The Principles of Chemistry, Volume II

by Dmitry Ivanovich Mendeleyev

[22 tri] It may be remarked that the black stain produced by the reduction of metallic silver disappears under the action of a solution of mercuric chloride or of potassium cyanide, because these salts act on finely-divided silver.

[23] Silver chloride is almost perfectly insoluble in water, but is somewhat soluble in water containing sodium chloride or hydrochloric acid, or other chlorides, and many salts, in solution. Thus at 100°, 100 parts of water saturated with sodium chloride dissolve 0·4 part of silver chloride. Bromide and iodide of silver are less soluble in this respect, as also in regard to other solvents. It should be remarked that silver chloride dissolves in solutions of ammonia, potassium cyanide, and of sodium thiosulphate, Na₂S₂O₃. Silver bromide is almost perfectly analogous to the chloride, but silver iodide is nearly insoluble in a solution of ammonia. Silver chloride even absorbs dry ammonia gas, forming very unstable ammoniacal compounds. When heated, these compounds (Vol. I. p. 250, Note [8]) evolve the ammonia, as they also do under the action of all acids. Silver chloride enters into double decomposition with potassium cyanide, forming a soluble double cyanide, which we shall presently describe; it also forms a soluble double salt, NaAgS₂O₃, with sodium thiosulphate.

Silver chloride offers different modifications in the structure of its molecule, as is seen in the variations in the consistency of the precipitate, and in the differences in the action of light which partially decomposes AgCl (see Note [25]). Stas and Carey Lea investigated this subject, which has a particular importance in photography, because silver bromide also gives photo-salts. There is still much to be discovered in this respect, since Abney showed that perfectly dry AgCl placed in a vacuum in the dark is not in the least acted upon when subsequently exposed to light.

[24] Silver bromide and iodide (which occur as the minerals bromite and iodite) resemble the chloride in many respects, but the degree of affinity of silver for iodine is greater than that for chlorine and bromine, although less heat is evolved (see Note [28 bis]). Deville deduced this fact from a number of experiments. Thus silver chloride, when treated with hydriodic acid, evolves hydrochloric acid, and forms silver iodide. Finely-divided silver easily liberates hydrogen when treated with hydriodic acid; it produces the same decomposition with hydrochloric acid, but in a considerably less degree and only on the surface. The difference between silver chloride and iodide is especially remarkable, since the formation of the former is attended with a greater contraction than that of the latter. The volume of AgCl = 26; of chlorine 27, of silver 10, the sum = 37, hence a contraction has ensued; and in the formation of silver iodide an expansion takes place, for the volume of Ag is 10, of I 26, and of AgI 39 instead of 36 (density, AgCl, 5·59; AgI, 5·67). The atoms of chlorine have united with the atoms of silver without moving asunder, whilst the atoms of iodine must have moved apart in combining with the silver. It is otherwise with respect to the metal; the distance between its atoms in the metal = 2·2, in silver chloride = 3·0, and in silver iodide = 3·5; hence its atoms have moved asunder considerably in both cases. It is also very remarkable, as Fizeau observed, that the density of silver iodide increases with a rise of temperature—that is, a contraction takes place when it is heated and an expansion when it is cooled.

In order to explain the fact that in silver compounds the iodide is more stable than the chloride and oxide, Professor N. N. Beketoff, in his ‘Researches on the Phenomena of Substitutions’ (Kharkoff, 1865), proposed the following original hypothesis, which we will give in almost the words of the author:—In the case of aluminium, the oxide, Al₂O₃, is more stable than the chloride, Al₂Cl₆, and the iodide, Al₂I₆. In the oxide the amount of the metal is to the amount of the element combined with it as 54·8 (Al = 27·3) is to 48, or in the ratio 112 : 100; for the chloride the ratio is = 25 : 100; for the iodide it = 7 : 100. In the case of silver the oxide (ratio = 1350 : 100) is less stable than the chloride (ratio = 304 : 100), and the iodide (ratio of the weight of metal to the weight of the halogen = 85 : 100) is the most stable. From these and similar examples it follows that the most stable compounds are those in which the weights of the combined substances are equal. This may be partly explained by the attraction of similar molecules even after their having passed into combination with others. This attraction is proportional to the product of the acting masses. In silver oxide the attraction of Ag₂ for Ag₂ = 216 × 216 = 46,656, and the attraction of Ag₂ for O = 216 × 16 = 3,456. The attraction of like molecules thus counteracts the attraction of the unlike molecules. The former naturally does not overcome the latter, otherwise there would be a disruption, but it nevertheless diminishes the stability. In the case of an equality or proximity of the magnitude of the combining masses, the attraction of the like parts will counteract the stability of the compound to the least extent—in other words, with an inequality of the combined masses, the molecules have an inclination to return to an elementary state, to decompose, which does not exist to such an extent where the combined masses are equal. There is, therefore, a tendency for large masses to combine with large, and for small masses to combine with small. Hence Ag₂O + 2KI gives K₂O + 2AgI. The influence of an equality of masses on the stability is seen particularly clearly in the effect of a rise of temperature. Argentic, mercuric, auric and other oxides composed of unequal masses, are somewhat readily decomposed by heat, whilst the oxides of the lighter metals (like water) are not so easily decomposed by heat. Silver chloride and iodide approach the condition of equality, and are not decomposed by heat. The most stable oxides under the action of heat are those of magnesium, calcium, silicon, and aluminium, since they also approach the condition of equality. For the same reason hydriodic acid decomposes with greater facility than hydrochloric acid. Chlorine does not act on magnesia or alumina, but it acts on lime and silver oxide, &c. This is partially explained by the fact that by considering heat as a mode of motion, and knowing that the atomic heats of the free elements are equal, it must be supposed that the amount of the motion of atoms (their vis viva) is equal, and as it is equal to the product of the mass (atomic weight) into the square of the velocity, it follows that the greater the combining weight the smaller will be the square of the velocity, and if the combining weights be nearly equal, then the velocities also will be nearly equal. Hence the greater the difference between the weights of the combined atoms the greater will be the difference between their velocities. The difference between the velocities will increase with the temperature, and therefore the temperature of decomposition will be the sooner attained the greater be the original difference—that is, the greater the difference of the weights of the combined substances. The nearer these weights are to each other, the more analogous the motion of the unlike atoms, and consequently, the more stable the resultant compound.

The instability of cupric chloride and nitric oxide, the absence of compounds of fluorine with oxygen, whilst there are compounds of oxygen with chlorine, the greater stability of the oxygen compounds of iodine than those of chlorine, the stability of boron nitride, and the instability of cyanogen, and a number of similar instances, where, judging from the above argument, one would expect (owing to the closeness of the atomic weights) a stability, show that Beketoff's addition to the mechanical theory of chemical phenomena is still far from sufficient for explaining the true relations of affinities. Nevertheless, in his mode of explaining the relative stabilities of compounds, we find an exceedingly interesting treatment of questions of primary importance. Without such efforts it would be impossible to generalise the complex data of experimental knowledge.

Fluoride of silver, AgF, is obtained by dissolving Ag₂O or Ag₂CO₃ in hydrofluoric acid. It differs from the other halogen salts of silver in being soluble in water (1 part of salt in 0·55 of water). It crystallises from its solution in prisms, AgFH₂O (Marignac), or AgF₂H₂O (Pfaundler), which lose their water in vacuo. Güntz (1891), by electrolising a saturated solution of Ag₂F, obtained polyfluoride of silver, Ag₂F, which is decomposed by water into AgF + Ag. It is also formed by the action of a strong solution of AgF upon finely-divided (precipitated) silver.

[24 bis] The changes brought about by the action of light necessitate distinguishing the photo-salts of silver.

[25] In photography these are called ‘developers.’ The most common developers are: solutions of ferrous sulphate, pyrogallol, ferrous oxalate, hydroxylamine, potassium sulphite, hydroquinone (the last acts particularly well and is very convenient to use), &c. The chemical processes of photography are of great practical and theoretical interest; but it would be impossible in this work to enter into this special branch of chemistry, which has as yet been very little worked out from a theoretical point of view. Nevertheless, we will pause to consider certain aspects of this subject which are of a purely chemical interest, and especially the facts concerning subchloride of silver, Ag₂Cl (see Note [19]), and the photo-salts (Note [23]). There is no doubt that under the action of light, AgCl becomes darker in colour, decreases in weight, and probably forms a mixture of AgCl, Ag₂Cl, and Ag. But the isolation of the subchloride has only been recently accomplished by Güntz by means of the Ag₂F, discovered by him (see Note [24]). Many chemists (and among them Hodgkinson) assumed that an oxychloride of silver was formed by the decomposition of AgCl under the action of light. Carey Lea's (1889) and A. Richardson's (1891) experiments showed that the product formed does not, however, contain any oxygen at all, and the change in colour produced by the action of light upon AgCl is most probably due to the formation of Ag₂Cl. This substance was isolated by Güntz (1891) by passing HCl over crystals of Ag₂F. He also obtained Ag₂I in a similar manner by passing HI, and Ag₂S by passing H₂S over Ag₂F. Ag₂Cl is best prepared by the action of phosphorus trichloride upon Ag₂F. At the temperature of its formation Ag₂Cl has an easily changeable tint, with shades of violet red to violet black. Under the action of light a similar (isomeric) substance is obtained, which splits up into AgCl + Ag when heated. With potassium cyanide Ag₂Cl gives Ag + AgCN + KCl, whence it is possible to calculate the heat of formation of Ag₂Cl; it = 29·7, whilst the heat of formation of AgCl = 29·2—i.e. the reaction 2AgCl = Ag₂Cl + Cl corresponds to an absorption of 28·7 major calories. If we admit the formation of such a compound by the action of light, it is evident that the energy of the light is consumed in the above reaction. Carey Lea (1892) subjected AgCl, AgBr, and AgI to a pressure (of course in the dark) of 3,000 atmospheres, and to trituration with water in a mortar, and observed a change of colour indicating incipient decomposition, which is facilitated under the action of light by the molecular currents set up (Lermontoff, Egoroff). The change of colour of the halogen salts of silver under the action of light, and their faculty of subsequently giving a visible photographic image under the action of ‘developers,’ must now be regarded as connected with the decomposition of AgX, leading to the formation of Ag₂X, and the different tinted photo-salts must be considered as systems containing such Ag₂X's. Carey Lea obtained photo-salts of this kind not only by the action of light but also in many other ways, which we will enumerate to prove that they contain the products of an incomplete combination of Ag with the halogens, (for the salts Ag₂X must be regarded as such). The photo-salts have been obtained (1) by the imperfect chlorination of silver; (2) by the incomplete decomposition of Ag₂O or Ag₂CO₃ by alternately heating and treating with a halogen acid; (3) by the action of nitric acid or Na₂S₂O₃ upon Ag₂Cl; (4) by mixing a solution of AgNO₃ with the hydrates of FeO, MnO and CrO, and precipitating by HCl; (5) by the action of HCl upon the product obtained by the reduction of citrate of silver in hydrogen (Note [19]), and ([6]) by the action of milk sugar upon AgNO₃ together with soda and afterwards acidulating with HCl. All these reactions should lead to the formation of products of imperfect combination with the halogens and give photo-salts of a similar diversity of colour to those produced by the action of developers upon the halogen salts of silver after exposure to light.

[25 bis] In order to determine when the reaction is at an end, a few drops of a solution of K₂CrO₄ are added to the solution of the chloride. Before all the chlorine is precipitated as AgCl, the precipitate (after shaking) is white (since Ag₂CrO₄ with 2RCl gives 2AgCl); but when all the chlorine is thrown down Ag₂CrO₄ is formed, which colours the precipitate reddish-brown. In order to obtain accurate results the liquid should be neutral to litmus.