It is a comparatively easy matter to determine the relative solubility of zinc and ferrous sulphides. If equal quantities of the equivalent solutions are mixed and a precipitant, ammonium sulphide, which would precipitate either sulphide, if its salt were present alone, is carefully and gradually added to the mixture, it will precipitate first the less soluble one (see p. [163]); and that one alone can be present permanently (i.e. in equilibrium) in contact with the solution containing the two salts. As a matter of fact,[409] [p206] we find that zinc sulphide is precipitated first, under conditions permitting the precipitation of ferrous sulphide if no zinc sulphate were present, and the precipitate of zinc sulphide remains unchanged in the presence of a mixture of the composition indicated (exp.).
It is clear, therefore, that the prediction, based on the conclusion drawn from the application of the principle of the solubility-product, is verified by experiment.
Now, closer examination of the solution of zinc sulphate, from which zinc sulphide has been precipitated by the action of hydrogen sulphide, shows, after the hydrogen sulphide has precipitated as much sulphide as it can and the solution has been passed through a filter, that a very considerable proportion of zinc salt is still present in the filtrate, and we must ask why hydrogen sulphide fails to precipitate the zinc completely. The concentration of the zinc-ion has grown somewhat smaller, but that is not the cause of the nonprecipitation of zinc sulphide under the new conditions, since hydrogen sulphide will precipitate the sulphide, if it is passed into a solution of zinc sulphate which contains even a smaller concentration of zinc-ion than the filtrate, in which it fails to give any further precipitate. If the filtrate is examined, it is found to be strongly acid, since sulphuric acid has been liberated by the action of hydrogen sulphide on zinc sulphate: ZnSO4 + H2S → ZnS ↓ + H2SO4. Sulphuric acid is a strong acid, which is very highly ionized, much more so than the exceedingly weak acid hydrogen sulphide, and consequently, as the precipitation of zinc sulphide proceeds, the concentration of hydrogen-ion in the solution rapidly grows larger and larger. But, the greater the concentration of the hydrogen-ion, the smaller is that of the sulphide-ion, since the product [H+]2 × [S2−] is a constant [equation (IV), p. [201]] for a solution kept saturated with hydrogen sulphide. The sulphide-ion is reduced in concentration very much more rapidly than is the zinc-ion.[410] As [S2−] is a factor in the [p207] solubility-product of zinc sulphide, it is clear that the value of this product must grow rapidly smaller during the precipitation of zinc sulphide from the solution, and that it may well, eventually, grow too small to surpass the value of the solubility-product constant KZnS. Precipitation of zinc sulphide will then cease. Obviously, the suppression of the sulphide-ion may be accomplished by the addition of hydrochloric, sulphuric or any other strong acid to the zinc sulphate solution in the first place, and then hydrogen sulphide should fail to precipitate any zinc sulphide at all. In fact, if to 50 c.c. of the 0.1 molar zinc sulphate solution 2 c.c. of hexamolar hydrochloric acid is added,[411] hydrogen sulphide does not precipitate even a trace of zinc sulphide (exp.).
We have found, then, that 0.1 molar zinc sulphate solution, acidified with a small excess of hydrochloric acid, fails to produce a precipitate of zinc sulphide, when it is saturated with hydrogen sulphide. We must conclude that, under these circumstances, the product of the ion concentrations is smaller than the solubility-product constant for zinc sulphide: ([Zn2+] × [S2−] / x) < KZnS, the new concentration of the sulphide-ion being represented by the symbol [S2−] / x.
It would follow, from these considerations, that the action of hydrochloric or sulphuric acid, in preventing the precipitation of zinc sulphide, depends on their producing a sufficiently high concentration of the hydrogen-ion, to keep the concentration of the sulphide-ion, in a mixture of zinc sulphate and hydrogen sulphide, below the point where the solution could become supersaturated with zinc sulphide. [p208] For exactly similar reasons, none of the sulphides of the zinc group is precipitated by hydrogen sulphide in (sufficiently) acid solutions.
It is evident, further, that, if a solution of zinc acetate (without the addition of any acid) is substituted for the zinc sulphate solution and is treated with hydrogen sulphide, an entirely different result, quantitatively considered, must be obtained. By the action of hydrogen sulphide on the acetate, acetic acid is liberated, according to Zn(CH3CO2)2 + H2S ⥂ ZnS ↓ + 2 CH3COOH. As a weak acid, acetic acid produces much less hydrogen-ion than is formed in equivalent solutions of sulphuric acid. Consequently, a much slighter suppression of the sulphide-ion and a much more complete precipitation of zinc sulphide from the acetate, than from the sulphate solution, must result. Such is the case. Zinc sulphide is, indeed, precipitated quantitatively by hydrogen sulphide from the acetate solution.
This behavior of zinc acetate[412]—and zinc salts of other weak acids show, of course, the same behavior—represents one of the pitfalls, into which the unwary analytical chemist is liable to fall, when he uses hydrogen sulphide. The separation of groups by hydrogen sulphide depends, as stated, on the fact, that, in the presence of a certain concentration of hydrogen-ion, hydrogen sulphide will not precipitate zinc sulphide and the remaining sulphides of the zinc group. To secure this concentration of hydrogen-ion, some hydrochloric acid is added to solutions, from which hydrogen sulphide is expected to precipitate none but sulphides of the copper and arsenic groups—and, as a rule, the purpose is accomplished, as desired. It is evident, however, that if a solution contains an acetate, say sodium acetate, or the salt of any other weak acid, e.g. a borate or a phosphate, the addition of hydrochloric acid will result, at least at first, in the liberation of the weaker acid and will not produce the excess of hydrogen-ion, required for the analysis. Zinc sulphide, and possibly nickel and cobalt sulphides,[412] may, under such conditions, be precipitated with the sulphides of the groups mentioned. Unless provision is made, therefore, to insure a certain excess of hydrogen-ion (p. [213]), or unless we are on our guard and look for zinc, nickel and cobalt in [p209] the analysis of the precipitate formed by hydrogen sulphide,[413] serious errors obviously could result. To add an inordinately large excess of hydrochloric acid to mixtures, in order to avoid this pitfall, will, as we shall presently see, only throw us more certainly into still another error, to which we are exposed in the use of this important reagent, hydrogen sulphide.
III. Precipitation of Cadmium Sulphide.
By the conditions of the experiment we started with a concentration of the cadmium-ion, [Cd2+], equal to the concentration, [Zn2+], of the zinc-ion in the zinc sulphate solution, and with the same concentration,[415] [S2−] / x, of sulphide-ion as was used when hydrogen sulphide failed to precipitate zinc sulphide. The corresponding factors of the products of the ion concentrations are equal, at the beginning of the two experiments, and we may put ([Cd2+] × [S2−] / x) = ([Zn2+] × [S2−] / x) = P′. We recall the fact, that we have already concluded, on the basis of the principle of the solubility-product, that [Zn2+] × [S2−] / x, or P′, is smaller than KZnS (p. [207]), and that [Cd2+] × [S2−] / x, or P′, is greater than KCdS.
P′ being smaller than KZnS and larger than KCdS, it is clear that [p210] KCdS is smaller than KZnS and that cadmium sulphide must be the less soluble of the two sulphides. As a matter of fact, if ammonium sulphide is carefully added to a mixture of equal quantities of the two salt solutions, cadmium sulphide is precipitated first, and when practically all of the cadmium is precipitated, a final precipitate of white zinc sulphide is obtained (exp.; see note, p. [205]). Or, if zinc sulphide is first precipitated by the addition of a little ammonium sulphide to 25 c.c. of the 0.1 molar zinc sulphate solution, care being taken to have zinc sulphate in excess, and if 25 c.c. of the 0.1 molar cadmium sulphate solution is then added to the mixture, the white zinc sulphide immediately gives way to the less soluble yellow cadmium sulphide (exp.; see p. [165]). Cadmium sulphide is thus proved to be the less soluble of the two sulphides, a result which confirms the prediction made above with the aid of the principle of the solubility-product, and we may indeed conclude that the solubility-product constant KCdS of cadmium sulphide must be smaller than the constant KZnS of zinc sulphide (see pp. [163]–[168], on fractional precipitation). We are, therefore, also justified in deciding that CdS may well be precipitated from acidulated solutions by hydrogen sulphide, when ZnS is not thus precipitated, simply because KCdS is sufficiently small (CdS is sufficiently insoluble) to make the product of the ion concentrations [Cd2+] × [S2−] / x, in spite of the extremely small value of [S2−] / x, greater than the constant KCdS, whereas the same small value of [S2−] / x makes it impossible for the product [Zn2+] × [S2−] / x to reach the value of the larger constant KZnS, required for the precipitation of ZnS. Since cadmium sulphide may be precipitated quantitatively under the conditions given, it is also evident that it may be precipitated even when the concentration of the cadmium-ion also has a rather small value. The relations, in regard to this point, will be discussed presently.