling-point 212, thus giving 180 between the two 
 points. The conversion of a reading on the Fahrenheit scale 
 to the corresponding reading on the Centigrade is easily 
 
 "Hi -t o f\ f\ f\ pr 
 
 effected, for ^ = VQQ = -= ; therefore F = C, and C = g F ; but 
 in converting F. into C. the 32 below the freezing-point must 
 
COMBUSTION 217 
 
 be subtracted before, and in the conversion of C. into F., added 
 after, the calculation is effected. This is given by the expres- 
 sions : 
 
 C = (F-32)J;F = f C + 32. 
 
 To determine the temperature of a given body, the ther- 
 mometer is brought into contact with it, the two allowed to 
 come into equilibrium, and the reading then taken. The 
 mass of the thermometer is kept as small as possible, so that 
 the quantity of heat entering or leaving it while coming into 
 equilibrium with the body may not sensibly interfere with 
 the temperature to be measured. 
 
 The construction of pyrometers for the measurement of high 
 temperatures is based upon various principles, such as the 
 expansion of solids and gases ; the increase in the electrical 
 resistance of a metal wire when heated ; the current generated 
 by a thermo-electric junction when its temperature is in- 
 creased, etc. ; but standard works on metallurgy, heat, and elec- 
 tricity must be consulted for a description of such instruments. 
 
 Specific Heat. It is a matter of common experience that 
 different weights of the same substance at a given tempera- 
 ture absorb or give out different quantities of heat when they 
 are raised or lowered through the same range of temperature, 
 and that these quantities are strictly proportional to the 
 masses of the substance under consideration. Thus, 2 pounds 
 of water at 50 C. would liberate twice as much heat in falling 
 to 20 C. as 1 pound would under the same conditions. Now, 
 if the heat given out by a given weight of water in cooling 
 through a given range of temperature is absorbed by a given 
 weight of iron, it is found that 3*57 times that weight of lead 
 is required to absorb the same quantity of heat ; or a quantity 
 of heat which would raise 100 grams of iron through, say, 10 C. 
 would raise 357 grams of lead through the same temperature. 
 This difference in the absorbing power of the two metals is 
 said to be due to the difference in their specific heats. The 
 
218 METALLURGICAL CHEMISTKY 
 
 heat-absorbing power of water is greater than that of any 
 other compound or element, and is taken as the standard. 
 The specific heat of water = 1, so that the specific heats of all 
 other pure substances is less than unity when compared with 
 this standard. 
 
 The method of determining the specific heats of solid bodies 
 such as metals may be thus described : A known weight of 
 water, W, is placed in a vessel called a calorimeter, and its 
 temperature, T, taken ; a known weight of the body w, which 
 has been raised to a temperature, f, is then immersed in the 
 water and moved about until the two come to the same 
 temperature, and the temperature of the water, Tj , a'gain 
 taken. Now, while the temperature of the water rises from 
 T to Tj , the temperature of the body falls from t to T^, 
 and the heat given out by the body is equal to that absorbed 
 by the water ; but the quantity of heat given out by the body 
 is proportional to its mass, specific heat, and fall in tempera- 
 ture, and the quantity absorbed by the water to the same 
 factors ; but these quantities of heat are equal, therefore 
 
 (t - T^) x w x s = (Ty-T ) x W x S, 
 
 where s and S are the specific heats of the body and water. 
 But as S = 1 i.e., the specific heat of water s, the specific 
 heat of the body, is the only unknown quantity, and its value 
 is easily found from the equation. 
 
 EXAMPLE. A lump of platinum weighing 2 ounces is raised to a 
 temperature of 410 C., and is then plunged into 20 ounces of water 
 at 16 C. ; the resulting temperature of the water is found to be 
 17'3 C. What is the specific heat of platinum ? 
 
 T 1 =410- 17-3 = 892-7, and T 1 -T = 17'3-16 = 1 3. 
 392-7x2xs = l'3x20xl; 
 1-3x20 nAOO 
 r S = S9^x2=- 088 - 
 
 A correction has to be made for the heat absorbed by the 
 calorimeter, but this and other points in connection with the 
 accurate determination of specific heats cannot be discussed 
 here. The specific heat of platinum between C. and 100 C. 
 
COMBUSTION 219 
 
 is 0*0323, and increases slightly at higher temperatures. 
 The definite relation between the specific heats and atomic 
 weights of metals is discussed on p. 128. 
 
 The unit of heat quantity is the calorie, which is the amount 
 of heat required to raise the temperature of 1 gram of water 
 from C. to 1 C. The great calorie or kilogram unit, which 
 is equal to 1,000 calories, is often used for practical work; 
 also the quantity of heat required to raise 1 pound of water from 
 32 F. to 33 F., or from C. to 1 C., may be taken as the 
 unit. 
 
 The definition of unit heat quantity makes it evident that 
 the measurement of heat energy depends upon the specific 
 heat of the absorbing body, and the temperature through 
 which it is raised, so that specific heat and temperature are 
 the factors to be considered in heat measurement. 
 
 When a solid changes to the liquid state or a liquid to the 
 gaseous state, heat is absorbed without change in temperature, 
 but the heat reappears when the inverse change takes place. 
 The heat absorbed by melting ice is 80 calories, and that 
 absorbed by boiling water is 537 calories. These numbers are 
 termed the latent heats of fusion of ice, and of vaporization 
 of water, for they represent the quantities of heat used up in 
 changing 1 gram of ice at C. to water, and 1 gram of water 
 at 100 C. to steam at the same temperatures. 
 
 SUMMARY. 
 
 Carbon and hydrogen are the principal combustible bodies, 
 and air is the important supporter of combustion. Carbon 
 dioxide and water are formed. The rate at which heat is 
 developed depends upon the rate at which the chemical 
 changes take place. A rapid current of heated air, or pure 
 oxygen, is used when a high local temperature is required. 
 Flame is the general result of burning a combustible gas. 
 The heat energy of combustion is the measurable equivalent of 
 the energy of chemical attraction which existed between the 
 
220 METALLUKGICAL CHEMISTRY 
 
 combustible body and the oxygen of the air used in its com- 
 bustion. This is generally true for all chemical changes by 
 which heat is developed. An exothermic change is when heat 
 is developed as a result of the change ; and an endothermic 
 change is when heat is absorbed as the change takes place. In 
 the one case energy is given out, and in the other it is 
 absorbed by the reacting bodies. Specific heat and tempera- 
 ture are the two factors of heat energy. 
 
 QUESTIONS. 
 
 1. What is understood by the ignition-point of a com- 
 bustible body ? 
 
 2. Give a short description of two processes for the prepara- 
 tion of oxygen. 
 
 3. What has been omitted from the equation : C + 2 = C0 2 ? 
 
 4. Describe and explain three characteristic experiments 
 with oxygen gas. 
 
 5. How would you prove that a combustible body is not 
 destroyed as it burns away 1 
 
 6. Explain the term "inversion of combustion." 
 
 7. Explain how you would localize or distribute the heat to 
 be obtained by burning a given sample of fuel. 
 
 8. Explain the terms "flame," "luminosity," and "flame 
 structure." 
 
 9. Give a description of the Bunsen burner. 
 
 10. Give a short account of the principle of the conservation 
 of energy. 
 
 11. Explain the terms "temperature" and "specific heat." 
 
 12. How is the specific heat of a metal determined ? 
 
CHAPTER XIII 
 PHOSPHOEUS AND ITS COMPOUNDS 
 
 WHEN bones are heated strongly out of contact with air, in a 
 similar manner to coal or wood (p. 178), liquid and gaseous 
 matters are given off, and a black residue is left in the retort. 
 This is bone-black, and is used for various purposes. If it 
 is completely burnt in a current of air, a white residue known 
 as bone-ash is obtained, which consists very largely of calcium 
 phosphate, Ca 3 (P0 4 ) 2 . This substance is used in the metal- 
 lurgical operation of cupellation to absorb molten oxide of 
 lead, and also for the extraction of phosphorus. Calcium 
 phosphate is a constituent of the older rocks, and, when they 
 are broken down in the slow process of soil formation, remains 
 in the soil. It is gradually converted into a soluble form, and 
 is then extracted by growing plants, which assimilate it, and 
 pass it on to animals using them for food. 
 
 Extraction of Phosphorus. Bone-ash when digested 
 with sulphuric acid is converted into calcium sulphate and 
 phosphoric acid. This acid is soluble in the remaining acid 
 liquor, but the calcium sulphate is practically insoluble, and 
 separates out in the solid state. After the acid liquid has been 
 run off it is evaporated to a syrup in an iron vessel, mixed 
 with charcoal powder, and further heated to a moderately 
 high temperature. The dry mixture is then transferred to 
 clay retorts arranged in a furnace, and there very strongly 
 heated. Phosphorus vapour distils over, and is collected in 
 vessels containing water. The crude phosphorus thus obtained 
 
 221 
 
222 METALLURGICAL CHEMISTEY 
 
 is purified and cast into sticks ready for use. The changes are 
 expressed by the equations : 
 
 Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 PO 4 . 
 
 Natural phosphates are found in some parts of the world in 
 considerable quantities. Apatite and phosphorite, both of 
 which contain calcium phosphate, are the most plentiful, and 
 may be used for the extraction of phosphorus. 
 
 Electric Smelting*. Recent improvements in the electric 
 furnace have made it possible to smelt either a natural 
 phosphate or bone-ash, when mixed with charcoal, for the 
 direct production of phosphorus. The preliminary treatment 
 with sulphuric acid is thus avoided, and a saving effected. 
 The extremely high temperature of the furnace makes the 
 reduction possible. 
 
 Properties Of PhOSphOPUS. Ordinary phosphorus is an 
 amber-yellow, waxy-looking solid, which melts at 44 C., and 
 boils at 290 C. It takes fire very readily, and should as a 
 rule be handled and kept under water. It is readily dissolved 
 by carbon bisulphide, CS 2 . When this solution is evaporated 
 the phosphorus is left in such a finely divided condition that 
 it bursts into flame spontaneously. It is also dissolved by a 
 solution of caustic soda or potash ; but in this case a reaction 
 takes place by which the gas phosphoretted hydrogen, PH 3 , 
 is formed and liberated. This gas when impure bursts into 
 flame directly it comes into contact with the air. It is 
 poisonous, and has the odour of decaying fish. 
 
 When yellow phosphorus is heated for some time to a 
 temperature about 240 to 250 C. out of contact with oxygen 
 it is gradually converted into red phosphorus, which is a 
 dark-red powder, insoluble in carbon bisulphide, CS 2 , and 
 perfectly safe to handle ; but when heated in contact with air 
 it bursts into flame. It is therefore a very different body in 
 appearance and properties from the yellow variety, although 
 
PHOSPHORUS AND ITS COMPOUNDS 223 
 
 it is the same chemical substance. The two are allotropic 
 modifications of the element. 
 
 Phosphorus unites readily with metals to form phosphides, 
 which are analogous to the sulphides in general composition ; 
 but they are not so readily obtained in the pure state. Some 
 oxides when heated are attacked by phosphorus, and the 
 phosphide of the metal is formed. This is the case with 
 red-hot lime. Two important commercial materials are phos- 
 phor tin and phosphor copper, which contain phosphides of 
 the metals. When lead is melted with an excess of phosphorus, 
 the metal absorbs more of the non-metal than it can hold 
 when in the solid state. This excess separates as crystals 
 when the mass solidifies. The lead is easily dissolved by 
 dilute nitric acid, and the phorphorus left behind as a black 
 residue. This resembles red phosphorus closely, except in 
 colour. It is, in fact, another allotropic modification of the 
 element. It may be noticed here that the association of 
 phosphorus with lead brings about a modification of the 
 former element ; and an analogy may be drawn between this 
 change and the conversion of amorphous carbon or diamond 
 carbon to the graphitic form by association with iron under 
 certain conditions (p. 167). 
 
 Phosphorus is one of the common impurities in commercial 
 iron, and passes into the metal from phosphates present in the 
 ores, and reduced by carbon at the high temperature of the 
 smelting furnace in which the metal is extracted. 
 
 When phosphorus burns in oxygen, or in a good supply of 
 air, phosphoric oxide, F 2 ^5' ^ s formed ; but when the air is 
 limited and the combustion slow, a large proportion of the 
 product consists of lower oxides, of which phosphorous oxide, 
 P 2 3 , is the most important. These oxides form solutions 
 which have a sour taste, and turn litmus red. They are, 
 therefore, acid-forming oxides. 
 
 Phosphoric oxide is a white, powdery, volatile solid, which 
 rapidly absorbs moisture, and deliquesces on exposure to the 
 air. When dissolved in cold water an acid is formed, which 
 
224 METALLURGICAL CHEMISTRY 
 
 unites with only one proportion of a base, and forms one series 
 of salts. It resembles nitric acid in general composition, and 
 is known as metaphosphoric acid, HP0 3 . But when the 
 solution is boiled another acid is formed, which is capable of 
 giving three well-defined series of salts. It is tribasic (p. 153), 
 and is known as orthophosphoric acid, H 3 P0 4 . It is the 
 normal or common phosphoric acid, and is readily prepared by 
 adding red phosphorus, a little at a time, to warm concentrated 
 nitric acid as long as the solid is taken up The solution is 
 then evaporated to a syrup, which is the practically pure acid. 
 It may be dissolved in water, and used for the preparation 
 of phosphates. The common prepared salt is disodium 
 hydrogen phosphate, Na 2 HP0 4 , 12H 2 0. When this salt is 
 heated it loses its water of crystallization first, and then at a 
 higher temperature parts with another molecule of water, and 
 is converted into a salt of pyrophosphoric acid, H 4 P 2 7 . The 
 last change is expressed thus : 
 
 2Na 2 HP0 4 = Na 4 P 2 7 + H 2 O. 
 
 Orthophosphate. Pyropho;-phate. 
 
 It would appear, then, that phosphoric oxide combines with 
 water in three proportions to form three distinct acids, thus : 
 
 1. H 2 + P 2 5 = 2HP0 3 . 
 
 Metaphos- 
 phoric acid. 
 
 2. 2H 2 + P 2 5 = H 4 P 2 7 . 
 
 Pyrophos- 
 phoric acid. 
 
 3. 3H0 + P0 = 2HP0 4 . 
 
 Orthophos- 
 phoric acid. 
 
 They all give corresponding salts with metallic bases. The 
 natural and prepared phosphates are for the most part ortho- 
 salts. The salts of the other acids are not often met with in 
 ordinary work. The phosphates of the alkaline metals and 
 ammonium are soluble in water, and three distinct salts of 
 each base can be obtained, if necessary. The three salts of 
 sodium are NaH 2 P0 4 , Na 2 HP0 4 , Na 3 P0 4 . Also, it is possible 
 
PHOSPHORUS AND ITS COMPOUNDS 225 
 
 to have two or three different bases in the same salt. Thus 
 the well-known microcosmic salt has the formula NaNH 4 HPO, t . 
 It is sodium ammonium hydrogen phosphate, and forms well- 
 defined crystals. 
 
 Phosphates of other metals are for the most part insoluble 
 in water, but soluble in acids, and can be obtained by mixing 
 together solutions of a soluble phosphate and of a soluble salt 
 of the other metal. The insoluble phosphate is precipitated, 
 and can be separated from the solution by filtration. 
 
 Phosphorous acid, H 3 P0 3 , gives a series of salts called phos- 
 phites, which are not of particular importance. 
 
 Phosphorus unites directly with chlorine, bromine, and 
 iodine, forming important compounds, which are treated fully 
 in advanced works on chemistry. 
 
 SUMMARY. 
 
 Phosphorus is not found in the free state in Nature. Its 
 principal natural compound is calcium phosphate, Ca 3 (P0 4 ) 2 . 
 The element exists in at least two distinct allotropic forms. 
 It combines readily with metals to form phosphides. Phosphor 
 tin and phosphor copper are largely used in the preparation 
 of alloys. The principal oxygen compounds of the element 
 are phosphoric oxide, P 2 O 5 , and phosphoric acid, H 3 P04. 
 Phosphates of the common metals are easily formed. 
 
 QUESTIONS. 
 
 1. Give a short description of the usual method for obtaining 
 phosphorus. 
 
 2. Name the common forms of phosphorus, and state how 
 they differ from each other. 
 
 3. What happens when phosphorus burns' in air ? 
 
 4. Name the three acids corresponding to phosphoric oxide. 
 To which do the common phosphates belong 1 ? 
 
 5. If you were required to obtain an insoluble phosphate of 
 a metal, how would you proceed 1 
 
 15 
 
CHAPTER XIV 
 
 SILICON AND ITS COMPOUNDS WITH OXYGEN 
 AND METALS 
 
 THE element silicon, although not often met with in the pure 
 state, is present in abundance in the earth's crust, combined 
 with oxygen and metals. Its most important compound is 
 the dioxide, silica, Si0 2 , which in the prepared state is a 
 perfectly white powder. One natural form of this oxide, 
 quartz, is often found in distinct crystals, which are sometimes 
 as transparent as glass. It is an acid-forming oxide, but as it 
 is quite insoluble in water under ordinary conditions, this 
 cannot be proved in the usual way. It is also very refractory 
 in character, requiring a very high temperature to soften and 
 fuse it. 
 
 Silica is not decomposed when strongly heated with carbon 
 unless a metal is present into which the reduced silicon can 
 pass as soon as it is liberated from the oxide. Thus, when 
 a mixture of silica, charcoal, and iron is very strongly heated, 
 silicon is reduced and passes into the metal. This is the source 
 of silicon in commercial iron, which always contains that 
 element, though sometimes in very small quantity. On the 
 other hand, a specially prepared iron may contain as much as 
 18 per cent, of silicon. The non-metal is no doubt combined 
 with part of the metal in the form of iron silicide, which 
 considerably modifies the properties of the remainder of the 
 metal with which it is associated. When iron containing 
 silicon is dissolved in an acid, and the solution evaporated to 
 dryness, a residue containing insoluble silica is obtained. 
 
 226 
 
SILICON AND ITS COMPOUNDS 227 
 
 Silica, when finely divided, is readily dissolved by a solution 
 of caustic soda or potash, and an alkaline salt of silicic acid 
 is formed. With caustic soda sodium silicate, Na 4 Si0 4 , is 
 obtained : 
 
 4NaHO + Si0 2 = Na 4 Si0 4 + 2H 2 O. 
 
 A similar change takes place when any form of silica in a 
 finely-divided state is fused with sodium carbonate in a 
 platinum crucible. The reaction proceeds slowly, and carbon 
 dioxide is liberated : 
 
 2Na 2 C0 3 + Si0 2 = Na 4 Si0 4 + 2C0 2 . 
 
 Fusion mixture, which is a mixture of sodium and potassium 
 carbonates, is more effective than the single carbonate. The 
 fused mass can be extracted with hot water, which dissolves 
 out the soluble silicate together with the excess of carbonate. 
 When dilute hydrochloric acid is added, in small quantities at 
 a time, to this solution the excess of carbonate is decomposed 
 without the silicic acid separating, and an excess of hydrochloric 
 may be thus introduced. This decomposes the silicate by 
 converting its metal into a chloride, and silica in combination 
 with water is retained in solution. 
 
 If this acid solution is poured into a tray fitted with a 
 bottom made of parchment paper, and floating on water in 
 a good-sized vessel, it is found that the excess of hydrochloric 
 acid, together with the whole of the sodium in the form of 
 chloride, passes through the parchment into the water under- 
 neath, so that nothing but silica is left in solution in the 
 tray. There is reason to believe that the silica is combined 
 with water in the form of silicic acid, which will not pass 
 through the parchment. But, although the solution can be 
 concentrated considerably, the pure acid has not been prepared. 
 In this respect it may be compared with the solution of 
 carbon dioxide, which is supposed to contain carbonic acid, 
 H 2 C0 3 . 
 
 Graham, who discovered and investigated the properties 
 of bodies taken advantage of in the above separation, gave 
 
 152 
 
228 METALLURGICAL CHEMISTRY 
 
 the name of colloids to those bodies which, when in solution, 
 will not pass through the parchment paper, and the name of 
 Crystalloids to those which will so pass. He also gave the 
 name of dialysis to the operation of separating crystalloids 
 and colloids when together in solution. Substances resembling 
 glue are colloids. 
 
 Silicic Acids. There is evidence in favour of the possible 
 existence of two silicic acids : orthosilicic acid, H 4 Si0 4 , and 
 metasilicic acid, H 2 Si0 3 . It is readily seen that the first is 
 tetrabasic, and is formed by the combination of silica with 2 
 molecules of water ; the second is dibasic, and contains only 
 1 molecule of water. 
 
 Silicates. The salts of sicilic acid are very numerous, and 
 are found in great variety in the earth's crust. They may 
 be formed by the direct combination of silica with basic 
 oxides. The facility with which a particular silicate is 
 formed depends upon its fusibility. But if a mixture of 
 silica and a basic oxide in the proper proportions to form a 
 certain silicate is raised to or a little above the melting-point 
 of the silicate, the combination will take place. In fact, in 
 some cases the combination takes place when the mixture is 
 exposed to a temperature below its fusing-point. This is known 
 as fritting" combination. In this case the two oxides must 
 be very finely divided and intimately mixed together. 
 
 EXP. 118. Weigh 15 grams of lead oxide, and mix with it the 
 necessary weight of fine, hard sand (silica sand), as calculated from 
 the equation given below. Put the mixture into a small clay 
 crucible, and raise it to a bright-red heat in a gas muffle, or wind 
 furnace. When thoroughly melted, pour the contents of the crucible 
 on to an iron plate. The resulting silicate when cold is an amber - 
 coloured glass. 
 
 The two oxides unite directly at a moderate temperature to 
 form lead silicate : 
 
 2PbO + Si0 2 =2PbO.Si0 2 . 
 2(207 + 16) 
 
SILICON AND ITS COMPOUNDS 229 
 
 The mode of writing the formula of a silicate adopted in the 
 equation is the one most used, and is justified by the way in 
 which the combination takes place. 
 
 EXP. 119. Mix 5 grains of finely-ground red copper scale, Cu 2 0, 
 with the necessary quantity of silica sand to form the silicate 
 Cu.jO.Si0 2 . Heat the mixture on a clay dish in a gas muffle to a 
 moderate red heat for half an hour. Allow the dish to cool, and 
 grind up a little of the residue in a mortar. When examined with 
 a lens the particles appear to be all alike. The particles of sand 
 and cuprous oxide fit together, and the combination takes place 
 without fusion. 
 
 FUSIBILITY OF SILICATES. 
 
 Single Silicates. Among the common silicates those of 
 the alkaline metals and lead have the lowest fusing-points ; 
 silicates of copper and iron take an intermediate place ; 
 while silicates of lime, magnesia, and alumina fuse only at 
 high temperatures. But, as no silicate will resist the excessive 
 temperature of an electric furnace, they may be described 
 generally as fusible compounds with a wide range of fusing- 
 points. 
 
 A normal silicate when in the fused state will, as a rule, 
 dissolve more of its basic oxide, or of silica, whichever is 
 presented to it. In the first case the normal silicate becomes 
 a basic salt, and in the second an anhydro-acid salt (p. 154). 
 The addition may either raise or lower the fusing-point 
 of the mass. There is no general rule with regard to this 
 change. 
 
 Complex Silicates. It may be stated generally that if 
 two chemical compounds, having well-defined melting-points, 
 are mixed together in certain proportions, the resulting 
 mixture will have a lower melting-point than either of the 
 constituents taken singly. Silicates form the best illustration 
 of this general principle, and every advantage is taken of it in 
 metallurgical operations. For example, the normal silicates 
 of lime and alumina are separately infusible at ordinary 
 furnace temperatures ; but if they are melted together, or 
 
230 METALLURGICAL CHEMISTRY 
 
 if arrangements are made for their simultaneous formation in 
 certain proportions, a readily fusible double silicate is the 
 result. 
 
 The formula for orthosilicic acid is H 4 Si0 4 , or 2H 2 O.Si0 1 ; 
 it is, therefore, tetrabasic. So that the normal salt of lime 
 is 2CaO.Si0 2 , and that of alumina 2Al 2 O 3 .3Si0 2 . Now, either 
 of these compounds is practically infusible at ordinary furnace 
 temperatures, but if they are heated together in certain pro- 
 portions a moderately fusible double silicate is obtained. 
 This may be expressed by the following formula, which is 
 usually aimed at when the double silicate is to be formed in a 
 metallurgical operation : 
 
 6(2CaO.Si0 9 ) + 2Al 2 3 .3Si0 2 
 6(2x56 + 60) 2x102 + 3x60 
 
 loST" ~~384~~ 
 
 The proportions given are easily verified, and the percentage 
 of each constituent can be calculated. If this is done, mixtures 
 of the oxides are readily made in convenient quantities to 
 give approximately the above composition. 
 
 EXP. 120. 1. Mix together 7 grams of quick-lime, and 3'8 grams 
 of fine sand ; put the mixture into a carbon crucible, and heat it 
 strongly in a wind furnace for twenty minutes. Put a lid on the 
 pot, and place it well down in the fire, so as to expose it to the 
 highest temperature the furnace will give. Remove the pot from 
 the fire at the end of stipulated time, and allow it to cool. 2. Mix 
 together 2'1 grams of alumina and 1/9 grams of fine sand, and 
 follow the instructions given for No. 1. On examination the mixtures 
 are found to be fitted together, and the one containing the lime 
 may show signs of incipient fusion. 3. Grind the two mixtures 
 thoroughly together in an iron mortar ; transfer the new mixture 
 so obtained to the carbon crucible, and heat it again for twenty 
 minutes under the same conditions as before. Remove the crucible 
 from the furnace, allow it to cool, and examine the result of the 
 experiment. 
 
 If the experiment is carefully conducted, a button of the 
 perfectly fused double silicate will be obtained. It is very 
 brittle and readily fractured. Very often it breaks up while 
 cooling, which is due to unequal contraction in different parts 
 
SILICON AND ITS COMPOUNDS 231 
 
 of the mass. The appearance of the fractured surface depends 
 somewhat upon the furnace conditions and on the rate of 
 cooling. It may vary from a bright glassy to a dull earthy 
 appearance. 
 
 A similar experiment may be made with the corresponding 
 compounds of metasilicic acid, H 2 O.Si0 2 . The composition 
 of the most fusible double silicate of lime and alumina is given 
 by the formula : 
 
 3(CaO.Si0 2 ) + Al 2 3 .3Si0 2 . 
 
 The metasilicate of lime will show more signs of fusion than 
 the corresponding orthosilicate ; and, if the temperature is 
 very high, the metasilicate of alumina will soften considerably. 
 It is, therefore, more fusible than the corresponding ortho- 
 silicate. The double silicate is also more fusible, and forms 
 the normal slag of furnaces working at a moderate tempera- 
 ture. Orthosilicate slags come from hotter furnaces. 
 
 If a small quantity of a basic oxide is added to a fused 
 silicate, it is dissolved, and exerts an influence upon (1) the 
 fusibility of the silicate and (2) its appearance. In most cases 
 the fusibility is increased, but this is not always so; an 
 addition of zinc oxide, ZnO, for example, decreases the 
 fusibility. Various colours are imparted to silicates by the 
 addition of oxides of iron (green), copper (red), cobalt (blue), 
 etc. Oxide of tin renders the silicate opaque, as in white 
 enamels. 
 
 Fireclay is an acid silicate of alumina containing com- 
 bined water, which, together with the fineness of its particles, 
 renders the clay plastic when mixed with water. Its com- 
 position may be represented by #Si0 2 .2Al 2 3 .3Si0 2 .2H 2 0, 
 when #Si0 2 represents a*variable proportion of the acid oxide. 
 This material, when approximately pure, is very refractory ; 
 but the addition of small quantities of basic oxides increases 
 its fusibility and renders it unfit for fire-resisting purposes. 
 This effect depends upon the nature and proportion of the 
 added oxide. Soda, potash, and oxide of lead exert the 
 
232 METALLUEGICAL CHEMISTRY 
 
 greatest influence ; lime and oxide of iron are also very 
 hurtful, the latter especially so if carbon or reducing gases 
 are present to reduce the ferric oxide, Fe 2 3 , to ferrous oxide, 
 FeO. These oxides unite with the excess of silica, and fusible 
 double silicates are formed, and if present as impurities in the 
 clay, are injurious to its refractory character. 
 
 Exp. 121. Mix a few grams of finely-powdered fireclay with 
 5 per cent, of dry sodium carbonate, which will carry about 3 per 
 cent, of soda, Na 2 0, into the mixture. Heat the mixture strongly 
 in a carbon crucible for twenty minutes, and examine the result. 
 
 Glass. If the experiments described above are carried out, 
 the tendency of fused silicates to form the transparent amor- 
 phous body known as glass is clearly demonstrated. Long 
 experience has taught the glass-maker the mixtures of silicates 
 which give the best results for particular purposes. Thus, 
 Window glass is a double silicate of lime and soda ; Bohemian 
 glass, which is a double silicate of lime and potash, will with- 
 stand a moderately high temperature without softening, and 
 changes in temperature without cracking: it is, therefore, 
 very useful for combustion-tube and for chemical apparatus 
 generally. Flint glass is a double silicate of lime and lead 
 oxide, and softens readily in an ordinary gas flame ; it is used 
 for making glass tube which is required to be easily bent. 
 Glass has to be carefully annealed by allowing it to cool very 
 slowly after it has been blown into shape. The plastic state 
 of hot glass, which extends through a wide range of tempera- 
 ture, renders it possible for the glass-blower to do with his 
 material what the potter does with plastic clay. The intro- 
 duction of small quantities of metallic oxides brings about the 
 variations in colour observed in ornamental glasses ; and a 
 very small quantity of free metal in an extremely finely divided 
 state may influence the colour, as is the case w r ith gold in ruby 
 glass. The opalescence in some varieties of glass is caused by 
 the added oxides not completely dissolving in the molten mass, 
 and the very fine solid particles being suspended in the fluid 
 matter while it sets. 
 
SILICON AND ITS COMPOUNDS 
 
 233 
 
 Nomenclature of Silicates. A system of naming, which has 
 been very largely adopted by metallurgists, depends upon the 
 proportion between the number of atoms of oxygen in the acid 
 and basic portions of the given silicate. This is clearly depen- 
 dent upon the assumption that a silicate is a combination of two 
 or more oxides, one of them always being silica. Take, for 
 example, the sodium salt of orthosilicic acid, 2Na 2 O.Si0 2 . The 
 relation here is 1 : 1, for there are 2 atoms of oxygen in each 
 part of the compound. But with the sodium metasilicate, 
 Na 2 O.Si0 2 , the relation is 2 : 1, for there are 2 atoms of 
 oxygen in the acid portion to 1 atom in the basic portion. 
 The number of atoms of the metal, which varies with its 
 valency, does not influence the character of the silicate. The 
 ratio 1 : 1 gives rise to monosilicates, and the ratio 2 : 1 
 bisilicates. But this does not account for the anhydro-acid 
 silicates, of which there are numerous examples. Thus in the 
 trisilicates the ratio is 3 : 1, and in the sesquisilicates it is 
 3:2. There are also silicates which contain an excess of the 
 basic oxide, and the most important of these are included in 
 the subsilicates in which the ratio is 1 : 2. 
 
 TABLE OF SILICATES. 
 
 RATIO OF NUMBER OF 
 
 
 
 
 ATOMS OFOxYGEM IN 
 ACID OXIDE TO NUM- 
 BER OF ATOMS OF 
 
 MONOVALENT 
 METAL, 
 
 DIVALENT METAL, 
 M"O. 
 
 TRIVALENT METAL, 
 M"'. 2 3 . 
 
 OXYGEN IN BASIC 
 
 
 
 
 OXIDE. 
 
 
 
 
 1 : 2 
 
 Sab 4Na 2 O.Si0 2 
 
 4CaO.Si0 2 
 
 4ALO s .3Si0 9 
 
 1 : 1 
 
 Mono 
 
 2Na 2 O.Si0 2 
 
 2CaO.Si0 2 
 
 2AUO s .3SiO., 
 
 2 : 1 
 
 Bi 
 
 Na 2 O.Si0 2 
 
 CaO.SiO., 
 
 Al 2 0,.3SiO., 
 
 3 : 1 
 
 Tri 
 
 2Na 2 O.Si0 2 
 
 2Ca0.3Si0 2 
 
 2Alo6 3 .9Si(X 
 
 q . Q 
 
 3:2 
 
 Sesqui 
 
 4Na. 2 0.3Si0 2 
 
 4Ca0.3Si0 2 
 
 4Al,0 3 .9Si0 2 
 
 The metals commonly present in silicates are : (1) Mono- 
 valent, general formula of oxide M' 2 0, where M represents an 
 atom of the metal ; (2) divalent, M"0 ; and (3) trivalent, 
 
234 METALLURGICAL CHEMISTEY 
 
 M'" 2 O 3 . Often all three of these types of oxides are present 
 in the same well-defined silicate. It is the readiness with 
 which acid, basic, and complex silicates seem to form that 
 makes these bodies so numerous. 
 
 When a natural or prepared silicate is in the form of a 
 glassy or stony-looking body, it is impossible to state for certain 
 that it is a well-defined compound, although its composition 
 can be accurately determined. But when it is found or 
 obtained in well-defined crystals, this difficulty disappears, and 
 an exact formula can be assigned to it. Such bodies often 
 occur in the earth's crust, and may also be prepared artificially. 
 In this respect it is to be borne in mind that slow cooling of a 
 molten mass favours the formation of crystalline structure 
 and the development of individual crystals. 
 
 Mono-, di-, and tri-valent metals are often associated in 
 the same natural silicate, and empirical formulae, as simple as 
 possible, are commonly used in such cases. Thus the well- 
 known mineral felspar is written empirically KAlSi 3 8 , and 
 this tells the beginner very little, except that it is a double 
 silicate of potassium and aluminium ; but if it is noticed that 
 the monovalent K' and trivalent Al"' give a tetravalent com- 
 bination (KA1)"", which is equivalent to 2 atoms of oxygen, 
 the formula may be written (KAl)0 2 .3Si0 2 , and the compound 
 is seen to be a trisilicate. Also, potassium may be more or 
 less completely replaced by sodium without altering the 
 crystalline form. In this case the different compounds are 
 isomorphous (p. 128). Isomorphism plays a very important 
 part in the mineral world. Na 2 and K 2 ; CuO, MgO, MnO, 
 and FeO ; A1 2 O 5 , Fe 2 3 , and MiiaOg replace each other without 
 altering the crystalline form of the minerals in which they 
 occur. 
 
 These natural bodies of definite composition are the 
 minerals which help to form the various rocks in the earth's 
 crust. Thus, granite is made up of three well-defined minerals : 
 felspar, mica, and quartz, and the crystals of these bodies 
 were formed during the extremely slow cooling of the molten 
 
SILICON AND ITS COMPOUNDS 235 
 
 rock. Granite is readily melted in an ordinary wind furnace, 
 but does not recover its crystalline structure on cooling. 
 The solidified mass resembles the slags formed in smelting 
 operations. 
 
 SUMMARY. 
 
 Silica, Si0 2 , the only oxide of silicon, is the most generally 
 occurring compound in the earth's crust. Quartz, flint, and 
 sand are its commonest forms. In combination with metallic 
 oxides it forms the natural and prepared silicates, which are 
 a very numerous class of bodies. There are two silicic acids, 
 which give rise to two well-defined series of compounds, 
 orthosilicates and metasilicates. These normal silicates, 
 when in the molten state, readily take up an excess of either 
 silica or metallic oxides, and form basic and anhydro-acid 
 silicates. Different silicates also mix together readily, and a 
 large number of complex bodies are thus formed. As a rule, 
 a complex silicate is more fusible than a single silicate. The 
 element silicon unites with metals to form silicides. It is 
 commonly found as an impurity in commercial iron. 
 
 QUESTIONS. 
 
 1. Give a short description of the physical and chemical 
 properties of silica. 
 
 2. What is the general effect of mixing silicates together, 
 upon the f using-point of the mixture 1 
 
 3. Describe an experiment in which a silicate is formed. 
 
 4. What is fireclay, and to what does it owe its plasticity 1 
 
 5. How are silicates classified for metallurgical purposes ? 
 
 6. What do you understand by the isomorphous substitution 
 of one metallic oxide for another in a silicate 1 
 
CHAPTER XV 
 WEIGHTS, MEASUEES, AND APPARATUS 
 
 THE Metric System of weights and measures, which is of 
 French origin, is almost universally used for scientific purposes. 
 It is a decimal system, and on that account is very easy to 
 work with. 
 
 Length. The unit of length is the metre, which is 
 divided into 10 parts, or decimetres; into 100 parts, or 
 Centimetres; and into 1,000 parts, or millimetres. The 
 Latin prefixes, deci-, centi-, and milli-, mean ^ T ^y, and 
 unrip respectively, and denote the submultiples of the 
 metre. 
 
 The Greek prefixes, deka-, hecto-, and kilo-, denote 10, 100, 
 and 1,000, respectively, and are used for multiples of the metre. 
 Thus, 1 kilometre = 1,000 metres. 
 
 Volume. The common measure of volume is the litre. 
 Its capacity is equal to that of a cube, 1 decimetre, or 10 
 centimetres on the side. The litre = 1 cubic decimetre = 1,000 
 cubic centimetres. Any volume less than a litre is usually 
 expressed in cubic centimetres. Thus, J litre = 500 cubic 
 centimetres = 500 c.c., c.c. being the contraction for cubic 
 centimetres. 
 
 Weight. The unit of weight is derived from the metre ; 
 it is the weight of a cubic centimetre of pure water at its 
 point of maximum density (4 C.), and is called the gramme 
 or gram. The submultiples are the decigram, or y 1 ^ gram ; 
 the centigram, or yj^ gram ; and the milligram, or 
 
 236 
 
WEIGHTS, MEASUEES, AND APPAKATUS 237 
 
 gram. The Greek prefixes are used for the multiples. Thus, 
 the kilogram = 1,000 grams. 
 
 ENGLISH EQUIVALENTS OF THE METRIC UNITS. 
 
 The metre = 39'37 inches. 
 ,, litre = 1*76 pints. 
 gram = 15'432 grains. 
 ,, kilogram = 2*2046 pounds avoirdupois. 
 
 The equivalents of the multiples or submultiples are easily 
 obtained by moving the decimal point to the right for 
 multiples, and to the left for submultiples. 
 
 Thus, 1 kilometre =39,370 inches. 
 1 centigram = 0*1 5342 grain. 
 
 The following will be found useful in making comparative 
 measurements : 
 
 1 pound avoirdupois = 7,000 grains. 
 
 i ounce ,, =437 *5 
 
 1 pound troy =5,760 
 
 1 ounce ,, = 480 ,, 
 
 1 pint = 20 fluid ounces. 
 
 1 pint of water at 15'5 C. weighs 20 ounces avoirdupois. 
 1 gallon of water at 15 '5 C. weighs 10 pounds avoirdupois. 
 1 cubic foot of water at 15'5 C. weighs 62-4 pounds avoirdu- 
 pois nearly. 
 
 The Balance. The most casual observer is familiar with 
 the ordinary beam scales used in everyday life ; and, as the 
 more sensitive balance used for scientific purposes is exactly 
 the same in principle, no difficulty should be experienced in 
 using it. Accuracy in weighing is only a matter of care and 
 experience. 
 
 The beginner should examine the balance he is using, and 
 notice that it consists of an accurately made beam, supported 
 by knife edges which rest on smooth horizonal plates or in 
 
METALLURGICAL CHEMISTRY 
 
 V-shaped grooves, and having two scale-pans suspended in a 
 similar manner, one from each end. The points of suspension 
 of the pans are at exactly the same distance from the point 
 of suspension of the beam, and the two halves of the beam 
 
 are exactly similar to each 
 other. When the balance 
 is properly set the pointer 
 should be perfectly vertical 
 if the beam is at rest, or 
 should swing to nearly 
 equal distances on each 
 side of the vertical division 
 on the scale at the foot of 
 the supporting pillar if it 
 is in motion. The smaller 
 the weight required to be 
 added to either one side or the other in order to disturb this 
 equality, the more sensitive the balance is, and the more 
 accurate the weighing which can be made upon it. A cheap 
 and sufficiently accurate form of balance is shown in Fig. 41. 
 
 The Weights. The Metric System of weights is almost 
 universally used for scientific purposes, and a box of gram 
 weights containing from 100 grams to 1 milligram will be 
 found sufficient for most purposes. The individual weights 
 are commonly marked as shown in the diagram. 
 
 FIG. 41. 
 
 Gramsj 
 
 Decigrams 
 
 Centigrams 
 
 Milligrams 
 
 100 
 10 
 0-5 
 0-05 
 
 50 
 5 
 
 0-2 
 0-02 
 
 20 
 2 
 
 0-1 
 0-01 
 
 20 
 1 
 
 0-1 
 0-01 
 
 0-005 0-002 0-001 0-001 
 
 Sometimes it is found convenient to use grain weights, and 
 sets varying from 1,000 grains to 0*5 grain would be found 
 sufficient for the work described in this book. In converting 
 the weight in grams given in the text into grains it will be 
 
WEIGHTS, MEASURES, AND APPARATUS 239 
 
 sufficient to multiply the weight in grams by 15*5. 1 gram = 
 15-432 grains. 
 
 Weighing. The body to be weighed is usually placed in 
 the left-hand pan, the weights in the right-hand one. The 
 weights should be used systematically, as weighings are made 
 more rapidly by so doing. The beam can be raised or lowered 
 by means of a lever arrangement worked from the front of the 
 stand, and bodies to be weighed should not be put into the 
 pan, nor should weights be added while the beam is swinging 
 freely. When weighing it is not necessary to wait for the beam 
 to come to rest, for when the pointer swings to very nearly 
 equal distances on each side of the perpendicular mark on 
 the graduated scale in front of which it moves, the weighing 
 is finished. Bodies which are at all likely to soil the scale-pan 
 should be weighed in a counterpoised watch-glass or crucible. 
 In fact, it is a good plan to make all weighings in this way, as 
 any little inaccuracy in the balance is thus neutralized. For 
 example, if a watch-glass is weighed, and a body then weighed 
 in it, the difference between the two weights is the weight of 
 the body. The beginner is strongly advised to get into the 
 habit of writing down the results of weighings before he leaves 
 the balance. A little practice, especially with tho assistance 
 of a teacher, will enable him to weigh accurately, and soon 
 convince him that methodical weighing is much more rapid 
 than an erratic putting on and pulling off of weights. 
 
 Measuring" Vessels. The volume measures of the Metric 
 System* are almost always used in volumetric work. The 
 cubic decimetre or litre is the most important, and contains 
 1,000 cubic centimetres. In the experiments described in 
 the text the volumes used are less than a litre, and are 
 generally expressed in cubic centimetres. 
 
 The burette is a long, narrow, glass tube, open at one end 
 and fitted at the other with a tap arrangement through which 
 
 * See the table of the Metric System of weights and measures at the 
 beginning of the chapter. 
 
240 
 
 METALLUKGICAL CHEMISTRY 
 
 small quantities of the contained liquid can be run. The tap 
 must be so under control that the liquid may be made to 
 leave the burette one drop at a time, if necessary. In cheap 
 burettes it is usual to connect the body with the nozzle by 
 means of a piece of rubber tube, and to close this by a spring 
 pinch-tap. A glass tap is shown in Fig. 42. A very simple 
 and effective substitute for the pinch-tap may be made by 
 introducing a piece of glass rod J inch long, rounded at the 
 ends, and of such diameter as to require gently forcing into 
 the rubber tube connection between the nozzle and the body 
 of the burette. This completely closes the 
 burette while at rest, but when the rubber 
 tube just over the rod is squeezed between the 
 finger and thumb a channel in the rubber is 
 made down the side of the rod by which the 
 liquid can escape. On removing the pressure 
 the burette is at once closed. The body of 
 the burette is graduated between A and B 
 into fifty divisions, each equal to 1 cubic centi- 
 metre, so that if a liquid is level with the zero 
 mark at A, and is run out until it is level with 
 B, exactly 50 cubic centimetres are delivered. 
 If the level is reduced to some point between 
 A and B, then the graduation mark at that 
 point gives the volume delivered. The sur- 
 face of the liquid in the burette is curved, and the read- 
 ings are best made by reference to the lower or convex 
 surface of the liquid. The burettes usually sold are not very 
 accurate ; but if any connected set of measurements are made 
 with the same burette a slight inaccuracy does not appreciably 
 affect the results. For liquids which corrode rubber tube, 
 burettes with glass taps must be used. 
 
 Measuring 1 Flasks are usually sold in four sizes i.e., 
 1,000 c.c., 500 c.c., 250 c.c., and 100 c.c. They sometimes 
 have two graduation marks on the neck, one to contain and 
 
 FIG. 42. 
 
WEIGHTS, MEASURES, AND APPARATUS 241 
 
 FIG. 43. 
 
 one to deliver. As these flasks expand or contract as they 
 
 are heated or cooled, the reading is only strictly true at the 
 
 temperature at which 
 
 the graduation was 
 
 made. This is usually 
 
 the normal temperature 
 
 of a laboratory, i.e., 
 
 15C. Thus, if a liquid 
 
 is poured into the dry 
 
 flask until the level 
 
 reaches the lower gradu- 
 ation mark />, Fig. 43, 
 
 the flask contains the 
 
 registered volume of 
 
 liquid ; but if the level 
 
 is brought to the upper graduation mark a, then it will deliver 
 
 the registered quantity if the liquid is poured into another 
 
 vessel. The extra liquid between the two marks compensates 
 
 for that which adheres to the sides of the flask after the main 
 
 bulk has been poured out. 
 
 Three very useful measures are shown in Fig. 44, which may 
 
 be used for rough measurements. They deliver 500, 250, and 
 
 50 c.c. respectively. For the 
 rapid measurement of a few 
 cubic centimetres of liquid a 
 pipette can be made by draw- 
 ing off to a point in the 
 gas flame a piece of glass 
 tube about T 8 F of an inch 
 internal diameter and 8 inches 
 long. To graduate it into 
 cubic centimetres, dip the 
 narrow end into water, and 
 when some of the liquid has 
 
 entered, remove the tube and let out as much of the liquid 
 
 as will run from the narrrow end. Without disturbing 
 
 16 
 
 
 c.c. 
 
 p 
 
 
 
 =-500 
 
 
 
 
 1-400 
 
 
 
 
 1-300 
 
 
 c.c. 
 
 p 
 
 
 
 H-300 
 
 
 " 201 
 
 
 c.c. 
 
 p 
 
 
 
 =- 
 
 
 
 
 rr 90 
 
 
 
 
 1-200 
 
 
 ^-150 
 
 
 80 
 
 70 
 
 
 
 
 
 
 
 
 
 6C 
 
 
 
 
 __ 
 
 
 10( 
 
 
 50 
 
 
 
 
 =-100 
 
 
 EE-JO 
 
 
 40 
 3C 
 -20 
 
 
 
 ^ 
 
 =i 
 
 S r 
 
 
 ^ ^ 
 
 10 
 
 s 
 
 FIG. 44. 
 
242 METALLURGICAL CHEMISTRY 
 
 the small quantity of liquid not delivered, close the narrow 
 end with a little stiff grease, and run in at the open end 1 c.c. 
 of water from a burette ; mark the level of the liquid by 
 scratching the side of the tube with a file ; run in a second cubic 
 centimetre without disturbing the first, and mark as before ; 
 continue this until about two-thirds of the tube has been 
 graduated. The pipette will then deliver the number of cubic 
 centimetres marked upon it. To use it, place the tube in the 
 liquid of which the measured quantity is required ; the liquid 
 enters the tube, and when the required quantity is in, place the 
 finger on the top of the tube and remove it. To deliver the 
 liquid, remove the finger, and it will run out. A very useful 
 pipette made in various sizes is shown in Fig. 44. 
 
 Wash-bottle. A very convenient form of wash-bottle for 
 quantitative work is shown in Fig. 45. It consists of a glass 
 flask, A, fitted up as shown in the figure. 
 The nozzle, B, is attached by means of a 
 piece of rubber tube, and the neck, C, is 
 bound with a layer of thick string. The 
 flexible nozzle allows of a stream of water 
 being sent in any required direction. If 
 the flask is thin the water in it may be 
 raised to boiling, either on the sand-bath or 
 on wire gauze, when the necessity of wash- 
 ing with hot water arises. 
 
 45^ Other Apparatus. In devising the 
 
 experiments described in the text, the chief 
 aim has been to use such apparatus as the student in an 
 elementary laboratory may be expected either to possess 
 or to have access to. Expensive apparatus of the easily- 
 broken type has been either excluded or attempts have 
 been made to imitate it in a simple manner. When more 
 accurate apparatus than that described can be used, better 
 results may be expected. But the student often gains some 
 knowledge by fitting up a piece of apparatus consisting 
 
WEIGHTS, MEASUEES, AND APPARATUS 243 
 
 of several parts, which by the ingenuity of the glass-blower 
 may be obtained complete. The principle involved in the 
 saying, " You press the button, we do the rest," is no doubt 
 useful for many purposes, but it is bad from an educational 
 point of view. 
 
 The table furnace shown in Fig. 8 was evolved from a 
 simpler form consisting of an iron stand carrying some loose 
 fireclay slabs, and having two ordinary Bunsen burners 
 underneath. In the first modification the movable burners 
 were replaced by three fixed burners, and in the last the 
 fireclay slabs were replaced by a rectangular fireclay chamber, 
 surmounted by a sheet-iron tube to increase the draught. 
 The temperature obtained in this chamber is very high for 
 a table furnace, and experiments which in the ordinary form 
 of furnace either fail or are only partially successful can be 
 carried out with it. 
 
 Conical flasks are much more convenient for general work 
 than the ordinary globular variety. The 8, 12, and 16 ounce 
 sizes are the most useful. 
 
 Porcelain Crucibles. The best Berlin crucibles are the 
 cheapest in the end, and Nos. 1 and 2 are the useful sizes. 
 
 Combustion-tube, with walls of medium thickness, should be 
 selected. Experiments are often a failure in thick-walled 
 tubes, because the heat from the ordinary Bunsen flame is 
 not able to penetrate sufficiently. On the other hand, a thin- 
 walled tube is liable to blow out when hot, if the gas pressure 
 inside is greater than that of the atmosphere outside. 
 
 Clay wasting-dishes are very handy for muffle work, and 
 Morgan's No. 1 is a useful size. 
 
 162 
 
APPENDIX 
 
 TABLE OF ELEMENTS, WITH SYMBOLS AND ATOMIC 
 WEIGHTS 
 
 NON-METALS. 
 
 NAME. 
 
 SVMBO,. Co*'?, j 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 Argon 
 
 A 
 
 40'0 
 
 Iodine 
 
 I 
 
 126-54 
 
 Arsenic ... 
 
 As 
 
 74-9 
 
 Nitrogen 
 
 N 
 
 14-01 
 
 Bromine 
 
 Br 
 
 7976 
 
 Oxygen 
 
 
 
 15-96 
 
 Boron B 109 
 
 Phosphorus ... ! P 
 
 30-96 
 
 Carbon 
 
 C 11-97 
 
 Silicon .. 
 
 Si 
 
 28-0 
 
 Chlorine 
 
 01 
 
 35-37 I Sclenion 
 
 Se 
 
 78-87 
 
 Fluorine 
 
 F 
 
 19-06 ;| Sulphur 
 
 S 
 
 31-98 
 
 Hydrogen 
 
 H 
 
 i-o 
 
 Tellurion 
 
 Te 
 
 127-7 
 
 COMMON METALS. 
 
 Aluminium 
 
 Al 
 
 27-04 
 
 Magnesium 
 
 Mg 
 
 23-94 
 
 Antimony 
 
 Sb 
 
 119-6 
 
 Manganese 
 
 Mn 
 
 54-8 
 
 Barium... 
 
 Ba 
 
 136-86 
 
 Mercury ... i Hg 199'8 
 
 Bismuth 
 
 Bi 
 
 207-5 
 
 Molybdenum ... ' Mo 95 -9 
 
 Cadmium 
 
 Cd 
 
 111-7 
 
 Nickel 
 
 Ni 
 
 58-6 
 
 Calcium 
 
 Ca 
 
 39-91 
 
 Platinum 
 
 Pt 
 
 194-3 
 
 Chromium 
 
 Cr 
 
 52-45 
 
 Potassium 
 
 K 39-03 
 
 Cobalt ... 
 
 Co 
 
 58-6 
 
 Silver 
 
 Ag i 107-66 
 
 Copper 
 
 Cu 
 
 63-18 
 
 Sodium 
 
 Na 22-995 
 
 Gold 
 
 Au 
 
 196-7 
 
 Strontium 
 
 Sr 
 
 87-3 
 
 Iron 
 
 Fe 
 
 55-88 
 
 Tin 
 
 Sn 
 
 117-35 
 
 Lead 
 
 Pb 
 
 206-92 
 
 Zinc 
 
 Zn 64-88 
 
 REMARKS. Sb from stibium ; Hg from hydrargyrum ; K from kalium ; 
 Cu from cuprum ; Ag from argentum ; Au from aurum ; Na from 
 natrium ; Fe from ferium ; Sn from stannum ; Pb.from plumbum. The 
 names of the more important elements are printed in thick type. 
 
 244 
 
APPENDIX 
 
 245 
 
 RARE METALS. 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 Beryllium 
 
 Be 
 
 9-08 
 
 Ruthenium 
 
 Ru 
 
 103-5 
 
 Caesium . 
 
 Cs 
 
 1327 
 
 Samarium 
 
 Sm 
 
 150-0 
 
 Cerium ... 
 
 Ce 
 
 142-2 
 
 Scandium 
 
 Sc 
 
 43-97 
 
 Didymium 
 Erbium ... 
 
 D 
 E 
 
 145-0 
 166-0 
 
 Tantalum 
 Terbium 
 
 Ta 
 Tr 
 
 182-0 
 160 
 
 Gallium 
 
 Ga 
 
 69-9 
 
 Thallium 
 
 Tl 
 
 203-7 
 
 Germanium 
 
 Ge 
 
 72-32 
 
 Thorium 
 
 Th 
 
 231-96 
 
 Indium ... 
 
 In 
 
 113-4 
 
 Thulium 
 
 Tu 
 
 169-4 
 
 Iridium 
 
 Ir 
 
 192-5 
 
 Titanium 
 
 Ti 
 
 48 
 
 Lanthanum 
 
 La 
 
 138-5 
 
 Tungsten 
 
 W 
 
 183-6 
 
 Lithium 
 
 Li 
 
 7-01 
 
 Uranium 
 
 U 
 
 239-8 
 
 Niobium 
 
 Nb 
 
 93-7 
 
 Vanadium 
 
 V 
 
 51-1 
 
 Osmium 
 
 Os 
 
 191-0 
 
 Ytterbium 
 
 Yb 
 
 172-6 
 
 Palladium 
 
 Pd 
 
 106-2 
 
 Yttrium 
 
 Y 
 
 89-9 
 
 Rhodium 
 
 Rh 
 
 103-1 
 
 Zirconium 
 
 Zr 
 
 90-4 
 
 Rubidium 
 
 Rb 
 
 85'2 
 
 
 
 
 CORRECTION OF THE VOLUME OF A GAS FOR TEMPERATURE, 
 PRESSURE, AND WATER VAPOUR. 
 
 Correction for Temperature. Gases expand Y |^ of 
 their volume for every increase of 1 C. in their temperature 
 and contract in the same degree when their temperature is 
 decreased (law of Charles). Now, suppose the gas to be cor- 
 rected were cooled to C., its volume divided into 273 parts, 
 and the whole again raised a degree at a time to the original 
 temperature. Then for a rise of 1 the increased volume 
 would be represented by 274, for 5 by 278, for t by 273 + /. 
 Therefore the ratio between the volume of a gas at / C. and 
 its volume at C. is expressed by (273 + ) : 273, and the 
 volume at C. is found by multiplying the volume at t C. by 
 
 273 
 >734-/- -^ the temperature of the gas is below C. t is 
 
 negative, and it is to be subtracted from, and not added 
 to, 273. 
 
246 METALLUEGICAL CHEMISTKY 
 
 Correction for Pressure. The volume of a gas is in- 
 versely as the pressure it sustains, or the volume decreases in 
 the same proportion as the pressure increases, and vice versa 
 (law of Boyle and Marriott). The pressure to which a gas 
 is subjected when standing over a liquid with the liquid at the 
 same level inside and outside the containing vessel, is exactly 
 that of the air outside, and can be determined in millimetres 
 of mercury by reading the barometer. If the normal pressure 
 be taken as equal to 760 millimetres of mercury, and the 
 pressure of the gas to be corrected is equal to p millimetres, 
 
 ^7 A r\ 
 
 then the ratio between the volumes at p and at 760 is . and 
 
 p ' 
 
 the corrected volume is found by multiplying the volume at 
 p by Tj^T)- p may be greater or less than 760. 
 
 Correction for Water Vapour. A gas standing over 
 water becomes saturated with the vapour of the liquid, but 
 the weight of the volume of water vapour present depends 
 upon the temperature of the gas. Thus, at a given tempera- 
 ture a definite weight of water gas will mix with the other 
 gas, and, by Dalton's law of partial pressures, will exert its 
 own share of the total pressure of the moist gas. But when 
 two gases are mixed together their partial pressures are pro- 
 portional to their original volumes. Therefore, if the partial 
 pressure of water vapour at the temperature of the moist gas 
 is known, the volume of the water vapour can be calculated. 
 A diagram showing the pressure of saturated water vapour 
 for a limited range of temperature is given on next page. 
 
 The temperature of the gas is to be taken in degrees Centi- 
 grade, and the tension of the vapour at that temperature is 
 expressed in millimetres of mercury. 
 
 The temperature is measured on the horizontal lines, and 
 the tension on the vertical ones. The curve is obtained by 
 marking off on the diagram the results of a number of experi- 
 ments for determining the vapour tension at varying tern- 
 
APPENDIX 
 
 247 
 
 peratures, between the two extreme temperatures, and joining 
 the point by a continuous line. 
 
 DIAGRAM SHOWING THE TENSION OF AQUEOUS VAPOUR IN 
 
 A GAS SATURATED WITH THE VAPOUR AT A 
 
 GIVEN TEMPERATURE. 
 
 24 
 
 Ul /J 
 
 5678 
 
 11 12 13 1* IS 16 17 1& 13 20 21 22 23 24 25 
 DEGREES CENTIGRADE 
 
 To use the diagram, read the temperature of the saturated 
 gas to the nearest tenth ; find this temperature on the hori- 
 zontal line ; estimate with the eye the point on the curve 
 vertically above it, and then the projection of this point on 
 the vertical line represents the required tension in milli- 
 metres. Example: If the temperature is equal to 16'3 C., 
 the tension is equal to 13'8 millimetres of mercury. 
 
INDEX 
 
 PAGE 
 
 Acm, hydrochloric .-. .. 69 
 
 nitric .-. .. ~v. 78 
 
 sulphuric . . . . . . 63 
 
 Acid-forming oxides . . 137 
 
 Acids, action of. on allovs . . 106 
 
 iron v. " , .. 105 
 
 basicity of .. .. 153 
 
 constitution of . . . . 152 
 
 phosphoric . . . . 224 
 
 silicic . . . . . . 228 
 
 Action, electro-chemical . . 107 
 
 Air gas, composition of . . 180 
 
 Alkalies, fixed . . . . 83 
 
 volatile 83 
 
 Allotropic modification 57, 166 
 
 Alloys, action of acids on . . 106 
 
 Aluminium 90 
 
 Amalgamated zinc . . . . 108 
 
 Ammonia . . . . . . 82 
 
 Ammonium chloride . . 81 
 
 Analysis . . . . . . 67 
 
 Anhydride 138 
 
 Apatite 222 
 
 Atom, definition of . . . . 42 
 
 Atomic heat 128 
 
 weights . . . . . . 125 
 
 Atomicity 131 
 
 Atoms and molecules, theory 
 
 of 51 
 
 Avogadro, law of . . . . 43 
 
 BALANCE 
 
 Basic oxides 
 
 Basicity of acids 
 
 Battery, voltaic 
 
 Black oxide of manganese 
 
 Bone ash 
 
 237 
 136 
 153 
 112 
 72 
 221 
 
 PAGE 
 
 Boyle and Marriott, law of . . 246 
 
 Brin's process . . . 202 
 
 Bromine . . . . . . 76 
 
 Bunsencell 112 
 
 Burette . . 240 
 
 CALORIE 
 Carbide of iron 
 Carbon 
 
 dioxide . . " Y 
 
 monoxide . . .'. 
 
 reduction by 
 Carbonates and acids 
 
 natural 
 
 prepared 
 
 219 
 105 
 163 
 172 
 174 
 185 
 171 
 167 
 168 
 
 Carbonic oxide, reduction by 190 
 Cast iron . . . . . . 105 
 
 Cell, Bunsen 112 
 
 Change, chemical . . . . 4 
 
 endothermic . . 177, 214 
 
 exothermic . . 177, 214 
 
 physical .. . . .. 3 
 
 Charcoal .. .. ..164 
 
 T combustion of . . . . 165 
 
 Charles, law of . . . . 245 
 
 Chemical change . . . . 4 
 
 combination, laws of 47, 52 
 
 equation . . . . 49, 54, 67 
 
 equivalents . . . . 116 
 
 table of . . . . 125 
 
 Chlorate of potassium . . 75 
 Chloride of ammonium . . 81 
 Coal ..177 
 
 gas 179 
 
 Colloids 228 
 
 Colouring of massicot .. 10 
 
 Combustion . . . . . . 200 
 
 249 
 
250 
 
 INDEX 
 
 Combustion, explosive 
 
 PACK 
 
 .. 206 
 
 FELSPAR 
 
 PAGE 
 
 .. 234 
 
 inversion of 
 
 .. 207 
 
 Fixed alkalies 
 
 .. 83 
 
 of charcoal 
 
 .. 165 
 
 Flame 
 
 .. 208 
 
 Common salt 
 
 37,69 
 
 luminosity of 
 
 .. 210 
 
 Complex oxides 
 
 .. 137 
 
 structure of 
 
 .. 209 
 
 silicates 
 
 .. 229 
 
 Flasks, conical 
 
 .. 243 
 
 Composition, percentage 
 
 .. 65 
 
 Fritting 
 
 .. 228 
 
 Compound, definition of 
 
 ..41,47 
 
 Fusibility of silicates 
 
 .. 229 
 
 Conical flasks 
 
 .. 243 
 
 
 
 Conservation of energy 
 
 211, 212 
 
 GAS, coal 
 
 .. 179 
 
 matter 
 
 .. 35 
 
 Gases 
 
 3 
 
 Constitution of acids 
 
 .. 152 
 
 diffusion of 
 
 .. 31 
 
 Copper, oxidation of 
 
 .. 10 
 
 Glass 
 
 .. 232 
 
 Crystallization, water of 
 
 .. 147 
 
 Gold, subdivision of . . 
 
 .. 36 
 
 Crystalloids 
 
 .. 228 
 
 Graham 
 
 . . 227 
 
 Crystals 
 
 .. 149 
 
 Graphite 
 
 .. 163 
 
 Crucibles, porcelain . . 
 
 .. 243 
 
 
 
 Current, electric 
 
 .. 109 
 
 HEAT, action of, on salts 
 
 .. 157 
 
 Cyanide, reduction by 
 
 .. 195 
 
 water 
 
 .. 30 
 
 
 
 atomic 
 
 .. 128 
 
 DENSITY 
 
 .. 33 
 
 definition of 
 
 .. 211 
 
 of vapour 
 
 .. 126 
 
 measurement of . . 
 
 .. 214 
 
 Dialysis 
 
 .. 228 
 
 reduction by 
 
 . . 182 
 
 Diffusion of gases 
 
 .. 31 
 
 specific 
 
 .. 217 
 
 Dioxide, carbon 
 
 .. 172 
 
 Hydrochloric acid . . 
 
 .. 69 
 
 sulphur 
 
 94 
 
 and metals 
 
 . . 86-91 
 
 Dissociation .. ..82 
 
 , 91, 101 
 
 Hydrogen, preparation of 
 
 24,94 
 
 Divisibility of matter 
 
 36 
 
 properties of 
 
 .. 25 
 
 Double salts .. 
 
 . . 155 
 
 reduction by 
 
 .. 188 
 
 Dressing 
 
 9 
 
 
 
 
 
 IGNITION point 
 
 .. 201 
 
 ELECTRIC current 
 
 .. 109 
 
 Indicator, methyl-orange 
 
 .. 141 
 
 action of, on water 
 
 .. 27 
 
 litmus 
 
 .. 136 
 
 Electricity, reduction by 
 
 .. 184 
 
 Inversion of combustion 
 
 .. 207 
 
 Electro-chemical action 
 
 .. 107 
 
 Iodine . . 
 
 .. 76 
 
 Electrolysis 
 
 111 
 
 Ions 
 
 . Ill 
 
 Electrolyte 
 
 .. Ill 
 
 Iron 
 
 .. 24 
 
 Element, definition of 
 
 41 
 
 action of acids on 
 
 . . 105 
 
 Elements, valency of 
 
 .. 131 
 
 carbide 
 
 .. 105 
 
 Endothermic change 
 
 177, 214 
 
 cast 
 
 .. 105 
 
 Energy 
 
 4 
 
 oxidation of 
 
 .. 12 
 
 conservation of . . 
 
 211, 212 
 
 pyrites 
 
 .. 61 
 
 kinetic 
 
 .. 211 
 
 rusting of . . 
 
 15 
 
 potential 
 
 .. 212 
 
 wrought 
 
 .. 105 
 
 Equation, chemical . . 4 
 
 9, 54, 67 
 
 Isomorphism 
 
 128, 156 
 
 thermo-chemical . . 
 
 .. 201 
 
 
 
 Equivalents, chemical 
 
 .. 116 
 
 KINETIC energy 
 
 .. 211 
 
 - table of 
 
 .. 125 
 
 
 
 Exothermic change . . 
 
 177, 214 
 
 LAW of Avogadro 
 
 .. 43 
 
 Explosive combustion 
 
 .. 206 
 
 Boyle and Marriott 
 
 . . 246 
 
 Extraction of phosphorus 
 
 .. 221 
 
 Charles 
 
 .. 245 
 

 INDEX 
 
 
 PAGE 
 
 
 Laws of chemical combina- 
 
 
 Oxidation of lead 
 
 tion 
 
 47,52 
 
 mercury 
 
 Lead, oxidation of . . 
 
 8 
 
 tin 
 
 - red 
 
 10 
 
 Oxide, nitric . . 
 
 Levigation 
 
 9 
 
 nitrous 
 
 Litmus indicator 
 
 136 
 
 Oxides, acid-forming 
 
 Liquids 
 
 2 
 
 basic 
 
 Luminosity of flame 
 
 210 
 
 complex 
 
 
 
 nomenclature of . . 
 
 MAGNESIUM 
 
 23 
 
 table of 
 
 Malleability .. 
 
 8 
 
 Oxygen 
 
 Manganese, black oxide of . . 
 
 72 
 
 preparation of 
 
 Mass 
 
 33 
 
 properties of 
 
 Massicot, colouring of 
 
 10 
 
 
 Matter, conservation of 
 
 35 
 
 PEROXIDE, nitric 
 
 definition of 
 
 33 
 
 Phosphoric acid 
 
 divisibility of 
 
 36 
 
 Phosphorus . . 
 
 physical states of . . 
 
 2 
 
 extraction of 
 
 Measurement of heat 
 
 214 
 
 red 
 
 temperature 
 
 215 
 
 yellow 
 
 Mechanical mixtures 
 
 50 
 
 Physical change 
 
 Mercury, oxidation of 
 
 17 
 
 states of matter . . 
 
 red oxide of 
 
 38 
 
 Preparation of oxygen 
 
 Metals, action of, on water . . 
 
 22 
 
 Prepared carbonates . . 
 
 and hydrochloric acid 86, 91 
 
 Porcelain crucibles . . 
 
 nitric acid 
 
 96 
 
 Potassium chlorate . . 
 
 sulphuric acid 
 reduction by 
 Methyl-orange 
 
 91 
 193 
 
 83 
 
 Potential energy 
 Properties of oxygen 
 Pyrites, iron 
 
 orange indicator 
 
 141 
 
 
 Metric System 
 
 236 
 
 REACTION, reduction by 
 
 Mixtures, mechanical 
 
 50 
 
 Red lead 
 
 Molecule, definition of 
 
 42 
 
 phosphorus 
 
 Molecules and atoms, theory 
 
 
 oxide of mercury . . 
 
 of . . 
 
 51 
 
 Reduction by carbon 
 
 Monoxide, carbon 
 
 174 
 
 carbonic oxide 
 
 NATURAL carbonates 
 
 167 
 
 cyanide 
 electricity 
 
 Nitric acid 
 
 78 
 
 heat 
 
 and metals 
 
 96 
 
 hydrogen 
 
 oxide 
 
 100 
 
 metals . . 
 
 peroxide 
 
 100 
 
 reaction 
 
 Nitrogen 
 
 20 
 
 Roasting 
 
 Nitrous oxide 
 
 99 
 
 Rusting of iron 
 
 Nomenclature of oxides 
 
 138 
 
 
 salts 
 
 151 
 
 SALT, common 
 
 silicates 
 
 233 
 
 Saltpetre 
 
 
 
 Salts, action of heat on 
 
 OXIDATION 
 
 19 
 
 double 
 
 of copper . . 
 
 10 
 
 nomenclature of . . 
 
 iron 
 
 12 
 
 table of 
 
 251 
 
 PAGE 
 8 
 
 . 17 
 
 . 12 
 
 . 100 
 
 99 
 
 . 137 
 
 . 136 
 
 . 137 
 
 . 138 
 
 . 160 
 
 . 20 
 
 . 202 
 
 . 204 
 
 . 100 
 
 . 224 
 
 . 222 
 
 . 221 
 
 . 222 
 
 . 222 
 
 3 
 
 2 
 
 . 202 
 
 . 168 
 
 . 243 
 
 . 75 
 
 . 212 
 
 . 204 
 
 . 61 
 
 . 196 
 
 . 10 
 
 . 222 
 
 . 38 
 
 . 185 
 
 . 190 
 
 . 195 
 
 . 184 
 
 . 182 
 
 . 188 
 
 . 193 
 
 . 196 
 
 . 61 
 
 . 15 
 
 37,69 
 
 . 78 
 
 . 157 
 
 . 155 
 
 . 151 
 160 
 
252 
 
 INDEX 
 
 Scales, thermometric 
 
 PAGE 
 216 
 
 Thermometric scales 
 
 PAGE 
 
 2L6 
 
 Silica 
 
 226 
 
 TSu nvirlfl tinn r\ 
 
 19 
 
 Silicates 
 
 228 
 
 1 Hi, UAlU.dLJ.UJJl Ul 
 
 J.Z 
 
 complex 
 
 229 
 
 VALENCY of elements 
 
 131 
 
 fusibility of 
 
 229 
 
 table of 
 
 133 
 
 nomenclature of . . 
 - table of 
 
 233 
 233 
 
 Vapour density 
 water, correction for . . 
 
 126 
 246 
 
 Silicic acids 
 
 228 
 
 Volatile alkali . . 
 
 83 
 
 Silicon 
 Solids 
 
 226 
 
 2 
 
 Voltaic battery 
 
 112 
 
 Speciiic heat 
 Structure of flame 
 Sulphides 
 
 217 
 209 
 61 
 
 WATER 
 action of electric current 
 
 21 
 
 97 
 
 Sulphur 
 dioxide 
 Sulphuric acid 
 and metals 
 
 55 
 94 
 63 
 91 
 
 on 
 heat on . . 
 metals on . . 
 of crystallization 
 vapour, correction for 
 
 At 
 
 30 
 22 
 
 147 
 
 246 
 
 TABLE of oxides 
 salts 
 silicates 
 
 160 
 160 
 233 
 
 Weighing 
 Weights, atomic 
 Wrought iron 
 
 239 
 125 
 105 
 
 valency 
 
 133 
 
 
 
 Temperature, measurement of 
 
 215 
 
 YELLOW phosphorus 
 
 222 
 
 Thermo-chemical equation . . 
 Thermometer 
 
 201 
 215 
 
 ZINC, amalgamated 
 
 108 
 
 THE END 
 
 BILLING AND SONS, LTD., PRINTERS, CUII.DFORD. 
 
#2)