DESCiMMi.c ilNORGANIC 
 
 GENERAL CHEMISTRY 
 
 Eext^lSooft for Colleges 
 
 BY 
 
 PAUL C. FREER, M.D., PH.D. (Munich} 
 
 PROFESSOR OF GENERAL CHEMISTRY, AND DIRECTOR OF THE LABORATORY OF 
 GENERAL CHEMISTRY, UNIVERSITY OF MICHIGAN 
 
 REVISED EDITION 
 
 i ITS asms? 
 
 
 Boston 
 
 ALLYN AND BACON 
 
 1895 
 
COPYRIGHT, 1894, 
 
 BY 
 ALLYN AND BACON. 
 
 I 
 
 C. J. FETEBS & SON, 
 
 TYPE-SETTEB6 AND ELECTBOTYPEBS, 
 
 145 HIGH STBEET, BOSTON. 
 

 PEE FACE. 
 
 IN compiling the descriptive portions of this work, I have chiefly 
 used GRAHAM-OTTO'S " Lehrbnch der Allgemeinen Chemie " (last 
 edition) and LADEXBURG'S " Handworterbuch der Chemie," although, 
 wherever any of the facts which it was necessary to incorporate 
 seemed doubtful, the original sources in the chemical periodicals 
 have been consulted ; of course, the discoveries which have been 
 brought forth in the last few years are taken entirely from the 
 latter. In the discussion of the double halides, of fluosilicic acid, 
 and of similarly constituted bodies, I have adopted the views which 
 have recently been brought into prominence by the publications of 
 Prof. Ira Eemsen, both in his larger text-book and in his contribu- 
 tions to current chemical literature. In those portions of the work 
 which refer to the application of physical methods in the study of 
 chemistry, OSTWALD'S " Outlines of General Chemistry " and LOTHAR 
 MEYER'S " Theoretische Chemie " have been consulted and not infre- 
 quently quoted. My views upon the subject of valence and the 
 use of structural formulae may possibly be regarded by many of 
 my colleagues as too conservative; but I have been led to adopt 
 these views by the growing conviction that the dogmatic use of 
 supposed laws of valence and of constitutional formulse founded 
 upon very incomplete experimental evidence, is causing more harm 
 than good to the advancement of chemical science. In discussing 
 chemical changes and reactions, I have endeavored to present the 
 various topics, not as a series of isolated facts, but as so connected, 
 the one with the other, that there is scarcely any one of the numer- 
 ous phenomena which are mentioned in this work which does not 
 find its analogon in some other portion of the field of chemical 
 study. The attempt has been made especially to call attention to 
 the influence exerted by the nature of the elements which make up 
 a chemical compound upon the character of that compound itself. 
 Tracing those connections may possibly have led me somewhat into 
 
 iii 
 
IV PREFACE. 
 
 the realm of speculation, notably so, perhaps, in my endeavor to 
 explain the behavior of the hydrogen compounds of the not-metals 
 by taking into consideration the relative influence exerted by the 
 masses of the atoms which go to make the molecules. I hope, 
 however, that the new arguments ventured on during the progress 
 of this work will not be condemned without a hearing. Of course, 
 a very complete knowledge of descriptive chemistry, both inorganic 
 and organic, is necessary before the study of so-called physical 
 chemistry can be pursued with profit ; nevertheless, wherever it 
 has seemed to me that some elementary facts from the realm of 
 physical chemistry would be comprehended by the pupil taking up 
 beginning chemistry, I have not hesitated to introduce the latter, at 
 the same time giving references to the best of the smaller text- 
 books on the subject. The atomic weights which I have used are 
 taken from the table recently prepared by F. W. Clarke, with the 
 atomic weight of oxygen " en as the standard) placed at 16. 
 
 The laboratory notes j appendix cover only the ground 
 
 taken by the not-metals ; tney are not intended as a laboratory man- 
 ual, but mainly as a guide to both teacher and pupil in compiling a 
 list of experiments. Every teacher prefers using his own methods 
 for laboratory instruction, with, of course, his own selection of the 
 work to be pursued ; in my own laboratory I follow a manual which 
 is made up of brief directions, accompanied by a very complete set 
 of questions, and all of the latter must be answered by the pupils. 
 I do not think, however, that pupils should be left in the laboratory 
 without other than a printed guide ; far from it, I like to see the 
 instructor always present in the room during laboratory hours, 
 guiding and assisting his pupils, and not infrequently working 
 with them. 
 
 Probably this work cannot be advantageously employed in the 
 secondary schools ; indeed, it is adapted for the use of students who 
 already have some knowledge of the elementary principles of the 
 science. Beginners should be taken through a course in which only 
 a few elements and compounds are discussed, with the purpose of 
 familiarizing the pupils with the fundamental laws which govern 
 chemical change. During the progress of such work as this, I 
 would not advise the use of chemical symbols or any reference to 
 the atomic theory. Our chemical symbols and equations are in 
 existence, in their present state, only because of the difficult experi- 
 mental work which has finally succeeded in establishing a consis- 
 
PREFACE. V 
 
 tent table of atomic weights. It is manifestly impossible to make a 
 student, without experimental knowledge, understand, in all its 
 bearings, a theory which it has taken some ninety years to place 
 upon its present footing. If an elementary course, in which every 
 stated fact has been proved by actual experiment, precedes the work 
 given in this book, the pupil will then be amply fitted to look at 
 chemical phenomena from the basis of the atomic theory. It is in 
 the hope that such a preparatory course has gone before, that I have 
 begun this text-book with the atomic theory. 
 
 PAUL C. FREER. 
 ANN ARBOR, June, 1894. 
 
 After the first five hundred copies were printed, the book was 
 subjected to a careful reading by several chemists, so that, it is 
 hoped, all misprints and errors have .<i entirely eliminated. 
 
 PAUL C. FREEH. 
 ANN ARBOR, May, 1895. j 
 

 \ 
 
 CHAPTER PAGE 
 
 I. Introductory 1 
 
 II. Oxygen 18 
 
 III. Hydrogen 27 
 
 IY. Water 36 
 
 V. Ozone and Hydrogen Dioxide 47 
 
 VI. The Halogen^ 53 
 
 VII. Fluorine and Hydrofluoric Acid 55 
 
 VIII. jChlorjjie 58 
 
 IX. Hydrochloric Acid 66 
 
 X. Bromine and Hydrobromic Acid 78 
 
 XL Iodine and Hydroiodic Acid 83 
 
 XII. fhTOxygen Family 89 
 
 XIII. Sulphur 91 
 
 XIV. Hydrogen Sulphide 96 
 
 XV. Selenium and Hydrogen Selenide 102 
 
 XVI. Tellurium and Hydrogen Telluride; Comparative Table of 
 
 the Elements of the Oxygen Family 104 
 
 XVII. Valence and the Oxygen Compounds 107 
 
 XVIII. The Compounds of Chlorine with Oxygen, and with Oxy- 
 gen and Hydrogen 119 
 
 XIX. Compounds of Bromine and of Iodine with Oxygen and 
 
 Hydrogen, the Compound of Iodine with Oxygen, and 
 
 the Compounds of the Halogens with each other . . 129 
 
 XX. The Compounds of the Elements of the Sulphur Family with 
 
 Oxygen, and with Oxygen and Hydrogen. Sulphur 
 
 Dioxide and Sulphurous Acid 134 
 
 XXI. Sulphur Trioxide, Sulphuric Acid, and the remaining Sul- 
 phur Acids 144 
 
 XXII. The Compounds of Selenium and Tellurium with Oxygen, 
 
 and with Oxygen and Hydrogen .160 
 
 XXIII. Nitrogen and the Atmosphere 163 
 
 XXIV. Compounds of the Elements of the Nitrogen Family . . . 175 
 
 vii 
 
Vlll 
 
 CONTENTS. 
 
 CHAPTER PAGE 
 
 XXV. Ammonia and the other Compounds of Nitrogen and Hy- 
 drogen 182 
 
 XXVI. The Compounds of Nitrogen with Oxygen, and with Oxygen 
 
 and Hydrogen 195 ' 
 
 XXVII. Phosphorus and Phosphine 211 
 
 XXVIII. The Compounds of Phosphorus with the Halogens, and with 
 
 Oxygen and the Halogens 219 
 
 XXIX. The Compounds of Phosphorus with Oxygen, and with 
 
 Oxygen and Hydrogen 223 
 
 XXX. Arsenic and Arsine 234 
 
 XXXI. The Compounds of Arsenic with the Halogens, with Oxygen, 
 
 and with Oxygen and Hydrogen 238 
 
 XXXII. The Compounds of Arsenic with Sulphur, and with Sulphur 
 
 and Hydrogen 244 
 
 XXXIII. Antimony and Stibine. The Compounds of Antimony with 
 
 the Halogens | 248 
 
 XXXIV. The Compounds of Antimony with Oxygen, and with Oxygen 
 
 and Hydrogen. The Sulphides of Antimdny .... 254 
 XXXV. Bismuth. The Compounds of Bismuth with the Halogens, 
 with Oxygen, with Oxygen and Hydrogen, and with 
 
 Sulphur 257 
 
 XXXVI. The Elements of the Carbon Family 265 
 
 XXXVII. Carbon 269 
 
 XXXVIII. The Compounds of Carbon with Hydrogen 274 
 
 XXXIX. The Compounds of Carbon with Chlorine, with Chlorine 
 
 and Oxygen, with Oxygen, and with Sulphur .... 285 
 XL. Compounds of Carbon with Nitrogen, with Nitrogen and 
 
 Hydrogen, and with Nitrogen, Oxygen, and Hydrogen, 294 
 XLI. Silicon, the Compounds of Silicon with Hydrogen, and 
 
 with the Halogens, the Oxide, and Acids of Silicon . 300 
 
 XLII. Germanium and its Compounds 309 
 
 XLIII. Tin and its Compounds 312 
 
 XLIV. Lead and its Compounds 320 
 
 XLV. The Elements of the Boron Family (The Earths) .... 326 
 
 XL VI. Boron and its Compounds 328 
 
 XLVII. Aluminium and its Compounds 333 
 
 > XLVIII. Gallium, Indium, and Thallium 343 
 
 ^ XLIX. The Determination of Atomic Weights. Dulong and Petit' s 
 
 Law. The Law of Isomorphism 347 
 
 L. The Periodic System of the Elements 361 
 
CONTENTS. 
 
 IX 
 
 XLIH. 
 
 \ LIV. 
 
 LY. 
 
 ^ 
 
 CHAPTER PAGE 
 
 ^*" LI. Neutralization. Double Decomposition. Dissociation of 
 
 Electrolytes ................ 375 
 
 The Alkali Metals .............. 383 
 
 Copper, Silver, and Gold ............ 396 
 
 The Family of the Alkaline Earths ......... 411 
 
 Zinc, Cadmium, and Mercury ...... ..... 423 
 
 The Elements belonging to the Primary Groups of the 
 
 Families III., IV., and V., of the Long Periods . . . 437 
 The Elements belonging to the Primary Group of the VI. 
 
 Family ................. 443 
 
 The Element forming the Primary Group of the VII. 
 
 Family ................. 459 
 
 Iron, Cobalt, and Nickel ............ 470 
 
 The Remaining Elements of the VIII. Family. (The Plati- 
 
 num Group. ) ............... 489 
 
 APPENDIX OF LABORATORY NOTES ...... 497 
 
 LVII. 
 LVIII. 
 
 LIX. 
 LX. 
 
GENERAL CHEMISTRY. 
 
 CHAPTEE I. 
 
 INTRODUCTORY. 
 
 THE ATOMIC THEORY AND THE COMPOSITION OF CHEMICAL 
 COMPOUNDS. 
 
 AMONG the most important theories of modern physical science 
 is the one which is based upon the supposition that all substances 
 are made up of small particles called atoms. The theory that 
 matter is not infinitely divisible, but that, upon attempted separa- 
 tion into smaller parts, a mass not capable of further subdivision 
 would result, was held by the Greek and Koman philosophers 
 by Democritus, Aristotle, Epicurus, and Lucretius and has been 
 transmitted to the present generation with many important modifi- 
 cations. The idea, for we could scarcely dignify by the name of 
 theory that which had so little foundation in experiment, was 
 partially lost sight of during the dark ages of chemistry, during 
 the time of the alchemists, when the sole aim of chemical study 
 was mercenary, when scepticism on the one hand and popular 
 superstition on the other had stifled all originality of thought and 
 co-ordination of theory in this field of knowledge ; it suffered no 
 better fate at the hands of those who succeeded the alchemists, for 
 they were men who used their small knowledge of chemical facts 
 for the purpose of discovering new drugs and remedies ; it could 
 expand into what it is only when chemistry, freed from the bane 
 of superstition, began to be followed for the sole purpose of increas- 
 ing human knowledge. 
 
 We trace the growth of a science of chemistry from the begin- 
 ning of the eighteenth century, for then chemists began to have 
 theories founded on experiment; these Avere undoubtedly often 
 false and misleading, but, nevertheless, scientific progress was inevi- 
 
 1 
 
2 ATOMIC THEORY; HISTORY. 
 
 table, because of the attempts to answer the problems arising. In 
 this century fell many of the greatest discoveries of modern chem- 
 ical science /notably the proof of the existence of more than one 
 variety of gas and that of the formation of the atmosphere from 
 two kinds of matter, oxygen and nitrogen. *It was learned that water 
 could be produced by the union of oxygen and hydrogen, and that 
 substances, in burning, absorb a constituent of the atmosphere, while 
 in so doing they gain in weight, the gain in weight of the burning 
 substance being exactly equal to the loss in weight sustained by 
 the atmosphere^ To the knowledge won at this time we owe our 
 understanding of a principle of nature upon which all chemical 
 speculations are based; that of the conservation of matter. The 
 English chemists, Black, Priestley, and Cavendish, were the men 
 whose efforts developed so many new facts; but it was owing to 
 the clear insight into the meaning of these discoveries obtained by 
 Antoine Laurent Lavoisier, that a greater service was rendered to 
 humanity, because by him a proper explanation of the phenomena 
 involved was given. Without the discovery of oxy_gen by Priestley, 
 or the composition of water by Cavendish, Lavoisier might not 
 have proved the law of the conser^ajioji_oj_jnatter nor have estab- 
 lished the theory of combustion held at the present time ; but it is 
 equally true that, without Lavoisier's genius the work of the Eng- 
 lish scientists would not have accomplished the result of preparing 
 chemistry for the unprecedented advance recorded of it in the nine- 
 teenth century. During the time of these great discoveries the 
 atomic jheory, though tacitly accepted, was not made the basis of 
 investigation ; but when the present century dawned, chemists 
 began to feel the need of some rational explanation of those 
 phenomena which most concerned them. 
 
 The first decade brought the discovery that when two substances 
 unite chemically, a compound is always formed in unvarying pro- 
 portions by weight. Thus, iron and sulphur unite to form iron 
 sulphide, a substance in which there are four parts by weight of 
 sulphur for every seven parts by weight of iron, no matter where 
 or how the combination takes place ; and there is also another com- 
 pound of iron and sulphur in which seven parts by weight of iron 
 accompany eight of sulphur. Under whatever conditions, or in 
 whatever place, either of these chemical bodies are produced, the 
 resulting proportions are always the same. If there is more sulphur 
 
LAW OF DEFINITE PROPORTIONS. 3 
 
 present than is necessary for combination, then the excess of sul- 
 phur remains unchanged, and if more iron is employed, then iron is 
 found after the union. We have two compounds of carbon and 
 oxygen, called oxides for reasons similar to those which gave the 
 name of sulphides to the compounds of iron and sulphur. In one 
 of these, six parts by weight of carbon are united with eight parts 
 of oxygen ; in the other, six parts of carbon are united with sixteen 
 parts of oxygen. What is true as regards iron, carbon, oxygen, and 
 sulphur characterizes the multitude of other substances which have 
 been studied with the object of ascertaining the relative proportions 
 by weight in which the constituent parts unite. To the discoveries 
 outlined above, have been added the demonstration that the relative 
 proportions of sulphur and oxygen, for instance, are preserved in 
 whatever compound they are encountered. Thus, sulphur and 
 oxygen form two compounds, called oxides of sulphur; in one of 
 these, two parts by w r eight of sulphur unite with two of oxygen, 
 in the other two of sulphur with three of oxygen ; but the relation- 
 ship in the weights of oxygen and sulphur in the various compounds 
 cited first becomes apparent if we calculate the weights, placing sul- 
 phur at 16, while preserving the proportion between the various 
 parts, thus : 
 
 Iron and sulphur 28 parts of iron unite with 16 of sulphur. 
 
 a 44 it no .. .. .. .. . . oo . . . . 
 
 Carbon and oxygen ...... 6 " " carbon " " 8 " oxygen. 
 
 It it it a 4. .. .. .. .. "JiJ 44 44 
 
 Sulphur and oxygen .... 16 " " sulphur " " 16 " " 
 " 16 " " " " " 24 " 
 
 These relations represent actual facts, whatever explanation we 
 may see fit to attach to them ; but such facts as these necessarily 
 give rise to speculations as to the underlying causes. Why should 
 it not be possible to have 28 parts of iron united, at one time with 
 16 parts of sulphur, at another with 17 parts, at yet another with 
 15 parts ? As chemists could see no reason for such regularity in 
 the composition of matter, the facts themselves were at first dis- 
 puted, until repeated experiment rendered them incontrovertible. 
 Assuming the constancy of proportion in chemical compounds, even 
 before such constancy was proved, the English chemist, John 
 Dalton, sought an explanation in the following hypothesis, which 
 has been accepted as a basis for chemical speculation ever since its 
 
4 MODERN ATOMIC THEORY. 
 
 establishment, and which is here given in the form at present 
 accepted. 
 
 Matter is not infinitely divisible, but is composed of very small 
 and discrete entities called atoms, there being as many different 
 kinds of atoms as there are varieties of substance which have never 
 been decomposed into two or more forms with differing properties. 
 The elements having weight, the atoms, being portions thereof, ne- 
 cessarily also have weight ; and we assume the weight of an atom 
 of a given element to be equal to that of each other atom of the 
 same element, but to differ from that of an atom of any other 
 element. The atoms of different elements unite to form the 
 smallest individual group of a compound, and atoms of the same 
 kind unite to produce the smallest individual group of an element. 
 These groups are known as molecules, and the agglomeration of 
 molecules forms tangible matter. The weight of any given 
 molecule, called its molecular weight, is therefore equal to the sum 
 of the weights of its constituent atoms. If I subdivide any com- 
 pound body, water for instance, I can continue the operation until 
 I arrive at the smallest individual particle thereof, a molecule ; if I 
 divide this, I no longer have water, but two different kinds of mat- 
 ter, hydrogen and oxygen. This illustration will also serve to show 
 the difference between a so-called chemical and a physical change. 
 Water can be decomposed into its molecules with comparative ease ; 
 by changing it into steam these particles are so far separated that 
 they travel in right lines independently of each other. A much 
 greater heat than is necessary for the production of steam (or the 
 application of some other form of energy, such as electricity) is 
 necessary to effect any further change, and this change brings with 
 it the destruction of the nature of the substance in question. Water 
 is no longer present ; but in its place we have two different kinds of 
 matter, hydrogen and oxygen, so that a chemical change has been 
 brought about. 
 
 With the atomic hypothesis in view, the constant composition 
 by weight of compound substances is readily explained. For 
 instance, a molecule of iron sulphide is composed of atoms of iron 
 and of sulphur, each molecule containing the same number of atoms 
 of the two elements. If all atoms of iron are alike in weight, and 
 if all atoms of sulphur bear the same relationship to each other, it 
 follows that every molecule of iron sulphide must have the same 
 
LAW OF MULTIPLE PROPORTIONS. 5 
 
 composition as every other molecule of the same substance; and 
 from this it follows that tangible quantities of iron sulphide, which 
 are simply agglomerations of the individual molecules, must have 
 the same proportional composition by weight as these individual 
 particles. By accepting the theory as it has been outlined, the 
 unvarying composition of purely chemical compounds necessarily 
 follows. Of course, two or more substances may be mixed in any 
 proportion, but such a mixture does not have the, characteristics 
 of a chemical compound. The various constituents of a mixture 
 can be separated with greater or less ease by simple mechanical 
 operations ; but as soon as a chemical compound is formed from the 
 various parts of any mixture, a substance having a definitely con- 
 structed molecule results. 
 
 We saw that iron and sulphur are, however, capable of forming 
 two compounds with each other. In comparing these two sulphides, 
 it has been discovered that the part by weight of sulphur which 
 unites with a given weight of iron in one of these substances is 
 exactly twice that which is found united to the same quantity of 
 iron in the other ; and repeated, painstaking experiments have dis- 
 covered a great number of similar cases, in which one element forms 
 tAvo or more distinct compounds with some other. In comparing 
 such compounds it has ahvays been shown that the parts by weight 
 of one of the elements, Avhen united to a fixed quantity by Aveight 
 of the other, bear a simple relationship to each other; for 
 example : 
 
 Mercury and oxygen form two oxides ; in one, for every 100 parts of mer- 
 cury there are 4 parts of oxygen ; in the other, for every 100 parts of mercury 
 there are 8 parts of oxygen. 
 
 Nitrogen and oxygen form five oxides ; in these compounds the parts by 
 weight of oxygen which are united with 100 parts by weight of nitrogen 
 are as 57. 1 : 114.2 : 171.3 : 228.4 : 285.5, which figures are to each other as 
 1 :2 :3 :"4 : 5. 
 
 The results of these discoveries can be summed up in the 
 folloAving law of .multiple proportions : 
 
 If two elements, a and b, unite in more than one proportion, the 
 parts by weight of b Avhich Avill unite Avith a definite quantity of a 
 Avill be in simple ratio to each other.* 
 
 * See John Dalton, New System of Chemical Philosophy (1808). 
 
6 STOICHIOMETEIC QUANTITIES. 
 
 This law of multiple proportions is readily explained by the 
 atomic hypothesis. For, let us suppose, using the two sulphides of 
 iron as an example, that the one composed of 28 parts of iron to 16 
 parts of sulphur has a molecule constructed of one atom of iron 
 and one atom of sulphur. In order to change this molecule into one 
 containing more sulphur, the only possible means is by the addition 
 of another atom of sulphur. But as the atoms of sulphur all have 
 the same weight, it follows that the amount of sulphur in the 
 newly formed molecule must be to that in the original as 2 : 1. We 
 might represent the change graphically as follows, using the black 
 circle to represent an atom of iron, the white one sulphur. 
 
 28 16 16 16 28 16 
 
 O + O 000 
 
 Molecule of iron sulphide + 1 atom of sulphur = Molecule No. 2. of iron 
 sulphide. 
 
 We could have come to the same conclusions, deducing the law 
 of multiple proportions as a necessary consequence of the atomic 
 structure of matter, had we used combinations of any other 
 elements ; for the law is universal in its application. What is true 
 of the individual molecule must also be true of tangible matter. 
 
 The proportions by weight in which the elements unite are 
 called their sto'ichiometric quantities ; these belong not only to 
 individual pairs of elements, but are universal. This truth at 
 once becomes apparent if we select a given weight of some one 
 element and then determine, experimentally, the parts by weight 
 in which all other elements will unite with it : by this means it 
 will be found that the quantities in which the various elements unite 
 with the standard are also the relative proportions, or multiples or 
 submultiples of the proportions, in which they will unite ivith each other. 
 These facts are in entire accordance with the atomic theory j for as 
 all chemical reactions take place between the atoms, and as the 
 atoms of a given element have a constant, equal mass, different from 
 the masses of the atoms of all other elements, it follows that the 
 relative proportions between the parts by weight in which the 
 elements combine must also be constant. Were all chemical com- 
 pounds to be of the simplest nature, so that each molecule would be 
 formed by the union of but two atoms, the determination of the 
 relative weights of the atoms would resolve itself into the simple 
 
STANDARD FOR ATOMIC WEIGHTS. 7 
 
 problem of determining the relative parts by weight in which the 
 elements unite ; that this is not the case, however, will at once be 
 seen from the fact that the same elements can enter into two or 
 more compounds. One thing must be true, however the stoichio- 
 metric quantities, of necessity, bear some simple relationship to the 
 relative weights of the atoms themselves ; and it must be equally 
 true that, if it is at all possible actually to determine the relative 
 numbers which are to be assigned to the atoms, such determinations 
 must be based primarily on a discovery of the stoichiometric quan- 
 tities.* The subsequent task is the one of calling to our aid all pos- 
 sible guides, physical or chemical in nature, which will lead us to 
 agree upon which ones of the various multiples, or sub-multiples, of 
 the stoichiometric quantities are to be selected as expressing the 
 true relative weights of the atoms. 
 
 The methods by which such an agreement has been brought 
 about, being of a somewhat complicated nature, are left for discus- 
 sion to a subsequent part of the book.t 
 
 The selection of a standard by which all other weights can be 
 compared is as necessary in dealing with atoms as it is in the men- 
 suration of distance, it being immaterial what standard is selected, 
 provided all of the weights can easily and accurately be compared 
 with this. During the first years of our atomic hypothesis, the 
 weight of the atom of hydrogen, being the smallest appertaining to 
 any element, was selected as unity ; but subsequently this practice 
 was abandoned in favor of oxygen, the weight of the atom of which 
 was placed at one hundred. Hydrogen once more resumed its 
 original position during the middle of the century and, until 
 recently, all weights of atoms (atomic weights), were compared with 
 this. If we call to our aid certain theories concerning the nature 
 of gases, a consideration which must be deferred until the pupil 
 has become acquainted with a larger number of chemical facts, we 
 can place the ratio between the atomic weights of hydrogen and 
 
 * The absolute weights of the atoms, being extremely small fractions of a 
 milligramme, are quantities not obtainable with any degree of accuracy. 
 
 t As these determinations involved the most painstaking and difficult 
 manipulations, as well as very advanced views as regards theoretical deduc- 
 tions from certain physical facts, it of necessity followed that wide differences 
 of opinion, only disappearing within the most recent times, were manifested ; 
 indeed, absolute certainty is not even now attained or attainable as regards 
 the atomic weights. 
 
8 
 
 ATOMIC WEIGHTS. 
 
 oxygen at 1 : 15.88, a number which very nearly coincides with 
 1 : 16. Considerable uncertainty exists as to the accuracy of this 
 ratio, for recent investigation has altered the relation repeatedly. 
 If atomic weights are referred to hydrogen as unity, a recalculation 
 of all atomic weights is necessary whenever investigation shows 
 the accepted ratio between oxygen and hydrogen to be untenable ; 
 for these constants have been determined (for the greater number 
 of elements), directly or indirectly, by an investigation of com- 
 pounds with oxygen. It seems more advisable, therefore, to adopt 
 \ oxygen as a standard, and, so as not to depart too far from numbers 
 -rendered familiar by accepted usage, to place the atomic weight of 
 this element at 16. By this means the atomic weight of hydrogen 
 becomes 1.008, a number which, for all practical purposes, can be 
 placed at unity. If any further correction in the ratio between 
 hydrogen and oxygen becomes necessary, such a change will in- 
 volve no further calculation. The methods by which the atomic 
 weights have been determined are not a subject for discussion at 
 the present time ; indeed, the great majority of them would be 
 entirely out of place in an elementary treatise ; suffice it to say, 
 that so complete has been their application that the weights which 
 are placed in the following table * are, with unimportant exceptions, 
 accepted as correct by all chemists : 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 ALUMINIUM 
 
 Al 
 
 27 
 
 Erbium 
 
 Er 
 
 166.3 
 
 ANTIMONY . . 
 ARSENIC. . . . 
 BARIUM 
 
 Sb 
 As 
 Ba 
 
 120. 
 75. 
 137 43 
 
 Fluorine .... 
 Gadolinium . . 
 Gallium 
 
 F 
 
 Gd 
 Ga 
 
 19. 
 
 156.1 
 69. 
 
 Beryllium . . . 
 BISMUTH 
 
 Be 
 Bi 
 
 9. 
 
 208 
 
 Germanium . . 
 GOLD 
 
 Ge 
 
 Au 
 
 72.3 
 197.3 
 
 BORON 
 
 B 
 
 11. 
 
 HYDROGEN. . . 
 
 H 
 
 1.008 
 
 BROMINE 
 
 Br 
 
 79 95 
 
 Indium 
 
 In 
 
 113.7 
 
 
 Cd 
 
 112 
 
 IODINE . . 
 
 I 
 
 126.85 
 
 Caesium. 
 
 Cs 
 
 132 9 
 
 Iridium 
 
 Ir 
 
 193.1 
 
 CAI CIUM 
 
 Ca 
 
 40 
 
 
 Fe 
 
 56. 
 
 CARBON .... 
 CERIUM 
 
 C 
 
 Ce 
 
 12. 
 
 140 2 
 
 Lanthanum . . 
 LEAD 
 
 La 
 Pb 
 
 138.2 
 206.95 
 
 CHLORINE . . . 
 CHROMIUM . . 
 Cobalt 
 
 Cl 
 Cr 
 Co 
 
 35.45 
 52.1 
 59. 
 
 LITHIUM .... 
 MAGNESIUM . . 
 MANGANESE . . 
 
 Li 
 Mg 
 Mn 
 
 7.02 
 24.3 
 55. 
 
 Columbium . . 
 COPPER .... 
 
 Cb 
 
 Cu 
 
 94. 
 63.6 
 
 MERCURY . . . 
 Molybdenum . . 
 
 Hg 
 Mo 
 
 200. 
 96. 
 
 * Compiled by F.W. Clarke; Jan. 1, 1894. Jour, of the Am.Chem.Soc.16; No. 3. 
 
NUMBER OF ELEMENTS. 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 NAME. 
 
 SYMBOL. 
 
 ATOMIC 
 WEIGHT. 
 
 
 Nd 
 
 140 5 
 
 SODIUM 
 
 Na 
 
 23.05 
 
 ]N ickel 
 
 Ni 
 
 58 7 
 
 Strontium . . . 
 
 Sr 
 
 87.6 /// 
 
 NITROGEN. . . 
 Osmium .... 
 OXYGEN .... 
 Palladium . . . 
 PHOSPHORUS . 
 Platinum . . . 
 POTASSIUM . . 
 Praseodymium 
 
 N 
 Os 
 
 Pd 
 P 
 Pt 
 K 
 Pr 
 
 14.03 
 190.8 
 16. 
 106.6 
 31. 
 195. 
 39.11 
 143 5 
 
 SULPHUR. . . . 
 Tantalum . . . 
 Tellurium . . . 
 Terbium .... 
 Thallium .... 
 Thorium .... 
 Thulium .... 
 Tin 
 
 S 
 Ta 
 Te 
 Tb 
 Tl 
 Th 
 Tm 
 Sn 
 
 32.06 
 182.6 
 125. * 
 160. 
 204.18 
 232.6 
 170.7 
 119. 
 
 Rhodium . . . 
 Rubidium . . . 
 Ruthenium . . 
 Samarium . . . 
 Scandium . . . 
 Selenium . . . 
 SILICON 
 
 Rh 
 Rb 
 Ru 
 
 Sm 
 
 Sc 
 Se 
 Si 
 
 103. 
 85.5 
 101.6 
 150. 
 44. 
 79. 
 28 4 
 
 Titanium .... 
 Tungsten .... 
 Uranium .... 
 Vanadium . . . 
 Ytterbium . . . 
 Yttrium .... 
 ZINC 
 
 Ti 
 W 
 U 
 V 
 Yb 
 Yt 
 Zn 
 
 48. 
 184.9 
 239.6 
 51.4 
 173. 
 89.1 
 65.3 
 
 SILVER .... 
 
 Ag 
 
 107.92 
 
 Zirconium . . . 
 
 Zr 
 
 90.6 
 t 
 
 We are acquainted with sixty^seven different kinds of matter, 
 no individual variety of which has been decomposed into two 
 or more simpler forms ; but whether such decomposition will 
 ever occur, it is impossible to state. By the union of these elements 
 all substances known to us are produced. By far the greater pro- 
 portion of matter, being composed of molecules containing two or 
 more atoms differing from each other in kind, is compound in its 
 nature. The individual atoms do not, except in rare instances, 
 exist as such; they are united to form molecules, the difference 
 between the molecule of the element and that of the compound 
 being that, while in the former the atoms are all of the same kind, 
 in the latter they differ. Atoms are grouped together to form 
 
 * According to Brauner (Monatshefte fiirChemie; 10, 445) tellurium has 
 an atomic weight of 127.6. Brauner observes, however, that the substance 
 which has heretofore been considered as an element probably contains other 
 elements as well. The true atomic weight of tellurium he believes will be 
 found to be between 125 and 126. 
 
 t The ratio between the atomic weights of hydrogen and oxygen is 
 1 : 15.88. The term glucinum is frequently used instead of beryllium. The 
 more important elements are in capitals. As a matter of expediency the 
 pupil should memorize the atomic weights of a few of the more important 
 elements. The acquirement of this knowledge is best deferred until the indi- 
 vidual elements are discussed, when the weights can be learned during the 
 progress of the study. 
 
10 CHEMICAL AFFINITY, CHEMICAL ENEKGY. 
 
 molecules which are more or less stable, and this stability of equi- 
 librium must be brought about by some force acting between the 
 individual atoms. This force has been compared to the attraction of 
 gravitation, and has by some been considered to be identical with 
 it. The attraction of gravitation, however, is capable of manifesta- 
 tion between bodies at a great distance from each other, while the 
 number of bodies acted on in this manner by any given body is 
 unlimited. The attraction between the atoms seems capable of 
 manifestation only through an extremely small interval of space, 
 and then only between a limited number of atoms. A new term is 
 therefore necessary to designate this force which holds the atoms 
 in equilibrium in the molecule, and, for want of better ones, the 
 expressions " chemical affinity," or " chemism," are used. Where 
 a very stable compound exists, the atoms composing it are said to 
 have a great affinity for each other. An inquiry into the relative 
 stability of chemical compounds is of the greatest importance and, 
 consequently, will frequently be made during the progress of this 
 work ; but it must not be forgotten that the term " chemical 
 affinity " is used simply to designate a force which has never been 
 resolved into simpler factors, and of the nature of which we are 
 consequently ignorant. 
 
 When a substance, either by reason of its position or of its 
 motion, is capable of performing work, it is said to possess energy, 
 and work is done when a resistance is overcome. The atoms 
 possess energy because they are, in uniting, capable of performing 
 work, by reason of their chemical affinity. Illustrations of this 
 performance of work by the union of atoms are familiar. The 
 motions of machinery driven by steam can, with the greatest ease, 
 be traced to the chemical union of the oxygen of the atmosphere 
 with the coal under the boilers; and the movements of animals 
 can in the same way be shown to be derived from the chemical 
 energy of the various substances which form the nutriment of the 
 body. The energy possessed by the individual atoms can be 
 likened to potential energy (energy of position), for by its means 
 the atoms are capable of performing work, just as is a stone when 
 raised above the level of the earth. 
 
 The measure of work is the force (P), which overcomes resistance, into 
 the distance (<S) through which this force acts (L= P S). The amount of 
 work which a body is capable of performing, by reason of its position, is called 
 
CHEMICAL ENERGY. 11 
 
 its potential energy. If the work (i) of the force (P) is neutralized by a 
 negative force (P 7 ), L is not lost, for P can perform its work () as soon as 
 P' ceases to act. A body of the weight P, which has been lifted through the 
 distance S, has an amount of potential energy equal to PS, for as soon as it 
 is dropped, it can perform the work PS. The energy possessed by the atoms 
 is to us different from potential energy, because it has never been resolved 
 into the factors PS. The amount of work which a body in motion is capable 
 of performing by reason of that motion, is called its kinetic energy (energy of 
 motion). This is equal to one-half the product of the mass of the body and 
 
 the square of its velocity (L = ~ ). If a part of the potential energy pos- 
 sessed by a body has been used, and thereby a certain amount of kinetic 
 energy has been produced, then the sum of the remaining capacity for work 
 and of the produced kinetic energy is equal to the original amount of potential 
 energy; or, what is the same thing, equal to the kinetic energy which would 
 have been produced if the entire store of potential energy had been used to 
 perform work. If the stone, which we had supposed to be suspended, had 
 been allowed to fall to the ground, the entire content of potential energy 
 would have been changed to heat, which is a form of kinetic energy. These 
 few remarks illustrate the principle of conservation of energy, which was 
 denned by Mayer in 1842. 
 
 What is true of the stone is also true of the atoms, for their 
 potential energy, or better, chemical energy (a term which will in- 
 volve us in no contradictions), can be transformed into kinetic 
 energy upon their union to form a compound. While we cannot 
 measure the amount of chemical energy possessed by the atoms, we 
 are able to measure the kinetic energy which is produced when 
 they unite by reason of their chemical affinity, for by far the 
 greatest amount of this kinetic energy is manifested as heat. If 
 the stone, which we have used as an illustration, falls through 
 a certain distance, its potential energy is converted into kinetic 
 energy, and to once more raise this stone to its original position 
 will require an equivalent expenditure of work. Just so with the 
 atoms. If two atoms have a great chemical affinity for each other, 
 they possess a great amount of chemical energy, which will be con- 
 verted into an equally great amount of kinetic energy when they 
 unite; and, in order to separate the molecule so formed into the 
 original atoms, we will have to apply the same amount of energy 
 in some form ; such a molecule is as a consequence stable, and pos- 
 sesses an amount of chemical energy much smaller than that of its 
 constituent atoms. On the other hand, two or more atoms may 
 possess very little chemical affinity for each other, so little, indeed, 
 
12 METALS AND NOT-MET ALS. 
 
 that an expenditure of energy is necessary to cause them to unite ; 
 therefore, the molecule caused by such a union will have a tendency 
 readily to convert its acquired chemical energy into some form of 
 kinetic energy, and will be unstable. Molecules which have been 
 formed with an evolution of heat, and which therefore possess less 
 energy than their constituent atoms, are said to be the product of 
 an exothermic reaction ; while those which have resulted from an 
 absorption of heat are the result of an endothermic one. -*The heat 
 given off or absorbed in these changes has, in many cases, been 
 measured, and is taken as an indication of the chemical affinity of 
 the atoms. The elements show great variations in this respect; 
 some have a great chemical affinity for each other, and thus form 
 stable compounds, others have very little, while some have none at 
 all, in which case no expenditure of energy will cause them to unite. 
 We must bear in mind that any given element varies much in its 
 affinity toward the various other elements, and also that what we 
 have explained as regards compounds formed between individual 
 atoms is true as regards compounds formed from two or more mole- 
 cules. The various phases of chemical reaction which we will sub- 
 sequently study, will, give abundant opportunity for returning to 
 this subject. 
 
 The elements are divided into two classes, the most marked 
 representatives of each of which exhibit the sharpest possible 
 chemical contrast toward the other; one of these classes is 
 familiar to all of us : it is that of the metals, with the superficial 
 qualities of which substances, such as metallic lustre, malleability 
 and ductility, we are tolerably well acquainted ; in addition to these 
 properties, metals are also, as a rule, good conductors of heat and 
 of electricity. At the opposite chemical extreme we find a class of 
 elements which can best be termed not-metals. A few of these are 
 gases, and thereby they differ from the metals,^ only one of which, 
 hydrogen, exists in this state at ordinary temperatures. The not- 
 metals, when solid, are colored bodies which are brittle, neither 
 malleable nor ductile, and either non-conductors of electricity and 
 poor conductors of heat, or at least they possess these two proper- 
 ties in a degree much inferior to that of the metals. A number of 
 not-metals, on superficial examination, will appear to have metallic 
 lustre, yet a closer inspection reveals the fact that this is only 
 .apparently the case, for, in forming a thin section of the element, 
 
NEGATIVE AND POSITIVE ELEMENTS. 13 
 
 we find the same to be capable of transmitting light.* Between the 
 two extremes we have a considerable number of elements which 
 may have the characteristics of both metals and not-metals ; 
 examples of such elements are arsenic and tellurium ; these sub- 
 stances, while possessing a distinctly metallic appearance, have 
 chemical characteristics, which would cause them to be more prop- 
 erly classed as not-metallic substances. The most pronounced metals 
 have a marked chemical affinity for the characteristic not-metals, 
 and form stable compounds with them ; the less apparent the 
 contrast in the elements, the more easily decomposed will, in many 
 cases, the compound resulting from the union of such elements be. 
 This statement must not, by any means, be taken as a general rule, 
 for we find the most stable substances resulting from the union of 
 atoms of the same element, as is the case in the molecules of chlo- 
 rine and of hydrogen, and also in molecules formed from elements 
 which closely resemble each other, as in the compounds of sulphur 
 and oxygen. If a compound formed of a metal and a not-metal is 
 subjected to the action of an electric current, the metal will separate 
 at the negative pole, the not-metal at the positive one ; and for this 
 reason the metals are called electro-positive, while the not-metals are 
 electro-negative. This rule was formerly supposed to be universal 
 in its application, and a system of chemistry was then established 
 which had for its basis a theory that all compounds are formed by 
 the union of a negative and a positive element to form a molecule, 
 and that the more complex substances are then produced by the 
 union of a negative and positive molecule. While the negative or 
 positive character of the elements composing a compound is undoubt- 
 edly of great influence upon the nature of the resulting body, yet 
 a system founded exclusively on these characteristics has proved to 
 be untenable ; convenience, however, causes us to retain the expres- 
 sions " negative " and positive " the term negative being used 
 to designate all of the properties which characterize the not-metal, 
 and the term positive those of the metal. 
 
 There are certain chemical characteristics which enable us to 
 draw the line between metal and not-metal with tolerable distinct- 
 ness. The metals all form compounds with oxygen, the not-metals, 
 (with two exceptions), fluorine and bromine, do the same ; these 
 
 * Some metals (gold foil) can imperfectly transmit light when hammered 
 to the very thinnest possible leaf. 
 
14 CHEMICAL SYMBOLS. 
 
 compounds are known as oxides. The oxides of the metals, in 
 most cases, are classed as bases, a term which is also applied to 
 compounds of the metals with oxygen and hydrogen ; the oxides of 
 the not-metals are termed the anhydrides of acids, if, on addition 
 of water, they are capable of forming acids. The contrast between 
 metal and not-metal becomes apparent when these oxides interact 
 chemically, for then a new compound, a salt, is produced. Thus 
 the oxide of potassium, a base, when brought in contact with the 
 .oxide of sulphur, an anhydride (sulphur trioxide) produces a salt, 
 the sulphate of potassium. We shall more thoroughly comprehend 
 the meaning of these terms as we become better acquainted with 
 chemical facts. 
 
 The writing of the innumerable chemical changes which take 
 place would be complicated and a co-ordination of the phenomena 
 rendered difficult, were not some system of notation employed 
 which would express the reactions without the trouble of writing in 
 full the names and atomic weights of the various elements. Feeling 
 this. need, some designation of the elements by signs has been em- 
 ployed since the time of the alchemists. The system introduced 
 by the Swedish chemist, Berzelius, is the one which has been found 
 to best answer all of the requirements, and is in use at the present 
 time. In this the various elements are designated by the first letter 
 of their English or Latinized name in capitals, or, where conflict 
 would arise, by the first letter of the name in capitals followed by 
 some other letter, usually the next following, in small type. Thus, 
 hydrogen has the symbol H; oxygen, 0; sodium, Na, from the 
 Latinized natrium; mercury, Hg, from the Latin hydrargyrum. 
 These symbols do not stand for the visible elements, but represent 
 the atoms, and hence include the idea of atomic weights. Thus O 
 stands for an atom of oxygen, and means that, as the atom of 
 oxygen is sixteen times as heavy as the atom of hydrogen, we have 
 sixteen times as much oxygen by weight as hydrogen, when the 
 latter is expressed by the symbol H. By writing the symbols side 
 by side, we express a chemical compound, as, for instance, H 2 0, 
 which stands for the compound water ; the number 2 placed after 
 and below the letter H meaning that in one molecule of water 
 there are two atoms of hydrogen, the whole combination meaning 
 that in H 2 we have two atoms of hydrogen united to one of 
 oxygen, in the ratio of 2 to 16 by weight ; H 2 being called the 
 
CHEMICAL EQUATIONS. 15 
 
 A 
 
 formula of water, and the sum of the atomic weights (represented 
 by 18), the formula weight. In the case of water, and in the case 
 of many other substances which can be obtained as gases, this 
 formula weight is also the weight of the molecule, and is therefore 
 the molecular weight. The combination of letters, NaCl, means 
 that one atom of sodium is united to one atom of chlorine to form 
 a formula weight of the compound (sodium chloride) ; at the same 
 time it indicates that twenty-three parts by weight of sodium are 
 united to thirty-five and a half parts of chlorine, these numbers 
 representing respectively the atomic weights of sodium and of 
 chlorine. Whether NaCl represents the molecular weight of 
 sodium 'chloride we cannot state, for the substance may be composed 
 of molecules formed by the union of a number of formula weights, 
 Na 01 ; however, it is extremely probable that, if such be the case, 
 the molecule of sodium chloride is produced by the union of a 
 number of entities which all have the composition N"aCl, so that 
 there is no great objection against using the terms molecular weight 
 and formula weight interchangeably in many cases.* 
 
 Qur present knowledge seems to show that the total amount of 
 matter contained in the universe never varies. The atoms compos- 
 ing chemical compounds may change their position or manner of 
 grouping, they may be transferred from one compound to another, 
 or the compound may be decomposed into its elements, but the 
 atoms can neither be created nor destroyed. The entire science of 
 chemistry is based upon the law of the indestructibility of matter 
 which was first understood by Lavoisier, and every chemical change 
 since his time has but served to prove its existence. As a conse- 
 quence of this law, the compounds produced by the interaction of 
 elements must equal in weight the amounts taken before the 
 reaction, and the sum of the weights of the elements or compounds 
 produced by the decomposition of a compound substance must equal 
 the weight of the substance originally employed. As our symbols 
 are used to express atomic weights as well as other characteristics 
 of the elements, it follows that we can bring any chemical reaction 
 into the form of an equation ; thus H -J- Cl = H Cl means that 
 hydrogen plus chlorine yields a compound of hydrogen and chlorine 
 (hydrochloric acid), and it also means that one part by weight of 
 
 * A number of cases where the formula weight and the molecular weight 
 are not identical will be encountered in the subsequent portions of this work. 
 
16 CHEMICAL EQUATIONS. 
 
 hydrogen plus 35.5 parts by weight of chlorine (the latter number 
 being the atomic weight of that element), equal 36.5 parts by 
 weight of hydrochloric acid. 
 
 The formula HgO = Hg -j- indicates that the oxide of mer- 
 cury is decomposed into mercury and oxygen, and that the weight 
 of the oxide of mercury used is equal to the sum of the weights of 
 oxygen and mercury produced. 
 
 C 2 4 H 2 = C0 2 + CO+H 2 
 
 means that a compound of carbon, hydrogen, and oxygen (oxalic 
 acid) can be decomposed into C0 2 (carbon dioxide), CO (carbon 
 monoxide), and H 2 (water), and that the weight of the oxalic acid 
 taken is equal to the sum of the weights of carbon dioxide, carbon 
 monoxide, and water produced. The atomic weights of carbon, 
 oxygen, and hydrogen are 12, 16 and 1 respectively ; hence 
 (2 X 12) -f (4 X 16) -f- 2 = 90 parts by weight (of oxalic acid) 
 yield, 44 -f- 28 + 18 = 90 parts by weight of carbon dioxide, carbon 
 monoxide, and water. 
 
 Some doubts will arise in the beginner's mind as to the legiti- 
 macy of the numbers used as atomic weights ; he will naturally ask, 
 why is the atomic weight of carbon 12 and not 6 ? Why, for 
 instance, should we suppose the formula of carbon dioxide to be 
 C0 2 and not C 2 2 ? These very doubts were entertained by chem- 
 ists during the first six decades of the century, and only by pains- 
 taking investigation, with the aid of a number of physical laws, 
 have they been removed. At the present period of our study we 
 must accept the atomic weights as they are given ; a discussion of 
 the various reasons for this acceptance would be far beyond the 
 scope of an elementary treatise ; but such as can be understood will 
 be introduced at a subsequent time. Chemists, as a rule, are scep- 
 tical to a degree, and the fact that our present atomic weights are 
 believed in by the great majority must be a sufficient guarantee for 
 their accuracy. The pupil must remember that chemical symbols 
 stand for material substances ; he must not be led to consider the 
 element nitrogen, for instance, as being simply the symbol N, nor 
 potassium, the symbol K, but should always bear in mind that these 
 letters stand for the atoms of existing elements, endowed with various 
 properties, which elements and properties would be in existence 
 even if no system of chemical notation had ever been established. 
 
ARRANGEMENT OF THE ELEMENTS. 17 
 
 The elements are not substances any one of which differs in 
 properties from all of the others; indeed, no one element exists 
 which does not exhibit characteristics which mark its resemblance 
 to a number of its fellows. Roughly speaking, the elements are 
 divided into metals and not-metals, the metals all bearing a certain 
 resemblance to each other, and the not-metals all being more or less 
 alike ; yet both metals and not-metals are naturally divided into 
 groups, of which the individual members have a strong family like- 
 ness, while these various groups also have a number of points in 
 common. We shall see, at a subsequent time, that, by arranging 
 the elements in the order of their atomic weights, beginning with 
 the one which has the smallest number, this resemblance will 
 become most apparent. By constructing a table of the elements 
 arranged in this manner, we have produced a system of classifica- 
 tion, a careful study of which has shown us that the characteristics 
 of any element are determined by the atomic weight of that 
 element ; so that if we are thoroughly familiar with this arrange- 
 ment, we can describe the properties of an element with tolerable 
 accuracy, without any previous knowledge of its chemical deport- 
 ment, provided that we are acquainted with its atomic weight. 
 This system will provide the order in which the various elements 
 will be discussed. 
 
 Having by this brief introduction learned some of the funda- 
 mental theories upon which our science is based, we will now go on 
 to the special descriptive portion of chemistry. In so doing we 
 shall find it expedient to become acquainted first with a typical note 
 metal, then with a metal, and immediately following this with the 
 compound produced by the chemical union of these two contrasting 
 elements. The not-metal will be oxygen, the metal, hydrogen, and 
 the compound, water. 
 
18 OXYGEN; OCCURRENCE, HISTORY. 
 
 CHAPTER II. 
 
 OXYGEN. 
 
 Symbol, ; atomic weight, 16 ; specific gravity, air = 1, is 1.10535, 
 H 2 = 2, is 31.76 ; 1 c.c. at and .76m. pressure weighs 
 .00142952 gram. 
 
 OXYGEN is the element which occurs in greatest quantity upon 
 our planet ; it forms about 47.3 per cent of the solid portion of the 
 earth, 85.8 per cent of the ocean, and, including its occurrence in 
 the atmosphere, about 50 per cent of the total substance of the 
 globe. It is present in the atmosphere as uncombined oxygen, 
 mixed with nitrogen and a few other gases, and constitutes one- 
 fifth of the entire gaseous envelope of the earth. United with 
 hydrogen it produces water, forming eight-ninths of that substance ; 
 by far the greater portion of the crystalline rocks contain oxygen 
 combined with other elements ; the soil and the various forms of 
 vegetable and animal life contain this element chemically united, 
 and also to some extent as free oxygen. 
 
 Oxygen was isolated on August 1st, 1774, by an English chem- 
 ist, Joseph Priestley, who prepared it by heating " red precipitate " 
 (red oxide of mercury) in the focus of a burning glass exposed to 
 the sun's rays. He termed the gas dephlogisticated air because, 
 according to the theory then held, a substance, in burning, gave off 
 an hypothetical principle called phlogiston, which was subsequently 
 thought to be identical with hydrogen. As oxygen is capable of 
 supporting combustion in a very energetic manner, it was supposed 
 that the element must necessarily be able to take the phlogiston 
 from the burning substance; and from this property the gas 
 received its first name. 
 
 In order to prepare oxygen, the following methods can be 
 resorted to : 
 
 Red oxide of mercury is heated in a hard glass tube ; ! * it then 
 
 .* The superior numbers in the text refer to the numbers of the notes in 
 the laboratory appendix. 
 
OXYGEN; PREPARATION. 
 
 19 
 
 breaks down into mercury and oxygen. The formula of this oxide 
 is Hg 0, which means that in one formula weight of mercuric oxide 
 
 ]D p ID ]p ID 
 
 Fig. 1. 
 
 there are united one atom of mercury and one of oxygen. The re- 
 action is represented as follows : Hg = Hg -|- 0. The atomic 
 weight of mercury is 200, that of oxygen 16 ; hence the equation 
 indicates that 216 parts by weight of the oxide of mercury yield 
 200 parts of mercury and 16 parts of oxygen upon heating. 
 
 Black oxide of manganese (manganese dioxide) 2 is heated in an 
 iron tube, the oxygen so prepared being collected over water. 3 
 
 The reaction is as follows: 
 
 3 Mn 2 = Mn 3 4 -f- 2 ; the substance formed is an oxide of 
 manganese containing less oxygen, in proportion to the quantity 
 of manganese, than does Mn 2 . As one formula weight of Mn 3 4 
 contains three atoms of manganese, it follows that, in writing the 
 equation, three times one formula weight of Mn 2 must be used, 
 and this is indicated by placing the coefficient 3 before the formula 
 Mn 2 . As the atomic weight of manganese is 55, it follows that 
 261 parts by weight of manganese dioxide yield 229 parts of the 
 oxide containing less oxygen and 32 parts of oxygen. From the 
 examples cited, the meaning of chemical equations will be suf- 
 ficiently understood, so that the student, by consulting the table of 
 atomic weights, will be able to determine the relations by weight 
 in all others. We shall therefore, in the future, be contented with 
 
20 
 
 OXYGEN; .PREPARATION. 
 
 writing equations without indicating the quantities of the sub- 
 stances reacting.* 
 
 Xeither of the methods which have been given is of 
 great practical use in the preparation of oxygen. Better 
 results are obtained by heating chlorate of potassium (a 
 white, crystalline substance, somewhat resembling com- 
 mon salt 4 ) in a flask connected with a 
 delivery tube, as in Fig. 2. Chlorate 
 of potassium breaks down as follows : 
 
 
 
 KC10 3 = KCl-f 30.f 
 
 (Chlorate of potassium yields potas- 
 sium chloride and oxy- 
 gen.) The most approved 
 method of preparing oxy- 
 gen for laboratory use, 
 is by heating a mixture 
 of chlorate of potassium 
 and manganese dioxide, 
 the apparatus used being F <9- 2 - 
 
 the same as for that of the previous experiment. 5 By this 
 means the oxygen passes off at a much lower temperature, and quite 
 regularly. The part which the manganese dioxide plays in the re- 
 action is not as yet definitely understood, for it remains as manga- 
 nese dioxide in the final result. It is not improbable that the 
 substance takes up oxygen from the chlorate of potassium, thus 
 forming an oxide of manganese containing relatively more oxygen 
 than does Mn0 2 (a so-called higher oxide), which then readily 
 breaks down into manganese dioxide and oxygen, the black oxide 
 of manganese thus acting as a conveyer of oxygen. Substances 
 
 * It is just as well for beginning students to neglect equations, and to fix 
 their attention more especially on learning the facts, and also upon familiariz- 
 ing themselves with the appearance and character of the substances used in 
 the reactions. The fact that these reactions take place only between definite 
 masses of matter must, however, never be lost sight of. 
 
 t The reaction, as represented, gives the final result obtained by heating 
 chlorate -of potassium. There is in reality an intermediary product formed. 
 For a full description of the change see Chapter xviii. The pupil should, by 
 means of the table of atomic weights, determine the relationships by weight 
 in this and in a large number of subsequent equations. 
 
OXYGEN; PROPERTIES, OXIDES. 21 
 
 which act in this manner, their influence not being clearly under- 
 stood, are said to act by catalysis. All of the methods for prepar- 
 ing oxygen which have been cited have one thing in common ; the 
 substances decomposed contain oxygen, which is driven off by 
 heat, there remaining either elements or compounds containing pro- 
 portionally less oxygen than the original material. As energy is 
 added to effect the separation, the substances produced contain a 
 greater amount of chemical energy than those employed; while all 
 of these reactions belong to the class termed chemical analyses.* 
 
 Oxygen is colorless, odorless, and tasteless ; it is converted to a 
 liquid by a pressure of 50 atmospheres at a temperature of 118. 
 A certain temperature exists, constant for any given gas, but vary- 
 ing with different gases, above which they cannot be liquefied by 
 any pressure. This temperature is called the critical temperature ; 
 for oxygen it is at 118. Liquid oxygen boils at 182 at 740 
 mm. pressure, and at 197 at 25 mm.f Oxygen is but slightly 
 soluble in water, 100 volumes of that liquid dissolving 4.1 volumes 
 of oxygen at 5. $ Small as this amount is, it is sufficient to fur- 
 nish the material required by fishes for their physiological functions, 
 and this solubility is therefore of the highest importance; water 
 containing no free oxygen is incapable of supporting marine life. 
 
 Oxygen forms chemical compounds (called oxides) with all other 
 elements excepting fluorine and bromine, and it will combine with 
 the latter element provided some metal is also a constituent of the 
 compound. In the formation of many of the oxides, a large amount 
 of heat is produced ; this may be so great that the substance, 
 whether a burning solid or a vapor, will become incandescent ; this 
 phenomenon takes place with greater intensity in oxygen than in 
 the air. That the latter statement is true may be proved by placing 
 a glowing pine chip in a jar of oxygen, for then the wood will 
 at once burst into flame. 
 
 The phenomena attendant upon the burning of bodies in oxygen 
 
 * Reactions in which a compound body is decomposed into two or more 
 simpler ones are termed chemical analyses; those in which a more complex 
 body is formed from two or more simpler ones are termed chemical 
 syntheses. 
 
 t Dewar and Fleming, Philosoph. Mag. ; 34, 326. 
 
 t By this expression is meant that 100 liters of water would dissolve 4.1 
 liters of the gas. 
 
22 OXYGEN; COMBUSTION. 
 
 are all essentially alike, so that a very few examples will serve to 
 illustrate their character : 
 
 If a piece of phosphorus is kindled and then placed in a jar of the gas, it 
 will continue to burn with a dazzling white light, 6 forming a dense smoke. This 
 smoke consists of fine particles of the solid oxide of phosphorus, produced by 
 the combustion;") this oxide contains eighty parts by weight of oxygen for every 
 sixty-two parts by weight of phosphorus ; its formula is therefore P 2 O 5 (see 
 table of atomic weights). The action can be represented as follows: 
 
 2 P+5 O=P 2 O 5 . 
 
 Sulphur, which has been ignited in the air, will continue to burn in oxygen 
 with a brilliant blue flame ; 6 in this case the product of the combination is a 
 gas, composed of equal parts by weight of sulphur and of oxygen ; in order to 
 have the foregoing composition, this oxide of sulphur must contain two atoms 
 of oxygen to one of sulphur; it is termed sulphur dioxide. 
 
 S+20=SO,. 
 
 A piece of glowing carbon becomes brightly incandescent when placed in 
 oxygen, the substance burning to form a gaseous compound, carbon dioxide, 
 (C0 2 ):- ^,0^ 
 
 C + 20 = C0 2 . 
 
 The three oxides produced by the above reactions are oxides of 
 not-metals, and hence bear the characteristics of the anhydrides 
 of acids. (See page 14.) 
 
 Even substances such as iron, which we do not ordinarily consider as 
 being combustible, burn readily in oxygen. If a steel watch spring is heated 7 
 and placed in a jar of the gas, the substance will continue to burn. The heat 
 produced is sufficient to cause the iron to burn, while sparks of white hot 
 iron oxide spatter around the jar in which combustion is taking place; 
 these particles may even be hot enough to become fused into the sides of 
 the jar. 
 
 3 Fe+4 O = Fe 3 O 4 * 
 
 From these few examples we can form some idea of the readi- 
 ness with which oxides are formed ; and what is true of the elements 
 referred to also holds good with a large number of other ones. 
 When a burning substance is a gas, or when, in burning, it either 
 gives off a gas or is converted into a gas, then a flame is produced 
 (such is the case in the burning of phosphorus and of sulphur) ; on 
 the other hand, when no vapor is present to be heated to incandes- 
 
 * The oxide Fe 3 O 4 , occurring as a mineral called magnetite, is the oxide 
 of iron always formed at high temperatures when the metal is in the presence 
 of oxygen. 
 
OXYGEN; KINDLING TEMPERATURE. 23 
 
 cence, then the burning substance simply glows (this is seen in the 
 combustion of carbon).^ In a restricted sense, the term combustion 
 refers only to the unidh of various substances with oxygen, with 
 the evolution of light and heat. During any chemical process in 
 which two elements directly unite, heat is liberated; but, if the 
 union of a body with oxygen takes place slowly, so that in any 
 given interval of time only very small amounts of the substances 
 under consideration were to combine, we might not be able to note 
 any increase of temperature, because the inconsiderable quantity of 
 heat given off during this period would be conducted away by the 
 surroundings. The rusting of iron consists of such an oxidation, 
 without any perceptible rise in temperature ; it is, nevertheless, just 
 as much a form of combustion as the more rapid and brilliant one 
 referred to above. Combustion may, therefore, take place either 
 slowly or rapidly. A substance must be heated to a certain point 
 before slow can be converted into rapid combustion, and this tem- 
 perature is called the kindling temperature. The kindling temper- 
 ature varies for different substances, but is constant for any given 
 substance, and if the burning body is cooled below that point, an 
 extinction of the fire results. AVe have daily evidence of this 
 when lamp flames are blown out, for this process consists simply in 
 cooling, below their kindling; temperature, the gases which form the 
 flame ; while, on the other hand, the application of a lighted match 
 to a gas burner serves only to heat the escaping gases to their kind- 
 ling temperature. If a substance is undergoing slow combustion, in 
 such a situation that the heat given off cannot be readily conducted 
 away, the temperature of the whole will gradually rise until the 
 kindling point is reached ; and in this manner spontaneous combus- 
 tion takes place in heaps of oiled rags, where the oil is being slowly 
 oxidized by the oxygen in the atmosphere. Substances which are 
 capable of oxidation possess a certain amount of chemical energy 
 in the presence of oxygen, and, when they unite with that element, 
 this energy is converted into heat ; but, as the amount of energy in 
 any given body must be constant before combustion, no matter 
 whether the process is to be slow or rapid, therefore, provided the 
 products of combustion do not vary, it follows that the amount of 
 heat liberated must be the same in either case. It sometimes 
 occurs, however, that the substances formed by slow combustion 
 may be chemically different from those by rapid, when, of course, 
 
24 OXYGEN; OXIDES. 
 
 the amount of heat liberated woiild depend upon the product 
 formed.* 
 
 It would be erroneous to suppose that the phenomena of combus- 
 tion are apparent only when oxygen is taking part in the perform- 
 ance; for, as the process of direct chemical union between two bodies 
 is always accompanied by heat, it follows that in the union of other 
 elements or compounds this may be so great that the substances re- 
 acting begin to glow or burst into flame. Thus, for instance, hydro- 
 gen can burn in chlorine just as readily as it can in oxygen; so 
 that, in the very broadest sense, combustion would refer to the 
 union of any substances with the evolution of heat. This very 
 broadening of the sense, however, would rob it of its significance, so 
 that the term combustion might, perhaps, more properly refer only 
 to the union of any substance with oxygen, and with the evolution 
 of light and heat; while, instead of slow combustion, the term slow 
 oxidation, or, in the case of other elements, some more particular 
 term, would be preferable. 
 
 Two elements which are capable of directly uniting must possess 
 a certain amount of chemical energy, just as a stone raised above 
 the ground possesses potential energy ; and the chemical energy is 
 converted into heat when they combine, just as the potential energy 
 of the stone is changed to kinetic on falling. Now, after the stone 
 has fallen, it requires just as much energy to bring it once more 
 back to its original position as was given off by it in its descent. 
 Just so with chemical compounds. In order to decompose them, as 
 much energy must be added as was given off in their formation. 
 Therefore, in decomposing a body formed by combustion or slow 
 oxidation, just as much heat, or other form of energy, must be 
 applied as was given off in the production of the compound. The 
 heat (or the quantity of any form of energy) necessary for the 
 decomposition of a substance, is therefore equal to the heat (or 
 energy) given off in its formation ; it follows that those bodies which 
 burn energetically in oxygen, with the evolution of much heat and 
 light, will necessarily form stable oxides. 
 
 The bodies produced by the union of the various elements with 
 oxygen are termed oxides, and we are acquainted with the oxides 
 of all elements excepting those of fluorine and bromine. The great- 
 
 * For example, phosphorus, when slowly oxidized, forms phosphorus 
 trioxide; when hurned it forms phosphorus pentoxide. 
 
OXYGEN; FORMATION OF OXY-SALTS. 25 
 
 est diversity of characteristics exists among these compounds, but, 
 as we have seen in the introduction, they can in most instances 
 be classed either as bases or as anhydrides of acids; the bases 
 are oxides of metals which exhibit the same contrast toward the 
 oxides of not-metals (called anhydrides) as the metal itself does 
 toward the not-metal, with this difference, however, that the con- 
 trast is intensified by the addition of oxygen to the not-metal in 
 the formation of the anhydride. The result of the chemical union 
 of base and anhydride is a salt (usually possessing somewhat indif- 
 ferent chemical properties), and the product of the union of metal 
 with not-metal not infrequently has the characteristics of a salt', 
 such an instance is found in common salt (sodium chloride, Na Cl). 
 Examples of the oxides of metals which are bases are : 
 
 Zinc oxide, Zn ; sodium oxide, Na 2 ; magnesium oxide, 
 MgO. 
 
 Examples of the anhydrides of acids are : 
 
 Sulphur trioxide, S0 3 ; phosphorus pentoxide, P 2 5 ; carbon 
 dioxide, C0 2 . 
 
 The production of salts from the union of these two contrasting 
 substances can be represented as follows : 
 
 ZnO + S0 3 =ZnS0 4 , 
 
 Zinc oxide + Sulphuric anhydride = Zinc sulphate. 
 
 3 Na 2 + Pr0 5 = 2 Na 3 P0 4 , 
 
 Sodium oxide + Phosphoric anhydride = Sodium phosphate. 
 
 MgO +C0 2 =MgC0 3 , 
 
 Magnesium oxide + Carbonic anhydride = Magnesium carbonate ; 
 
 and if we compare these reactions with the formation of common 
 salt from sodium and chlorine : 
 
 Na + Cl = Na Cl, 
 
 Sodium + Chlorine = Sodium Chloride, 
 
 the similarity in the formation of salts from the two contrasting 
 compounds and from two contrasting elements is quite evident. 
 
 There are, in addition to the oxides referred to above, a large 
 number which are neither basic nor acidic. To this class belong 
 the oxides known as hyperoxides, of which manganese dioxide is an 
 example. All of these have the metallic character of the metal 
 about neutralized by the not-metallic oxygen, so that where they 
 are capable of taking up more oxygen, upon such addition the 
 
26 OXIDES ; NOMENCLATIVE. 
 
 resulting compound bears the characteristics of the negative an- 
 hydride of an acid; and when they lose oxygen to form a lower 
 oxide, the resulting substance has more or less the characteristics of 
 a base. They are on the turning point from base to anhydride, 
 and all of them lose a part of their oxygen readily, as does manga- 
 nese dioxide. A number of the oxides of the not-metals, also, are 
 not anhydrides of acids, and hence do not directly unite with the 
 bases to form salts. All of these oxides contain but little oxygen, 
 as, for example, carbon monoxide, CO ; nitrous oxide, N 2 0. 
 
 Many of the elements are capable of forming a number of 
 oxides (nitrogen can even form five), and these vary in character 
 with the increase in the relative quantity of oxygen. A number of 
 metals can form two oxides, both of which are bases, and in such a 
 case the oxide containing the lesser amount of oxygen relatively to 
 the quantity of metal is designated by the suffix ous attached to the 
 name of the metal, the one containing the greater by ic ; thus, 
 copper forms two oxides, cuprous oxide, Cu 2 ; cupric oxide, Cu ; 
 iron has the same property, and the two oxides, Fe and Fe 2 3 
 are termed ferous and ferric oxides respectively ; mercury can 
 form mercurous and mercuric oxides (Hg 2 and Hg 0). The salts 
 derived by the union of these oxides with the anhydrides of acids 
 are designated in a similar manner ; for example : 
 
 Ferrous sulphate, Fe S 4 , is produced by the union of ferrous 
 oxide with the anhydride of sulphuric acid : 
 
 Fe + S0 3 = Fe S0 4 
 
 and ferric sulphate, Fe 2 (S 4 ) 3 , is formed by uniting ferric oxide 
 with sulphur trioxide : 
 
 Fe 2 3 +3S0 3 =Fe 2 (S0 4 ) 3 . 
 
 The suffix ous indicates, when attached to the name of the metallic 
 element in a salt, that this salt is derived from an oxide containing 
 less oxygen, in proportion to a fixed amount of the metal, than does 
 some other basic oxide of the same element. 
 
HYDROGEN; OCCURRENCE, HISTORY. 27 
 
 CHAPTER III. 
 
 HYDROGEN. 
 
 Symbol, H ; atomic weight, 1.008 ; specific gravity, air = 1, is .06960 ; 
 1 c.c. H weighs .00009001 gram at and .76 meters pressure. 
 
 FREE hydrogen is found in nature only in small quantities ; it 
 occurs in the gases which escape from petroleum wells, in natural 
 gas, and in the exhalations of some volcanoes. It is said to have 
 been found in a condensed state in some meteorites, and it is fre- 
 quently given off in processes of fermentation and decay. The 
 element is, however, contained in enormous quantities in the chro- 
 mosphere of the sun, the protuberances observed during eclipses 
 consisting, for the greater part, of hydrogen which is subjected to 
 intense local disturbances. The presence of so much free hydrogen 
 in the atmosphere of the sun is accounted for by the fact that the 
 temperature of that body is so high that it renders all chemical 
 union impossible. The fixed stars, Sirius for example, also contain 
 large quantities of uncombined hydrogen ; on the earth the total 
 amount of free and combined hydrogen is only about 1 per cent of 
 the entire mass. 
 
 Hydrogen was first described as a peculiar form of air by 
 Cavendish in 1766, although it had previously been observed by 
 Paracelsus, Boyle, and others, who obtained it by the action of 
 dilute acids on certain metals, but confused it with other combus- 
 tible gases.* By Cavendish, Priestley, and contemporaneous chem- 
 ists, it was at one time considered to be pure phlogiston, the ele- 
 ment which was supposed to be given off by substances in burning. 
 Lavoisier explained the composition of water at' a later date, and 
 proved that substance to be an oxide of hydrogen ; by this means he 
 demonstrated that hydrogen is a substance which can act like a 
 metal when it forms an oxide, and that therefore it could not be 
 identical with phlogiston. Lavoisier gave the element the name of 
 
 * See Kopp, Geschiclite der Chemie, 3, 260. 
 
28 
 
 HYDROGEN ; PREPARATION. 
 
 hydrogene (from i>8<ap water and the root yei/ to produce), as water 
 is produced by burning the gas in air or oxygen. 
 
 The methods for preparing hydrogen are as follows : 
 
 1. By the decomposition of water by means of the electric 
 current, when hydrogen separates at the negative pole and oxygen 
 at the positive one. The apparatus used in performing this experi- 
 ment will be seen in Fig. 3.* f 
 
 2. By decomposition of water by means of metals. 
 The most pronouncedly metallic elements, such as so- 
 dium and potassium, have such a great affinity for 
 oxygen that they will unite with the 
 
 oxygen combined in water, and expel 
 the hydrogen even at ordinary temper- 
 atures. If a piece of sodium, the size 
 of a large pea, is placed in water, a 
 violent reaction 8 takes place, the heat 
 evolved melts the sodium, which, being 
 specifically lighter than water, floats 
 on the surface ; hydrogen is at the 
 same time set at liberty. If the 
 water is thickened by starch paste, 
 so that the molten metal cannot move 
 freely on its surface, the hydrogen 
 will be heated to its kindling tem- 
 perature and take fire. The gas can 
 be collected by placing the 
 sodium in a small wire cage 
 which is placed under the 
 surface of the water in a 
 pneumatic trough, and by 
 then inverting a test-tube filled with water over the metal, bubbles 
 of hydrogen arise and soon fill the tube. 9 (Fig. 4.) The water 
 
 * Pure water is a not-conductor of electricity, or at least very nearly a 
 not-conductor, and hence a little sulphuric acid must be added to the water 
 before performing the experiment. 
 
 t It will be noticed during the experiment that for each cubic centimeter 
 of oxygen liberated at the positive pole, we have 2 c.c. of hydrogen formed at 
 the negative. 
 
 Fig. 3. 
 
HYDHOGEN : PREPARATION. 
 
 29 
 
 Fig. 4. 
 
 in which the sodium has been dissolved has acquired a soapy feel 
 and an alkaline taste ; it contains sodium hydroxide. 
 
 Na + HOH = Na OH -f H. 
 
 Sodium + water = Sodium hydroxide + hydrogen. 
 
 A reaction of this kind is 
 called one of substitution, 
 for the sodium hydroxide can 
 be considered as water in 
 which one atom of the metal 
 hydrogen has been replaced 
 by one of the metal sodium, 
 and similar reactions may be 
 expected when metals are 
 brought in contact with a 
 compound of hydrogen and 
 any negative element or 
 group of elements'. 
 
 Such changes result with 
 other elements, the metallic properties of which are as pronounced 
 as those of sodium ; for instance, potassium will react with water 
 as follows : 10 
 
 K +HOH = KOH + H. 
 
 Potassium + water = Potassium hydroxide + hydrogen. 
 
 The group of elements OH, when it is united with some element or 
 group of elements, is called the hydroxyl group, and substances 
 which contain this are called hydroxides ;' thus, water, 'HOH, could 
 be called hydrogen hydroxide, and by replacing the hydrogen with 
 some other element, such" as sodium, potassium, calcium, or mag- 
 nesium, we have sodium, potassium, calcium, or magnesium hy- 
 droxides. 
 
 NaOH, Sodium hydroxide, 
 
 K OH, Potassium hydroxide, 
 Ca (OH ) 2 , Calcium hydroxide. 
 Mg(OH) 2 , Magnesium hydroxide. 
 Fe (OH ) 2 , Ferrous hydroxide. 
 Zn (OH ) 2 , Zinc hydroxide. 
 
 These hydroxides vary in formula, by reason of the fact that the in- 
 dividual atoms of different metals are capable of uniting with differ- 
 
30 HYDROXIDES. 
 
 ent numbers of hydroxyl groups. In the examples before us we 
 have two elements, sodium and potassium, the atoms of which unite 
 each with one hydroxyl group ; two elements, calcium and magne- 
 sium, where they unite with two ; and we can cite a third class 
 of elements, of which iron and aluminium are representatives, in 
 which one atom of the metal can unite with three of these groups. 
 
 Al (OH) 3 , Aluminium hydroxide, 
 Fe (OH) 8 , Ferric hydroxide. 
 
 The hydroxides of the metals, as well as the oxides, are bases, 
 and with acids they yield salts and water; they therefore present 
 the salne chemical contrast toward anhydrides of acids, or acids, as 
 do the oxides (see page 25), thus : 
 
 K 2 +S0 3 = K 2 S0 4 
 
 Potassium oxide + Sulphuric anhydride = Potassium sulphate. 
 
 2KOH +S0 3 =K 2 S0 4 +H 2 
 
 Potassium hydroxide -f Sulphuric anhydride = Potassium sulphate + Water. 
 
 K 2 O + H 2 S0 4 =K 2 S0 4 ' +H 2 
 
 Potassium oxide + Sulphuric acid = Potassium sulphate -}- "Water. 
 
 2KOH + H 2 S0 4 =K 2 S0 4 +2H 2 
 
 Potassium hydroxide + Sulphuric acid = Potassium sulphate + Water. 
 
 The acid employed in the last two reactions is formed from sul- 
 phuric anhydride and water, S0 3 + H 2 = H 2 S0 4 ; so that both 
 the oxide and hydroxide of potassium, with either the anhydride of 
 sulphuric acid or the acid itself, will yield a salt (potassium sul- 
 phate) and water (excepting in the case of the oxide of potassium 
 and sulphuric anhydride when no water is formed). What is true 
 of potassium is true of other metals as well. The oxides and 
 hydroxides of the metals are therefore bases, if, with acids, they 
 yield salts and water. The hydroxyl group is not confined in its 
 union to metals alone, the not-metals likewise are capable of form- 
 ing hydroxides, but the hydroxides of not-metals are acids ; the 
 discussion of the character of these hydroxides is, however, better 
 deferred until the pupil becomes acquainted with a greater number 
 of chemical compounds. 
 
 The more pronouncedly metallic in its nature a metal is, the 
 more readily will it be able to decompose water, liberating hydro- 
 gen. Sodium, potassium, or calcium does so at ordinary tempera- 
 tures ; magnesium does so, provided the water is boiling ; iron, if 
 
HYDROGEN ; PREPARATION. 31 
 
 the metal is heated, for when steam is passed over red hot iron, 
 hydrogen is produced, as follows : 
 
 3 Fe + 4 H 2 = Fe 3 4 + 8 H. 
 
 Not-metals (excepting red hot carbon) do not generate hydrogen 
 from water. 
 
 3. By the decomposition of acids by metals. 
 
 A better method of preparing hydrogen for laboratory use than 
 the one by the actions of metals on water is by that of the metals 
 011 acids, by which means hydrogen and a salt are formed, thus : 
 
 As action of sodium on hydrochloric acid is not practically 
 available, it is necessary to substitute some other metal, such as 
 zinc or iron, when the chloride of the metal and hydrogen are 
 formed : 
 
 Zn + 2HC1 =ZnCl 2 + 2H 
 
 Zinc + Hydrochloric acid = Zinc chloride + Hydrogen.f 
 
 Fe + 2HC1 =FeCl 2 + 2H 
 
 Iron + Hydrochloric acid = Ferrous chloride + Hydrogen. 
 Instead of hydrochloric acid, dilute sulphuric acid could be used ; 
 
 tlms: Zn +H 2 S0 4 =ZnS0 4 +2H 
 
 Zinc + Sulphuric acid = Zinc sulphate -{- Hydrogen. 
 
 Fe +H 2 S0 4 =FeS0 4 + 2H 
 
 Iron + Sulphuric acid = Ferrous sulphate + Hydrogen. 
 
 The action of either sodium or potassium on water does not 
 differ in principle from that of zinc or iron on hydrochloric or sulv 
 phuric acid, for in the case of either water or the acids, the me- 
 tallic element, hydrogen, is attached to a negative element, or group 
 of elements, thus : 
 
 H(OH), H 2 (S0 4 ), H(C1), 
 
 * This reaction can take place between sodium and gaseous H Cl. (See 
 note 32 appendix. ) The solution of hydrochloric acid, which is the ordinary 
 laboratory preparation, is not available for the purpose, for the action is too 
 violent. 
 
 t On comparing the formulae of zinc chloride and ferrous chloride with 
 those of zinc hydroxide and ferrous hydroxide, the pupil will see that one 
 atom of the metal combines with the same number of chlorine atoms as it 
 does with hydroxyl groups. 
 
32 
 
 HYDROGEN ; PREPARATION. 
 
 and this hydrogen, when either the hydroxide or a salt is formed, 
 is displaced by another metal which possessed greater chemical 
 energy when in contact with these negative groups. This is true 
 of sodium and hydroxyl, zinc and the group S0 4 , or zinc and 
 chlorine ; for in each case the heat of formation of the hydroxide, 
 sulphate or chloride, is greater than is that of the corresponding 
 hydrogen compound ; as a consequence, the metal, when brought in 
 contact with those compounds, will tend to displace the hydrogen 
 to form a system possessing less energy than did the original one ; * 
 thus in the reaction, 
 
 the Zn -f- H 2 S0 4 possesses more chemical energy than Zn S0 4 -+- 2 H, 
 
 as is seen by the fact that heat is pro- 
 duced when the reaction takes place. 
 
 In order to prepare hydrogen for labo- 
 ratory use, it is best to place granulated 
 zinc n in a flask, with delivery and safety 
 tube (Fig. 5), and to pour in 
 dilute sulphuric 12 acid through 
 the thistle tube. The gas gen- 
 erated from zinc is purer than 
 that which is prepared from 
 ordinary iron, so that the latter 
 is not used for laboratory 
 purposes. 13 ' u 
 
 Hydrogen is a color- 
 less gas ; when pure it is 
 odorless and tasteless ; at 
 probable temperature 
 of 213 and a pressure 
 of 40 atmospheres, it is 
 changed to a colorless 
 liquid. It is not very 
 
 Fig. 5. 
 
 soluble in water, one hundred volumes of that substance dissolv- 
 
 * In discussing chemical energy, if the reaction under consideration is ex- 
 pressed in the form of a chemical equation, the bodies to the left of the sign of 
 equality are termed " a system," and those to the right are given the same 
 name. Pattison Muir, Principles of Chemistry, p. 247. 
 
HYDROGEN ; PROPERTIES. 33 
 
 ing 1.82 volumes of hydrogen at 20.* Hydrogen has the smallest 
 specific gravity of any substance with which we are acquainted ; it 
 occupies 14 times the space taken by an equal weight of air, or 
 11,160 times that of water, its specific gravity being .06960. In 
 consequence of this specific lightness, jars can be filled with hydro- 
 gen by placing them mouth downward, and allowing the gas to 
 enter through a tube extending upward to the bottom, when tjie 
 hydrogen will expel the air ; the collecting of gases in this manner 
 is termed " upward displacement." Balloons or soap bubbles filled 
 with hydrogen will rise in the air, and if two beakers are suspended, 
 mouth downward, on a balance, and then exactly tared, one arm of 
 the balance can be caused to rise by filling the corresponding 
 beaker with hydrogen, poured upward from another vessel. 
 
 Hydrogen passes readily through porous substances. It has the 
 smallest specific gravity, and the greatest rate of diffusion of any 
 gas, because the rate of diffusion of gases is inversely as the square 
 roots of their densities, f the law being approximately true. By 
 diffusion we mean the power which a gas has of completely mixing 
 with another gas, or gases, which, without stirring, are placed in 
 contact with it. This can take place even though the gases are 
 separated by a partition, provided this latter is sufficiently porous 
 to allow of the passage of the molecules. Hydrogen can pass 
 through substances such as unglazed earthenware, paper, or even 
 metals like platinum, when the latter are heated. 15 Some metals, for 
 example, palladium, have the power of condensing large quantities 
 of hydrogen. Palladium, when in the shape of wire, can absorb 
 872 times its own volume, and, when used as a negative electrode 
 in a battery, it can condense very nearly 1,000 volumes ; so that 1 
 c.c. of palladium can absorb about a liter of the gas.$ The palla- 
 dium gains in volume during the operation, and becomes specifically 
 lighter. Hydrogen so absorbed is said to be occluded. Occluded 
 hydrogen is chemically much more active than is the ordinary gas, 
 and reacts with substances on which the element ordinarily has no 
 effect. 16 Heat is evolved in the process of occlusion ; and a portion 
 of the hydrogen taken up by palladium is probably chemically com- 
 
 * Cl. Winckler, Berichte d. Deutsch. Chem. Gesell. 24; 97. 
 t The terms " density " and "specific gravity," when referring to gases, 
 are used interchangeably. 
 
 J M. Thoma, Zeitschrift fur Physikalische Chemie ; 3, 83. 
 
34 HYDROGEN ; DIFFUSION, OCCLUSION. 
 
 bined ; the metal liberates a portion of its occluded hydrogen even 
 at ordinary temperature, while at 130 to 140 the greater part 
 passes off. So energetically does the hydrogen occluded by palla- 
 dium act that, if the metal is saturated with the gas, it will glow 
 spontaneously in the air, the hydrogen uniting with oxygen to 
 form water. The condensation of gases by metals can be easily 
 shown by suspending a coiled platinum wire in a burning, not- 
 luminous gas flame until it is red hot, turning out the flame and 
 then instantly turning it on again. The platinum occludes the 
 gases escaping from the burner, and, being warm, they unite with 
 the oxygen of the atmosphere, so that the wire continues to glow 
 and finally kindles the gas. Occluded gases are frequently used in 
 chemical operations. 
 
 Animals, when placed in an atmosphere of hydrogen, die of 
 asphyxia ; hydrogen is not poisonous, per se, as is proved by the 
 fact that animals can live without discomfort in an atmosphere of 
 hydrogen to which has been added a sufficient supply of oxygen. 
 Hydrogen is not able to support combustion in the ordinary sense 
 of the term, a burning candle being extinguished when placed in it. 
 In order to be combustible in a gas, a body must have a great ten- 
 dency to unite with that gas ; necessarily the substances which are 
 ordinarily considered as combustible, such as wood or a candle, can- 
 not burn in hydrogen, for they themselves are largely composed of 
 hydrogen. Hydrogen burns quite readily in oxygen, with a very 
 hot flame, and vice versa, oxygen will burn in hydrogen with the 
 production of an equally hot flame ; in both instances water is pro- 
 duced (see page 36). Hydrogen will also burn in other gases for 
 which it has a great affinity, in chlorine, for instance. The hydrogen 
 compounds of the elements in which the gas is burned will, in each 
 instance, be produced. 
 
 In order to prove the formation of water by the combustion of 
 hydrogen in the air, we have but to dry thoroughly the gas passing 
 from a generator ; 17 ignite it at the tip of a burner, 18 and then hold 
 a cold beaker over the same. Drops of water soon collect and run 
 down the sides of the vessel. The flame of hydrogen burning in 
 oxygen has a very high temperature, and is very nearly not-lumi- 
 nous ; but it can be changed to a luminous one by dropping a little 
 powdered coal or lycopodium into it, for not-luminous flames are 
 simply glowing gases : they ara made luminous by glowing solids 
 
HYDROGEN; COMBUSTION OF. 35 
 
 within them. Hydrogen, like other substances, burns more energet- 
 ically in oxygen than in air, so that the temperature of a flame of 
 hydrogen burning in the former is much higher than that of the 
 same gas burning in the latter. This fact is taken advantage of in 
 the use of the oxy-hydrogen blow-pipe. This apparatus consists of 
 two concentric tubes drawn nearly to a point ; through the outer one 
 hydrogen is admitted and ignited, and then a stream of oxygen gas 
 is passed through the central tube, thus mixing the two gases at their 
 point of combustion. The most extreme heat of a flame of hydrogen 
 burning in oxygen is thus obtained. The illuminating power of cal- 
 cium lights, is due- to a piece of lime (an infusible substance), heated 
 to incandescence by the oxyhydrogen blow pipe ; zinc, iron, or tin, 
 burns readily in the flame ; a piece of platinum wire is instantly 
 melted; indeed, the fusing of platinum for the manufacture of 
 various utensils is accomplished by this means. 
 
36 WATER; CHEMICAL HISTORY. 
 
 CHAPTER IV. 
 
 WATER. 
 
 Symbol, H 2 0; specific gravity 1, at 4 C ; specific gravity of vapor, 
 air = 1, is .6208 ; H 2 = 2, is 17.88. 
 
 WHEN hydrogen is burned in oxygen, or oxygen in hydrogen, 
 water is formed. The study of the composition of this substance, 
 and the methods by which this has been advanced, might serve as a 
 type of all other similar investigations ; so that for the purpose of 
 becoming somewhat acquainted with the means of chemical research, 
 the subject will be discussed at some length. That hydrogen and 
 oxygen, in uniting, form water is easily*proved ; but the question at 
 once arises, is water alone produced, or may not some other sub- 
 stance originate simultaneously, when the combustion takes place ? 
 Indeed, when all of the products of combustion of hydrogen in the 
 air are collected, an acid substance, nitric acid, is also found to be 
 present. The formation of an acid by this means puzzled the 
 original investigators ; and not until they had exploded a mixture 
 of pure hydrogen and pure oxygen, as Cavendish did in 1780, and 
 so discovered that no acid whatever was produced under those cir- 
 cumstances, did they come to the conclusion that hydrogen and 
 oxygen, in burning, formed water only. Another problem as 
 regards water then presented itself. Before the last quarter of 
 the eighteenth century, it was generally supposed that water, by 
 boiling, was changed to an earthy substance, for on boiling even 
 pure distilled water in glass or earthen vessels, an earthy residue 
 remained after evaporation. Lavoisier took upon himself to prove 
 that the universally accepted notion was erroneous. In order to do 
 this, he sealed some pure water in a glass flask, weighed it, then 
 kept it heated for some weeks and found that the total weight was 
 unchanged. On evaporating the water, a solid substance remained, 
 but then Lavoisier found that the flask had lost in weight, and this 
 loss in weight was exactly equal to that of the earthy residue. As 
 a consequence, he came to the conclusion that the earthy substance, 
 
WATER ; COMPOSITION OF. 37 
 
 supposed to be formed by the boiling of water, was nothing more 
 than a dissolved portion of the glass, deposited by evaporation. 
 Scheele, at the same time, showed that this earthy residue had 
 exactly the same composition as the glass in which the water had 
 been boiled. Chemists had then proved two things by their inves- 
 tigations ; that hydrogen and oxygen, in uniting, form nothing but 
 water, and that this water, in evaporating, was volatilized unchanged. 
 We have seen that water is decomposed by the electric current. 
 It is not only necessary for the chemist to prove that hydrogen and 
 oxygen are produced by this means ; he must also show that nothing 
 but these gases is formed. Sir Humphry Davy noticed that, in 
 electrolyzing water, a certain amount of alkali separated at the 
 negative pole, while acid was formed at the positive one. There 
 were only three possible theories to hold : one was that these 
 substances were produced from the water; the second, that they 
 were formed by the decomposition of the glass, for glass is composed 
 of an acid and an alkaline substance ; and the third, that the sur- 
 rounding air took part in the change. In order to discover the 
 source of these impurities, Davy transferred the water to a gold 
 vessel, and then continued the process of electrolysis. Acid and 
 alkali were formed as before, but, while the production of acid con- 
 tinued, the amount of alkali became less and less, until it finally dis- 
 appeared, showing that it was owing to the decomposition of the 
 glass. Davy came to the conclusion that this latter was due to 
 nitric acid, produced from the atmosphere, and, as a consequence, 
 the gold vessel was placed under a bell-jar on the receiver of an air 
 pump, the air exhausted and replaced by pure hydrogen, when the 
 formation of acid ceased.* It was proved, therefore, that water 
 changed to nothing but hydrogen and oxygen on electrolysis, and 
 with that the chain of evidence was complete. There remained 
 now to determine what proportionate volumes of hydrogen and 
 oxygen were produced in the electrolysis of water, and what pro- 
 portionate volumes united. to form water. This matter had been 111 
 dispute for a number of years, until Gay Lussac and Humboldt 
 showed that exactly twice as much hydrogen as oxygen, by volume, 
 is contained in water ; 19 in other words, for every two cubic centi- 
 meters of hydrogen there is one cubic centimeter of oxygen,! and 
 
 * See Kopp, Geschichte der Chemie; 1, 377. 
 
 t See E. W. Morley, Amer. Journal of Science, [3] 41; 220, 276. 
 
38 WATER; COMPOSITION BY WEIGHT. 
 
 if hydrogen and oxygen are mixed in exactly these proportions, 20 
 the gas produced explodes violently when ignited by means of 
 a taper or electric spark. If the exact mixture is exploded in a 
 short heavy glass tube closed with mercury, over a trough filled 
 with that metal, then, after the explosion, the mercury will com- 
 pletely fill the tube, as the volume of water formed is insignificant 
 as compared to the volume of the gases used. The following facts 
 have then been proved as regards water. It is composed of hydro- 
 gen and oxygen only; it decomposes into nothing but hydrogen 
 and oxygen ; the volume of hydrogen is to the volume of oxygen as 
 2:1; and furthermore, two volumes of hydrogen mixed with one of 
 oxygen form an explosive mixture, from which only water is 
 produced, without a residue of either hydrogen or oxygen. 
 
 Having proved the composition of water by volume, our next 
 task is to discover the same by weight, and this was first most 
 accurately determined by Dumas.* By passing pure hydrogen 
 over heated copper oxide, the following reaction takes place : 
 
 CuO +2H = Cu +H 2 
 
 Copper oxide + hydrogen = Copper + water. 
 
 By this means we have a most ready method of discovering the 
 proportional parts by weight in which hydrogen and oxygen unite. 
 For, let us suppose the copper oxide to be perfectly dry and 
 accurately weighed, then, after the reaction, the weight of the 
 remaining copper subtracted from that of the copper oxide would 
 give us the weight of oxygen which has gone to form water ; if, by 
 some means, we can collect the water formed, and weigh the same 
 with equal accuracy, the difference between the weight of the 
 oxygen and that of the total water would give the weight of hy- 
 drogen. 
 
 a = weight of copper oxide. c = weight of water formed, 
 
 b = weight of copper after the re- c x = weight of hydrogen. 
 
 action, 
 a b = weight of oxygen = x and x : c x will represent the ratio 
 
 in which oxygen and hydrogen unite to form water. 
 
 This ratio was found by Dumas to be 1 : 7.98. Dumas took 
 great care to purify the hydrogen which he used. 21 The water was 
 
 * Berzelius and Dulong first undertook to determine the quantitative com- 
 position of water by this method. 
 
WATER ; FORMULA OF. 39 
 
 collected by passing the gas through an empty glass flask, in which 
 the greater part was condensed, and then through a series of tubes 
 filled with fused caustic potash and phosphorus pentoxide, which 
 substances absorbed all of the remainder. The tubes were filled 
 with perfectly dry and pure air before the operation, and then 
 weighed ; after the water had been formed they were once more 
 filled with dry and pure air and again weighed, when the gain 
 in weight gave the amount of water formed. This latter, in nine- 
 teen experiments, was 945.439 grams; the loss in weight of the 
 copper oxide, and consequently the amount of oxygen in the water, 
 was 840.161 grams, leaving for hydrogen (945.439 840.161) or 
 105.278 grams. The ratio between the combining weights of hydro- 
 gen and oxygen, therefore, is as 105.278 : 840.161, or as 1 : 7.98. Of 
 late Morley, by using a different method, has proved that the ratio 
 is more probably 1 : 7.94.* 
 
 The facts which we have discovered as regards water are then 
 as follows : 
 
 It can be decomposed into hydrogen and oxygen, and yields two 
 volumes of hydrogen to one of oxygen ; it is formed from two vol- 
 umes of hydrogen and one of oxygen, and when the water so formed 
 is heated above the boiling point (i.e., measured as steam), it will be 
 seen that 2 volumes of hydrogen -f 1 volume oxygen produce 2 vol- 
 umes water vapor ; and, furthermore, water contains the two ele- 
 ments in the proportional parts by weight of one of hydrogen to 
 eight of oxygen. AVe shall subsequently see that in equal volumes 
 of gases, under the same conditions of temperature and pressure, 
 there are equal numbers of particles ; so that if two volumes of 
 hydrogen unite with one of oxygen to form water, we are tolerably 
 safe in assuming that two atoms of hydrogen unite with one of 
 oxygen, and if we neglect the decimals, two parts by weight of 
 hydrogen unite with sixteen of oxygen ; from this it follows that in 
 all probability the atomic weight of oxygen is sixteen. This last 
 conclusion cannot be absolutely certain, as we shall subsequently 
 see ; but we are sure that one part by weight of hydrogen unites 
 with eight of oxygen to form water, and formerly chemists were of 
 the opinion that speculation should not extend beyond this certainty. 
 They wrote the formula of water HO, meaning by this that hydro- 
 
 * Very nearly the same ratio (7.935 to 7.945) has been found by Kayleigh, 
 Keiser, Noyes, Leduc, and others. 
 
40 WATER; PROPERTIES. 
 
 gen and oxygen unite to form water, in the proportion of one part 
 by weight of hydrogen to eight of oxygen ; and as eight parts by 
 weight of oxygen are equivalent in combining power with one of 
 hydrogen, eight was called the equivalent weight of oxygen. The 
 custom in respect to oxygen was followed in regard to other elements 
 as well. In despair of finding some guide by which to determine 
 the true atomic weights, equivalent weights were resorted to, as the 
 latter involved no speculation, but were simply founded upon 
 the known experimental facts. As chemical and physical theory 
 and investigation became more perfect, it was, however, seen that 
 we could form very definite conclusions as regards the atomic 
 weights, so that equivalent weights were abandoned, and our 
 present atomic weights came into general use. The investigations 
 on the composition of water will serve to illustrate the means 
 by which our present chemical facts have become known, and the 
 pupil should remember that when in the future the atomic weights 
 of elements or the formulae of compounds are mentioned, these have 
 been discovered by some equally painstaking means of investigation. 
 Water, at ordinary temperatures, is a nearly colorless liquid ; 
 when light is passed through a thick layer of the substance, it 
 will be seen to have a distinctly blue color. When heated to its 
 boiling point, which, with 760 m - m - pressure, is at 100 centigrade, 
 it is transformed into a colorless gas.. When water is cooled, the 
 substance contracts, following the general law ; when heated, it 
 expands. If, however, water having a temperature above 4 centi- 
 grade is cooled, it will contract until that temperature is reached, 
 and will then begin to expand until centigrade, at which point it 
 freezes. Water at 4 centigrade, then, has the greatest specific grav- 
 ity,* so that any body of water, on cooling toward that temperature, 
 will become specifically heavier ; that portion on the surface, because 
 it cools first r will sink ; the warmer and lighter water below will 
 rise, and in this manner a continuous circulation will be kept up 
 until the entire body has arrived at the temperature of 4 centi- 
 grade. Now, the water on cooling further will expand, and hence 
 the cooler portion will float on the surface until it arrives at 0, 
 when freezing begins, and the crust of ice formed, not being able 
 to sink, will act as a protection to the water below. As a con- 
 
 * If we place the density of water at 4 centigrade at 1., then water at 
 centigrade has a specific gravity of .99988. 
 
WATER ; SOLUTION. 41 
 
 sequence, water freezes on the surface and not from below upward. 
 In freezing, water expands, the specific gravity of ice being .9167 ; 
 on cooling below its freezing point, ice continuously contracts, fol- 
 lowing the usual law. The freezing point of water, or, rather, the 
 melting point of ice, at standard atmospheric pressure, is taken as 
 on the centigrade thermometer ; this point is lowered by com- 
 pression, so that ice at centigrade can be fused by increasing the 
 pressure ; as soon as this is relieved the liquid instantly freezes. This 
 fact can be experimentally proved by firmly pressing together two 
 pieces of ice, which will adhere as soon as the pressure is relieved. 
 
 Water is able to dissolve a large number of substances to form 
 solutions. Solutions are homogeneous mixtures of two or more 
 substances, one of which must be either gaseous or liquid. These 
 mixtures cannot be separated by simple mechanical means. Gases 
 can form these homogeneous mixtures in any proportion. When two 
 substances are liquid, the solution may take place in any proportion, 
 as between alcohol and water, alcohol and ether, acetic acid and 
 water ; or one liquid may partially dissolve another, as in the case 
 of water and ether ; * or, lastly, one liquid may be entirely insoluble 
 in another, as in the case of some oils and water. When a solid 
 dissolves in a liquid, the latter is only able to take up a certain 
 quantity, which, with any given liquid, varies with the nature of 
 the solid. When the liquid has dissolved as much of the solid as 
 it will, the solution is said to be saturated. As a rule, gases are 
 less soluble in liquids the higher the temperature ; they are there- 
 fore expelled from their solutions on heating. Solids, on the other 
 hand, are, as a rule, more soluble in hot liquids than in cold ones. 
 As a consequence, a saturated solution, at the boiling point of the 
 liquid, will contain more of the dissolved solid than it will at 
 a lower temperature, so that, on cooling, the dissolved substance 
 separates, frequently in a crystalline form. The process of dissolv- 
 ing crystalline solids in hot liquids and separating by cooling, is 
 called recrystallization, and is very frequently employed as a 
 means of purifying crystalline substances. The solubility of solids 
 
 * In this latter case we can scarcely say that one liquid dissolves the other, 
 for to take an example from the instance just cited, the ether will dissolve just 
 as much water as the water will dissolve ether. For a more extended element- 
 ary discussion of solutions see Ostwald's Outlines of General Chemistry, Walker's 
 translation; Macmillan, 1890. 
 
42 WATER OF CRYSTALLIZATION. 
 
 in water varies very greatly ; some substances are insoluble, others 
 but very slightly so, so that the solution may be saturated when 
 only a trace of dissolved substance is present ; again, we have solids 
 such as potassium hydroxide, calcium chloride, or magnesium chlo- 
 ride, which are soluble even in their own volume of water. The 
 solutions formed are clear liquids, which may have the color of the 
 dissolved solid. In making solutions we frequently have marked 
 changes of temperature. Where substances form no chemical 
 union with water in dissolving, the temperature is lowered, for 
 work must be done to change the crystallized solid to a liquid ; 
 thus potassium or ammonium nitrate, on dissolving in water, effect 
 a very marked cooling. 
 
 Quite a number of substances can chemically take up water to 
 form crystalline compounds of a definite form ; this is done in such 
 a way that a given quantity of the solid unites with a definite quan- 
 tity of water. The water is apparently not present as such, for the 
 crystals are perfectly dry, the water appearing only when the solid 
 is heated, for it then passes off. Such combined water is known as 
 water of crystallization ; the number of molecules of water of crys- 
 tallization united with one formula weight of the solid, as well as 
 the crystalline form of the substance, are always the same for any 
 given compound ; in some few cases, however, the same individual 
 can crystallize in different forms, with different amounts of water 
 of crystallization. To cite a few examples, one formula weight of 
 copper sulphate crystallizes with five molecules of water, as 
 Ou S0 4 , 5 H 2 ; ferrous sulphate with seven molecules as Fe S0 4 , 
 7 H 2 ; sodium carbonate with ten molecules, as Na 2 C0 3 , 10 H 2 ; 
 and alum with twenty-four molecules, as A1 2 K 2 (S0 4 ) 4 -f- 24 H 2 0. 
 Substances which are chemically similarly constituted will fre- 
 quently crystallize in the same, or at least in very similar, 
 crystalline forms ; such substances are said to be isomorphous.* A 
 
 * The student should consult some elementary text-book in regard to the 
 principles of crystallography. Isomorphous salts are, for instance, 
 
 MgSO 4 ,7H 2 O. CaCO 3 . 
 
 Zn SO 4 , 7 H 2 O. Sr CO 3 . 
 
 Two or more substances are not truly isomorphous, unless one can replace 
 the other without altering the crystalline form, thus, MgSO 4 , 7H 2 O and 
 Zn SO 4 , 7 H 2 O, both crystallize in the rhombic system, and zinc can replace 
 magnesium in crystals of Mg SO 4 , 7 H 2 O, without altering the crystalline form. 
 
WATEE ; CHEMICAL UNION OF. 43 
 
 number of compounds lose their water of crystallization on stand- 
 ing in the air, their crystalline form is destroyed, and they finally 
 disintegrate to form not-crystalline powders; such substances are 
 efflorescent. On the other hand, we have a large number of bodies 
 which are capable of taking up water from the atmosphere, and 
 dissolving in the moisture so concentrated ; these are deliquescent. 
 By far the greater number of substances are neither deliquescent 
 nor efflorescent at ordinary temperatures, but all substances with 
 water of crystallization, give the liquid up at a comparatively mod- 
 erate heat. When water of crystallization has been expelled from 
 a substance, the body is said to be anhydrous. Anhydrous salts 
 dissolve in water with the evolution of heat ; many, like calcium 
 chloride, are capable of absorbing moisture readily, and as a conse- 
 quence are most useful for drying gases. It must be borne in 
 mind, however, that the majority of substances which crystallize, 
 do not contain water of crystallization. 
 
 Water combines with many substances when it does not enter 
 as water of crystallization ; for example, the oxides of the metals, 
 when soluble in water, unite with that substance to form the 
 hydroxides : 
 
 K 2 -fH 2 = 2KOH 
 
 Potassium Oxide + Water = Potassium hydroxide. 
 
 Ca + H 2 = Ca (OH) 2 
 
 Calcium Oxide + Water = Calcium hydroxide.* 
 
 In such cases the water entering into combination with the 
 oxide is decomposed, it is no longer present as water but as 
 hydroxyl, to which group of atoms attention has already been 
 called on page 30. The hydroxides, of course, differ most mark- 
 
 * Note again the difference between the formulas K OH and Ca (OH) 2 . 
 Here, we see one atom of calcium has the power of retaining twice as many 
 hydroxyl groups as one of potassium does ; a relationship exactly similar to 
 that existing between potassium and magnesium hydroxides, and similar 
 to the relationship in the formula of the chlorides, Ca C1 2 , K Cl. If we were to 
 call chlorine X, hydroxyl Y, the relationship would become apparent in 
 Ca X 2 , Ca Y 2 ; KX, KY. These chemical changes are rendered more ap- 
 parent if the pupil will take the trouble to write out the formulae of reactions, 
 atom for atom, as has been done above. 
 
44 WATER; PURIFICATION OF. 
 
 edly in stability ; some are very readily decomposed by heat into 
 water and the oxide, while others, for example, potassium hy- 
 droxide, can be raised to a high red heat without changing to 
 the oxide. As a rule, the less pronouncedly metallic the element 
 forming the hydroxide is, the more readily will that hydroxide be 
 decomposed by heat. For example, calcium hydroxide is decomposed 
 as follows : 
 
 /OHJJaO/H Ca(QH)2 =Ca0 +H2Q 
 
 / / TT Calcium hydroxide = Calcium oxide -f- "Water. 
 
 \OH /OH 
 
 Water can unite with the anhydrides of acids to form hy- 
 droxides (called acids), as readily as it can enter into combination 
 with the oxides of metals. Attention will subsequently be called 
 to this class of reactions. 
 
 Pure water is very difficult to obtain, owing to the great capa- 
 city which the liquid has for dissolving various substances. The 
 impurities may be of two kinds, those mechanically suspended, and 
 those dissolved ; while the dissolved impurities may be classed under 
 three heads, gaseous, liquid, and solid. The mechanically suspended 
 impurities may be removed by nitration, that is, by passing the water 
 through some porous substance, such as unsized paper (so-called 
 filter paper), or through unglazed porcelain. Those dissolved must 
 be removed by distillation.* The gaseous impurities cannot all be 
 eliminated by this means, for the oxygen and nitrogen of the atmo- 
 sphere, as well as the other impurities in the air, are soluble in 
 water, and will therefore be found in the distilled water. Liquid 
 impurities may be removed by repeated distillation, provided their 
 boiling point is not too near that of water ; where the boiling points 
 of two liquids are within 10 of each other, a complete separation 
 by distillation is impossible. The solid impurities remain behind 
 in the vessel from which water has been distilled. If distilled 
 water is to be free from gases, it must be boiled for some time in 
 bottles, and then hermetically sealed before allowing to cool. 
 
 All naturally occurring waters are more or less impure, the 
 purest being rain-water and melted snow, but even these contain 
 
 * Distillation is simply the process by which the steam from boiling water 
 is collected in suitable vessels, and condensed to a liquid. For a description 
 of the process of distillation, any larger text-book can be consulted. 
 
WATER; NATURAL WATERS. 45 
 
 such solid and gaseous substances as they can collect in passing 
 through the atmosphere. Rain-water, in falling on the soil, takes 
 up such soluble substances as are contained therein, and in so doing 
 is changed to spring-water. Of course, the dissolved impurities of 
 spring-water vary greatly with the nature of the soil through which 
 the water has passed. When the amount of dissolved impurity is 
 not very great and is mainly calcium carbonate or sulphate, the 
 water is fresh water ; when quantities of salts are dissolved from 
 the soil, the water becomes a mineral water. Some springs are 
 heated as they issue from the earth ; these are called thermal 
 springs. The most frequent salts in mineral waters are sodium 
 and magnesium sulphates, sodium carbonate and carbonate of iron, 
 while gaseous constituents like carbon dioxide and sulphuretted 
 hydrogen are also met with. Eiver water necessarily must con- 
 tain the dissolved impurities which were in the springs from which 
 the stream has its source, with some modifications introduced by the 
 nature of the soil over which it has passed ; it is often rendered 
 unfit for drinking purposes by contaminations which have been in- 
 troduced, purposely or accidentally, through decaying animal or 
 vegetable substances. As streams pass through more or less thickly 
 inhabited regions, the proximity of dwellings and factories is apt to 
 cause dangerous impurities in the water, for as a rule very little care 
 is exerted to keep sewage from reaching the same stream from which 
 drinking water is taken. The water polluted will, on analysis, prove 
 to contain compounds of nitrogen which have their origin in putres- 
 cent animal substances ; the quantity of these compounds diminishes 
 as the water is exposed to the oxidizing action of the atmosphere ; 
 yet, in the water thus purified by nature, there may be a large number 
 of living micro-organisms which may cause disease, so that a drinking 
 water should be examined not only with the purpose of ascertaining 
 the quantity of decaying animal substance present, but also as to the 
 number and kind of micro-organisms contained in it. Water may 
 be contaminated by a considerable quantity of sewage and yet be 
 harmless, for it may contain no harmful disease germs ; while, on 
 the other hand, water considered as pure may be extremely danger- 
 ous, by reason of germs contained therein. 
 
 In concluding the chapter on this most important subject, it 
 is well once more to call attention to the necessity of thoroughly 
 understanding the changes of energy which take place in the forma- 
 
46 WATER; REVIEW or CHEMISTRY OF. 
 
 tion and decomposition of water. We must remember that hydro- 
 gen and oxygen (a metal and a not-metal) unite most readily, 
 provided some impulse (such as an electric spark or fire) is added 
 to a mixture of the gases, while, during the union, much heat is 
 given off, so that water possesses much less energy than the ele- 
 ments hydrogen and oxygen; and furthermore the application of 
 just as much kinetic energy is necessary to decompose water as has 
 been given off in its formation. When the energy is applied by 
 means of an electric current, two volumes of hydrogen and one of 
 oxygen are produced. 
 
 We have, by this time, become somewhat more intimately 
 acquainted with chemical equations. The pupil should practise 
 writing the chemical equations which have been given, until he is 
 entirely familiar with them, but he must always remember that the 
 mere memorizing of such equations is useless; that the reasons 
 why any reaction should take place, and why certain substances are 
 formed from any reaction, are of infinitely more importance than 
 the equations which, however, fix these reasons in the memory. 
 
OZONE; HISTORY. 41 
 
 CHAPTER V. 
 
 OZONE AND HYDROGEN DIOXIDE. 
 
 WHEX compared with the ^extraordinary chemical activity of 
 oxygen at higher temperatures, the tendency of that element to 
 unite with other substances under ordinary conditions is not very 
 marked. Iron, copper, and similar substances, which are oxidized 
 when heated in the gas, remain unchanged in perfectly pure oxygen 
 at ordinary temperatures, and we can readily see that, did oxygen 
 possess the capability of oxidizing under these circumstances, there 
 would result a complete alteration of the existing conditions upon 
 the surface of the earth. The element is not, however, limited to 
 the one form which we have discussed ; it is also capable of existing 
 in another character, in which its properties differ most remarkably 
 from those of ordinary oxygen. 
 
 An element capable of existing in two or more different physical 
 and chemical forms is said to possess the property of allotropism ; 
 and the different modifications of the same element are called its 
 allotropic forms. Oxygen, then, exists in two allotropic forms, 
 oxygen and ozone ; in the former one it has the properties which 
 have already been described ; in the latter it possesses a marked 
 odor; when large quantities are present, it is irritating to the 
 mucous membrane of the throat and nose, and it is a most active 
 oxidizer even at ordinary temperatures. We are acquainted with 
 ozone only when it is diluted with oxygen or air. 
 
 The fact that a room, in which a powerful generator of static 
 electricity is in action, becomes filled with an odor resembling that 
 of phosphorus has been known for some time,* and, in 1840, Schon- 
 bein found that the same odor is produced when moist phosphorus 
 is exposed to the atmosphere. The substance causing this odor 
 was, at first, owing to its resemblance to chlorine, supposed to be 
 an element ; at a later date investigation seemed to show that it 
 was simply an oxide of hydrogen differing from water, but, finally > 
 
 * Since 1785 (Von Marum). 
 
48 OZONE; RELATION TO OXYGEN BY VOLUME. 
 
 de la Rive proved that if oxygen, perfectly pure and dry, is passed 
 through a narrow glass tube in which are inserted two platinum 
 wires, between which electric sparks are passing, a quantity of 
 ozone is generated. This proved without a doubt that ozone was 
 generated from oxygen alone, and subsequently the proof was 
 brought that ozone, on heating, yields no.thing but oxygen. 
 
 Having discovered that ozone is simply oxygen in another 
 form, there remained to be .decided whether, in forming the former 
 substance from the latter, any change in the bulk of the gas 
 occurred. Further study showed that a diminution in volume 
 takes place; and this contraction was such that from 3 c.c. of 
 oxygen there result 2 c.c. of ozone, and conversely from two of 
 ozone there are formed three of oxygen. We have learned that in 
 equal volumes of gases there are equal numbers of particles ; it fol- 
 lows that if we were able to obtain pure ozone, there would be as 
 many particles of ozone in a given volume as there would be in the 
 same bulk of oxygen. Now, we have seen that in the formation 
 of ozone, oxygen contracts from 3 volumes to 2, it follows that a 
 given weight of ozone occupies only two-thirds the volume of the 
 same weight of oxygen ; hence, the weights of equal volumes of 
 oxygen and ozone must be to each other as 2:3; and hence, if 
 there are the same number of molecules in equal volumes of each 
 gas, the weight of a molecule of oxygen must be to that of ozone as 
 2 : 3. We will :learn,v empirically for the present, that the molecule 
 of oxygen is composed of two atoms, and that its molecular weight, 
 as it is the sum of the atomic weights of the atoms composing the 
 molecule, must be 32, it follows that the molecular weight of ozone 
 is 48, and that, if ordinary oxygen has a molecule composed of two 
 atoms of oxygen, ozone must have one consisting of three. In this 
 case, then, the cause of allotropism is evidently found in the differ- 
 ent molecular structure of the two modifications of the same 
 element, and from this the student will see that a change in the 
 composition of a molecule brings with it a change in the character 
 of the substance, regardless of whether that molecule is composed 
 of atoms of the same kind or of those of different kinds. The two 
 reactions, 
 
 S + 2 = S0 2 , Sulphur dioxide, 
 + 2 = 00 2 Ozone, 
 will serve to make this meaning more clear. By oxidizing sulphur 
 
OZONE; PROPERTIES. 
 
 we obtain sulphur dioxide, a body differing in properties from both 
 sulphur and oxygen ; by oxidizing oxygen we obtain ozone, a body 
 differing in properties from oxygen, but not, perhaps, as markedly 
 as sulphur dioxide does from sulphur. An addition of energy is 
 necessary in order to produce ozone from oxygen ; it is therefore an 
 endothermic compound, and, as a consequence, has a great tendency 
 to break down with the evolution of heat. The fact that it can 
 oxidize metals under ordinary conditions, has already been alluded 
 to ; * it also can oxidize a great many organic substances, such as 
 albumen, milk, shavings, corks, or india rubber ; if such substances 
 are placed in oxygen containing ozone, the odor of the latter 
 disappears at once. As ozone is formed in a great variety of ways, 
 for instance, by the evaporation of liquids or by discharges of elec- 
 tricity, it follows that more or less of the substance must occur in 
 the atmosphere at times, but, owing to the presence of oxidizable 
 substances, we should scarcely expect any ozone to be present in 
 the air of cities, t Large quantities of ozone would undoubtedly be 
 harmful if inhaled ; it has never been proved that small quantities 
 have any effect. J 
 
 Ozone is a gas which has a blue tint, which can be seen by look- 
 ing in the direction of a white paper, through a long tube contain- 
 ing ozone. If it could be obtained pure, it would undoubtedly be 
 easily condensed to a liquid, for, although it is always greatly 
 diluted with oxygen, it nevertheless forms an indigo-blue liquid at 
 temperatures above those required to liquefy oxygen. 23 
 
 Hydrogen and oxygen form two distinct compounds, in one of 
 which (water) there are two parts by weight of hydrogen united to 
 sixteen of oxygen ; in the other, two of hydrogen to thirty-two of 
 oxygen ; the existence of these two compounds forms an excellent 
 example of the law of multiple proportions. We have already 
 decided that the formula of water is H 2 0, and hence we must 
 assign the formula H 2 O 2 to hydrogen dioxide, remembering that, as 
 we have not been able to obtain this substance in the form of a gas, 
 
 * Some pure and bright mercury shaken with gas which contains even 
 traces of ozone, is instantly oxidized, the mercury losing its lustre and, in 
 part, adhering to the sides of the flask. 
 
 t The presence of ozone in the atmosphere has recently been denied. Ilos- 
 vay-Ilosva; Bulletin Soc. Chim.; [3] 2, 377. 
 
 + Labbe and Oudin; Comptes Rendus; 113, 141. 
 
50 HYDROGEN DIOXIDE; PREPARATION. 
 
 H 2 2 can represent only the formula weight ; the molecular weight 
 may be any multiple of this formula weight, or (n H 2 2 ).* 
 
 (Hydrogen dioxide is prepared by adding a dilute acid, preferably 
 sulphuric acid, to barium dioxide. 
 
 Ba0 2 +H 2 S0 4 =BaS0 4 +H 2 2 , 
 
 Barium Dioxide -f- Sulphuric acid = Barium sulphate + Hydrogen dioxide. 
 Barium sulphate is insoluble in water ; 24 it can therefore be 
 allowed to settle to the bottom of the vessel in which the dioxide 
 of hydrogen is prepared, and the clear supernatent liquid then 
 poured off; by allowing the excess of water to evaporate, 25 there 
 remains a very concentrated solution of hydrogen dioxide, having a 
 specific gravity of 1.45, and which does not freeze at 30. The 
 concentrated solution" must be preserved in ice, for on warming to 
 ordinary temperatures, a rapid evolution of oxygen takes place, and 
 nothing but water remains; a too rapid heating of the liquid to 
 the boiling point of water will even cause it to explode. Dilute 
 solutions of the dioxide have a bitter taste ; they are used for 
 purposes of bleaching. 
 
 Hydrogen dioxide owes its chief value to the readiness with 
 which it yields its oxygen, it resembling ozone in that particular ; 
 indeed, one of the most delicate tests for both of these substances 
 is the same, and owes its value to their oxidizing power. Hydrogen 
 dioxide or ozone, when either is added to a solution of potassium 
 iodide, yields iodine : 
 
 H 2 2 + 2KI =2KOH + 21. 
 
 Hydrogen dioxide + Potassium iodide = Potassium hydroxide + Iodine. 
 3 +2KI +H 2 = 2KOH + 2 +21 
 
 Ozone + Potassium iodide + Water = Potassium hydroxide -f Oxygen + Iodine. 
 
 In the latter case the addition of water is necessary; in the 
 former a compound of hydrogen and oxygen which yields water is 
 already present, but in both cases the oxygen given off changes the 
 iodide to the hydroxide of potassium. Iodine has the power of 
 turning starch paste to a deep blue color, so that the addition of 
 some of this substance will render even minute traces of iodine 
 visible. ) > 
 
 * The molecular weight of hydrogen dioxide is probably 34, and the mole- 
 cule H 2 O 2 , a fact which has been discovered by determining the freezing point 
 of hydrogen dioxide solutions. See Orndorff and White; Zeitschrift fiirPliysik. 
 Chemie; 12, 1. 
 
NASCENT STATE. 51 
 
 Both ozone and hydrogen dioxide owe their peculiar powers of 
 oxidation to the fact that they can yield oxygen in a condition 
 known as the nascent state. Oxygen, as well as a number of other 
 elements, exists as molecules, each molecule being formed of two 
 atoms. A considerable amount of energy is necessary to decompose 
 these molecules ; indeed, in the case of hydrogen, for instance, it is 
 doubtful whether any heat which we can command will be able to 
 decompose its molecules into individual atoms. As a consequence, 
 it follows that these atoms possess much more chemical energy than 
 do the molecules, and hence must have a greater tendency to unite 
 with some other atom or with some molecule. When an element is 
 liberated from any of its compounds, it must at first exist as indi- 
 vidual atoms, no matter how short a time is necessary for the atoms 
 to unite to form molecules. If, however, any substance is present 
 011 which the atoms can act, they will primarily react with that sub- 
 stance. It follows, therefore, that elements are chemically most 
 active at the very moment of their liberation from compounds (in 
 statu nascendi). If we pass hydrogen gas through nitric acid no 
 change will take place, no matter how long we may continue the 
 operation ; but if we generate hydrogen within the acid, as, for in- 
 stance, by placing a piece of zinc in nitric acid, the hydrogen will 
 rob the nitric acid of its oxygen, forming water and an oxide of 
 nitrogen containing less oxygen than does nitric acid. 1 * Instances 
 of the action of elements in the nascent state are extremely numer- 
 ous, but we are also aware of a number of cases where compounds 
 are more energetic, chemically, at the moment of their formation 
 than at any other time, and in such cases this explanation of the 
 nascent state is inadequate. The compound CO, carbon monoxide, 
 can act, under certain conditions, as if it were in the nascent state ; 
 but we have no reason to suppose that this compound ever exists 
 otherwise than as the molecule represented by the formula CO. 
 The above explanation of the nascent state, if correct, is probably 
 applicable only in a limited number of cases. The fact remains, 
 however, no matter what reason we see fit to assign to the 
 phenomenon, that elements frequently enter into reaction at the 
 moment of their liberation from compounds, where they would be 
 entirely indifferent under other circumstances. Oxygen in the 
 nascent state is liberated by hydrogen dioxide and by ozone, for 
 
 * See chapter xxvi. 
 
52 NASCENT STATE. 
 
 
 
 one molecule of ozone, | > (0 3 ) breaks down into one molecule of 
 
 
 
 
 oxygen and an atom of the same element | > = | +0, while 
 
 
 
 the same quantity of hydrogen dioxide decomposes into a molecule 
 
 H-0 H +0 
 of water and an atom of oxygen, | = > , and, as a coii- 
 
 H-0 H 
 
 sequence, both ozone and hydrogen dioxide are powerful oxidizers 
 and bleachers. The tendency of oxygen to unite with other atoms, 
 when it is liberated from ozone, is so great that it can even take 
 oxygen away from other compounds to form a molecule of oxygen ; 
 for instance, -when ozone is brought in contact with silver oxide, 
 the following reaction takes place : 
 
 and similar changes are produced with hydrogen dioxide.* 
 
 * The occurrence of hydrogen dioxide in the atmosphere is doubtful. 
 See Schone ; Berichte d. Deutsch. Chem. Gesell.; 26, 3011. Hydrogen dioxide 
 has lately been obtained in a pure state by distillation in a vacuum ; its boiling 
 point is 85 under 68 mm. pressure. R. Wolfenstein, Ber. d. Deutsch. Chem. 
 Gesell.; 27, 3307. 
 
THE HALOGENS ; COMPARISON OF. 53 
 
 CHAPTER VI. 
 
 THE HALOGENS. 
 
 WE have now studied a metal, a typical not-metal, and the pro- 
 duct formed by the union of the two ; and have gained an insight 
 into quite a number of chemical reactions, as well as into the mode 
 of action of the molecules and atoms ; so that we will now go to the 
 discussion of the elements by families, taking them up in the 
 natural order assigned to them by their atomic weights, remember- 
 ing that, as was said in the introduction, the properties of the 
 various elements are determined by their atomic weights. The first 
 group of elements which we will study includes those which are 
 most not-metallic in their characteristics, and the plan will be to 
 work from this family through others with a diminishing not- 
 metallic character, until finally we arrive at groups composed 
 entirely of metals. 
 
 The group of Halogens (salt producers) comprises four elements : 
 
 Fluorine, atomic weight 19. 
 
 Chlorine, 35.45. ' 
 
 Bromine, " " 79.95. <% 
 
 Iodine, " 126.85. , ., 
 
 With increasing atomic weight we have a decrease in the not-metallic 
 properties of the elements included in this family, and this change 
 is well shown in the decreasing stability of the compounds of the 
 halogens with the metals. If we examine the hydrogen compounds, 
 which are formed by the union of one atom. of hydrogen with one of 
 the halogen (HF, HC1, HBr, HI), we are at once impressed by this 
 change for hydroiodic acid (HI) decomposes most readily upon 
 heating, a hot wire introduced into the gas will change it to hydro- 
 gen and iodine ; hydrobromic acid (H Br) is less readily separated 
 into its elements ; hydrochloric acid (H Cl) is broken down, if at 
 all, only by the application of a very great amount of heat ; and we 
 have, so far as we know, never effected a decomposition of hydro- 
 
54 THE HALOGENS ; COMPARISON OF. 
 
 fluoric acid by heat alone. We can see this same difference illus- 
 trated by bringing chlorine into a solution containing sodium 
 "bromide or iodide; the chlorine will at once form sodium chloride, 
 liberating bromine or iodine ; while bromine, added to a solution of 
 sodium iodide, will set iodine free, and fluorine would, without 
 doubt, decompose the compounds of any of the other halogens with 
 the metals. This difference in the character of the halogens is 
 shown by the heat of formation of the hydrogen compounds given 
 in the table at the end of Chapter XI. With increasing atomic 
 weight we necessarily have changes in the physical properties of 
 the elements. Fluorine is a nearly colorless gas, chlorine a yellow- 
 ish-green gas, rather easily converted to a liquid, bromine is a dark 
 brown liquid at ordinary temperatures, while iodine is a solid of 
 almost metallic appearance. The halogens all have a peculiar odor, 
 and attack the skin and mucous membrane. Other points of resem- 
 blance will become apparent in the detailed study of the elements. 
 They all, with the . exception of fluorine and bromine, form oxides, 
 and, all but fluorine, acids with oxygen and hydrogen ; a study of 
 the latter will be deferred until a subsequent chapter is reached. 
 The halogens themselves are never found as free elements, but occur 
 united to some metal, as the fluoride, chloride, bromide, or iodide. 
 The metals most frequently found united with the halogens are 
 sodium, potassium, magnesium, or calcium, so that for instance, 
 sodium chloride (common salt) is the most frequently occurring 
 compound of chlorine. Having given a few of these general char- 
 acteristics, we will go to the study of the individual elements. 
 
FLUORINE; OCCURRENCE, PREPARATION. 55 
 
 CHAPTER VII. 
 
 FLUORINE AND HYDROFLUORIC ACID. 
 
 Symbol, F; atomic weight, 19.0^ Formula, HF; specific gravity of 
 liquid, .9879 ; of gas, air = 1, is 1.364 ; H 2 = 2, is 39.32 ; the 
 molecule is represented by the formula H 2 F 2 . 
 
 THIS element chiefly occurs in nature, combined with the metal 
 calcium, as fluorspar (fluorite, CaF 2 ), a crystalline mineral quite 
 frequent of occurrence ; in addition to this, cryolite, a fluoride of 
 sodium and aluminium (A1F 8 , 3 NaF), occurs in large masses in 
 Greenland, and is a considerable source of fluorine compounds ; 
 while small quantities of fluorides occur in the enamel of teeth and 
 blood, and traces of the same are found in sea-water. 
 
 Fluorine has, because of its great chemical affinity for other ele- 
 ments, be they metal or not-metal, until recently resisted all attempts 
 to isolate it ; this is due to the fact that it would combine with other 
 substances as soon as liberated from its compounds. Quite recently 
 a French chemist, Moissan, succeeded in preparing the element by 
 electrolysis of perfectly pure liquid hydrofluoric acid placed in a U 
 shaped platinum tube and cooled to a low temperature.* When the 
 electric current passes through hydrofluoric acid the latter is decom- 
 posed into hydrogen and fluorine, just as we decompose water into 
 hydrogen and oxygen, the hydrogen separating at the negative pole, 
 the fluorine at the positive. 
 
 The element is a pale yellow gas,f which does not attack plati- 
 num at low temperatures, but instantly unites with elements such 
 as silicon, boron, arsenic, sulphur, iodine, iron ; and with organic sub- 
 stances, such as cork, petroleum, etc. The substances so attacked 
 take fire in the gas, so that all the phenomena of combustion in 
 oxygen are repeated with fluorine in a more violent degree and 
 under ordinary conditions. If the gas is passed into water, the 
 
 * By liquid methyl-chloride, boiling point 22.5. 
 t Moissan; Comptes Rendus; 109, 937. 
 
56 FLUORINE; PROPERTIES. HYDROFLUORIC ACID. 
 
 latter is instantly decomposed, ozone * and hydrofluoric acid being 
 produced. This reaction is very interesting, for the power of 
 decomposing water which members of the halogen family possess, 
 diminishes with increasing atomic weight. The reaction is as 
 
 follows : 
 
 2 F + Ho = 2 HF + O. 
 
 Fluorine + Water = Hydrofluoric acid + Oxygen. 
 
 The atoms of nascent oxygen when liberated can combine with 
 each other to form ozone. 
 
 The compound of hydrogen and fluorine, hydrofluoric acid, was 
 first identified as a peculiar acid by Scheele (1771), although the 
 fact that a substance which would attack glass could be produced 
 by the action of sulphuric acid on fluorspar had been known for 
 some time (since 1670). The nature of hydrofluoric acid at first 
 was misunderstood, owing to the opinion formerly held by chemists 
 that all acids must contain oxygen, so that the existence of oxygen 
 was presupposed in hydrofluoric acid. We now know that no com- 
 pounds of fluorine and oxygen exist. 
 
 Hydrofluoric acid can, as we have seen, be produced by the 
 action of fluorine on water and by direct union of hydrogen and 
 fluorine, just as water is formed by direct union of hydrogen and 
 oxygen. So great is the tendency to form hydrofluoric acid that 
 fluorine will probably take hydrogen away from any other com- 
 pound containing that element. To prepare hydrofluoric acid for 
 the laboratory or for commercial purposes, other, less expensive, 
 means than the ones which have been given are resorted to. If a 
 fluoride, such as sodium fluoride or calcium fluoride, is treated with 
 sulphuric acid, the following reaction takes place : f 
 
 2NaF + H 2 S0 4 = Na 2 S0 4 -f 2 HF 
 
 Sodium fluoride + Sulphuric acid = Sodium Sulphate + Hydrofluoric acid. 
 
 CaF 2 +H 2 S0 4 = CaS0 4 + 2 HF 
 
 Calcium fluoride + Sulphuric acid = Calcium Sulphate + Hydrofluoric acid. 
 
 Reactions such as the above are very frequently met with ; the 
 hydrogen of the acid simply exchanging places with the metal of 
 
 * If a tube, filled with fluorine, has a few drops of water admitted to it, 
 ozone is formed in such quantity that the contents are temporarily changed to 
 a deep blue color. 
 
 t As hydrofluoric acid attacks glass, the pure substance must be prepared 
 in platinum vessels. 
 
HYDROFLUORIC ACID; PROPERTIES. 57 
 
 the salt to form a new salt and a new acid. Such reactions are 
 designated as double decompositions, and as a practical hint we can 
 say that they take place when an insoluble or a volatile substance 
 can be produced.* We shall inquire more closely into the nature 
 of double decompositions when we have studied a larger number of 
 chemical reactions. The method given above is one which is very 
 frequently employed in the preparation of the various acids from 
 their salts. 
 
 Hydrofluoric acid is a colorless, mobile liquid which freezes at 
 102, boils at -f- 19 (about the temperature of a warm room), 
 and which fumes strongly in the air because of its attraction for 
 moisture. The vapor of hydrofluoric acid is very irritating even 
 when inhaled in small quantities, while any considerable amount 
 can cause death. A drop of the acid put on the hand causes a most 
 painful blister, which ultimately changes to a slowly healing ulcer, 
 so that great care must be exercised in handling this acid. The 
 usual commercial acid is a solution of hydrofluoric acid in water. 
 It is a colorless, extremely acid liquid, it fumes in the air, and is 
 transported in bottles made either of paraffin or of guttapercha, 
 for the acid readily attacks glass. 
 
 * Hydrofluoric acid is a volatile substance. 
 
58 CHLORINE; OCCURRENCE, HISTORY. 
 
 CHAPTER VIII. 
 
 CHLORINE. 
 
 5 a 
 
 Symbol, Cl ; atomic weight, 35.45 ; specific gravity, air = 1 is, 2.46, 
 
 H 2 = 2, is 70.84. 1 c.c. iveighs .0031825 gram. 
 
 CHLORINE occurs in nature combined with various metals as 
 chlorides, never as the free element. The most important chloride 
 is that of sodium (Nad, sodium chloride, or common salt).,; This 
 substance forms the major portion of the solid residue left upon 
 evaporation of sea-water, and is consequently the larger part of the 
 salt beds of marine origin and of those composed of rock salt. 
 Chlorine also occurs in considerable quantities combined with 
 potassium as the mineral sylvin (KC1, potassium chloride), and as 
 a chloride of magnesium and of potassium ( K Cl, Mg C1 2 + 6 H 2 0, 
 called carnallite) ; the chlorides of iron, lead, silver, etc., occur in 
 small quantities, while chlorides are found in the tissues and fluids 
 of plants and animals and in their ashes. 
 
 The element was not discovered until 1774, at the beginning of 
 the period in which the great advances in chemistry recorded in 
 the introductory chapter were made ; the first chemist to prepare 
 chlorine was Scheele. He called it dephlogisticated muriatic acid, 
 for it was muriatic acid from which phlogiston (hydrogen) had been 
 extracted. Chlorine was for some time supposed to be a compound 
 of oxygen with an unknown element called murium ; Sir Humphry 
 Davy first definitely asserted that chlorine was an element, naming 
 the element chlorine from x\wp6s, greenish yellow, the gas having 
 that color. 
 
 1. Preparation of Chlorine by Electrolysis. 
 
 Evidently, in order to prepare chlorine, our method must be to 
 remove the metallic constituent from some chloride ; hydrogen chlo- 
 ride (hydrochloric acid) being the chloride easiest available. If we 
 subject a concentrated solution of hydrochloric acid to the action of 
 an electric current, 26 we shall decompose the substance in exactly 
 the same manner as we can hydrofluoric acid, with the difference 
 
CHLORINE; PREPARATION. 59 
 
 that we can readily perform this operation in glass vessels (which 
 are not attacked by chlorine), as the chemical energy displayed by 
 chlorine is not so great as that of fluorine. We can also remove 
 hydrogen from hydrochloric acid by other means ; for instance, 
 by some oxidizing agent, when water and chlorine are produced, as 
 follows : 
 
 2HC1 +0 =H 2 0+2C1. 
 
 Hydrochloric acid + Oxygen = Water + Chlorine. 
 
 The oxygen of the atmosphere can accomplish this under proper 
 conditions, and a process of commercial preparation of chlorine* 
 has its origin in this fact. If a mixture of hydrochloric acid and 
 oxygen is passed through a heated tube in which are placed pieces 
 of porcelain or fire-brick, saturated with a solution of copper sul- 
 phate, chlorine and water are formed. The copper sulphate re- 
 mains unchanged at the end of the reaction, so that the reason for 
 its peculiar action is not understood. 
 
 2. Preparation of Chlorine from Manganese Dioxide and Hydro- 
 chloric Acid. 
 
 Manganese dioxide is the most convenient oxidizer for the 
 preparation of chlorine ; when it is brought in contact with hydro- 
 chloric acid, the following reaction takes place : 27 
 
 1. Mn0 2 +4HC1 =MnCl 2 +2H a O+2Cl. 
 
 Manganese dioxide .+ Hydrochloric acid=Manganous chloride + Water + Chlorine. 
 
 The manganese dioxide furnishes the oxygen which changes 
 the hydrogen of hydrochloric acid to water, and at the same time 
 in all probability a chloride of manganese, having four atoms of 
 chlorine in the formula weight, is formed, thus : 
 
 2. 
 
 This chloride is, however, very unstable and breaks down into 
 manganous chloride and chlorine, as follows : 
 
 so that, combining equations 2 and 3, we obtain equation 1. As a 
 rule, chemical equations are expected only to represent the ultimate 
 
 * Deacon's process. 
 
60 CHLORINE; PREPARATION. 
 
 product of any chemical reaction, as does equation 1, but if we 
 wish thoroughly to understand chemical changes we must not be 
 contented with the mere equation, but should also inquire into all 
 of the intermediary stages which bring about chemical reactions. 
 Other oxidizing agents, as well as manganese dioxide, * are capable 
 of furnishing oxygen to form water and chlorine from hydro- 
 chloric acid. 
 
 3. Preparation of Chlorine from Sodium Chloride, Sulphuric Acid, 
 and Manganese Dioxide. 
 
 A mixture of common salt (Na Cl, sodium chloride) and sulphuric 
 acid yields hydrochloric acid, so that it is often convenient to pre- 
 pare chlorine by mixing sodium chloride and manganese dioxide in 
 a flask and then adding sulphuric acid. In this case it must be 
 remembered that the manganous chloride (MnCl 2 ) which might 
 be formed (equation 3, preceding page,) would also be acted upon 
 by the sulphuric acid, forming manganous sulphate and hydro- 
 chloric acid, which would further be converted by the manganese 
 dioxide, according to equation 2 : 
 
 Mn C1 2 + H 2 S0 4 = Mn S0 4 + 2 H Cl. 
 
 MnO 2 + 2 H 2 S0 4 + MnCl 2 = 2 MnS0 4 +2 H 2 0+2 Cl; 
 
 so that, when the usual mixture of sodium chloride and manganese 
 dioxide is used, the following complete reaction takes place : 
 
 Mn0 2 + 2 Na Cl + 2 H 2 S0 4 = MnS0 4 + Na 2 S0 4 + 2 H 2 + 2 Cl. 
 
 In all cases the principle of the reaction is the one given on the 
 previous page, the sodium chloride serving simply to furnish hydro- 
 chloric acid. ) 
 
 4. Other Methods for Preparing Chlorine. 
 
 Several other methods for the preparation of chlorine have been 
 devised ; some of these are of commercial value, as, for instance, 
 the preparation by heating magnesium chloride in a current of air, 
 
 * Such oxidizing agents are potassium permanganate (K Mn O 4 ), potas- 
 sium bichromate (K 2 Cr 2 O 7 ), nitric acid (HNO 3 ), etc. In all these cases the 
 principle is the same, the sole object being to remove hydrogen from hydro- 
 chloric acid, forming chlorine and water. Manganese dioxide is employed 
 because it is cheap. 
 
CHLORINE; PROPERTIES. 61 
 
 the reaction resembling that of the action of oxygen on hydrochloric 
 
 acid : 
 
 MgCl 2 +0 = MgO + 2C1 
 
 Magnesium chloride + Oxygen = Magnesium oxide + Chlorine. 
 
 The magnesium oxide formed can be dissolved in hydrochloric acid, 
 once more forming magnesium chloride, and thus the process can 
 be continued without serious loss of magnesium.* 
 
 (Chlorine, at ordinary temperatures, is a greenish yellow gas, the 
 color of which becomes darker upon heating. It has a peculiar, 
 irritating odor, which must be tested only when the gas is very 
 dilute, for any great quantity of chlorine entirely destroys the sense 
 of smell, causing inflammation of the mucous membrane of the 
 throat and lungs, coughing and hemorrhages. An annoying 
 catarrh follows the inhalation of the gas, so that great care must 
 be exercised in working with chlorine. The presence of any con- 
 siderable quantity of chlorine in the air may cause death. A pres- 
 sure of six atmospheres, at 0, condenses chlorine to a liquid ; at 
 ordinary atmospheric pressure it becomes liquid at 35 and it 
 freezes at 102 ; its specific gravity, air being one, is 2.46 at 
 temperatures up to 1200,f at higher temperatures the specific 
 gravity becomes less, being but 2.02 at 1400 ; this shows that the 
 chlorine molecules begin to decompose into the individual atoms at 
 about white heat.J This decomposition of molecules into simpler 
 ones or into atoms is termed dissociation ; the temperature of dis- 
 sociation varies with different substances according to the amount 
 of heat given off in their formation ; it being, of course, necessary 
 to add as much energy to decompose a substance as is given off in 
 its formation. At very high temperatures, such as exist in, the 
 chromosphere of the sun, dissociation of all complex substances 
 is complete, so that in such a situation no chemical compound is 
 possible. 
 
 Chlorine is tolerably soluble in water, one volume of that liquid 
 
 * A convenient laboratory method for preparing chlorine is by the decom- 
 position of pressed cubes of calcium hypochlorite (chloride of lime) by means 
 of diluted hydrochloric acid. (See chapter XVIII.) 
 
 t This, H = 2 would give a specific gravity of 70.84, the atomic weight of 
 chlorine is 35.45, which would give a molecular weight of C1 2 = 70.90, hence 
 below 1200 chlorine molecules consist of two atoms to the molecule. 
 
 $ Langer and V. Meyer; Berichte d. Deutsch. Cheni. Gesell. ; 15, 1721. 
 
62 
 
 CHLORINE; CHEMICAL REACTIONS. 
 
 absorbing 2.5 volumes of chlorine at ordinary temperatures. The 
 solution of the gas so produced, known as chlorine water, has the 
 odor of chlorine and many of the chemical properties of the gas. 
 When chlorine water is cooled nearly to the temperature of freezing 
 water it is changed to a transparent, crystalline substance (chlorine 
 hydrate,* having the composition 2 Cl -}- 8 H 2 0), which slowly 
 gives off chlorine at low temperatures, and 
 rapidly upon heating. If a few chlorine 
 hydrate crystals are placed in one end of a 
 bent glass 'tube, which is sealed at the other 
 extremity, cooled by snow and salt (such a 
 tube is shown by Fig. 6), chlorine will con- 
 dense to a liquid in the cold part of the 
 tube as soon as the crystals are gently 
 warmed. 
 
 Chemically, chlorine greatly resembles 
 oxygen, with this distinction ; while oxygen, 
 
 under common circumstances, Is inactive, chlorine, at ordinary 
 temperatures, unites with many elements, metallic or not-metallic, 
 to form chlorides, the formulae of which bear a great resemblance 
 to the corresponding oxides ; for example : 
 
 P 2 5 PC1 5 A1 2 8 Aids 
 
 Phosphorus pentoxide. Phosphorus pentachloride. Aluminium oxide. Aluminium chloride. 
 
 Fig. 8. 
 
 P 2 3 
 
 Phosphorus trioxide. 
 
 PC1 3 
 
 Phosphorus trichloride. 
 
 The difference between the formulae of chlorides and oxides, as 
 seen from the above, is that in the oxides one-half as many atoms 
 of oxygen unite with one atom of the other element entering into 
 the compound as do chlorine atoms in the chlorides ; this relation- 
 ship becomes clearer if we double the formulae of the chlorides for 
 purposes of comparison : 
 
 P 2 5 , 
 
 P 2 C1 10 . 
 P 2 C1 6 . 
 
 One oxygen atom is therefore capable of taking the place of 
 two chlorine atoms in chemical compounds, and in writing chemical 
 
 Prepared by passing chlorine into ice-water. 
 
CHLOKINE; COMBUSTION IN. 63 
 
 formulae this difference must always be borne in mind. What is 
 true of chlorine applies to the other halogens as well. 
 
 The union of the various elements with chlorine and the resem- 
 blance between these reactions and combustions in oxygen is made 
 clear by the following examples : 
 
 'Phosphorus, which has previously been ignited in the air, will 
 continue to burn in an atmosphere of chlorine with a pale greenish 
 flame : 
 
 P + 3 Cl = P C1 8 (Phosphorus trichloride). 
 
 Pronounced metals, such as sodium, when heated to their 
 kindling temperature, will burn in chlorine gas : 
 
 Na + Cl = NaCl (Sodium chloride). 
 
 Carbon, when heated in the presence of chlorine, will form 
 a chloride : 
 
 C + 4 Cl = C C1 4 (Carbon tetrachloride).* 
 
 Chlorine unites with hydrogen with such facility that it will 
 even extract hydrogen from its numerous compounds with carbon. 
 A piece of filter paper, saturated with turpentine and placed in 
 chlorine gas, will take fire ((sometimes with explosive violence)} form- 
 ing hydrochloric acid, (and separating, in the form of soot, the 
 carbon which was in the turpentine^f But it is not only from com- 
 pounds of carbon that chlorine will extract hydrogen ; it will do so 
 from a multitude of other substances containing this element ; for 
 example, chlorine will decompose sulphuretted hydrogen ( H 2 S), 
 ammonia (NH 8 ), or even water, in each case forming hydro- 
 chloric acid, and liberating the element previously combined with 
 hydrogen : 
 
 H 2 S + 2C1 =2HC1 +S, 
 
 Hydrogen sulphide + Chlorine = Hydrochloric acid + Sulphur, 
 NH 3 +3C1 =3HC1 +N, 
 
 Ammonia + Chlorine = Hydrochloric acid + Nitrogen, 
 
 H 2 + 2C1 =2HC1 +0. 
 
 Water + Chlorine = Hydrochloric acid + Oxygen. 
 
 * The pupil will note that in this formula two chlorine atoms take the 
 place of one oxygen atom, as will be seen by comparing the formulae C C1 4 
 and CO., . 
 
 t Turpentine is a compound of carbon and hydrogen. 
 
64 CHLORINE WATER. 
 
 The cause for these reactions in each case is that, in contact 
 with hydrogen, chlorine possesses greater chemical energy than the 
 three other elements. This is shown by the fact that hydrochloric 
 acid, when formed from its elements and then dissolved in water, 
 has a greater heat of formation than water, sulphuretted hydrogen, 
 or ammonia. 
 
 Thus, NH 3 formed from its elements, gives 204 K,* 
 Ho S " " 73 K, 
 
 H 2 " " " 684 K, 
 
 2HC1 " 786 K.f 
 
 Chlorine water placed in the sunlight will form hydrochloric 
 acid and liberate oxygen ; J but chlorine and water can yield oxygen 
 and form hydrochloric acid even without the aid of sunlight, pro- 
 vided some substance is present which can be oxidized. It is to 
 this property that chlorine owes its commercial value, its chief 
 industrial use being as a bleaching agent, its power of bleaching 
 depending, at least in the great majority of cases, upon its capability 
 of liberating nascent oxygen from water. That this is the case can 
 be proven by placing a piece of colored calico in a jar of dry chlo- 
 rine j no bleaching will take place until water is added, and then 
 
 * K stands for the quantity of heat which a gram of water loses when 
 cooled from 100 to 0. In speaking of heats of formation or of the thermal 
 changes during reactions, the quantities of substances reacting are taken at 
 as many grams as are expressed by the atomic or formula weights. Thus, 
 when we say the heat of formation of H 2 O is 680 K, we mean that two grams 
 of hydrogen uniting with sixteen of oxygen give 680 K. By heat of solution 
 we mean the heat given off by dissolving the formula weight of a substance, 
 in grams, in an unlimited quantity of water. Thus " the heat of solution of 
 hydrochloric acid is 172 K," means that 36.45 grams of hydrochloric acid give 
 off 172 K while dissolving in. an unlimited amount of water. The unit, K, 
 and the figures which are given here and in subsequent parts of the work, 
 correspond to those employed by Ostwald. See also Ostwald's Outlines of 
 General Chemistry, Walker's translation, Macmillan, 1890, p. 212. 
 
 t These figures refer to the heats of formation of the compounds from 
 their elements, and in the presence of an excess of water. 
 
 (t Chlorine water liberates oxygen and forms pure hydrochloric acid only 
 when the solution is placed in the direct sunlight. In diffused light, even in 
 bright daylight, in addition to oxygen, hypochlorous acid and chloric acid as 
 well as hydrochloric acid are produced. Pedler; Journal of the Chem. Soci- 
 ety ; 1890, 613/j 
 
CHLORINE; OXIDIZING POWERS. 65 
 
 the bleaching action of chlorine at once becomes apparent. 28 Chlo- 
 rine is very frequently employed as an oxidizing agent in laboratory 
 work ;\ an oxidizing agent being a body which can chemically furnish 
 oxygen, either per se, or by the decomposition of some oxide. ) Ex- 
 amples of the former class are such bodies as manganese dioxide, 
 nitric acid, potassium bichromate or potassium permanganate ; all of 
 these substances are direct oxidizers, and, while they oxidize, they 
 themselves are reduced. Examples of indirect oxidizers are chlorine, 
 bromine, or iodine, for these elements decompose water in order to 
 accomplish the same result. 
 
66 HYDROCHLORIC ACID; HISTORY. 
 
 CHAPTER IX. 
 
 HYDROCHLORIC ACID. 
 
 Formula, H Cl ; specific gravity, air = 1, is 1. 2658, H 2 =.-2, is 36.45 ; 
 1 c.c. at and .76 m. weighs .0016442 gram. 
 
 HYDROCHLORIC acid very seldom occurs free in nature, and 
 then only in the exhalations of some volcanoes and in the springs 
 arising from their craters ; for instance, the Eio Vinagre, arising in 
 the Andes, is said to contain .08 per cent, the Paramo de Ruiz, in 
 New Granada, .8 per cent of free hydrochloric acid. 
 
 The aqueous solution of the acid was first prepared by Basil 
 Valentine in the fifteenth century, by distillation of salt (NaCl, 
 sodium chloride) with ordinary green vitriol (FeS0 4 , ferrous sul- 
 phate) ; it was subsequently investigated by a number of alchemists, 
 but the pure gas was not obtained until 1772, when Priestley isolated 
 pure hydrochloric acid. The old name was spiritus sails, or acidum 
 muriaticum (from murias, sea salt), and at the present day the 
 aqueous acid is frequently termed muriatic acid. When first inves- 
 tigated, hydrochloric acid was supposed to contain oxygen; but 
 Davy, in 1810, proved that it was composed of nothing but hydrogen 
 and chlorine. 
 
 The acid is best prepared by treating the chloride of a metal, 
 preferably sodium chloride, with sulphuric acid, when the following 
 reaction takes place : 
 
 2 Nad +H 2 S0 4 =Na 2 S0 4 +2HC1 
 
 Sodium chloride + Hydrogen sulphate = Sodium sulphate + Hydrogen chloride. 29 
 
 The gas can be collected over mercury, or, being heavier than 
 air, by downward displacement, but not over water, as it is ex- 
 tremely soluble in that substance. 
 
 Hydrochloric acid is a colorless gas with an acid odor ; it fumes 
 in the air, owing to its power of condensing moisture from the 
 atmosphere to form an aqueous solution of hydrochloric acid ; it 
 cannot be breathed, as it causes violent coughing; it is neither 
 
HYDROCHLORIC ACID; PREPARATION. 67 
 
 combustible, nor will other substances burn in it. ) The stability of 
 union of hydrogen and chlorine is very great ; at about 1800 (high 
 white heat) hydrochloric acid begins slightly to decompose into 
 hydrogen and chlorine, but, as we have become aware of the great 
 tendency which hydrogen has to unite with chlorine, this stability 
 is not unexpected. Hydrochloric acid is very soluble in water ; at 
 one volume of water will dissolve 505 volumes of hydrochloric 
 acid gas ; the solution then contains 43 per cent of the acid. 30 
 
 (When chlorine is brought in contact with hydrogen, in the dark, 
 no reaction takes place ; if the mixture of the two gases is exposed 
 to the sunlight, or if a lighted taper is applied, or an electric spark 
 allowed to pass through the gases, a violent explosion takes place 
 and hydrc^lilcTic_aci^ is produced. On the other hand, if a current 
 of electricity is passed through a concentrated solution of hydro- 
 chloric acid, 81 the chlorine will separate at the positive pole, the 
 hydrogen at the negative ; this resembles the decomposition of 
 water, excepting that with hydrochloric acid equal volumes of 
 hydrogen and chlorine are produced, while in the case of water two 
 volumes of hydrogen and one of oxygen result. When hydrogen 
 and chlorine are mixed in equal volumes and then exploded, there 
 is no change in volume, but the mixture of gases is entirely con- 
 verted into hydrochloric acid; furthermore, sodium, when it is 
 placed in a closed volume of hydrochloric acid, will absorb the 
 chlorine (forming NaCl, sodium chloride) and the volume of 
 the gas will be diminished by one-half. 82 } 
 
 We have now proved that hydrochloric acid decomposes into 
 equal volumes of hydrogen and chlorine, and that equal volumes of 
 hydrogen and chlorine unite to form hydrochloric acid, the resulting 
 volume being equal to the sum of the volumes of hydrogen and 
 chlorine before union. The relationship between the volumes of 
 hydrogen and oxygen which are capable of producing water without 
 leaving a residue of either gas is equally simple ; so, indeed, is that 
 between the volumes of any gases uniting to form a gaseous com- 
 pound, so that Gay Lussac, who first accurately investigated the 
 -subject, was able to formulate the following law : 
 
 " Two gases always unite in such a way that their volumes bear 
 I a simple ratio to each other, and the volume of the resulting prod- 
 I uct, if it is a gas, also bears a simple relationship to the volumes 
 I of the original gases." 
 
68 KINETIC GAS THEORY. 
 
 According to the kinetic theory of the nature of gases, now uni- 
 versally held, these substances are composed of particles of matter 
 which are flying about in right lines, until they impinge on the 
 sides of the vessel in which the gas is contained, or on each other, 
 when, being perfectly elastic, they rebound. The particles of the 
 gas are separated by such distances that their own volume exer- 
 cises no influence on the volume of the gas as a whole. Now, the 
 weight of a given volume of gas is but the sum of the weights of the 
 individual particles making up that gas ; and the specific gravity of 
 any gas, if air is taken as unity, is equal to the weight of a given 
 volume of that gas as compared with that of the same volume of 
 air. Investigation has shown us that the ratios between the spe- 
 cific gravities of elementary gases, and those between their com- 
 bining weights, bear a simple relationship to each other ; thus, the 
 specific gravity of hydrogen in round numbers, is .07, of oxygen, 
 1.12 ; but .07 is to 1.12 as 1 : 16, and 1 part of hydrogen unites with 
 8 parts of oxygen (1 with ^) to form water. The specific gravity 
 of chlorine is 2.46, and the relationship is : .07 : 2.45 : : 1 : 35.1 ; 
 but the combining weight of chlorine is 35.45, as that part by 
 weight of chlorine unites with one part of hydrogen. If we use 
 hydrogen as the standard instead of air,* the relationship is seen 
 more readily, thus : 
 
 H = 1. Then the 
 
 combining weight of oxygen is 8. (or ^ 6 -) its specific gravity 16. 
 " " chlorine is 35.45 " " " 35.45 
 
 " " nitrogen is 4.66 (or *) its " " 14. 
 
 From these numbers it will be seen that equal volumes of gases 
 bear a similar relationship by weight to each other, as do the indi- 
 vidual particles of which they are composed, and, therefore, it is 
 reasonable to suppose that the numbers of particles in the gases 
 themselves must bear a simple ratio to each other, as 1 : 2, 1 : 3, or 
 1 : 1, indeed the last ratio (1 : 1) is the most probable one, for, by 
 constructing a theory that in equal volumes of gases there are 
 equal numbers of particles, we can most readily explain the sim- 
 ple relationship which exists between the combining volumes of 
 gases. This was the conclusion reached by Gay Lussac in the 
 
 * The ratio between the weights of hydrogen and of an equal volume of air, 
 is as .069:1 or as 1:14.4, hence, any specific gravity with air as unity can 
 be converted to hydrogen as unity, by multiplying by 14.4. 
 
RELATIONS OF GAS VOLUMES. 
 
 69 
 
 course of his investigations. Thus, as equal volumes of hydrogen 
 and chlorine unite to form hydrochloric acid, and if the ratio be- 
 tween the weights of equal volumes of the gases is the same as that 
 between the weights of the individual particles, it follows that the 
 volumes of the two gases contained equal numbers of these parti- 
 cles before their union. Dalton, however, soon pointed out a defect 
 in this reasoning.* Let us suppose we have a volume of hydrogen 
 containing 1000 atoms, then, according to the theory, an equal vol- 
 ume of chlorine will also contain 1000 atoms, the two unite, thus 
 forming 1000 molecules of hydrochloric acid. The natural result 
 would be, 
 
 or : 1 volume hydrogen + 1 volume chlorine = 1 volume hydrochloric acid, 
 
 for, as we have seen, the volume of the molecule exercises no 
 influence on the volume of the gas. Nature, however, contradicts 
 this reasoning, for we know that hydrogen and chlorine unite with- 
 out change of volume ; in other words, 1 volume of hydrogen -f- 1 
 volume chlorine = 2 volumes hydrochloric acid^ and it follows 
 that if hydrogen and chlorine have equal numbers of atoms in 
 equal volumes, then hydrochloric acid must have but one-half the 
 number in the same volume, for, 
 
 500 
 HC1 
 
 500 
 HC1 
 
 1 volume hydrogen + 1 volume chlorine = 2 volumes hydrochloric acid, 
 
 * The example cited by Dalton in order to refute Gay Lussac's argument, 
 was nitric oxide and not hydrochloric acid. Hydrochloric acid could not have 
 been referred to by Dalton, as he supposed that the compound contained oxy- 
 gen. The same relationship by volume exists between nitrogen, oxygen, and 
 nitric oxide as obtains between hydrogen, chlorine, and hydrochloric acid, so 
 that, for the purpose of the present argument, either gas will do equally well. 
 
 t Let the pupil, instead of using the expression " volume," substitute 
 u liter," and the whole subject will appear more clear, thus, 
 
 1 liter hydrogen + 1 liter chlorine = 2 liters hydrochloric acid. 
 
70 
 
 AYOGADRO'S HYPOTHESIS. 
 
 so 1000 molecules of hydrochloric acid must occupy twice the 
 volume previously taken by 1000 atoms of hydrogen, and hence in 
 one volume of hydrochloric acid there must be but 500 molecules. 
 It was left for an Italian physicist, Arnadeo Avogadro, to explain 
 this seeming discrepancy between theory and fact. Avogadro sup- 
 posed the elementary gases to be composed of molecules instead of 
 atoms. As a usual thing, these molecules are composed of two 
 atoms, so that, accepting this hypothesis, we would arrive at the 
 following result : 
 
 200 
 
 OHC1 
 
 1 volume hydrogen + 1 volume chlorine = 2 volumes hydrochloric acid. 
 
 A reaction of this kind would then consist simply of an inter- 
 change of the atoms composing the molecules, so that, whereas we 
 previously had molecules each of which was composed of atoms 
 of the same kind, we now would have molecules each of which is 
 composed of atoms which are of a different kind. This will be 
 more apparent if we write the equation as follows : 
 
 H H H H 
 
 + = I I 
 Cl Cl Cl Cl 
 
 What is true, then, of the volume is true of the individual mole- 
 cule, there being the same number of molecules in equal volumes ; 
 the terms volume and molecule can therefore be used interchange- 
 ably. 
 
 Simple as Avogadro's explanation was, it was not generally ac- 
 cepted by chemists, chiefly because he tried to apply it in cases 
 where substances which never were in a gaseous state were con- 
 cerned, so that it was not until many years later that it was uni- 
 versally adopted. It was then brought into prominence and is now 
 one of our fundamental hypotheses, furnishing to us the best means 
 of determining the molecular weights of chemical compounds and 
 elements. 
 
 Let us suppose that we have a volume of hydrogen weighing 
 
AVOGADRO'S HYPOTHESIS. 
 
 71 
 
 two grams, then an equal volume of chlorine must weigh 70.9 grams, 
 for the weights of equal volumes of elementary gases are to each 
 other as their atomic weights and 1 : 35.45 : : 2 : 70.9. 
 
 HC1 
 
 72. 
 
 HC1 
 
 9 grams. 
 
 1 volume hydrogen + 1 volume chlorine = 2 volumes hydrochloric acid. 
 
 It follows that a volume of hydrochloric acid equal to that of the 
 
 72 9 
 hydrogen taken must weigh ~ grams or 36.45 grams, and, as in 
 
 equal volumes of the gases there are equal numbers of molecules, 
 the ratio between the weights of the individual molecules of hy- 
 drogen and hydrochloric acid must be 2 : 36.45, and therefore, if the 
 molecule of hydrogen is two, then the molecule of hydrochloric acid 
 is 36.45 ; so that if hydrogen is taken as a standard and placed at 
 two, the specific gravity of hydrochloric acid is equal to its molec- 
 ular weight,, and we shall soon see that this rule can be made 
 general as follows : ***" 
 
 If hydrogen be placed at two, then the molecular weights and 
 specific gravities of gases are the same number.* 
 
 Hydrogen and oxygen unite to form water in the proportion of 
 two volumes of hydrogen to one of oxygen and, if the water so 
 formed is measured in the form of vapor, we shall find that two 
 volumes of this vapor are produced, thus : 
 
 H 
 
 1 volume 
 
 + 
 
 H 
 
 1 volume 
 
 + 
 
 
 
 1 volume 
 
 = 
 
 2 volumes 
 
 H 2 
 
 If, according to Avogadro's hypothesis, oxygen has two atoms to 
 the molecule, and supposing that we select volumes each of which 
 contain 1000 molecules, the reaction which takes place will be as 
 follows : 
 
 * In dealing with the specific gravity of gases it is not necessary to deal 
 with exact numbers. Thus, using the oxygen standard for atomic weights, 
 hydrogen is 1.008, but, for all practical purposes', the decimal can be neglected. 
 
72 
 
 SPECIFIC GRAVITIES OF GASES. 
 
 or, as molecule and volume can be used interchangeably, 
 H H H H H H H H 
 
 00 
 
 V V 
 
 
 
 1 Mol. H + 1 Mol. H + 1 Mol. = 2 Mols. H 2 0. 
 
 Now, if each volume of hydrogen weighs two grams, then the same 
 volume of oxygen will weigh thirty-two grams, and consequently : 
 
 2 grams 
 H 
 
 + 
 
 2 grams 
 H 
 
 + 
 
 32 grs. 
 
 
 = 
 
 18 grs. 
 H 2 
 
 + 
 
 18 grs. 
 H 2 
 
 Therefore, in the case of water also, if hydrogen is placed at 
 two, the molecular weight and the specific gravity are identical, and 
 similar methods of reasoning, backed by experiment, have shown 
 the same to be true in regard to the specific gravities of all gases.* 
 
 The whole of the preceding reasoning can be summed up as 
 follows : 
 
 If in equal volumes of all gases there are equal numbers of mol- 
 ecules, then the weights of these equal volumes must bear the same 
 relationship toward each other as do the weights of the individual 
 molecules. The relative weights of the molecules are therefore to 
 be ascertained by determining the relative weights of equal vol- 
 umes of gases ; and, for obvious reasons, the weight of a molecule 
 of hydrogen has been selected as the standard ; as the molecule of 
 hydrogen consists of two atoms, this standard for measuring molec- 
 ular weights is placed at two. If the specific gravities of gases have 
 been determined with air as unity, then, in order to recalculate them 
 so as to compare them with hydrogen as two, they must be multi- 
 plied by 28.8, for : - 
 
 spec. grav. of hydrogen : spec. grav. of air : : 2 : 28.8. 
 
 The value of the discoveries just cited as an assistance in determin- 
 ing the atomic weights of elements is apparent. By a determination 
 * For apparent exceptions see ammonium chloride. 
 
SPECIFIC GRAVITIES OF GASES. 73 
 
 of the specific gravity of a gas, we ascertain the relative .weight of 
 a molecule of that gas as compared with the weight of a molecule 
 of hydrogen ; as, for instance, in the case of water, the molecular 
 weight cannot be more or less than eighteen, and in this eighteen 
 parts by weight of water, quantitative analysis shows us that we 
 have sixteen parts by weight of oxygen and two of hydrogen. The 
 atomic weight of oxygen, therefore, cannot be more than sixteen, 
 unless we wish to accept the existence of a fraction of an atom of 
 oxygen in a molecule of water. The maximum number for the 
 atomic weight of oxygen is consequently fixed by experiment. 
 That it is not some fraction of sixteen we cannot state so definitely, 
 yet all evidence points against this assumption. In the first place, 
 two volumes of hydrogen unite with one of oxygen to form water ; 
 the presumption, therefore, is that water has the formula H 2 0, and 
 hence sixteen would be the atomic weight of oxygen; and in the 
 second, we never have encountered any compound of oxygen which 
 can be obtained in a gaseous state, and whose molecular weight we 
 therefore know, which contains relatively less than sixteen parts by 
 weight of that element. The magnitudes at present in use for our 
 atomic weights are the results of reasoning exactly similar to that 
 given in the case of oxygen, assisted in many cases by deductions 
 drawn from analogies existing between a number of elements and 
 by other less important methods of determining molecular weights, 
 and we shall subsequently see that the atomic weights at present in 
 use are the only ones by means of which the elements, when arranged 
 in the order of their atomic weights, naturally fall into series and 
 families which show the greatest resemblance to one another. This 
 existence of elements as molecules is used to explain the chemical 
 activity of elements in the nascent state. (See page 51.) 
 
 Chemical equations, expressing the changes which take place 
 when simple or compound substances are brought in contact, indi- 
 cate the initial bodies and the final result by formulae based upon 
 our atomic weights, taking no account of the changes of energy ; 
 they, as a rule, represent only the main course of a reaction, while 
 other minor reactions are often going on in a mixture of two or 
 more substances. Many equations are true only for certain condi- 
 tions of temperature, concentration, etc. A chemical equation is 
 simply an algebraic expression which can be constructed quite inde- 
 pendently of chemical facts, and when so constructed is entirely 
 
74 HYDROCHLORIC ACID; THKH MO-CHEMISTRY. 
 
 useless if not pernicious in its tendency. We must always bear in 
 mind that chemical facts and experiments are infinitely more valu- 
 able than chemical equations, the latter being useful only as a short 
 method of expressing those changes which we know to take place. 
 
 In uniting to form hydrochloric acid, hydrogen and chlorine 
 give 220 K. ( See page 64.) In dissolving in water an additional 
 173 K is evolved, so that the solution of hydrochloric acid in water 
 possesses much less chemical energy than does the gas. We should 
 therefore expect hydrochloric acid gas to enter into a number of re- 
 actions where the solution would be inert. If hydrochloric acid 
 gas and oxygen are passed through a heated tube (see page 59) the 
 following reaction takes place : 
 
 = H 2 + 2 Cl 
 
 while if chlorine water is allowed to stand in the sunlight : 
 2 Cl + H 2 = 2 H Cl -f 0. 
 
 This contrast is explained by the difference in energy between 
 gaseous hydrochloric acid and the solution, for : 
 
 H + Cl = H Cl gives 220 K, 
 2H + 2C1 = 2HC1 440 K, 
 2H-f = H 2 684 K, 
 
 hence, the heat of formation of water (the measure of the chemical 
 energy of H and 0) is greater than that of two molecules of hydro- 
 chloric acid gas, and therefore the reaction 2 H Cl + = H 2 + 
 .2 Cl would be accompanied by an evolution of heat. On the other 
 ;hand, 
 
 H + Cl = H Cl, dissolved in water gives 393 K, 
 2 H + 2 Cl = 2 H Cl 786 K, 
 
 and hence the heat of formation of dissolved hydrochloric acid is 
 greater than that of water. As a result the reaction, 
 
 2 C1 + H 2 0=2 HC1 + 0, 
 
 in the presence of water is accompanied by an evolution of heat. Of 
 course, the more concentrated a solution of hydrochloric acid is, the 
 nearer will it approach the condition of the gas, and hence the 
 greater will be its reactiveness. Similar studies with other bodies 
 show us that the difference in energy between dissolved substances 
 and those undissolved is often very marked, and in considering 
 
ACIDS; NATURE OF. 75 
 
 whether certain chemical reactions will take place we should take 
 tliis difference into account. 
 
 The compound of hydrogen and chlorine is called an acid 
 because it has certain distinctive properties possessed by every 
 substance entitled to be classed as such. 
 
 An acid is a compound containing hydrogen united to a negative 
 element or group of elements, which hydrogen can be replaced by a 
 metal to form a salt. 
 
 No definition of an acid is entirely satisfactory, as we have 
 a number of substances which contain hydrogen replaceable by a 
 metal, as, for instance, water in the reaction HOH -}- K = KOH -(- H, 
 yet we scarcely would call water an acid, nor KOH a salt, although, 
 essentially, there is no difference between this reaction and the 
 following : Zn + 2 H Cl = Zn C1 2 + 2 H ; in one case, it is the neg- 
 ative group of elements OH, in the other, the negative element Cl, 
 which is united to hydrogen, and the reactions take place because 
 K or Zn has a greater chemical energy when brought in contact 
 with OH or Cl, respectively, than has hydrogen ; they being more 
 metallic in their nature than is hydrogen, and hence presenting a 
 greater contrast to the not-metal. If we call hydrochloric acid an 
 acid, and water not one, it is simply because expediency shows us 
 that it is well to classify those hydrogen compounds, the hydrogen 
 of which is easily replaced by a large number of metals to form 
 salts, under the head of acids; an indication of the propriety of the 
 name being that the substance designated as " acid " has the power 
 of turning a vegetable dye '(blue litmus) to a red color. It is evi- 
 dent that this latter distinction is purely arbitrary, and unimportant 
 as regards the true chemical nature of acids. Many substances 
 which are not acids will turn blue litmus to a red color, while, on the 
 other hand, some substances decidedly acid * have no effect what- 
 ever upon litmus. 
 
 In forming chlorides we can employ four leading methods, three 
 of which illustrate general characteristics of acids ; they are : 
 
 1st. By direct union of the elements, as, for example : 
 
 P +3C1=PC1 3 , 
 
 Zn +2 Cl = ZnCl 2 , C +4C1 
 
 Organic substances acting as acids (aceto-acetic ether). 
 
76 CHLORIDES; FORMATION OF. 
 
 This method leads to the formation of chlorides of the not- 
 metals as well as those of the metals, the process being analogous 
 to that of combustion in oxygen. 
 
 2d. By the action of hydrogen chloride on a metal, by which 
 means the chloride is produced and hydrogen liberated, as : 
 
 Zn +2HCl = ZnCl 2 + 2 H, 
 Fe -f2 HCl = FeCl 2 + 2 H, 
 Mg + 2 H Cl = Mg C1 2 + 2 H. 
 
 This reaction takes place with metals only. 
 
 3d. By the action of hydrogen chloride on the oxides of metals, 
 when the chloride and water are produced, as : 
 
 ZnO +2HCl = ZnCl 2 +H 2 0, 
 CaO + 2HCl = CaCi 2 +H 2 0, 
 MgO +2 HC1 = MgCl 2 + H 2 0. 
 
 4th. By the action of hydrogen chloride on the hydroxides of the 
 metals, as : 
 
 OH HC1 Cl HOH, 
 
 Zn( + =Zn( + 
 
 X OH HC1 X C1 HOH, 
 
 Ca(OH) 2 +.2HCl =CaCl 2 +2H 2 0, 
 
 Mg(OH) 2 + 2HCl =MgCl 2 +2H 2 0. 
 
 When an oxide or hydroxide reacts in the above manner it is 
 the oxide or hydroxide of a metal, and is designated as a base ; 
 while the chemical process of forming a salt by addition of an acid 
 to a base or base to an acid is called neutralization (the acid or base 
 is neutralized). This term " base " is one dictated by expediency ; 
 and when we speak of a substance as basic in character, we mean it 
 can form an oxide or a hydroxide which will neutralize acids. 
 
 The reactions under 3 and 4 are general to all acids and bases ; 
 if we designate any metal by M, any acid by HX, then the general 
 rule will be that : 
 
 M 2 +2HX=2MX + H 2 0; M OH + HX=MX + H 2 0; 
 MO +2HX= MX 2 + H 2 0; M (OH) 2 + 2HX=MX 2 +2H 2 0; 
 M 2 3 -f 6HX=2MX 3 + 3H 2 0; M(OH) 3 -f 3HX=MX 3 +3H 2 0. 
 
NEUTRALIZATION. 77 
 
 Thus : - 
 
 K 2 +2HN0 3 = 2KN0 8 + H 2 0; 
 
 Zn + 2HBr =ZnBr 2 + H 2 0; 
 A1 2 3 + 6HNO. = 2Al(NO,),+3H a O; 
 
 KOH + HN0 3 = KN0 3 + H 2 0; 
 
 Zn(OH) 2 +2HM) 3 = Zn(N0 3 ) 2 +2H 2 0; 
 
 Al(OH), + 3Htf0 8 = Al(N0 8 ) 8 + 3H 2 0. 
 
 The metals differ among each other in their power of replacing 
 hydrogen in acids to form salts ; some replace one atom, some two, 
 some three, to form one molecule of the salt; but whatever the 
 acid, this number is always the same for any given metal, and it 
 can be ascertained from the formula of the chloride of the metal 
 (as KC1, ZnCl 2 , A1C1 3 ) ; the metal will replace as many hydrogen 
 atoms in any acid as there are chlorine atoms in the formula weight 
 of its chloride, and the hydroxide will contain as many hydroxyl 
 groups (OH ) as there are chlorine atoms in the chloride.* 
 
 The reactions under 2 take place between a number of metals 
 and acids, but the applications are much less general than 3 and 4, 
 and often, indeed, where the acid contains oxygen, no hydrogen is 
 evolved, but some other substance is produced from the acid by the 
 action of hydrogen in the nascent state. 
 
 * See pages 30, 32, and 43. 
 
78 BROMINE; OCCURRENCE, PREPARATION. 
 
 CHAPTER X. 
 
 BROMINE AND HYDROBROMIC ACID. 
 
 *. 
 
 Symbol, Br ; atomic weight, 79.95 ; specific gravity of fluid, 3.187 at 
 
 0; specific gravity of gas, below 900, tm- = l, is 5.54, H 2 = 2, 
 is 159.5. Formula, HBr; specific gravity, air = 1, is 2.81, 
 H 2 = 2,ts 80.95; 1 c.c. a* a^ .76 m. pressure weighs .00365 
 
 IN many respects bromine resembles chlorine ; indeed, such 
 modifications in chemical characteristics as it represents are due 
 simply to its larger atomic weight. It has the same tendency to 
 unite with metals to form salts, and hence is not free in nature, but 
 is always found combined with metals in the form of bromides ; its 
 compounds generally accompany those of chlorine, yet they are 
 always present in lesser quantity, bromine being one of the rarer 
 elements. The bromides of sodium and magnesium are found in 
 the great majority of salt springs, especially in those of Saratoga 
 Springs, in the Saginaw Valley, and in the southeastern portion of 
 Ohio, where the bromide of potassium also occurs. In Europe the 
 brines from the salt works of Kreuznach and Strassfurth are 
 especially rich in bromides. Marine fauna and flora also fre- 
 quently contain bromides. The brines of the various salt works 
 are evaporated, thus crystallizing the sodium chloride for the man- 
 ufacture of table salt, there remaining a not-crystallizable brine 
 (mother liquor) which is especially rich in bromides. In this 
 mother liquor Ballard discovered bromine in 1826, calling it 
 bromine from fip&fjLos, a stench. 
 
 Bromine is prepared from its compounds in a manner entirely 
 analogous to the method used in isolating chlorine. A bromide is 
 mixed with manganese dioxide and sulphuric acid (see pages 59, 60), 
 when the following reaction takes place : 83 
 
 H 2 S0 4 -f Mn 2 = Mn S0 4 + Na 2 S0 4 + 2 Br + 2 H 2 0. 
 
 Hydrobromic acid and manganese dioxide would yield bromine 
 
BROMINE; PROPERTIES. 79 
 
 just as hydrochloric acid and manganese dioxide give us chlo- 
 rine : 
 
 4 H Cl + Mn 2 = Mn C1 2 + 2 Cl + 2 H 2 
 
 but owing to the difficulty of preparing hydrobromic acid, the latter 
 substance is an expensive article, so that, from reasons of economy, 
 this method is not available. 
 
 Bromine is a dark brown liquid, almost black when any consid- 
 erable thickness is observed; it melts at 7. 3 and boils at 63. 05, 
 a point considerably below the boiling point of water. , The specific 
 gravity of the liquid at is 3.18. 'When the liquid is allowed to 
 stand in the air it evaporates rapidly, even at ordinary tem- 
 peratures, yielding reddish brown vapors very irritating to the 
 mucous membrane of the eyes, nose, and throat, and possessing an 
 odor much resembling that of chlorine. The specific gravity of the 
 vapor, air being 1, at 800, is 5.54, giving 159.5 as its density, H 
 = 2. This shows that at this temperature the molecule of bromine 
 is composed of two atoms like that of chlorine, for the atomic 
 weight is 80, so that 160 would be the molecular weight of Br 2 . 
 About 1200 the specific gravity of the vapor decreases until it 
 .becomes 3.7, showing that at high temperatures some of the Br 2 
 molecules have changed to individual atoms. 
 
 Bromine is soluble in water, the solution having a brownish 
 color and properties similar to those of chlorine water, one part of 
 bromine at 15 being soluble in 33 parts of water. If this solution 
 is cooled to the freezing point of water, crystals of a compound of 
 bromine with water of crystallization ( 2 Br + 10 H 2 ) separate. 
 (See page 62.) 
 
 The solution of bromine in water is an oxidizing agent, and hence 
 bleaches just as chlorine water does ; this property is due to the 
 same cause, the liberation of oxygen. 
 
 The formation of oxygen becomes apparent if a tube containing 
 bromine water is inverted over a water trough and exposed to the 
 sunlight ; oxygen separates as it does in the case of chlorine water, 
 although not with such great rapidity.* 
 
 * Owing to the greater ease with which bromine is handled, it is more fre- 
 quently in use as a laboratory oxidizing agent than is chlorine. 
 
80 HYDROBROMIC ACID; PROPERTIES. 
 
 The compounds of bromine resemble those of chlorine in every 
 particular, and the bromides and chlorides of the same metal are 
 isornorphous.* 
 
 Bromine does not unite with hydrogen as readily as does chlo- 
 rine, its higher atomic weight rendering its chemical character less 
 negative, and hence its tendency to unite with metals less pro- 
 nounced ; as a consequence, a mixture of bromine and hydrogen can 
 be allowed to stand in the sunlight for any length of time without 
 the formation of hydrobromic acid ; the union is only to be brought 
 about by more energetic means, such as the electric spark, or the 
 passing of a mixture of hydrogen and bromine over platinized 
 asbestos. 
 
 The heat of formation of hydrobromic acid gas f is only 121 K, 
 while that of gaseous hydrochloric acid is 220 K, so that we should 
 expect hydrobromic acid to be more easily decomposed than is 
 hydrochloric acid. The consequences of this instability are un- 
 pleasantly apparent in the difficulties encountered in the prepara- 
 tion of hydrobromic acid. 
 
 In preparing hydrochloric acid we had but to treat sodium chlo- 
 ride with sulphuric acid, as follows : 
 
 and a similar reaction takes place when a bromide is substituted 
 for a chloride : ^^ C^x>JM( v> ^> 
 
 2 ]STa Br + H 2 S0 4 = Na a S0 4 + 2 H Br ; 
 
 but hydrobromic acid (being so much less stable than hydrochloric), 
 owing to the heat of the reaction, breaks down into hydrogen and 
 bromine, after which decomposition the nascent hydrogen attacks 
 the sulphuric acid, reducing the latter to form sulphur dioxide and 
 water : - 1. 2 H Br = 2 H -f- 2 Br 
 
 2. H 2 S0 4 + 2H = 2H 2 + S0 2 ; 
 
 and, as a consequence, the hydrobromic acid produced by this 
 means is contaminated with sulphur dioxide. 
 
 Preparation of hydrobromic acid. 
 
 In order to prepare hydrobromic acid for laboratory use, advan- 
 tage is taken of the instability of the halogen compounds of the 
 not-metals. 
 
 * See page 42 and foot-note. t Using gaseous bromine. 
 
HYDROBROMIC ACID ; PKOPERTIES. 81 
 
 When phosphorus trichloride is added to water, the following 
 change takes place : 
 
 /jCl + HjOH /OH 
 P JC1 + H;0 H = P OH + 3HC1. 
 \C1 + HJOH \OH 
 
 Phosphorus trichloride + water = phosphorus hydroxide (phosphorous acid) 
 + hydrochloric acid. 
 
 P Cl, + 3 H 2 = P (OH ) 3 + 3 H Cl. 
 The same with phosphorus tribromide : 
 
 P Br 3 + 3 H 2 = P (OH ) 8 -f- 3 H Br. 34 
 
 In performing this operation it is not necessary to employ the 
 finished tribromide of phosphorus, for a mixture of phosphorus, 
 bromine, and water will answer the same purpose. 
 
 Hydrobromic acid is a colorless gas, with an acid odor resem- 
 bling that of hydrochloric acid ; it fumes strongly in the air, owing 
 to the absorption of moisture ; it is extremely soluble in water, one 
 part of that substance absorbing as much as 82 per cent of hydro- 
 bromic acid, and, as a consequence, the gas cannot be collected over 
 water. If a quantity is desired, it must either be separated by col- 
 lecting over mercury, or by the displacement of air in some vessel, 
 for, as the specific gravity of hydrobromic acid is 2.79, it can be 
 poured downward in the atmosphere. 
 
 As has already been stated, hydrobromic acid is much less 
 stable than is hydrochloric acid* ; and therefore, on adding chlorine 
 to the former, bromine is separated and hydrochloric acid is 
 
 In its chemical behavior, hydrobromic acid is like hydrochloric 
 acid. When brought in contact with bases, it 'neutralizes them to 
 form salts. Na QH + H Br = Na Br + HOH 
 KOH + HBr = KBr + HOH 
 CaO +2HBr = CaBr 2 +H 2 0. 
 
 * Hydrobromic acid is partly broken down into its constituents if it is ex- 
 posed to the sunlight; the same is true, in a much smaller degree, of hydro- 
 chloric acid, if it is in concentrated solution in the presence of oxygen. (A. 
 PJchardson; Journal Chem. Soc. ; 51, 801.) 
 
\ 
 
 82 BROMIDES ; FORMATION. 
 
 The methods of formation of the bromides are analogous to 
 those of the chlorides. They are : 
 
 1. Direct union of the element in question with bromine, this 
 applying to metal or not-metal. 
 
 2. The addition of a metal to hydrobromic acid, when the 
 bromide is formed and hydrogen liberated. 
 
 3. The action of hydrobromic acid on the oxides or hydroxides 
 of the-metals, when the bromides and water are produced. 
 
 Hydrobromic acid is formed of one volume of hydrogen and one 
 volume of bromine vapor, united to form 2 volumes of hydrobromic 
 acid ; the same considerations advanced under chlorine, show that 
 the bromine molecules, provided the temperature be not too high, 
 consist of two atoms to the molecule. (See pages 69, 70, 71.) 
 
IODINE; OCCURRENCE, HISTORY. 83 
 
 CHAPTER XL 
 
 IODINE AND HYDROIODIC ACID. 
 
 Symbol, I ; atomic weight, 126.85 ; specific gravity, 4.948 ; specific 
 weight of vapor, air = 1, is 8.84 (below 600), H 2 = 2, is 255.21. 
 Formula, HI ; specific weight, air = 1, is 4.44, H 2 = 2, is 127.9. 
 1 c.c. of the gas, and .76m., weighs .005767 gram. 
 
 (THIS last member of the halogen family is not found, as such, in 
 nature ; in that way it resembles fluorine, chlorine and bromine ; its 
 compounds, although they occur in company with those of chlorine 
 and bromine in almost all deposits in which the halides of the 
 metals are found, are not by any means present in such large quan- 
 tities. The element, in combination, occurs in sea water, both as the 
 iodide of sodium and of magnesium, although the quantities of these 
 salts are so small that their presence cannot be proved, excepting 
 by some special means. Sea plants, such as the algae, as well as 
 representatives of the animal kingdom (sea sponges, crabs, oysters, 
 etc.), can assimilate and concentrate traces of iodides so that, on 
 drying and burning, iodine can easily be proved to be present in 
 their ashes. The iodides occur in salt springs, in deposits of rock 
 salt, in a number of mineral springs, as at Kreuznach and Reichen- 
 hall, in river water, and also in some water plants growing in 
 flowing fresh water. 
 
 The element was discovered in 1811 by Courtois, a saltpetre 
 manufacturer in Paris, who found its compounds in the mother- 
 liquors, left after extracting the ashes of sea plants and crystalliz- 
 ing the less soluble portions. The name iodine is taken from 
 iwSr/s, violet. The weed which is washed up by the spring storms 
 on the islands on the western coast of Scotland or of Ireland, or on 
 the coast of Normandy, or that which grows upon the rocks, is dried 
 and burned, the fused mass remaining as the ash is brought into 
 the market under the Scotch name of kelp, or Normanic, varec. 
 The amount of iodine contained in this ash is very small, and is ex- 
 tracted from the last remaining mother-liquors, obtained by crystal- 
 
84 IODINE; PREPARATION, PROPERTIES. 
 
 lizing the aqueous extract of kelp. (The method of preparation of 
 iodine is identical with that of chlorine or bromine ; 35 the iodide is 
 treated with sulphuric acid and manganese dioxide : 
 
 rf 
 
 Mn O 2 + 2 KI + 2H 2 S0 4 = MnS0 4 + K 2 S0 4 + 2 H 2 + 2 1. 
 
 In order to purify the iodine it is sublimed, the iodine being heated 
 in retorts and collected as crystals in cold chambers. 
 
 v The element is almost black, grayish solid, with a lustre closely 
 resembling that of the metals ; when pure and fused it is entirely 
 black ; in thin plates it is translucent with a brownish-red color) 
 Its specific weight is 4.94 ; it melts at 113 115 and boils at about 
 200 ; * when heated in a vacuum it does not melt, but vaporizes 
 without fusion. The vapor of iodine has a beautiful violet color f 
 and a specific gravity, air being 1, of 8.84, below 600, giving a 
 density, H 2 = 2, of 255.21, showing that below this temperature 
 iodine has a molecule consisting of two atoms, I 2 ; but if the heat is 
 gradually increased, the specific gravity of the vapor diminishes, so 
 that at 1570 it is only (air = 1) 5.67, or (H 2 = 2) 163.69, indicat- 
 ing that the vapor of iodine at bright red heat consists of a mixture 
 of the individual atoms, and the molecules I 2 , dissociation beginning 
 above 600.$ Iodine is but very little soluble in water, about 7000 
 parts of water dissolving one part of iodine ; water containing salts 
 in solution has a greater solvent action; some solutions, such as 
 those of potassium iodide and of hydroiodic acid, have the power of 
 dissolving large quantities of iodine ; the element is also extremely 
 soluble in substances such as alcohol, ether, carbon bisulphide, or 
 chloroform. 
 
 i Chemically, iodine resembles chlorine or bromine ; it unites with 
 sulphur, phosphorus, and other not-metals with which the other 
 lialogens form compounds ; in combination with the metals it forms 
 iodides ; and if hydrogen, mixed with iodine vapors, is passed over 
 
 * Stas; Gesetze der Chemischen Proportionen ; 141. 
 
 t Observe the same by throwing some iodine on a hot stove or into a hot 
 porcelain crucible. When the vapor of iodine is saturated, it is pure blue in 
 color. 
 
 \ At 1600-1700, H. Biltz and Y. Meyer find the vapor density of iodine to 
 correspond with a molecular weight which indicates that at the temperature 
 of the experiment the molecules I 2 have completely broken down into the 
 individual atoms. (Berichte derDeutsch. Chem. Gesell. 22; 726.) 
 
HYDROIODIC ACID ; PREPARATION. 85 
 
 spongy platinum which is heated, some hyclroiodic acid is formed, 
 although hydroiodic acid is an endothermic compound, and there- 
 fore possesses more energy than the individual elements of which 
 it is composed. } 
 
 When the preparation of hydrobromic acid was mentioned, we 
 saw that the method available for hydrochloric acid was not feasi- 
 ble, owing to the relative instability of hydrobromic acid ; the acid 
 breaking down into hydrogen and bromine, while the hydrogen re- 
 duces the sulphuric acid to sulphurous acid ; as a consequence we 
 are compelled to use a reaction which depends on the decomposition 
 of the bromide of phosphorus by water. We could not, therefore, 
 expect to obtain hydroiodic acid by the action of sulphuric acid on 
 sodium iodide, for hydroiodic acid decomposes even more readily 
 than does hydrobromic ; in this case the reduction of sulphuric acid 
 takes place so energetically that sulphuretted hydrogen (H 2 S ) is 
 produced. If we write the equations representing the action of 
 hydrobromic and hydroiodic acid upon sulphuric acid, this relative 
 instability at once becomes apparent : 
 
 H 2 S0 4 + 2 H Br = 2 H 2 + S0 a + 2 Br, 
 H 2 S0 4 + 8HI =H 2 S + 4H 2 0+8I. 
 
 for in the latter case the decomposition" is much more energetic and 
 far-reaching.* 
 
 ; In the preparation of hydroiodic acid, therefore, (we are com- 
 pelled to resort to a round-about method similar to that employed 
 in the production of hydrobromic acid ; the reaction which is used 
 depends on the instability of the iodide of phosphorus in the pres- 
 cnceofwate, 
 
 * This reaction illustrates quite well the influence exerted by the mass of 
 the chemical reagents used. If a large excess of sulphuric acid and but little 
 hydroiodic are present, a considerable quantity of sulphur dioxide (SO 2 ) is pro- 
 duced, while generally more or less sulphur is deposited. These changes are 
 represented by the equations : 
 
 H 2 S0 4 + 6HI = 4H 2 Q + S +61. 
 
 H 2 SO 4 + 2 HI = 2 H., O + SO 2 + 21. 
 
 Under ordinary conditions, in a test tube, all of these changes take place at the 
 same time, so that the equation given in the text only represents one of the 
 various changes going on; a variation of temperature or mass of the reagents 
 can alter the proportions of hydrogen sulphide, sulphur, or sulphur dioxide 
 produced, 
 
86 HYDROIODIC ACID; PROPERTIES. 
 
 Of course, a mixture of iodine, water, and red phosphorus 
 
 answers the purpose. 36 (Hydroiodic acid can be collected in empty 
 
 "jars by displacement of air, or it can be collected over mercury ; its 
 
 extreme solubility in water renders the filling of vessels with the 
 
 gas impossible where water is present.) 
 
 Hydroiodic acid is a colorless gas, with the acid odor of hydro- 
 chloric or hydrobromic acid ; it fumes in the air, owing to absorption 
 of moisture/ Its specific weight (air being 1), is 4.44 ; H 2 being 2, it 
 is 127.9 ; the molecular weight of HI is 127.85, the slight discrep- 
 ancy between the observed specific gravity and the molecular weight 
 being due, undoubtedly, to the difficulties encountered in obtaining 
 pure hydroiodic acid. Hydroiodic acid is quite easily condensed to 
 a liquid ; at 17. 8 c. its vapor tension is but two atmospheres ; 
 it solidifies at 51 c. On heating, hydroiodic acid is easily de- 
 composed, the change into hydrogen and iodine beginning at 100, 
 " and being complete at 700.* (Chlorine decomposes hydroiodic 
 acid with almost explosive violence, forming hydrochloric acid and 
 iodine.fj * 
 
 Hydroiodic acid is composed of equal volumes of hydrogen and 
 iodine ; two volumes of hydroiodic acid yield one of hydrogen and 
 one of iodine, provided the temperature be kept sufficiently high to 
 vaporize the iodine formed. The gas is extremely soluble in water ; 
 the solution, when saturated at 12, contains 57.7 percent of hydro- 
 iodic acid ; it must be kept well stoppered and in the dark ; for, 
 when exposed to the air, a separation of iodine takes place, owing 
 to the formation of water, as follows : 
 
 ( Hydroiodic acid has a strongly acid reaction ; it turns blue lit- 
 mus solution red, and in all respects resembles hydrochloric and 
 hydrobromic acids, the methods of formation of the iodides being 
 exactly like those of bromides and chlorides. The heat of forma- 
 tion of gaseous hydroiodic acid from hydrogen and solid iodine is 
 
 *^Bodenstein ; Berichte d. Deutsch. Chem. Gesell. ; 26, 2603. 
 
 t The iodine first separates as the element, and this is then converted into 
 the trichloride of iodine, a yellow solid. As a consequence, the violet vapors 
 which at first appear when chlorine is added to hydroiodic acid subsequently 
 disappear, while the trichloride of iodine settles on the sides of the vessel in 
 which the reaction takes place. 
 
HALOGENS ; COMPARATIVE TABLE OF. 
 
 87 
 
 61 K ; from gaseous iodine and hydrogen, almost nil ; the heat of 
 solution is 192 K, so that the heat of formation of hydroiodic acid 
 in water is 131 K. The gaseous acid is therefore an endo thermic 
 compound, possessing a tendency to decompose into its constituent 
 parts ; the solution possesses much less chemical energy, and hence 
 a greater stability. In the following table a comparison of the 
 properties of the halogens has been undertaken, their differences 
 with increasing atomic weight thus being made more apparent : 
 
 THE HALOGENS. 
 
 
 F. 
 
 Cl. 
 
 Br. 
 
 I. 
 
 
 Pale yellow 
 
 Greenish yel- 
 
 Dark 
 
 Black 
 
 
 gas. 
 
 low gas. 
 
 brown 
 
 solid. 
 
 
 
 
 liquid. 
 
 
 Density of liquid. 
 
 
 
 1.33 
 
 3.18 
 
 4.97 
 
 Density < Air = 1 
 
 1.26 
 
 2.46 
 
 5.54 
 
 8.84 
 
 of vapor. ( H 2 = 2 
 
 36.28 
 
 70.84 
 
 159.5 
 
 255.21 
 
 Molecule of gaseous 
 
 
 
 
 
 element. 
 
 F 2 
 
 C1 2 
 
 Br 2 
 
 I* 
 
 The density of iodine vapor at 447 is 8.8 ; at 1570 the density 
 is 5.67, which yields 163.69 as the density, H 2 = 2.* Dissociation 
 is therefore far advanced at this temperature, so that at a higher 
 one the molecule and atom of iodine nearly correspond. 
 
 HF. HC1. HBr. HI. 
 
 Stability. 
 
 J=== 
 
 Heat of for- 
 
 
 
 
 
 mation. 
 
 
 
 H, Cl = 220 K. 
 
 H,Br=121K. 
 
 H, I = - 61 K. 
 
 Ditto, plus 
 
 
 
 
 
 water. 
 
 
 
 393 K. 
 
 320 K. 
 
 131 K. 
 
 * A determination by V. Meyer (Berichte d. Deutsch, Chem. Gesell. ; 13, 
 1310), taken at 1437 gives a density of 4.76 for the vapor of iodine; this, 
 H 2 = 2, would give a specific gravity of 137.4, or a molecular weight which very 
 nearly corresponds to the atomic weight of iodine. See also foot-note page 
 84. This refers to the last work on the subject. 
 
88 HALOGENS; COMPARATIVE TABLE OF. 
 
 Chlorine replaces bromine when brought in contact with bromides, and 
 iodine when brought in contact with iodides. 
 Bromine replaces iodine in the iodides. 
 Fluorine liberates oxygen from water, even in the dark. 
 
 2F + H 2 0=2HF + 0. 
 
 Chlorine liberates oxygen from w T ater in the sunlight. 
 2C1 + H 2 O=2HC1 + O. 
 
 Bromine liberates oxygen from water in the sunlight more slowly than 
 chlorine. 
 
 Iodine does not decompose water. 
 
THE OXYGEN FAMILY. 89 
 
 CHAPTER XII. 
 
 THE OXYGEN FAMILY. 
 
 THREE elements sulphur, selenium, and tellurium show a 
 relationship toward oxygen similar to that displayed by chlorine, 
 bromine, and iodine toward fluorine. With increasing atomic 
 weights we have similar changes in the physical properties of the 
 elements, illustrating the decreasing not-metallic characteristics of 
 the family ; oxygen is a colorless gas, sulphur a yellow solid, sele- 
 nium a dark red solid (in one of its allotropic forms), while tellurium 
 is an element having entirely the appearance of a metal. All of 
 the elements of this family form hydrogen compounds, which, with 
 the exception of water, are colorless gases at ordinary temperatures ; 
 in this they resemble the halogens, for in that family the hydrogen 
 compound of the element with the smallest atomic weight (fluorine) 
 is liquid at ordinary temperatures. The compounds with hydrogen, 
 which are produced by the members of the oxygen family, have 
 one atom of the element united to two of hydrogen ; the atoms of 
 the elements under consideration, therefore, display twice as great 
 a power of retaining hydrogen atoms in close proximity as the 
 atoms of the halogens do ; the formulae of these hydrogen compounds 
 are H 2 0, H 2 S, H 2 Se, H 2 Te; their stability and heat of formation 
 diminish with increasing atomic weight, accompanying the decreas- 
 ing not-metallic properties of the elements. This deportment is 
 exactly parallel to the similar changes with the members of the hal- 
 ogen family, so that hydrogen selenide decomposes at about the 
 same temperature as hydroiodic acid (150), while hydrogen tellu- 
 ride is not known in a pure state, and decomposes even at ordinary 
 temperatures if placed in contact with the air. The hydrogen 
 compounds of this family are by no means as acid in their proper- 
 ties as are those of the halogens,* indeed, water does not redden 
 litmus, while hydrogen sulphide, selenide, and telluride do not turn 
 
 * The acid properties of the hydrogen compounds of the not-metals dimin- 
 ish with an increase in the number of hydrogen atoms in the molecule. 
 
90 THE OXYGEN FAMILY. 
 
 blue litmus into a pronounced red. The hydrogen compounds of 
 sulphur, selenium, and tellurium are much less soluble in water 
 than hydrochloric, hydrobromic, and hydroiodic acids are. 
 
 In comparing hydrofluoric acid, water, and ammonia (H 3 N), we 
 see that hydrofluoric acid, as its name implies, is an acid ; water 
 acts like an acid only when it is brought in contact with the most 
 pronounced metals, such as sodium or potassium (sometimes it can 
 act as a base), while ammonia (H 8 K) is basic in its character. This 
 gradation of properties is possibly due to the successively increasing 
 predominance of the positive element in the three substances, for it 
 is obvious that the three hydrogen atoms in a molecule of ammonia 
 will have a much greater effect in determining the character of the 
 compound than the one present in a molecule of hydrofluoric acid ; 
 hence H 3 ISf as an entire compound partakes of the nature of a metal 
 and is positive. As the negative element in the hydrogen com- 
 pounds increases in atomic weight, its larger mass may cause its 
 character to become more prominent ; hence H 3 Sb is not basic, an- 
 timony having next to the highest atomic weight in the family of 
 which nitrogen is a member. 
 
 The study of the hydrogen compounds of the elements of the 
 oxygen family is less completely finished than is that of the corre- 
 sponding compounds of the halogens, because hydrogen selenide and 
 hydrogen telluride have rarely been prepared in a pure state. 
 
 The elements of this family, with the possible exception of tel- 
 lurium, exist in two allotropic forms. They form compounds with 
 oxygen which are anhydrides of acids ; the study of these will be 
 deferred until the oxygen compounds of the not-metals are consid- 
 ered. The compounds of sulphur, selenium, and tellurium with 
 the metals have formulae corresponding to those of the oxides, 
 thus : 
 
 Na 2 Na 2 S Na 2 Se Na 2 Te 
 
 Sodium oxide ; Sodium sulphide ; Sodium selenide; Sodium telluride. 
 
 FeO FeS FeSe FeTe 
 
 Ferrous oxide ; Ferrous sulphide ; Ferrous selenide ; Ferrous telluride. 
 
 Oxygen has already been discussed, so that the study of this 
 family will be continued with sulphur. 
 
SULPHUR; OCCURRENCE, HISTORY. 91 
 
 CHAPTER XIII. 
 
 SULPHUR. 
 
 Symbol, S ; atomic weight, 32.06 ; specific gravity of solid, 2.045 ; 
 specific gravity of vapor (above 1000), air = 1, is 2.2, H 2 = 2, 
 is 63.36. 
 
 SULPHUR is found, often in company with gypsum or limestone, 
 in volcanic regions such as those of Southern Italy and of Sicily. 
 In Europe the largest quantities come from the provinces of Gir- 
 genti and Catania in Sicily ; the crater of the volcano Purace in 
 South America, with a surface of about 1200 yards square, is cov- 
 ered with a layer of sulphur more than a yard in thickness ; while 
 in our own country considerable quantities of sulphur are found in 
 California. The element, when combined, is chiefly found in the 
 form of sulphides, the most important of which are given in the 
 following table : 
 
 Iron sulphides : Iron pyrites and markasite, Fe S 2 . 
 
 Lead sulphide : Galenite (galena), PbS. 
 
 Zinc Sulphide: Zinc blende, ZnS. 
 
 Copper and iron sulphide: Chalcopyrite (CuFe) S 2 . 
 
 Sulphur, united with oxygen and a metal, occurs in the sulphates 
 of calcium, CaS0 4 , barium, BaS0 4 , magnesium, MgS0 4 , etc. ; the 
 most important of these is gypsum, Ca S0 4 + 2 H 2 0. Many organic 
 compounds (such as mustard oils) contain sulphur. 
 
 Sulphur has been known from very ancient times, having been 
 used as a medicine by the Greeks and Romans. The alchemists 
 considered it an essential portion of all combustible substances, 
 while during the period in which credit was given to the phlogiston 
 theory, it was looked upon as a compound of sulphuric acid with 
 phlogiston. After our present chemical theories were initiated by 
 Lavoisier's studies on oxidation, sulphur was classed as an element. 
 
 Of course, as sulphur is found uncombined, any method for pre- 
 paring the element from its compounds is of no essential value in 
 the laboratory. Sulphur might be furnished, just as is chlorine, 
 
92 SULPHUR; PREPARATION. 
 
 bromine, or iodine, by treating its hydrogen compound with some 
 oxidizing agent ; for instance, by mixing a sulphide with manganese 
 dioxide and adding sulphuric acid (a process similar to that used in 
 the preparation of the halogens) ; but such a method would be of 
 no practical value, and would only be of use because it shows the 
 similarity of action between the hydrogen compounds of the ele- 
 ments of the family under discussion and those of chlorine, bromine, 
 and iodine. An interesting method for the preparation of sulphur, 
 which is important because it explains the occurrence of the ele- 
 ment in the neighborhood of volcanoes, is by the action of sulphur 
 dioxide on sulphuretted hydrogen. 37 The reaction takes place as 
 follows: S Q 2 + 2H 2 S = 3S + 2H 2 0* 
 
 Sulphur dioxide + Sulphuretted hydrogen = Sulphur + Water. 
 Both sulphur dioxide and sulphuretted hydrogen occur in the gases 
 exhaled from volcanoes. 
 
 If iron pyrites (FeS 2 ), is heated, the following reaction takes 
 place : 3 Fe S 2 = Fe 3 S 4 -J- 2 S,t so that, probably, by reason of this 
 fact there exists another source of natural sulphur. The heat 
 necessary to decompose iron pyrites might have been furnished 
 either directly by subterranean fires or by some process of oxida- 
 tion going on throughout the entire mass of sulphide. 
 
 In order to prepare commercial sulphur from the rock in which 
 the element occurs, the latter is broken into small pieces and placed 
 in heaps upon the side of a hill ; the heaps are then covered with 
 earth, and the sulphur lighted ; the heat from the burning portion 
 is sufficient to melt the remainder, which is allowed to run into a 
 receiver. This method, in vogue in the sulphur regions of Sicily, 
 necessarily involves considerable waste, so that a number of im- 
 provements, a description of which belongs in a more extended and 
 technical work, have been introduced. The crude sulphur is puri- 
 fied by distillation. The sulphur of commerce occurs in two forms, 
 in sticks or lumps, and as a yellow powder called flowers of sulphur. 
 The latter is formed by rapidly cooling the sulphur vapors upon the 
 walls of the receiver during the process of distillation. 
 
 * Compare this reaction with the one for the preparation of chlorine by 
 means of an oxidizing agent. 
 
 t Compare this reaction with the one for the preparation of oxygen by 
 heating manganese dioxide : 
 
 3 Mil O 2 = Mn 3 O 4 + 2 O. 
 
SULPHUR; PROPERTIES. 93 
 
 Sulphur in its ordinary form is a solid of a light yellow color. 
 When heated it melts at 114, forming an amber colored liquid ; on 
 increasing the temperature to 150 the molten sulphur darkens in 
 color and becomes viscid, so that at 200 it is nearly black, and so 
 thick that it can no longer be poured ; at about 340 it once more 
 becomes fluid, but remains of a dark color. Its boiling point lies at 
 440 centigrade, at which temperature it changes to a dark brown 
 vapor. 38 The specific gravity of this vapor, air being 1, at about 470, 
 is 7.8 ; but upon increasing the temperature this gradually dimin- 
 ishes until 1000 is reached, when the density is 2.3. At the latter 
 temperature the molecule of sulphur resembles that of oxygen, 
 hydrogen, and chlorine, for it consists of two atoms.* 
 
 Sulphur is found in two allotropic forms ; one soluble in carbon 
 bisulphide, the other insoluble. The soluble variety can be par- 
 tially changed into the insoluble one by melting or by exposure 
 to the sunlight. Both forms occur in the sulphur of commerce, 
 flowers of sulphur being especially rich in the insoluble variety. 
 
 Sulphur can exist in four crystalline forms, only two of which 
 are easily obtainable and of importance. f When de- 
 posited from its solutions in carbon bisulphide, it 
 crystallizes in octahedra belonging to the rhombic 
 system (Fig. 7), but if sulphur is melted and allowed 
 to cool slowly, it forms long prismatic needles (mono- 
 clinic system), as in Fig. S. 39 The latter form changes 
 into the former on standing ; the needles becoming 
 opaque, and, while apparently retaining their crystal- 
 line form, finally consist only 
 of an aggregation of rhombic 
 crystals. During this process 
 Fig - 8 - heat is evolved, and, therefore, 
 
 * It is generally considered that the molecule of sulphur consists of six 
 atoms at temperatures just above the boiling point; but the recent investiga- 
 tions of Biltz seem to show that sulphur has no definite vapor density below 
 1000. Molecules of greater complexity than S 2 may exist; these gradually, 
 with increasing temperature, decompose into simpler ones. The specific gravity 
 of 7.8 indicates (hydrogen being 2) a specific gravity of 224.6, which corre- 
 sponds to a molecule S 7 . Compare H. Biltz ; Zeitschrif t f iir Phys. Chemie ; 3, 228. 
 
 t A substance which exists in two crystalline forms is said to be dimor- 
 phous; one existing in three, trimorphous; one existing in many, polymor- 
 phous. For a description of the four crystalline modifications of sulphur, see 
 W. Muthmann; Zeitschrif t fur Krystallographie; 17, 336. 
 
94 SULPHUR; CHEMICAL BEHAVIOR OF. 
 
 the rhombic, stable variety of sulphur possesses less energy than 
 the other. 
 
 Sulphur, when heated to just below its boiling point and subse- 
 quently rapidly cooled by pouring into cold water, forms a plastic 
 mass somewhat resembling india rubber ; this gradually becomes 
 hard and changes to yellow sulphur. The soft sulphur is dark in 
 color, probably by reason of impurities ; it contains both soluble 
 and insoluble sulphur j sulphur crystals are sometimes entirely 
 soluble.* 
 
 In its chemical behavior, in many respects, sulphur closely re- 
 sembles oxygen; the sulphides, in formula, are parallel to the 
 oxides, as the following examples demonstrate : 
 
 OXIDES. SULPHIDES. 
 
 Na 2 0, Na 2 S, 
 
 K 2 0, K 2 S, 
 
 Fe 0, Fe S, 
 
 Fe 3 4 , Fe 3 S 4 , 
 
 -^ 2 ^35 -*2 ^s, 
 
 P 2 5 , P 2 S 5 . 
 
 Sulphur combines with the majority of the elements ; and the sul- 
 phides can be prepared by direct union of the elements, just as the 
 oxides can. As further examples of the resemblance between sul- 
 phur and oxygen, the formulae of the sulphhydrates, which corre- 
 spond to those of the hydroxides, are most interesting. 
 
 HYDKOXIDES. SULPHHYDRATES. 
 
 NaOH, NaSH, 
 
 KOH, KSH. 
 
 The sulphides and sulphhydrates of the metals resemble the bases ; 
 with acids they form salts and sulphuretted hydrogen, just as the 
 bases form salts and water, thus : 
 
 (NaSH+ HC1= 
 HC1= 
 
 * Plastic sulphur is often formed when sulphur is separated from its com- 
 pounds by chemical means, as, for instance, in the oxidation of many of the 
 sulphides of metals with nitric acid. 
 
 t Many sulphides are not attacked by acids under ordinary circumstances, 
 
SULPHIDES AS ACIDIC ANHYDRIDES. 95 
 
 A number of the sulphides of the not-metals resemble the anhy- 
 drides of acids, for they form salts with the sulphides of the metals ; 
 in such salts the oxygen has been replaced by sulphur, for 
 example : 
 
 CS 2 +Na 2 S = Na 2 CS 3 
 
 Carbon disulphide + Sodium sulphide = Sodium thiocarbonate.* 
 
 CO 2 -j-Na 2 = Na 2 C0 3 
 
 Carbon dioxide + Sodium oxide = Sodium carbonate. 
 
 In these reactions carbon disulphide and sodium sulphide bear the 
 same relationship to each other as sodium oxide and carbon dioxide 
 do. Sulphur also resembles chlorine, bromine, or iodine, for the 
 sulphides can be prepared in the same way as the halides are ; for 
 example, when phosphorus and sulphur are heated together, a sul- 
 phide of phosphorus is formed, just as is the case if phosphorus is 
 heated in chlorine, for in the latter event a chloride is produced. 
 Iron-filings, heated with sulphur, will form the sulphide of iron, 
 just as the same metal, heated with chlorine, will produce the 
 chloride. In addition to the resemblances which have been men- 
 tioned, the fact can be cited that many of the sulphides of the not- 
 metals are decomposed by water, forming an acid and hydrogen 
 sulphide, just as the chlorides of the same elements are decomposed, 
 forming an acid and hydrogen chloride, for example : 
 
 1. P 2 S 3 + 6 H 2 = 2 H 3 P0 3 + 3 H 2 S, 
 
 2. P 2 S 5 + 8 H 2 = 2 H 3 P0 4 + 5 H 2 S. 
 
 1. PC1 3 + 3H 2 O= H 3 P0 3 + 3HC1. 
 
 2. PC1 5 + 4H 2 0= H 3 P0 4 +5HC1. 
 
 The acid formed in the reactions which are numbered 1 is phosphor- 
 ous acid, in the reactions which are numbered 2 is phosphoric acid. 
 The tendency to form oxygen compounds of the not-metals, in most 
 cases, is greater than is the one to form compounds containing other 
 elements which are similar to oxygen. 
 
 and the same may be said of a number of oxides. Representatives of this 
 class are found as minerals. Whenever the sulphides are dissolved by acids, 
 they act like bases. 
 
 * The term thio is frequently used in place of sulpho; thus sodium sulpho- 
 carbonate is more often termed sodium tMocarbonate ; See chapter xxxix. 
 
96 HYDKOGEN SULPHIDE; OCCURRENCE ; HISTORY. 
 
 CHAPTER XIV. 
 
 HYDROGEN SULPHIDE. 
 
 Formula, H 2 S; specific gravity, air = 1, is 1.17, H 2 = 2, is 33.7 ; 1 
 c.c. of gas weighs .00152 grams. 
 
 HYDROGEN sulphide (sulphuretted hydrogen) occurs, mixed 
 with other gases and vapors, in some volcanic exhalations, and is 
 also occasionally present in coal and in other mines. As it is one 
 of the products of the decay of various animal and vegetable sub- 
 stances, it is necessarily frequently found in the atmosphere and in 
 water. Many sea plants, when exposed to the sun's rays, give off 
 sulphuretted hydrogen ; and when organic substances, which contain 
 sulphur (for example, bituminous coal), are heated without access 
 to the air,* sulphuretted hydrogen is found in the gases which are 
 given off. Many mineral waters contain large quantities of sul- 
 phuretted hydrogen ; these waters come from sulphur springs ; the 
 gas can be detected by its peculiar odor, which resembles that of 
 rotten eggs. 
 
 Sulphuretted hydrogen has certainly been known since the 
 sixteenth Or seventeenth century ; at a later time it was more accu- 
 rately studied by Scheele, who considered it to be composed of 
 heat, sulphur, and phlogiston. After Lavoisier's time its true nature 
 was explained. 
 
 1. Preparation of sulphuretted hydrogen by direct union of the 
 elements. 
 
 Sulphuretted hydrogen can be prepared, with some difficulty, by 
 passing hydrogen through molten sulphur. When we recall the 
 fact that hydrogen unites with explosive violence with either oxy- 
 gen or chlorine ; the diminished energy with which hydrogen and 
 sulphur or hydrogen and bromine combine becomes apparent. 
 
 2. Preparation of sulphuretted hydrogen for laboratory use. 
 
 * Called dry distillation. 
 
HYDROGEN SULPHIDE; PREPARATION. 97 
 
 Sulphuretted hydrogen is prepared for laboratory use by the 
 action of some acid on a sulphide ; for instance : 
 
 Fe S + H 2 S0 4 = Fe S0 4 + H 2 S, 
 FeS+2HCl = FeCl 2 +'H 2 S, 
 H 2 S0 4 = ZnS0 4 + H 2 S. 
 
 The oxide or the chloride would act in exactly the same way : 
 
 FeO + H 2 S0 4 = FeS0 4 + H 2 0, 
 Fe C1 2 + H 2 S0 4 = Fe S0 4 + 2 H Cl, 
 
 ZnO +H,S0 4 = ZnS0 4 + H 2 0, 
 Zn C1 2 + H 2 S0 4 = Zn S0 4 + 2 H Cl, 
 
 so that there is no essential difference between the action of a base 
 and that of other similarly constructed compounds in the presence 
 of sulphuric acid ; indeed, were sulphides and chlorides, and not 
 oxides, the most frequent chemical compounds, or were hydrochloric 
 acid and sulphuretted hydrogen, and not water, produced in many 
 reactions, the term base would never have been used to designate 
 the oxide. 40 
 
 Sulphuretted hydrogen is a colorless gas, with an intensely 
 disagreeable odor; it liquefies at a temperature of 11, with a pres- 
 sure of fifteen atmospheres. It boils at ordinary pressures at 61. 8 
 and becomes solid at 85 ; it is tolerably soluble in water, for at 
 ordinary temperatures one volume of water dissolves two volumes 
 of the gas. The solution slightly reddens litmus, but the red color 
 disappears on exposure to the atmosphere. Sulphuretted hydrogen 
 is very poisonous ; when inhaled in small quantities it causes head- 
 ache, loss of appetite, dizziness, and inflammation of the eyelids, 
 while persons who have been poisoned by sulphuretted hydrogen 
 often have fainting spells, at intervals, during some weeks. Death 
 may be caused by -gfa of a volume of sulphuretted hydrogen in the 
 atmosphere. 
 
 When a stream of the gas is lighted, it burns to form water and 
 sulphur dioxide ; a mixture of two volumes of sulphuretted hydro- 
 gen and three of oxygen is highly explosive ; where an insufficient 
 supply of oxygen is present during the combustion of the gas, 
 sulphur is deposited. Both chlorine and bromine decompose 
 sulphuretted hydrogen in a manner similar to the decomposition of 
 
98 HYDROGEN SULPHIDE; THERMOCHEMISTRY. 
 
 the same substance by oxygen, forming hydrobromic or hydrochloric 
 acid : 
 
 H 2 S + = H 2 + S, 
 
 H 2 S + 2 Cl = 2 H Cl + S, 
 
 H 2 S-f 2Br = 2HBr+S. 
 
 In each of these three cases the reason for the reaction is found in 
 the excess of the heat of formation of water, hydrochloric acid, and 
 hydrobromic acid over that of sulphuretted hydrogen. Iodine does 
 not decompose hydrogen sulphide when no water is present, 
 because, in the production of hydroiodic acid, heat is absorbed ; but 
 if the reaction takes place in contact with water, the heat of solution 
 of hydroiodic acid is sufficient to cause the following change to take 
 
 place: H 2 S+2I=2HI + S. 
 
 The foregoing changes will be understood if the heats of forma- 
 tion of these various compounds are placed in a column so as to 
 render comparison easier : 
 
 H 2 = 684K, 
 2HC1 = 440K, 
 2 H Br = 242 K, 
 2 HI = - 122 K, while 
 2 HI dissolved in water = 262 K. 
 
 Now, the heat of formation of sulphuretted hydrogen is but 27 K ; 
 the heat of solution of the same in water is 46 K ; therefore, the heat 
 of formation of hydrogen sulphide in water is 73 K, a number much 
 smaller than any of those given above.* 
 
 These relationships become apparent if we write the complete 
 reactions, including the thermal values of the various parts, thus : 
 
 * In each of the cases in which hydrogen sulphide is decomposed by a 
 halogen, the reaction, H 2 S + 2 X = 2 HX + S, takes place (X representing 
 either Cl, Br, or I). From the table given above, it follows that the system, 
 H 2 S + 2 X possesses more energy than 2 HX + S ; heat is therefore evolved 
 when chlorine, bromine, or iodine acts on hydrogen sulphide, and hence the 
 reaction takes place. In each of these cases the energy which must be added 
 is that which is necessary to decompose H 2 S into 2 H + S; for it is obvious that 
 hydrogen sulphide must be broken down into hydrogen and sulphur before the 
 rearrangement with the new, and more stable bodies (the halhydric acids) can 
 take place; the energy which is given off is from the reaction 2 H +2 X = 2 HX; 
 in each case the latter must exceed the former if a change is to take place. 
 
HYDROGEN SULPHIDE ; DECOMPOSITION. 99 
 
 1. H 2 S[27K] -f 2C1 = 2H,C1[440K]-H 2J S[27K] + S = 
 413 K. 
 
 2. H 2 S aq. [73 K] + 2 1 = 2 H, I aq. [262 K] - H 2 , S aq. [73 K] 
 + S = 189 K. 
 
 3. H 2 S [27 KJ + = H 2 ,0 [684 K] - H 2 , S [27 K] + S = 
 657 K. 
 
 The term aq. used above in equation (2) means that the react- 
 ing substances are dissolved in a quantity of water so great that a 
 further addition would have no effect on the thermal value of the 
 equation. The commas placed between the two portions of a 
 chemical formula indicate that the substance is to be formed from 
 the elements in question. 
 
 Hydrogen sulphide is readily decomposed into its elements, just 
 as hydrobromic and hydroiodic acids are, so that the former cannot 
 be prepared in the presence of concentrated sulphuric acid any more 
 than the two latter can, and it follows that sulphuretted hydrogen 
 cannot be dried by being passed through sulphuric acid. A hot 
 wire placed in the gas will dissociate it, 41 but no change in the 
 volume takes place, because the volume of the sulphur separated is 
 minimal when compared with the volume of the gas as a whole, 
 and, therefore, the hydrogen occupies the same space as that 
 previously taken by sulphuretted hydrogen : 
 
 H H 
 
 x S = + sulphur. 
 
 H H 
 
 1 volume hydrogen sulphide = 1 volume hydrogen. 
 
 In place of each molecule of hydrogen sulphide, therefore, we 
 have formed one molecule of hydrogen, or, in place of each volume 
 of hydrogen sulphide, an equal volume of hydrogen, so that this 
 fact proves, provided Avogadro's hypothesis is correct, that in a 
 molecule of sulphuretted hydrogen there are contained at least two 
 atoms of hydrogen. 
 
 We assume that there is but one atom of sulphur, because of the 
 analogy between water and sulphuretted hydrogen, and, also, because 
 while Ho S contains two parts by weight of hydrogen and thirty- 
 two of sulphur, in no compound, the molecular weight of which has 
 
100 HYDROGEN SULPHIDE; REACTIONS OF. 
 
 been determined and which. contains sulphur, is that element found 
 with a proportional weight smaller than thirty-two. 
 
 Hydrogen sulphide resembles the acids, because, in cases where 
 it reacts with the oxide or hydroxide of a metal, the corresponding 
 sulphide or sulphydrate is formed : 
 
 CaO + H 2 S =CaS + H 2 0, 
 Na OH + HSH = NaSH + HOH. 
 
 The reactions of the same bases with hydrochloric acid are as 
 
 follows : 
 
 CaO +2HCl = CaCl 2 + H 2 0, 
 NaOH-fHCl = NaCl + HOH. 
 
 Sulphuretted hydrogen contains two atoms of hydrogen, which 
 can be replaced by metals, so that we must distinguish two series of 
 compounds, in the first of which one of these atoms is substituted 
 as in Na SH ; in the second, two, as in Ca S. We shall subsequently 
 see that the same rule holds good with all acids containing two 
 replaceable hydrogen atoms. 
 
 The sulphides of a large number of metals are insoluble in dilute 
 acids, so that when sulphuretted hydrogen is passed into a solution 
 containing a salt of one or more of these metals, the corresponding 
 sulphide is precipitated : 
 
 Cu S0 4 + H 2 S = Cu S + H 2 S0 4 , 
 Pb (N0 3 ) 2 + H 2 S = PbS+2H N0 3 , 
 CdCl 2 + H 2 S = CdS+2HCl. 
 
 The sulphides of some metals are soluble in dilute acids, but 
 insoluble in water or in alkalis ; these will not be precipitated unless 
 provision is made to neutralize the acid formed. This can be accom- 
 plished either by passing sulphuretted hydrogen into a solution ren- 
 dered alkaline by the addition of ammonium or sodium hydroxide 
 solutions, or, better, by adding a soluble sulphide. 
 
 Na 2 S + Zn S0 4 = Zn S + Na 2 S0 4 . 
 
 Lastly, the sulphides of a third class of metals are soluble both 
 in acid or alkaline water; these will, of course, not be precipitated 
 by sulphuretted hydrogen. The action of sulphuretted hydrogen 
 on the salts of the metals is a ready means of detecting the presence 
 
HYDROGEN PERSULPHIDE. 101 
 
 of the metals in a solution, and many of the processes 1 of ' quali'tairVe' 
 analysis are founded upon these reactions, so that a further discus- 
 sion belongs to that branch of applied chemistry. 
 
 In addition to hydrogen sulphide there exists another compound 
 of hydrogen and sulphur called hydrogen persulphide. This latter 
 has the formula H 2 S 2 ; * a yellow, oily fluid with a penetrating odor 
 and corrosive action. It resembles its prototype, hydrogen dioxide, 
 (H 2 2 ), in the fact that it is stable in the presence of dilute acids, 
 and in the ease with which it decomposes into sulphur and sulphur- 
 etted hydrogen, just as hydrogen dioxide does into water and oxygen. 
 
 * This may be the formula of the compound, although this is not definitely 
 settled. There may exist more than one persulphide of hydrogen, or the reason 
 of varying amounts of sulphur which are found in the persulphide may be due 
 to sulphur dissolved by that substance. 
 
102 '' ' SELENIUM; OCCURRENCE, HISTORY. 
 
 CHAPTER XV. 
 
 SELENIUM AND HYDROGEN SELENIDE. 
 
 Symbol, Se ; atomic weight, 79.1 ; specific gravity, 4.5 ; specific grav- 
 ity of vapor, above 1400, air = 1, is 5.7 ; H 2 = 2, is 164 (calcu- 
 lated for Se 2 158.2 ). formula, H 2 Se ; specific gravity, air = 1, is 
 2.8,H 2 =2,^80.6. 
 
 SELENIUM occurs in selenides just as sulphur is found in sul- 
 phides; in quantity, however, it is found much more sparingly. 
 Free selenium is only very rarely found, and then in some volcanic 
 regions. The chief selenides are those of lead ( Pb Se) and of iron 
 (FeSe), while the selenide of silver (Ag 2 Se) also occurs. Lately, 
 considerable quantities of selenium combined with bismuth have 
 been found in some parts of South America. 
 
 Selenium was discovered by Berzelius, in 1817, in the lead cham- 
 bers used in the manufacture of sulphuric acid. He at first con- 
 fused the element with tellurium, but subsequently proved it to be 
 a hitherto unknown element, which he called selenium from o-cXyv-r), 
 moon, because the name of the other element was derived from 
 tellus, earth. 
 
 The occurrence of selenium in the sulphuric acid chambers, and 
 in the flues of furnaces in which sulphides are roasted, is due to the 
 presence of selenium in the ores (such as iron pyrites, copper pyrites, 
 or zinc blende). The selenium is burned to selenium dioxide, and 
 mechanically carried into the flues and chambers when the sul- 
 phides are roasted. Selenium dioxide is easily reduced to selenium 
 by means of reducing agents such as sulphur dioxide,* so that the 
 dust of the chambers contains selenium chiefly in the form of the 
 element. The isolation of selenium is quite a complicated process, 
 and a description of the methods must be left to a larger work. 
 
 Selenium exists in two allotropic forms, one soluble, the other 
 
 * Reducing agents are such substances as are capable of removing oxygen 
 or the equivalent of oxygen from chemical compounds, or they are such sub- 
 stances as can add hydrogen to elements or compounds. 
 
HYDROGEN SELENIDE. 103 
 
 insoluble in carbon bisulphide ; in that way it resembles sulphur. 
 When selenium is separated from its compounds by chemical means 
 it is, when moist, a crimson powder ; but it is dark red and soluble 
 in carbon bisulphide when dry. By heating the element above 80 
 it becomes iron gray in color. Selenium melts at 217, and boils at 
 665, and then, when cooled suddenly, is insoluble in carbon bisul- 
 phide. The selenium of commerce is formed by casting melted 
 selenium into sticks. It is almost metallic in its appearance, and 
 black in color. Selenium can exist in more than one crystalline 
 form, in that way resembling sulphur. 
 
 Chemically, the properties of selenium are closely akin to those 
 of sulphur. It burns in the air, forming selenium dioxide, just as 
 the latter forms sulphur dioxide. The selenides and selenium 
 compounds, in general, have formulae exactly like those of the cor- 
 responding sulphur compounds. The element is of little importance 
 excepting in a comparative study of the elements. 
 
 Hydrogen selenide is the complete analogon of hydrogen sul- 
 phide. It can be prepared with difficulty by the direct union of the 
 elements, being obtained by passing hydrogen over selenium heated 
 to its boiling point ; however, unless great care is taken to regulate 
 the temperature, the heat will decompose the hydrogen selenide so 
 formed. The gas can also be prepared by adding an acid to the 
 selenide of iron : 
 
 Fe Se + 2 H Cl = Fe C1 2 + H 2 Se, 
 
 in a manner analogous to the preparation of hydrogen sulphide. 
 
 Hydrogen selenide is a colorless gas, with a most penetrating 
 odor, somewhat resembling that of sulphuretted hydrogen ; it is. 
 extremely poisonous.* The gas, upon being heated, begins to. "de- 
 compose at a temperature of 150, but is only completely dissociated 
 at a considerably higher point. It burns even more readily than 
 does sulphuretted hydrogen, and, of course, is decomposed by chlo- 
 rine or bromine, or by iodine in the presence of water. It is more 
 soluble in water than is sulphuretted hydrogen, and has the same 
 action upon soluble salts of the metals as the latter. 
 
 * Care must be taken in working with the gas, for its odor clings to the 
 clothes for many days. 
 
104 TELLURIUM AND HYDROGEN TELLURIDE. 
 
 CHAPTER XVI. 
 
 TELLURIUM AND HYDROGEN TELLURIDE ; COMPARATIVE 
 TABLE OP THE ELEMENTS OF THE OXYGEN FAMILY. 
 
 Symbol, Te ; atomic weight, 125 j * specific gravity, 6.25 ; specific 
 gravity of vapor, air = l, is 9, H 2 =2, is 259 (above 1400). 
 Formula, H 2 Te. 
 
 TELLURIUM resembles both selenium and sulphur; it occurs as tel- 
 luride of silver, gold, lead, and also as tellurium. It was discovered 
 in 1782, and identified as an element in 1798. It is very rare and 
 of comparatively little importance. Its preparation from its ores is 
 a complicated process. It is a silver white, metallic appearing ele- 
 ment, \vhich melts at about 500, and boils at a high temperature, 
 forming an orange-colored vapor. The element is with difficulty 
 obtained free from selenium. Like sulphur and selenium, it exists 
 both as amorphous and crystalline tellurium. 
 
 Hydrogen telluride was discovered by Davy in 1810. It is best 
 prepared by adding hydrochloric acid to zinc or magnesium tel- 
 
 Zn Te + 2 H Cl = H 2 Te + Zn Cl a . 
 
 It is a colorless gas, which entirely resembles the hydrogen com- 
 pounds of sulphur and selenium. It burns readily, with a blue 
 flame, and is gradually decomposed into hydrogen and tellurium 
 even at ordinary temperatures. It is instantly changed on exposure 
 to the air : - H 2 Te + = H 2 + Te.t 
 
 The tellurides are with difficulty obtained pure, and are prepared 
 like the sulphides and selenides. 
 
 * Brauner ( Monatshefte fur Cheraie; 10, 411) maintains that the substance 
 which has heretofore been regarded as tellurium is really a mixture ; while the 
 pure element, if subsequently prepared, will probably have an atomic weight of 
 125 to 120, yet the only consistent numbers so far obtained give it an atomic 
 weight of 127.6. 
 
 t Compare this with the action of hydroiodic acid when exposed to the 
 air. 
 
ELEMENTS OF SULPHUR FAMILY ; TABLE OF. 
 
 105 
 
 A comparative table of the elements of the oxygen family will 
 serve to render the relationship between its members more appar- 
 ent, and will also make clear the resemblance between this family 
 and the halogens : 
 
 ELEMENTS. 
 
 
 ATOMIC 
 WEIGHT. 
 
 SPECIFIC 
 GRAVITY 
 OF SOLID. 
 
 MELTING 
 POINT. 
 
 BOILING 
 POINT. 
 
 APPEARANCE. 
 
 NOT- 
 METALLIC 
 PROP- 
 
 
 
 
 
 
 
 ERTIES. 
 
 o 
 
 16. 
 
 
 
 
 
 -182 
 
 Gas. 
 
 
 s 
 
 32.06 
 
 2.04 
 
 114 
 
 440 
 
 Yellow solid. 
 
 
 Se 
 
 79.1 
 
 4.5 
 
 217 
 
 665 
 
 Dark red powder, 
 
 
 
 
 
 
 
 black when fused. 
 
 
 Te 
 
 125. (?) 
 
 6.25 (?) 
 
 500 
 
 above 1000 
 
 Silver white, metal- 
 
 
 
 
 
 
 
 lic appearance. 
 
 
 
 SPECIFIC 
 
 SPECIFIC 
 
 
 
 
 GRAVITY 
 
 GRAVITY 
 
 
 
 
 OF 
 
 OF 
 
 MOLECULE. 
 
 
 
 VAPOR. 
 
 VAPOR. 
 
 
 
 
 AlR=l. 
 
 H 2 = 2. 
 
 
 
 o 
 
 1.1 
 
 31.76 
 
 2 
 
 * The molecular weights, Se 2 = 158, Te 2 = 
 
 
 
 
 
 250, are somewhat less than the specific gravi- 
 
 
 
 
 
 ties found, but near enough to show that the 
 
 s 
 
 2.2 
 
 63.3t 
 
 S 2 
 
 molecule exists as two atoms. 
 
 
 
 
 
 t The specific gravities of sulphur and se- 
 
 
 
 
 
 lenium vapors are not constant below 1000; 
 
 
 
 
 
 they gradually become larger as the boiling 
 
 Se 
 
 5.7 
 
 164.*t 
 
 Se 2 
 
 points are approached. The molecules S 2 and 
 
 i 
 
 
 
 
 Se 2 seem to form larger aggregations as the 
 
 
 
 
 
 elements approach the temperature of lique- 
 
 
 
 
 
 faction ; no definite formulae seem assignable 
 
 Te 
 
 9. 
 
 259.* 
 
 Te 2 
 
 to these molecules. 
 
 HYDROGEN COMPOUNDS. 
 
 
 HEAT OF 
 FORMA- 
 
 STABIL- 
 
 
 HEAT OF 
 FORMA- 
 
 STABIL- 
 
 
 
 
 
 
 
 
 
 
 TION. 
 
 
 
 TION. 
 
 
 
 H,0 
 
 684 K 
 
 
 2 HF 
 
 
 
 
 H 2 S 
 
 H 2 Se 
 
 27 K 
 - 111K 
 
 
 2HC1 
 2HBr 
 
 440 K 
 242 K 
 
 
 All of the hydro- 
 gen compounds are 
 colorless gases above 
 
 H 2 Te 
 
 
 
 
 2 HI 
 
 -122K 
 
 
 100 Centigrade. 
 
106 ELEMENTS OF SULPHUR FAMILY ; TABLE OF. 
 
 On comparing the atomic weights of the elements of the oxygen 
 family with those of the halogens, we see that the former are, 
 throughout, somewhat smaller for corresponding elements ; the dif- 
 ference, however, is but slight* 
 
 16 F 19 
 
 S 32.06 Cl 35.45 
 
 Se 79 Br 79.95 
 
 Te 125 I 126.85 
 
 * Tellurium, if its atomic weight is 127.6, is at present an exception. See 
 page 104. 
 
VALENCE AND THE OXYGEN COMPOUNDS. 107 
 
 CHAPTER XVII. 
 
 VALENCE AND THE OXYGEN COMPOUNDS OP THE NOT-METALS. 
 
 THE elements, the properties of which, we have studied, form 
 compounds with hydrogen, all of which, with two exceptions 
 hydrogen dioxide and the corresponding sulphur compound can 
 be obtained as vapors, the specific gravities, and hence the molecular 
 weights of which can, therefore, readily be ascertained. By this 
 means we arrived at the conclusion that one atom of chlorine, bro- 
 mine, or iodine could unite with but one atom of hydrogen, while 
 one of oxygen and of the remaining members of that family could 
 unite with two. These elements, therefore, differ among themselves 
 in their power of retaining hydrogen atoms. In addition to the 
 foregoing there are other elements, the hydrogen compounds of 
 which are composed of three atoms of hydrogen to one of the nega- 
 tive element ; these elements are nitrogen, phosphorus, arsenic, and 
 antimony. If we designate any one of these elements by Y, then 
 the formula of the hydrogen compounds would be YH 3 . Only two 
 other elements, carbon and silicon, form gaseous hydrogen com- 
 pounds; the general formula of these is ZH 4 All hydrogen 
 compounds can therefore be classed under four heads : 
 
 YH, XH 2 , YH 3 ,ZH 4 .* 
 
 From the outset we have considered chemical compounds as formed 
 by the conjunction of the atoms of elements ; the atoms themselves 
 are hypothetical, f but using this hypothesis as a basis, a chemical 
 
 * We can imagine all apparent variations from these types as formed by 
 the substitution of one or more of these hydrogen atoms by some other element, 
 or groups of elements. Thus, we have considered sodium hydroxide as water, 
 in which one atom of hydrogen has been replaced by sodium ; Na-O-H, H-O-H ; 
 hydrogen dioxide as water, in which one atom of hydrogen has been replaced 
 by hydroxyl H-O O-H, H O-H, and so on. 
 
 t Sir William Thomson considers them to be rings formed by vortical 
 motion of the ether; a visible example of such motion would be a smoke ring 
 blown by a locomotive. 
 
108 VALENCE; HYDROGEN COMPOUNDS. 
 
 theory productive of the greatest results has been developed. 
 " Every finite quantity of matter occupies a position in space which 
 is definable with regard to other material particles; the question as 
 to the relative position (or motion) of atoms in the molecule is 
 scientifically justified, and must be put sooner or later " * by persons 
 holding the atomic theory. This problem has been put by chemists 
 ever since the time of Berzelius ; and the great advance in organic 
 chemistry, which has been reflected in inorganic chemistry, is the 
 result of its successful solution. In the hydrogen compounds of the 
 not-metallic elements we are cognizant of the number of atoms in 
 the molecule, because we have been able to determine the quantitative 
 composition of these compounds and also the molecular weights. 
 
 As the mass of the atom of the not-metal, in hydrogen com- 
 pounds, is so much greater than that of the hydrogen, and as, in 
 the formation of new chemical compounds from those of hydrogen, 
 it is always the hydrogen which is replaced by some other element, 
 provided the resulting compound remains identical or similar in 
 character; therefore, it is more than probable that the hydrogen 
 atoms are joined to the not-metal. Whether the position of these 
 is fixed, or whether they are free to rotate around the not-metal, we 
 cannot decide ; but recent investigation all tends toward the former 
 theory. In the compounds VH, XH 2 , YH 3 , and ZH 4 . the numerical 
 capacity possessed by the not-metal of uniting with one, two, three, 
 and four atoms of hydrogen is termed the valence of the element, 
 and we can conveniently express this valence by Roman numerals 
 
 I II III IV 
 
 placed over the symbol of the not-metal, V, X, Y, Z, or by lines 
 drawn from them, as V , X , Y or Z . The element 
 
 I 
 
 which can unite with one atom of hydrogen is said to be univalent ; 
 that which can unite with two, bivalent ; with three, trivalent ; and 
 with four, quadrivalent; hydrogen is always univalent. All ele- 
 ments, the valence of which is more than one, may be called poly- 
 valent. When a univalent element has united with another element 
 or radicle,t its capacity for further union has ceased ; a bivalent 
 element, however, when united with an element or radicle by one 
 
 * W. Ostwald, Outlines of General Chemistry, Walker's translation, 
 t A group of elements which can, as a whole, replace an element in a 
 chemical compound, is called a radicle. 
 
VALENCE; CHLORINE COMPOUNDS. 109 
 
 valence has not lost its capacity for further union with other ele- 
 ments or radicles, it is unsaturated; for example, H X is in this 
 
 H 
 
 I 
 condition, and can act as a univalent radicle. Similarly H Y is 
 
 also unsaturated and univalent; H Y unsaturated and bivalent, 
 H H 
 
 I I I 
 
 H Z H, H Z , and H Z , unsaturated, and respectively uni-, 
 
 I i I 
 
 bi-, and trivalent. 
 
 Only the not-metals, however, form hydrogen compounds obtain- 
 able as gases, so that with other elements, if we desire a similar 
 means of determining valence, we must seek for gasifiable compounds 
 with some univalent element other than hydrogen. Many of the 
 metals and of the not-metals are capable of forming such compounds 
 with chlorine; the molecular weights of these can therefore be 
 determined. 
 
 The halogens form compounds with formulse analogous to those 
 of hydrogen, and such compounds can obviously be used to deter- 
 mine the valence of the elements ; for if the number of hydrogen 
 atoms with which an element is united in a molecule indicates the 
 valence of that element, so must the number of chlorine atoms in a 
 similar molecule. We can, therefore, construct a table containing 
 a series of chlorine compounds, just as we did with the hydrogen 
 compounds, and further investigation shows us that all of these 
 compounds can be brought under six heads ; using M as a general 
 term to denote an atom of an element with the capacity of uniting 
 with chlorine, we have : 
 
 MCI, MC1 2 , MCI., MC1 4 , MC1 6 , MC1 6 , and using Koman 
 numerals to designate the valence : 
 
 I II III IV V VI 
 
 M, M, M, M, M, M. 
 
 In many cases in which the hydrogen compound of a given ele- 
 ment exists, we can also study the chlorine compounds, so that 
 a determination of the valence of the elements by means of the 
 latter offers the advantage of being applicable in a greater number 
 of cases. Sometimes, as is the case with some metals, the bromide 
 
110 VALENCE IN COMPOUNDS OF POLYVALENT ELEMENTS. 
 
 or iodide is obtainable as a gas, while the chloride is not ; then the 
 former compounds answer just as well as a means of determining 
 the valence. In our considerations we have, so far, always been 
 able to appeal to experiment to answer any questions which may 
 arise ; in those which follow we shall have to indulge, more or less, 
 in speculation. 
 
 Can elements, polyvalent toward hydrogen, unite with each 
 other, and, if so, what is their valence ? The answer to the first 
 part of the question has already been given ; we have become aware 
 of compounds such as C0 2 , S0 2 , CS 2 , all of which are formed by the 
 union of polyvalent elements ; and in these cases, as in the vast 
 majority of those which fall within the scope of this book, one 
 atom of one of the elements in the molecule unites all the others ; 
 furthermore, substances such as C0 2 , S0 2 , or CS 2 can have their 
 molecular weights determined in the same way as can those of com- 
 pounds of hydrogen, so that the same reasoning will apply with the 
 former as with the latter. We could, therefore, suppose such com- 
 pounds to have a structure similar to that of water (H H) ; i.e., 
 C 0, S 0, S C S ; in such an event, carbon or sulphur 
 would be bivalent, as is oxygen in water. Indeed, any theory other 
 than this as regards the valence of the elements in those compounds 
 which contain bivalent elements, goes beyond the realm of facts 
 which are absolutely proved by experiment. In spite of this, the 
 majority of chemists have thought themselves justified in holding 
 other views, and a few of the reasons for their opinions may not be 
 out of place here. One atom of oxygen unites with two atoms of 
 hydrogen to form a molecule of water, so that in this compound it 
 is undoubtedly bivalent ; furthermore, oxygen is capable of uniting 
 with two atoms of any other univalent element, such as sodium or 
 potassium, the oxides of which are Na 2 0, K 2 0, and the element can 
 also unite two univalent radicles, or groups of elements, to form 
 compounds (calling any radicle q) of the formula q q ; so that 
 an atom of this element can serve as a link between groups of ele- 
 ments, a function which is evidently impossible for univalent ele- 
 ments. When oxygen replaces hydrogen in the compounds of that 
 element, then one atom of the former always takes the place of two 
 of the latter ; for instance, the compound CH 4 , in being oxidized, 
 
 forms, in addition to water, C j Q 2 at first, and then C -j Q : - 
 
VALENCE OF OXYGEN. Ill 
 
 CH 4 + 20 = CH 2 + H 2 0; arid 
 CH 2 H- 2 = C0 2 + H 2 0. 
 
 The valence of an element can, as we have seen, also be discovered 
 by a study of the formula of its chloride ; and when a chloride is 
 converted into an oxide, one atom of oxygen always replaces two of 
 chlorine : 
 
 CHLORIDES. OXIDES. CHLORIDES. OXIDES. 
 
 2NaCl Na 2 2A1C1 8 A1 2 3 
 
 2KC1 K 2 2FeCl s Fe 2 O 3 
 
 CaCl 2 CaO C C1 4 C 2 
 
 FeCl 2 FeO 2PC1 3 P 2 O 3 .* 
 
 From these considerations it is supposed that oxygen remains 
 bivalent wherever it enters into chemical combination. In assum- 
 ing this to be the case, we must consider an atom of oxygen as 
 united to other elements in a way unlike that in which one of hy- 
 drogen unites. Applying what we learned in regard to univalent 
 elements, where we saw that when one atom of an univalent ele- 
 ment is united to one of any other, its further power of union is 
 exhausted, we construct the following arbitrary rule : 
 
 One valence of any element in a chemical compound always calls 
 for and neutralizes a corresponding valence in the other element or 
 elements with which it is united. 
 
 The two valences of a bivalent element, therefore, are supposed 
 to neutralize two corresponding valences in any element or com- 
 pound with which it is united. The following examples will serve 
 to make this meaning more clear : 
 
 C\H 
 \H 
 
 {Two atoms of uni- ^ C Four atoms of uni- 
 valent hydrogen re- valent hydrogen re- 
 placed by one of bi- 1 I placed by two of bi- 
 valent oxygen, carbon f | valent oxygen, carbon 
 remaining quadriva- remaining quadriva- 
 lent. J 1^ lent. 
 
 * All of the chlorides in this list have been obtained as gases, their molec- 
 ular weights and formulae are certain; the first six corresponding oxides have 
 not, but having once determined the atomic weights of the elements, the for- 
 mulae of such compounds follow from their composition by weight. 
 
112 FORMULA OF OXIDES. 
 
 Where an atom of a polyvalent element has an odd number of 
 valences, it follows that these cannot be exactly neutralized by 
 those of an atom of a bivalent element; this fact will become 
 apparent by a study of the following formulae : 
 
 n /Cl ,0 
 
 P Cl with oxygen yields P (phosphorus remaining triva- 
 
 \ci 
 
 lent) this group will change to j ^ with one oxygen atom 
 
 unsaturated ; and the latter must therefore unite with some other 
 element or group of elements, so that : 
 
 
 
 (phosphorus remaining trivalent and oxygen uniting the two univa- 
 lent groups of atoms). By a similar application of the rule we can 
 come to the conclusion that, where five oxygen atoms unite with two 
 of some other polyvalent element, the latter is quinquivalent, and 
 where three unite with one it is hexavalent, so that the following 
 table of the formulae of oxides can be constructed (using X to de- 
 note an atom of any element) : 
 
 i 
 
 X 2 0, valence of X one ; denoted by X 2 
 
 ii 
 
 X 0, 
 
 " X two ; 
 
 
 
 X 
 
 
 
 
 III 
 
 X 2 3 , 
 
 " X three ; 
 
 a 
 
 " X 2 3 
 
 
 
 
 IV 
 
 X 2 , 
 
 " X four; 
 
 tt 
 
 " X 2 
 
 
 
 
 V 
 
 X 2 5 , 
 
 X five ; 
 
 tt 
 
 " X s0 5 
 
 
 
 
 VI 
 
 X 3 , 
 
 " X six; 
 
 a 
 
 " X 3 
 
 
 
 
 VII 
 
 X 2 7 , 
 
 " X seven ; 
 
 rt 
 
 <e X 
 
 
 
 
 VIII 
 
 X 4 , 
 
 " X eight ; 
 
 a 
 
 X 4 
 
 The formulae of all oxides, with a few exceptions, correspond 
 to these general formulae, and what has been said of oxygen applies 
 to all other bivalent elements. In endeavoring, therefore, to ascer- 
 
FORMULAE OF OXIDES. 113 
 
 tain the valence of an. element forming an oxide, we must consider 
 all of the oxygen atoms in the molecule to be held in position by 
 the atoms of that element, just as is the case in the structure of the 
 hydrogen and chlorine compounds we must ascertain the gravi- 
 metric composition by analysis, so that we can learn the number of 
 atoms united in the formula weight ; where we can obtain the com- 
 pound as a gas we must learn its specific gravity, and by this means 
 its molecular weight. For instance, there are two oxides of phos- 
 phorus, in one of which 62 parts by weight of phosphorus unite 
 with 48 of oxygen, in the other 62 of phosphorus unite with 80 of 
 oxygen. From the study of many compounds of oxygen and phos- 
 phorus, we have concluded that the atomic weights of these two 
 elements are 16 and 31 respectively, hence the composition by 
 weight of the first of- the two compounds leads us to a formula 
 P 2 O 3 , that of the second to a formula P 2 5 . Now we can construct 
 two formulae by writing out the individual atoms, grouping them 
 together so that connecting lines will express the valences, as 
 follows : 
 
 o=r-o-p = o, 
 
 Such diagrams, constructed of chemical symbols expressing a 
 theory regarding the manner in which atoms are grouped in a 
 chemical compound, are called structural formulas. Structural 
 formulae may be more or less complete, in order to express the prev- 
 alent ideas in regard to the relative positions of the atoms in a 
 molecule. The two given above are not intended to show more than 
 that the oxygen atoms are all united to those of phosphorus by 
 means of two valences apiece, and that the two phosphorus atoms 
 are united to each other by means of the interposed oxygen atom ; 
 other formulae, however, might be constructed illustrating a theory 
 regarding the relative position of all of the atoms, as has been done 
 with compounds of carbon, in which field recent investigation has 
 brought chemistry to a point where the relative positions in space 
 of atoms forming a molecule can be studied. It will fall within 
 the scope of this work to use only structural formulae, such as 
 those which are given above. 
 
 A trivalent element cafci unite with bivalent elements, with other 
 trivalent elements, or with a quadrivalent element; it can unite 
 
114 STRUCTURAL FORMULAE. 
 
 three univalent radicles, and so on ; its valence remaining three, one 
 valence, in uniting, neutralizes another, as in the case of uni- and 
 bivalent elements; we also could develop similar theories with 
 the other polyvalent elements. 
 
 Such are, briefly, the rules which, for the sake of uniformity, 
 are used by the majority of chemists. In applying such methods, 
 chemists are apt to allow themselves to forget how far experiment 
 has gone, and thus to confuse the theory and its application with 
 the phenomena of nature. The terms " valence " or " bond," used 
 to express the means of union between atoms, may become to them 
 material things, the lines used on paper to express the supposed 
 manner of union of atoms may simulate real linkings between exist- 
 ing atoms ; so that, at the present time, the science of chemistry is 
 in danger of being discredited by a too dogmatic conception of the 
 application of these methods ; and thus one of her greatest achieve- 
 ments of recent times, the theory of valence, may ultimately prove 
 a serious obstacle in the path of her development. In comparing 
 two compounds, such as H-O-H and 0-C-O, we cannot, without in- 
 dulging in speculations, assert more than that, in one case, oxygen 
 has the capacity of retaining two atoms of hydrogen in a molecule ; 
 and that, in the other, carbon presents the same relationship toward 
 oxygen ; the supposition that oxygen is retained in the molecule C0 2 
 by a force differing in its manifestations from the one retaining 
 hydrogen in H 2 is purely gratuitous, and based upon the action of 
 oxygen and hydrogen in entirely different compounds. In studying 
 substances, the molecular weights of which we have not, as yet, been 
 able to determine, as is the case with P 2 3 , experiment can go no 
 farther than to show us the composition of such compounds by 
 weight ; of the structure and size of the real molecules and the man- 
 ner of union of the atoms in these molecules, we have no knowledge, 
 and therefore, with such bodies, our rules of valence can be applied 
 only to the formula weights. If, however, we remember where the 
 boundary between speculation and experiment lies, we shall very 
 much simplify the study of chemistry by the application of these 
 rules. The valence of the atom of an element is, as much as any 
 other property of that element, determined by the family in which 
 that element belongs, and hence by its atomic weight. The position 
 of an element in the periodic system oi. the elements, therefore, 
 offers one of the best means of determining its valence. An atom 
 
ANHYDRIDES AND ACIDS. 
 
 115 
 
 may, under differing conditions, vary the number of oxygen atoms 
 with which it unites, thereby changing its valence ; the not-metals 
 are especially prone each to form a number of oxides. All of 
 the not-metals, with the exception of fluorine and bromine, have 
 oxides which are constructed according to certain types. The fol- 
 lowing is a table, containing the formulae of those oxides which 
 are most frequently met with ; as will be seen, the same general 
 structure often belongs to compounds in more than one family. X, 
 in each family, is used to represent a typical element : 
 
 HALOGENS. 
 
 OXYGEN FAMILY. 
 
 NITBOGEN FAMILY. 
 
 CASBON FAMILY. 
 
 I 
 
 IV 
 
 I 
 
 IV 
 
 X 2 
 
 X0. 2 
 
 x. 2 o* 
 
 X0 2 
 
 III 
 
 VI 
 
 III 
 
 
 X 2 3 
 
 X0 3 
 
 x. 2 o 3 
 
 
 v 
 
 
 V 
 
 
 X 2 5 
 
 
 X 2 5 
 
 
 VII 
 
 
 
 
 X,0 7 
 
 
 
 
 
 Highest valence 
 seven. 
 
 Highest valence 
 six. 
 
 Highest valence 
 five. 
 
 Highest valence 
 four. 
 
 When the oxide of a not-metal is the anhydride of an acid, and 
 is changed to the acid by the addition of water, the valence of the 
 atoms of the element which characterizes that anhydride is not 
 supposed to be changed; such a change can only be effected by 
 either oxidation or reduction. One or two examples will make this 
 apparent. C1 2 Of is a compound in which the two atoms of chlo- 
 rine are united to one of oxygen ; chlorine being univalent and the 
 structural formula of this oxide comparable with that of water : 
 
 H H 
 
 Water. 
 
 Cl Cl 
 
 Chlorine monoxide. 
 
 When chlorine monoxide is added to water the following change 
 takes place : ^ + H 2 0=2HOC1 
 
 Chlorine monoxide -{- Water = Hypochlorous acid. 
 
 The nature of this reaction becomes more apparent if it is written 
 as follows : 
 
 * This oxide of nitrogen does not form an acid on addition of water, but 
 a corresponding acid is known, which breaks down into water and this oxide, 
 t X 2 O under the head of halogens in the above table. 
 
116 FORMATION OF CHLORINE ACIDS. 
 
 H H H / H 
 
 H 2 + C1 2 0= + / / / 
 
 Cl Cl Cl / Cl 
 
 Whether we consider H Cl as water in which the' atom of 
 hydrogen is replaced by one of chlorine, or as chlorine monoxide in 
 which one atom of chlorine is replaced by one of hydrogen, is a 
 matter of indifference, chlorine obviously remains univalent in 
 the acid, as it was in the anhydride. Now, in order to change 
 C1 2 to C1 2 3 we must add oxygen.* 
 
 Cl Cl + 20 = = Cl Cl = 0. 
 
 The process of oxidation leaves the group of atoms Cl Cl 
 intact. Such a group as this is entirely independent of any added 
 valence assumed by chlorine. Now, when C1 2 3 reacts with 
 water : - ^ Og -f H 2 = 2 H0 2 Cl 
 
 Chlorine trioxide + Water = Chlorous acid, or 
 
 H H H 
 
 H 2 + C1 2 3 = + 
 
 = Cl Cl = = Cl Cl = ; 
 
 the valence of chlorine remains three, for the added oxygen atoms 
 have not changed the nature of the reaction. What is true of 
 C1 2 and C1 2 3 must be true of X 2 or X 2 3 , and by similar 
 structural formulae we can show that when the reactions 
 
 X 2 5 + H 2 = 2 H0 3 X, and X 2 7 + H 2 = 2 H0 4 X 
 
 take place, the valence of X remains unaltered, for the point of at- 
 tack for the water is the oxygen uniting the two univalent atoms, 
 or groups of atoms. With the other class of anhydrides given in the 
 table named, i.e., those formed by not-metals having an even num- 
 ber of valences, the addition of water has no different result. As 
 there is but one atom of the typical element present in these oxides, 
 the entire molecule of water must add itself to one molecule of the 
 oxide ; but a little consideration will show us that this change is 
 identical with those we have discussed, for an oxygen atom which 
 is present in the anhydride, with the addition of a molecule of 
 water, forms two hydroxyl groups : 
 
 * It is scarcely necessary to remind the pupil that the oxygen atom so 
 added must be fixed by chlorine; the oxygen in Cl O Cl, being bivalent, 
 is incapable of further union. 
 
FORMATION OF SULPHUR ACIDS. 
 
 117 
 
 a. 
 
 a. S0 2 +H 2 = H 2 S0 3 
 
 Sulphur dioxide + "Water = Sulphurous acid. 
 
 S0 3 + H 2 0=H 2 S0 4 
 
 6. Sulphur trioxide -f- Water = Sulphuric acid. 
 
 H 
 
 H 
 H 
 
 
 
 II 
 S 
 
 II 
 
 
 
 
 II 
 
 b. 0=S 
 
 + 
 
 H 
 
 
 
 H 
 
 I 
 H 
 
 H 
 = S H 
 
 
 
 When S0 2 is changed to S0 3 by oxidation, the valence of sulphur 
 is changed, but obviously the added oxygen atom plays no more 
 part in the reaction which takes place by addition of water, than 
 it does in the cases previously considered. The water which is 
 taken up by anhydrides is therefore decomposed by them, it breaks 
 down into H and -0-H. The hydrogen attaches itself to oxygen, 
 while the hydroxyl group unites with the atom of the not-metal, 
 which characterizes the anhydride. Water, therefore, does not enter 
 into the composition of the acids, but hydroxyl groups do. In the 
 following table are placed the general formulae of the oxides pre- 
 viously considered and the corresponding acids ; X, as heretofore, 
 denoting the not-metal : 
 
 ( ' 
 
 X 2 + H0 = 2lHO X 
 
 Ill 
 
 X 2 3 
 
 v 
 
 X 2 5 
 
 VII 
 
 X 2 7 
 
 IV 
 
 X 2 
 
 VI 
 
 X 0, 
 
 _ 9 
 
 m 
 
 H 2 = 2VH0 3 X 
 
 ( VII \ 
 + H 2 = 2^H0 4 Xy 
 
 IV 
 
 + H 2 = H 2 3 X 
 
 VI 
 
 oar 
 
118 FORMATION OF OXIDES AND ACIDS. 
 
 The formulae of these six oxides and of the corresponding acids 
 occur more frequently, and are of greater importance, than any 
 others which the pupil will encounter during his study of the not- 
 metals. The addition of oxides of metals other than hydrogen 
 to the anhydrides must necessarily produce reactions similar to 
 those we have just considered. 
 
 
OXIDES AND OXY-ACIDS OF CHLOKINE. 119 
 
 CHAPTER XVIII. 
 
 THE COMPOUNDS OF CHLORINE WITH OXYGEN, AND WITH 
 OXYGEN AND HYDROGEN. 
 
 CHLORINE forms the following compounds with, oxygen, and 
 with oxygen and hydrogen : 
 
 1. C1 2 0, chlorine monoxide, 1. HO Cl, hypochlorous acid, 
 
 2. C1 2 3 , chlorine trioxide, 2. H0 2 Cl, chlorous acid, 
 
 3. (Cl 2 , chlorine dioxide,)* 3. H0 3 Cl, chloric acid, 
 
 4. H0 4 Cl, perchloric acid. 
 
 The first and second of these oxides are respectively the an- 
 hydrides of hypochlorous acid and of chlorous acid ; the third does 
 not form any corresponding acid.* The anhydrides of the third 
 and fourth acids are not known ; whenever an attempt is made to 
 isolate them, complete decomposition takes place. The formulae 
 of the compounds would be C1 2 5 and C1 2 7 , were they capable of 
 existence. 
 
 Where there are four acids which contain oxygen, and which 
 are derived from the same element, the nomenclature given in the 
 following table has been adopted ; beginning with the one contain- 
 ing the least oxygen, we have : 
 
 1. Prefix hypo termination in ous. 
 
 2. No prefix- " -ous. 
 
 3. No prefix " - ic. 
 
 4. Prefix per- " " ic. 
 
 Where but two acids are known, the nomenclature is the same 
 as under 2 and 3. 
 
 Salts derived from an acid terminating in -ous, have their 
 names end in -ite; those from an acid terminating in -ic, in 
 ate. The prefix in the name of the acid is always retained in 
 the name of the salt. The salts obtained from the above acids by 
 
 * On being passed into a solution of potassium hydroxide, it forms potas- 
 sium chlorite, K Cl O 2 , and potassium chlorate, K Cl O 3 . 
 
120 CHLOEINE MONOXIDE. 
 
 replacing the hydrogen with some other metal would consequently 
 be named as follows, with potassium as the metal : 
 
 1. KO Cl, potassium hypocbjorite, 
 
 2. K0 2 Cl, potassium chlorite, 
 
 3. K0 3 Cl, potassium chlorate, 
 
 4. K0 4 Cl, potassium perchlorate. 
 
 Where there are but two acids, the nomenclature is the same as 
 under 2, and 3. 
 
 The oxides of chlorine all are unstable, some of them even 
 extremely explosive, so that chlorine and oxygen cannot be brought 
 to unite directly. In endeavoring to find a cause for this extremely 
 ready decomposition, we observe that these oxides are produced by 
 the union of two very similar elements, so that their instability is a 
 circumstance which is certainly not unexpected ; for oxygen and flu- 
 orine, which resemble each other even more than do oxygen and chlo- 
 rine, are incapable of combination. One might be tempted, from 
 such a circumstance, to argue that not-metals which are most like 
 each other would have the least chemical affinity ; but this is true 
 only in a very few cases. The oxides of sulphur and of phosphorus, 
 for instance, are extremely stable bodies; the different halogens can, 
 in some instances, combine with each other; while the existence of 
 molecules such as those of the elements (for example of chlorine and 
 of hydrogen), shows us that even the atoms of the same element 
 may be most firmly united. Any attempts to generalize regarding the 
 relative stability of these oxides are at the present time premature. 
 
 (Chlorine monoxide is not of itself important, but the acid formed 
 from it by the addition of water is of the greatest commercial value ; 
 its salts, more especially that of calcium, being the compounds most 
 often used for the purpose of bleaching. Chlorine monoxide can be 
 prepared by passing chlorine over mercuric oxide. 
 
 It is a reddish-yellow gas, with a penetrating odor resembling 
 that of chlorine. At low temperatures it is condensed to a red 
 liquid, which boils at 19-20 centigrade. It has a specific grav- 
 ity, air being 1, of 3.00, hydrogen being 2, of 86.4. Its molecular 
 weight is therefore 86.9, its formula C1 2 0.* It gradually dis- 
 
 * When slightly warmed, chlorine monoxide explodes; this decomposition 
 may even take place at ordinary temperatures. 
 
HYPOCHLOBITES. 121 
 
 solves in water, forming an orange-colored solution which contains 
 hypochlorous acid : - 
 
 H H H H 
 
 + / + / 
 
 Cl Cl Cl Cl 
 
 The salts derived from hypochlorous acid are easily prepared ; 
 they are really of greater importance and are more stable than is 
 the acid itself. 
 
 In 1788, Berthollet discovered that a bleaching solution could be 
 prepared by passing chlorine into a solution of an alkali. He at 
 once attempted to utilize his discovery commercially, and so estab- 
 lished a factory at Javelles, where bleaching water was prepared by 
 treating a solution of potashes (potassium carbonate) with chlorine ; 
 this solution then contained hypochlorous acid, and was called eau 
 de Javelles. At a later date, lime took the place of potashes and 
 the resulting bleaching powder became known as chloride of lime. 
 
 If chlorine is passed into a solution of potassium hydroxide, the 
 first product will necessarily, because of the great chemical af- 
 finity between the metal and the not-metal, be potassium chloride ; 
 therefore : -- K _ _ H + C l = K Cl + H. 
 
 VV G \\ t" Q*r ^ 
 
 The group 'O H is, however, unsaturated, so that it can 
 
 unite with another atom, in this case with one of chlorine, H 
 -|- Cl = H Cl. The first products obtained by passing chlo- 
 rine into a solution of potassium hydroxide would therefore be potas- 
 sium chloride and hypochlorous acid. Now, potassium hydroxide 
 is a base ; therefore, with hypochlorous acid, it will form a salt 
 (potassium hypochlorite), and water, as follows : 
 
 KOH + HO Cl = KO Cl + H 2 0. 
 
 , As a consequence, the ultimate products of the action of chlorine 
 upon potassium hydroxide are potassium chloride, potassium hypo- 
 chlorite, and water. The entire change is summed up in the follow- 
 ing reaction : 
 
 2 KOH 4- 2 Cl = K Cl + KO Cl + H 2 0.* 
 
 * That free hypochlorous acid is at first produced, and that this acid sub- 
 sequently reacts with potassium hydroxide to form potassium hypochlorite, is 
 proved by the fact that if a weak base, which is incapable of forming a hypo- 
 chlorite (zinc oxide, mercuric oxide), is used, then the chloride and hypochlo- 
 rous acid are the products of the reaction. See also Remsen, Chemistry, p. 113. 
 
122 HYPOCHLOROTJS ACID; DECOMPOSITION. 
 
 If calcium hydroxide is used in place of potassium hydroxide, 
 then calcium chloride, calcium hypochlorite, and water are produced ; 
 the formulation of the reaction would be modified by the bivalence 
 of the atoms of calcium, which have twice the capacity for repla- 
 cing hydrogen in acids that those of potassium have. This reac- 
 tion is therefore as follows : 
 
 H ( 01 H Cl 
 
 + 4Cl = Ca-3 + and 
 
 -0 H ( 01 H 01 
 
 _ H HOC1 (_0 C1 HOH 
 
 2. Ca^ + =Ca-} + 
 
 H HOC1 ( 01 HOH 
 
 uniting 1 and 2, we have : 
 
 2 Ca (OH), + 4 01 = Ca C1 2 + Ca (0 Cl) 2 + 2 H 2 0.* 
 
 A moderate heat changes hypochlorites into chlorates ; and there- 
 fore the above reactions must be performed in the cold ; the one 
 using potassium hydroxide takes place only with dilute solutions. 
 The hypochlorites bleach because, when acidified, they liberate 
 chlorine. The first change, upon addition of acids, is the produc- 
 tion of hypochlorous acid, as follows t : 
 
 a. K001 + HOI =KC1 +HOC1, 
 
 b. 2 KOC1 -f H 2 S0 4 == K 2 S0 4 + 2 01 OH, 
 
 c. Ca(OCl) 2 + 2HCl =CaCL> + 2HOC1. 
 
 This portion of the reaction is like the formation of hydrochloric 
 acid from the chlorides. However, hypochlorous acid is unstable, 
 and breaks down as follows : 
 
 HOC1 = 
 
 and hydrochloric acid, in the presence of nascent oxygen, forms 
 * There is some reason to suppose that chloride of lime contains the com- 
 
 / _ /-n 
 
 pound Ca j _ Q ^j a substance which would be partly chloride and partly 
 
 hypochlorite; the relative proportions of chloride and hypochlorite would 
 
 Cl 
 remain the same as that given above, for 2 Ca _ Q C1 would contain the same 
 
 percentage of calcium, chlorine, and oxygen as Ca C1. 2 + Ca (OC1) 2 . 
 
 t A solution of hypochlorous acid in water may, if dilute, be kept for 
 some time. 
 
HYPOCHLOKOUS ACID; DECOMPOSITION. 123 
 
 chlorine and water (see pages 59, 74) ; so that, if hydrochloric acid 
 has been used to liberate hypochlorous acid, chlorine, and not 
 oxygen, will be set free : 
 
 H N Cl H Cl 
 
 d. HOC1 + HC1 = H 2 + 2C1; + \ = \ + | 
 
 H \ Cl H Cl 
 
 Now, a chloride is always formed simultaneously with a hypo- 
 chlorite if chlorine acts upon such bases as potassium or calcium 
 hydroxide ; so that where sulphuric acid is employed to liberate 
 hypochlorous acid, the following changes take place : 
 
 KOC1 + KC1 + H 2 S0 4 = K 2 S0 4 + HOI + HOC1; 
 Ca(OCl)j + CaCl 2 + 2H 2 S0 4 = 2CaS0 4 + 2HC1 + 2HOC1; 
 
 after they are set free, the hypochlorous acid and hydrochloric acid, 
 can react as in equation d-, the complete change would therefore 
 be represented by the equations : ^fc 
 
 KOC1 + KC1 + H 2 S0 4 = K 2 S0 4 + H 2 +201; 
 Ca(OCl) 2 + CaCl 2 + 2 H 2 S0 4 = 2 CaS0 4 + 2 H 2 + 4 Cl ; 
 
 Chlorine, consequently, is liberated when hydrochloric or sulphuric 
 acid is added to a hypochlorite ; and the acidified hypochlorites 
 exercise their bleaching action by reason of the liberation of that 
 element.* 
 
 Hypochlorous acid is unknown in a pure state, but its solution 
 can be prepared by passing the anhydride (C1 2 0) into water, or 
 better still, by suspending mercuric oxide in water and then sub- 
 jecting it to the action of chlorine ; by this means the formation 
 of the anhydride is avoided. . Concentrated solutions of hypochlo- 
 rous acid possess the odor of chlorine and disintegrate in the dark, 
 although they change more rapidly in the daylight. If the acid is 
 quite dilute it is much more stable, and can then even be distilled 
 without great decomposition. A solution of hypochlorous acid is 
 an energetic oxidizer; we have seen that hydrochloric acid is 
 changed to chlorine by it, and other hydrogen compounds are 
 
 * Calcium hypochlorite (bleaching powder) continually gives off chlorine, 
 when it is exposed to the air. This is probably due to the action of the car- 
 bonic acid of the atmosphere. 
 
124 POTASSIUM CHLORATE. 
 
 similarly affected. Hydrogen sulphide, for instance, is acted upon 
 as follows by hypochlorous acid : 
 
 H 2 S + HOC1 = S + H 2 + HCL 
 
 Vegetable dyes, such as litmus and indigo, are instantly bleached 
 even by dilute solutions of the acid, and many other organic sub- 
 stances are destroyed by it, so that it is much used for bleaching 
 cotton and linen goods. With silks and woollens it is useless, for 
 these it colors yellow. 
 
 The most stable acids of hydrogen, oxygen, and chlorine are 
 those which contain relatively the most oxygen, and the same is 
 true of the corresponding salts, so that hypochlorous acid and the 
 hypochlorites will tend to change into compounds with more oxygen 
 in proportion to the same amount of chlorine ; one portion of the 
 salt or acid being oxidized at the expense of the other. As a conse- 
 quence, a potassium hypochlorite solution forms potassium chlorate 
 when it. is heated to boiling : 
 
 3KC10 = KC10 8 + 2KC1, or, 
 KC10 
 
 KC10 KC1 
 
 KC10 KCl^ 3 ' 
 
 Now, 2 KOH + 2 Cl = KC1 + KC10 + H 2 0, therefore, 
 
 1. 6 KOH + 6C1=3KC10 + 3KC1+3H 2 0; 
 and when the solution is hot and concentrated, 
 
 2. 3KC10 = 2KC1 + KC10 3 ; 
 
 therefore, when chlorine is passed into potassium hydroxide solution 
 under those conditions, the result (combining equations 1 and 2), is 
 as follows : 
 
 6 KOH + 6 Cl = 5 K Cl + K Cl 3 + 3 H 2 0. 
 
 In a similar way, the solution of calcium hypochlorite changes to 
 the chlorate of calcium on heating. Potassium chlorate is much 
 less soluble than is calcium chlorate, so that potassium chlorate can 
 also be prepared by adding the solution of a potassium salt to a 
 
POTASSIUM PERCHLORATE. 125 
 
 solution containing calcium chlorate ; this method is the one used 
 for preparing the salt on a large scale : 
 
 Ca(C10 3 ) 2 + 2KC1 = 2KC10 8 + CaCl 2 . 
 
 The chlorates, when heated to a temperature considerably higher 
 than that required to effect the change from hypochlorites to chlo- 
 rates, yield oxygen and are transformed into a mixture of chloride 
 and perchlorate, one portion of the salt being oxidized at the ex- 
 pense of the other : 
 
 C1K 
 
 ( 
 
 KCUO- 
 (0 
 
 ( 
 
 |-KCUO = 
 (0 
 
 (oio 
 
 i KCllOjO 
 
 (o ; 
 
 KC10, + KC10 3 = KC10 4 + KC1 + 20.* 
 
 Finally, as potassium perchlorate is not able to take up more oxy- 
 gen, it breaks down into potassium chloride and oxygen at a low 
 red heat : 
 
 KC10 4 = 
 
 When potassium chlorate is heated, the salt melts at a moderate 
 temperature ; when the latter is increased, oxygen begins to pass 
 off; the salt again solidifies when it has changed completely to a 
 mixture of the chloride and perchlorate ; finally, it once more melts 
 at a red heat and then the potassium perchlorate parts with all 
 of its oxygen, leaving potassium chloride in the flask. (See page 20 
 and foot-note.) 
 
 (The chlorates, especially that of potassium, are but little infe- 
 rior in commercial importance to the hypochlorites. They are used 
 chiefly for their oxidizing powers, while potassium chlorate is also 
 of medicinal value. In using a chlorate, care must be taken not to 
 have the salt mixed with any substance easily oxidized ; very seri- 
 ous explosions have, for instance, resulted from grinding potassium 
 chlorate and sugar simultaneously in the same mortar. Specimens 
 of chlorate to be used for preparing oxygen should always be first 
 
 * More complicated equations are frequently given for this reaction, but 
 although there is some variation in the amount of oxygen formed, and in the 
 relative proportions of potassium chloride ahd of perchlorate left in the flask, 
 yet the probability seems to be that in the great majority of cases the simplest 
 equation, which is that given above, is realized. 
 
126 CHLORIC AND PERCHLORIC ACID. 
 
 tested on a small scale, in order to insure their safety. A demon- 
 stration of these facts is readily supplied by rubbing a trace of 
 potassium chlorate over a very small bit of sulphur in a rough 
 mortar, or by mingling some powdered chlorate with red phospho- 
 rus, by gently brushing the two substances together with a feather ; 
 when the mixture is struck with a glass rod a sharp explosion will 
 result.* The commercial application of potassium chlorate lies 
 chiefly in the preparation of fireworks and of explosive matches. 
 
 Potassium perchlorate is very nearly insoluble in cold water, 
 and therefore is of value in qualitative analysis. Because of its 
 greater stability it is sometimes used for pyrotechnic purposes in 
 place of potassium chlorate. 
 
 Chloric and perchloric acids are more easily decomposed by heat 
 than are their salts, but they possess greater stability than do 
 hypochlorous or chlorous acids. Either can be prepared according 
 to the usual method, by the addition of sulphuric acid to the corre- 
 sponding salt : - 
 
 2KC10 3 + H 2 S0 4 = K 2 S0 4 + 2 HC10 3 1 
 
 but they also are the products of decomposition of those chlorine 
 acids which contain less oxygen; these, when heated, change to 
 chloric acid, but as hydrochloric acid is formed during this change, 
 a certain amount of chlorine must also be given off, because these 
 powerful oxidizers always destroy such compounds of hydrogen. 
 Chloric acid finally changes to perchloric acid upon being heated 
 above 40: 
 
 2 H Cl 3 = H Cl 4 + H Cl + 2 0. 
 
 Of course, the oxygen formed by this decomposition further acts 
 on the hydrochloric acid, forming water and liberating chlorine, so 
 that the reaction is more complicated than the equation. (Of all 
 these acids, perchloric acid is the most stable ; its aqueous solution 
 can be distilled without decomposition, so that by this means it can 
 be separated from the sulphuric acid used in its preparation. 
 It is an oily substance, and, because it can be obtained in a pure 
 
 * Only very small quantities must be used. 
 
 t It is better to use barium chlorate, for then sulphuric acid will form 
 insoluble barium sulphate, which can be filtered off. 
 
CHLORINE TRIOXIDE; CHLORINE DIOXIDE. 127 
 
 state, best illustrates the intense capacity for oxidation possessed 
 by these chlorine compounds. This is shown by placing a drop of 
 the acid upon paper or wood, for then a violent explosion ensues. 
 The acid itself, when kept for some time, decomposes spontaneously 
 with explosive violence. 
 
 Chlorine trioxide, chlorous acid, and chlorine dioxide are the 
 only chlorine and oxygen compounds which remain for discussion. 
 
 Chlorine trioxide, the anhydride of chlorous acid, is made by the 
 reduction of chloric acid by means of arsenic trioxide.* The sub- 
 stance is a green gas with a most penetrating and irritating odor. 
 In preparing the gas the temperature must be kept quite low, 
 otherwise a most dangerous explosion may result. It forms a dark- 
 brown liquid at the temperature of snow and salt, and this liquid 
 decomposes, even if it is kept in the dark. When dissolved in 
 water, chlorine trioxide produces chlorous acid : 
 
 C1 2 3 + H 2 = 2HC10 2 . 
 
 The solution is a powerful bleaching and oxidizing agent. On 
 standing, it changes to chloric acid and hydrochloric acid, which 
 latter is further oxidized to chlorine and water. The acid neutral- 
 izes bases very slowly. The salts formed by such neutralization are 
 called chlorites, and are powerful oxidizers ; many of them change 
 to the chlorates quite readily, the nature of this alteration being, 
 in principle, the same as that accompanying the transformation of 
 chlorates into perchlorates. Potassium chlorite is converted into 
 the chlorate at 160. 
 
 Chlorine dioxide is an unstable, greenish-yellow gas, formed 
 when concentrated sulphuric acid acts upon potassium chlorate. 
 In this reaction we might expect the production of chloric 
 anhydride : - 2 H C1 ^_^ Q = ^ ^ . 
 
 for the concentrated acid would remove water from chloric acid. 
 This is not the case, however, for the chlorine pentoxide, which 
 might result, is incapable of existence, so that a part of its oxygen 
 is used in oxidizing chloric acid to perchloric acid while chlorine 
 dioxide is given off : 
 
 HC10 3 + C1 2 5 = HC10 4 + 2 C10 2 . 
 
 * 2 H Cl 3 + As 2 3 = H 2 O + C1 2 O 3 + As 2 O 5 . The H Cl O 3 can be formed 
 by adding nitric acid to potassium chlorate : K Cl O 3 + H NO 3 = K NO 3 + 
 H Cl 3 ; when potassium nitrate and chloric acid result. The existence of 
 
 p.hlorinfi trirnrirlfi is donhtfnl. ' SP.P. T.iphio-'s Anrmlpn 900 RJ. 
 
128 OXY-ACIDS OF CHLORINE ; STRUCTURE. 
 
 The gas is a most powerful oxidizing agent ; combustible substances 
 burn in it with explosive violence. This may be shown by mixing 
 some potassium chlorate with sugar * and then adding a drop of con- 
 centrated sulphuric acid, when the mass will instantly take fire. If 
 a little chlorate of potassium is placed in a deep glass, covered with 
 water, a small piece of phosphorus dropped in and then sulphuric 
 acid carefully poured directly on the salt by means of a pipette, the 
 combustion of phosphorus by means of the chlorine dioxide liber- 
 ated can be seen to take place under the surface of the water. If 
 the gas is warmed, a dangerous explosion results, so that care must 
 be taken never to heat a mixture of sulphuric acid and potassium 
 chlorate. The specific gravity of the gas shows that it has the for- 
 mula C10 2 , so that if we consider oxygen as bivalent, chlorine is 
 quadrivalent in this compound. 
 
 The generally accepted theory regarding the constitution of 
 these acids is as follows. The hydrogen atom is not attached to 
 chlorine, but to oxygen, forming a part of the hydroxyl group; 
 and this hydrogen atom is replaced by metals when the salts are 
 formed : 
 
 H Cl H Cl 
 
 Hypochlorous acid. 
 
 H Cl = 
 
 Chlorous acid. 
 
 The existence of the hydroxyl group in acids containing oxygen 
 has already been discussed. (See pages 115, 116, 117.) 
 
 The most striking characteristics of the acids composed of 
 chlorine, oxygen, and hydrogen are their intense power of oxidiz- 
 ing, their extreme instability, and the tendency which those with a 
 lesser amount of oxygen have to change into those with a greater. 
 Although the salts are, as a rule, less easily decomposed, they 
 nevertheless display similar properties. 
 
 * Do not rub in a mortar. 
 
HYPOBROMOUS ACID. 129 
 
 CHAPTER XIX. 
 
 COMPOUNDS OF BROMINE AND OF IODINE WITH OXYGEN AND 
 HYDROGEN, THE COMPOUND OF IODINE WITH OXYGEN, 
 AND THE COMPOUNDS OF THE HALOGENS WITH EACH 
 OTHER. 
 
 ALL attempts to isolate oxides of bromine have proven futile; 
 unstable as the oxides of chlorine are, those of bromine are evi- 
 dently still more so. The acids containing bromine, oxygen, and 
 hydrogen are known only in aqueous solutions; their formulae 
 correspond to those of the chlorine-acids, but bromous acid is 
 unknown, and the existence of perbromic acid is very doubtful. 
 The only compounds of bromine, oxygen, and hydrogen with which 
 we have to deal are therefore hypobromous acid, H Br 0, and bromic 
 acid, H Br 3 , while in addition we must discuss the salts derived 
 from these. 
 
 Solutions of hypobromous acid in water are produced under cir- 
 cumstances exactly analogous to those which were observed in the 
 preparation of hypochlorous acid ; such solutions have powerful 
 bleaching properties, and are very readily decomposed even by 
 slight warmth. When bromine is added to very dilute potassium 
 hydroxide solution, a liquid having bleaching properties is pro- 
 duced; the reaction is similar to that encountered in studying the 
 action of chlorine on a dilute and cold solution of caustic potash. 
 (See page 121.) 
 
 2 KOH + 2 Br = KBr + KOBr + H 2 0. 
 
 When the potassium hydroxide solution is too concentrated, 
 bromate of potassium is produced even at ordinary temperatures, 
 the conversion of hypobromites into bromates being a change much 
 more readily produced than the corresponding one with chlorine, 
 but the principle of the action is the same : 
 
 6 KOH + 6 Br = 5 K Br + K 3 Br + 3 H 2 O * 
 
 * It is not necessary to enter into the explanation of the course of these 
 reactions; the pupil should undertake this by repeating the various phases 
 given under chlorine, while substituting bromine for the latter element. 
 
130 BKOMIC ACID ; IODIC ACID. 
 
 The bromate of potassium is not very soluble, so that it can be fil- 
 tered from the solution containing the bromide and then be recrys- 
 tallized from hot water ; bromate of barium can be prepared in a 
 similar manner. The latter will yield a solution of bromic acid 
 when exactly enough sulphuric acid to form barium sulphate is 
 added : 
 
 Ba (Br 3 ) 2 + H 2 S0 4 = Ba S0 4 + 2 H Br 3 . 
 
 The solution of bromic acid is colorless, and may be concentrated 
 by evaporating the excess of water in a vacuum, but when warmed 
 the acid breaks down completely into bromine, oxygen, and water. 
 Naturally, all of the compounds under consideration are powerful 
 oxidizers ; the bromates form very explosive mixtures with oxidiz- 
 able substances, while the bromate of ammonium may even explode 
 spontaneously. 
 
 Iodine forms the pentoxide I 2 5 ,and two acids, iodic acid, 
 HI0 8 , and per-iodic acid. If the anhydride of the latter acid 
 existed, it would have the formula I 2 7 , for in the previous chap- 
 ter we saw that the theoretical anhydride of perchloric acid would 
 
 VII 
 
 be C1 2 7 ; now, by the addition of water to these anhydrides, the 
 first products would be per-iodic and perchloric acids respectively, 
 as follows : 
 
 I,0 7 + H 2 = 2HI0 4 , 
 C1 2 7 +H 2 0=2HC10 4 . 
 
 If we recall the structural formulae of these acids, it seems 
 reasonable to suppose that the oxygen atoms contained in them 
 would be capable of adding the elements of water to form hydroxyl 
 groups, in conformity with the tendency manifested in the produc- 
 tion of such groups by the addition of water to the anhydrides, and 
 in this way more complicated compounds would result : 
 
 = + H-OH 
 
 = 
 
 U-H ~ TT O-H 
 
 0-H 
 
 O-H 
 
 O-H 
 O-H 
 O-H 
 O-H 
 O-H 
 O-H 
 O-H 
 
 First acid, Second acid, Third acid, Fourth acid. 
 
 X0 4 H + H 2 = X 5 H 3 + H 2 = X 6 H 6 - +H 2 = X0 7 H 7 . 
 
PER-IODIC ACID; IODINE PENTOXIDE. 131 
 
 With the addition of each molecule of water, one oxygen atom 
 of the acid can yield two hydroxyl groups, until finally all have 
 been converted, and a complete hydroxide has been produced. 
 This last acid is called the normal acid. By separating water from 
 the normal acid, the various other acids can be formed ; so that in 
 the end we arrive at an anhydride, which is obviously identical 
 with that from which we started. None of these changes involve 
 either an oxidation or a reduction, for, if such were to take place, 
 we should produce acids derived from different anhydrides, in which 
 the valence of the characterizing element would vary; therefore, we 
 can assume that in the change from the anhydride to the first acid, 
 and in the subsequent conversion of this to the normal acid, no 
 alteration in the valence of X, in the above formulae, has taken 
 place. The process of adding water to these anhydrides is called 
 hydration, and the acids, excepting the ones with least amount of 
 hydrogen, are called hydrated acids. Hydrated acids occur quite 
 frequently, but normal acids are extremely unstable ; their existence 
 even in solution is doubtful, for when a large number of hydroxyl 
 groups are attached to the same element, they will always show a 
 great tendency to separate water ; yet often the acids lying between 
 that having the least hydrogen and the normal acid have no tend- 
 ency to break down; in fact, they are sometimes the only ones 
 which we encounter. In many of the hydrated acids only a portion 
 of the hydrogen atoms can be replaced by metals to form salts, a 
 fact not surprising if we consider that, as soon as a hydrogen atom 
 in an acid is replaced by a more metallic element, the whole com- 
 pound is rendered more positive, and therefore has its tendency to 
 take up positive elements diminished. Per-iodic acid exists in the 
 hydrated form H 5 6 1, corresponding to the third acid of the series 
 given on the table above. 
 
 The pentoxide of iodine is a white powder, which melts at 300, 
 and then instantly decomposes into oxygen and iodine ; it is pro- 
 duced by oxidizing iodine with nitric acid, or by heating iodic acid 
 for some time at 170. The oxides of chlorine are all endothermic 
 substances, but 453 K are liberated in the formation of I 2 O 5 , so 
 that, while the former compounds are explosive, the latter is quite 
 stable. From this we see that, with the diminishing not-metallic 
 character of the halogens, there appears a diminishing stability of 
 the compounds of those elements with metals, while at the same 
 
132 IODIC ACID; PER-IODIC ACID. 
 
 time an increasing stability of the oxides is manifested; but when 
 we try to draw general conclusions from these facts, we must 
 remember that the oxides of bromine are less stable than those of 
 chlorine. 
 
 Iodic acid can be prepared by oxidizing iodine, suspended in 
 water, by means of chlorine, or by adding the anhydride, I 2 5 , to 
 water. It is a crystalline solid with powerful oxidizing properties ; 
 phosphorus and arsenic, for instance, are oxidized by it respectively 
 to phosphoric and arsenic acids, and it even changes graphite to 
 carbon dioxide. 
 
 The iodates are formed either by adding a base to iodic acid ; 
 
 MOH + HI0 3 , = MI0 3 + HOH,* 
 
 or by dissolving iodine in an alkali, a reaction which is parallel to 
 that which takes place with chlorine or bromine : 
 
 6KOH + 6I = 5KI-f KI0 8 4- 3 H 2 0. oj^ 
 
 Iodic acid can be liberated by adding a non-oxidizable acid to 
 the iodates. Many of the iodates form per-iodates when heated, 
 but the latter can also be produced from the former by the addition 
 of some oxidizer. Two of the per-iodates (AgI0 4 , KI0 4 ) cor- 
 respond to the perchlorates in formula, but the majority of the 
 salts of per-iodic acid are derived from the hydrated acids ; for in- 
 stance, Na 5 10 6 from H 5 10 6 , and Ag 3 I0 5 from H 3 10 5 . Salts of 
 per-iodic acid, in which only a part of the hydrogen atoms have been 
 replaced by other metals, are known ; an example of such a salt 
 would be Na^ H 3 10 6 . Per-iodic can be isolated from its salts by the 
 addition of some other acid, and when separated from its solutions 
 by slow evaporation, it is a crystalline solid of the formula H 5 10 6 , 
 and is a powerful oxidizer. Other more complicated per-iodic acids 
 and per-iodates exist, but for their study the pupil must be referred 
 to some larger work. 
 
 The halogens can form a number of compounds with each other. 
 Three of these are produced by the union of the elements, atom for 
 atom ; they are Br Cl, bromine monochloride, an unstable liquid 
 
 * It obviously makes no difference whether we write iodic acid HIO 3 or 
 HO 3 1, sulphuric acid H 2 O 4 S, or H 2 SO 4 , etc., excepting in cases where we in- 
 tend to convey some idea in regard to the structural formulae of the acids; but 
 the method of writing the formulae of acids with the symbol of oxygen as the 
 terminal letter is the one rendered more familiar by usage. Both systems 
 are employed in this book. 
 
OXY-ACIDS OF THE HALOGENS ; TABLE OF. 
 
 133 
 
 decomposing above 10 ; I Cl, iodine monochloride, a fluid which is 
 readily decomposed by water; and IBr, iodine monobromide, a more 
 stable, crystalline solid. In addition to these, a solid trichloride of 
 iodine, I C1 3 , and a liquid, pentafluoride IF 5 , are known. Com- 
 pounds having the formulae Br Cl, I Cl, and I Br are formed by the 
 direct union of the elements ; IC1 3 , by the addition of chlorine to 
 I Cl ; and IF 5 , by the action of iodine on the fluoride of silver. All 
 of these compounds are decomposed by the addition of water, 
 although I C1 3 can exist, provided but little water is present. The 
 reaction with I F 5 is as follows : 
 
 IF 5 + 3H 2 = HI0 3 + 5HF. 
 
 The decomposition of the fluoride and the chloride of iodine by 
 means of water is a change similar to those produced by the hydra- 
 tion of other halides of not-metals. For example, phosphorus tribro- 
 mide breaks down into phosphorous acid and hydrobromic acid, when 
 it is added to wa'ter. (See page 80.) lodic acid, which is derived 
 from an oxide, I 2 5 , in which iodine is quinquivalent, is therefore 
 also formed from a fluoride of iodine in which the element is like- 
 wise quinquivalent. In comparing the formulae of I Br, I C1 3 , and 
 I F 5 , we are impressed with the fact that the more atoms of an- 
 other halogen can combine with an atom of iodine, the greater the 
 difference between the atomic weights of the two uniting elements. 
 The following is a table of the formulae belonging to the com- 
 pounds discussed in the last two chapters : 
 
 CHLORINE. 
 
 BBOMINE. 
 
 IODINE. 
 
 Oxides C1 2 O 
 C1 3 3 
 Cl O 2 
 
 Acids HO Cl 
 HO 2 C1 
 
 Acids H Br 
 
 
 
 
 
 
 
 
 H0 3 C1 
 HO 4 C1 
 
 H0 3 Br 
 H 4 Br* 
 
 I 2 5 
 
 H0 3 I 
 H0 4 IJ 
 
 * Existence doubtful. 
 
 t The hypoiodite of potassium probably exists in solution immediately 
 after adding iodine to a cold and dilute solution of potassium hydroxide; it, 
 however, soon changes to the iodate on standing. Hypoiodous acid is un- 
 known. See Colischonn ; Zeitschrif t f iir Analyt. Cliem. ; 29, 566. 
 
 J This acid exists in its hydrated form, H 5 IO 6 . The stability of all of the 
 acids and of their salts increases with increasing number of oxygen atoms. 
 They are all powerful oxidizers. 
 
134 OXIDES OF ELEMENTS OF THE SULPHUR FAMILY. 
 
 CHAPTER XX. 
 
 THE COMPOUNDS OF THE ELEMENTS OF THE SULPHUR FAM- 
 ILY WITH OXYGEN, AND WITH OXYGEN AND HYDROGEN. 
 SULPHUR DIOXIDE AND SULPHUROUS ACID. 
 
 Sulphur dioxide and sulphurous acid. Sulphur dioxide Formula, 
 S0 2 ; specific gravity, air = 1, is 2.23, H 2 =2, is 64.22 j 1 c. c. 
 at and .76 m. pressure weighs .002896 gram. 
 
 THE oxygen compounds of the elements of the sulphur family, 
 while they resemble those of the halogens to a certain extent, nev- 
 ertheless differ widely from the latter, both in their formulae and 
 characteristics. The two series of compounds resemble each other 
 chiefly because the members of both are anhydrides ; they differ very 
 greatly, however, in the ease with which they are decomposed. The 
 oxides of chlorine are explosive compounds, while that of iodine 
 is disintegrated at 300 ; on the other hand, the oxides of the ele- 
 ments of the sulphur family are quite stable and have different for- 
 mulae. The oxides of the halogens (with the exception of Cl 2 ) are 
 formed by joining two atoms of the not-metal by means of a bi- 
 valent oxygen atom, as in Cl Cl and = Cl Cl = 0, 
 while the oxides of the sulphur group have no such linking, as 
 will be seen from the formulae of the following compounds : * 
 
 S j and S -5 = 
 
 ( = 
 
 Sulphur dioxide. Sulphur trioxide. 
 
 This difference in the formulae of the anhydrides produces a dif- 
 ference in those of the acids derived from them; for, in changing 
 the oxides of the halogens into acids, the water, in order to form 
 two hydroxyl groups, attacks the linking oxygen, and thus gives us 
 acids containing but one hydrogen atom ; but in converting the 
 oxides of the sulphur family no such separation can take place. 
 
 * Oxides of sulphur with the formulae S 2 O 3 and S 2 O 7 have been de- 
 scribed, but these are comparatively unimportant. 
 
OXY-ACIDS OF ELEMENTS OF THE SULPHUR FAMILY. 135 
 
 The following formulae will make this more apparent (X denotes 
 a halogen atom and Y an atom of any element of the sulphur 
 group) : - 
 
 X H X OH ,,(=0 / =0 
 
 and 1 X ^ , TT ^ \ Q H 
 
 X H X OH H ( H 
 
 X 2 + H 2 = 2X OH and Y0 2 + H 2 = Y0 3 H 2 
 
 The above makes it plainly evident that, while the first acid 
 derived from any of the oxides of the halogens must contain one 
 hydrogen atom, those derived from the sulphur group must contain 
 two. As a general rule, the most important acids in any chemical 
 family have as many replaceable hydrogen atoms as are contained 
 in the corresponding hydrogen compounds ; for instance, H Cl and 
 HC10 3 , HI and HI0 3 , H 2 S and H 2 S0 4 , H 2 Se and H 2 SeO 3 . 
 All of the hydrated acids must bear a simple relationship to these, 
 for they are formed therefrom merely by the addition of water. 
 (See page 131.) Subsequently we shall see that the same rule ap- 
 pertains to the elements of the nitrogen family. 
 
 The elements of the sulphur family form the following oxides 
 and acids : 
 
 50 2 sulphur dioxide + H 2 = H 2 S0 3 sulphurous acid. 
 
 50 3 sulphur trioxide -j- H 2 O = H 2 S0 4 sulphuric acid. 
 Se 2 selenium dioxide -f H 2 = H 2 Se 3 selenious acid. 
 Te 2 tellurium dioxide + H, = H 2 Te 3 tellurous acid. 
 Te 3 tellurium trioxide + H, = H 2 Te 4 telluric acid. 
 
 Two oxides, S 2 3 and S 2 7 , have also been made, while a num- 
 ber of sulphur acids of less importance than the above exist. 
 Mention of these will be made at the proper time. All the ele- 
 ments of the sulphur group, on burning in air or oxygen, form the 
 dioxides ; and these can be converted into the trioxides by oxida- 
 tion, excepting in the case of selenium, the trioxide of which has 
 never been prepared. 
 
 The natural occurrence of sulphur dioxide is limited to the gases 
 which escape from the craters of volcanoes ; but, as sulphur is gen- 
 erally found in the coals used as fuel, sulphur dioxide must be a 
 product of their combustion, and hence occurs in minute traces in 
 the atmosphere of cities, although, being moist, it is rapidly ox- 
 
136 SULPHUR DIOXIDE; HISTORY, PREPARATION. 
 
 idized. The history of sulphur dioxide is as ancient as that of 
 sulphur itself ; for, as it is produced by the combustion of the latter, 
 its properties could not fail to become an object of interest. The 
 Romans were well acquainted with the disinfecting powers of burn- 
 ing sulphur and used it in cleansing their wine-skins ; sulphur 
 dioxide was confounded with sulphuric acid by the alchemists ; 
 Stahl first proved its individuality ; Priestley obtained it pure by 
 collecting it over mercury, and Lavoisier explained its composition. 
 
 The preparation of sulphur dioxide. 
 
 Sulphur dioxide can be formed either by oxidizing sulphur or by 
 deoxidizing sulphuric acid. With the first method we are already 
 acquainted, for we saw that sulphur, or combustible substances 
 containing sulphur, yield sulphur dioxide when they are burned ; 
 the second method is best employed in preparing the gas for labo- 
 ratory use. Substances such as charcoal, sulphur, and some of the 
 metals, will reduce sulphuric acid while they themselves become 
 oxidized. If charcoal is heated with sulphuric acid, carbon dioxide 
 and sulphurous acid are produced : 
 
 2 H 2 S0 4 + C = C0 2 + 2 H 2 S0 8 ; 
 
 but the latter, being an acid the anhydride of which is a gas, breaks 
 down into water and that anhydride : 
 
 Sulphur acts in a manner similar to charcoal, with the difference 
 that with it only sulphur dioxide can be formed : 
 
 2 H 2 S0 4 -f S = S0 2 + 2 H 2 S0 3 , and 2 H 2 S0 3 = 2 H 2 + 2 S0 2 , 
 so that : 
 
 2 H 2 S0 4 + S = 3 S0 2 + 2 H 2 0. 
 
 Better than either of these methods is the preparation by means 
 of copper and sulphuric acid. Cold sulphuric acid has very little 
 action on copper ; but if copper shavings are heated with sulphuric 
 acid, sulphur dioxide will be given off. 4 ' 2 Some doubt exists as to 
 the mechanism of this reaction. One explanation which has been 
 offered is as follows. When dilute sulphuric acid acts on zinc, 
 hydrogen is produced : 
 
 Zn + H 2 S0 4 = Zn S0 4 + 2 H, 
 
 but if the sulphuric acid is hot and concentrated, not hydrogen, but 
 sulphur dioxide is formed. It is therefore reasonable to suppose 
 
SULPHUR DIOXIDE; PREPARATION. 137 
 
 that the first result of the contact of zinc and sulphuric acid is 
 always the liberation of hydrogen ; but when the acid is hot and 
 concentrated the conditions are so altered that it will give up its 
 oxygen very readily. The hydrogen which is being generated would 
 then form Water with the oxygen, so that sulphurous acid would be 
 set free : 
 
 1. Zn-fH 2 S0 4 = ZnS0 4 + 2H 
 
 2. 2 H + H 2 S0 4 = 2 H 2 + S0 2 , combining 1 and 2 we have, 
 
 3. Zn + 2 H 2 S0 4 = Zn S0 4 + 2 H 2 + S0 2 . 
 
 Now, although copper produces no hydrogen with sulphuric 
 acid, yet it can be conjectured that the reaction takes place in a 
 manner similar to that given above, substituting copper for zinc. 
 There is strong reason to suppose, however, that when metals pro- 
 duce sulphur dioxide from sulphuric acid, they act exactly as do 
 carbon or sulphur, by removing oxygen without having to call in 
 the aid of hydrogen : 
 
 H, S0 4 + Cu = H 2 S0 3 + Cu 0. 
 
 The copper oxide formed, being a base, would dissolve in sul- 
 phuric acid, forming copper sulphate and water : 
 
 Cu + H 2 S0 4 = Cu S0 4 + H 2 0, 
 so that the entire reaction would be : - 
 
 2 H 2 S0 4 + Cu = Cu S0 4 + S0 2 + H 2 0. 
 
 Reactions of this kind are, however, not so simple as we are apt to 
 believe is the case. 
 
 Another method quite frequently employed in the preparation 
 of sulphurous acid, is by an addition of an acid to a sulphite ; for 
 example, 
 
 Na, S0 3 + 2 H Cl = 2 Na Cl + H 2 S0 3 . 
 
 Na 2 S0 8 + H 2 S0 4 = Na 2 S0 4 + H 2 S0 8 , 
 
 the sulphurous acid so formed then breaks down into sulphur 
 dioxide and water. This way of preparing sulphur dioxide is often 
 very convenient, for the same apparatus which is employed in 
 the production of hydrochloric acid can be used. 
 
 Sulphur dioxide is a colorless gas, with the familiar odor of a 
 burning sulphur match. The application of the moderate cold of 
 salt and snow changes it to a clear liquid, which boils at 10. 43 
 If this liquid is evaporated rapidly under the air pump, the temper- 
 
138 SULPHUR DIOXIDE; PROPERTIES, SULPHUR TRIOXIDE. 
 
 ature sinks to 68, while the liquid freezes at 76. Sulphur 
 dioxide is poisonous; when it is present in small quantities it 
 causes irritation of the throat and violent coughing ; in larger quan- 
 tities, hemorrhages from the lungs, mouth, and nose occur. Work- 
 men who are continually exposed to the gas, are affected with loss 
 of appetite and headache. Vegetation is destroyed by sulphur 
 dioxide, so that in many places very stringent laws are passed 
 regulating the working of factories, from the chimneys of which 
 sulphur dioxide escapes. Sulphur dioxide is not combustible, a fact 
 which is self-evident when we consider that it is the only compound 
 of sulphur ever formed by burning that element in the air ; it is an 
 extremely stable body, and hence will not support combustion ; in 
 fact, only in a few instances, such as in its action on sulphuretted 
 hydrogen, does it appear as an oxidizer. (See page 92.) 
 
 When sulphur burns in oxygen, no change of volume of the gas 
 occurs ; this phenomenon is like that which we observed in the de- 
 composition of sulphuretted hydrogen by means of a hot iron, for 
 that also took place without alteration of volume. The explanation 
 is the same in both cases ; in the one, each molecule of H 2 S 
 yields a corresponding molecule of hydrogen, while the volume of 
 solid sulphur produced need not be taken in consideration ; in the 
 other each molecule of oxygen takes up an atom of sulphur to pro- 
 duce a molecule of S0 2 ; so that any number of molecules of H 2 S 
 would yield the same number of molecules of H 2 , and any number 
 of molecules of 2 would yield the same number of S0 2 . 
 
 Sulphur dioxide can, under proper conditions, readily add 
 oxygen to form sulphur trioxide ; so, for instance, sulphur trioxide 
 is produced by passing a mixture of sulphur dioxide and oxygen 
 through a heated tube containing a piece of platinized asbestos, or 
 by exposing sulphur dioxide to the action of ozone. The gas can 
 add chlorine, just as it can oxygen, for sulphur dioxide and chlo- 
 rine, mixed and placed in the sunlight, produce sulphuryl chloride, 
 S0 2 C1 2 , a compound which is of considerable importance to us 
 theoretically : - S0 2 + 2 Cl = SO, CU . 
 
 Sulphur dioxide is quite soluble in water ; one volume of water 
 absorbs 45 volumes of the gas at ordinary temperatures ; the solu- 
 tion has the odor and characteristics of sulphur dioxide, and con- 
 tains sulphurous acid.* Sulphurous acid is a much more reactive 
 
 * A hydrated sulphurous acid having a definite crystalline form has been 
 
SULPHUROUS ACID; OXIDATION OF. 139 
 
 substance, chemically, than the gaseous anhydride ; it readily absorbs 
 oxygen from the atmosphere, and therefore, if left exposed for any 
 considerable length of time, will contain nothing but sulphuric acid. 
 It is consequently self-evident that oxidizing agents, such as chlo- 
 rine, bromine, or nitric acid, will change sulphurous acid to sul- 
 phuric acid with the greatest ease. Two atoms of chlorine or 
 bromine liberate one atom of oxygen from one molecule of water : 
 
 H 2 0+2C1 = 2HC1 + 0, 
 
 and one formula weight of sulphurous acid requires one atom of 
 oxygen to change it to sulphuric acid : 
 
 H 2 S0 3 + = H 2 S0 4 , 
 therefore H 2 S0 3 + 2 X + H 2 = H 2 S0 4 + 2 HX, 
 
 where X is used to designate the halogen. The oxidation of sul- 
 phurous acid by means of nitric acid and the oxides of nitrogen, 
 which is used in one of the most important commercial processes 
 known, namely, in the preparation of sulphuric acid, will be 
 discussed in connection with that substance. Sulphurous acid is 
 one of the favorite reducing agents in the laboratory, and we shall 
 frequently have occasion to refer to it as such. 
 
 When heated in a sealed tube, sulphurous acid changes to 
 sulphuric acid and sulphur : 
 
 3 H 2 S0 3 = 2 H 2 S0 4 + H 2 + S, 
 
 and in the same way the sulphites, when heated, always form sul- 
 phates by using all of their oxygen for this purpose, for instance : 
 
 4 Na 2 S0 3 = 3 Na 2 S0 4 + Na 2 S. 
 
 These reactions remind us most forcibly of those which take place 
 with the oxygen acids of the halogens, for those acids and salts 
 of chlorine, bromine, or iodine which contain the most oxygen, are 
 also the most stable. 
 
 Acids which contain one atom of hydrogen replaceable by metals 
 are termed unibasic ; those with two, dibasic; those with three, 
 tribasic ; those with four, quadribasic, and so on ; while all acids 
 with more than one replaceable hydrogen atom are polybasic ; ac- 
 cording to this nomenclature sulphurous acid is a dibasic acid. We 
 
 isolated. This hydrate has the formula H 2 SO 3 4- 6 H 2 O. This hydrate liber- 
 ates sulphur dioxide even at ordinary temperatures; it is completely decom- 
 posed at 71. Graham-Otto, vol. 4 [2], 1471. 
 
140 DIBASIC ACIDS ; PRIMARY AND SECONDARY SALTS. 
 
 are acquainted with two series of salts derived from dibasic acids, 
 accordingly as the metal replaces one or both atoms of hydrogen 
 in a formula weight of the acid. If we designate a univalent 
 metal by M', then these salts of sulphurous acid would be M' HS0 3 
 and M' 2 S0 3 , respectively; if M" denotes a divalent metal, then 
 the formulae are : 
 
 H -S0 3 
 
 M" ( = M" ( HS0 3 > and M" S0 3 . 
 
 H _1S0 3 
 
 Salts formed by replacing one atom of hydrogen in a molecule of 
 an acid by a metal, are called primary; those by replacing two, 
 secondary ; those by replacing three, tertiary, and so on ; so that 
 NaHS0 3 would be primary, and JS"a 2 S0 3 secondary sodium sulphite.* 
 When a base acts upon sulphurous acid we can consider the first re- 
 action to be as follows : 
 
 _0 H+MOH ( M+H 2 
 
 S - 
 
 (_ H ( H 
 
 S0 3 H 2 + MOH = S0 3 HM + H 2 0. 
 
 Sulphurous acid -+- a base = A primary sulphite + water. 
 More of the base acting on the primary salt would then produce 
 the secondary : 
 
 f0 M ( M 
 
 S -o -S -o 
 
 ( H + MOH ( M+H 2 O. 
 
 S0 3 MH + MOH = S0 3 M 2 + H 2 0. 
 
 Primary sulphite -f a base = Secondary sulphite + water. 
 Adding a base to a primary salt will therefore produce a second- 
 ary one; and inversely, adding more of the acid to the secondary 
 salt will produce the primary : - 
 
 S0 3 M 2 + H 2 S0 3 = 2 S0 3 MH. 
 
 That these reactions must take place is evident when we 
 consider that the two hydrogen atoms in sulphurous acid belong to 
 different hydroxyl groups, and that they therefore are as independ- 
 ent of each other as if they belonged to different acids. 
 
 Salts which contain a portion of the replaceable hydrogen of the 
 
 * Sometimes called the acid and the neutral sodium sulphite; the primary 
 sulphite is also sometimes termed sodium bisulphite. 
 
RELATIVE STRENGTH OF ACIDS. 141 
 
 acid are frequently termed acid salts, and those which have ex- 
 changed all of their hydrogen, neutral salts ; but such designations 
 are frequently misleading, for we are acquainted with salts of the 
 former class, such as the primary carbonate of sodium, NaHC0 3 , 
 which have a neutral, or even a slightly alkaline, reaction; while 
 in some cases those of the latter are acid toward litmus, as in the 
 case with aluminium sulphate, A1 2 (S0 4 ) 3 . Where a metal like 
 sodium, which has most pronounced metallic properties, replaces 
 the hydrogen of a weak acid, the resulting salt is apt to have 
 an alkaline reaction ; and, conversely, where a metal which is 
 not strongly characterized, like aluminium, is found forming a 
 salt with a strong acid, then the reaction of the salt will probably 
 be acid. But what is a weak and what is a strong acid? The 
 question is much easier to ask than to answer ; but the following 
 must at present be taken as coming nearest the truth. When an 
 acid acts on a salt it partially expels the acid combined in the lat- 
 ter, and unites with the base, and if equivalent quantities are taken, 
 the distribution ratio of the base is a measure of the affinity. 
 By equivalent quantities we mean the amount in grams expressed 
 by the formula weights representing one atom of hydrogen in each 
 acid. Thus equivalent quantities of sodium chloride and sulphuric 
 acid would be : 
 
 Na Cl, formula weight 23 + 35.5 = 58.5 
 
 iH2S04j ii 2+32+ (4X16) =49 
 
 The equivalent quantities : 
 
 58.5 grams sodium chloride and 49 grams sulphuric acid, for Na Cl and H 2 SO 4 . 
 
 " Daily experience in the laboratory teaches us that the affinity 
 of acids for bases is of such a nature as to appear a specific property 
 of the acids. When we say that carbonic acid is a weak acid and 
 sulphuric acid a strong one, we do not thereby mean that it is so 
 with respect to this or the other base, but that it is so in general." * 
 Any metal replacing the hydrogen of acetic acid, for instance, will 
 be more feebly united in the resulting salt than it would be with 
 sulphuric acid, no matter if we compare the salts of the intensely 
 metallic potassium or of the weakly metallic aluminium. If we 
 mix sodium chloride and nitric acid in aqueous solution and in 
 
 * Ostwald, Outlines of General Chemistry, p. 337 and sub. 
 
142 RELATIVE STRENGTH OF ACIDS. 
 
 equivalent quantities, the following changes will take place, until 
 an equilibrium is reached : 
 
 NaCl-hHN0 3 = NaN0 8 + H Gland 
 Na N0 3 + H Cl = Na Cl + HN0 3 . 
 
 So that NaCl, HN0 3 , NaN0 3 , and HC1 will all be present. Now, 
 the two salts, sodium chloride and sodium nitrate, will be formed 
 in equal quantities, so that hydrochloric and nitric acids are acids 
 of about the same strength. If, on the other hand, we mix sodium 
 chloride with sulphuric acid in equivalent quantities, then about 
 two-thirds of the metal will remain in sodium chloride, while one- 
 third will go to form sodium sulphate, and accordingly, sulphuric 
 acid is a weaker acid than hydrochloric acid, contrary to what 
 is generally supposed. If we use molecular formulae instead of 
 equivalent weights, then the reaction between sodium sulphate 
 and hydrochloric acid could be expressed by the following equa- 
 tion : 
 
 3 Xa 2 S0 4 + 6 H Cl = 2 H 2 S0 4 + Na 2 S0 4 + 2 H Cl + 4 Na Cl, 
 
 or, as the primary sulphate is formed in the presence of an excess 
 of sulphuric acid : 
 
 3 Na 2 S0 4 + 6 HC1 = H 2 S0 4 + 2 XaHS0 4 + 2 HC1 + 4 Nad, 
 
 so that there will be twice, as much sodium in sodium chloride after 
 the reaction, as there is in sodium sulphate; or, using the total 
 sodium as a basis of calculation, two-thirds of the sodium will go to 
 form the chloride, and one-third to form the sulphate.* The dis- 
 tribution ratio of the metal used in the above reactions would 
 therefore be as 1:1 with nitric and hydrochloric acids, and as 
 2 : 1 with hydrochloric and sulphuric acids. This conclusion, at 
 first sight, seems very strange if we remember that when toler- 
 ably concentrated sulphuric acid is added to sodium chloride or 
 sodium nitrate, sodium sulphate is formed, and either hydro- 
 chloric acid or nitric acid is given off, but we must consider that, 
 after mixing salts and acids, the least volatile acid will, upon heat- 
 ing, finally expel the more volatile ones from their salts, and, we 
 
 * Of course, if no water is present, hydrochloric acid, being volatile, will 
 pass off, .-o that no equilibrium results until all hydrochloric acid has been 
 expelled. This equation is only true, therefore, where the solution of hydro- 
 chloric acid in water renders the latter not-volatile (see pages 57 and G7). 
 
RELATIVE STRENGTH OF ACIDS. 143 
 
 must also remember, that sulphuric is much less volatile than either 
 nitric or hydrochloric acid (see page 57). It is for this reason 
 that sulphuric acid is also finally expelled from its salts when 
 they are heated with much weaker not-volatile acids, such as phos- 
 phoric or silicic. A few of the more important acids can be writ- 
 ten in the following order, judging their strength by means of the 
 relative amounts of metal which they will take from a chloride, 
 when mixed with the latter in equivalent quantities : 
 
 1. Hydrochloric acid. 6. Selenic acid. 
 
 2. Nitric acid. 7. Phosphoric acid. 
 
 3. Hydrobromic acid. 8. Hydrofluoric acid. 
 
 4. Hydroiodic acid. 9. Silicic acid. 
 
 5. Sulphuric acid. 
 
 The heat produced by neutralizing an acid with a base has appar- 
 ently nothing to do with the strength of this acid, for the greatest 
 heat is produced by neutralizing hydrofluoric acid, and the next 
 greatest with sulphuric acid ; so that if we were to take the rela- 
 tive heats of neutralization as indicating the relative strengths of 
 acids, quite a different order from the one given above would re- 
 sult, and certainly those salts which give the greatest amount of 
 heat in their formation, would, in the dry state, require the most 
 energy for their decomposition. At some future time the relation 
 between the heat of neutralization and the strength of an acid 
 will undoubtedly be discovered. 
 
144 SULPHUR TKIOXIDE; PREPARATION, PROPERTIES. 
 
 CHAPTER XXI. 
 
 SULPHUR TRIOXIDE, SULPHURIC ACID, AND THE REMAINING 
 SULPHUR ACIDS. 
 
 Sulphur trioxide ; Formula, S0 3 ; specific gravity, 1.97 at 20 ; spe- 
 cific gravity of vapor, 2.76, H 2 =2, is 79.48. 
 
 SULPHUR dioxide does not unite with oxygen under ordinary 
 circumstances, but when a mixture of the two gases is passed over 
 heated platinized asbestos, union takes place and sulphur trioxide 
 is formed. When the latter is required for use in any quantity it is 
 better to heat fuming sulphuric acid, as this substance contains a 
 large quantity of the trioxide dissolved, or else to heat an easily 
 decomposed sulphate which will form a base and sulphur trioxide. 
 Ferric sulphate is best for this purpose ; the reaction takes place as 
 follows : 
 
 Fe 2 (S0 4 ) 3 = Fe 2 8 + 3S0 8 . 
 
 The action of ozone on sulphur dioxide also produces the trioxide 
 (page 138). 
 
 Pure sulphur trioxide is a colorless liquid at ordinary tempera- 
 tures, but when gradually cooled it forms colorless prismatic crystals 
 which melt at 15. The substance boils at 46, and forms a color- 
 less vapor which has a specific gravity of 2.76, air being one, and 
 therefore, hydrogen being two, of 79.48 ; its molecular weight is con- 
 sequently 80, and the formula S0 3 . The ordinary sulphur trioxide 
 of commerce is a substance crystallizing in felt-like crystals resem- 
 bling asbestos, and this form has been taken for a second modifica- 
 tion of the body; in all probability, however, this difference in 
 appearance is due to the presence of traces of water. Pure sul- 
 phuric anhydride does not redden litmus paper nor does it attack 
 the hands ; the acid and corrosive properties only appear when, by 
 the addition of water, the substance is converted into sulphuric 
 acid. Sulphur trioxide greedily absorbs moisture from the atmos- 
 phere ; when a little of the substance is placed in water it unites 
 with the latter with a hissing noise, like that produced in immers- 
 
SULPHURIC ACID; CONSTITUTION. 145 
 
 ing a red-hot iron. When heated to a red heat, sulphur trioxide is 
 dissociated, forming sulphur dioxide and oxygen in the proportion 
 of two volumes of the former to one of the latter : 
 =0 =) r=00=| 
 
 =o + o = U = s | = o o = ; s. 
 
 = = ) " 0" ="0" 
 
 Sulphur trioxide is the anhydride of sulphuric acid and yields the 
 latter on addition of water ; S0 3 -f H 2 = H 2 S0 4 . 
 
 The constitution of sulphuric acid will best be understood if we 
 consider its formation from sulphuryl chloride, S0 2 C1 2 . We have 
 seen that sulphur dioxide and chlorine unite to form sulphuryl chlo- 
 ride (page 138). This change can be represented by structural 
 formulae as follows : 
 
 ci 
 
 = 
 = 
 
 r; 
 
 s 
 
 = 
 
 = 
 
 Cl I Cl. 
 
 Now, sulphuryl chloride, when water is added, breaks down into 
 sulphuric acid and hydrochloric acid, and it is obvious that in such 
 a reaction the chlorine atoms must be replaced by hydroxyl groups, 
 so that sulphuric acid must contain two of the latter. The follow- 
 ing formulae will make this conclusion apparent : 
 
 _C1 + H H r H + HC1 
 
 = ' I =0 
 
 =0 5 1 =0 
 
 ^_ Cl + H H I H -f H Cl. 
 
 Sulphuric acid can successively take up one and two molecules 
 of water to form two hydrated acids, H 4 S0 5 (H 2 S0 4 4- H 2 0), and 
 H 6 S0 6 (H 2 S0 4 + 2H 2 0) ; it is in the latter form that the acid 
 probably exists when a large excess of water is present, so that the 
 three oxygen atoms of sulphuric anhydride, when the latter is dis- 
 solved, finally give place to six hydroxyl groups : 
 
 O H 
 O H 
 O H 
 O II 
 O H 
 
 =0 
 
 S\ =0 = S = + H 
 
 O H 
 
 l_O H 
 
 0-H 
 3 2 = H. 2 S0 4 , + H 2 = H 4 S0 5 , + H 2 O =H SO 6 . 
 
146 SULPHURIC ACID; HISTORY. 
 
 The two acids, H 2 S0 4 -f H 2 0, and H 2 S0 4 + 2 H 2 0, are known, 
 but by far the greater number of sulphates are derived from 
 H 2 S0 4 . 
 
 Sulphuric acid is one of the most important commercial pro- 
 ducts, and the part which it plays in modern civilization is 
 fundamental. Until the use of the ammonia process, it was essen- 
 tial for making soda, and soda is required to produce both soap 
 and glass.* It is true that, if we can measure the civilization 
 of a nation by the amount of soap which it uses, the quantity of 
 sulphuric acid consumed can with greater reason be taken as an 
 indication of the stage of development arrived at by a people. 
 
 The acid has been known since the time of the Arabian alche- 
 mists, but it was first accurately described by Basil Valentine, who 
 prepared it in the fifteenth century by heating green vitriol (ferrous 
 sulphate, Fe S0 4 ) with sand. The French emigrants of the reign of 
 Louis XIV. taught the English how to prepare the acid by oxidiz- 
 ing sulphur with nitre, and a quack doctor named Ward, using this 
 method, established the first sulphuric acid factory in England, at 
 Richmond. Dr. Ward partly filled a glass vessel of about 200 
 litres capacity with water, and then placed within this an earthen 
 pot containing an iron ladle, on which was burning a mixture of 
 sulphur and saltpetre, keeping the whole tightly covered until the 
 combustion was complete. Of course, this was all very crude, and 
 the product proportionally dear, but, nevertheless, it was a vast 
 improvement on the old alchemistic method, for the price of sul- 
 phuric acid was reduced from about three dollars and twenty-five 
 cents to sixty cents a pound. In 1746 the glass vessel was replaced 
 by a lead chamber, and after this improvement exportation to the 
 Continent began, so that the acid became known as English sul- 
 phuric acid, a name which it bears to the present day. The next 
 improvement consisted in substituting steam for water and in burn- 
 ing the mixture of saltpetre and sulphur outside the chamber, 
 while the water vapor was passed through a flue together with the 
 products of oxidation, the process by this means becoming continu- 
 ous. Finally, the sulphur was burned in a separate furnace, while 
 the sulphur dioxide so formed was oxidized in the chamber by 
 means of nitric acid; and so the present continuous process was 
 evolved from Dr. Ward's glass vessel, while the price of the acid 
 sank from sixty cents to about one cent a pound. 
 
 * See p. 390. 
 
SULPHURIC ACID; MANUFACTURE. 147 
 
 The manufacture of sulphuric acid is based upon the following 
 changes : - S0 2 + H 2 = H s S0 3 , 
 
 H 2 S0 3 + = H 2 S0 4 , 
 
 and, therefore, its cheapness depends upon the ease with which sul- 
 phur dioxide can be prepared, and upon the substance used as an 
 oxidizer. It might not be unreasonable to suppose that the oxygen 
 of the atmosphere would be most available for the purpose ; but, 
 unfortunately, sulphurous acid is oxidized much too slowly for com- 
 mercial purposes by oxygen alone. If, however, we could furnish 
 some ready means of conveying oxygen from the atmosphere to the 
 sulphurous acid by the intervention of some chemical compound, the 
 problem would be solved. This is done by means of nitric acid in 
 the present continuous process for the manufacture of sulphuric 
 acid, in which the main changes taking place are as follows : 
 
 1. In a mixture of nitric acid and sulphurous acid, each com- 
 ponent acts on the other ; the sulphurous acid is oxidized to sul- 
 phuric acid, and the nitric acid is reduced to nitrous acid. 
 
 Nitrous acid, like other acids the anhydrides of which are gases, 
 at once breaks down into water and nitrous anhydride : 
 
 so that the entire change can be represented by the equation : 
 
 2. Nitrogen trioxide (N 2 3 ) with water, sulphur dioxide, and 
 oxygen forms a compound known as nitrosyl-sulphuric acid, which 
 latter is simply sulphuric acid in which an hydroxyl group is sub- 
 stituted by the group of elements N0 2 .* 
 
 f OH 
 
 n I =0 n 
 
 Ol =0 
 
 L OH. 
 
 Sulphuric acid. Nitrosyl-sulphuric acid. 
 
 * The group NO 2 is called the nitro group ; it is univalent, just as is the 
 hydroxyl group. The formula of nitrosyl-sulphuric acid may possibly be re- 
 
 f O N = 
 presented by JO ^ o . This would be sulphuric acid, in which one 
 
 [ 0-H 
 hydrogen atom has been replaced by the univalent group, N = O (the nitroso 
 
148 SULPHUKIC ACID' MANUFACTURE. 
 
 The reaction can be represented as follows : 
 
 (OH 
 
 2 S0 2 + N 2 3 + 2 + H 2 = 2 Q J 2 
 
 (N0 2 . 
 
 3. Nitrosyl sulphuric acid breaks down into X 2 3 and sulphuric 
 acid on addition of water : 
 
 fOH (OH 
 
 a. S \ 2 = S ] O 2 
 
 1 :-- : (OH + N0 2 H, 
 
 U N0 2 + H I OH 
 
 b. 2 N0 2 H = N 2 8 + H 2 0, so that, combining a and b } we 
 have : 
 
 OH 
 
 Oo + H 2 0=2H 2 S0 4 + N 2 3 . 
 
 NO, 
 
 4. The nitrogen trioxide, so regenerated; can, with steam and 
 air, once more form nitrosyl sulphuric acid, which, with water, will 
 form sulphuric acid ; so that, theoretically, an infinitely small quan- 
 tity of nitric acid, introduced at the beginning of the operation, 
 would oxidize any amount of sulphur dioxide. That this is not 
 in reality the case is due to the fact that other minor reactions, 
 producing lower oxides of nitrogen, take place ; and also because 
 the nitrogen of the air, as it takes no part in the reaction, gradually 
 dilutes the gases to such an extent as to render loss inevitable. 
 The commercial production of sulphuric acid is carried on in works 
 a simple diagram of which is shown in Fig. 9. 44 Sulphur dioxide is 
 prepared by burning either sulphur or iron pyrites (FeS 2 ) in a 
 furnace with free access of air ; the gas enters the flue (V), and is 
 conducted to the top of the tower (G) which is filled with pieces of 
 fire-brick. Two vats (&), one containing dilute sulphuric acid, and 
 the other a concentrated acid in which oxides of nitrogen are dis- 
 solved, are constantly emptying their contents into G. The concen- 
 trated acid is supplied from the tower (G'), the object of which we 
 shall see later on. Concentrated sulphuric acid can dissolve large 
 quantities of the oxides of nitrogen, but dilute acid has no such 
 power, so that mixing the contents of the two receptacles at b 
 
 group). This interpretation is rendered probable by the ease with which 
 nitrosyl-sulphuric acid is broken down by water, as bodies containing the nitro 
 group are, as a rule, more stable. 
 
SULPHURIC ACID; MANUFACTURE. 
 
 149 
 
 liberates these oxygen compounds. The acid and the oxides of 
 nitrogen mingle with the sulphur dioxide entering at a ; the hot 
 gas serves to concentrate the dilute acid, while the latter cools the 
 gas before it passes into the leaden chambers. Sulphur dioxide, 
 mixed with oxides of nitrogen, now enters the bottom of the chamber 
 1, and, in this, comes in contact with steam and the vapors of nitric 
 acid (the latter prepared by heating a mixture of sulphuric acid and 
 sodium nitrate) ; the fuel being the sulphur burning to form sulphur 
 dioxide. Sulphuric acid is formed in the first lead chamber, while 
 
 Fig. 9. 
 
 the unused gases are passed into 2, and then into 3, in both of 
 which places they come in contact with more steam, so that, in 
 these, the changes are completed. The air, which is supplied by 
 the draught of a large chimney, gives up its oxygen in going through 
 the chambers, and therefore the gases become so diluted with nitro- 
 gen . as to be no longer capable of taking part in the reactions. 
 These diluted portions are passed irr at the top of the tower G', 
 which contains pieces of coke, over which concentrated sulphuric 
 acid is constantly trickling ; here the acid dissolves the remaining 
 oxides of nitrogen which pass from chamber 3, and, saturated with 
 these gases, it can be pumped to one of the vats b above the first 
 tower, to be used as was indicated above. Experience has shown 
 that the acid collected on the floors of the chambers must not con- 
 tain more than 65-66 per cent of H 2 S0 4 , and must not have a 
 specific gravity higher than 1.5 ; but as the commercial acid has a 
 
150 SULPHURIC ACID; PROPERTIES. 
 
 specific gravity of 1.83, and must contain 89-90 per cent of H 2 S0 4 , 
 the concentration of he chamber acid is carried farther by placing 
 it in flat lead pans and evaporating the excess of water until a 
 specific gravity of 1.75 is reached ; and then, because a stronger 
 acid attacks lead, by finally completing the evaporation in platinum 
 or glass vessels. 
 
 Commercial sulphuric acid is a liquid * which, on superficial exam- 
 ination, appears like an oil ; it is usually light brown in color, owing 
 to impurities. The latter consist of sulphate of lead (which is in- 
 troduced from the lead chambers, and which is never absent), of the 
 oxides of nitrogen ( N 2 3 and N0 2 ), of hydrochloric acid, and of 
 arsenious oxide (As 2 3 ) (the latter is present because arsenic is 
 found in the minerals roasted for the preparation of sulphur diox- 
 ide). In burning iron pyrites, the iron is changed to its oxide, and 
 the sulphur to the dioxide ; and, as the pyrites frequently contain 
 selenium, selenium dioxide and selenium collect in the flues and in 
 the muddy residue at the bottom of the chambers. The brownish 
 color of commercial sulphuric acid is caused by organic substances 
 (dust, etc.) which have fallen into it. Pure sulphuric acid is pre- 
 pared from the commercial article by distilling from platinum 
 vessels. Many so-called pure acids, however, contain arsenic, 
 because arsenious oxide is volatile, and will therefore pass over in 
 the distillation, unless care has been taken to previously oxidize it 
 to arsenic acid. 
 
 The pure acid, H 2 S0 4 , is a colorless, oily liquid, with a specific 
 gravity of 1.85 at 0. Upon being cooled to 0, it crystallizes in 
 large prismatic crystals, which melt at 10. 5 ; it boils at 338, but, 
 before that temperature is reached, the acid begins to decompose 
 into sulphur trioxide and water ; this separation is perfect if heated 
 somewhat above its boiling point : 
 
 H 2 S0 4 = S0 3 +H 2 0. 
 
 Sulphuric acid has a great inclination to take up water, and in 
 so doing can form two hydrated acids : H 4 S0 5 and H 6 S0 6 . The 
 first one of these is formed when a mixture of sulphuric acid, with 
 just enough water, is cooled to a low temperature, then prisms of 
 H 4 S0 5 separate. The second, or normal, hydrate is produced by 
 
 * The name " oil of vitriol " was given to sulphuric acid by the alche- 
 mists, because of its oily appearance, and because it was first prepared from 
 green vitriol (ferrous sulphate). 
 
SULPHURIC ACID ; REACTIONS. 151 
 
 adding the requisite quantity of water to H 4 S0 5 . A large amount 
 of heat is developed in the formation of these hydrates, yet the 
 heat production does not ^ase Avhen exactly enough water to pro- 
 duce the normal hydrate ha^ been added ; it will continue until the 
 proportions are expressed by H 2 S0 4 + 1600 H 2 0, when 178 K will 
 have been developed. Sulphuric acid has such a strong tendency 
 to unite with water that it can take the elements of that compound 
 from organic substances. If it is mixed with sugar, starch, pieces 
 of wood, or similar substances which contain hydrogen and oxygen 
 in exactly the proportions to form water, it will char them as if 
 they had been burned in an insufficient supply of oxygen; for, 
 after the hydrogen and oxygen have been taken from such bodies, 
 nothing but carbon remains. In a similar way sulphuric acid will 
 attack the skin or mucous membrane, so that the concentrated acid 
 is a violent poison. 
 
 Reducing agents readily change sulphuric acid to sulphur 
 dioxide, and even to sulphur, or sulphuretted hydrogen. We have 
 studied examples of such reduction in the changes which take place 
 when hydroiodic or hydrobromic acid acts upon sulphuric acid (see 
 pages 80 and 85). As a general rule, hydrogen compounds will 
 readily reduce sulphuric acid if, like hydroiodic and hydrobromic 
 acids, they are unstable ; for instance, hydrogen sulphide, selenide, 
 or telluride will act in the following way : 
 
 H 2 S0 4 + H 2 S = 2 H 2 + S0 2 + S. 
 
 This reaction is exactly like that taking place between hydro- 
 bromic and sulphuric acids : 
 
 H 2 S0 4 + 2 H Br = 2 H 2 + S0 2 + 2 Br, 
 
 only in the one case sulphur, and in the other bromine, is produced. 
 We have already discussed the reduction of concentrated sul- 
 phuric acid by metals (page 136), so that in this place it is only 
 necessary to add that, in addition to copper or zinc, silver, mercury, 
 and a number of other metals, will produce sulphur dioxide when 
 they are heated with sulphuric acid ; but we must remember that, 
 in cases where metals are attacked by the dilute acid, hydrogen is 
 .liberated, as we saw when we discussed the preparation of that ele- 
 ment (see page 31). Other easily oxidized substances, such as char- 
 coal and sulphur, will also readily reduce sulphuric acid. Diluted 
 sulphuric acid must possess much less chemical energy than does 
 
152 SULPHATES; PRIMARY AND SECONDARY. 
 
 the concentrated liquid ; as a consequence, it will take more energy 
 to decompose it. This fact is evident when we recall the great 
 amount of heat which is given off when sulphuric acid is dissolved 
 in water. It is because of its relative instability that concentrated 
 sulphuric acid is so easily reduced. 
 
 One of the chief laboratory uses for sulphuric acid is to prepare 
 other acids by its action on the salts of the latter, and we have 
 already encountered a number of cases in which the acid was used 
 for this purpose ; for example : 
 
 2KN0 3 +H 2 S0 4 = K 2 S0 4 + 2HN0 3 , 
 2 K Cl 4 + H 2 S0 4 = K 2 S0 4 + 2 H Cl 4 , 
 Na 2 S0 3 + H 2 S0 4 = Na 2 S0 4 + H 2 S0 3 , 
 
 2 Na Cl + H 2 S0 4 = Na 2 S0 4 + 2 H Cl. 
 
 Heretofore, in studying such reactions, we have always taken 
 the formation of the secondary sulphate for granted ; but this, in 
 reality, does not take place if an excess of sulphuric acid is present. 
 Sulphuric acid is dibasic, and can therefore form two series of salts, 
 the primary (MH S0 4 ), and the secondary (M 2 S0 4 ). Now, if we 
 compare the action of sulphuric acid on sodium nitrate or sodium 
 chloride with that of the same acid on sodium hydroxide, we shall 
 see that they are analogous processes ; and, if we can, in acting on 
 sulphuric acid with sodium hydroxide, replace first one and then 
 both hydrogen atoms with the metal, it follows that we should have 
 the same changes were we to substitute sodium chloride or sodium 
 nitrate : 
 
 1. XaOH + H O^O =O_H 9 O =Xa O ^0 =O 
 
 H O^O=O H 0^0=0 
 
 Xa OH + H 2 SO 4 = H. 2 O + Xa H SO 4 
 
 2. Xa O. =O = Xa O..O=O 
 XaOH + H O^O=O H 2 O + Xa O^O=O 
 
 Xa OH + Xa HSO 4 = H 2 O + Xa 2 SO 4 
 
 1. XaCl + H O. 0=0_HC1 +Xa O. 0=0 
 
 H O > 0=O H 0^0=0 
 
 Xa Cl + H 2 SO 4 = H Cl + Xa HSO 4 
 
 2. Xa O. 0=O = Xa O^O=O 
 XaCl + H 0-^0=0 HC1 +Xa 0^0=0 
 
 Xa Cl + Xa HS0 4 = H Cl + Xa 2 SO 4 . 
 
 The above reactions are a necessary result of the fact that, in 
 polybasic acids, the hydrogen atoms are entirely independent of 
 
SULPHATES; PRIMARY AND SECONDARY. 153 
 
 each other. If, therefore, in decomposing a salt, we use an excess 
 of sulphuric acid, the primary sulphate results ; if we use an excess 
 of the salt, we produce the secondary sulphate, for, comparing the 
 two reactions : 
 
 NaCl + H 2 S0 4 = NaHS0 4 + H Cl, and 2 NaCl + H 2 SO 4 = Na 2 SO 4 + 2 HC1, 
 
 we see that, in the second one, we have twice as much sodium chlo- 
 ride in proportion to the acid as in the first. In laboratory practice 
 it is expedient to calculate the relative quantities of salt and sul- 
 phuric acid so as always to produce the primary sulphate, because 
 it is easier to fuse and more convenient to handle than the secon- 
 dary. On heating a primary sulphate * we form the corresponding 
 secondary sulphate thus : 
 
 2 Na HS0 4 = Na 2 S0 4 + H 2 S0 4 , 
 and we can accomplish the same result by adding a base : 
 
 Na HS0 4 + Na OH = Na 2 S0 4 + H 2 0. 
 
 It is obvious that the hydroxide so added may contain a different 
 metal from that already present in the salt, so that secondary salts 
 containing two metals may be formed : 
 
 Na HS0 4 +KOH = Na KS0 4 + H 2 0. 
 
 By adding sulphuric acid to the secondary sulphate we can produce 
 the primary : 
 
 Na 2 S0 4 + H 2 S0 4 = 2 Na HS0 4 . 
 
 No salts formed by replacing all of the hydrogen atoms in 
 either of the two hydrated acids, H 4 S0 5 and H 6 S0 6 , exist, but 
 some are known in which two of these have been substituted by 
 a metal ; such salts are frequently considered as being ordinary sul- 
 phates, with the additional water attached in some mysterious way 
 known as "molecular union," and so their formulae are written 
 M"S0 4 4- H 2 0, and M"S0 4 + 2 H 2 ; but it is more rational to look 
 upon these as secondary salts, M"H 2 S0 5 and M"H 4 S0 6 of the hy- 
 drated acids, H 4 S0 5 and H 6 SO G . This theory is borne out by the 
 fact that many of these salts lose water only at temperatures con- 
 siderably above that necessary to expel water of crystallization, 
 which fact seems to indicate that water, as such, is not present in 
 them, but that it is in the form of hydroxyl groups. A number of 
 more complicated salts are supposed to be derived from normal sul- 
 
 * Also termed an acid or a bisulphate. 
 
154 DISTJLPHURIC ACID; THIOSULPHURIC ACID. 
 
 phuric acid, but for information regarding these a larger manual 
 must be consulted. 
 
 Sulphur trioxide is very soluble in sulphuric acid ; the solution 
 is a liquid, having an oil-like appearance which gives off dense 
 white fumes when it is exposed to the air : 
 
 This substance, which has a composition represented by the for- 
 mula H 2 S 2 7 , is termed fuming sulphuric acid ; it is derived from 
 sulphuric acid by separating water from hydroxyl groups belonging 
 to different molecules, so that its constitution would be represented 
 as follows : 
 
 r OH HO ] r OH HO i 
 
 nO In n On 
 
 o j o ..................... o ro Oio oro 
 
 _io_ H H| oj -o- 
 
 This acid, which is termed disulphuric or pyrosulpliuric acid, 
 is, therefore, dibasic ; it is formed by linking two monovalent 
 groups, S0 3 'H, by means of a divalent oxygen atom, and its name, 
 disulphuric acid, suggests this constitution. The union of two 
 such monovalent groups by means of a polyvalent atom is a phe- 
 nomenon of quite common occurrence. On adding water to disul- 
 phuric acid, sulphuric acid is formed, and on extracting water from 
 disulphuric acid, sulphur trioxide remains; so that this acid lies 
 between sulphuric acid and its anhydride, bearing the same rela- 
 tionship to sulphuric acid as the latter does to H 4 S0 5 : 
 
 H 2 S 2 7 + H 2 = 2 H 2 S0 4 , 
 
 H 2 S 2 7 -H 2 = 2S0 3 . 
 
 We have seen that, because of the great chemical similarity 
 between the two elements, sulphur can take the place of oxygen 
 in many acids. We are acquainted with the salts of one acid 
 (tliiosulphurie acid), derived from sulphuric acid by means of such 
 a substitution ; the acid itself is not known. Thiosulphate of so- 
 dium, Na 2 S 2 3 , the most common salt of this acid, can be consid- 
 ered as sulphate of sodium, in which one atom of oxygen has been 
 replaced by one of sulphur : 
 
 Na 2 S0 4 , sodium sulphate, and Na 2 SS0 8 , sodium thiosulphate. 
 
 This compound is frequently called the hyposulphite of sodium ; 
 
THIONIC ACIDS. 
 
 155 
 
 but, obviously, such a name is not advisable, because it suggests a 
 relationship to sulphurous acid similar to that sustained by hypo- 
 chlorous acid to chlorous acid, while such a parallelism does not in 
 reality exist. Thiosulphates are changed to the sulphates by heat- 
 ing, all of the oxygen present in the thiosulphate being used to form 
 the sulphate, while the excess of sulphur unites with the excess of 
 metal to form the sulphide.* 
 
 3 = 3 
 
 Na S 
 
 When thiosulphuric acid is liberated from its salts by the addition 
 of other acids, it at once breaks down into water, sulphur dioxide, 
 and sulphur, f 
 
 Na 2 S 2 3 + H 2 S0 4 = Na 2 S0 4 + H 2 + S0 2 + S. 
 
 In addition to the ones which have been mentioned, a series of 
 acids which contain sulphur, and which have the following for- 
 mulae, exists : 
 
 1. Dithionic acid, H-O-S 
 
 = 0=) 
 
 = O O = \ S-O-H = H 2 S 2 O 6 
 
 ,=0 0=) 
 
 2. Trithionic acid, H-O-S \ = O O = > S-O-H = H 2 S 3 O 6 
 
 = = 
 
 3. Tetrathionic acid, H-O-S < = O O = > S-O-H = H 2 S 4 O ( 
 
 , =0 = 
 
 4. Pentathionic acid, H-O-S ] = O O = J> S-O-H = H 2 S 5 O 6 . 
 
 ' S S S- 
 
 * The formulae of the sulphides of some metals, notably those of the 
 alkali metals, certainly bear a most remarkable resemblance to the oxygen 
 compounds we have just been studying. Thus we have sulphides of potassium, 
 K 2 S 2 , Ko S 8 , K 2 S 4 , and K 2 S 5 , called polysulphides, the last two of which when 
 written K 2 SS 8 and K 2 SS 4 might possibly be K 2 SO 3 and K 2 SO 4 , in which oxy- 
 gen is replaced by sulphur ; as, however, the parallelism does not extend 
 beyond the mere relationship in the number of atoms, and as we have no 
 knowledge of the structural formulae of the polysulphides, this interpretation 
 is purely speculative. (See Drechsel; Journal fur Praktische Chemie, 
 [2] 4, 20.) 
 
 t The liberation of sulphur on addition of acids distinguishes thiosul- 
 phates from sulphites. 
 
156 SULPHUR; HALIDES OF. 
 
 These acids, in all probability, contain the univalent group of 
 elements : 
 
 ( =0 
 
 -SJ-o 
 
 (_0 H, 
 
 which also occurs in disulphuric acid (see page 154). Two of 
 these groups are united in dithionic acid, while in the remaining 
 three they are joined by means of sulphur atoms, as is shown by 
 the formulae. Trithionic acid, therefore, is disulphuric acid in which 
 the linking oxygen atom is replaced by one of sulphur. A larger 
 text-book must be consulted for the methods of preparation and 
 general characteristics of these acids. The constitutional formulae 
 of the compounds which have been discussed in that part of the 
 work following sulphuric acid, express the present state of our 
 knowledge, but the whole subject will bear further investigation. 
 Constitutional formulae are constructed with the view of so arran- 
 ging the grouping of the atoms in them that all of the reactions 
 entered into by the substances which they represent can be ex- 
 plained by them, and further, they also frequently indicate the man- 
 ner in which these substances are formed. As soon as any facts 
 -contradicting a structural formula in general use are discovered, 
 the formula must either be abandoned or, at least, so modified- as to 
 agree with the new discovery. 
 
 As might be expected, sulphur can form unstable chlorides by 
 direct union with chlorine. The first product of the action of 
 chlorine on sulphur is the inonochloride, S 2 C1 2 , if the action of 
 chlorine is continued, the dichloride, SC1 2 , is formed, and lastly, 
 if a large excess of chlorine acts on the dichloride at a temperature 
 of 22, the tetrachloride, S C1 4 , results.* This latter compound 
 decomposes, when warmed above 22, liberating chlorine, while 
 the dichloride is only stable below 6 to 10. Corresponding com- 
 pounds, S 2 Br 2 and S 2 I 2 , exist ; and an iodide, S I e , is also described.! 
 
 Compounds of sulphur which contain both chlorine and oxygen 
 are derived from sulphurous and sulphuric acids, by substituting 
 chlorine for hydroxyl groups. They are termed acid chlorides, and 
 
 * Compare the behavior of this chloride with the unstable chloride of 
 manganese (MnCl 4 ), which is supposed to be formed when hydrochloric acid 
 acts on manganese dioxide. 
 
 t Lamers; Jour, fur Prakt. Chem.; 84, 349. 
 
SULPHUR; ACID CHLORIDES OF. 157 
 
 are SOC1 2 , thionyl chloride; S0 3 C1H, sulphury 1-hydroxyl chloride 
 (chlor-sulphonic acid); and S0 2 C1 2 , sulphuryl chloride. 
 
 _0 H rCl 
 
 H 
 
 Sulphurous acid. Thionyl chloride. Sulphuric acid. Chlor-sulphonic Sulphuryl 
 
 acid. chloride. 
 
 Following the law which we found to be general with the chlorides 
 of the not-metals, these compounds are decomposed into the corre- 
 sponding acids by addition of water ; thionyl chloride forming sul- 
 phurous acid, and the last two both yielding sulphuric acid. The 
 decomposition of sulphuryl chloride in this way not only has an 
 important theoretic bearing in the constitution of sulphuric acid, 
 (see page 138), but what is more, the formation of S0 3 Cl H from 
 hydrochloric acid and sulphuric anhydride : 
 
 = 
 
 = 
 
 = + HC1 (Z&~* 
 
 shows a resemblance between hydrochloric acid and water in chemi- 
 cal behavior ; for the following reaction ( 2 ) is clearly analogous to 
 reaction 1 : 
 
 c\ 
 
 The existence of these compounds illustrates the similarity be- 
 tween hydroxides of the metals and of the not-metals, for in both 
 classes of compounds the hydroxyl groups can be replaced by chlo- 
 rine ; though with metals this substitution is much more easily 
 brought about than with not-metals. The hydroxide of the metal 
 has only to be treated with hydrochloric acid in order to form the 
 very stable chloride : 
 
 K-0-H + HC1 = KC1 + H-0-H, 
 
 while in forming the acid chlorides some roundabout method, which 
 excludes the presence of water, must be resorted to ; for these com- 
 pounds are all decomposed by the latter substance. 
 
158 SULPHUROUS ACID; STRUCTURE OF. 
 
 Chemists have lately believed sulphurous acid to have the 
 constitution : 
 
 r = r H 
 
 I. H R-<=0 and not II. 
 
 U (_0 H, 
 
 -0 H, 
 
 so that, if this view is correct, an analogy between that acid and 
 thionyl chloride does not exist. This interpretation, which has 
 been given to some experiments made with compounds belonging 
 in organic chemistry, seems to be unnecessary ; it is more probable 
 that the replaceable hydrogen in all of the oxy-acids which we 
 have studied is present in the hydroxyl groups, so that the for- 
 mula of sulphurous acid would be as is shown above. (II.) 
 
 Three other oxides of sulphur, S 2 3 , S 2 7 , and S0 4 , are 
 known. The first two very readily decompose, the former into 
 sulphur and sulphur dioxide (2 S 2 3 = S -f 3 S0 2 ), and S 2 7 into 
 sulphur trioxide and oxygen (S 2 7 = 2 S0 3 -f- 0). 
 
 The following table will make the relationship between the sul- 
 phur acids more apparent : 
 
OXY-ACIDS OF SULPHUR ; TABLE OF. 
 
 159 
 
 .0000 
 
 s t 
 
 a 1 ii ii I 
 
 3-^000 
 
 -J I 
 
 PPPP 
 
 go 
 
 o 
 
 %% g 
 
 ffff 
 
 OOO 
 
 5 
 
 o * 
 ? 
 
 H 
 
 p ff s g 
 
 ; 
 
 P + 5 P + 
 ^ w $ w 
 
 S" 
 
 o 
 
 ?; 
 
 o a 2. 
 
 ^ 02, 
 5 
 5' 
 
 1 
 
 1 
 
 ?* 3" 1 
 
 -- n 
 
 2 
 
 c 
 
 GO 
 
 + g.0 
 
 
 3 3 
 
 3 
 
 'Sa 
 
 QQ 
 
 O 
 
 02 * 
 
 O + 
 
 5n I 
 
 ?:$?$ 
 
160 SELENIUM DIOXIDE; TELLURIUM DIOXIDE. 
 
 CHAPTER XXII. 
 
 THE COMPOUNDS OF SELENIUM AND TELLURIUM WITH 
 OXYGEN, AND WITH OXYGEN AND HYDROGEN. 
 
 SELENIUM and tellurium do not manifest so great a variety in 
 the formation of oxides and acids as sulphur 5 but such oxides as 
 they do form are constructed in a manner analogous to the structure 
 of the compounds of the latter element. They each, on burning, 
 yield the corresponding dioxide : 
 
 Se + 2 = Se 2 , and Te + 2 = Te 0,, 
 
 and both the dioxides are solid bodies ; in fact, all of the oxides of 
 the not-metals having higher atomic weights are solids. Sulphur 
 dioxide, in solution, is a powerful reducing agent, always showing 
 the greatest tendency to take up oxygen so as to form sulphuric 
 acid, while, on the other hand, the dioxides of selenium and tellu- 
 rium exhibit a behavior which is exactly the reverse ; they part with 
 their oxygen so readily that even the particles of dust which may 
 come in contact with them serve as reducing agents, and so liberate 
 selenium or tellurium. Of course, sulphurous acid readily reduces 
 the two oxides, while it is changed to sulphuric acid : 
 
 2H 2 S0 3 + Se0 2 = 2H 2 S0 4 -f Se, 
 2 H 2 S0 3 + Te 2 = 2 H 2 S0 4 -f Te, 
 
 and, as a consequence, selenium, and not selenium dioxide, is found 
 in the lead chambers of sulphuric acid works. (See page 102.) 
 
 Selenium dioxide is a white, snow-like solid, which is best pre- 
 pared by oxidizing selenium with nitric acid. It does not melt ; 
 but, on heating, it changes directly from a solid to a gas, at a 
 temperature below the boiling point of sulphuric acid. Selenium 
 dioxide greedily absorbs moisture from the air, forming selenious 
 acid. The latter is a dibasic acid, giving primary and secondary 
 salts, MH Se 3 and M 2 Se 3 . 
 
 Tellurium dioxide is produced by the oxidation of tellurium 
 with nitric acid. It is a colorless, crystalline solid, which is but 
 little soluble in water. As tellurium has such a high atomic 
 
SELENIC ACID; TELLURIC ACID. 161 
 
 
 
 weight, its lower oxides begin to bear some resemblance to the 
 oxides of the less pronounced metals ; the dioxide has therefore 
 scarcely any tendency to unite with water, nor is it dissolved by 
 weak bases ; an additional sign that tellurium is approaching the 
 metals in its character is the fact that the tetrachloride, Te C1 4 , is 
 not completely decomposed by cold water. Tellurous acid is a 
 white powder which, when warmed, readily decomposes into tellu- 
 rium dioxide and water. It is a dibasic acid, and forms primary 
 and secondary salts, MH Te 3 and M 2 Te 3 . 
 
 Selenium trioxide has never been prepared, but the correspond- 
 ing acid, Ho Se 4 , selenic acid, can be formed by oxidizing selenious 
 acid with chlorine or bromine, just as the corresponding sulphur 
 compound is oxidized by the same elements. Solutions of selenic 
 acid cannot be concentrated beyond the point at which they contain 
 95 per cent of H 2 Se 4 , because, after this percentage is reached, 
 the acid breaks down into selenious acid and oxygen. The rule 
 with the selenium acids is therefore exactly the reverse of that 
 which held true as regards the halogen and the sulphur compounds, 
 for in dealing with those, the acids containing the most oxygen are 
 the final products formed by heating the ones containing the lesser. 
 Selenic acid is dibasic, and forms primary and secondary salts, 
 MH Se 4 and M 2 Se 4 . 
 
 The properties of telluric acid are much like those of selenic 
 acid, with the exception that the anhydride Te 3 is known. The 
 latter compound is produced by heating telluric acid, which breaks 
 down into its anhydride and water, H 2 Te 4 = Te O 3 + H 2 0. The 
 oxygen compounds of the sulphur family are exactly parallel with 
 those of the halogen family as regards stability. Thus, fluorine 
 form no oxide, and the oxide of oxygen (ozone) is quite unstable ; 
 chlorine forms a greater variety of oxides than does any other 
 element of the halogens, while sulphur exhibits the same prop- 
 erty in its own family. The third members of either family, 
 selenium and bromine, are the elements which, next to the first 
 member of each family, have the least tendency to form oxides ; 
 while with iodine and tellurium the capacity to form stable oxides 
 is once more manifested. 
 
 Both selenium and tellurium can form compounds with sulphur ; 
 but, while those of selenium, excepting the simplest (the mono- 
 sulphide, Se S), are not very definitely understood, those of tellu- 
 
162 OXIDES OF SULPHUR FAMILY; THERMOCHEMISTRY OF. 
 
 rium exactly correspond to the oxygen compounds of that element, 
 so that we have a disulphide, Te S 2 , and a trisulphide, Te S 3 ; and 
 this latter compound resembles the trioxide, Te 3 , so far as to be 
 the anhydride of an acid, for with sulphides, like those of sodium or 
 potassium, it will form salts in which all of the oxygen is replaced 
 by sulphur : 
 
 K 2 S + TeS 3 =K 2 TeS 4 , 
 = K 2 Te0 4 . 
 
 The existence of such a compound as K 2 Te S 4 is another proof of 
 the great resemblance between sulphur and oxygen. 
 
 The chlorine compounds of selenium and tellurium are like 
 those of sulphur. 
 
 The following table contains the thermo-chemical relations which 
 exist between the oxides of the sulphur family and between the 
 acids, HX0 3 , of the halogen family. Unfortunately, the heats of 
 formation of the oxides of the halogens have not been accurately 
 determined, so that in this table we are compelled to compare ox- 
 ides with compounds of oxygen and hydrogen : 
 
 
 
 STABILITY. 
 
 S0 2 
 Se0 2 
 Te O 2 
 
 710 K 
 572 K 
 773 K 
 
 H Cl O 3 
 HBr0 3 
 HI0 3 
 
 239 K* 
 160 K* 
 582 K* 
 
 x 
 
 HEATS OF FORMATION OF THE CHLOBIUES. 
 
 STABILITY. 
 
 S 2 C1 2 
 Seo C1 2 
 
 143 K 
 222 K 
 
 
 A 
 
 SeCl 4 
 Te C1 4 
 
 462 K 
 774 K 
 
 
 
 The heats of formation and stability of the chlorine compounds 
 increase with the atomic weights, and therefore increase with the 
 decreasing not-metallic character of the elements. This rule is 
 reversed with the hydrogen compounds, so that, in this family, the 
 chlorine and the hydrogen compounds are the less easily decom- 
 posed, the greater the chemical contrast between the elements 
 which form them. 
 
 * Dissolved in water- 
 
NITROGEN; HISTORY. 163 
 
 CHAPTER XXIII. 
 
 NITROGEN AND THE ATMOSPHERE. 
 
 Nitrogen, symbol, 1ST; atomic weight 14.03; specific gravity, air 
 = 1, is .97209, H 2 = 2, is 28 ; 1 c.c. at emd .76 m. we^fo 
 .00125718 0r-ram. The Atmosphere, specific gravity, 1 ; specific 
 gravity, H 2 = 2, *s 28.8 ; 1 c.c. air at awd .76 m. pressure 
 weighs .00129327 
 
 As the atmosphere is simply a mixture of gases, with nothing 
 in its deportment from which general chemical comparisons could 
 be made, it seems advisable to treat of it in a separate chapter ; and 
 subsequently to introduce a discussion of the characteristics which 
 the elements of the nitrogen family have in common. A descrip- 
 tion of the element nitrogen must, however, be given before the 
 atmosphere is considered. 
 
 An English chemist, Rutherford, noticed in 1772 that a closed 
 space of air, in which an animal had breathed for a sufficient length 
 of time, was then unfit further to support either combustion or life, 
 even after care had been taken to remove all of the carbon dioxide 
 by means of lime water or caustic potash ; and from this fact he 
 drew the conclusion that the residuum contained a peculiar kind of 
 air. When, subsequently, it was discovered that elements which 
 form solid oxides, when burned in a closed space of air, leave 
 nothing but this same gas, the general consensus of opinion among 
 chemists was that this gas must be phlogisticated air j for the burn- 
 ing substance had given up its phlogiston^ and that portion of the 
 atmosphere which remained must, as a consequence, have taken it 
 up ; oxygen, in which substances burned so readily, was therefore 
 most certainly pure dephlogisticated air. These theories as regards 
 the nature of the atmosphere were unfortunate for the older school 
 of chemists, for it was not difficult to prove that carbon, in burning 
 in oxygen (or, as they said, in giving up its phlogiston), formed a 
 kind of phlogisticated air which was different from the gas which 
 
164 NITROGEN; OCCURRENCE, PREPARATION. 
 
 we now call nitrogen, but which was likewise unable to support 
 either life or combustion. Lastly, hydrogen, which was supposed 
 to be pure phlogiston, when burned in dephlogisticated air forms 
 water ; therefore that liquid should also be phlogisticated air, yet 
 it differs entirely from both of the other forms. As a consequence 
 of these facts, chemists were brought face to face with the neces- 
 sity of assuming the existence of three different kinds of phlogisti- 
 cated air. It was left for Lavoisier to prove the fallacy involved in 
 these theories and to show that the atmosphere contained two dis- 
 tinct gases, one of which he called oxygene, the other azote (from a 
 and L'amx6g, to sustain life). The name nitrogene (from nitrum, 
 saltpetre, and the root yfv, to produce) was introduced by Chaptal, 
 and from this the English name, nitrogen, has resulted. The word 
 azote is still used by French chemists. 
 
 Nitrogen occurs, as such, in the atmosphere, and small quanti- 
 ties of it are also dissolved in water ; combined, it is found in the 
 form of sodium nitrate or Chile saltpetre, large deposits of which 
 occur in the northern provinces of Chile ; and it also is present as 
 ammonium compounds, the latter being found in small quantities 
 in the air and in the soil ; nitrogen is also an essential constituent 
 of many organic substances of both animal and vegetable origin. 
 Despite its wide distribution, nitrogen forms scarcely one per cent 
 of the total substance of the globe, even if we include in this its 
 gaseous envelope, for very little, if any, nitrogen is found in the 
 older geologic formations. 
 
 The preparation of nitrogen. 
 
 While we are acquainted with no simple method by which to 
 remove the nitrogen from the atmosphere so as to obtain the 
 oxygen contained therein, the removal of the oxygen is a com- 
 paratively simple matter ; to do this it is only necessary to burn 
 some substance in a closed volume of air. In preparing nitrogen, 
 it is, of course, expedient to combust a body which will form a 
 solid oxide, and which will therefore leave no gaseous residuum, 
 excepting nitrogen. Phosphorus answers this purpose admirably, 
 for when it is burned it forms phosphorus pentoxide, a solid readily 
 soluble in water. 45 Another method for preparing nitrogen is by 
 passing air over copper foil heated in an infusible glass tube by 
 means of a combustion furnace ; the copper then becomes oxidized, 
 copper oxide, Cu 0, being formed, while the unchanged nitrogen 
 passes on, to be collected over water. 
 
THE ATMOSPHERE; COMPOSITION. 165 
 
 Nitrogen is a colorless, odorless gas, with a specific gravity, air 
 being one, of .97209, and hydrogen being two, of 28, so that, as 
 the atomic weight of the element is 14, the molecule of nitrogen 
 consists of two atoms, just as is the case with oxygen, hydrogen, 
 and chlorine. Its critical temperature is 146. 3 at a pressure 
 of 35 atmospheres. This means' that at 146. 3 nitrogen becomes 
 fluid if a pressure of 35 atmospheres is exerted, but that above this 
 temperature pressure cannot condense it to a liquid. The boiling 
 point of liquid nitrogen is at 194, with a pressure of 740 m. m. ; 
 the freezing point is at 214. Nitrogen is scarcely soluble in 
 water, for one volume of the latter dissolves only .015 volumes of 
 the gas at ordinary temperatures. 
 
 Chemically, nitrogen is a remarkably indifferent substance when 
 it is in the free state ; this circumstance is very striking when we 
 consider that the same element, when combined with others, can 
 take part in a great number of chemical reactions, and so manifest 
 the greatest chemical activity. Nitrogen will unite with other 
 elements only under the greatest provocation ; so, for instance, 
 nitrogen and hydrogen will form ammonia^ under the influence of 
 an electric current, and nitrous acid, nitric acid, and ammonium 
 nitrate are formed in small quantities when hydrogen burns in air. 
 Owing to its chemical indifference, nitrogen will neither burn nor 
 will it support combustion ; and, naturally, animals are asphyxiated 
 when placed in the gas. 
 
 Toward the latter part of the eighteenth century, when chemists 
 began to inquire into the nature of the atmosphere, the determina- 
 tion of the exact quantities of nitrogen and oxygen going to form 
 that gas became an interesting subject for research. Lord Caven- 
 dish first made a series of accurate investigations of the air in and 
 about London, and came to the conclusion that there were about 79 
 out of every 100 volumes of air left as phlogisticated air after 
 combustion had taken place. Lavoisier, of course, also studied the 
 subject, but with very uncertain results, for at one time he gave 
 the volume of oxygen as one-fourth, at another time as one-fifth of 
 the total atmosphere. Gay Lussac and Humboldt made some really 
 accurate determinations of air in the neighborhood of Paris, and 
 they found that there were between 20.9 and 21.2 volumes of oxy- 
 gen in every 100 volumes of air ; and, subsequently, a great number 
 of chemists, notably Dumas, Regnault, and Bunsen, made a long 
 
166 THE ATMOSPHERE ; CARBON DIOXIDE IN. 
 
 series of accurate determinations, the results of which showed that 
 the atmosphere contained 20.9 volumes of oxygen and 79.1 volumes 
 of nitrogen in every 100 volumes, but that these quantities were 
 subject to frequent, but very slight, variations. The following 
 figures will show the extent of these differences : 
 
 Lyons, Berlin, 20.9 volumes of oxygen in 100 volumes of air. 
 Algiers, 20.4 " " " " 100 " " " 
 
 Calcutta, 20.3 " " " " 100 " " " 
 
 Atlantic Ocean, 21.5 " " " " 100 " " " 
 
 These variations, which are too great to be within the range of 
 experimental error, indicate that the atmosphere is a mechanical 
 mixture, and not a chemical compound. Another proof of this 
 conclusion can be obtained by examining the difference in the solu- 
 bility of oxygen and of nitrogen in water. A given volume of 
 water will dissolve quite a little more oxygen than it will nitrogen, 
 so that, if we place some water which has been exposed to the air, 
 under the bell of an air pump and then exhaust, the bubbles of 
 gas which pass from the liquid will be a mixture of oxygen and 
 nitrogen, containing more of the former element than does the air. 
 The gas so formed can again be dissolved in water, and the water 
 once more exhausted ; and if the operation is repeated often enough, 
 with the same gas, the latter will finally be almost pure oxygen. 
 The oxygen and nitrogen of the air can therefore be separated by 
 simple, not-chemical, means. The discovery that oxygen and nitro- 
 gen can be mixed to form air, without a change either in the total 
 volume or in the amount of heat contained in the two gases, forms 
 a final argument against the conception of the atmosphere as a 
 chemical compound. 46 
 
 Oxygen and nitrogen form the bulk of the atmosphere, but 
 other substances are always present in minute quantities. The 
 most important of these impurities are carbon dioxide, ammonium 
 carbonate, nitrate and nitrite, and water vapor, as well as solid 
 particles of dust, which are both organic and inorganic in their 
 origin. Carbon dioxide, which is invariably found in the air, is as 
 important to living organisms as oxygen itself. It is of far greater 
 value than is the uncombined nitrogen, for the latest investigations 
 show that the admixture of the latter gas does not make the air 
 more adaptable for respiration, animals being able to live as com- 
 fortably in pure oxygen as they do in air; on the other hand, 
 
ATMOSPHERIC CARBON DIOXIDE; REACTIONS. 167 
 
 carbon dioxide is a plant food, and is absolutely necessary for 
 vegetable life. Carbon dioxide is being constantly added to the 
 atmosphere from burning fuel, from volcanic craters and fissures, 
 as a product of the breathing of animals and from decomposing or- 
 ganic matter, because it is produced by the combustion or decay of 
 carboniferous substances. If no means were provided for the 
 removal of the atmospheric carbon dioxide, the increase in the 
 amount of the latter would soon destroy all living organisms de- 
 pendent upon -respiration. Fortunately, plants growing in the sun- 
 light absorb carbon dioxide from the air, using for this absorption 
 a green coloring matter which is contained in the leaves, and which 
 can eliminate oxygen from, and add hydrogen to, carbon dioxide. 
 By this means a substance which is able to form all of the innu- 
 merable compounds of carbon, hydrogen, and oxygen occurring in 
 the vegetable kingdom, is produced.* 
 
 The carbon dioxide of the atmosphere is, therefore, continually 
 being removed while oxygen is being returned ; but were this pro- 
 cess to go on without any compensating production of carbon diox- 
 ide, plant life, and consequently animal life, would soon cease. 
 The supply of carbon dioxide is, however, renewed by one means 
 or another, so that the quantity in the atmosphere remains quite 
 constant. In former geologic periods the atmosphere was undoubt- 
 edly much warmer and contained much greater quantities of carbon 
 dioxide than it does now. Hot rains were continually pouring 
 
 * According to the theory proposed by Baeyer, and which is held by most 
 chemists, this substance is formic aldehyde, CH 2 O ; or 
 
 ( = H 2 
 
 . =0 ( = 
 
 Formic aldehyde can be considered as carbon dioxide in which one atom of 
 oxygen has been replaced by two of hydrogen ; it therefore contains the ele- 
 ments, hydrogen and oxygen, in exactly the proportions necessary to form 
 water, while sugars, starch, and cellulose also contain the elements in the 
 same proportions. Glucose has the formula C 6 H 12 O 6 = 6 (C H 2 O), so that it 
 can very possibly be formed by simple condensation of six formic aldehyde 
 molecules ; and, indeed, a substance very nearly identical with glucose has been 
 made artificially by this means. Cane sugar, starch, and cellulose are made 
 from C 6 H 12 O 6 by separation of water between the molecules; so that the 
 theory of the reduction of carbon dioxide by plants to form formic aldehyde 
 seems very reasonable. The hydrogen for this reduction is, probably, fur- 
 nished by the decomposition of water, the oxygen of which is eliminated. 
 
168 ATMOSPHEKIC CARBON DIOXIDE ; REACTIONS. 
 
 down and copiously watering the continents and islands which had 
 been formed, so that vegetation on a gigantic scale nourished 
 wherever the soil was favorable. Our coal beds were produced by 
 the destruction of the flora of that period ; by this means enormous 
 quantities of carbon dioxide were removed from the atmosphere, 
 and the carbon stored for use when the supply of that gas should 
 no longer be sufficient to support vegetation. After the atmosphere 
 had assumed the composition which it possesses at present, animal 
 life flourished. The plants take up carbon dioxide from the atmos- 
 phere, form their tissues therefrom, animals live on them or prey 
 on each other ; both plants and animals die, decay, and the carbon 
 dioxide once more finds its way into the atmosphere ; so that a con- 
 tinuous metamorphosis, with its energy given by the light and heat 
 of the sun, is in progress. Small quantities of carbon dioxide are, 
 however, lost in the formation of carbonates of the metals, because, 
 in disintegrating, the silicates, which form the main body of the 
 rocks, liberate the bases which they contain, so that these can unite 
 with other acids (of which carbonic acid is one), and thus a certain 
 amount of the latter substance is permanently removed from the 
 atmosphere. This loss has been supplied by the carbon dioxide of 
 volcanic origin and by that formed in burning the fuel which was 
 stored as coal in a former geologic era ; so that, as far as we know, 
 the total amount of carbon dioxide in the atmosphere is not dimin- 
 ishing ; if it is growing less, the rate of decrease is so very slow 
 that, in the short time during which chemists have been able to make 
 accurate observations, no change could be noted. The quantity 
 of carbon dioxide in a given volume of the air varies slightly ; but, 
 normally, it is about three parts in ten thousand, and it seems that 
 the proportion of carbon dioxide is greater at night than in the 
 daytime, and in summer than in winter. In the higher regions of 
 the atmosphere, where vegetation is impossible, the amount of the 
 gas may even increase to eleven parts in ten thousand,* while a con- 
 tinued rainstorm may diminish it to two and a quarter. The 
 amount of carbon dioxide in crowded rooms is increased by the 
 breathing of the people within the closed air space, yet this does 
 not generally take place to such an extent that the oppressive feel- 
 ing caused by such an atmosphere can be ascribed entirely to it; 
 the unpleasant effect is due to exhalations of organic matter which 
 pass from the lungs. The presence of carbon dioxide in the atmos- 
 
 * Doubtful. 
 
THE ATMOSPHERE; WATER VAPOR IN. 169 
 
 phere can be proved at any time by exposing some clear lime water 
 to the action of the air, for a white crust of the carbonate of cal- 
 cium will be formed in a short time. 47 
 
 Water vapor is always present in the atmosphere in quantities 
 varying with the temperature, season of the year, and locality ; it is 
 just as important as carbon dioxide to living organisms. The evap- 
 oration of oceans, lakes, and rivers furnishes a never-ending supply 
 of water, the amount of which is generally greater in hot than in 
 cold weather, and greater by day than by night. 
 
 The higher the temperature of a gas, the greater will be the 
 amount of water vapor which it can take up, for the quantity of the 
 latter which can be contained in a closed space (either a vacuum or 
 filled with gas) increases with the temperature, but is an unalterable 
 amount at any definite point. If as much water vapor as can possi- 
 bly be present at the existing temperature is contained in a gas, the 
 latter is said to be saturated with that vapor, and no further evapo- 
 ration of water can take place unless the temperature is increased ; 
 on the other hand, a decrease would diminish the amount of vapor 
 which can be present, so that a portion of the moisture would be 
 precipitated as water. From these considerations it follows that, if 
 the atmosphere is nearly saturated with moisture, any diminution 
 in the temperature will cause a fall of rain or the formation of dew, 
 while at the same time no evaporation can take place when such 
 a condition prevails. The amount of water in the atmosphere is 
 generally greatest near the seashore ; for, owing to changes in tem- 
 perature, much of the water will be precipitated before the moisture- 
 laden air can pass far inland. Drops of water collect on a cold sur- 
 face, because the air in the immediate neighborhood is cooled below 
 the point at which it is saturated with vapor. This point is called 
 the dew point ; and, as the exact amount of water vapor which can 
 be contained in a given volume of the atmosphere at any definite 
 temperature is known, the discovery of the dew point affords a 
 ready means of ascertaining the amount of moisture in the air. 
 The ratio between the tension of the water vapor which would be 
 found were the air fully saturated at the prevailing temperature, 
 and that tension which really exists, is called the relative humid- 
 ity.* The quantity of water present in the atmosphere can also be 
 
 * By vapor tension at a given temperature is meant the pressure, in milli- 
 meters of the barometer, which is exerted by a vapor at that temperature. 
 
170 THE ATMOSPHERE ; AMMONIA IN. 
 
 ascertained by passing a known volume of air over weighed tubes 
 filled with calcium chloride or any other substance which will 
 readily absorb moisture, for then the gain in weight will give the 
 exact amount of moisture which is present. 
 
 The water in the atmosphere is absolutely essential to plant life. 
 The liquid falls upon the soil as rain, and is then absorbed by the 
 radicles; afterward it circulates through the entire system of the 
 plant, taking part in various physiological changes, and finally 
 evaporates from the leaves. The amount of moisture which passes 
 from large areas covered by vegetation is enormous, so that wooded 
 districts cause an equitable distribution of rain. 
 
 Another impurity present in the atmosphere is ammonia; this 
 substance, however, is always found combined with acids, as ammo- 
 nium carbonate, nitrate, or nitrite. These salts are washed into the 
 soil by the rain, and are then taken up by plants to form those por- 
 tions of their tissues which, in addition to carbon, hydrogen, and 
 oxygen, also contain nitrogen, so that the ammonium compounds in 
 the atmosphere are essential to plant life ; yet the amount of these 
 is very small and variable ; the greatest quantity ever found has 
 been 47.6 parts by weight in one million of the atmosphere. Small 
 quantities of other impurities, such as sulphur dioxide and sul- 
 phuretted hydrogen, may occur in restricted areas where such gases 
 are being formed ; as, for instance, in districts where large quanti- 
 ties of sulphur-bearing coal are burned. Ozone is also at times 
 -supposed to be present in the atmosphere. 
 
 The solid particles floating in the air as dust may be of two 
 kinds, inorganic and organic. Sodium chloride is always present 
 in the inorganic particles ; the organic substances may be of the 
 greatest variety, and frequently contain micro-organisms which can 
 inaugurate disease. 
 
 The pressure which the atmosphere is exerting, by reason of its 
 weight, is measured by the barometer. In the seventeenth century 
 some Florentine pump-makers, wishing to convey water to a very 
 great height by means of a long suction pump, discovered to their 
 chagrin that, no matter how great their exertions, the water would 
 not follow the piston for more than thirty-two feet, and so Galileo 
 Galilei was appealed to for an explanation. The cause assigned to 
 this phenomenon by the great man was, however, entirely a wrong 
 one, for he maintained that a column of water longer than thirty- 
 
THE BAROMETER. 171 
 
 two feet would be broken by its own weight, just as would a bar of 
 iron of sufficient length ; therefore, water could never be pumped 
 to any great height. Torricelli, a pupil of Galileo, soon after (1643) 
 found the right explanation. He reasoned that, when the pump 
 created a vacuum, the water was pressed upward by the weight of 
 the atmosphere; if this were so, the height of a column of a spe- 
 cifically heavier liquid, such as mercury, which the atmosphere would 
 be capable of sustaining, should be proportionally less. Accord- 
 ingly, Torricelli rilled a glass tube (sealed at one end) with mer- 
 cury, closed the open end with his thumb and, inverting the tube, 
 placed this in a vessel filled with the same metal ; the column of 
 mercury then sank until its upper surface was between 28 and 29 
 inches above the level of the liquid in the trough, so that a vacuum 
 was produced in the upper part of the tube. The experimenter 
 now concluded that the weight of the column of mercury in the 
 tube must be equal to the weight of a column of water with an 
 equal cross section and a height of thirty-two feet, while both col- 
 umns exerted a pressure which was opposed by an equal one pro- 
 duced by the atmosphere acting on the surface of the liquid in the 
 open vessels in which the tubes were placed. If this conclusion 
 was correct, the height of a column of mercury which could be sus- 
 tained by the atmosphere would be less on the mountain-tops than 
 on the low lands, and so Pascal, hearing of Torricelli's experiment, 
 induced his brother-in-law, Perrier, to ascertain, experimentally, if 
 this theory was really the correct one. Perrier took a barometer to 
 an altitude of 5,000 feet, and reported that at that elevation the 
 mercury stood three inches lower than it did in Paris. The whole 
 matter was now clear ; the atmosphere exerted a pressure which 
 could be measured by the height of the column of mercury it could 
 sustain, while the instrument constructed with this end in view 
 subsequently became known as the barometer. 
 
 The distance from the upper level of the mercury in the barome- 
 ter tube to that in the vessel underneath, is the height of the barom- 
 eter ; at the level of the sea the average is 760 m. m., and in all 
 scientific work this is taken as the standard for barometric meas- 
 urements ; as the weight of a column of mercury 760 m. m. long 
 and with one square inch cross section is 15 pounds, it follows that 
 the pressure exerted by the atmosphere is 15 pounds to the square 
 inch. We do not feel this enormous weight because the air presses 
 
172 VOLUMES OF GASES; CALCULATIONS. 
 
 equally in all directions, and because the pressure from within our 
 bodies counterbalances that from without. 
 
 The volume occupied by any gas is inversely as the pressure 
 exerted on it. Double the pressure and you halve the volume, 
 quadruple the pressure and the volume will be one-fourth, and so 
 on. If V and V, be the volumes of a gas at the same temperature 
 but under different pressures, P and P, , then : 
 
 1. V : V,::P,:P , 
 
 2. V P =V,P,, 
 
 now if P is the standard barometric pressure of 760 m. m., 
 then : - 
 
 3. V = -l^-< 
 760' 
 
 or, the volume of a gas at standard pressure is equal to the volume 
 at the existing pressure multiplied by that pressure in millimetres 
 and divided by 760 ; but the existing pressure is the pressure of 
 the atmosphere measured by the height of the barometer, h, 
 (P, = h)sothat:- 
 
 4 Vn ' 
 
 ' 760 
 
 All gases, when not very near the point at which they become 
 liquid, expand ^fa of their volume for each degree of temperature,* 
 so that -a litre of gas at will be 1 -f ^^ litres at + 1 and 
 1 -j_ jfy at 10, and ten litres would be 10 + 2 W at + 1, conse- 
 quently : 
 
 t 6. V=Vl t. 7. V = _L 
 
 where V is the volume of any gas at 0, V, the volume of the 
 same gas at any temperature, t is that temperature, and = ^73. 
 Uniting 3 and 7 in one equation, we have : 
 
 8 ' Vo = 76oJ'+ u Q 
 
 so that the volume of any gas observed at and 760 m. m. is 
 equal to the volume multiplied by the height of the barometer and 
 divided by 760 times one plus ^3- of the temperature. If the gas 
 
 * More accurately .00367. 
 
THE ATMOSPHERE; DEPTH OF. 173 
 
 to be measured is in a eudiometer tube,* partially filled with, mer- 
 cury, the pressure under which it is, is naturally not that of the 
 atmosphere, but is atmospheric pressure minus the height of the 
 column of mercury in the tube. If this height in millimetres be 
 called w, then : 
 
 9 y = V,(h-w) 
 
 760 (1 + t) 
 
 and, furthermore, if the gas is saturated with water-vapor, the 
 
 tension of water-vapor in millimetres, at the temperature of obser- 
 vation, must also be deducted from the barometric measurement, 
 
 so that : , T . , . - 
 
 10 V _ V, [h - (w + fr)] 
 
 760(1 -fat) 
 
 Where V ( = volume observed. 
 
 h = observed barometric pressure. 
 
 w = height of mercury column above trough. 
 
 4> = tension of water vapor in millimetres. 
 
 a = .00367. 
 
 t = observed temperature. 
 
 When the relations between volumes of gases are considered, these 
 are supposed to be under standard conditions, at and 760 in. m. 
 pressure. 
 
 The height of the barometer at any place is constantly undergo- 
 ing variations, for the atmosphere is always subject to more or less 
 serious local disturbances, which affect the pressure exerted by it. 
 The depth of the atmosphere is uncertain ; it has been variously 
 given at from thirty to two hundred miles. As the pressure is 
 greatest on the surface of the earth, the air must be densest at this 
 point, and must diminish in density the higher the altitude. Prob- 
 ably at an elevation of ten miles the pressure of the atmosphere 
 would be imperceptible. The temperature of the air becomes less 
 the greater the distance from the earth. The specific weight of the 
 air at 760 m. m, and 0, being the most frequently used standard 
 of measurement, is generally taken as unity. Sometimes, however, 
 gases are measured with hydrogen as the standard, when the specific 
 gravity of air becomes 14.38 (hydrogen = 1), or 28.76 (hydrogen 
 = 2).t In order, therefore, to find the specific gravity of any gas, 
 
 * See note 20 of the Appendix. t In round numbers, 14.4 and 28.8. 
 
174 THE ATMOSPHERE ; RELATION TO LIFE. 
 
 with hydrogen as unity, we must multiply the specific gravity, 
 referred to air, by 14.38. 
 
 All living organisms upon the earth are dependent upon the 
 atmosphere. Its oxygen, its carbon dioxide, its ammonia, and its 
 water-vapor are necessary to all forms of life, for those constituents 
 which cannot be directly used by animals, indirectly find their way 
 into their systems. The air is inspired by animals, its oxygen is 
 absorbed, comes in contact with every tissue in the body and is ex- 
 haled, charged with carbon dioxide and water-vapor, after its oxi- 
 dizing action is completed. The plants make use of the carbon 
 dioxide which finds its way into the atmosphere ; when in the sun- 
 light they assimilate it and thus form the greater portion of their 
 tissues, but plants, as well as animals, require oxygen for their 
 existence; neither can plants live without the presence of com- 
 pounds of nitrogen, for many of their most essential chemical con- 
 stituents, such as the albumens, are composed chiefly of carbon, 
 hydrogen, oxygen, and nitrogen. A little of this nitrogen may 
 possibly be taken directly from that which is contained in the 
 atmosphere, but certainly the major portion is furnished by com- 
 pounds of nitrogen found in the soil. These compounds would 
 soon be entirely used up were it not for their constant renewal by 
 the addition of those substances which, originally in the atmo- 
 sphere, are washed to the ground by rains, and by such nitrogenous 
 products as are produced in the soil by the decay of organic sub- 
 stances. The plants are thus able to build their tissues from the 
 simplest inorganic materials from carbon dioxide, water, ammo- 
 nium nitrates and nitrites. Animals have no such power ; they 
 destroy, where plants create. Some live upon plant substances 
 and assimilate the ready-formed, complicated organic compounds ; 
 others prey upon each other, so as to get these constituents second- 
 hand ; they all die and, by decaying, once more return to the soil 
 and air those substances which the plants had used ; thus a cease- 
 less rotation of the life-supporting constituents of the atmosphere 
 is going on ; while the energy necessary to cause these metamor- 
 phoses is furnished by the heat and light of the sun. 
 
ELEMENTS OF NITROGEN FAMILY. 
 
 175 
 
 
 CHAPTER XXIV. 
 
 COMPOUNDS OF THE ELEMENTS OF THE NITROGEN FAMILY. 
 
 The elements of the nitrogen family are : 
 
 1. Nitrogen, atomic weight 14.03 
 
 2. Phosphorus, 
 
 3. Arsenic, 
 
 4. Antimony, 
 
 5. Bismuth, 
 
 The same changes in the physical characteristics of the ele- 
 ments which, with increasing atomic weight, are observed in the 
 halogen and in the sulphur families, are repeated in that group of 
 elements of which nitrogen is a representative. This will be seen 
 from the following table : 
 
 O, N, colorless gases. 
 
 Cl, yellow gas, 
 
 P, yellow, easily fused solids. 
 
 Br, brown liquid, Se, As, metallic appearing solids. 
 
 [, black solid, Te, Sb, silver-white appearing solids 
 
 - Bi, reddish metallic solid. 
 
 Specific Melting 
 
 gravities. points. 
 
 HYDROGEN COMPOUNDS. 
 
 
 
 
 
 General formula of hydrogen compounds of 
 
 
 
 
 
 halogens, H X. 
 
 
 
 
 
 General formula of hydrogen compounds 
 
 FH 
 
 H 2 
 
 N H 3 
 
 
 of sulphur group, H 2 X. 
 
 
 
 
 
 General formula of hydrogen compounds of 
 
 C1H 
 
 S H 2 
 
 P H 3 
 
 
 nitrogen group, H 3 X. 
 
 
 
 
 
 Valence of the elements of the halogen 
 
 Br 11 
 
 SeH a 
 
 AsH 3 
 
 
 family toward hydrogen is 1. 
 
 
 
 
 
 Valence of the elements of the sulphur fam- 
 
 I H 
 
 TeH a 
 
 SbH 3 
 
 
 ily toward hydro'gen is 2. 
 Valence of the elements of the nitrogen fam- 
 
 
 
 
 
 ily toward hydrogen is 3. 
 
 In the nitrogen family, as well as in the other two which we 
 have studied, the elements, as the atomic weights increase, change 
 
 * An element, not as yet discovered, belongs in the interval between anti- 
 mony and bismuth ; as a consequence the differences between antimony and 
 bismuth are much more marked than are the differences between any other 
 neighboring pair of elements we have so far considered. 
 
176 
 
 ELEMENTS OF NITROGEN FAMILY. 
 
 into substances entirely metallic in appearance, and the alteration 
 is even more fundamental in its character than it is in the family 
 of the halogens or in the oxygen family, for the chemical as well as 
 the physical properties of the last two elements in this group are 
 more metallic than they are not-metallic. Once more we come in 
 contact with a family of elements in which the one having the 
 smallest atomic weight is a colorless gas, that with next higher a 
 yellow solid, easily fused and easily burned, and the next a grayish 
 white solid, with the appearance of a metal and the chemical be- 
 havior of a not-metal (for in this respect arsenic is entirely like 
 tellurium). Antimony and bismuth are metals both in their phys- 
 ical and chemical properties, but antimony shows a not-metallic 
 nature in some of its compounds, while bismuth is always a 
 metal. The melting points of the elements in this family increase 
 with the atomic weights, exactly as is the case in the sulphur and 
 halogen families. 
 
 All of the elements belonging to the nitrogen group, with the 
 exception of bismuth (which is too much of a metal to do so), form 
 gaseous hydrogen compounds which, following the general rule, 
 diminish in stability as the atomic weight of the characterizing 
 element increases. These compounds are formed by union of three 
 hydrogen atoms to one of the nitrogen-like element so that their 
 general formula is H 3 X, while the valence of the typical element 
 is three. By comparing all the elements in the halogen, sulphur, 
 and nitrogen groups, we can see that, as the atomic weights of 
 the families, as a whole, diminish, the valences toward hydrogen 
 
 increase. 
 
 F. 0. N. 
 At. wt. 19 16 14 
 
 F H. O H 2 N H 3 
 
 As we pass from one element 
 to another corresponding one to 
 the right of it, there is a dimi- 
 nution of not more than five 
 units in the atomic weight and 
 an increase of one in the val- 
 ence. No elements exist, the 
 atomic weights of which lie be- 
 tween any two on the horizontal 
 lines, so that these twelve ele- 
 ments form a portion of the 
 natural grouping obtained by 
 arranging the elements in the 
 order of their atomic weights. 
 
 Cl. S. P. 
 At. wt. 35.5 32 31 
 
 C1H. SH 2 PH 3 
 
 Br. Se. As. 
 At. wt. 80 79 73 
 
 Br H. Se H a As H 3 
 
 I. Te. Sb. 
 
 At, wt. 127 125 120 
 
 IH. TeH 2 SbH 3 
 
 
 Valence I. II. III. 
 
 Atomic wts. ^~^^> 
 
 Valence. 
 
ELEMENTS OF NITROGEN FAMILY. 177 
 
 The hydrogen compounds of the elements of the nitrogen fam- 
 ily show a chemical character' which differs from that found in the 
 other two groups which have been studied, and the reason for this 
 difference is to be found in the greater number of hydrogen atoms 
 which are joined to one atom of the typical element contained in 
 them. The hydrogen compounds of the halogens are acids, be- 
 cause one atom of each of those elements can unite with but one 
 atom of hydrogen; the positive character of the latter, therefore, 
 is not sufficient to counterbalance the negative properties of the 
 halogen. The elements of the sulphur family form hyrlr.Qgfm.-p*>m- 
 pounds which are but slightly acid ; for the two hydrogen atoms 
 contained in these have twice the effect of the one in the chlorine 
 group, while, lastly, the hydrngp.Ti p.n^pnnnrls of the elements of 
 the nitrngftTT_jFfljm1y a.rp. p.it.hftT ha.sip. nr^Tiftiitrfl.l, for here the three 
 hydrogen atoms entirely counterbalance the chemical character of 
 that of the element with which they are united. However, an 
 atom of hydrogen has a small mass, so that its influence on the 
 character of a compound would become less as the mass of the 
 atom of the element with which it is united becomes greater. 
 This connection between mass and chemical character can be seen 
 in the sulphur and nitrogen families ; thus, water is nearly neutral 
 in behavior, being both basic and acid ; sulphuretted hydrogen 
 (where the mass of the sulphur atom is twice that of an oxygen 
 atom) has the power of reddening litmus, while its basic character 
 is much less than that of water ; lastly, hydrogen selenide and 
 telluride are also weakly acid. In the nitrogen family, ammonia 
 is a pronounced base, being capable of uniting with almost all acid 
 substances to form stable compounds ; PH 3 (phosphine) yields salts 
 with but a few acids, such as H Br and HI, while As H 3 (arsine) 
 and SbH s (stibine) can form no salts ; so that the increase in the 
 mass of the not-metallic elements has gradually counterbalanced 
 the effect of the positive hydrogen, even in spite of the fact that 
 the not-metallic character of the elements themselves has dimin- 
 ished. Were antimony as negative as nitrogen, stibfrie would, in 
 that event, undoubtedly be an acid ; it being the diminishing not- 
 metallic character of the elements in a family, as we pass to those 
 members with higher atomic weights, which influences the character 
 of the hydrogen compounds which are formed. Should iodine, to 
 cite another example, be as not-metallic as chlorine, we would 
 
178 HYDROGEN COMPOUNDS OF NOT-METALS ; COMPARISON. 
 
 expect hydroiodic acid to be a much stronger acid than hydrochloric, 
 for the mass of an atom of iodine is greater than that of an atom 
 of chlorine ; that this is not the case is due to the fact that the in- 
 creased mass of iodine renders that element less not-metallic in its 
 character than is chlorine. Organic chemistry teaches us that if 
 we substitute a more positive group of elements for the hydrogen 
 contained in ammonia, the resulting compound becomes more like a 
 metal ; and that by substituting a less positive group, it becomes 
 less like a metal in its character; and also that arsine and stibine, 
 which are neutral, can be made to act like metals if we only 
 replace their hydrogen atoms by groups of atoms having a more 
 positive chemical character.* 
 
 Whatever truth there may be in the above considerations, the 
 facts themselves are to be remembered ; so that : 
 
 Compounds HX of the halogens are acid. 
 
 Compounds H 2 X of the sulphur group are acid (and sometimes 
 basic). 
 
 Compounds H 3 X of the nitrogen group are basic or neutral. 
 
 In calling these compounds acid, we mean that they react with 
 bases to form salts or salt-like bodies, in calling them basic, that 
 they can unite with acids to form similar substances. The follow- 
 ing reactions will make this clear : 
 
 HC1+ KOH = KC1 + H 9 0. 
 H 2 S + 2 KOH = K 2 S + 2 H 2 0. 
 Acid. Base. 
 HC1+ H 3 N^=NH 4 C1. 
 H 2 S + H 3 N = NH 4 SH. 
 Acid. Base. 
 
 Water can sometimes act like a base. This fact is seen from its 
 behavior when it is brought in contact with anhydrides and is made 
 apparent by comparing the following two equations : 
 
 S0 3 -+-H 2 = H 2 S0 4 , 
 
 S0 3 +K 2 = K 2 S0 4 .t 
 Anhydride. Base. Salt. 
 
 * A "positive" element is one which in a given chemical compound 
 behaves like a metal; a negative element is one which behaves like a not- 
 metal ; while positive and negative groups are such as can chemically show the 
 same contrast to each other as do a metal and a not-metal. 
 
 t See pages 30 and 31. 
 
ELEMENTS OF NITROGEN FAMILY; OXIDES. 
 
 179 
 
 One great difficulty in attempting a classification is found in the 
 necessity of using such arbitrary terms as " basic " and " acidic," 
 which are often differently applied by different chemists ; this diffi- 
 culty becomes less, for the present, if we remember that only the 
 general and important chemical characteristics shown by the hydro- 
 gen compounds of the not-metals are at present being considered ; 
 so that when the student subsequently discovers, for instance, that 
 the hydrogen atoms in ammonia can be replaced by potassium, he 
 need not, for that reason, look upon ammonia as an acid he need 
 remember merely that this is an individual case, which is an ex- 
 ception to the rule. (See page 75.) 
 
 All of the elements of the nitrogen family form oxides ; the gen- 
 eral formulae of the most important of these, together with those of 
 the acids derived from them, are given on the following table ; the 
 oxides of the corresponding halogen compounds are also tabulated 
 for purposes of comparison : 
 
 HALOGENS. 
 
 VAL- 
 ENCE. 
 
 NITROGEN FAMILY. 
 
 In the nitrogen family, 
 the acids H Y O 2 and H Y O 3 
 can (excepting the nitrogen 
 acids) form stable hydrated 
 acids : 
 HYO 2 + H 2 O = H 3 YO 3 
 HY0 3 + H 2 O = H 3 YO 4 
 
 OXIDES. 
 
 ACIDS. 
 
 OXIDES. 
 
 ACIDS. 
 
 NOMENCLATURE. 
 
 X 2 
 X 2 3 
 
 x,p. 
 
 (X. 7 ) 
 
 HXO 
 HX0 2 
 HXO 3 
 HX0 4 
 
 1 
 3 
 5 
 
 7 
 
 
 
 Hypo-ous acid 
 ous acid 
 ic acid 
 
 Y 2 3 
 Q, 
 
 II Y0 2 
 
 H Y o 3 
 
 
 
 
 
 In this table X represents any halogen, Y any member of the nitrogen 
 family. 
 
 The trioxides and the pentoxides, as well as the acids derived 
 from them by the addition of water, have corresponding formulae 
 in both families, and the nomenclature of the acids is also parallel, 
 thus : 
 
 C1 2 3 yields HC10 2 (chlorous acid), and N 2 3 yields HN0 2 , 
 (nitrous acid). 
 
 I 2 5 yields HI0 3 (iodic acid), and jST 2 5 yields HN0 3 (nitric 
 acid). 
 
 There are no very important acids in the nitrogen family corre- 
 sponding to hypochlorous and hypobromous acids ; and, as the 
 highest valence of the elements of this group toward oxygen is 
 five, there can foe none corresponding to perchloric acid, which 
 is derived from chlorine with a valence of seven. All of the ele- 
 ments of the nitrogen family, with the exception of nitrogen and 
 
180 ELEMENTS OF NITROGEN FAMILY; OXIDES. 
 
 bismuth, have a great tendency to form hydrated acids which are 
 much less readily decomposed than are those of the corresponding 
 class in the other two families. As a consequence, phosphorous 
 acid exists, not as HP0 2 , but as H 3 P0 3 (HP0 2 + H 2 0), and 
 phosphoric acid as H 3 P0 4 (HP0 3 -+- H 2 0), more often than as 
 HP0 3 , so that the acids most frequently met with contain the 
 same number of hydrogen atoms in a formula weight as there are 
 hydrogen atoms in a molecule of ammonia. The oxide Y 2 3 be- 
 comes more basic as the atomic weight of Y increases, so that, 
 while N 2 8 and P 2 3 are the anhydrides of acids, As 2 3 is both a 
 base and an anhydride, Sb 2 3 is more basic than acidic, while 
 Bi 2 3 acts altogether as a base. This change in the nature of the 
 oxides is the natural result of the change from not-metal to metal 
 which takes place in this family, as we pass from member to 
 member in the direction of increasing atomic weights. 
 
 The pentoxides, Y 2 5 , all are anhydrides 
 of acids with no basic properties ; * therefore, 
 an addition to the amount of oxygen present 
 in the compounds which are basic, changes 
 the latter into more negative and, as a conse- 
 quence, acidic bodies. The general rule is 
 that, where several oxides of the same metal 
 exist, the character of these becomes less 
 basic as the number of oxygen atoms in a 
 
 molecule increases; so that frequently the lowest oxide may be a 
 strong base, forming most stable salts with acids, while the highest 
 may be the anhydride of an acid, and, as a consequence, the oxides 
 of the same element may not resemble each other so much as they 
 do the corresponding oxides of some other element. 
 
 Nitrogen, in addition to the two oxides, N 2 3 and oST 2 5 , forms 
 three others with the formulae, N 2 0, NO, and N0 2 ; these will be 
 discussed at the proper time. The chlorides of the elements of 
 the nitrogen family have the general formulae Y C1 3 and Y C1 5 ; 
 their stability increases as the metallic character of the elements 
 becomes more pronounced : 
 N C1 3 (?), explosive. 
 
 * Bi 2 O 5 shows scarcely any of the properties of an anhydride ; the com- 
 pound H Bi O 3 , corresponding to HNO 3 , has been isolated, but it forms no 
 salts and shows a tendency to break down, giving off oxygen. 
 
 BASIC. 
 
ELEMENTS OF NITROGEN FAMILY; CHLORIDES. 181 
 
 P C1 3 , P C1 5 , decomposed by water to phosphorous and phospho- 
 ric acids. 
 
 As C1 3 , , exists in the presence of little water, entirely 
 
 decomposed by much water. 
 
 Sb C1 8 , Sb C1 5 , partially decomposed by water. 
 
 BiClg, , partially decomposed by water. 
 
182 AMMONIA; OCCURRENCE, HISTORY. 
 
 CHAPTER XXV. 
 
 AMMONIA AND THE OTHER COMPOUNDS OF NITROGEN 
 AND HYDROGEN. 
 
 Ammonia; formula, NH 3 . Specific gravity, air = \, is .589, H 2 
 =2, is 16.96 ; molecular iveight, 17 ; 1 cc. at and . 76 m. pres- 
 sure, weighs .0007651 gram. Hydrazin ; formula, N 2 H 2 . 
 Azoimid; formula, N 3 H. 
 
 AMMONIA is by far the most important of the three compounds 
 of nitrogen and hydrogen. It is never found, as such, in nature, 
 but always occurs as an ammonium salt, in combination with some 
 acid ; in the atmosphere and soil it is found as ammonium carbon- 
 ate, nitrite, and nitrate ; in mineral waters and in volcanic regions, 
 as the sulphate. Ammonium compounds occur in almost all plants, 
 in the air exhaled from the lungs, and in the urine of animals. 
 Those ammonium compounds which are found in the soil and in 
 clay are produced from outside sources. 
 
 Salts of ammonium were first introduced into Europe from 
 Eastern countries, especially from Armenia, so that the name given 
 by the Arabian alchemists to the chloride, the salt which earliest 
 came into prominence, was sal armeniacum ; but this term was sub- 
 sequently altered to sal ammoniacum, which had been given to so- 
 dium chloride imported from the neighborhood of the temple of 
 Jupiter Ammon in the Libyan Desert. Ammonium compounds 
 were first prepared in the East by the distillation of camels' dung, 
 and were much prized as universal remedies ; subsequently other 
 animal refuse was used in the preparation of the carbonate ; at one 
 time the substance prepared by the dry distillation of harts' horns 
 was considered a most potent medicine, so that the name, spirits of 
 hartshorn, is still used to designate a solution of ammonia in 
 water. No distinction was made between ammonia and the car- 
 bonate of ammonium ; both were called volatile alkali, until 
 Priestley collected pure ammonia gas over mercury. The term 
 ammoniaque was introduced by the French chemists at the close of 
 the last century, and from this our term ammonia has its origin. 
 
AMMONIA ; PREPARATION. 183 
 
 Preparation of ammonia by reduction of the oxides of nitrogen. 
 
 Under ordinary circumstances nitrogen and hydrogen have no 
 tendency whatever to unite ; only if electric sparks are allowed to 
 pass for some time through a mixture of the two gases does union 
 take place. If, however, hydrogen acts upon some oxide of nitro- 
 gen under proper conditions, especially if the hydrogen is in the 
 so-called nascent state, ammonia is formed with the greatest ease. 
 For instance, if a mixture of nitric oxide (NO) and hydrogen are 
 passed through a tube heated to dull redness and containing a few 
 pieces of spongy platinum, ammonia and water are produced : 
 
 NO + 5H = NII 3 -hH 2 0. 
 
 If pieces of iron or zinc act upon very dilute nitric acid, am- 
 monium nitrate is produced ; the hydrogen which first results from 
 the action of the metal on the acid not only takes away all of the 
 oxygen from the compound of nitrogen which it attacks, but in 
 addition, even unites with that element to form ammonia. (See 
 nitric acid.) 
 
 Preparation of ammonia from organic substances. 
 
 Decaying organic substances, which contain nitrogen, give off 
 ammonia, and this substance, uniting with the carbon dioxide 
 formed at the same time, produces ammonium carbonate. Owing 
 to its formation in this way the odor of ammonia can always be 
 detected in the neighborhood of heaps of manure or in stables. 
 The compounds of nitrogen, which are an essential constituent of 
 existing plants and animals, must have been just as necessary in 
 past geologic eras, so that the vegetable remains, which are found 
 in the form of bituminous coal, are rich in compounds of carbon, 
 hydrogen, oxygen, and nitrogen; the dry distillation of this coal 
 at present produces nearly all of the ammonia in use. When 
 bituminous coal is heated without the access of air, a large num- 
 ber of products of great commercial importance are given off. 
 They are : 
 
 Gaseous Illuminating gas, ammonia, sulphuretted hydrogen. 
 
 Liquid Water, benzol, toluol, phenol, etc. 
 
 Solid Naphthalene, anthracene, etc. 
 
 The highest boiling products are substances like asphalt, while 
 the carbon remains behind in the form of coke. 
 
 The ammonia which passes from the gas manufacturing retorts 
 
184 
 
 PREPARATION. 
 
 is separated from the mixture of gases (obtained by the dry distil- 
 lation of coal) by passing the gases upon surfaces of wood over 
 which a continuous stream of water is trickling ; * the solution so 
 formed is a dark-colored liquor from which it is an easy matter 
 to obtain ammonia in a pure state, t To do this, the dissolved 
 ammonia can be converted either into ammonium chloride or sul- 
 phate by the addition of hydrochloric or sulphuric acid. The so- 
 lution of ammonia in water, we may assume, for purposes of 
 comparison, contains the gas combined with water as ammonium 
 hydroxide : - NH 3 + H 2 = KH 4 OH. 
 
 When acids are added to this hypothetical hydroxide, an ammonium 
 salt is produced, just as a potassium or sodium salt would be formed 
 under similar circumstances : 
 
 KOH+ HC1 = KC1 + H 9 0, 
 
 2KOH + H 2 S0 4 = K 2 S0 4 +2H 2 0, 
 
 2s(NH 4 )OH + HC1 = (NH 4 )C1 + H 2 0, 
 
 2 (NH 4 ) OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2 H 2 0, 
 
 (The group of elements represented by the formula NH 4 very much 
 resembles potassium.) In order to liberate the ammonia from these 
 salts, it is only necessary to treat them with some such base as so- 
 dium or calcium hydroxide, for these not-volatile bases would expel 
 a volatile one from its salts, so that : - 
 
 NH 4 C1 + NaOH = NaCl + NH 4 OH, 
 
 (NH 4 ) 2 S0 4 -f- 2 Na OH = Na, S0 4 + 2 NH 4 OH, 
 
 2(NH 4 )C1 + Ca(OH) 2 = Ca C1 2 + 2 NH 4 OH. 
 
 The ammonium hydroxide which would be formed breaks down 
 into ammonia and water, NH 4 OH = NH 3 -f H 2 0, so that pure am- 
 monia gas can be liberated from the chloride or sulphate of ammo- 
 nium. The gas is then collected by passing it into water, and the 
 solution so formed is the liquor ammonia of commerce. A similar 
 method can be employed in the production of small quantities of 
 
 * The greater portion of the illuminating gas is insoluble, or nearly in- 
 soluble in water. The ammonia, on the other hand, is quite readily soluble, 
 so that it can be separated from the other constituents of the gas by solution 
 in water. 
 
 t The solution contains the least part of the ammonia in the form of free 
 ammonia; the greater part is in the form of ammonium salts, such as ammo- 
 nium sulphide, cyanide, or sulphocyanide. 
 
AMMONIA; PROPERTIES. 
 
 185 
 
 ammonia for laboratory use. 48 One objection to the method of 
 obtaining pure ammonia, which has been detailed, is found in the 
 fact that the sulphate and chloride so produced are very impure 
 and permeated by a tar-like substance. As a consequence, of late 
 years, the crude liquors are directly distilled, after the addition of 
 lime (the latter substance being used to decompose the ammonium 
 salts present in the liquid), the tarry matter being removed from 
 the ammonia by filtering through pieces of paraffine. By this 
 process of distillation the pure ammonia passes off first, and is 
 collected in water. 
 
 Ammonia is a colorless gas with a penetrating alkaline odor. 
 At 40 and under a pressure of one atmosphere, it condenses to 
 a colorless liquid which boils at 38. 5 and which changes into a 
 crystalline solid at 75. The gas has a specific gravity of .589 
 at and 760 m.m., which, with hydrogen as two, gives a density of 
 16.96, so that the molecular weight of this compound is 17. As 
 a consequence it is specifically lighter than air, and can be col- 
 lected by upward displacement, exactly as is done with hydrogen. 
 Ammonia is quite a stable compound, just as are the correspond- 
 ing substances in the oxygen and halogen families, namely, water 
 and hydrofluoric acid ; it follows that ammonia does not support 
 combustion. It also burns with difficulty, even when it is mixed 
 with oxygen and ignited ; 49 the products of combustion are water, 
 nitrogen, and traces of the oxides of nitrogen. Naturally it does 
 not support life. Ammonia can be completely decomposed into 
 nitrogen and hydrogen by passing an electric spark through the 
 gas for some length of time; and, when this is done, two vol- 
 umes of ammonia yield three of hydrogen and one of nitrogen. 
 
 
 = 
 
 
 
 
 H 
 
 1 
 H 
 
 ( H 
 
 KVH 
 
 (H 
 
 I 
 
 3 
 
 s^ 
 
 r 
 
 H 
 
 1 
 H 
 
 ( H 
 
 N^H 
 (H 
 
 
 
 I 
 I 
 
 I 
 I 
 
 
 or 2 NH 3 = N 2 + 3 
 
186 
 
 AMMONIA ; COMPOSITION BY VOLUME. 
 
 Ammonia, when it is heated to 780 in a porcelain or iron tube, is 
 almost completely dissociated into hydrogen and nitrogen. 
 
 Ammonia can, under proper circumstances, burn in oxygen to 
 form water and nitrogen, and it is acted 011 by chlorine in a similar 
 way, producing hydrochloric acid and nitrogen : 
 
 3 + 3 Cl = N + 3 H Cl. 
 
 By means of this reaction we can readily determine the relative 
 volumes of nitrogen and hydrogen which go to make ammonia. If 
 we fill a long glass tube with chlorine, and then add ammonia 
 water, without admitting any air, hydrochloric acid and nitrogen 
 will be formed. We learned that one volume of chlorine unites 
 with one volume of hydrogen to form hydrochloric acid, so that the 
 volume of chlorine in the vessel must have united with an equal 
 volume of hydrogen, which was contained in the ammonia; and, 
 therefore, the volume of hydrogen which was present in the ammo- 
 nia is given by the volume of chlorine in the vessel. Now, by 
 cautiously admitting very dilute sulphuric acid, we can absorb the 
 excess of ammonia, while the hydrochloric acid has already formed 
 ammonium chloride with the ammonia water. Nothing but nitro- 
 gen will, therefore, remain in place of the chlorine originally used, 
 and its volume will exactly fill one-third of the tube. It follows 
 that one volume of nitrogen and three of hydrogen unite to form 
 ammonia. The following diagram will make this clear : 
 
 in solution = 
 
 The hydrochloric acid is removed by dissolving in ammonia water, 
 so that the nitrogen remains, and this is one third of the original 
 volume of chlorine. 
 
 We have seen that when ammonia is decomposed by the electric 
 spark, two volumes of the gas yield three of hydrogen and one of 
 nitrogen, and that, conversely, ammonia is formed from three vol- 
 umes of hydrogen and one of nitrogen. From the first of these two 
 discoveries we can conclude, as we do in respect to oxygen and 
 
AMMONIA; PROPERTIES. 
 
 187 
 
 Jf 
 
 chlorine, that nitrogen contains two atoms in its molecule (see page 
 71) ; from both we conclude that the formula of ammonia is NH 3 . 
 In addition, the specific gravity of ammonia shows us that its mo- 
 lecular weight is 17; and quantitative analysis shows that in 17 parts 
 by weight of ammonia there are fourteen of nitrogen and three of 
 hydrogen. The atomic weight of nitrogen is therefore presumably 
 14, and must remain so unless some compound of nitrogen should 
 be discovered, the molecular weight of which is known, and which 
 contains relatively less than fourteen parts by weight of nitrogen. 
 
 Ammonia is very soluble in water ; one volume of water dissolves 
 813 times its own volume of ammonia gas. 50 The solution has the 
 odor of ammonia gas, is alkaline (for it changes red litmus to blue), 
 and is generally considered as containing the ammonia combined 
 with water in the form of ammonium hydroxide, NH 4 OH.* Upon 
 warming the solution, ammonia is expelled, and this fact is made 
 use of in the preparation of artificial ice. An iron vessel (B), con- 
 taining ammonia water, is connected by pipes with another one 
 constructed in the form of 
 a double-walled hollow cyl- 
 inder (A), all of the con- 
 nections being such that 
 no gas can escape (Fig. 
 10). Ammonia water is 
 placed in the first vessel 
 and is warmed. The es- 
 caping gas passing into the 
 space between the double 
 walls of the second vessel, 
 which is cooled by means 
 of cold water, the pressure 
 of ammonia and the cold 
 combined condense the gas 
 to a liquid. The conditions are now reversed, the cylinder (A) has 
 
 * According to recent investigations, it seems probable that ammonium 
 hydroxide is not formed when ammonia dissolves in water, for all other hydrox- 
 ide solutions conduct electricity with as great ease as salt solutions; but the 
 solution of ammonia in water has only -^ the conducting power of a salt of the 
 same molecular weight. As all of the salts of ammonium correspond to those 
 of potassium, it is, however, convenient to consider the solution of ammonia 
 in water as ammonium hydroxide. 
 
 Fig. 10. 
 
188 AMMONIUM SALTS. 
 
 its jacket of cold water removed ; while the first vessel (B) is cooled; 
 the condensed gas boils and is absorbed by the water contained in 
 B, and when the liquid ammonia boils, enough heat is absorbed to 
 freeze a can of water placed within the hollow cylinder. 
 
 Ammonia gas can add itself to acids to form salts; a few 
 examples of such additions are represented by the following 
 equations : 
 
 NH 8 + H Cl = NH 4 Cl (ammonium chloride), 
 2 NH 3 + H 2 S0 4 = (KH 4 ) 2 S0 4 (ammonium sulphate), 
 NH 3 + H N0 3 = NH 4 ]tf0 3 (ammonium nitrate). 
 
 These salts are termed ammonium salts. The univalent group of 
 elements NH 4 , which, chemically, resembles potassium, is called 
 ammonium ; it takes the place of hydrogen in the acids, and the 
 salts so produced have a character similar to that of the salts 
 of the true metals. Where a group of elements acts in this man- 
 ner, being transferred from one compound to another without 
 decomposition, just as an element would be, it is termed a radicle. 
 Ammonium is the unchanging constituent of a large number of 
 compounds ; if the grouping of elements represented by the formula 
 NH 4 is destroyed, then ammonium salts lose their identity as such. 
 Ammonium is a radicle which can act like a metal ; we are ac- 
 quainted with other radicles which are composed entirely of iiot- 
 metals, and which act as much like the latter as ammonium does 
 like the former ; indeed, it is very difficult to say just which com- 
 pounds shall be called radicles and which shall not be. It is, per- 
 haps, best to limit the term to such groups of elements as are very 
 frequently met with as the unchanging constituents of a large num- 
 ber of compounds, and which can be transferred from one compound 
 to another as a whole, and without alteration. 
 
 When ammonia dissolves in water, its solution, for convenience 
 in chemical notation, is generally considered as containing the 
 hydroxide of ammonium: 
 
 NH 3 + H 2 = NH 4 OH. 
 
 With this hypothesis in view, the parallelism between ammo- 
 nium and the metals becomes more apparent, for : 
 
AMMONIUM SALTS. 
 
 189 
 
 (NH 4 )OH + HC1 
 
 Ammonium hydroxide + Hydrochloric acid 
 
 KOH + H Cl 
 
 Potassium hydroxide + Hydrochloric acid 
 
 2 f (NH 4 )OH + H 2 S0 4 
 
 Ammonium hydroxide 4- Sulphuric acid ; 
 
 2 KOH + H 2 S0 4 
 
 Potassium hydroxide + Sulphuric acid = 
 
 C1 + H 2 0. 
 
 Ammonium chloride + Water. 
 
 *C1 + H 2 0. 
 
 Potassium chloride + Water. 
 
 :(KH 4 ) 2 S0 4 + 2H 2 0. 
 
 Ammonium sulphate -f- Water. 
 
 K 2 S0 4 + 2 H 2 0. 
 
 Potassium sulphate + Water. 
 
 Ammonium salts, upon heating, decompose into ammonia and 
 the corresponding acid.* For example, ammonium chloride, when 
 vaporized, yields ammonia and hydrochloric acid : 
 
 !STH 4 Cl = NH 3 + H Cl. 
 
 The specific gravity of ammonium chloride vapor, if hydrogen = 
 2, provided no decomposition had taken place, should be the same 
 as its molecular weight, or 53.5. In reality it is only one-half of 
 this number, or 26.75. Let us suppose a volume of hydrogen 
 weighs two grams, tlien an equal volume of ammonium chloride 
 vapor weighs 26.75 grams ; but were it vaporized without decompo- 
 sition then it would weigh 53.5 grams ; it follows that the ammo- 
 nium chloride has decomposed into a molecule of ammonia and one 
 of hydrochloric acid.f 
 
 1 vol. H. 1 vol. NH 4 C1 vapor. 
 
 2 vols. (NH. + HC1). 
 
 A 
 
 / 53.5 
 U T H 3 
 
 grams \ 
 + HClJ 
 
 26.75 gr. 
 
 26.75 gr. 
 
 If undecomposed. If decomposed. 
 
 * For exceptions to this rule see next page. 
 
 t Recently, ammonium chloride lias been vaporized under less than 
 atmospheric pressure, in which condition no decomposition took place, as the 
 specific gravity of the vapor indicated. The molecules of the gas had the 
 formula NH 4 C1. This proves that nitrogen is quinquivalent in ammonium 
 chloride. The vapor density of ammonium chloride in an atmosphere of am- 
 monia gives numbers much larger than those which would have been obtained 
 had complete dissociation taken place. See Pullinger and Gardner; Proc. 
 Chem. Soc.; 1891, 2. 
 
190 AMMONIUM SALTS; DECOMPOSITION OF. 
 
 These conclusions are a necessary result of Avogadro's hypothe- 
 sis (page 71). Quite a number of bodies dissociate into two simpler 
 ones on vaporizing. All of these give abnormal specific gravities 
 for their vapors, but the explanation is always similar to that which 
 has just been given in connection with ammonium chloride. The 
 fact that the specific gravity of ammonium chloride vapor, if hy- 
 drogen is two, does not correspond to the molecular weight, does not 
 invalidate the method of obtaining the molecular weights of gases 
 by means of their specific gravities ; it only shows us that we must 
 be careful to ascertain if the gas, the specific gravity of which we 
 are about to determine, is identical in chemical constitution with 
 the liquid or solid from which it is produced. 
 
 The atoms of nitrogen are trivalent in ammonia, but the element 
 is unsaturated (see page 108), although it is not capable of taking 
 up any more positive atoms, such as hydrogen, unless these elements 
 are joined to a negative element or group of elements. The pres- 
 ence of the large number of hydrogen atoms which are contained 
 in a molecule of ammonia has rendered that compound positive ; it 
 therefore has no tendency to further unite with positive substances ; 
 but when a negative compound such as hydrochloric (or any 
 other) acid is added to ammonia, the latter separates the former 
 into two parts ; namely, into hydrogen and the not-metallic ele- 
 ment or group of elements with which hydrogen was united; these 
 two parts add separately, so that nitrogen becomes quinquivalent : 
 
 C-K f -H 
 
 in r _H v _H in ( H v _H 
 
 W ] -H + HC1=WJ -H JI . -H + HNO.=|n --H 
 
 1 ( H H ' ( H H 
 
 L ci L N0 3 . 
 
 We have seen that water is similarly decomposed when it adds it- 
 self to oxides in order to form hydroxides (see page 117). 
 
 All ammonium salts are decomposed by heat, with the results 
 which are detailed under 1 and 2 : 
 
 1. The ammonium salts may be entirely disintegrated, as is the 
 case with ammonium nitrate ; for, when that substance is heated, 
 neither ammonia nor nitric acid is produced. 
 
 2. Ammonia and the acid are produced. 
 
 a. The acid may be volatile, then nothing remains. 
 
HYDROXYLAMIN. 191 
 
 b. The acid may be not-volatile, then it remains. 
 
 c. The acid may be decomposed by heat, then its decomposition 
 
 product remains. 
 
 The radicle NH 4 , ammonium, has never been isolated, but its 
 metal-like nature is shown by the fact that it forms an amalgam 
 with mercury. If an ammonium salt, like ammonium chloride, is 
 decomposed by sodium in the presence of mercury, the ammonium 
 liberated will form an amalgam with the latter.* 
 
 NH 4 Cl + Na .= NH 4 + Ka CL_ 
 
 The mercury expands and becomes of the consistency of soft 
 butter. Ammonium amalgam gradually decomposes in the air, am- 
 monia and hydrogen are given off, and the mercury shrinks to its 
 former size. 51 
 
 It has been mentioned (page 183) that, when diluted nitric acid 
 is reduced by means of finely divided zinc or iron, ammonium ni- 
 trate is produced. Under other circumstances (for example when 
 the reducing metal is tin) reduction does not go so far and a par- 
 tially oxidized ammonia, termed hydroxylamin, is formed. The 
 same result can be accomplished by reducing nitric oxide with tin 
 and hydrochloric acid. 
 
 Hydroxylamin is ammonia in which one hydrogen atom has 
 been replaced by hydroxyl, as the following structural formulae 
 will indicate : 
 
 H (OH 
 
 Ammonia. Hydroxylamin. 
 
 Hydroxylamin is basic in character, its salts resemble ammonium 
 salts, and they are probably similarly constructed : 
 
 < 
 
 * Amalgams are solutions of metals in mercury. They sometimes have 
 definite crystalline forms or definite quantities of metal and mercury. The 
 mercury may in these compounds possibly play a part similar to that of water 
 of crystallization. 
 
 t The structural formula of hydroxylamin is not, as yet, absolutely cer- 
 tain ; this doubt is due to the fact that, heretofore, it has been difficult to 
 isolate any quantity of the pure base. This has, however, of late, been suc- 
 cessfully accomplished, so that we may soon expect to have all doubt removed. 
 
192 HYDRAZIN. 
 
 OH 
 H 
 -Cl 
 
 Hydroxylarnin hydrochloride. 
 
 Free hydroxylamin is a solid, crystalline substance which melts at 
 32-35 and which boils at 56-57 at 22 m.m. pressure. When 
 heated at ordinary pressures it explodes violently at 130. It can- 
 not be preserved for any length of time without decomposing.* 
 On the other hand, the salts of hydroxylamin are perfectly stable. 
 
 Nitrogen forms two other compounds with hydrogen ; namely, 
 hydrazin, N 2 H 4 , and azoimid, N 3 H; both of these substances were 
 recently discovered by Curtius. Hydrazin can be considered as 
 analogous to hydrogen dioxide, for the latter is water in which one 
 atom of hydrogen is replaced by hydroxyl : 
 
 H-O-H and HO-OH, 
 
 while the former is ammonia in which one atom of hydrogen has 
 been replaced by the univalent group NH 2 , 
 
 H N = H 2 , and H 2 =N N-=H 2 ; 
 
 the group of atoms to which hydrogen rs attached is, in the one 
 case, , and in the other = N N= ; as oxygen is bi- 
 valent, each oxygen atom in the above group will be capable of 
 uniting with only one hydrogen atom ; while, as nitrogen is trivalent, 
 each nitrogen atom is capable of further union with two hydrogen 
 atoms. (See pages 107, 108.) 
 
 Hydrazin has, in all probability, never been isolated in a pure 
 state ; in reactions where it would be expected, not it, but a mix- 
 ture of nitrogen and ammonia is formed. Hydrazin is basic in its 
 character, so that it, like ammonia, unites with acids to form salts. 
 These salts have the same composition as those of ammonium, with 
 the exception that one atom of hydrogen is replaced by the univalent 
 radicle (the amido group) NH 2 f : 
 
 * J. W. Briihl; Berichte d. Deutsch. Chem. Gesell.; 26, 2508. 
 
 t Salts of hydrazin in which both NH 2 groups were supposed to be united 
 with acids (C1H 3 N NH 3 C1) were formerly described. Such salts are now 
 known not to exist; only one of the NH 2 groups can unite with acids. 
 Curtius; Berichte d. Deutsch. Chem. Gesell.; 26, 409. 
 
AZOIMID. 
 
 193 
 
 fH 
 H 
 H 
 H 
 
 LCI 
 
 Ammonium chloride. 
 
 N 
 
 H 
 H 
 H 
 
 Id 
 
 Hydrazoniurn chloride 
 
 Hydrazin is soluble in water ; its solutions have an alkaline reaction. 
 Azoimid, or hydrogen nitride, N 3 H, is a colorless gas with a 
 peculiar, very penetrating odor ; it is quite poisonous, and its solu- 
 tion in water has an extremely irritating effect upon the skin. A 
 solution of azoimid can be prepared by oxidizing a solution of hy- 
 drazin in water by means of nitrous acid.* The most interesting 
 fact in regard to this compound is that it is a strong acid, greatly 
 resembling hydrochloric or hydrobromic acid, but a short consid- 
 eration will show us the reason for this chemical behavior, for in 
 azoimid the mass of the three nitrogen atoms entirely overbalances 
 that of the one hydrogen atom, and consequently the compound, as 
 a whole, is negative ; we would therefore expect azoimid to be acid 
 in its nature. f Azoimid produces dense white fumes when brought 
 in contact with ammonia, just as hydrochloric acid does ; the sub- 
 stance formed in the one case is ammonium nitride, just as in the 
 other it is* ammonium chloride : 
 
 = NH 4 C1. 
 
 The solution of azoimid attacks copper, aluminium, zinc, and other 
 metals, forming the nitrides and liberating hydrogen; it dissolves 
 oxides and hydroxides of metals. The nitrides formed, in all cases, 
 resemble the corresponding chlorides. Hydrogen nitride differs 
 markedly from hydrogen chloride by being very unstable, for even 
 
 * Curtius; Berichte d. Deutsch. Chem. Gesell.; 26, 1263. 
 
 t As we increase the relative number of hydrogen atoms as compared to 
 each nitrogen atom in the molecules of the three compounds of nitrogen and 
 hydrogen, we pass from an acid to bodies with an entirely opposite chemical 
 character. This change reminds us of the transition from not-metals to met- 
 als which we encounter in the natural families formed by the elements which 
 we have studied with the difference that with the increase of the mass of 
 the not-metallic element in these compounds, the negative properties increase, 
 while in a natural family they diminish. 
 
194 AZOIMID. 
 
 a shock or a slight increase in temperature will cause it to explode 
 with terrific force. It follows from this that the grouping of three 
 nitrogen atoms in this molecule takes place only under great ten- 
 sion, so that the molecule is subjected to a constant strain, just as 
 a wound-up watch-spring is. The structural formula assigned to 
 hydrogen nitride by its discoverer is : 
 
 N. 
 
 II )N-H. 
 
 x/ 
 
 The ammonium salt of azoimide (NH 4 N 3 ), and the similar com- 
 pound of hydrazine (N 5 H 5 ), may be mentioned as being two other 
 compounds composed exclusively of nitrogen and hydrogen. 
 
NITROGEN ; COMPOUNDS WITH OXYGEN. 195 
 
 CHAPTEK XXVI. 
 
 THE COMPOUNDS OF NITROGEN WITH OXYGEN, AND WITH 
 OXYGEN AND HYDROGEN. 
 
 Nitrous oxide; formula, N 2 ; specific gravity, air = 1, is 1.527, 
 H 2 = 2, is 43.98 ; molecular weight is 44. 1 c.c. of the gas weighs 
 .0019835 gram at and .76 m. pressure. Nitric oxide ; for- 
 mula, N ; specific gravity, air = 1, is 1.0384, H 2 = 2, is 29.9 ; 
 molecular weight is 30. 1 c.c. of the gas weighs .0013488 gram. 
 Nitrogen peroxide ; formula, N0 2 ; specific gravity, air = 1, is 
 1.58, H 2 = 2, is 45.5 (at 130 ), molecular weight is 46 (at 140 ). 
 
 NITROGEN forms the following compounds with oxygen and 
 hydrogen : 
 
 1. No 0, Nitrous oxide, HNO, hyponitrous acid, 
 
 2. NO, Nitric oxide, -, 
 
 3. N 2 3 , Nitrogen trioxide, HN0 2 , nitrous acid, 
 
 4. NO* , Nitrogen dioxide, , 
 
 5. N 2 5 , Nitrogen pentoxide, HN0 3 , nitric acid. 
 
 Of these compounds, N 2 0, N 2 3 , and N 2 5 are similar to those en- 
 countered in the study of the halogens, for there we have C1 2 0, 
 C1 2 3 , and L 5 , so that, provided we consider oxygen as being 
 uniformly bivalent, the valence of nitrogen in these oxides is one, 
 three, and five ; the acids derived from these three oxides, HNO, 
 hyponitrous acid, HN0 2 , nitrous acid, HN0 3 , nitric acid, also have 
 formulae like those of the halogen acids ; but no per-nitric acid exists, 
 so that the highest valence displayed by the nitrogen family is five. 
 Nitric oxide does not act like the anhydride of an acid ; it is but 
 little soluble in water, and is not attacked by bases ; neither has 
 nitrous oxide the characteristics of an anhydride, it has no tendency 
 to form hyponitrous acid with water or hyponitrites with alkalies ; 
 but, on the other hand, it is produced when a solution of hyponi- 
 trous acid is warmed, so that it must be looked upon as the anhy- 
 dride of that acid. We will begin the discussion of the oxides of 
 
196 NITROUS OXIDE; PREPARATION. 
 
 nitrogen with nitrous oxide, following with nitric oxide, nitrogen 
 trioxide, nitrogen peroxide, and nitrogen pentoxide in the order 
 named. 
 
 Nitrous oxide never occurs as such ; it is solely a product of the 
 laboratory. The gas was discovered by Priestley in 1776, and was 
 first called dephlogisticated nitric gas ; its composition was not ex- 
 plained until some time after its discovery, when Davy proved it to 
 be an oxide of nitrogen. It is best prepared by heating ammonium 
 nitrate, when water and nitrous oxide are formed as follows : 52 
 
 NH 4 N0 3 = N 2 + 2 H 2 0,* 
 
 but the gas can also be produced by the reduction of nitric oxide by 
 means of finely divided metals, such as zinc, iron, or lead. 
 
 Nitrous oxide is a colorless gas, with a very slight odor and 
 sweBti^h~tg,ste. Its specific gravity, air = 1, is 1.527, which, with 
 hydrogen as two, would give 43.98, so that the molecular weight 
 is, in round numbers, 44. In this weight, analysis shows that there 
 are twenty-eight parts by weight of nitrogen and sixteen of oxygen, 
 so that nitrous oxide contains, in its molecule, one atom of oxygen 
 and two of nitrogen ; for, by means of the study of water and other 
 compounds of oxygen, we have concluded that the atomic weight of 
 oxygen is sixteen, provided that of hydrogen is 1.008 ; and from 
 our study of the composition of ammonia and other nitrogen com- 
 pounds, it follows that the atomic weight of nitrogen is 14. Nitrous 
 oxide has, therefore, a structure similar to that of water, as will 
 be seen by comparing the formulae : 
 
 N N and H H. 
 
 Nitrogen is, therefore, univalent in nitrous oxide, just as hydrogen 
 is in water. From a further study of the composition of nitrous 
 oxide, we see that two volumes of nitrogen will unite with one of 
 oxygen to form two volumes of nitrous oxide, just as two volumes 
 of hydrogen unite with one of oxygen to form two of water ; the 
 conclusions regarding the composition of water, at which we arrived 
 on page 73, are consequently equally applicable to nitrous oxide. 
 
 Nitrous oxide is, to a considerable extent, soluble in water; 
 one volume of water will absorb about 1.3 volumes of the gas at 
 
 * A similar reaction takes place when ammonium nitrite is heated ; only 
 then nitrogen and not nitrous oxide is formed : 
 
 NH 4 NO 2 - 2 N + 2 H. 2 O. 
 
NITROUS OXIDE ; PROPERTIES. 197 
 
 Nitrous oxide will support combustion almost as readily as oxygen ; 
 a glowing pine chip will take fire in the gas, and phosphorus, 
 as well as sulphur, which have been ignited in the air, will con- 
 tinue to burn brilliantly in nitrous oxide. The great tendency 
 to give off oxygen which is displayed by nitrous oxide, is readily 
 understood when we consider that it is an endothermic compound, 
 in the formation of which work which is equivalent to 180 K must 
 be done; the gas, therefore, possesses more energy than its con- 
 stituents, and will break down at the first opportunity. That, 
 however, considerable impulse is required to inaugurate this de- 
 composition is shown by the fact that feebly burning sulphur is 
 extinguished in the gas, while that which is combusting with 
 some energy will continue to burn in nitrous oxide with almost 
 the same brilliancy as if it were placed in oxygen. When a sub- 
 stance which, like phosphorus, forms a solid oxide, burns in nitrous 
 oxide, there is no change in volume, for the molecule of N 2 O sim- 
 ply loses oxygen, while a molecule of K is left in its place ; one 
 thousand molecules of nitrous oxide would therefore yield the same 
 number of molecules of nitrogen, or x molecules of nitrous oxide 
 would yield x molecules of nitrogen ; it therefore follows that the 
 volume of nitrogen which is formed has the same number of parti- 
 cles as the volume of gas from which it is produced, provided the 
 gases have the same temperature and are under the same pressure, 
 but, when tivo gases, under the same temperature and pressure, con- 
 tain equal numbers of molecules, they have equal volumes. 
 
 Nitrous oxide is quite readily condensed to a liquid ; at it 
 becomes fluid under a pressure of thirty atmospheres ; its boiling 
 point, under atmospheric pressure, is 88, it becomes solid at 
 - 100. 
 
 Although nitrous oxide can give up its oxygen so readily to 
 burning substances, it cannot do the same thing in order to support 
 respiration. If the gas is inhaled, the first effect is loss of con- 
 sciousness, accompanied by a rumbling in the ears, while the person 
 undergoing treatment experiences an involuntary tendency to laugh ; 
 as a consequence of this effect Davy named this substance J.aughinj 
 gas. Small animals are very rapidly killed by nitrous oxide. TEe 
 ^ r etfects"of the inhalation of the gas disappear soon after pure air is 
 taken into the lungs and as a consequence it is extensively used 
 as an anaesthetic in place of chloroform or ether. The nitrous 
 
198 NITRIC OXIDE ; PREPARATION. 
 
 oxide used for this purpose is condensed and transported in iron 
 bottles. 
 
 Nitric oxide is the oxide of nitrogen which contains the next 
 greater quantity of oxygen. It results from the action of many 
 metals, or, indeed, of other oxidizable substances on diluted nitric 
 acid ; it does not occur in a free state, for the oxygen of the atmos- 
 phere converts it into a mixture of the two higher oxides, N 2 3 
 and N0 2 . The most convenient method of preparing the gas is by 
 the action of nitric acid on copper. 53 When copper is treated with 
 concentrated nitric acid (spec. grav. 1.4) nothing but nitrogen diox- 
 ide ( N0 2 ) and nitrogen trioxide ( N 2 3 ) are produced as reduction 
 products of the acid.* These gases are formed in the proportion 
 of about 90 per cent of the former and 10 per cent of the latter. 
 As the nitric acid is diluted, the relative quantity of dioxide dimin- 
 ishes and that of trioxide increases, when the specific gravity of 
 the acid has reached 1.25, nitric oxide begins to appear, and, when, 
 more water is added, nitric oxide (mixed with a very little nitrous 
 oxide) is the sole product of the reaction. These changes are very 
 readily understood if the following facts are remembered : 
 
 1. Nitric oxide is oxidized to nitrogen dioxide and to nitrogen 
 trioxide by means of concentrated nitric acid.f 
 
 2. Nitrogen dioxide, on the addition of water, changes to nitric 
 acid and nitric oxide : 
 
 3 N0 2 + H, = 2HN0 3 + NO. 
 
 It is very evident, therefore, that no nitric oxide can be formed 
 from concentrated nitric acid and metals, so that the first reaction 
 between the acid and copper consists in a simple oxidation of the 
 metal, the nitric acid giving up the least possible quantity of 
 oxygen, as will be seen from the equations : 
 
 1. 2 HN0 3 = 2 N0 2 + H 2 + 04 
 
 2. Cu + = CuO. 
 Combining 1 and 2 we have : 
 
 3. 2 HN0 3 + Cu = CuO + 2 N0 2 + H 2 0. 
 
 * Freer and Higley; Amer. Chem. Journ. ; 15, 71. 
 
 t Veley; Proc. Royal Soc.; 52, 27. 
 
 \ When nitric acid oxidizes in such a way that nitrogen dioxide is the 
 reduction product, then two formula weights of nitric acid are capable of fur- 
 nishing one oxygen atom. This fact must be kept in mind in constructing 
 equations. 
 
NITRIC OXIDE; PREPARATION. 199 
 
 One further change must also take place, because copper oxide is a 
 base, and, therefore, in the presence of nitric acid, must form copper 
 nitrate and water : 
 
 CuO + 2 HN0 3 = Cu( N0 3 ) 2 + H, 0. 
 
 The complete reaction between copper and concentrated nitric acid 
 is, therefore, as follows : 
 
 Cu + 4HN0 3 = Cu(N0 3 ) 2 + 2 H 2 + 2NO 2 . 
 Now, if a considerable quantity of water is present, then every 
 three molecules of nitrogen dioxide change two of nitric acid and 
 one of nitric oxide ; and so the following reactions take place : 
 
 a. (3 Cu + 12 HN0 3 = 3 Cu(N0 3 ) 2 + 6 H 2 + 6 N0 2 ) 
 
 b. (6 N0 2 + 2 H 2 = 4 HN0 3 + 2 NO) 
 
 />! V 37 1 C* S^ (/ 
 
 Combining a and b we have : 
 
 3 Cu+12 HN0 3 +2 H 2 = 3 Cu ( N0 3 ) 2 + 4 HN0 3 + 2 NO + 6 H 2 0. 
 
 What is true of copper is true of the other metals as well, with 
 the difference that some of the latter (lead for example) yield 
 nitric oxide, mixed with considerable quantities of nitrous oxide, 
 when they act upon diluted nitric acid. 
 
 As a result of these experiences it is expedient to prepare nitric 
 oxide by the action of copper on dilute nitric acid. 
 
 Reducing agents, which are not of a metallic nature, such as 
 sulphur dioxide, will also produce nitric oxide from nitric acid. 
 For instance, dilute nitric acid oxidizes, dry sulphur dioxide to sul- 
 phuric acid, while at the same time nitric oxide is produced. 
 When nitric acid acts in this way, it is convenient to ignore the 
 intermediary production of nitrogen dioxide and its subsequent 
 decomposition to nitric acid and nitric oxide, and to consider the 
 oxidizing action as being always produced as follows : 
 
 a. 2 HN0 3 = H 2 + 2 NO + 3 0. 
 
 Two formula-weights of nitric acid have, therefore, three atoms 
 of oxygen at their disposal, and so they can oxidize three molecules 
 of sulphur dioxide to sulphuric acid as follows : 
 
 b. 3 S0 2 + 3 H 2 + 3 = 3 H 2 S0 4 . 
 Combining a and b, we have : 
 
 c. 2 HN0 3 + 3 S0 2 + 3 H 2 = 3 H 2 S0 4 + 2 NO + H 2 0. 
 In all other cases where nitric oxide is produced from nitric acid, 
 
200 NITRIC OXIDE ; PROPERTIES. 
 
 the reaction can be regarded as taking place in a manner similar 
 to the above.* 
 
 Nitric oxide is a colorless gas, which instantly turns dark brown 
 on exposure to the atmosphere, nitrogen trioxide and nitrogen di- 
 oxide being formed ; as a consequence there can be no experiment 
 showing whether the gas is tasteless and od.orless as well as color- 
 less. The two gases produced by contact of nitric oxide with the 
 air are poisonous. The specific gravity of nitric oxide is 1.038 
 when air is the standard ; this corresponds to a density of 29.9, 
 H = 2, or to a molecular weight of 30. It follows, as in this molec- 
 ular weight there are fourteen parts by weight of nitrogen and 
 sixteen of oxygen, that the formula of nitric oxide is NO. The 
 gas is composed of equal volumes of nitrogen and oxygen, just as 
 hydrochloric acid is of hydrogen and chlorine. If we wish to con- 
 sider oxygen as bivalent, then nitrogen must also be bivalent in 
 this compound. The specific gravity of nitric oxide does not change 
 even at a temperature as low as 70. At 153. 6 and at atmos- 
 pheric pressure, nitric oxide changes to a colorless liquid which 
 solidifies at 167, forming a snow-like mass. The gas is very 
 stable ; it can be heated to 1200 without alteration ; at white heat 
 it is completely broken down into nitrogen and oxygen. One hun- 
 dred volumes of water dissolve about five volumes of nitric oxide 
 at ordinary temperatures. 
 
 Nitric oxide does not allow substances to burn in it as readily 
 as does nitrous oxide. For instance, phosphorus does not take fire 
 in the gas unless it is heated to a point considerably above its 
 melting temperature ; in the latter event, it will unite with the 
 oxygen of nitric oxide with the greatest energy. On the other 
 hand, sulphur, a burning candle, or burning hydrogen is extin- 
 guished by nitric oxide. A mixture of carbon disulphide and nitric 
 oxide burns with an exceedingly brilliant flame. Metals like zinc 
 or iron, which are easily oxidized, will, if moist, readily remove a 
 part of the oxygen from nitric oxide and in that way produce nitrous 
 oxide : 
 
 * All of the methods which have been given produce nitric oxide mixed 
 with some nitrous oxide. In order to produce pure nitric oxide, a solution of 
 nitrogen trioxide in sulphuric acid is treated with mercury. The mercury 
 will then reduce the nitrogen trioxide to pure nitric oxide. (See Emich; 
 Monatshefte f iir Chemie ; 13, 74. ) 
 
NITROGEN TRIOXIDE ; NITROGEN DIOXIDE. 201 
 
 2 NO + Fe = N 2 + Fe 0. 
 
 Priestley first prepared the latter gas by this method.* 54 
 
 The existence of gaseous nitrous anhydride, N 2 3 , is doubtful, 
 there being strong reason to suppose that, in all cases where chem- 
 ists have endeavored to obtain pure nitrous anhydride, they have 
 only succeeded in producing a mixture of nitric oxide (NO) and 
 nitrogen dioxide ( N0 2 ) ; such a mixture, obviously, would contain 
 the same proportion of nitrogen and oxygen by weight as nitrogen 
 trioxide : 
 
 NO + N0 2 = N 2 3 
 
 The brown gas which results when nitric oxide is mixed with 
 an excess of oxygen, is nearly pure nitrogen dioxide (N0 2 ); but that 
 which is formed by mixing nitric oxide with ah amount of oxygen 
 not sufficient to produce the peroxide undoubtedly consists of a 
 mixture of nitrogen dioxide and nitrogen trioxide. Both of these 
 gases are easily condensed to the liquid form, in which state nitro- 
 gen trioxide certainly can exist. Fluid nitrogen trioxide is an in- 
 digo-colored liquid which is condensed at 10 and which boils 
 below 0, giving off dark brown vapors which change, in part at 
 least, into nitric oxide and nitrogen peroxide. 
 
 Nitrogen dioxide is produced when nitric oxide is exposed to 
 the atmosphere : 
 
 or when the nitrates of certain metals are heated, lead nitrate being 
 the most convenient for this purpose : 55 
 
 Pb (N0 3 ) 2 = Pb + 2 N0 2 + O.f 
 
 * Nitric oxide disolves in a solution of ferrous sulphate, giving a dark brown 
 color. It has been said that, by heating this solution, pure nitric oxide can 
 be driven off. This is not the case, however, as ferrous sulphate reduces a. 
 portion of the nitric oxide to nitrous oxide, while it is itself, in part, oxidized 
 to ferric sulphate. (Emich; Monatshefte fiir Chemie; 13, 73.) 
 
 t The nitrate of lead first, undoubtedly, breaks down into lead oxide and 
 nitric anhydride: 
 
 Pb (NO 8 ) 2 = Pb O + No O 5 , 
 just as nitric acid would break down into water and nitric anhydride: 
 
 but the latter, at the temperature of the reaction, forms oxygen and nitrogen 
 dioxide : 
 
 N 2 5 = 2 N0 2 + O. 
 
202 NITROGEN DIOXIDE ; PROPERTIES. NITROGEN PENTOXIDE. 
 
 The gas has a dark brown color which deepens as the tempera- 
 ture is increased ; it has a corroding action, giving a saff ron coloring 
 to the skin, and other nitrogen-bearing organic compounds. A 
 moderate cold condenses the gas to a yellow liquid, which becomes 
 lighter in color the lower the temperature, and which solidifies at 
 _ 9 to 15 and boils at about 22. Nitrogen dioxide, when it is 
 at a temperature just above the boiling point of the liquid, has a 
 vapor density which indicates that its molecule has the formula 
 N 2 4 ; but these molecules, as the heat is increased, begin to break 
 down into those having the composition N0 2 , so that the specific 
 gravity of this substance diminishes until 140 is reached, when 
 the dissociation of N 2 4 into N0 2 is complete ; at 600 the gas has 
 become entirely colorless and has decomposed into nitric oxide and 
 .oxygen.* Nitrogen dioxide is a powerful oxidizer; carbon and 
 strongly heated phosphorus burn in it, and the presence of this gas 
 dissolved in fuming nitric acid probably gives rise to the powerful 
 oxidizing action of the latter substance. Nitrogen dioxide is 
 changed into nitric acid and nitric oxide when it is dissolved- in 
 water (see page 198), so that the same tendency to form the acids 
 with the greatest possible amount of oxygen, which we observed 
 existing in the halogen and sulphur families, is once more en- 
 countered in the case of the compounds under discussion. 
 
 Nitrogen pentoxide is the anhydride of nitric acid and is best 
 prepared by removing the wafer t'rdm~~concentrated nitric acid by 
 means of phosphoric anhydride : 56 
 2 
 
 r\ "'_": NOo 
 
 D I H J = >0 + H,0 
 iOHi N0 2 
 
 o 2 
 
 2HN0 3 = N 2 5 +H 2 0. 
 
 The compound is a crystalline solid which melts at 30 and boils at 
 45.5, while it is at the same time partially decomposed ; it cannot 
 
 * Nitric oxide, at a temperature above red heat, seems to decompose into 
 nitrogen peroxide and nitrogen, and the resulting nitrogen peroxide is only 
 broken down completely at a temperature produced by a white hot platinum 
 wire. The final result is complete decomposition into nitrogen and oxygen. 
 It follows that the above temperature (600) is open to question. (See Emich; 
 Monatshefte fur Chemie; 13, 79.) 
 
NITRIC ACID; HISTORY. 203 
 
 be kept for any length of time because of its tendency to break 
 down into nitrogen dioxide and oxygen : 
 
 N 2 5 = 2N0 2 + 0. 
 
 Dangerous explosions may be the result of this decomposition, if the 
 pentoxide has been kept in a sealed tube. Nitric anhydride forms 
 nitric acid when it is added to water : 
 
 N 2 5 + H 2 0=2HN0 3 . 
 
 Nitric acid has been known ever since the time of the Arabian 
 alchemists. The first authentic account of its preparation is given 
 by Geber, who, in the ninth century, made it by distilling a mixture 
 of saltpetre (potassium nitrate), blue vitriol (copper sulphate), and 
 alum (aluminium and potassium sulphate). The first samples of 
 nitric acid were undoubtedly an impure article. The name given 
 to it was aqua dissolutiva or aqua fortis, while the term aqua regia, 
 was used to designate a mixture of nitric and hydrochloric acids. 
 Nitric acid, or aqua fortis, the alchemists discovered, had the power 
 of dissolving all known metals with the exception of gold, while 
 aqua regia would attack even this, so-called, noblest of all metals 
 almost nothing could withstand its corrosive action ; surely, 
 thought they, this liquid must be closely allied to the " alcahest," 
 the universal solvent which they were seeking. At the beginning 
 of the eighteenth century, nitric acid was extensively made by the 
 action of sulphuric acid on nitre (saltpetre) ; to this method of 
 preparation it owes its present name, which is derived from spiritus 
 nitri. Lavoisier first proved that nitric acid contained oxygen, and 
 its definite composition was ascertained during the present century. 
 
 Nitric acid can be produced by the direct union of nitrogen, 
 oxygen, and water. Such a synthesis takes place when electric 
 sparks are passed through moist air. 67 In all probability nitrogen 
 peroxide, N0 2 , is at first generated,* however, the latter, when in 
 
 * This method for the preparation of nitrogen dioxide reminds us of the 
 similar one used in forming ozone (see page 48). Nitrogen and oxygen are 
 both the first members of their respective families, there is but little difference 
 between their atomic weights, and hence they should show points of resem- 
 blance, as indeed they do, for they are both colorless gases. Ozone can be 
 considered as the oxide of oxygen, OO 2 ; it then corresponds to the oxides of 
 sulphur, selenium, and tellurium, SO 2 , Se O. 2 , Te O 2 . In the manner of its 
 formation and in its formula it is analogous to nitrogen peroxide, NO 2 ; fur- 
 
204 NITRATES ; FORMATION OF. 
 
 contact with water, breaks down into nitric oxide and nitric acid. 
 (See page 198.) Oxides of nitrogen are also produced during com- 
 bustion in the air ; * these oxides, in contact with moisture, are 
 further converted into nitric acid. As ammonia is generally pres- 
 ent in the atmosphere, this substance, uniting with the nitric acid, 
 produces ammonium nitrate, so that this salt occurs in the air. 
 Nitric and nitrous acids are, however, much more readily formed by 
 the oxidation of ammonia or of the oxides of nitrogen than by the 
 direct union of the elements. When organic substances (which con- 
 tain nitrogen) decay, the nitrogen passes off as ammonia, and this 
 substance, with the acids present in the air and with the carbon 
 dioxide formed at the same time, produces ammonium carbonate, 
 nitrate, and nitrite (page 170 ). When bases are present in the soil, 
 an oxidation of the nitrogen takes place so that nitrates are produced 
 instead of ammonium salts. Calcium nitrate is, as a consequence, 
 frequently found on the walls of stables and cellars, while in the 
 neighborhood of East Indian villages, where the surface soil con- 
 tains potash, potassium nitrate is extensively met with; the col- 
 lecting of this substance forms the exclusive occupation of a 
 number of natives. Large deposits of sodium nitrate occur in 
 the province of Tarapaca.in the northern part of Chile, this sub- 
 stance is known as Chile saltpetre or nitre ; its presence is probably 
 due to the decay of marine vegetation which flourished on what is 
 now terra firma, during the period when a portion of the South 
 American coast was submerged. This supposition is sustained by 
 the fact that sodium chloride and salts containing bromine and 
 iodine are found mixed with the nitre. 
 
 Nitric acid is best prepared for laboratory use by the addition 
 of sulphuric acid to a nitrate, a method which we have frequently 
 employed in the isolation of other acids (see page 152). The reac- 
 tion may be represented by the following equation : 
 
 Na N0 3 + H 2 SO, = Na HS0 4 + HN0 3 , 
 or, if comparatively little sulphuric acid is used : ' 
 
 thermore, being an endothermic compound, it has a great tendency to give 
 up one atom of oxygen, OO. 2 = OO + O, just as nitrogen peroxide does, 
 N0 2 = NO + O. 
 
 * This formation takes place in greatest quantity when hydrogen is 
 burned in air. 
 
NITKIC ACID; PREPARATION, PltOPER/TIES. 205 
 
 2 Na N0 3 + H 2 S0 4 = Na^ S0 4 + 2 HN0 3 , 
 
 for, when the quantity of salt to be decomposed is relatively great, 
 as compared with the amount of acid used, then the secondary and 
 not the primary sulphate results. (See page 153.) 58 
 
 Nitric acid is a colorless liquid which has probably never been 
 prepared entirely free from water. It boils at 86, and becomes 
 solid at 47 ; if it contains water enough to have a specific weight 
 of 1.3 it congeals at 19 ; the purest acid known has a specific 
 gravity of 1.55 ; * it fumes in the air and turns yellow when ex- 
 posed to the sunlight, because it breaks down into nitrogen peroxide, 
 water, and oxygen. The same change takes place when nitric acid 
 is distilled, for the distillate from a colorless, pure acid is colored 
 because of decomposition. This behavior reminds us forcibly of 
 the chlorine acids. At temperatures just above the boiling point 
 of nitric acid, the specific gravity of the vapor shows that but little 
 decomposition has taken place, for, while the molecular weight of 
 HN0 3 would be 63, the specific gravity of the vapor, H = 2, is 59.3 ; 
 the vapor density of the acid diminishes as the temperature is 
 increased, so that at 250 it is 36, therefore, at that temperature, 
 the following change has taken place : 
 
 4 HIS T 3 = 2 H 2 + 4 N0 2 + 2 .f 
 
 Considerable heat is evolved when nitric acid is dissolved in water, 
 so that the dilute acid possesses less chemical energy, and is there- 
 fore more stable, than the concentrated one. It is doubtful if defi- 
 nite hydrated acids, such as are encountered with sulphuric acid, are 
 derived from nitric acid ; certainly the heat of solution of the latter $ 
 is much less than that of the former. (See pages 150, 151.) 
 
 * The pupil must remember that the specific gravities of liquids and solids 
 are taken with water as unity. 
 
 t If a volume of hydrogen weighs two grams, then the same volume of 
 nitric acid weighs 63 grams, 4 volumes of nitric acid would therefore weigh 
 252 grams. These decompose into two volumes of water vapor weighing 36 
 grams, 4 of nitrogen peroxide weighing 184 grams, and 1 of oxygen weighing 
 32 grams. The 4 volumes of nitric acid therefore yield 7 volumes of the de- 
 composition products; these 7 volumes weigh, 2 52 grams, 1 volume equal to 
 that of 2 grams of hydrogen therefore weighs 36 grams; in other words, pro- 
 vided this decomposition takes place, the specific gravity of the mixed gases 
 must be 36, if H = 2. (See pages 72 and 189.) 
 
 J 71 K as compared with 178 K. Some investigators (Wislicenus, Ber- 
 thellot) have maintained that definite hydrated nitric acids exist in solution. 
 
206 NITRIC ACID ; OXIDIZING ACTION. 
 
 i 
 Nitric acid has a great tendency to give up its oxygen when it 
 
 is brought in contact with reducing substances. Examples of this 
 oxidizing effect we have seen in the preparation of sulphuric from 
 sulphurous acid (page 147), and in the formation of nitric oxide 
 from copper and nitric acid (page 199). Nitric acid will attack 
 many organic substances, oxidizing them, while at the same time 
 the acid itself is reduced; when the substance attacked is like 
 starch or sugar and the acid is tolerably concentrated, then the 
 oxides, N 2 3 and N0 2 , are the main reduction products. The 
 organic substance is often completely destroyed, yet in quite a 
 number of cases the body attacked is so changed that the nitro- 
 group, N0 2 (see nitrosyl sulphuric acid, page 147, foot-note), is 
 substituted for hydrogen ; such an action is produced when, under 
 certain circumstances, nitric acid reacts with glycerine, forming 
 nitro-glycerine. Concentrated nitric acid violently attacks the skin 
 and mucous membrane ; that portion with which it has come in con- 
 tact turns yellow, blisters, and finally forms an ulcer ; if the acid is 
 somewhat dilute, the yellow color will appear without the blistering. 
 Nitric acid also attacks silk in the same way, turning it yellow, 
 and, if the acid is concentrated, destroying it; vegetable dyes are 
 destroyed by it, so that cloth upon which nitric acid has accidentally 
 been dropped cannot be restored to its original color by neutraliza- 
 tion with ammonia water. 69 
 
 As we have seen, many metals dissolve in nitric acid to form 
 the corresponding nitrates, a reduction product being produced at 
 the same time. These reactions can practically be classed under 
 two heads. 
 
 a. Those in which ammonia is produced, the ammonia at once 
 uniting with nitric acid to form ammonium nitrate ; this change 
 takes place when dilute nitric acid is added to zinc, tin, or to some 
 other metals; the reaction can be represented in two stages, as 
 follows : 
 
 1. 8 HN0 3 + 4 Zn = 4 Zn (N0 3 ) 2 + 8 H (formation of hydro- 
 gen). 
 
 2. HN0 3 -f 8 H=3 H 2 + NH 3 ; NH 3 + HN0 3 = NH 4 NO $ 
 (reduction of nitric acid and formation of ammonium nitrate). 
 
 Uniting 1 and 2 we have : 
 
 3. 4 Zn + 10 HN0 3 =4 Zn ( N0 3 ) 2 + NH 4 N0 3 + 3 H 2 (com- 
 plete reaction). 
 
NITRIC ACID ; REDUCTION OF. 207 
 
 b. Those in which nitric oxide is formed by the action of dilute 
 nitric acid : 
 
 3 Cu + S HN0 3 = 3 Cu ( N0 3 ) 2 + 2 NO + 4 H 2 0.* ( Page 199.) 
 
 These two classes of reactions, however, only represent what 
 most frequently takes place ; it is known, for instance, that when 
 zinc acts 011 a mixture of nitric and sulphuric acids, a partially 
 oxidized ammonia known as hydroxylamine, NH 2 OH (see page 191), 
 results ; and when copper is dissolved in diluted nitric acid, nitrous 
 oxide, and even nitrogen, may be given off, the production of 
 nitrous oxide increasing with the amount of copper nitrate present. 
 In attempting to construct equations for such reactions we there- 
 fore generally represent only the principal changes which take 
 place. 
 
 The reduction of nitric acid by metals is generally attributed to 
 the action of nascent hydrogen, and, certainly, in some cases this 
 theory is well founded. We cannot enter into the subject very deeply 
 in this text-book, f but the following facts may not be out of place. 
 When a piece of magnesium is dissolved in dilute nitric acid, hy- 
 drogen is at first given off ; the production of hydrogen, however, 
 soon ceases, while the oxides of nitrogen make their appearance ; 
 it also seems very probable that the hydrogen which has been 
 occluded by palladium (page 34) passes through nitric acid, un- 
 changed, until that hydrogen which is supposed to be chemically com- 
 bined with the palladium begins to be liberated ; then the evolution 
 of hydrogen stops, while the lower oxides of nitrogen make their 
 appearance. Apparently, then, hydrogen which is just in the act 
 of being liberated from its compounds has a greater chemical activ- 
 ity than has ordinary hydrogen, so that, whether we regard this 
 hydrogen as acting by reason of its existence as individual atoms 
 or not, there is reason to suppose that in some cases we should con- 
 sider the reduction of nitric acid by metals which are dissolving 
 therein as being caused by hydrogen. The equations given above 
 are intended to illustrate this conclusion. In the reactions which 
 have been described as taking place between copper and nitric acid, 
 
 * The reaction here given represents the simplest form of the combined 
 equations a and b on page 199. 
 
 t For a more complete account of nascent reactions the pupil may refer to 
 M. M. Pattison Muir, Principles of Chemistry. See also page 51. 
 
208 NITRATES ; NITRITES ; HYPONITRITES. 
 
 however, the evidence all seems to point toward a direct oxidation 
 of the metal by the acid (see page 198). 
 
 Nitric acid is a monobasic acid ; it has in its formula weight but 
 one hydrogen atom which can be replaced by metals. The nitrates 
 are all decomposed by heat, the change taking place in one of three 
 ways : 
 
 a. The nitrate is entirely decomposed, as is ammonium nitrate 
 (page 196). 
 
 /?. The nitrate breaks down into the oxide of the metal, oxygen 
 and nitrogen peroxide : Pb ( N0 3 ) 2 = Pb -}- 2 N0 2 -f (see page 
 201). If the oxide of the metal is decomposed by heat, of course 
 nothing but the metal remains. 
 
 y. The nitrate gives off oxygen, leaving the nitrite. This 
 decomposition is confined to the nitrates of very pronounced metals, 
 such as potassium or sodium. 
 
 In these decompositions nitrates differ from chlorates, for when 
 the latter are heated, the perchlorates are quite often produced ; the 
 reason is obvious, nitrogen cannot take up more oxygen than is 
 necessary to form the oxide N 2 5 , which is the anhydride of nitric 
 acid ( page 197), so that no pernitrates can be formed. 
 
 The existence of nitrous acid is doubtful, although the nitrites 
 are stable and well-characterized compounds. When an acid is 
 added to a nitrite, the nitrous acid which is formed at once breaks 
 down into its anhydride and water, and the anhydride is further 
 decomposed, so that N0 2 and NO are produced. By passing im- 
 pure nitrogen trioxide (formed by the reduction of nitric acid) into 
 ice-cold water, a blue liquid, which possibly is nitrous acid, is pro- 
 duced, but the slightest increase in temperature causes the latter to 
 change into nitric acid and nitric oxide. 
 
 We have already studied the manner in which nitrites are 
 formed by heating nitrates, so nothing more need be added except- 
 ing the statement that many nitrites can best be prepared from 
 potassium nitrite by double decomposition ( page 57), while potas- 
 sium nitrite is produced by heating nitrate of potassium with lead. 
 
 The hyponitrites and hyponitrous acid alone remain for dis- 
 cussion. 
 
 The hyponitrite of potassium can be prepared by reducing ni- 
 trate of potassium with sodium amalgam* the latter substance, 
 
 * For sodium amalgam see foot-note, reference. 32 of appendix. For the 
 preparation of hyponitrous acid from hydroxylamin and nitrous acid see Wisli- 
 cenus, Ber. d. Deutsch. Chem. Gesell. 26 ; 771, and Paal, ibid. 1027. 
 
NITROGEN OXIDES AND ACIDS; TABLE OF. 
 
 209 
 
 when in contact with water, forms sodium hydroxide and hydro- 
 gen, while hydrogen in the nascent state robs the potassium nitrate 
 
 of its oxygen : 
 
 KN0 3 + 4 H = KN T + 2 H 2 . 
 
 The hyponitrite of silver, which is insoluble in water, can be pre- 
 pared by adding silver nitrate to potassium hyponitrite : 
 
 Ag N0 3 + KNO == KN0 8 + Ag NO, 
 
 and the free acid can be formed from the latter by the addition of 
 hydrochloric acid : 
 
 Ag NO + H Cl = Ag Cl + HNO.* 
 
 Hyponitrous acid exists only in very dilute solutions; when warm 3d 
 or when allowed to stand, it decomposes, yielding nitrous oxide (its 
 anhydride) and water : 
 
 2 HNO = H 2 + N 2 . 
 The acid is of no practical importance. 
 
 In the following table the nitrogen compounds are compared with 
 those of chlorine : 
 
 OXIDES. 
 
 ACIDS. 
 
 NAMES. 
 
 OXIDES. 
 
 ACIDS. 
 
 NAMES. 
 
 C1 2 O 
 
 HO Cl 
 
 Hypochlorous acid 
 
 N 2 O 
 
 HO N* 
 
 Hyponitrous acid 
 
 C1 2 3 
 
 HO 2 C1 
 
 Chlorous " 
 
 N 2 3 
 
 HO 2 N 
 
 Nitrous 
 
 (Cl 2 6 )t 
 
 HO 3 C1 
 
 Chloric " 
 
 N 2 6 
 
 HO 3 N 
 
 Nitric " 
 
 (Cl 2 7 )t 
 
 H0 4 C1 
 
 Perchloric " 
 
 
 
 
 
 
 
 * Silver chloride is insoluble in water. The molecules of hyponitrous acid 
 have, probably, a molecular weight which is double that of the formula weight 
 ( Ho N 2 O 2 ). This conclusion is rendered probable from a study of some of the 
 organic derivatives of hyponitrous acid. See Meyer and Jacobson; Lehrbuch 
 der Organischen Chemie ; I. 207. 
 
 t .Oxides C1. 2 O 5 , C1. 2 O 7 do not exist, the corresponding acids, HO 3 Cl, 
 HO 4 C1, do. 
 
 | Hyponitrous acid breaks down into its anhydride, N 2 O, and water, but 
 cannot be formed by dissolving N 2 O in water. It probably has a molecular 
 weight represented by the formula ( HON ) 2 . 
 
 Nitrous acid is stable only in very cold water. The existence of the 
 anhydride N 2 O 3 is doubtful. 
 
 The salts of these acids are much more stable than the acids themselves. 
 All of the acids are powerful oxidizers, all of the oxides are unstable. Those 
 of chlorine are explosive, those of nitrogen support combustion. 
 
 Cl O 2 and NO 2 are not the anhydrides of acids. The former on addition 
 of water forms chloric and chlorous acids, Cl O 2 + H 2 O = H Cl O 2 +H Cl O 3 ; 
 
210 
 
 NITROGEN OXIDES; HEATS OF FORMATION. 
 
 The heats of formation of the oxides of nitrogen, as far as they 
 have been ascertained, are given in the following table : 
 
 N 2 
 
 N 
 
 N 2 
 
 -180K 
 210 K 
 - 77 K 
 131 K 
 
 Nitric oxide should be less stable than nitrous oxide ; as a 
 consequence, NO is changed to N 2 O by moist iron filings, 
 zinc dust, etc. NO 2 is more stable than NO, and is produced 
 therefrom readily by the addition of oxygen. The higher ox- 
 ides are more stable than those with less oxygen. N 2 O s is 
 an exothermic compound ; it is a crystalline solid which can 
 easily be formed from nitric acid by extracting water. 
 
 HNO 2 
 HN0 3 
 
 308 K* 
 491 K* 
 
 Nitric acid has a greater heat of formation than has ni- 
 trous acid; it is therefore the acid of oxygen and nitrogen 
 which is most easily formed. The same rule is observed in 
 the halogen and oxygen families, where those acids which 
 contain the most oxygen are the most stable. 
 
 the latter nitric acid and nitric oxide, 3 NO 2 + H 2 O = 2 H NO 3 + NO. At 
 low temperatures NO 2 becomes N 2 O 4 , and probably Cl O 2 becomes C1 2 O 4 . 
 
 The acids are all unibasic. 
 
 * Acids in solution. 
 
PHOSPHORUS; OCCURRENCE. 211 
 
 CHAPTER XXVII. 
 
 PHOSPHORUS AND PHOSPHINE. 
 
 Phosphorus ; symbol, P ; atomic weight, 31 ; specific gravity of yellow 
 phosphorus, 1.83, of red phosphorus, 2.1. Specific gravity of 
 vapor, air = 1, is 4.16, H 2 = 2, is 119.80 ; molecular weight, 124 ; 
 molecule, P 4 . Phosphine ; formula, PH 3 ; specific weight, air 
 = 1, is 1.185, H 2 = 2, i's 34.12 ; molecular weight, 34 ; 1 c.c. o/ 
 
 A,e gas at 0and .76 m. weighs .0015276 gram. 
 PHOSPHORUS never occurs, as such, in nature ; indeed, such a 
 possibility is precluded by the chemical nature of the element, as 
 an example of which we have but to recall the energy with which it 
 burns in oxygen. The compounds of phosphorus which are most 
 frequently found are : 
 
 Apatite, a combination of calcium phosphate and calcium chloride (or 
 fluoride), Ca 3 ( PO 4 ) 2 , Ca C1 2 . 
 
 Osteolite, calcium phosphate, Ca 3 ( PO 4 ) 2 . 
 
 Vimanite, ferrous phosphate, Fe 3 ( PO 4 ) 2 + 8 H 2 O. Phosphates of alu- 
 minium and of lead also occur. 
 
 Phosphates are always present in the soil ; they are essential to 
 the growth of plants and are taken up by the roots, so that plant 
 ashes, especially those of the cereals, often contain large quantities 
 of the phosphates of calcium and magnesium ; the latter Wd their 
 way into the animal organism (calcium phosphate is the most im- 
 portant inorganic constituent of the bones) ; the waste products 
 are returned to the soil by means of the solid excrements and the 
 urine; so that, as manure, they are once more brought into the 
 proper condition to play their part in plant growth. These changes 
 taking place with phosphoric dteid remind us forcibly of the similar 
 ones encountered with carbon dioxide and ammonia ; none of these 
 necessary substances are the permanent property of any one organ- 
 ism ; they are simply borrowed for a time, and must be returned to 
 the place from which they came. 
 
 Phosphorus was discovered by a Hamburg alchemist named 
 Brand, who accidentally prepared the element while searching for 
 
212 PHOSPHORUS; HISTORY, PROPERTIES. 
 
 the philosopher's stone. Subsequently Kunkel published an account 
 in which he described a method of obtaining the substance, but 
 until the middle of the last century the supply was so small that 
 phosphorus was a very expensive article ; it was exclusively pre- 
 pared from decaying urine, and the price in England was ten ducats 
 an ounce. At a later date a method was discovered by which phos- 
 phorus could be obtained from the calcium phosphate procured 
 either from mineral deposits or from bones ; but even then it was 
 mainly kept as a curiosity until the introduction of matches ren- 
 dered its cheap production necessary. At the present time phos- 
 phorus is prepared from bones by first burning the latter in order 
 to destroy the organic matter contained in them ; the calcium phos- 
 phate is then changed to the primary phosphate of calcium 
 (Ca (H 2 P0 4 ) 2 ) by means of sulphuric acid. Primary phosphates 
 are soluble in water, so that a solution can be formed which is 
 further evaporated and heated, by which means the primary phos- 
 phate of calcium loses water and is converted into calcium meta- 
 phosphate (Ca (P0 3 ) 2 ) ; and the latter substance, when heated 
 with charcoal and sand, yields phosphorus. 
 
 Phosphorus exists in two allotropic forms,* the most common of 
 which is a slightly yellow, wax-like solid, which becomes brittle 
 when cold, and which is readily soluble in carbon bisulphide ; it 
 melts at 44 and boils at 250, forming a colorless vapor which has 
 a specific gravity of 4.16 at red heat. This, with hydrogen as two, 
 gives 119.8, while the molecular weight of P 4 would be 124. The 
 observed specific gravity is therefore somewhat less than the molec- 
 ular weight, 124, a fact which probably finds its explanation in 
 the decomposition of some of these complex molecules into simpler 
 ones. At white heat the specific gravity of phosphorus vapor has 
 diminished to 3.14, so that at this temperature nearly all of the P 4 
 molecules have dissociated into those having the composition P 2 . 
 
 When a solution of ordinary phosphorus in carbon bisulphide is 
 exposed to the sunlight, the other, insoluble, red, amorphous modifi- 
 cation of the element gradually separates. This change can be 
 accomplished more quickly and effectually by heating phosphorus 
 to about 300 in a closed vessel ; 60 the same transformation also 
 
 * A form of phosphorus resembling flowers of sulphur has been prepared 
 by rapidly cooling phosphorus vapors. This may be a third allotropic form of 
 phosphorus. 
 
PHOSPHORUS ; ALLOTROPIC FORMS OF. 213 
 
 occurs through the influence of electricity. Amorphous phosphorus 
 is a dark red substance which is generally produced in the form of 
 a powder, the specific gravity of which is 2.1. When heated to 
 a temperature of 358 in a vacuum, and even more rapidly at 445, 
 amorphous phosphorus is changed back to the yellow variety ; the 
 kindling temperature of the former about coincides with this point. 
 The transformation of the element from its ordinary crystalline 
 form into the amorphous one is accomplished only when the phos- 
 phorus is under a pressure of several atmospheres and at a higher 
 temperature. 
 
 Red phosphorus is perfectly insoluble in carbon bisulphide, 
 ether, and similar substances by which the other allotropic form is 
 readily dissolved. It can be exposed to the atmosphere for any 
 length of time without change* while the other variety will absorb 
 oxygen, melt, and, under proper conditions, may take fire spontane- 
 ously. Yellow phosphorus, when placed in warm, moist air and in 
 the dark, emits a pale white light f which is in part due to the 
 slow oxidation of the element. \ 
 
 Yellow phosphorus is an intense poison ; even small doses cause 
 local inflammations in various organs of the body, and have a sec- 
 ondary effect on the nervous system. The serious symptoms caused 
 by poisoning with phosphorus first become apparent some hours 
 after taking ; they manifest themselves by intense pain in the gas- 
 tric region, finally extending throughout the entire abdomen; the 
 vomit will contain phosphorus, have a peculiarly garlic-like odor, 
 and will be luminous in the dark; the patient is restless, fearful, 
 and trembling. The post-mortem examination reveals inflammation 
 of the mucous membrane of the stomach, accompanied by fatty de- 
 generation of the liver, kidneys, and heart. Fatal doses are from 
 .2 to .5 gram. Cases of phosphorus poisoning are not uncommon, 
 as phosphorus mixed in a dough made of cold water and flour is 
 frequently used as a rat-poison ; this has especially been the case 
 since the element has become quite cheap by reason of its use in 
 the manufacture of matches, the heads of a number of varieties of 
 which are made of a mixture of gum arabic and phosphorus. 
 
 Yellow phosphorus is soluble in carbon bisulphide, ether, and 
 
 * This statement is denied by Pedler, Journ. Chem. Soc. 1890, 608. 
 t So-called phosphorescence. 
 
 J That this is not entirely so is proven by the fact that phosphorus is not 
 luminous in dry oxygen below 20 C., or in that gas under pressure. 
 
214 PHOSPHINE; PREPARATION. 
 
 ethereal oils ; it is insoluble in alcohol and water, but volatile in 
 the vapors of the latter. When slowly oxidized in moist air it 
 changes to phosphorous acid : 
 
 2P+3 = P 2 3 , 
 
 P 2 3 + 3 H 2 0=2 P0 3 H 3 , 
 
 this oxidation is supposed to be the cause of the phosphorescence of 
 the element. When burned, both yellow and red phosphorus yield 
 phosphorus pentoxide, which can further unite with water to form 
 phosphoric acid : 
 
 2P + 50 = P 2 5 , 
 
 P 2 5 + 3 H 2 = 2 P0 4 H 3 . 
 
 The element will combine with chlorine, bromine, or any of the 
 halogens, just as it will with oxygen. (See page 63.) Both of 
 the oxides of phosphorus are anhydrides, and all of the halogen 
 compounds are decomposed by water. (See pages 80, 84.) 
 
 Phosphorus forms three compounds with hydrogen, PH 3 , phos- 
 phine, P 2 H 4 , liquid hydrogen phosphide, and P 4 H 2 , solid hydrogen 
 phosphide. 
 
 Phosphine is a gas which bears the same resemblance to ammo- 
 nia that hydrogen sulphide does to water. As ammonia is formed 
 with difficulty by the direct union of nitrogen and hydrogen, we 
 would scarcely expect phosphine to be produced in a similar way, 
 and yet the compound seems to be readily procured as a result of 
 the action of nascent hydrogen upon phosphorus.* This unex- 
 pected result is possibly due to the fact that the breaking stress 
 of the molecules of phosphorus is less than the same for those of 
 nitrogen. The best method of preparing phosphine for laboratory 
 use is by decomposing calcium phosphide with water or dilute acids. 
 The formula of calcium phosphide has not been definitely ascer- 
 tained, but we can compare this reaction with similar ones in which 
 hydrochloric acid or hydrogen sulphide has been produced by the 
 action of an acid upon a chloride or a sulphide. Another way, 
 which has less to recommend it, but which is more frequently used, 
 
 * By throwing small pieces of phosphorus into a flask in which zinc and 
 dilute sulphuric acid, or tin and sulphuric acid, are generating hydrogen, the 
 temperature being about 70. Compare J. Brossler; Berichte der Deutschen 
 Chemischen Gesellschaf t ; 1881, 1757. 
 
PHOSPHINE; PROPERTIES. 215 
 
 is by heating small pieces of phosphorus in a solution of potassium 
 hydroxide. 61 * 
 
 The gas formed by either of these methods is a mixture of 
 hydrogen compounds having the formulae of PH 3 and P 2 H 4 (unless 
 concentrated hydrochloric acid is used to decompose the calcium 
 phosphide, in which event the substance corresponding in structure 
 to ammonia is alone produced). This mixture of gases takes fire 
 spontaneously when it comes in contact with the air, while pure 
 phosphine, PH 3 , does not possess this property. The spontane- 
 ously inflammable gas can be altered in this respect by passing 
 it through a tube cooled with snow and salt, for by this means 
 the liquid hydrogen phosphide (P 2 H 4 ) is condensed, while the 
 phosphine passes on, to be used as occasion requires. 
 
 Phosphine is a colorless gas with an intensely disagreeable, 
 garlic-like odor. Its specific gravity, air = 1, is 1.185, which, 
 H = 2, is 34.12. The molecular weight of PH 3 is therefore 
 34.024, for analysis has proven that in phosphine there are 31 
 parts of phosphorus and 3.024 of hydrogen by weight. It fol- 
 lows that 31 represents the maximum value for the atomic weight 
 of phosphorus, for, as the molecular weight of phosphine is known, 
 we cannot imagine any atomic weight for phosphorus greater than 
 this number without believing that we have a fraction of an atom 
 of phosphorus in PH 3 . When phosphine is heated, or when elec- 
 tric sparks are passed through it, the gas breaks down into phos- 
 phorus and hydrogen; in this case two volumes of hydrogen 
 phosphide yield three of hydrogen, the phosphorus, being solid, 
 when separated exerts no influence on the volume of the gas as a 
 whole. (Compare pages 99 and 138.) 
 
 H 
 
 r = TT + phosphorus. 
 
 HP 1S H * 
 
 H 
 
 * The reaction is said to take place as follows : 
 
 3 KOH + 4 P + 3 H 2 =3 KH 2 PO. 2 + PH 3 . 
 
 The primary hypophosphite of potassium (KPO + H. 2 O) would thus be 
 formed. The phosphine generated always contains hydrogen, so that its for- 
 mation is probably due to that element acting in the nascent state. 
 
216 PHOSPHINE; PHOSPHONIUM COMPOUNDS. 
 
 From this equation it is evident that two molecules of phosphine 
 produce three of hydrogen, and the terms " volume " and " mole- 
 cule " can be used interchangeably, as we saw on page 70. 
 
 Phosphine can be mixed with pure oxygen without taking fire, 
 but if the pressure is suddenly diminished the gases will explode. 
 In the air the kindling temperature is 149,* the products of the 
 combustion are phosphoric anhydride and water : 
 
 2PH 3 + 80 = P 2 5 +3H 2 0, 
 
 and these two substances naturally combine to form phosphoric 
 acid : 
 
 P 2 & + 3H a O = 2H 3 P0 4 . 
 
 Of course, chlorine, bromine, or iodine would act on phosphine in 
 a manner parallel to the action of oxygen,f the products of the 
 reaction being halhydric acids and the corresponding halogen com- 
 pounds of phosphorus ; for example : 
 
 PH 3 + 6C1 = PC1 3 + 3HC1. 
 
 So readily is phosphine decomposed, that even sulphur when 
 placed in contact with that substance, changes it to the sulphide 
 of phosphorus and hydrogen sulphide, so that here we have an 
 instance in which sulphur causes changes similar to those produced 
 by oxygen and the halogens. It need scarcely be added that con- 
 centrated sulphuric or nitric acid will decompose phosphine just as 
 they would hydrobromic acid, hydroiodic acid, or sulphuretted hy- 
 drogen. When phosphine is passed into solutions of metallic salts 
 it, in many cases, produces the corresponding phosphides of the 
 metals (compare sulphuretted hydrogen, page 99). Phosphine is 
 sparingly soluble in water, and is very poisonous, so that all opera- 
 tions in which it is generated must be conducted either in the open 
 air, or under a hood with a strong draught. The gas changes to a 
 liquid at 85 and becomes solid at 133. 
 
 Phosphine can unite with the halhydric acids to form phospho- 
 nium compounds exactly as ammonia does in the production of 
 ammonium salts, the group PH 4 being termed phosphonium for the 
 same reason that NH 4 is ammonium : 
 
 * The friction of the glass-stopper in the neck of a bottle filled with phos- 
 phine may be sufficient to ignite the gas. 
 
 t Chlorine and phosphine, when mixed, explode very violently. 
 
LIQUID HYDROGEN PHOSPHIDE. 217 
 
 PH 3 + HI = PH 4 1, phosphonium iodide, 
 NH 3 + HI = NH 4 1, ammonium iodide. 
 
 As phosphine is readily oxidized and is much less basic in its 
 character than is ammonia,* it will not unite with acids containing 
 oxygen ; in this respect it differs from ammonia. The phosphonium 
 compounds are readily decomposed by water or by alkalies : t 
 
 PH 4 I + H 2 = PH 3 + HI+H 2 0, 
 PH 4 I + KOH = PH 3 + KI + H 2 0, 
 
 the latter reaction being exactly like those observed with ammonium 
 
 salts : 
 
 + KOH =NH 3 + KI + H 2 0. 
 
 It follows from the above that phosphonium salts cannot be formed 
 where water is present. 
 
 The compound P 2 H 4 is a liquid at ordinary temperatures. It is 
 isolated, as was stated above, by cooling the mixture of gases 
 obtained by one of the ordinary methods in use for the production 
 of phosphine. It is a colorless, highly refractive liquid, which boils 
 at about 35, and which takes fire spontaneously when exposed to 
 the air. The determination of the specific gravity of the vapor 
 shows it to have a molecular weight of 66 ; its formula is, there- 
 fore, PH 2 PH 2 ; the analogous compound of nitrogen is hydra- 
 zin, NH 2 NH 2 ; but, unlike the latter, liquid hydrogen phosphide 
 is not basic, and can, therefore, form no salts. This fact is not sur- 
 prising if the same diminution of basic properties takes place with 
 the hydrogen compounds of phosphorus, as was observed in the case 
 of the similar ones containing nitrogen. ( Page 193, foot-note.) 
 
 A solid compound, P 2 H, is formed by treating phosphine with 
 chlorine which has been highly diluted with carbon dioxide. By 
 this means a part of the hydrogen of phosphine is removed, and a 
 yellow powder is produced, which, when dry, can be heated as high 
 as 150 without taking fire. 
 
 * See page 177. 
 
 t Phosphonium bromide or phosphonium iodide is more easily formed 
 than is phosphonium chloride. The latter salt is prepared by subjecting a 
 mixture of equal volumes of phosphine and of hydrochloric acid to a pressure 
 of 20 atmospheres; phosphonium iodide is produced by direct union at ordinary 
 pressure. 
 
218 
 
 HYDROGEN PHOSPHIDES; TABLE OF. 
 
 The following table shows the relationship between the com- 
 pounds just discussed and the corresponding ones containing 
 nitrogen : 
 
 NH 3 , ammonia, 
 N 2 H 4 , hydrazin, 
 
 N, H, azoimid. 
 
 PH 3 , phosphine. 
 P 2 H 4 , liquid hydrogen : 
 phosphide. 
 
 NH 3 + HX = NH 4 X, Ammonium salts. 
 PH 3 + HX = PH 4 X, Phosphonium salts. 
 
 Phosohonium fluoride has not. as vet. 
 
 P a H, solid hydrogen phos- 
 phide. 
 
 been prepared; with this exception 
 X represents any halogen. 
 
 The compounds of nitrogen are all gases; they are not spontaneously 
 inflammable, while those of phosphorus are either gaseous, liquid, or solid, 
 and burn with the greatest ease. 
 
PHOSPHORUS; HALIDES OF. 
 
 219 
 
 CHAPTEE XXVIII. 
 
 THE COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS, 
 AND WITH OXYGEN AND THE HALOGENS. 
 
 As the atomic weights in this family increase, an increasing 
 stability of the compounds formed with the halogens is observed. 
 Those of nitrogen are very explosive substances ; but in the case 
 of the element under consideration a number of quite stable chlo- 
 rides, bromides, and iodides have been accurately studied ; indeed, 
 some of these can be classed among our most important laboratory 
 reagents. They are given in the following table : 
 
 FLUORIDES. 
 
 CHLORIDES. 
 
 BROMIDES. 
 
 IODIDES. 
 
 PF 3 , phosphorus trifluoride, 
 PF 8 , phosphorus pentafluo- 
 ride, 
 An iodide of phosphorus, PI 
 
 P C1 3 , trichloride, 
 P Cl s , pentachloride, 
 
 , phosphorus di-iodide, 
 
 P Br 3 , tribromide, 
 P Br 6 , pentabro- 
 mide, 
 
 exists. 
 
 PI 3 , triiodide. 
 
 
 The above compounds are all substances which, because they 
 are the halides of a not-rnetal, are readily decomposed by water to 
 form the corresponding acid of phosphorus, together with the 
 hydrogen compound of the halogen which was united with that 
 element. We have seen that this instability in the presence of 
 water has been made use of in the preparation of hydrobromic and 
 hydroiodic acids. (Pages 80 and 85.) 
 
 The character of the trihalogen compounds changes somewhat 
 w r ith the nature of the halogen, the boiling point increases with the 
 increase of the molecular weight, just as it does in the case of the 
 free elements, while tne readiness with which these substances are 
 decomposed is also, apparently, greater in the bromide and iodide 
 than it is in the fluoride and chloride. 
 
 P F 3 is a gas, liquid at 10 under a pressure of 40 atmospheres. 
 P C1 3 is a liquid which boils at 76, heat of formation 755 K. 
 P Br 3 is a liquid which boils at 175, heat of formation 448 K. 
 P I 3 is a solid which melts at 55 and is decomposed by boiling. 
 PF 3 , specific gravity of vapor, air = 1, is 3.02, which, H 2 = 2, is 88; the 
 molecular weight is 88. 
 
220 PHOSPHORUS ; TRIHALIDES J PENTAHALIDES. 
 
 P C1 3 , specific gravity of vapor, air = 1, is 4.8, which, H 2 = 2, is 138; the 
 molecular weight is 137.35. 
 
 PBr 3 , specific gravity of vapor, air = 1, is 9.7, which, H 2 = 2, is 279.3; the 
 molecular weight is 270.88. 
 
 From these determinations of the specific gravities of the vapors it 
 is evident that the general formula of all of the trihalides of phos- 
 phorus is PX 3 , so that phosphorus is trivalent in these compounds, 
 just as it is in phosphine, or just as nitrogen is in ammonia. The 
 trihalides of phosphorus are all formed by treating phosphorus with 
 an amount of halogen insufficient to produce the compounds PX 5 , * 
 and when they are decomposed by water they break down as 
 follows : 
 
 X + HOH f OH 
 
 X + HOH P ] OH + 3 HX, 
 
 X + HOH ( OH 
 
 so that phosphorous acid results in all cases. 
 
 Phosphorus is unsaturated in the trihalogen compounds, and it 
 is therefore capable of a further addition to form pentahalides, the 
 valence of the element increasing from three to five : 
 PX 3 + 2 X = PX 5 . 
 
 P F 6 is a gas which liquefies at 16, 46 atmospheres pressure. 
 
 P C1 6 is a solid which melts at 148 (under diminished pressure), and which 
 boils at about 160 165, with partial decomposition. 
 
 P Br 5 is a solid which decomposes into bromine and phosphorus tribromide 
 .at 100. 
 
 These compounds, when added to water, yield phosphoric acid, 
 as will be seen from the following : 
 
 X + HOH r_0 H 
 
 X + HOH __0 H 
 
 X + HOH = P^l H+5HX. 
 
 X + HOH |_0 H 
 
 X + HOH [ H 
 
 From the above equation we would expect the production of normal 
 phosphoric acid, but, as we have already seen, the normal acids 
 have the greatest tendency to separate water, by this means yield- 
 ing more stable compounds (pages 130 and 131), so that P (OH) 6 
 breaks down to form ordinary phosphoric acid : 
 
 * Excepting the trifluoride which is formed by a somewhat complicated 
 process. 
 
PHOSPHORUS; OXY-HALIDES OF. 221 
 
 P(OH) 5 = P0 4 H 3 + H 2 0. 
 
 The difference between the tri- and pentachloride of phosphorus 
 lies in the different amounts of chlorine as compared to the 
 quantity of phosphorus contained in each ; in the trichloride 
 phosphorus is trivalent, in the pentachloride it is quinquivalent; 
 the one yields phosphorous acid, the other phosphoric acid by 
 the addition of water, so that, plainly, the same difference in the 
 valence of phosphorus exists in these acids as is found to exist 
 in the valence of the element in the two chlorides. 
 
 Compounds containing both bromine and chlorine can be formed 
 by adding bromine to the trichloride, or chlorine to the tribromide 
 of phosphorus. Under proper conditions a portion of the chlorine 
 or bromine in the pentachloride or bromide of phosphorus can be 
 replaced by oxygen ; the result is the production of an oxy-chloride 
 or bromide of phosphorus. Phosphorus oxy-chloride has a chemical 
 character which is analogous to that of the similar compounds of 
 sulphur discussed 011 page 145. 
 
 POC1 3 phosphorus oxychloride, liquid, boils at 110. 
 
 PO Br 3 phosphorus oxybromide, solid, melts at 55, and boils at 193. 
 
 These substances are produced by adding the calculated amount 
 of water to the pentachloride or bromide of phosphorus : 
 
 PX 5 + H 2 = POX 3 + 2 HX. 
 
 In effecting this change, one atom of oxygen has taken the 
 place of two atoms of chlorine or bromine, so that the constitution 
 of these compounds is as follows : 
 
 -x 
 
 p- 
 
 n- 
 
 t-x 
 
 and this structural formula is further supported by the vapor densi- 
 ties of the oxychloride and oxybromide, which exactly agree with 
 the theory. 
 
 Both phosphorus oxychloride and bromide are converted into 
 ordinary phosphoric acid by the addition of water, therefore the 
 latter compound contains three hydroxyl groups and is formed as 
 follows : 
 
222 PHOSPHORUS OXYCHLOKIDE. 
 
 =o r=o 
 
 -Cl + HOH I -OH 
 
 p - _ p - 
 
 -Cl + HOH OH" 
 
 Cl + HOH L OH 
 
 All of the halogen compounds of phosphorus fume in the air, 
 because moisture decomposes them while liberating hydrogen chlo- 
 ride, bromide, or iodide. Phosphorus oxychloride and oxybromide 
 have an indescribably unpleasant odor, so that all work with these 
 substances, as well as with the pentahalogen compounds, must be 
 so conducted that the vapors cannot be inhaled. 
 
PHOSPHOKUS ; TKIOXIDE. 223 
 
 CHAPTER XXIX. 
 
 THE COMPOUNDS OF PHOSPHORUS WITH OXYGEN, AND WITH 
 OXYGEN AND HYDROGEN. 
 
 PHOSPHORUS forms four oxides ; two, P 2 8 and P 2 5 , correspond 
 to nitrogen trioxide, N 2 3 , and nitrogen pentoxide, N 2 6 , one, P0 2 , 
 is analogous to N0 2 , and the fourth oxide probably has the for- 
 mula P 4 0. Some other oxides of phosphorus have been described, 
 but further investigation must establish their identity. Phosphorus 
 trioxide and pentoxide are both acidic anhydrides, the one of phos- 
 phorous, the other of phosphoric acid. The most common forms of 
 these acids differ from the corresponding ones of nitrogen by being 
 hydrated, so that phosphorous acid is not HP0 2 but H 3 P0 3 , and 
 phosphoric acid not HP0 3 but H 3 P0 4 . 
 
 Phosphorus trioxide is produced by slowly oxidizing phosphorus 
 in a stream of oxygen diluted with carbon dioxide. It is a crystal- 
 line solid which melts at 22.5 and which boils at 173. 3, being 
 changed to a colorless vapor which has a specific gravity, air = 1, 
 of 7.6; this indicates a molecule of the formula P 4 6 , as the spe- 
 cific gravity calculated for P 2 3 is 3.8.* The oxide is completely 
 decomposed when heated to 300, at which temperature phosphorus 
 and the oxide P 2 4 are produced. Phosphorus trioxide is oxidized 
 when brought in contact with oxygen ; the action may even become 
 so violent as to cause spontaneous combustion to ensue ; f the oxi- 
 dation product is phosphorus pentoxide ; the latter substance, as we 
 have seen, is also produced when phosphorus is burned in air or 
 oxygen (page 22). 62 Phosphorus pentoxide is a flaky, not crystal- 
 line powder ; under certain circumstances it can be obtained in a 
 
 * Thorpe and Tutton, Journ. Qhem. Soc. ; 1890 ; 545. 
 
 t This phenomenon is possibly due to the fact that a mixture of phosphorus 
 trioxide, pentoxide, and phosphorus is present after slow oxidation of phos- 
 phorus ; the phosphorus would then cause the spontaneous combustion of the 
 mass. Pure phosphorus trioxide unites with oxygen and becomes luminous 
 when placed in the gas and under diminished pressure ; the glowing ceases 
 when the pressure is increased. This fact reminds us of the similar phenom- 
 enon observed with phosphorus and oxygen. 
 
224 PHOSPHOROUS ACID. 
 
 crystalline form. Phosphorus pentoxide greedily absorbs moisture 
 from the air; it is therefore deliquescent; its tendency to unite 
 with water is so great that, if a little of it is placed in that liquid, 
 it dissolves with a hissing noise similar to that which is heard when 
 a red-hot iron is immersed. Phosphoric anhydride, because it is 
 able to absorb perfectly all moisture from gases, is a favorite labora- 
 tory reagent for drying. It is the more valuable, because the usual 
 drying agent, calcium chloride,* does not completely remove water 
 from substances with which it is brought in contact. Phosphorus 
 pentoxide is quite volatile if it is carefully heated to 250 ; but 
 above that temperature it changes into another, so-called poly- 
 meric f form, which evaporates very slowly below bright red heat. 
 
 Phosphorous acid is produced when the trioxide or trichloride 
 is dissolved in water ; it is analogous to nitrous acid, although it is 
 much more stable. When phosphorus trioxide unites with water 
 the first product which we should expect would be H P0 2 , f or : 
 
 P 2 3 + H 2 = 2 H P0 2 (see page 117). 
 
 The compound so produced, however, takes up one more molecule 
 of water to form the hydrated acid H 3 P0 3 : 
 
 HP0 2 + H 2 = H 3 P0 3 , 
 
 in which condition only, the acid is capable of existence. 
 
 Phosphorous acid contains three hydrogen atoms in each for- 
 mula weight, but no more than two of these can be replaced by 
 metals at the same time. The following explanation of this phe- 
 nomenon seems the most reasonable. The character of any chemi- 
 cal compound is influenced by all of the elements in that compound ; 
 no one element or group of elements is able entirely to suppress 
 any one of the others with which it is united. When a hydroxide 
 
 * It has been shown that a glass tube four inches in length, filled with 
 phosphorus pentoxide, will entirely dry a gas which is slowly passing through. 
 
 t The polymeric form of a substance is supposed to be produced by the 
 union of simpler molecules of that substance to form a more complicated 
 molecule. Thus, ordinary P 2 O 6 , let us suppose, is formed of molecules each 
 of which is composed of x times the formula weight P 2 O 5 , or x (P 2 O 5 ), each 
 molecule of the polymeric form would then contain a number of these simpler 
 molecules, or n (x[P 2 O 5 J). Such polymeric forms are quite frequently met 
 with in organic chemistry, and, possibly, the phenomenon of allotropism may 
 in many cases be caused by a union of simpler molecules to form more com- 
 plex ones. 
 
PHOSPHITES. 225 
 
 (for instance, that of potassium) reacts with phosphorous acid it 
 is to be presumed that the first product will be the primary salt 
 
 (pae-e 140) : 
 
 OH + KOH ( OK 
 
 OH = PJOH + H 2 0. 
 
 OH (OH 
 
 The metallic element which is present in the salt after this reac- 
 tion, renders the whole compound less negative, so that the next 
 reaction : /- K OK 
 
 P 
 1 
 
 ( OH (OH 
 
 would take place less readily than the first. The secondary salt has 
 now entirely lost all acid properties, owing to the increased mass of 
 metal present, so that all attempts to replace the third hydrogen 
 atom will fail.* Many chemists think that experimental evidence 
 has proven the formula of phosphorous acid to be : 
 
 v in 
 
 = _ / OH 
 
 H _P OH and not P OH 
 
 OH \ OH 
 
 so that it would contain only two hydroxyl groups. One of the 
 hydrogen atoms would then be joined to phosphorus, thus rendering 
 the not-metal quinquivalent ; by means of this hypothesis they have 
 sought to explain the fact that phosphorous acid will only form pri- 
 mary and secondary salts. The remarks on page 158 will apply 
 equally well in this case. 
 
 Phosphorous acid, like sulphurous and nitrous acid, is easily 
 oxidized, and when so acted on it forms phosphoric acid. Chlorine, 
 bromine, iodine, nitric acid, and even sulphurous acid can bring 
 about this oxidation. The following equations will serve as illus- 
 trations (see page 139) : 
 
 a. 2C1 + H 2 0=2HC1 + 0, 
 
 b. H 3 P0 3 + =H 3 P0 4 . 
 Combining a and b, we have : 
 
 c. H 3 P0 3 + 2 Cl + H 2 = H 3 P0 4 + 2 HCl.f 
 
 * A number of polybasic acids which are encountered in organic chemis- 
 try show this same character; they present the phenomenon of having one 
 atom of the metal in their salts more reactive than the others. 
 
 t The pupil should practise the writing of a large number of equations in 
 
226 pHOSPHonic ACIDS. 
 
 Phosphorous acid follows out the general rule which we ob- 
 served to be in force with the chlorine and sulphur acids ; namely, 
 it changes into the acid with greater amount of oxygen when it is 
 heated. The change can be represented by the following equa- 
 
 tion : - 4 H 3 P0 3 = 3 H 3 P0 4 + PH 3 . 
 
 The evolution of phosphine, if the phosphorous acid is heated too 
 rapidly, may sometimes become so violent that an explosion takes 
 place. Phosphorous acid is a colorless, crystalline solid which 
 melts at 70 to 74, and which is very soluble in water. 
 
 When phosphoric anhydride is exposed to the air, it deliquesces 
 and is converted into a phosphoric acid, which, however, is not the 
 one usually encountered, but is the less hydrated acid correspond- 
 ing to nitric or chloric acid ; the following will make the parallel- 
 ism clear : 
 
 P 2 5 + H 2 0=2HP0 3 , 
 
 Phosphoric anhydride -1- Water = Phosphoric acid. 
 
 N 2 5 +H 2 = 2HN0 3 , 
 
 Nitric anhydride + Water = Nitric acid. 
 
 C1 2 5 +H 2 0=2 HC10 3 , 
 
 Chloric anhydride + Water = Chloric acid. 
 
 When the phosphoric acid so obtained is dissolved in an excess of 
 water and allowed to stand for some time, or when it is boiled with 
 water, it takes up more of that liquid to produce the ordinary form 
 of the acid ; in effecting this change, one oxygen atom together with 
 the elements of water forms two hydroxyl groups : 
 
 (=o - ho 
 
 = + HOH -p _g_S 
 
 <_o H _o H 
 
 The acid with the formula HP0 3 is termed metaphosphoric, while 
 H 3 P0 4 has the name of orthophosphoric acid. Nitric acid, HK0 3 , 
 and chloric acid, H Cl 3 , are therefore really meta nitric and meta 
 
 which cases of oxidation and reduction occur. In all cases he must consider, 
 first, the substance to be oxidized; second, the amount of oxygen which it 
 will take up; third, the oxidizing agent; and, fourth, the amount of oxygen 
 which the latter will yield and the products which it forms when it oxidizes. 
 Examples of oxidation have been frequently given on the previous pages of 
 this book. 
 
PHOSPHORIC ACIDS : NOMENCLATURE. 
 
 22-7 
 
 chloric acids, but as the corresponding ortho acids (H 3 !N"0 4 and 
 H 3 Cl 4 ), are not known, there is no necessity of applying a special 
 designation to the less hydrated and common form of these sub- 
 stances. 
 
 The nomenclatwe which is applied to the acids derived from 
 the other elements of the nitrogen family corresponds to that used 
 with phosphorus, as the following table will demonstrate : 
 
 FORMULA. 
 
 NOMENCLATURE. 
 
 H As0 2 
 
 Meta-arsenious acid. 
 
 H 3 As0 3 
 
 Ortho-arsenious acid. 
 
 H As0 3 
 
 Meta-arsenic acid. 
 
 H 3 As0 4 
 
 Ortho-arsenic acid. 
 
 A similar system of nomenclature is also frequently applied for the 
 designation of acids derived from elements belonging to other 
 groups than that of nitrogen, for instance : 
 
 H 2 Si 3 is the formula of meta silicic acid, 
 H 4 Si0 4 " " " ortho " . 
 
 Unfortunately, this system is not rigidly carried out with all of the 
 various hydrated acids which have been discovered, and, further- 
 more, acids have been distinguished by the prefixes ortho and meta 
 when the difference between them has nothing whatever to do with 
 their hydration, so that these distinctive names can only be used 
 in cases where they have the sanction of custom. 
 
 Orthophosphoric acid changes into metaphosphoric acid at a red 
 heat, and if this temperature is maintained for a sufficient length 
 of time, the latter substance will finally evaporate. When ortho- 
 phosphoric is changed into metaphosphoric acid, a third acid is first 
 produced as an intermediary product. This acid, having its place 
 between the other two, is (owing to the fact that it is produced by 
 heating) termed pyrophosphoric acid. The changes which ortho- 
 phosphoric acid undergoes can be represented by the following : 
 
 2 H 3 P 4 = H 4 P 2 7 + H 2 0; and 
 
 Orthophosphoric 
 acid. 
 
 H, P 2 0, 
 
 Pyrophosphoric 
 acid. 
 
 Pyrophosphoric 
 acid. 
 
 = 2 H P0 3 + H 2 0. 
 
 Metaphosphoric 
 acid. 
 
228 METAPHOSPHORIC ACID. 
 
 The structural formulae, representing the formation of pyrophos- 
 phoric acid, are as follows : 
 
 P 
 
 
 OH HO 
 OH HO 
 
 P=P 
 
 
 
 OH HO 
 OH HO 
 
 
 
 Por,2H 3 P0 4 =H 4 P 2 7 +H 2 0. 
 
 I 1 OHH 10 
 
 n //0 
 
 It is supposed that two univalent groups P OH are united by 
 
 \OH 
 
 means of oxygen, just as are two similar ones in disulphuric acid 
 (described on page 154) ; such complicated acids, produced by the 
 separation of water and the joining of more complicated groups 
 of elements by means of oxygen atoms, are of quite frequent 
 occurrence.* 
 
 Metaphosphoric acid is a colorless, glass-like substance (acidum 
 phosphoricum glaciale) which greedily absorbs moisture from the 
 air. It is a monobasic acid, and, according to the rules which have 
 been dwelt upon in the former portions of this book, should only 
 form one class of salts. This is not the case, however, as a number 
 of metaphosphates of the same metal are known ; the existence of 
 these various metaphosphates is explained by the theory that meta- 
 phosphoric acid itself can exist in several polymeric forms. (See 
 page 224.) f Metaphosphoric acid precipitates egg albumen from 
 its solutions, a property which distinguishes it from ortho and pyro- 
 
 * From its formation by heating pyrophosphoric acid, it seems likely that 
 ordinary metaphosphoric acid has the formula : 
 
 r=o o=n 
 
 r\ 
 
 p -Q_ HO H-O- ip 
 
 I o j 
 
 or, H 2 P 2 O 6 = 2 HPO 3 ; for any other formula could be produced only if we 
 suppose the pyrophosphoric acid to be split asunder by the very means which 
 generally unites groups of atoms, namely, by the separation of water between 
 two hydroxyl groups. Of course, this is merely speculation, but the existence 
 of more than one sodium metaphosphate seems to warrant the belief that meta- 
 phosphoric acid can appear in polymeric forms. 
 
 t The metaphosphoric acids are supposed to have the formulae : 
 
 HP0 3 
 
 2(HP0 3 )=H 2 P 2 6 , 
 3(HP0 3 )=H 3 P 3 9 , 
 or, in general, to be n (HPO 3 ). 
 
OKTHOPHOSPHATES. 229 
 
 phosphoric acid, and it changes to orthophosphoric acid when it is 
 dissolved in water and allowed to stand ; it is rapidly converted by 
 boiling the solution. 
 
 Orthophosphoric acid is tribasic ; it forms primary, secondary, 
 and tertiary salts. Representatives of each of these classes are 
 known, and all are of importance. If M' represents any univalent 
 metal, then : 
 
 M' H 2 P0 4 is the primary phosphate, 
 
 M' 2 HP0 4 is the secondary phosphate, and 
 
 M' 3 P0 4 is the tertiary phosphate ; 
 and if M" is any bivalent metal, then : 
 
 M" ( H 2 P0 4 ) 2 is the primary phosphate, 
 
 M' v ( HP0 4 ) is the secondary phosphate, and 
 
 M" 3 ( P0 4 ) 2 is the tertiary phosphate. 
 
 The formulae of the last three salts are self-evident if we consider 
 that a bivalent metal replaces two hydrogen atoms of an acid, while 
 a primary salt has one, a secondary two, and a tertiary three such, 
 atoms replaced by that metal. 
 
 When the secondary phosphates are heated to a sufficiently high 
 temperature, the pyrophosphate very frequently results, thus : 
 
 2 Na2 HP0 4 = Na 4 P 2 7 + H 2 0, 
 
 and when the primary ones are similarly treated the metaphosphate 
 is produced : - Na HS P Q 4 = Na P o 3 + H 2 0. 
 
 All secondary and tertiary phosphates, excepting those of the alkali 
 metals and of ammonium, are insoluble in water, while, on the other 
 hand, all primary phosphates are soluble. The tertiary and second- 
 ary phosphates, therefore, are dissolved by acids, because the latter 
 produce the primary phosphates, thus : 
 
 Ca 3 ( P0 4 ) + 4 H Cl = Ca ( H 2 P0 4 ) 2 + 2 Ca C1 2 ,* 
 
 ( Insoluble tertiarv ) , -, ( soluble primary ) . -,, 
 4 -i -I i ", > 4- acid = < , . r T r > -4- salt. 
 ( calcium phosphate ) ' ( calcium phosphate ) ' 
 
 By the addition of soluble bases to the soluble primary phosphates, 
 
 the insoluble tertiary ones are precipated : 
 
 3 Ca (H 2 P0 4 ) 2 + 12 NH 4 OH =Ca 3 (P0 4 ) 2 + 4 (NH 4 ) 3 P0 4 + 12 H 2 0. f 
 
 * See pages 140 and 153. 
 
 t The formulation of a reaction such as this is really not as formidable as 
 
230 OKTHOPHOSPHATES. 
 
 The soluble secondary phosphates of the alkali metals are the 
 salts of those elements which are usually encountered ; the insoluble 
 secondary phosphates of pronounced metals like calcium, barium, 
 or strontium, which form such salts, are readily produced from 
 these by precipitation : 
 
 Na 2 HP0 4 + Ca C1 2 = Ca HP0 4 + 2 Na Cl,* 
 
 Soluble. Soluble. Insoluble. Soluble. 
 
 The soluble tertiary phosphates of the alkali metals have a 
 strongly alkaline reaction,f and show the greatest tendency to 
 separate a portion of the metal contained in them as the metallic 
 hydroxide : 
 
 Na 3 P0 4 + H 2 = Na 2 HPO 4 + Na OH. 
 
 This characteristic bears out the theory which we developed when 
 discussing phosphorous acid (pages 224, 225). Phosphoric acid, 
 being a stronger acid than phosphorous acid, will have a greater 
 tendency to form sodium salts, so that, while the tertiary sodium 
 phosphate, although unstable, really does exist, the tertiary phos- 
 phite is not known. Only primary and secondary salts of very pro- 
 nounced metals, such as sodium, potassium, and calcium, are known ; 
 
 at first glance it would appear to be. The process consists in simply neutral- 
 izing the excess of acid hydrogen atoms in the primary salt by means of 
 ammonia. As the tertiary phosphate of calcium is always formed, we must 
 necessarily, in endeavoring to picture by atomic symbols the changes which 
 really take place, use three times the formula weight of calcium primary phos- 
 phate, because there are three atoms of calcium in Ca 3 (PO 4 ) 2 . The remainder 
 of the work simply consists in counting up how much ammonia it will take 
 to replace all of the remaining hydrogen atoms by the group NH 4 . A study 
 of this reaction in the laboratory reveals that exactly the proportion by 
 weight of Ca 3 (PO 4 ) 2 and (NH 4 ) 3 PO 4 , expressed in this equation are, in 
 reality, produced by nature. 
 
 * In reality a mixture of secondary and tertiary phosphate of calcium is 
 at first precipitated : 
 
 3 Na 2 HPO 4 + 4 Ca C1 2 = Ca HPO 4 + Ca 3 (PO 4 ) 2 + 2 HC1 + 6 Na Cl, 
 and in consequence of the free hydrochloric acid, the fluid above the precipi- 
 tate assumes an acid reaction. The hydrochloric acid, however, gradually 
 reacts with the tertiary phosphate to produce the primary: 
 
 Ca 3 (P0 4 ) 2 + 2 HC1 = 2 Ca HPO 4 + Ca Cl.,; 
 
 so that, in the end, the reaction given above is realized. The same is true of 
 barium phosphate. 
 
 t i.e., they turn red litmus of a blue color (page 75). 
 
METAPHOSPHATES ; PYKOPHOSPHATES. 231 
 
 in other cases the tertiary salt alone exists; therefore when, for 
 example, silver nitrate is added to the secondary phosphate of 
 sodium, the insoluble tertiary phosphate of silver is precipitated : 
 
 Na 2 HP0 4 + 3 Ag N0 3 = Ag 3 P0 4 + 2 Na N0 3 + HN0 3 . 
 
 This and similar reactions puzzled the chemists of former days not 
 a little ; for, by mixing a solution of two salts which were neutral, 
 they obtained a liquid of acid reaction (because of the free acid 
 formed). 
 
 All metaphosphates, provided they are not those of a volatile 
 metal-like substance, such as ammonium, are unchanged even by 
 quite a high heat ; it follows that these salts will be formed under 
 conditions which render the existence of the salts of the great 
 majority of other acids impossible; as a consequence, phosphoric 
 acid will, if the temperature is sufficiently increased, ultimately 
 decompose the salts of much stronger acids (such as sulphuric), 
 leaving a phosphate in their place ; while, on the other hand, in 
 solution, or in the cold, the exact reverse takes place, i.e., the other 
 acids decompose the phosphates. 
 
 Pyrophosphoric acid is quadribasic, but forms only two classes 
 of salts ; namely, those with M 4 P 2 7 and those with M 2 H 2 P 2 7 as 
 their general formulae. On being heated with water the pyrophos- 
 phates change into the secondary orthophosphates. 
 
 Phosphoric acid is necessary for animal and vegetable life ; 
 tertiary calcium phosphate forms the major portion of the inorganic 
 constituents of the bones and teeth ; but phosphoric acid, combined 
 in some form, is also found in the blood, in the muscle and nerve 
 tisues, and in the brain. The phosphates which are found in the 
 soil are generally of an insoluble variety ; however, the chemical 
 action of the various substances which are present, aided by water, 
 renders them finally partially soluble and absorbable by plants. 
 The primary calcium phosphate has, when mixed with other ingre- 
 dients, an extensive sale as superphosphate ; it is used as a fertilizer. 
 
 Two other oxides of phosphorus, P0 2 and P 4 0, are known; 
 they are of little importance except that P0 2 is analogous to N0 2 . 
 A larger work must be consulted for their description. 
 
 One more acid of phosphorus, hypophosphorous acid, H 3 P0 2 , 
 remains to be considered. Salts of this acid are produced by the 
 action of phosphine on a solution of an alkaline hydroxide. The 
 
232 HYPOPHOSPHOROUS ACID. 
 
 free acid is formed by acidifying the barium salt with the calcu- 
 lated quantity of diluted sulphuric acid. It is a crystalline sub- 
 stance which melts at 17. 4, and which oxidizes to phosphorous acid 
 when in contact with the air. When heated it changes to phosphine 
 and phosphoric acid. 
 
 Hypophosphorous acid, like phosphorous acid, is a hydrated acid, 
 which is derived from the hypothetical meta acid by the addition 
 of one molecule of water to each formula weight : 
 
 Only one atom of hydrogen in a molecule of hypophosphorous 
 acid is capable of being replaced by metals to form a salt ; this 
 fact is exactly what would be expected from the behavior of 
 phosphorous acid. (See page 224.) 
 
 The following table shows the connection between the compounds 
 discussed in the last two chapters : 
 
OXY-ACIDS OF PHOSPHOKUS ; TABLE OF. 
 
 233 
 
 ^ 
 o 
 >d 
 
 I 
 
 s 
 
 I 
 
 ZJ 
 
 
 1.1 
 
 c-o 
 ' 
 
 I 
 o 
 
 WWW 
 ooo 
 
 WWW 
 
 oooo 
 WWW 
 
 If If! 
 
 B||I| 
 
 n q o 
 'SE-S'S ^ 
 
 95 o <=: sr o 
 
 8 !*.& 
 
 3g S? * 
 If^lJ: 
 
 gjaSs- a 
 
 r+ *^ rf" ^^ 
 
 O 
 
 *' 3 
 
 * B * 
 
 C) - 02 
 
 I 
 
 I 
 
 i 
 
 ^f 
 
 2 
 
 W + 
 
 g 
 
 B- 
 
234 . ARSENIC ; OCCURRENCE. 
 
 CHAPTER XXX. 
 
 ARSENIC AND ARSINE. 
 
 Arsenic ; symbol, As ; atomic iv eight, 75 ; specific gravity, 5.7. Spe- 
 cific gravity of vapor at red heat, air = 1, is 10.3, H 2 = 2, is 
 '296.6 ; molecular weight of As 4 = 300, of As 2 = 150. Arsine, 
 As H 3 ; specific gravity, air = 1, is 2.7, H 2 = 2, is 78.02 ; 
 'molecular weight, 78.021 ; 1 c.c. of the gas at and .76 m. 
 pressure weighs .003499 gram. 
 
 ARSENIC occurs quite frequently in the form of the uncombined 
 element, especially in formations which contain the metallic sul- 
 phide ; the native arsenic is found in grayish black, reniform masses. 
 Combined arsenic occurs in the arsenides of many metals, the chief 
 examples of which are : 
 
 Arsenopyrite, FeAsS (corresponding to FeS 2 , iron pyrites, one atom of 
 sulphur being replaced by one of arsenic). 
 
 Smaltite, (Co, Fe Ni) As 2 (corresponding to Fe S 2 , iron pyrites, both atoms 
 of sulphur being replaced by arsenic). 
 
 Cobaltglance, Co As S. 
 
 Two sulphides of arsenic are also not infrequently found as 
 
 minerals ; they are : 
 
 As 2 S 2 (realgar). 
 As 2 S 3 (orpiment). 
 
 In addition to these occurrences, arsenic also appears as the 
 oxide As 2 3 , and, in some arsenates, these latter probably the 
 result of the oxidation of the arsenides. The arsenic of commerce 
 is either the naturally occurring element, or it is prepared from the 
 arsenides which are found as minerals. 
 
 The two sulphides referred to above and the oxide As 2 3 have 
 been known since ancient times ; Aristotle mentions the former, but 
 the term arsenicon seems first to have been used by Dioscorides. 
 The element arsenic does not seem to have been isolated until the 
 latter part of the seventeenth century; of course, its chemical 
 nature was not understood until much later. 
 
ARSENIC ; PREPARATION, PROPERTIES. 235 
 
 The commercial preparation of arsenic depends upon the forma- 
 tion of ferrous sulphide and free arsenic, when arsenopyrite is 
 heated : - Fe As S = Fe S + As. 
 
 The impure arsenic so prepared is purified by sublimation ; for when 
 heated, it volatilizes without previously melting, and then collects 
 in crystals on the colder parts of the retorts. 
 
 Arsenic has a steel-gray color and metallic lustre; when frac- 
 tured, it resembles white pig iron. Arsenic, like sulphur and phos- 
 phorus, exists in two allotropic forms, one of which is crystalline,* 
 and the other amorphous. When the element is heated in a glass 
 tube, amorphous arsenic is deposited on the walls near the flame as 
 a black mirror, while the bright, shiny crystals of the other variety 
 appear upon the cooler portions. The amorphous form is changed 
 into the crystalline one by heating to 360. Crystallized arsenic 
 has a specific gravity of 5.76, amorphous of 4.71. 
 
 When heated in the air, arsenic burns to form the trioxide 
 As 2 3 ; in this particular the element differs from phosphorus, which, 
 under similar circumstances, forms the pentoxide P 2 5 . When 
 heated or burned, arsenic assumes a peculiar odor, somewhat resem- 
 bling that of phosphine. The element unites directly with chlorine, 
 bromine, or iodine to form the corresponding halogen compounds, 
 and arsenic also readily combines with a number of metals, thus 
 yielding arsenides. 
 
 Arsenic volatilizes at 450 ; the vapors have a lemon yellow 
 color and a most disagreeable garlic odor. Their specific gravity, 
 air = 1, at a low red heat, is 10.3, which, H 2 = 2, gives 296.6 the 
 molecular weight at this temperature is therefore 300 and the mole- 
 cule consists of four atoms (As 4 ). The density of the arsenic 
 vapors, however, gradually diminishes as the heat is increased, 
 until it reaches 5.37 at 1736. f The molecules, As 4 , which corre- 
 spond to those of phosphorus, P 4 , therefore begin to dissociate as a 
 white heat is reached, so that they very nearly change into As 2 ; $ 
 the value for the specific gravity would be 5.22, air = 1, were 
 complete dissociation into As 2 to result. 
 
 * Hexagonal system, rhombohedra. 
 
 t Latest determinations by H. Biltz and Y. Meyer. (Ber. d. Deutsch. 
 Chem. Gesell.; 22, 726). 
 
 t Possibly they partially dissociate into the individual atoms; there is no 
 means of determining this point. 
 
236 AESINE; PREPARATION. 
 
 Arsenic is an element which is on the boundary line between 
 metal and not-metal. Crystallized arsenic is entirely metallic in 
 appearance ; although it is neither malleable nor ductile, it conducts 
 heat and electricity quite readily. Chemically, arsenic is almost 
 entirely a not-metal, but its approach to a metallic character is 
 evinced by the instability of its hydrogen compound. 
 
 Arsine, AsH 3 , is the analogon of phosphine, PH 3 , and of am- 
 monia, NH 3 . It is produced, similarly to phosphine, by the action 
 of an acid upon some arsenide, thus : 
 
 As 2 Zn 3 + 6H Cl = 2 AsH 3 +3 ZnCl 2 . 
 
 Zinc arsenide. Arsine. 
 
 Such a reaction is exactly like the ones which were studied when 
 the preparation of hydrogen sulphide and other hydrogen compounds 
 was discussed, thus : 
 
 Fe S + 2 H Cl = Fe C1 2 + H 2 S. 
 
 Zn 3 As 2 + 6 H Cl = 3 Zn C1 2 + 2 As H 3 . 
 
 This method furnishes pure arsine. Another and more important 
 way of preparing arsine is by the action of nascent hydrogen upon 
 an acid solution of a soluble arsenic compound. For instance, a 
 solution of arsenic trioxide in hydrochloric acid, when added to zinc 
 which is covered with dilute sulphuric acid and which is therefore 
 generating hydrogen, will develop arsine, the latter, however, 
 naturally is mixed with hydrogen : 
 
 As,0 3 +12H = 2AsH 3 + 3 H 2 0. 
 
 The most delicate test for arsenic (Marsh's test) is based upon this 
 chemical fact. 63 
 
 Arsine is a colorless gas, with a most disagreeable odor. It is 
 an intense poison, so that it is imperatively necessary to take every 
 precaution when experimenting with the gas, especially when the 
 latter is obtained pure, as it is from zinc arsenide.* The specific 
 gravity of arsine is 2.7, air == 1, or 78.02, H 2 = 2; the molecular 
 weight is, therefore, 78.024 (As = 75, 3 H = 3.024 ; As H 3 = 78.024) 
 and the formula AsH 3 . Arsine changes to a liquid at 102, 
 and forms a white crystalline mass at 119. When ignited in 
 the air the gas burns with a bluish white flame, forming water and 
 arsenic trioxide, the latter compound appearing in the form of a 
 
 * Arsine should never be generated otherwise than under a hood or in the 
 open air. 
 
ARSINE; PROPERTIES. 237 
 
 white smoke. If the air supply in which arsine is burning is 
 limited, or if the flame is cooled, for example, by holding a porce- 
 lain plate in it, the arsenic formed by the decomposition of the gas 
 into its elements will separate and will form a dark spot on the 
 surface in contact with it. A mixture of arsine and oxygen explodes 
 violently when ignited. 
 
 Arsine is much less stable than phosphine. When it is passed 
 through a heated glass tube it readily decomposes into hydrogen 
 and arsenic, the latter being deposited as a black mirror. Of 
 course, arsine is a powerful reducing agent ; sulphuric acid is readily 
 decomposed by it, just as that acid is by hydroiodic acid or hydro- 
 gen sulphide, but in this case the acid is robbed of all of its oxygen 
 while the sulphide of arsenic is formed ; * other acids and even water 
 or an alkaline solution can likewise decompose hydrogen arsenide. 
 Arsine reduces silver nitrate in solution ; . metallic silver is precipi- 
 tated and arsenic trioxide is formed ; this fact is of importance in 
 the detection of arsenic. The basic properties manifested by ammo- 
 nia and phosphine are absent in arsine; the latter forms no 
 arsonium compounds ; it can be made to act as a base only when 
 the hydrogen atoms contained in it are substituted by some more 
 positive (so-called organic) radicle, like methyl, f 
 
 * This would appear to be an example of the nascent action of a solid 
 element; arsenic, at the moment of its liberation from arsine, readily com- 
 bines with the hydrogen sulphide, or, perhaps, with the sulphur formed by 
 the reduction of the sulphuric acid. 
 
 t See methane, chapter on compounds of carbon and hydrogen. 
 
238 ARSENIC ; HALIDES OF. 
 
 CHAPTER XXXI. 
 
 THE COMPOUNDS OF ARSENIC WITH THE HALOGENS, WITH 
 OXYGEN, AND WITH OXYGEN AND HYDROGEN. 
 
 THE halogen compounds of arsenic are not complicated by the 
 existence of two series, for only the trihalogen derivatives have 
 with certainty been prepared ; their chief characteristics are as 
 follows : 
 
 As F 3 , arsenic trifluoride, liquid, boils at 63.* 
 AsCl 3 , arsenic trichloride, liquid, boils at 130.2, solid at -18. 
 'AsBr 3 , arsenic tribromide, solid, melts at 20-25, boils at 220. 
 As I 3 , arsenic tri-iodide, solid, sublimes when heated. 
 
 Arsenic shows its resemblance to the metals and at the same 
 time its connection with the not-metals nowhere better than in the 
 chemical behavior of its chlorine compound. The latter substance 
 can be formed, as are the chlorides of metals, by the action of hydro- 
 chloric acid on the trioxide of arsenic, so that in this case arsenic 
 trioxide is a base and arsenic a metal : - 
 
 As 2 3 + 6HCl = 2 AsCl 3 + 3H 2 0. 
 
 This formation of the chloride of arsenic is not, however, 
 possible unless very little water is present ; it takes place, for in- 
 stance, when dry hydrochloric acid is passed over the oxide of 
 arsenic, or when the latter substance is distilled with a mixture of 
 sulphuric and hydrochloric acid.t On the other hand, when an 
 excess of water is added to the trichloride, it is entirely decomposed 
 into hydrochloric acid and arsenic trioxide : 
 
 * A pentafluoride of arsenic (AsF 5 ) has been isolated by Moissan. 
 This compound was prepared by electrolyzing the trifluoride, when arsenic 
 and the pentafluoride are formed. The compound is interesting because it is 
 the only pentahalide of arsenic, and therefore brings the halides of arsenic in 
 line with those of phosphorus. 
 
 t The trichloride of arsenic is also formed when the trioxide is warmed 
 with a concentrated solution of hydrochloric acid ; it then separates as an oil, 
 insoluble in the excess of the acid. 
 
ARSENIC TRIOXIDE. 239 
 
 a. 2AsCl 3 + 6H 2 0=2As(OH) 3 + 
 
 b. 2As(OH) 3 =As 2 3 -j-3H 2 0.* 
 
 An arsenious acid which corresponds to phosphorous acid does 
 not exist; where its formation is to be expected, not it, but its 
 anhydride, As 2 3 , results. The easy decomposition of the trichlo- 
 ride of arsenic by an excess of water, therefore, is a phenomenon 
 classing arsenic as a not-metal, for the chlorides of the pronounced 
 metals are unchanged by the action of water. f The addition of a 
 little water only partially decomposes arsenic trichloride : 
 
 rei + HOH (OH 
 
 As ] Cl + HOH = As ] OH + 2 HC1. 
 
 ( ci I ci 
 
 The compound so produced is a so-called basic salt, i.e., a salt which 
 is in part hydroxide and in part chloride (see antimony trichlo- 
 ride). The chloride, bromide, or iodide of arsenic can readily be 
 formed by the direct union of the elements. These compounds are 
 all extremely poisonous substances ; the fluoride, chloride, and 
 bromide fume in the air, absorb moisture, and decompose, leaving 
 the trioxide. 
 
 Arsenic forms two oxides, As 2 3 ,$ and As 2 O 5 ; they are the 
 anhydrides respectively of arsenious and of arsenic acid ; arsenious 
 acid is, however, known only in its salts, for we have seen that when 
 it is liberated from these it at once breaks down into its anhydride 
 and water. Owing to the few reactions in which arsenic acts like 
 a metal, the oxides are sometimes named arsenious and arsenic 
 oxides in conformity with the nomenclature usually adopted where 
 a metal forms two such compounds (see page 26). 
 
 Arsenious oxide (arsenic trioxide, As 2 3 ) is the most common 
 preparation of arsenic, having been known to the ancients, and hav- 
 ing been a familiar substance ever since the time of the Romans ; 
 it is popularly known by the name of arsenic or white arsenic ; 
 when cases of poisoning by arsenic occur, the substance used is 
 
 * An intermediary product, As(OH) 2 Cl, is formed, in all probability 
 before complete decomposition is accomplished. This compound is a basic 
 chloride of arsenic (see page 181). t See page 79. 
 
 t The specific gravity of this substance, taken by Victor Meyer, above 1,500, 
 corresponds to the molecule of the formula As 4 O 6 . The smallest molecule of 
 the trioxide of arsenic with which we are acquainted is therefore As 4 O 6 . 
 
240 ARSENIC TKIOXIDE; ACTION AS A POISON. 
 
 generally " white arsenic." As its name implies, it is a white solid, 
 resembling ordinary flour. The commercial product is formed by 
 roasting the arsenical sulphides which occur as mineral deposits ; 
 arsenico-pyrite is especially advantageous for this purpose ; in 
 roasting this, care must be taken to have a sufficient supply of air, 
 otherwise arsenic (formed as shown on page 235) sublimes together 
 with the oxide. Quantities of white arsenic are also produced 
 while burning cobalt, nickel, tin, and silver ores ; the arsenious 
 oxide is collected in cold chambers and is purified by sublimation. 
 The white arsenic of commerce generally contains traces of the 
 corresponding oxide of antimony. 
 
 Sublimed arsenic trioxide appears in two forms, in the first as a 
 dimorphous substance, crystallizing both in the regular and in the 
 monoclinic system, and in the second as an amorphous, glass-like 
 body, which gradually changes into a porcelain-like mass. The 
 oxide volatilizes at '200.* 
 
 Arsenious oxide is a poison to all animals, and even to plants. 
 Owing to its t^stelessness and the readiness with which it can be 
 mixed with foods, it is frequently used as a poison, both intention- 
 ally and accidentally. It acts in two ways. 
 
 a. As a corrosive substance, attacking the organic surfaces with 
 which it comes in contact, it therefore causes local inflammation in 
 the stomach and intestinal tract. 
 
 b. It has a destructive effect on the medullas of the nerves. 
 The more rapidly the poison is absorbed the less the action under a 
 and the more the one under b is observed ; it follows that the latter 
 is more prominent when the arsenic is administered in solution, 
 and, of course, it is of material influence whether the stomach is full 
 or empty at the time of taking. .005 gram has a marked effect, 
 and if continued can cause death; .06 to .12 gram may cause death 
 in a few days, and .2 to .3 gram in a few hours. Symptoms: 
 nausea, salivation, burning in the gastric region, vomit yellow or 
 greenish, and possibly streaked with blood, while traces of white 
 arsenic may be visible therein, great thirst, colic, and sensitiveness 
 of the abdomen. The symptoms of the action under b are : intense 
 fear, convulsive movements, trembling and cramps in the extremi- 
 ties, fainting spells, and delirium ; where large quantities have been 
 
 * According to Selmi, it already begins to vaporize at 100. 
 
ARSENITES; ARSENIC PENTOXIDE. 241 
 
 rapidly absorbed the patient may be entirely unconscious. On post 
 mortem examination, fatty degeneration of the liver and heart, as 
 well as in many other organs, the effects of the poisoning being much 
 like those resulting from the taking of phosphorus. An antidote is 
 a mixture of freshly precipitated ferric hydroxide and magnesium 
 oxide, the endeavor being to form the insoluble ferric arsenite. 
 Glass-like arsenic trioxide is much more soluble in water than is the 
 crystallized variety, 100 parts of water at ordinary temperatures 
 dissolve four parts of the former and 1.2 to 1.3 parts of the latter. 
 
 The chemical action of arsenic trioxide is twofold, for it can 
 appear both as a base and as the anhydride of an acid. 
 
 a. It acts like a base, because it can dissolve in a number of 
 acids. One or two of the salts produced, for instance, the chloride, 
 AsCl 3 , and a sulphate (As 2 (S0 4 ) 3 -f- S0 3 ), have been isolated; 
 these substances are decomposed by water. 
 
 ft. It acts like the anhydride of an acid, because it dissolves in 
 bases to form arsenites : 
 
 As, 3 + 2 KOH = 2 As 2 K + H 2 0, 
 
 The meta-arsenites ( like potassium meta-arsenite, As 2 K ) are the 
 most common salts of this acid; they correspond to the nitrites 
 N0 2 M ; a few orthoarsenites, As 3 M 3 , are also known. The 
 arsenites of the alkali metals are soluble in water ; the others are 
 either soluble with difficulty or entirely insoluble ; they are decom- 
 posed by hydrochloric acid. The alkaline solutions of arsenic 
 trioxide are most powerful reducing agents ; they have the greatest 
 tendency to take up oxygen to form arsenates, but, on the other 
 hand, arsenious oxide is quite readily reduced, and a method for 
 detecting arsenic is based upon the ease with which this substance 
 gives up oxygen. 64 
 
 Arsenic pentoxide, As 2 R , is the anhydride of arsenic acid; it 
 corresponds to N 2 5 and P 2 5 ; three acids, analogous to those of 
 phosphorus, are derived from this oxide ; they are : 
 
 H As 3 , meta-arsenic acid, 
 H 3 As 4 , orthoarsenic acid, 
 H 4 As 2 7 , pyroarsenic acid. 
 
 Arsenic pentoxide is produced when arsenic acid is heated to a 
 low red heat ; a higher temperature produces decomposition into 
 
242 ARSENIC ACIDS. 
 
 arsenic trioxide and oxygen. It is a colorless, amorphous mass, 
 which greedily absorbs moisture and finally deliquesces. 
 
 Orthoarsenic acid, H 3 As 4 , the only arsenic acid which exists 
 in aqueous solutions, is produced by oxidizing arsenic trioxide (for 
 instance, by chlorine, bromine, or nitric acid*). When the solu- 
 tion produced by oxidizing arsenic trioxide is evaporated to dryness, 
 orthoarsenic acid separates in needle-like crystals ; when these are 
 heated to 180 they separate water, and change into pyroarsenic 
 
 which latter, at 206, changes into meta-arsenic acid : 
 H 4 As 2 7 = 2HAs0 3 + H 2 0, 
 
 and this substance finally, at red heat, decomposes to yield the an- 
 hydride, As 2 5 . Each one of these compounds is converted into 
 orthoarsenic acid by solution in water. 
 
 Orthoarsenic acid is tribasic, and forms primary, secondary, and 
 tertiary salts ; the formulae of these are, of course, exactly parallel 
 to those of the corresponding derivatives of phosphoric acid (see 
 page 229). Primary and secondary arsenates, when they are heated 
 to redness, undergo the same changes which take place with primary 
 and secondary phosphates : 
 
 MH 2 As 4 = M As 3 + H 2 0. 
 
 Primary arsenate. Meta-arsenate. 
 
 M 2 H As 4 = M 4 As 2 T + H 2 0. 
 
 Secondary arsenate. Pyroarsenate. 
 
 The arsenates of the alkalies are soluble in water, but only 
 the primary arsenates of the other metals dissolve ; the latter 
 are, however, dissolved by mineral acids (see page 229) ; it will be 
 remembered that this is also the case with the phosphates. 
 
 Arsenic acid is a tolerably good oxidizer ; quite a number of 
 reducing agents reduce it to arsenic trioxide. It has of late years 
 been extensively used in the manufacture of aniline dyes, because, 
 while it certainly gives up its oxygen, it does not do so with such 
 facility as to destroy the organic substance it acts upon. Sulphur 
 dioxide reduces arsenic acid as follows : 
 
 * The pupil should write these equations, using the knowledge acquired 
 in studying the previous oxidizing action of these substances. 
 
AKSENIC ACIDS. 243 
 
 H 3 As 4 + H 2 + S0 2 = H 3 As 3 + H 2 S0 4 , 
 
 the arsenious acid so formed, of course, breaking down into arsenic 
 trioxide and water. This reaction serves to distinguish arsenic acid 
 from phosphoric acid, for the latter substance, although in every 
 respect like the former, has no tendency whatever to give up oxygen 
 and change to phosphorous acid. Arsenic acid is a powerful poison. 
 
244 ARSENIC; TEISULPHIDE. 
 
 CHAPTER XXXII. 
 
 THE COMPOUNDS OF ARSENIC WITH SULPHUR, AND WITH 
 SULPHUR AND HYDROGEN. 
 
 ARSENIC forms three sulphides, two of which correspond in 
 formula to the two oxides ; they are As 2 S 3 and As 2 S 5 , while the 
 third is a ruby red mineral substance (realgar), As 2 S 2 . 
 
 Arsenic trisulphide and arsenic pentasulphide are especially 
 interesting, because they act much like anhydrides of oxy-acids, and 
 therefore illustrate the marked chemical resemblance between sul- 
 phur and oxygen. 
 
 Arsenic trisulphide is found as a natural, yellow colored mineral 
 which bears the name of orpiment. It is produced, as are the sul- 
 phides of many metals, by the action of hydrogen sulphide upon 
 an acidified solution of arsenic trioxide, for the sulphide of arsenic 
 is one of the sulphides which are insoluble in dilute acids. (See 
 page 100.) 2 As C1 3 + 3 H 2 S = A 2 S 3 + 6 H Cl. 
 
 When precipitated, it is a bright yellow powder. Arsenic trisul- 
 phide can also be formed by direct union of the elements, just as is 
 the trioxide. Arsenic trioxide dissolves in alkaline solutions to 
 form a meta-arsenite : 
 
 As 2 3 + 2 KOH = 2 As 2 K + H 2 0, 
 
 and the trisulphide is dissolved by hydrosulphides in exactly the 
 same way:- As 2 S 3 -f 2KSH = 2 As S 2 K + H 2 S ; 
 in this case potassium metasulpharsenite is formed. Metasulphar- 
 senious acid may be considered as derived from an orthosulphar- 
 senious acid by the separation of hydrogen sulphide, just as 
 meta-arsenious acid might be formed from orthoarsenious acid. 
 The following equations will make this clear : 
 
 r SH (S 
 
 As < SH = As < OTT r H 2 S, 
 
 (SH 
 Orthoaulpharsenious acid = Metasulpharsenious acid + Hydrogen sulphide ; 
 
ARSENIC ; PENTASULPHIDE. 245 
 
 H 
 
 Orthoarsenious acid = Meta-arsenious acid + Water. 
 
 We saw, however, that meta-arsenious acid is incapable of exist- 
 ence ; for, when liberated from its salts, it at once breaks down as 
 follows : - 2 As 2 H = As 2 3 + H 2 : 
 
 metasulpharsenious acid acts in exactly the same way : 
 
 2 AsS 2 H = As 2 S 3 + H 2 S; 
 
 and, therefore, when acids are added to solutions of the sulpharsen- 
 ites, arsenic trisulphide is precipitated : 
 
 2 AsS 2 K + 2 HC1 = As 2 S 3 + H 2 S+2 KC1. 
 
 Arsenic trisulphide can readily be dissolved in the hydroxides of 
 the alkalies, for then a mixture of the arsenite and sulpharsenite is 
 produced ; * it is also dissolved by ammonium sulphhydrate, but in 
 the latter event a salt of pyrosulpharsenious acid is the result ; the 
 free acid corresponding to this ammonium salt does no't exist, but it 
 can be supposed to be formed by the separation of one molecule of 
 hydrogen sulphide from two formula weights of an hypothetical 
 orthosulpharsenious acid : 
 
 SH HS SH HS 
 
 (SH HS) (_._S--) 
 
 2 As (SH ) 3 = As 2 S 5 H 4 + H 2 S.f ( See page 228.) 
 
 Some salts derived from orthosulpharsenious acid, As ( SH ) 3 , are 
 also known ; none of the acids exist in the free state ; they at once 
 break down into the trisulphide of arsenic and hydrogen sulphide 
 when liberated from their salts. 
 
 Arsenic pentasulphide can be produced from the trisulphide by 
 fusing the latter with sulphur : - 
 
 As 2 S 3 -f 2 S = As 2 S 5 ; 
 
 this sulphurization corresponding to the oxidation of arsenic tri- 
 oxide: - As 2 3 + 2 0=As 2 5 . 
 
 Arsenic pentasulphide cannot be precipitated from cold solutions of 
 pure arsenic acid by means of hydrogen sulphide, because the ar- 
 
 * 2 A&J S 3 + 4 KOH = 3 As S 2 K + As O. 2 K + 2 H 2 O. 
 t As 2 S 3 + 4 NH 4 SH = As 2 S 5 ( NH 4 ) 4 + 2 H 2 S. 
 
246 SULPHAK SENATES. 
 
 senic acid is first reduced to arsenious acid by means of sulphuretted 
 hydrogen, which latter substance is oxidized to sulphur ; as a conse- 
 quence, a mixture of arsenic trisulphide and sulphur is produced ; * 
 the pentasulphide is, however, produced by the action of hydrogen 
 sulphide on a boiling solution of arsenic acid. 
 
 Arsenic pentasulphide dissolves in the sulphides, sulphhydrates, 
 or hydroxides of the alkalies to form sulpharsenates ; it therefore 
 acts as the anhydride of sulpharsenic acid : 
 
 As 2 S 5 + 2 K SH =2 As S 3 K + H 2 S. 
 
 The sulpharsenates are derived from three sulpharsenic acids, 
 orthosulpharsenic acid, As S 4 H 3 , corresponding to orthoarsenic 
 acid, As 4 H 3 ; metasulpharsenic acid, As S 3 H, corresponding to 
 meta-arsenic acid, As 3 H ; and pyrosulpharsenic acid, As 2 S 7 H 4 , 
 corresponding to pyroarsenic acid, As 2 7 H 4 . When an acid is 
 added to a solution of a sulpharsenate, orthosulpharsenic acid is 
 precipitated : 
 
 As S 4 K 3 + 3 H Cl = As S 4 H 3 + 3 K Cl. 
 
 This orthosulpharsenic acid, when boiled, changes to the pentasul- 
 phide of arsenic and hydrogen sulphide, just as the orthoarsenic 
 acid changes to arsenic pentoxide and water at a red heat : 
 
 2 As S 4 H 3 = As 2 S 5 + 3 H 2 S, 
 2 As0 4 H 3 =As 2 5 -f 3H 2 0. 
 
 When arsenic trisulphide is dissolved in the sulphides of the alkalies 
 which contain an excess of dissolved sulphur,! then the sulpharsen- 
 ates are found in the solution ; the superfluous sulphur sulphurizes 
 the trisulphide just as oxygen oxidizes the trioxide ; the penta- 
 sulphide so formed, of course, produces the sulpharsenate by union 
 with the metallic sulphide : 
 
 As 2 S 5 +K 2 S=2 AsS 3 K. 
 
 The following table shows the most important facts discussed in 
 this chapter : 
 
 * H 3 As O 4 + H 2 S = H 3 As O 3 + H 2 O + S. If the solution of arsenic acid 
 is strongly acid with hydrochloric acid, then arsenic pentasulphide is precipi- 
 tated by means of hydrogen sulphide (Brauner and Tomicek; Monatshefte 
 fur Chemie ; 1887, 607). In any event, even if the arsenic acid is perfectly pure, 
 only a partial reduction takes place. See also McCay, Am. Chern. Journ. 12, 547. 
 
 t See foot-note, page 155. 
 
ARSENIC ; TABLE OF OXIDES AND SULPHIDES. 
 
 247 
 
 OXIDES. 
 
 As 2 O 3 , forms no acids, forms meta-arsenites, M As O 2 , and orthoarsenites, 
 
 M 3 AsO 3 . 
 
 ( As O 3 H, meta-arsenic acid, forms meta-arsenates. 
 As 2 O 5 < As O 4 H 3 , orthoarsenic acid, forms orthoarsenates. 
 
 ( As 2 O 7 H 4 , pyroarsenic acid, forms pyroarsenates. 
 
 HALOGEN COMPOUNDS. 
 
 Arsenic forms trihalides, As X 3 , but no pentahalides, As X 5 , excepting the fluoride. 
 The trihalides are decomposed by an excess of water, forming arsenic trioxide 
 and halhydric acids : 2 As X 3 + 3 H 2 O = As 2 O 3 + 6 HX. 
 
 SULPHIDES. 
 
 As 2 S 3 , forms no acids, forms metasulpharsenites, M As S 2 , orthosulpharsen- 
 ites, M 3 As S 3 , and pyrosulpharsenites, M 4 As 2 S 5 . 
 
 As,S 5 , forms orthosulpharsenic acid, ^ M As S 3 , metasulpharsenates. 
 As S 4 H g , which yields : ) ^ As S * < orthosulpharsenates. 
 
 J ( M 4 As 2 S 7 , pyrosulpharsenates. 
 
 
 AESENITE8 AND SULPIIAKSEMITES. 
 
 AB8ENATE8 AND 8ULPHAE8ENATE8. 
 
 Meta-salts. 
 Ortho-salts. 
 Pyro-salts. 
 
 M As O 2 , M As 4 S 2 . 
 M 3 As O 3 , M 3 As S 3 . 
 
 AT \ " S 
 
 M As O 3 , M As S 3 . 
 
 M 3 As O 4 , M 3 As S 4 . 
 M 4 As 2 O 7 , M 4 As 2 S 7 . 
 
 , JJ. 4 AB] 5 . 
 
248 ANTIMONY; OCCURRENCE, PREPARATION. 
 
 CHAPTER XXXIII. 
 
 ANTIMONY AND STIBINE. THE COMPOUNDS OP ANTI- 
 MONY WITH THE HALOGENS. 
 
 Antimony ; symbol, Sb ; atomic weight, 120 ; specific gravity, 6.7. 
 Specific gravity of vapor, at 1640, air = 1, is 9.78 ; H 2 = 2 is 
 281.66 ; molecular iveight of Sb 2 = 240. Stibine ; formula, 
 Sb H 3 ; specific gravity not determined. 
 
 BUT little antimony is found as the native * element ; it is most 
 frequently encountered, combined with sulphur, in the mineral stib- 
 nite, Sb 2 S 3 , a lead gray substance with metallic lustre, from which 
 most of the antimony of commerce is obtained. Stibnite has been 
 known since the most ancient times. It is mentioned by Diosco- 
 rides as trrt/x/xt and by Pliny as stibium. It was mainly used in 
 medicine as an external application, but it also formed a pigment for 
 blackening the eyebrows. The name antimonium was applied to it 
 at a later date. The element and its compounds interested the 
 immediate successors of the alchemists greatly, for they thought 
 them to be most wonderful and potent medicinal remedies. 
 
 Antimony is prepared from its sulphide by one of two common 
 metallurgical processes ; the compound is either melted with iron, 
 by which means ferrous sulphide and antimony are produced ; or it 
 is roasted in a draught of air, the sulphur burned off, and the re- 
 sulting oxide of antimony ( Sb 2 4 ) further heated with charcoal, 
 whereupon antimony and carbon monoxide are formed. 
 
 Antimony is much more metallic in its nature than is arsenic ; 
 its appearance indicates this, for it is silver white, with a metallic 
 lustre ; the metal is neither malleable nor ductile, has a crystalline 
 structure, f is brittle and easily pounded into a fine powder ; it 
 melts at 425 and is vaporized at a high red heat, t The specific 
 
 * A " native" element is an element occurring uncombined as a mineral, 
 t An amorphous form of antimony has also been described. 
 J This vaporization is much retarded if the element is covered with a 
 layer of oxide. 
 
ANTIMONY; ALLOYS. 249 
 
 gravity of vapor of antimony at 1640 is 9.78. * This number is 
 somewhat greater than the one which should be found were the 
 molecule of antimony Sb 2 , so that antimony certainly has no stable 
 molecules of the formula Sb 4 , corresponding to those of arsenic and 
 phosphorus, As 4 and P 4 . Probably, were it possible to ascertain 
 the specific gravity of antimony at about 1800, we would find a 
 value which would indicate molecules formed of Sb 2 , or perhaps 
 even of the individual atoms. 
 
 Antimony, when heated to a high red heat in air or in oxygen, 
 burns to form the trioxide, Sb 2 3 ; it unites with the halogens in 
 the same way, powdered antimony even burns vigorously when 
 dropped into a flask containing chlorine, while the trichloride, 
 SbCl 3 , is produced. Antimony shows its metallic nature, chemi- 
 cally, by dissolving in hydrochloric or sulphuric acid; with the 
 former, it produces the trichloride and hydrogen, with the latter the 
 sulphate of antimony and sulphur dioxide, for, as dilute sulphuric 
 acid does not attack the element, reduction of the hot and concen- 
 trated acid takes place exactly as it does when that substance is 
 heated in contact with copper. (See pages 75, 137 and 151.) 
 
 Nitric acid oxidizes antimony, as it does phosphorus or arsenic, 
 but the reaction varies according to the temperature, concentration, 
 and mass of the acid ; where the latter is cold and/dilute, antimony 
 trioxide is formed ; as the temperature is increased, or the acid 
 becomes m^re concentrated, antimonic acid begins to be produced 
 until, under proper conditions, this may be the entire result of the 
 oxidation ; of course, the nitric acid is reduced at the same time. 
 ( See pages 199 and 206.) 
 
 Antimony is an ingredient of a number of commercially very 
 important alloys. 
 
 A combination produced by fusing two or more metals together 
 is termed an alloy. Some meWla can be alloyeH. with each other in 
 ail proportions, others only partially, while sdme wiM not mix at 
 all ; an exactly parallel case is* found in the behavior of ordinary 
 liquids, some of which, like alcohol .and water, can foe mixed in any 
 ratio, others, like water and ether, will only partially dissolve each 
 other, while lastly, some oils and water remain" entirely separate. 
 The question as to whether alloys are mere* mechanical mixtures of 
 
 * H. Biltz and Y. Meyer; Ber. d. Deutsch. Chem. Gesell.; 22, 726. 
 
250 ALLOYS; CHEMICAL NATUKE OF. 
 
 fused metals, or whether they have the character of chemical com- 
 pounds, has been the subject of continued discussion, and, indeed, 
 the same may be said of solutions, which latter certainly have not 
 the characteristics of mere mechanical mixtures. The facts sustain- 
 ing the theory of chemical combination are as follows. An alloy 
 has a specific gravity different from the mean of the specific gravi- 
 ties of the component metals and a melting point which also is not 
 the mean of the melting points of constituents. The melting point 
 of a metal is generally diminished by being alloyed ; some alloys 
 become liquid at a temperature much below that at which any of 
 the constituents fuse. Heat is given off in the formation of alloys,* 
 while many also have a definite crystalline form. ^Nevertheless, those 
 alloys which are produced by metals which mix in any proportion, 
 do not have that characteristic which is supposed to be the sine qua 
 non of a chemical compound the definite composition by weight 
 and yet they cannot be separated into their constituent parts by sim- 
 ple mechanical means. These discrepancies are explained by the 
 theory that certain definite compounds of the metals are really 
 formed, but that these are further dissolved in the excess of one or 
 the other of the constituents. This theory is borne out by the fact 
 that many alloys, on being slowly cooled, allow mixtures of a defi- 
 nite crystalline form and gravimetric composition, to separate in the 
 same way as ordinary solid chemical compounds can be ciystallized 
 from solutions. f Those metals which are chemically most like 
 each other, generally mix most easily to form alloys. The majority 
 of metals are white or gray, and most alloys are also white or gray ; 
 copper and gold are red and yellow respectively ; their alloys pre- 
 sent modifications of these colors, unless the admixture of the 
 a,dded metal is so great in quantity as to entirely conceal this char- 
 acteristic. $ Amalgams, as we have already seen, are alloys con- 
 taining mercury. They are easily fused if, indeed, they are not 
 .soft at ordinary temperatures ; the mercury evaporates when the 
 amalgam is heated to the boiling point of that element. Some 
 
 * Fused zinc and fused copper when mixed with each other liberate so 
 much heat that the mass may spatter out of the containing crucible, and a 
 large part of the zinc evaporates. It must also be remembered that heat is 
 not infrequently given off in the formation of mere solutions. 
 
 t Rudberg; Poggendorff s Annalen 18, page 240. 
 
 t An admixtureof 30% of tin, to copper, will entirely destroy the red color. 
 
 In that way amalgams resemble crystals with water of crystallization. 
 
STIBINE ; PREPARATION, PROPERTIES. 251 
 
 amalgams have a definite crystalline form and chemical composition. 
 If, as seems unavoidable, we regard alloys as chemical compounds, 
 we must, nevertheless, believe that they certainly cannot be classed 
 with substances ordinarily considered as such ; 011 the other hand, 
 their existence is a constant argument against a too dogmatic con- 
 ception of the laws of definite proportions and of valence. 
 
 The most important alloy of antimony is composed of one part 
 of that metal to four of lead ; this substance is used as type metal, 
 copper and bismuth are occasionally added to this ; the metal used 
 for stereotyping also contains tin. 
 
 The hydrogen compound of antimony is called stibine. It is 
 produced, like arsine, by the action of acids on the alloy of anti- 
 mony and zinc,* of by the action of nascent hydrogen on soluble 
 compounds of antimony ; the gas is therefore produced by the same 
 means which furnish arsine, as a consequence the latter may con- 
 tain stibine unless care is taken to exclude that gas ; in fact, stibine 
 may be mistaken for arsine unless especial precautions are taken to 
 distinguish between the two. 65 
 
 Stibine is a colorless and odorless gas, scarcely soluble in water. 
 As would be expected, owing to the metal-like nature of antimony, 
 this gas is much less stable than arsine ; it can be compared to 
 hydrogen telluride (page 104) in this respect, for it partially de- 
 composes into antimony and hydrogen, even at 56. From this 
 instability it follows that stibine cannot be obtained free from 
 hydrogen at ordinary temperatures. Stibine changes to a liquid at 
 -91.5; the liquid boils at 18, and solidifies at 102.5, form- 
 ing a snow-like mass ; when passed through a heated glass tube it 
 decomposes into antimony and hydrogen, just as arsine does into 
 arsenic and hydrogen ; the antimony is deposited in the form of a 
 mirror, which is more metallic in appearance than is the one formed 
 of arsenic, and which, furthermore, is not readily volatilized. When 
 ignited, stibine burns with a white flame, giving off a dense smoke 
 of the trioxide of antimony ; a spot of metallic antimony will form 
 on a cold porcelain plate held in the flame. Of course chlorine, 
 bromine, or iodine attacks stibine, producing the corresponding 
 halhydric acid and the halogen compound of antimony. 
 
 As the basic properties of the hydrogen compounds of the ele- 
 
 * The alloy consists of three parts zinc to two of antimony. 
 
 v 
 
 Of THS 
 
252 ANTIMONY; HALIDES OF. 
 
 merits of this family have already disappeared when arsine is 
 reached, it follows that stibine can form 110 salts corresponding to 
 those of ammonium and phosphonium. 
 
 Another (solid) compound of antimony and hydrogen has been 
 described ; to this the formula Sb 4 H has been ascribed ; its exist- 
 ence, however, is doubtful. 
 
 Antimony forms two series of compounds with the halogens, 
 Sb X 3 and Sb X 5 ; they correspond to those of phosphorus. 
 
 SbF 3 , solid. 
 
 Sb C1 3 , solid ; soft crystalline mass, melts 73, boils 223. 
 
 Sb Br 3 , solid ; melts 90, boils 280. 
 
 Sb I 3 , solid ; melts 167, boils 401. 
 
 Sb F 5 , not crystalline, gum-like mass. 
 
 Sb C1 5 , fluid; melts 6, can be distilled without'decomposition in a vacuum; 
 boils at 68, 14 m.m. pressure. 
 
 Sb I 5 , solid ; melts 78, and decomposes when beated to a higher tempera- 
 ture. 
 
 The trihalogen compounds can be produced by direct union of 
 the elements, and the pentahalogen compounds by addition of halo- 
 gen to the trihalides. The trichloride has been most thoroughly 
 studied, and its chemical behavior will serve as an example for that 
 of the other haloid compounds. The trichloride is, as would be ex- 
 pected, much less readily decomposed by water than the correspond- 
 ing trichloride of arsenic ; indeed, it can even be dissolved in 
 tolerably dilute hydrochloric acid without decomposing. When 
 ifc is added to water, it does not completely break down, but changes 
 into a so-called basic salt. 
 
 A basic salt is one which is formed by the interaction of less of 
 the acid and more of the base than is necessary to produce the normal 
 salt. 
 
 We can consider all oxides of the metals which act as bases as 
 being derived from the corresponding hydroxides by loss of water, 
 the manner of formation being similar to that of the anhydrides : 
 
 H 2 S0 4 - H 2 = S0 3 ; Ca(OH) 2 - H 2 = CaO; 
 2P (OH) 8 - 3 H 2 = P 2 3 ; 2Fe (OH) 3 - 3 H 2 = Fe 2 3 . 
 
 With this in mind we can define a basic salt as one in which 
 only a portion of the hydroxyl groups of a base have reacted with 
 an acid to form a salt and water ; it is therefore the reverse of a 
 so-called acid salt (page 140 and foot-note), in which only a portion 
 
ANTIMONY ; BASIC SALTS OF. 253 
 
 of the hydrogen in an acid has been replaced by a metal. The 
 same nomenclature can therefore be adopted with both classes of 
 salts, and the term primary, secondary, and tertiary basic salts can 
 be used. When water is added to the trichloride of antimony the 
 following reaction takes place : 
 
 f Cl + HOH ( OH 
 
 Sb-3 Cl + HOH = Sb-5 OH + 2HC1. 
 
 ( Cl ( Cl 
 
 The compound Sb ( OH ) 2 Cl is therefore the primary basic 
 chloride of antimony ; the two hydroxyl groups contained therein 
 can afterward separate water, as follows : 
 
 (OH 
 
 Sb^OH= 
 (Cl 
 
 The compound Sb Cl is therefore also a basic salt. In this 
 case the group of elements Sb is a radicle which chemically re- 
 sembles a monovalent metal, as a comparison of the following 
 formulae will make clear : 
 
 (SbO)Cl,* 
 Nad, 
 KC1. 
 
 The monovalent radicle Sb is sometimes called stibionyl, and 
 is quite frequently encountered in the basic salts of antimony ; such 
 an instance is found in the formula of stibionyl sulphate (Sb 0) 2 S0 4 . 
 
 Those metals which have not a very pronounced metallic char- 
 acter are the ones which form basic salts; metals like sodium, 
 potassium, or calcium do not so produce them. 
 
 Antimony trichloride can combine with the chlorides of a num- 
 ber of metals to form double chlorides with formulae like the fol- 
 lowing, Sb C1 3 , 3 K Cl ; more extended mention of these will be 
 made in the chapter on aluminium. 
 
 The pentachloride of antimony is an unexpectedly stable com- 
 pound ; it can be boiled in a vacuum without change,! and, further- 
 more, it is not decomposed by cold water, but forms a crystalline 
 substance containing water of crystallization. 
 
 * More complicated basic chlorides than this one also exist; for their 
 study the student must refer to a larger work. 
 
 t Anschiitz and Evans; Liebig's Annalen; 239, 285. 
 
254 ANTIMONY ; TRIOXIDE. 
 
 CHAPTER XXXIV. 
 
 THE COMPOUNDS OF ANTIMONY WITH OXYGEN, AND WITH 
 OXYGEN AND HYDROGEN. THE SULPHIDES OF ANTIMONY. 
 
 ANTIMONY forms the following oxides, Sb 2 3 , Sb 2 4 , Sb 2 5 ; 
 only the last one of these is a pronounced acidic anhydride ; indeed, 
 the trioxide acts as a base when brought in contact with pronounced 
 acids. 
 
 Antimony trioxide is found in nature as the mineral senarmon- 
 tite. It can be formed by burning antimony, or by oxidizing the 
 element with dilute boiling nitric acid. Chemically, it acts just 
 like the corresponding oxide of arsenic ; it dissolves in alkaline 
 hydroxides to form meta-antimonites : 
 
 Sb 2 3 + 2 KOH = 2 Sb(XNa + H 2 0. 
 
 Two hydroxides derived from this oxide are known ; one, which 
 might be called pyro-antimonous acid, has the formula, Sb 2 5 H 4 ; 
 the other is the normal hydroxide, Sb(OH) 3 , corresponding to 
 orthophosphorous acid, P (OH ) 3 . Antimony trioxide dissolves in 
 pronounced acids to form salts in which antimony is the metal ; so, 
 for instance, it produces the trichloride with hydrochloric acid : 
 
 Sb 2 3 + 6H Cl= 2 Sb C1 3 + 3H 2 0. 
 
 The salts formed with other than the halhydric acids are basic 
 ones,* insoluble in water ; while the trichloride, formed by the 
 above reaction, is also converted into the insoluble basic chloride by 
 adding water to the acid solution (see page 252). Antimony triox- 
 ide is easily reduced to the metal by heating with charcoal or in 
 hydrogen. When heated in the air it takes up oxygen and changes 
 to the tetroxide, Sb 2 4 , which substance is likewise formed from 
 the pentoxide by heating, so that Sb 2 4 is the most stable of the 
 oxides of antimony. The trioxide is but little soluble in water. 
 
 TJie pentoxide of antimony, Sb 2 5 , is produced by oxidizing anti- 
 mony, with fuming nitric acid (page 202), or with aqua regia (page 
 
 * A neutral and an acid sulphate have been described by Adie ; Journ. 
 Chem. Soc. ; 1890 ; 540. 
 
ANTIMONIC ACIDS. 255 
 
 203) ; it loses oxygen at red heat and changes into the tetroxide, 
 Sb 2 4 . 
 
 The basic properties of Sb 2 3 have been entirely destroyed by 
 adding the two oxygen atoms necessary to form Sb 2 5 , so that 
 the latter substance forms no salts with acids. 
 
 The following acids derived from antimony pentoxide are known ; 
 they correspond to those of phosphorus and of arsenic, and are : 
 Sb 3 H, meta-antimonic acid, 
 
 Sb 4 H 3 , ortho-antimonic acid, 
 
 Sb 2 7 H 4 , pyro-antimonic acid, 
 
 An acid having the formula, Sb 2 9 H 8 , is also known ; this is a 
 hydrated pyro-antimonic acid formed as follows : 
 
 Sb 2 7 H 4 + 2H 2 = Sb a 9 H 8 . 
 
 The acid H 8 Sb 2 9 changes into ortho-antimonic acid on standing 
 with water ; it is converted into pyro-antimonic acid at 100 ; pyro- 
 antimonic acid forms meta-antimonic acid at 200 ; and this is finally 
 converted into the anhydride Sb 2 5 at 300. 
 
 Only two sulphides of antimony, Sb 2 S 3 and Sb 2 S 5 , are known 
 with certainty. They are entirely analogous to the corresponding 
 arsenic compounds. 
 
 The trisulphide, Sb 2 S 3 , has already been mentioned as the min- 
 eral antimonite. On roasting in a current of air it is changed, 
 first into the trioxide, Sb 2 3 , and then into the tetroxide, Sb 2 4 . 
 The amorphous sulphide is produced by the addition of hydrogen 
 sulphide to an acidified solution of antimony trioxide, just as the 
 corresponding arsenic compound is formed. It is dark orange- 
 colored powder, insoluble in water or dilute acids,* soluble in 
 alkaline hydroxides or sulphides ; by disolving in the latter it forms 
 a pyro-sulphantimonite ; in this way antimony differs from arsenic, 
 which produces the meta-sulpharsenite under similar circumstances : 
 
 Sb 2 S 3 + 4 KSH = Sb, S 5 K 4 -f 2 H 2 S.f 
 
 (The reaction for the formation of antimonite and sulphantimonite 
 by dissolving Sb 2 S 3 in alkaline hydroxides is given under arsenic, 
 
 * Under certain conditions antimony trisulphide can be obtained in a form 
 which is soluble in water. Schulze ; Journal fiir Praktische Chemie ; [2] 27, 
 320. The statement that a meta-sulphantimonite is also produced is made Ly 
 Ditte; Compt. Rend.; 102, 168. 
 
 t Compare with the following : 
 
 Sb 2 O 3 + 6 KOH = Sb 3 K 3 + 3 H 2 O. 
 
256 ANTIMONY ; PENTASULPHIDE. 
 
 page 245, foot-note.) When the alkaline sulphide contains dissolved 
 sulphur, the sulphaiitimonate is produced ; the same * is the case with 
 arsenic. 
 
 Antimony pentasulphide can be precipitated by adding sulphu- 
 retted hydrogen to a solution containing antimonic acid ; f it bears 
 a complete resemblance to the corresponding arsenic compound, dis- 
 solving in alkaline sulphides or hydroxides to form sulphantimon- 
 ates, the majority of which are salts of the ortho acid, H 3 SbS 4 . 
 Only the salts of the alkalies and alkaline earths $ are soluble in 
 water. The pentasulphide of antimony is precipitated and hydro- 
 gen sulphide is formed when acids are added to solutions contain- 
 ing the sulphantimonates. 
 
 It is scarcely necessary to add a table of the antimony compounds, 
 as they correspond so completely to those of arsenic ; one fact can be 
 reiterated, namely, that antimony can form pentahalogen derivatives, 
 while arsenic is unable to do so. 
 
 * Cl. Zimmerman (personal information). 
 
 t Difference between antimony and arsenic. (See page 246 and foot-note. ) 
 
 J The alkaline earths are calcium, barium, and strontium. 
 
 
BISMUTH ; OCCURRENCE, PREPARATION. 257 
 
 CHAPTEE XXXV. 
 
 BISMUTH. THE COMPOUNDS OF BISMUTH WITH THE HALO- 
 GENS, WITH OXYGEN, WITH OXYGEN AND HYDROGEN, AND 
 WITH SULPHUR. 
 
 Bismuth ; Symbol, Bi ; Atomic iveight, 208.9 ; Specific gravity 
 (electrolytic), 9.74; specific gravity of vapor, air=l, is 10.1, H 2 
 =2, is 290.8 ; molecular weight of Bi 2 is 417.8. 
 
 BISMUTH, the element having the highest atomic weight in this 
 family, is, as would be expected, entirely a metal, although its com- 
 pounds correspond in formula to those of the other members. 
 
 The element occurs in nature as : 
 
 Native bismuth, accompanying cobalt, nickel, and lead ores in gneiss and 
 other crystalline rocks and in clay slate. 
 
 Bismuth trioxide, Bi 2 O 3 ; an earthy mass called bismite. 
 Bismuth trisulphide, Bi 2 S 3 ; bismuthinite. 
 Bismuth telluride, Bi 2 ( Te, S) 8 ; tetradymite. 
 
 It has not been ascertained with certainty how early in the his- 
 tory of metallurgy bismuth became known ; it was formerly con- 
 fused with a number of minerals, all of which went by the name of 
 marcasite. It first became known as a new metal in the fifteenth 
 century, and was called bisemat or ivisemutum, but was even then 
 confounded with antimony ; at the end of the eighteenth century, 
 however, it was universally considered to be a metal. 
 
 Bismuth, like antimony, is prepared from its ores by roasting 
 the sulphide with iron or the oxide with charcoal. The metal which 
 finds its way into commerce contains small quantities of arsenic and 
 iron; the natural metal is nearly chemically pure. The bismuth 
 which is to be used in the preparation of compounds intended for 
 pharmaceutical purposes is subjected to special processes of purifi- 
 cation, for all arsenic must be removed from it. 
 
 Bismuth is a reddish metal with a pronounced metallic lustre 
 and a coarse crystalline structure. Its specific gravity is 9.74. The 
 metal is brittle, neither malleable nor ductile, and can readily be 
 
258 BISMUTH ; PROPERTIES ; HALIDES OF. 
 
 pounded into a powder. It melts at 270, and, in solidifying, ex- 
 pands just as water does in forming ice. It boils at about 1400 ; 
 the specific gravity of its vapor at white heat, (1600 -1700) is 
 10.1.* The calculated specific gravity, were the molecule Bi 2 , is 
 14.4, and is 7.2 if the same were composed simply of individual 
 atoms ; as the number 10.1 is considerably less than that which 
 would be found were the molecule diatomic and to be expressed by 
 the formula Bi 2 , it follows that the vapor of bismuth consists, in 
 part at least, of the individual atoms. At a slightly higher temper- 
 ature than that of the experiment, the molecular and atomic weights 
 of bismuth would probably be identical. 
 
 Dilute hydrochloric or sulphuric acids do not attack bismuth ; 
 concentrated hydrochloric acid has but little effect ; hot and concen- 
 trated sulphuric acid dissolves it, forming the sulphate of bismuth 
 and sulphur dioxide (see page 137) ; nitric acid quite readily forms 
 the nitrate, while it is itself reduced. (See page 206.) 
 
 A number of alloys of bismuth are commercially important ; 
 they all have low melting points ; for example, a mixture of five 
 parts of bismuth, three of lead, two of tin, and three of cadmium 
 fuses at 65. One of the chief uses of these alloys is in the copy- 
 ing of wood-cuts and in stereotyping. 
 
 The halogens form the following compounds with bismuth : 
 
 Bi F 3 , bismuth trifluoride. 
 
 Bi C1 3 , "bismuth trichloride. 
 
 Bi Br 3 , bismuth tribromide. 
 
 Bi I 3 , bismuth tri-iodide. 
 
 All of these compounds are solid bodies at ordinary temperatures. Other 
 halogen derivatives of bismuth are described ; they are unstable bodies which 
 have probably not been obtained in a pure state. 
 
 The trihalogen compounds can either be formed by the direct 
 union of bismuth with chlorine, bromine, or iodine, or by dissolving 
 the trioxide, Bi 2 3 , in concentrated halhydric acids, this oxide of 
 bismuth being entirely basic in its character. The chloride or 
 bromide of bismuth can be dissolved in very little water or in acids 
 without change ; these salts, however, decompose on the addition of 
 an excess of water just as do the corresponding antimony com- 
 pounds (see page 253), while at the same time the insoluble basic 
 chloride or bromide is formed : 
 
 * H. Biltz and V. Meyer; Ber. d. Deutsch. Chem. Gesell.; 22, 725. 
 
BISMUTH; PROPERTIES; OXIDES OF. 259 
 
 ( Cl (OH 
 
 BiJ C1 + 2H 2 = BiJ OH + 2HC1, 
 (01 (01 
 
 (OH 
 
 BiJOH= B^Xi -fH 2 0. 
 (01 
 
 The iodide of bismuth is decomposed only by boiling water ; it is 
 insoluble in cold water. The halogen compounds of bismuth form 
 a number of double salts which will be discussed later j * the for- 
 mulae of a few of these salts are as follows : 
 
 K Cl, Bi C1 3 + H 2 O ; 2 Na Cl, Bi C1 3 + H 2 O. 
 
 2 K Cl, Bi C1 3 + 2 H 2 O ; 2 Na Cl, Bi C1 3 + 3 H 2 O. 
 
 An acid composed of the fluoride of bismuth and of hydrofluoric acid, 
 Bi F 3 , 3 HF is also known. 
 
 The following oxides of bismuth are known : 
 
 Bi O, bismuth monoxide, brownish black, crystalline powder, 
 Bi 2 O 3 , bismuth trioxide, light yellow powder, 
 Bi 2 O 4 , bismuth tetroxide, dark brown powder, 
 Bi 2 O 5 , bismuth pentoxide, brown powder. 
 
 None of these oxides dissolve in alkaline hydroxides to form 
 salts. The one with least oxygen, Bi 0, is produced by the reduc- 
 tion of salts derived from the oxide Bi 2 O 3 ; f it is insoluble in 
 water, and, when dry, oxidizes so readily in the air that the com- 
 pound burns like a piece of tinder ; the trioxide is produced by this 
 combustion. 
 
 Bismuth burns at a high red heat, forming the trioxide, Bi 2 3 . 
 This substance can be more readily prepared by the addition of 
 potassium hydroxide to a solution of a bismuth salt. The first 
 product formed is the hydroxide, Bi (OH) 3 , which is insoluble in 
 water : - Bi ( N o 3 ) 8 + 3 KOH = Bi (OH) 3 + 3 KNO,. 
 
 This reaction is a usual one with most metals, for by far the 
 greater number of metallic hydroxides are insoluble in water. The 
 following reactions are typical : 
 
 * For a discussion of the nature of double salts, see chapter on Aluminium, 
 and also Remsen; Amer. Chem. Journ.; 14, 165, and Inorganic Chemistry 
 (Holt & Co., 1890), page 461. 
 
 t By treating solution of chloride of bismuth, in hydrochloric acid, with 
 an alkaline solution of stannous chloride, Sn C1 2 . See chapter on Tin. 
 
260 BISMUTH NITRATE, SUBNITRATE. 
 
 Fe S0 4 + 2 KOH = K 2 S0 4 + Fe (OH) 2 , 
 Cu S0 4 + 2 KOH = K 2 S0 4 + Cu (OH),, 
 Mg (N0 3 ) 2 + 2 KOH = K 2 (N0 8 ) 2 + Mg (OH) 2 . 
 
 In each of these cases the insoluble hydroxide is precipitated by 
 the addition of a soluble hydroxide to the salts of the metals. The 
 hydroxides of the alkalies are the only ones which are readily solu- 
 ble in water ; those of the metals calcium, barium, and strontium are 
 soluble with some difficulty, all others are insoluble. Those hy- 
 droxides of the metals which are not capable of being dissolved by 
 water are generally converted into the corresponding oxides by the 
 application of a very moderate heat. The hydroxide of bismuth 
 loses water below 100, and is converted into a compound having 
 the formula Bi 2 H : * 
 
 (OH 
 
 Bi^OH = Bi 
 (OH 
 
 and this, lastly, above boiling heat, is changed to the trioxide : 
 
 2 Bi 2 H = Bi, 3 + H 2 0, 
 just as nitrous acid changes into water and N 2 3 . 
 
 The various salts of bismuth are produced by dissolving the 
 oxide or the hydroxides in acids. The chloride has already been 
 discussed, the only other ones we need mention are the nitrate and 
 sulphate. . 
 
 BISMUTH NITRATE. Bi ( NO 3 ) 8 + 5 H 2 O.t Prepared by dissolving 
 either the element bismuth, or the oxide, or one of the hydroxides in 
 nitric acid. Bi 2 O 3 + 6 HNO 3 = 2 Bi ( NO 3 ) 3 + 3 H 2 O. Clear, large 
 crystals, which melt at 80 in their water of crystallization, and which 
 then give off nitric acid, forming the basic nitrate at 120. 
 
 BASIC BISMUTH NITRATE. ( Subnitrate of bismuth ) Bi ONO 3 + H 2 O,f 
 
 or (OH 
 
 Bi^OH 
 
 (N0 3 . 
 
 Formed by decomposing the nitrate with hot water. Insoluble in 
 water. It is a primary basic salt in which one hydroxyl group has 
 been replaced by the group NO 3 ; when heated in a platinum dish, it 
 
 * This hydroxide is the one corresponding to nitrous acid NO 2 H. 
 t Yvon; Bull. Soc. Chim. ; [2], 27, 491. 
 
 J More complicated formulae have been assigned ; for these a larger man- 
 ual must be consulted. 
 
BISMUTH PENTOXIDE ; SULPHIDES. 261 
 
 changes to the oxide Bi 2 O 3 , giving off nitric acid. The salt Bi ONO 3 
 can be considered as derived from the hydroxide Bi O 2 H, by replacing 
 the hydroxyl group with NO 8 . 
 
 Bi | H + HN0 3 . Bi <[ n_ + H 2 O. 
 
 The subnitrate of bismuth is the most important salt of bismuth, 
 
 being extensively used in medicine. 
 
 BISMUTH SULPHATE. Bi 2 (SO 4 ) 3 + 7 H 2 O. Prepared by dissolving 
 the oxide in concentrated sulphuric acid and then adding water. At 
 100 it changes to Bi 2 (SO 4 ) 3 + 3 H 2 O. Boiling water changes it into 
 the insoluble basic sulphate. 
 
 fOBiO 
 
 J OBiO = (BiO) 2 S0 4 . 
 
 01 o 
 
 to 
 
 The univalent group Bi O corresponds to stibionyl (page 253). 
 
 The pentoxide of bismuth, Bi 2 5 , is produced by oxidizing the 
 trioxide by means of hydrogen peroxide. It is an orange or brown 
 colored powder, which forms but one hydroxide, the latter corre- 
 sponding to the meta-acids. This hydroxide, Bi 3 H, is changed to 
 the oxide Bi 2 5 on heating. Neither the oxide nor hydroxide has 
 basic or acid properties.* 
 
 Bi 2 4 has been considered as a compound formed by replacing 
 the hydrogen of Bi 3 H by the univalent group of elements Bi 0, 
 acting as a metal : 
 
 Bi Bi 3 = Bi 2 4 . 
 
 The sulphides of bismuth correspond to the oxides Bi and 
 Bi 2 3 . Bismuth monosulphide, Bi S, is of little importance. Bis- 
 muth trisulphide is found in nature as the mineral bismuthinite. 
 It can be produced by fusing bismuth and sulphur together, or it 
 can be formed by the action of hydrogen sulphide on a solution of 
 a bismuth salt : 
 
 2 Bi C1 3 + 3 H 2 S = Bi 2 S 3 + 6 H Cl. 
 
 The sulphide of bismuth, like the oxide, has no resemblance in 
 its action to the acidic anhydrides ; it cannot be dissolved in the 
 hydroxides or sulphides of the alkalies to form sulpho-salts. On 
 the other hand, it is readily attacked by concentrated hydrochloric 
 acid to form the trichloride of bismuth and hydrogen sulphide : 
 
262 
 
 ELEMENTS OF NITROGEN FAMILY ; 
 
 Nitric acid dissolves it to form the nitrate of bismuth ; while the 
 hydrogen sulphide which is set at liberty is, of course, oxidized by 
 the excess of nitric acid to form sulphur, and, finally, sulphuric 
 acid. 
 
 The elements of the nitrogen family, as we have seen, show the 
 same graduation in properties with increasing atomic weights as was 
 displayed by the elements belonging to the sulphur group. This is 
 best brought to light by the following table : 
 
 ATOMIC 
 WEIGHTS. 
 
 USUAL PHYSICAL 
 CONDITION. 
 
 MELTING 
 POINTS. 
 
 BOILING 
 POINTS. 
 
 SPECIFIC 
 GRAVITY. 
 
 PBOPEETIES. 
 
 N 14.03 
 
 A gas. 
 
 203 
 
 193 
 
 
 
 not-metallic. 
 
 P 31. 
 
 Yellow, solid, eas- 
 ily fused.* 
 
 44 
 
 250 
 
 1.82 
 
 
 As 75. 
 
 Steel gray, brittle, 
 crystalline. 
 
 About 450 t 
 
 800" t 
 
 5.7 
 
 
 Sb 120. 
 
 Silver white, brit- 
 tle, crystalline. 
 
 About 425" 
 
 About 1300 
 
 6.7 
 
 
 Bi 208.9 
 
 Reddish, metallic, 
 brittle, crystal- 
 line. 
 
 270 
 
 About 1600 
 
 9.8 
 
 metallic. 
 
 SPECIFIC GRAVITY OF VAPORS. 
 
 
 AIR = 1. 
 
 H 2 = 2. 
 
 TEMP. 
 
 MOLECULES. 
 
 N. 
 
 .9713 
 
 27.9 
 
 Ordinary 
 
 X 2 
 
 P. 
 
 4.16 
 
 119.8 
 
 800 
 
 P 4 
 
 
 3.14 
 
 90.4 
 
 1708 
 
 P 2 and P 4 
 
 As. 
 
 10.3 
 
 296.6 
 
 900 
 
 As 4 
 
 
 5.37 
 
 154.0 
 
 1736 
 
 As 2 
 
 Sb. 
 
 9.78 
 
 281.6 
 
 1640 
 
 Sb 2 
 
 Bi. 
 
 10.1 
 
 290.8 
 
 1650 
 
 Bi! and Bi 2 
 
 * Red phosphorus in part is then changed to yellow at 261. Specific 
 gravity, 2.08 to 2.14. 
 
 t Arsenic volatilizes without previously melting, unless it is under pres- 
 sure. 
 
COMPARATIVE TABLE OF COMPOUNDS. 263 
 
 HYDROGEN COMPOUNDS XH 3 , X 2 H 4 , X 3 H. 
 
 
 
 HEAT 
 
 
 
 
 
 
 
 
 
 OF 
 
 STABIL- 
 
 
 
 BASIC PBOPERTIES. 
 
 
 
 
 
 FORMA- 
 
 ITY. 
 
 
 
 
 
 
 
 
 TION. 
 
 
 
 
 
 
 
 N 
 
 NH 3 
 
 120 K 
 
 
 N a H 4 
 
 N 3 H 
 
 NH 3 + HI = NH 4 I \ 
 
 
 The stability 
 
 P 
 
 PH 3 
 
 43 K 
 
 
 P a H 4 
 
 P 4 H 2 
 
 \ 
 PH 3 +HI=PH 4 I \ 
 
 
 and basic prop- 
 e r t i e s d i- 
 
 
 
 
 
 
 
 \ 
 
 
 m i n i s h with 
 
 As 
 
 As H 3 
 
 -441 K 
 
 
 
 As~H, 
 
 As H 3 
 
 
 increasing me- 
 
 Sb 
 
 ^b H 
 
 
 
 
 
 Sb H 
 
 
 tallicproperties 
 of elements; 
 
 
 
 
 
 
 
 
 
 only ammonia 
 
 Bi 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 combine with 
 
 
 
 
 
 
 
 
 
 acids to form 
 
 
 
 
 
 
 
 
 
 salts.* 
 
 
 
 
 
 
 
 
 
 
 CHLORINE COMPOUNDS, X C1 3 , XC1 5 . 
 
 N 
 
 N C1 3 
 
 Liquid, explosive, 
 
 
 
 t Completely decomposed 
 
 
 P 
 
 P C1 3 
 
 " decomposed by H 2 O t 
 
 PC1 5 
 
 Solid 
 
 by water as follows : 
 X C1 3 + 3 H 2 O = X (O H ) 3 
 
 As 
 
 As Cl 
 
 i 4. 
 
 
 
 + 3 H Cl. 
 
 Sb 
 
 ^\.S V^1 3 
 
 SbCl 3 
 
 Solid, " " t 
 
 SbCl 6 
 
 Solid. 
 
 $ Partially decomposed by 
 water, forming basic chlo- 
 
 Bi 
 
 Bi Cl 
 
 U <( ( (C 4- 
 
 
 
 ' 1 o -T 11 
 
 
 1>1 V^1 3 
 
 
 
 
 riUtrS US lOUOWS . 
 
 XC1 3 + 2H 2 O = X(OH) 2 
 
 
 
 
 
 
 Cl ; X (OH ) 2 Cl = X O Cl + 
 
 
 
 
 
 
 H 2 O. 
 
 
 
 
 
 
 The pentachlorides with 
 
 
 
 
 
 
 water yield phosphoric or 
 
 
 
 
 
 
 antimonic acid. 
 
 OXIDES, HYDROXIDES, SULPHIDES, AND SULPHO SALTS. 
 
 OXIDES. X 2 O, X 2 O3,X 2 O6. 
 
 SULPHIDES. 
 
 N 2 O 
 
 N O 
 N 2 3 
 N 2 4 
 N 2 6 
 
 
 
 
 BiO 
 
 Bi 2 3 
 Bi 2 4 
 Bi 2 5 
 
 
 As 2 S 2 
 As 2 S 3 
 
 
 
 
 As 2 O 3 
 
 
 
 
 
 1%0 3 
 P0 4 
 
 P 2 6 
 
 Sb 2 3 
 Sb 2 4 
 Sb 2 6 
 
 F 2 S 3 
 
 Sb 2 S 3 
 
 Bi 2 S 3 
 
 As 2 5 
 
 1%S B 
 
 As 2 S 6 
 
 Sb 2 S B 
 
 
 
 * Phosphine forms only a limited number of phosphonium salts; these 
 are decomposed by water. The compound N 2 H 4 forms salts with acids ; P 2 H 4 
 does not. N 3 H is acid in its character, resembling H Cl, H Br, and H I. 
 
264 ELEMENTS OF NITROGEN FAMILY. 
 
 The acids derived from the oxides X. 2 O 3 are formed according to the types 
 X (OH) 3 , ortho acids; XO 2 H, meta acids; H 4 X 2 O 5 , pyro acids. Nitrogen 
 forms only the meta acids and salts derived from this ; phosphorus forms only 
 the ortho acids, but salts of the two other ones are known ; arsenic forms 
 no hydroxides, the salts of the meta arsenious acid are the most frequent; 
 antimony forms the hydroxides corresponding to the ortho and pyro acid, 
 the meta-antimonites are the most frequent; bismuth forms the hydroxide 
 Bi(OH) 3 , it has no acid properties. The trioxides of nitrogen and phos- 
 phorus are acidic only, those of arsenic and antimony are both basic and 
 acidic, that of bismuth is basic only. 
 
 The pentoxides are all acidic with the exception of that of bismuth, which 
 is neither acidic nor basic ; the acids derived from the oxides X 2 O 5 are formed 
 according to the types XO 4 H 3 , ortho acids ; XO 3 H, meta acids; X 2 O 7 H 4 , 
 pyro acids. Nitrogen forms only the meta acid, NO 3 H; bismuth forms one 
 hydroxide Bi O 3 H, not acid in its nature. 
 
 The sulphides As 2 S 8 , As 2 S 5 , Sb 2 S 3 , Sb 2 S 5 are insoluble in water; they 
 dissolve in alkaline sulphides to form salts of sulpho acids. These acids are 
 exactly like the oxygen acids with the exception, that in them sulphur has 
 taken the place of oxygen, atom for atom. The sulphides P 2 S 3 , P 2 S 5 , are 
 decomposed by water; the sulphide Bi 2 S 3 has no acidic properties. 
 
ELEMENTS OF CARBON FAMILY. 265 
 
 CHAPTER XXXVI. 
 
 THE ELEMENTS OF THE CARBON FAMILY. 
 
 THE elements of the carbon family are carbon, silicon, germa- 
 nium, tin, and lead. Changes similar to those observed in the pre- 
 ceding family are caused by the increase in the atomic weights 
 belonging to the elements in this one ; but, as the whole family is 
 less not-metallic than is that of which nitrogen is the representa- 
 tive, only two of the elements, namely, carbon and silicon, can form 
 hydrogen compounds ; the transition from not-metal to metal takes 
 place after the second member of the group. 
 
 The alterations in the physical properties of the isolated ele- 
 ments show this increasing metallic character ; for while carbon 
 and silicon are found either in the forms of amorphous black 
 elements or crystalline bodies which, in the case of diamond, may 
 even be transparent, germanium, tin, and lead have a brilliant 
 metallic lustre. Germanium and tin, however, have a crystalline 
 structure,* while lead is the only member of the family which is 
 perfectly malleable and ductile. The fusing points of the elements 
 under discussion diminish with increasing atomic weights, while the 
 specific gravities increase. 
 
 Carbon; infusible, possibly softens in the heat of the electric arc; specific 
 gravity (as graphite) 2.2. 
 
 Silicon; fuses in the heat of the electric arc; specific gravity (graphitoi- 
 dal) 2.49. 
 
 Germanium; melts at about 900; specific gravity 5.46. 
 
 Tin; melts at 230; specific gravity 7.29. 
 
 Lead; melts at 325; specific gravity 11.44. 
 
 The general formulae for the hydrogen compounds of the ele- 
 ments of the chlorine, oxygen, and nitrogen families are respect- 
 ively, XH, XH a , and XH 3 , in the carbon family the corresponding 
 compounds are XH 4 . The power of fixing hydrogen atoms which is 
 
 * Compare with arsenic, antimony, and bismuth in the preceding, less 
 metallic, family. 
 
266 
 
 ELEMENTS OF CARBON FAMILY; COMPARISON. 
 
 possessed by any one atom of a not-metal is exhausted when the 
 number four is reached ; indeed, no elements other than those which 
 we have considered are capable of forming gaseous hydrogen com- 
 pounds with definite formulae.* It follows that the maximum val- 
 ence which any element displays toward hydrogen alone is four (see 
 pages 107, 108). The relationship between the groups of elements 
 and the formulae of the hydrogen compounds becomes more apparent 
 if we arrange, in the order' of their atomic weights, the symbols of 
 those individuals which we have studied, reversing the order ob- 
 served in the table on page 175 : 
 
 ELEMENTS. 
 
 HYDROGEN COMPOUNDS. 
 
 RELATION TO ATOMIC WEIGHTS. 
 
 C N O F 
 
 CH 4 NH 3 OH 2 FH 
 
 As we pass from left to right, 
 
 12 14 16 19 
 
 
 the atomic weights of the ele- 
 
 Si P S Cl 
 
 SiH 4 PH 3 SH 2 C1H 
 
 ments on any horizontal line 
 increase, while their valence to- 
 
 28 31 32 35.5 
 
 
 ward hydrogen diminishes. No 
 
 As Se Br 
 
 AsH 3 SeH 2 BrH 
 
 elements with atomic weights 
 
 75 79 80 
 
 
 lying between those of any two 
 
 Sb Te I 
 
 Sb H 3 Te H 2 I H 
 
 on any horizontal line are 
 known. These elements are 
 
 120 125 127 
 
 , 
 
 therefore a section of that table 
 
 
 
 which would be obtained by 
 
 
 
 arranging all of the elements in 
 
 
 
 the order of their increasing 
 
 
 
 atomic weights. (See page 17.) 
 
 The elements given on the above table are the only ones which 
 are capable of forming gaseous hydrogen compounds. The hydro- 
 gen compound of silicon is more easily decomposed than is that of 
 carbon ; for the rule is without exception that, with increasing atomic 
 weight, in any given family, there is a diminution in the stability of 
 the hydrogen compounds with the members of that family. 
 
 Carbon possesses in the most eminent degree that property which 
 we observed in a rudimentary form in the hydrogen compounds of 
 nitrogen ; namely, the element can form an almost unlimited number 
 of complicated hydrogen compounds derived from a nucleus of 
 carbon atoms, united one with the other, just as the two nitrogen 
 atoms are joined in hydrazin (page 192). As many as sixty 
 carbon atoms are known to be thus united in a long and simple 
 chain, while the variety of compounds may be almost indefinitely 
 
 * See chapter on boron and its compounds. 
 
ELEMENTS OF CARBON FAMILY; COMPARISON. 267 
 
 increased by branching side chains, or by the formation of rings of 
 atoms, each of which can serve as a nucleus for further substitution 
 or addition. The study of these compounds forms, at present, a 
 separate branch of chemistry, which is generally termed organic 
 chemistry, although, of course, there is no real distinction between 
 this and so-called inorganic chemistry. A few of the simpler car- 
 bon and hydrogen compounds will be taken up in the course of this 
 work. 
 
 All of the elements of the carbon family form dioxides, X0 2 , 
 and all but silicon are capable of producing a monoxide, XO. The 
 dioxides of carbon, silicon, germanium, and tin have the character 
 of acidic anhydrides ; the dioxide of tin, however, like the trioxide 
 of arsenic, can be both acidic and basic, for it dissolves as well 
 in acids as in bases to form salts. The monoxide of carbon is 
 neither acidic nor basic, that of germanium is slightly basic, the 
 other monoxides are all basic in their character and form well- 
 defined salts with acids. 
 
 The acids derived from the oxides have the general formulae 
 H 2 XO 3 and H 4 X0 4 ; those of the first class being the meta-acids, 
 those of the second, the ortho-acids; these compounds all are 
 unstable, readily losing water and leaving the corresponding anhy- 
 dride X0.j ; indeed, it is doubtful if carbonic acid exists at all, even 
 in aqueous solution. The silicic acids, both ortho and meta, are 
 changed to silicon dioxide when heated ; they lose a large amount 
 of water even when dried at ordinary temperatures, so that the 
 existence of hydrated silicic acids of definite formula is doubtful. 
 Germanium dioxide apparently forms no hydrates ; the stannic 
 acids, both ortho and meta, are completely dehydrated when heated 
 to redness. Lead dioxide has no acidic properties. 
 
 All inorganic carbonates are derived from a meta-carbonic acid, 
 ILCOg; only a few organic derivatives of ortho-carbonic acid, 
 H 4 CO 4 are known. Both ortho- and meta-silicates exist and form 
 two classes of frequently occurring minerals, while salts of much 
 more complicated silicic acids (formed by the separation of water 
 between two or more formula weights of the ordinary acids) are 
 quite common. Both ortho- and meta-stannic acid, H 4 Sn 4 and 
 H 2 Sn 3 , are known ; although all salts are derived from the 
 latter compound. The relationships between the compounds dis- 
 cussed above are made more apparent by the following table : 
 
268 
 
 ELEMENTS OF CARBON FAMILY; COMPARISON. 
 
 OXIDES. 
 
 META-ACID8. 
 
 OKTHO-ACIUS. 
 
 -M ETA-SALTS. 
 
 ORTHO-SALTS. 
 
 C/-v 
 
 HP A # 
 
 4. 
 
 Mr* c\ 
 
 
 O 2 
 
 Si O 2 
 f^ _ /-\ 
 
 H 2 Si0 3 * 
 
 H 4 Si 4 * 
 
 2 *> ^3 
 
 M 2 Si 3 
 
 MPo O + 
 
 M 4 Si O 4 
 
 SnO 2 
 
 H 2 SnO 3 
 
 H 4 SnO 4 
 
 2 <^e U 3 j 
 M 2 Sn O 8 
 
 M 4 Sn O 4 
 
 "b O 2 
 
 
 
 
 
 The monoxides, with the exception of that of carbon, are bases ; 
 they dissolve in acids to form a number of well-defined salts, which, 
 when compared with the salts formed from the basic lower oxides of 
 the elements of the preceding family, show a similar, though not 
 quite so well marked, tendency to change into basic salts on the 
 addition of water. 
 
 The following sulphides, corresponding to the oxides, have been 
 studied : 
 
 Carbon monosulphide, CS (?) 
 
 Carbon disulphide, C S 2 . 
 Silicon disulphide, Si S 2 . 
 
 Germanium monosulphide, GeS; Germanium disulphide, GeS 2 . 
 Stannic monosulphide, SnS; Stannic disulphide, Sn S 2 . 
 
 Plumbic monosulphide, Pb S; 
 
 The disulphide of silicon is too unstable to enter into other com- 
 pounds ; indeed, it is decomposed even by the moisture of the air ; 
 but the other disulphides dissolve in the sulphides of the alkali 
 metals to form sulpho-salts which, in formula, correspond to the 
 oxy-salts. 
 
 The chlorine compounds and chlorine and oxygen compounds 
 are formed after the general formulae X C1 2 , X C1 4 , and X C1 2 . 
 Of course, a representative of each of these classes is not known for 
 every element in the family ; the most important ones will be indi- 
 vidually discussed in the succeeding chapters. 
 
 * Existence as acids doubtful. 
 
 t Only orthocarbonates of organic compounds are known, 
 t More thorough investigation of germanium salts is necessary. 
 Carbon also forms a number of other, not very well-defined sulphides 
 "with more or less complicated formulae. 
 
CARBON ; OCCURRENCE. 269 
 
 CHAPTER XXXVII. 
 
 CARBON. 
 
 Carbon ; symbol, C ; atomic weight, 12 ; specific gravity, as diamond, 
 3.5, as graphite, 2.14 to 2.35. 
 
 THE element carbon occurs in three modifications, two of which, 
 diamond and graphite, are of crystalline structure, while the third 
 is amorphous carbon, and occurs widely distributed in the form of 
 coal. By far the greater quantity of carbon, however, is found 
 combined in the numerous compounds of that element. Carbon 
 dioxide (as was mentioned on page 167) is an essential constituent 
 of the atmosphere. The carbonate of calcium forms limestone, 
 chalk, marble, and the two crystalline minerals, calcite and arragon- 
 ite ; the combined carbonates of calcium and magnesium, under the 
 name of dolomite, are the principal structure of great masses of 
 rock ; and, furthermore, the carbonates of iron, zinc, barium, man- 
 ganese, and lead are important additions to the mineral wealth of 
 the world. Carbon is also invariably present in all of the innumer- 
 able organic compounds with which we are acquainted ; and, further- 
 more, the products of vegetable disintegration which are classed 
 under the head of coal are, in the main, composed of the element 
 under discussion. 
 
 Carbon is dimorphous ; as diamond it crystallizes in the regular 
 system, while as graphite it is monosymmetric.* 
 
 The greater number of diamonds occur in the older alluvial 
 deposits, but some have been found imbedded in a laminated gran- 
 ular quartz rock called itacolumite ; they are also sometimes present 
 in a species of conglomerate, composed of rounded, siliceous pebbles, 
 quartz, and chalcedony. Diamonds were originally imported into 
 Europe from the East Indies, from which portion of the world and 
 from Borneo the only specimens were procured until the year 1727, 
 when large diamond fields were discovered in Brazil. In 1867 the 
 
 * Formerly supposed to be hexagonal. 
 
270 DIAMOND; OCCURRENCE, PROPERTIES. 
 
 diamond fields of South Africa were opened. Some diamonds are 
 also found in the Urals, in New South Wales, and in the United 
 States. 
 
 The diamond is distinguished by its extreme hardness, its 
 great power of refraction and brilliant lustre ; its specific gravity is 
 3.5 ; it is a poor conductor of electricity and of heat. The mineral 
 also occurs in black pebbles or masses known as carbonado, having 
 a specific gravity, of from 3.01 to 3.4. A coarse variety of diamond 
 which, owing to imperfections in structure, is unfit for jewelry, is 
 sold for glass-cutting purposes, under the name of bort. The weight 
 of the diamond is measured in carats ( 1 carat = .205 gram) ; the 
 price per carat increases in geometric ratio, although always modi- 
 fied by the quality of the stone. The largest diamond is about the 
 size of half a hen's egg ; it originally weighed 800 carats, but was 
 greatly reduced in weight by cutting ; the Pitt or Regent diamond 
 weighs 136.75 carats, and is of unblemished transparency and lustre. 
 When not in contact with the air, diamond can be heated to a white 
 heat without alteration ; when heated between the carbon points of 
 an arc light, it swells and changes to a grayish mass with an almost 
 metallic lustre ; in this form it resembles ordinary coke. When 
 heated in the air, diamond takes fire at about 1000, and then burns 
 to form carbon dioxide (C0 2 ), leaving only a very slight trace of 
 ash. Oxidizing agents, such as fused potassium nitrate, or potas- 
 sium bichromate and sulphuric acid, can oxidize diamond to carbon 
 dioxide. Sir Humphry Davy was the first to prove that diamond 
 consisted of nearly pure carbon. 
 
 Graphite, also called plumbago or black lead, is the second crys- 
 talline form of carbon. It occurs in beds and imbedded masses in 
 the oldest geologic formations, in granite, gneiss, micaceous schists, 
 and crystalline limestone. It is probably, in some instances, the 
 resul^ of the alteration of deposits of coal by heat, although its 
 origin is as yet imperfectly understood. In some places the graph- 
 ite is found quite pure ; for instance, in the " Eureka Black Lead 
 Mine " at Sonora, California, there is a bed from twenty to thirty feet 
 in thickness, which contains the substance in so pure a state that it 
 can be cut in blocks and shipped without further preparation.* The 
 ash left on burning this graphite is only about five per cent of the 
 whole. Sometimes the graphite is, of course, much more impure, 
 so that it may be entirely unfit for use. The chief occurrences of 
 * Practically no graphite is now mined in California. 
 
GRAPHITE; OCCURRENCE, PROPERTIES. 271 
 
 the mineral are in the Urals in Siberia ; in Borrowdale, Cumber- 
 land ; * in Arendal, Norway ; and in some parts of Austria, Russia, 
 and France ; while large quantities are also found in the East Indies. 
 In the United States, the, mineral occurs quite frequently, notably 
 in California ; at Sturbridge, Mass. ; at Ticonderoga, and in the 
 northern part of Michigan. Graphite can be artificially prepared 
 by crystallization of carbon which has been dissolved in melted 
 iron ; for, when gray pig iron is dissolved in acids, the insoluble 
 graphite remains in the form of small, delicate scales ; a similar 
 form of the substance has also been discovered in some meteorites. 
 
 Graphite is used in the manufacture of lead pencils, infusible 
 crucibles and as a lubricator ; it is adapted to the latter purpose 
 because the substance is soft and scaly. It is grayish black, and 
 has almost a metallic lustre. When burned in the air, it forms car- 
 bon dioxide, and leaves an ash, which, when it is derived from the 
 purer varieties, consists mainly of ferric oxide and silica. When 
 heated with concentrated nitric acid for some length of time, graph- 
 ite changes to a yellow, crystalline body, which contains carbon, 
 hydrogen, and oxygen. This substance is known as graphitic acid. 
 Graphitic acid, when heated, disintegrates almost with explosive 
 violence, leaving a voluminous black residuum, which apparently 
 consists of very finely divided graphite. This latter form of the 
 substance is applied as a covering to the moulds used in electroplat- 
 ing ; for, as graphite is a good conductor of electricity, it renders 
 the surfaces of the non-conducting substances from which these are 
 made, capable of conducting electricity, the conductivity of the sur- 
 face being an essential preliminary to forming a metallic deposit. 
 
 Those compounds which are formed in animal and vegetable 
 organisms, and which are classed under the general head of organic 
 substances, are produced by the union of a very few elements ; 
 namely, carbon, hydrogen, oxygen, nitrogen, sulphur, and pjjos- 
 phorus. When such substances decompose" in the open air they 
 break down completely, changing for the most part into gaseous 
 products ; but when the vegetable fibres are protected by a layer of 
 water, as is the case in peat-bogs, the process of decomposition goes 
 on slowly ; certain portions of the constituents of the organic sub- 
 stances, especially oxygen and hydrogen, generally pass off in other 
 combinations, while the vegetable substance becomes changed, first 
 into peat and then into bituminous coal, and, at the same time, the 
 * The mines of Borrowdale are now exhausted. 
 
272 
 
 COAL; FORMATION. 
 
 percentage of contained carbon increases. Peat, brown coal, bitu- 
 minous coal, and anthracite coal are successive steps in the process 
 of floral decomposition ; when the anthracitic stage is reached the 
 changes have become so complete that a black, shiny, homogeneous 
 mass has resulted ; in this mass the original vegetable structure has 
 entirely disappeared, or is, at least, so indistinct that special means 
 must be taken for its detection. The pressure to which the dead 
 organic structures are subjected is of material influence on the 
 rapidity with which a peat formation is changed to anthracite ; in- 
 deed, in districts of Russia where the coal has not been placed 
 under very great pressure, a brown coal (lignite), which can 
 scarcely be distinguished from peat, is found in places where the 
 age of the deposit would lead one to expect anthracite. A compari- 
 son of the approximate composition of the combustible portions 
 of some of the varieties of coal will show the changes which 
 the vegetable matter undergoes during its decomposition more 
 clearly : 
 
 
 CARBON, PEE CENT. 
 
 HYDROGEN, PER CENT. 
 
 OXYGEN ANU NITROGEN, 
 PER CENT. 
 
 Wood 
 
 50 
 
 6 
 
 44 
 
 Peat 
 
 60 
 
 5.75 
 
 34.25 
 
 Lignite 
 Bituminous coal 
 
 67 
 
 87 
 
 5.3 
 
 5.6 
 
 27.7 
 7.4 
 
 Anthracite 
 
 94 
 
 3.4 
 
 2.6 
 
 The various forms of coal are amorphous, and therefore differ 
 markedly from diamond or graphite, both of which are crystalline. 
 
 When organic substances are heated without access of air they 
 undergo a process of carbonization; the volatile products of this 
 destructive distillation pass off as gases and liquids (see page 183), 
 while amorphous carbon is left behind. Similar changes take place 
 during the destructive distillation of bituminous coal, leaving coke.** 
 Coke is a porous, shiny form of amorphous carbon ; it conducts elec- 
 tricity and heat about as well as graphite. Coke may contain as 
 much as 91.5 per cent of its total weight in the form of carbon. 
 
 Gas coal, which collects on the walls of retorts in which bitu- 
 minous coal is heated to form illuminating gas, is a product of the 
 decomposition of gaseous compounds of carbon and hydrogen. It 
 
CARBON; AMORPHOUS. 273 
 
 has almost a metallic lustre, resembling very dense coke ; it is diffi- 
 cult to ignite, and conducts heat and electricity quite well. 
 
 Wood charcoal is produced by the imperfect combustion of wood 
 sticks, animal charcoal by a similar treatment of animal refuse, 
 such as bones or blood. The finer forms of bone charcoal are 
 termed bone black and ivory black. All forms of charcoal, but 
 especially the varieties of animal charcoal, have a remarkably pro- 
 nounced tendency to absorb coloring matters from solutions. 67 
 These colored substances are apparently deposited within the 
 porous substance of the coal, for they can be extracted unchanged 
 therefrom by means of the proper solvents.* The property of 
 absorbing coloring matter does not belong to charcoal alone ; all 
 insoluble porous substances can perform the same office in a greater 
 or less degree ; f in rare instances it may happen that the charcoal 
 exercises a reducing action on the absorbed matter. Crude sugar is 
 decolorized by means of charcoal. 
 
 The purest form of amorphous carbon is lamp-black, which 
 results from the combustion of carbon and hydrogen compounds 
 where an imperfect supply of oxygen is provided, or where ' the 
 flame is cooled before perfect combustion has taken place ; lamp- 
 black is therefore deposited on a cold porcelain or metal plate 
 placed within a luminous gas or lamp flame. The lamp-black of 
 commerce is obtained by burning resinous pine wood, tar, or some 
 kinds of bituminous coal. The substance is collected on coarse 
 cloths hung over the burning wood placed in suitable chambers. 
 Lamp-black is used in the manufacture of printers' and Indian ink. 
 
 * Indigo, which has been dissolved in sulphuric acid and absorbed from 
 this solution by charcoal, can be extracted from the charcoal by alkalies. Me- 
 tallic oxides, absorbed by charcoal, can be extracted by strong acids. 
 
 t Aluminium hydroxide, ferric hydroxide, or precipitated sulphide of lead 
 can absorb coloring matter. 
 
274 METHANE; OCCURRENCE. 
 
 CHAPTER XXXVIII. 
 
 THE COMPOUNDS OF CARBON WITH HYDROGEN. 
 
 Methane ; formula, CH 4 ; specific gravity, air = 1, is .5531, H 2 = 2, 
 is 15.93 ; molecular weight, 16.032. 1 c.c. of the gas at and 
 760 m.m. weighs .0007153 gram. 
 
 THE simplest hydrogen compounds of the carbon family have 
 the formula XH 4 , where X represents an atom of some element of 
 that family; as a consequence the valence toward hydrogen pos- 
 sessed by the atoms of the elements of this group is greater by one 
 than is the valence of those in the preceding (nitrogen) family. 
 
 Methane, or marsh gas (the hydrogen compound of carbon which 
 corresponds to ammonia in the nitrogen family), occurs quite fre- 
 quently in nature as a product of the decay of vegetable tissues. 
 The muddy bottom of any stagnant, marshy pool, when stirred, 
 emits bubbles of marsh gas, which, however, always contain from 
 ten to twenty per cent of carbon dioxide, as well as a small amount 
 of nitrogen. The metamorphoses which resulted in the formation 
 of coal beds, having been similar to the changes taking place in 
 marshes, must necessarily also have produced methane, so that, as a 
 consequence, pockets of the gas (sometimes subjected to great pres- 
 sure) not infrequently occur in coal mines ; the escaping gas, when 
 a pocket is tapped, forms a dangerously explosive mixture with the 
 air. Methane, in some places, escapes from openings in the ground ; 
 the gas which is passing off is sometimes either intentionally or 
 accidentally ignited ; the burning gas wells so produced are, in some 
 instances, regarded with superstitious reverence, as is the case with 
 the holy fire at Baku on the Caspian Sea. 
 
 The natural gas which is used so extensively for illuminating 
 and heating purposes in a number of places in the United States 
 consists, for the most part, of methane. Methane is always pro- 
 duced in large quantity by the dry distillation of coal ; it therefore 
 forms the major portion of illuminating gas. 
 
METHANE; PREPARATION. 275 
 
 Methane cannot be prepared by direct union of the elements 
 carbon and hydrogen, yet if hydrogen can be brought to act upon 
 carbon when the latter is in what may be considered the nascent 
 state, then the circumstances are such that union can take place. 
 An example of such a production of methane is found in the reac- 
 tion between carbon monoxide (CO) and hydrogen, under the in- 
 fluence of a strong discharge of electricity from an induction coil : 
 
 The Preparation of Methane. 
 
 The best method to prepare methane for laboratory use is by the 
 dry distillation of some organic substance, and, as we have seen, 
 our choice of these is not very limited. Experience has shown, 
 however, that a mixture of some dry acetate with a strong base 
 will yield methane ; an example of the production of the gas in this 
 way can be found in the decomposition of the acetate of sodium, when 
 that substance is heated in the presence of sodium hydroxide : 
 
 CH.COONa+NaOH = CH 4 +Na 2 C0 3 
 
 Sodium acetate + Sodiurn hydroxide = Methane + Sodium carbonate. 
 Sodium acetate can be considered as methane in which one hydro- 
 gen atom has been replaced by the univalent group COONa. 
 
 Methane is a colorless gas, without odor or taste. Its specific 
 gravity is .5531. Under a pressure of one atmosphere, methane 
 boils at 164 ; if evaporated quickly under diminished pressure 
 the liquid will be cooled to below its freezing point, and will form 
 a snow-like mass. The gas burns in oxygen or air with a nearly 
 colorless flame which is much like that of hydrogen, while carbon 
 dioxide and water are formed : 
 
 Methane is decomposed into its constituent elements only at 
 quite a high heat ; when passed through a white-hot tube it breaks 
 down into carbon and hydrogen ; it is, therefore (with the exception 
 
 * A similar and most interesting production of methane is by the action of 
 copper on a mixture of the vapors of carbon bisulphide and hydrogen sulphide : 
 
 8Cu + 2H 2 S + CS 2 =4Cu2S + CH 4 . 
 
 Here the carbon and the hydrogen may both be considered to be in the nascent 
 state ; the copper simply removes sulphur and leaves the carbon and hydrogen 
 to rearrange themselves into the most stable configuration under existing con- 
 ditions. (See page 51.) 
 
276 
 
 METHANE ; VOLUMETEIC COMPOSITION. 
 
 of ammonia), more stable than any of the hydrogen compounds of 
 the preceding family. Methane, if it is mixed with exactly enough 
 oxygen to form carbon dioxide,* can be exploded by means of an 
 electric spark ; if care is taken to keep the water vapor produced 
 by this reaction in the form of a gas, the result is that one volume 
 of methane with two of oxygen forms one volume of carbon dioxide 
 and two of water; the total volume of the mixture of gases is 
 therefore the same after the explosion as it was before : 
 
 1 vol. methane, + 2 vols. oxygen, = 1 vol. carbon dioxide + 2 vols. water vapor. 
 
 From these results it follows that one molecule of methane is 
 able to form two molecules of water vapor, and, consequently, as 
 two molecules of water vapor contain four atoms of hydrogen, 
 methane must also have four hydrogen atoms in its molecule. The 
 specific gravity of methane, hydrogen = 2, is 15.9; this shows that 
 its corrected molecular weight must be 16, f for when methane is 
 analyzed we find it composed of 12 parts by weight of carbon and 4 
 parts of hydrogen, which form 16 parts of methane. The 4 parts of 
 hydrogen, as we have seen, represent four atoms; that 12 parts of 
 carbon represent one atom we presume to be the case because, in no 
 compound of carbon, the specific gravity, and hence the molecular 
 weight, of which is known, has carbon ever been found to enter with 
 a less proportional weight than twelve. After considering these 
 experimental facts, we conclude that the formula of methane is CH 4 . 
 ( See page 187.) 
 
 When methane is mixed with chlorine and placed in the dark, 
 no reaction takes place, but when the mingled gases are exposed to 
 the sunlight, a violent explosion results ; hydrochloric acid and car- 
 bon being produced. This action is exactly parallel to the action 
 of chlorine on all other hydrogen compounds (with the exception of 
 hydrofluoric and hydrochloric acids) : 
 
 * A eudiometer tube is used for this purpose. See Note 20, laboratory 
 appendix. 
 
 t In exact numbers 16.032, because the atomic weight of hydrogen is 
 1.008; in considerations of this kind it is better to neglect the decimal. 
 
METHANE; SUBSTITUTION. 277 
 
 CH 4 + 4 Cl = C + 4 H Cl, 
 NH 3 + 3Cl"=N +3HC1, 
 SH 2 + 2 Cl = S + 2 H Cl, 
 BrH + Cl = Br + H Cl. 
 
 When, however, the chlorine is allowed to attack methane 
 slowly, as it does in diffused light, substitution of hydrogen for 
 chlorine results, so that the following changes successively take 
 place : 
 
 CH 4 + 2 Cl = CH 3 Cl + HC1, 
 
 CH 4 + 4 Cl = CH 2 C1 2 + 2 H Cl, 
 
 CH 4 + 6 Cl = CH Cl s + 3 H Cl, 
 
 CH 4 + 8C1= CC1 4 + 4HC1. 
 
 When one hydrogen atom has been removed from methane, the un- 
 saturated univalent radicle (see page 108) is termed methyl, when 
 two hydrogen atoms are removed the bivalent radicle is methylen, 
 and when three hydrogen atoms are removed the trivalent radicle is 
 methin : 
 
 CH 3 , methyl; CH 3 C1, methyl chloride. 
 
 CH 2 =, methylen ; CH 2 C1 2 , methyleii chloride. 
 
 CH ==, methin; CHC1 3 , methin chloride (chloroform). 
 
 These chlorinated substances can, therefore, all be considered as 
 methane in which one, two, or three atoms of hydrogen have re- 
 spectively been replaced by chlorine ; they partake more or less of 
 the nature of methane, although the introduction of successive 
 chlorine atoms causes the resulting compound to depart more and 
 more from the character of the type, thus : 
 
 Methyl chloride is a gas, colorless, becomes liquid at 23. 7. 
 Methylen chloride, liquid, boils at 40. 
 Methin chloride, liquid, boils at 61.2. 
 Carbon tetrachloride, liquid, boils at 76.5. 
 
 With the introduction of each chlorine atom the boiling point in- 
 creases, and, therefore, each of these changes brings the character of 
 the chlorine substituted methane farther from that of the colorless 
 gas from which it is derived. 
 
 When hydrogen is removed from methane, the resulting unsat- 
 urated monovalent group, methyl, cannot exist alone ; in this respect 
 it is like a free atom of hydrogen, and, therefore, methyl seeks the 
 first opportunity of uniting with some atom or radicle to form a 
 
278 HIGHER HYDROCARBONS. 
 
 saturated compound. We have seen that, when chlorine is present, 
 the methyl reacts with that element to form methyl chloride ; if no 
 such other substance with which methyl is capable of union can be 
 found, the radicle will join with itself to form dimethyl (ethane) ; 
 CH 3 CH 3 .* Dimethyl, like methane, is capable of having from 
 one to six of its hydrogen atoms substituted by chlorine. The first 
 reaction which takes place between ethane and chlorine is as 
 follows : 
 
 CH 3 CH 3 + 2 Cl = CH 3 CH 2 Cl + H Cl. 
 
 The compound CH 3 CH 2 C1 (C 2 H 5 C1) is termed ethyl chloride 
 (the same system of nomenclature applies in this case as it does 
 with the methyl compounds,). If we remove one hydrogen atom from 
 ethane to form the monovalent, unsaturated radicle ethyl, the latter, 
 if no other substance is present with which it can unite, will form 
 diethyl ( butane) just as methyl forms dimethyl (ethane) : 
 
 CH 3 CII 3 - H = CII 3 CH 2 ; CH 3 CH 2 + CII 3 CH 2 = CH 3 CH a CH a CH 3 . 
 
 Ethane Hydrogen = ethyl. Ethyl + ethyl = diethyl (butane). 
 
 Diethyl can likewise have its hydrogen atoms substituted by chlo- 
 rine, and, by the loss of one atom of hydrogen, can be converted into 
 the monovalent, unsaturated radicle (butyl), which further unites 
 with itself to form dibutyl or octane. It is possible, however, so to 
 modify the above reaction as to bring ethyl and methyl together, in 
 which case the two radicles will unite to form ethyl-methyl (pro- 
 pane) t ; i n the same way, a mixture of propyl and ethyl will yield 
 propylethyl (pentane). By a judicious combination of the iodides 
 of organic radicles, carbon and hydrogen compounds containing as 
 many as sixty carbon atoms in a molecule t have been prepared. These 
 substances, as they contain only carbon and hydrogen, are termed 
 
 * Methyl iodide boiled with zinc dust or with sodium forms ethane. The 
 reaction takes place as follows : 
 
 CH 3 I + Zn + CH 3 I=ZnI 2 + CH 3 + CH 3 . CH 8 + CH 3 =CH 3 CH 3 . 
 This reaction is also applicable in the formation of the more complicated 
 compounds which follow. Of course, did we wish to prepare some of the latter, 
 we would not use methyl iodide and zinc, but would employ the iodides of those 
 radicles which we wish to unite. 
 
 t By boiling a mixture of ethyliodide and methyliodide with zinc-dust 
 thus : 
 
 CH 3 CH 2 I + Zn +CH 3 I = ZnI 2 + CH 3 CH 2 CH 3 . 
 
 Ethyliodide + methyliodide = zinc iodide + ethylmethyl (propane). 
 
 J Hell and Hagele ; Berichte der Deutsch. Chem. Gesell. ; 22, 502. This 
 hydrocarbon (C m H 122 ) was produced by heating C 30 H 61 1 with sodium. 
 
PETllOLEUM. 
 
 , 
 
 hydrocarbons, the particular class of saturated hydrocarbons now 
 under discussion being called paraffins. The first seven represen- 
 tatives of this class are given in the following table : 
 
 
 BOILING POINT. 
 
 SPECIFIC GRAVITY 
 
 OF LIQUID. 
 
 Methane, CH 4 
 
 -164 
 
 0.415 (at - 
 
 164) 
 
 Ethane, C. 2 H 6 
 
 
 
 
 
 
 Propane, L/ 3 1 8 
 Butane, C 4 H 10 
 
 + 1 
 
 0.6 
 
 
 Pentane, C 5 H 12 
 
 + 37 
 
 0.627 
 
 
 Hexane, C 6 H 14 
 
 + 69 
 
 0.658 
 
 
 Heptane, C 7 H 16 
 
 + 98 
 
 0.683 
 
 
 A general formula for these compounds is C n H 2n + 2 , where n rep- 
 resents the number of carbon atoms in the chain ; an increase of this 
 number by one in any given paraffin of the series under discussion 
 raises the boiling point of that paraffin by about 19. The com- 
 pounds which begin the series are gases, those with from five to 
 sixteen carbon atoms are liquids at ordinary temperatures, the re- 
 mainder are solids with melting points ranging from 18 to 74, * 
 the specific gravities of the hydrocarbons increase with the number 
 of carbon atoms in the molecule, but are always less than unity. 
 
 The hydrocarbons C n H 2 w + 2 are found in coal oil ; the latter is 
 generally technically divided as follows : 
 
 Petroleum ether, boiling point 50 to 70; pentane and hexane. 
 Benzine, u 70 to 90; hexane and heptane. 
 
 Ligroine, " 90 to 120; heptane and octane. 
 
 Petroleum, (kerosene) " 150 to300; octane to hexadecane (C^H^). 
 The higher boiling portions are vaseline and paraffin. 
 
 If one hydrogen atom is removed from ethane, there results an 
 un saturated radicle (ethyl) which cannot exist alone ; we saw that it 
 unites with itself to form diethyl ( butane). Experience has shown 
 us, however, that an entirely different result may be expected if we 
 simultaneously remove a hydrogen atom from each of the carbon 
 atoms contained in ethane ; the molecule containing the pair of 
 neighboring carbon atoms, which have thus become unsaturated, is 
 then capable of independent existence and is called ethylene : 
 
 * A few compounds recently discovered may prove to be an exception to 
 this rule. The hydrocarbon with sixty atoms of carbon in one molecule melts 
 at 102. 
 
280 ETHYLENE. 
 
 CH 3 CH 3 , ethane, CH 2 CH 2 , ethylene. 
 
 What is true of ethylene remains true when the hydrogen atoms 
 of that substance are substituted by other atoms or groups of atoms ; 
 we can, therefore, beginning with ethylene as a nucleus, by repla- 
 cing the hydrogen atoms with ethyl, methyl propyl, etc., construct a 
 new series of unsaturated carbon compounds which would have the 
 general formula C n H 2 n . A further discussion of these complicated 
 substances belongs in the domain of organic chemistry. The fact 
 that carbon atoms are never known to be unsaturated in organic 
 compounds unless the unsaturated atoms are side by side * and the 
 consideration that carbon is tetravalent in methane, have led chem- 
 ists to regard the carbon atoms in ethylene as being joined to each 
 other in a different way from that in which they are in ethane ; for, 
 if we suppose the carbon atoms to be always quadrivalent, then the 
 pair of carbon atoms in ethane are joined by one valence of each 
 atom, where in ethylene they are united by two. The following 
 diagram will make this more clear : 
 H H 
 
 H H 
 
 H C C H c C x 
 
 I I K' X H; 
 
 H H 
 Ethane. Ethylene. 
 
 These structural formulae express the theory that each valence of 
 any carbon atom (which is of necessity tetravalent) must neutralize 
 a corresponding valence of some other atom. All we can really 
 know regarding these combinations is that the carbon atoms in any 
 such compound as ethylene are held together by a certain force, of 
 the nature of which we are ignorant, and which we call chemism or 
 chemical affinity. It is usually stated that " carbon has four points 
 of affinity, or four valences;" of course, provided we consider 
 chemism as a force, such a theory is not tenable, because no force 
 can act unless it has something to act upon ; when the carbon atoms 
 are united as they are in ethylene a certain amount of residual force 
 
 * A few compounds in organic chemistry admit of a different interpreta- 
 tion for their structural formulae ; it seems not improbable, for example, that 
 a hydrocarbon of the formula CH 2 CH 2 CH 2 is capable of existence. 
 In such a compound the two carbon atoms at the ends of the chain would be 
 unsaturated and trivalent. 
 
ACETYLENE. 281 
 
 (beyond that required to hold these two atoms together) acts upon 
 and retains the four hydrogen atoms ; this much we know from 
 experimental evidence ; but if we further suppose that the carbon 
 atoms are each joined by two points of affinity, we must then ac- 
 cept the proposition that the force which unites the two atoms is 
 manifest only from four distinct spots upon their surfaces, an hy- 
 pothesis which is not in accord with what we know as regards the 
 attraction which one mass exercises toward another. A statement 
 which would more nearly accord with our experimental knowledge 
 would be that in ethylene we have two trivalent, and hence unsat- 
 urated, carbon atoms joined to each other.* If by some means we 
 remove a hydrogen atom from each of the neighboring carbon atoms 
 in ethylene, there results a compound CH CH=(C 2 H 2 ) which 
 is termed acetylene. The carbon atoms in acetylene are supposed 
 to be united by a so-called " triple linking," for the theory which 
 was used in explaining the constitution of ethylene must compel the 
 supposition that the two additional unsaturated valences in acetylene 
 must neutralize each other ; the formula of the latter substance is, 
 therefore, written CH = CH. In this case we also must indulge in 
 speculation if we wish to go farther than to assume any more than 
 the existence of two divalent unsaturated carbon atoms in acetylene. 
 Acetylene can act as the nucleus of another series of hydrocarbons 
 with the general formula C n H 2n _ 2 , for, by substituting the hydro- 
 gen atoms in acetylene by organic radicles (methyl, ethyl, etc.), we 
 can produce long chains of atoms. The hydrocarbons of the series 
 C u H 2n and C n H 2n _ 2 , being unsaturated, can readily add substances 
 such as chlorine, bromine, hydrobromic acid, etc. 
 
 Ethylene is a colorless gas which is poisonous ; it has a specific 
 gravity of .9852, is tolerably soluble in water, and when heated 
 breaks down into methane and acetylene : 
 
 * Recent investigations in organic chemistry seem to show that carbon 
 atoms react as if their attractive force were exerted along four lines con- 
 necting the centre of a sphere with four points symmetrically grouped upon 
 its surface; these four points would then correspond to the angles of a regular 
 tetrahedron. The hydrogen atoms in methane would then be at the points of 
 the tetrahedron, and the two carbon atoms in ethylene would be joined along 
 the line connecting two of these points. By this hypothesis the difference 
 between saturated and unsaturated carbon chains can be explained. See also 
 M. M. Pattison Muir; Principles of Chemistry. The discussion here given 
 has been set forth in detail by Lossen; Liebig's Annalen; 204, 265. 
 
282 THE FLAME. 
 
 Owing to this latter decomposition ethylene burns with a lumi- 
 nous flame, because acetylene is a gas which undergoes decomposition 
 into methane and carbon : 
 
 2C 2 H 2 = CH 4 + 3C, 
 
 and therefore it emits a luminous flame, for in each case the glow- 
 ing particles of carbon emit the light. 
 
 Illuminating gas, prepared by the distillation of soft coal, is 
 composed chiefly of hydrogen, methane, ethylene and acetylene, 
 carbon monoxide and carbon dioxide, and nitrogen. The quality of 
 the flame is determined by the amount of ethylene and acetylene 
 present, for a gas which contains these unsaturated hydrocarbons 
 burns with a luminous flame, while methane or carbon monoxide 
 scarcely gives any light during combustion. 
 
 A flame can be observed wherever a gas, in consequence of 
 chemical action, is heated sufficiently to cause it to glow. In most 
 cases this chemical reaction is caused by chemical union ; that it is 
 possible, however, to have a flame resulting from the heat of decom- 
 position of an endothermic compound is proved by the appearance 
 of a flash accompanying the explosion of nitrogen chloride. The 
 simplest case of flame production can be illustrated by mixing two 
 gases, which are capable of giving off a large amount of heat in 
 their union, and igniting the mixture by means of an electric spark 
 or by means of a flame. In this mixture the molecules of the two 
 gases are intimately intermingled, the reaction takes place almost 
 simultaneously at all points throughout the volume of the gas, and 
 is, therefore, accompanied by an explosion and the formation of a 
 homogeneous flame. f 
 
 The most common form of flame is produced by a stream of gas 
 pouring into a volume of another with which it can chemically 
 unite ; if the entering gas is heated to its kindling temperature, 
 then union will take place along the boundary where the two gases 
 touch. The conical shape assumed by the flame of a gas escaping 
 from a round vent is caused by the diminution in the quantity of 
 that gas as the distance from the opening increases, this diminution 
 being caused by the consumption of the gas in burning. The 
 structure of the flame of a gas burning in air can be taken as a 
 type of all others. Such a flame consists of a number of zones 
 which can be easily distinguished by their appearance. The flame 
 
THE FLAME. 
 
 283 
 
 of a candle, for instance, exhibits a dark centre which is surrounded 
 
 by a conical, luminous zone ; a piece of paper placed over this, as 
 
 is shown in Fig. 11, will be charred in the form of a circle, this 
 
 experiment showing that the gases in 
 
 the centre of the flame are not heated 
 
 to a high temperature. It is also true 
 
 that a small piece of phosphorus 
 
 placed in the dark centre will not 
 
 burn ; so there can be no oxygen pres- 
 
 ent. Outside of this central cone 
 
 there is a luminous zone of some 
 
 thickness where the oxygen is unit- 
 
 ing with the escaping gases ; but, as 
 
 oxygen cannot enter into this por- Fig. n. 
 
 tion of the flame in excess, the carbon, separated from the glowing 
 
 gases* by reason of the high temperature of the flame, is not com- 
 
 pletely burned, but is only heated to a white heat. Oxygen is pres- 
 
 ent in excess on the outer surface of the luminous zone, and therefore 
 
 combustion is most energetic in this division of the flame, so that 
 
 an enveloping mantle, which is scarcely visible, results ; this is the 
 
 hottest part of the flame. The glowing carbon which causes a 
 
 flame to become luminous is produced by the breaking down of 
 
 the heated hydrocarbons present in the gas, a change similar to 
 
 that undergone by acetylene taking place : 
 
 From this it follows that if the temperature of the flame can be 
 lowered to such a point as to prevent this decomposition, the flame 
 will become non-luminous ; such an alteration can be brought about 
 by diluting the gases (before burning) 69 with some indifferent sub- 
 stance, such as carbon dioxide, and the same result can be accom- 
 plished by providing a supply of air to the illuminating gas before 
 the vent at which the flame is lighted is reached. The Bunsen 
 burner attains this end by causing the gas which is escaping from 
 a small central opening to traverse a wider brass tube before igni- 
 
 * See ethylene and acetylene. It seems probable that ethylene breaks down 
 into acetylene and methane, that acetylene condenses to form more complicated 
 hydrocarbons, that the latter finally yield carbon and hydrogen. Methane, at 
 high temperatures, also yields carbon and hydrogen. 
 
284 THE FLAME. 
 
 tion. At the bottom of this brass tube two holes are pierced, allow- 
 ing the entrance of a limited supply of air. This air mingles with 
 the escaping gases, and thus provides for complete combustion 
 before the decomposition of the hydrocarbons takes place. Un- 
 doubtedly the non-luminous character of the Bunsen flame is also, 
 in part, brought about by the addition of the indifferent gas nitro- 
 gen, which must necessarily enter the burner in company with 
 oxygen. If the gases composing a flame can be cooled below their 
 kindling temperature, the flame will be extinguished. 
 
 From what has been said regarding the formation of a flame it 
 follows that it is a matter of indifference which of the two gases 
 uniting to form the flame is entering, and which forms the surround- 
 ing medium, for the phenomena are caused solely by the union 
 of the two. The terms " combustible " and " a supporter of com- 
 bustion," as applied to gases, are therefore used simply because it 
 is more usual to see gases burning in oxygen or air than it is to 
 see oxygen or air burning in other gases. Of course, the phenom- 
 ena attendant upon union with oxygen also appear with other 
 gases which (chlorine, for example) act like oxygen. 
 
CAIIBON TETRACHLORIDE. 285 
 
 CHAPTER XXXIX. 
 
 THE COMPOUNDS OP CARBON WITH CHLORINE, WITH CHLO- 
 RINE AND OXYGEN, WITH OXYGEN, AND WITH SULPHUR. 
 
 Carbon dioxide ; formula, C0 2 ; specific gravity, air = 1, is 1.529, 
 H 2 = 2, is 44 ; molecular weight, 44 ; one c.c. of the gas at and 
 .76 m. weighs .001986 gram. Carbon monoxide ; formula, CO ; 
 specific gravity, air = 1, is .96744, H 2 = 2, is 27.86 ; molecular 
 weight, 28 ; 1 c.c. of the gas at and .76 m. weighs .0012511 
 gram. 
 
 THE only chloride of carbon which need be mentioned in this 
 work is the tetrachloride, C C1 4 ; some more complicated chlorine 
 derivatives of carbon chains are known, but a work on organic 
 chemistry must be consulted in regard to their properties. Carbon 
 tetrachloride is derived from methane by replacing all of the hydro- 
 gen atoms with chlorine ; and it can be prepared, as we have seen, 
 from the latter substance by the action of chlorine ; it is not, how- 
 ever, practically expedient to commence with methane in order to 
 produce the tetrachloride, because methin chloride (C H C1 3 , chloro- 
 form) is easily procured by other means ; and then, beginning with 
 this chlorinated methane, we can, by passing chlorine into the boil- 
 ing liquid, finally substitute the remaining hydrogen atom : 
 
 Carbon tetrachloride is a colorless liquid which boils at 76.5, 
 and which, unlike most of the chlorides of the not-metals, is but 
 slowly decomposed by cold water ; on warming with an excess of 
 water it is, however, readily converted into carbonic and hydro- 
 chloric acids ; but carbonic acid, like all acids whose anhydrides are 
 gases, at once breaks down into its anhydride and water, so that, 
 although we may consider orthocarbonic acid to be the first result of 
 the reaction, carbon dioxide is the only tangible carbon compound 
 produced : * 
 
 * Organic derivatives of orthocarbonic acid; namely orthocarbonic acid, 
 
 rOC 2 H 5 
 
 in which all of the hydrogen atoms are replaced by ethyl, C 
 
286 CARBON MONOXIDE. 
 
 r ci + HOH r OH 
 
 Cl HOH 
 
 Cl + HOH OH 
 
 2. C (OH ) 4 = CO (OH ) 2 + Ho 0, and 
 
 3. CO (OH ) 2 = C0 2 + H 2 0. 
 
 The action of alkalies differs from that of water, for when carbon 
 tetrachloride is treated with potassium or sodium hydroxide, the 
 stable potassium or sodium carbonate is produced. 
 
 Carbontetrabromide, C Br 4 , and tetra-iodide, C I 4 , are also known ; 
 the former is a solid, which melts at 92, and boils at 189. 5 ; the 
 latter is a solid, which breaks down into carbon and iodine when 
 heated. 
 
 The two important oxides of carbon are carbon 'monoxide, CO, 
 and carbon dioxide, C0 2 ; only the latter acts like the anhydride of 
 an acid.* 
 
 Carbon dioxide was formerly supposed to be the only oxide of 
 carbon ; for carbon monoxide, even as recently as the beginning of 
 this century, was believed to be identical with hydrogen, or, at least, 
 to contain hydrogen. Woodhouse, of Philadelphia, in the year J 800, 
 first proved that the combustible gas obtained by reducing metallic 
 oxides with charcoal was not hydrogen, and demonstrated that it 
 contained carbon; but carbon monoxide was not recognized as a 
 combustible oxide of carbon until after the year 1802. | 
 
 Carbon monoxide is produced by the incomplete combustion 
 and also by the dry distillation of bituminous coal and of organic 
 matter ; for this reason it occurs in illuminating gas. Carbon mon- 
 oxide is likewise always formed when reducible metallic oxides, such 
 as those of iron or zinc, are heated with charcoal : 
 
 ZnO + C = Zn + CO. 
 
 The heating of metallic oxides with charcoal is a general method 
 of preparing metals from their ores. These reactions are, how- 
 ever, like many others, reversible, so that carbon dioxide is re- 
 duced to carbon monoxide by metals such as iron or zinc, when 
 
 * A few cases in which carbon monoxide exhibits acidic properties are 
 known. For instance, it unites with solid caustic potash when heated, form- 
 ing potassium formate : CO + KOH = CHO 2 K. 
 
 t In 1801 Cruikshank demonstrated that carbon monoxide is an oxide of 
 carbon. 
 
CARBON MONOXIDE. 287 
 
 these are heated to a high temperature. At 1300 carbon monoxide 
 is partially decomposed into carbon dioxide and carbon. When 
 steam is passed over red-hot charcoal, carbon monoxide and hydro- 
 gen are produced : 
 
 H 2 + C=CO + 2H, 
 
 and the mixture of combustible gases so obtained, after being 
 passed through volatile hydro-carbons, is quite extensively used as 
 illuminating gas.* Carbon dioxide can also enter into a similar 
 reaction ; for, when it is passed through a layer of hot coal, charcoal, 
 or coke, carbon monoxide results, as is shown by the following 
 equation : 70 
 
 C0 2 + C=2CO. 
 
 It is for this reason that the carbon dioxide, produced by the 
 free combustion of the coal just above the grate in a stove, is 
 changed to carbon monoxide by passing over the hot coals above. 
 In many cases carbon monoxide acts as a reducing agent ; for 
 instance, the gas passed over red-hot ferric oxide reduces the latter 
 to metallic iron : 
 
 Fe 2 3 + 3 CO = 2 Fe +3 C0 2 . 
 
 In order to prepare carbon monoxide for laboratory use advan- 
 tage is taken of the decomposition of oxalic acid by heat or by 
 means of concentrated sulphuric acid.f The acid breaks down 
 as follows : 
 
 CO OH co :OH 
 
 = rwvTT r,C 2 4 H 2 = C0 2 + CO + H 2 0.t 
 CO OH 
 
 The sulphuric acid, which is added, assists the operation by reason 
 of its great tendency to take up water. 71 The carbon dioxide, which 
 is formed simultaneously with carbon monoxide, can be removed 
 by passing the mixture of gases through a solution of potassium 
 hydroxide, by which means potassium carbonate is formed, while 
 the carbon monoxide can be collected over water. 
 
 Carbon monoxide is a colorless gas which, when it is pure, has 
 scarcely any odor. Its specific gravity, air = 1 , is .96744. Carbon 
 
 * So-called water gas. 
 
 t This reaction is common to a number of other dibasic organic acids, 
 t Oxalic acid can be derived from ethane by replacing four of the hydro- 
 gen atoms in one molecule by oxygen and the remaining two by hydroxyl. 
 
288 CARBON MONOXIDE; PROPERTIES. 
 
 monoxide is one of the gases which is with difficulty condensed 
 to a liquid, it does not become fluid at 136 and under 150 atmos- 
 pheres pressure; however, the application of a still greater cold 
 changes it to a colorless liquid which boils at 190 under 760 m. m. 
 pressure, and which becomes solid at 207. One volume of water 
 dissolves about .023 of a volume of carbon monoxide ; the gas is, how- 
 ever, quite soluble in a hydrochloric or an ammoniacal solution of 
 cuprous chloride. 
 
 Carbon monoxide burns readily in oxygen or in air ; the product 
 of the combustion is carbon dioxide : 
 
 The pale blue flames observed above an anthracite coal fire are 
 caused by carbon monoxide. 
 
 The gas acts as a poison ; it replaces the oxygen which is chemi- 
 cally combined in the blood by an equal volume of carbon monox- 
 ide, each molecule of carbon monoxide must, therefore, take the 
 place of a molecule of oxygen in oxy haemoglobin ; the oxidizing 
 powers of the blood are thereby destroyed, for, as carbon monoxide 
 forms a more stable compound with haemoglobin than does oxygen, 
 it is obvious that, once the carbon monoxide-haemoglobin is formed, 
 this compound cannot ,be broken up by oxygen. Blood which has 
 been saturated with carbon monoxide retains its red color on ex- 
 posure to the air for a longer time than that which has been 
 oxygenated.* 
 
 Carbon monoxide, when mixed with chlorine and placed in the 
 sunlight, unites directly with that element to form carbonyl chloride 
 (phosgen): 
 
 CO + 2C1 = COC1 2 ; 
 
 this reaction being exactly like the similar one observed in the case 
 of sulphur dioxide (see page 138). Carbonyl chloride, at ordinary 
 temperatures, is a colorless gas with a most peculiar, penetrating 
 odor. By means of snow and salt it can be condensed to a liquid 
 which boils at 8. Water readily decomposes carbonyl chloride, 
 forming hydrochloric acid and carbon dioxide. The reaction can be 
 considered as taking place in two phases : 
 
 * Carbon monoxide can readily be detected in the blood by means of the 
 peculiarity of the absorption spectrum of blood saturated with the gas. 
 
CARBON DIOXIDE; OCCURRENCE/ 289 
 
 / 
 
 f Cl + HOH ( OH / 
 
 1 C-}0 = QJO -f/^ H Cl. 
 
 ( Cl + HOH ( OH 
 
 f OH 
 
 (OH 
 
 The substance reacts in a similar way with ammonia, forming car- 
 bonyl diamide (urea) and hydrochloric acid : * 
 
 Cl NH 8 ( NH 2 
 
 + = C-3 +2HC1. 
 
 Cl NH 3 ( NH 2 
 
 A study of these reactions shows us that carbonic acid, H 2 C0 3 , can 
 be considered as carbonyl chloride in which two chlorine atoms in 
 each molecule have been replaced by hydroxyl groups ; the acid itself, 
 however, does not exist ; it is only known as its anhydride C0 2 . 
 From the above it is also evident that urea is carbonic acid in which 
 all hydroxyl groups have been replaced by NH 2 . 
 
 Carbon dioxide is of far greater importance than carbon monox- 
 ide. Its occurrence in the atmosphere and the manner and sources 
 of its production were discussed on pages 166 and 167. 
 
 Phenomena by which carbon dioxide are produced were known 
 in the earliest times, but the gas itself escaped observation. It was 
 not until the close of the sixteenth century that a peculiar gas, 
 which we now know to be carbon dioxide, was observed escaping 
 from some mineral waters ; Van Helmont (1577-1644) first distin- 
 guished this gas from others and gave to it the name of gas sylvestre; 
 he showed that this substance was produced by the action of acids 
 on alkalies or lime, by the burning of coals, and, in addition, was 
 also formed during the processes of fermentation. Black, in 1757, 
 showed the difference between the so-called caustic alkalies (now 
 known as hydroxides) and mild alkalies (carbonates), and found 
 that a peculiar kind of air (carbon dioxide), which he called fixed 
 air, was expelled from the latter by the addition of acids. Lavoisier 
 first explained the true nature of carbon dioxide, and gave to it the 
 name of acide carbonique. 
 
 Pure carbon dioxide can best be prepared by the addition of an 
 acid to some carbonate : 
 
 * The monovalent group NH 2 is called the amido group ; see page 192. 
 
290 CARBON DIOXIDE; PROPERTIES. 
 
 DHa, C0 3 + 2 H Cl = 2 Na Cl + H, -f C0 2 , 
 Ca C0 3 + 2 H Cl = Ca C1 2 + Ho + C0 2 , 
 Na^ C0 3 + H 2 S0 4 = Na 2 S0 4 -f H 2 + C0 2 . 
 
 Because carbonic acid is one of the weak acids, and because it so 
 readily breaks down into water and gaseous carbon dioxide (see 
 page 285), it follows that almost any other acid will liberate carbon 
 dioxide from the carbonates, so that the general formula : 
 
 M 2 C0 3 + 2 H X = 2 M X + H 2 + C0 2 , 
 
 (where M represents a univalent metal, and H X a monobasic acid) 
 will hold good with very few exceptions.* The most convenient 
 method of preparing carbon dioxide for laboratory use, is by the 
 action of dilute hydrochloric acid on marble (carbonate of calcium). 72 
 Carbon dioxide is a colorless gas, which neither burns nor sup- 
 ports combustion. It has a specific gravity of 1.529, air = 1, or 44, 
 H 2 = 2 ; its molecular weight is therefore 44, and its formula C0 2 . 
 This formula is further sustained by the fact that there is no change 
 in the volume of the gas when carbon burns in pure oxygen, so that 
 each molecule of carbon dioxide must contain one molecule, or two 
 atoms, of oxygen. Because carbon dioxide has such a high specific 
 gravity, it can be poured downward from any vessel containing it, 
 and it is for this reason that carbon dioxide collects at the bottom 
 of wells and mines into which the gas is escaping. 73 Cold and pres- 
 sure combined condense carbon dioxide to a liquid which boils at 
 78. 2 ;f the vapor tension of fluid carbon dioxide is 36 atmospheres 
 at 0, and 73 atmospheres at 30; the critical point is 30. 9 ; above this 
 temperature no pressure can convert the gas into a fluid. When 
 carbon dioxide rapidly evaporates in a vacuum, the temperature 
 sinks to 97 ; the liquid, when allowed to escape from a small 
 opening, condenses to a white, snow-like mass, the temperature of 
 which, at atmospheric pressure, is 78. Liquid carbon dioxide is 
 colorless, and has a specific gravity of .995 at 10. Liquid carbon 
 dioxide is extensively used in commercial operations ; for instance, 
 
 * Some few acids, like hydrocyanic acid, are unable to decompose carbon- 
 ates, while some few carbonates which occur as minerals (for instance, dolo- 
 mite) are not readily attacked by dilute acids. Quite a number of organic 
 substances which act like acids are unable to decompose carbonates. 
 
 t Doubtful. The temperature of melting carbon dioxide, mixed with 
 ether, is 80 (Dewar and Fleming; Philosph. Mag.; 34, 329). 
 
CARBONATES, PRIMARY. 291 
 
 in the manufacture of soda water, in fire extinguishers, and in oper- 
 ations where the great pressure exerted by it can be used. It is 
 transported in thick-walled steel tubes. 
 
 Carbon dioxide is the anhydride of carbonic acid, but the latter 
 substance is extremely unstable. It is probably formed as a white 
 mass when the pressure is suddenly removed from carbon dioxide 
 which, in the presence of water, has been nearly condensed to a 
 liquid at 0. There is no probability that water which is saturated 
 with carbon dioxide at ordinary temperatures contains the acid 
 H 2 C0 3 as such, for the solution behaves physically like an ordinary 
 gas solution, and not like that of an acid ; on the other hand, the 
 carbonates of the pronounced metals are extremely stable sub- 
 stances. With a diminution of the metallic nature of the salt- 
 forming element, the -carbonates become less stable, and very weak 
 bases, like the oxide of aluminium or ferric oxide, cannot react with 
 carbonic acid at all. 
 
 The carbonates are all derived from a dibasic acid H 2 C0 3 .' The 
 secondary ' carbonates, M 2 C0 3 , are, as a rule, insoluble in water ; 
 only those of the alkali metals and of ammonium * dissolve ; the 
 other carbonates can therefore, be obtained from these .by precipita- 
 tion with the soluble salt of some other metal, for example : 
 
 Na* C0 3 + Ba C1 2 = Ba C0 3 + 2 Na Cl. 
 
 Soluble. Soluble. Insoluble. Soluble. 
 
 The carbonates of the alkali metals can be fused without change ; 
 all other carbonates are more or less readily decomposed into car- 
 bon dioxide and the metallic oxide by heating, thus : 
 
 CaC0 3 = CaO + C0 2 , 
 
 and this decomposition takes place the more readily, the less basic 
 the metallic oxide is, so that many carbonates are even decomposed 
 on boiling with water. This increasing stability of the carbonates 
 with the increase in the metallic character of the salt-forming ele- 
 ment, is exactly parallel to the same gradation observed in the 
 chlorides of the elements of the phosphorus family (see page 181). 
 
 The primary carbonates, with the exception of those of the 
 alkalies, exist only in aqueous solution ; they can be obtained, where 
 their existence is possible, by treating a solution of a secondary 
 carbonate, or even a finely divided insoluble secondary carbonate 
 
 * Lithium carbonate is soluble with difficulty. 
 
292 CARBONATES; SECONDARY. 
 
 suspended in water, with carbon dioxide ; they are unstable and are 
 readily broken down' by heat : 
 
 2 Na HC0 3 = Na 2 C0 3 -f H 2 + C0 2 ; 
 
 where they exist they are soluble. The solution of calcium car- 
 bonate in temporary hard water is caused by the formation of the 
 primary calcium carbonate by means of the carbon dioxide con- 
 tained in the air or added to the water by decaying substances : 
 
 CaC0 3 + H 2 C0 3 = Ca (HC0 3 ) 2 . 
 
 This soluble primary carbonate is decomposed when the water 
 evaporates * or when it is heated. The temporary hard waters for 
 this reason deposit their calcium carbonate as a white coating on 
 the walls of the kettle in which they are boiled. 
 
 Secondary carbonates, when soluble, have a strongly alkaline 
 reaction, the primary ones are neutral. f 
 
 The following table gives a few of the most important naturally 
 occurring carbonates : 
 
 Calcium carbonate ) Massive varieties; chalk, limestone, marble. 
 
 (CaCO) > Crystallized varieties; calcite (Iceland spar), arra- 
 
 gonite. 
 
 Calcium and magnesium carbonate (Ca, Mg), CO 3 , dolomite. 
 Ferrous carbonate; Fe CO 3 , siderite. 
 Barium carbonate; Ba CO 3 , witherite. 
 Strontium carbonate ; Sr CO 3 , strontianite. 
 
 The above carbonates are frequently found as isomorphous mix- 
 tures. The carbonates of lead, zinc, and manganese are also found, 
 as well as basic carbonates of copper, bismuth, and zinc. Carbonates 
 of sodium with more or less water of crystallization occur as soda 
 ( Na 2 C0 3 + 10 H 2 0) and trona ( Na 4 H 2 [C0 3 ] 3 3 H 2 0) . t 
 
 The compound of carbon and sulphur which corresponds to C0 2 
 is CS 2 , carbon disulphide. This liquid can be formed by heating 
 carbon in sulphur vapor, so that the method of its production cor- 
 responds to that of carbon dioxide. Carbon disulphide is a colorless, 
 mobile liquid, which, when pure, has a pleasant ethereal odor. Its 
 specific gravity is 1.29 ; the specific gravity of its vapor is 2.626 
 (air = 1) ; it boils at 48, and is very little soluble in water. When 
 
 * Formation of stalactites. 
 
 t They have no effect on litmus or turmeric paper, but do have an alkaline 
 reaction toward rosolic acid. 
 
 t (2 Na HC0 3 + Na 2 CO 3 + 3 II, O. ) 
 
THIOCARBONATES. 293 
 
 gently warmed with the alkaline sulphides, carbon disulphide is 
 dissolved while the salts of sulpho-acids are formed : 
 
 These salts of trithiocarbonic acid (sulpho-dithio carbonic acid) * 
 correspond to those of carbonic acid, the oxygen atoms in the latter 
 having been replaced by sulphur. When acids are added to its 
 salts, sulpho-dithio carbonic acid, H 2 CS 3 , separates as an oil: 
 
 K 2 CS 3 + 2HC1==H 2 CS 3 + 2KC1, 
 
 but the thio acid so produced, although it is not as unstable as 
 the corresponding oxy-acid, nevertheless, gradually breaks down as 
 
 follows : 
 
 H 2 CS 3 =H 2 S + CS 2 , 
 
 just as carbonic acid decomposes into water and carbon dioxide. 
 These reactions of carbon disulphide remind us forcibly of the simi- 
 lar ones encountered with the sulphides of arsenic and antimony. 
 (See pages 245, 255.) A compound of carbon, oxygen, and sulphur, 
 having the formula COS (lying between the dioxide and sul- 
 phide of carbon), as well as acids derived from it, are also known ; 
 and several sulphides of carbon, differing from carbon di-sulphide 
 (for example, C 2 S 3 ) have also been described. 
 
 * The name trithiocarbonic acid is derived from Qftov, sulphur, and the 
 name thio-acids is frequently employed for sulpho-acids and thio-compounds 
 for sulpho-eompounds. An endeavor is made to establish the following dis- 
 tinction: where sulphur is attached to carbon only, it is called sulpho, where 
 it is attached to carbon on the one hand and a metal, or a group of elements 
 
 < SH ^ CL 
 
 acting like a metal, on thej>ther, it is called thio; thus: C j O / is( dttkiit 
 
 ( SH 
 
 carbonic acid, and C ] S_ is sulphothioc&vltomc acid. This nomenclature 
 I OH * B' 
 
 is, however, logically carried out only h tne compoundsjjkf carbon. 
 
 * 
 
 sjjkf c 
 
294 CYANOGEN. 
 
 CHAPTER XL. 
 
 COMPOUNDS OF CARBON WITH NITROGEN, WITH NITROGEN 
 
 AND HYDROGEN, AND WITH NITROGEN, OXYGEN, 
 
 AND HYDROGEN. 
 
 ONLY a very few of the more important of these compounds can 
 be strictly considered as belonging to the realm of inorganic chem- 
 istry, and these only will be briefly considered in this work. The 
 most prominent of the substances to be discussed are derivatives of 
 the monovalent group of elements, cyanogen,* CN. This group can 
 be attached to other elements or groups of elements in two ways, 
 either through the element nitrogen, by which means substances 
 having the general structural formula MNC are formed ; or through 
 carbon, the general structural formula of the latter class of com- 
 pounds being MCN ; in the first case isocyanides are formed, in the 
 second true cyanides (also called nitriles). Representatives of both 
 classes of compounds are known, f 
 
 All nitrogen-bearing organic compounds, or, indeed, the nitro- 
 genous coals derived from these, yield the cyanide of sodium or of 
 potassium when heated with one or the other of those metals ; the 
 cyanide of potassium is also formed wh.en^coal which contains nitro- 
 gen compounds is heated with common pbtaSff ; by this process the 
 greater quantity of the cyanide of potassium which finds commer- 
 cial %application is prepared. Free cyanogen can be formed by 
 heating the cyanide of mercury to a dark red heat : - 
 
 The group CN is as incapable of individual existence as is 
 methyl, CH 3 ; two of the radicles therefore unite to form dicya- 
 nogen, CN CN, just as methyl unites to form dimethyl. | 
 
 * From x i ' avo '>i blue. 
 
 t The pupil must refer to some larger work on organic chemistry for the 
 reactions which characterize the isocyanides. 
 
 J Cyanogen can be considered as methyl in which three atoms of hydro- 
 gen are replaced by one atom of trivalent nitrogen; viz., 
 
HYDROCYANIC ACID. 295 
 
 Cyanogen is a colorless gas, extremely poisonous, with an irri- 
 tating odor. Its specific gravity is 1.804, air = 1, or 51.9, H = 2, 
 while the molecular weight of (CN ) 2 is 52. The gas is quite solu- 
 ble in water, and is combustible, burning with a characteristic 
 purple flame. It liquefies at 20.7 and becomes solid at 34. 4. 
 Aqueous solutions of cyanogen gradually decompose. 
 
 Chemically, cyanogen greatly resembles the halogens ; the mono- 
 valent group GN", when united with hydrogen, forms hydrocyanic 
 acid, just as chlorine, similarly united, forms hydrochloric acid; 
 furthermore, hydrogen, sodium, or potassium will unite directly 
 with cyanogen to form hydrocyanic acid, sodium cyanide, or potas- 
 sium cyanide, just as the same elements will unite with chlorine to 
 form hydrochloric acid, sodium, or potassium chloride. The group 
 CNj cyanogen, is therefore a compound radicle acting like a not- 
 metal, and is chemically the opposite of the metal-like radicle 
 ammonium. 
 
 Hydrocyanic acid * is prepared, according to the general method, 
 by adding an acid to a cyanide : 
 
 2 KCN + H 2 S0 4 = 2 HCN + K 2 S0 4 .f 
 
 Or, better, by distilling commercial potassium ferrocyanide with 
 diluted sulphuric acid (concentrated acid cannot be used for this 
 purpose, as it generates carbon monoxide). When pure, hydrocy- 
 anic acid is a colorless, mobile liquid, which boils at 26 and melts 
 at 14 ; it is without acid reaction toward litmus, and has a 
 peculiar odor, somewhat resembling that of bitter almonds. It is 
 intensely poisonous even when inhaled in small quantities ; one 
 drop placed on the tongue of a dog will cause instant death ; the 
 poison can also act through contact with abrasions of the skin. 
 The acid mixes with water in*all proportions, and solutions of vary- 
 ing strength form the commercial prussic acid ; the aqueous solution 
 decomposes on standing. The acid was first prepared in a pure 
 state by Gay Lussac, in 1811. 
 
 Compare the preparation of cyanogen with the preparation of oxygen by heat- 
 ing mercuric oxide. 
 
 * Also called prussic acid because it is the source of prussian blue. 
 
 t The usual method is to decompose potassium ferrocyanide (see latter) 
 with sulphuric acid. 
 
296 POTASSIUM FERROCYANIDE. 
 
 Two structural formulae are possible for hydrocyanic acid; 
 namely, C = N H and H C = N ; in the first one of these the 
 hydrogen atom is attached to nitrogen, in the second to carbon ; 
 organic derivatives of both forms are known ; the cyanides of the 
 metals are most probably derivatives of the first form, CNH.* 
 
 Hydrocyanic acid is a very weak acid, the cyanides of the 
 alkali metals and of barium, calcium, or strontium are even decom- 
 posed by moist carbon dioxide, so that these substances, when in 
 contact with ' the air, emit the odor of hydrocyanic acid. The 
 cyanides of most of the heavy metals are insoluble, an exception 
 to this rule is the cyanide of mercury ; some of these cyanides are 
 not readily decomposed by acids ; t on the other hand, the alkaline 
 cyanides, which are so readily decomposed by weak acids, are ex- 
 tremely stable when heated ; they can even be fused without under- 
 going a chemical change. Almost all cyanides which are soluble in 
 water are converted into so-called double cyanides when treated 
 with the cyanides of the alkalies. Many of these double cyanides 
 are stable, crystalline bodies which have the nature of chemical 
 compounds Q for example, ferrous cyanide when brought in contact 
 with potassium cyanide forms a double cyanide of the formula 
 
 Fe(CN) a + 4KCN = K 4 (CN) 6 Fe. 
 
 Ferrous cyanide + Potassium cyanide = Potassium ferrocyanide. 
 
 Potassium ferrocyanide bears no resemblance either to potassium 
 cyanide or to ferrous cyanide ; it is, indeed, the salt of a tolerably 
 strong crystalline acid, hydroferrocyanic acid, H 4 (CN) 6 Fe. This 
 acid can be prepared by the addition of strong hydrochloric acid to 
 potassium ferrocyanide; the ferric salt of this acid, ferric ferrous 
 cyanide, can be obtained from potassium ferrocyanide by adding a 
 ferric salt : 
 
 3 K 4 (CN ) 6 Fe + 4 Fe C1 3 = Fe 4 [(CN ) 6 Fe] 3 + 12 K Cl. 
 This substance is the insoluble blue dye known as prussian blue. 
 
 Ferric cyanide is also able to form a double salt with potassium 
 cyanide : 
 
 * The experiments which seem to indicate that cyanides are derived from 
 both forms of hydrocyanic acid can easily be explained on the assumption 
 that they are all derived from the first form. Compare Nef ; Liebig's Anna- 
 len; 270, 329. 
 
 t Cyanides of mercury, silver, and gold. 
 
POTASSIUM FERRICYANIDE ; CYANIC ACID. 297 
 
 Fe (CN ) 3 + 3 KCN = K 3 (CN ) 6 Fe, 
 
 and this substance, on addition of a ferrous salt, also forms an in- 
 soluble blue compound, Turnbull's blue.* The cyanides of other 
 metals which are chemically closely allied to iron can form similar 
 double cyanides ; a larger work must be consulted in regard to the 
 properties of those compounds. 
 
 When the cyanide of potassium is oxidized,t it changes to the 
 cyanate of potassium, CNOK. This oxidation is easily explained if 
 we consider hydrocyanic acid to be analogous to carbon monoxide : 
 
 C=N H, C=0, 
 
 Hydrocyanic acid, Carbon monoxide ; 
 
 for then it would naturally follow that the former would be oxidized 
 as readily as is the latter : 
 
 C=N H + = C | ^ T ~ H and C=0 + = C j 
 
 
 It seems highly probable, however, that the salts of this cyanic 
 acid assume another form, namely, one in which the metallic atom 
 is attached to oxygen : t 
 
 j = N H , H J^N 
 
 1 = 'K , = |-0. 
 
 K 
 
 Such a difference between the structure of the free acid and of 
 the salts derived from it seems not unlikely, if the great tendency 
 which is displayed by potassium or sodium to unite with oxygen is 
 remembered. The acid having the formula CONH is termed isocy- 
 anic acid, and the one with the structure CNOH, normal cyanic acid. 
 Ordinary cyanic acid is at present known only in one form ; it is a 
 very mobile, volatile liquid, which is stable only below 0, its odor 
 resembles that of sulphur dioxide. A polymeric form of cyanic 
 acid (CNOH ) 3 , cyanuric acid, is also known, it is a solid substance. 
 
 * Probably identical with prussian blue. 
 
 t This oxidation even takes place upon exposing potassium cyanide to the 
 air. 
 
 J Organic derivatives obtained from the cyanates undoubtedly are derived 
 from cyanic acid of the first formula, CONH ; but this fact does not seem to 
 prove that the metal in the cyanates is attached to nitrogen. 
 
 It cannot, as yet, be absolutely decided which of the two forms, CNOH 
 or CONH is represented by ordinary cyanic acid. Organic derivatives of 
 both acids (those in which hydrogen is replaced by ethyl, etc.) are known, 
 
298 UREA; CARBAMIC ACID. 
 
 Ammonium cyanate, on standing, or more rapidly on heating, 
 changes into urea, a substance of the greatest physiological 
 importance : - CNONH 4 = CO ( NH 2 ), 
 
 Urea can be considered as carbonic acid in which both hydroxyls 
 have been replaced by the amido group, NH 2 . 
 
 (OH ( NH 2 
 
 (OH ( NH 2 
 
 Carbonic acid. Urea. 
 
 A compound derived from carbonic acid, but in which but one 
 hydroxyl group is replaced by the amido group (NH 2 ) is also 
 known ; this substance is termed carbamic acid, C0 2 HNH 2 : 
 
 (0 ro 
 
 C < OH carbonic acid and C -j OH carbamic acid. 
 ( OH ( NH 2 
 
 The ammonium salt of this acid, 
 
 ^-\ 
 
 is found in commercial carbonate of ammonium ; on heat- 
 ing to 130 it changes into urea : 
 
 CO 
 
 NH 2 
 
 Another interesting method by which urea can be formed is by 
 the action of ammonia on carbonyl chloride : 
 
 ( Cl NH 3 ( NH 2 
 
 (VO+ = CJ + 2HC1, 
 
 ( Cl NH 8 ( NH 2 
 
 for this change is analogous to the reaction which takes place when 
 water acts on the same substance : 
 
 ( Cl + HOH ( OH 
 
 C^O =(OO+ 2 HC1 (see page 289). 
 
 ( Cl + HOH ( OH 
 
 The oxygen atom in a formula weight of urea can be replaced by 
 the divalent group = NH, and the resulting compound (guanidine), 
 which has the formula : 
 
 although those of normal cyanic acid (CNOH ) have never been obtained in a 
 
 I = O 
 pure state. It seems not unlikely that the free acid has the structure C ~ 
 
 while the salts are derived from C 
 
GUANIDINE ; CYANAMIDE. 
 
 299 
 
 (HH,, 
 
 differs but little from urea in its character. 
 
 The foregoing compounds are extremely interesting from a theo- 
 retical standpoint, because they illustrate the close chemical resem- 
 blance between the amido and hydroxyl groups ; for the former can 
 take part in the formation of compounds which are similar in char- 
 acter to the more familiar ones derived from the latter, and ammonia 
 can, therefore, play much the same role as water in the reactions 
 entered into by the carbon compounds which have just been dis- 
 cussed. This similarity between nitrogen and oxygen compounds 
 serves forcibly to illustrate the fact that no element is isolated in 
 properties and chemical character, and it should be the endeavor of 
 every student of chemistry to detect and understand the resem- 
 blances which are in reality present. 
 
 The connection between carbonic acid and the above compounds 
 is shown by the following table : 
 
 2, 
 
 , ( OH 
 
 b. (OH c. 
 
 (OH 
 
 d. ( OH c. r NH 2 
 
 1 (0 OH 
 
 L " 1 OH 
 
 r J OH n 
 
 Ut j OH 
 
 j OH 
 J NH 
 
 C ' J NHo C ' j NH 2 
 
 I OH 
 
 ^NH 2 
 
 CKH 
 
 2 I NH 2 I NH 2 
 
 
 less water give, 
 
 
 less ammonia give, 
 
 '( 
 
 (0 
 
 
 (0 ( NH 
 
 C < OH 
 
 C ] OH 
 
 
 C ] NH 2 C ] NH 2 
 
 (OH 
 
 ( ^H 2 
 
 
 ( J^H 2 ( NH 2 
 
 carbonic acid 
 
 , carbamic acid, 
 
 urea, guanidine, 
 
 H 2 O 
 
 -H 2 N 
 
 H 3 
 
 -H 2 -NH 3 -NI 
 
 o C (0 
 
 c io ^I^H ^|o "iNHo ^INH 
 
 carbon dioxide, cyanic acid, carbon dioxide, cyanamide, cyanic acid, cyanamide. 
 
 The hypothetical compounds &, c, cZ, e, are supposed to be derived from the 
 hypothetical normal carbonic acid (a) by replacing hydroxyl groups by amido 
 groups (OH by NH 2 ) . In this table the hydroxyl group is considered as anal- 
 ogous to the amido group and of equal valence, and the group NH (imid 
 group) as analogous to a divalent oxygen atom. Of the compounds under (2), 
 carbonic acid and carbamic acid are known only in their derivatives. 
 
 * Cyanamide is a solid, melting at 40 and differing from urea by one 
 molecule of water. 
 
300 SILICON; PREPARATION. 
 
 CHAPTER XLI. 
 
 SILICON, THE COMPOUNDS OF SILICON WITH HYDROGEN, 
 
 AND WITH THE HALOGENS, THE OXIDE, AND 
 
 ACIDS OF SILICON. 
 
 Silicon; symbol, Si ; atomic weight, 28.4 ; specific gravity of solid 
 (graphitoidal ), 2.49. 
 
 SILICON never occurs in nature as the uncombined element, but 
 silicon compounds are among the most important and widely dis- 
 tributed constituents of the crust of the earth. The primitive crys- 
 talline rocks are in greater part either silicon dioxide or else salts 
 of the various silicic acids. It has been estimated that 27.2 per 
 cent of the globe (excluding the atmosphere) consists of silicon. 
 Despite the abundance of silicon compounds, the element itself 
 was not discovered until 1823, in which year it was isolated by 
 Berzelius. 
 
 Silicon is best prepared by the reduction of some of the halo- 
 gen compounds of the element by means of sodium or potassium ; 
 for instance, by passing the vapors of the tetrachloride of silicon 
 over heated sodium : - 
 
 SiCl 4 + 4Na=Si + 
 
 The element so prepared is an amorphous brown powder, which 
 does not conduct electricity ; it is readily ignited in the air, burning 
 to form silicon dioxide. Silicon is produced when potassium or 
 sodium fluosilicates * are fused with aluminium ; the aluminium, 
 uniting with the fluorine, in part forms aluminium fluoride, while 
 the unchanged, molten metal dissolves the silicon which is liberated ; 
 when the mass is cooled a portion of the silicon separates in needle- 
 shaped crystals.f Silicon crystallizes in grayish black, regular 
 octahedra which have a metallic lustre. The silicon which is 
 
 * K 2 Si F 6 or Na 2 Si F 6 , salts of fluosilicic acid, KU Si F 6 . 
 
 t Melted zinc or iron can also dissolve silicon ; when they cool, the silicon 
 separates in crystals ; as silicon is formed from silicon dioxide, carbon, and iron 
 at a high heat, it follows that ordinary pig iron must contain silicon. 
 
SILICON ; HYDRIDE, TETKACHLORIDE. 301 
 
 formed from molten iron or zinc may also have the appearance 
 of graphite, although the crystalline form is, apparently, the same 
 as that of the needle-like crystals ; graphitoidal silicon has a specific 
 gravity of 2.49 ; crystallized silicon conducts electricity readily and 
 cannot be ignited in the air. When heated to a high white heat, 
 silicon can be fused and even cast into sticks ; when cooled it solidi- 
 fies to form a mass which somewhat resembles a piece of pure 
 crystalline graphite. The amorphous and crystalline varieties of 
 silicon remind us of the similar forms displayed by carbon. 
 
 Only one compound of silicon and hydrogen, silicon hydride, 
 Si H 4 , is known. This substance bears a great resemblance to the 
 corresponding compound in the nitrogen family, namely, to phos- 
 phine, for it takes fire spontaneously when brought into the air, 
 and it is formed by the action of an acid on magnesium silicide ; 
 in principle this is the same method as that employed for the prepa- 
 ration of phosphine from calcium phosphide. The preparation of 
 pure silicon hydride is a difficult process. Silicon hydride is a 
 colorless gas which, when pure, does not take fire spontaneously, 
 but which has such a low kindling temperature that it can readily 
 be ignited by a warm glass rod ; the gas is liquefied at 11 by a 
 pressure of 50 atmospheres. Silicon hydride is, of course, readily 
 decomposed by chlorine or bromine, and it resembles hydrogen 
 sulphide and phosphine by producing precipitates with quite a 
 number of metallic salts. 
 
 The compounds of silicon with the halogens possess general for- 
 mulae similar to the corresponding compounds of carbon ; silico- 
 chloroform, Si H C1 3 , corresponding to ordinary chloroform, CHC1 3 , 
 and silicon tetrachloride, SiCl 4 , corresponding to CC1 4 , serve to 
 illustrate this resemblance. Silicon tetrachloride is a colorless, 
 mobile liquid which boils at 58, and which is energetically de- 
 composed by water. It has a most penetrating and irritating odor 
 which resembles that of phosphorus pentachloride. The halogen 
 compounds of silicon are all unstable bodies which are decomposed by 
 water to form silicic acid and the corresponding halhydric acid : 
 
 Si C1 4 + 3 H 2 = Si 3 H 2 + 4 H Cl . 
 
 The most important halogen compound of silicon is undoubtedly 
 the tetrafluoride, Si F 4 . This substance can readily be produced 
 by the action of hydrofluoric acid on silicon dioxide : 
 
302 SILICON AND TETKAFLUOUIDE. 
 
 Si 2 + 4 HF = Si F 4 + 2 H 2 . 
 
 In this reaction the silicon dioxide is a base, for it yields a salt and 
 water when brought in contact with an acid. Silicon tetrafluoride 
 is a colorless gas, which fumes strongly in the air, because when in 
 contact with water vapor it decomposes and forms silicic acid ; it 
 has a most penetrating and irritating odor, arid it reddens litmus 
 paper, even when dry. At a temperature of 105 and at 9 atmos- 
 pheres pressure silicon tetrafluoride is converted into a colorless 
 liquid, which becomes solid at 140. The specific gravity of the 
 vapor, air =1 , is 3.6 , a number which would agree with a calcu- 
 lated molecular weight of 103.6, so that according to this the mole- 
 cule of silicon tetrafluoride is SiF 4 . As silicon tetrafluoride is 
 readily produced by the action of hydrofluoric acid on silicon 
 dioxide, it follows that hydrofluoric acid will attack glass, for that 
 substance contains a large proportion of silicon dioxide. 74 Water 
 instantly decomposes silicon tetrafluoride, forming silicic acid and 
 fluosilicic acid : 
 
 2 SiF 4 + 3 H 2 = H 2 Si 3 + Si F 6 H 2 + 2 HF. 
 
 Fluosilicic acid is especially interesting because it shows the chem- 
 istry of fluorine to us in an entirely new light ; for, if we compare 
 the following two formulae : 
 
 (0 (F 2 
 
 Si-} OH and Si^F 2 H 
 (OH (F 2 H 
 
 Silicic acid and Fluosilicic acid, 
 
 we see that fluosilicic acid is constructed similarly to silicic acid, 
 but with this difference ; for each oxygen atom in silicic acid we 
 have two fluorine atoms in fluosilicic acid. Two fluorine atoms in 
 some chemical compounds are therefore able to take the place of 
 one oxygen atom without materially altering the chemical nature 
 of those compounds. Substances which are constructed in a manner 
 similar to fluosilicic acid are not infrequent ; but, as their resem- 
 blance to oxygen compounds is not generally so marked as in the 
 case under discussion, their true nature is often misunderstood and 
 concealed under the names of double salts (see aluminium).* The 
 
 * Hydro-ferrocyanic acid, H 4 Fe (CN) 6 , may be cited as an instance 
 where a complex compound containing hydrogen (and in which cyanogen has 
 taken the place of oxygen) can act as an acid. 
 
FLUOSILICIC ACID. 303 
 
 wider the range of our acquaintance with chemical compounds be- 
 comes, the more do we see that the most various substances, which 
 may or may not contain oxygen, can act as acids, provided only 
 they contain hydrogen attached to a not-metallic element or group 
 of elements, and which hydrogen can be replaced by metals to 
 form salts (see page 75). The reason for the salt formation is 
 found in the simple fact that the salts produced by the replace- 
 ment of the acid hydrogen possess less chemical energy than do the 
 acids themselves, so that the reactions by which these salts are 
 formed are exothermic. One consideration forces itself upon us 
 when we study the structural formula of fluosilicic acid, and that 
 is the impossibility of maintaining the theory of the constant 
 monovalence of fluorine ; for, unless we wish to take the untenable 
 position that fluosilicic acid is a compound formed of finished 
 molecules of silicon fluoride and hydrofluoric acid (is a so-called 
 " molecular compound "), we must look upon the fluorine atoms as 
 being divalent in fluosilicic acid : 
 
 Even if we call this acid a molecular compound (SiF 4 , 2 HF), such 
 a supposition does not help matters in the least ; for then we must 
 regard Si F 4 and HF as still having chemical affinity at their dis- 
 posal, which supposition is contrary to the theory that silicon is 
 only tetravalent and fluorine only monovalent, for then neither of 
 these compounds should be capable of further union after Si F 4 and 
 HF have been formed. So long, therefore, as our present theories 
 of valence are maintained, we must regard fluorine as being both 
 uni- and bivalent. This conclusion is strengthened by the fact 
 that the specific gravity of hydrofluoric acid shows that substance 
 to have a molecule corresponding to the formula H 2 F 2 . 
 
 Fluosilicic acid is known only in solution ; when evaporated be- 
 yond a certain concentration it breaks down into silicon tetrafluo- 
 ride and hydrofluoric acid, much as ordinary silicic acid does into 
 silicon dioxide and water.* A similar change takes place with the 
 fluosilicates ; for these salts, when heated, break down into silicon 
 
 * A hydrate of fluosilicic acid, H 2 SiF 6 , 2 H 2 O, is known; it is a crystal- 
 line body which melts at 19 and which is extremely hygroscopic. 
 
304 SILICON DIOXIDE. 
 
 tetrafluoride and the fluoride of the metal entering into the salt 
 formation : - Si F 6 K 2 = Si F 4 + 2 KF. 
 
 Almost all of the silicofluorides are soluble in water ; the fluosilicate 
 of potassium is, however, nearly insoluble ; and, as almost all potas- 
 sium salts are dissolved by water, it is evident that fluosilicic acid 
 is a very welcome reagent for the detection of potassium compounds 
 in solution. 
 
 Silicon forms but one well-known oxide, the dioxide Si 2 . This 
 substance occurs in three forms : crystalline, cryptocrystalline, 
 and amorphous. Crystallized silicon dioxide is dimorphous, being 
 found as quartz * and as tridymite.t The quartz crystals are often 
 colored more or less by impurities ; when the color so produced is 
 purple or bluish violet the crystal is called amethyst. The crystal- 
 lized variety of quartz frequently occurs in large masses, displaying 
 no crystalline faces, while smaller fragments of the mineral are 
 found as a constituent of the granitic rocks. The cryptocrystalline $ 
 varieties of quartz show the greatest diversity of color and appear- 
 ance ; they generally contain more or less water and are more read- 
 ily acted on by hydrofluoric acid than the crystallized varieties. 
 Examples of cryptocrystalline quartz are chalcedony, carnelian, 
 agate, onyx, and flint. Sea sand consists, for the most part, of 
 quartz finely ground by the action of the water. 
 
 Amorphous silicon dioxide can be prepared by the addition of an 
 acid to a soluble silicate : 
 
 i 3 + 2 H Cl = H 2 Si 3 + 2 Na Cl ; 
 Sodium silicate. Silicic acid. 
 
 and by then heating the silicic acid until all water is expelled ; the 
 silicon dioxide formed in this way is a white, impalpable powder 
 which can be readily dissolved in alkalies ; when heated to a high 
 white heat this variety of the dioxide becomes crystalline and can 
 then no longer be dissolved by cold alkalies. 
 
 The two simplest theoretical hydrates of silicon dioxide are ortho- 
 silicic acid, Si (0 H) 4 , and meta-silicic acid, Si 3 H 2 ; neither 
 
 * Hexagonal, tetartohedral ; combinations of pyramid and prism. 
 
 t Asymmetric. 
 
 \ Varieties of crystalline minerals in which the crystals are so small as 
 not to be detected by the eye are called cryptocrystalline. Such rocks are fre- 
 quently erroneously termed amorphous. 
 
DIALYSIS. 305 
 
 of these acids is known with certainty, but a solution which prob- 
 ably contains ortho-silicic acid can be obtained by the following 
 means: by adding cold dilute hydrochloric acid to a very dilute 
 cold solution of sodium silicate the following reaction presumably 
 takes place : 
 
 i 3 + 2 H Cl = H 2 Si 3 + 2 Na Cl ; 
 
 the meta-silicic acid so formed then unites with water to form ortho- 
 silicic acid : 
 
 H 2 Si 3 + H 2 = H 4 Si 4 . 
 
 Under the conditions of the reaction there is no separation of insol- 
 uble silicic acid, as there is when the solutions are more concen- 
 trated, or'when they are heated. If the clear liquid containing 
 sodium chloride and silicic acid is put in a vessel the bottom of 
 which is formed of a membrane such as parchment, and this vessel 
 is then placed in pure water, the sodium chloride and excess of hy- 
 drochloric acid will pass out into the water, while the silicic acid 
 will remain behind in the solution. The process by which this sep- 
 aration takes place is osmosis, and the silicic acid is said to be sepa- 
 rated by dialysis; substances which are able to pass through such a 
 membrane are called crystalloids ; those which cannot pass through, 
 colloids. As the process of osmosis is one of extreme importance 
 in animal and plant life, a brief discussion of some of the principal 
 facts which have been learned regarding it may not be out of place 
 here.* 
 
 When a layer of water is carefully poured over any aqueous so- 
 lution, the two liquids will not remain in this condition, for diffu- 
 sion will take place just as it does between layers of different gases 
 (see page 34), so that the solution will begin to rise in a direction 
 contrary to the force of gravity and will finally completely mix with 
 the pure water, the motion only ceasing when the substance in solu- 
 tion is uniformly distributed throughout the mass of water. This 
 motion can be arrested by placing a septum between the water and 
 the solution ; and if this septum is of such a material as to allow 
 water to pass through, but not the dissolved substance, and if, fur- 
 thermore, the septum is in the shape of a cell which can be covered 
 by an air-tight cap which is so constructed that it can be connected 
 
 * See Ostwald, Outlines of General Chemistry (Walker), for a more com- 
 plete description of this topic. 
 
306 OSMOTIC PRESSURE. 
 
 with a manometer by means of a glass tube, an increase of pressure 
 will be observed in the interior of the cell, because the water will 
 force its way in, while the substance in solution cannot escape. 
 Now, a remarkable fact is observed in regard to this pressure, for if 
 the temperature is kept constant the pressure will be proportional 
 to the strength of the solution; thus, with a solution of sugar, the fol- 
 lowing pressures were observed, 
 
 A 1 per cent solution gave a pressure of 535 m. m. 
 n 2 " " u " " " " 1016 " " 
 
 u 4 tt u a a a a a 2082 " " 
 
 These observed pressures in millimeters are nearly in the pro- 
 portion of 1:2:4. This law of osmotic pressure, which is true of 
 all dissolved substances, is exactly like that regulating the pressure 
 of gases, for these are also proportional to the densities (i.e., con- 
 centrations). It has further been observed that temperature has 
 the same influence on osmotic pressure as it has on the pressure of 
 gases ; for the pressure increases proportionally to the absolute tem- 
 perature, and in the same ratio for all dissolved substances. The 
 increase in pressure for each degree of temperature is 5 | 5 ,* a frac- 
 tion identical with that obtained as an increase for each degree in the 
 pressure of gases which are kept at constant volume ; this relation for 
 osmotic pressure may therefore be expressed in the same way as it is 
 for gases. If we know the osmotic pressure ( P ) at 0, then at t it 
 will be P + P .00367 t = P ( 1 -f .00367 t) (see pages 172, 173). 
 The osmotic pressure of a substance in solution has the same value 
 as the pressure that substance would exert were it a gas occupying 
 the same volume as the solution. It seems reasonable to suppose, 
 therefore, as the laws which govern the pressures of gases also hold 
 good for osmotic pressure, that the substances which are contained 
 in solution are present in such solution in the same condition as that 
 in which they occur in gases, i.e., as the individual molecules. 
 
 The membranes through which the various fluids in living 
 organisms must find their way by osmosis act on the same princi- 
 ple as the septa which are artificially prepared ; and, as the same 
 increase of temperature causes a like increase of osmotic pressure 
 in all fluids, it follows that solutions which are in osmotic equilib- 
 rium between the contents of a living cell and the liquid with- 
 
 * Exactly .00367. 
 
ORTHO- AND META SILICATES. 307 
 
 out at any given temperature, say 0, are also in equilibrium at 
 38. 
 
 When the solution containing dialized silicic acid is evaporated, 
 the acid congeals to a gelatinous mass, which is then no longer 
 soluble in water ; and when this is separated and dried the remain- 
 ing amorphous powder has the formula, approximately, of H 2 Si 3 ; 
 this, when heated, loses water and forms amorphous silicon dioxide. 
 
 The silicates are either orthosilicates, M 4 Si 4 , or metasilicates, 
 M 2 Si 3 , or they are derived from more complicated silicic acids 
 which, according to the number of silicon atoms in their formula 
 weights, are called di-, tri-, or, in general, polysilicates. All of these 
 salts have numerous representatives in the mineral deposits of the 
 earth. The formulae of a few of these are given in the following 
 tabular statement : 
 Orthosilicates : f Olivin? Mgg g . ^ ^ g . ^ 
 
 I Garnet, Ca 3 Fe 2 (Si 4 ) 3 .* 
 
 Acid, Si J | Mica; the various forms of this mineral are complicated 
 
 orthosilicates. 
 
 f Wollastonite, Ca Si O 3 . 
 Leucite, KA1 (SiO 3 ) 2 . 
 Beryll, Be 3 Al 2 (Si0 3 ) 6 . 
 Related to wollastonite, but of more complicated struc- 
 
 Metasilicates : 
 
 J 
 i OJ 
 
 Acids, Si 4 OH 
 I OH 
 
 ture, are the important minerals, hornblende and 
 augite. 
 
 Disilicic acid is formed by the separation of one molecule of 
 water from two formula weights of metasilicic acid, just as disul- 
 phuric acid is derived from two of sulphuric ( page 150). 
 
 (0 0) (0 O) 
 
 SU OH-f HO V-Si = Si^ OH HO V- Si. 
 
 (OH HO) ( -0- ) 
 
 Only two or three examples of disilicates are known. Trisilicic 
 acid is formed by separating four molecules of water from three 
 formula weights of Si (OH ) 4 ; 3 Si (OH ) 4 4 H 2 O = Si 3 O 8 H 4 . 
 
 * Of the twelve hydrogen atoms in three formula weights of orthosilic 
 acid, six are replaced by trivalent iron and six by bivalent calcium. In this 
 silicate, and in others, two or more formula weights of the acid are united by 
 an atom of a polyvalent element, which replaces hydrogen atoms belonging in 
 part to one, and in part to another formula weight of the acid ; a simple 
 example of such a case we have encountered in the formula of the tertiary 
 phosphate of calcium, Ca 3 (PO 4 ) 2 , (see page 229). 
 
308 
 
 TRISILICATES. 
 
 Trisilicates: 
 
 
 (O 
 
 
 Si] OH 
 
 
 f 
 
 
 (OH 
 
 Acid, 
 
 Si JOH 
 
 
 ' O 
 
 
 SijOH 
 
 Orthoclase (feldspar), KAlSi 3 O 8 . 
 
 Oligoclase (soda, lime feldspar), NaAl(Si 3 O 8 ), CaAl 
 (AlSi 2 )0 8 . 
 
 (In the important group of minerals known as feldspars, it 
 not infrequently happens that a portion of the sili- 
 con is replaced by aluminium; this is seen in the 
 formula of oligoclase. ) 
 
 The quantitative composition of the silicates shows that every 
 one of them can be considered as derived from one of the above 
 mentioned acids. A number of basic and acid silicates also exist. 
 Among acid silicates, kaolin (clay) H 2 AL> (Si0 4 ) 2 + H 2 0, may be 
 mentioned. 
 
 The silicates are such extremely important minerals, their com- 
 position is so varied and their distribution so far reaching, that the 
 study of their structure forms one of the most important branches 
 of modern mineralogy. All silicates, excepting those of the alkali 
 metals, are insoluble in water. 
 
GERMANIUM. 309 
 
 CHAPTER XLIL 
 
 GERMANIUM AND ITS COMPOUNDS. 
 Germanium; symbol, Ge ; atomic weight, 72.3. 
 
 THIS element was discovered in the year 1886, by the German 
 chemist, Clemens Winckler, and is especially interesting from the 
 fact that it is one of the elements the existence of which was pre- 
 dicted before its discovery. This prediction was based upon the 
 fact that, when the elements were arranged in the order of their 
 atomic weights (page 17), an unfilled gap appeared to exist be- 
 tween gallium (atomic weight 69) and arsenic (atomic weight 75), 
 (see third table, chapter xlix.), which gap, as the nature of the then 
 known elements showed, should be filled by a representative of the 
 carbon family (see table of the periodic system). Germanium was 
 subsequently discovered in a silver ore which was formerly con- 
 founded with silver sulphide (argentite) and which has the formula 
 3 Ag 2 S, Ge S 2 and is termed argyrodite ; the properties of the new 
 element were found to be in accord with the predictions in regard 
 to its nature. The isolation of germanium is a very complicated 
 process. 
 
 Germanium, owing to its higher atomic weight, must be much 
 more metallic in its nature than silicon ; and, indeed, this difference 
 in its character is shown by the non-existence of a hydrogen com- 
 pound of the element. The metal has a brilliant metallic lustre, 
 and is readily formed from its oxide by reduction with charcoal ; and 
 like its fellows, carbon and silicon, it crystallizes in crystals be- 
 longing to the regular system. Its specific gravity is 5.47, and its 
 melting point less than that of silver (954). The metal is neither 
 malleable nor ductile ; it is quite brittle and can be readily pounded 
 to a powder; in this respect it resembles arsenic. When heated 
 to a high heat in the air, the metal, after fusing, oxidizes to form 
 the dioxide Ge 2 . Germanium does not dissolve in hydrochloric 
 acid ; it is oxidised to Ge O 2 by nitric acid or aqua regia. Hot and 
 concentrated sulphuric acid dissolves it to form the sulphate, while 
 the acid is itself reduced to sulphur dioxide (see page 136). 
 
810 GERMANIUM; COMPOUNDS OF. 
 
 Germanium combines with all of the halogens to form com- 
 pounds of the general formula GeK 4 , where E represents an atom 
 of any halogen. Germanium tetrachloride, formed by passing 
 chlorine over heated germanium, is a colorless liquid which very 
 much resembles the corresponding chloride of silicon, Si C1 4 ; it is 
 decomposed by water, and boils at 86 ; the specific gravity of its 
 vapor (between 300 and 740) is 7.43, this, H 2 = 2, is 213.9 ; the cal- 
 culated molecular weight for Ge C1 4 is 214, (Ge = 72.3, 4 01 = 141.8) ; 
 it follows from this that the formula of the chloride of germanium 
 corresponds to those of the chlorides of silicon or carbon, and that the 
 maximum atomic weight of germanium is 72.3 (see page 72). As, 
 before the discovery of germanium, no element was known to exist 
 having an atomic weight between that of gallium (69) and arsenic 
 ( 75 ), and as an element belonging to the carbon family and having 
 an atomic weight of approximately 72 would evidently find a fitting 
 place in the system obtained by arranging the elements in the order 
 of their atomic weights, therefore, the gravimetric quantity of ger- 
 manium ( 72.3 ), which unites with 141.8 (or 4 X 35.45 ) parts by 
 weight of chlorine, is probably the correct atomic weight of the 
 element in question. If, at any time, a compound of germanium 
 should be discovered which, with a known molecular weight, should 
 contain proportionally less than 72.3 parts by weight of the ele- 
 ment, then the atomic weight which is at present accepted will have 
 to be abandoned. A compound of germanium and chlorine, GeCl 2 , 
 corresponding to stannous chloride, Sn C1 2 , has also been mentioned, 
 but its properties have not, as yet, been thoroughly investigated. 
 The other halogen compounds of the element need not be described, 
 although the existence of a germanium-chloroform, Ge H C1 3 (a 
 compound corresponding to chloroform, CHC1 3 , and to silico-chlo- 
 roform, Si H C1 3 ), should be emphasized as showing the relationship 
 between germanium and the preceding elements of this family. 
 Germanium-chloroform is a colorless liquid which boils at 72 and 
 which is readily decomposed by water. 
 
 Germanium forms two oxides, Ge 2 and Ge 0. The former is 
 produced either by burning the powdered element in a stream of 
 oxygen, by oxidizing it with nitric acid, or by decomposing the 
 chloride with water. It is a white powder, somewhat soluble in 
 water, the solution probably containing the hydroxide H 2 Ge 3 , 
 corresponding to metasilicic acid, H 2 Si 3 . Germanium dioxide 
 
GERMANIUM; COMPOUNDS OF. 311 
 
 acts as a weak acidic anhydride, dissolving in the hydroxides of the 
 alkali metals ; it has no basic properties. The second oxide, Ge 0, 
 is an unstable substance which oxidizes in the air, has a hydroxide, 
 Ge (OH) 2 , derived from it, and is weakly basic in its character. 
 It is a powerful reducing agent. 
 
 Two sulphides of germanium, Ge S 2 and Ge S, are known. 
 These correspond to the sulphides of tin. The disulphide, Ge S 2 , 
 dissolves in the sulphides of the alkali metals to form sulpho-salts ; 
 it therefore has the character of an acidic anhydride, and resembles 
 the sulphides of arsenic, antimony, tin, and carbon. 
 
312 TIN; OCCURRENCE, PREPARATION. 
 
 CHAPTER XLIII. 
 
 TIN AND ITS COMPOUNDS. 
 
 Tin ; symbol, Sn ; atomic weight, 119. 
 
 TIN, the third element of the carbon family, has its metallic 
 properties so decidedly pronounced that, disregarding its many 
 resemblances to the not-metals, it is generally classed with the 
 metals. In reality it bears about the same relation to metals and 
 not-metals as antimony does ; however, both of its oxides (stannous 
 oxide, Sn 0, and stannic oxide, Sn 2 ) have basic properties, and, 
 furthermore, tin cannot form a hydrogen compound ; on the other 
 hand, both oxides, when brought in contact with strong bases, can 
 act like acidic anhydrides. 
 
 The time of the discovery of tin is not known. A knowledge 
 of the metal has been attributed to the Hebrews, Greeks, and Phoe- 
 nicians, but no certainty exists as to this. Undoubtedly, Pliny 
 distinctly mentions tin under the name of plumbum candidum, and 
 moreover, the metal was used by the Romans for covering iron in 
 order to k^ep that metal from rusting. The term stannum dates 
 from the fourth century. 
 
 Tin is one of the comparatively rare elements, and its occurrence 
 in the free state as a mineral is somewhat doubtful. It is chiefly 
 found as the crystallized dioxide, Sn 2 ; this substance (which is 
 termed cassiterite or tinstone), is deposited in crystals and in the 
 massive form in veins traversing granite, gneiss, and mica schist in 
 Cornwall and Devonshire, the Malay peninsula, New South Wales, 
 and Queensland. Some tin has also been discovered in the United 
 States. An impure sulphide of tin (stannite), SnS 2 , occasionally 
 appears in mineral deposits. 
 
 The tin of commerce is exclusively prepared from tinstone. The 
 mineral is crushed and washed and then heated with charcoal ac- 
 cording to the usual metallurgical process (see page 287). The 
 tin which melts and is collected at the bottom of the furnace is gen- 
 erally quite impure, for it contains copper, iron, arsenic, antimony, 
 and lead. These foreign substances are, for the most part, removed 
 
TIN; PROPERTIES. 313 
 
 by heating the crude tin to a temperature jfist above its melting 
 point, and then allowing the pure metal to run off from its higher 
 melting alloys. 
 
 Tin is nearly silver-white, with a metallic lustre ; it is scarcely 
 corroded when exposed to the air ; it is soft and can be hammered 
 and rolled into thin sheets (tin foil) ; its specific gravity is 7.3 ; it 
 melts at 228, and evaporates at a temperature between 1600 and 
 1700. The metal has a great tendency toward crystallization, 
 either when it is separated from its compounds, * or when it congeals 
 after fusion. Tin, like carbon, is dimorphous, occurring both in 
 tetragonal and in rhombic crystals. If block tin is cooled to a 
 very low temperature, or even if it is allowed to stand for a long 
 time, it changes into a grayish powder ; this powder will reassume 
 a metallic appearance only upon being fused. This amorphous 
 form of tin corresponds to amorphous carbon or silicon. Ordinary 
 tin, cast into forms, assumes a crystalline structure; if a stick 
 formed of the metal is bent, a peculiar crepitation is felt, and if the 
 operation is rapidly repeated several times, the piece of tin will 
 become quite hot at the place of bending ; both the noise and the 
 heat are caused by the friction of the minute crystals, one upon the 
 other. 
 
 Tin is attacked by acids with considerable ease; hydrochloric 
 acid dissolves it to form stannous chloride : 
 
 Sn-f 2HCl=SnCL,+2H, 
 
 and in this way the element shows its metallic nature ; that this is 
 not very pronounced, however, is shown by the fact that the re- 
 action between tin and hydrochloric acid takes place much more 
 slowly than it does between the same -acid and iron or zinc. Hot 
 and concentrated sulphuric acid dissolves tin to form stannous sul- 
 phate, while the acid is reduced to sulphur dioxide (see page 137 ). 
 
 * By electrolysis of stannous chloride, or by placing a piece of zinc in a 
 solution of stannous chloride, the zinc then takes the place of tin in the salt: 
 
 Zn + Sn C1 2 = Zn C1 2 + Sn. 
 
 Such substitutions of one metal for another in salts are quite frequently met 
 with, but are not surprising if we compare salts with acids, for, as we know, 
 zinc can readily replace hydrogen in hydrogen chloride, and why not tin 
 in stannous chloride? The only essential is that stannous chloride should have 
 more chemical energy than zinc chloride, and that heat should be given off in 
 the reaction. 
 
314 STANNOUS CHLORIDE. 
 
 Cold and dilute nitric acid dissolves tin without any evolution of 
 gas ; the tin forms stannous nitrate, while the nitric acid is reduced 
 to ammonia (page 206 [a] ) ; this reaction is made clear by the fol- 
 lowing equation : 
 
 4 Sn + 10 HN0 3 = 4 Sn ( N0 3 ) 2 + NH 4 N0 3 + 3 H 2 0.* 
 
 On the other hand, hot and concentrated nitric acid oxidizes tin to 
 insoluble metastannic acid, H 2 Sn 3 , and is itself reduced to the 
 lower oxides of nitrogen (page 207 [&] ).f Tin shows its relation- 
 ship to the not-metals by dissolving in alkaline hydroxides to form 
 salts of stannic acid, M 2 Sn 3 . 
 
 Tin forms two series of compounds ivith the halogens ; the first 
 of these, with the general formula SnX 2 (where X represents any 
 halogen), can be formed, as are the salts of other metals, by dissolv- 
 ing the corresponding oxide in halhydric acids : 
 
 Sn + 2 HX = SnX 2 + H 2 0, 
 
 while the compounds SnX 2 can be converted into those of the 
 second series, Sn X 4 , by the addition of the corresponding halogen. 
 The two chlorides of tin, Sn C1 2 , stannous chloride, and Sn C1 4 , 
 stannic chloride (see page 26), are the most important of these 
 halogen compounds. 
 
 Stannous chloride, Sn C1 2 , can be formed by dissolving tin or 
 stannous oxide, Sn 0, in hydrochloric acid ; when anhydrous, it is a 
 crystalline substance which melts at 250 and boils at 606. Its 
 vapor density was formerly supposed to correspond to a molecular 
 weight calculated from the formula Sn 2 C1 4 ; but later investigations 
 have shown that no definite specific gravity can be assigned to it. 
 Stannous chloride dissolves in small quantities of water without 
 change ; an excess of the solvent, however, partially converts it 
 into an insoluble basic chloride (see page 253) : 
 
 Stannous chloride is a powerful reducing agent ; when exposed to 
 the air it absorbs oxygen and changes into a mixture of stannic 
 chloride and the basic chloride just mentioned. Stannous chloride 
 instantly reduces mercuric chloride to mercurous chloride : 
 
 * Hydroxylanim is also produced by this reduction, especially if hydro- 
 chloric acid is present. 
 
 t Compare with antimony, page 249. 
 
STANNIC CHLORIDE. 315 
 
 2 Hg C1 2 + Sn C1 2 = Sn C1 4 + 2 Hg Cl,* 
 
 and the mercurous chloride, by further action, is even finally 
 changed to mercury. Ferric chloride is reduced to ferrous chloride 
 by stannous chloride : 
 
 2 Fe C1 3 + Sn C1 2 = Sn C1 4 + 2 Fe C1 2 , 
 
 while arsenic trioxide is reduced to metallic arsenic by the same 
 substance, the compound of tin being, in this case, oxidized to stan- 
 nic acid. 
 
 When alkaline hydroxides are added to a solution of stannous 
 chloride, insoluble stannous hydroxide is at first precipitated : 
 
 Sn C1 2 + 2 KOH = Sn (OH ) 2 + 2 K Cl, 
 
 but this substance, because it presents in a slight degree the charac- 
 ter of an acid, is dissolved by an excess of the alkaline solution to 
 form stannites which are salts of a stannous acid having the formula 
 H 2 Sn 2 3 , formed, as are a number of acids which we have already 
 discussed, by the separation of one molecule of water from two 
 formula weights of the hydroxide : 
 
 g (OH Sn-OH 
 
 OH Sn-OH 
 
 The potassium compound, formed by dissolving stannous hydrox- 
 ide in an excess of potassium hydroxide, therefore, has the formula 
 K 2 Sn 2 3 ; this salt, when heated, breaks down into tin and potassium 
 stannate : 
 
 Sn 2 3 K 2 = Sn 3 K 2 -f Sn.t 
 
 Stannic chloride can be formed from stannous chloride by heating 
 the latter substance and then passing dry chlorine over it. It is a 
 colorless liquid which fumes in the air, and which boils at 114; it 
 greedily absorbs moisture, and then produces crystals having the 
 formula SnCl 4 -f3H 2 0. Stannic chloride forms a series of 
 
 * For this reason stannous chloride is used as a reagent for soluble salts 
 of mercury, for, as mercurous chloride is insoluble in water, a precipitate of 
 the latter is formed when stannous chloride is added to a solution containing 
 mercury. 
 
 t See pages 139 and 155. This reaction is similar to previous ones which 
 we have studied. 
 
316 STANNOUS OXIDE. 
 
 double salts with the chlorides of the alkali metals,* these double 
 salts have the general formula SnCl 6 M 2 , where M represents an 
 atom of an alkali metal ; they therefore correspond to the silico- 
 fluorides, Si F 6 M 2 (see page 302 ), which we look upon as being 
 salts of fluosilicic acid ( H 2 Si F 6 ) ; there is consequently no reason 
 why the conclusions regarding the nature of fluorine in the fluosili- 
 cates should not be equally applicable to chlorine in these com- 
 pounds of tin; indeed, a substance H 2 SnCl 6 + 6 H 2 0, which must 
 be considered as the acid from which these double salts are de- 
 rived, has in all probability been isolated. 
 
 Tin forms two oxides, the monoxide, SnO, and the dioxide, 
 Sn 2 ; these correspond to the oxides of carbon, CO and C0 2 ; the 
 former of these oxides is almost altogether basic in its character, 
 while the latter most frequently acts as the anhydride of an acid. 
 
 Stannous oxide can best be prepared by heating the correspond- 
 ing hydroxide without access of air. It is a dark brown substance 
 which dissolves in acids to form stannous salts, or in alkalies to 
 form stannites ; this latter reaction has already been fully described 
 under stannous chloride. The stannous salts are colorless when 
 formed from a colorless acid, and are readily oxidized when in con- 
 tact with the air ; those insoluble in water are nearly all soluble in 
 dilute hydrochloric acid, and show a tendency to change into basic 
 salts on the addition of water. 
 
 When tin is heated to a sufficiently high temperature in air or 
 in oxygen, it burns to form the dioxide, Sn 2 ; this substance when 
 cold is a white powder, but when hot assumes a yellowish color; 
 after being exposed to a high temperature for some time it becomes 
 insoluble both in acids and alkalies. The crystallized variety of 
 the oxide, found as the mineral tinstone, is also insoluble ; the only 
 means by which this substance can be brought into solution is by 
 fusion with potassium or sodium hydroxide, when the respective 
 stannates, M 2 Sn 3 , of the metals are formed. 
 
 Two stannic acids, identical in gravimetric composition, but dif- 
 fering in physical and chemical properties, are derived from the 
 anhydride Sn0 2 ; both have the formula H 2 Sn0 3 .t Ordinary 
 
 * Stannous chloride also forms double halides with the chlorides of the 
 alkali metals ; the double halides of potassium are K Sn C1 3 4- H 2 O and K 2 Sn 
 C1 4 + 2 H 2 O . 
 
 t Two orthostannic acids (Sn (OH ) 4 ) have also been described. See 
 Neumann; Monatshefte fur Chemie; 12, 515. 
 
STANNIC ACIDS. 317 
 
 stannic acid can be formed by adding exactly enough potassium 
 hydroxide solution to a solution of stannic chloride to precipitate 
 stannic hydroxide : 
 
 Sn C1 4 + 4 KOH = Sn (OH) 4 + 4 K Cl ; 
 
 this substance separates as a jelly which resembles silicic acid. 
 When this is dried it loses water and changes to a gum-arabic like 
 mass which approximately has the formula H 2 Sn 3 . This variety 
 of stannic acid is readily soluble both in acids and alkalies. The 
 other form of stannic acid, generally called metastannic acid, is 
 presumably a polymeric form of ordinary stannic acid, so that if 
 the molecule of the latter had the formula H 2 Sn 3 , that of the 
 former would be expressed by n (H 2 Sn 3 ). A more correct system 
 of nomenclature is one by which ordinary stannic acid is termed 
 a stannic acid, while metastannic acid is called ft stannic acid. 
 ft stannic acid is produced in the form of an insoluble white powder 
 when tin is oxidized by means of strong and hot nitric acid ; when 
 carefully dried in a vacuum it has the formula H 2 Sn 3 ; it is 
 insoluble in acids, and when heated to redness, loses water and 
 changes into the dioxide of tin. If ft stannic acid is digested with 
 hydrochloric acid for some time, the hydrochloric acid then poured 
 off and pure water added, the stannic chloride so formed will dis- 
 solve ; on addition of alkalies to this solution, however, ft stannic 
 acid is once more precipitated. When ft stannic acid is boiled with 
 sodium hydroxide it is converted into the ft stannate of sodium, 
 which can be dissolved by pouring off the excess of caustic soda 
 solution and then adding pure water. Fusion with solid caustic 
 alkalies converts ft stannic acid into salts of ordinary stannic acid. 
 Stannic chloride, or what amounts to the same thing, the solu- 
 tion of either stannic acid or ft stannic acid in hydrochloric acid, 
 resembles the chlorides of the not-metals in. so far as it is converted 
 into the corresponding acid by boiling with water ; this decompo- 
 sition is, however, only partial, for if a solution of ordinary stan- 
 nic acid in hydrochloric acid is boiled in a retort, the volatile 
 stannic chloride passes over unchanged in company with the water 
 and hydrochloric acid, while very little stannic acid will remain 
 behind ; on the other hand, a solution of stannic chloride derived 
 from ft stannic acid is completely decomposed into that acid by 
 boiling, while no stannic chloride whatever will, pass over. 
 
318 TIN; SULPHIDES. 
 
 The stannates of the alkali metals are the only ones which are 
 soluble in water ; in that way these salts resemble those of silicic 
 and carbonic acids. 
 
 Tin forms two sulphides, with the formulae Sn S and Sn S 2 ; in 
 structure these compounds correspond to the oxides. The monosul- 
 phide, Sn S, can be produced by adding hydrogen sulphide to an 
 acidulated solution of a stannous salt. The sulphide is a brownish 
 black powder, which is insoluble in dilute acids, but is dissolved by 
 concentrated hydrochloric acid or by aqua regia ; in the latter case 
 stannic chloride is formed. Simple sulphides of the alkalies scarcely 
 attack it ; it is dissolved by the polysulphides (see page 155, foot- 
 note), because the latter sulphurize it to form salts of sulphostannic 
 acid, in a manner exactly similar to their action on the trisulphide 
 of antimony (see page 256). The disulphide of tin, Sn S 2 , can be 
 precipitated from a weakly acid solution of stannnic chloride by 
 means of hydrogen sulphide ; it forms a yellow precipitate which is 
 not dissolved by dilute acids, but which is soluble in strong hydro- 
 chloric acid or aqua regia ; with the latter reagent it forms stannic 
 chloride. As would be expected, owing to the weakly metallic na- 
 ture of tin, the disulphide acts as if it were an acidic anhydride ; 
 it is, therefore, readily attacked by either the hydroxides or sul- 
 phides of the alkalies (see page 256 and foot-note), for in the former 
 case a mixture of stannate and sulphostannate is formed,* while in the 
 latter the sulphostannate alone is produced, f This behavior of tin 
 is very much like that of antimony or arsenic under similar circum- 
 stances (see pages 245 and 255). On addition of acids to the solu- 
 tion of sulphostannates, stannic sulphide is precipitated, Sn S 3 K 2 -f 
 2 H. Cl = 2 K Cl + H 2 S + Sn S 2 , for the sulphostannic acid which 
 would be formed at once breaks down into hydrogen sulphide and 
 the disulphide of tin. (It will be noticed that the formulae of the 
 sulphostannates correspond to those of the salts of dithio-sulpho- 
 carbonic acid [see page 293].) 
 
 The compounds of tin show an almost perfect concordance with 
 those of carbon and silicon, when the formulae alone are considered ; 
 chemically, however, the acid nature of the substances in question 
 is materially reduced because of the metallic character belonging 
 to tin. This metallic character becomes much more pronounced in 
 
 * 3 Sn S 2 + 6 KOH = Sn O 3 K^ + 2 Sn S 3 K 2 + 3 H. 2 O. 
 t SnS 2 
 
TIN AND CARBON ; COMPARISON OF. 
 
 319 
 
 the next element of the family (lead), so that the oxides of that 
 element are for the most part basic. 
 
 The relationship between the compounds discussed in this chapter 
 will be seen from the following table : 
 
 TIN AND CARBON. 
 
 OXIDES. 
 
 CHLORIDES. 
 
 ACIDS. 
 
 SULPHIDES. 
 
 
 CO 
 CO,' 
 
 SnO 
 SnO 2 
 
 CC1 4 
 
 SnCl 2 
 SnCl 4 
 
 
 H 2 Sn 2 O 3 
 
 H 2 SnO 3 
 
 
 SnS 
 
 Sn S 2 
 
 Forms no sulpho salts. 
 Forms sulpho salts, M 2 X S 3 . 
 
 H 2 CO 3 
 
 CS 2 
 
 The oxide Sn O is both basic and acidic. 
 
 It dissolves in acids as follows : Sn O + 2 HX = Sn X 2 + H 2 O. 
 
 It dissolves in bases as follows : 2 Sn O + 2 MOH = Sn 2 O 3 M 2 + H 2 O. 
 
 The oxide Sn O 2 is both basic and acidic. 
 
 It dissolves in halhydric acids as follows : 
 
 SnO., + 4HX = SnX 4 + 2H 2 0. 
 It dissolves in bases as follows : 
 
 Sn 2 + 2 MOH = Sn O 3 M 2 + H 2 O. 
 
 a STANNIC ACID, H 2 Sn O 3 . 
 
 STANNIC ACID, H 2 Sn O 3 . 
 
 Derived from an oxide soluble in acids 
 and alkalies. 
 
 The chloride, SnCl 4 , derived from it is 
 volatile in the vapors of dilute hydrochloric 
 acid. 
 
 Sodium hydroxide readily dissolves it in 
 the cold, forming ordinary stannate of 
 sodium. 
 
 Ordinary stannic chloride, on addition of 
 alkalies, precipitates ordinary stannic acid. 
 
 Derived from an oxide insoluble in acids 
 and alkalies. 
 
 The chloride, SnCl 4 , derived from it is 
 not volatile with water vapors (metastan- 
 nic chloride); it decomposes into hydrochlo- 
 ric acid and ft stannic acid when the solution 
 is boiled. 
 
 Sodium hydroxide, when boiling, forms 
 ft stannate of sodium, which is soluble in 
 water. 
 
 ft stannic chloride, on addition of alkalies, 
 precipitates ft stannic acid. 
 
320 LEAD; OCCURRENCE, PREPARATION. 
 
 CHAPTER XLIY. 
 
 vj 
 
 LEAD AND ITS COMPOUNDS. 
 
 Lead ; symbol, Pb; atomic weight, 206.95. 
 
 THE element is seldom found as the native metal. Its most 
 important natural compound is the sulphide, Pb S, which occurs 
 widely distributed as galena (also termed galenite). The carbon- 
 ate (cerassite, PbC0 3 ), the sulphate (anglesite, PbS0 4 ), the chro- 
 mate, phosphate, and molybdate are also not infrequently met 
 with. 
 
 Lead is one of the metals which has been known since the old- 
 est times, having been familiar to the Israelites. The Eomans 
 made much the same use of the metal as we do at the present time, 
 for they constructed water-pipes of it, and prepared a solder com- 
 posed of two parts of lead and one of tin. 
 
 The lead which is met with as a commercial product usually 
 contains copper, iron, and traces of silver. It is prepared by heat- 
 ing the sulphide with finely divided iron : 
 Pb S + Fe = Fe S + Pb. 
 
 This operation is conducted in tall furnaces, which, in shape, re- 
 semble the blast furnaces for the manufacture of pig iron. Another 
 method for the production of the metal consists in roasting the 
 sulphide in a current of air, by which means it is in part oxidized, 
 so that a mixture of the sulphate, oxide, and sulphide are formed, 
 and when this mixture is heated to a higher temperature the sul- 
 phate and oxide are finally reduced by the sulphide, which is still 
 present; the sulphur passes off in the form of sulphur dioxide, 
 while the lead remains behind. The crude lead contains silver, 
 antimony, arsenic, copper, iron, and zinc; the oxidizable impurities 
 are removed by melting the metal in the air. When silver is pres- 
 ent in sufficient quantity to pay for its isolation, the entire mass is 
 melted and subjected to a blast of air, by which means lead oxide 
 is produced ; the latter melts and is run off from the surface, while 
 the silver remains behind unchanged. The lead oxide so formed 
 
LEAD ; PROPERTIES. 321 
 
 can then be once more reduced to lead by means of charcoal. 
 Another method for removing the silver consists in melting the 
 silver-bearing lead and then allowing the mass to cool slowly; pure 
 lead crystallizes at first ; this can be removed by means of a ladle, 
 while the molten mass remaining, which is very rich in silver, can 
 be treated according to the method mentioned above ; or the silver 
 can be removed by melting the lead with zinc, for zinc mixes with 
 lead in a small proportion only, while it is able to dissolve all of the 
 silver. 
 
 Lead is a metal with a bluish-gray color and metallic lustre; 
 it is malleable and easily fused, its melting point is 330, and it 
 can be boiled at a high white heat.* The metal crystallizes in 
 octahedra ; its specific gravity is 11.4 ; when freshly cut it has a 
 bright, metallic surface, which, however, soon becomes covered with 
 a layer of the oxide, and this protects the remainder from further 
 corrosion. If a piece of zinc is placed in a solution of a lead salt, 
 the lead will separate in a crystalline form, while the zinc takes its 
 place : 
 
 . Zn + Pb (N0 8 ) 2 = Zn (X0 3 ) 2 + Pb; 
 
 similar substitutions are not infrequently met with in the chemis- 
 try of other metals (see page 313 and foot-note). 
 
 When lead is covered with water which is in contact with the 
 air, it becomes covered with a layer of lead hydroxide ; the latter 
 substance is to a certain extent soluble in water ; as a consequence 
 water which has passed through new lead pipes contains more or 
 less of the hydroxide in solution, and may, for this reason, prove 
 to be highly poisonous. Hard water gradually changes the hy- 
 droxide into the entirely insoluble carbonate, so that, in time, the 
 pipes become covered with a protective coating. In dealing with 
 lead pipes, however, care must be taken to have no decaying organic 
 substances present, for such impurities may remove the carbonate 
 and greatly increase the solubility of the lead. 
 
 Lead is not readily attacked by hydrochloric f or cold sulphuric 
 acid ; hot and concentrated sulphuric acid has some effect on it, as 
 
 * The boiling point lies between 1450 and 1600 (Carnelley and Williams; 
 Journ. Chem. Soc. ; 35, 563). 
 
 t Hot hydrochloric acid attacks lead to a considerable extent (Sharpies; 
 hem. News; 50, 126). 
 
. 
 
 322 LEAD MONOXIDE. 
 
 is evinced by the fact that commercial sulphuric acid always con- 
 tains lead ; diluted nitric acid readily dissolves the metal to form 
 lead nitrate. Concentrated nitric acid has but little effect. 
 
 Lead enters quite readily into a number of alloys, some of 
 which have already been mentioned (see page 251) ; the metal is 
 easily amalgamated by mercury. 
 
 Lead forms four oxides, Pb 2 0, suboxide of lead ; Pb 0, lead 
 monoxide ; Pb 2 O 3 , lead trioxide ; and Pb 2 , lead dioxide, or lead 
 hyperoxide. Of these, the oxides PbO and Pb 2 are the most 
 important; the monoxide, PbO, corresponds to carbon monoxide, 
 and the dioxide, Pb 2 , to carbon dioxide. 
 
 The oxide of lead, Pb 0, can easily be produced by heating the 
 nitrate (see page 201), or, like other oxides of weakly pronounced 
 metals, it can be formed by heating the hydroxide : 
 
 Pb(OH) 2 = PbO + H 2 0. 
 
 This oxide of lead is ordinarily a yellow powder, which is easily 
 melted to an orange-colored mass (litharge) ; but another, red, 
 modification of the oxide can be obtained by heating lead hydrox- 
 ide to 150. The hydroxide of lead separates as a white precipi- 
 tate when a base is added to a solution of a lead salt : 
 
 Pb ( N0 3 ) 2 + 2 Na OH = Pb (OH ) 2 + 2 Xa NO S , 
 
 and, like stannous hydroxide, it is both a base and an acid. As a 
 result of its basic properties, it readily dissolves in acids to form 
 salts : 
 
 Pb (OH) 2 + 2 HN0 3 = Pb (N0 3 ) 2 + 2 H 2 ; 
 
 and because of its acid properties it dissolves in pronounced 
 alkalies : 
 
 Pb (OH) 2 + 2 KOH = Pb (OK) 2 + 2 H 2 0; 
 
 but the hot saturated solution deposits lead oxide on cooling (com- 
 pare with stannous hydroxide, page 315). The salts of lead can 
 be formed by dissolving either the oxide or hydroxide in acids, and, 
 being salts of a pronounced metal, they are more or less stable; 
 they are poisonous when in a soluble form. Among the most im- 
 portant salts of lead are the chloride, PbCl 2 , sulphateJPb S0 4 , and 
 chromate, Pb Cr 4 , which are insoluble or nearly insoluble in 
 water; they can therefore be produced from the soluble lead salts 
 by the addition of a soluble chloride, sulphate, or chromate. The 
 
LEAD DIOXIDE. 323 
 
 soluble carbonates of the alkalies when added to the solution of a 
 lead salt cause a precipitate of basic carbonate of lead, Pb(OH) 2 
 Pb C0 3 ; this substance is white lead. Among the important soluble 
 lead salts, the acetate (Pb(C, H 8 0,) 2 , sugar of lead) may be men- 
 tioned. One interesting fact as regards lead salts is the isomor- 
 phism which the sulphate displays with the sulphates of the very 
 pronouncedly metallic alkaline earths (calcium, barium, and stron- 
 tium), and the isomorphism of the carbonate with the carbonates 
 forming the arragonite group.* This relationship shows us that 
 the pronounced metallic properties of lead have caused it to depart 
 so far from the family type that its salts resemble those of the most 
 characteristic divalent metals ; this isomorphism is also displayed in 
 the case of some other lead compounds. 
 
 The oxide of lead next in importance to the monoxide, Pb 0, is 
 the dioxide or hyperoxide, Pb 2 ; this substance has only very weakly 
 basic or acid properties. It belongs to the class of hyperoxides of 
 which manganese dioxide is, perhaps, the best known representa- 
 tive. The hyperoxides are all neutral, or nearly neutral bodies, 
 which, when heated with sulphuric acid, give off oxygen and 
 change to the sulphate derived from the. oxide MO as a base, and 
 which, when treated with hydrochloric acid, liberate chlorine. The 
 dioxide of lead is occasionally found in nature ; it can be formed 
 in the laboratory by treating the oxide Pb 3 4 f with nitric acid, 
 or by oxidizing the acetate of lead with a solution of chloride of 
 lime. 
 
 Lead dioxide is a dark brown powder, which is a powerful 
 oxidizing agent ; indeed, it can oxidize sulphur dioxide so readily 
 that the heat of the reaction may even cause it to glow, provided it 
 is finely divided and placed in an atmosphere of the gas. When 
 in contact with strong bases it dissolves to form salts of an acid, 
 plumbic acid, H 2 Pb 3 , which, in formula, .corresponds to carbonic 
 acid, but which, like the latter, does not exist in the free state ; 
 only a very few salts of this acid are known. On the other hand, 
 
 * The carbonates of calcium, barium, manganese, and iron. 
 t The red oxide of lead, Pb 3 O 4 , may be looked upon as a mixture of the 
 oxides Pb O and Pb O 2 : 
 
 2PbO + PbO 2 = Pb 3 O 4 ; 
 
 the nitric acid dissolves out the monoxide Pb O and leaves the dioxide Pb O 2 ; 
 compare this formula with Mn 3 O 4 and Fe 3 O 4 . 
 
324 MINIUM. LEAD SULPHIDE. 
 
 the dioxide is soluble in some acids, although the salts which are 
 presumably formed by this action have not been isolated. 
 
 The red oxide of lead ( Pb 3 _0 4 , so-called minium) is produced by 
 carefully oxidizing finely powdered litharge at 300 to 400 at 
 a higher temperature it breaks down with the liberation of oxygen 
 and the regeneration of litharge. This oxide is extensively used as 
 a pigment under the name of red lead. The suboxide of lead, Pb 2 0, 
 and the sesquioxide, Pb 2 O 3 , are of little importance. 
 
 The metallic character of lead is so predominant that the nature 
 of its oxides is very much at variance with that of the oxides of 
 the elements at the beginning of this family ; indeed, in the forma- 
 tion and character of the compounds Pb 3 4 and Pb 2 , lead very 
 much resembles manganese, while the isomorphism of its salts with 
 those of calcium, barium, and strontium, brings it in close connection 
 with the alkaline earths ; on the other hand, it is like carbon, sili- 
 con, and tin, for its oxide Pb 2 shows weakly acid properties. In 
 fact, the chemical characteristics of lead are not very marked in any 
 direction, nor, indeed, is this neutral behavior unexpected, for we 
 find it to be quite a general fact that the elements with high atomic 
 weights and specific gravities * display no very pronounced chemical 
 properties ; the crowding of a large mass into a small space, as is 
 the case with these elements, is therefore unfavorable for the 
 manifestation of striking chemical phenomena. 
 
 Lead forms only one sulphide which has been accurately stud- 
 ied, the monosulphide Pb S. This substance is found as the min- 
 eral galena, crystallized in cubes of metallic appearance. In the 
 laboratory it can be produced either by direct combination of the 
 elements, or by precipitation from acid solutions of lead salts by 
 means of hydrogen sulphide ; when so prepared it is a black, amor- 
 phous powder which is not attacked by cold hydrochloric acid, but 
 which is attacked by that substance when it is hot and concen- 
 trated. Oxidizing agents, such as nitric acid, change it into the 
 insoluble sulphate of lead, and a similar transformation is brought 
 about by roasting in the air. 
 
 In the following table the most important compounds which 
 have representatives in the chemistry of a number of elements of 
 this family are placed side by side; those acids which are not 
 
 * I.e., with small atomic volumes (see chapter xlix.). 
 
COMPOUNDS OF CARBON FAMILY. 
 
 325 
 
 known in a free state, but salts of which exist, are placed in 
 parentheses : 
 
 
 OXIDES. ACIDS. 
 
 Monoxides. 
 Dioxides. 
 
 CO, , Ge O, Sn O, Pb O. , , , H 2 Sn 2 O 3) Pb (OH) 2 . 
 C0 2 , Si O 2 , Ge O 2 , Sn O 2) Pb O 2 . (H 2 CO 3 ), H 2 SiO 3 , ,H 2 Sn O 3) (H 2 PbO 3 ) 
 
 SALTS. 
 
 M 2 CO 3 , M 2 Si O 3 , , M 2 Su O 3 , M 2 Pb O 3 . 
 
 
 CHLOBIDES. 
 
 
 SULPHIDES. 
 
 Dichlorides. 
 Tetrachlorides. 
 
 , ,GeCl 2 ,SnCl 2 , 
 Pb C1 2 . 
 C C1 4 , Si C1 4 , Ge C1 4 , 
 SnCl 4 , PbCl 4 .* 
 
 Monosulphides. 
 Bisulphides. 
 
 CS, , , SnS, 
 PbS. 
 CS 2 , SiS 2 , GeS 2 , SnS 2 , 
 
 Presumably formed because the dioxide dissolves in acids. 
 
326 ELEMENTS OF BORON FAMILY. 
 
 CHAPTER XLV. 
 
 THE ELEMENTS OF THE BORON FAMILY (THE EARTHS). 
 
 THE elements of the family of which boron is the member with 
 the smallest atomic weight, are the last which will be considered 
 where any purely not-metallic element occurs. As has been re- 
 peatedly mentioned, the groups of elements become, as a whole, 
 more metallic in their nature as the atomic weights diminish ; this 
 fact is readily recognized by comparing the various families in the 
 order in which they have been studied, as is shown in the table on 
 page 266. In the nitrogen group there are four elements, nitrogen, 
 phosphorus, arsenic, and antimony, which could be classed with the 
 not-metals ; in that of carbon there are but two, carbon and silicon ; 
 while in the one under consideration, boron alone appears to us with 
 pronouncedly not-metallic characteristics, while even this element 
 is unable to form a gaseous hydrogen compound of sufficient sta- 
 bility to render an accurate study of its properties practicable. 
 
 The elements comprising the boron family are given in column 
 1, those of the carbon family are, for purposes of comparison, placed 
 in column 2. 
 
 1. Boron, atomic weight, 11 2. Carbon, atomic weight, 12 
 Aluminium, " " 27 Silicon, " " 28.4 
 
 Gallium, " " 69 Germanium, " " 72.3 
 
 Indium, " " 113.7 Tin, " " 119. 
 
 Thallium, " " 204.18 Lead, " " 206.95 
 
 The highest valence toward oyxgen displayed by the elements 
 of the nitrogen family is five, as is shown by the existence of the 
 pentoxides, X 2 5 ; in the carbon family the power of uniting with 
 oxygen is exhausted when the dioxide, X0. 2 , in which the element 
 is quadrivalent, is reached ; while in the boron group, with a return 
 to the type of oxide shown by nitrogen and its fellows, the highest 
 valence displayed toward oxygen is only three, so that the character- 
 istic oxides of this group have the formula X 2 3 . These oxides, 
 of course, suffer a diminution in their acidic character the greater 
 the atomic weight of the element forming them is ; so that boron 
 trioxide, B 2 3 , acts as an acidic anhydride under all circumstances ; 
 
ELEMENTS OF BOKON FAMILY; COMPARISON. 327 
 
 aluminium trioxide, A1 2 3 , is both basic and acidic ; the oxide of 
 gallium, Ga 2 3 , displays the same character; while the oxide of 
 indium, In 2 3 , is almost exclusively basic, for it with difficulty 
 dissolves in caustic alkalies, and the unstable compound so formed 
 is broken down by warming the solution ; finally, the trioxide of thal- 
 lium, T1 2 3 , is not affected by the reagents in question ; from the 
 above comparative statement it follows that the oxides X 2 O 3 are 
 more basic the greater the atomic weight of X. As the elements in 
 the family increase in atomic weight, they display the same ten- 
 dency to form a number of oxides which is observed in the case of 
 lead in the carbon family ; this fact will become apparent by a 
 study of the following table : 
 
 Boron forms one oxide, B. 2 O 3 ; 
 
 Aluminium forms one oxide, A1 O 3 ; 
 
 Gallium forms two oxides, Ga O and Ga 2 O 3 ; 
 
 Indium forms two oxides, InO and In 2 Og; 
 
 Thallium forms three oxides, T1 2 O, T1 2 O 3 , and Tl O 2 . 
 
 The oxides with least amount of oxygen, derived from any given 
 element in this family, capable of forming more than one oxide are, 
 without exception, basic in their character, while the trioxides, with 
 the exception of that of thallium, are both basic and acidic. The 
 majority of the salts which contain an element which is a member 
 of this family are derived from the trioxide X 2 3 . 
 
 What is true of the oxides is also true of the halogen com- 
 pounds ; the elements with high atomic weights each are capable of 
 forming more than one chloride, bromide, or iodide, while alumin- 
 ium and boron are confined to one apiece. The trihalide, like the 
 trioxide, is the compound common to all of the members of the 
 family. The relationship between these compounds can be seen by 
 examining the following table ; as will be noticed, the rule which 
 holds good in all the preceding families, namely, that the boiling 
 points of the trichlorides are higher the greater the molecular 
 weight of the compound, is also without exception in this group : 
 
 BC1 3 , liquid, boils at 17; 
 
 Al C1 3 , solid, melts at 190, boils at 183; 
 
 Ga C1 3 , solid, melts at 75.5, boils at 220; 
 
 In C1 3 , solid, volatilizes at red heat without melting; 
 
 Tl C1 3 , gives off chlorine when heated, and changes to the chloride Tl Cl. 
 
 In addition to the chlorides given above, gallium forms a compound with 
 the formula Ga C1 2 , indium the chlorides, In Cl and In C1 2 , while thallium also 
 has a monochloride, Tl Cl. 
 
328 BORON; OCCURRENCE, PREPARATION. 
 
 CHAPTER XLVI. 
 
 BORON AND ITS COMPOUNDS. 
 
 Boron y symbol, B ; atomic weight, 11. 
 
 IN its physical characteristics boron bears a marked resemblance 
 to carbon (the element having the next highest atomic weight to its 
 own) ; but in the chemistry of its oxides and chlorides, boron is 
 very much like the members of the nitrogen family. 
 
 Boron is never found uncombined in mineral deposits. Its com- 
 pounds, which not infrequently occur in nature, are either salts of 
 boric acid or the acid itself. The most important of these minerals 
 
 are: 
 
 Borax (tinkal) NaH (BO 2 ) 2 , 4 H 2 O; 
 
 Borocalcite, CaH 2 (BO 2 ) 4 , 5 H 2 O. 
 
 The borates of other metals, for instance, of iron and magnesium, are 
 also found, while solutions of boric acid sometimes occur in lagoons of volcanic 
 regions. 
 
 Although the element was not isolated until 1807,* and was not 
 accurately described until 1824,t its compounds, especially tinkal, 
 occurring as they do in mineral deposits, were known in very early 
 times, the natural borax having become familiar to Europeans by 
 importations from India. 
 
 Boron can be isolated by. heating the oxide ( B 2 3 ) with sodium, 
 or potassium fluoborate (KBF 4 ) with potassium. When so obtained 
 it is an amorphous, brownish-black powder which greatly resembles 
 silicon in appearance ; it is quite readily dissolved by melted alu- 
 minium, and, when the metal containing the boron is cooled, the ele- 
 ment separates in the form of reddish-yellow, diamond-like crystals, 
 which are very hard and lustrous. When the aluminium has been 
 heated to a temperature only just above its melting point, then the 
 dissolved boron appears in a graphitoidal form. Elementary boron, 
 therefore, displays modifications similar to those belonging to carbon. 
 The specific gravity of boron is 2.68 ; it is infusible. 
 
 * By Guy Lussac and Thenard, 
 t By Berzelius. 
 
BORON; HALIDES. 329 
 
 When heated in the air, amorphous boron burns to form the 
 trioxide B 2 3 ; the same modification of the element is readily 
 oxidized by nitric acid, or even by concentrated sulphuric acid,* 
 boric acid, B 3 H 3 , being produced. Boron can unite directly with 
 chlorine, bromine, with some metals, and with nitrogen. 
 
 A gaseous hydrogen compound of boron was not known until 
 quite recently, f It was then prepared in an impure state by treat- 
 ing the boride of magnesium with hydrochloric acid ; this method 
 corresponds to the one by which silicon hydride, Si H 4 , and pure 
 arsine and stibine are produced. Hydrogen boride is very unstable ; 
 it burns in air or in oxygen with a bright green flame, and it is 
 slightly soluble in water. The formula assigned to the gas is B H 3 , 
 but a more extended investigation of its composition is necessary. 
 
 The halogen compounds of boron correspond to the general 
 formula of B X 3 , so that, in structure, they are identical with the 
 trihalides of the nitrogen family. The trichloride and trifluoride are 
 the only representatives of these compounds which we need consider. 
 The trichloride is formed by the direct union of chlorine and boron ; $ 
 when first discovered it was supposed to be a gas at ordinary tem- 
 peratures, but subsequent investigations proved it to be a liquid 
 with a boiling point at 17. Being the chloride of a not-metal, 
 boron trichloride is, of course, readily decomposed by water, and 
 normal or ortho boric acid results from this decomposition : 
 
 Cl HOH OH 
 
 CI + HOH OH 
 
 It will be noticed that this change is parallel to the one undergone 
 by phosphorus trichloride when in the presence of water (see pages 
 81 and 220). 
 
 Boron trifluoride is a colorless gas, and is interesting because its 
 chemical properties are much like those of the tetrafluoride of sil- 
 icon (see page 302). It can be prepared by treating the dry 
 
 * Compare the action of amorphous carbon (charcoal) on nitric acid 
 (note 59 of appendix) and on sulphuric acid (page 136). 
 
 t In 1881, by Jones and Taylor. See also Sabatier; Comptes Rendus; 112, t 
 865. 
 
 t By passing chlorine over an intimate mixture of boron and carbon heated 
 to redness. This method of preparation is exactly like the one employed in 
 the formation of the chloride of silicon, Si C1 4 . 
 
330 BORIC ACID. 
 
 trioxide of boron with hydrofluoric acid,* so that, in this reaction, 
 boron trioxide is a base : 
 
 B 2 3 + 6HF=2BF 3 + 3H 2 0. 
 
 Boron trifluoride, when passed into water, undergoes a decomposition 
 similar to that experienced by silicon tetrafluoride, for it breaks 
 down into boric and fluoboric acids (see page 303) : 
 
 4 BF 3 + 3 H 2 = 3 HBF 4 + B (OH ) 3 . 
 
 Fluoboric acid bears a very close resemblance to fluosilicic acid. 
 It forms fluoborates with the general formula of MBF 4 , when it is 
 brought in contact with the hydroxides of the alkali metals ; and in 
 these compounds fluorine must necessarily be considered as a biva- 
 lent element for reasons identical with those brought forward in the 
 discussion of silicofluorides on page 303. 
 
 The only oxide of boron is the trioxide, B 2 3 . The latter can 
 be formed either by burning amorphous boron, or, as is more expe- 
 dient, by heating the hydroxide B (OH ) 3 ( boric acid) to redness ; 
 the trioxide is a glass-like mass which is soluble in water and which, 
 as it volatilizes only at a very high temperature, will, when heated 
 with the salts of other volatile acids, finally decompose those salts 
 and form borates. For the same reason, fused boric acid is able to 
 dissolve the great majority of metallic oxides. 75 
 
 Boric acid ( B [OH ] 3 ) not infrequently occurs in natural deposits, 
 it being found in a crystalline state in the neighborhood of the 
 iumaroles t of Tuscany ; the acid so found is called " sassolin." 
 'The water of the lagoons in this region contains about one-tenth per 
 cent of boric acid ; this amount is increased to as much as one per 
 'Cent by collecting the water in cisterns and then allowing the vapors, 
 charged with boric acid, to condense in these receptacles ; the cistern 
 water is finally evaporated in flat leaden pans, which are warmed by 
 the steam which is escaping from the earth ; the solid residue which 
 remains, containing about 75 per cent of boric acid, is purified by 
 
 * By mixing the trioxide Bg Og with fluorspar, fusing, and then adding 
 sulphuric acid to the mass when cool. Compare this formation with that of 
 silicon tetrafluoride. 
 
 t Fumaroles are jets of water vapor which escape from fissures in the 
 earth in volcanic regions ; these vapors condense on the surface and form small 
 lagoons, which are kept boiling by the continued injection of hot vapors. 
 
BOKATES. 331 
 
 recrystallization. Boric acid is also prepared for commercial use by 
 decomposing natural borax * and borocalcite by means of acids. 
 
 Boric acid is a white, crystalline, flaky solid which, to the touch, 
 has a peculiar fatty feeling ; it is tolerably soluble in water, one 
 part of the acid being taken up by twenty-six parts of the solvent 
 at ordinary temperatures. Boron trioxide is also quite soluble in 
 alcohol; when this solution is lighted, the solvent burns with a 
 characteristic green flame. 76 f 
 
 The usual form of boric acid is orthoboric acid ; when this sub- 
 stance is heated to 100, water is given off and metaboric acid is 
 produced : 
 
 OH 
 
 (OH 
 B(OH) 3 =B0 2 H + H 2 0.* 
 
 When metaboric acid is heated to 160 it is changed to tetraboric 
 acid, which has the formula H 2 B 4 7 , its structure being on a plan 
 similar to those of di- and trisilicic acids (see pages 307 and 308). 
 The most common salt of boric acid, commercial borax, is derived 
 from tetraboric acid. Finally, when tetraboric acid is fused, all of 
 the hydroxyl groups are separated in the form of water, and the 
 anhydride, B 2 3 , is formed. 
 
 The metaborates, MB0 2 , and the tetraborates, M 2 B 4 7 , are 
 the most stable salts of boric acid. The orthoborates, M 3 B 3 , are 
 easily decomposed ; the organic derivatives of orthoboric acid are, 
 however, tolerably stable substances. 
 
 The not-metallic properties of the elements of the boron family 
 are not very p'ronounced; indeed, the element with next higher 
 atomic weight, namely, aluminium, is almost invariably metallic in 
 its nature ; so that, as a consequence, we would expect boron triox- 
 ide to act as a base under some circumstances ; and this is found to 
 be the case when we consider that the oxide dissolves in hydrofluo- 
 ric acid to form boron trifluoride, BF 3 , and is algo made apparent 
 by the existence of a phosphate of boron, BPO 4 . 
 
 * Found in large quantities in some alkaline lakes, for instance in Borax 
 Lake, California. 
 
 t This green flame is due to the formation of the triethyl ester of acid, i.e. ; 
 boric acid in which the three hydrogen atoms have been replaced by ethyl. 
 
 | Compare with the table on page 227. 
 
332 BORON NITRIDE. 
 
 Only one oxy-, or acid-chloride of boron, having a character 
 similar to the chlorides of sulphur and of phosphorus (see pages 
 156 and 221) exists. This substance has the formula B Cl ; it is 
 derived from metaboric acid by replacing the hydroxyl group with 
 chlorine ; it is decomposed by water. 
 
 The sulphide of boron in formula corresponds to the oxide ; it 
 can be formed by direct union of the elements. No sulpho-salts 
 derived from this compound are known; indeed, it is decomposed 
 with the greatest violence when brought in contact with water, so 
 that in the latter respect it resembles the sulphide of silicon. 
 
 Boron is one of the few elements that is capable of direct union 
 with nitrogen. The nitride of boron, BN, * is a white solid, formed 
 by heating an intimate mixture of borax and ammonium chloride. 
 Like cyanogen, it is a stable substance ; it is not attacked either by 
 acids or alkalies, but when heated in a current of steam it breaks 
 down into boric acid and ammonia : -^ 
 
 BN + 3 H 2 = B (OH ) 3 + NH 3 . 
 
 The occurrence of boric acid in the fumaroles of Tuscany is attrib- 
 uted to the subterranean decomposition of the nitride of boron by 
 means of water vapor. 
 
 * This nitride is therefore analogous to cyanogen, (CN)2 . 
 
ALUMINIUM ; OCCURRENCE. 333 
 
 CHAPTER XLVII. 
 
 ALUMINIUM AND ITS COMPOUNDS. 
 
 Aluminium; symbol, Al; atomic weight, 27 ; specific gravity, 2.56. 
 
 ALUMINIUM never occurs as the uncombined metal. Its oxide, 
 hydroxides, fluoride, and silicates are, however, very widely dis- 
 tributed in mineral deposits. The chief aluminium compounds 
 which are found in the form of mineral individuals are : 
 
 Coftindum; A1 2 O 3 , named ruby when found in red, transparent crystals. 
 
 Diaspor; A1O 2 H, this hydroxide, in formula, corresponds to metaboric 
 acid. 
 
 Beauxite; A1. 2 O (OH) 4 . 
 
 Hydrargyllite ; A1(OH) 3 , this substance is the normal aluminium hy- 
 droxide. 
 
 Cryolite; 3NaF,AlF 3 . 
 
 Spinell ; Mg A1 2 O 4 , a magnesium salt of the hydroxide Al O 2 H . 
 
 In addition to these oxides and hydroxides and the salts derived from 
 them, aluminium forms a large number of silicates. Among the orthosilicates 
 which contain aluminium are : 
 
 Garnet; Ca 3 A1 2 (Si O 4 ) 3 . 
 
 Muscovite; KH 2 A1 3 (SiO 4 ) 3 ; and allied to muscovite are the various 
 micas. 
 
 Kaolin (clay); H 4 Al 2 Si 2 O 9 . 
 
 Some of the most important meta-silicates, for example, hornblende and 
 augite, have already been mentioned, as have also the very important polysili- 
 cates which belong to the feldspar group (see page 308). 
 
 So extended is the distribution of aluminium in the mineral 
 kingdom that it can safely be asserted -that the element is a 
 constituent of the greater number of natural silicates. Basic 
 sulphates and phosphates of aluminium are also, not infrequently, 
 met with. 
 
 Although its compounds are so widely distributed, aluminium 
 itself was not discovered until 1827, when Wohler prepared the 
 metal by heating powdered aluminium chloride with potassium. 
 At a later date, St. Claire Deville introduced . the use of sodium, 
 which is comparatively cheap, in place of the very dear metal 
 
334 ALUMINIUM; PREPARATION. 
 
 potassium and, at a still later date, Kose improved the process by 
 using cryolite, Al F 3 , 3 KF, instead of aluminium chloride. Until 
 very recently, however, all of the aluminium of commerce was 
 prepared by heating the chloride of that metal with sodium, by 
 which reaction sodium chloride and aluminium are formed : 
 
 Al C1 3 + 3 Na = Al + 3 Na Cl. 
 
 In the last four years this expensive process has given way to 
 the electrolytic production of aluminium.* The material operated 
 on is commercially pure aluminium oxide (A1 2 3 ) , which is ob- 
 tained by chemical methods from beauxite or other aluminous 
 minerals. A bath of melted cryolite is maintained in a fused 
 condition by passing an electric current of several thousand 
 amperes through carbon cylinders hung in the melted cryolite, 
 which is itself contained in an iron tank lined with carbon. The 
 suspended rods constitute the anode and the tank the cathode. 
 When pulverized aluminium oxide is stirred into this bath it is 
 promptly dissolved, and then becomes the electrolyte (see page 13), 
 which is acted on by the current. The metallic aluminium, in a 
 melted condition, goes to the cathode, while the oxygen which is 
 at the same time separated from the decomposing oxide, burns up 
 the carbon anodes, which, in consequence, require frequent renewal. 
 By reason of this improvement in the manufacture of the metal, 
 the price of aluminium has fallen very greatly of late years. 
 
 The specific gravity of aluminium is 2.56, t its melting point is 
 700 ; the metal cannot be volatilized. Aluminium is a good con- 
 ductor of heat ; it conducts electricity almost eight times as well as 
 iron does. When covered with concentrated nitric acid, the metal is 
 transferred into a condition in which it is not further attacked by 
 the acid ; when the metal is in this state it will generate an electric 
 current when placed in contact with ordinary aluminium. A metal 
 acting in this way is said to be in the " passive state." No adequate 
 explanation of this phenomenon has as yet been given. $ Alu- 
 minium is easily attacked by hydrochloric acid. 
 
 * Personal communication from Prof. J. W. Langley, whom I wish to 
 take this opportunity of thanking most heartily.- 
 
 t This specific gravity is for cast aluminium ; hammered aluminium has a 
 specific gravity of 2.67. 
 
 J It is supposed that, in the case of aluminium, the metal becomes covered 
 with a layer of hydrogen which protects it from further attack. Diluted sul- 
 
ALUMINIUM, PROPERTIES. 335 
 
 Aluminium, because of its small specific gravity, its toughness, 
 and the difficulty with which it is attacked by the corroding agents 
 which ordinarily come in contact with a metal in general use, will, 
 in the future, have its commercial usefulness limited only by the 
 cost of its production, and, as we have seen, the latter is constantly 
 diminishing. A number of aluminium alloys are finding extended 
 application; perhaps the most important of these is aluminium 
 bronze, composed of about ten parts of aluminium to ninety parts of 
 copper ; this composition is more easily worked than ordinary bronze, 
 is tougher, is tarnished with difficulty, and has the color of gold. 
 Perfectly pure aluminium is not tarnished, either in dry or moist 
 air. 
 
 Aluminium forms but one series of compounds with the halo- 
 gens ; these compounds have the general formula A1X 3 , and of 
 these the chloride (Al C1 3 ) is the most important. Aluminium chlo- 
 ride is produced by heating powdered aluminium in a current of 
 dry chlorine. Although a solution of aluminium oxide or hy- 
 droxide in hydrochloric acid undoubtedly contains aluminium chlo- 
 ride, just as is the case in a similar solution of arsenic trioxide 
 (see page 181), nevertheless, the salt cannot be isolated by evapo- 
 rating the liquid because, as the trichloride is the halogen com- 
 pound of a metal with very weakly pronounced metallic properties, 
 it at once breaks down into the hydroxide of aluminium and hydro- 
 chloric acid : 
 
 ( Cl + HOH f OH 
 
 Al ] Cl + HOH = Al ] OH + 3 HC1. 
 
 ( Cl + HOH ( OH 
 
 Aluminium chloride is a white, crystalline solid which fumes when 
 in contact with the air ; it greedily absorbs moisture, and, while 
 giving off hydrochloric acid, changes into aluminium hydroxide ; it 
 boils at 180. The vapor density of aluminium chloride was 
 formerly supposed to be 9.34, if air is taken as unity ; this specific 
 gravity corresponds to a molecular weight of 267, and a formula 
 A1 2 C1 6 . This experimental evidence inaugurated a theory of the 
 quadrivalence of aluminium in its trichloride, for the structure of 
 
 phuric acid scarcely attacks aluminium under ordinary atmospheric pressure, 
 but if the metal is placed in a vacuum, then bubbles of gas are given off and 
 the action goes on. Similar phenomena are observed with nitric acid. (See 
 Ditte; Comptes Rendus; 110, 573.) 
 
336 ALUMINIUM; HALIDES. 
 
 the latter compound, were the molecule to have the formula A1 2 C1 6 
 would be as follows : 
 
 Cl ) ( Cl 
 
 Cl V- Al Al J Cl 
 
 Cl ) ( Cl. 
 
 A similar constitution was assigned to the trichloride of iron, 
 which substance likewise was supposed to have a molecule corre- 
 sponding to the formula Fe 2 Cl 6 . The latest investigations of 
 Nilsson and Pettersson* have shown, however, that the vapor 
 density of 9.34 found for aluminium chloride is only incidental 
 to a certain temperature,! and this specific gravity steadily dimin- 
 ishes as the heat is increased, until it reaches 4.6 at 800 ; this 
 latter number remains constant up to 1500 ; from this it is evident 
 that the molecular weight of aluminium chloride is 133.5; the 
 formula of this compound is consequently A1C1 3 ,$ so that alumin- 
 ium must be regarded as a trivalent element. Even if no vapor 
 density determinations of the chloride of iron had been made, it 
 would seem probable that the trichloride has the formula of FeCl 3 , 
 for iron can replace aluminium in isomorphous mixtures, so that a 
 difference in the valence of the two elements in compounds derived 
 from the trioxide M 2 3 seems scarcely probable. It is unnecessary, 
 however, to call to our aid any such reasoning, because the final 
 investigations regarding the vapor density of ferric chloride have 
 definitely decided that that substance, like aluminium chloride, has 
 a molecule represented by the formula Fe C1 3 . 
 
 The halogen compounds of aluminium possess the power in a 
 marked degree of forming double salts with the halides of other 
 metals. Formerly these double salts were looked upon as molec- 
 ular additions, formed of a finished molecule of some halogen 
 salt of aluminium united with the halide of an alkali metal. 
 Such a theory really means that we have no knowledge of the 
 structure of such compounds, although, in maintaining it we must 
 believe that the molecules have a residuum of chemism at their dis- 
 posal. The theory of simple molecular addition would no longer 
 
 * Zeitschrift fur Physikalische Chemie, iv., 206. t 357. 
 
 I This discovery is further borne out by the fact that certain organic deriv- 
 atives of aluminium, which have been obtained as gases, undoubtedly contain 
 that element in a trivalent form (aluminium triethyl, aluminium trimethyl, 
 aluminium acetylaceton). 
 
ALUMINIUM; DOUBLE SALTS. 337 
 
 be tenable if any of the double salts could be obtained as gases 
 with unchanged composition, and in the case of the double chloride 
 of sodium and aluminium, Deville states that the compound can be 
 vaporized without separating it into molecules of Al C1 3 and Ka Cl. 
 A theory which is of late being regarded with considerable favor is 
 the one which supposes that, in the double halides, the halogen 
 compounds of such weakly metallic elements as arsenic, antimony, 
 bismuth, or aluminium assume the role of acidic anhydrides, while 
 the halides of the alkali metals are the bases ; and, in order to 
 maintain such a theory, the assumption is inevitable that the atoms 
 of the halogens can, under certain circumstances, become divalent ; 
 such a belief is strengthened by the existence of acids like fluosi- 
 licic acid, H 2 Si F 6 , and fluoboric acid, HBF 4 (see pages 303 and 330). 
 That structures similar to those of the salts of oxy-acids are pos- 
 sessed by the double halide salts seems more than probable if we 
 consider the following: "AYhen a halide of any element combines 
 with the halide of an alkali metal to form a double salt, the num- 
 ber of molecules of the alkali salt which are added to one molecule 
 of the other halide is never greater, and is generally less, than the 
 number of halogen atoms contained in the latter." * It must be 
 confessed, however, that the rule has a number of exceptions. The 
 chlorides of aluminium and potassium, and of aluminium and 
 sodium have the formulae A1C1 3 , KC1, and A1C1 3 , XaCl, and if we 
 regard A1C1 3 as analogous to an acidic anhydride, and KC1 and 
 Nad as analogous to bases, the structure of these compounds 
 would be as follows : 
 
 ^Jv and 
 C1 2 K 
 
 the parallelism between these formulae and those of the correspond- 
 ing oxy-compounds becomes apparent when we consider the structure 
 of the latter ; namely : 
 
 The double fluorides of aluminium are of two kinds, the first of 
 
 ( -p 
 which, with the general formula Al -j -p 2 iyp are constructed similarly 
 
 to the chlorine compounds ; while, second, with the general formula 
 
 * Remsen; American Chemical Journal; 14, 85. See also Remsen; Chem- 
 istry, p. 461. 
 
338 ALUMINIUM OXIDE. 
 
 (F 2 M (OM* 
 
 Al J F 2 M, correspond to the theoretical ortho-aluminates, Al J OM 
 
 (F 2 M (OM. 
 
 In addition to the substances which have just been discovered, a 
 number of compounds, in jwhich aluminium chloride is a base, are 
 known; an example would be the compound A1C1 3 , PCl 5 .f The 
 existence of these double chlorides once more reminds us of the 
 great resemblance between the chemistry of the halogens and that 
 of oxygen (see pages 62, 63, and 157 ). 
 
 Aluminium forms but one oxide, the trioxide A1 2 3 . This sub- 
 stance occurs as the mineral corundum, which, when finely divided 
 and mixed with oxide of iron, is called emery. The transparent, 
 red crystals of the oxide are called ruby. Aluminium trioxide is 
 produced when aluminium is heated to a high temperature in air or 
 in oxygen, or when the hydroxide is heated ; the latter substance, 
 because it is insoluble in water, is precipitated from solutions of 
 aluminium salts by adding ammonia water : 
 
 AlCl 3 -t-3NH 3 + 3H 2 = Al (OH ) 3 + 3 NH 4 Cl ; 
 and when heated : - 
 
 2 Al (OH ) 3 = A1 2 3 + 3 H 2 0. 
 
 The oxide which has been heated to redness, or the naturally occur- 
 ring crystalline varieties, are insoluble both in acids, water, or solu- 
 tions of the alkalies ; they can be brought into solution by fusing 
 with caustic alkalies. The oxide which has not been heated is both 
 basic and acidic, for it dissolves in acids to form the salts of alu- 
 minium, and in bases to form aluminates. 
 
 Several hydroxides are derived from aluminium trioxide. The 
 first of these, the normal hydroxide Al (OH ) 3 , is precipitated from 
 solutions of aluminium salts by ammonia water or caustic alkalies ; 
 in using the latter reagents, however, care must be taken not to add 
 an excess, otherwise solution of the hydroxide with the formation 
 of an aluminate takes place. $ The remaining hydroxides are de- 
 
 * In generalizing in regard to these double halides we must except the 
 double cyanides (page 296). 
 
 t Compare the formulae Al P C1 8 and Al PO 4 . 
 
 | Excepting with excess of ammonia solution, for aluminium hydroxide 
 is but little soluble in that reagent. Precipitation of aluminium hydroxide 
 from aluminium salts by means of an excess of ammonia water is complete, if 
 the precipitate and liquid are boiled until the odor of ammonia disappears. 
 
ALUMINATES; SPIKELLS. 339 
 
 rived from the normal compound by loss of water ; the most impor- 
 tant of these is the meta-hydroxide, Al 2 H,* which is found as the 
 mineral diaspor. The salts of this substance, formed by replacing 
 the hydrogen with a metal, are the types of the important group of 
 minerals known as the spinells. The spinells are a group of isomor- 
 phous compounds, each one of which is derived from a hydroxide, 
 
 X -j QTT in which hydroxide X can be either trivalent aluminium, 
 
 iron, or chromium ; the hydrogen of this hydroxide is replaced by a 
 divalent metal, M", so that the general formula for these minerals 
 would be M" (X 2 ) 2 ; M" is either divalent iron, magnesium, man- 
 ganese, or zinc. The typical spinell is the aluminate of magnesium, 
 Mg (A10 2 ) 2 . When meta-aluminium hydroxide is dissolved in 
 caustic alkalies, the solution, unless it is a very concentrated one 
 containing an excess of the solvent, contains the meta-aluminate 
 of the particular alkali metal used : 
 
 Al 2 H + MOH = Al 2 M + H 2 0. 
 
 The aluminates of the alkali metals can be obtained in a crys- 
 tallized state by evaporating a solution of aluminium hydroxide in 
 an excess of concentrated caustic alkali. 
 
 In addition to the ortho- and meta-hydroxides, another hydroxide 
 of aluminium, A1 2 5 H 4 , is frequently met with as the mineral 
 beauxite ; this substance is formed by the separation of water 
 between two formula weights of ortho-aluminium hydroxide (see 
 page 307 ). 
 
 The normal hydroxide of aluminium, when freshly precipitated, 
 is a gelatinous substance which readily dissolves in acids to form 
 the salts of aluminium ; among these, perhaps, the most important 
 are the sulphates. 
 
 ALUMINIUM SULPHATE. A1 2 (SO 4 ) 3 + 8 H 2 O; formed by dissolving alu- 
 minium oxide or hydroxide in sulphuric acid and evaporating to dry- 
 ness. When heated it loses its water of crystallization at 100, and at 
 red heat gives off sulphur trioxide, leaving aluminium oxide behind : 
 A1 2 (SO 4 ) 3 = A1 2 O 3 + 3 SO 3 . 
 
 ALUMINIUM AND POTASSIUM SULPHATE. A1 2 (SO 4 ) 3 , K 2 SO 4 + 24 H 2 O. 
 This substance is commonly known as alum. The alums are double 
 salts composed of one formula weight of the sulphate of a monovalent 
 alkali metal or of ammonium, combined with the sulphate of a trivalent 
 
 * Corresponding in formula to metaboric acid, BO 2 H. 
 
340 ALUMINIUM; SALTS OF. 
 
 metal, the general formula being M'" 2 (SO 4 ) 3 , M 2 SO 4 + 24 H 2 O. 
 M'" can be either aluminium, iron, chromium, or one of the rarer 
 metals belonging to the boron family; M can be any one of the alkali 
 metals, or ammonium. The alums are all isomorphous and crystallize 
 in octahedra belonging to the regular system. They can be formed 
 by evaporating to dryness a mixture of the solutions of the sulphates 
 of any one of the trivalent metals mentioned and of one of the alkalies. 
 
 A number of basic sulphates of aluminium are known ; some of 
 these are found in the form of mineral deposits. Several neutral 
 and basic phosphates of aluminium occur in nature ; perhaps the 
 most important of these is wavellite, 2 Al P0 4 , Al (OH) 8 + 9 H 2 0. 
 When sodium phosphate is added to the neutral solution of an 
 aluminium salt, the tertiary phosphate of aluminium, A1P0 4 , is 
 produced in the ibrm of an insoluble precipitate ; this is changed to 
 a basic phosphate by boiling with ammonia water. 
 
 As has been mentioned, aluminium is a constituent of a very 
 large number of silicates, some of which have already been dis- 
 cussed. Of these the most important is undoubtedly the hydrated 
 tertiary orthosilicate which approximately has the composition 
 expressed by the formula H 2 A1 2 (Si0 4 ) 2 -j- H 2 and which is 
 known as clay or kaolin. This substance is the result of the dis- 
 integration of feldspar, or of rocks which contain a large proportion 
 of that mineral (as some granites do) ; owing to the destructive 
 action of the weather, the feldspar decomposes into aluminium sili- 
 cate, silicon dioxide, and the silicates of the alkali metals ; the 
 latter are washed away, brought into the soil, and there, after in- 
 teracting with other chemical constituents with which they come in 
 contact, they are in a proper condition to be absorbed by plants. 
 The kaolin which remains on the spot where disintegration occurs, 
 owing to the formation of silicon dioxide during the process of de- 
 struction, is necessarily mixed with that substance and not infre- 
 quently contains mica ; some clays, however, are washed to some 
 distance from their place of formation, and these may have taken 
 up the most varied impurities, such as the carbonates of calcium 
 and magnesium, the oxides of iron and of manganese, quartz sand, 
 and other materials. 
 
 An impure clay, not infrequently colored by ferric oxide, is 
 used in the manufacture of bricks ; it is pressed into moulds and 
 baked in a kiln until it becomes hard. Pure kaolin is white in 
 color and is used in the manufacture of porcelain ; when it is moist 
 
CLAY ; PORCELAIN. 341 
 
 it forms a very plastic mass. Kaolin which is entirely free from 
 iron is alone useful in the manufacture of porcelain because, as the 
 mass is heated to the point where it softens and becomes glassy, 
 any foreign substances would make themselves apparent by their 
 color. The clay is purified by being agitated with water ; when the 
 coarser portion separates at the bottom, the finer parts are moulded 
 into forms, dried, and finally heated to a red heat. This latter 
 treatment makes the article undergoing the process of manufacture 
 strong, but leaves it porous; in order to finish the same it is covered 
 with a mixture of silicon dioxide, aluminium oxide, and sodium car- 
 bonate, which ingredients form an easily fusible glass, and it is then 
 heated to a temperature at which the clay begins to soften and at which 
 the glazing has been converted into a coating of transparent glass. 
 Fayence is made of clay of somewhat coarser structure than that 
 used in the manufacture of porcelain ; the thickness of the dishes 
 is greater, and the ware is not heated to a temperature high enough 
 to convert it into the glass-like mass which forms porcelain. Fay- 
 ence is covered with a glaze which is much like that given above, 
 with the difference that some lead oxide is added. Common stone 
 ware is made from clay which is even more impure than that from 
 which fayence is prepared ; the glazing is either put on by covering 
 the ware with the mixture to be used before burning, or it is made 
 by throwing common salt on the utensils while they are being 
 heated in the furnace ; the salt evaporates, comes in contact with the 
 surface of the materials used, and covers it with a fusible soda glass. 
 But one sulphide of aluminium, with a formula A1 2 S 3 , corre- 
 sponding to the oxide, is known. This compound can be produced 
 by heating a mixture of powdered aluminium and sulphur, but, like 
 many sulphides of the not-metals, it is readily decomposed by water, 
 yielding the hydroxide of aluminium and hydrogen sulphide : 
 
 A1 2 S 3 + 6 H 2 = 2 Al (OH ) 3 -f 3 H 2 S ; 
 
 from this it follows that when an alkaline sulphide is added to a 
 solution of an aluminium salt, aluminium hydroxide is precipitated ; 
 so, for instance, the following reaction takes place between alumin- 
 ium sulphate and ammonium sulphide : 
 
 3 (NH 4 ), S + A1 2 ( S0 4 ) 3 -f 6 H 2 = 2 Al (OH ) 3 + 3 (NH 4 ) 2 S0 4 
 
 + 3 H 2 S. 
 A similar reaction takes place when a soluble carbonate is added to 
 
342 ALUMINIUM SULPHIDE. 
 
 the solution of an aluminium salt, for, owing to the extreme insta- 
 bility of aluminium carbonate, the hydroxide and not the carbonate 
 is precipitated : 
 
 3 (NH 4 ) 2 00 3 + A1 2 (S0 4 ) 3 + 6 H 2 = 2 Al (OH) 3 + 3 (NH 4 ) 2 
 S0 4 + 3 H 2 + 3 C0 2 . 
 
 All of the reactions which have just been mentioned illustrate 
 the weakly basic character of aluminium oxide and hydroxide and 
 show the close relationship existing between aluminium and the not- 
 metals. 
 
GALLIUM. 
 
 CHAPTER XLVIII. 
 
 GALLIUM, INDIUM, AND THALLIUM. 
 
 Gallium; symbol, Ga ; atomic weight, 69 ; 
 Indium; symbol, In ; atomic weight, 113.7 ; 
 Thallium; symbol, Tl ; atomic weight, 204.18. 
 
 GALLIUM, indium, and thallium are very sparingly represented in 
 nature ; they are of scarcely any commercial importance, so that the 
 interest in them is purely theoretical in its character, and is taken 
 because they complete the family of elements of which boron and 
 aluminium are the chief representatives. 
 
 Gallium was discovered by Lecocq de Boisbaudran in zinc- 
 blende.* The metal is hard, brittle, and crystalline in its structure ; 
 it is scarcely malleable or ductile ; it melts at 30. 15 and is not 
 volatile even at a high temperature ; its specific gravity is 5.95. The 
 metal scarcely oxidizes when exposed to the air, and it is readily 
 obtained by electrolysis of a solution of the oxide in alkalies ; it de- 
 composes steam and liberates hydrogen (see page 30) ; like alumin- 
 ium, it is soluble in hot, caustic alkalies. The chief characteristics 
 of the compounds of gallium are given in the following table : 
 
 OXIDES, Ga O, Ga 2 O 3 . The former is the least stable of the oxides; it is 
 basic; the latter is white, infusible, reduced to the metal at white 
 heat by a current of hydrogen. It is both basic and acidic, but dis- 
 solves only in the most concentrated caustic alkalies. 
 
 HYDROXIDE, Ga (OH ) 3 , formed by precipitating the solutions of soluble 
 gallium salts with ammonia water; it is somewhat soluble in an excess 
 of the reagent. When heated, the hydroxide readily loses water and 
 forms the oxide. 
 
 CHLORIDES, Ga C1 2 , Ga C1 3 . The former is a solid which melts at 164 
 and boils at 535; its vapors have a specific gravity which corresponds 
 to a molecular weight represented by the formula Ga C1 2 . The latter 
 was formerly supposed to have the formula Ga 2 C1 6 (see page 336), but 
 recent determinations of the specific gravity of the vapor of gallium 
 trichloride show that body to have a molecular weight corresponding 
 
 * Zinc sulphide, Zn S. 
 
344 INDIUM. 
 
 to the formula Ga C1 3 * at 440. The trichloride melts at 75. 5 and 
 boils at 215 to 220. It dissolves in water, but when the solution is 
 evaporated it is, in part, changed into the basic chloride, just as is the 
 case with the trichloride of antimony (page 253). 
 
 THE SALTS OF GALLIUM are produced by dissolving the hydroxide in the 
 various acids; the sulphate, when evaporated with the sulphates of 
 the alkali metals or of ammonium, forms alums. (See page 339.) 
 
 The discovery of gallium in 1875 was of especial interest because, 
 in the periodic system of the elements as arranged by Mendelejeff 
 a few years previous to that time, an element belonging to the 
 family of which aluminium is a representative was found to be 
 missing.t An element, the chemical and physical properties of 
 which should lie between those of aluminium and indium, and 
 which would have an atomic weight of approximately 69, was 
 therefore predicted by Mendelejeff, and this prediction was subse- 
 quently brilliantly verified by Lecocq de Boisbaudran. 
 
 The next element of this family is indium. Like gallium, it 
 occurs in some specimens of zinc-blende. The element was dis- 
 covered in 1863 by fceich and Eichter. It is white, with a metallic 
 lustre lying between that of platinum and silver ; it is softer than 
 lead, and is very malleable and ductile. Its specific gravity is 7.4 ; 
 it melts at 176, and is somewhat volatile at red heat ; when 
 heated to redness in the air, it burns to form In 2 3 . The chief 
 properties of its compounds are given in the following table : 
 
 OXIDES, In O and In. 2 Os . The former is made by reducing the trioxide 
 in a current of hydrogen ; it burns in the air to form In 2 Oa . The 
 trioxide, In 2 O 8 , is the most stable oxide of indium, and is the one 
 corresponding to the typical oxide of the family ; it is produced when 
 indium is burned in the air, or when the hydroxide In (OH ) 3 is heated ; 
 this latter substance is obtained by precipitation from solutions of in- 
 dium salts by means of ammonia water. The oxide is easily reduced 
 to the metal by heating the same in a current of hydrogen, the monox- 
 ide appearing as an intermediary stage in this reduction. The trioxide 
 and the corresponding hydroxide are mainly basic in their character; 
 they dissolve in acids to form the salts of indium ; they are, however, 
 also weakly acidic, for they are dissolved by the hydroxides of potas- 
 sium or sodium. 
 
 CHLORIDES, In C1 2 , In C1 3 . The latter is formed by the action of chlo- 
 
 * Nilssen and Petterson; Comptes Rendus; 107, 572. 
 t A similar gap was found to exist in the carbon family; this was subse- 
 quently filled by the discovery of germanium (see page 309). 
 
THALLIUM. 345 
 
 rine on indium ; it sublimes at 440 without melting; its vapor den- 
 sity corresponds to a molecule of the formula In Cls ; it dissolves in 
 water without change, but, on heating the solution, hydrochloric acid 
 passes off and a basic chloride is formed. Indium chloride readily 
 unites with the chlorides of the alkali metals to produce double salts 
 corresponding in formula to those of aluminium (see page 337). 
 
 SULPHIDE, In 2 S 3 , is formed by direct union of indium and sulphur at red 
 heat. A sulphohydrate of indium, In(SH)g, is precipitated from 
 neutral or weakly acid solutions of indium salts b$ hydrogen sul- 
 phide; in this respect the character of indium approaches that of the 
 most pronounced metals of the preceding (carbon) family. 
 
 THE SULPHATE OF INDIUM, when evaporated with the sulphate of an 
 alkali metal or of ammonium, produces an alum (see page 
 
 Thallium, the element having the highest atomic weight in this 
 family, was discovered by Crookes in 1861, that investigator find- 
 ing it in the residues covering the floors of the channels of certain 
 sulphuric acid works ; since that time it has been discovered in 
 zinc-blende, iron pyrites, and copper pyrites. The metal is white, 
 of crystalline structure, and greatly resembles tin ; it is malleable 
 and ductile, has a specific gravity of 11.9, melts at 290, and boils 
 at a white heat. Thallium is readily oxidized in the air, and dis- 
 solves in sulphuric or nitric acid without much difficulty. The 
 high atomic weight of thallium is unfavorable to the expression of 
 a very pronounced chemical character, so that, as is the case with 
 lead, it appears with a number of oxides which each, individually, 
 resemble a different group of elements ; for instance, the monoxide, 
 T1 2 0, is, in its chemical behavior, very much like the oxides of the 
 monovalent elements of the alkali family, while the trioxide, T1 2 3 , 
 falls into line with the similar compounds of the aluminium group. 
 The characteristics of the most important thallium compounds are 
 given below : 
 
 OXIDES, T1 2 O, T1 2 O 3 , Tl O 2 . The first, thallium monoxide or thallous 
 oxide, is formed by the slow oxidation of the metal in the air; it is a 
 brownish-black powder, which is soluble in water, forming the hydrox- 
 ide T1OH; this remarkable solubility shows the resemblance between 
 this oxide and the oxides of the alkali metals, for the latter are like- 
 wise soluble in water; a solution of thallous hydroxide has a strongly 
 alkaline reaction; it neutralizes acids to form salts, which, for the 
 most part, are soluble in water [resemblance to the salts of the alkalies 
 (see latter)]. Thallium trioxide (thallic oxide), formed by heating 
 thallium to a red heat in oxygen, is insoluble in water; the hydroxide, 
 Tl O. 2 H, is formed by precipitating from a solution of a thallium salt 
 
346 THALLIUM. 
 
 by means of ammonia water; neither the oxide nor hydroxide has 
 acidic properties; both are oxidizing agents, having a great tendency 
 to change into thallous oxide. A higher oxide of thallium, Tl 0% , so- 
 called thallic acid, is also said to exist. 
 
 CHLORIDES, Tl Cl, Tl C1 3 . The first, thallous chloride, is insoluble in 
 water, is precipitated from solutions of thallous salts by hydrochloric 
 acid, and very much resembles the chloride of silver in appearance. 
 (A larger work must be consulted for a description of the thallous 
 salts.) 'The trichloride, T1C1 8 , formed by treating thallium with 
 chlorine, is decomposed into thallous chloride and chlorine when 
 heated. 
 
 SULPHIDES. Two sulphides of thallium, T1 2 S and T1 2 S 3 , are known. 
 
ATOMIC WEIGHTS. 347 
 
 CHAPTER XLIX. 
 
 THE DETERMINATION OF ATOMIC WEIGHTS. DULONG AND 
 PETIT'S LAW. THE LAW OF ISOMORPHISM* 
 
 THE investigations into the gravimetric composition of chemical 
 compounds, which were undertaken at the beginning of the century 
 and which finally developed the laws of definite and multiple pro- 
 portions, succeeded not only in establishing these purely empirical 
 laws, but, as the spirit of inquiry in man leads him to seek a cause 
 behind every regularly recurring phenomenon or law, naturally an 
 explanation for the laws of definite and multiple proportions was 
 looked for, and, as we have seen, found in the atomic hypothesis. 
 In spite of the subsequent almost universal acceptance of these 
 laws, some chemists continued to doubt their exactness ; indeed, the 
 gravimetric determinations of Dalton's time were too unsatisfac- 
 tory and varying to inspire much confidence ; when, at a later date, 
 Berzelius subjected the work, which had been done with the purpose 
 of establishing the laws relating to the definite composition of 
 matter, to a more exact revision, and so became a firm believer in 
 their existence, no adequate reason for their non-acceptance by 
 other chemists could be advanced ; but, when the subsequent discov- 
 ery of an error in Berzelius's determination of the relative weight 
 with which carbon enters into combination with other elements, 
 shook confidence in all of the established rules, the atomic theory 
 was left in a. most unsatisfactory condition, and it was not until 
 1860 that the painstaking and accurate work of Stas succeeded in 
 showing that the laws of definite and multiple proportions are not 
 merely approximate, but are, in reality, mathematically exact. How- 
 ever, the atomic hypothesis had existed in its present form before 
 Stas's time, and it naturally was the endeavor of chemists to deter- 
 mine, not only the mere fact that, for instance, a parts by weight 
 
 * See also Lothar Meyer; Die Grundziige der Theoretischen Chemie; Leip- 
 zig, 1890. (Outlines of Theoretical Chemistry, trans, by Bedson and Williams; 
 Longmans.) 
 
348 ATOMIC WEIGHTS BY AVOGADKO'S HYPOTHESIS. 
 
 of chlorine always unite with b parts by weight of silver to form 
 the chloride of silver, they also endeavored, as we have seen in the 
 preceding portions of the work, by calling to their aid various hy- 
 potheses and theories of greater or less plausibility, to fix exactly 
 the relative weights of the atoms of silver and chlorine ; these 
 atomic weights must necessarily bear such a relationship to each 
 other that, in uniting to form silver chloride, they would always 
 produce that substance with the proportion of a parts by weight of 
 chlorine to b of silver. It is evident that we can only determine 
 the atomic weights from the stoi'chiometric quantities, provided we 
 have some means of knowing the number of atoms united to form 
 the molecules. This, however, it is not possible to do by direct 
 observation, so that, in selecting atomic weights, we must resort to 
 more or less probable hypotheses. Gravimetric determinations 
 alone, therefore, can do no more than give the relative parts by 
 weight with which two or more substances unite to form a chem- 
 ical compound; thus, in studying the composition of water by 
 weight, we could but determine that eight parts by weight of oxy- 
 gen unite with one part of hydrogen ; it was only after chemists 
 combined the lesson taught by the gravimetric composition of water 
 with the phenomena attendant on its formation and decomposition, 
 i.e., with the facts that two volumes of hydrogen always unite with 
 one volume of oxygen to form water, and that water, when decom- 
 posed, always yields two volumes of hydrogen and one of oxygen, 
 and after these experimental facts were explained by Avogadro's 
 hypothesis (page 70 and sub.), that the conclusion was definitely 
 reached that each molecule of water, is, in reality, composed of two 
 ;atoms of hydrogen united to one of oxygen, by reason of which 
 conclusion the atomic weight of oxygen was fixed at 16 and not at 8. 
 As has been repeatedly mentioned, the determinations of the 
 specific gravities of gases, if we accept Avogadro's hypothesis, give 
 us the magnitudes of their molecular weights, and, when these are 
 fixed, provided, in each case the accurate sto'ichiometric * compo- 
 sition of the substance in question is known, t we can determine the 
 
 * The relative proportions by weight in which substances unite to form 
 chemical compounds are called the stoichiometric quantities. (See page 6.) 
 
 t It is evident that the determination of the vapor density of a substance 
 is of no value unless the stoichiometric composition is known. For instance, 
 it is of no influence on the determination of the atomic weight of nitrogen to 
 
MAXIMUM ATOMIC WEIGHTS. 
 
 349 
 
 maximum atomic weights of the elements entering into the struc- 
 ture of the molecules of the various gases. (See page 73.) The 
 following table, which illustrates the method by which maximum 
 atomic weights are determined from a comparison of the vapor 
 densities of gases, will serve to more clearly fix these facts in the 
 mind of the pupil : 
 
 NAME OF GAS. 
 
 d. 
 
 d X 28.8. 
 
 M. 
 
 ANALYSIS BY WEIGHT. 
 
 Nitric oxide. 
 
 1.039 
 
 30. 
 
 30.03 
 
 14.03 nitrogen' + 16 oxygen. 
 
 Nitrogen dioxide. 
 
 1.58 
 
 45.5 
 
 46.03 
 
 14.03 " + 32 " 
 
 Phosphorus trichloride. 
 
 4.88 
 
 140.9 
 
 137.35 
 
 31. phosphorus + 106.35 chlorine. 
 
 Phosphorus tri-iodide. 
 
 14.46 
 
 417.1 
 
 411.55 
 
 31. " + 380.55 iodine. 
 
 Hydrochloric acid. 
 
 1.247 
 
 36. 
 
 36.458 
 
 1.008 hydrogen + 35.45 chlorine. 
 
 Hydroiodic acid. 
 
 4.443 
 
 128. 
 
 127.858 
 
 1.008 " + 126.85 iodine. 
 
 Water. 
 
 0.623 
 
 17.99 
 
 18.016 
 
 2.016 " + 16 oxygen. 
 
 Sulphur dioxide. 
 
 2.247 
 
 64.9 
 
 64.06 
 
 32.06 sulphur + 32 
 
 In this table d is the specific gravity, air = 1 ; d x 28.8 is the specific grav- 
 ity, H 2 = 2 ; * M is the molecular weight, found by adding the figures given in 
 the last column. 
 
 After a study of the above table we can see that, were chlorine, 
 for example, to occur in but one compound, and that one the tri- 
 chloride of phosphorus, evidently the atomic weight of the element 
 would be placed at a maximum of 106.35 ; it could be no greater, 
 for the molecular weight of the chloride of phosphorus is known. 
 However, if we glance further down the column, we discover ny- 
 drochloric acid with a molecular weight of 36.457 ; each molecule 
 of this contains but 35.45 parts by weight of chlorine ; it follows 
 therefore that the weight of 106.35, which is the amount of chlo- 
 rine contained in the chloride of phosphorus, really represents 
 three atoms of chlorine, provided the molecule of hydrochloric acid 
 contains lut one of these. The number 35.45 must therefore be 
 fixed upon as the atomic weight of chlorine, and must remain so, 
 
 determine the specific gravity of ammonia and find this to be .589, and that 
 the molecular weight is 17; we must also know that this 17 parts by weight of 
 ammonia contains 14 parts by weight of nitrogen ; when, however, the two facts 
 are combined, we can say that the atomic weight of nitrogen cannot be more 
 than 14, for, after we have ascertained the specific gravity of ammonia, we 
 then have proof that there is a substance, the molecular weight of which is 
 not greater than 17, which contains but 14 parts of nitrogen. 
 
 * The specific gravity of air (H. 2 = 2) is 28.8, hence specific gravities taken 
 with air as unity are converted to those with H 2 = 2 by multiplying by 28.8. 
 
350 DETERMINATION OF MOLECULAR WEIGHTS. 
 
 unless, at some future time, we were to discover a compound of 
 chlorine, the molecular weight of which is known and which con- 
 tains relatively less than this quantity. In the latter event, a 
 molecule of hydrochloric acid would necessarily contain more than 
 one atom of chlorine. Similar considerations will help us to select 
 the number 126.85 as representing the atomic weight of iodine, 
 while a comparison of the figures in the table given above will 
 further show us that, by means of the determinations even of the 
 very few substances mentioned there, the maximum atomic weights 
 of nitrogen, phosphorus, sulphur, oxygen, chlorine, and iodine are 
 given. These same methods of investigation have been applied in 
 every case where the study of elements and compounds in the 
 gaseous state has been possible, so that the maximum atomic 
 weights of the greater number of elements have been ascertained 
 with reasonable certainty. 
 
 The determination of the specific gravities of gases, although by 
 far the most important, is not the only method for ascertaining 
 molecular weights.* In 1882 F. M. Raoult demonstrated that 
 aqueous solutions of organic substances, provided they contain the 
 dissolved compounds in quantities proportional to their molecular 
 vughts, have identical freezing points, and, subsequently, the same 
 law was found to hold good for other substances as well, although 
 the amount of depression differs for each solvent. If A is the low- 
 ering of the freezing point of a solvent, brought about by the solu- 
 tion of n molecular weights of a certain substance in g grams of the 
 solvent, then : 
 
 1. A = r-, 
 
 9 
 
 where r is a constant depending only on the nature of the solvent. 
 When the molecular weight of the substance is not known, this can 
 be ascertained by experimentally determining the lowering of the 
 freezing point of the solvent brought about when p grams of the 
 
 substance are dissolved in g grams of that solvent, for then n=^-, 
 
 M 
 where M is the molecular weight, so that equation 1 becomes : - 
 
 2. A = ^,or, 
 
 Mg 
 
 * See Ostwald; Outlines of General Chemistry (Walker, page 136). The 
 description of the method in that book is here given. 
 
DETERMINATION OF MOLECULAR WEIGHTS. 351 
 
 3. M^ZL. 
 
 *</ 
 
 The constant, r, can be determined once and for all, for any 
 given solvent, by dissolving one or two substances of known molec- 
 ular weight therein, and observing the depression of the freezing 
 point ; for, from equation 1 : 
 
 The objection to this method of determining molecular weights 
 lies in the fact that it is applicable only in the limited number of 
 cases where substances are soluble in a medium which is capable 
 of being frozen at a convenient temperature, while, furthermore, the 
 law has not held good in a number of cases which have been 
 observed.* The application of this method is, nevertheless, very 
 valuable to confirm molecular weights determined by some other 
 means, and to ascertain the same in many cases where the determi- 
 nation of the specific gravities of gases is impracticable. Similar 
 methods, based upon the lowering of the vapor pressure of solutions 
 by reason of dissolved substances, have also been applied to the 
 determination of molecular weights ; for their study the pupil must 
 refer to some text-book more especially devoted to these subject.- Y 
 
 The atomic weights which, by the application of Avogadro's 
 hypothesis ( page 70 ) and the determination of the vapor densities 
 of elementary and compound substances, have been selected as those 
 which really represent the relative weights of the individual atoms, 
 would be much more worthy of confidence if other physical meas- 
 urements could, when correctly interpreted, indicate that this 
 selection had been properly made. Such an aid is found in the 
 application of a law discovered by Dulong and Petit in the early 
 part of this century. These two investigators proved that the 
 
 * Salts are in part decomposed by solution in the same direction as they 
 are by the electric current; so, for example, each molecule of potassium, so- 
 dium, or ammonium chloride is dissociated into a univalent not-metallic atom 
 (an anion), and a univalent metallic atom (or group such as NH 4 ) (a cathion) 
 by great dilution ; the dissociation in the case of these salts is almost a complete 
 one; it is less in the case of sulphates, etc. Of course the lowering of the 
 freezing point with solutions of such substances [so-called electrolytes (see 
 page 13 and 28)] does not give results indicative of the molecular weights. See 
 Harry C. Jones; Zeitschrift fiir Physikal. Chemie; xi., 535. 
 
 t Ostwald; Outlines of General Chemistry; Walker's translation. 
 
352 
 
 LAW OF DTJLONG AND PETIT. 
 
 greater the combining weight of any given element was found to 
 be, the less was the specific heat (capacity for heat) of that element 
 in the solid form, so that the product of the combining weight of any 
 given element and its specific heat, was found to be very nearly equal 
 to the same product for any other element. This law is susceptible 
 of a very simple physical explanation. The specific heat of a body 
 is the quantity of heat necessary to increase the temperature of the 
 unit weight of that body by 1 ; it follows that the product of the 
 specific heat and the combining weight of an element represents 
 the quantity of heat necessary to warm the combining weight 
 through 1, and if we select those numbers which, by using Avoga- 
 dro's hypothesis, we have decided upon as being the atomic weights, 
 then the above-mentioned product is, in the great majority of cases 
 very nearly 6.4 ; this product can be termed the " atomic heat " of 
 the elements, so that Dulong and Petit' s law assumes the following 
 simple form : 
 
 The atomic heats (capacities for heat of the individual atoms) 
 of all elements are very nearly equal* 
 
 This law is true without exception with all perfect (ductile) 
 metals ; it holds good with almost all metals which, like antimony or 
 tin, also have not-metallic properties (which are brittle, but which 
 have metallic lustre), and also applies to the greater number of not- 
 metals. A few examples will serve as an illustration : 
 
 ELEMENT. 
 
 c 
 
 a 
 
 ac 
 
 
 Lithium. 
 
 .941 
 
 7.02 
 
 6.6 
 
 In this table: 
 
 Magnesium. 
 Chromium. 
 Iron. 
 
 .250 
 .121 
 .114 
 
 24.3 
 52.1 
 
 56. 
 
 6.07 
 6.3 
 6.38 
 
 c = specific heat, 
 a = atomic weight, 
 ac = atomic heat. 
 
 Cobalt. 
 
 .107 
 
 59. 
 
 6.3 
 
 
 Nickel. 
 
 .108 
 
 58.7 
 
 6.3 
 
 
 Bromine. 
 
 .084 
 
 79.95 
 
 6.7 
 
 
 Gold. 
 
 .032 
 
 197.3 
 
 6.3 
 
 
 If we are, therefore, acquainted with an element, the maximum 
 
 * For a more extended exposition of the results obtained by Dulong and 
 Petit's law, see Lothar Meyer; Die Grundzuge der Theoretischen Chemie; 
 Leipzig, 1890. (Outlines of Theoretical Chemistry, trans, by Bedson and 
 Williams; Longmans.) 
 
EXCEPTIONS TO LAW OF DULONG AND PETIT. 353 
 
 atomic weight of which we have never been able to determine by 
 means of the vapor densities of some of its compounds, we can, as a 
 next resort, select as its true atomic weight that stoichiometric 
 quantity which, when multiplied by the specific heat of the ele- 
 ment in question, will give us a number approximately equal to 6.4. 
 Of course, were no method available for fixing atomic weights as 
 valuable as that given us by the determinations of the specific grav- 
 ities of gases, this law of Dulong and Petit would lead to no 
 definite results ; for it is obvious, when we consider the above table, 
 that were we to select as atomic weights numbers exactly one-half 
 as large as those given above, the product ac would still remain 
 constant and would be approximately equal to 3.2 ; so that only if, 
 by the use of other physical and chemical means, we have deter- 
 mined a certain number of atomic weights, and have definitely de- 
 cided that in the case of the elements in question, that product must 
 be 6.4, and not 3.2, it then follows that in all undetermined cases 
 
 the atomic weights must be nearly equal to the quotients of . 
 
 G 
 
 Naturally, the numbers so obtained are not the exact weights ; they 
 can only be close enough to the true numbers to show us that the 
 atomic weights to be selected are not one-half or twice or three 
 times the stoichiometric quantities which come closest to those 
 indicated by dividing 6.4 by the respective specific heats. 
 
 The few marked exceptions to Dulong and Petit's law are not 
 calculated to give any great difficulty, for, in cases where they 
 occur, the elements in question form a number of gasifiable com- 
 pounds, and consequently there is no reason why the specific grav- 
 ities of these compounds should not definitely fix their maximum 
 atomic weights. One example of such an exception will serve to 
 illustrate this conclusion. The chemical equivalent weight of car- 
 bon is 3,* its specific heat as diamond is. 1.47, as graphite, 1.98. 
 The atomic weight of carbon must be some rational multiple of its 
 
 . * By the chemical equivalent weight of an element is meant that quantity 
 of the element which will enter into combination with one part by weight 
 (one atom) of hydrogen, or which will take the place of one atom of hydrogen 
 in a chemical compound. The equivalent weight of carbon is 3, because three 
 parts by weight of carbon unite with one of hydrogen. The atomic weight of 
 carbon is, however, 4x3 = 12, for, by a determination of the specific gravity of 
 the gas, it has been decided that a molecule of the simplest hydrogen compound 
 of carbon has the formula CH 4 . 
 
354 MOLECULAR HEATS OF COMPOUNDS. 
 
 equivalent weight, so that, according to the rule, we must select a 
 number as the atomic weight of carbon which, in addition to being 
 a rational multiple of the equivalent weight, will also, when multi- 
 plied by the specific heat of carbon, give a product approximately 
 equal to 6.4. This product is reached in the case of diamond when 
 we take fourteen times the equivalent weight, and therefore it would 
 follow that the atomic weight of carbon is 42 ; in the case of 
 graphite, however, we would come to a different result, for then 
 ac = 6.5 when a = 33. These atomic weights are entirely impos- 
 sible, however, for we are acquainted with a number of carbon com- 
 pounds in which (the vapor density and hence the molecular weight 
 being known) we relatively have only 12 parts by weight of carbon, 
 so that there remains no alternative but to believe that carbon pre- 
 sents an exception to 'Dulong and Petit' s law. However, the spe- 
 cific heat of carbon is much greater at high temperatures than it is 
 at low ones, at 900 it is .459 (ac would then be equal to 5.51), so 
 that, at a white heat, carbon would probably follow Dulong and 
 Petit's law. These facts make it apparent that specific heat deter- 
 minations are of no value in the selection of atomic weights unless 
 they are made through large intervals of temperature, and are then 
 found to be constant. Boron, silicon, beryllium, phosphorus, and 
 sulphur have, like carbon, very small specific heats. These ele- 
 ments, which are exceptions to Dulong and Petit's law, have small 
 atomic weights, and are (all but beryllium) not-metals. 
 
 Dulong and Petit's law is applicable in other cases besides those 
 in which the specific heats of solid elements are to be considered, 
 for, as the capacity for heat of a solid compound is very nearly equal 
 to the sum of the capacities of its constituent parts, it is evident that, 
 in many cases, the individual atomic heats can be obtained by a de- 
 termination of the specific heats of compounds. This method is 
 especially useful in cases where the elements themselves are not to 
 be obtained as solids ; for example : * - 
 
 Silver; atomic weight, 107.9; specific heat, .056; ac = 6.04 
 Iodine; " " 126.8; " " .054; ac = 6.8 
 
 The sum of the atomic heats of silver and iodine = 12.84 
 
 * Kopp (Liebig's Annalen, Suppl. 3, 290 ) made a large number of accurate 
 determinations of the specific heats of compound bodies, and to him and Reg- 
 nault is due the development of the proof of the truth of law in regard to 
 compound bodies. The figures here given are taken from L. Meyer's Grund- 
 ziige der Theoretischen Chemie, but the original source is to be found in 
 Kopp's researches, from which the figures are taken. 
 
MOLECULAR HEATS OF COMPOUNDS. 355 
 
 The observed specific heat of silver iodide is .0616. This, multi- 
 plied by the sum of the atomic weights of silver and iodine, gives 
 14.5, a number which is but little greater than the one calculated 
 above, namely, 12.84. The same is the case with silver bromide, 
 for:- 
 
 Silver ; atomic weight, 107.9; specific heat, .056; ac = 6.04 
 
 Bromine; " " 79.9; " " .084; ac = 6.7 
 
 Calculated for Ag Br = 12.8 
 
 Silver bromide; formula weight, 187.8; specific heat, .074; ac = 13.9 
 
 The two examples given above show us, therefore, that in the 
 case of substances the formulae of which are made up of two atomic 
 weights, the capacity for heat of the formula weight of the com- 
 pound is very nearly equal to twice that of either of the individual 
 atoms, and further investigation would demonstrate that compounds 
 made up of three atomic weights ( Pb Br 2 and Pb I 2 ) have about 
 three times the capacity of each individual atom. Having settled, 
 by experimental evidence, that Dulong and Petit's law holds good 
 for compounds, all of the constituent parts of which can be obtained 
 in the solid state, we can extend the method so as to indirectly de- 
 termine the atomic heats of elements which are gases at any tem- 
 perature which permits of a detailed study of their properties. For 
 example, we wish to discover the atomic heat of chlorine : 
 
 Silver chloride ; formula weight, 143.35; specific heat, .0911; ac = 13.1 
 
 Silver; atomic weight, 107.9 ; " " .056; ac = 6.04 
 
 Capacity for heat of 35.45 parts of chlorine = 6.96 
 
 The difference, 6.9, obtained as the atomic heat of chlorine is, 
 however, very close to the average of 6.4 which is found by direct 
 observation in a large number of solid elements ; so that if we regard 
 the number 35.45, as the true atomic weight of chlorine which num- 
 ber is selected by reason of a large number of specific gravity deter- 
 minations of gasifiable chlorine compounds, then the atomic heat of 
 chlorine very nearly coincides with the numbers obtained by direct 
 observation on a large number of solid elements, and it therefore 
 follows that the maximum atomic weight of 35.45 is confirmed by an 
 application of Dulong and Petit's law. Applications of this method 
 with other elements which, like chlorine, have such low melting 
 points that their respective specific heats cannot be measured when 
 they are in the solid state, have led to like results, although there 
 are a few exceptions, nearly all of which belong to compounds con- 
 taining not-metals with small atomic weights.* 
 
 * Compounds of hydrogen, nitrogen, fluorine, oxygen. 
 
356 LAW OF ISOMORPHISM. 
 
 Dulong and Petit's law nelps us to confirm numbers already 
 decided upon as atomic weights, and when a decision has been 
 reached in regard to a large number of elements by using some other 
 method as a guide, it can then help us to determine which numbers, 
 representing multiples or simple fractions of the equivalent weights, 
 are the true atomic weights of elements which form no gasifiable 
 compounds and which are themselves not volatile. In no case can 
 the law be of assistance in determining molecular weights. 
 
 A third method which has been of use in determining the num- 
 bers to be selected as atomic weights has been a study of the laws 
 of isomorphism. In 1819, Mitscherlich discovered that certain ele- 
 ments can replace others in chemical compounds without thereby 
 materially altering the crystalline forms of those compounds. 
 These, as well as the elements replacing each other in them, are 
 said to be isomorphous (see page 42). The replacing of one ele- 
 ment by another always takes place in definite stoichiometric quan- 
 tities ; for example, in comparing the members of the isomorphous 
 group, represented by the formulae Na Cl, Na Br, Na I, we find that 
 35.45 parts by weight of chlorine are always replaced by 80 parts 
 of bromine and 126.85 parts of iodine, and it is very evident that, 
 provided the atomic weight of one of these elements has been decided 
 upon by some other method, then the atomic weights of the others 
 will be given by the stoichiometric quantities which can replace 
 that element without altering the crystalline form of the compound. 
 This conclusion is correct, provided that in isomorphous mixtures 
 elements replace each other atom for atom, and this is the case with 
 the vast majority of isomorphous compounds.* This method for 
 determining atomic weights becomes very far reaching when we 
 consider that the same element may belong to two or three isomor- 
 phous groups of compounds, so that, when a decision as to the 
 atomic weights of the elements in one of these groups has been ar- 
 rived at, we can then cross over to another, and, perhaps, in that one 
 discover an element which also belongs to two or three new isomor- 
 
 * Mere similarity or even identity of crystalline form, even when two 
 bodies which show such similarity or identity are similarly constituted, does 
 not mean that the bodies in question are isomorphous. In order to conform 
 with the laws of isomorphism they must be able to replace each other with- 
 out thereby materially altering the form of the crystal. Such isomorphons 
 compounds are the alums (page 340), the vitriols (see magnesium), the calcite 
 and arragonite groups of carbonates. 
 
LAW OF ISOMORPHISM. 357 
 
 phous groups, and so on. It follows, therefore, that the selection of 
 the atomic weight of one element according to some other method than 
 that using the law of isomorphism may lead to the determination of 
 the atomic weights of a large number of others. For example, the 
 atomic weight of zinc, obtained by a determination of the vapor 
 density of some of the compounds of that element which can be 
 studied as gases,* has been placed at 65.3. This proportional part 
 by weight of zinc when crystallized in the sulphate is, however, 
 replaced by : 
 
 58.7 parts of nickel ; 56 parts of iron ; 
 59. " " cobalt ; 55 " " manganese, and 
 24.3 parts of magnesium. 
 
 Iron, manganese, and magnesium are further isomorphous with 
 40 parts of calcium on the one hand and with 27 parts of alumin- 
 ium and 52.1 parts of chromium on the other. Now, 40 parts of 
 calcium can replace 87.6 parts of strontium, 137.43 parts of barium, 
 and 206.95 parts of lead, so that this example shows that the selec- 
 tion of the maximum atomic weight of zinc can bring about the 
 determination of the atomic weights of a large number of other 
 elements, which are connected with zinc by isomorphous compounds. 
 In using the law of isomorphism one fact must, however, be borne 
 in mind. A number of instances are known in which groups of 
 atoms can replace individual atoms isomorphously ; this is the case 
 with ammonium and potassium salts, for the radicle NH 4 takes 
 the place of the atom K ; and if this can be true of ammonium and 
 potassium, it is, without doubt, true of a large number of other 
 radicles and elements; the law can, therefore, be applied only 
 where no possible doubt can exist. All of the methods which have 
 been outlined a t bove have for their object the selection of those 
 multiples or fractions of the chemical equivalent weights which 
 seem to us to represent the true atomic weights. The term " equiv- 
 alent weights " was suggested at the beginning of this century by 
 Wollaston, because that investigator, as well as a large number of 
 others, was of the opinion that no theories having sufficient plaus- 
 ibility would ever be advanced to enable us correctly to determine 
 the true atomic weights. If we understand by the expression 
 " equivalent weights " those quantities of the elements which 
 
 * For example, zinc ethyl, Zn (C 2 H 5 ) 2 . 
 
358 
 
 EQUIVALENT WEIGHTS. 
 
 unite with one equivalent, or one part by weight of hydrogen, or 
 which can take the place of one part by weight of hydrogen in 
 compounds, the determination is comparatively easy ; the true 
 atomic weights must then be multiples of these equivalent weights, 
 for one atom of an element cannot unite with less than one atom of 
 hydrogen; this conclusion becomes evident after a study of the 
 following table : 
 
 One part by weight of hydrogen * unites with 
 
 19. parts of fluorine, 
 35.45 " " chlorine, 
 79.95 " " bromine, 
 126.85 " " iodine, 
 
 n. 
 
 8. parts of oxygen, 
 16.03 " " sulphur, 
 39.5 " " selenium, 
 62.5 " " tellurium, 
 
 in. 
 
 4.7 parts of nitrogen, 
 10.3 " " phosphorus, 
 25. " " arsenic, 
 40. " " antimony, 
 
 and one part by weight of hydrogen is replaced by 
 
 7.02 parts of lithium, 
 23.05 " " sodium, 
 
 39.1 
 
 potassium, 
 
 12.15 parts of magnesium, 
 20. " " calcium, 
 43.8 " " strontium, 
 
 9 parts of aluminium. 
 
 In table I., the atomic weights and the equivalent weights are 
 alike; in table II., the atomic weights are twice the equivalent 
 weights ; in table III., they are three times the equivalent weiglits. 
 
 If every element were to combine with hydrogen to form a 
 well-defined hydrogen compound, or if we could easily replace 
 hydrogen in hydrogen compounds by every other element, the 
 determination of the various equivalent weights would be compara- 
 tively simple. This is not, by any means, the case ; so that at 
 one time, equivalent weights were defined as those weights which 
 would combine with one equivalent (eight parts by weight) of 
 oxygen. 
 
 With this interpretation matters become much more complicated, 
 as the following table will show : 
 
 8 parts by weight of oxygen (one equivalent) unite with: 
 
 8.68, or 17.37, or 26.05 parts by weight of chromium, 
 
 or with : 
 
 18.6, or 21 or 28 parts by weight of iron, etc. 
 
 Among these weights, which is to be selected as the equivalent 
 weight of iron or of chromium ? One thing only appears certain, 
 the various equivalent weights which we have to choose from in 
 
 * Although the true atomic weight of hydrogen is 1.008, oxygen being 16, 
 the decimal can be neglected, and hydrogen be placed as = 1. 
 
ELECTROLYTIC EQUIVALENTS. 359 
 
 the case of any given element, are in simple ratio to each other, so 
 that it follows that the true atomic weight of any given element 
 must be either a simple multiple or a fraction of one of its equiva- 
 lent weights ; but which multiple or which fraction ? It was from 
 such confusion, when many chemists had abandoned all hope of 
 ever selecting the true atomic weights of the elements, that the 
 logical application of the results obtained from Avogadro's hypo- 
 thesis rescued chemistry, by giving to it a uniform theoretical basis 
 for determining atomic weights. 
 
 One method, which gave tolerably uniform results in the deter- 
 mination of equivalent weights, was inaugurated by Faraday when 
 he discovered the fact that, when an electric conductor of the sec- 
 ond class (an electrolyte) * is made a conductor in a circuit, the 
 amount of the electrolyte which will be decomposed is always pro- 
 portional to the strength of the current. If, then, we allow the 
 same current to pass through two electrolytes, the constituents 
 which are separated from these are, electrically and chemically, 
 equivalent. In order, therefore, to determine the equivalent weight 
 of an element by this means we have but to decompose some com- 
 pound of that element by an electric current which, at the same 
 time, traverses some conducting hydrogen compound (dilute sul- 
 phuric acid for instance), and then to discover, experimentally, what 
 part by weight of the element is separated simultaneously with one 
 part by weight of hydrogen. This method is, however, also sub- 
 ject to error. For example, if we decompose the chloride of iron 
 or of copper, we are puzzled in the selection of the equivalent weights 
 of these elements because each forms two chlorides, which latter 
 compounds separate different quantities of metal in the same time 
 and with the same current ; the quantities of iron separated are, how- 
 ever, in simple ratio to each other, and the same is true of the quan- 
 tities of copper. To give an example, if we electrolyze ferrous 
 chloride, 28 parts of iron will separate simultaneously with one part 
 of hydrogen ; but if we electrolyze ferric chloride, 18.6 parts of iron 
 will be deposited during the same time ; these quantities of iron are 
 to each other as 3 : 2. Another objection to the method of deter- 
 mining equivalent weights by electrolysis is that the compounds 
 of many elements are non-conductors of electricity. 
 
 * A substance which conducts electricity while being at the same time 
 decomposed, for instance, acidulated water (see page 28). 
 
360 PRESENT ATOMIC WEIGHTS. 
 
 After the equivalent weights have been determined by accurate 
 methods of quantitative analysis, and after all hypotheses and ex- 
 perimental facts have been considered, we can select the maximum 
 atomic weights (which are multiples of these equivalent weights) be- 
 cause of the true understanding of the meaning of the numbers 
 obtained by the determination of the specific gravities of gases, or 
 of molecular weight determinations brought about by other means 
 (and hence by the methods founded upon Avogadro's hypothesis). 
 These atomic weights have been universally accepted as the basis 
 of our present system. This solution seems to be the correct one if 
 we remember that only our present atomic weights, when arranged 
 in order, beginning with that belonging to the element with least 
 atomic weight, and ending with that with the greatest, form a natu- 
 ral system, in which elements with similar properties recur after 
 stated intervals. This arrangement, known as the periodic system 
 of the elements, will be discussed in the next chapter. 
 
PERIODIC SYSTEM. 361 
 
 CHAPTER L. 
 
 THE PERIODIC SYSTEM OP THE ELEMENTS. 
 
 THE theoretical basis for the previous discussions in this book 
 has- been the assumption that the properties of compounds are, in 
 reality, the properties of the molecules which, when massed to- 
 gether, make up those compounds ; and we have also come to the 
 conclusion that the character of any individual molecule depends on 
 the nature and also upon the relative position of the atoms which, 
 in uniting, produce that molecule. If now, we further wish to dis- 
 cover whether any connection exists between the characters of the 
 various elements, we must first attempt to compare those funda- 
 mental constants which appertain to the atoms and which are also 
 definitely determinable or *to be ascertained with great accuracy. 
 Such constants are both the atomic weights and the specific gravi- 
 ties of elementary bodies. 
 
 In the year 1868 both Mendelejeff and Lothar Meyer* demon- 
 strated that, when the elements are arranged in the order of their 
 increasing atomic weights, each element will differ in properties 
 from those immediately preceding and following it, but, after certain 
 definite intervals, elements will recur which possess very similar 
 characteristics. This interval they found, when the elements are 
 arranged as indicated, to be after every seventh element if the 
 first fourteen members alone are considered, but after every seven- 
 teenth in the remainder of the series. The elements were therefore 
 divided into two " short periods " of seven each, and into five " long 
 periods " of seventeen individuals ; the periods being so selected that 
 each begins with an alkali metal and ends with a halogen ; for, by 
 this arrangement, the extremities are formed of sharply contrasting 
 elements. The long periods are, however, in reality double, for in 
 each one the first seven elements resemble the last seven in many 
 
 * Newlands also called attention to the connection between atomic weights 
 and chemical properties of elements. 
 
362 PERIODIC SYSTEM; ARRANGEMENT. 
 
 important chemical characteristics, while the three remaining ele- 
 ments which are in the middle * form a separate group. 
 
 This arrangement of the elements, including the division into 
 short and long periods, is shown in the table on the opposite page. 
 
 The first two (short) periods and the third and fourth (long) 
 periods are complete with one exception, that of an element with 
 an atomic weight of about 100 and corresponding to manganese. 
 The remaining periods have only a few known representatives in 
 each, and all of the latter, with possibly one or two exceptions, may be 
 classed with the very rare elements. As a consequence, all general- 
 izations must be made from a study of the complete periods, while 
 the conclusions drawn from the latter can only be cautiously ex- 
 tended so as to cover the ones in which a number of representatives 
 are missing. 
 
 The short periods are the types of all ; the others are constructed 
 after those models with this exception: While the short periods, 
 beginning as they do with a pronounced metal and ending with a 
 pronounced not-metal, will show all gradations of character which 
 lie between the two extremes, while traversing the five intervening 
 elements ; the long periods, although they begin and end each with 
 an equally pronounced metal and not-metal, nevertheless require the 
 interposition of fifteen elements to effect the gradation which is pro- 
 duced by five in the short periods ; from this it follows that the 
 elements in the middle of the long periods must be further removed 
 in character from those in the middle of the short periods than are 
 the elements at either extremity ; we would, therefore, expect ele- 
 ments like chromium, manganese, iron, cobalt, nickel, and copper to 
 have no counterparts in the typical periods. This conclusion is not 
 strictly correct, however, for, as has been mentioned, the long 
 periods are each really composed of two shorter ones. The last 
 members of the first half of the long periods therefore resemble in 
 many ways the last members of the typical short periods, while the 
 first elements of the second half are much like the elements which 
 begin the short periods. Any given long period begins with a most 
 pronounced metal like potassium or caesium; following this are 
 six elements, each one of which is less metallic in its nature than 
 the one immediately preceding it ; the eighth, ninth, and tenth ele- 
 
 * Elements No. 8, 9, and 10 of any long period. 
 
PERIODIC SYSTEM ; ARRANGEMENT. 
 
 363 
 
 H- * 
 
 
 
 
 
 
 
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364 
 
 PERIODIC SYSTEM; ARRANGEMENT. 
 
 ments then form a transition from the first to the last seven of 
 the long period ; they are successively 'more metallic in their nature. 
 The eleventh element is once more a tolerably pronounced metal, 
 although it has by no means so metallic a nature as the alkali 
 metals which begin a period. From the eleventh to the seventeenth 
 element the not-metallic properties again become more and more 
 pronounced, until the last elements of the periods (bromine or 
 iodine) are among the most intensely not-metallic elements with 
 which we are acquainted. The gradation in properties shown by 
 the periods can, perhaps, be graphically represented as follows : 
 
 Short period, 
 7 elements ; metallic properties 
 
 Long period, 
 17 elements; metallic properties 
 
 not-metallic 
 
 properties. 
 
 The elements which find their places in the vertical columns 
 of the periodic system form the natural families or groups of ele- 
 ments ; when we compare the short periods with the long ones we 
 find that the first two elements in each of the short periods belong 
 to the same family as the first two in each of the longer ones, while 
 the last five in the former correspond to the last five in the latter ; 
 as has been pointed out, the intervening ten elements in the long 
 periods differ more or less from any which occur in the typical pe- 
 riods. This relationship is clearly shown by the following table : 
 
 Li, 
 
 Be, 
 
 
 
 
 
 
 
 
 
 
 
 B, 
 
 c, 
 
 N, 
 
 o, 
 
 F, 
 
 Na, 
 
 Mg, 
 
 
 
 
 
 
 
 
 
 
 
 Al, 
 
 Si, 
 
 P, 
 
 s, 
 
 Cl, 
 
 K, 
 
 Ca, 
 
 Sc, 
 
 Ti, 
 
 V, 
 
 Or, 
 
 Mn, 
 
 Fe, 
 
 Co, 
 
 Ni, 
 
 Cu, 
 
 Zn, 
 
 Ga, 
 
 Ge, 
 
 As, 
 
 Se, 
 
 Br, 
 
 Rb, 
 
 Sr, 
 
 Y, 
 
 Zr, 
 
 Cb, 
 
 Mo, 
 
 , 
 
 Itu, 
 
 Rh, 
 
 Pd, 
 
 Ag, 
 
 Cd, 
 
 In, 
 
 Sn, 
 
 Sb, 
 
 Te, 
 
 I, 
 
 Cs, 
 
 Ba, 
 
 La, 
 
 Ce, 
 
 Di, 
 
 f 
 
 f 
 
 t 
 
 } 
 
 t 
 
 , 
 
 f 
 
 } 
 
 t 
 
 , 
 
 9 
 
 , 
 
 
 
 Yb, 
 
 , 
 
 Ta, 
 
 W, 
 
 , 
 
 Os, 
 
 Ir, 
 
 Ft, 
 
 Au, 
 
 Hg, 
 
 Tl, 
 
 Pb, 
 
 Bi, 
 
 , 
 
 , 
 
 
 
 
 Th, 
 
 , 
 
 u, 
 
 , 
 
 
 
 
 
 
 
 
 
 
 
 As we pass from member to member along the complete series 
 of the elements we encounter, at one time, a gradual, at another, an 
 abrupt change in the character of the elements ; the gradual changes 
 between the two extremities of any period, the abrupt changes as 
 we pass from one period to another, and these changes are brought 
 about, with close resemblance, in each one of the periods, so that 
 nearly every property of any given element is repeated in one or 
 more subsequent ones; this repetition is found, not only in the 
 chemical character of the elements and their compounds, but is also 
 
ATOMIC VOLUMES. 365 
 
 apparent, even in a greater degree, in their physical properties, the 
 properties of any individual element are therefore determined by the 
 position of that element in the periodic system. 
 
 One of the easily determined constants which appertains to each 
 element is its specific gravity in the solid state, and that property 
 periodically increases and decreases as we pass in regular order from 
 element to element along the entire series. This fact becomes very 
 apparent if, instead of comparing the specific gravities themselves 
 (i.e., the quantities of matter contained in the unit volume), we 
 compare the volumes which are occupied by weights of the respec- 
 tive elements, which weights are so taken that the number of 
 grams correspond to the atomic weights. By this means we can 
 arrive at the volumes occupied by the same number of atoms in 
 each case, and these, necessarily, must bear the same relationship to 
 each other as do the volumes occupied by the individual atoms. 
 These volumes can appropriately be termed " atomic volumes," and 
 they are readily ascertained, in the case of each element, by dividing 
 the atomic weight by the specific gravity, so that : 
 
 V =-. 
 c 
 
 In this equation V represents the atomic volume, a the atomic 
 weight, and c the specific gravity of the solid element. For in- 
 stance, the atomic weight of lithium is 7.02, its specific gravity 
 
 7 02 
 
 is .59, its atomic volume is therefore = 11.9 ; the quotient 
 
 .59 
 
 11.9 means that 7.02 grams of lithium occupy 11.9 cubic centi- 
 metres of space ; the atomic weight of manganese is 55, its specific 
 gravity is 8, its atomic volume 6.9, therefore 55 grams of manga- 
 nese take up 6.9 cubic centimetres. The table on following page 
 demonstrates the relationship between the atomic volumes and the 
 periodic system. 
 
 In each period, whether it be short or long, the specific gravity 
 begins with a minimum (with the specifically light alkali metals), 
 advances to a maximum at the middle, and then once more dimin- 
 ishes to a minimum at the opposite, not-metallic extremity ; each 
 period, therefore, represents a complete wave in regard to specific 
 gravities, the beginning being in tfre trough, the middle at the crest, 
 and the end in the succeeding trough. The reverse is true in regard 
 to atomic volumes ; these begin with their maximum at the alkali 
 
366 
 
 ATOMIC VOLUMES. 
 
 metals, diminish to a minimum at the centre of the periods, and 
 then once more increase to a maximum at the other extremity ; the 
 
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 TH 
 TH 
 
 > 
 
 iq 
 
 CO 
 05 
 
 O 
 
 CO 
 
 o 
 
 CO 
 
 CO 
 
 TH 
 
 
 
 
 
 10 
 
 00 
 
 H 
 
 b- 
 
 CO 
 
 CO 
 
 TH 
 
 N 
 
 1 1 
 
 TH 
 
 o 
 
 CO 
 
 . 
 
 
 
 
 
 00 
 
 CO 
 
 CXI 
 
 TH 
 
 ** 
 
 CO 
 
 CXI 
 
 
 b- 
 
 ^44 
 
 
 
 
 
 O 
 
 10 
 
 TH 
 
 eS 
 
 CO 
 
 iq 
 
 c? 
 
 * 
 
 CO 
 CO 
 
 II 
 
 id 
 II 
 
 a 
 
 1O 
 
 T 1 
 
 CO 
 
 10 
 
 II 
 
 changes in the atomic volumes in each period, therefore, may be 
 compared to a wave the crest of which corresponds to the beginning 
 
ATOMIC VOLUMES. 367 
 
 alkali metal, the trough to the middle of the period, and the succeed- 
 ing crest to the next following alkali metal. The atomic volumes 
 of the alkali metals, however, increase rapidly with increasing atomic 
 weight, so much so, indeed, that the atomic volume of sodium is 
 twice that of lithium, and the atomic volume of potassium is twice 
 that of sodium ; each succeeding wave which represents the changes 
 in the atomic volumes of the elements forming one period has a 
 greater amplitude than the one preceding it, and a lesser amplitude 
 than the one following. If, then, we take the atomic volumes as our 
 guide, representing these as ordinates and the atomic weights as 
 abscissae, we can represent the periodic system in the form of succes- 
 sive wave-like curves, the relative position of any element upon 
 these curves determining the properties of that element. Those 
 elements which find their places on a descending branch of one of 
 these curves, and which immediately follow a maximum, and the 
 next following elements down to a minimum and even a little 
 beyond this point, are difficult to fuse, and are not volatile, and, 
 furthermore, they are less fusible the nearer they approach a mini- 
 mum ; those elements in the ascending curves are easily melted, and, 
 with few exceptions, are volatile ; of the elements belonging to the 
 latter class, nitrogen, oxygen, and fluorine in the first period are 
 gases ; in the second, chlorine only is a gas, phosphorus and sulphur, 
 however, melt at a low temperature and are easily volatilized ; in the 
 first long period the volatile elements begin with zinc, in the next 
 following with silver, and in the next (incomplete) one with mercury. 
 Lothar Meyer, in consideration of these facts, has established the 
 following rule : In that portion of the series in which the atomic 
 volumes are decreasing with increasing atomic weights, the elements 
 are not volatile and are fusible with difficulty ; on the other hand, 
 where the atomic volumes are increasing the elements are easily 
 melted and are volatile. All of the other properties of the elements 
 vary twice in all of the long periods ; so, for instance, the alkali 
 metals at the maxima of the curves as well as the metals immedi- 
 ately following are malleable and ductile ; then, as the minima are 
 approached, there follow brittle, crystalline metals ; these, in turn, 
 at the minima give way to malleable and ductile ones ; and succeed- 
 ing the latter, as the next maximum is approached, the not-metals, 
 which are neither malleable nor ductile, find their places. 
 
 If we designate as positive those elements whose oxides, in the 
 greater number of cases, act as bases, and as negative those ele- 
 
368 
 
 VALENCE AND THE PERIODIC SYSTEM. 
 
 ments whose oxides are anhydrides,* then the periods all begin 
 with strongly positive alkali metals ; next following, on descending 
 curves of the atomic volumes, are a number of less positive elements ; 
 as a minimum is approached the latter give way to one or two neg- 
 ative (or at least in greater part negative) individuals ; succeeding 
 these, at the minimum and at the beginning of the ascent toward 
 the maximum are a number of positive elements ; and finally the 
 curves are completed by elements which are entirely negative. 
 
 That the chemical properties of the elements are altogether in 
 harmony with the periodic system has been repeatedly noted, as 
 certain sections of that system have been discussed during the 
 progress of this work ; and, as these relationships have already been 
 explained at some length at those portions of the work where the 
 various families have been taken up, and as they will be further 
 noted in those which are to come, it seems scarcely necessary at 
 this place to do more than to state briefly sonle few connections. 
 
 By determining the specific gravities of the gases obtained by 
 volatilizing the halogen and hydrogen compounds of the elements, 
 we are able to determine the respective valences of the atoms (see 
 page 107). If we compare the halogen and hydrogen compounds 
 in the first period with those in the second, we find the changes in 
 valence, as we go from left to right, to be identical in both. The 
 beginning elements (alkali metals) are invariably univalent. Pass- 
 ing from these to the right, we find that the valence is increased by 
 one with each pair of successive elements until a maximum of four 
 is reached in connection with the members of the carbon family ; it 
 then successively diminishes until it once more reaches a minimum 
 of one (halogen family). These changes are made apparent by the 
 following table : 
 
 (H represents an atom either of hydrogen or of chlorine.} f 
 
 Compounds, 
 
 Li, 
 LiK, 
 
 Be, 
 BeR 2 , 
 
 B, 
 BR 8 , 
 
 c, 
 
 CR 4 , 
 
 N, 
 NR 3 , 
 
 o, 
 
 OR 2 , 
 
 F, 
 FR. 
 
 Valence, 
 
 1, 
 
 2, 
 
 3, 
 
 4, 
 
 3, 
 
 2, 
 
 1. 
 
 Compounds, 
 
 Na, 
 NaR, 
 
 Mg, 
 MgRs, 
 
 Al, 
 A1R 3 , 
 
 Si, 
 SiR 4 , 
 
 P, 
 PR 3 , 
 
 s, 
 
 SR 2 , 
 
 01, 
 
 C1R. 
 
 * See page 13. 
 
 t This comparison is legitimate only in the first two periods, and then 
 with the understanding that, when a hydrogen compound of an element does 
 exist, we study that substance and not the chloride. 
 
VALENCE AND TH PERIODIC SYSTEM. 
 
 369 
 
 A different result becomes apparent in comparing the compounds 
 which the elements form with oxygen. The valence toward that 
 element (page 111) begins with one in the family of the univalent 
 alkali metals and increases to a maximum of seven in that of the 
 halogens. This change becomes evident if we represent a single 
 valence of one of the atoms of oxygen by r, = 2r, and then group 
 the elements as shown in table : 
 
 
 
 1 
 
 
 
 
 
 
 Li, 
 
 Be, 
 
 B, 
 
 c, 
 
 H, 
 
 o, 
 
 F, 
 
 Oxides, 
 
 Li. 2 0, 
 
 BeO, 
 
 B 2 3 , 
 
 CO,, 
 
 N 2 5 , 
 
 i 
 
 > 
 
 
 Lir, 
 
 Ber 2 , 
 
 Br 3 , 
 
 Cr 4 , 
 
 Nr 5 , 
 
 ~~i 
 
 . 
 
 
 Na, 
 
 Mg, 
 
 Al, 
 
 Si, 
 
 P, 
 
 s, 
 
 Cl, 
 
 Oxides, 
 
 Na 2 0, 
 
 MgO, 
 
 A1 2 3 , 
 
 Si O, , 
 
 P 2 5 , 
 
 S0 3 , 
 
 (C1 2 7 ), 
 
 
 Nar, 
 
 Mgr. 2 , 
 
 Alr 3 , 
 
 Sir 4 , 
 
 Pr 5 , 
 
 Sr 6 , 
 
 Clr 7 . 
 
 Valence, 
 
 1, 
 
 2, 
 
 3, 
 
 4, 
 
 5, 
 
 6, 
 
 7. 
 
 In making this comparison we must remember that carbon, 
 nitrogen, phosphorus, sulphur, and chlorine each forms a number 
 of oxides, so that it is only by comparing those compounds which 
 contain the greatest amount of oxygen that the elements in a short 
 period present a regular increase of valence from one to seven. 
 
 In the long periods there are but few hydrogen compounds, 
 and these belong only to the last few elements of the periods ; so 
 that if we wish to compare the long periods with the short ones, we 
 shall be compelled to resort almost exclusively to the chlorine com- 
 pounds. If the halides of the elements in the short periods were 
 to correspond exactly to the hydrogen compounds, we should be 
 able to construct the following: table : 
 
 Hydrogen compounds, N H 3 , 
 PH 3 , 
 
 Halogen compounds, N C1 3 , 
 PCL, 
 
 OH 2 , 
 SH 2 , 
 
 FH, 
 C1H. 
 
 OC1 2 , FC1, 
 SC1 2 , C1C1. 
 
 All of these chlorine compounds do, in reality, exist ; but, in 
 addition, phosphorus forms a pentachloride, P C1 5 , and sulphur a 
 tetrachloride, S C1 4 , and a monochloride, S 2 C1 2 ; and the difficulty 
 of a systematic comparison of the halides is further enhanced 
 by the discovery of chlorides having the formulae Mo C1 5 and W C1 6 . 
 From these facts we are, perhaps, justified in drawing the conclu- 
 sion that, probably, in the two short periods the valence toward 
 
370 
 
 VALENCE AND THE PERIODIC SYSTEM. 
 
 chlorine would increase from one to seven, just as it does toward 
 oxygen, if it were not for the fact that the not-metallic elements 
 in those periods are too negative to enable them to retain any 
 great number of chlorine atoms in a stable molecule. This objec- 
 tion does not appertain to molybdenum and tungsten, and with 
 these, as the above formulae show, the valence toward chlorine can 
 be five and six. The dual nature of the long periods is plainly 
 demonstrated by a study of the valence which the elements belong- 
 ing therein develop toward oxygen. The initial alkali metal is al- 
 ways univalent ; advancing from this the valence steadily increases 
 with each successive element, until a maximum of seven is reached ; 
 it then diminishes with the eighth, ninth, and tenth ; is again equal 
 to one with the eleventh, and from this to the seventeenth increases 
 to a maximum of seven. In order to illustrate these changes, we 
 will select the first long period, and, in order to demonstrate more 
 clearly the existing resemblance between its first and second halves, 
 will place the one under the other : 
 
 I. 
 
 II. 
 
 III. 
 
 IV. 
 
 V. 
 
 VI. 
 
 VII. 
 
 VI. 
 
 IV. 
 
 III. 
 
 g K, 
 
 Ca, 
 
 Sc, 
 
 Ti, 
 
 V, 
 
 Cr, 
 
 Mn, 
 
 Fe, 
 
 Co, 
 
 Ni, 
 
 3 K a O, 
 
 CaO, 
 
 Sc 2 3 , 
 
 Ti O 2 , 
 
 V 2 5 , 
 
 Cr0 3 , 
 
 Mn 2 O 7 , 
 
 FeO,, 
 
 Co0 2 , 
 
 Ni 2 3 , 
 
 & Kr, 
 
 Car 2 , 
 
 Scr 3 , 
 
 Tir 4 , 
 
 Vr B , 
 
 Crr e , 
 
 Mn r 7 , 
 
 Fer 6 , 
 
 Cor 4 , 
 
 Ni r a . 
 
 8 Cu, 
 
 Zn, 
 
 Ga, 
 
 Ge, 
 
 As, 
 
 Se, 
 
 Br, 
 
 
 
 
 2j Cu 2 0, 
 
 ZnO, 
 
 Ga 2 3 , 
 
 GeO 2 , 
 
 As 2 O 8 , 
 
 Se0 3 , 
 
 (Br 2 7 ), 
 
 
 
 
 Cur, 
 
 Znr 2 , 
 
 Gar 3 , 
 
 Ger 4 , 
 
 Asr e , 
 
 Ser e , 
 
 Brr 7 . 
 
 
 
 
 The elements in the first half of this period (beginning with 
 a very intensely metallic alkali metal and ending with a metal 
 [manganese]), are necessarily much more positive in their char- 
 acter than are those of the second half, which ends with a pro- 
 nounced not-metal (bromine); nevertheless, the two sections bear a 
 striking resemblance to each other. This period is, therefore, 
 formed of a primary and secondary short period and of three ele- 
 ments, iron, cobalt, and nickel, which connect the two halves and 
 form a gradual transition from one to the other. Both the primary 
 and secondary short periods resemble the typical short periods. This 
 relationship is made apparent in the following table, which is the 
 one in general use : 
 
MENDELEJEFF'S TABLE. 
 
 371 
 
 o 
 
 DC 
 
 o 
 
 to 
 
 era 
 
 Ci 
 
 OP 
 
 Or 
 
 to 
 
 O 
 
 OP 
 
 o 
 
 p. Or -j .^ OQ 
 
 bO CO CO 
 
 to 
 
 00 
 
 tO 
 
 CO ij 
 ^ 3 fcO 
 
 o ^ o >fl 
 
 pro* 
 
 CD 
 
 CO 
 
 OO 
 
 to . 
 
 03 e 
 
 * o 
 
 co 
 
 01 
 
 tO 
 
 
 
 SO 5^ Cn 
 
 ^ Oi & fcO 
 
 to H ^ 
 
 Ci 
 
 01 
 
 0) O CD 
 
 O 
 
 CO 
 
 o 
 
 00 
 
 00 l-l 
 
 Or 
 
 s 
 
 Ci ^ 
 
 * Now separated into neodymium (140.5) and praseodymium (143.6). 
 Atomic weight of tellurium is doubtful (see page 104 
 
372 
 
 TABLE OF OXIDES. 
 
 The elements which are placed between the vertical lines consti- 
 tute the natural families, those in the horizontal lines are the 
 periods or series ; those periods which have even numbers constitute, 
 with the exception of the first short period (number 2), the first 
 sections of the long periods; those with the odd numbers, with the 
 exception of the second short period (number 3) form the second sec- 
 tions ; the series having even numbers, therefore, bear the closest 
 resemblance to each other, while those having odd numbers also show 
 a great similarity of characteristics ; on the other hand, the periods 
 numbered with odd numbers bear a much less marked resemblance 
 to those with even ones. These peculiarities will be demonstrated 
 more at length during the discussion of the individual families of 
 metals. If a table of the oxides of the elements is constructed, 
 while following out the arrangement given above, the remarkable 
 regularity displayed in the formation of those compounds is made 
 apparent.* 
 
 2. 
 3. 
 4. 
 5. 
 6. 
 7. 
 8. 
 9. 
 10. 
 11. 
 12. 
 
 i. 
 
 ii. 
 
 III. 
 
 IV. 
 
 V. 
 
 VI. 
 
 VII. 
 
 Li 2 0, 
 Na 2 0, 
 K 2 0, 
 Cu 2 O, 
 Rb 2 0, 
 Ag 2 0, 
 Cs 2 0, 
 ? 
 
 Be 2 2 , 
 Mg 2 2 , 
 Ca 2 2 , 
 Zn 2 2 , 
 Sr 2 2 , 
 Cd 2 O 2 , 
 Ba 2 O 2 , 
 
 B 2 3 , 
 A1 2 3 , 
 Sc 2 3 , 
 Ga 2 3 , 
 Y 2 3 , 
 In 2 3 , 
 La 2 O 3 , 
 
 C 2 4 , 
 Si 2 O 4 , 
 Ti 2 4 , 
 Ge 2 4 , 
 Zr 2 4 , 
 Sn 2 O 4 , 
 Ce 2 O 4 , 
 i 
 
 Pb 2 4 , 
 
 No O 5 , 
 
 P 2 " 5 , 
 Y 2 6 , 
 
 As 2 5 , 
 Cb 2 5 , 
 Sb 2 5 , 
 Di 2 5 , 
 
 Ta 2 5 , 
 Bi 2 5 , 
 
 82 6f 
 Cr 2 6 , 
 Se 2 6 , 
 Mo 2 6 , 
 Te 2 6 , 
 
 W 2 6 , 
 U 2 , 
 
 C1 2 7 , 
 Mn 2 O 7 , 
 Br, 7 , 
 
 I 2 7 , 
 
 i 
 
 Yb,0 3 , 
 Tl a 3 , 
 
 ^ 
 
 > 
 
 5 
 
 Au. 2 0, 
 
 5 
 
 Hg 2 O 2 , 
 
 5 
 
 ! 
 
 In constructing this table the formulae of the oxides in the 
 families of beryllium, carbon, and sulphur have been doubled, so as 
 to render the increase in the valence toward oxygen, as the series 
 proceed from left to right, more apparent. Of course, a number of 
 the elements form oxides with formulae differing from those given 
 in the table ; only those oxides have been selected for purposes of 
 comparison, which are, in any given family, common to all of the 
 members of that family, or, in cases where the oxides themselves f 
 
 * See Lothar Meyer; Grundziige der Theoretischen Chemie; p. 65. 
 t For instance, oxides of the formulae L> O 7 and Br 2 O 7 do not exist; but 
 the acids derived from these, and the salts of these acids, are known. 
 
PREDICTION OF ELEMENTS. 
 
 373 
 
 are not known, their existence has been considered as theoretically 
 possible because some derivatives of the missing oxides have been 
 described. In any given family, any one of the oxides given on 
 the above table in the vertical column belonging to that family, 
 may be termed the typical oxide of that group. 
 
 By a skilful combination of the connections which have been 
 emphasized in the last chapter, Mendelejeff was able to predict the 
 existence of a number of elements, unknown at the time of the dis- 
 covery of the periodic system ; for the purpose of illustrating the 
 method adopted by that investigator, one example will be given here. 
 
 No element fitting into the fourth series, group three, was 
 known at the time when the periodic system was discovered ; yet, 
 Avere such a one to be isolated in the future, it should, in its proper- 
 ties, be related* to aluminium in the same way as calcium is to 
 magnesium, or as titanium is to silicon. Its atomic weight should 
 be about 44, inasmuch as it would follow K (39), Ca (40), and be 
 followed by Ti (48) and V (51). In predicting the properties of 
 this element, Mendelejeff reasoned that it would be as much 
 more metallic than aluminium, as calcium is than magnesium, or 
 as titanium is than silicon. This unknown element Mendelejeff 
 called ekaboron, with a symbol Eb ; and for the purposes of com- 
 parison, the predicted properties of ekaboron and the real proper- 
 ties of scandium, the element which was subsequently discovered, 
 are placed side by side : * 
 
 SCANDIUM. 
 
 Atomic weight about 44. 
 
 Oxide, Eb^ O 3 , soluble in acids, analogous 
 to Al a O 3 , but more basic; insoluble in alka- 
 lies. 
 
 Salts of Eb, colorless, yield gelatinous pre- 
 cipitates with Na OH, Na 2 Co 3 . 
 
 Sulphate, #& 2 (S0 4 ) 3 , will form a double 
 salt with K 2 SO 4 , which will not be isomor- 
 phous with the alums. 
 
 Atomic weight, 44. 
 
 Oxide, Sc a O 3 , soluble in strong acids, anal- 
 ogous to Al a O 3 , but decidedly more basic; 
 insoluble in alkalies. 
 
 Salts of Sc are colorless, and yield gelatin- 
 ous precipitates with NaOH, Na 2 CO 3 . 
 
 Sulphate, Sc 2 (S0 4 ) 3 , forms a double salt 
 with K 2 SO 4 , which is not isomorphous with 
 the alums. 
 
 It seems scarcely necessary to enter into a more detailed 
 description of the periodic system at this place ; the elements 
 which have already been considered have been discussed in their 
 relation to the natural groups of which they are members, so that 
 
 * See Pattison Muir; Principles of Chemistry; p. 201. 
 
374 PREDICTION OF ELEMENTS. 
 
 their individual connections have been sufficiently pointed out ; 
 those which are to follow will be described in the order given by 
 the periodic system, while attention will be called to the character 
 of the various families at the proper place. It does not fall 
 within the scope of this work to give a detailed description of 
 each individual metal, as has been done with the not-metals ; for 
 the ground is very abundantly covered by the large number of 
 works on qualitative analysis, which discuss many of the chemical 
 reactions peculiar to the metals because the latter are mainly inter- 
 esting from an analytical standpoint ; and, moreover, much of the 
 chemistry of the salts of the metals is merely a repetition of what 
 has already been taken up. 
 
NEUTRALIZATION. 375 
 
 CHAPTER LI. 
 
 NEUTRALIZATION. DOUBLE DECOMPOSITION. DISSOCIATION 
 OF ELECTROLYTES. 
 
 THE phenomena attending the neutralization of an acid by a base 
 are of such importance that a brief discussion of their nature is 
 necessary. If, to use potassium hydroxide as an example, that base 
 is brought in contact with hydrochloric acid, the following change, 
 as expressed by our chemical equations, takes place : 
 
 KOH+ HC1=KC1+H 2 0. 
 
 Of the two systems, KOH + H Cl and K Cl + H 2 0, the former is 
 in unstable equilibrium ; the two systems correspond to two differ- 
 ent quantities of energy, so that when the former is converted into 
 the latter, energy is conducted away in the form of heat. The 
 sum of the energy thus conducted away, and of that remaining in 
 the system K Cl + H 2 0, must be equal to that originally contained 
 in KOH -f H Cl. We are also acquainted with chemical reactions 
 in which energy must be conducted to a system in order to change it 
 into a second one ; the former class of reactions are exothermic, 
 the latter are endothermic (page 12). In how many portions the 
 energy may be communicated or may pass off has evidently no 
 effect on the final value.* 
 
 In the reaction cited above there must be an expenditure of 
 energy sufficient to decompose KOH into K -f OH and H Cl into 
 H -f Cl before the rearrangement into K C1.+ H 2 can take place ; 
 but during the complete reaction energy passes off in the form of 
 heat, and the amount of the latter is evidently independent of any 
 intermediate changes. The reaction between potassium hydroxide 
 and hydrochloric acid, taking into account all of the thermal values> 
 would be expressed as follows : 
 
 KOH + H Cl = K, Cl [1012 K] + H 2 , [684 K] (K, O, H [1165 K] 
 + H Cl [393 K]) ; 
 
 * See Ostwald, Outlines of General Chemistry, p. 368. 
 
376 HEAT OF NEUTRALIZATION. 
 
 and from this it will be seen that the sum of the heats of forma- 
 tion of a formula weight of potassium chloride and one of water is 
 greater by 138 K than that of the sum of the heats of formation of 
 similar quantities of potassium hydroxide and hydrochloric acid. 
 This may be expressed in the following terms : " The heat of for- 
 mation of K, Cl + H 2 , is greater than that of K, 0, H -f H, Cl," * 
 and, as we may disregard the intermediary changes, we may express 
 the final result as follows : 
 
 KOH aq + HC1 aq = K Claq + H 2 + aq + 138 K, 
 
 the symbol aq signifying that the constituents are dissolved in a 
 quantity of water so large that the addition of any more of the 
 reagents will not affect the thermal value ; or, as the formation of 
 water takes place in all neutralizations of acids with bases, and as 
 its mixture with the salt solution can produce no thermal effect, we 
 may write the equation : 
 
 KOH aq + H Cl aq = K Cl aq + 138 K. 
 
 Now, it has been shown that different acids, as well as different 
 bases, evolve different amounts of heat when neutralized, but the 
 difference between the amounts of heat given off by any two bases 
 when neutralized by a series of acids or between any two acids when 
 neutralized by a series of bases is always the same. The strong 
 monobasic acids (hydrochloric acid, hydrobromic acid, hydroiodic 
 acid, nitric acid, chloric acid, bromic acid, perchloric acid, iodic 
 acid), all give off very nearly the same amount of heat when neu- 
 tralized by an equimolecular quantity of caustic soda ; this amount 
 is very nearly 138 K. Among dibasic acids, on the contrary, a 
 different behavior is observed. Although some, like the monobasic 
 acids, liberate 139 K for each equivalent ; t others, on the other 
 hand, liberate more. For instance, if increasing quantities of sul- 
 phuric acid are added to a weight in grams of sodium hydroxide 
 equivalent to one combining weight of the hydroxide, an evolution 
 of heat takes place until sufficient acid has been added to form the 
 secondary sulphate, Na 2 S0 4 . This amounts to 157 K for one equiv- 
 
 * The comma introduced between the two portions of a chemical formula 
 indicates that the portions of that formula separated are to interact chemi- 
 cally to form the complete substance. 
 
 t By equivalent is meant one-half the formula weight, or that proportion 
 by weight of the acid which would contain one part by weight of hydrogen. 
 
STRONG AND WEAK BASES. 377 
 
 alent weight in grams, and to 314 K for one formula weight in 
 grams, of sulphuric acid. If sulphuric acid is further added, an 
 absorption of heat is observed until a limit of 33 K is reached. 
 This absorption takes place during formation of the primary sul- 
 phate (see page 153), so that we here observe the phenomenon of an 
 endothermic reaction inaugurated spontaneously. The hydroxides 
 of the alkali metals and of the alkaline earths form a group of bases 
 which act similarly toward the monobasic acids ; when neutralized 
 with hydrochloric acid they give off, for each equivalent weight in 
 grams, about 139 K. Other hydroxides, such as aluminium or ferric 
 hydroxide, have much smaller heats of neutralization. The fact 
 that so much heat is evolved when acids are brought in contact 
 with bases explains why the reactions of neutralization take place 
 so readily, and why the salts formed are so often among the most 
 stable chemical bodies. Of course, some very weak acids are able 
 only partially to neutralize bases, and very weak bases to partially 
 neutralize acids. The heats of neutralization in these cases will be 
 small, and the salts easily decomposed by the addition of water or 
 of acids, or even by slight warming. Such examples have been en- 
 countered in the study of the chlorides of bismuth and antimony 
 and in the study of hydrocyanic acid.* 
 
 When we use the expressions " strong " and " weak " bases, it, 
 however, becomes as necessary to define the meaning of those terms 
 as it was when we used similar designations in regard to acids (see 
 page 141). But little work, as compared with that done in the 
 study of acids, has been accomplished in regard to the affinity of 
 bases, yet we can place the ratio between the strengths of two bases 
 as being given by the relative rapidity with which the two are able 
 to decompose a salt of a third, not very pronouncedly metallic, sub- 
 stance, f Experiments in this direction have shown that the hy- 
 droxides of the alkali metals all act about equally in this respect, 
 
 * See pages, 253, 259, 296. 
 
 t Such a substance is methylacetate, CH 3 COOCH 3 . In this substance 
 the monovalent radicle methyl (page 277) takes the place of a metal. The 
 reaction between methylacetate and potassium hydroxide could be expressed 
 by the following equation : 
 
 CH 3 COOCH 3 + KOH = CH 3 COOK + CH 3 OH, 
 potassium acetate and methyl hydroxide (methyl alcohol) being formed. 
 
378 MASS ACTION. 
 
 the alkaline earths are but little behind the alkalies, while ammonia 
 develops a very slow action. 
 
 Our previous chemical study has shown that chemical changes 
 depend upon the quantity of heat which is produced or absorbed 
 during the various reactions, and that they also depend greatly upon 
 the temperature and other external conditions under which these 
 reactions take place. Now, it is also a matter of common experi- 
 ence that the mass of the active chemical reagents has exactly the 
 same influence as the temperature, in such a way that an increase 
 of the mass * may bring about the same effect as a diminution of 
 temperature, and a diminution of the mass has the same effect as 
 an increase of temperature, and vice versa. This relationship is 
 made clear in reactions in which dissociation takes place ; so, for in- 
 stance, if molecules of N" 2 4 are dissociated by heat to form NO 2 , 
 the reunion of molecules of N0 2 to form N 2 4 will take place the 
 more frequently in the unit of time, the more often contact between 
 the molecules take place; but this contact will take place the of tener, 
 the greater the density of the gas, i.e., the smaller the space which 
 a given quantity has at its disposal ; the density of a gas increases 
 with the pressure upon it, and by this means the active mass in 
 the unit volume becomes greater. From the foregoing it follows 
 that the amount of N0 2 present is smaller the greater the density, 
 but it is also true that the amount of N0 2 will be smaller the lower 
 the temperature. As a result of this and many similar investiga- 
 tions, the law has become well established that the amount of chem- 
 ical action which a substance can exert in any case is proportional 
 to the active mass of that substance which is present in the unit vol- 
 ume. Some of the best illustrations of this law are found in the 
 domain of organic chemistry, and, therefore, are out of place in 
 a book of this kind ; but, nevertheless, as an understanding of some of 
 these cases is necessary for the thorough comprehension of the de- 
 ductions which are to follow, the simple results of one such inves- 
 tigation will be detailed. It has been shown that alcohols can 
 react with acids in much the same way as inorganic hydroxides do. 
 Methyl alcohol (methyl hydroxide)! when brought in contact with 
 acetic acid will produce methyl acetate and water, in a manner simi- 
 
 * I.e., of the quantity of reagent contained in the unit volume. 
 j The alcohols are hydroxides of organic radicles and, in some respects, 
 correspond to inorganic bases. Thus, methyl alcohol is methyl hydroxide; 
 
DOUBLE DECOMPOSITION. 379 
 
 lar to the production of potassium acetate and water from acetic 
 acid and potassium hydroxide ; the difference between the two reac- 
 tions being that methyl alcohol and acetic acid react slowly and (if 
 no provision is made to remove the water produced) incompletely, 
 while potassium hydroxide and acetic acid neutralize each other at 
 once. Now, in such a reaction as that which takes place between 
 methyl alcohol and acetic acid, the first change can be represented 
 as follows : 
 
 MeOH + CH 3 COOH = CH 3 COO Me + H 2 
 
 Methyl hydroxide -\- Acetic acid = Methyl acetate -f Water ; 
 
 however, after a certain amount of reaction has gone on, the mass 
 of the water which is formed, becomes so great that it begins to de- 
 compose methylacetate into methyl alcohol and acetic acid : 
 
 CH 3 COO Me + H 2 = CH 3 COOH + Me OH, 
 
 so that a point is finally reached at which exactly as many mole- 
 cules of methyl acetate are decomposed, as are formed in a given 
 interval of time ; the entire system is then in a state of equilibrium. 
 Obviously, an increase in the mass of water will bring with it an 
 increased decomposition of acetate, and vice versa. 
 
 Similar conditions of equilibrium are encountered when two 
 salts in solution are mixed. 
 
 Let us suppose two substances, A B and C D to be in solution, 
 and let us suppose that A B acts on CD to produce two new sub- 
 stances, A D and B C. At the first instant of the reaction the 
 substances in solution can be expressed by the following : 
 
 AB CD 
 AD BC. 
 
 These changes will go on until a point is reached in which A D and 
 B C will be present in such mass that A D and B C, reacting on 
 each other, will reproduce exactly as many molecules of A B and 
 C D as A B and C D will produce of A D and B C. The solution is 
 then in a state of equilibrium, which can be disturbed, however, by 
 
 ethyl alcohol is ethyl hydroxide, etc. If we designate methyl by the symbol 
 Me, ethyl by the symbol Et, then the relationship between the structure of 
 alcohols and inorganic bases is made plain by the following formulae: 
 
 Me OH, Et OH, KOH, 
 
 Methyl alcohol; Ethyl alcohol; Potassium hydroxide. 
 
380 DISSOCIATION BY SOLUTION. 
 
 increasing the mass of one or the other of the constituents. Let us 
 suppose, however, that A D, which is produced, is either insoluble 
 or volatile, it is then removed from the solution as fast as it is 
 formed, and let us further suppose that B C can have no effect on 
 the insoluble or volatile substance produced. A familiar example 
 of such a change would be the action of silver nitrate on hydro- 
 chloric acid, by which means silver chloride, which is insoluble, and 
 nitric acid are produced : 
 
 AgN0 3 + H Cl = AgCl + HNO. 
 
 In such a case, then, as the silver chloride is removed from the solu- 
 tion as fast as it is produced, no equilibrium can result until the 
 entire mass of silver nitrate has been converted into silver chloride. 
 The same would be true if the substance produced were a gas, and 
 as a consequence, would as certainly be removed as if it were an in- 
 soluble solid. These reactions, known as double decompositions, are 
 among the most common which are to be considered in the chemis- 
 try of the metals. We have already encountered a number of them, 
 an instance being the action of hydrogen sulphide on soluble salts 
 of the heavy metals (see page 100) and the empiric rule has come 
 about as the result of experience that a complete double decomposi- 
 tion takes place when an insoluble or volatile substance is 
 produced.* 
 
 In discussing the determination of molecular weights by means 
 of the lowering of the freezing point of solutions (see page 350), we 
 learned that if quantities of substances, so selected that they are 
 proportional to their molecular weights, are dissolved in equal 
 amounts of the same solvent, then the lowering of the freezing 
 point of that solvent will be the same in each case ; when one 
 molecular weight of a body is dissolved in one hundred molecular 
 weights of water, this depression amounts to 0.62. This rule has, 
 however, not been found to hold good with many salt solutions, for, 
 with these, the depression is much greater than would be expected ; 
 indeed, with sodium chloride or potassium chloride, it is very nearly 
 twice the calculated amount. A solution of sodium chloride, there- 
 fore, behaves as if the quantity of salt represented by the formula 
 
 * The laws -of mass action have of late been greatly developed by a num- 
 ber of prominent chemists ; their discussion is, however, out of place in an ele- 
 mentary text-book. For more detailed information the pupil can refer to 
 Ostwald's Outlines of General Chemistry, Walker's translation. 
 
DISSOCIATION BY SOLUTION. 381 
 
 Na Cl were not one, but were in reality two molecules : it seems 
 probable, therefore, that the mere act of solution dissociates the 
 sodium chloride (NaCl) into its ions (Na-j- Cl), and the same is 
 true, although in a lesser degree, of other salts. It has been 
 pointed out that exactly those substances which (like sodium chlo- 
 ride) are electrolytes (see page 359), are the ones which exhibit this 
 anomalous behavior when in solution. Substances which are not 
 electrolytes follow the usual rule. At first glance it seems rather 
 startling that it is those bodies which, owing to the great energy 
 displayed in their formation, we have regarded as being among the 
 most stable compounds, should be the very ones which are so 
 readily decomposed on solution; but the theory becomes more 
 probable if we remember that salts and acids (the substances disso- 
 ciated by solution) enter into a large number of chemical reactions 
 without difficulty, and that, therefore, their constituent parts are 
 not so firmly united as we would imagine. It must not be thought, 
 however, that we can prove the presence of free chlorine or free 
 sodium in a solution of sodium chloride, unless we have collected 
 the elements at two electrodes by means of an electric current. The 
 dissociated ions themselves, when in solution, are possibly charged 
 with negative and positive electricity, and in this they differ from 
 the free elements. According to the theory which has been out- 
 lined, the neutralization of a base by an acid in solution might be 
 represented as follows (MOH representing the hydroxide of any 
 metal, and HA any acid) : * 
 
 M + OH +H+A = M + A + H 2 
 
 Dissociated base + Dissociated acid = Dissociated salt + water. 
 A glance at the above equation will demonstrate that the change 
 in condition in the entire system consists simply in the formation 
 of water from its ions, H and OH, and this formation is obviously 
 independent of the nature of the acid (AH) and of the base (MOH). 
 This view explains fully the fact that the heat of neutralization is 
 independent of the nature of the base and of the nature of the acid 
 (see page 376). 
 
 Now, it has been demonstrated that all electrolytes which man- 
 
 * See Ostwald; Zeitschrift fur Physikal. Chemie; 7, 423; and also Meyer; 
 Grundziige der Theoret. Cliemie; p. 360 and following 
 
382 DISSOCIATION BY SOLUTION. 
 
 ifest the abnormal lowering of the freezing point which has been 
 mentioned (which are therefore partially dissociated on solution) 
 conduct electricity the more readily the greater the dissociation. 
 Dissociation of an electrolyte seems, therefore, to be essential for 
 its conducting electricity, and it is probably for this reason that 
 pure water and pure liquid hydrochloric acid, for example; do not 
 conduct electricity, but that a solution of hydrochloric acid in water 
 does. At high temperatures, however, many pure substances are 
 electrolytes (pure sodium chloride, pure potassium chloride, for 
 example) ; but then, at high temperatures, the tendency toward 
 decomposition is already a great one. 
 
 With the hypothesis of dissociation in view, it is not necessary 
 to assume that, as has been stated above, two salts, on mixing their 
 solutions, in reality yield four salts by double decomposition. Let 
 us take the following case for an example. Suppose we have a 
 solution of sodium chloride ; this will contain free ions of sodium 
 and of chlorine, and now let us add to this a solution of potassium 
 nitrate (ions, K and N0 3 ) ; obviously, if dissociation is nearly com- 
 plete, such a solution will be identical in every respect with the 
 one produced by mixing sodium nitrate with potassium chloride, a 
 result which is identical with that which would be arrived at were 
 double decomposition (see page 379) to take place. A more com- 
 plete discussion of these topics belongs in a text-book more espe- 
 cially devoted to physical chemistry. It will probably be shown, in 
 time, that all reactions which take place at ordinary temperatures 
 and in solution, are, of necessity, preceded by dissociation. 
 
THE ALKALI METALS. 383 
 
 CHAPTER LII. 
 
 THE ALKALI METALS. 
 
 Lithium, ; symbol, Li ; atomic weight, 7.02. 
 Sodium ; symbol, Na ; atomic weight, 23.05. 
 Potassium. ; symbol, K ; atomic weight, 39.11. 
 Rubidium ; symbol, Eb ; atomic weight, 85.5. 
 Caesium, ; symbol, Cs ; atomic weight, 132.9. 
 
 THE alkali metals are the chemical opposites of the halogens ; 
 the latter are the most negative (not-metallic) elements with wtdch 
 we are acquainted, while the former are the most posit^er^ Tme 
 metallic properties of the alkalies increase with increasing atomic 
 weight, just as the not-metallic properties in the halogen group 
 diminish from fluorine to iodine ; these changes can, perhaps, best 
 be studied by a comparison of the readiness with which the indi- 
 vidual members of both families decompose water. This decompo- 
 sition is of an exactly opposite character, accordingly as the element 
 in question is a member of the halogen or of the alkali group. As 
 a reference to page 75 will recall, the halogens decompose water, 
 liberating oxygen, and they do this the more readily the more neg- 
 ative they are; on the other hand, the alkalies decompose water, 
 liberating hydrogen, and this reaction takes place the more readily 
 the more positive the metal in question is. Fluorine, when brought 
 in contact with water, instantly forms hydrofluoric acid, and sets 
 free oxygen, even in the absence of light; chlorine does so only 
 when the solution of that gas is placed in the sunlight ; bromine 
 enters into this reaction more slowly than chlorine, while iodine 
 has no effect. In the case of the alkali metals, lithium decomposes 
 water, forming lithium hydroxide and hydrogen, but the metal 
 does this quietly, without itself melting or without generating 
 sufficient heat to cause the hydrogen to take fire ; sodium attacks 
 water energetically ; the metal is heated to its melting point, but 
 the hydrogen which is being evolved does not burst into flame ; 
 potassium melts, while sufficient heat is developed to ignite the 
 
384 
 
 THE ALKALI METALS. 
 
 hydrogen ; while both rubidium and caesium enter into the decom- 
 position with explosive violence. The above relationships are more 
 clearly shown by the following table : 
 
 THE HALOGENS. 
 
 DECOMPOSITION OF WATER. 
 
 NOT-METALLIC CHARACTER. 
 
 ELEMENTS. 
 
 ATOMIC WEIGHTS. 
 
 
 
 F, 
 
 19. 
 
 
 - 
 
 Cl, 
 
 35.45 
 
 
 
 
 Br, 
 
 79.95 
 
 Liberating oxygen. 
 
 
 I, 
 
 126.85 
 
 THE ALKALI METALS. 
 
 DECOMPOSITION OF WATEB. 
 
 METALLIC CHARACTER. 
 
 ELEMENTS. 
 
 ATOMIC WEIGHTS. 
 
 M 
 
 H 
 
 Li, 
 
 7.02 
 
 ^r 
 
 y 
 
 Na, 
 
 23.05 
 
 vL 
 
 vL 
 
 K, 
 
 39.11 
 
 T 
 
 T 
 
 Kb, 
 
 85.5 
 
 Liberating hydrogen. 
 
 
 Cs, 
 
 132.9 
 
 The decomposition of water takes place with increasing readi- 
 ness in the direction of the arrows ; the not-metallic character in 
 the halogen family, the metallic character in the alkali family, 
 increase in the direction of the arrows. Of course, the halogens 
 and the alkali metals unite with the greatest energy; the latter 
 burn in an atmosphere of the former with a brilliant light, while 
 most stable halides are produced. 
 
 The alkali metals are soft, malleable, and ductile, and possess a 
 brilliant metallic lustre. When they are exposed to the air, they 
 almost instantly become coated with a layer of oxide ; this oxide 
 absorbs moisture and carbon dioxide, and is soon converted into a 
 mixture of the carbonate and hydroxide ; as a consequence of these 
 changes, pieces of the alkali metals which are in contact with the 
 atmosphere become the centres of small pools of deliquescent hy- 
 droxide. Because of this capacity for oxidation, the metals are 
 kept under some liquid hydrocarbon, such as petroleum naphtha. 
 
 Contrary to the rule observed with not-metallic elements (which 
 form the opposite extremity of the periods), the metals comprising 
 a family which finds its place at the beginning of these periods, 
 
ALKALI METALS ; OXIDES. 385 
 
 show a diminution of their melting points, with an increase in their 
 atomic weights ; this change is evident from the following table : 
 
 Li, specific gravity, .589 atomic volume, 11.9 melting point, 180 
 
 Na, " " .972 " " 23.6 " " 95.6 
 
 K, " " .865 " " 45.2 " " 62.5 
 
 Kb, " " 1.52 *' " 56.2 " " 38.5 
 
 Cs, " " 1.88 " " 70.6 " " 26 . 
 
 The alkaline oxides are soluble in water, and when so dissolved 
 they produce the corresponding hydroxides ; from this it follows 
 that the former cannot be made excepting by processes during 
 which water is rigidly excluded. Each of the elements can be 
 burned in oxygen to form the corresponding oxide, which is con- 
 structed according to the general formula M 2 ; the reaction by 
 which the oxides change to the hydroxides can be represented by 
 the following equation : 
 
 M 2 + Ho = 2 MOH. 
 
 During the change from oxide to hydroxide, and during the subse- 
 quent solution, a large amount of heat is given off ; for instance, in 
 the case of Na 2 0, this amounts to 550 K. The hydroxides are, 
 without exception, soluble ; this solubility increases with the metal- 
 lic nature, and, hence, with the atomic weights and volumes of the 
 alkali metals, the more positive the metal, therefore, the more solu- 
 ble the hydroxide. The rule is so far without exception that, in 
 no other family of metals do we encounter hydroxides which are 
 readily soluble in water.* If we examine the next family to the 
 right of the alkalies in the periodic system, we shall discover that 
 the hydroxides of the two members with the smallest atomic 
 weights (beryllium and magnesium) are insoluble,! while those of 
 calcium, strontium, and barium, although in no case so readily sol- 
 uble as the hydroxides of the alkali metals, have their solubility 
 increased with increasing atomic weight, so that the same rule, hold- 
 ing good with the alkali metals, appertains to this family also. 
 With the families mentioned (the alkalies and the alkaline earths), 
 the list of soluble hydroxides of the metals is practically exhausted, 
 for by far the greater number of hydroxides of the purely metallic 
 elements are insoluble in pure water. 
 
 * The hydroxide of lithium, in the alkali family, is not soluble with great 
 readiness. Lithium is also the least positive of the alkali metals. 
 t That of magnesium is very nearly insoluble. 
 
386 ALKALI METALS; HYDROXIDES. 
 
 / 
 
 The hydroxides of the alkali metals (the caustic alkalies) can be 
 prepared by the action of the lespective metals on water. 
 
 M -f HOH = MOH + H, 
 
 or by covering slaked lime with a solution of an alkaline carbonate, 
 allowing the mixture to stand, and then filtering, when the following 
 reaction has taken place 
 
 M 2 C0 3 + Ca (OH ) 2 = Ca C0 3 + 2 MOH, 
 
 Soluble. Soluble. Insoluble. Soluble. 
 
 The filtered liquid is evaporated, at first in porcelain, and finally in 
 iron or silver dishes.* All of the alkaline hydroxides can be fused 
 without decomposing into the corresponding oxide and water, and 
 the solution of any one of them absorbs carbon dioxide when ex- 
 posed to the air, so that, if it is to be kept for any length of time, 
 it must be placed in closed flasks. The solutions are strongly alka- 
 line in taste and in reaction toward litmus, and neutralize acids with 
 the greatest readiness. 
 
 The sulphides of the alkali metals bear a great resemblance to 
 the oxides ; they are all quite soluble in water, and are formed, as 
 a general rule, by reduction of the corresponding sulphates by heat- 
 ing with charcoal : 
 
 M 2 S0 4 + 4 C = M 2 S + 4 CO, 
 
 or by addition of hydrogen sulphide to a solution of the correspond- 
 ing hydroxide, by which means the sulphhydrate is produced : 
 
 MOH -f H 2 S = MSH -f H 2 0. 
 
 An addition of an equal amount of hydroxide to the sulphhydrate 
 then produces the sulphide : 
 
 MSH + MOH = M 2 S + H 2 0. 
 
 (In this case H 2 S acts like a dibasic acid ; see page 139.) 
 The solutions of alkaline sulphides are able to dissolve sulphur to 
 form so-called polysulphides (see page 155, foot-note). In the 
 cases of sodium and potassium, polysulphides having the following 
 formulae have been isolated : 
 
 Na 2 S 2 , Na 2 S 4 , K 2 S 2 , K 2 S 4 , 
 
 Na 2 S 3 , Na 2 S 5 , K 2 S 3 , K 2 S 5 . 
 
 * Platinum vessels must not be heated with concentrated solutions of 
 alkaline hydroxides, for they are readily attacked by caustic alkalies. 
 
ALKALI METALS; SULPHIDES. 387 
 
 All of these, on addition of acids, form the corresponding salt, 
 hydrogen sulphide, and sulphur, for instance : 
 
 K 2 S 5 + 2 H Cl = 2 K Cl + H 2 S + 4 S. 
 
 When exposed to the air they are oxidized, forming salts of 
 thiosulphuric acid, while sulphur is liberated : 
 
 K 2 S 2 + 30=K 2 S 2 3 , 
 
 K 2 S 3 + 3 = K 2 S 2 3 + S, 
 
 K 2 S 4 + 3 = K 2 S 2 3 + 2 S, 
 
 K 2 S 5 + 3 = K 2 S 2 3 + 3 S, 
 
 The formulae of these sulphides can, perhaps, be best explained on 
 the supposition that they are salts of thio or sulpho acids ; so that 
 the action of sulphur on the monosulphides would be analogous to 
 that of oxygen on the same substances. For example, the sulphide 
 of potassium would be oxidized to the sulphate : 
 
 K 2 S + 40 = K 2 S0 4 ," 
 while the sulphide would be sulphurized to the pentasulphide : 
 
 K 2 S + 4 S = K 2 S 5 . 
 
 According to this theory, the pentasulphide would be the potassium 
 salt of dithio-disulpho sulphuric acid, or : 
 
 SK 
 
 S 
 S 
 SK. 
 
 This theory is, however, sustained only by the fact that the 
 sulphides of the alkalies can take up no more than four atoms of 
 sulphur. The acids corresponding to these sulphides have never 
 been isolated.* 
 
 The sulphhydrates of the alkali metals, corresponding to the 
 hydroxides in formula, are soluble in water, and are produced by 
 the action of hydrogen sulphide on the hydroxides : 
 
 KOH + H 2 S = KSH + H 2 0. 
 
 Both the sulphides and the sulphhydrates are bases ; with acids they 
 form salts and hydrogen sulphide (the oxides and hydroxides form 
 salts and water ; see page 78) : - 
 
 * This theory as to the constitution of the polysulphides was first proposed 
 by Drechsel (Journal fur Prakt. Chemie; [2] 4, 20). It has also been adopted 
 by Remsen in his Inorganic Chemistry, p. 207. 
 
388 ALKALI METALS; HALIDES. 
 
 K 2 +2HC1 = 2KC1 + H S 
 K 2 S + 2HC1 = 2KC1 + H 2 S 
 KOH + H Cl = K Cl + H 2 
 KSH -f H Cl = K Cl + H 2 S. 
 
 The halogen derivatives of the alkalies are extremely stable, and 
 are, throughout, soluble in water. Of these compounds, the most 
 common in occurrence are the chlorides, of which sodium chloride 
 is by far the most often met with. Sodium chloride, or common 
 salt, occurs in extensive beds in rocks of various ages, associated 
 with gypsum, calcite, clay, and sandstone. It frequently occurs in 
 solution in salt springs, and is always found in the sea, of which 
 it forms 2.5 per cent. The salt of commerce is often obtained by 
 evaporating sea water in lagoons by means of the heat of the sun, 
 as is done in France. Lime, gypsum, and ferric hydrate separate 
 at first; afterward the salt begins to crystallize and can be raked 
 out; at last there is left a mother liquor which contains sodium 
 chloride, magnesium chloride, potassium chloride, and magnesium 
 sulphate. In many cases, where the brine obtained from salt 
 springs is evaporated, the mother liquors contain the bromides and 
 iodides of the alkalies. The simple halides of the alkali metals 
 crystallize in the regular system, most frequently in cubes. As we 
 have seen (page 337), the halides of the alkali metals have a 
 tendency to crystallize with other halides in the form of double 
 salts, in which the alkaline halide presumably plays the part of the 
 base. A number of these double halides occur as natural minerals. 
 Among the most important of them is cryolite, Al F 3 , 3 Na F (see 
 page 333). None of the bromides or iodides of the alkalies occur 
 as crystalline mineral individuals. 
 
 As would be expected, the heat of formation of the alkaline 
 halides increases with the increasing metallic nature of the alkali 
 metal forming such a halide. This relationship is readily seen from 
 the following table : 
 
 HEAT OF FORMATION OF THE CHLORIDES. 
 
 Li Cl, 938 K, 
 Na Cl, 976 K, 
 K Cl, 1043 K. 
 
 A distinctive feature of the chemistry of the alkali metals lies 
 in the fact that the salts of these metals are almost without excep- 
 
SODIUM CARBONATE. 389 
 
 tion soluble in water ; they can be produced by neutralizing the 
 hydroxide solutions with the various acids. 
 
 The carbonates of the alkalies are soluble in water, differing in 
 this way from those of the alkaline earths ; however, being the car- 
 bonates of the most positive metals, they are not decomposed by 
 heat, as are the same salts of all other elements ; the carbonate of 
 the least metallic element of the alkali family ( lithium) is also the 
 least soluble in water ; the solubility of the carbonates, as we pass 
 from member to member in this group, increases with the increase 
 in the metallic character, and hence of the atomic weight of the 
 alkali metal producing the salt. The most important carbonates 
 are those of sodium and of potassium, and, as sodium carbonate 
 (common soda) is of great commercial importance, it is advisable to 
 enter into a brief description of the process of its manufacture. 
 
 Sodium carbonate occurs in some mineral waters (Karlsbad) and 
 as a remainder after evaporating the water of alkaline lakes, it 
 is also a constituent of the ashes of sea-plants,* and it was from 
 these that the soda of commerce was made until the end of the last 
 century. During the period of the French Revolution, a large 
 reward was offered for the discovery of a process by means of which 
 sodium carbonate could be prepared from the chloride, as the latter 
 substance was a product which was both cheap and easily purified. 
 Nicholas LeBlanc,f owing to this inducement, discovered a process 
 which has been used with but little modification up to the present 
 day. The chemical changes upon which this method depends are 
 as follows : 
 
 Sodium chloride is treated with sulphuric acid) when hydrochlo- 
 ric acid and primary sodium sulphate are formed : 
 
 Na Cl + H 2 S0 4 = Na HS0 4 + H 01. 
 
 The primary sulphate is then heated with sodium chloride, produ- 
 cing hydrochloric acid and the secondary sulphate : 
 
 Na Cl + Na HS0 4 = Na 2 S0 4 + H Cl. 
 
 The hydrochloric acid which passes off is absorbed by water and is 
 used as ordinary commercial hydrochloric acid. 
 
 The sodium sulphate is next converted into crude soda by heating 
 
 * The ashes of land plants consist mainly of potassium carbonate, 
 t Physician to the Due <T Orleans. 
 
390 SODIUM CARBONATE. 
 
 with anthracite coal and chalk (calcium carbonate), the tempera- 
 ture reaching 1000. Despite the extended attention given to the 
 soda manufacture in the last few years, the chemical processes 
 taking place have not been definitely settled, yet the following are 
 most frequently accepted as nearly correct : Sodium sulphate first is 
 reduced to sodium sulphide by the coal : 
 
 Na 2 S0 4 + 2 C = Na 2 S + 2 C 2 ; 
 
 and the sodium sulphide, together with chalk, then changes to 
 sodium carbonate and calcium sulphide : 
 
 Na 2 S -f> Ca C0 3 = Na 2 C0 3 + Ca S. 
 
 The crude soda is extracted with water and the liquors are evapo- 
 rated, when tolerably pure sodium carbonate crystallizes. This, when 
 slowly crystallized from water, separates with ten molecules of that 
 substance and forms crystals of commercial soda, Na 2 C0 3 -f- 10 H 2 . 
 The latter effloresce when in contact with the air, losing 9 molecules 
 of water and changing into a powder, which has the composition 
 
 A modern process, known as the ammonia-soda piwcess, has of late 
 succeeded to a large extent in taking the place of the older method. 
 This late improvement depends on the fact that primary sodium 
 carbonate is not very readily soluble in cold water. Ammonia 
 solution, saturated with an excess of carbon dioxide, contains pri- 
 mary ammonium carbonate (NH 4 ) HC0 3 , and when this is added to 
 a solution of common salt the following change takes place : - 
 
 ( NH 4 ) HC0 3 + Na Cl = Na HC0 3 + NH 4 Cl ; 
 
 the primary sodium carbonate (sodium bicarbonate) separates as a 
 crystalline powder, and the latter, when heated, gives off water and 
 carbon dioxide, leaving the secondary carbonate (see page 291) : 
 
 2NaHC0 3 = Na 2 CO + H 2 + C0 2 . 
 
 The nitrates of sodium and of potassium are of importance in the 
 manufacture of gunpowder. Sodium nitrate is found in the north- 
 ern part of Chile,* where it occurs in extensive deposits, accom- 
 panied by sodium chloride and other salts, the presence of which 
 seems to indicate that the formation of the nitrate is due to the 
 decay of marine plants, the occurrence of these deposits in this 
 place being one of the proofs of the theory that this portion of 
 
 * Province of Tarapaca ; the nitrate is called Chile saltpetre. 
 
SODIUM AND POTASSIUM NITRATES. 391 
 
 South America was at one time submerged. Sodium nitrate is puri- 
 fied by washing with water and recrystallization ; the mother liquors, 
 which are left, contain considerable quantities of the iodides and are 
 used in the manufacture of iodine. Unfortunately, sodium nitrate, 
 because it is hygroscopic, cannot be used in the preparation of gun- 
 powder ; so that, preliminary to the production of that explosive, the 
 nitrate of sodium must be converted into nitrate of potassium. This 
 is accomplished by treating a saturated solution of sodium nitrate, 
 at boiling heat, with a solution of potassium chloride, when a double 
 decomposition, accompanied by the formation of the less soluble 
 chloride of sodium, takes place : 
 
 Na N0 3 + K Cl = Na Cl + KN0 3 , 
 
 The solution of potassium nitrate is filtered and allowed to crystal- 
 lize. Potassium nitrate occurs as a mineral deposit in many places, 
 where nitrogenous organic matter is decaying in the presence of pot- 
 ash (see page 204) ; localities in which this natural production of 
 potassium nitrate assumes commercial importance are found in Spain, 
 Egypt, Peru, and especially India, from which latter country potas- 
 sium nitrate is exported in considerable quantities. 
 
 The use of potassium nitrate in the manufacture of gunpowder 
 depends on its oxidizing powers. When potassium nitrate is mixed 
 with charcoal and ignited, the following reaction takes place. 
 
 2 KN0 3 + 3 C =C0 2 + CO + K 2 C0 3 + 2 N ; 
 
 the carbon dioxide, which is left in combination as potassium carbon- 
 ate, can be further liberated by the previous addition of sulphur : 
 
 2 KN0 3 + 3 C + S == 3 C0 2 + 2 N + K 2 S. 
 
 The formation of such a large amount of gaseous material from the 
 small volume of solid causes the explosion. The reaction given is, 
 however, only approximately correct, for other changes, not defi- 
 nitely understood, also take place. 
 
 The other salts of the alkalies have been sufficiently mentioned 
 in the course of the chapters in which the various acids have been 
 discussed. The following table will make the formulae and solubility 
 of some of the salts most apparent : 
 
 ALKALI METALS. 
 
 OXIDES, M 2 O, converted to the hydroxides by addition of water. 
 HYDROXIDES, MOH, soluble in water; least soluble is Li OH; solubility 
 increases with increasing atomic weight of the alkali metal. 
 
392 SPECTROSCOPE. 
 
 SULPHIDES, M 2 S, soluble, probably converted into MSH + MOH by 
 
 addition of water. 
 
 SULPHHYDRATES, MSH, soluble in water; MOH + H 2 S = MSH + H 2 O. 
 CARBONATES, M 2 CO 3 and MHCO 3 , soluble in water; least soluble is 
 
 Li 2 COg ; solubility increases with increasing atomic weight. 
 NITRATES, MNO 3 , soluble in water, change to nitrites and oxygen when 
 
 heated. (See page 208.) 
 
 SULPHATES, M 2 SO* and MHSO 4 . Soluble in water. 
 PHOSPHATES, MH 2 PO 4 , M 2 HPO 4 , M 3 PO 4 , tertiary phosphates change 
 
 to secondary phosphate and hydroxide of alkali metals on addition of 
 
 water. (See page 230. ) 
 SILICATES, M 2 SiO 3 , soluble in water. 
 POTASSIUM PERCHLORATE, KC1O 4 , potassium fluosilicate, K 2 SiF 6 , are 
 
 soluble with difficulty, and hence are precipitated from solutions of 
 
 potassium salts by addition of the corresponding acids, sodium pyro- 
 
 antimonate, Na 2 H 2 Sb 2 O 7 ; insoluble in cold water. 
 
 The salts of ammonium correspond entirely to the salts of potas- 
 sium, and hence are frequently discussed in connection with the 
 alkali metals. Their nature has been sufficiently explained on page 
 188 and sub. 
 
 Detection of metals by means of the spectroscope. 
 
 The readiest means for the detection of the various alkali metals 
 is by means of the spectroscope ; indeed, caesium and rubidium 
 were not known to exist until the examination, by the spectroscope, 
 of the residues left by evaporation of certain mineral waters re- 
 vealed their presence. The principle upon which the use of the 
 the spectroscope depends is as follows : 
 
 Light which contains waves of only one wave-length is mono- 
 chromatic (homogeneous). When a ray of such light, passing 
 through the air, comes in contact with a transparent medium of 
 greater density, it is refracted toward a line normal to the surface 
 of the latter, and the smaller the wave-length, the greater is the 
 refraction ; for each wave-length there is a corresponding index of 
 refraction, provided the media through which each kind of light 
 passes remain the same. If a ray of light contains waves of various 
 lengths, then each kind of wave will be refracted according to its 
 refractive index, so that the whole will be separated into as many 
 monochromatic rays as it contained different wave-lengths. Such 
 a ray of light, falling from a narrow slit upon a prism, the edge of 
 which is parallel to the slit, produces a series of parallel images of 
 the opening; and if the light consists- of all of the colors between 
 
THE SPECTROSCOPE. 
 
 393 
 
 two determined extremes, the image obtained appears as a continu- 
 ous spectrum, produced by a number of different colored images of 
 the slit, which merge the one into the other. Light which is 
 emitted by the sun or by white-hot bodies, contains an infinite 
 number of waves which differ in length ; this light is white ; when 
 passed through a prism it gives a spectrum containing all the colors 
 of the rainbow, beginning with red and passing through the various 
 modifications of color (orange, yellow, green, blue) to violet, at the 
 opposite extremity. 
 
 Fig. 12. 
 
 The spectroscope (Fig. 12) consists of a telescope (A) which throws 
 parallel rays of light admitted through a small slit at (S) upon a prism (P); 
 the spectrum formed is observed by the telescope at B, which is so focussed as 
 to give a sharp image of the same; at the same time a mirrored image of a 
 millimeter scale, photographed and placed in C, is so reflected as to be visible 
 above the spectrum when the observer glances through B. The manner of 
 opening and closing the slit by means of the thumb screw (e) is shown by the 
 small figure (d). 
 
 The spectrum of a white-hot solid, when so observed, is contin- 
 uous, but this is not the case with, glowing gases. These, when 
 examined by the spectroscope, show a number of bright, colored 
 lines upon a black or nearly black background. The colors and 
 relative position of the lines are definite ones for each individual 
 glowing gas, and are always different for gases of differing chem- 
 ical composition. The reason for the appearance of these lines is 
 that the glowing gases emit light only of certain determined wave- 
 lengths, the varieties, and hence the colors, of which are generally 
 
 
394 
 
 SPECTRUM ANALYSIS. 
 
 few in number ; as a consequence, light of each wave-length, being 
 refracted according to its index of refraction, appears in a different 
 place on the spectrum as a sharp line of the color belonging to that 
 particular wave-length. 
 
 The flame of a Bunsen burner is not luminous, but if a platinum 
 wire is placed in it, the latter becomes heated and emits a white 
 light. If, now, the wire is coiled as in Fig. 13, and is (after moist- 
 ening with a little hydrochloric acid) 
 dipped into a little sodium chloride, 
 the adhering salt, when brought Fig ' 13 ' 
 
 into the flame, will vaporize and will emit a pure yellow light. In 
 the same way, the light emitted by potassium salts will be violet ; * 
 by lithium or strontium, red ; by copper, barium, or thallium, green ; 
 by zinc, blue, etc.f If the yellow sodium light, obtained as above, 
 is observed by means of a spectroscope, a bright yellow line on a 
 dark background is seen. The position of this yellow line corre- 
 sponds to the position occupied by yellow in the continuous spec- 
 trum. Lithium will show a red line and a less marked yellow one, 
 
 potassium a red line and a blue 
 one. In fact, each individual metal 
 displays characteristic lines in defi- 
 nite parts of the spectrum, while 
 the lines of no two metals corre- 
 spond exactly. In order, then, to 
 discover the presence of any metal 
 or metals in a mixture, it is only 
 necessary to place volatile com- 
 pounds of those elements in the 
 not-luminous flame and then to ob- 
 The minutest traces of the metals 
 in question can be detected in this way. 
 
 If a not-luminous flame is placed in front of a luminous back- 
 ground which emits a white light, and if then some sodium salt is 
 
 * Best seen when the potassium flame is observed through a piece of blue 
 glass. When sodium is also present, the blue absorbs the yellow rays, while 
 permitting the violet ones to pass through. 
 
 t The contrasting colors of the various flames can be best observed by 
 using a lamp with four or five burners (Fig. 14) and, after fixing platinum 
 wires in the stands, as shown in the figure, bringing the entire number simul- 
 taneously into the lighted burners. 
 
 Fig. 14. 
 
 serve the spectrum produced. 
 
SPECTRUM ANALYSIS. 395 
 
 volatilized in this not-luminous flame, the spectroscope will show a 
 continuous spectrum, due to the white light, with this difference, 
 however, that in the place where the yellow band of light belonging 
 to the sodium spectrum usually occurs, there is now seen a black 
 band. This phenomenon is due to the fact that when a ray of 
 white light containing all colors is passed through a glowing gas 
 which emits light rays only of certain definite colors, this glowing 
 gas is able to absorb from the white light the rays of exactly the 
 same color as those which itself emits. It follows that, when 
 white light is passed through the glowing vapor of a potassium 
 compound, there will appear (on the continuous spectrum) a dark line 
 in the red and one in the blue ; in the case of lithium, a dark line 
 in the red and one in the yellow, etc. Such spectra are called ab- 
 sorption spectra. The spectra of the sun and of the fixed stars are 
 not perfectly continuous, but are traversed by a series of fine, dark 
 lines which have been proven to correspond exactly to the absorp- 
 tion spectra of the glowing vapors of the elements with which we 
 come in contact on the earth ; the spectra of the sun and the fixed 
 stars are, therefore, absorption spectra, caused by the white light 
 of the glowing central mass passing through the surrounding chro- 
 mosphere. The gaseous envelope of the sun must therefore contain 
 the glowing vapors of a number of elements ; and those elements 
 are identical with the ones encountered on the earth.* By means 
 of the spectroscope we have, therefore, been able, to a great extent, 
 to analyze the composition of the gases which surround the sun and 
 the fixed stars. 
 
 * Some lines appear in the spectrum of the sun which do not correspond 
 to those emitted by any known element on the earth. 
 
396 COPPER, SILVER, AND GOLD. 
 
 CHAPTEK LIIL 
 
 COPPER, SILVER, AND GOLD. 
 
 Copper ; symbol, Cu; atomic weight, 63.6; 
 Silver; symbol, Ag; atomic weight, 107.92 ; 
 Gold ; symbol, Au ; atomic weight, 197.3. 
 
 THESE three elements find their places at the beginning of the 
 second section of the long periods, and, because the second portion 
 of the latter shows many points of resemblance to the first, we must 
 expect copper, silver, and gold to be in some respects like the alkali 
 metals, of which family they form the secondary group (see page 
 371). Naturally, as copper, silver, and gold are much nearer to the 
 not-metallic end of the long periods than are the alkali metals, we 
 must not, in the chemical behavior of the former, look for metallic 
 properties by any means so pronounced as are encountered with the 
 latter ; this difference is manifested in marked degree by the fact 
 that neither copper, silver, nor gold decomposes water; the resem- 
 blances between the elements of this group and the alkalies are 
 confined chiefly to their univalence, by reason of which each element 
 forms halogen derivatives having the formula MX (corresponding 
 to those of the alkali metals), and to the isomorphism between the 
 crystalline form of some of the compounds derived from these ele- 
 ments in their univalent condition and similar compounds of the 
 alkalies (of the latter, especially those of sodium) ; on the other 
 hand, copper and gold differ very widely from the alkalies by being 
 able to form higher oxides and salts derived from these ; in the case 
 of copper, this higher oxide has the formula Cu ; of gold, Au 2 3 . 
 The oxides, M 2 0, and the hydroxides and halides derived from 
 these are, therefore, the typical compounds belonging to the entire 
 family comprising the alkali metals, as well as copper, silver, and 
 gold ; but those members of this family which find their places 
 cut the beginning of the secondary division of the long periods are also 
 able to form oxides in which the valence of the elements is more 
 than one. In the formation of the latter oxides, the elements in 
 
COPPER : OCCURRENCE. 
 
 397 
 
 question (Cu, Ag, Au) appear as connecting links between the ele- 
 ments of the eighth group (Fe, Co, Ni, etc.) and the second halves of 
 the long periods. Copper, silver, and gold find their places as the 
 minima and the beginning of the second portions of the curves of 
 atomic volumes (see page 365) ; they are, therefore, malleable, duc- 
 tile, fusible, electropositive, and good conductors of electricity. 
 
 Their chief physical constants are given* in the following 
 table : 
 
 
 ATOMIC WEIGHT. 
 
 SPECIFIC GBAVITY. 
 
 ATOMIC VOLUME. 
 
 MELTING POINT. 
 
 Copper, 
 
 63.6 
 
 8.8 
 
 7.1 
 
 1050 
 
 Silver, 
 
 107.92 
 
 10.5 
 
 10.1 
 
 950 
 
 Gold, 
 
 197.3 
 
 19.3 
 
 10.2 
 
 1030 
 
 All of these elements are volatilized when heated in the flame 
 of the oxyhydrogeii blowpipe. They crystallize, as do the alkali 
 metals, in forms belonging to the regular system. 
 
 The most important mineral forms in which these elements occur 
 are given in the following table : 
 
 COPPER. As native copper, Lake Superior region, Siberia, Chile, Aus- 
 tralia; as chalcocite (cuprous sulphide), Cu 2 S, in Cornwall, Siberia, 
 Saxony, Western Montana ; as chalcopyrite (copper pyrites), Cu Fe S 2 , 
 similar localities to chalcocite.; as cuprite, Cu 2 O, in Lake Superior 
 regions; as melaconite, CuO, in Lake Superior regions. 
 
 Copper also frequently occurs, combined with arsenic and anti- 
 mony trisulphides, in the mineral tetrahedrite (gray copper ore, 
 fahl-ore), which has approximately the formula, 4 (Cu 2 , Ag 2 , Fe, 
 Zn) S, Sb 2 S 3 , or 4 (Cu 2 , Fe, Zn) S, As 2 S 3 . As the formulae will 
 show, iron and zinc accompany copper in these ores, while silver is 
 also frequently encountered therein.* 
 
 In addition to the above sulphur compounds, the basic carbon- 
 ates of copper, malachite, Cu 2 (HO) 2 C0 3 , and lazurite, Cu 3 (H0) 2 
 (C0 3 ) 2 , are important minerals. 
 
 * In this mineral two atoms of copper replace one of zinc or of iron iso- 
 morphously. The possibility of this replacement is probably due to the fact 
 that the size of the molecule prevents an undue influence on the crystalline 
 form by one or two atoms; nevertheless, it serves admirably to illustrate the 
 necessity of great caution in using the laws of isomorphism for the purpose of 
 determining atomic weights (see page 356). 
 
398 COPPER; METALLURGY. 
 
 SILVER. As native silver, in the United States, Mexico, Peru, Norway, 
 Saxony, Bohemia, Siberia, etc. 
 
 As argentine, Ag 2 S, isomorphous with chalcocite, Cu 2 S ; silver 
 sulphide also occurs in conjunction with lead sulphide and in nu- 
 merous minerals in which it is combined with the sulphides of 
 antimony, Sb 2 S 3 , arsenic, As 2 S 3 , and iron, Fe 2 S 3 . 
 
 GOLD. As native gold in quartz veins in conjunction with iron pyrites, 
 chalcopyrite, galena, and other sulphides. Gold particles also occur 
 in the gravel or sand of rivers or valleys in auriferous regions, or on 
 the slopes of mountains or hills whose rocks contain in some part 
 auriferous veins. 
 
 Compounds of gold are very infrequent as minerals. The tellu- 
 ride of gold and silver, (Ag, Au) 2 Te, is sometimes found. 
 
 METALLURGY OF COPPER. The copper ores, the most valuable 
 of which are the oxides and sulphides, are roasted; by this means 
 the volatile compounds of arsenic and antimony are removed, while 
 the sulphides of iron, which are present, are easily converted into the 
 oxide, the sulphur passing off as sulphur dioxide. If the material 
 used in the preparation of copper should contain large quantities of 
 quartz, the latter substance, in melting, will attack the ferric oxide 
 in order to form the silicate of iron, which can be run off in the 
 form of slag ; if the ores used do not already contain silicon diox- 
 ide, the latter must be added as quartz or sand. The product ob- 
 tained after the smelting is much richer in copper than the original 
 ore, and contains both cuprous and cupric oxides as well as the 
 corresponding sulphides. This material is again roasted and the 
 remaining iron separated as slag, while the oxides and sulphides of 
 copper mutually reduce each other as follows : 
 
 Cu 2 S + 2 Cu a == 6 Cu + S0 2 , 
 
 Cu 2 S+2CuO = 4Cu-fS0 2 . 
 
 It is frequently necessary to repeat the roasting process several 
 times while adding the coal and sand. Finally, the fused copper is 
 stirred with poles of green wood. Not infrequently the copper 
 ores contain considerable quantities of silver, so that the separation 
 of the latter becomes profitable. The processes used are, however, 
 somewhat complicated and need not be entered into in this work. 
 
 SILVER. Considerable quantities of silver are obtained from 
 galena (PbS, lead sulphide), and separated from the lead by the 
 process of cupellation, a description of which was given on pages 
 
SILVER; METALLURGY. 399 
 
 320 and 321. The relative amounts of silver in the lead ore may 
 vary to such an extent that, at one time, the silver is the chief 
 product of the process while, at another, it is present in such small 
 quantities that lead forms the only commercially valuable substance. 
 In some cases the ores are ground ; the richer ones are roasted with 
 common salt so as to convert the silver compounds into silver chlo- 
 ride ; the crushed mixture is then treated with mercury and hot 
 water, either in barrels or in cast-iron pans, the mercury taking the 
 place of silver in the chloride : 
 
 AgCl+Hg = Ag + HgCl, 
 
 while the excess of the fluid metal dissolves the silver to form an 
 amalgam. This amalgam is then washed, strained, and carefully 
 distilled from cast-iron retorts ; the mercury passes off and the re- 
 maining silver is cast into ingots. Several processes of silver ex- 
 traction by the so-called " wet way " depend on the conversion of 
 the silver ores into the chloride by means of a solution of sodium 
 chloride, the extraction of the chloride of silver by means of a solu- 
 tion of sodium hyposulphite (thiosulphate), in which substance sil- 
 ver chloride is soluble, and the subsequent precipitation of silver 
 sulphide from this solution by means of the sulphide of sodium. 
 The sulphide of silver is separated and roasted, by means of which 
 process the sulphur burns off and metallic silver remains. Chemi- 
 cally pure silver is prepared by boiling pure silver chloride with a 
 mixture of potassium hydroxide and grape sugar. 
 
 GOLD. As gold almost exclusively occurs in the form of the 
 native element, the process of its extraction consists simply in a 
 separation of the various impurities which accompany the metal. 
 Alluvial washing (or placer digging) is done by placing the aurifer- 
 ous deposit found on the banks of rivers or -in the valleys in shal- 
 low pans, and then washing off the lighter portions, while the 
 specifically heavier gold remains behind, mixed with pebbles and 
 stones. From the latter it can be mechanically separated. The 
 former gold beds, having become to a great extent exhausted, the 
 process of hydraulic mining is now frequently resorted to. The sides 
 of the hills which contain gold-bearing conglomerate are washed out 
 by means of powerful streams of water ; the washings are conducted 
 through a channel containing a number of sluice boxes which collect 
 the heavier particles. Mercury is placed in each of these boxes 
 
400 COPPER; SILVER; GOLD; CHEMICAL ACTION. 
 
 (because that metal is capable of forming an amalgam with the 
 smaller particles of gold) ; the sluice boxes are opened from time to 
 time, and the metal contained in them is mechanically separated. 
 Where the gold occurs in veins imbedded in quartz, the material is 
 mined, crushed, and the gold extracted by means of mercury ; the 
 amalgam so formed is treated as is that of silver. 
 
 Copper, silver, and gold all have a metallic lustre and are 
 malleable and ductile ; neither silver nor gold will burn ; indeed, the 
 oxides of these metals decompose when heated ; copper, on the 
 other hand, burns when heated in oxygen to a high temperature, 
 the compound formed being cupric oxide. Chlorine and bromine 
 attack all three of the metals, forming the corresponding halides, 
 while iodine attacks copper and silver. 
 
 Nitric acid readily dissolves either copper or silver ; forming 
 cupric nitrate, Cu (N0 3 ) 2 , and silver nitrate, AgN0 3 , respectively; 
 the acid does not attack gold (page 207, b). 
 
 Aqua regia (see page 203) attacks either copper, silver, or gold, 
 producing the corresponding chlorides. 
 
 Sulphuric acid, when hot and concentrated, dissolves either 
 copper or silver, producing cupric sulphate, Cu S0 4 , and argentic 
 sulphate, Ag 2 S0 4 ; sulphur dioxide is at the same time given off 
 (see page 136). 
 
 Caustic alkalies, in solution, hot and concentrated, dissolve gold. 
 
 The alloys of the three metals are quite important. Two parts 
 of copper alloyed with one part of zinc form a yellow metal ( brass ) ; 
 alloys of copper and tin are known as bell metal, gun metal, and 
 bronze, according to the proportions of the ingredients; ninety 
 parts of copper united with ten parts of aluminium is the most 
 common form of aluminium bronze. Commercial silver is always 
 alloyed with copper, as pure silver is too soft for the ordinary pur- 
 poses of coinage and the manufacture of jewelry ; the silver coins 
 in use contain from 7.5 to 10 per cent of copper ; gold is also 
 invariably alloyed with copper or silver, the resulting alloys being- 
 much harder than pure gold.* 
 
 * The fineness of gold is measured in carats; the number of carats used 
 in designating a particular alloy of gold indicate the number of parts of pure 
 gold contained in twenty-four parts of alloy; thus, 18 carat gold has IS parts 
 of gold in 24 parts of alloy. 
 
CUPROUS COMPOUNDS. 401 
 
 In their chemical behavior, copper, silver, and gold differ quite 
 markedly, and for this reason it will be necessary to discuss the 
 chemistry of each metal separately. 
 
 Copper forms two oxides,* cuprous oxide, Cu 2 0, and cupric 
 oxide, Cu 0. Cuprous oxide is found in nature as the mineral cu- 
 prite, occurring in octahedra or in cubes ; in the laboratory the oxide 
 can be produced by heating a mixture of copper and cupric oxide to 
 aredheat:- C u + CuO = Cu 2 O. 
 
 A hydroxide corresponding to this oxide is unknown, but if 
 a solution of copper sulphate, mixed with glucose, is warmed with 
 alkalies, a yellow precipitate which has the formula 4 Cu 2 -f- H 2 
 is produced ; this substance is not completely dehydrated until a tem- 
 perature of 360 is reached ; when exposed to the air this hydrate 
 rapidly takes up oxygen, forming blue cupric hydroxide, Cu (OH) 2 . 
 Cuprous oxide dissolves in ammonia to produce a colorless fluid, 
 which, however, soon absorbs oxygen from the atmosphere and 
 assumes a deep blue color. Cuprous oxide is very easily decom- 
 posed by many oxy-acids, and for this reason very few cuprous 
 salts are known; dilute nitric or sulphuric acids attack it, liber- 
 ating metallic copper and producing the corresponding cupric 
 
 Cu 2 + H 2 S0 4 = Cu S0 4 -f H 2 + Cu. 
 
 The most important cuprous salts are the cuprous halides. 
 
 Cuprous chloride, Cu Cl, is a white solid which is with difficulty 
 soluble in water. In this respect it resembles the corresponding 
 chloride of silver. Like the latter it is readily dissolved by 
 ammonia, with which it forms a substance having the formula 
 NH 3 CuCl; the latter compound possibly consists of ammonium 
 chloride, NH 3 H Cl, with the difference that one atom of hydrogen 
 in the formula weight has been replaced by- one of copper, so that 
 in the formation of this body cuprous chloride would act as does 
 hydrochloric acid in ammonium chloride. Cuprous chloride is 
 readily oxidized when exposed to the air. 
 
 Cuprous iodide is the only iodide of copper which exists. It is 
 formed by adding the solution of an iodide to a copper sulphate 
 solution ; the cupric iodide, which we should expect to be formed, 
 at once breaking down into cuprous iodide and iodine : 
 
 * An oxide Cu 4 O has also been described. 
 
402 CUPRIC COMPOUNDS. 
 
 Cu S0 4 + 2 KI = K 2 S0 4 + Cu I + I.* 
 
 Cupric oxide, Cu 0, is the most stable oxide of copper. It can 
 readily be formed by heating copper in a current of air or of oxygen, 
 or by decomposing cupric nitrate, Cu (N0 3 ) 2 , (page 208, /?). It is 
 a black substance which readily loses oxygen when it is heated in a 
 current of hydrogen (see page 39) or with charcoal : 
 
 Cu -f C = Cu -f CO. 
 
 Cupric hydroxide, Cu (OH) 2 , is formed in a manner parallel to 
 the formation of the hydroxides of most metals i.e., by precipitation 
 from the solution of a copper salt upon the addition of a soluble 
 hydroxide. It appears as a blue, flaky precipitate, which readily 
 loses water when it is warmed ; it then turns black and forms cupric 
 oxide. Ammonia water dissolves both the oxide and hydroxide; 
 the solution has a deep blue color which can be observed even when 
 only very little copper is present. Both cupric oxide and hydroxide 
 are bases ; they dissolve in acids to form stable cupric salts. The 
 latter, when they contain water of crystallization, are generally blue 
 or green. 
 
 Cupric chloride, Cu C1 2 , is formed by the action of chlorine on 
 copper, or by dissolving the oxide or hydroxide in hydrochloric 
 acid, and then evaporating the solution and drying at 100 ; when 
 anhydrous, cupric chloride is a brown powder ; when crystallized 
 from water it forms green crystals of the composition Cu C1 2 -f-2 H 2 ; 
 it is readily soluble in water. The chloride forms double salts with 
 the chlorides of the alkali metals. The bromide resembles the 
 chloride in every respect. , 
 
 Cupric sulphate, Cu S0 4 + 5 H 2 ( blue vitriol), is formed by 
 dissolving copper in sulphuric acid (page 137). t The commercial 
 product is prepared by heating copper sulphide in a current of air, 
 extracting with water and recrystallizing. It forms large, blue 
 crystals belonging to the triclinic system ; it is readily soluble in 
 water. When exposed to the air the crystals effloresce and lose 
 two molecules of water of crystallization ; at 100 two more mole- 
 
 * The iodine which is liberated can be removed by the addition of a 
 reducing agent such as SO 2 (see page 139). 
 
 + In this reaction several secondary products [cuprous sulphide, Cu. 2 S, 
 and compounds, (CuaS, CuO,) and (CuS, CuO)] are produced. 
 
CUPRIC SALTS. 403 
 
 cules pass off, so that a salt of the composition Cu S0 4 + H 2 is 
 left ; the latter is probably a secondary salt of the hydrated sul- 
 phuric acid having the formula H 4 S0 5 , and should therefore be 
 written CuH 2 S0 5 . At 230 the two hydroxyl groups present in 
 this salt (see page 153) finally separate water, and leave a white 
 powder which has a composition expressed by the formula Cu SO 4 ; 
 as soon as water is added to this, a blue solution containing 
 Cu S0 4 -|-5H 2 O is produced; the latter belongs to a class of sul- 
 phates known as the vitriols (see magnesium). Copper sulphate can 
 unite with ammonia to form compounds in which molecules of 
 ammonia take the place of molecules of water of crystallization,* 
 for instance, salts having the composition Cu S0 4 + 5 NH 3 , 
 Cu S0 4 + 4 ISTHg + H 2 0, Cu S0 4 + 3 NH 3 + 2 H 2 0, etc., are capa- 
 ble of existence. These substances furnish examples of cases where 
 the compound of nitrogen and hydrogen plays the same role as the 
 compound of oxygen and hydrogen (see page 299). An adequate 
 chemical explanation of the existence of salts with water of crys- 
 tallization and of these ammonium compounds, provided we adhere 
 to our present theories of valence, is as yet lacking. f A number of 
 basic sulphates of copper have been described. 
 
 Cupric nitrate, Cu (N0 3 ) 2 -f 3 H 2 0, is formed by dissolving 
 copper in nitric acid (see page 198) or by adding nitric acid to 
 cupric oxide or hydroxide. The salt exists in blue, prismatic crys- 
 tals which are soluble in water, and which break down completely 
 into cupric oxide and nitrogen peroxide when heated (see page 208). 
 
 Basic carbonates of copper. The secondary, normal carbonate, 
 Cu C0 3 , is unknown. A basic carbonate which has the same com- 
 position as the mineral lazurite is produced by adding the solution 
 of the carbonate of an alkali metal to a solution of a copper salt. 
 
 * The same is true of cupric chloride, Cu CL> + 2 H 2 O, for a substance of 
 the composition Cu C1 2 + 2 NH 3 is known. 
 
 t The nitrogen atom in ammonia is unsaturated, as is proved by the easy 
 formation of ammonium compounds. The resemblance between ammonia and 
 water in these salts would lead us to suspect that the oxygen atom is also un- 
 saturated in water. This conclusion is not very startling if we take into con- 
 sideration the great resemblance between oxygen and sulphur, the latter 
 element being able to take part in compounds in which its atoms are hexa- 
 valent. It has also been shown that methyl ether (CH 3 O CH 3 ) can add 
 hydrochloric acid, just as ammonia would; the unstable compound so formed 
 probably contains tetravalent oxygen. 
 
404 COPPER; SULPHIDES. 
 
 The carbonate is insoluble in water, and has the composition 
 Cu(OH) 2 CuC0 3 or:- Cu _ OH 
 
 )C0 3 
 Cu OH. 
 
 Other basic carbonates (malachite, for example) occur in the form 
 of minerals. All copper carbonates, when heated to 300, lose car- 
 bon dioxide and leave cupric oxide. The green coating formed 
 on metallic copper which has been exposed to moist air, is due to 
 the formation of a basic carbonate. 
 
 Sulphides of copper. Copper forms two sulphides which, in 
 formula, correspond to the oxides. Cuprous sulphide, Cu 2 S, occurs 
 as the mineral chalcocite and is also, most probably, a constituent of 
 chalcopyrite, for the latter substance is regarded as a sulpho-salt in 
 which cuprous sulphide is the base and ferric sulphide the acidic 
 anhydride:- C ^ + ^ = Gn ^ t * 
 
 Cuprous sulphide is a substance which is formed with remarkable 
 readiness, for it can be produced by merely pressing copper and sul- 
 phur firmly together,! or by heating or even rubbing copper with a 
 sufficient quantity of sulphur to form cuprous sulphide. 
 
 * Cupric sulphide.! Cu S, is produced by precipitating a slightly 
 acid solution of a copper salt by means of hydrogen sulphide (see 
 page 100). Cupric sulphide dissolves in hot, concentrated hydrochlo- 
 ric acid, forming hydrogen sulphide and cupric chloride. Nitric 
 acid dissolves it, producing cupric nitrate, Cu (N0 3 ) 2 , while the 
 hydrogen sulphide, which is liberated, is oxidized to sulphur (see 
 page 206). The precipitate is nearly black, and, when moist, is 
 readily oxidized to cupric sulphate when it is exposed to the air : $ 
 
 Cu S + 4 = Cu S0 4 . 
 
 When heated to 200, in a current of hydrogen, it is changed into 
 cuprous sulphide. 
 
 * One-half the formula weight of the above salt would lead to the usual 
 formula of chalcopyrite, namely, Cu Fe S 2 . 
 
 t This sulphurizing action reminds us forcibly of the oxidation of copper 
 in the air. 
 
 f This oxidation is taken advantage of in the commercial formation of 
 blue vitriol from sulphides of copper. 
 
SILVER ; SALTS OF. 405 
 
 Compounds of silver. Silver is univalent in nearly all of its 
 compounds. Its most stable oxide is Ag 2 0, but oxides having the 
 formulae Ag 4 * and Ag are also known. The most stable oxide, 
 Ag 2 0, is produced by adding a soluble- hydroxide to a solution of a 
 silver salt. It is a black precipitate which readily dissolves in an 
 excess of ammonia (compare with cuprous oxide). When heated it 
 readily breaks down into silver and oxygen. 
 
 Silver chloride, Ag Cl, one of the most characteristic silver salts, 
 is insoluble in water and is therefore produced as a white precipitate 
 when hydrochloric acid, or the solution of a chloride, is added to a 
 solution of a silver salt : 
 
 AgN0 3 + H Cl = HN0 3 +AgCl,t 
 Ag N0 3 + Na Cl = .Na N0 3 + Ag Cl. 
 Soluble. Soluble. Soluble. Insoluble. 
 
 Owing to this insolubility, the formation of silver chloride affords 
 a ready means of detecting the presence either of silver or of a 
 chloride in a solution. When silver chloride is exposed to the light, 
 it rapidly changes color, becoming violet at first ; that color however 
 soon becomes darker until the entire mass turns black. The same 
 change also takes place with the equally insoluble bromide or iodide 
 of silver, which latter salts can be formed by precipitation exactly 
 as is the chloride. The chloride of silver is sometimes found as a 
 mineral, the name of which is cerargyrite. Silver chloride is very 
 readily dissolved by an aqueous solution of ammonia ; silver bromide 
 is less soluble in that medium, while silver iodide is entirely insolu- 
 ble. When heated with hydrochloric acid, both the bromide and 
 iodide of silver are converted into the chloride, but, on the other 
 hand, the chloride of silver changes into the bromide when treated 
 with cold hydrobromic acid, and both the chloride and bromide are 
 converted into the iodide when covered with cold hydroiodic acid 
 and allowed to stand. $ 
 
 One of the chief uses for the silver halides is in the art of pho- 
 tography ; their application depends on the fact that light effects 
 
 * The existence of this oxide is doubtful ; it is possibly a mixture of metal- 
 lic silver and of Ag 2 O. 
 
 t See page 380. 
 
 t These changes afford an excellent example of the reversability of many 
 chemical reactions when the conditions are changed. 
 
406 PHOTO-CHEMICAL ACTION. 
 
 such a change in these compounds as to cause the subsequent for- 
 mation of a film of finely divided metallic silver upon those portions 
 of a glass plate which have been covered with a thin layer of silver 
 halide and which have been acted upon by the light ; this film being 
 produced when the silver halides are treated with certain reducing 
 solutions ; * the unchanged silver halide is subsequently removed 
 by solutions which have a solvent action on the unchanged salts but 
 which leave metallic silver untouched. A solution of this nature 
 is, for instance, one of sodium thiosulphate (hyposulphite), Na 2 S 2 O a 
 (see page 154). | 
 
 A plate covered with a thin layer of gelatine, containing some sil- 
 ver halide (preferably the iodide) very evenly distributed through- 
 out the mass, is exposed to the light in such a manner that a 
 perfectly clear image of some object to be photographed is thrown 
 upon it by means of a camera. The changes in the silver salt then 
 take place according to the intensity of the light which is cast upon 
 the object. When the silver is finally deposited on the plate by 
 means of the processes which have been outlined above, a perfectly 
 clear image of the object to be photographed is left. The plate so 
 produced is a " negative ; " the photograph is prepared from this by 
 covering paper with a sensitive film similar to that which was used 
 on the negative, and then exposing a piece of this, placed under the 
 negative upon which the image has been fixed, to the action of the 
 light ; the picture is then produced in such a manner that, wherever 
 a dark spot appears on the negative, a light spot will be seen on the 
 photograph. The latter is subsequently developed and fixed by a 
 process similar to that used in the preparation of the negative. 
 
 The chemical action of light on the silver halides is not as yet 
 definitely understood. The chemical changes which are useful in 
 the art of photography are not, however, by any means the only 
 ones which are caused by light and which we encounter. Hydro- 
 gen and chlorine, it will be remembered, are entirely without action 
 on each other when in the dark, but unite with the greatest readi- 
 ness in the sunlight ; the compounds of carbon and hydrogen (see 
 page 277) are substituted by halogens when placed in the light, 
 
 * So-called developers ferrous sulphate, pyrogallic acid, liydrochinon, etc. 
 t The cause of the solution is the formation of the soluble double thiosul- 
 phate of silver and sodium, Ag 2 S 2 O 3 , Na 2 S 2 O 3 . 
 
SILVER ; NITRATE, SULPHIDE. 407 
 
 but are not attacked in the dark, while many vegetable dyes bleach 
 wii^k the greatest rapidity when under the influence of the sunlight. 
 The action of light in bringing about chemical changes is identical 
 with that of heat ; and, probably, all light rays are capable of causing 
 such changes. It was formerly supposed that only certain light rays 
 were capable of causing chemical reactions and these were desig- 
 nated as "actinic" rays; in view of later developments the distinc- 
 tion between " actinic " and " not-actinic " rays has, of course, 
 disappeared. 4 
 
 Silver nitrate, Ag N0 3 , is formed by dissolving silver, or silver 
 oxide, in nitric acid ; it is soluble in water, and is separated from 
 its solution by evaporation.* The salt crystallizes in plates belong- 
 ing to the rhombic system, and can be fused at 224 without chan- 
 ging its composition ; the fused salt is cast into sticks, and is popu- 
 larly termed " lunar caustic." Silver nitrate is not changed when 
 exposed to the light, unless some organic substances (dust, etc.) are 
 brought in contact with it ; it then turns black, owing to reduction 
 and deposition of metallic silver.f Dry silver nitrate absorbs 
 ammonia quite readily ; the same is true of the dry chloride. The 
 latter substance forms a compound, 2 Ag Cl + 3 NH 3 , $ which loses 
 ammonia when heated. This power of absorbing ammonia is one 
 of the most marked resemblances between silver and copper salts. 
 
 Sulphide of silver, Ag 2 S, is formed by fusing sulphur and silver 
 together, or by precipitation from the solution of a soluble silver 
 salt by means of hydrogen sulphide ; the natural mineral argentite, 
 Ag 2 S, has a metallic appearance much like that of lead ; the pre- 
 cipitated sulphide is black. The compound is readily decomposed 
 by heating with lead or iron ; by this process the sulphides of the 
 latter metals and free silver are produced. The blackening of silver 
 when exposed to the air is brought about by the action of hydrogen 
 sulphide ; the minute traces of that gas which are present in the 
 atmosphere attack the metal and form silver sulphide. 
 
 Many other salts of silver have been investigated. A number of 
 
 * Silver nitrate in a pure form can be precipitated from its concentrated 
 aqueous solutions by the addition of concentrated nitric acid, in which sub- 
 stance it is insoluble, or nearly insoluble. 
 
 t Silver nitrate, when in contact with the skin, produces a black stain. 
 
 t Other compounds, AgCl + 3NH 3 and Ag Cl + 2 NH 3 , have also been 
 described. 
 
408 GOLD ; SALTS OF. 
 
 these display a tendency to form double salts with other silver com- 
 pounds, or with the salts of the alkali metals. For their descrip- 
 tion a larger work must be consulted. 
 
 Compounds of gold. Gold forms three oxides, aurous oxide, 
 Au 2 0, aurous-auric oxide, Au 4 4 , and auric oxide, Au 2 O s . The 
 first of these is produced with great difficulty, and is of no great 
 importance excepting as an illustration of the resemblance between 
 the compounds of gold, silver, and copper. The last one, Au 2 3 is 
 both a base and an acid. The hydroxide, Au(OH) 3 , loses water 
 when allowed to stand, changing to a metahydroxide of the formula 
 AuO(OH); the latter changes to auric oxide at 150; at 220 
 auric oxide decomposes completely into gold and oxygen. The 
 greatest difficulty which is encountered in the study of gold com- 
 pounds lies in the readiness with which they break down and sepa- 
 rate metallic gold. Auric oxide is readily dissolved by potassium 
 or sodium hydroxide ; the salts produced are derived from meta- 
 auric hydroxide, Au 2 H, so that they have the formula, M Au 2 ; 
 in this respect auric oxide resembles the oxide of aluminium (see 
 page 338). 
 
 Chlorides of gold. Three chlorides of gold, corresponding to 
 the oxides, are known. They are AuCl, aurous chloride; Au 2 Cl 4 , 
 aurous-auric chloride ; and Au C1 8 , auric chloride. The first of 
 these, Au Cl, is insoluble in water. When the substance is covered 
 with water and allowed to stand, it breaks down into auric chloride 
 (which is soluble) and metallic gold : 
 
 3 Au Cl = Au C1 3 + 2 Au ; 
 
 but, on the other hand, a solution of auric chloride, when evapo- 
 rated, breaks down into aurous chloride and free chlorine. 
 
 Auric chloride. Gold is dissolved by aqua regia ; the substance 
 contained in solution is auric chloride, but the latter cannot be 
 isolated by evaporation, because, as was just mentioned, it decom- 
 poses into aurous chloride and chlorine. Gold, when finely divided 
 and treated with dry chlorine at a temperature of 180, forms aurous- 
 auric chloride, Au 2 C1 4 ; this salt, when heated to 220, breaks 
 down into gold and auric chloride ; the auric chloride sublimes and 
 collects on the cooler surfaces. Chlorine does not attack gold at 
 300. A solution of auric chloride can also be prepared by allow- 
 ing aurous chloride (covered with water) to stand (see above) ; 
 
CHLOKAURIC ACIDS. 
 
 409 
 
 when this solution is carefully evaporated, crystals having the 
 composition Au C1 3 -f- 2 H 2 are formed. 
 
 Chlorauric acids. When a solution of gold in aqua regia, to 
 which concentrated hydrochloric acid has been added, is evaporated, 
 crystals of an acid, H Au C1 4 -j- 4 H 2 O, separate. This compound 
 is another one of the class of substances of which fluosilicic and 
 fluoboric acids (pages 303, 330) and the double salts of aluminium 
 are examples (see page 337), two chlorine atoms taking the place 
 of one oxygen atom ; the parallelism becomes clear when we com- 
 pare the formula of potassium aurate with those of chlorauric acid 
 and of its potassium salt : 
 
 . fCl, ...((31, 
 
 AU JOK Au jci*H Au ( ^ 
 
 Potassium aurate. Chlorauric acid. Potassium chloraurate. 
 
 Potassium chloraurate, KAuCl 4 , can be produced by mixing 
 solutions of auric chloride and potassium chloride and evaporating 
 to crystallization. Potassium bromoaurate, KAuBr 4 , and potas- 
 sium iodoaurate, K Au I 4 , are also known. 
 
 Two sulphides of gold, aurous sulphide, Au 2 S, and auric sul- 
 phide, Au 2 S 3 , are known. The former is produced by treating a so- 
 lution of potassium aurous cyanide (Au C N, KCN = K Au (CN ) 2 ) 
 with hydrogen sulphide ; the latter by precipitation from a cold 
 neutral solution of auric chloride by means of hydrogen sulphide ; 
 if the solution is hot, nothing but metallic gold separates.* 
 
 The chief compounds discussed in the last chapter are given in 
 the following tables : 
 
 COMPOUNDS TYPICAL OF THE FAMILY, CORRESPONDING TO COM- 
 POUNDS OF THE ALKALI METALS. 
 
 
 OXIDES. 
 
 CHLORIDES. 
 
 6DLPHIDE8. 
 
 
 Copper, 
 
 Cu 2 O 
 
 CuCl t 
 
 Cu 2 Sf - 
 
 Cuprous compounds. 
 
 Silver, 
 
 A&O: 
 
 AgClf 
 
 Ag 2 Sf 
 
 Argentic compounds. 
 
 Gold, 
 
 Au 2 OJ 
 
 AuClf 
 
 Au 2 Sf 
 
 Aurous compounds. 
 
 The compounds of copper and silver unite with ammonia. 
 The salts of silver are derived from the oxide Ag 3 O ; for instance, Ag NO 3 , 
 Ag 3 S0 4 . 
 
 t Insoluble in water and in dilute acids. 
 
 t Decomposed into oxygen and the metal when heated. 
 
 A gold sulphide, Au S, is also described. 
 
410 COPPER ; SILVER ; GOLD ; COMPOUNDS OF. 
 
 COMPOUNDS NOT TYPICAL. 
 
 
 OXIDES. 
 
 HYUBOXIDES. 
 
 CHI.OBIDE8. 
 
 SULPHIDES. 
 
 Cupric compounds, 
 Auric compounds, 
 
 CuO 
 Au 2 3 t 
 
 Cu(OH) 2 
 
 Au(OH) 3 t 
 
 CuCl. 2 * 
 AuCl 3 *t 
 
 CuS 
 
 Au 2 S 3 t 
 
 The salts of copper are derived from cupric oxide; for instance Cu (NO 3 ) 2 , 
 Cu SO 4 . 
 
 Auric oxide is acidic in its character. It is also basic, for the few gold 
 salts which are known (except AuCl) are derived from it. The aurates are 
 derived from the hydroxide Au O 2 H ; the salts derived from the alkalies are 
 Au O 2 M. Auric chloride has the character of an acidic anhydride ; chlorau- 
 ric acid, Au CU H, chloraurates, Au CU M. 
 
 * Soluble in water. 
 
 t Decomposed by heat, leaving gold behind. 
 
 'if i 
 
THE ALKALINE EARTHS. 411 
 
 CHAPTER LIV. 
 
 THE FAMILY OF THE ALKALINE EARTHS. 
 
 Beryllium (Glucinum) ; symbol, Be ; atomic weight, 9 ; 
 Magnesium ; symbol, Mg ; atomic weight, 24.3 ; 
 Calcium ; symbol, Ca ; atomic weight, 40 ; 
 Strontium ; symbol, Sr ; atomic weight, 87.6 ; 
 Barium ; symbol, Ba ; atomic weight, 137.43. 
 
 THE elements comprising the primary group of the elements of 
 the family of alkaline earths are : 
 
 Beryllium, atomic weight, 9 specific gravity, 1.99, atomic volume, 4.52 
 
 Magnesium, " " 24.3 " " 1.74 " " 13.8 
 
 Calcium, " " 40 " " 1.57 " " 25.4 
 
 Strontium, " " 87.6 " " 2.50 " " 34.9 
 
 Barium, " " 137.43 " " 3.75 " " 36.5 
 
 Of these, the first two (beryllium and magnesium) belong to the 
 typical short periods, and therefore resemble the following three 
 (calcium, barium, and strontium), but they also are closely allied 
 with the three elements (zinc, cadmium, and mercury) which com- 
 prise the secondary group belonging to this family. They resemble 
 calcium, barium, and strontium, because their oxides, hydroxides, 
 and the salts derived from them are formed according to the same 
 formulae; they differ from those three elements and fall into line 
 with zinc, cadmium, and mercury, by reason of the solubility of 
 their sulphates * and because of their tendency to form double 
 salts when their salts are brought in contact with those of ammo- 
 nium. 
 
 Beryllium f and magnesium are prepared as metals by heating 
 the chlorides with metallic sodium : 
 
 * The sulphates of calcium, strontium, and barium, are insoluble or 
 nearly insoluble in water. 
 
 t Beryllium chiefly occurs as a metasilicate of beryllium and aluminium, 
 known as beryl, Be 3 A1 2 (Si O 3 ) 6 . When these crystals are transparent and 
 colored green by chromic oxide they are termed emerald. 
 
412 ALKALINE EARTHS; PROPERTIES OF. 
 
 MCl 2 +2Na==2NaCl + M. 
 
 This method is exactly parallel to the one formerly employed in 
 the preparation of aluminium (see page 333). The name glucinum 
 was first given to beryllium, owing to the sweetish taste of the 
 salts of this metal. 
 
 Beryllium is white with a silver-like lustre, malleable and ductile; its 
 melting point is lower than that of silver; it is slightly oxidized when heated 
 to a high temperature ; when finely powdered and heated it Tmrns with a bril- 
 liant light; it readily burns in an atmosphere of chlorine; it dissolves in aque- 
 ous hydrochloric acid. 
 
 Magnesium is silver white, malleable and ductile, melts at 700, is volatile 
 .at a high heat; the metal does not oxidize in dry air, but it readily corrodes 
 when in contact with water; when heated above its melting point in the air it 
 burns with a most brilliant white light;* the metal also readily burns in 
 chlorine; it is easily soluble in dilute acids. 
 
 Calcium, barium, and strontium are isolated by electrolyzing the 
 fused chlorides in a crucible from which air is excluded ; the posi- 
 tive electrode is made of gas carbon, which is not attacked by 
 chlorine, the negative electrode consists of iron wire. This method 
 is also employed in preparing the alkali metals. Calcium can also 
 be made by reducing the iodide by means of metallic sodium. 
 
 Calcium, strontium, and barium are yellow with metallic lustre. The 
 freshly cut surfaces of the metals soon become covered with a layer of oxide ; 
 the metals must therefore be preserved under petroleum (see page 384). 
 When heated in the air they burn with a brilliant light. They all energeti- 
 cally decompose water at ordinary temperatures, liberating hydrogen. 
 
 The changes which are brought about by the increasing atomic 
 weights and volumes, as we pass downward in the list given above, 
 are shown in the following table : 
 
 Beryllium does not decompose water. 
 
 Magnesium decomposes boiling water. 
 
 Calcium decomposes water at ordinary temperatures. 
 
 Strontium decomposes water very readily at ordinary temperatures. 
 
 Barium decomposes water as readily as the alkali metals. 
 
 The oxides and hydroxides which are typical of the family have 
 the formulae MO and M(OH) 2 ; the metals are therefore divalent 
 and replace two atoms of hydrogen in acids. 
 
 * This magnesium light has a marked chemical effect on silver halides, 
 and is used as a flash light in photography. 
 
CALCIUM; OXIDE, HYDROXIDE. 
 
 413 
 
 
 OXIDES. 
 
 
 HYDROXIDES. 
 
 Beryllium 
 
 BeO 
 
 Not changed to hydroxide by the ad- 
 dition of water 
 
 Be (OH a ) 
 
 Magnesium 
 
 MgO 
 
 Slowly changed to hydroxide by the 
 addition of water 
 
 Mg(OH) 2 
 
 Calcium 
 
 CaO 
 
 Readily changed to hydroxide by the 
 addition of water 
 
 Ca(OH) 2 
 
 Strontium 
 
 SrO 
 
 Readily changed to hydroxide by the 
 addition of water 
 
 Sr(OH) 2 
 
 Barium 
 
 BaO 
 
 Readily changed to hydroxide by the 
 addition of water 
 
 Ba(OH) 2 
 
 The hydroxides of beryllium and of magnesium are insoluble in 
 water. The solubility of the other three increases with increasing 
 atomic weight (see page 385) while the stability increases with the 
 solubility. 
 
 Be (OH) 2 decomposes at about 300; Be (OH) 2 = Be O + H 2 O. 
 
 Mg (OH) a " " a low red heat; Mg (OH) 2 = Mg O + H. 2 O. 
 
 Ca(OH) 2 " " a highredheat; Ca(OH) 2 = Ca O +H 2 O. 
 
 Sr(OH) 2 " "a white heat; Sr(OH) 2 = SrO+H 2 O. 
 
 Ba (OH) 2 , which crystallizes from water in crystals of the formula 
 Ba (OH) 2 + 8 H 2 O, and which is quite soluble in water, can be fused without 
 change. 
 
 The oxides and hydroxides are all strong bases; the solutions of calcium, 
 strontium, and barium hydroxides have an alkaline reaction. The hydroxide 
 of beryllium, being that of the least metallic of all the elements, can also dis- 
 solve in caustic alkalies, so that, under certain circumstances, it acts as an acid. 
 The oxide and hydroxMe of calcium, known respectively as quick and slaked 
 lime, are the most important of these compounds. 
 
 Calcium oxide is prepared by heating the carbonate in " lime- 
 kilns " until it decomposes into calcium oxide and carbon di- 
 
 Ca C0 3 = Ca + C0 2 . 
 
 The quick-lime so prepared is more or less impure, according to the 
 condition of the limestone or marble used ; when brought in contact 
 with water it unites with that liquid to form calcium hydroxide 
 (or slaked lime) ; at the same time a large amount of heat is 
 
 developed : 
 
 CaO+H 2 = Ca(OH) 2 . 
 
 If quick-lime is exposed to the air, it absorbs carbon dioxide and 
 water. It then crumbles and is said to be " air slaked." 
 
414 BARIUM SUPEROXIDE. 
 
 Slaked lime finds its chief use in the preparation of mortar. 
 Mortar is prepared by stirring together slaked lime and sand until 
 the mass assumes the consistency of thick porridge. When placed 
 between bricks the mixture gradually hardens, the calcium hydrox- 
 ide absorbing carbon dioxide and changing into calcium carbonate. 
 The sand which is added in all probability serves the purpose of 
 rendering the mortar porous, and it thereby facilitates the absorp- 
 tion of carbon dioxide ; it certainly does not, as was formerly sup- 
 posed, render the product hard by forming calcium silicate. A 
 mixture of quick-lime, aluminium oxide, and silicon dioxide forms 
 Portland cement. The latter, when brought in contact with water, 
 gradually hardens, owing to the union of the calcium and alumin- 
 ium oxides with the silicon dioxide which is present. The chemi- 
 cal process depends upon the formation of hydrated calcium silicates 
 as well as of calcium aluminate. 
 
 Barium superoxide. In addition to the ordinary oxide, BaO, 
 barium is also able to form a hyperoxide, Ba 2 (see pages 323 and 
 324). This compound is prepared by passing oxygen over barium 
 oxide which is heated to redness, or by heating a mixture of barium 
 oxide and potassium chlorate. The substance is a white powder 
 which loses oxygen at a bright red heat. It is a powerful oxidizer : 
 hydrogen, boron, carbon, sulphur, etc., are changed to the corre- 
 sponding oxides when heated with it ; in many cases the tempera- 
 ture of the mass even spontaneously increases to redness during the 
 process. Barium hyperoxide, when mixed with cold water, forms a 
 hydrate, Ba 2 + 6 H 2 ; boiling water decomposes it, liberating 
 oxygen and leaving barium hydroxide. Barium hyperoxide forms 
 hydrogen peroxide (see page 50) when treated with dilute acids : 
 
 Ba0 2 + 2 HCl=BaCl 2 + H 2 2 . 
 
 In this respect barium hyperoxide differs radically from other hy- 
 peroxides, for the latter liberate chlorine when in contact with 
 hydrochloric acid (see pages 60 and 323). The cause of this differ- 
 ence lies in the fact that the hyperoxides of manganese and lead 
 form intermediary compounds before setting free the halogen. 
 
 The chlorides of the elements of this family (MC1 2 ) are all 
 soluble in water; those of beryllium and magnesium decompose 
 when heated in a current of air, giving off chlorine and leaving the 
 oxide ; the chlorides of calcium, barium, and strontium are more 
 stable. The chloride of calcium, CaCl 2 -|-6 H 2 0, melts in its water 
 
ALKALINE EARTHS; HALIDES, CARBONATES. 415 
 
 of crystallization at 29 ; at 100 it becomes anhydrous and then 
 again melts; this fused form of calcium chloride is deliquescent, 
 and, because it greedily absorbs moisture, is frequently used as a 
 drying agent. The chloride of strontium is not deliquescent, while 
 the chloride of barium slowly takes up water from the air. The 
 bromides and iodides are like the chlorides in every respect. Both 
 the chlorides of calcium and of magnesium occur as minerals in 
 some salt deposits, the former as chlorocalcite, the latter as the 
 extremely deliquescent mineral bischofite. Carnallite is a double 
 chloride of potassium and magnesium, K Cl, Mg C1 2 + 6 H 2 0, which 
 is found quite frequently in the Strassfurth salt regions. The in- 
 creasing metallic character of the elements of this family, as we 
 pass from the member with the smallest atomic weight to that with 
 the largest, is very well illustrated by the behavior of the chlorides 
 in the presence of water : 
 
 Be C1 2 + 4 H 2 O, completely decomposed into Be O + 2 H Cl when its 
 solution is evaporated. 
 
 Mg C1 2 + 6 H 2 O, completely decomposed into Mg O + 2 H Cl when its 
 solution is evaporated. 
 
 Ca C1 2 + 6 H 2 O, partly decomposed into basic calcium chloride when heated 
 with water. 
 
 Sr C1 2 + 6 H 2 O, not decomposed. 
 
 Ba C1 2 + 2 H 2 O, not decomposed. 
 
 All of the carbonates of the members of this family are insolu- 
 ble in water. They are the more stable the more positive the metal 
 forming them is, so that their stability increases as the atomic 
 weights of the elements, counting from above downward, become 
 greater. Beryllium carbonate is only capable of existence as a 
 normal salt when it is in an atmosphere of carbon dioxide ; when 
 exposed to the air it breaks down, giving off carbon dioxide and 
 leaving a basic carbonate; magnesium carbonate begins to break 
 down at 100, calcium carbonate at a low red heat, while neither 
 strontium nor barium carbonate decomposes until a white heat is 
 reached. All of the carbonates can be prepared by precipitation 
 from solutions of the salts of the respective metals by addition of a 
 soluble carbonate, such as that of sodium. The following will serve 
 as examples : 
 
 Na 2 C0 3 +BaCl 2 =2 Nad -f BaCO 3 
 
 Soluble. Insoluble. 
 
416 ALKALINE EARTHS; OCCURRENCE. 
 
 K 2 C0 3 + CaCl 2 = 2KC1 + CaC0 3 
 
 ) 2 C0 3 + Mg C1 2 = 2 NH 4 Cl + Mg C0 3 
 
 Soluble. Insoluble. 
 
 The carbonates of magnesium, calcium, strontium, and barium, 
 and the double carbonate of magnesium and calcium, form an ex- 
 tremely important dimorphous and isomorphous group of minerals ; 
 in some localities entire mountain ranges are made up of these com- 
 pounds, while the various amorphous and cryptocrystalline varieties 
 of calcium carbonate (known as chalk, limestone, and marble) con- 
 stitute deposits of astonishing magnitude (see page 269). The fol- 
 lowing table gives the relationship between the crystalline carbonates 
 of the elements belonging to this family : 
 
 HEXAGONAL SYSTEM, KHOMBOHEDRA. KHOMBIC SYSTEM. 
 
 Calcite group. Arragonite group. 
 
 Calcite (calcspar) Ca CO 3 Arragonite. 
 
 Magnesite Mg CO 3 
 
 Dolomite Ca CO 3 , Mg CO 3 
 
 BaCO 3 Witherite. 
 
 Sr CO 3 Strontianite. 
 
 While magnesium carbonate is not, by itself, capable of crystallizing 
 in the same form as arragonite, yet, when mixed with the carbonate 
 of manganese or of calcium, it assumes that form ; on the other 
 hand, calcite crystals which contain barium and strontium are also 
 found. This isodimorphous group of minerals is very far reaching, 
 for" the carbonates of zinc, iron, manganese and cobalt also belong to 
 the calcite group, while those of manganese, iron and lead are also 
 found crystallizing in the form of arragonite. 
 
 The sulphates of the elements of this family may be divided into 
 two classed, those of beryllium and magnesium, ^which are soluble 
 in water, and those of calcium, strontium, and barium, which are in- 
 soluble, or nearly insoluble. Of the latter class, that of calcium is 
 soluble with difficulty, that of strontium is less soluble than that of 
 calcium, while the sulphate of barium is insoluble ; these sulphates 
 can therefore be produced as white precipitates on the addition of 
 a soluble sulphate to the solutions of salts of the respective metals.* 
 
 * Calcium sulphate will not be precipitated if the solution of the calcium 
 salt is too dilute. A solution of calcium sulphate will precipitate strontium 
 and barium salts; a solution of strontium sulphate will precipitate barium 
 salts. 
 
THE VITKIOLS. 
 
 417 
 
 Magnesium sulphate is the representative of a large number of 
 sulphates which are known as vitriols. The vitriols, with the 
 exception of copper sulphate (which contains five molecules of 
 water) all crystallize with seven molecules of water of crystalliza- 
 tion and form a typical isomorphous group of compounds ; the one 
 exception, copper sulphate, Cu S0 4 + 5 H 2 0, can, however, crystal- 
 lize with seven molecules of water when it is present in an isomor- 
 phous mixture with some other vitriol. The vitriols, when heated to 
 100, change to salts having the composition MS0 4 + H 2 ; the 
 last remaining molecule of water passes off at a higher temperature, 
 and for this reason these substances are commonly regarded as 
 being secondary salts of the hydrated sulphuric acid, H 4 S0 5 , so 
 that their formulae would be MH 2 S0 5 + 6 H 2 0. The following is 
 a list of the compounds comprising this isomorphous group : 
 
 Be H 2 SO 5 + 6 H 2 O Beryllium sulphate. 
 
 Mg H 2 SO 5 + 6 H 2 O Magnesium sulphate. 
 
 ZnH 2 SO 5 +6H 2 O. Zinc sulphate (white vitriol). 
 
 FeH 2 SO 5 +6H 2 O Ferrous sulphate (green vitriol). 
 
 Ni H 2 So 5 +6H 2 O Nickel sulphate. 
 
 Co H 2 SO 5 + 6H 2 O. Cobalt sulphate. 
 
 CuH 2 SO 6 + 4H 2 O Copper sulphate (blue vitriol). 
 
 Copper sulphate crystallizes with six molecules of water when in an iso- 
 morphous mixture with one or more of the other vitriols. All of the vitriols 
 can crystallize with one formula weight of potassium or of ammonium sulphate 
 to form double salts of the general formula M' 2 SO 4 , M"SO 4 + 6 H 2 O. 
 
 The sulphates of calcium, strontium, and barium occur quite fre- 
 quently as minerals ; they constitute an isomorphous group, which 
 crystallizes in the rhombic system ; the representatives of this group 
 are given in the following table : 
 
 NAME OF MINERAL. 
 
 FORMULA. 
 
 1. Anhydrite 
 
 1. 
 
 CaSO 4 
 
 2. Celestine 
 
 2. 
 
 SrS0 4 
 
 3. Barite 
 
 3. 
 
 BaSO 4 
 
 4. Barytocelestine 
 
 4. 
 
 (Ba, Sr)SO 4 
 
 5. Anglesite 
 
 5. 
 
 PbSO 4 
 
 Calcium sulphate also occurs as gypsum, in which mineral it 
 crystallizes with two molecules of water of crystallization, Ca S0 4 
 -j- 2 H 2 ; gypsum is frequently found as a massive variety, in 
 
418 CALCIUM SULPHATE, PHOSPHATES. 
 
 which condition it bears the name of alabaster. When gypsum is 
 heated to a little above 100 it loses its water of crystallization, 
 and is converted into a fine white powder which is termed plaster 
 of Paris ; this substance, when mixed with water, once more unites 
 with that liquid, and then changes into a firm mass which has the 
 composition of alabaster; this process is termed the " setting" of 
 plaster of Paris. Many natural waters contain calcium sulphate in 
 solution ; such waters are known as permanent hard waters, because 
 the calcium salt is not removed from them by boiling. Temporary 
 hard waters contain primary calcium carbonate, Ca (HC0 3 ) 2 ; the 
 latter substance, however, breaks down into carbon dioxide and 
 secondary calcium carbonate (insoluble) when the solution is boiled ; 
 the calcium which is contained in the primary carbonate is, there- 
 fore, entirely removed by this process. Calcium is able to form an 
 insoluble compound with soap ; hard waters, therefore, form a pre- 
 cipitate when brought in contact with the latter substance, and will 
 consequently form a lather only after all the calcium salts have 
 been removed (see page 43). 
 
 The tertiary and secondary phosphates of the elements belonging 
 to this family are all insoluble in water ; the primary phosphates 
 are soluble ; as a consequence, the tertiary phosphates are dissolved 
 on addition of the stronger acids, such as hydrochloric or nitric. 
 The tertiary phosphate of calcium is the only one of these which 
 occurs as a mineral, the name of which is osteolite, Ca 3 (P0 4 ) 2 ; 
 this substance is often found in massive deposits, which are espe- 
 cially extensive in Florida ; guano is tertiary calcium phosphate 
 which is mixed with a number of impurities, such as calcium car- 
 bonate, magnesium carbonate, gypsum, etc. A double salt of cal- 
 cium phosphate and calcium chloride is found in crystals belonging 
 to the hexagonal system, and is termed apatite, Ca 3 (P0 4 ) 2 -f Ca C1 2 . 
 In addition to its occurrence as mineral deposits, calcium phosphate 
 is the chief constituent of the inorganic portions of the bones (see 
 page 211).* The reactions relating to the conversion of the terti- 
 ary phosphate into the soluble primary one are given on pages 229 
 and 230. The soluble primary phosphate of calcium, mixed with 
 gypsum and other impurities, goes by the name of superphosphate ; 
 this substance is used as a foundation for the mixtures which find 
 their way into the market as artificial fertilizers. Of course, the 
 
 * Bone-ash contains 85 per cent of calcium phosphate. 
 
MAGNESIUM PHOSPHATE. 419 
 
 phosphate in a fertilizer must, in part at least, be in a soluble 
 form so that it can be readily absorbed by plants ; the conversion 
 of the insoluble tertiary into the soluble primary phosphate is 
 effected by means of sulphuric acid : 
 
 The value of a fertilizer is determined by the amount of soluble 
 phosphate which it contains. 
 
 Magnesium, phosphate behaves exactly as does calcium phos- 
 phate, the tertiary, secondary, and primary phosphate being known ; 
 the latter is soluble in water. The most important phosphate of 
 magnesium is the insoluble ammonium-magnesium phosphate, 
 Mg(NH 4 )P0 4 ; the formation of this salt as a precipitate may be 
 used as a test for the presence either of magnesium or of phos- 
 phoric acid in a solution. The salt is produced by adding a soluble 
 phosphate to a mixture of a soluble magnesium salt with ammonia 
 and ammonium chloride. If ammonia alone were added to a solu- 
 tion containing a magnesium salt, a portion of the latter would be 
 decomposed and the base precipitated as magnesium hydroxide, 
 while a part would remain in solution as a double salt of magnesium 
 and ammonium, for magnesium salts have the power of forming 
 compounds with ammonium salts, and these compounds are not 
 decomposed by ammonia : 
 
 2 Mg S0 4 + 2 NH 3 + 2 H 2 = Mg (OH ) 2 + ( KH 4 ) 2 S0 4 , Mg S0 4 . 
 
 The previous addition of an ammonium salt to a solution containing 
 a salt of magnesium therefore prevents any precipitation of mag- 
 nesium hydroxide by means of ammonia. When ammonium-mag- 
 nesium phosphate is heated, it loses ammonia and changes into the 
 secondary phosphate of magnesium : 
 
 MgKH 4 P0 4 = Mg HP0 4 + NH, , 
 
 and the latter, on further heating, again loses water, and finally 
 leaves magnesium pyrophosphate (see page 231) : 
 
 2 Mg HP0 4 = Mg 2 P 2 7 + H 2 0. 
 
 Arsenic acid reacts similarly to phosphoric acid. It forms an 
 ammonium-magnesium arsenate, MgKB 4 As0 4 . The precipitate 
 formed is not to be distinguished from ammonium-magnesium phos- 
 phate ; if both arsenic and phosphoric acid are present in any solu- 
 
420 CALCIUM SILICATES. 
 
 tion, then the arsenic acid must, previous to the precipitation of 
 the phosphoric acid, be reduced to arsenious acid by means of sul- 
 phur dioxide, for in this form it is not precipitated by the mixture 
 of magnesium salts. 
 
 The hypochlorite of calcium, when mixed with calcium hydroxide 
 and calcium chloride, is known as chloride of lime. Because of 
 the invariable occurrence of calcium chloride in conjunction with 
 calcium hypochlorite, the theory is not unfrequently maintained 
 that calcium hypochlorite is in reality a mixed chloride and hypo- 
 chlorite of calcium having the formula : 
 
 Ca 
 
 (0( 
 JC1. 
 
 C1 
 
 Such a body would have the same composition by weight as an 
 equimolecular mixture of calcium chloride and calcium hypochlo- 
 rite, Ca C1 2 + Ca (0 Cl) 2 = 2 Ca (0 Cl) Cl. The reactions peculiar 
 to calcium hypochlorite have been explained on pages 122 and 123. 
 
 The chlorates and nitrates of calcium, barium, and strontium are 
 soluble in water and are extensively used in the manufacture of 
 Greek fire ; the chlorate of strontium, when mixed with oxidizable 
 substances and ignited, gives an intensely red light, while that of 
 barium produces a green one. 
 
 The silicates of calcium are of the greatest importance because 
 of the fact that they are essential in the manufacture of glass. 
 Calcium metasilicate, CaSi0 3 , occurs as the mineral wollastonite, 
 and a great many other naturally occurring silicates contain cal- 
 cium (see page 307) ; these silicates are, however, crystalline in 
 their structure, while the artificial silicates are, as a rule, amor- 
 phous. Glass consists of a vitreous mixture of the silicates of the 
 alkalies and of calcium, with silicon dioxide; in some forms of 
 glass, however, lead oxide may replace calcium oxide. Ordinary 
 window glass is produced by fusing sand, calcium carbonate, and 
 sodium carbonate together; the silicate of calcium and sodium so 
 formed is, in reality, of a crystalline structure, but this structure 
 is concealed by the vitreous mass of silicon dioxide ; such a glass 
 is readily attacked by laboratory reagents or even by the continued 
 action of water, and, after it has been acted on for some time, the 
 crystalline condition becomes apparent.* Window glass is first 
 
 * All kinds of glass are more or less attacked by water, or alkalies; dilute 
 acids do not appear to have any effect ; concentrated acids prevent the solution 
 
GLASS. 421 
 
 blown and then cut into suitable pieces ; for that reason it is more 
 or less irregular in thickness and does not present a perfectly 
 smooth surface. Plate glass has essentially the same composition, 
 but is cast on flat plates and finally polished. Bohemian glass is 
 made by fusion of potassium carbonate, a little sodium carbonate, 
 silicon dioxide, and calcium carbonate. The replacing of the 
 sodium carbonate by potassium carbonate, with the resulting for- 
 mation of a potassium-calcium silicate, renders the glass difficult 
 to fuse. Bohemian glass is used for the manufacture of chemical 
 apparatus, a further advantage belonging to this variety being in 
 the fact that it is not readily attacked by chemical reagents. 
 Bottle glass contains more calcium silicate than either of the 
 above varieties ; it is frequently colored green by the presence 
 of ferrous silicate. The commoner bottles are made from impure 
 materials. Flint glass is of similar composition to ordinary lime- 
 soda glass, with the exception that the lime is replaced by lead 
 oxide ; it is characterized by having a very high index of refrac- 
 tion,* great lustre and high specific gravity ; it is the most fusible 
 variety of glass. | Flint glass is used in the manufacture of opti- 
 cal instruments and in some chemical apparatus. Strass is flint 
 glass which is very rich in lead; it is used for making artificial 
 gems. Glass is stained by adding inorganic coloring matter to the 
 
 of the glass to a certain extent. The solubility of glass in solutions of sodium 
 hydroxide or sodium carbonate (containing -fa of a grain-molecular weight 
 in one litre of water) is about three to five times the solubility of the same 
 glass in pure water. This relationship varies with different kinds of glass. 
 The solubility is greater in sodium carbonate than it is in sodium hydroxide 
 solution. A glass which is made up as follows has the best composition for 
 general laboratory purposes : 
 
 K 2 6.2 per cent. Mn O 0.2 per cent. 
 
 Na 2 O 6.4 per cent. A1. 2 O 3 .."... 0.4 per cent. 
 
 Ca O 10.0 per cent. Si O 2 76.8 per cent. 
 
 Such a glass loses .0037 gram per 100 sq. c. m. to sodium hydroxide solution 
 of the above strength and .0059 gram per sq. c. m. to sodium carbonate. (See 
 Foerster; Berichte d. Deutsch. Chem. Gesell.; 26, 2915.) 
 
 * The index of refraction for flint glass is 1.8, while that of window glass 
 is 1.53. 
 
 t In working with lead glass care must be taken not to bring the same 
 into the reducing flame (see page 283), which is that portion immediately out- 
 side of the central zone, otherwise a part of the lead silicate will be reduced 
 and lead will separate ; the latter renders the glass black and opaque. 
 
422 GLASS. 
 
 colorless varieties and fusing; thus, blue glass is produced by add- 
 ing a little cobalt salt, green glass by copper and chromium, etc. 
 The various glass utensils which are used must be previously an- 
 nealed by a very slow cooling process ; if this precaution is not 
 taken, the outer surfaces, cooling more rapidly than the remainder 
 of the mass, establish such a tension that the slightest scratch upon 
 the surface will cause the entire object to be shattered. This con- 
 dition is best shown in the so-called Prince Rupert's drops. The 
 latter are made by fusing glass and allowing the drops to fall into 
 water; when the end of the small pear-shaped mass so formed is 
 broken, the drop disintegrates into a sandy mass with explosive 
 violence. 
 
ZLN T C; CADMIUM; MEECUKY. 
 
 423 
 
 CHAPTER LV. 
 
 ZINC, CADMIUM, AND MERCURY. 
 
 Zinc ; symbol, Zn ; atomic weight, 65.3 ; 
 Cadmium ; symbol, Cd ; atomic weight, 112 ; 
 Mercury ; symbol, Hg ; atomic weight, 200. 
 
 ZINC, cadmium, and mercury form the secondary group of the 
 family of the alkaline earths. They are the second elements of 
 the second halves of the long periods, while calcium, strontium, and 
 barium occupy the same position in the first halves. As a conse- 
 quence they bear much the same relationship to the alkaline earths 
 as copper, silver, and gold bear to the alkalies. The typical oxides 
 and hydroxides, as well as the salts, derived from zinc, cadmium, and 
 mercury, are, therefore, of the same formula as they are with cal- 
 cium, strontium, and barium. In both divisions of the family the 
 metals are divalent, so that the oxides have the general formula 
 MO, and the hydroxides M (OH) 2 
 
 In their physical characteristics the three elements are all, most 
 certainly, metallic in their nature ; but zinc, the one with the least 
 atomic weight,* is less positive than the other two. The change in 
 physical character brought about by the increasing atomic weight, 
 as we pass from zinc to mercury, is shown in the following table : 
 
 METAL. 
 
 ATOMIC WEIGHT. 
 
 SPECIFIC GRAVITY. 
 
 ATOMIC VOLUME. 
 
 MELTING POINT. 
 
 Zinc 
 
 65.3 
 
 7.15 
 
 9.1 
 
 417 
 
 Cadmium 
 
 112. 
 
 8.65 
 
 12.9 
 
 317 
 
 Mercury 
 
 200. 
 
 13.59 
 
 14.7 
 
 39 
 
 All three of the metals are volatile ; their boiling points decrease 
 with increasing atomic weight, just as their melting points do ; this 
 
 * Zinc is also the element with the next lower atomic weight to that of 
 gallium, which element has many of the characteristics of a not-metal; we 
 should therefore scarcely expect zinc to present very marked metallic properties. 
 
424 ZINC ; PROPERTIES. 
 
 phenomenon, which is most strikingly illustrated in the case of the 
 trio of metals under discussion, is exactly the reverse of the changes 
 in the melting points taking place in the not-metallic families at the 
 right hand extremities of the periods. It will be remembered, also, 
 that the melting points of the alkali metals diminish as we pass 
 from member to member in the direction of increasing atomic 
 weights, and the same is true of the metals constituting the first 
 portion * of the family under discussion ; the elements, therefore, 
 which comprise the first two families in the periodic system show 
 decreasing melting points with increasing atomic weights ; whether 
 the same is true of the boiling points cannot be stated, as many of 
 the elements cannot be volatilized by any of the means at our com- 
 mand ; it certainly is true of the three metals under consideration, 
 for: 
 
 Zinc boils at 927. 
 
 Cadmium boils at 772. 
 
 Mercury boils at 357. 
 
 Zinc is an element with a brilliant metallic lustre which pos- 
 sesses a bluish tint and a crystalline structure, f It is malleable 
 only at temperatures between 100 and 150 ; at ordinary/ 
 temperatures it is easily fractured ; at 200 the metal can readily 
 be pounded into a powder; when heated to its point of vaporization 
 in the air, zinc burns with a bluish white flame, producing zinc 
 oxide, Zn O. The metal easily dissolves in dilute acids (see pages 
 33, 206). Solutions of the caustic alkalies, when warmed, attack 
 zinc, forming zincates and liberating hydrogen ; in this respect zinc 
 resembles aluminium and a number of other metals which can dis- 
 play both metallic and not-metallic properties ; this behavior is not 
 unexpected when we consider that zinc is the element with least 
 metallic properties in the group we are considering, and that every 
 other element following it in the same period can also display certain 
 not-metallic properties. 
 
 Cadmium is a glistening, tin-like metal ; it is . soft, though 
 harder than tin, and it has a crystalline structure ; when heated to 
 
 * Beryllium, magnesium, calcium, strontium, and barium have decreasing 
 melting points as we pass along the line in the order named. The melting 
 points of calcium and strontium are not accurately determined. 
 
 t When pieces of zinc are bent, a peculiar crepitation, similar to that with 
 crystallized tin, is observed (see page 313). 
 
CADMIUM ; MERCURY ; PROPERTIES. 425 
 
 its boiling point in the air, it burns like zinc, forming cadmium 
 oxide, Cd 0. The metal dissolves in acids less readily than does 
 zinc, but with this exception shows the same behavior. The metal 
 does not dissolve in solutions of the caustic alkalies. 
 
 Mercury is a bluish white metal which is fluid under ordinary 
 circumstances. The solid substance (formed at 39) is soft and 
 malleable when pure. Mercury is not changed in the air at ordi- 
 nary temperatures ; if heated for some time near its boiling point 
 and in the presence of oxygen, it is changed to red mercuric oxide 
 ( Hg 0) ; ozone readily attacks it without the necessity of heating. 
 Hydrochloric acid, or dilute sulphuric acid, does not dissolve mer- 
 cury ; concentrated sulphuric acid, when hot, attacks the metal and 
 liberates sulphur dioxide ; nitric acid, concentrated or dilute, acts 
 upon mercury ; if the acid is dilute, mercurous nitrate is formed, if 
 concentrated, mercuric nitrate is produced ; alkalies do not attack 
 mercury. 
 
 As has been stated, zinc, cadmium, or mercury can be easily vol- 
 atilized. The specific gravities of their vapors are as follows : 
 
 Zinc, specific gravity of vapor, air = 1, 2.36, H 2 = 2, 67.96; molecular weight, 65.3. 
 
 Cadmium, specific gravity of vapor, air = 1, 3.95, H 2 = 2, 113.47; molecular weight, 112. 
 Mercury, specific gravity of vapor, air = 1, 6.83, H 2 = 2, 196.7; molecular weight, 200. 
 
 The above determinations show that the molecular weight and 
 the atomic weight of each of the three elements, when they are in 
 the state of a vapor, are identical. Zinc, cadmium, and mercury, 
 therefore, in changing from the liquid to the gaseous state, separate 
 at once into the individual atoms, provided our decision, as to what 
 the relative weights of these atoms are, is the correct one ; the 
 only other elements which exhibit the same phenomenon are so- 
 dium and potassium (and iodine when heated above 1600 ; on de- 
 creasing that temperature the atoms of iodine gradually unite to 
 form diatomic molecules [see page 84] ) . In the course of our study 
 we have therefore encountered elements with one, with two, with 
 three, and with four atoms united to form a molecule of the gas ; the 
 molecules having three and four atoms are, however, dissociated at 
 high temperatures and then change into those having two, while some 
 of the latter have already been dissociated into the individual atoms. 
 Undoubtedly, were we able to command a sufficiently high tempera- 
 ture in the apparatus used for determining the vapor densities, we 
 should be able to discover that all diatomic molecules can be changed 
 
426 
 
 ZINC : CADMIUM : OCCURRENCE. 
 
 into monatomic ones. The following gives a list of elements of 
 which it has been possible to obtain specific gravities while they 
 were in the gaseous state ; in the cases of selenium and antimony 
 some doubt exists as to whether the molecules are in reality tria- 
 tomic, or whether the vapor density numbers which have been 
 obtained are only accidental : 
 
 MONATOMIC.* 
 
 DIATOMIC. 
 
 TKIATOMIC. 
 
 TETKATOMIC. 
 
 Sodium 
 
 Hydrogen 
 
 Ozone 
 
 Phosphorus 
 
 Potassium 
 
 Chlorine 
 
 Selenium (?) 
 
 Arsenic 
 
 Zinc 
 
 Bromine 
 
 Antimony (?) 
 
 
 Cadmium 
 
 Iodine (below 600) 
 
 
 
 Mercury 
 
 Oxygen 
 
 
 
 Iodine (above 1600) 
 
 Sulphur (above 1000) 
 
 
 
 Bismuth (?) 
 
 Selenium 
 
 
 
 
 Tellurium (?) 
 
 
 
 
 Nitrogen 
 
 
 
 
 Phosphorus f , a ? 
 4 white 
 Arsenic j heat 
 
 
 
 The fact that so many elementary gases are formed of complex 
 molecules was not understood when the theory that equal volumes 
 of gases contain equal numbers of particles was first advanced ; the 
 discrepancy frequently observed between the specific gravity of 
 the gases and the atomic weights of elements determined by other 
 means, therefore, led to a disbelief in Avogadro's hypothesis, and to 
 considerable confusion in the determination of atomic weights. 
 
 The principal minerals in which zinc, cadmium, and mercury 
 occur are as follows : 
 
 Zinc and cadmium. The occurrence of uncombined zinc as a 
 mineral is doubtful. Zinc and cadmium occur as the sulphides, 
 Zn S and Cd S, in an isomorphous and dimorphous group which also 
 includes the sulphides of manganese, iron, and nickel. The sulphide 
 of zinc occurs in crystals belonging to the regular system termed 
 sphalerite or zinc-blende, and in hexagonal crystals (wurtzite). The 
 sulphide of cadmium is isomorphous with wurtzite, and is known as 
 greenockite. Zinc is sometimes found as zincite, which is the ox- 
 
 * It seems scarcely necessary to state that the term monatomic means ex- 
 isting as molecules formed of one atom, diatomic of two, etc. 
 
ZINC; CADMIUM; MERCURY; METALLURGY. 427 
 
 ide, ZnO, colored red by means of manganese. The carbonate of 
 zinc, Zn C0 3 (smithsonite), is isomorphous with calcite (see page 
 416). An aluminate of zinc, Zn (Al 2 ), isomorphous with spinell 
 (see page 339), is also met with. 
 
 Mercury is sometimes found in small, fluid globules in places 
 where the most important mineral containing mercury (namely, the 
 red sulphide, cinnabar [HgS],) also occurs. In addition, amal- 
 gams of mercury with silver and gold are sometimes met with. 
 
 Zinc is obtained from its ores by roasting the sulphide in a 
 draught of air, and by subsequently heating with charcoal the oxide 
 produced by this means. As zinc is volatile at no very high tem- 
 perature, the mixture of oxide and charcoal is placed in earthen- 
 ware retorts which are gradually heated ; carbon monoxide passes 
 off during this process : 
 
 Zn -f C = Zn + CO. 
 
 Finally, the temperature is increased to a point at which zinc begins 
 to distil; earthenware receivers are then placed before the open- 
 ings of the retorts, and the metal is collected therein. The first 
 portions which pass over are deposited on the walls of the receivers 
 in the form of a fine dust which always contains more or less oxide 
 of zinc ; this product, which is known as zinc-dust, is very fre- 
 quently used as a reducing agent in the laboratory. The fused zinc 
 which finally collects is generally impure, containing lead, iron, and 
 cadmium ; it is separated from those metals by repeated distillation. 
 The cadmium, having a lower boiling point, passes over first, while 
 the lead and iron remain behind. The preparation of cadmium is 
 like that of zinc. As cadmium ores generally contain zinc, the 
 metal is separated from the latter by distillation. Mercury is ob- 
 tained by roasting the sulphide, the mercury . which passes off being 
 collected in receivers which are connected with the oven. The ad- 
 dition of charcoal is unnecessary during this process, because the 
 oxide of mercury, which would be formed by roasting the sulphide, 
 is further decomposed into mercury and oxygen by heat (see page 
 18). 
 
 As has been mentioned, zinc is a constituent of the alloy known 
 as brass. When sheet-iron is coated with zinc, it is known as gal- 
 vanized iron. Zinc readily forms an amalgam with mercury ; an 
 extensive commercial use of this fact is made in amalgamating zinc 
 
428 ZINC; CADMIUM; OXIDES. 
 
 battery plates, the latter being cleaned, dipped in acid, and rubbed 
 with mercury so as to produce a thin layer of amalgam. Cadmium, 
 zinc, and mercury form an amalgam which readily hardens ; this sub- 
 stance is used in filling teeth. Alloys of cadmium with lead and bis- 
 muth are used where a very low-melting metal is required. Alloys 
 of mercury are termed amalgams ; a number of these have definite 
 composition and crystalline form. The nature of amalgams has 
 been discussed on page 250. 
 
 Zinc and cadmium form but one oxide apiece ; these oxides, in 
 formulae, correspond to the typical oxide of the family, M 0. The 
 hydroxides, M(OH) 2 , can be produced by adding soluble hydrox- 
 ides to solutions containing salts of cadmium or zinc : 
 
 MC1 2 + 2KOH = 2KC1 + M(OH) 2 . 
 
 However, in adding a caustic alkali to a zinc salt, care must be 
 taken not to use an excess of the reagent ; for zinc hydroxide acts 
 as an acid when in the presence of strong bases, dissolving in the 
 latter to form zincates : 
 
 Zn(OH) 2 + 2KOH = Zn (OK) 2 + 2 H 2 0, 
 
 (see pages 315 and 322) ; the zincate so formed is, however, de- 
 composed by boiling, zinc hydroxide being precipitated ; ammonia 
 water has the same effect as solutions of the caustic alkalies, an 
 excess of that reagent dissolving the precipitated hydroxide while 
 producing ammonium zincate.* Acid solutions of zinc salts, or 
 mixtures of ammonium and zinc salts, are not affected by ammonia. 
 In the case of cadmium, caustic alkalies precipitate the hydroxide, 
 which, however, is not dissolved by an excess of the reagent, cad- 
 mium hydroxide having no acid properties ; on the other hand, 
 ammonia, when in the presence of ammonium salts, produces no 
 precipitate, for cadmium resembles zinc and magnesium in the facil- 
 ity with which its salts form compounds with those of ammonium. 
 
 Zinc oxide is a white powder ; yellow when heated. 
 
 Cadmium oxide is a brown powder. 
 
 Zinc hydroxide is a white powder; changes to zinc oxide and 
 water when heated. 
 
 * Difference between zinc and magnesium, for ammonia precipitates a 
 portion of the magnesium as magnesium hydroxide. The latter is insoluble 
 in an excess of the reagent (see page 419). 
 
ZINC ; CADMIUM ; SALTS OF. 429 
 
 Cadmium hydroxide is a white powder ; changes to cadmium 
 oxide and water when heated. 
 
 In addition to the methods for the preparation of the oxides 
 which have been given above, those compounds can also be pro- 
 duced by heating the respective carbonates or nitrates. 
 
 The chloride of zinc can be produced by heating zinc in a cur- 
 rent of chlorine, or by dissolving zinc, or the oxide or hydroxide of 
 the metal, in hydrochloric acid, evaporating the solution and dis- 
 tilling. The fused salt is cast into sticks which are extremely 
 deliquescent. The salt crystallizes from a concentrated solution of 
 hydrochloric acid in crystals having the formula Zn C1 2 + H 2 ; 
 but when an aqueous solution of zinc chloride is heated, a partial 
 decomposition into the basic chloride takes place : 
 
 Zn 
 
 and this substance, when boiled with water, finally loses all chlo- 
 rine and changes to the hydroxide (see magnesium chloride, pages 
 414, 415). The chloride of cadmium, Cd C1 2 , is similar to that of 
 zinc ; like zinc chloride, it is volatile, but is not decomposed when 
 the solution is evaporated. 
 
 Zinc sulphate, formed by dissolving either the hydroxide, oxide, 
 carbonate, or the metal in sulphuric acid, belongs to the class of sul- 
 phates which are termed vitriols (see page 417). Like all of the 
 sulphates belonging to this group, it is soluble in water and crystal- 
 lizes with seven molecules of water, six of which it loses at 100, 
 while the seventh passes off only at a higher temperature. 
 
 Cadmium, sulphate does not belong to che group of vitriols. Its 
 crystals have the formula 3 Cd S0 4 -j- 8 H 2 0. 
 
 When a solution of the carbonate of .an alkali metal or of 
 ammonium is added to the solution of a zinc salt, an insoluble basic 
 carbonate of zinc is precipitated. This substance has a varying 
 composition, according to the conditions under which it is produced. 
 The normal secondary carbonate of zinc, Zn C0 3 , occurs as the min- 
 eral smithsonite, belonging to the calcite group. The carbonate 
 is easily decomposed into zinc oxide and carbon dioxide when it is 
 heated. 
 
 Cadmium carbonate is precipitated from solutions containing a 
 cadmium salt by addition of a soluble carbonate. Owing to the 
 
430 MERCURY; OXIDES OF. 
 
 more metallic nature of cadmium, the precipitate so formed con- 
 sists of the secondary carbonate, Cd C0 3 . 
 
 Zinc sulphide is precipitated from the neutral or alkaline solu- 
 tions of zinc salts by addition either of hydrogen sulphide or of a 
 soluble alkaline sulphide.* When so precipitated, it forms a white 
 powder ; in a crystalline state it is found as zinc-blende, a substance 
 which has much the same color as ordinary resin. 
 
 Cadmium sulphide is precipitated from solutions of cadmium 
 salts,, even if the latter are slightly acid, for the substance is insolu- 
 ble in dilute acids. The precipitate is yellow in color. 
 
 Mercury forms two classes of compounds, mercurous compounds, 
 derived from the metal in its monovalent state, and mercuric com- 
 pounds, derived from divalent mercury. The same distinction, it 
 will be remembered, existed between cuprous and cupric derivatives. 
 
 Oxides of mercury. Mercurous oxide ( Hg 2 0), mercuric oxide 
 ( Hg 0). The former is produced by adding potassium hydroxide 
 to a solution of mercurous nitrate, the hydroxide, which would be 
 expected, at once breaking down into water and the oxide : 
 
 1. 2 Hg N0 3 + 2 KOH = 2 KST0 3 + 2 Hg OH, 
 
 2. 
 
 Mercurous oxide is a black powder which is quite unstable ; when 
 exposed to the light, it breaks down into mercury and mercuric 
 
 oxide: - Hg 2 = Hg + HgO.:f 
 
 * If the zinc salt is the salt of a strong acid, such as hydrochloric, nitric, 
 or sulphuric, only a portion of the zinc is precipitated as the sulphide, by 
 means of hydrogen sulphide; for, as will be seen from the following reaction, 
 a portion of the acid is set free during the reaction, and the acid which is 
 formed decomposes the precipitated sulphide in order to form once more a 
 soluble salt of zinc : 
 
 Zn SO 4 + H,, S = Zn S + H 2 SO 4 , 
 
 Zn S + H 2 SO* = Zn SO 4 +~H. 2 S. 
 
 This difficulty is not encountered if the zinc salt of a weak acid is used, or 
 if ammonia is previously added so as to neutralize any acid which may be 
 liberated. 
 
 t It will be remembered that the same is true of the formation of the 
 oxide of silver; when a soluble hydroxide is added to the solution of a silver 
 salt, not the hydroxide, but the oxide, is precipitated : 
 
 2 Ag NO 3 + 2 KOH = 2 KNO 3 + Ag 2 O + H. 2 O. 
 
 J This change reminds us of the ones \vhich we encountered with many 
 of the oxides and acids of the not-metals. 
 
MERCUEOUS CHLORIDE. 431 
 
 Addition of acids produces mercurous salts ; oxidizing agents change 
 mercurous compounds into mercuric compounds. 
 
 Mercuric oxide exists in two forms, according to the method of 
 its preparation; the one is red and of crystalline structure, the 
 other is yellow and amorphous. The red oxide can be prepared 
 either by heating mercury to just below its boiling point in the 
 presence of oxygen, when, after a long time, it becomes covered with 
 crystalline scales of the substance sought ; or by heating mercuric 
 nitrate, which salt breaks down into mercuric oxide and nitrogen 
 peroxide (see page 201) : 
 
 The yellow oxide of mercury is produced by precipitation from 
 solutions of mercuric salts after the addition of a soluble hydroxide, 
 the hydroxide at once breaking down into the oxide and water : 
 
 1. Hg(N0 3 ) 2 + 2 KOH = Hg (OH ) 2 + 2 KNO , , 
 
 2. Hg(OH) 2 = HgO + H 2 0. 
 
 Both varieties of mercuric oxide turn black when heated ; they 
 resume their usual color after cooling; at a dull red heat they 
 decompose into mercury and oxygen (see page 18) ; sunlight has 
 the same effect as heat. When mercuric oxide is treated with an 
 acid it produces mercuric salts. All soluble mercury compounds, as 
 well as the metal itself, are extremely poisonous. Under certain 
 circumstances mercuric oxide displays slightly acidic properties ; 
 for instance, it is attacked, in small quantities, by fused potassium 
 hydroxide.* 
 
 Chlorides of mercury. Mercurous chloride (calomel), HgCl, 
 mercuric chloride (corrosive sublknate), Hg C1 2 . 
 
 Mercurous chloride sometimes occurs in a crystalline form as a 
 mineral. Mercurous chloride can be prepared either by heating 
 mercuric chloride with a sufficient quantity of mercury : f 
 
 HgCl 2 + Hg = 2HgCl, 
 
 or by adding hydrochloric acid, or a soluble chloride, to a solution 
 
 * Precipitated (yellow) mercuric oxide is also slightly soluble in cold 
 solutions of caustic alkalies. 
 
 t This operation must be carried on in large flasks stoppered with chalk. 
 The calomel then sublimes from the lower part of the flask and collects on the 
 cooler portions. If the vessel is too small, a portion of the mercurous salt 
 will evaporate. 
 
432 MERCUROUS CHLORIDE. 
 
 containing a mercurous salt, for mercurous chloride is insoluble in 
 water or dilute acids : 
 
 H Cl = HgCl-f-HN0 3 , 
 Hg N0 3 + Na Cl = Hg Cl + Na N0 3 . 
 
 In this respect, then, mercurous compounds are much like those of 
 silver ; indeed, all of the monovalent heavy metals act alike in pro- 
 ducing insoluble chlorides.* 
 
 Mercurous chloride, when heated, evaporates without previously 
 melting. The specific gravity of the vapor is 8.01 ; the calculated 
 specific gravity for a gas composed of molecules of the formula 
 HgCl is 8.14; in this respect mercurous chloride differs from 
 cuprous chloride, for the vapor density of the latter compound leads 
 to a formula Cu 2 C1 2 . If mercurous chloride is treated with a solu- 
 tion of ammonia, it turns black and produces an insoluble compound 
 which has the formula NH 2 Hg 2 Cl : 
 
 2 Hg Cl + 2 NH 3 = NH 2 Hg 2 Cl + NH 4 Cl ; 
 
 this substance is regarded as being ammonium chloride in which 
 two atoms of hydrogen have been replaced by two of mercury : f 
 
 1ST 
 
 fH 
 H 
 
 H ammonium chloride, and 
 H 
 
 H 
 H 
 
 Hg mercurous chloramide. 
 
 Hg 
 
 Cl 
 
 ,61 
 
 If mercurous chloride is boiled with hydrochloric acid it is converted 
 into mercuric chloride and mercury : 
 
 2HgCl=Hg+HgCl 2 . 
 
 Sulphuric acid, hot and concentrated, changes it into a mixture of 
 mercuric sulphate and mercuric chloride, nitric acid into mercuric 
 nitrate and mercuric chloride. Chlorine readily converts calomel 
 into mercuric chloride. 
 
 Mercuric chloride. This salt can be produced either by heating 
 
 * Cuprous chloride, Cu Cl, argentic chloride, Ag Cl, aurous chloride, Au Cl, 
 mercurous chloride, Hg Cl, and thallous chloride, Tl Cl, are insoluble. 
 
 t It will be remembered that the same power of replacing hydrogen in 
 ammonium chloride is found in the case of cuprous chloride ; the formula of 
 this ammonium compound is, however, Cu Cl, NH 3 or NH 3 Cu Cl. Argentic 
 chloride, Ag Cl, also possesses the power of absorbing ammonia. 
 
MERCURIC CHLORIDE. 433 
 
 mercury in a current of chlorine, * by dissolving the metal in aqua 
 regia, or by dissolving mercuric oxide in hydrochloric acid. The 
 corrosive sublimate of commerce is usually prepared by heating 
 mercuric sulphate with sodium chloride in a wide-mouthed retort ; 
 the mercuric chloride sublimes and collects on the cold portions of 
 the vessel, while sodium sulphate remains behind : 
 
 Hg S0 4 + 2 Na Cl = Hg C1 2 + Na 2 S0 4 . 
 
 Corrosive sublimate is a transparent, crystalline mass which is 
 soluble in water and which crystallizes from that solvent when 
 evaporated ; it crystallizes from aqueous solution in thin prisms ; 
 one hundred parts of water at dissolve 5.73 parts, and at ordinary 
 temperatures about 7 parts of mercuric chloride. The solutions 
 gradually decompose when they are exposed to the light, while an 
 insoluble basic chloride of mercury is formed. No change takes 
 place if they are kept in the dark. Mercuric chloride has a great 
 tendency to form double salts with the chlorides of other metals ; 
 so, for instance, the compounds Na Cl, Hg C1 2 + 3 H 2 ; K Cl, Hg 
 C1 2 + H 2 are known ; in this respect mercuric chloride has 
 strongly marked acidic properties, although, when heated with phos- 
 phorus pentachloride, it can act as a base, for it produces the 
 double salt 3 Hg C1 2 , 2 P C1 5 .f Mercuric chloride is so volatile 
 that a portion passes off when its solution is evaporated with hy- 
 drochloric acid ; this loss mav be prevented by adding potassium or 
 sodium chloride, for the double chloride which is formed is not vol- 
 atile. Reducing agents readily convert mercuric chloride into mer- 
 curous chloride. Stannous chtoride changes corrosive sublimate 
 into calomel : 
 
 2 Hg C1 2 + Sn C1 2 = 2 Hg Cl -h Sn C1 4 , 
 
 and if an excess of the reagent is added, the mercurous chloride is 
 finally changed to mercury (see page 314) : 
 
 2 Hg Cl + Sn C1 2 = 2 Hg + Sn C1 4 . 
 
 * The mercury burns with a pale flame, and forms a white sublimate of 
 mercuric chloride. 
 
 t 3 Hg C1 2 , 2 P C1 5 = Hg 3 (P C1 8 ) 2 . Compare this formula with that of the 
 tertiary phosphate ; viz., Hg 3 ( PO 4 ) 2 . 
 
434 MERCUROUS NITRATE. 
 
 Ammonia produces a white precipitate of mercuric chloramide, 
 NH 2 Hg Cl. This substance is analogous to the corresponding 
 mercurous compound with the exception that one atom of bi- 
 valent mercury takes the place of two atoms of hydrogen. The 
 distinction between the two is made apparent by the following 
 formulse : 
 
 HI 
 
 H 
 
 Hg mercurous chloramide and N-[ TT mercuric chloramide* 
 
 Her M S 
 
 PI 
 
 Cl 
 
 The iodides of mercury are similar to the chlorides. Mercurous 
 iodide, formed by the direct union of mercury and iodine, readily 
 breaks down into mercury and mercuric iodide : 
 
 (the chloride suffers the same change when heated). Mercuric 
 iodide displays even a greater tendency to form double salts (in 
 which it plays the part of an acidic anhydride) than does mercuric 
 chloride. The fact that the halides of so many metals have acidic 
 properties, while the oxides do not, is not difficult to comprehend if 
 we remember that chlorine, bromine, and iodine belong to the most 
 not-metallic family with which we are acquainted, so that the halo- 
 gen compounds should be more negative than are those of oxygen. 
 Some compounds are known in which the oxide of a metal is the 
 base, and the chloride, bromide, or iodide is the acidic anhydride. 
 Such compounds are the double' oxychlorides of mercury. An 
 example of one of these compounds is Hg 0, 2 Hg C1 2 , the insoluble 
 precipitate formed by the gradual decomposition of a mercuric 
 chloride solution. 
 
 Mercurous nitrate is produced by the action of dilute nitric acid 
 upon mercury. When diluted with water, it forms basic nitrates 
 which are soluble with difficulty. These basic nitrates vary in com- 
 
 * A number of similar compounds derived from ammonia are known ; for 
 example, dimercuric ammonium chloride ( NHg. 2 Cl) and similar salts derived 
 from other acids have been described, but for these a larger text-book must be 
 consulted. 
 
MERCURIC SALTS. 435 
 
 position according to the temperature or the amount of water used. 
 The simplest compound has the formula HgN0 3 , Hg 2 -f- H 2 0, 
 and may possibly be a tertiary salt of an hypothetical ortho-nitric 
 acid, H 3 N0 4 (see page 205) : - 
 
 Mercuric nitrate is produced by dissolving mercury in concen- 
 trated nitric acid ; when dissolved in an excess of water it forms an 
 insoluble basic nitrate. 
 
 Mercuric cyanide, Hg (CN ) 2 , is of importance because it is the 
 only cyanide of the heavy metals which is soluble in water. For 
 this reason it is a very useful reagent in chemical analysis. The 
 cyanide can be readily produced by dissolving mercuric oxide in a 
 solution of hydrocyanic acid. When heated, it decomposes into 
 mercury and cyanogen (see page 294). Mercuric cyanide readily 
 forms double salts with the cyanides of other metals. 
 
 Mercurous sulphide has not, as yet, been prepared. Hydrogen 
 sulphide, when passed through a solution of a mercurous salt, pre- 
 cipitates a mixture of mercuric sulphide and mercury. Mercuric 
 sulphide is the chief ore of mercury ; it is a red, crystalline mineral, 
 termed cinnabar. When hydrogen sulphide is added to the slightly 
 acid solution of a mercuric salt, mercuric sulphide is precipitated in 
 the form of a black powder ; the latter changes into the red variety 
 by heating, or by treating for a long time with solutions of caustic 
 alkalies. Exposure to the light gradually changes the red sulphide 
 into the black variety. Black mercuric sulphide is also produced 
 by rubbing mercury and sulphur together. Mercuric sulphide is 
 not attacked by dilute acids ; concentrated nitric acid in part dis- 
 solves it, and in part converts it into a white, insoluble double com- 
 pound of mercuric nitrate and mercuric sulphide; this compound 
 has the formula Hg ( N 3 ) 2 2 Hg S. Cinnabar, when heated, turns 
 black, and, unless the temperature was too high, resumes its original 
 red color on cooling. 
 
 All mercury compounds, if they are salts of volatile acids, are 
 volatile ; if, on the other hand, they are salts of not-volatile acids, 
 either the acids themselves or their decomposition products remain 
 after heating (see page 190). 
 
 * See Remsen ; Chemistry, p. 627. 
 
436 
 
 ZINC; CADMIUM; MERC UK Y ; TABLE OF. 
 
 TYPICAL COMPOUNDS IN THIS FAMILY. 
 
 
 OXIDES. 
 
 HYDBOXIDES. 
 
 CHLOBIDEB. 
 
 SULPHATES. 
 
 SULPHIDES. 
 
 Zinc 
 Cadmium 
 Mercury 
 
 ZnO 
 CdO 
 HgO* 
 
 Zn (OH) 2 t 
 Cd (OH) 2 
 
 Zn C1 2 J 
 Cd C1 2 J 
 HgCl 2 
 
 Zn SO 4 + 7 H 2 O 
 3 Cd S0 4 + 8 H 2 O 
 HgS0 4 
 
 ZnS** 
 CdS 
 HgS 
 
 
 * Exists in two varieties, yellow and red. 
 
 t Both basic and acidic in its character: Zn (OH) 2 + 2HC1 = ZnCl 2 + 
 2 H 2 O and Zn (OH) 2 + 2 KOH = Zn (OK) 2 + 2 H 2 O. 
 
 J Readily form double chlorides with ammonium chloride. The latter are 
 not decomposed by ammonia. 
 
 Forms mercuric chloramide (NH 2 HgCl) with ammonia. 
 
 ** Soluble in dilute acids. 
 
 All of the salts of the elements belonging to this group show a great ten- 
 dency to produce double salts. 
 
 MEKCUROUS COMPOUNDS (not typical). 
 
 Hg 2 O, mercurous oxide. 
 
 Hg Cl, mercurous chloride. 
 
 HgNO 3 , mercurous nitrate. 
 
 Mercurous chloride is insoluble in water. When covered with ammonia 
 solution it forms mercurous chloramide (NH 2 Hg 2 Cl), the nitrate forms basic 
 salts on addition of water. 
 
THE GADOLLNTTE EARTHS. 437 
 
 CHAPTER LVI. 
 
 THE ELEMENTS BELONGING TO THE PRIMARY GROUPS OP THE 
 FAMILIES HE., IV., AND V., OF THE LONG PERIODS. 
 
 THE elements comprising the primary group of the third family 
 bear somewhat the same relation to boron and aluminium that cal- 
 cium, strontium, and barium do to beryllium and magnesium. The 
 elements are : 
 
 Scandium ; symbol, Sc ; atomic weight, 44 ; 
 Yttrium; symbol, Y; atomic weight, 89.1; 
 Lanthanum ; symbol, La ; atomic weight, 138.2 ; 
 Ytterbium; symbol, Yb; atomic weight, 173. 
 
 All of these elements are extremely rare; they occur in an 
 ortho-silicate known as gadolinite, the most common formula of 
 which is Be 2 ( Y0) 2 Fe (Si 4 ) 2 , and also in an extremely complicated 
 salt of titanic acid known as euxenite. Considerable uncertainty 
 has existed as to the number of elements really contained in this 
 and in the following group ; for some investigators, by reason of 
 the peculiarities of the absorption spectra of some of the salts 
 of these metals, have undertaken to prove that the usually accepted 
 number must be largely increased. 1 * 
 
 Scandium, it will be remembered, was one of the elements pre- 
 dicted by Mendelejeff (see page 373). The typical oxide of these 
 elements is M 2 3 , corresponding to B 2 3 and A1 2 O 3 ; this oxide is 
 basic in its character ; it does not dissolve in caustic alkalies ; so that 
 these elements are more metallic than is aluminium. The hydrox- 
 ides have the formula M (OH) 3 , the sulphates, M 2 (S0 4 ) 3 , and, un- 
 like the sulphates of the secondary group of this family ; t they do 
 
 * Elements occurring in gadolinite and belonging to this group are gado- 
 linium, 156.1; samarium, 150; terbium, 160; erbium, 166.3. For the descrip- 
 tion of these elements the original literature must be consulted. For methods 
 of separation, see Kriiss ; Liebig's Annalen ; 265, 1. Didymium also was among 
 the elements of the gadolinite group; this substance has recently been sepa- 
 rated into two elements, neodymium and praseodymium. 
 
 t Sulphates of aluminium, gallium, indium, thallium. 
 
438 ELEMENTS OF TITANIUM GKOUP. 
 
 not form alums ; this distinction is similar to that existing between 
 the sulphates of calcium, strontium, and barium, and those of zinc, 
 cadmium, and mercury; for the sulphates of the first three are 
 insoluble, while those of the second three are soluble, and have 
 a great tendency to form double salts with the sulphates of the 
 alkalies. 
 
 The elements comprising the primary group of family IV. are : 
 
 Titanium ; symbol, Ti ; atomic weight, 48 ; 
 Zirconium ; symbol, Zr ; atomic weight, 90.6 ; 
 Cerium ; symbol, Ce ; atomic weight, 140.2 ; 
 Thorium ; symbol, Th ; atomic weight, 232.6. 
 
 Of the compounds of these four elements, those of titanium are 
 by far the most common ; indeed, compounds containing titanium 
 are not at all infrequent, for the element occurs in many iron ores ; 
 titanic iron ore is looked upon as being ferrous titanate, Fe Ti 3 ; 
 the compound is, however, isomorphous with ferric oxide, and fre- 
 quently occurs in company with that extremely important sub- 
 stance ; furthermore, titanic oxide is often a constituent of magnetic 
 iron ore, Fe 3 O 4 , which latter substance is a member of the spinell 
 group (see page 339), so that it seems not improbable that titanic 
 iron is really derived from a hydroxide, Ti (OH ), which is analo- 
 gous to Al (OH ) ; if this relationship is granted, then titanium 
 can, under certain circumstances, act as a trivalent element, and 
 this behavior would bring it in line with cerium (see below). 
 Titanium dioxide, Ti 2 , is found in two mineral forms, brookite 
 and anatas,* and as a polymeric form, Ti 2 4 , which is called rutile, 
 and which is isomorphous with an ortho-silicate of zirconium, 
 Zr Si O 4 , known as zircon, and with tinstone (see page 312). The 
 relationship between these three oxides, all of which belong to ele- 
 ments in the same family and are isomorphous, becomes apparent 
 if we double the formula of the oxide of tin, thus : 
 
 SnSn0 4 , tinstone, 
 Ti Ti 4 , rutile, 
 Zr Si 4 , zircon. 
 
 * These two oxides of titanium are not isomorphous with quartz and 
 tridymite, yet the form of brookite is so close to that of tridymite that isomor- 
 phism is considered possible (Groth). (See page 304.) 
 
TITANIUM GROUP ; COMPOUNDS OF. 
 
 439 
 
 Cerium occurs in gadolinite and also in a silicate termed cerite. 
 
 The compounds of the elements of this group are analogous to 
 those of silicon ; this connection will be seen from the following 
 table : 
 
 
 OXIDES. 
 
 CHLORIDES. 
 
 FLUORIDES. 
 
 HYDROXIDES. 
 
 METAHYDBOXIDES. 
 
 Silicon 
 Titanium 
 
 Si0 2 
 TiO 2 
 
 7r O 
 
 SiCl 4 * 
 TiCl 4 * 
 
 7r Cl t 
 
 SiF 4 
 TiCl 4 
 Zr PI 8 
 
 Si(OH) 4 
 Ti(OH) 4 
 Zr (OH)4 
 
 Si0 3 H 2 
 
 Ti0 8 H 2 (?) 
 
 
 Th O 
 
 Th Cl t 
 
 Th CL S 
 
 Th (OH)4 
 
 
 
 
 
 
 
 
 * Decomposed by water. 
 
 t Partially decomposed by water; forming a basic chloride: 
 
 ( Cl + HOH 
 -I Cl + HOH 
 
 Zr ) Cl 
 
 lei 
 
 OH 
 
 "1 C1 
 
 lei 
 
 Zr 
 
 OH 
 
 ATT 
 
 X; i 
 
 Cl 
 
 =zr 
 
 Cl 
 
 t Not decomposed except by hot water. 
 
 All of these fluorides behave exactly as does silicon tetrafluoride ; they 
 form compounds analogous to fluosilicic acid and the fluosilicates (see page 
 303): 
 
 H 2 Si F 6 , fluosilicic acid ; K 2 Si F 6 , potassium fluosilicate. 
 H 2 Ti F 6 , fluotitanic acid ; K 2 Ti F 6 , potassium fluotitanate. 
 K 2 Zr F 6 , potassium fluozirconate. 
 K 2 Th F 6 , potassium fluothorate. 
 
 The compounds of cerium are omitted from this table because the chemistry 
 of this element is not yet sufficiently clear for purposes of comparison. 
 
 The above table and the explanatory notes very plainly show 
 the intimate family connection between silicon and the three fol- 
 lowing elements ; the oxides of titanium and zirconium resemble 
 silicon dioxide in the fact that, after they have been heated, they 
 are insoluble in water and even in the strongest acids or alkalies ; to 
 be brought in solution they must be heated with concentrated sul- 
 phuric acid for a long time, or they must be fused with alkalies ; the 
 oxide of thorium is somewhat less obstinate. The last three ele- 
 ments in the above table, having higher atomic weights than silicon, 
 are also more metallic in their nature ; their hydroxides are, conse- 
 quently, both weak acids and weak bases. The salts in which they 
 act as acids are but little known ; indeed, it is doubtful if thorium 
 hydroxide has acidic properties. The salts of the alkalies, so far 
 as known, correspond to the metasilicates, and hence have the 
 
440 ELEMENTS OF VANADIUM GROUP. 
 
 general formula M 2 X0 3 . Among the salts derived from the hy- 
 droxides, acting as bases, the sulphates, with the general formula 
 M(S0 4 ) 2 , are perhaps the most prominent. The chemistry of 
 cerium is, as yet, uncertain in many respects ; the element forms 
 two series of compounds, in one of which, presenting compounds 
 such as Ce 2 3 , CeCl 3 , Ce(N0 3 ) 3 , it is trivalent, and resembles 
 lanthanum, an element in the preceding family ; in the other it is 
 tetravalent, and by means of compounds Ce F 4 , Ce 2 , Ce ( S0 4 ) 2 , it 
 falls in line with the family numbered IV. It may be added that 
 titanium likewise forms more than one oxide ; for a compound, 
 Ti 2 3 , and a sulphate, Ti 2 (S0 4 ) 3 , have been described. 
 
 The elements comprising the primary group of family V. are : 
 
 Vanadium ; atomic weight, 51.4 ; symbol, V. 
 
 Columbium ; atomic weight, 94; symbol, Cb. 
 
 Neodymium ; atomic weight, 140.5 ; symbol, Nd. 
 
 Praseodymium ; atomic weight, 143.5 ; symbol, Pr. 
 
 Tantalum, atomic weight, 182.6 ; symbol, Ta. 
 
 Although more metallic in their nature than the elements form- 
 ing the secondary group of this family,* the four elements, with 
 perhaps the exception of neodymium and praseodymium, bear many 
 points of resemblance to this group. In some respects the two lat- 
 ter substances are very much like cerium and lanthanum, being 
 trivalent in most of their compounds. Vanadium is the best known 
 of all of these elements. This element is as much like arsenic or 
 antimony as titanium is like silicon ; in very many respects it is, in- 
 deed, like the typical element of the family, nitrogen, for it forms 
 as many oxides as the latter, and these oxides have similar for- 
 mulae, thus : 
 
 OXIDES OF NITROGEN. OXIDES OF VANADIUM. 
 
 N 2 V 2 
 
 NO(N 2 2 ) V 2 2 
 
 N 2 3 V 2 3 
 
 N0 2 ,(N 2 4 ) V 2 4 
 
 N 2 5 V 2 5 
 
 Vanadium occurs in nature chiefly as the lead, zinc, or bismuth 
 salt of vanadic acid. Vanadinite is a double salt composed of lead 
 
 * Arsenic, antimony, and bismuth. 
 
VANADIC ACIDS. 
 
 441 
 
 vanadate and lead chloride, having the formula 3 Pb 3 (V0 4 ) 2 -j- 
 PbCl 2 *. The element forms three chlorides, V C1 2 , V C1 3 , and 
 VC1 4 . The most important compounds of vanadium are the va- 
 nadic acids, which correspond to those of phosphorus and of arsenic, 
 the latter being elements belonging to the same family : - 
 
 
 OXIDES. 
 
 META-ACIDS. 
 
 OETHO-ACIDS. 
 
 PYKO-ACIDS. 
 
 Phosphorus 
 Arsenic 
 Vanadium 
 
 P 2 5 
 
 As 2 5 
 V 2 5 
 
 HPO 3 
 HAsO 3 
 HV0 3 
 
 H 3 P0 4 
 
 H 3 As O 4 
 H 3 V0 4 
 
 H 4 P 2 7 
 H 4 As., O 7 
 H 4 V 2 7 
 
 Vanadates of all these acids are known ; for instance, we have : 
 
 Na V0 3 , sodium metavanadate ; 
 Na 3 V0 4 , sodium orthovanadate ; 
 Na 4 V 2 7 sodium pyrovanadate. 
 
 In addition to the above, however, more complicated salts of 
 polyvanadic acids, which are formed in the same manner as the 
 polysilicic acids, are known. Free metavanadic acid is a golden 
 yellow, crystalline compound. Reducing agents readily change 
 vanadic acid into the lower acid, V 2 4 . 
 
 Columbium and tantalum are very rare elements which differ 
 from vanadium just as much as antimony does from arsenic; for 
 they are able to form pentachlorides and pentabromides, while the 
 pentahalides of vanadium have not as yet been prepared. Colum- 
 bium is also frequently termed niobium. It occurs in the min- 
 eral columbite, which is a metacolumbate of iron, having the 
 formula Fe (Cb 3 ) 2 ; tantalum is found as a metatantalate of iron, 
 Fe(Ta0 3 ) 2 . 
 
 The brief mention of the very rare metals which have been 
 discussed in this chapter is sufficient to demonstrate the family 
 relationship existing between them and the much more common 
 elements which were described in the first portion of the book ; of 
 course, they form a great number of compounds, some of them very 
 complicated, which cannot be taken up in a book of this kind ; for 
 this study a large manual must be consulted. 
 
 * Yanadinite is isomorphous with apatite, which has a formula 3 Ca 3 
 ( PO 4 ) 2 + Ca C1 2 , calcium replacing lead, and vanadium replacing phos~ 
 phorus isomorphously. 
 
442 GADOLINITE METALS AXD PERIODIC SYSTEM. 
 
 The uncertainty which exists as to the number of elements 
 which can be isolated from gadolinite, euxenite, and allied minerals 
 leaves that portion of the periodic system in which these elements 
 find their places in a very chaotic condition. Certainly, if all of 
 these supposed elements should finally prove themselves to be such, 
 we should then have to conclude that several elements, with very 
 nearly the same atomic weight, must occupy the same position in 
 the periodic system. This discovery need not overthrow existing 
 relations with well-known elements, for the following are facts 
 which cannot be controverted. 
 
 The periodic system, as it now stands, is undoubtedly a natural 
 grouping of the elements. The fact that certain portions of the 
 system are now completely filled with well-known elements does 
 not prove that other elements, intermediate to known groups, will 
 never be discovered. Indeed, were we acquainted with seven hun- 
 dred elements, instead of seventy, these seven hundred would no 
 more be unconnected individuals than are our present number. 
 They too would fall into an arrangement, in periods and families, 
 as the elements now do ; only, with such a large number of indi- 
 viduals, the change from family to family would not present such 
 an abrupt transition as at present. 
 
ELEMENTS OF CHROMIUM GKOUP. 
 
 443 
 
 CHAPTER LVII. 
 
 THE ELEMENTS BELONGING TO THE PRIMARY GROUP OP THE 
 
 VI. FAMILY. 
 
 Chromium ; symbol, Or ; atomic weight, 52.1 ; 
 
 Molybdenum ; symbol, Mo ; atomic iveight, 96 j 
 
 Tungsten ( Wolfram) ; symbol, W ; atomic weight, 184 ; 
 
 Uranium ; symbol, U ; atomic weight, 239.6. 
 THE typical elements belonging to the sixth family, in the short 
 periods, are oxygen and sulphur ; and, as has been shown by the 
 arrangement of the periodic system given on page 363, the individ- 
 uals more immediately connected with those two elements are sele- 
 nium and tellurium, while chromium, molybdenum, tungsten, and 
 uranium, having their positions near the middle of the long periods, 
 vary much more from the character of the types in the short periods 
 than do the metals which have been discussed in the preceding chap- 
 ters. The metallic nature of the elements forming the primary 
 group of the sixth family is most apparent in the behavior of the 
 lower oxides and in the salts derived from these ; the highest oxide 
 of each element is the typical one, X0 3 ; this compound displays 
 the character of an acidic anhydride, although, in the case of the 
 most metallic element of the family (uranium), it is also basic under 
 some circumstances. The salts derived from the typical oxide, in 
 formulae, correspond to the sulphates, selenates, and tellurates, and 
 in some instances to the pyrosulphates (see page 154); this relation- 
 ship is made apparent by the following table : 
 
 OXIDES. 
 
 ACIDS. 
 
 SALTS. 
 
 DI-8ALT8. 
 
 OXIDES. 
 
 ACID8.- 
 
 SALTS. 
 
 PI-SALTS. 
 
 S0 3 
 
 H 2 S0 4 
 
 M, S0 4 
 
 M 2 S 2 7 
 
 Cr0 3 
 
 H 2 CrO 4 * 
 
 M 2 Cr0 4 
 
 M 2 Cr 2 7 
 
 (Se0 3 ) 
 
 H 2 Se0 4 
 
 M 3 SeO 4 
 
 
 MoO 3 
 
 H 2 MoO 4 t 
 
 M 2 Mo O 4 
 
 M 2 Mo 2 O 7 
 
 Te0 3 
 
 H 2 TeO 4 
 
 M 2 Te 4 
 
 
 WO 3 
 
 H a W0 4 t 
 
 M 3 WO 4 
 
 M 2 W 2 7 
 
 
 
 
 
 UO 3 
 
 
 
 M 2 U 2 O 7 
 
 * The acid is not known ; when liberated from its salts it breaks down 
 into its anhydride, Cr O 3 , and water. 
 
 t The ortho-acid, H 4 XO 5 , is also known. 
 
 Chromium, molybdenum, and tungsten further form a number of very 
 complicated salts derived from polyacids which are produced in a manner 
 analogous to the polysilicates (see -page 307). 
 
444 CHROMIUM GROUP; OCCURRENCE. 
 
 The most marked characteristic of these elements, and one they 
 share with the others having their position afc the middle of the long 
 periods, lies in the power which the individuals possess of forming 
 several series of compounds, in each of which they exhibit a differ- 
 ent valence ; so, for instance, molybdenum forms the chlorides 
 Mo C1 2 , Mo C1 3 , Mo C1 4 , and Mo Cl 6 , while tungsten exhibits W C1 2 , 
 WC1 4 , WC1 5 , and WC1 6 ; these two elements possess chlorides, 
 therefore, in which they are respectively quinquivalent and hexava- 
 lent; and, passing backward from these, we find a series of com- 
 pounds in which the valence successively diminishes by one until a 
 minimum is reached at divalence. It will be remembered that sul- 
 phur shows some resemblance to molybdenum and tungsten by 
 forming three chlorides of the formulae S 2 C1 2 , S C1 2 , and S C1 4 re- 
 spectively, but the latter is capable of existence only at very low 
 temperatures; it is not inconceivable that, were the proper condi- 
 tions attainable, a penta- and hexa-chloride of sulphur might also be 
 produced (see page 370). 
 
 None of the elements of this group occur uncombined as natural 
 minerals ; the principal compounds which are found are given in the 
 following table : 
 
 Chromium. Found as chromite (chromic iron) isomorphous with spinell, 
 formula Fe (CrO 2 ) 2 ;* when the chromium is replaced in part by aluminium 
 and by ferric iron, and the ferrous iron by magnesium, the mineral is called 
 chromspinell; chromite forms veins or imbedded masses in serpentine rock. 
 
 Crocoite. Lead chromate, Pb Cr 4 ,t is sometimes found. 
 
 Molybdenum. Found as the sulphide MoS 2 , called molybdenite; as the 
 molybdate of lead, Pb Mo O 4 , called wulfenite;J and as molybdite, Mo O 3 , 
 the anhydride of molybdic acid. 
 
 * This compound is ferrous chromite, derived from a hydroxide of the 
 formula CrO(OH), analogous to AIO(OH). The ferrous iron in chromite 
 can be replaced isomorphously by divalent chromium, the trivalent chromium 
 in CrO (OH) by ferric iron (see page 339). 
 
 t The chromates of the alkali metals are, without exception, isomorphous 
 with the corresponding sulphates ; naturally occurring lead chromate is, how- 
 ever, not isomorphous with anglesite (PbS0 4 ), but artificially prepared crys- 
 tals of lead chromate have proven to be isomorphous 'with the latter. The 
 isomorphism of the chromates and sulphates clearly demonstrates the family 
 connection between chromium and sulphur. 
 
 | Wulfenite is not isomorphous with crocoite, but it is sometimes found 
 with a contents of chromium replacing molybdenum. It is isomorphous with 
 the corresponding salt of tungstic acid. 
 
CHROMIUM GEOUP; METALLURGY. 445 
 
 Tungsten. As scheelite, Ca WO 4 ; reinite, Fe WO 4 ; and stoltzite, Pb WO 4 . 
 All of these minerals are isomorphous with wulf enite. Tungsten is also found 
 as tungstite, W0 3 , the anhydride of tungstic acid. 
 
 Uranium. As pitchblende, (UO 2 ,Pb) W 2 9 ; and as the sulphate of ura- 
 nium, which exists as an impure mineral sometimes called uranocher. 
 
 The elements under discussion are not of any commercial impor- 
 tance when isolated from their compounds. Chromium can be pre- 
 pared by electrolyzing the fused chloride, CrCl 3 , in a manner 
 analogous to the preparation of the alkali metals, the alkaline 
 earths, and of aluminium ; or the metal can be obtained by heating 
 the chloride with sodium or with zinc in the absence of the air, the 
 process being like that formerly used for obtaining aluminium (see 
 page 333). Sodium amalgam, when treated with chromic chloride, 
 forms sodium, chloride and liberates chromium, which latter element 
 then forms an amalgam with mercury. It can be separated from 
 this by distilling the mercury in a current of hydrogen. Tungsten, 
 molybdenum, and uranium can be prepared by reducing the oxides 
 of these metals by means of hydrogen at red heat, the elements in 
 question being much more easily separated from their oxides than is 
 chromium from its corresponding compounds. In the case of ura- 
 nium, the element can even be obtained by heating its oxides with 
 charcoal. The most important physical properties of these elements 
 are given in the following table : 
 
 Chromium, specific gravity 6.8, crystalline, of metallic appearance, in- 
 fusible. 
 
 Molybdenum, specific gravity 8.6, silver white, infusible. 
 
 Tungsten, specific gravity 18.1, steel-colored plates, fusible at a high tem- 
 perature. 
 
 Uranium, specific gravity 18.4, white, metallic lustre, fusible at a high 
 temperature. 
 
 Chromium, molybdenum, tungsten, and uranium all have small 
 atomic volumes. As was mentioned on page 367, they have their 
 places on the descending branches and hear the minimum of the 
 curves formed by using the atomic volumes as ordinates, and the 
 atomic weights as abscissae ; they are therefore infusible, or at least 
 fusible with difficulty, and they form colored salts. 
 
 Chromium is slowly oxidized when heated in the air, more 
 rapidly in a current of oxygen ; the oxide which is formed has the 
 formula Cr 2 3 . The metal is dissolved by hydrochloric acid, or by 
 hot sulphuric acid, the chloride, CrCl 3 , or the sulphate, Cr 2 (S0 4 ) 3 , 
 
446 CHROMIC OXIDE ; HYDROXIDE. 
 
 being produced, according to the acid used. Potassium nitrate and 
 potassium chlorate, when fused with chromium, give up their oxy- 
 gen and form potassium chrornate. 
 
 Molybdenum is slowly attacked when heated in the air. It is 
 readily converted into the trioxide, Mo 3 , by oxygen at a high 
 temperature. Nitric acid or aqua regia attacks the metal to form 
 molybdic acid. Tungsten behaves as does molybdenum, while 
 uranium even dissolves in dilute hydrochloric or sulphuric acid, 
 hydrogen being at the same time evolved. 
 
 The most important compounds of chromium are derived from 
 two oxides, chromic oxide, Cr 2 3 , which is mainly basic in its char- 
 acter, and chromium trioxide., Cr0 3 , which acts as an acidic anhy- 
 dride, and which is the oxide typical of the family. In addition to 
 these tw r o, there exists a chromous hydroxide, Cr (01? ) 2 (the oxide 
 corresponding to which is not known), and a chromous-chrornic oxide 
 of the formula Cr 3 4 . 
 
 Chromic oxide, Cr 2 3 , is a dark green powder which is insoluble 
 in acids when it has been heated to a high temperature (see page 
 338) ; after it has been subjected to such treatment, it can be 
 brought into solution by fusion with caustic alkalies, or with the 
 primary sulphate of potassium.* Chromic oxide dissolves in fused 
 glass, and imparts a fine green color to the substance ; for this rea- 
 son it is used as a green paint for tinting porcelain. 
 
 Chromic hydroxide can be precipitated from solutions of chromic 
 salts by the addition of ammonia water. When dry, it has the com- 
 position represented by the formula Cr ( OH ) 3 + 2 H 2 0. When 
 the latter is heated to 200, it changes to metachromic hydroxide, 
 CrO(OH). This substance corresponds to the similar aluminium 
 compound (see page 339). Chromic oxide and hydroxide are both 
 basic and acidic in their character. The freshly precipitated hy- 
 droxide is readily dissolved by caustic alkalies, forming deep green 
 solutions of the chromites' of the respective metals ; if the solutions 
 so formed are allowed to stand, or if they are boiled, the hydroxide 
 is once more precipitated, f A number of chromites occur as min- 
 
 * In this process potassium-chrome alum, K 2 SO 4 , Cr 2 (SO 4 ) 3 + 24 H 2 O, is 
 produced. 
 
 t Differing from aluminium hydroxide, the solution of which in alka- 
 lies is not decomposed by boiling. When chromic oxide is present, alkalies 
 are also able to dissolve considerable quantities of ferric oxide. 
 
CHROMIC SALTS. 447 
 
 erals ; the latter are derived from metachromic hydroxide and are 
 isomorphous with spinell ; these compounds can also be artificially 
 prepared by fusi-ng chromic oxide with the metallic oxide, which is 
 to be used as the base, boron trioxide, B 2 3 , being used as a flux.* 
 The chromic salts, in which chromic oxide acts as a base, are pro- 
 duced by dissolving the oxide or hydroxide in acids. The most 
 important of these are described below : 
 
 CHROMIC CHLORIDE, Cr Cl 3 ,t is produced by burning chromium in an atmos- 
 phere of chlorine, or by heating an intimate mixture of chromic oxide 
 and charcoal in a current of dry chlorine. The salt so produced sub- 
 limes in the form of pink plates ; the chloride which is formed by dis- 
 solving the hydroxide in water crystallizes in dark green needles of the 
 formula Cr C1 3 + 6 H 2 O ; anhydrous chromic chloride cannot be pre- 
 pared from this, as, upon heating, the salt loses hydrochloric acid and 
 changes into a basic chromic chloride. The dry chloride is nearly in- 
 soluble in water or in acids; when heated in air, it gives off chlorine 
 and leaves chromic oxide. The vapor density of gaseous chromic 
 chloride shows that the substance has a molecule corresponding to 
 the formula Cr C1 3 . 
 
 CHROMIC SULPHATE, Cr 2 (SO 4 ) 3 + 15 H 2 O, is formed by dissolving chromic 
 hydroxide in concentrated sulphuric acid, and then allowing the solu- 
 tion to absorb moisture from the air; the salt is reddish violet in color; 
 when heated to 100 it loses water of crystallization, and changes to a 
 green salt having the composition Cr 2 (SO 4 ) 3 + 5 H 2 O. When a solu- 
 tion of chromic sulphate is mixed with a solution of an alkaline sul- 
 phate and evaporated, an alum is formed in which chromium has 
 taken the place of aluminium ; an example of such a salt is K 2 SO 4 , 
 Cr 2 (SO 4 ) 3 + 24 H 2 O (see page 339). 
 
 When solutions of caustic alkalies are added to solutions of 
 chromic salts, a precipitate of chromic hydroxide is produced ; the 
 latter is soluble in an excess of the precipitating medium ; on the 
 other hand, chromic hydroxide does not dissolve in ammonia solu- 
 tion, and can, as a consequence, be precipitated from solutions of 
 chromic salts by the addition of that reagent, even in excess; t al- 
 
 * Zinc chromite, Zn(CrO 2 )o, and manganous chromite, Mn(CrO 2 ) 2 , 
 have been prepared in this way. 
 
 t The formula Cr 2 C1 6 was formerly assigned to this compound ; but the 
 latest determinations of the specific gravity of this substance, while in the 
 state of a vapor, show it to be Cr C1 3 (see page 335). Nilsson and Pettersson; 
 Comptes Rendus; 107, 529. 
 
 J The precipitation of chromic hydroxide is very much retarded, and may 
 be entirely prevented, by the presence of not-volatile organic acids such as 
 citric acid, tartaric acid, oxalic acid, etc. 
 
448 CHROMIC ACID. 
 
 kaline sulphide solutions or solutions of the carbonates precipitate 
 chromic hydroxide for reasons identical with those mentioned in 
 the chapter on aluminium (see page 341).* When in alkaline solu- 
 tion, chromic hydroxide is completely oxidized to a chromate by the 
 addition of chlorine or bromine ; the same change can also be brought 
 about by the addition of other oxidizing agents, f or by fusion with 
 potassium nitrate or chlorate. 
 
 Chromic acid, H 2 Cr0 4 , is not known; its anhydride, Cr0 3 , 
 chromium trioxide, is produced when the acid is liberated from its 
 salts ; this is best accomplished by adding tolerably concentrated 
 sulphuric acid to a solution of potassium or sodium dichromate : 
 
 K 2 Cr 2 7 + H 2 S0 4 = K 2 S0 4 + H 2 + 5 Cr 3 . 
 
 The anhydride crystallizes in beautiful carmine-red needles which 
 melt at 193, forming a dark red fluid which loses oxygen at 250, 
 and changes to green chromic oxide ; the oxide is readily soluble in 
 water, the solubilfty being diminished by the addition of sulphuric 
 acid.J 
 
 Chromium trioxide is a most energetic oxidizing agent ; even 
 dilute solutions instantly change sulphurous acid into sulphuric 
 acid: L 2Cr0 3 + 3H 2 S0 3 = Cr 2 3 +3H a S0 4 ; 
 2. Cr 2 3 + 3 H 2 S0 4 = O 2 (S0 4 ), + 3 H 2 0. 
 
 In the same way hydrogen sulphide is oxidized, while sulphur is 
 separated. Many organic substances are also readily attacked by 
 chromium trioxide, that substance being at the same time reduced 
 to chromic oxide (respectively to the chromic salts which would be 
 formed by the acids which may be present). 
 
 Although chromic acid is unknown, an acid chloride of chromium 
 called chromyl chloride, Cr 2 C1 2 , which may be considered as being 
 chromic acid in which the two hydroxyl groups have been replaced 
 
 * Chromium is completely separated from solutions of chromic salts by 
 the addition of freshly precipitated barium carbonate. The chromium sepa- 
 rates as chromic hydroxide mixed with basic chromic salt. 
 
 t For instance, lead superoxide, Pb O. 2 , when lead chromate is produced. 
 
 J Chromium trioxide is least soluble in sulphuric acid of about 85 per cent. 
 
 For instance, alcohol is oxidized to aldehyde and to acetic acid ; the re- 
 lationship between these three compounds can be seen from the following 
 structural formulae: 
 
 CH 3 - CH 2 OH + O = CH 3 - COH + H 2 O. CH 3 COH + O = CH 3 COOH. 
 Alcohol, Aldehyde, Acetic acid. 
 
CHROMYL CHLORIDE ; CHROMATES. 449 
 
 by chlorine, can be produced by adding concentrated sulphuric acid 
 to an intimate mixture of sodium chloride with potassium dichro- 
 mate. Chromyl chloride is a dark red fluid which boils at 118, and 
 which instantly decomposes into chromium trioxide and hydrochloric 
 acid on the addition of water : 
 
 Cr 2 C1 2 + H 2 = 2 H Cl + Cr 3 . 
 In structure, this compound is analogous to sulphuryl chloride : 
 
 
 Cl I Cl 
 
 Chromyl chloride and Sulphuryl chloride (see page 157.) 
 Salts of another acid chloride of chromium, in which only one 
 hydroxyl group in a formula weight has been replaced by chlorine, 
 are also described. These salts correspond to those of chlor-sul- 
 phonic acid (see page 157). 
 
 The chromates are derived from a hypothetical dibasic acid 
 analogous in formula to sulphuric acid : 
 
 fOH fOH 
 
 s Jo 
 
 S |0 
 
 tOH lOH 
 
 Sulphuric acid and Chromic acid; 
 
 and the dichromates form a dichromic acid, H 2 Cr 2 7 (also hypo- 
 thetical), which is analogous to disulphuric acid, H 2 S 2 O 7 (see page 
 154).* The most important chromates and dichromates are those 
 of potassium and of sodium. 
 
 POTASSIUM CHROMATE is a yellow, crystalline salt, which is readily soluble 
 in water, and which is isomorphous with potassium sulphate. Upon 
 addition of dilute acids it is converted into the dichromate : 
 
 2 K 2 Cr O 4 + 2 HNO 3 = K 2 Cr 2 O 7 + 2 KNO 3 + H 2 O.t 
 
 * Tri- and poly-chromates have been described (see page 307). A di-sul- 
 phuric acid in which a portion of the sulphur has been replaced by chromium 
 is also known. 
 
 t In this reaction the primary chromate of potassium, KHCr O 4 , may be 
 considered to be the first product : 
 
 K 2 Cr O 4 + HNO 3 = KHCr O 4 + KNO 3 . 
 
 Two formula weights of this primary salt would then separate water, leaving 
 the dichromate : - 2 KHCr ^ = ^ ^ ^ + ^ Q 
 
450 BICHROMATES ; LEAD CHIIOMATE. 
 
 On the other hand, potassium dichromate is changed to the chromate 
 by alkalies, thus : K 2 Cr. 2 O 7 + 2 KOH = 2 K 2 Cr O 4 + H 2 O. 
 POTASSIUM DICHROMATE crystallizes in red prisms or plates which are 
 soluble in water. It melts at a low red heat without decomposition, 
 and it loses oxygen and is converted into a mixture of chromic oxide 
 and potassium chromate at a bright white heat : 
 
 2 K 2 Cr 2 O 7 = 2 K 2 Cr O 4 + Cr 2 O 3 + 3 O. 
 
 It also liberates oxygen when it is heated with non-oxidizable mineral 
 acids (such as sulphuric acid). With sulphuric acid, chromic sulphate 
 (respectively potassium-chrome alum) is formed: 
 
 K 2 Cr 2 O 7 + 4 H 2 SO 4 = K 2 SO 4 , Cr 2 (SO 4 ) 8 + 4 H 2 O + 3 O.* 
 
 Of course, an acid solution of potassium dichromate can oxidize the 
 hydrogen compounds of the not-metals; these reactions are discussed 
 on pages 60 and 97. Potassium dichromate is extensively used for the 
 preparation of battery fluids ; however, of late years, the more soluble 
 sodium dichromate is taking the place of the potassium salt. 
 
 The reactions shown by sodium dichromate are identical with 
 those of potassium dichromate. 
 
 THE CHROMATE OF LEAD is insoluble in water, and is produced by adding 
 a soluble chromate or dichromate to the solution of a lead salt, as 
 follows : 
 
 1. K 2 Cr 2 O 4 + Pb (NO 3 ) 2 = 2 K NO 3 + Pb Cr O 4 . 
 
 2. K 2 Cr. 2 O 7 + 2 Pb (NO 3 ) 2 + H 2 O = 2 KXO 3 + 2 HNO 3 + 2 Pb Cr O 4 . 
 
 In the second reaction, therefore, free acid is produced; the same is 
 true in other cases where, by double decomposition, an insoluble 
 
 * In this reaction the oxygen present in the chromate, in excess of that 
 necessary to form chromic oxide, passes off. The same is true of the other 
 reactions in which potassium dichromate, or dichromates in general, are used 
 as oxidizing agents. One formula weight of potassium dichromate, therefore, 
 has three atoms of oxygen at its disposal for oxidizing purposes. In formulat- 
 ing reactions this fact is the essential one to be taken into consideration. One 
 formula weight of potassium dichromate will therefore oxidize three of sul- 
 phurous acid to sulphuric acid, or three molecules of alcohol to aldehyde, etc. 
 Concentrated hydrochloric acid is oxidized to chlorine by the dichromate; in 
 the latter case, of course, six molecules of hydrochloric acid are changed to 
 chlorine by the dichromate, for: 
 
 6HC1 + 30=3H 2 0+ 6C1; 
 
 however, the excess of hydrochloric acid will subsequently form potassium 
 chloride and chromic chloride with the bases present, so that the complete 
 reaction would be represented as follows : 
 
 K 2 Cr 2 O 7 + 14 HC1 = 2 KC1 + 2 Cr C1 3 +-G Cl + 7 H 2 O. 
 
CHROMOUS COMPOUNDS. 451 
 
 cliromate can be formed, for the precipitation takes place by the 
 addition of either a soluble cliromate or dichromate to the solution 
 containing the salt of the metal capable of forming such an insoluble 
 cliromate. Lead cliromate possesses a bright yellow color, which 
 makes it useful as a paint (chrome yellow*); addition of potassium 
 or sodium hydroxide to chrome yellow changes that substance into 
 an insoluble basic lead chroniate which is termed chrome red : 
 
 2 Pb Cr O 4 + 2 KOH = (Pb OH) 2 Cr O 4 + K 2 Cr O 4 . 
 
 BARIUM CHKOMATE, Ba Cr O 4 , is insoluble in water, but soluble in hydro- 
 chloric or nitric acid. In this way the salt differs from the equally 
 insoluble barium sulphate, for the latter is insoluble both in water 
 and in acids. The solubility of barium cliromate is not increased by 
 the presence of other salts in solution in the supernatent fluid, but the 
 solubility of the sulphate is increased. 
 
 The chromous compounds, which, are derived from chromous 
 hydroxide (CrOH) 2 , are of far less importance than either the 
 chromic salts or the chromates. 
 
 CHROMOUS CHLORIDE, CrCl 2 , is produced when metallic chromium is 
 heated in a current of dry hydrochloric acid, or when chromic chlo- 
 ride, Cr C1 3 , is reduced by means of a current of dry hydrogen. The 
 substance produces a blue solution in water ; the latter, however, 
 rapidly turns green, owing to the absorption of oxygen and the forma- 
 tion of a basic chromic chloride; addition of alkaline hydroxides to 
 this solution precipitates brownish-yellow chromous hydroxide. One 
 of the chief characteristics of all chromous compounds is the extreme 
 ease with which they take up oxygen in order to produce chromic 
 salts. A number of chromous salts of acids other than hydrochloric 
 acid have also been prepared. The specific gravity of the vapor 
 of chromous chloride shows that it has a molecule corresponding to 
 the formula CrCl 2 . (ISTilsson and Pettersson; Comptes Rend.; 107, 
 529.) 
 
 An oxide of chromium having the formula Cr 3 4 is also known. 
 This substance is regarded as being composed of chromous and 
 chromic oxides : - Cr + C r 2 3 = O 3 4 ; 
 
 so that, in this compound, chromous oxide plays the part of a base, 
 and chromic oxide that of an acidic anhydride. 
 
 The compounds of chromium which are used in the arts are pre- 
 pared from chromic iron. The latter substance is finely ground, 
 washed with water, and then intimately mixed with potash-lime ; * 
 
 * A mixture of potassium and calcium hydroxides produced by slaking 
 quick-lime with a potassium hydroxide solution; soda-lime is produced by 
 using sodium instead of potassium hydroxide. 
 
452 MOLYBDENUM; COMPOUNDS OF. 
 
 it is afterward dried at 150, and finally heated in reverberatory 
 furnaces, the oxygen so supplied changing the chromic compound 
 into potassium chroniate and calcium chromate ; * after heating for 
 a sufficient length of time, the mass is extracted with the least 
 quantity of boiling water, and the solution so produced is then 
 treated with potassium sulphate. By this means the calcium 
 chromate is converted into potassium chromate ; the latter salt is 
 finally changed into potassium dichromate by the addition of sul- 
 phuric acid. Potassium and sodium dichromates are used for the 
 preparation of various paints (chrome yellow, chrome red, etc.) ; in 
 a number of processes of dyeing; in the preparation of chromic 
 oxide, which is used for porcelain painting ; as oxidizing agents ; 
 and in a number of other ways. 
 
 The compounds of molybdenum, tungsten, and uranium are not, 
 by any means, so important as are those of chromium. 
 
 Molybdenum forms the following oxides : 
 
 1. Molybdenum monoxide, MoO; brown (nearly black) in color. 
 
 2. Molybdic oxide, Mo. 2 O 3 ; black in color. 
 
 3. Molybdenum dioxide, MoO 2 ; dark brown in color, t 
 
 4. Molybdenum trioxide, Mo O$ ; white in color. 
 
 The last of these is produced by roasting finely powdered molyb- 
 denite, by which means the sulphur is burned off, and impure yellow 
 molybdenum trioxide is left ; the latter is extracted with ammonia 
 water, which produces ammonium molybdate. The ammonium 
 molybdate is recrystallized, and is then converted into molybdic acid 
 by gently heating it : 
 
 ( NH 4 ) 2 Mo 4 = H 2 Mo 4 + 2 NH 3 . 
 
 The trioxide is produced by dehydrating the latter compound. The 
 oxide is difficultly soluble in water, is white and of crystalline 
 structure ; reducing agents change it to the second oxide, Mo 2 3 ; and 
 this, when gently heated, takes up oxygen and forms the dioxide, 
 Mo 2 ; the first oxide is formed by adding a concentrated solution 
 of potassium hydroxide to molybdenum dichloride, Mo C1 2 . Hydrox- 
 ides, Mo (OH ) 8 and Mo (OH ) 4 , corresponding to Mo 2 3 and M0 2 
 are also known. 
 
 * Of course, in the preparation of sodium dichromate, soda-lime is used, 
 t A blue oxide, Mo 3 O 8 , is also described. 
 
MOLYBDIC ACIDS. 453 
 
 Molybdenum forms the following chlorides : 
 
 Mo C1 5 , produced by passing dry chlorine over heated molybdenum. 
 Mo CU , produced by heating the trichloride in a current of carbon dioxide.* 
 Mo C\s , produced by heating the pentachloride in a current of hydrogen 
 at 250. 
 
 Mo C1 2 , produced by heating the trichloride in a current of carbon dioxide.* 
 
 Molybdenum pentachloride boils at 268, changing into a dark 
 brown gas which has a specific gravity of 9.4 at 350 ; this vapor 
 density corresponds to a molecule having the formula Mo C1 5 , so 
 that molybdenum affords an example of an element belonging to the 
 sulphur family which enters into the formation of a compound in 
 which the individual is pentavalent (see page 370). 
 
 Unlike chromium trioxide, the trioxide of molybdenum is capa- 
 ble of forming a number of hydrated acids which, in formula, corre- 
 spond to the hydrated sulphuric acids ; this relationship is made 
 plain by the following table : 
 
 MOLYBDIC ACIDS. SULPHURIC ACIDS (see page 145). 
 
 Mo O 3 4- H 2 O = H 2 Mo O 4 . SO 3 + H 2 O = H 2 SO 4 . 
 
 H 2 MoO 4 + H 2 O = H 4 Mo O 5 . H 2 SO 4 + H 2 O = H 4 SO 6 . 
 
 H 4 Mo0 5 + H~ O = H 6 Mo O 6 . H 4 SO 5 + H 2 O = H 6 SO 6 . 
 
 The acid having the formula H 4 Mo0 5 is formed by adding 
 nitric acid to a solution of sodium or potassium molybdate. It is 
 nearly insoluble in water, and when dried in vacua loses water while 
 changing to H 2 Mo 4 . H 6 Mo 6 is soluble in water, and is pro- 
 duced by dialysis of a solution of molybdic acid in a manner similar 
 to the separation of soluble silicic acid (see page 305). Salts of a 
 number of complicated molybdic acids are also known ; the formulae 
 of a few of these are given in the following table : 
 
 Na 2 Mo 2 O 7 , sodium dimolybdate.t 
 Na 2 Mo3O 10 , sodium trimolybdate. 
 ]STa 2 Mo 4 O 18 , sodium tetramolybdate. 
 ]STa 2 Mo 8 O 25 , sodium octomolybdate. 
 
 The method of formation of these polymolybdates is similar to that 
 of the polysilicates. 
 
 Acid solutions of molybdic acid are readily reduced by means of 
 metallic zinc or tin. During such a reduction the color of the solu- 
 
 * 2 Mo C1 3 = Mo C1 2 + Mo C1 4 . These chlorides must be produced in the 
 absence of free oxygen ; otherwise the oxy-chlorides are formed. 
 t Corresponding to the disulphate and dichromate. 
 
454 MOLYBDIC ACIDS. 
 
 tion at first becomes blue, then green, and finally black, at which 
 stage of the reaction the monoxide Mo is precipitated. 
 
 When nitric acid is added to a solution of ammonium molybdate, 
 ammonium tetramolybdate [(NH 4 ) 2 Mo 4 13 ] is produced. The 
 latter substance is of great importance in analytical chemistry, for 
 the reason that, upon addition of phosphoric acid or a soluble phos- 
 phate, the phosphoric acid is completely separated as a constituent 
 of a yellow precipitate known as ammonium phospho-molybdate. 
 The latter substance has the composition expressed by the formula 
 ( NH 4 ) 3 P0 4 , 11 Mo 3 . * When this salt is treated with aqua regia 
 the free phospho-niolybdic acid, H 3 P0 4 , 11 Mo0 3 , goes into solu- 
 tion. The power which molybdic acid possesses of uniting with 
 other acids to form complicated compounds is most important in 
 the case of phosphoric acid ; however, it is not confined to that 
 substance alone, for similar unions of molybdic acid with arsenic 
 and silicic acids are also known. 
 
 Molybdenum forms three sulphides, a disulphide, Mo S 2 , a 
 trisulphide, MoS 3 , and a tetrasulphide, MoS 4 . If, in the latter 
 compound, we regard sulphur as having a valence of two, then 
 molybdenum may possibly be octovalent. In speculating as regards 
 the valence of molybdenum in such a compound, we must always 
 bear in mind, however, that, as it cannot be vaporized, its molecular 
 weight is unknown. 
 
 The tetrasulphide of molybdenum can act as an acidic anhydride, 
 for it forms a potassium salt of the formula K 2 Mo S 5 . 
 
 Tungsten produces only two oxides, a dioxide, W0 2 , and a 
 trioxide, W0 3 , f the latter being the anhydride of tungstic acid ; 
 the element, however, is able to enter into four chlorides, namely : 
 
 Tungsten dichloride, W C1 2 . 
 Tungsten tetrachloride, W C1 4 . 
 Tungsten pentachloride, W C1 5 . 
 Tungsten hexachloride, W C1 6 . 
 
 * More complicated formulae have recently been assigned to ammonium 
 phospho-molybdate and the phospho-molybdic acid. They are ( NH 4 ) 8 PO 4 , 
 12 Mo O 3 + ( NH 4 ) 2 HPO 4 , 12 Mo O 3 + 8 H 2 O and (HN 4 ) 2 HPO 4 , 12 Mo O 3 + 
 29 H 2 O. 
 
 t An oxide, W 3 O 8 , corresponding to Mo 3 O 8 , is known. This oxide is 
 probably a combination of W O 2 acting as a base and W O 3 acting as an anhy- 
 dride: 
 
TUNGSTIC ACIDS. 
 
 455 
 
 The last compound is produced by heating tungsten in a cur- 
 rent of chlorine ; it boils at 346, and has a vapor density which 
 corresponds to the molecular weight expressed by the formula 
 W C1 6 ; so that tungsten must be hexavalent in the hexachloride. 
 
 Provided we regard the atoms of oxygen as always being diva- 
 lent, then the highest valence of the elements of the sulphur group, 
 when in combination with oxygen, is six ; and in one instance at 
 least, as is shown by the existence of W C1 6 , the valence toward 
 chlorine also reaches that number. It seems not improbable, there- 
 fore, that, were the proper conditions attainable, the remaining ele- 
 ments of this family would also be able to produce compounds 
 which, in each molecule, would contain six atoms of chlorine. We 
 should then have a series of chlorides, derived from members of the 
 first six families, which would exactly correspond to the oxides, two 
 chlorine atoms taking the place of one of oxygen. This will be 
 made clear from the following general formulae : 
 
 FAMILY. 
 
 1. 
 
 2. 
 
 3. 
 
 4. 
 
 5. 
 
 6. 
 
 Chlorides 
 
 RC1 
 
 RC1 2 
 
 RC1 3 
 
 RC1 4 
 
 RC1 5 
 
 RC^ 
 
 Oxides 
 
 R,0 
 
 RO 
 
 R 2 3 
 
 RO 2 
 
 R 2 5 
 
 R0 3 
 
 The oxides of the seventh family, K 2 7 , have as yet no corre- 
 sponding halide, but it seems not impossible that some of the 
 missing compounds will ultimately be discovered. In the first six 
 families, however, the highest valence of the elements toward oxy- 
 gen and toward chlorine is given by the number of the family to which 
 respectively each group of elements belongs. 
 
 Two tungstic acids, which correspond to H 2 S0 4 and H 4 S0 5 , 
 are known ; they are H 2 W0 4 and H 4 W0 5 . The first of these is 
 a yellow powder, which is produced by decomposing the aqueous 
 solution of an alkaline tungstate with an excess of hot acid : 
 
 Na 2 W0 4 
 
 H 2 W0 4 . 
 
 The second is produced by using cold instead of hot acid. A num- 
 ber of polytungstates which are similar to the polymolybdates are 
 also known. Tungstic acid has the same ability of uniting with 
 other acids to form complicated compounds as is possessed by molyb- 
 
456 URANIUM; OXIDES OF. 
 
 die acid. We are acquainted, for instance, with phosphotungstic 
 acid, arsenotungstic acid, silicotungstic acid, etc. Of these com- 
 pounds, perhaps the most important is silicotungstic acid, the sodium 
 salt of which is formed by boiling a polytungstate of sodium * with 
 precipitated silicic acid ; the latter dissolving to form sodium silico- 
 tungstate, having a formula Na 8 Si W i2 42 + 22 H 2 0. This salt 
 is extremely soluble in water, and its solution has a high specific 
 gravity. f The free acid is soluble in ether. 
 
 Uranium has the highest atomic weight, and hence the most 
 metallic nature of any of the elements under discussion. As a con- 
 sequence, its trioxide can act both as a base and as an acid. 
 
 The oxides of uranium correspond exactly to those of tungsten. 
 They are U0 2 , uranous oxide, U0 3 , uranic oxide, and U 3 8 , which 
 is considered to be uranous-uranic oxide. Only three chlorides of 
 uranium, namely, a trichloride, U C1 3 , a tetrachloride, U C1 4 , and a 
 pentachloride, U C1 5 , are known. 
 
 Uranous oxide, U 2 , is basic in its properties and forms salts of 
 the general formula U X 4 , where X represents the remainder of a 
 monobasic acid after the removal of hydrogen. The uranous salts 
 are colored green, and are easily oxidized to compounds derived from 
 uranic oxide. 
 
 When uranic oxide enters into combination with acids, it forms 
 basic salts in which the divalent radicle =U0 2 plays the part of a 
 divalent metal. This radicle is called uranyl, and its relationship 
 to its salts is similar to that of the univalent radicle stibionyl, 
 SbO , which was described on page 253. The resemblance 
 between the radicle uranyl and the atoms of divalent metals can 
 be seen by comparing the following formulae of uranyl and calcium 
 salts : 
 
 (U0 2 ) (N0 3 ) 2 , uranyl nitrate. Ca (NO 3 ) 2 , calcium nitrate. 
 
 (U0 2 ) S0 4 , uranyl sulphate. Ca S0 4 , calcium sulphate. 
 
 (U0 2 ) 3 (P0 4 ) 2 , uranyl phosphate. Ca 3 (P0 4 ) 2 , calcium phosphate. 
 
 * Sodium paratungstate, Na 2 W 12 O 41 . This salt is formed hy fusing to- 
 gether reinite (FeWO 4 ) and sodium carbonate. It finds extensive applica- 
 tion in the manufacture of a fireproof sizing for inflammable materials. 
 
 t The specific gravity is 3. when the solution is saturated at ordinary tem- 
 peratures. 
 
URANYL HYDROXIDE ; URANATES. 457 
 
 The hydroxide from which the uranyl salts are derived can be com- 
 pared to calcium hydroxide, the divalent group =U0 2 , taking the 
 place of one atom of calcium : 
 
 U0 2 (OH) 2 and Ca(OH) 2 
 
 Uranyl hydroxide and calcium hydroxide. 
 
 Uranyl hydroxide can, therefore, dissolve in acids (for example, 
 nitric acid), and it then forms uranyl nitrate, exactly as calcium 
 hydroxide can dissolve in the same reagent to form calcium 
 nitrate ; the two reactions may consequently be expressed as fol- 
 lows : 
 
 U0 2 (OH) 2 + 2 HM) 3 = U0 2 (N0 3 ) 2 + 2 H 2 ; 
 Ca (OH) 2 + 2HN0 3 = Ca (N0 3 ) 2 + 2H 2 0. 
 
 The uranyl salts are yellow, with a green fluorescence. 
 
 Uranyl hydroxide, in addition to being a base, is, however, also 
 an acid ; it dissolves in strong bases to form uranates ; and, by writ- 
 ing the formula in a manner slightly different from that given above, 
 it will be seen that uranyl hydroxide is also uranic acid, and there- 
 fore it corresponds to sulphuric acid : 
 
 H 2 S0 4 , sulphuric acid, and H 2 U0 4 , uranic acid. 
 
 The uranates, however, in formula resemble the disulphates and 
 dichromates, and not the sulphates and chromates. When a uranyl 
 salt is treated with a solution of a caustic alkali, the first change 
 would be the formation of uranyl hydroxide (uranic acid) : 
 
 U0 2 (N0 3 ) 2 + 2 KOH = U0 2 (OH) 2 
 this, however, reacts with the alkali to form a diuranate : 
 2 U0 2 (OH ) 2 + 2 KOH = K 2 U 2 7 + 3 H 2 0.* 
 
 The chief compounds discussed in the last chapter are given in 
 the following table : 
 
 * Sodium uranate, when fused with glass, imparts a yellow tint with green 
 fluorescence to the same. 
 
458 ELEMENTS OF CHKOMIUM FAMILY; TABLE OF. 
 
 CHLOBIDES. 
 
 OXIDES. 
 
 CrCLj 
 CrCl 3 
 
 MoCl 2 
 MoCl 3 
 MoCl 4 
 MoCl 6 
 
 WC1 2 
 
 
 (CfO> 
 
 Cr 2 3 f 
 
 MoO 
 
 Mo 2 O 3 * 
 MoO 2 
 
 
 
 
 
 
 WC1 4 
 
 WC1 5 
 WC1 6 
 
 UC1 4 
 UC1 5 
 
 WO 2 
 
 U0 2 * 
 
 
 Cr0 3 t 
 
 Mo0 3 t 
 
 W0 3 t 
 
 U0 3 J 
 
 
 
 
 
 
 Cr 3 4 
 
 Mo 3 8 
 
 W 3 8 
 
 U 3 8 
 
 The oxides on the last line are considered to be combinations of two other 
 oxides. 
 
 ACIDS. 
 
 DI- AND POLY-ACIDS. 
 
 CrO 8 
 
 Mo0 8 
 
 W0 3 
 
 TJ0 3 
 
 All of these elements also form salts 
 
 
 H 2 MoO 4 
 H 4 Mo O 5 
 
 H 2 W0 4 
 H 4 W0 5 
 
 H 2 U0 4 tt 
 H 4 U0 5 ** 
 
 derived from a di-acid having the gen- 
 eral formula H 3 X 2 O 7 , and they also 
 form salts derived from complicated 
 
 
 
 H 6 Mo0 6 ** 
 
 
 
 
 poly-acids; the latter are formed by 
 
 
 
 
 
 
 uniting 3, 4, 5, etc., formula weights of 
 
 
 
 
 
 the acids H 2 XO 4 , and then separating 
 
 
 
 
 
 water until a dibasic acid is left. 
 
 The above acids are all dibasic (see page 140). 
 
 * Basic. t Acidic. J Basic and acidic. 
 
 The solution of Mo O 2 reddens litmus paper, but possesses no other acid 
 properties. 
 
 ** Existence doubtful. 
 
 tt Also acts as a basic hydroxide, uranyl hydroxide, U0 2 (OH) 2 . 
 
MANGANESE ; OCCURRENCE. 459 
 
 CHAPTER LVIIL 
 
 THE ELEMENT FORMING THE PRIMARY GROUP OF THE 
 SEVENTH FAMILY. 
 
 I 
 
 Manganese ; symbol, Mn ; atomic weight, 55. 
 
 ONLY one element which should undoubtedly have its place in 
 the primary group of the seventh family has, as yet, been discov- 
 ered ; and that element is manganese. Manganese, having its place 
 at the middle of one of the long periods, must necessarily differ 
 very markedly from the typical elements of the family (i.e., from the 
 halogens) ; and, indeed, a great variation from the properties of the 
 latter is to be expected even without any such consideration, for 
 manganese is metallic in its nature, while fluorine, chlorine, bromine, 
 and iodine are the most negative of all elements. As a conse- 
 quence, we should expect the greatest resemblance between the halo- 
 gens and manganese to lie in the derivatives of the highest oxides. 
 In these compounds, the metallic nature of manganese is almost 
 entirely overshadowed by the negative elements with which it is 
 combined, so that the permanganates, R Mn 4 ,* in many respects 
 (such as isomorphism, solubility, etc.), are very much like the per- 
 chlorates. The lower oxides of manganese, on the other hand, 
 bear no resemblance to the halogen oxides ; in fact, their nearest 
 prototypes are to be found among the oxides of iron, chromium, 
 cobalt, nickel, or lead, while in many respects Mn acts very much 
 like the oxides of calcium, magnesium, or zinc. The typical oxide 
 of the seventh family, therefore, is X 2 7 ; in no case, excepting 
 that of manganese, has it been isolated ; it is, however, known in 
 its derivatives (permanganates, perchlorates, periodates, and the 
 corresponding acids). 
 
 Manganese is never found as the uncombined element. The 
 chief minerals in which it occurs are given in the following 
 table : - 
 
 * R represents a univalent metal. 
 
460 MANGANESE; PROPERTIES. 
 
 Braunite, Mn 2 3 . Pyrolusite (polianite), Mn 2 . 
 
 Hausmannite, Mn 3 4 . Manganite, MiiO (OH).* 
 
 Pyrolusite and inanganite are the most important ores of man- 
 ganese. They both occur in large beds and in veins. Manganous 
 oxide, Mn 0, is also sometimes found as a mineral termed manga- 
 nosite. The carbonate, rhodochrosite, Mn C0 3 , belongs to the calcite 
 group (see page 416), while manganocalcite, (Mn, Ca) C0 3 , probably 
 is isomorphous with arragonite. 
 
 The element itself is very difficult to separate from its ores ; for 
 such powerful reducing agents as red-hot charcoal or hydrogen are 
 able to change the higher oxides into manganous oxide only, but 
 not into manganese.! Manganese can, however, be isolated either 
 by heating manganous chloride with sodium, or by electrolysis of 
 the fused chloride or fluoride. 
 
 Manganese is a grayish-white metal, which somewhat resembles 
 cast iron ; it is crystalline in structure and brittle, although it pos- 
 sesses a certain amount of toughness. The specific gravity is about 
 8, and its atomic value 6.9. Manganese is, therefore, at the mini- 
 mum of one of the curves of atomic volumes ; the element with next 
 smaller atomic weight has a larger atomic volume, and, as a conse- 
 quence, manganese is difficult to fuse, and forms colored salts. The 
 melting point of manganese lies at about 1900 ; this is probably 
 somewhat higher than that of iron. Pure manganese, after polishing, 
 rapidly becomes dull when exposed to the air, owing to oxidation. 
 The metal is energetically attacked by acids. $ Pure manganese 
 has no technical application ; an alloy of manganese and iron 
 (f erro-manganese, spiegeleisen) is, however, of the greatest commercial 
 importance for the manufacture of Bessemer steel. 
 
 Manganese forms the following oxides : 
 
 Mn O, manganous oxide. 
 Mn 2 O 3 , manganic oxide. 
 Mn 3 O 4 , manganous-manganic oxide. 
 
 MnO 2 , manganese dioxide (manganese hyperoxide, black oxide of man- 
 ganese). 
 
 Mn 2 O 7 , permanganic anhydride. 
 
 * Corresponding to Al O (OH). Trivalent manganese can replace alu- 
 minium and chromium isomorphously in the spinells. 
 
 t The conversion of the oxides into metallic manganese by means of char- 
 coal takes place only at a high white heat. 
 
 } Impure manganese decomposes even water very readily. 
 
MANGANOUS COMPOUNDS. 461 
 
 Manganous oxide is basic in its character ; it readily dissolves in 
 acids to form the manganous salts ; this oxide can be produced 
 through reduction of any one of the higher oxides by heating 
 in a current of hydrogen. It is green, or grayish-green, in color, 
 and when, exposed to the air, it readily absorbs oxygen to form 
 Mn 3 4 .* Manganous hydroxide, Mn (OH ) 2 , is separated as a 
 white precipitate when alkaline solutions or ammonia water are 
 added to a solution containing a manganous salt ; precipitation b}^ 
 means of ammonia is, however, entirely prevented by the presence 
 of ammonium salts, for manganous salts have a tendency to form 
 double salts with the compounds of ammonium, identical with that 
 displayed by the similar compounds of zinc or of magnesium (see 
 pages 419 and 438). When exposed to the air, manganous hy- 
 droxide rapidly turns brown, because it absorbs oxygen and is con- 
 verted into manganous-manganic oxide, Mn 3 4 ; when dissolved in 
 acids, the corresponding manganous salts are produced. The latter 
 do not. spontaneously oxidize when exposed to the air. 
 
 MANGANOUS CHLORIDE is contained in the colorless solutions obtained by 
 dissolving any one of the oxides of manganese in hydrochloric acid.t 
 When slowly evaporated the solutions deposit pinkish colored tablets 
 of the formula Mn C1 2 + 4 H 2 O. The anhydrous salt cannot be ob- 
 tained from these by heating, because the chloride, at a temperature 
 high enough to drive off water of crystallization, loses chlorine, ab- 
 sorbs oxygen, and in part changes into Mn 3 O 4 . In order to obtain 
 the chloride in an anhydrous condition, the water of crystallization 
 must be driven off in a current of dry hydrochloric acid gas. Man- 
 ganous chloride is extremely soluble in water, and shows the greatest 
 tendency to form double chlorides with the corresponding salts of 
 other metals. 
 
 MANGANOUS SULPHATE is produced by dissolving any one of the oxides of 
 manganese in hot sulphuric acid,t or better, by dissolving the carbon- 
 ate in the diluted acid. Upon evaporating and cooling to below 6, 
 crystals having the formula Mn SO 4 + 7 H 2 O separate. The latter 
 are isomorphous with the vitriols (see page 417). If the temperature 
 of crystallization is between 7 and 20, then the crystals contain but 
 
 * This action may become so violent as to cause the whole mass to glow. 
 
 t Should an oxide containing more oxygen than Mn O be dissolved, the 
 excess of oxygen will oxidize the hydrochloric acid, liberating chlorine (see 
 page 58). 
 
 \ When an oxide of manganese containing more oxygen than Mn O is 
 dissolved, then the excess of oxygen passes off. 
 
462 MANGANIC COMPOUNDS. 
 
 five molecules of water, and are then isomorphous with ordinary sul- 
 phate of copper (blue vitriol; see page 403). The sulphate is readily 
 soluble in water. 
 
 MANGANOUS CARBONATE is insoluble in water, and is therefore precipi- 
 tated from solutions of manganous salts by the addition of a soluble 
 carbonate. The naturally occurring salt is isomorphous with calcite. 
 
 MANGANOUS SULPHIDE is insoluble in water, but soluble in dilute acids * 
 (see page 100) ; it is therefore precipitated from solutions containing 
 manganous salts by the addition of an alkaline sulphide solution : 
 
 Mn SO 4 + (KH 4 ) 2 S = (NH 4 ) 2 SO 4 + Mn S. 
 
 Manganous sulphide is a flesh-colored precipitate which readily ab- 
 sorbs oxygen from the air, while, at the same time, it turns of a 
 brown color. 
 
 Manganic oxide, Mn 2 3 , occurs in nature as the mineral brau- 
 nite, which is the hardest ore of manganese. In the laboratory it 
 may be produced by heating manganous oxide, manganese dioxide, 
 or manganous-manganic oxide to red heat in a current of hydrogen. 
 The oxide is black in color, insoluble in water, and, when heated to 
 a white heat, changes to Mn 3 4 . The hydroxide Mn 2 H, corre- 
 sponding to A10 2 H, occurs as the mineral manganite ; this com- 
 pound can also be formed by slow oxidation of manganous hydroxide, 
 Mn (OH ) 2 , in the air, but, if the manganous hydroxide is covered 
 with ammonia solution, then the product of oxidation is the normal 
 manganic hydroxide, Mn (OH ) 3 . Both the oxide and hydroxides 
 are weakly basic in character. The salts derived from them are 
 unstable, and resemble those derived from the oxide of aluminium, 
 A1 2 3 . They are decomposed by the addition of an excess of water. 
 Their solutions are dark brown in color, and on addition of alkalies 
 precipitate manganic hydroxide. 
 
 MANGANIC CHLORIDE, Mn C1 3 , is produced by dissolving manganic hydrox- 
 ide in cold hydrochloric acid. The solution has a dark-brown color, 
 and, on standing, liberates chlorine, leaving manganous chloride. 
 
 MANGANIC SULPHATE, Mn 2 (SO 4 ) 3 , is a dark-green, amorphous powder, 
 which is produced by heating finely divided manganese dioxide with 
 concentrated sulphuric acid to 110. A portion of the oxygen of the 
 dioxide then passes off, while manganese sulphate remains. The re- 
 action may be considered as taking place in two stages : 
 
 1. 2MnO 2 = Mn 2 O 3 + O; 
 
 2. Mn 2 3 
 
 * Even in acetic acid (difference from zinc sulphide). 
 
MANGANESE DIOXIDE. 463 
 
 If the heating be carried too far, more oxygen will be evolved, and 
 manganous sulphate, Mn SO 4 , will remain. Manganic sulphate is in- 
 teresting because it forms compounds with the sulphates of the alkali 
 metals, which are isomorphous with the alums; this fact illustrates 
 the close resemblance between trivalent manganese, aluminium, 
 chromium, and ferric iron. Potassium-manganic sulphate has the 
 formula K 2 SO 4 , Mn 2 (SO 4 ) 3 + 24 H 2 O (see page 339). 
 
 Manganous-manganic oxide, Mn 3 4 , occurs as a brownish-black, 
 crystalline mineral known as hausmannite ; it is produced by heat- 
 ing any of the other oxides of manganese, when in contact with 
 the air, to a red heat.* Manganous-manganic oxide is considered 
 as being a manganous salt of or^Ao-manganous acid (the hydroxide 
 of manganese dioxide, Mn (OH ) 4 , being designated as manganous 
 acid).-\ This theory is expressed by the following formula: 
 
 o 
 
 n = n 
 
 Mn 
 
 Mn(OH) 4 -f 2MnO = Mn 3 4 + 2H 2 0; 
 
 Mn ; + 2MnO = Mn^ + 2H 2 0; 
 
 [OH [o[ 
 
 and is borne out by the fact that manganoms-manganic oxide, when 
 treated with dilute nitric or sulphuric acid, forms manganous nitrate 
 or sulphate, while manganese dioxide is left behind : 
 
 1. Mn 4 Mn 2 + 4 HN0 3 = Mn (OH ) 4 + 2 Mn ( N0 3 ) 2 . 
 
 2. Mn(OH) 4 = 
 
 This reaction is similar to that encountered with Pb 3 4 (see page 
 323). 
 
 Manganese dioxide, Mn 2 , is probably the most important com- 
 pound of manganese. It occurs in large quantities as a mineral 
 which is named pyrolusite. The latter has a steel-gray color, 
 metallic lustre, and crystallizes in prisms belonging to the rhombic 
 system. A considerable amount of this oxide is mined in the New 
 England States, and in California. Manganese dioxide can be pre- 
 
 * The oxide Mn 2 O 3 is stable when in an atmosphere of oxygen, if the 
 temperature is no higher than that of a Bunsen burner. At white heat it is 
 also converted into Mn 3 O 4 . 
 
 t Manganous acid would thus be parallel with sulphurous acid : SO 2 , 
 sulphur dioxide; H 2 SO 3 , sulphurous acid; H 4 SO 4 , orthosulphurous acid. 
 Mn O 2 , manganese dioxide ; H 2 Mn Os , manganous acid ; H 4 Mn O 4 , ortho- 
 manganous acid. 
 
464 MANGANESE DIOXIDE; REACTIONS. 
 
 pared artificially by oxidizing manganous carbonate with an al- 
 kaline solution of chlorine (see pages 122 and 123). 
 
 Manganese dioxide, when heated to a high red heat, loses one- 
 third of its oxygen, and changes into nianganous-manganic oxide 
 (see page 19) : - 
 
 During this decomposition, the dioxide first changes into manganic 
 oxide, Mn 2 3 , and then the latter compound loses the quantity 
 of oxygen necessary to produce Mn 3 4 , as the temperature is in- 
 creased to a high red heat. Acids decompose manganese dioxide. 
 When acting in the cold they not infrequently produce manganic 
 salts, while the surplus of oxygen is liberated (see page 462) ; on 
 the other hand, hot acids leave manganous salts behind. Of course, 
 if any oxidizable substance is present, the liberated oxygen does not 
 pass off as such, but is used up in the work of oxidation. The 
 reaction, when warm sulphuric acid is brought in contact with man- 
 ganese dioxide, is as follows : 
 
 Mn 2 + H 2 S0 4 = Mn S0 4 + H 2 + 0, 
 
 but, on the other hand, hydrochloric acid, because it is readily 
 oxidized, liberates chlorine : 
 
 Mn 2 + 4 H Cl = Mn C1 2 + 2 H 2 -f 2 Cl. 
 
 In the latter case it is not at all improbable that Mn C1 4 is at first 
 formed, and that the latter salt subsequently breaks down into man- 
 ganous chloride and chlorine (see pages 60 and 61).* The reactions 
 of manganese dioxide are very similar to those of the corresponding 
 compound of lead (see page 323). Several hydroxides related to 
 
 * Recent investigations render it probable that the chloride Mn C1 4 , when 
 formed, assumes the part of an acidic anhydride, and, uniting with the excess 
 of hydrochloric acid which is present, forms an acid of the formula H 2 Mn C1 6 , 
 analogous to H Si F 6 , and to other similar compounds which we have encoun- 
 tered (see pages 303, 316 and 330). The compound H 2 MnCl 6 then breaks 
 
 down as follows : 
 
 H 2 Mn C1 6 = 2 H Cl + Mn C1 2 + 2 Cl. 
 
 It is, however, very certain that the compound is not alone present in the 
 beginning of the reaction between hydrochloric acid and manganese dioxide, 
 for manganese trichloride, Mn C1 3 , is also produced : 
 
 MnO 2 + 4 H Cl = Mn C1 3 + Cl + 2 H 2 O. 
 The whole matter may therefore be regarded as not as yet definitely settled. 
 
MANGANATES. 465 
 
 manganese dioxide are known. The simplest of these is Mn (OH ) 2 . 
 These hydroxides have acidic properties,* and form salts which are 
 designated as manganites. However, none of the latter are derived 
 from the simple ortho- or meta-hydroxides Mn (OH ) 4 , or 
 Mn (OH ) 2 but, like the salts of so many acids which we have 
 already studied, they are produced by complicated poly-manganous 
 acids. Two examples of the latter are H 2 Mn 2 5 and H 2 Mn 5 O n ; 
 their formation might be imagined as taking place as follows : 
 
 1. 2MnO(OH) 2 = H 2 Mn 2 5 + H 2 0. 
 
 2. 5 Mn (OH ) 2 = H 2 Mn 5 O n + 4 H 2 0. 
 
 POTASSIUM PENTAMANGANITE, K 2 Mng O n , is formed by passing carbon 
 dioxide into a solution of potassium manganate. 
 
 CALCIUM DIMANGANITE, Ca Mn 2 O 5 , is important because, when treated 
 with hydrochloric acid, it generates chlorine : 
 
 Ca O, 2 Mn O 2 + 10 H Cl = Ca C1 2 + 2 Mn C1 2 + 5 H 2 O + 4 Cl. This 
 salt is readily produced by heating a mixture of manganous hydroxide 
 and calcium hydroxide in a current of air, so that a method t of util- 
 izing the waste manganous chloride, which was formerly lost during 
 the commercial preparation of chlorine, has been founded on this 
 reaction. 
 
 When manganese dioxide is fused with a caustic alkali in the 
 presence of an oxidizing agent, or even in the air, a manganate is 
 produced. The manganates, in chemical composition, are analogous 
 to the sulphates, chromates, and molybdates, etc. This will be seen 
 from the following formulae : 
 
 K 2 Mn0 4 , potassium manganate. 
 K 2 CrO 4 , potassium chromate. 
 K 2 Mo 4 , potassium molybdate. 
 
 However, those manganates which are soluble in water differ very 
 markedly from their prototypes in one particular; they are very 
 readily decomposed by the addition of an excess of the solvent and 
 are stable only in alkaline solution. In so decomposing, they change 
 into permanganates and manganese dioxide. The manganates, when 
 
 * The hydroxide of manganese dioxide, MnO(OH) 2 , is sufficiently acid 
 in its properties to redden blue litmus paper, and to expel carbonic acid from 
 the carbonates of the alkalies. Several of the hydroxides are found as min- 
 erals; they are termed " wad." 
 
 t Weldon's process. 
 
466 PERMANGANIC ACID. 
 
 dry, have a deep red color, with a metallic lustre. They are power- 
 ful oxidizers, and the salts of the alkali metals only are soluble in 
 water. The solutions are green, but change to red on addition of 
 an acid.* Neither manganic acid, H 2 Mn0 4 , nor its anhydride, 
 Mn 8 , has been isolated. 
 
 The anhydride of permanganic acid, Mn 2 7 , is the only one of 
 'the oxides, which is the anhydride of an acid of the general formula 
 H X 4 (and which consequently is related to perchloric and per-iodic 
 acids), which has been isolated. It is a dark-green (almost black), 
 oily liquid, produced by adding potassium permanganate, in small 
 quantities, to concentrated sulphuric acid. The liquid must be 
 cooled by means of a mixture of snow and salt during the process, 
 and, after the operation is completed, warmed to ^60, when the 
 anhydride distils. The highest oxide of manganese is extremely 
 unstable ; if allowed to stand, it spontaneously liberates oxygen 
 and leaves manganic oxide, Mn 2 3 ; it is a most powerful oxidizer ; 
 paper and alcohol are instantly ignited by it. When added to water, 
 it forms permanganic acid : 
 
 Mn 2 7 + H 2 = 2 H Mn 4 . 
 
 Permanganic acid is entirely analogous to perchloric acid. It 
 can be produced by decomposing a solution of barium permanganate 
 with exactly the requisite amount of sulphuric acid. By means of 
 the ensuing double decomposition, insoluble barium sulphate and 
 permanganic acid are produced : 
 
 Ba (Mn 4 ) 2 + H 2 S0 4 = Ba S0 4 + 2 H Mn 4 ; 
 
 the red solution so formed is then evaporated to dryness, when per- 
 manganic acid remains in the form of a reddish-brown, crystalline 
 substance. Permanganic acid is quite unstable ; it breaks down 
 when exposed to the light, and its solutions, like those of perchloric 
 acid, are powerful oxidizers. 
 
 * Due to the formation of a permanganate : 
 
 5 K 2 Mn O 4 + 4H 2 SO 4 = 4 K Mn O 4 + Mn SO 4 + 3 K 2 SO 4 + 4 H 2 O. 
 When water, and not acid, is added to the manganate, a hydroxide derived 
 from Mn O 2 is formed : 
 
 3 K 2 Mn O 4 + 3 H 2 O = 2 K MnO 4 + Mn O (OH) 2 + 4 KOH. 
 The potassium hydroxide will then react with the hydroxide of manganese to 
 form a manganite. Very weak acids (such as carbonic acid) facilitate the 
 change. 
 
PERMANGANATE OF POTASSIUM. 467 
 
 The most important permanganate is the permanganate of potas- 
 sium. This salt is produced by fusing manganese dioxide with a 
 mixture of potassium hydroxide and an oxidizing salt (such as 
 potassium nitrate or potassium chlorate *) ; the dark-green flux 
 then contains potassium manganate, which is converted into the 
 permanganate by dissolving in water, and then passing carbon 
 dioxide into the solution, f Potassium permanganate crystallizes in 
 long prisms, belonging to the monoclinic system; the crystals are 
 dark green, almost black, in color; their solution in water has an 
 intense reddish-purple color ; the salt is a most powerful oxidizing 
 agent. In oxidizing with potassium permanganate, there is an 
 essential difference between the action of the salt in acid or in 
 alkaline solution ; in the former event the permanganate changes 
 to a manganous salt; in the latter to manganese dioxide, subse- 
 quently to a manganite. 
 
 1. Acid solution. 
 
 2KMn0 4 +3H 2 S0 4 ==K 2 S0 4 + 2MnS0 4 + 50+3 H 2 0. 
 2. Alkaline solution. 
 
 a. 2 KMnO^rf 2 KOH = 2 K 2 Mn 0/+ H 2 + ; 
 
 b. 2K 2 Mn0 4 <f 2 H 2 0=4 KOH + 2 Mn0 2 + 2 0. 
 
 The first change, in alkaline solution, is, therefore, from the per- 
 manganate to the manganate, the second from the manganate to 
 manganese dioxide; of course, the latter substance subsequently 
 produces a manganite with the excess of caustic potash which is 
 present. From these equations it follows that, for every two for- 
 mula-weights of potassium permanganate in acid solution, there 
 are five atoms of oxygen to be used in oxidation, while for every 
 two formula-weights in alkaline solution, there are but three. The 
 following two equations illustrate the application of these rules : 
 
 2 K Mn 4 + 3 H 2 S0 4 + 5 H 2 S0 3 = K 2 S0 4 + 2 Mn S0 4 + 
 
 5 ff 2 S0 4 + 3 H 2 0. 
 2 KMn0 4 + 8 Na 2 S0 3 + H 2 = 2 KOH + 2 Mn0 2 + 
 
 * The oxygen of the atmosphere is also able to effect the change. 
 
 t The reaction is as follows : 
 
 3 K 2 Mn O 4 + 2 CO 2 = 2 K 2 CO 3 + 2 K Mn O 4 + Mn O 2 ; 
 
 the reactions taking place on the addition of water or of sulphuric acid are 
 given on page 466. foot-note. 
 
468 MANGANESE; COMPOUNDS OF. 
 
 Owing to its oxidizing powers, potassium permanganate is frequently 
 used as a disinfecting agent. 
 
 Manganese, with its diversity of compounds in which it displays 
 such a great difference in valence, serves admirably to illustrate the 
 fact that the chemical behavior of an element is, to a large extent, 
 relative, and that the properties of its compounds depend just as 
 much upon the elements with which it is united, and upon the manner 
 of such union, as they do upon the individual characteristics of the 
 element itself. Manganese, when entering into the formation of 
 manganous salts, is far more like magnesium or zinc than it is like 
 manganese in the inanganates or permanganates ; when manganic 
 oxide, as a base, has united with acids to form manganic salts, then 
 manganese resembles trivalent chromium, iron, or aluminium, even 
 to such an extent that its sulphate will form alums ; in the man- 
 ganates the element may be compared to sulphur in the sulphates ; 
 while, lastly, the permanganates are analogous to the perchlorates 
 or per-iodates. One essential distinction, however, remains as exist- 
 ing between the various compounds of manganese and the elements 
 with which it has been compared : the compounds of manganese can 
 be converted the one into the other, while, of course, those of zinc, 
 for example, can never be changed into those of chromium, or those 
 of iron into those of chlorine. Oxidizing agents transform manga- 
 nous salts into manganic salts, manganese dioxide, manganates, and 
 permanganates successively ; while, on the other hand, reducing 
 agents, beginning with the permanganates, produce exactly the 
 opposite result. 
 
 The relationship between the compounds of manganese and those 
 of a few other elements may be seen from the following table : 
 
 Manganous compounds ( Sulphate, Mn SO 4 * 7 H 2 O ) ( Zn SO 4 + 7 H, O 
 
 (Oxide Mn<m \ Nitrate, Mn (NO 3 ) 2 > l*e < Mg(NO,) a 
 
 ( Chloride, Mn C1 2 ) ( CaCl 2 
 
 Manganic compounds ( Sulphate, Mn a (SO 4 ) 3 
 
 (Oxide, Mn 2 O,) ) Alum, Mn 3 (SO 4 ) 3 K 2 SO 4 , 24 H 2 O 
 
 {. 
 
 Manganates ( Potassium manganate, K 2 Mn O 4 | like i K 2 SO 4 or K 2 Cr O 4 
 (Acid, H 2 MnO 4 ) ( Barium manganate, Ba Mn O 4 f \ BaSO 4 or BaCrO 4 
 
 Permanganates ( Potassium permanganate, K Mn O 4 | j.. ( K Cl O 4 
 (Acid, HMnO 4 ) ( Barium permanganate, Ba(MnO 4 ) 2 j ( Ba(ClO 4 ) 2 
 
 The oxides of manganese are, perhaps, most like those of lead ; 
 but, in formula, they also resemble those of the type of the family, 
 chlorine. 
 
 ... 
 llke 
 
MANGANESE; COMPOUNDS OF. 469 
 
 OXIDES OF CHLORINE. LEAD. MANGANESE. 
 
 C1 2 
 
 Cl Pb O Mn O 
 
 C1 2 3 Pb 2 8 Mn 2 3 
 
 C10 2 Pb0 2 Mn0 2 
 
 (C1 2 7 ) 'Mn 2 7 
 
 Pb 3 4 Mn 3 4 
 
 This process of comparison, were space to permit, could be car- 
 ried much farther ; and, indeed, the formation of tables like the 
 above would be a most instructive exercise for the pupil. 
 
470 IRON ; COBALT ; NICKEL. 
 
 CHAPTER LIX. 
 
 IRON, COBALT, AND NICKEL. 
 
 Iron ; symbol, Fe ; atomic weight, 56 ; 
 Cobalt ; symbol, Co ; atomic weight, 59 ; 
 Nickel ; symbol, Ni ; atomic weight, 58.7. 
 
 IRON, cobalt, and nickel are members of the eighth family of 
 elements, which consists of three groups, each of which contains 
 three individuals, namely : 
 
 1. Iron, cobalt, nickel. 
 
 2. Euthenium, rhodium, palladium. 
 
 3. Osmium, iridium, platinum. 
 
 The properties of these elements are such that they form a 
 gradual transition from the last elements of the first halves of the 
 long periods, to the first ones in the second halves ; so that we should 
 expect iron to be very much like manganese, and nickel to bear 
 marked resemblance to copper ; and such is the case. The valence 
 of the elements in their highest oxides, passing from manganese 
 (through iron, cobalt, and nickel) to copper, diminishes with each 
 individual as the atomic weight increases ; manganese, in the per- 
 manganates, has a valence of seven ; iron, in ferric acid, a valence 
 of six ; cobalt a valence of three in the oxide Co 2 3 ; * nickel, almost 
 without exception, forms compounds derived from Ni ; while, 
 lastly, copper can appear as a univalent metal in its cuprous form. 
 
 Mn. Fe. Co. Ni. Cu. 
 
 Highest valence, VII VI III II t I ( II ) 
 
 Oxides, Mn 2 7 Fe 2 6 Co 2 3 Ni 2 2 Cu 2 0. 
 
 Iron, cobalt, and nickel are near the minimum of the curve of 
 
 * An oxide of cobalt, Co O 2 , probably also exists. 
 
 t Nickel can form an oxide Ni 2 O 3 ; but the latter forms no salts, and is 
 decomposed with the greatest of ease. 
 
IRON ; COBALT ; NICKEL. 
 
 471 
 
 atomic volumes formed in the period of which they are members, 
 while the elements which follow in the same period show a rapid 
 increase in their atomic volumes as we pass along the series in the 
 direction of increasing atomic weights ; the three individuals in 
 question are therefore malleable and ductile, have high melting 
 points,* and form colored salts. The physical constants mentioned 
 in this connection are given in the following table : 
 
 
 ATOMIC WEIGHT. 
 
 SPECIFIC GEAVITY. 
 
 ATOMIC VOLUME. 
 
 MELTING POINT. 
 
 Iron 
 
 56. 
 
 7.8 
 
 7.2 
 
 1770 (?) 
 
 Cobalt 
 
 59. 
 
 8.5 
 
 6.9 
 
 1750 
 
 Nickel 
 
 58.7 
 
 8.8 
 
 6.7 
 
 1570 
 
 Nickel, owing to its chemical reactions, specific gravity, atomic 
 volume, and melting point, apparently has its position in the periodic 
 system immediately following that of cobalt, although its atomic 
 weight is somewhat less than that of the latter element ; it seems 
 probable, therefore, that, at some future time, more exact study will 
 prove the atomic weights of the two elements in question to have 
 been inaccurately determined ; f if this should not prove to be the 
 case, however, then cobalt and nickel certainly form a most remark- 
 able exception to Meiidelejeff's rule, in the arrangement of that 
 system. 
 
 The principal minerals in which iron, cobalt, and nickel occur 
 are as follows : 
 
 * The melting points are, however, lower than those of the elements im- 
 mediately preceding which have diminishing atomic volumes with increasing 
 atomic weights. 
 
 t Gerhardt Kriiss has recently published some work in which he under- 
 takes to show that what has hitherto been regarded as pure nickel in reality 
 contains an admixture of one, or more, hitherto undiscovered elements, and 
 that the same is probably true of cobalt. Some of the fractions into which 
 Kriiss divided nickel have an atomic weight lying between 56 and 58, the others 
 between 60 and 100; in view of these results, the atomic weights of cobalt and 
 nickel are as yet undetermined. Clemens Winckler, however, by reason of his 
 previous investigations on the atomic weight of nickel and by reason of a 
 review on some parts of his former work, which he instituted with apparently 
 pure materials, doubts Kriiss's results; so that, until further light is thrown on 
 the subject, the old theories as regards cobalt and nickel must be maintained. 
 See Kriiss and Schmidt ; Berichte d. Deutsch. Chem. Gesell. ; 22, 11 and 2026 ; 
 Clemens Winckler, ibid. 890. See also Mond, Journ. Chem. Soc. 1890, 753. 
 Mond has prepared chemically pure nickel. 
 
472 IRON ; COBALT ; NICKEL ; OCCURRENCE. 
 
 Native iron. The occurrence of masses of iron of terrestrial origin has 
 been mentioned several times, but is not beyond doubt. Meteoric iron 
 is not infrequently found; it usually contains from 1 to 20 per cent of 
 nickel. These meteorites contain the metal arranged in striae with a dif- 
 fering contents of nickel, so that, as they offer a differing resistance to 
 acids, meteoric iron, when polished and subjected to the corroding 
 action of reagents, will show a surface marked by regular etchings. 
 
 Iron pyrites, Fe 82 , occurs in rocks of all ages : it is isomorphous with arsen- 
 ical pyrites, FeAsS, with the sulphide of manganese, MnS 2 (haiie- 
 rite), and with the sulphides and arsenides of cobalt and nickel, having 
 the general formulae MS 2 , M As2 , or M As S. Iron pyrites is dimor- 
 phous, for a mineral of the same formula, belonging to a different crys- 
 talline system, is known; this mineral is called markasite. 
 
 Ferrous sulphide, Fe S, is found as troilite. 
 
 Ferric sulphide, Fe 2 83 , frequently plays the part of an acidic anhydride, 
 and with bases, such as Cu2 S, Ag 2 S, or Cu S, forms minerals of which 
 chalcopyrite, Cu 2 S, Fe 2 S 3 = 2 Cu Fe S 2 , is an example (see page 397). 
 
 Hematite (specular iron) is ferric oxide, Fe 2 Og . It is one of the most im- 
 portant iron ores, and occurs in rocks of all ages. 
 
 Magnetite (magnetic iron ore) is ferrous-ferric oxide, Fe O, Fe 2 O 3 = Fe 3 O 4 . 
 This oxide is isomorphous with the spinells (page 339), and is probably 
 similarly constituted. 
 
 Various hydroxides of ferric oxide are also frequently met with. 
 The chief representative of this most important class of minerals is 
 limonite ( brown hematite), Fe 4 9 H 6 = 2 Fe 2 3 -f 6 H 2 0. 
 
 Siderite (spathic iron) is ferrous carbonate, Fe COs . It occurs in many 
 rock strata, in gneiss, mica slate, clay slate, and with the coal forma- 
 tion. It is isomorphous with calcite. 
 
 Iron is also found as a constituent of a large number of silicates, 
 In consequence of the disintegration of the rocks in which it occurs, 
 it finds its way into the soil and into the natural waters. It is an 
 invariable constituent of chlorophyll (the green coloring matter of 
 leaves), and it is always found in the haemoglobin of the blood. 
 
 Cobalt chiefly occurs as cobaltite, Co As S (CoS. 2 4- Co As 2 ), isomorphous 
 with iron pyrites; * as smaltite, Co (Fe Ni) As 2 , also isomorphous with 
 iron pyrites; and as danaite, (Fe, Co) (AsS) 2 , isomorphous with 
 markasite. Cobaltous carbonate, (sphserocobaltite) CoCO 3 ,isalso 
 
 * As might be expected, cobalt replaces iron isomorphously; but in this 
 mineral we have another phenomenon which is not so self-evident, namely, ar- 
 senic replaces sulphur isomorphously. This substitution is not infrequent in 
 the group of minerals of which iron pyrites is the representative (see page 234). 
 
. . 
 
 IRON ; METALLURGY. 473 
 
 sometimes found, as well as the arsenate, Co 3 ( As O 4 ) 2 + 8 H 2 O, which 
 is called erythrite. 
 
 Nickel occurs in meteorites as an alloy of iron ; as millerite, Ni S, isomor- 
 phous with niccolite, as gersdorfite, Ni As S, isomorphous with cobalt- 
 ite and iron pyrites; as the arsenide, niccolite, Ni As (isomorphous with 
 zinc-blende, page 426); as a basic carbonate, as the arsenate, and as a 
 double sulphide of iron and nickel (Fe, Ni) S. This sulphide occurs 
 as a massive variety, and is termed pentlandite. 
 
 Iron, cobalt, and nickel are all easily reduced from their oxides 
 by means of charcoal. The metallurgy of iron is among the most 
 important commercial operations of the present time. In the prep- 
 aration of this metal, the most important ores are the oxides, hy- 
 droxides, and the carbonate. The ores are crushed when necessary, 
 and sometimes roasted for the purpose of expelling the water and 
 carbon dioxide, and of changing the oxides as much as possible into 
 ferric oxide, Fe 2 3 .* They are then reduced in a blast furnace. 
 The latter consists of a shaft, varying in height from fifty to ninety 
 feet, with a maximum diameter of from twenty to twenty-three feet, 
 the shaft being shaped like two truncated cones united at their 
 bases ; below these is a circular chamber or hearth which is built of 
 firebrick, and in which two small openings can be made, the one 
 higher up than the other ; the lower one is used for drawing off the 
 molten metal, the upper one for the slag. The remainder of the 
 furnace is constructed of firebrick and encased in boiler iron. The 
 blast is introduced through from six to eight openings, termed tuy- 
 eres, at from four to six feet above the bottom of the furnace, and 
 the air which is forced in through these is heated by means of the 
 waste gases passing from the furnace. The furnace is charged with 
 alternate layers of iron ore, coke or charcoal, and limestone. The 
 latter substance, uniting with the siliceous matter f which is present 
 in the ore, forms a fusible glass called the slag.$ The molten slag 
 collects at the bottom of the furnace, with the metal, and, being of 
 
 * Some sulphur-bearing ores are roasted to burn off the sulphur, but these 
 ores form a very small proportion of the total iron compounds used. I am 
 indebted to Prof. E. D. Campbell for a review of the pages relating to the 
 metallurgy of iron. 
 
 t Feldspar, slate, quartz, etc. 
 
 J If the ore contains limestone in sufficient or excessive quantity, it may 
 be necessary to add siliceous matter such as feldspar. The lime prevents the 
 formation of a ferruginous slag, which would entail a loss of iron. 
 
474 IKON; METALLURGY. 
 
 less specific gravity, floats upon the surface ; it can therefore be 
 drawn off at the upper small opening (termed the cinder notch) in 
 the base of the furnace. The slag is not formed until almost com- 
 plete reduction has taken place, so that it exercises its protective 
 action only upon the molten metal at the bottom of the furnace. 
 
 The chemical changes which take place in a blast furnace are 
 quite complicated, and all of them are not, as yet, definitely under- 
 stood ; however, the most important reactions in the production of 
 cast iron are as follows. The carbon, uniting with the oxygen en- 
 tering from the tuyeres, forms carbon dioxide immediately above 
 those openings, but the carbon dioxide is almost instantly and com- 
 pletely reduced to carbon monoxide in passing over the red-hot coke 
 or charcoal (see page 287) ; hot carbon monoxide now comes in con- 
 tact with the descending charges of ore, and reduces the oxide to a 
 spongy form of iron : 
 
 Ee 2 3 + 3 CO = 2 Fe + 3 C0 2 . 
 
 The portion of the furnace in which this reduction occurs has a tem- 
 perature of from 450 to 900. The spongy metal passes downward 
 in the furnace, the temperature increasing to 950 at the widest 
 part of the furnace ; at this point the iron takes up carbon to form 
 a chemical combination with that element, and at a lower zone, 
 when the temperature is about 1400, the mass, which has been in 
 a pasty condition, melts and runs down into the hearth. The latter 
 is tapped from time to time, and the iron cast into semi-cylindrical 
 moulds called " pigs." 
 
 Pig iron is quite impure ; it contains carbon, silicon, sulphur, 
 phosphorus, and manganese, and is divided into two chief classes, 
 white cast iron * and gray cast iron ; the former contains the greater 
 3>art of its carbon chemically united with the iron-, the latter, in 
 consequence of the presence of silicon, has separated the major por- 
 tion of its carbon in the form of graphite. The proportion of carbon 
 in white iron is from 3 to 4.4 per cent ; in gray iron, the total 
 carbon is about the same as in white iron, but, because by far the 
 greater proportion of it is in the graphitic for,m, the iron is of a 
 darker color. If the iron ore contained a considerable quantity of 
 manganese, the latter is reduced with the iron, and the alloy so 
 
 * White cast iron is formed at a lower furnace temperature than gray cast 
 iron or spiegeliron. 
 
. 
 
 WKOUG^IT IRON ; STEEL. 
 
 formed is capable of taking up a considerably greater quantity of 
 carbon (as high as 6.9 per cent); this form of iron is known as 
 spiegeliron. Cast iron is brittle, easily fusible (its melting point is 
 about 1050 ) ; it cannot be welded or tempered ; when treated with 
 hydrochloric acid it dissolves, and the combined carbon passes off in 
 the form of hydrocarbons, which possess a most disagreeable odor, 
 while the graphite remains undissolved. 
 
 Wrought iron is produced from cast iron by puddling. Pud- 
 dling consists of melting cast iron in a furnace lined with ferric 
 oxide. The oxygen of this lining, combining with the carbon, sili- 
 con, phosphorus, and manganese, forms the oxides of those elements. 
 The carbon passes off as carbon monoxide ; the silicon dioxide, 
 phosphoric acid, and manganous oxide unite with the excess of the 
 oxides of iron to form a slag, so that a metal which is nearly pure 
 is formed. Wrought iron contains less than .3 per cent of carbon* 
 and very small amounts of silicon and phosphorus ; it possesses a 
 fibrous texture, is malleable and ductile, and melts at about 1900. 
 
 Steel contains more carbon than wrought iron, and always less 
 than cast iron ; it is produced by the Bessemer, open-hearth, and 
 crucible processes. 
 
 The Bessemer steel process produces steel directly from cast 
 iron; it consists, briefly, of first burning out the impurities in 
 melted cast iron, by placing the latter in a large converter and 
 forcing air in through the bottom,f and then, after stopping the 
 blast, of adding spiegeliron $ until the requisite amount of carbon is 
 present. Bessemer steel is used in the manufacture of rails and of 
 other large steel implements ; it contains from .1 to 1. per cent of 
 combined carbon. 
 
 The open-hearth process consists in melting (in large regenerative 
 furnaces) a mixture of steel scraps and a small proportion of cast 
 iron, and of then adding a sufficient quantity of spiegeliron to the 
 molten metal to give the desired percentage of carbon and man- 
 
 * When it contains more than .3 per cent it is steel. 
 
 t The oxidizing action of the air maintains the high temperature. 
 
 % If the cast iron contains a large amount of phosphorus the crucibles 
 (converters) are lined with burned dolomite or magnesite (see page 416), the 
 phosphorus is then oxidized, and forms calcium phosphate with lime which is 
 added with the metal. This latter process is termed the "basic Bessemer 
 process." 
 
476 IRON; PROPERTIES. 
 
 ganese. Open-hearth steel has the same composition as Bessemer 
 steel. 
 
 Crucible steel is made by melting (in small covered crucibles) 
 purest wrought iron, after adding a sufficient supply of charcoal, 
 with pure pig iron, and a small amount of manganese dioxide. 
 This steel contains from .6 to 2. per cent of carbon ; and the purity 
 of the materials from which it is made, as well as its freedom from 
 dissolved oxides and gases, renders this class of steel the only one 
 suitable for the manufacture of the highest grade of tools. 
 
 When steel is heated, and rapidly cooled by plunging into water, 
 it becomes very hard and brittle ; this hardened steel, when once 
 more heated and allowed to cool slowly, becomes elastic ; this pro- 
 cess is called tempering. Steel is capable of taking a very high 
 polish. 
 
 Chemically pure iron can be prepared by reducing either the pure 
 oxide, oxalate, or chloride of iron in a current of hydrogen ; it has a 
 specific gravity of 7.84, and a melting point which is probably not 
 less than 1800 ; * it is bluish-gray, almost white, in color, and is 
 malleable and ductile ; one of the most striking physical properties 
 that it possesses is that of magnetism. Pure iron loses its magnet- 
 ism as soon as a magnet, which has been placed in its neighbor- 
 hood, is removed ; steel, however, is able to retain the property. 
 Pure iron is not attacked by dry oxygen at ordinary temperatures ; 
 when exposed to moist air it undergoes slow oxidation, f forming 
 ferric oxide, Fe 2 3 ; if it is heated in oxygen it burns, forming 
 Fe 3 4 , mixed with Fe 2 3 (see page 22) ; it unites with the halo- 
 gens in the same manner. Iron will rust when placed under water 
 which contains dissolved oxygen ; this action is accelerated by the 
 presence of acids and retarded by .the presence of alkalies. $ Dilute 
 hydrochloric or sulphuric acids dissolve iron, liberating hydrogen 
 and forming ferrous chloride and ferrous sulphate respectively (see 
 page 32) ; concentrated sulphuric acid, when cold, is without action, 
 
 * The melting point of pure iron has been variously given as being 1560, 
 1587, 1600, 1800, and it has even been stated that absolutely pure iron is 
 infusible. 
 
 t It is stated that moist, pure air does not attack iron ; the rusting tak- 
 ing place only if carbon dioxide is present. 
 
 J The rusting is also assisted by the presence of salts, especially of those 
 of ammonium. 
 
COBALT; NICKEL; METALLURGY. 477 
 
 when heated with iron it liberates sulphur dioxide and produces 
 ferric sulphate (see page 136) ; concentrated nitric acid has the same 
 effect on iron as it has on aluminium (see page 334), the metal does 
 not dissolve but is transferred to the " passive state ; " when in this 
 condition it is no longer attacked by the dilute nitric acid, nor will 
 it separate copper from a solution of copper sulphate, a reaction 
 into which ordinary iron very readily enters : 
 
 Cu S0 4 + Fe = Fe S0 4 + Cu.* 
 
 Several explanations as to the reason of this condition have been of- 
 fered. One of these is that the iron becomes covered with a very 
 thin layer of ferrous-ferric oxide, which is insoluble in nitric acid ; 
 this theory, however, is without absolute experimental proof, f 
 Dilute nitric acid dissolves iron, forming ferrous nitrate, while a por- 
 tion of the acid is reduced to ammonium nitrate (see page 206, a). 
 
 Cobalt is quite difficult to obtain in a pure state, because its ores always 
 contain iron and nickel, from which latter element the metal is not easy 
 to separate; copper, bismuth, lead, or silver may also be present. The 
 chief points in the separation are : first, the burning away of the sul- 
 phur and arsenic present in the ores; and secondly, the separation of 
 the copper, bismuth, etc., by means of sulphuretted hydrogen, after 
 the oxides produced by the roasting have been dissolved in acids. 
 The solutions which remain after the sulphides (which have been 
 precipitated) have been filtered off, are treated with chlorine and cal- 
 cium hypochlorite ; by this means the cobalt salts, which are present, 
 are oxidized to insoluble cobaltic oxide before those of nickel are 
 attacked; the cobaltic oxide, when the operation has gone just far 
 enough, is separated, dried, and reduced to metallic cobalt by heating 
 in a current of hydrogen. 
 
 Nickel is more easily obtained than cobalt, because it is present in greater 
 quantity, and because its ores are, as a rule, purer. The sulphides or 
 arsenides of nickel are roasted in the same manner as those of cobalt ; 
 the oxide so obtained is mixed with charcoal and heated, by which 
 means reduction to metallic nickel takes place. 
 
 Cobalt forms crystalline metallic plates which have a specific 
 gravity of 8.5, and which melt at a somewhat lower temperature 
 
 * See page 313, foot-note. 
 
 t In support of this theory is the fact that the passive state can also be 
 produced by other oxidizing agents, such as chloric, bromic, iodic, or chromic 
 acids. Another theory of considerable plausibility is that the iron becomes 
 covered with a thin layer of gas; both hypotheses are borne out by the fact 
 that passive iron, when rubbed, returns to its normal state. 
 
478 COBALT; NICKEL; PROPERTIES. 
 
 than iron. Like the latter, cobalt is capable of attracting the 
 magnet. The metal is susceptible of a very high polish, is malle- 
 able and very ductile. The metal, after it has been cast into solid 
 pieces, is entirely unaltered by exposure to the air ; at white heat, 
 however, it burns to form cobaltous-cobaltic oxide, Co 3 4 . When 
 heated and then plunged into concentrated nitric acid it becomes 
 "passive." Hydrochloric or sulphuric acids slowly dissolve the 
 metal, forming cobaltous chloride and cobaltous sulphate respectively. 
 Dilute nitric acid readily dissolves cobalt to produce cobaltous 
 nitrate. 
 
 Nickel. It is somewhat doubtful if chemically pure nickel has 
 ever been obtained* (see page 471) ; that which has hitherto been 
 regarded as such is produced by reduction from the pure oxalate 
 of nickel. When cast into cubes it is a lustrous, almost silver- 
 white metal which is nearly as malleable and ductile as iron. It 
 melts at a temperature lower than the melting point of either 
 cobalt or iron. It is attracted by the magnet at ordinary temper- 
 atures, but loses this property when heated to 350. Nickel which 
 has been cast into solid pieces is not oxidized in the air, and it 
 scarcely burns even when heated white hot in an atmosphere of 
 oxygen. The metal is but slowly attacked by hydrochloric or sul- 
 phuric acid; it readily dissolves in nitric acid to form nickelous 
 nitrate. Concentrated nitric acid renders the metal "passive." 
 The specific gravity of nickel is 8.9. 
 
 Alloys of iron. Alloys of manganese, tungsten, or nickel, with 
 steel, possess great toughness, and are of growing commercial im- 
 portance. The power which zinc and tin possess, of adhering to the 
 surfaces of iron sheets, is probably due to the formation of alloys. 
 
 Alloys of nickel are used in the preparation of coins,f in the 
 manufacture of German silver (which contains copper, zinc, and 
 nickel) and in a number of other ways. Nickel-plating is accom- 
 plished by electrolyzing a solution of nickel-ammonium sulphate, | 
 the metal to be covered being the negative electrode. 
 
 Compounds of iron. Iron forms two series of compounds, fer- 
 
 * Mond, Journ. Chem. Soc. 1890, 753, asserts that chemically pure nickel 
 can be prepared by the decomposition of a liquid compound of nickel and carbon- 
 monoxide, Ni(CO) 4 . The atomic weight of this nickel is 58.69. 
 
 t The nickel coins of the United States contain 75% copper and 25% nickel. 
 
 J Steel can be nickel-plated by simply plunging into a bath of nickel-ammo- 
 nium sulphate. 
 
FERROUS COMPOUNDS. 479 
 
 rous compounds, in which the metal is divalent, and ferric com- 
 pounds, in which it is trivalent. In addition to these two stages of 
 oxidization, there exists a ferric acid (the anhydride of which would 
 be Fe 8 ), in which iron is hexavalent. 
 
 Ferrous oxide, Fe 0, is produced by reducing ferric oxide in a 
 current of hydrogen ; the corresponding hydroxide has the formula 
 Fe (OH) 2 . This compound is precipitated by adding ammonia 
 water to a solution of a ferrous salt ; the precipitate is white, but 
 it rapidly turns brown on exposure to the air, because it changes 
 into ferric hydroxide. Both ferrous oxide and hydroxide are bases ; 
 they dissolve in acids to form ferrous salts. 
 
 FERROUS CHLORIDE is produced by dissolving iron in hydrochloric acid, 
 while excluding the air. The dry salt, having the formula Fe C1 2 , can 
 then be isolated by evaporating the solution in a current of hydrogen. 
 It is a white, crystalline mass which volatilizes at a high red heat. The 
 vapor density of ferrous chloride at white heat, air = 1, is 4.39. This 
 number corresponds to a molecular weight given by the formula 
 Fe G1 2 . At a lower temperature the molecules are probably Fe 2 C1 4 . 
 When exposed to the air, ferrous chloride rapidly oxidizes to a mix- 
 ture of ferric chloride and ferric oxide : 
 
 6 Fe C1 2 + 3 O = 4 Fe C1 3 + Fe 2 O 3 . 
 
 When in solution, provided hydrochloric acid is present, ferric chlo- 
 ride alone is produced : 
 
 2 Fe Clo + 2 H Cl +O = 2 Fe C1 3 + H 2 O . 
 
 On the other hand, if an excess of hydrochloric acid is not present, 
 ferric chloride and an insoluble basic ferrous chloride are formed : 
 
 4 FeCl 2 + H 2 + O = 2 Fe j ^ H + 2 FeCl 3 . 
 
 Of course, the usual oxidizing agents (nitric acid, potassium chlorate, 
 and hydrochloric acid, chlorine, bromine, etc.) instantly change 
 ferrous chloride into ferric chloride. Ferrous chloride, like the chlo- 
 rides of the majority of other divalent elements, forms double salts 
 with the chlorides of the alkalies, as well as with the chlorides of a 
 number of other metals. 
 
 FERROUS SULPHATE (green vitriol, copperas), Fe S O 4 + 7 H 2 O, is isomor- 
 phous with the vitriols (see page 417). It is produced, commercially, 
 by the spontaneous oxidation of iron pyrites, or by dissolving iron in 
 dilute sulphuric acid (see page 32). The salt loses six molecules of 
 water at 100, and is completely dehydrated at 300; it then forms a 
 white powder. Like the other vitriols, ferrous sulphate forms double 
 salts with the sulphates of the alkali metals ; the latter contain six 
 
480 FERROUS COMPOUNDS. 
 
 molecules of water of crystallization, and have the general formula, 
 Fe SO 4 , M 2 SO + 6 H 2 O; they are not as easily oxidized as the pure 
 sulphate of iron. When exposed to moist air, ferrous sulphate is 
 oxidized to a mixture of ferric sulphate and ferric hydroxide: * 
 
 6 Fe SO 4 + 3 H 2 O + 3 O = 2 Fe 2 (SO 4 ) 3 + 2 Fe (OH) 3 , 
 
 while, in the presence of sulphuric acid, ferric sulphate alone is pro- 
 duced : 
 
 2 Fe SO 4 + H 2 SO 4 + O = Fe 2 (SO 4 ) 3 + H 2 O. 
 
 The usual laboratory oxidizing agents have the same effect. Ferrous 
 sulphate absorbs nitric oxide (NO), forming an unstable chemical 
 compound with the latter. The solution is dark brown in color, and 
 parts with the dissolved gas when heated, t 
 
 FERROUS SULPHIDE, Fe S, sometimes occurs as a mineral. It is formed 
 by heating together iron and sulphur or by precipitation, as a black 
 powder, when the solution of an alkaline sulphide is added to a solu- 
 tion of a ferrous salt: 
 
 Fe S0 4 + (NH 4 ) 2 S = Fe S + (NH 4 ) 2 SO 4 . 
 
 Ferrous sulphide belongs to the class of sulphides which are dissolved 
 by dilute acids ; it is, therefore, not precipitated by hydrogen sulphide 
 in acid solution (see page 100). When an alkaline sulphide is added 
 to a solution containing a ferric salt, not ferric sulphide, but ferrous 
 sulphide mixed with sulphur, is precipitated, this phenomenon being 
 due to the instability of ferric sulphide in aqueous solution; the change 
 may be represented by the following equations : 
 
 1. Fe 2 S 3 = 
 
 2. Fe 2 (SO 4 ) 3 + 3Na 2 S = 2 FeS + S + 3Na 2 SO 4 . 
 
 FERROUS CARBONATE, Fe CO 3 , occurs as the mineral siderite, isomorphous 
 with calcite. It can be formed in the laboratory by adding a soluble 
 carbonate to a solution containing a ferrous salt : 
 
 Fe S0 4 + Na 2 CO 3 = Na 2 SO 4 + Fe CO 3 . 
 Soluble. Soluble. Insoluble. 
 
 When moist, ferrous carbonate is readily oxidized by the air, ferric 
 hydroxide remaining, for ferric hydroxide (like the hydroxide of alu- 
 
 * Possibly a basic sulphate. 
 
 t The formation of this solution is a delicate test for nitric acid. Nitric 
 acid oxidizes ferrous sulphate as follows : 
 
 GFeS0 4 +2HN0 3 + 3H 2 S0 4 =3Fe 2 (S0 4 ) 3 + 4H 2 + 2NO; 
 
 the nitric oxide, which is liberated, colors the excess of ferrous sulphate. Of 
 course, this test may also be used for detecting the presence of a nitrate, for the 
 latter, with sulphuric acid, forms a sulphate and free nitric acid. 
 
FERRIC OXIDE ; HYDROXIDE. 481 
 
 minium and chromic hydroxide) is too weakly basic to form a carbon- 
 ate (see page 342). Ferrous carbonate is easily decomposed into 
 ferrous oxide and carbon dioxide by heat. 
 
 Ferric oxide, Fe 2 3 , is found in nature as hematite; in the 
 laboratory it can be produced by heating the corresponding hydrox- 
 ide, Fe (OH ) 8 : 
 
 2 Fe_(OH ) 8 = Fe 2 3 + 3 H 2 0. 
 
 The oxide so prepared has a fine red color ; it is known as rouge, and 
 is used as a polish for metals and glass. Ferric hydroxide is pre- 
 cipitated by adding an alkaline hydroxide or ammonia water to a 
 solution containing a ferric salt : 
 
 Fe 2 ( S0 4 ) 3 + 6 KOH = 2 Fe (OH ) 3 + 3 K 2 S0 4 ; 
 
 the ferric hydroxide so formed probably loses water spontaneously, 
 so that the precipitate is really the metahydroxide, Fe (OH). A 
 number of the ferric hydroxides, which are formed by the separa- 
 tion of water between two or more formula weights of the normal 
 hydroxide, occur as minerals which are commercially important; 
 such compounds are limonite, (4 Fe (OH ) 3 = Fe 4 3 (OH ) 6 -f- 3 H 2 0, 
 and xanthosiderite (bog iron ore), Fe 2 (OH ) 4 ; * the latter isomor- 
 phous with beauxite (see page 333). Ferric hydroxide can also be 
 obtained, in a soluble form, by dialyzing an extremely dilute solu- 
 tion of ferric chloride mixed with ammonium carbonate (see pages 
 305, 306) ; the solution has properties similar to those of dialyzed 
 silicic acid. Ferric hydroxide is distinguished from the hydroxide 
 of aluminium and from chromic hydroxide by the fact that it is in- 
 soluble in an excess of caustic alkali ; nevertheless, it does possess 
 acidic properties, as is proved by the existence of minerals like 
 franklinite, for the latter is a zinc ferrite, Zn ( Fe 2 ) 2 , derived from 
 meta ferric hydroxide, Fe (OH ) ; | f ranklinite is isomorphous 
 with spinell (see page 339), the latter being a salt of meta alumin- 
 ium hydroxide, Al O (OH ), acting as an acid. Besides being 
 acquainted with zinc ferrite, we know that magnetite is, in all 
 probability, a ferrous salt of ferric hydroxide, for it has the formula 
 Fe ( Fe 2 ) 2 , and is isomorphous with spinell ; and, lastly, calcium 
 
 * 2 Fe (OH% - Fe 2 O (OH) 4 + H 2 O. 
 
 t In franklinite a part of the zinc is replaced by ferrous iron and by man- 
 .ganous manganese. 
 
482 FERRIC SALTS. 
 
 and barium ferrites can be produced in the laboratory by heating 
 ferric oxide with calcium or barium oxide to a high red heat ; the 
 latter salts have the composition : 
 
 FejO FejO 
 
 ^ > Ca and ^ > Ba. 
 
 '"' T71 \ {J 
 
 On the other hand, ferric oxide and ferric hydroxide are bases, for 
 they dissolve in acids to form ferric salts.* 
 
 FERRIC CHLORIDE, Fe C1 3 , can be produced by dissolving ferric hydroxide 
 in hydrochloric acid, or by passing chlorine into a solution of ferrous 
 chloride : 
 
 Fe C1 2 + Cl = Fe C1 3 . 
 
 Upon evaporation in the cold, crystals of the composition Fe Cl 3 -f 3 H 2 O 
 are deposited; another hydrate of the formula Fe C1 3 + 6 H 2 O is formed 
 if the solution is allowed to stand at ordinary temperatures ; a solution 
 of ferric chloride cannot be heated, because then it breaks down into 
 hydrochloric acid and an insoluble oxychloride ; t reducing agents, such 
 as hydrogen sulphide, zinc and hydrochloric acid, sulphurous acid, etc., 
 change ferric chloride into ferrous chloride. Anhydrous ferric chloride 
 is prepared by passing chlorine over powdered iron ; the chloride evap- 
 orates at the temperature of boiling sulphur (448), and its vapor den- 
 sity, even at that temperature, is too low for a substance consisting of 
 molecules having the formula Fe 2 C1 6 , the specific gravity of the gas- 
 eous chloride rapidly decreases as the temperature is raised, so that 
 no doubt can exist as to the trivalence of iron in ferric chloride, the 
 molecule of which, in a gaseous state, is expressed by the formula 
 FeCl 8 , (see page 336).} Ferric chloride combines with other chlo- 
 rides to form double salts similar to those produced by aluminium 
 chloride (see page 337 ). 
 
 FERRIC SULPHATE, Fe 2 (SO4 ) 3 is formed by oxidizing ferrous sulphate 
 (see page 340) or by dissolving ferric hydroxide in sulphuric acid; it 
 forms alums with the sulphates of the alkali metals (see page 340 ). 
 
 Ferric sulphide, Fe 2 S 3 , analogous to ferric oxide, Fe 2 3 , can be 
 produced by heating ferrous sulphide and sulphur to a red heat : 
 
 * Ferric oxide, like the oxides of aluminium and chromium, is not dis- 
 solved by acids after it has been heated to a red heat. 
 
 t If the solution is quite dilute, complete decomposition and formation of 
 soluble ferric hydroxide takes place: 
 
 Fe C1 3 + 3 H 2 O = Fe (OH) 8 + 3 H Cl. 
 
 t W. Grunewald and V. Meyer, Berichte der Deutsch. Chem. Gesell. ;21, 
 687. 
 
FERROUS-FERRIC OXIDE. 483 
 
 it has the character of an acidic anhydride, and forms a number of 
 salts which occur as natural minerals ; an example of these com- 
 pounds is chalcopyrite (copper pyrites), in which cuprous sulphide 
 is the base : 
 
 Cu 2 S + Fe 2 S 3 = Cu 2 Fe 2 S 4 . 
 
 These sulphoferrites are analogous to the ferrites (examples of 
 the latter occur in the spinell group) j when heated, ferric sulphide 
 is converted into the compound Fe 3 S 4 ; this change is similar to 
 that undergone by ferric oxide. 
 
 Ferrous-ferric oxide, Fe 3 4 , occurs as magnetite, a mineral which 
 has the power of being attracted by the magnet. The oxide is pro- 
 duced when iron is burned in an excess of oxygen, or when steam is 
 passed over red-hot iron (see page 31 ) ; the black coating which 
 forms on iron which is heated to a high red heat consists of a mix- 
 ture of Fe 8 4 and Fe 2 3 . Ferrous-ferric oxide is considered to be 
 constituted similarly to the spinells ; its structure would, therefore, 
 be represented as follows : 
 
 Ho 
 
 > Fe. 
 
 The sulphide, Fe 3 S 4 , which, in formula, corresponds to magnetite, 
 is produced when a current of the hydrogen sulphide is passed over 
 red-hot iron : 
 
 3 Fe + 4 H 2 S = Fe 3 S 4 + 8 H. 
 
 Like the oxide, it has magnetic properties. 
 
 Ferric acid is not known in the free state ; the ferrates, which in 
 formula are analogous to the sulphates, chromates, and manganates, 
 however, exist in limited numbers ; the best known of these is potas- 
 sium ferrate, K 2 Fe 4 ; the latter is produced, just as is potassium 
 manganate, by fusing the metal with potassium nitrate ; the same 
 salt is also formed by passing chlorine into a solution of potassium 
 hydroxide which contains ferric hydroxide in suspension.* Potas- 
 
 * The oxidizing agent is the potassium hypochlorite which is produced 
 (see page 121) ; the solution of ferrate of potassium has a violet color. Ordi- 
 nary potassium hydroxide not infrequently contains some ferric hydroxide; 
 and the formation of potassium ferrate, with its violet color, is sure to puzzle 
 the student when he is preparing potassium hypochlorite. 
 
484 COBALTOUS COMPOUNDS. 
 
 slum ferrate, as well as the other ferrates, is unstable ; when allowed 
 to stand it breaks down, giving off oxygen and leaving ferric 
 hydroxide. 
 
 The disulphide of iron, Fe S 2 , has no analogon among the oxygen 
 compounds of the metal ; it is, however, extremely important, be- 
 cause it occurs so widely distributed as the mineral iron pyrites. 
 The disulphide is dimorphous ; as pyrites it crystallizes in the regu- 
 lar system, and as markasite it is rhombic ; it has a metallic lustre, 
 and, on a superficial examination, has somewhat the appearance of 
 gold. 
 
 The cyanides of iron, the ferro and ferricyanides, as well as the 
 corresponding acids, were discussed on page 296. 
 
 Compounds of Cobalt. Cobalt forms two series of compounds, 
 cobaltous compounds, derived from cobaltous oxide, Co 0, and cobaltic 
 compounds, derived from cobaltic oxide, Co 2 3 ; the cobaltous salts 
 are the ones which are most frequently met with, cobaltic oxide 
 having only very weakly basic properties. 
 
 Cobaltous oxide, Co 0, can be prepared by decomposing the corre- 
 sponding hydroxide, Co(OH) 2 , by heat, air being excluded; the 
 hydroxide is precipitated by adding an excess of caustic alkali to 
 the solution of a cobaltous salt : 
 
 Co ( N0 3 ) 2 + 2 KOH = Co (OH ) 2 + 2 KN0 8 . 
 
 Cobaltous hydroxide is rose colored, the oxide is olive green ; both 
 the oxide and hydroxide are strong bases, dissolving in acids to 
 form cobaltous salts. ' 
 
 COBALTOUS CHLORIDE, Co C1 2 , can be formed by dissolving the carbonate or 
 hydroxide in hydrochloric acid; when the solution is evaporated, red 
 crystals of the composition Co C1 2 + 6 H 2 O separate ; when heated 
 to 100 the salt becomes violet and has the composition Co C1 2 + 
 H 2 O; at 110-120 it loses all its water of crystallization, and is con- 
 verted into the anhydrous chloride, Co C1 2 , which is blue ; the same 
 rule appertains to all cobalt salts; when hydrated they are red or 
 rose-colored, when anhydrous they are blue.* 
 
 * This property of cobalt salts caused the use of those substances as far 
 back as 1757 for the preparation of sympathetic ink; this consists of cobalt 
 chloride in solution, and is, therefore, pale red, and almost invisible on the 
 paper. When the paper is heated, however, the salt becomes anhydrous, and 
 the writing appears in plain, blue characters. 
 
COBALTIC COMPOUNDS. 485 
 
 COBALTOUS NITRATE, Co (NO 3 ) 2 + 6 H 2 O, is the most common cobalt salt; 
 it loses water at 100, and at a higher temperature decomposes, giving 
 off nitrogen dioxide and leaving cobaltic oxide. 
 
 A COBALT GLASS, formed by fusing silicon dioxide, potassium carbonate, 
 and some salt of cobalt, has an intensely blue color; when finely ground, 
 it is known as smalt ; the glass has approximately the same composition 
 as ordinary potash-glass (see page 421) with the exception that the cal- 
 cium is replaced by cobalt. Fused borax will dissolve cobalt oxide, or 
 salts of cobalt, leaving an intensely blue, glass-like substance which 
 contains cobaltous metaborate ; the same is true of sodium metaphos- 
 phate. The formation of these two compounds is a ready means for 
 detecting the presence of cobalt. When heated with aluminium 
 oxide, cobalt compounds form a blue cobalt aluminate which is known 
 as Thenard's blue. 
 
 COBALTOUS SULPHIDE, Co S, is a black precipitate, formed by adding an 
 alkaline sulphide to a solution containing a cobalt salt : 
 
 Co (N0 3 ) 2 + (NH 4 ) 2 S = Co S + 2 NH 4 NO 3 ; 
 
 it differs from ferrous sulphide by not being soluble in very dilute 
 acids. 
 
 Cobaltic oxide, Co 2 8 , has very weakly basic properties ; its salts, 
 which are formed by dissolving the oxide in cold acids, are easily 
 decomposed, and are tolerably stable only in solution ; a number of 
 double salts, derived from a cobaltic salt united with a salt of some 
 other base, are known, which are more stable than the pure cobaltic 
 salts. 
 
 COBALTIC NITRITE combines with potassium nitrite to form a double salt 
 of the composition Co(NO 2 ) 3 , 3KNO 2 , which is of importance be- 
 cause it is insoluble in water and in cold, dilute acids ; its precipitation 
 may, therefore, be used in separating cobalt from nickel. It is produced 
 by adding potassium nitrite to a slightly acid -solution of a cobalt salt; 
 the nitrous acid, which is liberated from the potassium nitrite, then 
 oxidizes the cobaltous compound to a cobaltic one, after which reaction 
 the double nitrite of cobalt and potassium is precipitated. 
 
 A solution of cobaltous chloride, when in the presence of am- 
 monia,* and exposed to the air, is slowly oxidized ; by this means 
 
 * The addition of an excess of ammonia to cobaltous salts does not produce 
 a precipitate of hydroxide ; the same, it will be remembered, is true of cupric 
 salts. 
 
486 COBALT AMINES. 
 
 there is then formed a peculiar series of compounds composed of 
 cobaltic chloride united with a varying number pf molecules of am- 
 monia. These compounds are known as cobalt amines; the structural 
 composition of the latter, despite the extended amount of work 
 which has been done upon them, is, as yet, not understood ; analysis 
 shows us that one atom of cobalt can unite with three, four, five, and 
 six molecules of ammonia to form different radicles which can play 
 the part of individual trivalent bases ; these bases, themselves, exist 
 only in aqueous solution, the salts derived from them are, however, 
 stable ; the ammonia in these compounds cannot, therefore, be com- 
 pared to water of crystallization, as it is in the case of the ammonia 
 compounds of copper sulphate (see page 403). Only a few of the 
 simpler connections can be traced here ; for a more extended review 
 of this complicated subject the pupil must refer to a larger text- 
 book.* 
 
 r 1. Dichrocobaltic chloride, Co(NH 3 ) 3 C1 3 + H 2 O. 
 
 2. Praseocobaltic chloride, Co (NH 3 ) 4 C1 3 + H 2 O. 
 
 From solution of Co C1 2 J . Roseocobaltic ch ioride, Co (NH 3 ) 5 C1 3 + H 2 O. 
 
 in air are produced ( Purpure ocobaltic chloride, Co (XH,) 5 C1 3 t. 
 
 1 4. Luteocobaltic chloride, Co (NH 3 ) 6 C1 3 . 
 
 The radicles which assume metallic functions in this series of salts 
 may be compared with metallic cobalt ; for from each, in addition to 
 the chlorides, a number of other salts, such as the nitrate and sul- 
 phate, are derived. The relationship is demonstrated by the follow- 
 ing table : 
 
 Co Cl 3 , cobaltic chloride; Co (NO 3 ) 3 , nitrate; Co 2 (SO 4 ) 3 , sulphate; ' 
 
 Co(NH 3 ) 3 Cl 3 , dichro chloride ; Co(NH 3 ) 3 (NO 3 ) 3 , nitrate; [Co(NH 3 ) 3 ] 2 (SO 4 ) 3 , sulphate; 
 Co(NH 3 ) 4 Cl 3 , praseo chloride; Co(NH 3 ) 4 (NO 3 ) 3 , nitrate; [Co(NH 3 ) 4 ] 2 (SO' 4 ) 3 , sulphate; 
 Co(NH,). C1 3 , roseo chloride; Co(NH 3 ) 6 (NO 3 ) 3 , nitrate; [Co(NH 3 ) 5 ] 2 (SO 4 ) 3 , sulphate; 
 Co(NH 3 ) 6 Cl 3 , luteo chloride; Co(NH 3 ) 6 (NO 3 ) 3 , nitrate; [Co (NH 3 ) 6 ] 2 (SO 4 ) 3 , sulphate. 
 
 The cyanides of cobalt correspond to those of iron (see page 
 296). 
 
 The compounds of nickel are almost exclusively derived from 
 nickelous oxide. The latter can readily be produced by heating the 
 hydroxide, Ni (OH ) 2 , which is precipitated as a grass-green sub- 
 stance on adding a soluble hydroxide to a solution containing a salt 
 
 * Ladenburg; Handworterbuch der Chemie ; volume 5. 
 
 t Roseocobaltic chloride is formed from cobaltic chloride and concentrated 
 ammonia in the cold; purpureocobaltic chloride, by boiling the same with con- 
 centrated hydrochloric acid. The difference is in the water of crystallization. 
 
NICKEL; COMPOUNDS OF. 487 
 
 of nickel.* Nickel salts, when hydrated, are green; when anhy- 
 drous, yellow. 
 
 NICKELOUS CHLORIDE, Ni C1 2 + 6 H 2 O, is formed by dissolving the 
 oxide or hydroxide in hydrochloric acid. Nickelous sulphate, pro- 
 duced by substituting sulphuric for hydrochloric acid, is isomorphous 
 with the vitriols (see page 417). Nickelous sulphide is precipitated 
 as a black powder when a solution of an alkaline sulphide is added to 
 a solution containing a nickelous salt: 
 
 Ni S0 4 + (NH 4 ) 2 S = Ni S + (NH 4 ) 2 SO 4 . 
 Nickelous sulphide is insoluble in dilute acids. 
 
 The cyanides of nickel do not correspond to those of iron and of 
 cobalt, with the exception of nickelous cyanide. The latter, with a 
 formula of Ni (CN ) 2 , is precipitated when potassium cyanide is 
 added to a solution containing a salt of nickel ; on addition of an 
 excess of the reagent, it forms a double cyanide of the formula 
 Ni (CN ) 2 , 2 KCN, which, therefore, does not correspond to the 
 potassium salts of ferrocyanic and cobaltocyanic acids : 
 
 Ni (CN ) 2 , 2 KCN. K 4 Fe (CN ) 6 . K 4 Co (CN ) 6 . 
 
 Double salt of nickel. Potassium ferrocyanide. Potassium cobaltocyanide. 
 
 Potassium ferrocyanide and cobaltocyanide can be readily oxidized 
 to the cobaltic and ferric compounds, K 3 Fe (CN ) 6 and K 3 Co (CN ) 6 , 
 which are salts of cobalticyanic acid and ferricyanic acid respec- 
 tively ; but the nickel cyanide is not capable of oxidation. Nickel 
 can be precipitated from its double cyanide by the usual reagents, 
 such as ammonium sulphide ; but the iron or cobalt cannot be 
 separated by any such means, when either has entered into the 
 formation of the peculiar acids from which its double cyanides 
 are derived. In the behavior of its cyanides, in the isomorphism of 
 its sulphate with the vitriols, and in a number of other respects, 
 nickel resembles the first element in the second half of the long 
 period, i.e., copper. 
 
 The relationship between the formulae of the compounds of iron, 
 cobalt, nickel, chromium, manganese, and sulphur are given in the 
 following table : 
 
 * Excess of ammonia water dissolves the hydroxide ; this is similar to the 
 action of cobaltous hydroxide and cupric hydroxide. 
 
488 IRON ; COBALT ; NICKEL ; TABLE OF. 
 
 OXIDES WHICH ARE BASIC, AND SALTS DERIVED FROM THEM. 
 
 -ous 
 
 OXIDES. 
 
 -Oils SAI.'IS. 
 
 -ic 
 
 OXIDES. 
 
 -1C SALTS. 
 
 CrO 
 MnO 
 FeO 
 CoO 
 NiO 
 
 CrCl a 
 MnCl, 
 FeCl 2 
 CoCl 2 
 NiCl 2 
 
 chlorides 
 
 CrSO 4 * 
 MnSO 4 
 FeSO 4 
 CoSO 4 
 NiSO 4 
 
 sulphates 
 
 
 Cr 2 3 $ 
 Mn 2 O 3 
 Fe 2 3 * 
 Co 2 3 
 Ni 2 3 
 
 CrCl 3 
 Mn Cl 3 t 
 Fe Cl 3 t 
 CoCl 3 t 
 
 Cr 2 fS0 4 ) 3 
 Mn 2 (S0 4 ) 3 | 
 Fe 2 (S0 4 ) 3 t 
 
 Cr(NO 3 ) 3 
 
 Mn(N0 3 ) 2 
 Fe(N0 3 ) 2 
 Co(N0 3 ) 2 
 Ni(N0 3 ), 
 nitrates 
 
 Fe(N0 3 ) 3 t 
 
 
 chlorides 
 
 sulphates 
 
 nitrates 
 
 * Compounds marked * are unstable, very readily oxidized ; those marked t 
 readily change to -ous compounds. Ni 2 O 3 is known as its hydroxide Ni (OH) 3 , 
 formed by oxidizing a nickelous salt with potassium hypochlorite or hypobro- 
 mite; it forms no salts. 
 
 The -ous compounds are converted into -ic compounds by oxidation, the 
 -ic compounds into -ous compounds by reduction. 
 
 J Both basic and acidic, Cr 2 O 3 dissolves in caustic alkalies to form chro- 
 mites; Fe 2 O 3 forms some compounds which are analogous to the chroinites and 
 aluminates; these salts are derived from a hydroxide MO (OH). 
 
 OXIDES WHICH ARE COMPOSED OF THE TWO GIVEN ABOVE. 
 
 -ous-ic- OXIDES. 
 
 Cr 3 4 , Mn 3 4 , Fe 3 O 4 , Co 3 O 4 . 
 
 Of the elements in this table, manganese alone forms a dioxide, Mn O 2 ; how- 
 ever, sulphides or arsenides of the other elements, analogous in structure, 
 are known. 
 
 OXIDES WHICH ARE ACIDIC ONLY. 
 
 a. Acids and salts of the general formula H 2 XO 4 ; type H 2 SO 4 
 
 ANHYDRIDES. 
 
 CrO 3 
 
 M 2 Cr O 4 , chromates ; 
 M 2 Mn O 4 , manganates ; 
 M 2 Fe O 4 , ferrates. 
 
 b. Acids and salts of the general formula HXO 4 ; type H Cl O 4 
 
 ANHYDRIDE. 
 
 SALTS. 
 
 Mn. 2 O 7 
 
 M Mn O 4 , permanganates. 
 
ELEMENTS OF PLATINUM GROUP. 489 
 
 CHAPTER LX. 
 
 THE REMAINING ELEMENTS OF THE EIGHTH FAMILY. 
 (THE PLATINUM GROUP.) 
 
 Ruthenium ; symbol, Ru ; atomic weight, 101.6 ; 
 Rhodium ; symbol, Rh ; atomic weight, 103 ; 
 Palladium ; symbol, Pd ; atomic weight, 106.6 ; 
 Osmium ; symbol, Os ; atomic weight, 190.8 ; 
 Iridium ; symbol, Ir ; atomic weight, 193.1 ; 
 Platinum ; symbol, Pt ; atomic weight, 195. 
 
 THE six elements named above are all extremely rare, and, with 
 the exception of platinum, of little practical importance ; they fall 
 into two groups represented by the horizontal lines in the following 
 table, and also into three groups, represented by the vertical col- 
 umns, as follows : Ku> Kh pd 
 
 Atomic weight, 101.6 103. 106.6 
 
 Os. Ir. Pt. 
 
 Atomic weight, 190.8 193.1 195. 
 
 In passing from left to right in each of the two horizontal lines, 
 we find a gradation in properties similar to that observed in the group 
 composed of iron, cobalt, and nickel ; thus, ruthenium and osmium 
 are able to form higher oxides than the remaining four, just as iron 
 is able to form a higher oxide than either cobalt or nickel ; and the 
 same distinction in the formulae of the double cyanides which was 
 observed as existing between those of iron and cobalt on the one 
 hand and nickel on the other, is observed between ruthenium 
 rhodium, osmium iridium, and palladium platinum ; this dis- 
 tinction is made clear by the following table : 
 
 K 4 Fe(CN) 6 
 K 4 Co (CN) 6 
 K 4 Ru(CN) 6 
 K 4 Os(CN) 6 
 
 K 4 Fe(CN) 6 
 K 4 Co(CN) 6 
 K 4 Rh(CN) 6 
 K 4 Ir (CN) 6 ^ 
 
 Fe, Co; Rh, Ir. 
 
 Ni(CN) 2 , 2KCN 
 Pd(CN) 2 , 2KCN 
 Pt(CN) 2 , 2KCN 
 
 Fe, Co; Ru, Os. 
 
 Ni;Pd, Pt. 
 
490 
 
 ELEMENTS OF PLATINUM GKOUP. 
 
 The difference existing between the oxides may also, perhaps, 
 be best given in the form of a table. The highest oxides belong 
 to ruthenium and osmium; indeed, these are the only two ele- 
 ments which are able to form oxides of the formula M0 4 * in 
 which the valence of the element is presumably eight, provided 
 we regard oxygen as having a valence of two. The highest val- 
 ence toward oxygen displayed by any element, therefore, appears 
 in the eighth family. In the following table the formulae of the 
 oxides of the six elements under discussion are compared with 
 those of manganese, the element belonging to the preceding 
 (VII.) family; as will be noticed, there is a great resemblance, 
 although the highest valence of manganese toward oxygen is only 
 seven : 
 
 Mn. 
 
 Ru and Os. 
 
 Rh and Ir. 
 
 Pd and Ft. 
 
 Mn 2 O 7 
 (Mn0 3 )J 
 MnO 2 
 Mn 2 3 
 MnO 
 
 RuO 4 ; OsO 4 
 (Ru. 2 7 )t;- 
 (BuOa)t; (Os0 3 )t 
 RuO 2 ; OsO 2 
 Ru 2 3 ; Os 2 3 
 RuO; OsO 
 
 
 
 
 
 (IrO g ) 
 RhO 2 ; IrO 2 
 Rh 2 3 ; Ir 2 3 
 RhO; IrO 
 
 Pd0 2 ; Pt0 2 
 PdO; PtO 
 
 t The anhydride is not known; the salts, the per-ruthenites, M Ru O 4 , 
 corresponding to the permanganates, are known. 
 
 | The anhydrides are not known, but the salts, manganates, ruthenites, 
 and osmites are known ; their general formula is M 2 X O 4 , corresponding to the 
 .sulphates. An oxide Pd 2 O is also known. 
 
 One salt, K 2 Ir O 4 , derived from such an anhydride, has been described. 
 
 In the periodic system, the element following palladium is sil- 
 ver, and the one following platinum is gold ; as a consequence, we 
 find considerable resemblance between palladium and silver (for 
 instance, in the formation of the oxide Pd 2 0), and between plati- 
 num and gold. 
 
 The metals of the platinum group all possess a grayish-white 
 color and brilliant metallic lustre. They fuse at a high white heat, 
 the melting points decreasing in the two groups from ruthenium to 
 
 * An oxide of sulphur, S O 4 , is also mentioned. 
 
ELEMENTS OF PLATINUM GROUP ; OCCURRENCE. 491 
 
 palladium, and from osmium to platinum ; their specific gravities 
 and atomic volumes are as follows : 
 
 Eu. Eh. Pd. 
 
 Specific gravities, 12.261 12.1 11.4 
 Atomic volumes, 8.06 8.5 9.3 
 
 Os. Ir. Pt. 
 
 Specific gravities, 22.4 22.42 21.46 
 Atomic volumes, 8.55 8.6 9.08. 
 
 The elements of the platinum family are, therefore, at the mini- 
 mum of the curves of atomic volumes belonging to their respective 
 periods ; the next following elements, with larger atomic weights, 
 have larger atomic volumes. The platinum metals, are (like 
 iron, cobalt, and nickel) malleable and ductile, and form colored 
 salts. All of the elements which have small atomic volumes and 
 high specific gravities display several oxides and series of com- 
 pounds ; in none of these, however, is the chemical character of any 
 such element with small atomic volume so pronounced as to over- 
 shadow the character of the other elements in combination ; as has 
 been previously remarked, therefore, the crowding of a large amount 
 of matter into a small space seems to be unfavorable for manifesta- 
 tion of very decided metallic or not-metallic properties. All of the 
 elements belonging to the eighth family possess the power of con- 
 densing and transmitting hydrogen (see page 34). 
 
 The members of the platinum family all occur as the uncom- 
 bined metals, they are never, however, pure in their natural condi- 
 tion ; platinum may be found with only a slight admixture of iron, 
 but, as a general rule, alloys of the various metals forming the plat- 
 inum group (which may also contain silver and gold) occur. 
 
 The metals of this group are very difficult to dissolve in acids ; 
 ruthenium and iridium are not attacked by any acid ; platinum is 
 changed only by aqua regia ; osmium is dissolved by nitric acid as 
 well ; while palladium is soluble both in nitric and hydrochloric 
 acids ; all of the elements are very easily reduced and isolated from 
 their oxides or halides ; they all manifest a great tendency to unite 
 with chlorine. In their chlorine compounds, ruthenium, rhodium, 
 and palladium resemble iron, cobalt, and nickel, for they form 
 chlorides with the formulae M CL and M C1 3 ; on the other hand, 
 osium, iridium, and platinum have their most stable chlorides de- 
 
492 PLATINIC CHLORIDE. 
 
 rived from a tetravalent metal, with the general formula M C1 4 ; the 
 latter compounds have a great tendency to produce double salts 
 with the chlorides of the alkali metals ; these chlorides have the 
 composition M 2 R C1 6 ,* and may be regarded as derived from an acid 
 M 2 E, C1 6 , in which R C1 4 acts as an acidic anhydride ; the structure 
 of these acids would, therefore, be analogous to that of fluosilicic 
 acid (see page 303). 
 
 The higher oxides of this group (mentioned in the table, page 
 490) are acidic in their character, the lower ones are, at best, weakly 
 basic. The two tetroxides, Ru 4 and Os 4 ,f are quite volatile, and 
 are produced by powerful oxidation of the respective metals ; they 
 have a most peculiar, penetrating odor, which has been compared to 
 that of ozone ; $ the vapor densities correspond to molecules of the 
 formula R 4 . 
 
 Derived from a next lower, hypothetical, oxide, in which ruthe- 
 nium would be heptavalent, we have one salt, K Eu 4 ; osmium does 
 not furnish a parallel compound. Next to this we have salts of two 
 hypothetical acids, H 2 Ru 4 and H 2 Os 4 , which correspond to 
 manganic, chromic, and sulphuric acids ; the anhydrides, were they 
 known, w r ould have the formulae Ru 3 and Os 3 ; the salts are 
 termed ruthenites and osmites ; only the potassium salt of ruthenious 
 acid is known, but the potassium, sodium, and barium salts of osmous 
 acid have been isolated. 
 
 The dioxide, M0 2 , is the highest oxide which is common to all 
 of the platinum metals ; it is also the most important of the oxides 
 of this group, because the principal salts are derived from it, the 
 chief chlorides of this family having the formula M C1 4 . The dioxide 
 is basic in its character, although that of platinum is also slightly 
 acidic, for it dissolves in concentrated alkalies to form platinates. 
 The only chloride which need be mentioned in detail is that of 
 platinum ; its characteristics may serve as a type for all of the rest. 
 
 PLATINIC CHLORIDE, Pt C1 4 + 5 H 2 O, is prepared by adding silver nitrate 
 solution to the substance which is ordinarily termed platinic chloride, 
 but which is, in reality, an acid of the formula H 2 Pt C1 6 ; the first reac- 
 tion is as follows : 
 
 * M represents a univalent alkali metal, R either osmium, iridium, or 
 platinum. 
 
 t The latter is sometimes called osmic acid. 
 
 t The odor is almost exactly like that of chinon. 
 
CHLORPLATLNTC ACID. 493 
 
 H. 2 Pt C1 6 + 2 Ag NO 3 = Ag 2 Pt Cle + 2 HNO 3 ; 
 
 the silver salt of chlorplatinic acid, which is precipitated, is decomposed 
 by boiling water as follows : 
 
 Ag 2 Pt C1 6 = 2 Ag 01 + Pt C1 4 ; 
 
 the insoluble silver chloride is filtered, and the solution evaporated to 
 crystallization. 
 
 Chlorplatinic acid, H 2 Pt C1 6 + 6 H 2 0, is obtained by dissolving 
 platinum in aqua regia and cautiously evaporating ; it then forms 
 ochre-colored, deliquescent prisms. The chlorplatinates are of im- 
 portance in chemical analysis, because ammonium and potassium 
 chlorplatinates are nearly insoluble in water, and entirely insoluble 
 in alcohol. Potassium chlorplatinate is precipitated as a yellow, 
 crystalline substance, when potassium chloride, or any other potas- 
 sium salt, is added to a solution of chlorplatinic acid ; its formula is 
 K 2 Pt C1 6 . Similar results are obtained with ammonium salts. Am- 
 monium chlorplatinate, when heated, decomposes entirely, leaving 
 platinum in a very finely divided, spongy condition in which it is 
 known as " spongy platinum." Spongy platinum in a marked degree 
 possesses the power of occluding gases, and its use has been mentioned 
 in several of the preceding chapters (see, for instance, the prepara- 
 tion of sulphur trioxide, page 144).* The use of spongy platinum 
 depends on the fact that occluded gases are chemically much more 
 active than they are in the ordinary condition (see page 33). Pla- 
 tinic bromide or iodide forms acids similar to chlorplatinic acid. The 
 structure of these acids is similar to that of fluosilicic acid, H 2 Si F 6 , 
 which is also derived from a tetrahalide, Si F 4 ; so that in these 
 compounds, also, it is best to accept a theory of the divalence of the 
 halogen atoms (see pages 303 and 337). 
 
 Platinic chloride is able to unite with ammonia to form platinic- 
 amines, in which the chemical behavior of ammonia is identical with 
 that of the same substance in the cobaltamine salts. Platinic chlo- 
 
 * " Platinized asbestos " is formed by dipping asbestos in platinum chlo- 
 ride solution and then in one of ammonium chloride; ammonium chlorplati- 
 nate is thus precipitated on the fibres, and can be changed to spongy platinum 
 by heating. Palladium asbestos may be prepared in the same way. "Plati- 
 num black" is very finely divided platinum, formed by reducing a solution of 
 platinum chloride, to which potassium hydroxide has been added, by means 
 of alcohol. 
 
494 CIILOR-PLATINAMINES. 
 
 ride forms a number of these compounds, only two of which is it 
 necessary to mention here. 
 
 Platinic chloride can unite with two molecules of ammonia to 
 form chlor-platinamine chloride, which has the formula C1 2 Pt ( NH 3 ) 2 
 C1 2 : the radicle Pt ( NH 3 ) 2 being tetravalent and uniting with four 
 atoms of chlorine, just as one atom of platinum does in platinic 
 chloride : 
 
 Pt C1 4 , and Pt ( NH 3 ) 2 C1 4 , 
 
 Platinum chloride, and chlor-platinamine chloride. 
 
 However, in the latter compound, two atoms of chlorine can be 
 replaced by other acid groups (-N0 3 , or =S0 4 ) very readily, 
 while the other two are much less reactive; this compound may 
 therefore also be considered as the chloride of a divalent radicle 
 Cl 2 Pt (ISTH 3 ) 2 =. Because of the difficulty with which they are 
 replaced, it is supposed that the two less reactive chlorine atoms are 
 attached to platinum, while the other two are united to nitrogen in 
 the same way as one atom of chlorine is in ammonium chloride. 
 The structure of the ammonium compound is therefore better repre- 
 sented by the following formula : 
 
 ri p ( NH 3 Cl 
 
 ^1 2 r l> < T^-TT pi 
 
 and the nitrate and sulphate by the following : 
 N0 < UKlCLI 
 
 This theory is borne out by the fact that platinous chloride, Pt C1 2 , 
 can form a similar compound with ammonia, with the difference 
 that in this case two chlorine atoms, attached to platinum, are 
 missing ; this compound therefore has the structure : 
 
 (NH 3 C1 
 
 * JNH 3 C1. 
 
 In the second series of platinic amines, which will be described, 
 there are four molecules of ammonia for every atom of platinum. 
 The chloride in this series is therefore chlor-platindiainine chloride, 
 
 and has the formula : 
 
 , p <(NH 3 ) 2 C1 
 C1 Pt j(NH 3 ) 2 Cl. 
 
PLATINUM ; USES. 495 
 
 In this compound there exists the divalent radicle C1 2 Pt ( NH 3 ) 4 = , 
 and the similar combination derived from platinous chloride (having 
 two less chlorine atoms), has the formula : 
 
 p J(NH 3 ) 2 C1 
 t t(NH 3 ) 2 Cl. 
 
 One other oxide of this family should be mentioned. This is the 
 monoxide, MO, which, as well as the dioxide, is common to all the 
 members of the platinum group ; it is basic in its character, and from 
 it are derived the -ous salts. 
 
 Platinous chloride has the formula Pt C1 2 . It is formed by reduc- 
 tion of a solution of platinic chloride by means of sulphur dioxide ; 
 like platinum chloride, it readily forms double salts, and unites with 
 ammonia. 
 
 The commercial uses of the metals of the platinum group are 
 confined chiefly to the preparation of standard weights and meas- 
 ures, and to the manufacture of chemical apparatus. Most of the 
 platinum which occurs in crucibles and other utensils is not pure ; it 
 contains as much as 2 per cent of iridium, and, by very reason of 
 this admixture, is more valuable, because less readily attacked by 
 acids. Platinum vessels are extremely useful because of their high 
 melting point, because they are not attacked by oxygen, and are not 
 dissolved by the ordinary laboratory reagents. In using platinum 
 ware, care must be taken not to fuse caustic alkalies therein, nor 
 to heat with any metal or easily reducible compound containing a 
 metal, because platinum readily forms alloys. Furthermore, platinum 
 unites with silicon or phosphorus, the compounds so formed being 
 very brittle. Charcoal, or coal which contains silicon compounds in 
 its ash, should never be ignited in a crucible of platinum ; and the 
 .heating of compounds which contain phosphates in the presence of 
 a reducing agent (such as charcoal) should also be avoided. Lumi- 
 nous gas flames render platinum rough and brittle, the same is true 
 of the reducing flame of a Bunsen burner. Unless care is taken to 
 remove this roughness each time after using,* it will ultimately 
 penetrate the crucible, and, after rendering the utensil brittle, will 
 crack it. 
 
 By rubbing with a little fine sand until smooth. 
 
APPENDIX OF LABORATORY NOTES. 
 
 BEFORE entering upon laboratory work, the pupil should read and remem- 
 ber the following cautions : 
 
 Burns, Stains, and Fire. Yellow phosphorus should never be handled 
 excepting with a pair of tongs or pincers. When exposed to the air in a 
 warm room, it may take fire spontaneously; if touched by the hand, it will 
 take fire; the burns so produced are extremely painful, and may become 
 dangerous by reason of phosphorus poisoning. Sodium and potassium are 
 kept under naphtha; they should never be placed in water when the pieces 
 are larger than beans, and, in any event, a very small piece should first be 
 tested. Sodium which has not a clear and bright surface when cut, should be 
 rejected ; in all cases the outer coating of oxide should be cut away before 
 placing the metal in water. Burns are best treated by covering the spot with 
 a solution of cocaine in sweet oil, and then with an emulsion of lime-water, 
 glycerine, and sweet oil. Nitric acid stains the skin yellow; when concen- 
 trated, it will cause an ulcer to form. Bromine stains the skin brown; unless 
 instantly removed, it will cause a painful ulcer. Iodine stains the skin dark 
 brown ; nitrate of silver, black. Of course, every precaution should be taken 
 to keep the above substances from touching the hands or face ; but, in case of 
 accident, dilute sodium carbonate solution, followed by washing with clean 
 water, will be best to apply in the case of nitric acid and bromine. A solution 
 of sodium hyposulphite followed by water will remove iodine; cyanide of 
 potassium solution will remove silver stains.* Concentrated sulphuric acid 
 and solutions of chromic acid will attack the skin, not so rapidly, however, as 
 nitric acid ; in case of an accident they can often be removed by washing with 
 water or sodium carbonate solution before serious results have followed. Hot 
 sulphuric acid will instantly produce the most painful burns; any test-tube in 
 which sulphuric acid or anything else is being heated should be held by a test- 
 tube holder, and with its mouth pointing away from the manipulator or from 
 any one standing near. Ether or carbon bisulphide must not be used on a 
 desk within at least six feet of a burning Bunsen burner, as these liquids take 
 fire with the greatest readiness. Matches should be kept in a tin box, which 
 is never to be placed in the drawer of the desk, but should always be kept 
 outside. 
 
 * The teachers should not keep cyanide of potassium in a place where it is accessible 
 to any one but themselves; it should be handled under a hood, and should be applied in 
 removing silver stains, to the hands only, and then be instantly washed off after its work is 
 accomplished. 
 
 497 "* THS 
 
 UJUYIRSITYl 
 
498 APPENDIX OF 
 
 Inhalation of Fumes and Gases. Chemical experiments which will develop 
 poisonous or irritating gases should always be performed under a hood with a 
 good draught.* Chlorine, bromine, phosphorus pentachloride, phosphorus tri- 
 chloride, attack the mucous membrane of the eyes, throat, and nose ; continued 
 inhalation will give rise to bronchial inflammation; chlorine or bromine will 
 also cause nausea to ensue. If, by accident, the pupil should take an exces- 
 sive quantity of chlorine into the lungs, the quickest remedy is probably the 
 inhalation of the fumes of alcohol. The gaseous oxides of nitrogen are 
 poisonous, causing violent headache and nausea; phosphine, arsine, stibine, 
 are very poisonous ; ammonia is quite irritating. Work in which these sub- 
 stances are generated or used must be done under the hood.t Sulphuric acid 
 should not be heated to above 150 unless the apparatus is under the hood. 
 (The acid will break down, partly, into H 2 O and SO 3 ; the vapors of SO 3 are 
 irritating to the lungs.) Liquids containing hydrochloric, nitric, or hydro- 
 fluoric acids, etc., should be evaporated under the hood. Sulphuretted hydro- 
 gen is poisonous and disagreeable ; continued inhalation of even small quantities 
 will cause headache, and may have serious results. It is, therefore, imperatively 
 necessary, unless a room is especially provided in which to generate this gas, 
 that all work with hydrogen sulphide should be performed under the hood. 
 
 Explosions. The majority of accidents result from carelessness, there- 
 fore the invariable rule by which the student should govern himself in the 
 laboratory is, "never be careless, for carelessness may result in permanent 
 disfigurement or loss of sight." Hydrogen and oxygen, hydrogen and air, 
 hydrogen and chlorine, gaseous hydrocarbons and oxygen, phosphine and 
 oxygen, or phosphine and air, as well as the other not very stable hydrogen 
 compounds of the not-metals when mixed with oxygen or air, will, unless one or 
 the other constituent is present in proportionally small quantity, cause violent 
 explosions when they are ignited. In generating the gases, extreme care must 
 be taken not to bring a flame near the exit tube of the apparatus until the 
 pupil is sure that a brisk current of the generated gas has traversed the appa- 
 ratus for sufficient length of time to expel all air. Of course, no definite time 
 rule can be established, because this will vary with the size of the apparatus; 
 but when using the ordinary generating flasks of from 300 to 500 c.c. the pupil 
 should wait at least 7 to 10 minutes. Chlorate of potassium, permanganate of 
 potassium, and similar powerful oxidizers, must not be rubbed in a mortar when 
 in contact with substances which are readily oxidized (sugar, starch, sulphide 
 of antimony, sulphur, phosphorus [yellow or red], etc.). 
 
 In a well conducted laboratory, desks and apparatus are always kept as 
 clean as possible, and reagent bottles returned to their proper places as soon 
 as the occasion requiring their use is over. Bunsen burners can be cleaned by 
 
 * So urgent is this rule that pupils should be forbidden even to heat test-tubes or small 
 evaporating dishes with reagents which will give off fumes of hydrochloric acid, nitric acid, 
 hydrogen sulphide, bromine, chlorine, nitric oxide, etc., unless they do so under the hood. 
 A good hood is as necessary as a good burner. 
 
 t In the case of ammonia the precaution may be omitted, if only small quantities of the 
 reagent are to be used. 
 
LABORATORY NOTES. 499 
 
 unscrewing the outer tube and brushing the nipple with a dry, stiff test-tube 
 brush.* 
 
 The numbers of the following notes correspond to the reference 
 numbers in the text. 
 
 1. PREPARATION OF OXYGEN BY HEATING MERCURIC OXIDE. The 
 oxide decomposes at a low red heat. A little should be placed in a glass tube 
 300 m.m. in length, closed at one end and made of so-called infusible glass ; 
 outside of the tube, place a cylinder of copper wire gauze to prevent cracking; 
 a triple gas-burner is most convenient for heating. The tube is connected 
 with the trough of water, over which the gas is collected, by means of a safety 
 bottle ; the latter consisting of an empty 4 oz. w r ide-mouthed bottle, fitted with 
 a double bored rubber stopper and connecting glass tubes, the ends of which 
 must not extend below the bottom of the stopper. When such a safety bottle is 
 present, the water in the trough cannot be forced back into the hot tube, if, by 
 any accident, the flame should be extinguished ; for the cold water cannot get 
 beyond the safety bottle ; if it were to strike the hot tube, an accident would 
 be sure to follow. Such a safety bottle should always be interposed where a 
 pneumatic trough is used to collect gases which are generated in an apparatus 
 which is to be heated to a high temperature. (In the apparatus depicted, 
 Fig. 1 (Page 19), the place of a triple burner is supph'ed by a "combustion 
 furnace," which is a long oven heated by a number of flames.) 
 
 2. As BLACK OXIDE OF MANGANESE is sometimes adulterated with 
 charcoal, it is always necessary to test the chemical by heating a very little 
 of it in a test-tube before using a larger quantity; if no explosion results, it is 
 safe to use. 
 
 3. PREPARATION OF OXYGEN BY HEATING MANGANESE DIOXIDE. 
 The apparatus to be used is identical with that used in the experiment (Note 
 1), excepting that an iron tube, made of ordinary gas-pipe, capped and 18 
 inches in length, is substituted for the glass one. Fill this tube one-fourth full 
 of manganese dioxide, broken to the size of a pea.t Place the tube flat upon 
 the desk, and pound sharply after filling; this is for the purpose of making a 
 canal for the passage of the gas above the load in the tube. J Heat to a red 
 heat in a combustion furnace (Fig. 1), and use a safety bottle (Note 1). 
 
 4. THE CHLORATE OF POTASSIUM should be tested in the same way as 
 the black oxide of manganese (Note 2). 
 
 * Any instruction in glass bending or blowing which is necessary should be given by 
 the teacher before beginning laboratory work; after the instruction, practice alone will 
 make perfect. The pupil should buy W. A. Shenstone, Methods of Glass Blowing, Riving- 
 ton's, London, 1886. 
 
 t If this cannot be procured, take powdered magnanese dioxide. 
 
 This precaution must never be omitted in charging tubes with solid substances which 
 are to generate gases on heating; otherwise the apparatus will certainly explode. 
 
500 
 
 APPENDIX OF 
 
 5. PREPARATION OF OXYGEN BY HEATING CHLORATE OF POTASSIUM. 
 Take a flask of 200 c.c. capacity,* fitted with a single bored rubber stopper and 
 glass delivery tube (Fig. 15) and heat to a low red heat, collecting the gas 
 over water in a pneumatic trough ; it is best to insert a safety bottle between 
 the generating flask and the water (Note 1). As the flask in which this oper- 
 ation has been performed is always incapacitated for future use, and as the 
 operation is not used for the practical preparation of oxygen, it is better to 
 
 substitute the following: Heat 
 chlorate of potassium in a hard 
 glass test-tube until all of the 
 oxygen is expelled ; prove that 
 oxygen is present by placing a 
 glowing pine chip in the tube. 
 The most approved method of 
 preparing oxygen for laboratory 
 use is by heating a mixture of 
 chlorate of potassium and man- 
 ganese dioxide. Mix, in a mor- 
 tar, 25 grams of potassium 
 chlorate and 5 grams of man- 
 ganese dioxide; place the mix- 
 ture in a flask like the one 
 indicated in the first part of 
 this note and shown by Fig. 15 ; 
 heat gently with a Bunsen 
 burner until the gas conies off 
 slowly and regularly; collect 
 all of the gas in bottles over t 
 water in the pneumatic trough, 
 
 Fig. 15. 
 
 and save for future use. 
 
 6. COMBUSTION IN OXYGEN. Burn the substances mentioned in the 
 text, cutting phosphorus, sulphur, and carbon to the size of a pea. The de- 
 flagrating spoon in which these substances are burned (Fig. 16) should have 
 its handle thrust through a piece of sheet-zinc larger than the mouth of 
 the jar containing the gas and pierced with a small hole in the centre. It 
 is well to perform the burning of phosphorus in a large globe, care being 
 taken to have the deflagrating spoon sink so far into the vessel as to reach 
 the centre; if the jar is too small, the heat is apt to crack it. 
 
 7. COMBUSTION OF A STEEL WATCH-SPRING. The watch-spring should 
 
 * Round-bottomed flasks made of hard Bohemian glass have lately been brought into 
 the market; they are in every way more desirable than flat-bottomed or Erlenmeyer flasks. 
 
 t 32 oz. wide-mouthed (so-called salt mouth) common bottles are cheapest and best for 
 collecting gases. The pupil should have a number of square pieces of window glass larger 
 than the mouth of the bottle; when necessary to remove the latter, filled with gas, from the 
 trough, cover the mouth with one of the pieces of glass, by pressing the same against 
 the mouth of the bottle, raise the bottle up and invert it, still covered with the glass. 
 
LABORATORY NOTES. 
 
 501 
 
 be heated in the Bunsen burner and then straightened ; a small piece of cotton 
 is now tied at one end and dipped into molten sulphur; on igniting the sul- 
 phur, and plunging the spring into 
 a jar of oxygen, the heat given off 
 by the burning sulphur will cause 
 the iron to take fire. A cheap jar 
 should be employed ; because, during 
 the combustion, the temperature rises 
 high enough to melt the oxide of 
 iron, the small particles of which 
 fly off and become fused into the 
 walls of the vessel, and may even 
 break it. To illustrate the kindling 
 temperature, place a few drops of 
 dry carbon bisulphide in a test-tube, 
 warm until the tube becomes filled 
 with the vapors of the liquid, and 
 then place the end of a glass rod, 
 which has previously been warmed 
 in a Bunsen burner, within the 
 mouth of the test-tube; the carbon 
 bisulphide will take fire, thus furnish- 
 ing an illustration of a low kindling 
 temperature; the pupil can find ex- 
 
 amples of high kindling temperature 
 without suggestion. 
 
 Fig. 16. 
 
 8. PREPARATION OF HYDROGEN BY MEANS OF SODIUM AND WATER. 
 The method of preparation is described on page 28 of the text, and the 
 arrangement of the apparatus is made clear by figure 4 (page 29). If no wire 
 spoon is at hand, the piece of sodium, which should be cut to about the size 
 of a bean, may be wrapped in a small piece of copper gauze,* and this piece 
 then taken up by a pair of crucible tongs and slid under the mouth of the 
 test-tube, which has previously been filled with water and inverted in the 
 trough. Care should be taken to test the sodium to be used in any of these ex- 
 periments by placing a small piece on water and then standing aside ; for, un- 
 less the metal is clean, there is great danger of an explosion. Scraps of sodium 
 which have been kept in the laboratory for some time should never be used. 
 
 9. The method for preparing hydrogen by means of sodium and water is 
 expensive, and not to be recommended for ordinary purposes. It has been 
 used, however, where perfectly pure hydrogen is required. 
 
 * Filter-paper will even answer the purpose. If copper gauze is used, the sodium and 
 gause will sink to the bottom, the sodium will melt after a time, and, escaping through the 
 meshes of the gauze, will rise; care must be taken to place the tube so as to catch the parti- 
 cles of metal as they rise. If paper is used, the sodium and paper will rise to the top of the 
 tube, and the generation of hydrogen will go on at the point. 
 
502 
 
 APPENDIX OF 
 
 10. THE DECOMPOSITION OF WATER BY POTASSIUM. The piece of 
 potassium to be used should be cut about the size of a pea (handle neither 
 potassium nor sodium with the fingers!), and thrown on the water in the 
 pneumatic trough. The decomposition of water by potassium is so violent 
 that the heat generated sets fire to the hydrogen evolved, the latter burning 
 with a violet flame. Care must be taken to stand at some distance from the 
 water on which the potassium is floating, as an explosion may occur by means 
 of which the pieces of the metal will be thrown about. Such pieces may cause 
 painful burns, and, if in contact with the eyes, possible loss of sight. 
 
 11. GRANULATED ZINC. Prepared by fusing zinc in a stoneware cru- 
 cible, and pouring the melted metal in a thin stream into cold water, care 
 being taken to stir the water vigorously during the operation. 
 
 12. To PREPARE HYDROGEN BY MEANS OF ZINC AND DILUTE SUL- 
 PHURIC ACID.* Place 5 grams of zinc in a gas-generating flask fitted with 
 
 a double bored rubber stop- 
 per and delivery and safety 
 tube (Fig. 17) ;t now pour 
 dilute sulphuric acid through 
 the safety tube on to the 
 zinc, adding acid from time 
 to time as occasion requires. 
 Dilute sulphuric acid is pre- 
 pared by adding one part of 
 commercial acid to six parts 
 of water; in diluting sul- 
 phuric acid, pour the acid 
 gradually into the water, but 
 do not pour the water into 
 the acid ; cool the acid before 
 using, by placing the flask 
 under the hydrant. 
 
 13. IN ORDER TO PU- 
 RIFY THE HYDROGEN pre- 
 pared as in note 12, it should 
 be passed through a train of 
 wash-bottles (Fig. 18). In 
 the first is a solution of po- 
 Fig, 17. tassium permanganate to re- 
 
 move gases which can be 
 oxidized; viz., hydrocarbons, hydrogen sulphide, hydrogen arsenide, etc.; in 
 
 * Be sure to use zinc which is free from arsenic, or which, at best, contains only a trace 
 of that element. 
 
 t A thick walled flask of 300 c.c. capacity can be used for this and subsequent opera- 
 tions for generating a gas, provided there is no necessity of heating the flask. 
 
LABORATORY NOTES. 
 
 503 
 
 the second, potassium hydroxide, to remove acid vapors; and in the third, 
 concentrated sulphuric acid, to 
 remove moisture. Such a train 
 of bottles can be applied wher- 
 ever it is necessary to purify a 
 gas, provided, always, the sub- 
 stances used in them are varied 
 according to the nature of the 
 latter, so that no decomposition 
 of the gas can result. A judi- 
 cious choice of these washing 
 agents comes with larger expe- 
 rience. 
 
 14. WHEN LARGER QUAN- 
 TITIES OF GAS ARE REQUIRED, Fig. 18. 
 it is better to use a form of gas 
 
 generator depicted in Fig. 19. A vessel, 6 d, constricted in the middle, is 
 fitted with a glass stopcock delivery tube, e, and a wide globe funnel, with 
 long stem, is placed in the upper opening. The zinc 
 is placed in 6, and the acid added from above until 
 
 the apparatus is filled to about 
 
 the middle of the funnel; on 
 
 opening the stopcock the acid 
 
 ascends to the metal ; on clos- 
 ing, the generated hydrogen 
 
 once more expels the acid 
 
 from the central globe. In 
 
 this way the metals can be 
 
 kept indefinitely out of con- 
 tact with the acid, and need 
 
 only be acted on by it when 
 
 the stopcock e is opened. This 
 
 form of gas-generating ap- 
 paratus is known as Kipp's 
 
 gas generator. Several other 
 
 forms have been devised, but 
 
 this one seems to be the most 
 
 satisfactory ; it is useful for 
 
 generating any gas which does 
 
 Fig. 19. 
 
 Fig. 20. 
 
 not require heat in its manufacture. Never pour hot 
 acid into a Kipp generator, nor ever pick it up other- 
 wise than by grasping it with both hands around the 
 central globe. 
 
 Gases which are desired for future use are stored in a gasometer (Fig. 20). 
 A lower metal or glass tank, a, holding about forty litres, is connected with an 
 upper one, 6, by means of two tubes, one of which reaches to the bottom of a. 
 
504 
 
 APPENDIX OF 
 
 The gasometer is filled with water, all stopcocks closed, the cap covering the 
 bottom opening is removed, and the gas to be stored is run in through this; 
 when all the water has been replaced, the cap is screwed on. When the gas is 
 to be used, the upper tank is filled with water, the stopcock on the tube lead- 
 ing to the bottom is opened, and the gas allowed to escape through the 
 upper side opening, as required. 
 
 15. DIFFUSION OF HYDROGEN THROUGH A 
 POROUS SUBSTANCE. Construct an apparatus 
 such as is shown in Fig. 21. This consists of a clean, 
 dry, porous cup,* which is fastened at the end of 
 the tube c by means of a funnel ; this funnel has 
 exactly the diameter of the mouth of the cup, and 
 the two are fastened together, air-tight, by means 
 of rubber cement placed around the rims. The 
 tube c is connected with one opening of a double- 
 necked flask t by means of a single bored rubber 
 stopper, while a glass tube, drawn to a point and 
 reaching to the bottom of the flask, is connected 
 with the other opening, also by means of a rubber 
 stopper. When all the connections are air-tight, 
 place the porous cup in an atmosphere of hydro- 
 gen, by inverting over the cup a glass bell, open at 
 one end and connected with a hydrogen gener- 
 ator; the hydrogen will rush through the porous 
 cup much more rapidly than the air can escape, 
 for the specific gravity of air is 14.4 times that of 
 hydrogen. As a consequence, the air in the cup 
 and in c is forced down as if by a piston; this 
 causes a pressure on the water J which has previ- 
 ously been placed in the flask, and a fountain en- 
 sues. For an apparatus to illustrate effusion quan- 
 titatively, see Freer; Zeitschrift fur Physikalische 
 Chemie; IX. 669. 
 
 16. OCCLUSION OF HYDROGEN. Palladium 
 is very expensive ; one gram will suffice to show the 
 occlusion of hydrogen. Attach the palladium to a 
 platinum wire, heat red-hot in a Bunsen burner, 
 
 allow to cool slightly, and then place in a current of hydrogen passing from 
 a generator ; the occluded hydrogen will be oxidized by the air, the palladium 
 will once more begin to glow, and will finally heat the hydrogen to its kindling 
 temperature. So-called platinized asbestos is cheaper and just as available for 
 this experiment. Prepare platinized asbestos by heating asbestos which has 
 been dipped in a solution of platinum chloride and then into ammonium 
 
 * A small porous cup from a Bunsen battery. 
 
 t So-called Woulff's bottle. 
 
 t For lecture-room demonstrations, color the water with blue litmus. 
 
 Take care to have all of the air expelled from the generator! 
 
 Fig. 21. 
 
LABORATORY NOTES. 
 
 505 
 
 chloride. A little of this may be fastened, by means of fine platinum wire, within 
 a small loop made of iron wire, to which a handle about four inches in length 
 is attached. Be sure to perform the simple experiment given on the bottom of 
 page 34, and the last one on the first paragraph of the same page. 
 
 17. TO PROVE THAT HYDROGEN FORMS WATER WHEN BURNED IN AlR. 
 
 Dry the gas passing from the generator by means of the train (Note 13), or 
 by passing it through a glass U-tube filled with coarse fragments of brick soaked 
 with concentrated sulphuric acid, or through a similar tube filled with frag- 
 ments* of granulated calcium chloride (Fig. 22). 
 
 18. Care should be taken to expel all of the air from the apparatus before 
 lighting. To find out if this has been accomplished, test the gas by collecting a 
 test-tube full over water in a pneumatic trough, and lighting it; if it burns 
 quietly, the apparatus is safe. The burner can be made by drawing a glass tube 
 nearly to a point, inserting a small cylinder of rolled platinum foil in the 
 small end so produced, and then fusing the glass around the platinum ; after 
 
 the hydrogen burner has 
 
 been lighted, place a large 
 
 cold beaker over the flame, 
 and water will collect there- 
 in. 
 
 19. AN EXPLOSIVE 
 MIXTURE OF HYDROGEN 
 AND OXYGEN, This can 
 
 be prepared with safety by Fig 2 2. 
 
 using a small bottle, with- 
 out a bottom and with a narrow neck; to the latter, by means of a single bored 
 rubber stopper and short glass tube, a long rubber tube is fitted. This tube 
 can be closed by means of a pinchcock. The bottle, after the rubber tube is 
 closed by the pinchcock, is filled with water in the pneumatic trough, and then 
 hydrogen is run in until | of the water has been expelled ; the remaining J is 
 then similarly replaced by oxygen. The mixture of gases can be expelled by 
 lowering the bottle in the pneumatic trough and opening the pinchcock. By 
 placing the end of the rubber delivery tube under some soap-suds in a small 
 tin dish, a few soap-bubbles filled with oxy-hydrogen can be produced, and 
 then exploded by touching with a lighted taper. Care should be taken to 
 remove this dish to a safe distance from the bottle before exploding, and what- 
 ever oxy-hydrogen is left should be allowed to escape as soon as the experi- 
 ment requiring its use has been completed. If possible, some experiments 
 with the oxy-hydrogen blow-pipe (page 35) should be performed. 
 
 20. THE VOLUMETRIC COMPOSITION OF WATER. Use a eudiometer 
 holding 50 c.c. This instrument is a heavy glass tube (Fig. 23), which is closed 
 at one end and graduated in millimetres (sometimes in cubic centimetres), and 
 has two platinum wires inserted near the closed tip in such a manner that an 
 
 The size of a pea. 
 
506 
 
 APPENDIX OF 
 
 Fij. 23. 
 
 electric spark can pass from one to the other.* This tube is filled with mer- 
 cury and inverted over a mercury trough. In order to prove the volumetric 
 composition of water, slant the tube to one side, and admit about lOc.c. of 
 
 hydrogen, then add 7 c.c. of oxygen ; 
 clamp the tube tightly with its open 
 end forced against a leather washer 
 in the bottom of the mercury trough, 
 taking care to note the exact volume 
 of hydrogen admitted, the volume of 
 oxygen, the temperature, barometer, 
 and height of the mercury in the 
 tube. After all of these prepara- 
 tions are completed, a spark from 
 an induction coil is passed through 
 the gas. After the explosion is 
 completed t and the apparatus has 
 cooled, note the volume of gas re- 
 maining, the temperature, and the 
 barometric pressure, as well as the 
 height of the column of mercury in 
 the tube. Reduce the gas before and 
 after explosion to standard condi- 
 tions by means of the formula on 
 page 173, remembering that the height of the column of mercury in the tube, 
 before and after explosion, must be deducted from the barometric pressure. If 
 the pupil used 10 c.c. of hydrogen and 7 c.c. of oxygen, then the 10 c.c. of 
 hydrogen will have united with 5 c.c. of oxygen to form water, and 2 c.c. of 
 oxygen will remain. The decomposition of water by the electric current is made 
 clear by Fig. 3. The two electrodes are pieces of platinum foil ; the water to 
 be decomposed is acidulated with sulphuric acid. 
 
 21. THE COMPOSITIOX OF WATER BY WEIGHT. Dumas passed hydro- 
 gen through a series of U-tubes (Fig. 22) filled as follows, counting from his 
 hydrogen generator (or gasometer): 
 
 U-tube No. 1 ; glass fragments moistened with lead nitrate solution, to remove 
 hydrogen sulphide. 
 
 U-tube No. 2; glass fragments moistened with silver sulphate solution, to 
 remove hydrogen arsenide. 
 
 U-tube No. 3; pumice-stone moistened with caustic potash solution, to remove 
 carbon dioxide, and other acids, such as sulphurous acid, hydrochloric acid, etc. 
 
 U-tubes No. 4, 5; pieces of solid caustic potash, to remove acids. 
 
 U-tubes No. 6, 7, 8; phosphoric anhydride, to remove moisture. 
 
 The last tube he weighed before and after the operation ; if it had 
 
 * Not infrequently the eudiometers come to the laboratory with the wires so close 
 -together that the spark will not be large enough to ignite the gases; if such is the case, force 
 the ends apart with a long glass rod. 
 
 ~t Wrap a towel around the eudiometer tube before exploding. 
 
LABORATORY NOTES. 
 
 507 
 
 changed in weight at all, the experiment was rejected, as the hydrogen was 
 
 not pure. The hydrogen 
 
 so purified was passed 
 through a tube of hard 
 glass, containing copper 
 oxide ; this tube was evac- 
 uated by an air-pump 
 before beginning the op- 
 eration, and weighed 
 
 Fig. 22. 
 
 carefully ; it was then 
 connected at one end 
 with the apparatus delivering hydrogen which was purified as above ; at the 
 other end, with two U-tubes filled with phosphorus pentoxide ; the tubes were 
 carefully weighed before the operation. The hydrogen was now passed over 
 the copper oxide in the tube, which was heated in a combustion furnace 
 (Fig. 1) until the oxide was reduced to metallic copper. After the reduction 
 was complete, and the tube had completely cooled, it was once more evacuated 
 and weighed ; the loss in weight was equal to the weight of oxygen used to 
 form water. The two phosphorus pentoxide tubes, which were placed after the 
 copper oxide, were also weighed; the gain in weight was equal to the amount 
 of water which had been formed. The pupil should perform this operation 
 with a simpler apparatus, which can consist of the train of wash-bottles 
 (Note 13, Fig. 18) filled (counting from the hydrogen generator), No. 1 with 
 potassium permanganate solution, No. 2 with lead nitrate solution, No. 3 
 with caustic potash solution,* and lastly a U-tube filled with pieces of brick 
 soaked in sulphuric acid. The copper oxide tube should be of hard glass, with 
 a bulb blown on the middle, and one end drawn to a narrow opening. To this 
 end, after filling the bulb with granulated copper oxide (the copper oxide 
 should previously be dried at 150 and introduced into the tube while still 
 warm), and weighing the tube on the balance (weigh tOy^ of a milligram), 
 attach a U-tube filled with pieces of brick soaked in sulphuric acid (weigh 
 this tube also to T a ^ of a milligram). Now pass the hydrogen through the 
 apparatus, while heating the copper oxide to a low red heat. After reduction, 
 is complete, continue to pass hydrogen for some time, so that every trace of 
 water is carried over into the last U-tube ; allow to cool, weigh, and calculate 
 the results as given on page 38. Unless the ratio is within of .05 a unit of the 
 proportion 1 : 8, the experiment should be repeated. The connections be- 
 tween the various parts of the apparatus are made by means of rubber 
 tubing ; be sure to have this small enough to fit air-tight over the ends of the 
 tubes, and make the connections as short as possible. 
 
 To PROVE THE PRESENCE OF WATER OF CRYSTALLIZATION. Heat a 
 crystal of copper sulphate in a test-tube, and see if water passes off, whether 
 at a high or low temperature ; next dissolve the anhydrous salt in water, 
 
 * The wash-bottles cannot be filled with very much of the required solutions, otherwise 
 the pressure in the hydrogen generator will not be sufficient to overcome the hydrostatic 
 pressure in the wash-bottles ; the gas would, in consequence, not flow through the bottles. 
 
508 APPENDIX OF 
 
 evaporate in a porcelain dish on a water-bath * until crystallization begins, 
 set aside to cool, and observe the form of the crystals. Note the change in 
 temperature caused by dissolving sodium nitrate in water, and by dissolving 
 fused calcium chloride in water; take a small piece of quick-lime and place it 
 in a porcelain dish, then pour on a little water and note the change in 
 temperature. (See pages 43, 44.) An example of an efflorescent salt can be 
 obtained by procuring a few "soda crystals" (Na 2 CO 3 -f 10 H 2 O) ; a deli- 
 quescent substance is fused calcium chloride. 
 
 22. PREPARATION OF OZONE. The apparatus, Fig. 24, is the one 
 generally employed for obtaining considerable quantities of ozone. An outer 
 
 coating, a, is 
 of tinfoil, sur- 
 rounding a 
 glass tube; 
 within an 
 inner glass 
 
 L ' tube (c) is a 
 
 Fig . 24 piece of thin 
 
 copper foil ; 
 
 the two glass tubes are fused together at &, so that c is within a and reaches 
 nearly to the end of the latter farthest from 6. Dry oxygen is admitted to 
 the space between a and c by the tube at 6; it passes out at the farther end; a 
 is connected by means of a metal strip with one pole of an induction coil ; c, 
 by means of a similar metal strip placed at 6, with the other pole. When a 
 current passes, a silent discharge takes place between the tin foil surround- 
 ing a and the copper contained in c. This discharge must necessarily traverse 
 the oxygen which is passing in at 6, and by this means ozone is generated. 
 The preparation of ozone by the student can be accomplished by placing a 
 few pieces of phosphorus in a good sized bottle with a wide neck, covering 
 them partly with a very dilute aqueous solution of potassium dichromate and 
 .-sulphuric acid, warming the whole slightly (to about 24 about the tempera- 
 ture of a hot room in summer), and shaking gently from time to time. The 
 bottle will soon contain ozone ; the latter can be detected by the odor and by 
 hanging a strip of paper, soaked with a mixture of starch paste and a little 
 potassium iodide, within the neck of the bottle (page 50). 
 
 23. The blue color of ozone may be seen by filling a glass tube, 1 metre in 
 length, with ozonized oxygen and then looking through this against a white 
 background. 
 
 24. PREPARATION OF A SOLUTION OF HYDROGEN DIOXIDE. Add finely 
 powdered barium dioxide, gradually, to dilute sulphuric acid until a sufficient 
 quantity has been added to convert all of the sulphuric acid into barium sul- 
 
 * A copper basin, hemispherical in shape, with the top formed of concentric, movable, 
 copper rings. This vessel is filled with water and heated with a Bunsen burner, substances 
 to be evaporated are placed in porcelain evaporating dishes upon it. Remember that a 
 flame merely large enough to cause the water to boil is just as good for heating purposes as 
 one which will cause violent ebullition, and which will necessitate frequent refilling of the 
 water bath. 
 
LABORATORY NOTES. 
 
 509 
 
 phate,* taking care to keep the liquid quite cool during the operation, as 
 warmth destroys hydrogen hyperoxide ; the barium sulphate is allowed to 
 settle; the nearly clear, supernatent liquid is poured off and finally filtered. 
 To a portion of the colorless liquid so formed, add a mixture of starch paste 
 and a little iodide of potassium solution, and observe the blue color of iodine- 
 starch (page 50); to another portion add a dilute solution of potassium di- 
 chromate and a little sulphuric acid; an intensely blue color (ascribed to 
 the formation of perchromic acid) will be seen; shake the blue solution in the 
 test-tube with a little ether, t and the blue substance will be dissolved by 
 the latter. 
 
 25. The solution of hydrogen peroxide is concentrated by placing it in 
 an open dish under the bell of an air-pump which also has under it an open 
 dish containing concentrated sulphuric acid; the air is then exhausted, and the 
 solution allowed to evaporate. This is a common method for concentrating 
 liquids which are easily decomposed. 
 
 26. ELECTROLYSIS OF A SOLUTION OF HYDROCHLORIC ACID. The 
 apparatus to be used is shown by Fig, 25 ; it explains itself, with the exception 
 that the two elec- 
 trodes within the 
 
 U-shaped tube must 
 be made of gas car- 
 bon and not of plat- 
 inum. Such an 
 apparatus may be 
 especially ordered 
 for this experiment, 
 or the apparatus' 
 given by Fig. 3 can 
 be used, the two 
 platinum electrodes 
 being removed, and 
 in their places two 
 small pieces of gas 
 carbon, the con- 
 nections made by 
 platinum wire, can 
 be inserted. In- 
 stead of an ordi- 
 nary solution of 
 hydrochloric acid, 
 a solution of com- Fig. 25. 
 
 mon salt, saturated 
 
 * This point can be ascertained by allowing the solution to settle, taking out a drop of 
 the clear liquid, and adding this to a drop of barium chloride solution on a watch-glass; if no 
 insoluble barium sulphate is formed, the sulphuric acid has all been converted to barium 
 sulphate. 
 
 t By pouring a little ether over the liquid, placing the thumb over the mouth of the test- 
 tube, and shaking. 
 
510 
 
 APPENDIX OF 
 
 at ordinary temperatures, to which ^ of its volume of concentrated hydro- 
 chloric acid has been added, is used. When the electric current is turned on, 
 the chlorine will at first be absorbed by the liquid which fills the apparatus; 
 after the latter is entirely saturated with the gas, then equal volumes of 
 hydrogen and chlorine will be produced simultaneously. Care must be taken 
 not to open the stopcocks too frequently while the operation is going on, 
 otherwise the liquids in the two arms of the U-shaped tube become mixed, and 
 the action irregular. Too much liquid must not be placed in the apparatus, as, 
 otherwise, the increase in pressure, caused by the liberation of the gases and 
 consequent rising of liquid in the rear tube, will cause a greater absorption of 
 chlorine. 
 
 27. THE PREPARATION OF CHLORINE. The apparatus best adapted 
 for the preparation of all gases which are either noxious or poisonous is 
 shown by Fig. 26. When the gas is desired for use, the stopcock B is opened 
 
 and A is closed ; the gas 
 can then pass from the 
 generating-flask into the 
 vessel or apparatus where 
 it is to be used. When the 
 
 gas is not required, the 
 stopcock B is closed and 
 A is opened ; then it passes 
 from the generating-flask 
 into a vessel filled with 
 some liquid which will 
 completely absorb it. For 
 the preparation of chlorine 
 use a generating - flask 
 (round bottomed preferred) 
 of about 500 c.c. capacity; 
 place in this 50 grams of 
 manganese dioxide (pow- 
 dered); unite all parts of 
 the apparatus, and then 
 pour in a sufficient quan- 
 tity of hydrochloric acid 
 through the safety tube to 
 cover the manganese dioxide ; then warm gently with a Bunsen burner ; or, 
 mix in a mortar 50 grams of manganese dioxide, 50 grams of sodium chlo- 
 ride, place in the generating-flask, connect all parts of the apparatus, and 
 then add 150 c.c. of sulphuric acid (two parts sulphuric acid to one of water), 
 and warm gently. The pupil should prepare chlorine by both of these 
 methods. Collect the chlorine in dry bottles by displacement of air, as is 
 shown in the cut, and, when full, cover the bottles with glass plates, and place 
 aside for future use ; afterwards pass chlorine into 300 c.c. of water until 
 the liquid is saturated with the gas. Form chlorine hydrate crystals, as indi- 
 
 Fig. 26. 
 
LABORATORY NOTES. 
 
 511 
 
 cated on page 62 (paragraph 1). Into the jars filled with chlorine intro- 
 duce : 
 
 a. A little powdered antimony. 
 
 b. A few pieces of heated copper foil. 
 
 c. A piece of moist litmus paper. 
 
 d. A piece of filter paper, moistened with turpentine. 
 
 Try the bleaching power of your chlorine water on a piece of colored calico. 
 Experiments with chlorine must be conducted under a hood ! 
 
 28. BLEACHING BY MEANS OF CHLORINE. This can best be shown by 
 passing dry chlorine * through a flask of 1 litre capacity which has been fitted 
 with a triple bored rubber stopper; connect one of the glass tubes introduced 
 through this with the chlorine generator; connect the second with a small 
 flask containing water, so that the latter can be boiled when desired, so as to 
 force steam into the apparatus ; place a glass tube in the third hole of the 
 triple bored stopper (this last is for the escape of the superfluous gases); 
 it can, if desired, be so bent as to open into a small jar containing caustic 
 soda solution. Introduce a piece of colored calico into the 1 litre flask, and 
 admit dry chlorine: no bleaching action will 
 
 be observed. Now heat the water in the small 
 flask and force in steam: the calico will then 
 be bleached. 
 
 29. PREPARATION OF HYDROCHLORIC 
 ACID. The apparatus to be used is the 
 same as that for the preparation of chlorine 
 (Fig. 26). The generator is charged with 20 
 grams of sodium chloride, and then 30 grams 
 of sulphuric acid (two parts of acid to one of 
 water) are added through the safety tube; 
 heat very gently, and collect two jars of the 
 gas as was done with chlorine, and pass the 
 remainder of the gas into a beaker contain- 
 ing water. 
 
 Hydrochloric acid is produced by burning 
 chlorine in an atmosphere of hydrogen. In 
 a small flask (A, Fig. 27) of 150 c.c. capacity, 
 place 10 grams of powdered manganese diox- 
 ide. Fit the flask with a single bored rubber 
 stopper into which is inserted a tube, widened 
 at the centre, and with its widened part filled 
 with pieces of granulated calcium chloride 
 about the size of a pea; after pouring con- 
 centrated hydrochloric acid on the manganese 
 
 Fig, 27. 
 
 dioxide, put this stopper into the flask, allow the chlorine to escape for a time, 
 and then invert a jar, #, of hydrogen over the escaping chlorine, taking care to 
 
 * Dry the chlorine by passing it through a U-tube containing pieces of brick saturated 
 with cone, sulphuric acid (Note 21); 
 
512 
 
 APPENDIX OF 
 
 light the hydrogen just before it reaches the chlorine jet; the chlorine will 
 then burn in the hydrogen which fills the jar. The reverse, burning hydrogen 
 in chlorine, can be performed by filling the flask A with zinc and dilute sul- 
 phuric acid, and lengthening and bending the delivery tube, so that it may 
 form a burner which can extend down into a jar of chlorine ; after hydrogen 
 has expelled all of the air from the apparatus, light the jet, and lower the 
 flame into the chlorine jar; it will continue to burn. In both experiments 
 fumes of hydrochloric acid will be observed. 
 
 30. EXPERIMENTS WITH HYDROCHLORIC ACID. Place a strip of moist- 
 ened blue litmus paper in one of the jars of hydrochloric acid gas ; place a 
 lighted taper in the other, and see if the gas supports combustion. To a little of 
 your solution in a test-tube add some iron filings ; some pieces of zinc ; dilute 
 the solution and taste it; try its effects on blue litmus solution. To some dilute 
 hydrochloric acid, to which you have added a few drops of litmus solution, 
 add a solution of sodium hydroxide, drop by drop from a burette,* until the 
 red color of the litmus just turns into blue ; one drop of acid will then turn the 
 blue litmus red, a drop of sodium hydroxide will turn the color back to blue; 
 the solution is then neutral. If it is evaporated, nothing but sodium chloride 
 will remain ; that sodium chloride is formed can be proved by adding sulphuric 
 acid to the crystallized remainder. (Pages 76 and 77.) 
 
 31. The electrolysis of hydrochloric acid is described in Note 20. 
 
 The extreme solubility of hydrochloric acid in water may be shown by an 
 apparatus such as is depicted by Fig. 28. The upper flask is well filled with 
 dry hydrochloric acid gas, and is stoppered with a single bored rubber stopper ; 
 
 through the stopper there runs a 
 glass tube, drawn to a point, and 
 sealed without the flask, and nar- 
 rowed within. The large beaker 
 is filled with blue litmus solution. 
 When all is ready, the sealed point 
 is broken under the water. As the 
 fountain is somewhat slow in start- 
 ing, a small bulb of water, sealed, 
 may be placed within the flask, 
 and, when the apparatus is to be 
 used, can be broken by a quick 
 shake. The water will absorb hy- 
 drochloric acid, create a partial 
 vacuum, and the blue litmus solu- 
 tion will run in. 
 
 32. DECOMPOSITION OF HY- 
 DROCHLORIC ACID INTO ONE VOL- 
 UME OF CHLORINE AND ONE OF 
 HYDROGEN. The apparatus to 
 
 Fig. 28. 
 
 * A burette is a graduated glass tube about 500 m.m. in length, one end of which termi- 
 nates in a narrow tube, closed with a glass stopcock; a measured quantity of liquid can there- 
 fore be run out of this instrument by opening the stopcock at the bottom. 
 
LABORATORY NOTES. 
 
 513 
 
 be used is shown by Fig. 29; pure, dry hydrochloric acid gas is introduced 
 through the glass stopcock of the apparatus (which has previously been filled 
 with dry mercury) until one arm of the U-shaped tube is about two-thirds full, 
 the mercury being allowed to run off at the lower tap as fast as hydrochloric acid 
 enters through the upper stopcock; when enough gas has 
 been admitted, the upper and lower stopcocks are closed, 
 the mercury being placed at the same level in both arms. 
 The hydrochloric acid gas is now under atmospheric pres- 
 sure; the level of the mercury is therefore marked by a 
 rubber ring. Now sodium amalgam * is poured into the 
 open arm until the arm is quite full, a rubber stopper is 
 inserted tightly, and then the apparatus is shaken from 
 side to side, so that the sodium amalgam comes in con- 
 tact with the gas; the hydrochloric acid is decomposed 
 after the lapse of about one minute. The gas which 
 remains is now carefully brought back into the arm of the 
 apparatus which it originally occupied, the rubber stop- 
 per removed, and the mercury run out through the lower 
 tap until on the same level in both arms; it will then be 
 seen that the volume of hydrogen left is exactly one-half 
 of that which was occupied by the hydrochloric acid. 
 Take care to have the apparatus, sodium amalgam, and 
 hydrochloric acid perfectly dry before beginning the 
 experiment; for sodium amalgam, acting on water, will 
 generate hydrogen. Fi 9- 2g - 
 
 33. THE PREP- 
 ARATION OF BRO- 
 MINE. Take a 
 tubulated and 
 glass-stoppered re- 
 tort ( Fig. 30) of 200 
 c.c. capacity; place 
 in this 10 grams of 
 manganese dioxide 
 and 20 grams of 
 potassium bromide 
 (these constituents 
 previously mixed 
 in a mortar), clamp 
 the retort to a re- 
 tort stand, as shown 
 
 * To prepare sodium amalgam, take 500 grams of dry mercury, place in a large clay cru- 
 cible, and cover with an iron dish; now cut 7 grams of clean sodium into pieces the size of a 
 hickory nut, and throw these into the crucible; to start the reaction, heat a little mercury in 
 a test-tube, raise the cover on the crucible a little, and pour in the hot mercury; an instan- 
 taneous reaction, accompanied by a flash of light, will occur; during this, stand aside so as 
 not to inhale fumes of mercury; after all is over, stir the sodium amalgam with an iron wire; 
 place, when cool, in a wide-mouthed glass-stoppered bottle, and keep for future use. Pre- 
 pare sodium amalgam under the hood. 
 
 Fig 30. 
 
514 
 
 APPENDIX OF 
 
 in the figure, and thrust the neck of the retort far into the neck of a receiver of 
 500 c.c. capacity. This receiver is cooled by means of a stream of water, the 
 escape of which is provided for by a battery jar with a syphon placed under the re- 
 ceiver; the syphon should be connected with a sink. When all is ready, add 
 sulphuric acid (one part sulphuric acid to one part of water) to the manganese 
 dioxide and potassium bromide, put in the glass stopper of the retort, and seal 
 the latter tightly by means of a little plaster of paris and water; allow the 
 retort to stand for ten minutes, and then heat gently; the bromine will distil 
 and collect in the receiver. The operation should be performed under the 
 hood! Make a solution of bromine in water, and try its bleaching power as 
 you did with chlorine; fill a tube about 500 m.m. in length, closed at one end, 
 with chlorine water, and invert over a beaker glass which is partly filled with 
 the same liquid, and, after properly supporting the tube, place the whole in 
 the sunlight; do the same with bromine water, placing the apparatus filled 
 with the latter beside that containing chlorine water, and, after 24 hours, 
 notice the amount of oxygen separated by each halogen (see page 79). 
 
 Fig. 31. 
 
 34. PREPARATION OF HYDROBROMIC ACID. To a few particles of so- 
 dium or potassium bromide, in a test-tube, add a few drops of concentrated 
 sulphuric acid, and warm very gently; note color and odor of the gas which is 
 passed off. Pure hydrobromic acid cannot be prepared in this way. The 
 apparatus for the preparation of hydrobromic acid for laboratory use is shown 
 by Fig. 31. The flask A, which has a tube fused into the side of the neck,* 
 * A so-called fractional distilling-flask. 
 
LABORATORY NOTES. 515 
 
 has a capacity of 300 c.c. It is fitted with a single bored rubber stopper, into 
 which is inserted a drop funnel. The side tube connects with a glass tube, B, 
 which is about 300 m.m. in length, and which is filled with pieces of brick which 
 have been moistened and rolled in red phosphorus ; the upward slant to this 
 tube is necessary to prevent pieces of phosphorus and impure water from 
 being carried over into the funnel tube, C. The latter is interposed between the 
 tube B and the water which is in the beaker glass, solely to prevent the latter 
 from "sucking back" as soon as the current of hydrobromic acid becomes so 
 feeble that solution of the gas in water takes place so rapidly that the gas, 
 which is being generated, is unable to keep the liquid out of the apparatus. 
 By reason of the interposition of the funnel tube, the water can be forced back 
 no farther than the tube, provided its mouth is placed so as to just touch the 
 surface of the liquid in the beaker. Charge the generating-flask, A, with 25 
 grams of red phosphorus, and then add just barely enough water to cover the 
 latter* fill the drop funnel above half full of bromine (hood !), connect all parts 
 of the apparatus, and allow the bromine to fall on the phosphorus slowly, drop 
 by drop (care!); each drop of bromine will cause a flash of light and the for- 
 mation of hydrobromic acid. After all of the water in the flask, A, and the 
 tube, J5, has been saturated with hydrobromic acid, the acid will escape into the 
 funnel tube, C, and can be collected in a beaker of water, or, as the dry gas, by 
 downward displacement, as is done with chlorine. White crystals of phospho- 
 nium bromide (see page 217) will ultimately clog the apparatus if too little 
 water is present. This may become so serious as entirely to prevent the flow of 
 gas through B, in which case an explosion will inevitably result. Prevent this 
 accident by adding a few drops of water to the generator A as soon as you see 
 phosphonium bromide crystals forming. Perform the same experiments with 
 the solution of hydrobromic acid as you did with hydrochloric acid. Fill one 
 jar with dry hydrobromic acid gas, and invert a jar of chlorine of the same 
 size over it; observe the separation of bromine and the rate of diffusion of 
 the gases. 
 
 35. PREPARATION OF IODINE. Perform this operation with the same 
 apparatus and same proportions as you used in the preparation of bromine; 
 you can omit the cooling of the receiver ; of course, substitute potassium 
 iodide or sodium iodide for the bromide which you used in Note 33. 
 
 36. PREPARATION OF HYDROIODIC ACID. The apparatus is the same 
 as that used for the preparation of hydrobromic acid (Fig. 31), with the ex- 
 ception that the bricks, covered with moist red phosphorus, can be omitted 
 from the tube J?, as all of the iodine which may pass over from the generating 
 flask A will be condensed by the cold glass tube. In charging the apparatus, 
 place 50 grams of iodine in the generating flask A, add to this 10 grams of 
 water ; replace the drop funnel used in the hydrobromic experiment by an 
 ordinary funnel, which can be stoppered by a glass rod thrust into the neck 
 and ground into the tip so as to be water-tight ; through this funnel, after all 
 
 * In preparing hydrobromic acid the greatest care must be taken not to add too much 
 water, otherwise the liquid will dissolve the hydrobromic acid as fast as it is generated. 
 
516 
 
 APPENDIX OF 
 
 the connections of the apparatus are tight, gradually drop upon the iodine 
 red phosphorus, which has been stirred with water to a thick paste. By 
 following these directions the generation of phosphonium iodide, which 
 invariably results if the apparatus is arranged exactly as in the preparation 
 of hydrobroinic acid, is avoided. Care must be taken not to add the phos- 
 phorus too rapidly, otherwise an explosion would result. The proportions 
 which are most successful for the preparation of hydroiodic acid are : 10 parts 
 of iodine, 5 parts of phosphorus, and 2 parts of water. 
 
 37. THE PREPARATION OF SULPHUR. The formation of sulphur by 
 the action of hydrogen sulphide on sulphur dioxide. The apparatus is shown 
 
 Fig. 32. 
 
 by Fig. 32. The flask C, with two lateral tubulures, has 500 c.c. capacity ; in 
 the flask, A place about 10 grams, of copper shavings, stopper with a double 
 bored rubber stopper into which is inserted a safety tube and delivery tube, 
 add 100 grams of concentrated sulphuric acid through the safety tube, and 
 connect with C, as shown in the cut. The double-necked (Woulff) wash- 
 bottle -B, which contains a little water, is connected with a hydrogen sulphide 
 generator. The latter is the same as that used for hydrogen (Note 12), with 
 the exception that the zinc is replaced by 10 grams of ferrous sulphide ; the 
 flask C has its mouth emptying in a beaker, containing a solution of sodium 
 hydroxide for the purpose of absorbing the excess of gases. When all is 
 
LABORATORY NOTES. 
 
 517 
 
 ready, connect all parts of the apparatus, and heat the copper and sulphuric 
 acid in A, until the ebullition indicates that the formation of sulphur dioxide 
 has begun, then remove the flame ; next add dilute sulphuric acid to the ferrous 
 sulphide in the generating-flask connected with B. The sulphur dioxide 
 and hydrogen sulphide will meet in C ; after a time, small drops of plastic 
 sulphur will collect ; the latter soon become opaque and yellow, changing into 
 ordinary rhombic sulphur.* 
 
 38. THE DISTILLATION OF SULPHUR. Bend a test-tube, and clamp it 
 to a retort stand, as is shown by Fig 33. Place 3 or 4 grams of sulphur in the 
 test-tube and heat; collect the vapors in a 
 
 beaker of cold water. Endeavor to ob- 
 serve all of the phenomena mentioned on 
 page 92, and observe the formation of 
 flowers of sulphur on the surface of the 
 cold water. 
 
 39. THE CRYSTALLIZATION OF SUL- 
 PHUR. Rhombic sulphur. Take some 
 dry carbon bisulphide, t and place it in a 
 test-tube ; add roll sulphur to this and then 
 stopper; allow to stand until the carbon 
 bisulphide has taken up as much sulphur 
 as it will ; now pour the clear solution into 
 a second clean test-tube, stopper the latter 
 with some cotton, and put aside in a quiet 
 place. After a time, fine, transparent 
 crystals of rhombic sulphur will separate. 
 
 Monoclinic sulphur. Melt 100 grttms of sulphur in a flat porcelain evaporat- 
 ing-dish, and then allow to cool until a crust has formed over the surface; 
 perforate this crust by means of a glass rod before the entire mass becomes 
 solid, and then pour off the sulphur which has not solidified, through the hole 
 which has been formed. The bottom of the evaporating-dish will be covered 
 with colorless, transparent needles of monoclinic sulphur crystals. 
 
 40. THE PREPARATION OF HYDROGEN SULPHIDE. The apparatus 
 used is the same as that employed to generate hydrogen (Fig. 34). In the 
 generating-flask place 20 grams of ferrous sulphide, broken to about the size 
 of a bean; connect the delivery tube with one of the bottles from the train 
 used in drying hydrogen (this bottle should contain a little water, so as to 
 retain any acid fumes which may pass over) ; after all connections are made, 
 pour dilute sulphuric acid through the thistle tube. Collect the sulphuretted 
 hydrogen which passes off, as you did with chlorine and hydrobromic acid, by 
 displacement of air, in dry bottles (such as you used for hydrogen and oxygen); 
 after three bottles have been filled in this way, pass the remainder of the gas 
 
 Fig. 33. 
 
 * It takes very nearly an hour to form a good deposit of sulphur. 
 
 t Dry the carbon bisulphide by shaking it with fused calcium chloride. 
 
518 
 
 APPENDIX OF 
 
 into a bottle filled with water. Apply a lighted taper to one of the bottles 
 
 filled with dry gas ; place a 
 strip of filter paper which 
 is soaked with a solution of 
 lead acetate in the second ; 
 pass chlorine from the small 
 chlorine generator, which 
 you prepared for the experi- 
 ment in Note 29, into the 
 third. Take the solution of 
 hydrogen sulphide which 
 you have prepared, and add 
 a little of it to a solution of 
 copper sulphate ; to a solu- 
 tion of cadmium nitrate; to 
 a hydrochloric acid solution 
 of arsenic trioxide ; to a 
 slightly acid solution of stan- 
 nous chloride* (see page 09 
 of the text). Experiments 
 with hydrogen sulphide 
 must be conducted under a 
 hood ! 
 
 41. THE DECOMPOSI- 
 TION OF HYDROGEN SUL- 
 PHIDE BY HEAT. This 
 
 decomposition can best be shown by the apparatus depicted in Fig. 35. 
 This consists of a flask of about 1 litre capacity, well 
 stoppered with a double bored rubber stopper ; into 
 this two glass tubes are fitted, through which run two 
 pieces of tolerably thick copper wire, the tips of the 
 glass tubes being fused around these wires; the ex- 
 tremities of the copper wires are connected by a 
 platinum wire as shown in the cut. Fill the flask 
 completely with gaseous hydrogen sulphide, insert the 
 stopper, and pass an electric current through the plat- 
 inum wire by attaching the free ends of the copper 
 wires to the two poles of a battery (the current should 
 be just sufficient to cause the wire to glow) ; the 
 hydrogen sulphide will decompose at the line of con- p . 
 
 tact of the hot wire. 
 
 The experiments leading to the preparation of the oxygen acids and salts 
 of chlorine, bromine, and iodine are made sufficiently clear in the text, and 
 need give no difficulty, as the pupil is now able to prepare the halogens and 
 
 * The pupil can select solutions of the salts of other metals, and study the solubility in 
 acids and in alkalies of the sulphides produced by the addition of hydrogen sulphide. 
 
 Fig. 34. 
 
LABORATORY NOTES. 
 
 519 
 
 can use them in the formation of the various compounds mentioned in 
 chapters xviii. and xix. In performing these experiments, he should confine 
 himself to the preparation and reactions of those salts which are formed by add- 
 ing chlorine, bromine, and iodine to solutions of potassium hydroxide, and to the 
 formation of calcium hypochlorite by passing chlorine over slaked lime, and to 
 the decomposition of calcium hypochlorite by hydrochloric and sulphuric acids. 
 
 42. THE PREPARATION OF 
 SULPHUR DIOXIDE. The appa- 
 ratus is shown in Fig. 36. The 
 generating-flask should be of 500 
 c.c. capacity; in this place about 
 twenty grams of copper shavings, 
 connect all parts of the apparatus, 
 add about 100 c.c. of concentrated 
 sulphuric acid through the safety 
 tube, and heat by means of a Bun- 
 sen burner ; when the gas begins 
 to pass off, lower the flame so 
 as to secure a regular evolution. 
 The wash-bottle contains concen- 
 trated sulphuric acid. Collect three 
 jars of the gas by displacement of 
 air, and pass the remaining gas into 
 water. See if the gas will burn or 
 will support combustion; put a moist 
 strip of colored calico into one of the 
 jars, and a small red rose,* which 
 has been moistened, into another. 
 To portions of the solution of sul- 
 phur dioxide in water add, succes- 
 sively, a dilute alcoholic solution of Fig. 36. 
 iodine ; a solution of bromine ; a solution of ferric chloride ; a solution of po- 
 tassium dichromate. (All of these reagents 
 will illustrate the reducing power of sulphur 
 dioxide.) 
 
 43. LIQUID SULPHUR DIOXIDE. The 
 most convenient form in which to use sulphur 
 dioxide in the laboratory is as a liquid. The 
 apparatus, shown by Fig. 37, is placed, with its 
 stopcocks open, in a dish filled with a mixture 
 of pounded ice and salt,t and then a slow cur- 
 rent of sulphur dioxide, generated as explained 
 in Note 42, is passed through. In order to cool 
 the gas perfectly it should, after leaving the 
 drying-flask, traverse a spiral glass tube which 
 
 37. 
 
 * Roses are best to use for bleaching with sulphur dioxide; many other red flowers 
 bleach very slowly, some not at all. 
 
 t In making a freezing mixture, do not spare the salt. 
 
520 
 
 APPENDIX OF 
 
 is placed in a jar and cooled with snow and salt. When a sufficient quantity 
 of the gas has been liquefied, all of the stopcocks are closed, and the apparatus 
 is set aside until wanted. When gaseous sulphur dioxide is required, the 
 stopcock c is opened, and a portion of the gas run into the small bulb; c is 
 then closed, and by opening b the gas can be used without interfering with 
 the liquid in the large bulb. 
 
 44. THE MANUFACTURE OF SULPHURIC ACID. A laboratory apparatus 
 for illustrating the manufacture of sulphuric acid is shown in Fig. 38. A 
 
 glass globe (Y!) of about 5 
 litres capacity is fitted with a 
 rubber stopper through which 
 five holes are bored. Three 
 small flasks (a, b, c) are con- 
 nected with this by means of 
 glass tubes extending to the 
 middle of the large one, and 
 through the remaining holes 
 two tubes, also extending to 
 the middle of the large flask, 
 are fitted. Sulphur dioxide 
 is generated in a by heating 
 copper shavings and sulphu- 
 ric acid, nitric oxide in b by 
 Fig ' 38 ' means of copper and dilute 
 
 nitric acid (no heat is required), and steam is supplied, as wanted, by heating 
 water in c. Air can be forced in at d by connecting the tube with a bellows, 
 and the tube e, connected with a hood, is left as a vent-hole for the escape of 
 gases. When nitric oxide comes in contact with air, it is oxidized to a mix- 
 ture of NO 2 and N 2 O 3 ; so that, after that change has taken place, we have 
 the gas present which is necessary for the formation of sulphuric acid from 
 sulphurous acid. If very little steam is admitted, we can easily see the forma- 
 tion of nitrosyl-sulphuric acid; for the large globe becomes covered with frost- 
 like crystals of that substance. If an excess of steam is admitted, these 
 crystals disappear and sulphuric acid is formed, the latter collecting as an oily 
 liquid. By varying the amounts of nitric oxide, sulphur dioxide, air, and 
 steam, we can study all of the phases of sulphuric acid manufacture. 
 
 EXPERIMENTS WITH SULPHURIC ACID. Take some concentrated sulphu-/ 
 ric acid, and add it to a little cane sugar which you have placed in a test-tube; 
 stir the mixture with a glass rod, and allow to stand under the hood ; add 9.8 
 grams of concentrated sulphuric acid to 1.8 grams of water, place the liquid in 
 a small flask, and surround the latter with crushed ice and salt; crystals of 
 H 4 SO 3 will form; warm the H 4 SO 5 until melted, and once more add 1.8 grams 
 of water, place in the freezing mixture, and crystals of H 6 SO 6 will separate ; 
 melt the H G SO 6 which you have made, and then gradually add more water, 
 and notice if there is any increase of temperature.* Add some dilute sulphu- 
 
 * In this experiment it will be necessary to add rooter to sulphuric acid ; be sure to hold 
 the flask at a safe distance while pouring in the water, otherwise the heat which is generated 
 might cause the water to boil and spatter drops of sulphuric acid. 
 
LABORATORY NOTES. 
 
 521 
 
 ric acid to a solution of barium chloride; to a solution of strontium chloride; 
 
 and to two solutions of calcium chloride, the first of which is very dilute, the 
 
 second tolerably concentrated, and note the result (see page 416). 
 
 45. To ISOLATE NITROGEN FROM THE ATMOSPHERE. The apparatus is 
 
 shown by Fig. 30. Take a bell jar of 3 litres capacity, invert it over a glass 
 
 basin, or over your pneumatic trough; prepare 
 a float made of a flat cork, on one side of which 
 you have fastened a porcelain crucible cover by 
 forcing the handle firmly into the cork; place 
 this cork, with the crucible cover up, in your 
 pneumatic trough ; place a piece of phosphorus 
 the size of a bean on the cover, and light the 
 phosphorus with a hot wire (care in handling 
 phosphorus!); now invert the bell jar over the^ 
 float, and slightly raise the stopper at the top, 
 so that the gas within* 
 which necessarily ex- 
 
 Fig. 39. 
 
 pands very greatly (owing to the heat given off by the 
 burning phosphorus), can quietly escape. If this pre- 
 caution is omitted, the air will be forced out at the bot- 
 tom of the jar in large bubbles; the disturbance may 
 even tip over your phosphorus float. After the violent 
 combustion is over, insert the stopper of the bell jar and 
 allow the gas to cool ; water will rise in the jar to take 
 the place of the oxygen which has gone to form phos- 
 phorus pentoxide, and also to take the place of the air 
 which has been expelled. To test the gas remaining, 
 add enough water to the pneumatic trough to make the 
 level within and without the bell jar alike, and then in- 
 troduce a lighted taper.* 
 
 46. THE COMPOSITION OF THE ATMOSPHERE. A 
 crude method of determining the volumetric composi- 
 tion of the atmosphere is by means of the apparatus 
 shown by Fig. 40. Divide a long glass tube, closed at 
 one end, into five equal parts by means of rubber rings. 
 Invert this over a long cylinder containing water so that 
 the level without and within is at the first ring, and then 
 clamp the tube in place, Now fix a piece of phosphorus 
 on a long, sharp-pointed copper wire, taking care not to 
 touch the phosphorus with the hands, bend the wire as 
 shown in the cut, thrust the phosphorus up into the 
 tube, and set the apparatus aside for two days. The 
 oxygen of the enclosed air will then be entirely ab- Flg " 40 " 
 
 sorbed; and, by sinking the tube so that the level of the water without and 
 within is the same, the amount of nitrogen in the air can be ascertained. By 
 
 * The sides of the bell jar become coated with a red amorphous solid during this opera- 
 tion; this substance is a sub-oxide of phosphorus with the probable formula P 4 O; it is not 
 red phosphorus, as is generally supposed. 
 
522 APPENDIX OF 
 
 noting the height of the barometer before and after the experiment, and then 
 applying the necessary corrections, quite accurate results can be obtained ; but, 
 if such are required, a carefully graduated tube must be substituted for the 
 crudely divided one indicated. 
 
 In order to measure accurately the relative amounts of oxygen and nitro- 
 gen in the atmosphere, the eudiometer (Fig. 23, Note 20) is employed. The 
 instrument should have a capacity of 100 c.c. ; it is partially filled with mer- 
 cury, and inverted over the mercury trough so that about 25 c.c. of air will 
 remain enclosed ; about 14 c.c. of hydrogen are now run in, by slanting the 
 tube and placing the delivery tube of a hydrogen apparatus (which is gen- 
 erating pure and dry hydrogen) under its mouth; take all of the precautions 
 mentioned in Note 20, and ignite the mixture of gases with an electric spark; 
 be sure to read accurately the volume of air and the volume of hydrogen 
 before the explosion, and also to measure the height of the column of mercury 
 as indicated in Note 20; after the explosion read the volume of remaining gas 
 and reduce to standard Conditions, exactly as*was done before. The hydrogen 
 will have united with the oxygen to form water; therefore, one-third of the 
 volume, by which the mixture of the gases which were enclosed in the eudio- 
 meter has diminished, must have been oxygen.* 
 
 47. THE PRESENCE OF CARBON DIOXIDE IN THE ATMOSPHERE. 
 Take one of the train of wash-bottles which was used in drying hydrogen, 
 clean it, and fill it with clear lime-water, and attach to a Bunsen aspirator 
 which you have fastened to the hydrant,! and the suction tube of which is 
 connected with your wash-bottle in such a way that air will be drawn through 
 the apparatus in the same direction as hydrogen is forced through it, as indi- 
 cated in Fig. 18. The lime-water will soon become turbid, owing to the 
 formation of calcium carbonate ; if you use two wash-bottles, each of which 
 contains clear lime-water, then the one into which the air first passes will 
 become turbid, while the second will remain clear. The presence of moisture 
 in the atmosphere can be shown by exposing an open beaker, containing con- 
 centrated sulphuric acid, to the air; after some days the volume will be 
 observed to have increased, while the acid has become diluted with water. 
 
 48. THE PREPARATION OF AMMONIA. The apparatus for the prepara- 
 
 * If the temperature and barometer are the same before and after the experiment, no 
 correction need be made excepting for the difference caused by the change in the pressure 
 on the enclosed gas, brought about by the differing height of the column of mercury in the 
 eudiometer; the latter must be carefully noted, 1st, when air alone is in the eudiometer, 2d, 
 when hydrogen has been admitted, 3d, after the explosion. As the air contains an unknown 
 amount of moisture, saturate it with water vapor by admitting a drop of water above the 
 mercury in the eudiometer; and then, unless temperature and barometric pressure are differ- 
 ent after the experiment than they were before, there will be no necessity for paying atten- 
 tion to the amount of water present, for the gases will be saturated with water before and 
 after the experiment (see page 173). 
 
 t A Bunsen glass aspirator is a cheap instrument which aspirates air by using the water 
 pressure of a hydrant; it is indispensable for laboratory work, and should be kept in the 
 desk, ready for use. It is attached by means of rubber tubing, and its mechanism will 
 explain itself when the instrument is handled. A more effective, but also more expensive, 
 instrument is a Chapman brass aspirator; this can also be fastened to the hydrant; both 
 forms of aspirator can be purchased of any instrument dealer. 
 
LABORATORY NOTES. 
 
 523 
 
 tion of ammonia is shown by Fig. 41. The generating-flask is of 500 c.c. 
 
 capacity; in it are 
 placed 50 grams of 
 ammonium chloride 
 and 100 grams of 
 slaked lime (pre- 
 pare slaked lime by 
 slowly pouring 
 water on quick-lime 
 until the latter fi- 
 nally crumbles to a 
 powder) ; connect 
 the generating-flask 
 with a drying-cylin- 
 der, the latter being 
 filled with small 
 pieces of quick- 
 1 i m e, which are 
 present for the pur- 
 pose of drying the 
 gas; collect ammo- 
 Fig. 41. nia over mercury or 
 
 by displacement of 
 
 air in a jar which is held mouth downward, for ammonia is specifically lighter 
 than air. When all connections are made, add enough water to the mixture 
 of ammonium chloride and slaked lime in the flask to cause the latter to roll 
 into lumps on shaking; now heat gently and ammonia will pass off. The 
 drying-cylinder is a so-called " Fresenius " drying-tower, at the bottom open- 
 ing of which the gas enters; at the top, after traversing the intermediate space 
 filled with pieces of quick-lime, it escapes. These drying-towers are very 
 convenient for laboratory use, and, if possible, should be kept on hand. 
 Ammonia cannot be dried over calcium chloride, because it combines with that 
 substance ; it is self-evident that the gas cannot be dried by sulphuric acid. 
 
 49. EXPERIMENTS WITH AMMONIA. The combustion of ammonia may be 
 shown by filling a large test-tube with oxygen, then passing in some ammonia 
 gas, so as partially to displace the oxygen, and then quickly approaching the 
 mouth of the test-tube to a gas flame ; a weak explosion will follow.* Pass 
 some of the ammonia, which was generated by the experiment mentioned in 
 Note 48, into water, and, with the solution of ammonia so produced, perform 
 the following experiments : Place some ammonia in a beaker glass; add a 
 few drops of blue litmus solution, and then carefully add hydrochloric acid 
 from a burette, until the solution is neutral (see Note 30). Pour the liquid in 
 the beaker into an evaporating-dish, and evaporate to dryness on a water bath 
 (Note 21); heat a little of the salt which separates, on a piece of platinum 
 
 * Take care to wrap a towel around the test-tube before bringing its mouth to the 
 flame. 
 
524 
 
 APPENDIX OF 
 
 Fig. 42. 
 
 foil; heat some of the salt with slaked lime and water in a test-tube; repeat 
 the same experiments, substituting nitric acid and sulphuric acid for hydro- 
 chloric acid. 
 
 The absorption of ammonia by charcoal can be demonstrated by the appara- 
 tus shown by Fig. 42. This is a test-tube filled with dry ammonia and 
 inverted over a basin of mercury; a small piece of 
 charcoal, which has previously been glowed out in a 
 Bunsen burner, is introduced, and the test-tube is then 
 clamped with its mouth under the surface of the mer- 
 cury; the ammonia will be absorbed, and the mercury 
 will rise in the tube. 
 
 50. THE SOLUBILITY OF AMMONIA IN WATER. 
 Use' the same apparatus which you employed for 
 demonstrating the solubility of hydrochloric acid (Note 
 31, Fig. 28), filling the flask with dry ammonia gas by 
 displacement of air. 
 
 51. AMMONIUM AM A La AM. In a narrow cylinder 
 of 225 c.c. capacity, place 20 grams of ammonium chlo- 
 ride, add enough water so as just to cover 
 
 the salt, and then pour on sodium amalgam, 
 prepared as indicated in the foot-note to 
 
 Note 32: the ammonium amalgam will begin to form at once, and 
 the operation can be hastened by gently stirring with a glass rod. 
 
 52. THE PREPARATION OF NITROUS 
 OXIDE. The apparatus which it is best to 
 use for the preparation of this gas is shown 
 by Fig. 43. This consists of a 300 c.c. flask, 
 stoppered with a single bored rubber stopper 
 which is connected with a delivery tube and 
 safety bottle. In this flask, place 10 to 15 grams 
 of crystallized ammonium 
 nitrate,* and heat to a tem- 
 perature just sufficiently high 
 to cause a regular flow of the 
 gas ; collect over the pneumatic 
 trough by displacement of 
 water. Introduce a lighted 
 taper into a jar filled with 
 nitrous oxide, and see whether 
 it burns or supports combus- Fig. 43. 
 
 tion; repeat the experiments 
 
 given in Notes 6 and 7, using nitrous oxide instead of oxygen ; inhale a little 
 of the gas. 
 
 * Test a little of your ammonium nitrate by heating in a test-tube before you proceed to 
 the decomposition of larger quantities of the salt. 
 
 
 
 
 
LABORATORY NOTES. 
 
 525 
 
 53. PREPARATION OF NITRIC OXIDE. Use the apparatus shown by 
 Fig. 44, charge the flask with 20 grams of copper shavings, cover with water, 
 and slowly add ordinary nitric acid, waiting for the reaction to begin after 
 
 each addition; after all the 
 brown fumes which are at 
 first developed have disap- 
 peared, collect the gas over 
 the pneumatic trough by dis- 
 placement of water. Do not 
 heat. 
 
 54. EXPERIMENTS WITH 
 NITRIC OXIDE. Take one 
 of the bottles which you have 
 filled with nitric oxide and 
 turn it mouth upward; the 
 brown fumes of the higher 
 oxides of nitrogen will ap- 
 pear where contact with the 
 air takes place; repeat the 
 experiments illustrating com- 
 bustion which you performed 
 with nitrous oxide and with 
 oxygen, and satisfy yourself 
 as to whether nitric oxide 
 supports combustion as read- 
 ily afis either of those two 
 gases; place a few drops of 
 carbon bisulphide in one of your bottles of nitric oxide, after you have re- 
 moved it from the pneumatic trough, and covered its mouth with a piece of 
 glass plate, taking care not to 
 admit any more air than is ab- 
 solutely necessary while pouring 
 in the carbon bisulphide; now 
 shake the bottle back and forth 
 three or four times, keeping it 
 closed ; carefully allow the ex- 
 cess of carbon bisulphide to leak 
 out at the place where you hold 
 the glass cover on the bottle, and 
 then, while quickly removing the 
 cover, approach the mouth of 
 the bottle to a gas flame. (Pages 
 200 and 201.) 
 
 55. PREPARATION OF NI- 
 TROGEN PEROXIDE FROM NI- 
 TRIC OXIDE. The apparatus Fig ' 45 ' 
 
 Fig. 44. 
 
526 APPENDIX OF 
 
 is shown by Fig. 45. Fill with nitric oxide a glass flask of 500 c.c. capacity, 
 tubulated at one side and having a long neck ; invert the flask over a basin 
 of water so that its mouth is under the liquid ; connect the tube which is fitted 
 to the side tubulure with an oxygen gasometer (Note 14), and, by opening the 
 wire pinchcock which closes the rubber tube, admit a little oxygen. Brown 
 fumes of nitrogen peroxide will at once appear; the latter gas is, however, 
 rapidly absorbed by the water in the neck of the flask, the following reaction 
 taking place : 
 
 3 NO 2 + H 2 O = 2 HNO 3 + NO (see page 198). 
 
 After the gas in the glass globe has become colorless, owing to the disap- 
 pearance of the peroxide and the regeneration of a portion of the nitric 
 oxide, add a little more oxygen, wait for the absorption of the brown fumes 
 again, and repeat the experiment until all of the nitric oxide has been used 
 up; the water from the trough will then have completely filled the globe.* 
 
 The preparation of nitrogen peroxide by heating lead nitrate. The appa- 
 ratus is the same as that used for the preparation of oxygen by heating mercuric 
 oxide (Note 1, Fig. 1), excepting that the gas must be collected by displacement 
 of air, and not over water. Fill the hard glass tube one-quarter full of a mixture 
 of equal parts of sand and lead nitrate (the latter ground fine in a mortar), 
 pound the side of the tube sharply on the desk so as to form a canal for the 
 escape of the gas (Note 3), and heat in the combustion furnace, taking care 
 to decompose the lead nitrate at the rear end of the tube first; and, when this 
 has yielded all of the gas which it is capable of doing, advance the flame grad- 
 ually toward the mouth of the tube. By passing the nitrogen peroxide through 
 the apparatus used in condensing sulphur dioxide (Note 43, Fig. 37), it can be 
 obtained as a straw-colored liquid. Nitrogen trioxide can be liquefied by the 
 same means, the fluid being indigo blue in color. Prepare the trioxide by heating 
 a mixture of ordinary nitric acid and a little arsenious oxide t in the apparatus 
 which you used for the preparation of sulphur dioxide, after removing the 
 wash-bottle and replacing the latter by the gas-condensing apparatus, which is 
 well cooled by means of pounded ice and salt. 
 
 56. THE PREPARATION OF NITRIC ANHYDRIDE. The preparation of 
 this substance is not infrequently attended with danger; the reaction leading to 
 its formation is, therefore, scarcely to be attempted either in the laboratory 
 or in the lecture-room. Place concentrated nitric acid in a tubulated retort of 
 about 500 c.c. capacity, and then gradually add phosphoric anhydride until the 
 mixture of that solid with the nitric acid has formed a jelly-like mass; cool 
 the retort during this operation so that the temperature never exceeds 0; now 
 heat very gently on the water-bath, never allowing the temperature to exceed 
 60; distil the liquid which passes off into a receiver (arranged as in the 
 
 * Of course, if a globe which is tubulated is not at hand, an ordinary flask can be used; 
 the oxygen can then be run in under the water by using a bent delivery tube. 
 
 t If you have none but powdered arsenic trioxide, be careful to add it very gradually to 
 the nitric acid, and heat very gently, as the reduction of the nitric acid is apt to become 
 quite violent; if the porcelain-like variety of arsenious oxide is at hand, it is much better to 
 use that form. 
 
LABORATORY NOTES. 
 
 
 
 \ 
 
 ' 
 
 527 
 
 preparation of bromine, Fig. 30), which is kept cool by means of ice and salt; 
 the nitric anhydride will then solidify. Do not keep nitric anhydride for any 
 length of time ! 
 
 57. THE PREPARATION OF NITRIC ACID. By passing electric sparks 
 through moist air. Take a eudiometer tube (Fig. 23, Note 20), stopper it with 
 a rubber stopper, connect its platinum wires with a battery and induction coil, 
 and allow electric sparks to pass through for about an hour; the tube will 
 then be seen to be filled with brown fumes if it is held against a white back- 
 ground, and a little blue litmus solution introduced into the tube will be 
 turned red.* 
 
 58. THE SAME. By heating sodium nitrate and sulphuric acid. The 
 apparatus to be used is shown by Fig. 46. It is identical with that employed 
 
 Fig. 46. 
 
 for the preparation of bromine (Fig. 30). t In the retort, place 50 grams of 
 sodium nitrate, make all connections, and then add fifty grams of concen- 
 trated sulphuric acid; warm gently until drops of liquid begin to pass over, 
 and then endeavor to keep the retort about the temperature of distillation. 
 After a time a crystalline remainder will form in the retort; wash this out, 
 evaporate in an open dish, and investigate the nature of the crystals by heat- 
 ing in a test-tube; if they consist of the primary sulphate, they will separate 
 sulphuric acid on heating; repeat the experiment, using 25 grams of sulphuric 
 acid to 50 grams of potassium nitrate ; heat until no more nitric acid passes 
 off; and then, after cooling and re-crystallizing the remainder, test as you did 
 the primary sulphate (the secondary sulphate will not liberate sulphuric acid 
 on heating in a test-tube). 
 
 * Be careful to have the blue litmus solution as nearly neutral as possible, and be sure 
 to add no more than one or two drops. 
 
 t Nitric acid attacks rubber; it is therefore essential to have some high melting paraf. 
 fin in the laboratory; melt a little of this, and coat the rubber stopper by dipping it into the 
 liquid; unless the temperature of any reaction becomes high enough to melt the paraffin, 
 the latter will afford a perfect protection to the rubber. A good method is to use a glass- 
 stoppered retort, and to lute the stopper with plaster of paris. 
 
528 
 
 APPENDIX OF 
 
 59. EXPERIMENTS WITH NITRIC ACID. Make a solution of indigo by 
 dissolving a little indigo in concentrated sulphuric acid, warming slightly, 
 and then diluting with water ; to this solution add nitric acid until it is 
 bleached ; take a piece of white silk ribbon, and dip it into tolerably con- 
 centrated nitric acid; wash with clean water and put aside; after a time 
 examine its color and texture. Prepare fuming nitric acid by placing 
 100 grams of ordinary nitric acid in the apparatus used for preparing 
 bromine (Fig. 46), slowly adding 50 grams of concentrated sulphuric acid 
 and gently distilling ; by this means the nitric acid is deprived of nearly 
 all water with which it was mixed ; now clean the retort, put the dis- 
 tillate back into it, and then add a few pieces of starch, connect the ap- 
 paratus and slowly distil again ; the starch will generate lower oxides of 
 nitrogen (N2O 3 and NOo) while it is itself being oxidized, and these lower 
 oxides will be dissolved by the nitric acid in the receiver ; fuming nitric acid 
 is, therefore, nitric acid which contains lower oxides of nitrogen (see 
 page 206). Place some of the fuming nitric acid in a test-tube, as is shown 
 by Fig. 47, place the test-tube inside of a beaker, in order to render an 
 accidental cracking harmless, warm the nitric acid slightly, and then drop a 
 
 red-hot piece of charcoal, which is 
 cut the size of a pea, into the acid 
 (perform this experiment under the 
 hood!). Try the solubility of various 
 metals (iron, zinc, copper, platinum) 
 in nitric acid, and note the gases 
 which pass off. Take some pieces of 
 zinc, and dissolve them in very dilute 
 cold nitric acid ; evaporate the re- 
 mainder to dryness, and then see if 
 you can discover the presence of an 
 ammonium salt (Note 49) ; do the 
 same with a piece of magnesium 
 (page 206). Prepare aqua regia by 
 mixing one part of nitric acid with 
 three parts of hydrochloric acid; allow 
 
 to stand, and notice if the odor of chlorine is perceptible ; dissolve a small 
 piece of platinum in aqua regia. To a solution of ferrous sulphate, add a solu- 
 tion of potassium hydroxide, and note the appearance of the precipitate ; then 
 heat a solution of ferrous sulphate with nitric acid, add potassium hydroxide, 
 and note the appearance of the precipitate. (Ferrous hydroxide is precipi- 
 tated in the first case, ferric hydroxide in the second ; nitric acid oxidizes -ous 
 compounds to -ic compounds [prove this also by adding the acid to a solution 
 of sulphur dioxide in water].) 
 
 THE DECOMPOSITION OF THE NITRATES. The decomposition of the 
 nitrate of a heavy metal was illustrated in Note 55. The nitrates of the 
 alkalies decompose into the nitrite and oxygen when heated. Take some 
 potassium nitrate, place in a hard glass test-tube, and heat for some time to a 
 
LABORATORY NOTES. 529 
 
 bright red heat; bubbles of oxygen will pass off. Allow the test-tube to cool, 
 dissolve the remainder in water, add a little iodide of potassium solution 
 mixed with starch paste (Note 22), and then a drop of sulphuric acid: iodine 
 will at once be liberated. Do the same with some pure potassium nitrate 
 dissolved in water, and note the difference. (Nitrous acid at once liberates 
 iodine from iodide of potassium because it is a very quick oxidizer, just as 
 ozone and hydrogen peroxide are [page 50] ; nitric acid liberates iodine only 
 after a considerable interval of time. ) 
 
 60. THE PREPARATION OF RED PHOSPHORUS. Take a piece of hard 
 glass tubing, seal one end, place a piece of phosphorus the size of a pea in 
 the tube, and seal the other end by heating in a glass-blower's flame and 
 allowing the sides to fall together;* place the tube upright in an iron crucible 
 of its own length, and then fill the crucible with sand. Heat by means of a 
 triple burner. By this method the glass tube will be hot below and tolerably 
 cool above. At some point in its length the proper temperature of 300 will 
 be reached; red phosphorus will deposit at that place. (Be sure to perform 
 this experiment under a hood, so that if the tube should explode the flying 
 glass can do no damage.) Never handle the glass tube unless it is cold. 
 Open the tube by wrapping a towel around it, and exposing the tip of the long 
 sealed end to the flame; after a little air has been admitted in this way, you 
 can break open the tube. Never attempt to break open a sealed glass tube 
 unless you have taken this precaution. 
 
 The low kindling temperature of ordinary phosphorus can be shown by 
 dissolving a little in carbon bisulphide, and then pouring a few drops of this 
 solution on a piece of filter paper. After the carbon bisulphide has evapo- 
 rated, the finely divided phosphorus which remains on the filter paper will 
 take fire spontaneously. (Throw the solution of phosphorus in carbon 
 bisulphide down the sink as soon as you are through with it, and then wash 
 the test-tube. Be careful while handling phosphorus. ) 
 
 61. THE PREPARATION OF PHOSPHINE. The apparatus is shown by 
 Fig. 48. (See p. 530.) The small generating-flask (lOOc.c.) is fitted with a 
 double bored rubber stopper, a delivery tube 6, and a tube, A, which connects 
 with a hydrogen generator. In the generating-flask, place a solution of 20 
 grams of potassium hydroxide in 40 c.c. of water and two pieces of yellow 
 phosphorus as large as a bean ; now pass a current of hydrogen through the 
 flask until all of the air is expelled, shut off the hydrogen, and then heat the 
 generating-flask in a sand-bath. t Phosphine will pass off, and will take fire 
 spontaneously when it reaches the air. Demonstrate this by placing the end of 
 your delivery-tube under water before beginning the experiment; the individual 
 
 * Substances frequently must be heated in sealed tubes. The end which is closed be- 
 fore filling the tube should be round, like the bottom of a test-tube, and as thick as the walls 
 of the glass tube. After the substance to be heated is added, the other end of the tube is 
 sealed to a long point, not by heating the glass and drawing out, but by heating and allow- 
 ing the sides to fall together, so that the sealed point is as thick as the rest of the tube. 
 This operation requires considerable practice. 
 
 t A shallow iron dish containing sand. 
 
530 
 
 APPENDIX OF 
 
 bubbles will then rise to the surface and burn. A simpler and better way of 
 preparing phosphine is to place a piece of calcium phosphide in a basin of 
 dilute hydrochloric acid ; bubbles of phosphine will at once pass off, and take 
 
 Fig. 48. 
 
 fire spontaneously when they reach the air. The gas can be collected in a test- 
 tube filled with water, and inverted over the piece of calcium phosphide. 
 Phosphine which is not spontaneously combustible can be prepared by pass- 
 ing the gas, generated by either of the above methods, through the apparatus 
 used to condense sulphur dioxide (Note 43); the liquid phosphine which is 
 spontaneously combustible will then be condensed, while the gaseous phos- 
 phine will pass on. 
 
 62. THE PREPARATION OF PHOSPHORUS PENTOXIDE. Place a piece 
 of phosphorus on a porcelain plate, ignite it, and cover with a glass bell ; 
 phosphorus pentoxide will collect on the walls of the latter and on the plate. 
 Dissolve some of the pentoxide in water, and test the solution with blue litmus ; 
 expose some of the pentoxide to the air; to the solution of phosphoric anhy- 
 dride, add silver nitrate; boil another portion of the solution for some time, 
 nearly neutralize with ammonia, and then add silver nitrate. Take some 
 secondary sodium phosphate (ordinary sodium phosphate), and heat a little of 
 the salt in a hard glass test-tube until water of crystallization is driven off; 
 now heat more strongly (formation of pyrophosphate). To a solution of 
 secondary sodium phosphate, add a solution of calcium chloride, and then 
 perform the experiments suggested by the text on page 230. Take a little 
 sodium-ammonium hydrogen phosphate, and heat in a hard glass test-tube 
 (notice the odor!); finally heat until the substance forms a transparent glass; 
 place some of this transparent glass on the end of a platinum wire, and heat 
 with a little cobalt nitrate, manganese chloride, ferric chloride ; using a fresh 
 drop of fused sodium metaphosphate and a clean wire for each one of the salts. 
 
LABORATORY NOTES. 
 
 531 
 
 Phosphorus trichloride, when dissolved in water, produces phosphorous 
 acid. Place a few drops of the phosphorus trichloride in a test-tube, add 
 water, and notice how the oily liquid dissolves; test the solution by means of 
 blue litmus paper; evaporate the solution to dryness on a water-bath until all 
 of the hydrochloric acid has passed off, and then add an alcoholic solution of 
 iodine (page 225, paragraph 2, c). 
 
 Phosphorus pentachloride, when dissolved in water, produces phosphoric 
 acid ; take a little of the pentachloride on a spatula, and add it to water in a 
 test-tube. The reaction will take place at once, with a hissing noise; the 
 solution of phosphoric acid will not reduce iodine to hydroiodic acid, as 
 the solution of phosphorous acid does. 
 
 63. MARSH'S TEST FOR ARSENIC. The accurate details of this method 
 belong in works especially devoted to analytic chemistry. Take your hydro- 
 gen generating-flask, attach a U-shaped drying-tube to this, and connect the 
 latter with a hard glass tube, which has been drawn to a point, and which is 
 constricted at two places by being drawn out in the flame of a blast-lamp 
 (Fig. 49). 
 
 Fill the U-shaped tube with granulated calcium chloride. 
 
 Place 20 grams of pure zinc in your generating-flask, add dilute sulphuric 
 acid, and allow a brisk current of hydrogen to traverse the apparatus ; when 
 all is safe, ignite the jet at the drawn-out point of the hard tube, place 
 a cold porcelain plate in the flame, and see if the flame leaves a spot. Now 
 add a solution of arsenic trioxide in 
 hydrochloric acid to the generating-flask 
 through the thistle-tube; in a few min- 
 utes arsine will be developed, and the 
 hydrogen flame will assume a violet 
 color, with a white smoke ; now heat the 
 hard glass tube at a point just before one 
 of the constrictions; amorphous arsenic 
 will be deposited on the cold portions of 
 the tube in the form of a mirror; this 
 mirror is volatile, and can be driven from 
 place to place along the tube by heating it 
 with a Bunsenburner. Hold a cold porce- 
 lain plate in the flame at the tip of the 
 hard glass tube; this will cool the escap- 
 ing gases to a point below the kindling 
 temperature of arsenic; the latter element will therefore be deposited on the 
 plate as a black spot. This spot, when touched with a drop of sodium hypo- 
 chlorite solution on the end of a glass rod, will be instantly dissolved 
 (3 NaOCl + 2 As = As. 2 O 3 + 3 NaCl); when touched with nitric acid, it is 
 dissolved, owing to oxidation and solution of the arsenious oxide formed. 
 Arsine, when passed into a solution of silver nitrate, precipitates metallic 
 silver, and forms arsenic trioxide. This reaction can be obtained by extin- 
 guishing the flame at the tip of the hard glass tube, and passing the mixture 
 
 Fig. 49. 
 
532 APPENDIX OF 
 
 of gases, which are being generated, into a test-tube containing silver nitrate 
 solution ; black, metallic silver will be precipitated ; filter this off, and add am- 
 monia to the clear filtrate, until neutral; yellow arsenite of silver will be 
 precipitated. If the gas which is passing from the generator is stibine, and 
 not arsine, the mirror spots on the tube will have a grayish and more metallic 
 appearance, will form nearer to the flame, and will be very nearly not-volatile. 
 The spot on the porcelain mirror will not disappear on addition of sodium 
 hypochlorite ; it will turn white on addition of nitric acid, because the oxide 
 of antimony is insoluble in that substance. If stibine is passed into a solution 
 of silver nitrate, black silver antimonide, SbAg 3 , is precipitated; if this pre- 
 cipitate is boiled with a concentrated solution of tartaric acid, then the anti- 
 mony will be dissolved, and its presence can subsequently be readily proved by 
 precipitating antimony sulphide by means of sulphuretted hydrogen. For 
 further directions, consult some work on analytical chemistry. (See page 255.) 
 (All work with arsine or stibine must be done under the hood!) 
 
 64. A QUICK METHOD FOR DETECTING ARSENIC. Draw a piece of 
 hard glass tube to a point, as is shown in Fig. 50; place a little arsenious 
 oxide, or a small particle of the substance which you suspect to be arsenious 
 
 oxide, in the tip; above this, 
 place a small piece of char- 
 coal ; heat the tube at the spot 
 where the coal has lodged 
 until the coal is red-hot, and 
 then gradually draw the tube 
 through the flame until the 
 arsenic trioxide is heated ; the 
 
 f 50 latter will sublime, and will, 
 
 in passing over the hot char- 
 coal, be reduced to metallic arsenic, which will form a mirror similar to 
 that observed in Marsh's test. 
 
 The experiments which can be performed with the oxides and sulphides of 
 arsenic will suggest themselves during the study of chapters xxxi. and xxxii. 
 
 65. STIBINE. The preparation of stibine is exactly like that of arsine, 
 so that the pupil may follow the directions given in Note 63, excepting that he 
 must substitute a hydrochloric acid solution of antimony trioxide for that of 
 arsenious oxide. The distinctions existing between the spot produced on a 
 cold plate by burning arsine, and that produced by stibine, are given in the last 
 part of Note 63; the experiments showing such distinction should be followed 
 out. As antimony and arsenic are elements which are commonly met with in 
 analytical work, and as their reactions are fully described in all of the direc- 
 tions for qualitative analysis which are published, it is scarcely necessary to 
 enter into a more detailed discussion of experiments to be performed with 
 antimony at this place. By following the text, the teacher can easily select 
 experiments which need no detailed description (pages 252 to 256, and chapter 
 xxxiv. ). The same is true as regards bismuth. 
 
LABORATORY NOTES. 
 
 533 
 
 66. DESTRUCTIVE DISTILLATION OF COAL AND WOOD. Unless a lab- 
 oratory is especially well appointed, and unless the pupil has plenty of time at 
 his disposal, it will not be possible to carry out the dry distillation of bitumi- 
 nous coal or of wood ; when such work is attempted it belongs, more prop- 
 erly, in a course on organic chemistry. The pupil should, however, place a 
 small piece of wood in a hard glass test-tube, and heat, endeavoring, as much 
 as possible, to ascertain the nature of the products evolved ; he should also see 
 if the gases which pass off will burn. The same experiments should be per- 
 formed with a piece of bituminous coal. 
 
 67. THE ABSORPTION OF COLORING MATTER BY CHARCOAL. Prepare 
 a solution of indigo (Note 59), place in a 300 c.c. flask, and warm, after add- 
 ing two tablespoonfuls of animal charcoal ; * filter, and, if the solution is not 
 colorless, repeat the operation. Do the same with a solution of iodine in 
 iodide of potassium. The absorption of gases by charcoal was described in 
 Note 49. 
 
 69. THE FLAME. The fact that the centre of a flame is cold, while the 
 outer zone is hot, can be demonstrated by turning the flame of a Bunsen 
 burner down, shutting off the air supply so that the flame becomes luminous, 
 and then quickly placing a sheet of filter paper upon it so that about one-half 
 of the flame is below the surf ace. of the paper (Fig. 11, page 283); as soon as 
 the brown, burnt circle appears on the upper side of the paper, withdraw the 
 latter quickly. 
 
 Diluting illuminating gas with a non-combustible gas renders the flame non- 
 luminous. Arrange an apparatus as is shown by Fig. 51. Attach one arm of 
 the T-shaped glass tube to a 
 tap of illuminating gas, and 
 attach the other to an appa- 
 ratus generating dry carbon 
 dioxide; the remaining arm 
 of the T-tube is fitted with a 
 brass pipe which can be heat- 
 ed by a Bunsen burner. Turn 
 off the supply of carbon 
 dioxide by the pinchcock, 
 turn on the gas, and light the 
 burner; the flame will of 
 
 course be luminous. Now Fi 
 
 turn on the carbon dioxide 
 
 so that the latter gas mingles with and dilutes the illuminating gas; the flame 
 will at once become larger and non-luminous. However, if the brass exit-tube 
 which serves as a burner is heated to a point at which the decomposition of 
 ethylene into methane and carbon takes place, the flame will once more be 
 rendered luminous, although still diluted by carbon dioxide. 
 
 Animal charcoal can be obtained at any chemical supply house. 
 
534 
 
 APPENDIX OF 
 
 A flame can be extinguished by cooling below the kindling temperature. 
 Light a Bunsen burner, and then place a piece of copper wire gauze upon the 
 flame so that the latter will 
 about be bisected ; the flame ji 
 will continue to burn beneath 
 the gauze; but above the lat- 
 ter, because the wire conducts 
 the heat away too rapidly, no 
 flame will be seen. The re- 
 verse of this experiment (Fig. 
 52) is shown by placing the 
 gauze above an unlighted bur- 
 ner, turning on the gas, and 
 
 then lighting above the gauze; the flame will then not 
 form below. 
 
 To show that oxygen will burn in illuminating 
 gas. Construct an apparatus as shown by Fig. 53. 
 The tube B is connected with an illuminating gas tap, 
 the gas turned on, and allowed to run until all air is 
 expelled from the glass bulb; now light the gas at Fig. 53. 
 
 the top, the double bored stopper containing tubes A 
 
 and C being removed for the purpose. Attach A to a gasometer which will 
 furnish oxygen, turn on a slow stream of that gas, and then bring the stopper 
 which holds A and C back into position. As the current of oxygen comes 
 in contact with the burning gas escaping from the bulb, it will be ignited, and 
 will continue to burn in the illuminating gas. 
 
 70. THE PREPARATION OF CARBON MONOXIDE BY PASSING STEAM 
 OVER RED-HOT CHARCOAL. Take an iron gas-pipe 700 m.m. in length; 
 attach one end of it, by means of a rubber stopper, to a flask of 300 c.c. 
 capacity, which is so arranged on a retort stand that you can boil water in it, 
 and that the steam must pass through the iron tube. Connect the other end 
 of this iron tube with a safety bottle (Note 1), and put a delivery-tube, bent so 
 that it will open under the water in your pneumatic trough, into the bottle ; 
 place the iron tube into a combustion furnace (Fig. 1, Note 1), having previ- 
 ously filled the tube with pieces of charcoal which are broken to the size of a 
 pea; heat to a red heat, and then pass steam over the charcoal. After all of 
 the air has been expelled from the apparatus by means of the current of steam, 
 collect the escaping gas over the pneumatic trough by displacement of water. 
 This gas will be carbon monoxide mixed with hydrogen, as can be proven by 
 inverting one of the cylinders filled with it, and touching a lighted taper to 
 the mouth.* 
 
 * In using the combustion furnace for this experiment, take care that the iron tube is 
 sufficiently long to extend some little distance beyond either end of the furnace; otherwise 
 the rubber stoppers will become too hot and will melt. You should feel of the stoppers from 
 time to time; and if there is danger of their fusing, cool them by pouring on water. 
 
LABORATORY NOTES. 
 
 535 
 
 71. THE PREPARATION OF CARBON MONOXIDE FROM OXALIC ACID. 
 The apparatus is the same as that used for nitrous oxide (Note 52, Fig. 43). 
 In the retort, place 10 grams of crystallized oxalic acid, add 60 grams of con- 
 centrated sulphuric acid, and heat until a regular evolution of gas takes place; 
 pass this gas through a wash-bottle (Fig. 18) containing a solution of potas- 
 sium hydroxide (one part potassium hydroxide to two parts of water) before 
 you collect over the pneumatic trough. By means of this wash-bottle the 
 carbon dioxide, which is generated simultaneously with the carbon monoxide, 
 is absorbed (it forms potassium carbonate), while pure carbon monoxide is 
 collected; the latter gas will burn with a pale blue flame. 
 
 72. THE PREPARATION OF CARBON DIOXIDE. The apparatus is the 
 same as that used for the preparation of hydrogen or of hydrogen sulphide 
 (Fig. 34). Charge the generating-flask with 20 grams of marble which has 
 teen broken to the size of a hickory nut, pour on 100 c.c. of hydrochloric acid, 
 -which you have diluted with an equal amount of water, and pass the gas 
 which is generated through a wash-bottle containing water (if dry carbon 
 dioxide is required, pass the gas through a second bottle containing sulphuric 
 acid); collect the carbon dioxide by displacement of the air as you did chlo- 
 rine and hydrobromic acid. 
 
 Take a number of other carbonates (sodium carbonate, potassium carbon- 
 ate, barium carbonate, etc.), place a little of each in test-tubes, and add 
 hydrochloric acid to each test-tube ; ascertain if carbon dioxide is given off. 
 
 73. EXPERIMENTS WITH CARBON DIOXIDE. Fill 5 or 6 cylinders with 
 carbon dioxide. Construct an apparatus such as is shown by Fig. 54; a num- 
 ber of small candles are placed on wires which are 
 attached in an upright position t$ a larger wire which 
 is bent in the form of a flight of steps ; light all of 
 the candles, place them in a beaker glass, and pass 
 carbon dioxide into the latter by means of a tube 
 extending to the bottom; the lights will be extin- 
 guished successively from below upward ; this experi- 
 ment will also demonstrate the fact that carbon 
 dioxide neither burns, nor supports combustion. Pass 
 some carbon dioxide from your generator into a solu- 
 tion of barium hydroxide and into one of calcium 
 hydroxide : filter the precipitates formed in each case, 
 wash them from the filter papers into test-tubes, and 
 add hydrochloric acid ; prove that carbon dioxide is 
 passing off by holding a glass rod which has been Flg 54 
 
 dipped into lime-water (a solution of calcium hydroxide) just within the 
 mouths of the test-tubes; if carbon dioxide is being generated, the lime-water 
 tvill become turbid. Pass carbon dioxide into a concentrated solution of 
 potassium hydroxide and into one of sodium hydroxide; take a little of the 
 solutions so formed, and add hydrochloric acid to them; prove that carbon 
 dioxide passes off. Pass carbon dioxide into a solution of lime-water until 
 the precipitate which at first forms is re-dissolved (formation of the primary 
 
536 APPENDIX OF LABORATORY NOTES. 
 
 carbonate), and then boil the solution; add hydrochloric acid to the solution 
 before it is boiled and to the precipitate which is formed by boiling, and see 
 if carbon dioxide is given off in each case. 
 
 74. THE ETCHING OF GLASS BY MEANS OF HYDROFLUORIC ACID. Take 
 a shallow lead dish 3 inches long, 2 inches wide, and 1 inch deep; put about 
 10 grams of powdered calcium fluoride into this, and cover the fluoride with 
 sulphuric acid ; prepare a piece of window glass by dipping it into melted 
 paraffin,* allowing the latter to cool, and etching some figures by scratch- 
 ing away the paraffin coating with the point of a knife ; place this prepared 
 glass over the lead dish in the form of a cover, put the whole into a warm 
 place under the hood, and allow to stand for 6 or 8 hours; the glass will 
 then be etched where the paraffin was scratched away. (In working with 
 hydrofluoric acid, be extremely careful not to allow it to come in contact with 
 the hands nor to inhale the fumes. The ulcers caused by hydrofluoric acid 
 burns are very painful and slow to heal, and the results of inhalation of the 
 vapors may be dangerous.) 
 
 Preparation of silicon tetrafluoride and of Jluo silicic acid. Take a flask 
 of 300 c.c. capacity, put into this 10 grams of quartz sand mixed with 10 
 grams of powdered calcium fluoride ; fit a single bored rubber stopper to the 
 flask, and in this rubber stopper place a glass delivery-tube which is bent 
 with two right angles ; the end of this delivery tube will then point downward ; 
 to this end attach a funnel by means of the stem in such a manner that the 
 escaping gas must pass through the funnel, and place a beaker of water under 
 the funnel so that its rim just touches the surface of the water. Pour 100 grams 
 of concentrated sulphuric acid on the mixture of sand and calcium fluoride in 
 the generating-flask, connect the apparatus and warm gently; silicon tetra- 
 flouride will pass off, and will come in contact'with the water in the beaker; by 
 this means fluosilicic acid will be formed. While the solution of the latter is 
 being formed, it will at the same time become filled with flakes of silicic acid, so 
 that, if the delivery-tube had not been widened by the funnel, it would soon have 
 become clogged and an accident would have followed. Filter the fluosilicic acid 
 from the precipitated silicic acid, and test the reaction of the solution towards lit- 
 mus; add just enough potassium hydroxide solution to neutralize the fluosilicic 
 acid and allow to stand; nearly insoluble potassium fluosilicate will separate. 
 
 The elements which follow silicon in the text are nearly all of a metallic na- 
 ture, and none of the experiments which will fix their character in the mind of 
 the pupil will present any practical difficulty. It is better, therefore, for the 
 teacher to select such experiments as he deems proper from the text; indeed, it 
 often seems advisable to study some of the chemical relations of the metals in 
 the laboratory by the methods of qualitative analysis, while, at the same time, 
 becoming familiar with the more general aspect of the chemistry of those ele- 
 ments by following the text-book on general chemistry; the small manuals 
 containing directions for qualitative analysis are, however, so numerous that 
 it seems unnecessary to add a list of experiments, which practically cover the 
 same ground, to a text-book of general chemistry. 
 
 * Use so-called high melting paraffin. 
 
INDEX. 
 
 The numbers refer to the pages. 
 
 Acetylene 
 
 281 Acid, phosphoric 95, 226, 232 
 
 Acid, Boric 330 
 
 Bromic 130 
 
 Carbamic 298,299 
 
 Chlor-auric 409 
 
 Chlorides of sulphur 157 
 
 Chlorous 119 
 
 Chlor-platinic 493 
 
 Chlor-sulphonic 157 
 
 Cyanic 297 
 
 Cyanuric 297 
 
 Disilicic 307 
 
 Disulphuric 154 
 
 Dithionic 155 
 
 Ferric 483 
 
 Fluoboric 330 
 
 Fluosilicic 302,303 
 
 Hydrobromic 80-82 
 
 Hydrochloric 66-77 
 
 Hydrocyanic 295 
 
 Hydrofluoric 56,57 
 
 Hydroiodic 85-87 
 
 Hypobromous 129 
 
 Hypochlorous 115, 119, 121 
 
 Hyponitrous 208 
 
 Hypophosphorous 231 
 
 lodic 
 
 Meta-arsenic . . . 
 Meta-arsenious . . 
 Meta-antimonic . . 
 Meta-phosphoric 
 Meta-silicic . . . 
 Meta-vanadic . . . 
 
 Nitric 
 
 Nitrosyl-sulphuric . 
 Nitrous . 
 
 . . . 130, 132 
 . . 227,241,242 
 ..... 227 
 255 
 
 . . . 226,228 
 227, 304, 305, 307 
 
 441 
 
 . . . 203-210 
 
 147 
 
 . . 147, 148, 208 
 
 Ortho-arsenic 227, 241, 242 
 
 Ortho-arsenious 227 
 
 Ortho-antimonic 255 
 
 Ortho-phosphoric 227, 229 
 
 Ortho-silicic 227, 304, 305 
 
 Ortho-vanadic 441 
 
 Pentathionic 155 
 
 Perchloric 119, 126 
 
 Per-iodic .130 
 
 Phosphorous 95, 224 
 
 Plumbic 323 
 
 Pyro-antimonic 255 
 
 Pyro-antimonous 254 
 
 Pyro-arsenic 241 
 
 Pyro-phosphoric 227, 231 
 
 Pyro-vanadic 441 
 
 Selenic 161 
 
 Selenious 160 
 
 Stannic a 317 
 
 Stannic ft 317 
 
 Stannous 317 
 
 Sulpharsenic 245,246 
 
 Sulpharsenious 244, 245 
 
 Sulpho-dithio carbonic 293 
 
 Sulphuric, 56, 66, 80, 117, 136, 137, 145, 150 
 
 Sulphurous 138, 158 
 
 Telluric 161 
 
 Tellurous 161 
 
 Tetrathionic 155 
 
 Thiosulphuric 154 
 
 Trithionic 155 
 
 Trisilicic 307 
 
 Trithiocflrbonic (see sulpho-dithio 
 carbonic). 
 
 Acid reactions 76, 177, 178 
 
 Acids, action on salts 142 
 
 Affinity for bases 141 
 
 Cause of character 177, 178 
 
 Decomposition by metals .... 31 
 
 Definition of . . . 75 
 
 Formation of , from anhydrides . . 116 
 
 General formulae of 117 
 
 Hydrated 131, 135, 145 
 
 Nature of 32 
 
 Polybasic 139 
 
 Relative avidity of 141, 142 
 
 Strong and weak 141 
 
 Uni-, di-, and tribasic 139 
 
 Valence in 116, 117 
 
 Affinity, chemical ....... 10 
 
 Agate 304 
 
 Alabaster 418 
 
 Alkali metals, carbonates of . . . .389 
 
 537 
 
538 
 
 INDEX. 
 
 Alkali metals, comparative table of . 391 
 Decomposition of water by . . 383, 384 
 
 Double halides of 388 
 
 Halidesof 388 
 
 Halides of, heats of formation of . 388 
 
 Hydroxides of 385, 386 
 
 Oxides o? 385 
 
 Properties of 384, 385 
 
 Relations of 383 
 
 Sulphhydrates of 386 
 
 Sulphides of 386 
 
 Alkaline earths, carbonates of . 415, 416 
 
 Chlorides of 414* 
 
 Comparison of properties of ... 412 
 
 Metallurgy of 411 
 
 Oxides and hydroxides of . . 412, 413 
 
 Sulphates of 416, 4}7 
 
 Table of 411 
 
 Allotropism 48 
 
 Alloys 249,250 
 
 Aluminium, basic sulphates of ... 340 
 
 Halides of 336, 337 
 
 Halides, double salts of 337 
 
 Hydroxides of 30, 338 
 
 Occurrence of 333 
 
 Phosphates of ........ 340 
 
 Preparation of 333,334 
 
 'Properties of 334 
 
 Trichloride 336 
 
 Trichloride, molecular weight of . 336 
 
 jTrioxide 338 
 
 Trisulphide 341 
 
 Alum 339,340 
 
 Amalgams 191 
 
 Amido group 298 
 
 Ammonia, decompos'n of, by chlorine, 63 
 
 History of 182 
 
 In atmosphere 170 
 
 Preparation of 183, 184 
 
 Properties of i. . . . 185 
 
 Solubility of I ... 187 
 
 > Table of compounds of 263 
 
 Volumetric composition of . 185, 186 
 
 Ammonia liquor 184 
 
 Ammonia-soda process 390 
 
 Ammonium amalgam 191 
 
 Carbamate 298 
 
 Carbonate 298 
 
 Cyanate 298 
 
 Chloride 188,189 
 
 Molybdate . . . 452 
 
 Nitrate 188 
 
 Phospho-molybdate 454 
 
 Sulphate 188, 189 
 
 Chloride, vapor density of .... 189 
 Ammonium salts, Decomposition of 
 
 184, 189, 190 
 Formation of 189, 190 
 
 Ammonium salts, nature of .... 189 
 
 Occurrence of 182 
 
 Table of 263 
 
 Anatas 438 
 
 Anglesite 320, 417 
 
 Anhydrides . 14, 22, 25, 26, 30, 95, 130, 131 
 Conversion to acids . 115, 116, 117, 178 
 
 Valence in 115 
 
 Anhydrite 417 
 
 Antimonic acids 255 
 
 Antimonous acids 254 
 
 Antimony, acids of 254, 255 
 
 Alloys of ,.. 249-251 
 
 Basic salts of . 253 
 
 Double salts of 253 
 
 Halides of 252 
 
 History of 248 
 
 Metallurgy of 248 
 
 Occurrence of 248 
 
 Oxides of 254 
 
 Pentachloride 181, 253 
 
 Antimony, pentasulphide of .... 256 
 
 Pentoxide 254 
 
 Properties of 248, 249 
 
 Sulphides of 255, 256, 264 
 
 Tetroxide .' 254 
 
 Tribromide 252 
 
 Trichloride . . . 181, 249, 252, 253, 263 
 
 Trifluoride 252 
 
 Tri-iodide 252 
 
 Trioxide 249,254 
 
 Trisulphide 255 
 
 Apatite 211,418 
 
 Arragonite 416 
 
 Arsenates 242, 247 
 
 Arsenic acid, oxidizing action of 242, 243 
 
 Salts of . 241 
 
 Table of 247 
 
 Arsenic, halides of 238, 239 
 
 History of 234 
 
 Metallurgy of 235 
 
 Occurrence of 234 
 
 Properties of 235 
 
 Table of acids of 247 
 
 Table of oxides of 247 
 
 Table of sulphides of . . . . 247, 264 
 
 Pentafluoride 238 
 
 Pentasulphide 245,246 
 
 Pentoxide 239, 241 
 
 Tribromide 238 
 
 Trichloride 181,238,263 
 
 Trifluoride 238 
 
 Tri-iodide 238 
 
 Trioxide 235, 239, 240, 241 
 
 Trisulphide 244, 245, 247 
 
 Arsenious oxide 239, 247 
 
 Poisonous effects of 240 
 
 Chemical action of 241 
 
INDEX. 
 
 539 
 
 Arsenates 
 
 Arsenites 241-217 
 
 Arsenopyrite 2 <*4 
 
 Arsine, preparation of 236 
 
 Properties of 236,237 
 
 Artificial ice 187 
 
 Atmosphere, the, ammonia in 170 
 
 Carbon dioxide in 167, 168 
 
 Composition of, history of .... 165 
 
 Pressure of 170, 171 
 
 Pressure of, measurement of . 171, 372 
 Quantitative composition of ... 166 
 
 Relation of, to life 174 
 
 Solids in 170 
 
 Specific gravity of 173 
 
 Water vapor in 169 
 
 Atomic heats of compounds, rela- 
 tions of 353,354 
 
 Atomic heats of elements, relations 
 
 of 352-354 
 
 Atomic hypothesis, the ... 4, 6, 347 
 Atomic volumes, relations of . 365-367 
 
 Atomic weights, absolute 7 
 
 Determination of 349-361 
 
 Maximum 73,349 
 
 Of oxygen family 106 
 
 Standard of 7 
 
 Table of 8 
 
 Auric chloride 408 
 
 Hydroxide 408 
 
 Oxide 408 
 
 Sulphide 409 
 
 Aurous-auric oxide 406 
 
 Chloride 408 
 
 Aurous chloride 408 
 
 Cyanide 409 
 
 Sulphide 409 
 
 Avidity of acids . . . . .Ill, 142, 113 
 
 Avogadro 70 
 
 Avogadro's hypothesis, 71, 72, 73, 348 
 
 349, 360 
 Azoimid 193, 194 
 
 Barium, isolation of 412 
 
 Properties of 412 
 
 Carbonate 292^415,416 
 
 Chlorate 420 
 
 Hydroxide . . . . 413 
 
 Nitrate 420 
 
 Oxide 413 
 
 Permanganate 466 
 
 Sulphate 416,417, 
 
 Superoxide ill 
 
 Barite 417 
 
 Barometer, height of 171 
 
 History of 170 
 
 Variations of 173 
 
 Barytocelestine 417 
 
 14,25,26,30,31,177 
 
 Neutralization of 375-377 
 
 Relative strength of 377 
 
 Beauxite 333,334 
 
 Benzine 279 
 
 Beryll 307 
 
 Berylli um > isolation of 411 
 
 Properties of 412 
 
 Carbonate 415 
 
 Chloride 414 
 
 Hydroxide 413 
 
 Oxide 413 
 
 Sulphate 416,117 
 
 Berzelius M 
 
 Bismite 257 
 
 Bismuth, alloys of 258 
 
 Basic halides of 258 
 
 Halidesof 259 
 
 Metallurgy of 257 
 
 Native 257 
 
 Occurrence 257 
 
 Oxides of 259 
 
 Properties of 258 
 
 Salts of 260, 261 
 
 Hydroxide 259, 260 
 
 Monosulphide 261 
 
 Monoxide 259 
 
 Nitrate 260 
 
 Nitrate, basic 260 
 
 Pentoxide 259 
 
 Subnitrate (see basic nitrate). 
 
 Sulphate 261 
 
 Telluride 257 
 
 Tetroxide 259 
 
 Tribromide 258 
 
 Trichloride 258 
 
 Trifluoride 258 
 
 Tri-iodide 258 
 
 Trioxide 257, 259 
 
 Trisulphide 257,261 
 
 Bismuthinite 257 
 
 Blast furnace, construction of ... 473 
 
 Changes in 474 
 
 Borax . 328, 331 
 
 Berates, ortho-, meta-, and tetra- . . 331 
 
 Occurrence of 328 
 
 Boric acid, occurrence of 330 
 
 Ortho- and meta 331 
 
 Properties of 331 
 
 Borocalcite 328 
 
 Boron, acids of 329 
 
 Occurrence of 328 
 
 Preparation of 328 
 
 Properties of 329 
 
 Hydride 329 
 
 Nitride 332 
 
 Oxy-chloride 332 
 
 Phosphate 331 
 
540 
 
 INDEX. 
 
 Boron, trichloride 329 
 
 Trifluoride 329 
 
 Trioxide 330 
 
 Boron family, elements of 326 
 
 Oxides of elements of 326 
 
 Valence of elements of . . > . . 326 
 
 Halidesof ' . . 327 
 
 Braunite 460 
 
 Bromides, formation of 82 
 
 Bromine, hydrate of 80 
 
 Occurrence of 79 
 
 Oxy-acids of 129, 130 
 
 Preparation of 79, 80 
 
 Properties of 79 
 
 Monochloride 132 
 
 Water 80 
 
 Brookite 438 
 
 Butane (see diethyl). 
 
 Bunsen burner, the 283 
 
 Cadmium, alloys of 428 
 
 Metallurgy of 427 
 
 Occurrence of 426 
 
 Properties of 424 
 
 Carbonate 429 
 
 Chloride 429 
 
 Hydroxide 428 
 
 Oxide 428 
 
 Sulphate 429 
 
 Sulphide 430 
 
 Calcite 292,416 
 
 Group of minerals 416 
 
 Calcium, isolation of 412 
 
 Properties of 412 
 
 Carbonate 292, 413, 415, 416 
 
 Carbonate, primary 418 | 
 
 Chlorate 125,420 
 
 Chloride 414,420 
 
 Dimanganite 465 
 
 Hydroxide 30, 44, 413 
 
 Hypochlorite 122, 123 
 
 Oxide 413 
 
 Phosphate 212,418 
 
 Phosphate, primary .... 229,418 
 Phosphate, secondary . . . 230, 418 
 
 Phosphate, tertiary 229 
 
 Silicates 420, 421 
 
 Tungsten 445 
 
 Calcium and magnesium carbon- 
 ate 292,416 
 
 Calomel (see mercurious chloride). 
 
 Carbamic acid 298, 299 
 
 Carbon, allotropic forms of . . 269-273 
 
 Amorphous 272,273 
 
 Halidesof 285,286 
 
 Hydrogen compounds of . . 274-284 
 Hydrogen compounds of, table of . 279 
 .Nitrogen compounds of ... 294-299 
 
 Carbon, nitrogen and oxygen com- 
 pounds of 297-299 
 
 Nitrogen and oxygen compounds 
 
 of, table of . 299 
 
 Occurrence of <;69 
 
 Oxides of 286-291 
 
 Table of compounds of 319 
 
 Carbonates 291 
 
 Natural, table of 292 
 
 Primary and secondary . . .- 291, 292 
 
 Carbon dioxide . 22 
 
 Changes in atmosphere . . . 167, 168 
 
 History 289 
 
 In atmosphere 166-168 
 
 Preparation 290 
 
 Properties 290,291 
 
 Carbon disulphide 95, 292 
 
 Carbonic acid, meta- 291, 299 
 
 Ortho 285, 286 
 
 Carbonic acids 267 
 
 Carbon monoxide 286 
 
 Preparation 286, 287 
 
 Properties 287 
 
 Poisonous action 288 
 
 Tetrabromide 286 
 
 Tetrachloride 75, 285 
 
 Carbon family, acids of 267 
 
 Elements of, comparison of ... 265 
 Hydrogen compounds of . . 265, 266 
 Hydrogen compounds of, table of . 266 
 
 Oxides of 267 
 
 Oxides and acids of, table of ... 268 
 
 Sulphides of, table of 268 
 
 Table of compounds of 325 
 
 Carbonyl chloride . . . .288, 289, 298 
 
 Cassiterite 312,438 
 
 Cavendish 27, 36, 165 
 
 Celestine 417 
 
 Cement 414 
 
 Cerium, compounds of 439 
 
 Halidesof 440 
 
 Occurrence of 439 
 
 Oxides of 440 
 
 Cerussite 320 
 
 Chalcedony 304 
 
 Chalcocite 397 
 
 Chalcopyrite 91, 397 
 
 Charcoal 273 
 
 Chemism 10 
 
 Chlorates, formation of 124 
 
 Properties of 125 
 
 Chlo-rauric acid 409 
 
 Chloric acid 119 
 
 Decomposition of 126 
 
 Properties of 126 
 
 Chlorides, conversion of, into oxides 112 
 
 Formation of 76, 77 
 
 Formulae of 63, 64 
 
INDEX. 
 
 541 
 
 Chlorides, valence of elements in . . 109 
 
 Of sulphur 156 
 
 Chlorine, action on hydrogen com- 
 pounds 63 
 
 Bleaching action 64 
 
 Combustion in 63 
 
 History of 58 
 
 Occurrence 58 
 
 Oxides of 119,120 
 
 Oxidizing action of 65 
 
 Oxy-acidsof 119,128 
 
 Preparation of 59, 60, 61, 62 
 
 Properties of 61, 62 
 
 Chlorine dioxide 119, 127 
 
 Hydrate 62 
 
 Monoxide ....... 115, 119, -120 
 
 Trioxide 119, 127 
 
 Water 64 
 
 Chloroform (see methiu chloride). 
 
 Chlorous acid 119 
 
 Chlor-platin amines 494 
 
 Chlorsulphonic acid 157 
 
 Chromates 449, 450, 451 
 
 Chrome-alum 447 
 
 Chromic acid 448 
 
 Chloride 447 
 
 Hydroxide 446 
 
 Oxide 446 
 
 Chromite 444 
 
 Chromites 446 
 
 Chromium, Acid chlorides of . 448, 449 
 Compounds of, preparation . 451, 452 
 
 Compounds of, uses , 452 
 
 Occurrence of 444 
 
 Preparation of 445 
 
 Properties of 445, 446 
 
 Chromium family, elements of . . 443 
 Elements of, comparison .... 443 
 
 Halides of 444 
 
 Tables of compounds of . . 443, 458, 488 
 
 Chromous chloride 451 
 
 Hydroxide 451 
 
 Chromous-chromic oxide .... 451 
 
 Chromspinell 444 
 
 Chromyl chloride 448 
 
 Cinnabar 427, 435 
 
 Clay 340 
 
 Purification of 341 
 
 Coal, formation of 168, 271, 272 
 
 Cobalt, metallurgy of 477 
 
 Occurrence of 472 
 
 Properties of 472 
 
 Sulphides of 472 
 
 Cobalt amines . . . . ^ 486 
 
 Cobaltglance 234 
 
 Cobaltic nitrate 485 
 
 Oxide 485 
 
 Cobaltite 472 
 
 Cobaltous chloride 484 
 
 Nitrate 485 
 
 Oxide 484 
 
 Sulphate 417 
 
 Sulphide 485 
 
 Columbates 441 
 
 Columbite 441 
 
 Columbium, compounds of .... 441 
 
 Occurrence of 441 
 
 Combustion 22, 23, 24 
 
 In chlorine 24 
 
 In oxygen 22, 23 
 
 Rapid and slow 23 
 
 Constancy of matter, law of . . . 2, 15 
 
 Copper, alloys of 400 
 
 Chlorides of 401,402 
 
 Metallurgy of 398 
 
 Occurrence of 397 
 
 Oxides of 401, 402 
 
 Properties of 398 
 
 Salts of 401,402 
 
 Sulphides of 404 
 
 Copper, silver, and gold, alloys of . 400 
 
 Properties of 397, 400 
 
 Resemblance to alkalies .... 396 
 Table of compounds of ... 409, 410 
 Corrosive sublimate (see mercuric 
 chloride). 
 
 Corundum 333 
 
 Crocoite 444 
 
 Cryolite 333, 334 
 
 Cupric carbonates, basic .... 403, 404 
 
 Chloride 402 
 
 Hydroxide 260, 402 
 
 Nitrate 403 
 
 Oxide 38,402 
 
 Sulphate 42, 402, 417 
 
 Sulphate, formation of ... 137, 402 
 
 Sulphide 404 
 
 Cuprous chloride 401 
 
 Iodide 401 
 
 Oxide 401 
 
 Sulphide 404 
 
 Cyanates 297,298 
 
 Cyanic acid 297,298 
 
 Cyanamide 299 
 
 Cyanides 294 
 
 Cyanogen 294, 295 
 
 Dalton, John 3, 69 
 
 Dalton's hypothesis 3 
 
 Davy 37, 58, 104 
 
 Deliquescence 43 
 
 Dialysis 305 
 
 Diamond 269, 270 
 
 Diaspor 333 
 
 Dichromates 450 
 
 Diethyl 278,279 
 
542 
 
 INDEX. 
 
 Dimethyl 278,279 
 
 Disilicates : . . 307 
 
 Dissociation 61, 351 
 
 In solution 351, 380-382 
 
 Disulphuric acid 154 
 
 Dithionic acid 155 
 
 Dolomite 292,416 
 
 Double decomposition 57 
 
 Laws of 379,380 
 
 Double salts, structure of 337 
 
 Dulong and Petit, law of ... 351-356 
 Dumas 38, 165 
 
 Efflorescence 43 
 
 Ekaboron 373 
 
 Elements, division of 12 
 
 Electro-negative 13,368 
 
 Electro-positive 13, 368 
 
 Number of 9 
 
 Polyvalent, compounds of . . . . 114 
 
 Prediction of 373 
 
 Relation between 17 
 
 Specific heat of, relation of ... 352 
 Table of molecular weights of . . 426 
 
 Electrolysis 13, 37 
 
 Energy, chemical . 10, 11, 23, 24, 46, 74, 76 
 
 Kinetic 11, 12 
 
 Potential . 10, 11 
 
 Transformation of 11 
 
 Ethane (see Dimethyl). 
 
 Ethylene 280 
 
 Decomposition of 282, 283 
 
 Properties of 281,282 
 
 Structure of 280 
 
 Equations, chemical .... 16, 19, 20, 74 
 
 Equivalent quantities 141 
 
 Weights 39, 357, 358 
 
 Euxenite . . . . 437 
 
 Faraday's law 359 
 
 Ferric acid 483 
 
 Chloride 482 
 
 Hydroxide 30, 481 
 
 Oxide 481 
 
 Sulphate 482 
 
 Sulphide 482 
 
 Ferrites 481,483 
 
 Ferrous carbonate 292,480 
 
 Chloride 479 
 
 Columbate 441 
 
 Hydroxide 29, 260, 479 
 
 Oxide 479 
 
 Selenide 103 
 
 Sulphate 42,417,479,480 
 
 Sulphide 97,480 
 
 Tantalate 441 
 
 Tungstate 445 
 
 Ferrous-ferric oxide ... 22, 438, 483 
 
 Flame, the 22, 282, 283 
 
 Flint 304 
 
 Fluorine, combustion in 55 
 
 History of 53 
 
 Occurrence of 55 
 
 Preparation of 55 
 
 Properties of , 55 
 
 Valence of 303 
 
 Fluosilicic acid, constitution of . . 303 
 
 Preparation of S02 
 
 Properties of ....... 303, 304 
 
 Formic aldehyde 167 
 
 Formulae, chemical 15, IP, 20 
 
 Structural 108, 111, 112, 113 
 
 Gadolinite . . . 437,439 
 
 Earths 437, 442 
 
 Galenite 91,320 
 
 Gallium 343, 344 
 
 Garnet 307, 333 
 
 Gases, calculation of, to standard vol- 
 ume 172,173 
 
 Kinetic theory of 68, 69 
 
 Relation of pressure and volume of, 172 
 Ratio between specific gravity and 
 
 combining weight of .... 68 
 
 Relation by volume 70, 71 
 
 Relation of specific gravities to mo- 
 lecular weights of .... 72, 73 
 
 Gay Lussac 37, 68, 69, 165 
 
 Law of 68 
 
 Germanium 309, 310, 311 
 
 Chloroform 310 
 
 Dioxide 310 
 
 Disulphide 311 
 
 Monosulphide 311 
 
 Monoxide 310 
 
 Tetrachloride 310 
 
 Gersdorfite 473 
 
 Glass 420,421,422 
 
 Glucinum (see beryllium). 
 
 Gold, alloys of 400 
 
 Chlorides of 408 
 
 Hydroxides of 408 
 
 Metallurgy of 399, 400 
 
 Occurrence of 398 
 
 Oxides of 408 
 
 Graphite, occurrence 270 
 
 Uses 271 
 
 Greenockite 426 
 
 Guanidine 299 
 
 Guano 418 
 
 Gunpowder 391 
 
 Gypsum 417 
 
 Halhydric acids, relative stability of 
 
 53, 54, 88 
 Heats of formation of .... 87, 105 
 
INDEX. 
 
 543 
 
 Halhydric acids, tables of . . 87, 105, 266 
 
 Halogens, the 53 
 
 Anhydrides and acids of .... 134 
 
 Comparison of 53, 54 
 
 Compounds with each other . 132, 133 
 Oxy-acids of, comparison . . 179, 384 
 Oxy-acids of, heats of formation of, 162 
 Oxy-acids of, nomenclature of . . 119 
 Oxy-acids of, tables of ... 119, 133 
 
 Hausmannite 460 
 
 Hematite . 472 
 
 Heptane 279 
 
 Hexane 279 
 
 Hornblende 307,333 
 
 Humboldt 37, 165 
 
 Hydrargyllite 333 
 
 Hydrated acids 131, 135 
 
 Nomenclature of 227 
 
 Hydration 131 
 
 Hydrazin 192, 193 
 
 Hydro bromic acid 80,81 
 
 Heat of formation of 87 
 
 Hydrocarbons in coal-oil .... 279 
 Hydrochloric acid ...... fifi, fiz. 
 
 Action on metals 76 
 
 Action on oxides 76 
 
 Action on hydroxides 76 
 
 Heat of formation of . l . . . . 74, 87 
 Volumetric composition of . . . 68, 72 
 
 Hydrocyanic acid - 295, 296 
 
 Hydrofluoric acid 56, 57 
 
 Structure of 303 
 
 Hydrogen 27 
 
 Combustion of 35 
 
 Diffusion of 34 
 
 History of 27 
 
 Nascent 137, 207 
 
 Occlusion of 34 
 
 Occurrence of 27 
 
 Preparation of .... 28, 29, 31, 35- 
 
 Properties of 33, 35 
 
 Hydrogen and oxygen, ratio between 
 
 the atomic weights of ... 8, 39 
 Hydrogen antimonide .... 251, 263 
 
 Arsenide 236, 263 
 
 Boride 329 
 
 Hydrogen compounds of not-met- 
 
 als, comparison of . 90, 176, 177, 178 
 Comparative tables of ... 175, 176 
 
 Decompositions of 98 
 
 Heats of formation of .... 98, 108 
 
 Structure of 108 
 
 Unsaturated 109 
 
 Valence in 108, 109 
 
 Hydrogen hy peroxide 50, 51 
 
 Oxidizfng action of 51 
 
 Hydrogen persulphide 101 
 
 Phosphides 214, 217, 263 
 
 Hydrogen, selenide 103 
 
 Silicide 301 
 
 Hydrogen sulphide 96, 100 
 
 Composition of 99 
 
 Decomposition of 98, 99 
 
 Heat of formation of 98 
 
 Preparation of ' 96, 97, 183 
 
 Reactions of 100 
 
 Hydrogen Telluride 104 
 
 Hydroiodic acid 85, 86 
 
 Heat of formation of 87 
 
 Hydroxides 29,30,43 
 
 Decomposition of, by heat .... 260 
 
 Formation of 260 
 
 Of metals and not-metals, resem- 
 blances between 157 
 
 Hydroxyl . . 30, 31, 43, 117, 128, 131, 140 
 
 Hydroxylamin 191 
 
 Hydrochloride 192 
 
 Hyperoxides 25 
 
 Nature of 323 
 
 Hypobromous acid 129 
 
 Hypochlorites, formation of .... 121 
 
 Decomposition of 122, 123 
 
 Changes of, by heat 124 
 
 Hypochlorous acid 115, 121 
 
 Decomposition of 123 
 
 Oxidizing action of .... 123, 124 
 
 Properties of 123 
 
 Salts of 121 
 
 Hypophosphorous acid 232 
 
 Illuminating gas, formation of . 183, 282 
 
 Indium, properties of 344 
 
 Compounds of 345 
 
 lodates 132 
 
 lodic acid 130, 132 
 
 Iodine, history of 83 
 
 Oxy-acids of 130 
 
 Preparation, properties of ... 83, 84 
 
 Iodine monobromide 133 
 
 Monochloride 133 
 
 Pentafluoride 133 
 
 Pentoxide 131 
 
 Trichloride 133 
 
 Iodine starch 50 
 
 Iridium 489, 490, 492 
 
 Iron 31 
 
 Alloys of 478 
 
 Cast, gray 474 
 
 Cast, white 474 
 
 Chemically pure 476 
 
 Metallurgy of 473, 474 
 
 Occurrence of 472 
 
 Passive 477 
 
 Properties of 476 
 
 Spiegel 475 
 
 Sulphides of 31 
 
544 
 
 INDEX. 
 
 Iron, wrought 475 
 
 Iron, cobalt, and nickel, comparison 
 
 of 470 
 
 Properties of 470 
 
 Relation of, to periodic system 470, 471 
 
 Table of 488 
 
 Iron pyrites 91,472,484 
 
 Isomorphism 42, 356 
 
 Application to atomic weights 356, 357 
 
 Examples of 357, 358 
 
 Law of . 356 
 
 Kaolin 318, 333, 340 
 
 Lampblack 273 
 
 Lanthanum 437 
 
 Lavoisier 2, 15, 27, 36, 165 
 
 Law of definite proportions . . 3, 347, 348 
 3Iultiple proportions .... 3, 5, 347 
 
 Lazurite 397 
 
 Lead, metallurgy of 320 
 
 Occurrence of 320 
 
 Oxides of 322 
 
 Properties of 321 
 
 Red oxide of 324 
 
 Salts of 322 
 
 Table of compounds of 325 
 
 Acetate 323 
 
 Carbonate 323 
 
 Chloride 322 
 
 Chromate 322, 444, 450 
 
 Dioxide 323 
 
 Hydroxide 322 
 
 Monoxide 322 
 
 Molybdate 444 
 
 Nitrate 201 
 
 Sulphate 322, 417 
 
 Sulphide 91,324 
 
 Tungstate 445 
 
 Le Blanc soda process .... 389, 390 
 
 Leucite 307 
 
 Light, chemical action of ... 406, 407 
 
 Ligroine 279 
 
 Lime (see calcium oxide). 
 
 Limonite 472 
 
 Magnesite 416 
 
 Magnesium, isolation of 412 
 
 Properties of 412 
 
 Magnesium-ammonium arsenate 419, 420 
 Phosphate 419 
 
 Magnesium carbonate .415 
 
 Chloride 61, 415 
 
 Hydroxide 29, 76, 413 
 
 Oxide 413 
 
 Phosphate 419 
 
 Pyrophosphate 419 
 
 Secondary phosphate 419 
 
 Magnesium, sulphate .... 416, 417 
 
 Magnetite 472,483 
 
 Malachite 397 
 
 Manganates 465 
 
 Manganese, alloys of 460 
 
 Compounds of 460, 469 
 
 Occurrence of 459, 4GO 
 
 Oxides of 460 
 
 Properties of 460 
 
 Relation of compounds of .... 468 
 
 Relation of, to periodic system . . 459 
 
 Tables of compounds of ... 468, 469 
 
 Manganese dioxide 19, 20, 25, 60, 79, 84, 460 
 
 Chemistry of 460, 463 
 
 Hydroxides of 465 
 
 Occurrence of 460, 463 
 
 Manganite 460, 462 
 
 Manganites 465 
 
 Manganic alum 463 
 
 Chloride 462 
 
 Hydroxides 462 
 
 Oxide 460, 462 
 
 Sulphate 462 
 
 Manganous carbonate 462 
 
 Chloride 461 
 
 Hydroxide 461 
 
 Oxide 460 
 
 Sulphate 461 
 
 Sulphide 462 
 
 Manganous-manganic oxide 460, 461, 4.63 
 
 Markasite 91, 472 
 
 Marsh gas (see methane). 
 
 Mass action, laws of 378, 379 
 
 Melaconite 397 
 
 Mendelejeff 361,373 
 
 "Men dele Jeff's table of the periodic 
 
 system 371 
 
 Mercuric chloramide 434 
 
 Chloride 432, 433 
 
 Cyanide 294,435 
 
 Iodide 434 
 
 Nitrate 435 
 
 Oxide 19, 431 
 
 Sulphide 435 
 
 Mercurous chloramide .... 432, 434 
 
 Chloride 431 
 
 Iodide 434 
 
 Nitrate . . 434 
 
 Oxide 430. 
 
 Sulphide 435 
 
 Mercury, metallurgy of ...... 427 
 
 Occurrence of 427 
 
 Oxides of 430 
 
 Properties of 425 
 
 Salts of 431,435 
 
 Metals, characteristics of .... 12, 13 
 
 Oxides of 14 
 
 Metaphosphoric acid .... 228,233 
 
INDEX. 
 
 545 
 
 Methane, action of chlorine on ... 276 
 
 Occurrence of . . * 274 
 
 Preparation and properties of . . 275 
 
 Table of 279 
 
 Volumetric composition of .... 276 
 
 Methin 277 
 
 Chloride 277 
 
 Methyl 277 
 
 Methyl chloride 277 
 
 Methylen 277 
 
 Methylen chloride 277 
 
 Meyer, Lothar 361, 367 
 
 Mica 307 
 
 Millerite 473 
 
 Minium 324 
 
 Mitscherlich 356 
 
 Mixtures 5 
 
 Moissan 55 
 
 Molybdates 453,454 
 
 Molybdenites 444,452 
 
 Molybdenum, acids of 453 
 
 Halidesof 453 
 
 Hydroxides of 453 
 
 Occurrence of 444 
 
 Oxides of 452 
 
 Preparation and properties of . . 446 
 
 Sulphides of 454 
 
 Molybdenum dioxide 452 
 
 Disulphide 444 
 
 Monoxide 452 
 
 Pentachloride 453 
 
 Molybdic acids 453 
 
 Molybdite 444 
 
 Molecular weights, determination 
 
 of 349 
 
 Molecular weights of gases . . . 72, 73 
 
 Molecules 9 
 
 Endothermic and exothermic . . 12 
 
 Stability of 11 
 
 Muscovite 333 
 
 Nascent action 51, 137, 207 
 
 Neutralization, laws of .... 375-382 
 Relation to dissociation . . . 381, 382 
 
 Neutralization of acids, heat of . 143, 376 
 
 375 
 Laws of 375, 377 
 
 Nickel, alloys of 478 
 
 Cyanides of 487 
 
 Metallurgy of 477 
 
 Occurrence of 473 
 
 Properties of 478 
 
 Sulphate 417 
 
 Sulphides of 473 
 
 Nickelous chloride 487 
 
 Cyanide 487 
 
 Hydroxide 486 
 
 Niobium (see columbium). 
 
 Nitrates, decomposition of 208 
 
 Formation of 204,208 
 
 Nitric acid, anhydride of 202 
 
 History of 203 
 
 Preparation of 204 
 
 Production of 202, 203 
 
 Properties of 205,206 
 
 Reduction of .... 183, 198, 199, 207 
 Saltsof 208 
 
 Nitric bxide, oxidation of . . 198, 200, 201 
 
 Preparation of 198, 199 
 
 - Properties of 200 
 
 Nitrites 208 
 
 Nitro group 147 
 
 Nitrogen, acids of 195 
 
 History of 163, 164 
 
 In atmosphere 166 
 
 Occurrence of 164 
 
 Oxides of 180, 195 
 
 Preparation of 164 
 
 Properties of 165 
 
 Nitrogen dioxide 198,201 
 
 Formation of 202 
 
 Properties of 202 
 
 Nitrogen oxides, heats of formation of 210 
 Table of 209 
 
 Nitrogen pentoxide 202 
 
 Trichloride 180 
 
 Trioxide 147, 148, 201 
 
 Nitrogen family, comparison of mem- 
 bers 176 
 
 Halides of, comparison of . . 180, 181 
 
 Halides of, table of 263 
 
 Hydrogen compounds of 176, 177, 263, 266 
 Oxides of, comparison of . . 180, 263 
 Sulphides and sulpho-salts, table of, 
 
 263, 264 
 
 Table of atomic weights of . . 175, 262 
 
 Table of hydrogen compounds of . 175 
 
 218, 263 
 
 Table of oxides and acids of, 179, 263, 264 
 Table of properties of ... 175, 262 
 Table of specific gravities of vapors 
 
 of 262 
 
 Valence in oxides of 115 
 
 Nitrosyl-sulphuric acid, decomposi- 
 tion of 148 
 
 Formation of 147 
 
 Nitrous oxide 196 
 
 Composition of 197 
 
 Effects of 197, 198 
 
 Preparation of 196 
 
 Properties of 196, 197 
 
 Not-metals, characteristics of . . 12, 13 
 Oxides of 14 
 
 Occlusion of gases 34 
 
 Oligoclase 308 
 
546 
 
 INDEX. 
 
 Olivin 307 
 
 Orpiment 234 
 
 Orthoclase 308 
 
 Ortho-phosphoric acid . . . 229,233 
 
 Salts of 229 
 
 Osmium 489 
 
 Dioxide of 492 
 
 Oxides of 490 
 
 Osmosis 305 
 
 Osmotic pressure 306 
 
 Osteolite 211,418 
 
 Oxide of manganese, black (see man- 
 ganese dioxide). 
 
 Oxides 21,24 
 
 Oxides, classification of 25, 26 
 
 Formulae of, and valence of . 112, 113 
 
 Nomenclature of 26 
 
 Valence in 110 
 
 Oxygen 18 
 
 Combustion in 22 
 
 Discovery of 18 
 
 In atmosphere 166 
 
 Occurrence of 18 
 
 Preparation of 18, 19, 20 
 
 Properties of 21 
 
 Oxygen family, atomic weights of . 106 
 
 Comparative table of 105 
 
 Comparison of 89, 90 
 
 Hydrogen compounds of, 89, 90, 105, 266 
 
 Melting points of 105 
 
 Oxides and oxy-acids of 135 
 
 Specific gravities of 105 
 
 Specific gravities of gases .... 105 
 
 Valence in oxides of 115 
 
 Oxy-hydrogen blow-pipe .... 36 
 
 Ozone 47 
 
 Composition of 48 
 
 History of 47 
 
 Oxidizing action of .... 50, 51, 52 
 
 Preparation of 47, 49, 56 
 
 Properties of 49 
 
 Palladium 33, 489, 490 
 
 Cyanide 489 
 
 Dioxide N . . 492 
 
 Palladium-hydrogen 33,207 
 
 Pentathionic acid 155 
 
 Pentane 279 
 
 Pentlandite 473 
 
 Pentoxide of iodine ..... 130, 131 
 
 Nitrogen 203 
 
 Phosphorus 22 
 
 Perchlorate of potassium, 
 
 Decomposition of 125 
 
 Formation of 125 
 
 Properties of 126 
 
 Perchloric acid 119 
 
 Properties of 120 
 
 Per-iodates 132 
 
 Per-iodic acids 130 
 
 Hydrated 131 
 
 Periodic system , the .... 361-374 
 
 Arrangement of 362, 363 
 
 Relations of atomic volumes to 365, 366 
 
 Relations of families in 364 
 
 Relations of halides to ..... 370 
 
 Relations of hydrides to 370 
 
 Relations of oxides to ... 370, 372 
 Relations of valence to 369 
 
 Permanganic acid 466 
 
 Anhydride 460, 466 
 
 Petroleum 279 
 
 Ether 279 
 
 Phlogiston 18,27 
 
 Phosgen (see carbonyl chloride). 
 
 Phosphates, ortho, meta, and pyro . 228 
 
 229, 231 
 
 Phosphine, liquid 217 
 
 Preparation of 214 
 
 Properties of 215, 216 
 
 Volumetric composition of .... 215 
 
 Phosphites 227 
 
 Phosphonium compounds . . . 216, 215 
 Iodide 217 
 
 Phosphorescence 213 
 
 Phosphoric acids, 95, 214, 220, 221, 224, 226 
 
 Nomenclature of 222, 226 
 
 Oxidation of 225 
 
 Reactions of 226 
 
 Salts of 225 
 
 Phosphorus, acids of, table of ... 233 
 
 Allotropic forms of . 212 
 
 Halides of, table of 219 
 
 Halides of", reactions of 220 
 
 History of 211,212 
 
 Hydrogen compounds of .... 214 
 
 Occurrence of 211 
 
 Oxidation of 214 
 
 Oxides of, table of .... 223, 231, 233 
 
 Oxy-halides of 221 
 
 Poisonous action of 213 
 
 Red 213 
 
 Yellow 213,214 
 
 Phosphorus oxybromide 221 
 
 Oxychloride 221 
 
 Pentabromide 219, 220 
 
 Pentachloride . . . 181~ 219, 220, 263 
 
 Pentafluoride 219, 220 
 
 Pentasulphide 95 
 
 Pentoxide 22,214,223 
 
 Tribromide 81, 219, 220 
 
 Trichloride . . 75, 81, 181, 219, 220, 263 
 
 Trifluoride 219 
 
 Tri-iodide 85,219 
 
 Trioxide 223 
 
 Trisulphide 95 
 
INDEX. 
 
 547 
 
 Photography 406 
 
 Pitchblende A. . . 445 
 
 Plaster of Paris 418 
 
 Platinic chloride 492,493 
 
 Platinous chloride 495 
 
 Platinum, cyanides of 489 
 
 Dioxide of 492 
 
 Oxides of 490 
 
 Platinum group, comparison of . . 489 
 
 Cyanides of 489 
 
 Halidesof 491 
 
 Occurrence of 491 
 
 Properties of 490, 491 
 
 Utensils of 495 
 
 Plumbago (see graphite).. 
 
 Polysulphides 155, 386 
 
 Structure of 387 
 
 Porcelain 341 
 
 Potassium 30 
 
 Bromate 129 
 
 Brom-aurate 409 
 
 Chlorate 20, 120 
 
 Chlorate, formation of 124 
 
 Chlorate, decomposition of ... 125 
 
 Chlor-aurate 409 
 
 Chlorite 120 
 
 Chromate 449,465 
 
 Bichromate 450 
 
 Ferricyanide 297 
 
 Ferrocyanide 296 
 
 Fluosilicate 304 
 
 Hydroxide 30, 43, 387 
 
 Hypobromite 129 
 
 Hypochlorite .... 120, 121, 122, 123 
 Hypochlorite, decomposition of, by 
 
 heat 124 
 
 Hyponitrite 208 
 
 Iodide 50 
 
 lodo-aurate 409 
 
 Manganate 465 
 
 Nitrate 391 
 
 Pentamanganite 465 
 
 Permanganate 467, 468 
 
 Permanganate, decomposition of . 467 
 Permanganate, oxidizing action of, 467 
 
 Perchlorate 120, 125, 126 
 
 Sulphides 386,387 
 
 Uranate. 457 
 
 Precipitate, red 18, 19 
 
 Priestley 18, 27, 182, 196, 201 
 
 Primary salts, change to secondary, 140, 152 
 
 Prince Rupert's drops 422 
 
 Pyrolusite 460 
 
 Pyr phosphoric acid . . . . 231, 233 
 Pyrosulphuric acid (see disulphuric acid). 
 
 Quartz 304 
 
 Raoult .350 
 
 Raoult's Law \ 350, 351 
 
 Reactions, endothermic and exother- 
 mic 12 
 
 Realgar 234 
 
 Recrystallization 41 
 
 Reinite 445 
 
 Rhodium 489 
 
 Cyanides of 489 
 
 Dioxide of 492 
 
 Oxides of 490 
 
 Rutile 438 
 
 Saltpetre, Chile 204, 390 
 
 Salts, formation of ... 25, 30, 32, 77, 78 
 
 Acid and neutral 140 
 
 Basic 252 
 
 Primary and secondary HO 
 
 Reactions of 141 
 
 Scandium, chemistry of 437 
 
 Predicted properties of ... 373, 437 
 
 Scheele 58 
 
 Scheelite 445 
 
 Secondary salts, change to primary, 140, 152 
 
 Selenious acid 160 
 
 Selenites, primary and secondary . . 160 
 Selenium, allotropic forms .... 102 
 
 Dioxide of 160 
 
 History of 102 
 
 Isolation of 102 
 
 Monosulphide of 161 
 
 Occurrence of ... t 102 
 
 Properties of 103 
 
 Trioxideof 161 
 
 Siderite . ' 292, 472 
 
 Silicates, di 307 
 
 Disintegration of 340 
 
 Meta 267,307 
 
 Ortho 267, 307 
 
 Tri 308 
 
 Silicic acids, di- 307 
 
 Dialyzed 305 
 
 Meta- 304 
 
 Ortho- 304 
 
 Tri- 307 
 
 Silico-chloroform 301 
 
 Silicon, occurrence of 300 
 
 Preparation of 300 
 
 Properties of 301 
 
 Halides of 301, 302 
 
 Hydride 301 
 
 Oxides of 304 
 
 Dioxide 304 
 
 Dioxide, amorphous 304 
 
 Tetrachloride :XX),,301 
 
 Tetrafluoride 301, 302 
 
 Silver, alloys of 400 
 
 Halides of 405 
 
 Halides of, uses in photography . . 406 
 Metallurgy of 398, 399 
 
548 
 
 INDEX. 
 
 Silver, occurrence of . .... 398 
 
 Oxides of 405 
 
 Chloride 405 
 
 Hyponitrite 209 
 
 Nitrate 407 
 
 Sulphide 407 
 
 Slag 473 
 
 Smaltite 234,472 
 
 Sodium 29 
 
 Acetate 275 
 
 Sodium carbonate 43, 290, 291, 292, 389, 390 
 
 Preparation of 389, 390 
 
 Chloride . .... 7ft, 142 
 
 Dimolybdate 453 
 
 Hydroxide 29 
 
 Hyposulphite (see sodium thiosulphate). 
 
 Metaphosphate 229 
 
 Metavanadate 441 
 
 Nitrate 204,390 
 
 Octomolybdate 453 
 
 Orthovanadate 441 
 
 Phosphate, primary 229 
 
 Phosphate, secondary 229 
 
 Pyrovanadate 441 
 
 Silicate 304 
 
 Sulphate 142 
 
 Sulphate, action of hydrochloric 
 
 acid on 142 
 
 Sulphate, primary 152 
 
 Sulphate, secondary 152 
 
 Tetramolybdate 453 
 
 Thiosulphate 154,406 
 
 Trimolybdate 453 
 
 Tungstate ; 455 
 
 ions 41 
 
 ;cific gravities of gases ... 72, 73 
 
 jpe, the 393 
 
 jpic analysis 393-395 
 
 im , the 393 
 
 Absorption 395 
 
 Sphserocobaltite 472 
 
 Spinells, the 333, 339, 438, 444, 447, 460, 481 
 
 Stannic acid, ortho 
 Meta . 
 
 267 
 
 267 
 
 a 317 
 
 317 
 
 a and ft, table of 319 
 
 Stannic chloride 315, 317 
 
 Double salts with 316 
 
 Oxide 316 
 
 Sulphide 318 
 
 Stannous acid 315 
 
 .Chloride 313, 314, 433 
 
 'Hydroxide 315 
 
 Oxide 316 
 
 Sulphide 318 
 
 Status nascendi 51, 137 
 
 Steel . . . . 476 
 
 Steel, preparation of 475, 476 
 
 Stibine 49 251 
 
 Stibnite 248 
 
 Stibionyl 253 
 
 Stoichiometric quantities .... 6 
 
 Stoltzite 445 
 
 Strass 421 
 
 Strontianite 292,416 
 
 Properties of 412 
 
 Isolation of 412 
 
 Strontium carbonate 292, 416 
 
 Chlorate 420 
 
 Chloride 415 
 
 Hydroxide 413 
 
 Nitrate 420 
 
 Oxide 413 
 
 Sulphate 416, 417 
 
 Structure of hydrogen compounds, 108 
 
 Substitution 29,277 
 
 Sulphantimonites 255 
 
 Sulphantimonates 256 
 
 Sulpharsenates 245, 246, 247 
 
 Sulpharsenites 244, 245, 247 
 
 Sulphates, primary and secondary, 152, 153 
 
 Acid and neutral 152 
 
 Sulphides, formulae of 94, <J5 
 
 Formation of 100 
 
 Solubility of 100 
 
 Sulphites, change by heat 139 
 
 Sulphocarbonates (see thiocarbonates). 
 
 Sulphoferrites 483 
 
 Sulphostannates 318 
 
 Sulphur, allotropic forms of .... 93 
 
 Crystalline forms of 93 
 
 Halidesof ., 156 
 
 History of 91 
 
 Isolation of 92 
 
 Occurrence of 91 
 
 Oxides and oxy-acids of ... 134, 156 
 Oxy- and sulpho- acids of, table . . 159 
 
 Plastic 94 
 
 Properties of 93 
 
 Rare oxides of ... ./.' ... 158 
 Resemblances to chlorine .... 95 
 
 Sulphur dioxide 22, 117, 134 
 
 History of 136 
 
 Occurrence of 135 
 
 Oxidation of 138 
 
 Preparation of 136, 137, 148 
 
 Properties of 137, 138 
 
 Sulphur trioxide 117, 134, 144 
 
 Preparation of 144 
 
 Properties of 144 
 
 Dissociation of 145 
 
 Sulphur family, comparison with 
 
 / chromium family 443 
 
 Heats of formation of compounds 
 
 of. 162 
 
INDEX. 
 
 549 
 
 Sulphur family, oxides and acids 
 
 of 135, 162 
 
 Sulphuric acid 117 
 
 Action of, on metals 137 
 
 Chambers 149 
 
 Commercial preparation of . . 148, 149 
 
 Concentration of 150 
 
 Constitution of 145 
 
 Continuous process of manufacture 
 
 of 147 
 
 Heat of solution of .... 151, 152 
 
 History of 146 
 
 Hydrated 145, 151 
 
 Hydrated, salts of 153 
 
 Hydration of 150, 151 
 
 Manufacture of 147 
 
 Properties of 150 
 
 Reduction of 136, 151 
 
 Uses of, in laboratory 152 
 
 Sulph'hydrates, formulae of .... 94 
 Sulphuretted hydrogen .... 96-101 
 
 Sulphurous acid 139 
 
 Change of, on heating 139 
 
 Oxidation of , m 139 
 
 Structure of 158 
 
 Sulphuryl-chloride 138, 157 
 
 Decomposition of 115 
 
 Sun, chromosphere of 395 
 
 Superphosphates 418 
 
 .Jiymbols, chemicul 14tJLfi, 19, 20 
 
 Tantalates 441 
 
 Tantalum 441 
 
 Tellurates 161 
 
 Telluric acid 161 
 
 Tellurium 104 
 
 Occurrence of 104 
 
 Properties of 104 
 
 Dioxide 160 
 
 Disulphide 162 
 
 Trisulphide 162 
 
 Tetradymite 257 
 
 Tetrahedrite 397 
 
 Tetrathionic acid 155 
 
 Thallium 345 
 
 Chlorides of 346 
 
 Oxides of 345 
 
 Properties of 345 
 
 Sulphides of 346 
 
 Theory, dualistic 13 
 
 Thiocarbonates 293 
 
 Thionyl-chloride 157 
 
 Thionic acids, constitution of ... 156 
 
 Thiosulphates 154, 155 
 
 Thorium 439 
 
 Tin ] ' 312 
 
 Dioxide 316 
 
 Halidesof 314,315 
 
 Tin, history of 312 
 
 Hydroxides of 315 
 
 Metallurgy of 312 
 
 Occurrence of 312 
 
 Oxides of 316 
 
 Properties of 313 
 
 Sulphides of 318 
 
 Table of compounds of 319 
 
 Monoxide 316 
 
 Tin-stone (see cassiterite). 
 
 Titanium 438,439 
 
 Torricelli 171 
 
 Tridymite 304 
 
 Trisilicates 308 
 
 Trithionic acid 155 
 
 Troilite 472 
 
 Tungsten 445 
 
 Acids of 455 
 
 Halides of 454, 455 
 
 Occurrence of 445 
 
 Preparation of 445 
 
 Properties of 446 
 
 Dioxide 454 
 
 Hexachloride 455 
 
 Trioxide 444, 445 
 
 Tungstic acids 455 
 
 Tungstite 445 
 
 Unsaturated hydrogen compounds, 108 
 
 Uranates 457 
 
 Uranic oxide 456 
 
 acid 457 
 
 Uranium, occurrence of 445 
 
 Oxides of 456 
 
 Preparation of 
 
 Properties of 
 
 Uranous oxide 4 
 
 Uranous-uranic oxide ...... 4 
 
 Uranyl 456 
 
 Chemical relations of 457 
 
 Salts of 456 
 
 Hydroxide 456, 457 
 
 Urea 298,299 
 
 Valence 107 
 
 Of elements towards chlorine . . . 109 
 Of elements towards hydrogen . 107, 108 
 
 281 
 
 Of elements towards oxygen ... Ill 
 Of elements towards polyvalent 
 
 ones no, 112 
 
 Law of 111,112,113,114 
 
 Relation to atomic weights . 114, 369 
 
 Varying , 115 
 
 Vanadates 441 
 
 Vanadinite 440 
 
 Vanadium, acids of 441 
 
 Chlorides of , .441 
 
550 
 
 INDEX. 
 
 Vanadium, comparison with phospho- 
 
 rus and arsenic ...... 441 
 
 Occurrence of ........ 440 
 
 Oxides of .......... 440 
 
 Vitriols, the .......... 417 
 
 Vivianite ........... 211 
 
 Volumes of gases, law of combina- 
 
 tion of .......... G& 
 
 "Water, chemical history of ... 30, 37 
 Composition of, by volume . 37, 72M3 
 Composition of, by weight .... f 38 
 
 Crystallization of ...... 42, 43 
 
 Decomposition of, by chlorine . . 63 
 Decomposition of, by metals . 29, 30, 31 
 Energy of formation of ..... 45 
 
 Formation of ........ 34, 36 
 
 Properties of ......... 40 
 
 Quantitative composition of ... 39 
 Solution in .......... 41 
 
 Vapor in atmosphere ...... 169 
 
 "Waters., natural ......... 45 
 
 Pollution of ......... 45 
 
 Mineral ............ 45 
 
 Wavellite ........... 340 
 
 Witherite ......... 292,416 
 
 Wollastonite 
 
 "Work 
 
 Measure of 
 
 Wulf enite 
 
 Wurtzite ....... 
 
 307 
 
 10 
 
 11 
 
 444 
 
 426 
 
 Ytterbium 
 Yttrium . 
 
 437 
 437 
 
 Zinc 
 
 Alloys of . . 
 
 Hydroxide of . 
 
 Metallurgy of . 
 
 Occurrence of. 
 
 Oxide .... 
 
 Properties of . 
 Zinc blende . 
 
 32 
 
 ...... 427 
 
 . . ' . .30, 76, 428 
 
 427 
 
 426 
 
 428 
 
 428 
 
 ... 91, 344, 426 
 Zinc, Cadmium, Mercury, compari- 
 son of 423, 424 
 
 Compounds of 423 
 
 Relations of, to periodic system . . 423 
 Specific gravities of vapors .... 426 
 
 Table of compounds of 436 
 
 Zincates 428 
 
 Zinc carbonate 429 
 
 Chloride . 76, 429 
 
 Sulphate 417, 429 
 
 Sulphide 91, 97, 430 
 
 Telluride 104 
 
 Zincite 426 
 
 Zircon 438 
 
 Zirconium, compounds of 439 
 
 Occurrence of ... . .^ . . 438 
 Tetrachloride v*?T . . . 439 
 
 Appendix of Laboratory Notes, 497 to 536 
 
 UNIVERSITY 
 
U. C. BERKELEY LIBRARIES 
 
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