Chapter 36

Footnotes:[1]But it is impossible to foretell all the compounds formed by an element from its atomicity or valency, because the atomicity of the elements is variable, and furthermore this variability is not identical for different elements. In CO2, COX2, CH4, and the multitude of carbon compounds corresponding with them, the C is quadrivalent, but in CO either the carbon must be taken as bivalent or the atomicity of oxygen be accounted as variable. Moreover, carbon is an example of an element which preserves its atomicity to a greater degree than most of the other elements. Nitrogen in NH3, NH2(OH), N2O3, and even in CNH, must be considered as trivalent, but in NH4Cl, NO2(OH), and in all their corresponding compounds it is necessarily pentavalent. In N2O, if the atomicity of oxygen = 2, nitrogen has an uneven atomicity (1, 3, 5), whilst in NO it is bivalent. If sulphur be bivalent, like oxygen, in many of its compounds (for example, H2S, SCl2, KHS, &c.), then it could not be foreseen from this that it would form SO2, SO3, SCl4, SOCl2, and a series of similar compounds in which its atomicity must be acknowledged as greater than 2. Thus SO2, sulphurous anhydride, has many points in common with CO2, and if carbon be quadrivalent then the S in SO2is quadrivalent. Therefore the principle of atomicity (valency) of the elements cannot be considered established as the basis for the study of the elements, although it gives an easy method of grasping many analogies. I consider the four following as the chief obstacles to acknowledging the atomicity of the elements as a primary conception for the consideration of the properties of the elements: 1. Such univalent elements as H, Cl, &c., appear in a free state as molecules H2, Cl2, &c., and are consequently like the univalent radicles CH3, OH, CO2H, &c., which, as might be expected, appear as C2H6, O2H2, C2O4H2(ethane, hydrogen peroxide, oxalic acid), whilst on the other hand, potassium and sodium (perhaps also iodine at a high temperature) contain only one atom, K, Na, in the molecule in a free state. Hence it follows thatfree affinitiesmay exist. Granting this, nothing prevents the assumption that free affinities exist in all unsaturated compounds; for example, two free affinities in NH3. If such instances of free affinities be admitted, then all the possible advantages to be gained by the application of the doctrine of atomicity (valency) are lost. 2. There are instances—for example, Na2H—where univalent elements are combined in molecules which are more complex than R2, and form molecules, R3, R4, &c.; this may again be either taken as evidence of the existence of free affinities, or else necessitates such primary univalent elements as sodium and hydrogen being considered as variable in their atomicity. 3. The periodic system of the elements, with which we shall afterwards become acquainted, shows that there is a law or rule for the variation of the forms of oxygen and hydrogen compounds; chlorine is univalent with respect to hydrogen, and septavalent with respect to oxygen; sulphur is bivalent to hydrogen, and sexavalent to oxygen; phosphorus is trivalent to hydrogen and pentavalent in respect to oxygen—the sum is in every case equal to 8. Only carbon and its analogues (for example, silicon) are quadrivalent to both hydrogen and oxygen. Hence the power of the elements to change their atomicity is an essential part of their nature, and therefore constant valency cannot he considered as a fundamental property. 4. Crystallo-hydrates (for instance, NaCl,2H2O, or NaBr,2H2O), double salts (such as PtCl4,2KCl,H2SiF6, &c.), and similar complex compounds (and, according to Chap.I., solutions also) demonstrate the capacity not only of the elements themselves, but also of their saturated and limiting compounds, of entering into further combination. Therefore the admission of a definite limited atomicity of the elements includes in itself an admission of limitation which is not in accordance with the nature of chemical reactions.[2]The primary formations are those which do not bear any distinct traces of having been deposited from water (have not a stratified formation and contain no remains of animal or vegetable life), occur under the sedimentary formations of the earth, and are everywhere uniform in composition and structure, the latter being generally distinctly crystalline. If it be assumed that the earth was originally in a molten condition, the first primary formations are those which formed the first solid crust of the earth. But even with this hypothesis of the earth's origin, it is necessary to admit that the first aqueous deposits must have caused a change in the original crust of the earth, and therefore under the head of primary formations must be understood the most ancient of the products of decomposition (mostly by atmospheric, aqueous, and organic agency, &c.), from which all the rocks and substances of the earth's surface have arisen. In speaking of the origin of one or another substance, we can only, on the basis of facts, descend to the primary formations, of which granite, gneiss, and trachyte may be taken as examples.[2 bis]Chloride of sodium has been found to occur in the atmosphere in the form of a fine dust; in the lower strata it is present in larger quantities than in the upper, so that the rain water falling on mountains contains less NaCl than that falling in valleys. Müntz (1891) found that a litre of rain water collected on the summit of the Pic du Midi (2,877 metres above the sea level) contained 0·34 milligram of chloride of sodium, while a litre of rain collected from the valley contained 2·5–7·6 milligrams.[3]The extraction of the potassium salts (or so-called summer salts) was carried on at the Isle of Camarga about 1870, when I had occasion to visit that spot. At the present time the deposits of Stassfurt provide a much cheaper salt, owing to the evaporation and separation of the salt being carried on there by natural means and only requiring a treatment and refining, which is also necessary in addition for the ‘summer salt’ obtained from sea-water.[4]The double salt KCl,MgCl2is a crystallohydrate of KCl and MgCl2, and is only formed from solutions containing an excess of magnesium chloride, because water decomposes this double salt, extracting the more soluble magnesium chloride from it.[5]Owing to the fundamental property of salts of interchanging their metals, it cannot be said that sea water contains this or that salt, but only that it contains certain amounts of certain metals M (univalent like Na and K, and bivalent like Mg and Ca), and haloids X (univalent like Cl, Br, and bivalent like SO4, CO3), which are disposed in every possible kind of grouping; for instance, K as KCl, KBr, K2SO4, Mg as MgCl2, MgBr2, MgSO4, and so on for all the other metals. In evaporation different salts separate out consecutively only because they reach saturation. A proof of this may be seen in the fact that a solution of a mixture of sodium chloride and magnesium sulphate (both of which salts are obtained from sea water, as was mentioned above), when evaporated, deposits crystals of these salts, but when refrigerated (if the solution be sufficiently saturated) the salt Na2SO4,10H2O is first deposited because it is the first to arrive at saturation at low temperatures. Consequently this solution contains MgCl2and Na2SO4, besides MgSO4and NaCl. So it is with sea water.[6]The salt extracted from water is piled up in heaps and left exposed to the action of rain water, which purifies it, owing to the water becoming saturated with sodium chloride and then no longer dissolving it, but washing out the impurities.[7]When the German savants pointed out the exact locality of the Stassfurt salt-beds and their depth below the surface, on the basis of information collected from various quarters respecting bore-holes and the direction of the strata, and when the borings, conducted by the Government, struck a salt-bed which was bitter and unfit for use, there was a great outcry against science, and the doubtful result even caused the cessation of the further work of deepening the shafts. It required a great effort to persuade the Government to continue the work. Now, when the pure salt encountered below forms one of the important riches of Germany, and when those ‘refuse salts’ have proved to be most valuable (as a source of potassium and magnesium), we should see in the utilisation of the Stassfurt deposits one of the conquests of science for the common welfare.[7 bis]In Western Europe, deposits of rock salt have long been known at Wieliczka, near Cracow, and at Cardona in Spain. In Russia the following deposits are known: (a) the vast masses of rock salt (3 square kilometres area and up to 140 metres thick) lying directly on the surface of the earth at Iletzky Zastchit, on the left bank of the river Ural, in the Government of Orenburg; (b) the Chingaksky deposit, 90 versts from the river Volga, in the Enotaeffsky district of the Government of Astrakhan; (c) the Kulepinsky (and other) deposits (whose thickness attains 150 metres), on the Araks, in the Government of Erivan in the Caucasus; (d) the Katchiezmansky deposit in the province of Kars; (e) the Krasnovodsky deposit in the Trans-Caspian province; and (f) the Bardymkulsky salt mines in Kokhand.[8]The fracture of rock salt generally shows the presence of interlayers of impurities which are sometimes very small in weight, but visible owing to their refraction. In the excellently laid out salt mines of Briansk I counted (1888), if my memory does not deceive me, on an average ten interlayers per metre of thickness, between which the salt was in general very pure, and in places quite transparent. If this be the case, then there would be 350 interlayers for the whole thickness (about 35 metres) of the bed. They probably correspond with the yearly deposition of the salt. In this case the deposition would have extended over more than 300 years. This should be observable at the present day in lakes where the salt is saturated and in course of deposition.[9]My own investigations have shown that not only the sulphates, but also the potassium salts, are entirely removed by this method.[10]According to the determinations of Klodt, the Briansk rock salt withstands a pressure of 340 kilograms per square centimetre, whilst glass withstands 1,700 kilos. In this respect salt is twice as secure as bricks, and therefore immense masses may be extracted from underground workings with perfect safety, without having recourse to brickwork supports, merely taking advantage of the properties of the salt itself.[11]To obtain well-formed crystals, a saturated solution is mixed with ferric chloride, several small crystals of sodium chloride are placed at the bottom, and the solution is allowed to evaporate slowly in a vessel with a loose-fitting cover. Octahedral crystals are obtained by the addition of borax, urea, &c., to the solution. Very fine crystals are formed in a mass of gelatinous silica.[12]If a solution of sodium chloride be slowly heated from above, where the evaporation takes place, then the upper layer will become saturated before the lower and cooler layers, and therefore crystallisation will begin on the surface, and the crystals first formed will float, having also dried from above, on the surface until they become quite soaked. Being heavier than the solution the crystals are partially immersed under it, and the following crystallisation, also proceeding on the surface, will only form crystals along the side of the original crystals. A funnel is formed in this manner. It will be borne on the surface like a boat (if the liquid be quiescent), because it will grow more from the upper edges. We can thus understand this at first sight strange funnel form of crystallisation of salt. In explanation why the crystallisation under the above conditions begins at the surface and not at the lower layers, it must be mentioned that the specific gravity of a crystal of sodium chloride = 2·16, and that of a solution saturated at 25° contains 26·7 p.c. of salt and has a specific gravity at 25°/4° of 1·2004; at 15° a saturated solution contains 26·5 p.c. of salt and has a sp. gr. 1·203 at 15°/4°. Hence a solution saturated at a higher temperature is specifically lighter, notwithstanding the greater amount of salt it contains. With many substancessurface crystallisationcannot take place because their solubility increases more rapidly with the temperature than their specific gravity decreases. In this case the saturated solution will always be in the lower layers, where also the crystallisation will take place. Besides which it may be added that as a consequence of the properties of water and solutions, when they are heated from above (for instance, by the sun's rays), the warmer layers being the lightest remain above, whilst when heated from below they rise to the top. For this reason the water at great depths below the surface is always cold, which has long been known. These circumstances, as well as those observed by Soret (Chapter I., Note19), explain the great differences of density and temperature, and in the amount of salts held in the oceans at different latitudes (in polar and tropical climes) and at various depths.[13]By combining the results of Poggiale, Müller, and Karsten (they are evidently more accurate than those of Gay-Lussac and others) I found that a saturated solution att°, from 0° to 108°, contains 35·7 + 0·024t+ 0·0002t2grams of salt per 100 grams of water. This formula gives a solubility at 0° = 35·7 grams (= 26·3 p.c.), whilst according to Karsten it is 36·09, Poggiale 35·5, and Müller 35·6 grams.[14]Perfectly purefusedsalt is not hygroscopic, according to Karsten, whilst the crystallised salt, even when quite pure, attracts as much as 0·6 p.c. of water from moist air, according to Stas. (In the Briansk mines, where the temperature throughout the whole year is about +10°, it may be observed, as Baron Klodt informed me, that in the summer during damp weather the walls become moist, while in winter they are dry).If the salt contain impurities—such as magnesium sulphate, &c.—it is more hygroscopic. If it contain any magnesium chloride, it partially deliquesces in a damp atmosphere. The crystallised and not perfectly pure salt decrepitates when heated, owing to its containing water. The pure salt, and also the transparent rock salt, or that which has been once fused, does not decrepitate. Fused sodium chloride shows a faint alkaline reaction to litmus, which has been noticed by many observers, and is due to the presence of sodium oxide (probably by the action of the oxygen of the atmosphere). According to A. Stcherbakoff very sensitive litmus (washed in alcohol and neutralised with oxalic acid) shows an alkaline reaction even with the crystallised salt.It may be observed that rock salt sometimes contains cavities filled with a colourless liquid. Certain kinds of rock salt emit an odour like that of hydrocarbons. These phenomena have as yet received very little attention.[15]By cooling a solution of table salt saturated at the ordinary temperature to -15°, I obtained first of all well-formed tabular (six-sided) crystals, which when warmed to the ordinary temperature disintegrated (with the separation of anhydrous sodium chloride), and then prismatic needles up to 20 mm. long were formed from the same solution. I have not yet investigated the reason of the difference in crystalline form. It is known (Mitscherlich) that NaI,2H2O also crystallises either in plates or prisms. Sodium bromide also crystallises with 2H2O at the ordinary temperature.[16]Notwithstanding the great simplicity (Chapter I., Note49) of the observations on the formation of ice from solution, still even for sodium chloride they cannot yet be considered as sufficiently harmonious. According to Blagden and Raoult, the temperature of the formation of ice from a solution containingcgrams of salt per 100 grams of water = -0·6ctoc= 10, according to Rosetti = -0·649ctoc= 8·7, according to De Coppet (toc= 10) = -0·55c- 0·006c2, according to Karsten (toc= 10) - 0·762c+ 0·0084c2, and according to Guthrie a much lower figure. By taking Rosetti's figure and applying the rule given in Chapter I., Note49we obtain—i= 0·649 ×58·5/18·5= 2·05.Pickering (1893) gives forc= 1 - 0·603, forc= 2 - 1·220; that is (cup to 2·7) about - (0·600 + 0·005c)c.The data for strong solutions are not less contradictory. Thus with 20 p.c. of salt, ice is formed at -14·4° according to Karsten, -17° according to Guthrie, -17·6° according to De Coppet. Rüdorff states that for strong solutions the temperature of the formation of ice descends in proportion to the contents of the compound, NaCl,2H2O (per 100 grams of water) by 0°·342 per 1 gram of salt, and De Coppet shows that there is no proportionality, in a strict sense, for either a percentage of NaCl or of NaCl,2H2O.[17]A collection of observations on the specific gravity of solutions of sodium chloride is given in my work cited in Chapter I., Note50.Solutions of common salt have also been frequently investigated as regards rate ofdiffusion(ChapterI.), but as yet there are no complete data in this respect. It may be mentioned that Graham and De Vries demonstrated that diffusion in gelatinous masses (for instance, gelatin jelly, or gelatinous silica) proceeds in the same manner as in water, which may probably lead to a convenient and accurate method for the investigation of the phenomena of diffusion. N. Umoff (Odessa, 1888) investigated the diffusion of common salt by means of glass globules of definite density. Having poured water into a cylinder over a layer of a solution of sodium chloride, he observed during a period of several months the position (height) of the globules, which floated up higher and higher as the salt permeated upwards. Umoff found that at a constant temperature the distances of the globules (that is, the length of a column limited by layers of definite concentration) remain constant; that at a given moment of time the concentration,q, of different layers situated at a depthzis expressed by the equation B - Kz= log.(A -q), where A, B, and K are constants; that at a given moment the rate of diffusion of the different layers is proportional to their depth, &c.[18]IfS_0 be the specific gravity of water, andSthe specific gravity of a solution containingpp.c. of salt, then by mixing equal weights of water and the solution, we shall obtain a solution containing ½pof the salt, and if it be formed without contraction, then its specific gravityxwill be determined by the equation2/x=1/S0+1/S, because the volume is equal to the weight divided by the density. In reality, the specific gravity is always found to be greater than that calculated on the supposition of an absence of contraction.[19]Generally the specific gravity is observed by weighing in air and dividing the weight in grams by the volume in cubic centimetres, the latter being found from the weight of water displaced, divided by its density at the temperature at which the experiment is carried out. If we call this specific gravity S1, then as a cubic centimetre of air under the usual conditions weighs about 0·0012 gram, the sp. gr. in a vacuum S = S1+ 0·0012 (S1- 1), if the density of water = 1.[20]If the sp. gr. S2be found directly by dividing the weight of a solution by the weight of water at the same temperature and in the same volume, then the true sp. gr.Sreferred to water at 4° is found by multiplying S2by the sp. gr. of water at the temperature of observation.[21]According to Schiff 100 grams of alcohol, containingpp.c. by weight of C2H6O, dissolves at 15°—p=102040608028·522·613·25·91·2 grams NaCl.[22]Amongst the double salts formed by sodium chloride that obtained by Ditte (1870) by the evaporation of the solution remaining after heating sodium iodate with hydrochloric acid until chlorine ceases to be liberated, is a remarkable one. Its composition is NaIO3,NaCl,14H2O. Rammelsberg obtained a similar (perhaps the same) salt in well-formed crystals by the direct reaction of both salts.[23]But it gives sodium in the flame of a Bunsen's burner (see Spectrum Analysis), doubtless under the reducing action of the elements carbon and hydrogen. In the presence of an excess of hydrochloric acid in the flame (when the sodium would form sodium chloride), no sodium is formed in the flame and the salt does not communicate its usual coloration.[23 bis]There is no doubt, however, but that chloride of sodium is also decomposed in its aqueous solutions with the separation of sodium, and that it does not simply enter into double decomposition with the water (NaCl + H2O = NaHO + HCl). This is seen from the fact that when a saturated solution of NaCl is rapidly decomposed by an electric current, a large amount of chlorine appears at the anode and a sodium amalgam forms at the mercury cathode, which acts but slowly upon the strong solution of salt. Castner's process for the electrolysis of brine into chlorine and caustic soda is an application of this method which has been already worked in England on an industrial scale.[24]If MX and NY represent the molecules of two salts, and if there beno third substancepresent (such as water in a solution), the formation of XY would also be possible; for instance, cyanogen, iodine, &c. are capable of combining with simple haloids, as well as with the complex groups which in certain salts play the part of haloids. Besides which the salts MX and NY or MY with NX may form double salts. If the number of molecules be unequal, or if the valency of the elements or groups contained in them be different, as in NaCl + H2SO4, where Cl is a univalent haloid and SO4is bivalent, then the matter may be complicated by the formation of other compounds besides MY and NX, and when a solvent participates in the action, and especially if present in large proportion, the phenomena must evidently become still more complex; and this is actually the case in nature. Hence while placing before the reader a certain portion of the existing store of knowledge concerning the phenomena of double saline decompositions, I cannot consider the theory of the subject as complete, and have therefore limited myself to a few data, the completion of which must be sought in more detailed works on the subject of theoretical chemistry, without losing sight of what has been said above.[24 bis]When the mixture of potassium nitrate and sodium acetate was heated by Spring to 100°, it was completely fused into one mass, although potassium nitrate fuses at about 340° and sodium nitrate at about 320°.[25]H. Rose is more especially known for his having carefully studied and perfected several methods for the exact chemical analysis of many mineral substances. His predecessor in this branch of research was Berzelius, and his successor Fresenius.[25 bis]Historically the influence of the mass of water was the first well-observed phenomenon in support of Berthollet's teaching, and it should not now be forgotten. In double decompositions taking place in dilute solutions where the mass of water is large, its influence, notwithstanding the weakness of affinities, must he great, according to the very essence of Berthollet's doctrine.As explaining the action of the mass of water, the experiments of Pattison Muir (1879) are very instructive. These experiments demonstrate that the decomposition of bismuth chloride is the more complete the greater the relative quantity of water, and the less the mass of hydrochloric acid forming one of the products of the reaction.[26]From the above it follows that an excess of acid should influence the reaction like an excess of alcohol. It is in fact shown by experiment that if two molecules of acetic acid be taken to one molecule of alcohol, 84 p.c. of alcohol is etherified. If with a large preponderance of acid or of alcohol certain discrepancies are observed, their cause must be looked for in the incomplete correspondence of the conditions and external influences.[27]As an example two methods may be mentioned, Thomsen's and Ostwald's. Thomsen (1869) applied a thermochemical method to exceedingly dilute solutions without taking the water into further consideration. He took solutions of caustic soda containing 100H2O per NaHO, and sulphuric acid containing ½H2SO4+ 100H2O. In order that these solutions may be mixed in such quantities that atomic proportions of acid and alkali would act, for forty grams of caustic soda (which answers to its equivalent) there should be employed 49 grams of sulphuric acid, and then +15,689 heat units would be evolved. If the normal sodium sulphate so formed be mixed withnequivalents of sulphuric acid, a certain amount of heat is absorbed, namely a quantity equal ton.1650/(n+ 0·8)heat units. An equivalent of caustic soda, in combining with an equivalent of nitric acid, evolves +13,617 units of heat, and the augmentation of the amount of nitric acid entails an absorption of heat for each equivalent equal to -27 units; so also in combining with hydrochloric acids +13,740 heat units are absorbed, and for each equivalent of hydrochloric acid beyond this amount there are absorbed -32 heat units. Thomsen mixed each one of three neutral salts, sodium sulphate, sodium chloride and sodium nitrate, with an acid which is not contained in it; for instance, he mixed a solution of sodium sulphate with a solution of nitric acid and determined the number of heat units then absorbed. An absorption of heat ensued because a normal salt was taken in the first instance, and the mixture of all the above normal salts with acid produces an absorption of heat. The amount of heat absorbed enabled him to obtain an insight into the process taking place in this mixture, for sulphuric acid added to sodium sulphate absorbs a considerable quantity of heat, whilst hydrochloric and nitric acids absorb a very small amount of heat in this case. By mixing an equivalent of sodium sulphate with various numbers of equivalents of nitric acid, Thomsen observed that the amount of heat absorbed increased more and more as the amount of nitric acid was increased; thus when HNO3was taken per ½Na2SO4, 1,752 heat units were absorbed per equivalent of soda contained in the sodium sulphate. When twice as much nitric acid was taken, 2,026 heat units, and when three times as much, 2,050 heat units were absorbed. Had the double decomposition been complete in the case where one equivalent of nitric acid was taken per equivalent of Na2SO4then according to calculation from similar data there should have been absorbed -2,989 units of heat, while in reality only -1,752 units were absorbed. Hence Thomsen concluded that a displacement of only about two-thirds of the sulphuric acid had taken place—that is, the ratiok:k′ for the reaction ½Na2SO4+ HNO3and NaNO3+ ½H2SO4is equal, as for ethereal salts, to 4. By taking this figure and admitting the above supposition, Thomsen found that for all mixtures of soda with nitric acid, and of sodium nitrate with sulphuric acid, the amounts of heat followed Guldberg and Waage's law; that is, the limit of decomposition reached was greater the greater the mass of acid added. The relation of hydrochloric to sulphuric acid gave the same results. Therefore the researches of Thomsen fully confirm the hypotheses of Guldberg and Waage and the doctrine of Berthollet.Thomsen concludes his investigation with the words: (a) ‘When equivalent quantities of NaHO, HNO3(or HCl) and ½H2SO4react on one another in an aqueous solution, then two-thirds of the soda combines with the nitric and one-third with the sulphuric acid; (b) this subdivision repeats itself, whether the soda be taken combined with nitric or with sulphuric acid; (c) and therefore nitric acid has double the tendency to combine with the base that sulphuric acid has, and hence in an aqueous solution it is a stronger acid than the latter.’‘It is therefore necessary,’ Thomsen afterwards remarks, ‘to have an expression indicating the tendency of an acid for the saturation of bases. This idea cannot be expressed by the wordaffinity, because by this term is most often understood that force which it is necessary to overcome in order to decompose a substance into its component parts. This force should therefore be measured by the amount of work or heat employed for the decomposition of the substance. The above-mentioned phenomenon is of an entirely different nature,’ and Thomsen introduces the termavidity, by which he designates the tendency of acids for neutralisation. ‘Therefore the avidity of nitric acid with respect to soda is twice as great as the avidity of sulphuric acid. An exactly similar result is obtained with hydrochloric acid, so that its avidity with respect to soda is also double the avidity of sulphuric acid. Experiments conducted with other acids showed that not one of the acids investigated had so great an avidity as nitric acid; some had a greater avidity than sulphuric acid, others less, and in some instances the avidity = 0.’ The reader will naturally see clearly that the path chosen by Thomsen deserves to be worked out, for his results concern important questions of chemistry, but great faith cannot be placed in the deductions he has already arrived at, because great complexity of relations is to be seen in the very method of his investigation. It is especially important to turn attention to the fact that all the reactions investigated are reactions of double decomposition. In them A and B do not combine with C and distribute themselves according to their affinity or avidity for combination, but reversible reactions are induced. MX and NY give MY and NX, and conversely; therefore the affinity or avidity for combination is not here directly determined, but only the difference or relation of the affinities or avidities. The affinity of nitric acid not only for the water of constitution, but also for that serving for solution, is much less than that of sulphuric acid. This is seen from thermal data. The reaction N2O5+ H2O gives +3,600 heat units, and the solution of the resultant hydrate, 2NHO3, in a large excess of water evolves +14,986 heat units. The formation of SO3+ H2O evolves +21,308 heat units, and the solution of H2SO4in an excess of water 17,860—that is, sulphuric acid gives more heat in both cases. The interchange between Na2SO4and 2HNO3is not only accomplished at the expense of the production of NaNO3, but also at the expense of the formation of H2SO4, hence the affinity of sulphuric acid for water plays its part in the phenomena of displacement. Therefore in determinations like those made by Thomsen the water does not form a medium which is present without participating in the process; it also takes part in the reaction. (Compare Chapter IX., Note14.)Whilst retaining essentially the methods of Thomsen, Ostwald (1876) determined the variation of the sp. gr. (and afterwards of volume), proceeding in the same dilute solutions, on the saturation of acids by bases, and in the decomposition of the salts of one acid by the other, and arrived at conclusions of just the same nature as Thomsen's. Ostwald's method will be clearly understood from an example. A solution of caustic soda containing an almost molecular (40 grams) weight per litre had a specific gravity of 1·04051. The specific gravities of solutions of equal volume and equivalent composition of sulphuric and nitric acids were 1·02970 and 1·03084 respectively. On mixing the solutions of NaHO and H2SO4there was formed a solution of Na2SO4of sp. gr. 1·02959; hence there ensued a decrease of specific gravity which we will term Q, equal to 1·04051 + 1·02970 - 2(1·02959) = 0·01103. So also the specific gravity after mixture of the solutions of NaHO and HNO3was 1·02633, and therefore Q = 0·01869. When one volume of the solution of nitric acid was added to two volumes of the solution of sodium sulphate, a solution of sp. gr. 1·02781 was obtained, and therefore the resultant decrease of sp. gr.Q1= 2(1·02959) + 1·03084 - 3(1·02781) = 0·00659.Had there been no chemical reaction between the salts, then according to Ostwald's reasoning the specific gravity of the solutions would not have changed, and if the nitric acid had entirely displaced the sulphuric acid Q2would be = 0·01869 - 0·01103 = 0·00766. It is evident that a portion of the sulphuric acid was displaced by the nitric acid. But the measure of displacement is not equal to the ratio between Q1and Q2, because a decrease of sp. gr. also occurs on mixing the solution of sodium sulphate with sulphuric acid, whilst the mixing of the solutions of sodium nitrate and nitric acid only produces a slight variation of sp. gr. which falls within the limits of experimental error. Ostwald deduces from similar data the same conclusions as Thomsen, and thus reconfirms the formula deduced by Guldberg and Waage, and the teaching of Berthollet.The participation of water is seen still more clearly in the methods adopted by Ostwald than in those of Thomsen, because in the saturation of solutions of acids by alkalis (which Kremers, Reinhold, and others had previously studied) there is observed, not a contraction, as might have been expected from the quantity of heat which is then evolved, but an expansion, of volume (a decrease of specific gravity, if we calculate as Ostwald did in his first investigations). Thus by mixing 1,880 grams of a solution of sulphuric acid of the composition SO3+ 100H2O, occupying a volume of 1,815 c.c., with a corresponding quantity of a solution 2(NaHO + 5H2O), whose volume = 1,793 c.c., we obtain not 3,608 but 3,633 c.c., an expansion of 25 c.c. per gram molecule of the resulting salt, Na2SO4. It is the same in other cases. Nitric and hydrochloric acids give a still greater expansion than sulphuric acid, and potassium hydroxide than sodium hydroxide, whilst a solution of ammonia gives a contraction. The relation to water must be considered as the cause of these phenomena. When sodium hydroxide and sulphuric acid dissolve in water they develop heat and give a vigorous contraction; the water is separated from such solutions with great difficulty. After mutual saturation they form the salt Na2SO4, which retains the water but feebly and evolves but little heat with it, i.e., in other words, has little affinity for water. In the saturation of sulphuric acid by soda the water is, so to say, displaced from a stable combination and passes into an unstable combination; hence an expansion (decrease of sp. gr.) takes place. It is not the reaction of the acid on the alkali, but the reaction of water, that produces the phenomenon by which Ostwald desires to measure the degree of salt formation. The water, which escaped attention, itself has affinity, and influences those phenomena which are being investigated. Furthermore, in the given instance its influence is very great because its mass is large. When it is not present, or only present in small quantities, the attraction of the base to the acid leads to contraction, and not expansion. Na2O has a sp. gr. 2·8, hence its molecular volume = 22; the sp. gr. of SO3is 1·9 and volume 41, hence the sum of their volumes is 63; for Na2SO4the sp. gr. is 2·65 and volume 53·6, consequently there is a contraction of 10 c.c. per gram-molecule of salt. The volume of H2SO4= 53·3, that of 2NaHO = 37·4; there is produced 2H2O, volume = 36, + Na2SO4, volume = 53·6. There react 90·7 c.c., and on saturation there result 89·6 c.c.; consequently contraction again ensues, although less, and although this reaction is one of substitution and not of combination. Consequently the phenomena studied by Ostwald depend but little on the measure of the reaction of the salts, and more on the relations of the dissolved substances to water. In substitutions, for instance 2NaNO3+ H2SO4= 2HNO3+ Na2SO4, the volumes vary but slightly: in the above example they are 2(38·8) + 53·3 and 2(41·2) + 53·6; hence 131 volumes act, and 136 volumes are produced. It may be concluded, therefore, on the basis of what has been said, that on taking water into consideration the phenomena studied by Thomsen and Ostwald are much more complex than they at first appear, and that this method can scarcely lead to a correct interpretation as to the distribution of acids between bases. We may add that P. D. Chroustcheff (1890) introduced a new method for this class of research, by investigating the electro-conductivity of solutions and their mixtures, and obtained remarkable results (for example, that hydrochloric acid almost entirely displaces formic acid and only ⅔ of sulphuric acid), but details of these methods must be looked for in text-books of theoretical chemistry.[28]G. G. Gustavson's researches, which were conducted in the laboratory of the St. Petersburg University in 1871–72, are among the first in which the measure of the affinity of the elements for the halogens is recognised with perfect clearness in the limit of substitution and in the rate of reaction. The researches conducted by A. L. Potilitzin (of which mention will be made in Chapter XI., Note66) in the same laboratory touch on another aspect of the same problem which has not yet made much progress, notwithstanding its importance and the fact that the theoretical side of the subject (thanks especially to Guldberg and Van't Hoff) has since been rapidly pushed forward. If the researches of Gustavson took account of the influence of mass, and were more fully supplied with data concerning velocities and temperatures, they would be very important, because of the great significance which the case considered has for the understanding of double saline decompositions in the absence of water.Furthermore, Gustavson showed that the greater the atomic weight of the element (B, Si, Ti, As, Sn) combinedwith chlorinethe greater the amount of chlorine replaced by bromine by the action of CBr4, and consequently the less the amount of bromine replaced by chlorine by the action of CCl4on bromine compounds. For instance, for chlorine compounds the percentage of substitution (at the limit) is—BCl3SiCl4TiCl4AsCl3SnCl410·112·543·671·877·5It should he observed, however, that Thorpe, on the basis of his experiments, denies the universality of this conclusion. I may mention one conclusion which it appears to me may be drawn from the above-cited figures of Gustavson, if they are subsequently verified even within narrow limits. If CBr4be heated with RCl4, then an exchange of the bromine for chlorine takes place. But what would be the result if it were mixed with CCl4? Judging by the magnitude of the atomic weights, B = 11, C = 12, Si = 28, about 11 p.c. of the chlorine would be replaced by bromine. But to what does this point? I think that this shows the existence of a motion of the atoms in the molecule. The mixture of CCl4and CBr4does not remain in a condition of static equilibrium; not only are the molecules contained in it in a state of motion, but also the atoms in the molecules, and the above figures show the measure of their translation under these conditions. The bromine in the CBr4is,within the limit, substituted by the chlorine of the CCl4in a quantity of about 11 out of 100: that is, a portion of the atoms of bromine previously to this moment in combination with one atom of carbon pass over to the other atom of carbon, and the chlorine passes over from this second atom of carbon to replace it. Therefore, also, in the homogeneous mass CCl4all the atoms of Cl do not remain constantly combined with the same atoms of carbon, andthere is on exchange of atoms between different molecules in a homogeneous medium also. This hypothesis may in my opinion explain certain phenomena of dissociation, but though mentioning it I do not consider it worth while to dwell upon it. I will only observe that a similar hypothesis suggested itself to me in my researches on solutions, and that Pfaundler enunciated an essentially similar hypothesis, and in recent times a like view is beginning to find favour with respect to the electrolysis of saline solutions.[29]Berthollet's doctrine is hardly at all affected in principle by showing that there are cases in which there is no decomposition between salts, because the affinity may be so small that even a large mass would still give no observable displacements. The fundamental condition for the application of Berthollet's doctrine, as well as Deville's doctrine of dissociation, lies in the reversibility of reactions. There are practically irreversible reactions (for instance, CCl4+ 2H2O = CO2+ 4HCl), just as there are non-volatile substances. But while accepting the doctrine of reversible reactions and retaining the theory of the evaporation of liquids, it is possible to admit the existence of non-volatile substances, and in just the same way of reactions, without any visible conformity to Berthollet's doctrine. This doctrine evidently comes nearer than the opposite doctrine of Bergmann to solving the complex problems of chemical mechanics for the successful solution of which at the present time the most valuable help is to be expected from the working out of data concerning dissociation, the influence of mass, and the equilibrium and velocity of reactions. But it is evident that from this point of view we must not regard a solvent as a non-participant space, but must take into consideration the chemical reactions accompanying solution, or else bring about reactions without solution.[30]Common salt not only enters into double decomposition with acids but alsowith every salt. However, as clearly follows from Berthollet's doctrine, this form of decomposition will only in a few cases render it possible for new metallic chlorides to be obtained, because the decomposition will not be carried on to the end unless the metallic chloride formed separates from the mass of the active substances. Thus, for example, if a solution of common salt be mixed with a solution of magnesium sulphate, double decomposition ensues, but not completely, because all the substances remain in the solution. In this case the decomposition must result in the formation of sodium sulphate and magnesium chloride, substances which are soluble in water; nothing is disengaged, and therefore the decomposition 2NaCl + MgSO4= MgCl2+ Na2SO4cannot proceed to the end. However, the sodium sulphate formed in this manner may be separated by freezing the mixture. The complete separation of the sodium sulphate will naturally not take place, owing to a portion of the salt remaining in the solution. Nevertheless, this kind of decomposition is made use of for the preparation of sodium sulphate from the residues left after the evaporation of sea-water, which contain a mixture of magnesium sulphate and common salt. Such a mixture is found at Stassfurt in a natural form. It might be said that this form of double decomposition is only accomplished with a change of temperature; but this would not be true, as may be concluded from other analogous cases. Thus, for instance, a solution of copper sulphate is of a blue colour, while a solution of copper chloride is green. If we mix the two salts together the green tint is distinctly visible, so that by this means the presence of the copper chloride in the solution of copper sulphate is clearly seen. If now we add a solution of common salt to a solution of copper sulphate, a green coloration is obtained, which indicates the formation of copper chloride. In this instance it is not separated, but it is immediately formed on the addition of common salt, as it should be according to Berthollet's doctrine.The complete formation of a metallic chloride from common salt can only occur, judging from the above, when it separates from the sphere of action. The salts of silver are instances in point, because the silver chloride is insoluble in water; and therefore if we add a solution of sodium chloride to a solution of a silver salt, silver chloride and the sodium salt of that acid which was in the silver salt are formed.[31]The apparatus shown in fig.46(Chapter VI., Note12) is generally employed for the preparation of small quantities of hydrochloric acid. Common salt is placed in the retort; the salt is generally previously fused, as it otherwise froths and boils over in the apparatus. When the apparatus is placed in order sulphuric acid mixed with water is poured down the thistle funnel into the retort. Strong sulphuric acid (about half as much again as the weight of the salt) is usually taken, and it is diluted with a small quantity of water (half) if it be desired to retard the action, as in using strong sulphuric acid the action immediately begins with great vigour. The mixture, at first without the aid of heat and then at a moderate temperature (in a water-bath), evolves hydrochloric acid. Commercial hydrochloric acid contains many impurities; it is usually purified by distillation, the middle portions being collected. It is purified from arsenic by adding FeCl2, distilling, and rejecting the first third of the distillate. If free hydrochloric acid gas be required, it is passed through a vessel containing strong sulphuric acid to dry it, and is collected over a mercury hath.Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler, 1888; 2P2O3+ 3HCl = POCl3+ 3HPO3) at the ordinary temperature, and therefore the gas cannot he dried by this substance.[31 bis]In chemical works where sulphuric acid of 60° Baumé (22 p.c. of water) is employed, 117 parts of sodium chloride are taken to about 125 parts of sulphuric acid.[32]As in works which treat common salt in order to obtain sodium sulphate, the hydrochloric acid is sometimes held to be of no value, it might be allowed to escape with the waste furnace gases into the atmosphere, which would greatly injure the air of the neighbourhood and destroy all vegetation. In all countries, therefore, there are laws forbidding the factories to proceed in this manner, and requiring the absorption of the hydrochloric acid by water at the works themselves, and not permitting the solution to be run into rivers and streams, whose waters it would spoil. It may be remarked that the absorption of hydrochloric acid presents no particular difficulties (the absorption of sulphurous acid is much more difficult) because hydrochloric acid has a great affinity for water and gives a hydrate which boils above 100°. Hence, even steam and hot water, as well as weaker solutions, can be used for absorbing the acid. However, Warder (1888) showed that weak solutions of composition H2O +nHCl when boiled (the residue will be almost HCl,8H2O) evolve (not water but) a solution of the composition H2O + 445n4HCl; for example, on distilling HCl,10H2O, HCl,23H2O is first obtained in the distillate. As the strength of the residue becomes greater, so also does that of the distillate, and therefore in order to completely absorb hydrochloric acid it is necessary in the end to have recourse to water.As in Russia the manufacture of sodium sulphate from sodium chloride has not yet been sufficiently developed, and as hydrochloric acid is required for many technical purposes (for instance, for the preparation of zinc chloride, which is employed for soaking railway sleepers), therefore salt is often treated mainly for the manufacture of hydrochloric acid.[33]Thus the metallic chlorides, which are decomposed to a greater or less degree by water, correspond with feeble bases. Such are, for example, MgCl2, AlCl3, SbCl3, BiCl3. The decomposition of magnesium chloride (and also carnallite) by sulphuric acid proceeds at the ordinary temperature; water decomposes MgCl2to the extent of 50 p.c. when aided by heat, andmay be employedas a convenientmethod for the production of hydrochloric acid. Hydrochloric acid is also produced by the ignition of certain metallic chlorides in a stream of hydrogen, especially of those metals which are easily reduced and difficultly oxidised—for instance, silver chloride. Lead chloride, when heated to redness in a current of steam, gives hydrochloric acid and lead oxide. The multitude of the cases of formation of hydrochloric acid are understood from the fact that it is a substance which is comparatively very stable, resembling water in this respect, and even most probably more stable than water, because, at a high temperature and even under the action of light, chlorine decomposes water, with the formation of hydrochloric acid. The combination of chlorine and hydrogen also proceeds by their direct action, as we shall afterwards describe.[34]According to Ansdell (1880) the sp. gr. of liquid hydrochloric acid at 0° = 0·908, at 11·67° = 0·854, at 22·7° = 0·808, at 33° = 0·748. Hence it is seen that the expansion of this liquid is greater than that of gases (Chapter II., Note34).[35]According to Roscoe and Dittmar at a pressure of three atmospheres the solution of constant boiling point contains 18 p.c. of hydrogen chloride, and at a pressure of one-tenth atmosphere 23 p.c. The percentage is intermediate at medium pressures.[36]At 0° 25 p.c., at 100° 20·7 p.c.; Roscoe and Dittmar.[37]This crystallo-hydrate (obtained by Pierre and Puchot, and investigated by Roozeboom) is analogous to NaCl,2H2O. The crystals HCl,2H2O at -22° have a specific gravity 1·46; the vapour tension (under dissociation) of the solution having a composition HCl,2H2O at -24° = 760, at -19° = 1,010, at -18° = 1,057, at -17° = 1,112 mm. of mercury. In a solid state the crystallo-hydrate at -17·7° has the same tension, whilst at lower temperatures it is much less: at -24° about 150, at -19° about 580 mm. A mixture of fuming hydrochloric acid with snow reduces the temperature to -38°. If another equivalent of water be added to the hydrate HCl,2H2O at -18°, the temperature of solidification falls to -25°, and the hydrate HCl,3H2O is formed (Pickering, 1893).[38]According to Roscoe at 0° onehundredgrams of water at a pressurep(in millimetres of mercury) dissolves—p=1002003005007001000Grams HCl65·770·773·878·281·785·6At a pressure of 760 millimetres and temperaturet, onehundredgrams of water dissolvest=08°16°24°40°60°Grams HCl82·578·374·270·063·356·1Roozeboom (1886) showed that att° solutions containingcgrams of hydrogen chloride per 100 grams of water may (with the variation of the pressurep) be formed together with the crystallo-hydrate HCl,2H2O:t=-28°·8-21°-19°-18°c=84·286·892·698·4101·4p=—3345809001,073 mm.The last combination answers to the melted crystallo-hydrate HCl,2H2O, which splits up at temperatures above -17°·7, and at a constant atmospheric pressure when there are no crystals—t=-24°-21°-18°-10°-0°c=101·298·395·789·884·2From these data it is seen that the hydrate HCl,2H2O can exist in a liquid state, which is not the case for the hydrates of carbonic and sulphurous anhydrides, chlorine, &c.According to Marignac, the specific heatcof a solution HCl +mH2O (at about 30°, taking the specific heat of water = 1) is given by the expression—C(36·5 +m18) = 18m- 28·39 + 140/m- 268/m2ifmbe not less than 6·25. For example, for HCl + 25H2O, C = 0·877.According to Thomsen's data, the amount of heatQ, expressed in thousands of calories, evolved in the solution of 36·5 grams of gaseous hydrochloric acid inmH2O or 18mgrams of water is equal to—m=241050400Q=11·414·316·217·117·3In these quantities the latent heat of liquefaction is included, which must be taken as 5–9 thousand calories per molecular quantity of hydrogen chloride.The researches of Scheffer (1888) on the rate of diffusion (in water) of solutions of hydrochloric acid show that the coefficient of diffusionkdecreases with the amount of watern, if the composition of the solution is HCl,nH2O at 0°:—n=56·99·81427·1129·5k=2·312·081·861·671·521·39It also appears that strong solutions diffuse more rapidly into dilute solutions than into water.[39]If it be admitted that the maximum of the differential corresponds with HCl,6H2O, then it might be thought that the specific gravity is expressed by a parabola of the third order; but such an admission does not give expressions in accordance with fact. This is all more fully considered in my work mentioned in Chapter I., Note19.[40]As in water, the coefficient of expansion (or the quantitykin the expression St= S0-kS0t, or Vt= 1/(1 -kt)) attains a magnitude 0·000447 at about 48°, it might be thought that at 48° all solutions of hydrochloric acid would have the same coefficient of expansion, but in reality this is not the case. At low and at the ordinary temperatures the coefficient of expansion of aqueous solutions is greater than that of water, and increases with the amount of substance dissolved.[41]The figures cited above may serve for the direct determination of that variation of the specific gravity of solutions of hydrochloric acid with the temperature. Thus, knowing that at 15° the specific gravity of a 10 p.c. solution of hydrochloric acid = 10,492, we find that att° it = 10,530 -t(2·13 + 0·027t). Whence also may be found the coefficient of expansion (Note40).[42]Thus, for instance, with feeble bases they evolve in dilute solutions (Chapter III., Note 53) almost equal amounts of heat; their relation to sulphuric acid is quite identical. They both form fuming solutions as well as hydrates; they both form solutions of constant boiling point.[42 bis]Pybalkin (1891) found that copper begins to disengage hydrogen at 100°, and that chloride of copper begins to give up its chlorine to hydrogen gas at 230°; for silver these temperatures are 117° and 260°—that is, there is less difference between them.[43]When an unsaturated hydrocarbon, or, in general, an unsaturated compound, assimilates to itself the molecules Cl2, HCl, SO3, H2SO4, &c., the cause of the reaction is most simple. As nitrogen, besides the type NX3to which NH3, belongs, gives compounds of the type NX5—for example, NO2(OH)—the formation of the salts of ammonium should be understood in this way. NH3gives NH4Cl because NX3is capable of giving NX5. But as saturated compounds—for instance, SO3,H2O, NaCl, &c.—are also capable of combination even between themselves, it is impossible to deny the capacity of HCl also for combination. SO3combines with H2O, and also with HCl and the unsaturated hydrocarbons. It is impossible to recognise the distinction formerly sought to be established between atomic and molecular compounds, and regarding, for instance, PCl3as an atomic compound and PCl5as a molecular one, only because it easily splits up into molecules PCl3and Cl2.[44]Sal-ammoniac is prepared from ammonium carbonate, obtained in the dry distillation of nitrogenous substances (ChapterVI.), by saturating the resultant solution with hydrochloric acid. A solution of sal-ammoniac is thus produced, which is evaporated, and in the residue a mass is obtained containing a mixture of various other, especially tarry, products of dry distillation. The sal-ammoniac is generally purified by sublimation. For this purpose iron vessels covered with hemispherical metallic covers are employed, or else simply clay crucibles covered by other crucibles. The upper portion, or head, of the apparatus of this kind will have a lower temperature than the lower portion, which is under the direct action of the flame. The sal-ammoniac volatilises when heated, and settles on the cooler portion of the apparatus. It is thus freed from many impurities, and is obtained as a crystalline crust, generally several centimetres thick, in which form it is commonly sold. The solubility of sal-ammoniac rises rapidly with the temperature: at 0°, 100 parts of water dissolve about 28 parts of NH4Cl, at 50° about 50 parts, and at the ordinary temperature about 35 parts. This is sometimes taken advantage of for separating NH4Cl from solutions of other salts.[45]The solubility of sal-ammoniac in 100 parts of water (according to Alluard) is—0°10°20°30°40°60°80°100°100°28·4032·4837·2841·724655647377A saturated solution boils at 115°·8. The specific gravity at 15°/4° of solutions of sal-ammoniac (water 4° = 10,000) = 9,991·6 - 31·26p- 0·085p2, wherepis the amount by weight of ammonium chloride in 100 parts of solution. With the majority of salts the differentialds/dpincreases, but here it decreases with the increase ofp. For (unlike the sodium and potassium salts) a solution of the alkaliplusa solution of acid occupy a greater volume than that of the resultant ammonium salt. In the solution ofsolidammonium chloride a contraction, and not expansion, generally takes place. It may further be remarked that solutions of sal-ammoniac have an acid reaction even when prepared from the salt remaining after prolonged washing of the sublimed salt with water (A. Stcherbakoff).

Footnotes:

[1]But it is impossible to foretell all the compounds formed by an element from its atomicity or valency, because the atomicity of the elements is variable, and furthermore this variability is not identical for different elements. In CO2, COX2, CH4, and the multitude of carbon compounds corresponding with them, the C is quadrivalent, but in CO either the carbon must be taken as bivalent or the atomicity of oxygen be accounted as variable. Moreover, carbon is an example of an element which preserves its atomicity to a greater degree than most of the other elements. Nitrogen in NH3, NH2(OH), N2O3, and even in CNH, must be considered as trivalent, but in NH4Cl, NO2(OH), and in all their corresponding compounds it is necessarily pentavalent. In N2O, if the atomicity of oxygen = 2, nitrogen has an uneven atomicity (1, 3, 5), whilst in NO it is bivalent. If sulphur be bivalent, like oxygen, in many of its compounds (for example, H2S, SCl2, KHS, &c.), then it could not be foreseen from this that it would form SO2, SO3, SCl4, SOCl2, and a series of similar compounds in which its atomicity must be acknowledged as greater than 2. Thus SO2, sulphurous anhydride, has many points in common with CO2, and if carbon be quadrivalent then the S in SO2is quadrivalent. Therefore the principle of atomicity (valency) of the elements cannot be considered established as the basis for the study of the elements, although it gives an easy method of grasping many analogies. I consider the four following as the chief obstacles to acknowledging the atomicity of the elements as a primary conception for the consideration of the properties of the elements: 1. Such univalent elements as H, Cl, &c., appear in a free state as molecules H2, Cl2, &c., and are consequently like the univalent radicles CH3, OH, CO2H, &c., which, as might be expected, appear as C2H6, O2H2, C2O4H2(ethane, hydrogen peroxide, oxalic acid), whilst on the other hand, potassium and sodium (perhaps also iodine at a high temperature) contain only one atom, K, Na, in the molecule in a free state. Hence it follows thatfree affinitiesmay exist. Granting this, nothing prevents the assumption that free affinities exist in all unsaturated compounds; for example, two free affinities in NH3. If such instances of free affinities be admitted, then all the possible advantages to be gained by the application of the doctrine of atomicity (valency) are lost. 2. There are instances—for example, Na2H—where univalent elements are combined in molecules which are more complex than R2, and form molecules, R3, R4, &c.; this may again be either taken as evidence of the existence of free affinities, or else necessitates such primary univalent elements as sodium and hydrogen being considered as variable in their atomicity. 3. The periodic system of the elements, with which we shall afterwards become acquainted, shows that there is a law or rule for the variation of the forms of oxygen and hydrogen compounds; chlorine is univalent with respect to hydrogen, and septavalent with respect to oxygen; sulphur is bivalent to hydrogen, and sexavalent to oxygen; phosphorus is trivalent to hydrogen and pentavalent in respect to oxygen—the sum is in every case equal to 8. Only carbon and its analogues (for example, silicon) are quadrivalent to both hydrogen and oxygen. Hence the power of the elements to change their atomicity is an essential part of their nature, and therefore constant valency cannot he considered as a fundamental property. 4. Crystallo-hydrates (for instance, NaCl,2H2O, or NaBr,2H2O), double salts (such as PtCl4,2KCl,H2SiF6, &c.), and similar complex compounds (and, according to Chap.I., solutions also) demonstrate the capacity not only of the elements themselves, but also of their saturated and limiting compounds, of entering into further combination. Therefore the admission of a definite limited atomicity of the elements includes in itself an admission of limitation which is not in accordance with the nature of chemical reactions.

[2]The primary formations are those which do not bear any distinct traces of having been deposited from water (have not a stratified formation and contain no remains of animal or vegetable life), occur under the sedimentary formations of the earth, and are everywhere uniform in composition and structure, the latter being generally distinctly crystalline. If it be assumed that the earth was originally in a molten condition, the first primary formations are those which formed the first solid crust of the earth. But even with this hypothesis of the earth's origin, it is necessary to admit that the first aqueous deposits must have caused a change in the original crust of the earth, and therefore under the head of primary formations must be understood the most ancient of the products of decomposition (mostly by atmospheric, aqueous, and organic agency, &c.), from which all the rocks and substances of the earth's surface have arisen. In speaking of the origin of one or another substance, we can only, on the basis of facts, descend to the primary formations, of which granite, gneiss, and trachyte may be taken as examples.

[2 bis]Chloride of sodium has been found to occur in the atmosphere in the form of a fine dust; in the lower strata it is present in larger quantities than in the upper, so that the rain water falling on mountains contains less NaCl than that falling in valleys. Müntz (1891) found that a litre of rain water collected on the summit of the Pic du Midi (2,877 metres above the sea level) contained 0·34 milligram of chloride of sodium, while a litre of rain collected from the valley contained 2·5–7·6 milligrams.

[3]The extraction of the potassium salts (or so-called summer salts) was carried on at the Isle of Camarga about 1870, when I had occasion to visit that spot. At the present time the deposits of Stassfurt provide a much cheaper salt, owing to the evaporation and separation of the salt being carried on there by natural means and only requiring a treatment and refining, which is also necessary in addition for the ‘summer salt’ obtained from sea-water.

[4]The double salt KCl,MgCl2is a crystallohydrate of KCl and MgCl2, and is only formed from solutions containing an excess of magnesium chloride, because water decomposes this double salt, extracting the more soluble magnesium chloride from it.

[5]Owing to the fundamental property of salts of interchanging their metals, it cannot be said that sea water contains this or that salt, but only that it contains certain amounts of certain metals M (univalent like Na and K, and bivalent like Mg and Ca), and haloids X (univalent like Cl, Br, and bivalent like SO4, CO3), which are disposed in every possible kind of grouping; for instance, K as KCl, KBr, K2SO4, Mg as MgCl2, MgBr2, MgSO4, and so on for all the other metals. In evaporation different salts separate out consecutively only because they reach saturation. A proof of this may be seen in the fact that a solution of a mixture of sodium chloride and magnesium sulphate (both of which salts are obtained from sea water, as was mentioned above), when evaporated, deposits crystals of these salts, but when refrigerated (if the solution be sufficiently saturated) the salt Na2SO4,10H2O is first deposited because it is the first to arrive at saturation at low temperatures. Consequently this solution contains MgCl2and Na2SO4, besides MgSO4and NaCl. So it is with sea water.

[6]The salt extracted from water is piled up in heaps and left exposed to the action of rain water, which purifies it, owing to the water becoming saturated with sodium chloride and then no longer dissolving it, but washing out the impurities.

[7]When the German savants pointed out the exact locality of the Stassfurt salt-beds and their depth below the surface, on the basis of information collected from various quarters respecting bore-holes and the direction of the strata, and when the borings, conducted by the Government, struck a salt-bed which was bitter and unfit for use, there was a great outcry against science, and the doubtful result even caused the cessation of the further work of deepening the shafts. It required a great effort to persuade the Government to continue the work. Now, when the pure salt encountered below forms one of the important riches of Germany, and when those ‘refuse salts’ have proved to be most valuable (as a source of potassium and magnesium), we should see in the utilisation of the Stassfurt deposits one of the conquests of science for the common welfare.

[7 bis]In Western Europe, deposits of rock salt have long been known at Wieliczka, near Cracow, and at Cardona in Spain. In Russia the following deposits are known: (a) the vast masses of rock salt (3 square kilometres area and up to 140 metres thick) lying directly on the surface of the earth at Iletzky Zastchit, on the left bank of the river Ural, in the Government of Orenburg; (b) the Chingaksky deposit, 90 versts from the river Volga, in the Enotaeffsky district of the Government of Astrakhan; (c) the Kulepinsky (and other) deposits (whose thickness attains 150 metres), on the Araks, in the Government of Erivan in the Caucasus; (d) the Katchiezmansky deposit in the province of Kars; (e) the Krasnovodsky deposit in the Trans-Caspian province; and (f) the Bardymkulsky salt mines in Kokhand.

[8]The fracture of rock salt generally shows the presence of interlayers of impurities which are sometimes very small in weight, but visible owing to their refraction. In the excellently laid out salt mines of Briansk I counted (1888), if my memory does not deceive me, on an average ten interlayers per metre of thickness, between which the salt was in general very pure, and in places quite transparent. If this be the case, then there would be 350 interlayers for the whole thickness (about 35 metres) of the bed. They probably correspond with the yearly deposition of the salt. In this case the deposition would have extended over more than 300 years. This should be observable at the present day in lakes where the salt is saturated and in course of deposition.

[9]My own investigations have shown that not only the sulphates, but also the potassium salts, are entirely removed by this method.

[10]According to the determinations of Klodt, the Briansk rock salt withstands a pressure of 340 kilograms per square centimetre, whilst glass withstands 1,700 kilos. In this respect salt is twice as secure as bricks, and therefore immense masses may be extracted from underground workings with perfect safety, without having recourse to brickwork supports, merely taking advantage of the properties of the salt itself.

[11]To obtain well-formed crystals, a saturated solution is mixed with ferric chloride, several small crystals of sodium chloride are placed at the bottom, and the solution is allowed to evaporate slowly in a vessel with a loose-fitting cover. Octahedral crystals are obtained by the addition of borax, urea, &c., to the solution. Very fine crystals are formed in a mass of gelatinous silica.

[12]If a solution of sodium chloride be slowly heated from above, where the evaporation takes place, then the upper layer will become saturated before the lower and cooler layers, and therefore crystallisation will begin on the surface, and the crystals first formed will float, having also dried from above, on the surface until they become quite soaked. Being heavier than the solution the crystals are partially immersed under it, and the following crystallisation, also proceeding on the surface, will only form crystals along the side of the original crystals. A funnel is formed in this manner. It will be borne on the surface like a boat (if the liquid be quiescent), because it will grow more from the upper edges. We can thus understand this at first sight strange funnel form of crystallisation of salt. In explanation why the crystallisation under the above conditions begins at the surface and not at the lower layers, it must be mentioned that the specific gravity of a crystal of sodium chloride = 2·16, and that of a solution saturated at 25° contains 26·7 p.c. of salt and has a specific gravity at 25°/4° of 1·2004; at 15° a saturated solution contains 26·5 p.c. of salt and has a sp. gr. 1·203 at 15°/4°. Hence a solution saturated at a higher temperature is specifically lighter, notwithstanding the greater amount of salt it contains. With many substancessurface crystallisationcannot take place because their solubility increases more rapidly with the temperature than their specific gravity decreases. In this case the saturated solution will always be in the lower layers, where also the crystallisation will take place. Besides which it may be added that as a consequence of the properties of water and solutions, when they are heated from above (for instance, by the sun's rays), the warmer layers being the lightest remain above, whilst when heated from below they rise to the top. For this reason the water at great depths below the surface is always cold, which has long been known. These circumstances, as well as those observed by Soret (Chapter I., Note19), explain the great differences of density and temperature, and in the amount of salts held in the oceans at different latitudes (in polar and tropical climes) and at various depths.

[13]By combining the results of Poggiale, Müller, and Karsten (they are evidently more accurate than those of Gay-Lussac and others) I found that a saturated solution att°, from 0° to 108°, contains 35·7 + 0·024t+ 0·0002t2grams of salt per 100 grams of water. This formula gives a solubility at 0° = 35·7 grams (= 26·3 p.c.), whilst according to Karsten it is 36·09, Poggiale 35·5, and Müller 35·6 grams.

[14]Perfectly purefusedsalt is not hygroscopic, according to Karsten, whilst the crystallised salt, even when quite pure, attracts as much as 0·6 p.c. of water from moist air, according to Stas. (In the Briansk mines, where the temperature throughout the whole year is about +10°, it may be observed, as Baron Klodt informed me, that in the summer during damp weather the walls become moist, while in winter they are dry).If the salt contain impurities—such as magnesium sulphate, &c.—it is more hygroscopic. If it contain any magnesium chloride, it partially deliquesces in a damp atmosphere. The crystallised and not perfectly pure salt decrepitates when heated, owing to its containing water. The pure salt, and also the transparent rock salt, or that which has been once fused, does not decrepitate. Fused sodium chloride shows a faint alkaline reaction to litmus, which has been noticed by many observers, and is due to the presence of sodium oxide (probably by the action of the oxygen of the atmosphere). According to A. Stcherbakoff very sensitive litmus (washed in alcohol and neutralised with oxalic acid) shows an alkaline reaction even with the crystallised salt.It may be observed that rock salt sometimes contains cavities filled with a colourless liquid. Certain kinds of rock salt emit an odour like that of hydrocarbons. These phenomena have as yet received very little attention.

If the salt contain impurities—such as magnesium sulphate, &c.—it is more hygroscopic. If it contain any magnesium chloride, it partially deliquesces in a damp atmosphere. The crystallised and not perfectly pure salt decrepitates when heated, owing to its containing water. The pure salt, and also the transparent rock salt, or that which has been once fused, does not decrepitate. Fused sodium chloride shows a faint alkaline reaction to litmus, which has been noticed by many observers, and is due to the presence of sodium oxide (probably by the action of the oxygen of the atmosphere). According to A. Stcherbakoff very sensitive litmus (washed in alcohol and neutralised with oxalic acid) shows an alkaline reaction even with the crystallised salt.

It may be observed that rock salt sometimes contains cavities filled with a colourless liquid. Certain kinds of rock salt emit an odour like that of hydrocarbons. These phenomena have as yet received very little attention.

[15]By cooling a solution of table salt saturated at the ordinary temperature to -15°, I obtained first of all well-formed tabular (six-sided) crystals, which when warmed to the ordinary temperature disintegrated (with the separation of anhydrous sodium chloride), and then prismatic needles up to 20 mm. long were formed from the same solution. I have not yet investigated the reason of the difference in crystalline form. It is known (Mitscherlich) that NaI,2H2O also crystallises either in plates or prisms. Sodium bromide also crystallises with 2H2O at the ordinary temperature.

[16]Notwithstanding the great simplicity (Chapter I., Note49) of the observations on the formation of ice from solution, still even for sodium chloride they cannot yet be considered as sufficiently harmonious. According to Blagden and Raoult, the temperature of the formation of ice from a solution containingcgrams of salt per 100 grams of water = -0·6ctoc= 10, according to Rosetti = -0·649ctoc= 8·7, according to De Coppet (toc= 10) = -0·55c- 0·006c2, according to Karsten (toc= 10) - 0·762c+ 0·0084c2, and according to Guthrie a much lower figure. By taking Rosetti's figure and applying the rule given in Chapter I., Note49we obtain—i= 0·649 ×58·5/18·5= 2·05.Pickering (1893) gives forc= 1 - 0·603, forc= 2 - 1·220; that is (cup to 2·7) about - (0·600 + 0·005c)c.The data for strong solutions are not less contradictory. Thus with 20 p.c. of salt, ice is formed at -14·4° according to Karsten, -17° according to Guthrie, -17·6° according to De Coppet. Rüdorff states that for strong solutions the temperature of the formation of ice descends in proportion to the contents of the compound, NaCl,2H2O (per 100 grams of water) by 0°·342 per 1 gram of salt, and De Coppet shows that there is no proportionality, in a strict sense, for either a percentage of NaCl or of NaCl,2H2O.

i= 0·649 ×58·5/18·5= 2·05.

Pickering (1893) gives forc= 1 - 0·603, forc= 2 - 1·220; that is (cup to 2·7) about - (0·600 + 0·005c)c.

The data for strong solutions are not less contradictory. Thus with 20 p.c. of salt, ice is formed at -14·4° according to Karsten, -17° according to Guthrie, -17·6° according to De Coppet. Rüdorff states that for strong solutions the temperature of the formation of ice descends in proportion to the contents of the compound, NaCl,2H2O (per 100 grams of water) by 0°·342 per 1 gram of salt, and De Coppet shows that there is no proportionality, in a strict sense, for either a percentage of NaCl or of NaCl,2H2O.

[17]A collection of observations on the specific gravity of solutions of sodium chloride is given in my work cited in Chapter I., Note50.Solutions of common salt have also been frequently investigated as regards rate ofdiffusion(ChapterI.), but as yet there are no complete data in this respect. It may be mentioned that Graham and De Vries demonstrated that diffusion in gelatinous masses (for instance, gelatin jelly, or gelatinous silica) proceeds in the same manner as in water, which may probably lead to a convenient and accurate method for the investigation of the phenomena of diffusion. N. Umoff (Odessa, 1888) investigated the diffusion of common salt by means of glass globules of definite density. Having poured water into a cylinder over a layer of a solution of sodium chloride, he observed during a period of several months the position (height) of the globules, which floated up higher and higher as the salt permeated upwards. Umoff found that at a constant temperature the distances of the globules (that is, the length of a column limited by layers of definite concentration) remain constant; that at a given moment of time the concentration,q, of different layers situated at a depthzis expressed by the equation B - Kz= log.(A -q), where A, B, and K are constants; that at a given moment the rate of diffusion of the different layers is proportional to their depth, &c.

[17]A collection of observations on the specific gravity of solutions of sodium chloride is given in my work cited in Chapter I., Note50.

Solutions of common salt have also been frequently investigated as regards rate ofdiffusion(ChapterI.), but as yet there are no complete data in this respect. It may be mentioned that Graham and De Vries demonstrated that diffusion in gelatinous masses (for instance, gelatin jelly, or gelatinous silica) proceeds in the same manner as in water, which may probably lead to a convenient and accurate method for the investigation of the phenomena of diffusion. N. Umoff (Odessa, 1888) investigated the diffusion of common salt by means of glass globules of definite density. Having poured water into a cylinder over a layer of a solution of sodium chloride, he observed during a period of several months the position (height) of the globules, which floated up higher and higher as the salt permeated upwards. Umoff found that at a constant temperature the distances of the globules (that is, the length of a column limited by layers of definite concentration) remain constant; that at a given moment of time the concentration,q, of different layers situated at a depthzis expressed by the equation B - Kz= log.(A -q), where A, B, and K are constants; that at a given moment the rate of diffusion of the different layers is proportional to their depth, &c.

[18]IfS_0 be the specific gravity of water, andSthe specific gravity of a solution containingpp.c. of salt, then by mixing equal weights of water and the solution, we shall obtain a solution containing ½pof the salt, and if it be formed without contraction, then its specific gravityxwill be determined by the equation2/x=1/S0+1/S, because the volume is equal to the weight divided by the density. In reality, the specific gravity is always found to be greater than that calculated on the supposition of an absence of contraction.

[19]Generally the specific gravity is observed by weighing in air and dividing the weight in grams by the volume in cubic centimetres, the latter being found from the weight of water displaced, divided by its density at the temperature at which the experiment is carried out. If we call this specific gravity S1, then as a cubic centimetre of air under the usual conditions weighs about 0·0012 gram, the sp. gr. in a vacuum S = S1+ 0·0012 (S1- 1), if the density of water = 1.

[20]If the sp. gr. S2be found directly by dividing the weight of a solution by the weight of water at the same temperature and in the same volume, then the true sp. gr.Sreferred to water at 4° is found by multiplying S2by the sp. gr. of water at the temperature of observation.

[21]According to Schiff 100 grams of alcohol, containingpp.c. by weight of C2H6O, dissolves at 15°—p=102040608028·522·613·25·91·2 grams NaCl.

[21]According to Schiff 100 grams of alcohol, containingpp.c. by weight of C2H6O, dissolves at 15°—

[22]Amongst the double salts formed by sodium chloride that obtained by Ditte (1870) by the evaporation of the solution remaining after heating sodium iodate with hydrochloric acid until chlorine ceases to be liberated, is a remarkable one. Its composition is NaIO3,NaCl,14H2O. Rammelsberg obtained a similar (perhaps the same) salt in well-formed crystals by the direct reaction of both salts.

[23]But it gives sodium in the flame of a Bunsen's burner (see Spectrum Analysis), doubtless under the reducing action of the elements carbon and hydrogen. In the presence of an excess of hydrochloric acid in the flame (when the sodium would form sodium chloride), no sodium is formed in the flame and the salt does not communicate its usual coloration.

[23 bis]There is no doubt, however, but that chloride of sodium is also decomposed in its aqueous solutions with the separation of sodium, and that it does not simply enter into double decomposition with the water (NaCl + H2O = NaHO + HCl). This is seen from the fact that when a saturated solution of NaCl is rapidly decomposed by an electric current, a large amount of chlorine appears at the anode and a sodium amalgam forms at the mercury cathode, which acts but slowly upon the strong solution of salt. Castner's process for the electrolysis of brine into chlorine and caustic soda is an application of this method which has been already worked in England on an industrial scale.

[24]If MX and NY represent the molecules of two salts, and if there beno third substancepresent (such as water in a solution), the formation of XY would also be possible; for instance, cyanogen, iodine, &c. are capable of combining with simple haloids, as well as with the complex groups which in certain salts play the part of haloids. Besides which the salts MX and NY or MY with NX may form double salts. If the number of molecules be unequal, or if the valency of the elements or groups contained in them be different, as in NaCl + H2SO4, where Cl is a univalent haloid and SO4is bivalent, then the matter may be complicated by the formation of other compounds besides MY and NX, and when a solvent participates in the action, and especially if present in large proportion, the phenomena must evidently become still more complex; and this is actually the case in nature. Hence while placing before the reader a certain portion of the existing store of knowledge concerning the phenomena of double saline decompositions, I cannot consider the theory of the subject as complete, and have therefore limited myself to a few data, the completion of which must be sought in more detailed works on the subject of theoretical chemistry, without losing sight of what has been said above.

[24 bis]When the mixture of potassium nitrate and sodium acetate was heated by Spring to 100°, it was completely fused into one mass, although potassium nitrate fuses at about 340° and sodium nitrate at about 320°.

[25]H. Rose is more especially known for his having carefully studied and perfected several methods for the exact chemical analysis of many mineral substances. His predecessor in this branch of research was Berzelius, and his successor Fresenius.

[25 bis]Historically the influence of the mass of water was the first well-observed phenomenon in support of Berthollet's teaching, and it should not now be forgotten. In double decompositions taking place in dilute solutions where the mass of water is large, its influence, notwithstanding the weakness of affinities, must he great, according to the very essence of Berthollet's doctrine.As explaining the action of the mass of water, the experiments of Pattison Muir (1879) are very instructive. These experiments demonstrate that the decomposition of bismuth chloride is the more complete the greater the relative quantity of water, and the less the mass of hydrochloric acid forming one of the products of the reaction.

As explaining the action of the mass of water, the experiments of Pattison Muir (1879) are very instructive. These experiments demonstrate that the decomposition of bismuth chloride is the more complete the greater the relative quantity of water, and the less the mass of hydrochloric acid forming one of the products of the reaction.

[26]From the above it follows that an excess of acid should influence the reaction like an excess of alcohol. It is in fact shown by experiment that if two molecules of acetic acid be taken to one molecule of alcohol, 84 p.c. of alcohol is etherified. If with a large preponderance of acid or of alcohol certain discrepancies are observed, their cause must be looked for in the incomplete correspondence of the conditions and external influences.

[27]As an example two methods may be mentioned, Thomsen's and Ostwald's. Thomsen (1869) applied a thermochemical method to exceedingly dilute solutions without taking the water into further consideration. He took solutions of caustic soda containing 100H2O per NaHO, and sulphuric acid containing ½H2SO4+ 100H2O. In order that these solutions may be mixed in such quantities that atomic proportions of acid and alkali would act, for forty grams of caustic soda (which answers to its equivalent) there should be employed 49 grams of sulphuric acid, and then +15,689 heat units would be evolved. If the normal sodium sulphate so formed be mixed withnequivalents of sulphuric acid, a certain amount of heat is absorbed, namely a quantity equal ton.1650/(n+ 0·8)heat units. An equivalent of caustic soda, in combining with an equivalent of nitric acid, evolves +13,617 units of heat, and the augmentation of the amount of nitric acid entails an absorption of heat for each equivalent equal to -27 units; so also in combining with hydrochloric acids +13,740 heat units are absorbed, and for each equivalent of hydrochloric acid beyond this amount there are absorbed -32 heat units. Thomsen mixed each one of three neutral salts, sodium sulphate, sodium chloride and sodium nitrate, with an acid which is not contained in it; for instance, he mixed a solution of sodium sulphate with a solution of nitric acid and determined the number of heat units then absorbed. An absorption of heat ensued because a normal salt was taken in the first instance, and the mixture of all the above normal salts with acid produces an absorption of heat. The amount of heat absorbed enabled him to obtain an insight into the process taking place in this mixture, for sulphuric acid added to sodium sulphate absorbs a considerable quantity of heat, whilst hydrochloric and nitric acids absorb a very small amount of heat in this case. By mixing an equivalent of sodium sulphate with various numbers of equivalents of nitric acid, Thomsen observed that the amount of heat absorbed increased more and more as the amount of nitric acid was increased; thus when HNO3was taken per ½Na2SO4, 1,752 heat units were absorbed per equivalent of soda contained in the sodium sulphate. When twice as much nitric acid was taken, 2,026 heat units, and when three times as much, 2,050 heat units were absorbed. Had the double decomposition been complete in the case where one equivalent of nitric acid was taken per equivalent of Na2SO4then according to calculation from similar data there should have been absorbed -2,989 units of heat, while in reality only -1,752 units were absorbed. Hence Thomsen concluded that a displacement of only about two-thirds of the sulphuric acid had taken place—that is, the ratiok:k′ for the reaction ½Na2SO4+ HNO3and NaNO3+ ½H2SO4is equal, as for ethereal salts, to 4. By taking this figure and admitting the above supposition, Thomsen found that for all mixtures of soda with nitric acid, and of sodium nitrate with sulphuric acid, the amounts of heat followed Guldberg and Waage's law; that is, the limit of decomposition reached was greater the greater the mass of acid added. The relation of hydrochloric to sulphuric acid gave the same results. Therefore the researches of Thomsen fully confirm the hypotheses of Guldberg and Waage and the doctrine of Berthollet.Thomsen concludes his investigation with the words: (a) ‘When equivalent quantities of NaHO, HNO3(or HCl) and ½H2SO4react on one another in an aqueous solution, then two-thirds of the soda combines with the nitric and one-third with the sulphuric acid; (b) this subdivision repeats itself, whether the soda be taken combined with nitric or with sulphuric acid; (c) and therefore nitric acid has double the tendency to combine with the base that sulphuric acid has, and hence in an aqueous solution it is a stronger acid than the latter.’‘It is therefore necessary,’ Thomsen afterwards remarks, ‘to have an expression indicating the tendency of an acid for the saturation of bases. This idea cannot be expressed by the wordaffinity, because by this term is most often understood that force which it is necessary to overcome in order to decompose a substance into its component parts. This force should therefore be measured by the amount of work or heat employed for the decomposition of the substance. The above-mentioned phenomenon is of an entirely different nature,’ and Thomsen introduces the termavidity, by which he designates the tendency of acids for neutralisation. ‘Therefore the avidity of nitric acid with respect to soda is twice as great as the avidity of sulphuric acid. An exactly similar result is obtained with hydrochloric acid, so that its avidity with respect to soda is also double the avidity of sulphuric acid. Experiments conducted with other acids showed that not one of the acids investigated had so great an avidity as nitric acid; some had a greater avidity than sulphuric acid, others less, and in some instances the avidity = 0.’ The reader will naturally see clearly that the path chosen by Thomsen deserves to be worked out, for his results concern important questions of chemistry, but great faith cannot be placed in the deductions he has already arrived at, because great complexity of relations is to be seen in the very method of his investigation. It is especially important to turn attention to the fact that all the reactions investigated are reactions of double decomposition. In them A and B do not combine with C and distribute themselves according to their affinity or avidity for combination, but reversible reactions are induced. MX and NY give MY and NX, and conversely; therefore the affinity or avidity for combination is not here directly determined, but only the difference or relation of the affinities or avidities. The affinity of nitric acid not only for the water of constitution, but also for that serving for solution, is much less than that of sulphuric acid. This is seen from thermal data. The reaction N2O5+ H2O gives +3,600 heat units, and the solution of the resultant hydrate, 2NHO3, in a large excess of water evolves +14,986 heat units. The formation of SO3+ H2O evolves +21,308 heat units, and the solution of H2SO4in an excess of water 17,860—that is, sulphuric acid gives more heat in both cases. The interchange between Na2SO4and 2HNO3is not only accomplished at the expense of the production of NaNO3, but also at the expense of the formation of H2SO4, hence the affinity of sulphuric acid for water plays its part in the phenomena of displacement. Therefore in determinations like those made by Thomsen the water does not form a medium which is present without participating in the process; it also takes part in the reaction. (Compare Chapter IX., Note14.)Whilst retaining essentially the methods of Thomsen, Ostwald (1876) determined the variation of the sp. gr. (and afterwards of volume), proceeding in the same dilute solutions, on the saturation of acids by bases, and in the decomposition of the salts of one acid by the other, and arrived at conclusions of just the same nature as Thomsen's. Ostwald's method will be clearly understood from an example. A solution of caustic soda containing an almost molecular (40 grams) weight per litre had a specific gravity of 1·04051. The specific gravities of solutions of equal volume and equivalent composition of sulphuric and nitric acids were 1·02970 and 1·03084 respectively. On mixing the solutions of NaHO and H2SO4there was formed a solution of Na2SO4of sp. gr. 1·02959; hence there ensued a decrease of specific gravity which we will term Q, equal to 1·04051 + 1·02970 - 2(1·02959) = 0·01103. So also the specific gravity after mixture of the solutions of NaHO and HNO3was 1·02633, and therefore Q = 0·01869. When one volume of the solution of nitric acid was added to two volumes of the solution of sodium sulphate, a solution of sp. gr. 1·02781 was obtained, and therefore the resultant decrease of sp. gr.Q1= 2(1·02959) + 1·03084 - 3(1·02781) = 0·00659.Had there been no chemical reaction between the salts, then according to Ostwald's reasoning the specific gravity of the solutions would not have changed, and if the nitric acid had entirely displaced the sulphuric acid Q2would be = 0·01869 - 0·01103 = 0·00766. It is evident that a portion of the sulphuric acid was displaced by the nitric acid. But the measure of displacement is not equal to the ratio between Q1and Q2, because a decrease of sp. gr. also occurs on mixing the solution of sodium sulphate with sulphuric acid, whilst the mixing of the solutions of sodium nitrate and nitric acid only produces a slight variation of sp. gr. which falls within the limits of experimental error. Ostwald deduces from similar data the same conclusions as Thomsen, and thus reconfirms the formula deduced by Guldberg and Waage, and the teaching of Berthollet.The participation of water is seen still more clearly in the methods adopted by Ostwald than in those of Thomsen, because in the saturation of solutions of acids by alkalis (which Kremers, Reinhold, and others had previously studied) there is observed, not a contraction, as might have been expected from the quantity of heat which is then evolved, but an expansion, of volume (a decrease of specific gravity, if we calculate as Ostwald did in his first investigations). Thus by mixing 1,880 grams of a solution of sulphuric acid of the composition SO3+ 100H2O, occupying a volume of 1,815 c.c., with a corresponding quantity of a solution 2(NaHO + 5H2O), whose volume = 1,793 c.c., we obtain not 3,608 but 3,633 c.c., an expansion of 25 c.c. per gram molecule of the resulting salt, Na2SO4. It is the same in other cases. Nitric and hydrochloric acids give a still greater expansion than sulphuric acid, and potassium hydroxide than sodium hydroxide, whilst a solution of ammonia gives a contraction. The relation to water must be considered as the cause of these phenomena. When sodium hydroxide and sulphuric acid dissolve in water they develop heat and give a vigorous contraction; the water is separated from such solutions with great difficulty. After mutual saturation they form the salt Na2SO4, which retains the water but feebly and evolves but little heat with it, i.e., in other words, has little affinity for water. In the saturation of sulphuric acid by soda the water is, so to say, displaced from a stable combination and passes into an unstable combination; hence an expansion (decrease of sp. gr.) takes place. It is not the reaction of the acid on the alkali, but the reaction of water, that produces the phenomenon by which Ostwald desires to measure the degree of salt formation. The water, which escaped attention, itself has affinity, and influences those phenomena which are being investigated. Furthermore, in the given instance its influence is very great because its mass is large. When it is not present, or only present in small quantities, the attraction of the base to the acid leads to contraction, and not expansion. Na2O has a sp. gr. 2·8, hence its molecular volume = 22; the sp. gr. of SO3is 1·9 and volume 41, hence the sum of their volumes is 63; for Na2SO4the sp. gr. is 2·65 and volume 53·6, consequently there is a contraction of 10 c.c. per gram-molecule of salt. The volume of H2SO4= 53·3, that of 2NaHO = 37·4; there is produced 2H2O, volume = 36, + Na2SO4, volume = 53·6. There react 90·7 c.c., and on saturation there result 89·6 c.c.; consequently contraction again ensues, although less, and although this reaction is one of substitution and not of combination. Consequently the phenomena studied by Ostwald depend but little on the measure of the reaction of the salts, and more on the relations of the dissolved substances to water. In substitutions, for instance 2NaNO3+ H2SO4= 2HNO3+ Na2SO4, the volumes vary but slightly: in the above example they are 2(38·8) + 53·3 and 2(41·2) + 53·6; hence 131 volumes act, and 136 volumes are produced. It may be concluded, therefore, on the basis of what has been said, that on taking water into consideration the phenomena studied by Thomsen and Ostwald are much more complex than they at first appear, and that this method can scarcely lead to a correct interpretation as to the distribution of acids between bases. We may add that P. D. Chroustcheff (1890) introduced a new method for this class of research, by investigating the electro-conductivity of solutions and their mixtures, and obtained remarkable results (for example, that hydrochloric acid almost entirely displaces formic acid and only ⅔ of sulphuric acid), but details of these methods must be looked for in text-books of theoretical chemistry.

Thomsen concludes his investigation with the words: (a) ‘When equivalent quantities of NaHO, HNO3(or HCl) and ½H2SO4react on one another in an aqueous solution, then two-thirds of the soda combines with the nitric and one-third with the sulphuric acid; (b) this subdivision repeats itself, whether the soda be taken combined with nitric or with sulphuric acid; (c) and therefore nitric acid has double the tendency to combine with the base that sulphuric acid has, and hence in an aqueous solution it is a stronger acid than the latter.’

‘It is therefore necessary,’ Thomsen afterwards remarks, ‘to have an expression indicating the tendency of an acid for the saturation of bases. This idea cannot be expressed by the wordaffinity, because by this term is most often understood that force which it is necessary to overcome in order to decompose a substance into its component parts. This force should therefore be measured by the amount of work or heat employed for the decomposition of the substance. The above-mentioned phenomenon is of an entirely different nature,’ and Thomsen introduces the termavidity, by which he designates the tendency of acids for neutralisation. ‘Therefore the avidity of nitric acid with respect to soda is twice as great as the avidity of sulphuric acid. An exactly similar result is obtained with hydrochloric acid, so that its avidity with respect to soda is also double the avidity of sulphuric acid. Experiments conducted with other acids showed that not one of the acids investigated had so great an avidity as nitric acid; some had a greater avidity than sulphuric acid, others less, and in some instances the avidity = 0.’ The reader will naturally see clearly that the path chosen by Thomsen deserves to be worked out, for his results concern important questions of chemistry, but great faith cannot be placed in the deductions he has already arrived at, because great complexity of relations is to be seen in the very method of his investigation. It is especially important to turn attention to the fact that all the reactions investigated are reactions of double decomposition. In them A and B do not combine with C and distribute themselves according to their affinity or avidity for combination, but reversible reactions are induced. MX and NY give MY and NX, and conversely; therefore the affinity or avidity for combination is not here directly determined, but only the difference or relation of the affinities or avidities. The affinity of nitric acid not only for the water of constitution, but also for that serving for solution, is much less than that of sulphuric acid. This is seen from thermal data. The reaction N2O5+ H2O gives +3,600 heat units, and the solution of the resultant hydrate, 2NHO3, in a large excess of water evolves +14,986 heat units. The formation of SO3+ H2O evolves +21,308 heat units, and the solution of H2SO4in an excess of water 17,860—that is, sulphuric acid gives more heat in both cases. The interchange between Na2SO4and 2HNO3is not only accomplished at the expense of the production of NaNO3, but also at the expense of the formation of H2SO4, hence the affinity of sulphuric acid for water plays its part in the phenomena of displacement. Therefore in determinations like those made by Thomsen the water does not form a medium which is present without participating in the process; it also takes part in the reaction. (Compare Chapter IX., Note14.)

Whilst retaining essentially the methods of Thomsen, Ostwald (1876) determined the variation of the sp. gr. (and afterwards of volume), proceeding in the same dilute solutions, on the saturation of acids by bases, and in the decomposition of the salts of one acid by the other, and arrived at conclusions of just the same nature as Thomsen's. Ostwald's method will be clearly understood from an example. A solution of caustic soda containing an almost molecular (40 grams) weight per litre had a specific gravity of 1·04051. The specific gravities of solutions of equal volume and equivalent composition of sulphuric and nitric acids were 1·02970 and 1·03084 respectively. On mixing the solutions of NaHO and H2SO4there was formed a solution of Na2SO4of sp. gr. 1·02959; hence there ensued a decrease of specific gravity which we will term Q, equal to 1·04051 + 1·02970 - 2(1·02959) = 0·01103. So also the specific gravity after mixture of the solutions of NaHO and HNO3was 1·02633, and therefore Q = 0·01869. When one volume of the solution of nitric acid was added to two volumes of the solution of sodium sulphate, a solution of sp. gr. 1·02781 was obtained, and therefore the resultant decrease of sp. gr.

Q1= 2(1·02959) + 1·03084 - 3(1·02781) = 0·00659.

Had there been no chemical reaction between the salts, then according to Ostwald's reasoning the specific gravity of the solutions would not have changed, and if the nitric acid had entirely displaced the sulphuric acid Q2would be = 0·01869 - 0·01103 = 0·00766. It is evident that a portion of the sulphuric acid was displaced by the nitric acid. But the measure of displacement is not equal to the ratio between Q1and Q2, because a decrease of sp. gr. also occurs on mixing the solution of sodium sulphate with sulphuric acid, whilst the mixing of the solutions of sodium nitrate and nitric acid only produces a slight variation of sp. gr. which falls within the limits of experimental error. Ostwald deduces from similar data the same conclusions as Thomsen, and thus reconfirms the formula deduced by Guldberg and Waage, and the teaching of Berthollet.

The participation of water is seen still more clearly in the methods adopted by Ostwald than in those of Thomsen, because in the saturation of solutions of acids by alkalis (which Kremers, Reinhold, and others had previously studied) there is observed, not a contraction, as might have been expected from the quantity of heat which is then evolved, but an expansion, of volume (a decrease of specific gravity, if we calculate as Ostwald did in his first investigations). Thus by mixing 1,880 grams of a solution of sulphuric acid of the composition SO3+ 100H2O, occupying a volume of 1,815 c.c., with a corresponding quantity of a solution 2(NaHO + 5H2O), whose volume = 1,793 c.c., we obtain not 3,608 but 3,633 c.c., an expansion of 25 c.c. per gram molecule of the resulting salt, Na2SO4. It is the same in other cases. Nitric and hydrochloric acids give a still greater expansion than sulphuric acid, and potassium hydroxide than sodium hydroxide, whilst a solution of ammonia gives a contraction. The relation to water must be considered as the cause of these phenomena. When sodium hydroxide and sulphuric acid dissolve in water they develop heat and give a vigorous contraction; the water is separated from such solutions with great difficulty. After mutual saturation they form the salt Na2SO4, which retains the water but feebly and evolves but little heat with it, i.e., in other words, has little affinity for water. In the saturation of sulphuric acid by soda the water is, so to say, displaced from a stable combination and passes into an unstable combination; hence an expansion (decrease of sp. gr.) takes place. It is not the reaction of the acid on the alkali, but the reaction of water, that produces the phenomenon by which Ostwald desires to measure the degree of salt formation. The water, which escaped attention, itself has affinity, and influences those phenomena which are being investigated. Furthermore, in the given instance its influence is very great because its mass is large. When it is not present, or only present in small quantities, the attraction of the base to the acid leads to contraction, and not expansion. Na2O has a sp. gr. 2·8, hence its molecular volume = 22; the sp. gr. of SO3is 1·9 and volume 41, hence the sum of their volumes is 63; for Na2SO4the sp. gr. is 2·65 and volume 53·6, consequently there is a contraction of 10 c.c. per gram-molecule of salt. The volume of H2SO4= 53·3, that of 2NaHO = 37·4; there is produced 2H2O, volume = 36, + Na2SO4, volume = 53·6. There react 90·7 c.c., and on saturation there result 89·6 c.c.; consequently contraction again ensues, although less, and although this reaction is one of substitution and not of combination. Consequently the phenomena studied by Ostwald depend but little on the measure of the reaction of the salts, and more on the relations of the dissolved substances to water. In substitutions, for instance 2NaNO3+ H2SO4= 2HNO3+ Na2SO4, the volumes vary but slightly: in the above example they are 2(38·8) + 53·3 and 2(41·2) + 53·6; hence 131 volumes act, and 136 volumes are produced. It may be concluded, therefore, on the basis of what has been said, that on taking water into consideration the phenomena studied by Thomsen and Ostwald are much more complex than they at first appear, and that this method can scarcely lead to a correct interpretation as to the distribution of acids between bases. We may add that P. D. Chroustcheff (1890) introduced a new method for this class of research, by investigating the electro-conductivity of solutions and their mixtures, and obtained remarkable results (for example, that hydrochloric acid almost entirely displaces formic acid and only ⅔ of sulphuric acid), but details of these methods must be looked for in text-books of theoretical chemistry.

[28]G. G. Gustavson's researches, which were conducted in the laboratory of the St. Petersburg University in 1871–72, are among the first in which the measure of the affinity of the elements for the halogens is recognised with perfect clearness in the limit of substitution and in the rate of reaction. The researches conducted by A. L. Potilitzin (of which mention will be made in Chapter XI., Note66) in the same laboratory touch on another aspect of the same problem which has not yet made much progress, notwithstanding its importance and the fact that the theoretical side of the subject (thanks especially to Guldberg and Van't Hoff) has since been rapidly pushed forward. If the researches of Gustavson took account of the influence of mass, and were more fully supplied with data concerning velocities and temperatures, they would be very important, because of the great significance which the case considered has for the understanding of double saline decompositions in the absence of water.Furthermore, Gustavson showed that the greater the atomic weight of the element (B, Si, Ti, As, Sn) combinedwith chlorinethe greater the amount of chlorine replaced by bromine by the action of CBr4, and consequently the less the amount of bromine replaced by chlorine by the action of CCl4on bromine compounds. For instance, for chlorine compounds the percentage of substitution (at the limit) is—BCl3SiCl4TiCl4AsCl3SnCl410·112·543·671·877·5It should he observed, however, that Thorpe, on the basis of his experiments, denies the universality of this conclusion. I may mention one conclusion which it appears to me may be drawn from the above-cited figures of Gustavson, if they are subsequently verified even within narrow limits. If CBr4be heated with RCl4, then an exchange of the bromine for chlorine takes place. But what would be the result if it were mixed with CCl4? Judging by the magnitude of the atomic weights, B = 11, C = 12, Si = 28, about 11 p.c. of the chlorine would be replaced by bromine. But to what does this point? I think that this shows the existence of a motion of the atoms in the molecule. The mixture of CCl4and CBr4does not remain in a condition of static equilibrium; not only are the molecules contained in it in a state of motion, but also the atoms in the molecules, and the above figures show the measure of their translation under these conditions. The bromine in the CBr4is,within the limit, substituted by the chlorine of the CCl4in a quantity of about 11 out of 100: that is, a portion of the atoms of bromine previously to this moment in combination with one atom of carbon pass over to the other atom of carbon, and the chlorine passes over from this second atom of carbon to replace it. Therefore, also, in the homogeneous mass CCl4all the atoms of Cl do not remain constantly combined with the same atoms of carbon, andthere is on exchange of atoms between different molecules in a homogeneous medium also. This hypothesis may in my opinion explain certain phenomena of dissociation, but though mentioning it I do not consider it worth while to dwell upon it. I will only observe that a similar hypothesis suggested itself to me in my researches on solutions, and that Pfaundler enunciated an essentially similar hypothesis, and in recent times a like view is beginning to find favour with respect to the electrolysis of saline solutions.

Furthermore, Gustavson showed that the greater the atomic weight of the element (B, Si, Ti, As, Sn) combinedwith chlorinethe greater the amount of chlorine replaced by bromine by the action of CBr4, and consequently the less the amount of bromine replaced by chlorine by the action of CCl4on bromine compounds. For instance, for chlorine compounds the percentage of substitution (at the limit) is—

It should he observed, however, that Thorpe, on the basis of his experiments, denies the universality of this conclusion. I may mention one conclusion which it appears to me may be drawn from the above-cited figures of Gustavson, if they are subsequently verified even within narrow limits. If CBr4be heated with RCl4, then an exchange of the bromine for chlorine takes place. But what would be the result if it were mixed with CCl4? Judging by the magnitude of the atomic weights, B = 11, C = 12, Si = 28, about 11 p.c. of the chlorine would be replaced by bromine. But to what does this point? I think that this shows the existence of a motion of the atoms in the molecule. The mixture of CCl4and CBr4does not remain in a condition of static equilibrium; not only are the molecules contained in it in a state of motion, but also the atoms in the molecules, and the above figures show the measure of their translation under these conditions. The bromine in the CBr4is,within the limit, substituted by the chlorine of the CCl4in a quantity of about 11 out of 100: that is, a portion of the atoms of bromine previously to this moment in combination with one atom of carbon pass over to the other atom of carbon, and the chlorine passes over from this second atom of carbon to replace it. Therefore, also, in the homogeneous mass CCl4all the atoms of Cl do not remain constantly combined with the same atoms of carbon, andthere is on exchange of atoms between different molecules in a homogeneous medium also. This hypothesis may in my opinion explain certain phenomena of dissociation, but though mentioning it I do not consider it worth while to dwell upon it. I will only observe that a similar hypothesis suggested itself to me in my researches on solutions, and that Pfaundler enunciated an essentially similar hypothesis, and in recent times a like view is beginning to find favour with respect to the electrolysis of saline solutions.

[29]Berthollet's doctrine is hardly at all affected in principle by showing that there are cases in which there is no decomposition between salts, because the affinity may be so small that even a large mass would still give no observable displacements. The fundamental condition for the application of Berthollet's doctrine, as well as Deville's doctrine of dissociation, lies in the reversibility of reactions. There are practically irreversible reactions (for instance, CCl4+ 2H2O = CO2+ 4HCl), just as there are non-volatile substances. But while accepting the doctrine of reversible reactions and retaining the theory of the evaporation of liquids, it is possible to admit the existence of non-volatile substances, and in just the same way of reactions, without any visible conformity to Berthollet's doctrine. This doctrine evidently comes nearer than the opposite doctrine of Bergmann to solving the complex problems of chemical mechanics for the successful solution of which at the present time the most valuable help is to be expected from the working out of data concerning dissociation, the influence of mass, and the equilibrium and velocity of reactions. But it is evident that from this point of view we must not regard a solvent as a non-participant space, but must take into consideration the chemical reactions accompanying solution, or else bring about reactions without solution.

[30]Common salt not only enters into double decomposition with acids but alsowith every salt. However, as clearly follows from Berthollet's doctrine, this form of decomposition will only in a few cases render it possible for new metallic chlorides to be obtained, because the decomposition will not be carried on to the end unless the metallic chloride formed separates from the mass of the active substances. Thus, for example, if a solution of common salt be mixed with a solution of magnesium sulphate, double decomposition ensues, but not completely, because all the substances remain in the solution. In this case the decomposition must result in the formation of sodium sulphate and magnesium chloride, substances which are soluble in water; nothing is disengaged, and therefore the decomposition 2NaCl + MgSO4= MgCl2+ Na2SO4cannot proceed to the end. However, the sodium sulphate formed in this manner may be separated by freezing the mixture. The complete separation of the sodium sulphate will naturally not take place, owing to a portion of the salt remaining in the solution. Nevertheless, this kind of decomposition is made use of for the preparation of sodium sulphate from the residues left after the evaporation of sea-water, which contain a mixture of magnesium sulphate and common salt. Such a mixture is found at Stassfurt in a natural form. It might be said that this form of double decomposition is only accomplished with a change of temperature; but this would not be true, as may be concluded from other analogous cases. Thus, for instance, a solution of copper sulphate is of a blue colour, while a solution of copper chloride is green. If we mix the two salts together the green tint is distinctly visible, so that by this means the presence of the copper chloride in the solution of copper sulphate is clearly seen. If now we add a solution of common salt to a solution of copper sulphate, a green coloration is obtained, which indicates the formation of copper chloride. In this instance it is not separated, but it is immediately formed on the addition of common salt, as it should be according to Berthollet's doctrine.The complete formation of a metallic chloride from common salt can only occur, judging from the above, when it separates from the sphere of action. The salts of silver are instances in point, because the silver chloride is insoluble in water; and therefore if we add a solution of sodium chloride to a solution of a silver salt, silver chloride and the sodium salt of that acid which was in the silver salt are formed.

The complete formation of a metallic chloride from common salt can only occur, judging from the above, when it separates from the sphere of action. The salts of silver are instances in point, because the silver chloride is insoluble in water; and therefore if we add a solution of sodium chloride to a solution of a silver salt, silver chloride and the sodium salt of that acid which was in the silver salt are formed.

[31]The apparatus shown in fig.46(Chapter VI., Note12) is generally employed for the preparation of small quantities of hydrochloric acid. Common salt is placed in the retort; the salt is generally previously fused, as it otherwise froths and boils over in the apparatus. When the apparatus is placed in order sulphuric acid mixed with water is poured down the thistle funnel into the retort. Strong sulphuric acid (about half as much again as the weight of the salt) is usually taken, and it is diluted with a small quantity of water (half) if it be desired to retard the action, as in using strong sulphuric acid the action immediately begins with great vigour. The mixture, at first without the aid of heat and then at a moderate temperature (in a water-bath), evolves hydrochloric acid. Commercial hydrochloric acid contains many impurities; it is usually purified by distillation, the middle portions being collected. It is purified from arsenic by adding FeCl2, distilling, and rejecting the first third of the distillate. If free hydrochloric acid gas be required, it is passed through a vessel containing strong sulphuric acid to dry it, and is collected over a mercury hath.Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler, 1888; 2P2O3+ 3HCl = POCl3+ 3HPO3) at the ordinary temperature, and therefore the gas cannot he dried by this substance.

Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler, 1888; 2P2O3+ 3HCl = POCl3+ 3HPO3) at the ordinary temperature, and therefore the gas cannot he dried by this substance.

[31 bis]In chemical works where sulphuric acid of 60° Baumé (22 p.c. of water) is employed, 117 parts of sodium chloride are taken to about 125 parts of sulphuric acid.

[32]As in works which treat common salt in order to obtain sodium sulphate, the hydrochloric acid is sometimes held to be of no value, it might be allowed to escape with the waste furnace gases into the atmosphere, which would greatly injure the air of the neighbourhood and destroy all vegetation. In all countries, therefore, there are laws forbidding the factories to proceed in this manner, and requiring the absorption of the hydrochloric acid by water at the works themselves, and not permitting the solution to be run into rivers and streams, whose waters it would spoil. It may be remarked that the absorption of hydrochloric acid presents no particular difficulties (the absorption of sulphurous acid is much more difficult) because hydrochloric acid has a great affinity for water and gives a hydrate which boils above 100°. Hence, even steam and hot water, as well as weaker solutions, can be used for absorbing the acid. However, Warder (1888) showed that weak solutions of composition H2O +nHCl when boiled (the residue will be almost HCl,8H2O) evolve (not water but) a solution of the composition H2O + 445n4HCl; for example, on distilling HCl,10H2O, HCl,23H2O is first obtained in the distillate. As the strength of the residue becomes greater, so also does that of the distillate, and therefore in order to completely absorb hydrochloric acid it is necessary in the end to have recourse to water.As in Russia the manufacture of sodium sulphate from sodium chloride has not yet been sufficiently developed, and as hydrochloric acid is required for many technical purposes (for instance, for the preparation of zinc chloride, which is employed for soaking railway sleepers), therefore salt is often treated mainly for the manufacture of hydrochloric acid.

As in Russia the manufacture of sodium sulphate from sodium chloride has not yet been sufficiently developed, and as hydrochloric acid is required for many technical purposes (for instance, for the preparation of zinc chloride, which is employed for soaking railway sleepers), therefore salt is often treated mainly for the manufacture of hydrochloric acid.

[33]Thus the metallic chlorides, which are decomposed to a greater or less degree by water, correspond with feeble bases. Such are, for example, MgCl2, AlCl3, SbCl3, BiCl3. The decomposition of magnesium chloride (and also carnallite) by sulphuric acid proceeds at the ordinary temperature; water decomposes MgCl2to the extent of 50 p.c. when aided by heat, andmay be employedas a convenientmethod for the production of hydrochloric acid. Hydrochloric acid is also produced by the ignition of certain metallic chlorides in a stream of hydrogen, especially of those metals which are easily reduced and difficultly oxidised—for instance, silver chloride. Lead chloride, when heated to redness in a current of steam, gives hydrochloric acid and lead oxide. The multitude of the cases of formation of hydrochloric acid are understood from the fact that it is a substance which is comparatively very stable, resembling water in this respect, and even most probably more stable than water, because, at a high temperature and even under the action of light, chlorine decomposes water, with the formation of hydrochloric acid. The combination of chlorine and hydrogen also proceeds by their direct action, as we shall afterwards describe.

[34]According to Ansdell (1880) the sp. gr. of liquid hydrochloric acid at 0° = 0·908, at 11·67° = 0·854, at 22·7° = 0·808, at 33° = 0·748. Hence it is seen that the expansion of this liquid is greater than that of gases (Chapter II., Note34).

[35]According to Roscoe and Dittmar at a pressure of three atmospheres the solution of constant boiling point contains 18 p.c. of hydrogen chloride, and at a pressure of one-tenth atmosphere 23 p.c. The percentage is intermediate at medium pressures.

[36]At 0° 25 p.c., at 100° 20·7 p.c.; Roscoe and Dittmar.

[37]This crystallo-hydrate (obtained by Pierre and Puchot, and investigated by Roozeboom) is analogous to NaCl,2H2O. The crystals HCl,2H2O at -22° have a specific gravity 1·46; the vapour tension (under dissociation) of the solution having a composition HCl,2H2O at -24° = 760, at -19° = 1,010, at -18° = 1,057, at -17° = 1,112 mm. of mercury. In a solid state the crystallo-hydrate at -17·7° has the same tension, whilst at lower temperatures it is much less: at -24° about 150, at -19° about 580 mm. A mixture of fuming hydrochloric acid with snow reduces the temperature to -38°. If another equivalent of water be added to the hydrate HCl,2H2O at -18°, the temperature of solidification falls to -25°, and the hydrate HCl,3H2O is formed (Pickering, 1893).

[38]According to Roscoe at 0° onehundredgrams of water at a pressurep(in millimetres of mercury) dissolves—p=1002003005007001000Grams HCl65·770·773·878·281·785·6At a pressure of 760 millimetres and temperaturet, onehundredgrams of water dissolvest=08°16°24°40°60°Grams HCl82·578·374·270·063·356·1Roozeboom (1886) showed that att° solutions containingcgrams of hydrogen chloride per 100 grams of water may (with the variation of the pressurep) be formed together with the crystallo-hydrate HCl,2H2O:t=-28°·8-21°-19°-18°c=84·286·892·698·4101·4p=—3345809001,073 mm.The last combination answers to the melted crystallo-hydrate HCl,2H2O, which splits up at temperatures above -17°·7, and at a constant atmospheric pressure when there are no crystals—t=-24°-21°-18°-10°-0°c=101·298·395·789·884·2From these data it is seen that the hydrate HCl,2H2O can exist in a liquid state, which is not the case for the hydrates of carbonic and sulphurous anhydrides, chlorine, &c.According to Marignac, the specific heatcof a solution HCl +mH2O (at about 30°, taking the specific heat of water = 1) is given by the expression—C(36·5 +m18) = 18m- 28·39 + 140/m- 268/m2ifmbe not less than 6·25. For example, for HCl + 25H2O, C = 0·877.According to Thomsen's data, the amount of heatQ, expressed in thousands of calories, evolved in the solution of 36·5 grams of gaseous hydrochloric acid inmH2O or 18mgrams of water is equal to—m=241050400Q=11·414·316·217·117·3In these quantities the latent heat of liquefaction is included, which must be taken as 5–9 thousand calories per molecular quantity of hydrogen chloride.The researches of Scheffer (1888) on the rate of diffusion (in water) of solutions of hydrochloric acid show that the coefficient of diffusionkdecreases with the amount of watern, if the composition of the solution is HCl,nH2O at 0°:—n=56·99·81427·1129·5k=2·312·081·861·671·521·39It also appears that strong solutions diffuse more rapidly into dilute solutions than into water.

[38]According to Roscoe at 0° onehundredgrams of water at a pressurep(in millimetres of mercury) dissolves—

At a pressure of 760 millimetres and temperaturet, onehundredgrams of water dissolves

Roozeboom (1886) showed that att° solutions containingcgrams of hydrogen chloride per 100 grams of water may (with the variation of the pressurep) be formed together with the crystallo-hydrate HCl,2H2O:

The last combination answers to the melted crystallo-hydrate HCl,2H2O, which splits up at temperatures above -17°·7, and at a constant atmospheric pressure when there are no crystals—

From these data it is seen that the hydrate HCl,2H2O can exist in a liquid state, which is not the case for the hydrates of carbonic and sulphurous anhydrides, chlorine, &c.

According to Marignac, the specific heatcof a solution HCl +mH2O (at about 30°, taking the specific heat of water = 1) is given by the expression—

C(36·5 +m18) = 18m- 28·39 + 140/m- 268/m2

ifmbe not less than 6·25. For example, for HCl + 25H2O, C = 0·877.

According to Thomsen's data, the amount of heatQ, expressed in thousands of calories, evolved in the solution of 36·5 grams of gaseous hydrochloric acid inmH2O or 18mgrams of water is equal to—

In these quantities the latent heat of liquefaction is included, which must be taken as 5–9 thousand calories per molecular quantity of hydrogen chloride.

The researches of Scheffer (1888) on the rate of diffusion (in water) of solutions of hydrochloric acid show that the coefficient of diffusionkdecreases with the amount of watern, if the composition of the solution is HCl,nH2O at 0°:—

It also appears that strong solutions diffuse more rapidly into dilute solutions than into water.

[39]If it be admitted that the maximum of the differential corresponds with HCl,6H2O, then it might be thought that the specific gravity is expressed by a parabola of the third order; but such an admission does not give expressions in accordance with fact. This is all more fully considered in my work mentioned in Chapter I., Note19.

[40]As in water, the coefficient of expansion (or the quantitykin the expression St= S0-kS0t, or Vt= 1/(1 -kt)) attains a magnitude 0·000447 at about 48°, it might be thought that at 48° all solutions of hydrochloric acid would have the same coefficient of expansion, but in reality this is not the case. At low and at the ordinary temperatures the coefficient of expansion of aqueous solutions is greater than that of water, and increases with the amount of substance dissolved.

[41]The figures cited above may serve for the direct determination of that variation of the specific gravity of solutions of hydrochloric acid with the temperature. Thus, knowing that at 15° the specific gravity of a 10 p.c. solution of hydrochloric acid = 10,492, we find that att° it = 10,530 -t(2·13 + 0·027t). Whence also may be found the coefficient of expansion (Note40).

[42]Thus, for instance, with feeble bases they evolve in dilute solutions (Chapter III., Note 53) almost equal amounts of heat; their relation to sulphuric acid is quite identical. They both form fuming solutions as well as hydrates; they both form solutions of constant boiling point.

[42 bis]Pybalkin (1891) found that copper begins to disengage hydrogen at 100°, and that chloride of copper begins to give up its chlorine to hydrogen gas at 230°; for silver these temperatures are 117° and 260°—that is, there is less difference between them.

[43]When an unsaturated hydrocarbon, or, in general, an unsaturated compound, assimilates to itself the molecules Cl2, HCl, SO3, H2SO4, &c., the cause of the reaction is most simple. As nitrogen, besides the type NX3to which NH3, belongs, gives compounds of the type NX5—for example, NO2(OH)—the formation of the salts of ammonium should be understood in this way. NH3gives NH4Cl because NX3is capable of giving NX5. But as saturated compounds—for instance, SO3,H2O, NaCl, &c.—are also capable of combination even between themselves, it is impossible to deny the capacity of HCl also for combination. SO3combines with H2O, and also with HCl and the unsaturated hydrocarbons. It is impossible to recognise the distinction formerly sought to be established between atomic and molecular compounds, and regarding, for instance, PCl3as an atomic compound and PCl5as a molecular one, only because it easily splits up into molecules PCl3and Cl2.

[44]Sal-ammoniac is prepared from ammonium carbonate, obtained in the dry distillation of nitrogenous substances (ChapterVI.), by saturating the resultant solution with hydrochloric acid. A solution of sal-ammoniac is thus produced, which is evaporated, and in the residue a mass is obtained containing a mixture of various other, especially tarry, products of dry distillation. The sal-ammoniac is generally purified by sublimation. For this purpose iron vessels covered with hemispherical metallic covers are employed, or else simply clay crucibles covered by other crucibles. The upper portion, or head, of the apparatus of this kind will have a lower temperature than the lower portion, which is under the direct action of the flame. The sal-ammoniac volatilises when heated, and settles on the cooler portion of the apparatus. It is thus freed from many impurities, and is obtained as a crystalline crust, generally several centimetres thick, in which form it is commonly sold. The solubility of sal-ammoniac rises rapidly with the temperature: at 0°, 100 parts of water dissolve about 28 parts of NH4Cl, at 50° about 50 parts, and at the ordinary temperature about 35 parts. This is sometimes taken advantage of for separating NH4Cl from solutions of other salts.

[45]The solubility of sal-ammoniac in 100 parts of water (according to Alluard) is—0°10°20°30°40°60°80°100°100°28·4032·4837·2841·724655647377A saturated solution boils at 115°·8. The specific gravity at 15°/4° of solutions of sal-ammoniac (water 4° = 10,000) = 9,991·6 - 31·26p- 0·085p2, wherepis the amount by weight of ammonium chloride in 100 parts of solution. With the majority of salts the differentialds/dpincreases, but here it decreases with the increase ofp. For (unlike the sodium and potassium salts) a solution of the alkaliplusa solution of acid occupy a greater volume than that of the resultant ammonium salt. In the solution ofsolidammonium chloride a contraction, and not expansion, generally takes place. It may further be remarked that solutions of sal-ammoniac have an acid reaction even when prepared from the salt remaining after prolonged washing of the sublimed salt with water (A. Stcherbakoff).

[45]The solubility of sal-ammoniac in 100 parts of water (according to Alluard) is—

A saturated solution boils at 115°·8. The specific gravity at 15°/4° of solutions of sal-ammoniac (water 4° = 10,000) = 9,991·6 - 31·26p- 0·085p2, wherepis the amount by weight of ammonium chloride in 100 parts of solution. With the majority of salts the differentialds/dpincreases, but here it decreases with the increase ofp. For (unlike the sodium and potassium salts) a solution of the alkaliplusa solution of acid occupy a greater volume than that of the resultant ammonium salt. In the solution ofsolidammonium chloride a contraction, and not expansion, generally takes place. It may further be remarked that solutions of sal-ammoniac have an acid reaction even when prepared from the salt remaining after prolonged washing of the sublimed salt with water (A. Stcherbakoff).

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