CHAPTER XI

«THE COPPER AND SILVER GROUPS. PRECIPITATION WITH HYDROGEN SULPHIDE»

The sulphides of the metal ions of the zinc group are readily precipitated by ammonium or sodium sulphide, but hydrogen sulphide, in the presence of a small excess of a strong acid, such as hydrochloric acid, does not precipitate any of these sulphides (or any of the sulphides of the aluminium, the alkaline earth and the alkali groups). Under the same conditions ‹the sulphides of the metal ions of the silver group›, Ag^{+}, Hg^{+}, Pb^{2+}, of ‹the copper group›, Hg^{2+}, Pb^{2+}, Bi^{3+}, Cu^{2+}, Cd^{2+}, and also of ‹the arsenic group›, As^{3+}, As^{5+}, Sb^{3+}, Sb^{5+}, Sn^{2+}, Sn^{4+}, Pt^{2+}, Pt^{4+}, Au^{+}, and Au^{3+}, are precipitated. Advantage is taken of these relations in the following way, in systematic analysis: after the separation of the silver group, by precipitation of the difficultly soluble chlorides, hydrogen sulphide, in the presence of an excess of acid, is used to precipitate the sulphides of the ions of the copper and the arsenic groups, the two groups being precipitated together. Hydrogen sulphide, under these conditions, does not precipitate any sulphides of the zinc group or those of any of the remaining groups. Hydrogen sulphide is used, in this way, as one of the most valuable reagents in analytical work, enabling the analyst to separate whole groups of metal ions from other groups. There is also no other agent, equally important, which is more likely to be used in a wrong way and to lead to error.

«The Ionization of Hydrogen Sulphide.»—Before taking up the theory of the separation of these groups by precipitation with hydrogen sulphide in the presence of a strong acid, a discussion of some of the characteristics of the reagent will be in place. Hydrogen sulphide, like carbonic acid and other dibasic acids, ionizes in two stages; it produces first hydrogen-ion and hydrosulphide-ion, HS^{−}, and this ion in turn dissociates, producing hydrogen-ion and sulphide-ion, S^{2−}. [p200]

For the first dissociation, HSH ⇄ H^{+} + HS^{−}, we have

[H^{+}] × [HS^{−}] / [H_{2}S] = K_{1} (I)

The value of the constant[395] for this primary ionization of hydrogen sulphide is 0.91E−7. It is apparent that hydrogen sulphide, even in the primary ionization, is a very weak acid and produces a very small concentration of hydrosulphide-ion. In a solution, saturated at 25°, the total concentration of hydrogen sulphide is approximately 0.1 molar, and the concentration of hydrosulphide-ion, therefore, in the absence of any foreign acid, at most[396] 0.95E−4. The concentration of the dissolved, nonionized hydrogen sulphide, [H_{2}S], is practically a constant, if solutions saturated with hydrogen sulphide under a given pressure, say under atmospheric pressure, are considered. For such solutions, then, we may put more simply[396]

[H^{+}] × [HS^{−}] = ‹k› = (0.95E−4)^2 = 0.9E−8. (II)

The concentration of hydrosulphide-ion is, therefore, inversely proportional to the concentration of hydrogen-ion. It is clear that the addition of a strong acid, readily yielding concentrations of hydrogen-ion very much greater than 0.95E−4 (the value of [H^{+}] in a saturated aqueous solution of hydrogen sulphide) will, as the result of the greatly increased total hydrogen-ion concentration, reduce the concentration of hydrosulphide-ion to correspondingly low values. For instance, the presence of 0.1 molar hydrochloric acid will increase the concentration of hydrogen-ion close to a thousandfold and will reduce the concentration of hydrosulphide-ion to 0.9E−7.

For the secondary ionization (see p. 101) of hydrogen sulphide, HS^{−} ⇄ H^{+} + S^{2−}, we have

[H^{+}] × [S^{2−}] / [HS^{−}] = K_{2}. (III)

The value of this constant has recently been determined[397] and found to be 1.2E−15. Recalling the fact that the concentrations [H^{+}] and [HS^{−}] of hydrogen-ion and hydrosulphide-ion, [p201] respectively, resulting from the primary ionization, are each[398] 0.95E−4, we have for the concentration of the sulphide-ion, in aqueous solution saturated with hydrogen sulphide at atmospheric pressure and 25°, [S^{2−}] = 1.2E−15.

Combining equations (II) and (III), we have, further:[399]

[H^{+}]^2 × [S^{2−}] = ‹k› × K_{2} = ‹k›_{2} = 1.1E−23, (IV)

‹which shows, directly, the relation between the concentration of the sulphide-ion and that of the hydrogen-ion, the relation of primary importance in considering the precipitation of metal sulphides in acid solutions.› The ‹concentration of the sulphide-ion› is, thus, ‹inversely proportional› to the ‹square› of the ‹concentration› of ‹the hydrogen-ion›. A thousandfold increase in the concentration of the latter, which is very nearly the effect produced by the presence of 0.1 molar hydrochloric acid ([H^{+}] = 0.091), reduces the concentration of sulphide-ion in the saturated aqueous solution a millionfold: If we call [S^{2−}]_{Ac.} the concentration of the sulphide-ion in the acid solution, [S^{2−}]_{Ac.} = (1.1E−23) / (0.091)^2 = 1.3E−21, whereas, in the absence of acid, as found above, [S^{2−}] = 1.2E−15.

On the other hand, the addition of alkali to hydrogen sulphide, by neutralizing and suppressing the hydrogen-ion, and by forming the ‹salts› MeSH and Me_{2}S, will very greatly increase the concentrations of the hydrosulphide-ion and of the sulphide-ion. Since the constant for the secondary ionization of hydrogen sulphide shows that HS^{−} is an ‹exceedingly weak› acid, its salts, Me_{2}S, are very largely hydrolyzed, the constant for water being somewhat greater[400] than its own. According to Knox, in a 0.1 molar solution of Na_{2}S about 99% of the sulphide is hydrolyzed: Na_{2}S + H_{2}O ⥂ NaSH + NaOH. In spite of this almost complete hydrolysis, sufficient sodium sulphide remains in a solution of this [p202] substance, to yield a concentration of the sulphide-ion that is far greater than that obtained from a solution of hydrogen sulphide. In a 0.1 molar solution of sodium sulphide the concentration of the sulphide-ion is, approximately, [S^{2−}]_{alk.} = 0.9E−3, as compared with [S^{2−}] = 1.2E−15 in a saturated solution of hydrogen sulphide (25°, 760 mm.), and with 1.3E−21 in the same solution in the presence of 0.1 molar hydrochloric acid.

Ammonium sulphide (NH_{4})_{2}S is the salt of an extremely weak acid with a much weaker base than sodium hydroxide, and it is correspondingly more completely decomposed by water. In a 0.1 molar solution of the sulphide (NH_{4})_{2}S, we find the approximate concentration[401] of the sulphide-ion [S^{2−}]_{am.} = 1.8E−6, as compared with 0.9E−3 in a similar solution of Na_{2}S. But the concentration of sulphide-ion is still enormously greater than its concentrations in hydrogen sulphide in the absence and in the presence of acids (see above).

The following table[402] contains a ‹summary of the concentrations of sulphide-ion› in the various solutions discussed, as well as its concentration in the presence of 0.2 molar hydrochloric acid. In the separation of the copper and arsenic groups from the zinc and aluminium groups, a concentration of hydrogen-ion corresponding to the presence of 0.15 to 0.25 molar hydrochloric acid is satisfactory for an accurate separation for ordinary purposes.

Solution. [S^{2−}] 0.1 molar Na_{2}S: 0.9E−3 0.1 molar NaSH: 0.8E−5 0.1 molar (NH_{4})_{2}S: 1.8E−6 0.1 molar (NH_{4})SH: 1.4E−8 0.1 molar H_{2}S, sat. aq. sol., 25°: 1.2E−15 0.1 molar H_{2}S, 0.1 molar HCl:[403] 1.3E−21 0.1 molar H_{2}S, 0.2 molar HCl: 3.5E−22

[p203]

«Precipitation of Sulphides by Hydrogen Sulphide.»—Although the hydrosulphide-ion most likely takes part in the precipitation of sulphides, the main effect appears to be due to the sulphide-ion.[404] The sulphides seem to be both more stable and less soluble, than the hydrosulphides. It is very likely that hydrosulphides, mixed with the sulphides, are precipitated to a certain extent, but they are unstable, lose hydrogen sulphide and go over into sulphides much more readily than hydroxides, as a rule, change into oxides.[405] The final condition of equilibrium, if one waited, in a given case, until equilibrium was established, would depend, under these conditions, rather on the concentration of the sulphide-ion than on that of the hydrosulphide-ion, although the latter is present in much greater concentrations than the former. For the sake of simplifying the discussion of the theory of the separation of metal ions by precipitation with hydrogen sulphide, the discussion will be limited to the consideration of the sulphide-ion as the active precipitating agent. As stated, this seems to be in accordance with the known facts.

The absolute values of the solubilities of the various sulphides, which are involved in the discussion, are known, with any degree of accuracy, only in a few cases. The aim of the discussion will be, therefore, to develop, rather, the relations in the values involved, which may be readily determined. Wherever absolute quantities can be given, they also will be referred to.

«Theory of the Separation of Sulphides by Precipitation with Hydrogen Sulphide. I. Precipitation of Ferrous Sulphide.»—If we prepare a solution of ferrous sulphate, containing 27.8 grams of the salt, FeSO_{4}, 7 H_{2}O, in one liter (0.1 molar), we may call [Fe^{2+}] the concentration of the ferrous-ion in the solution. If such a solution is saturated with hydrogen sulphide, under atmospheric pressure, (‹exp.›), ferrous sulphide is not precipitated. We would decide, on the basis of the principle of the solubility-product, that the reason no precipitate of ferrous sulphide is formed is, that the product of the ion concentrations is ‹smaller› than the [p204] solubility-product constant characteristic of the sulphide (p. 151); [Fe^{2+}] × [S^{2−}] < K_{FeS}, in which [S^{2−}] is the concentration of sulphide-ion in the solution thus saturated with hydrogen sulphide. If an alkali, sodium or ammonium hydroxide, is added to the solution, a heavy precipitate of ferrous sulphide is immediately formed (‹exp.›). The salts of hydrogen sulphide, like all common salts, as we have just seen, are very much more highly ionized than is hydrogen sulphide itself, and the addition of the alkali has the effect of increasing enormously[406] the concentration of sulphide-ion, say to ‹x›[S^{2−}] (p. 202). Under these conditions, the product of the ion concentrations has evidently grown larger than the constant: [Fe^{2+}] × ‹x›[S^{2−}] > K_{FeS}.

«II. Precipitation of Zinc Sulphide.»—Now, if a solution of zinc sulphate is prepared so as to contain 28.7 grams of ZnSO_{4}, 7 H_{2}O, per liter, which would make it of the same molar concentration as the ferrous sulphate solution, we may, in view of the fact that analogous salts ionize approximately to the same extent in solutions of the same concentration, consider the concentration, [Zn^{2+}], of zinc-ion to be practically the same as the concentration, [Fe^{2+}], of the ferrous-ion in the ferrous sulphate solution. We may put [Zn^{2+}] = [Fe^{2+}].

If the solution of zinc sulphate is saturated with hydrogen sulphide, under the same conditions as were used with the ferrous salt solution (‹exp.›), or if we add the zinc sulphate solution to the mixture of ferrous sulphate and hydrogen sulphide (‹exp.›), we immediately obtain heavy white precipitates of zinc sulphide. We would decide, therefore, on the basis of the principle of the solubility-product, that in this case [Zn^{2+}] × [S^{2−}] > K_{ZnS}. Since we have used hydrogen sulphide under practically the same conditions, we may consider that [S^{2−}], in this experiment,[407] is the same as in the [p205] test with ferrous sulphate, and, by the conditions of the experiment, we have also made [Zn^{2+}] = [Fe^{2+}]. The two factors of the product are, therefore, the same, for the first moment, and we may put [Fe^{2+}] × [S^{2−}] = [Zn^{2+}] × [S^{2−}] = P.

Since P is smaller than K_{FeS} and larger than K_{ZnS}, it is clear that the ‹solubility-product constant for zinc sulphide must be smaller than that for ferrous sulphide›. The solubility-product constants, for similar salts, are a measure of their solubilities in water. We may obtain their values by determining the solubilities of salts in pure water, whenever the solubility is not affected by other chemical changes. In the present instance, the quantitative measurements, that have been made in this way, are open to question, owing to the considerable hydrolysis which sulphides, as salts of a very weak acid, undergo in solutions of such extreme dilution.[408] Until such relations have been taken into account quantitatively, it is better to limit ourselves for the present to the more accessible question of ‹relative› solubility.

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: ZnSO_{4} + H_{2}S → ZnS ↓ + H_{2}SO_{4}. 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 × [S^{2−}] 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 [S^{2−}] 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 K_{ZnS}. 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: ([Zn^{2+}] × [S^{2−}] / ‹x›) < K_{ZnS}, the new concentration of the sulphide-ion being represented by the symbol [S^{2−}] / ‹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(CH_{3}CO_{2})_{2} + H_{2}S ⥂ ZnS ↓ + 2 CH_{3}COOH. 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.»—Now, if a 0.1 molar solution of cadmium sulphate (22.6 grams of the salt,[414] CdSO_{4}, H_{2}O, ‹per› liter) is prepared, equivalent, in concentration, to the solutions of ferrous and zinc sulphates used previously, we may consider that the concentration of cadmium-ion is, approximately, the same as the concentrations of the ferrous-ion and zinc-ion in the solutions of their sulphates. If the solution of cadmium sulphate is added to the acidified solution of zinc sulphate, which was saturated with hydrogen sulphide but from which no zinc sulphide could be precipitated (‹exp.›, p. 207), a heavy precipitate of cadmium sulphide is at once obtained (‹exp.›). Or, if 2 c.c. of hexamolar hydrochloric acid is added to 50 c.c. of the cadmium sulphate solution and the mixture is saturated with hydrogen sulphide,[415] cadmium sulphide is precipitated readily (‹exp.›) and, in fact, quantitatively. We would decide, on the basis of the principle of the solubility-product, that, in this mixture, precipitation results because ([Cd^{2+}] × [S^{2−}] / ‹x›) > K_{CdS}.

By the conditions of the experiment we started with a concentration of the cadmium-ion, [Cd^{2+}], equal to the concentration, [Zn^{2+}], of the zinc-ion in the zinc sulphate solution, and with the same concentration,[415] [S^{2−}] / ‹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 ([Cd^{2+}] × [S^{2−}] / ‹x›) = ([Zn^{2+}] × [S^{2−}] / ‹x›) = P′. We recall the fact, that we have already concluded, on the basis of the principle of the solubility-product, that ([Zn^{2+}] × [S^{2−}] / ‹x›), or P′, is ‹smaller› than K_{ZnS} (p. 207), and that ([Cd^{2+}] × [S^{2−}] / ‹x›), or P′, is ‹greater› than K_{CdS}.

P′ being smaller than K_{ZnS} and larger than K_{CdS}, it is clear that [p210] K_{CdS} is smaller than K_{ZnS} 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 K_{CdS} of cadmium sulphide must be smaller than the constant K_{ZnS} 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 K_{CdS} is ‹sufficiently small› (CdS is sufficiently insoluble) ‹to make the product of the ion concentrations› [Cd^{2+}] × [S^{2−}] / ‹x›, ‹in spite of the extremely small value of› [S^{2−}] / ‹x›, ‹greater than the constant› K_{CdS}, whereas the same small value of [S^{2−}] / ‹x› makes it impossible for the product [Zn^{2+}] × [S^{2−}] / ‹x› to reach the value of the ‹larger constant› K_{ZnS}, 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.

«The Separation of the Copper and Arsenic Groups from the Zinc Group.»—Ferrous-ion and zinc-ion may be taken as typical representatives of the ions of the ‹zinc group, whose sulphides are not precipitated by hydrogen sulphide in the presence of a definite concentration of hydrogen-ion›, cadmium-ion as a representative of the ‹copper›, ‹silver› and ‹arsenic› groups, ‹whose sulphides are precipitated under the same conditions›. The separation of the groups depends, therefore, on the different solubilities of the sulphides of these [p211] groups, the sulphides of the zinc group being the most soluble. Since even these sulphides are precipitated quantitatively by ammonium sulphide and are very difficultly soluble, the separation is a kind of fractional precipitation of difficultly soluble salts, in which the fractionation is made possible and convenient by the use of an agent, hydrogen sulphide, the concentration of whose active precipitating component, the sulphide-ion, S^{2−}, is easily regulated and readily made sufficiently small, not to precipitate even the difficultly soluble sulphides of the iron group.

Solubilities vary from salt to salt, and we have already found that, in the zinc group, zinc sulphide is less soluble than the sulphide of a second member of the group, ferrous sulphide, and that the difference is revealed in a somewhat different behavior of their salts toward hydrogen sulphide, when the action is studied in some detail. Similar differences must be expected to exist among the sulphides of the groups that hydrogen sulphide precipitates even in the presence of an excess of hydrochloric acid. As these differences are the sources of some of the most common and most serious errors which analysts are liable to commit, the detailed study of the action of hydrogen sulphide must be continued a little further.

«The Effect of a Large Excess of Acid.»—The precipitation of cadmium sulphide depends on the relation of the product of the concentrations of the cadmium-ion and the sulphide-ion to the solubility-product constant for cadmium sulphide (see the equation, p. 209). Now, it is clear that if a ‹larger› and ‹larger excess of hydrogen-ion› is introduced by the addition of more concentrated hydrochloric acid to the cadmium sulphate solution, the concentration of sulphide-ion is correspondingly ‹reduced›.[416] The point might be reached, where the sulphide-ion factor becomes so small, that the product of the ion concentrations remains smaller than the value K_{CdS}, required for precipitation of cadmium sulphide.

In fact, if a large excess[417] of hydrochloric acid is added to 50 c.c. [p212] of the 0.1 molar solution of cadmium sulphate, hydrogen sulphide fails to precipitate any of the sulphide (‹exp.›).

But, if a few cubic centimeters of a 0.1 molar solution of cupric sulphate (25.0 grams of CuSO_{4}, 5 H_{2}O, per liter) are added to the solution from which hydrogen sulphide fails to precipitate cadmium sulphide, cupric sulphide is at once precipitated. And, if 15 c.c. of concentrated hydrochloric acid are added to 50 c.c. of the 0.1 molar cupric sulphate solution, there results a mixture corresponding to the cadmium sulphate solution from which hydrogen sulphide fails to precipitate CdS; we find that hydrogen sulphide will ‹precipitate› the ‹sulphide of copper› very readily, even under these adverse conditions (‹exp.›). Cupric sulphide must be even less soluble in water than cadmium sulphide,[418] and there is no difficulty in showing that such is the case. If ammonium sulphide, or hydrogen sulphide, is gradually introduced into a mixture of 25 c.c. each of the 0.1 molar sulphate solutions, cupric sulphide is precipitated first, and, if the precipitate is collected in fractions, [p213] pure yellow cadmium sulphide is precipitated last.[419] Or, if 25 c.c. of 0.1 molar cupric sulphate is added to the mixture in which a precipitate of cadmium sulphide displaced the more soluble zinc sulphide (p. 210), the yellow sulphide will, in turn, give way to the less soluble black sulphide of copper (‹exp.›).

We find thus that the precipitation of cadmium sulphide, by hydrogen sulphide in acid solution, ‹can be prevented by the presence of an excess of hydrochloric acid›, which does not prevent the precipitation of the less soluble cupric sulphide.[420] The fact, then, that, in an analysis of some unknown mixture, hydrogen sulphide produces a precipitate in acid solution, must not be considered as evidence that the conditions are such as to insure the precipitation of all the sulphides of the groups, which we intend to precipitate. To avoid error, conditions must be such as to insure the complete precipitation of the more soluble as well as the less soluble sulphides. The sulphides of ‹cadmium› and ‹lead›, in particular, and, to a lesser degree, the sulphides of antimony and tin, are most liable to remain unprecipitated and thus escape detection in systematic analysis. This is a matter of special importance, also, in detecting traces of the ions of these metals, especially of lead, which is a slow cumulative poison, even when absorbed in minute amounts, and which analysts must therefore be able to detect, even in traces, with absolute certainty. It is clear, from a consideration of the product of the ion concentrations, as affecting the precipitation or nonprecipitation of such a sulphide, that a much smaller excess of acid will prevent the precipitation of the last traces of lead sulphide, and, therefore, of all of it, if only traces are present, than will interfere with the precipitation of the sulphide in bulk.

«The Desirable Concentration of Acid (of Hydrogen-ion) and an Indicator for Correct Acidification.»—Summarizing the conclusions reached in regard to the conditions necessary for a successful separation of the copper and arsenic groups, by means of hydrogen sulphide, from the zinc and aluminium groups, we find [p214] that the concentration of the hydrogen-ion, in the solution to be treated with hydrogen sulphide, is the most important factor. Too great a concentration, as has just been shown, will prevent the precipitation of all, or part, of the more soluble sulphides of the former groups, notably of the sulphides of cadmium and lead, which is a common error in the laboratory. Too small a concentration, which may result when a salt of a weak acid, such as an acetate, borate or phosphate is present, may lead, as was shown above, to the precipitation of part of the zinc group, notably of zinc and possibly of nickel and cobalt, with the copper and arsenic groups, and thus lead to other errors. For the ordinary purposes of analysis, requiring the precipitation of say one milligram of any ion from 100 c.c. solution (one-tenth per cent, if one gram of substance is used for analysis), a concentration of hydrogen-ion of 0.1 to 0.3 gram-ion per liter forms a good basis for work.[421] The presence of this concentration of hydrogen-ion, irrespective of the possible presence of weak organic or inorganic acids, may be readily insured by a simple test with an appropriate indicator. ‹Methyl-violet›[422] is suitable for such a purpose, because it is sensitive only to the rather high concentrations of hydrogen-ion required: 0.1 c.c. or two drops of 0.05 to 0.1 molar hydrochloric acid, added to an equal volume of a very dilute solution (1 : 12,500) of the indicator, changes its color to a ‹pure blue›; 0.2 to 0.25 molar acid, used similarly, turns the indicator to a ‹blue-green› tint, and 0.33 molar acid produces a ‹yellow› or ‹yellow-green› hue. The ‹blue-green tint›, with which one becomes easily familiar, and which can, indeed, always be prepared for matching tints, may be used to recognize speedily, and with sufficient accuracy, a concentration of the hydrogen-ion of the strength desired, irrespective of its source.[423]

If an analyst aims to find even smaller quantities of a particular metal ion, ‹e.g.› traces of lead, the ordinary method of analysis [p215] may be modified, the source of error in the precipitation of traces of lead sulphide being kept in mind.[424]

Besides the complications mentioned, and provided against in the way discussed, there is still one more complication in the use of hydrogen sulphide: this is in the matter of the precipitation of ‹arsenic› sulphide from solutions containing ‹arsenic in the pentavalent condition›. Since the interpretation of this complication and the explanation of the methods for avoiding the errors, which may arise therefrom, are necessarily intimately connected with the chemical behavior of arsenic acid, this subject will be considered in the discussion of the arsenic group (Chapter XIII).

FOOTNOTES:

[395] Auerbach, ‹Z. phys. Chem.›, «49», 220 (1904).

[396] See footnote 1, p. 201.

[397] Knox (in Abegg's laboratory), ‹Trans. Faraday Soc.›, «4», 44 (1908).

[398] The concentration of the hydrogen-ion is really a little greater than that of the hydrosulphide-ion, as a result of the ionization of the latter, but the amount of hydrogen-ion formed in this way (about 1E−15) is so minute, compared with that formed by the primary ionization, that it is negligible.

[399] We can obtain the relation, directly, from H_{2}S ⇄ 2 H^{+} + S^{2−} and [H^{+}]^2 × [S^{2−}] / [H_{2}S] = K = 1.1E−22. For a given pressure of the hydrogen sulphide, [H_{2}S], expressing its solubility (about 0.1 molar at 25°), is constant, and therefore [H^{+}]^2 × [S^{2−}] = a constant, as given in equation (IV). Putting [H_{2}S] = 0.1, we have [H^{+}]^2 × [S^{2−}] = 0.1 × 1.1E−22 = 1.1E−23.

[400] On account of the great mass of water, we compare (see equation, p. 176) [H^{+}] × [HO^{−}] = 1.2E−14 (at 25°) with [H^{+}] × [S^{2−}] / [HS^{−}] = 1.2E−15.

[401] The calculation was made by the method used by Knox (‹loc. cit.›) for a molar solution. The degree of ionization of the salt was not considered and the correct ionization constant for ammonium hydroxide was used, 1.8E−5 in place of 2.3E−5. The latter, evidently, was used by Knox as the result of overlooking a correction, which Bredig made in his (Bredig's) first calculations of the constant; ‹cf.› Bredig, ‹Z. phys. Chem.›, «13», 293, footnote. The same erroneous constant is found in Kohlrausch and Holborn, ‹loc. cit.›, p. 194.

[402] For further values and for the method of calculation, see Knox, ‹loc. cit.›

[403] [S_{2−}] = 1.1E−23 / [H^{+}]^2, according to equation (IV), p. 201.

[404] Knox's work leads to that conclusion.

[405] The precipitation of sulphides, from a solution containing much more of the hydrosulphide-ion than of the sulphide-ion, is comparable with the precipitation of mercuric oxide, HgO, and of silver oxide, Ag_{2}O, by sodium or potassium hydroxide.

[406] On account of the presence of a small, unknown amount of sulphuric acid in the original solution, resulting from the hydrolysis of ferrous sulphate, the exact value of [S^{2−}] in the first solution cannot be calculated without further examination; but, according to the values given in the table on page 202, the value of ‹x›, indicating the growth in the concentration of S^{2−}, is ‹at least› 10^{12}, if 2 equivalents of NaOH, and 10^9, if 2 equivalents of NH_{4}OH are used to convert the 0.1 molar hydrogen sulphide into the corresponding sulphide Me_{2}S, of 0.1 molar concentration.

[407] [S^{2−}] is exactly the same in the two products, when equal volumes of the zinc and ferrous sulphate solutions are mixed and the mixture is saturated with hydrogen sulphide; ‹zinc sulphide› is precipitated.

[408] The difference in the values obtained, when hydrolysis is considered or neglected, is very considerable. ‹Vide› Bodländer, on the solubility of calcium carbonate, ‹Z. phys. Chem.›, «35», 23 (1900), and Stieglitz, ‹Carnegie Institution Publications›, No. «107», 249 (1909).

[409] In carrying out this ‹fractional precipitation› a very dilute solution of ammonium sulphide is used, so as to prevent the mechanical enclosure of black ferrous sulphide, which would discolor the white sulphide. The ammonium sulphide solution should be saturated with hydrogen sulphide, to prevent the precipitation of green ferrous-ferric oxide by an excess of free ammonia. It is best to prepare a set of the precipitates and to preserve them in well-stoppered vessels, and not to try to take the time and care necessary to effect a perfect fractionation as a lecture experiment. The presence of the ferrous sulphate, in the supernatant liquid above the first precipitate of zinc sulphide, may be readily demonstrated by pouring off some of the solution and adding an excess of ammonium sulphide to it. Of course, it is also perfectly legitimate, and easier, to precipitate first zinc sulphide from a pure zinc sulphate solution and then to add ferrous sulphate solution to the mixture and to preserve the mixture. If the zinc sulphide were not the less soluble, it would be rapidly converted into the black ferrous sulphide. (See p. 165, and see below, pp. 210, 213, where similar transformations are carried out as lecture experiments.)

[410] When 10% of the zinc in a 0.1 molar solution has been precipitated, 0.01 molar sulphuric acid has been formed. For the sake of a rough approximation, the acid may be considered completely ionized and then [H^{+}] = 0.02, which is 200 times the value of [H^{+}] in a saturated H_{2}S solution (p. 200); if the presence of a little sulphuric acid in the original zinc sulphate solution, resulting from a slight hydrolysis of the salt, is ignored, the concentration of the sulphide-ion is decreased roughly (200)^2 or forty thousandfold, while the concentration of zinc-ion falls 10%. The corrections, that have been indicated, would change the quantities involved, but they would not modify the character of the result.

[411] This proportion of acid, making the concentration of the hydrogen-ion, approximately, [H^{+}] = 0.2, is used, not because it represents the minimum concentration of the hydrogen-ion, which will prevent the precipitation of zinc sulphide in 0.1 molar zinc sulphate solution, but because it represents the practical conditions under which the precipitation of zinc sulphide is avoided, when the copper and arsenic groups are precipitated in qualitative analysis (see p. 213).

[412] Nickel and cobalt sulphides are also precipitated by hydrogen sulphide in the presence of free acetic acid, if sodium or potassium acetate is added, to suppress the hydrogen-ion of the acetic acid (p. 112).

[413] They would be found in the copper group.

[414] The sulphate, of this composition, is obtained by drying the crystallized sulphate in an air bath at 100–105°.

[415] ‹Cf.› the corresponding experiment with zinc sulphate, p. 207.

[416] [S^{2−}] = ‹k› / [H^{+}]^2. See equation (IV), p. 201.

[417] Any immediate precipitation of cadmium sulphide will be prevented by the addition of 10 c.c. of concentrated acid (sp. gr. 1.19) to 50 c.c. of the 0.1 molar solution, and 15 c.c. will completely prevent any precipitation of the sulphide. Of course, a smaller excess would prevent the precipitation of small quantities of the sulphide (‹e.g.› a half milligram of cadmium), which should easily be found in 50 c.c. (see p. 214).

[418] The value of the solubility-product constant for cupric sulphide, at 25°, was determined by Knox (‹loc. cit.›): [Cu^{2+}] × [S^{2−}] = 1.2E−42, corresponding to a concentration of 1.1E−21 of cupric-ion. Mercuric sulphide was found even less soluble: [Hg^{2+}] × [S^{2−}] = 2.8E−54, and its behavior agrees with such a relation (Lab. Manual, p. 50, § 2). The solubility-product constant for lead sulphide, which resembles cadmium sulphide in the fact that a large excess of acid prevents its precipitation, was found to be [Pb^{2+}] × [S^{2−}] = 2.6E−15, the constant being about 10^{27} times as large as the constant for cupric sulphide. This value for the solubility-product constant for lead sulphide must either be considerably larger than the true value or lead must be easily precipitated as a hydrosulphide, Pb(SH)_{2}, since solutions in which the product of the ion concentrations, [Pb^{2+}] × [S^{2−}], is very much smaller than the constant given, readily precipitate lead sulphide. Thus Noyes and Bray [‹J. Am. Chem. Soc.›, «29», 137 (1907)] report it possible to precipitate 1 to 2 milligrams of lead-ion in 100 c.c. of solution (say [Pb^{2+}] = 1E−4) with hydrogen sulphide in the presence of 4 c.c. of hydrochloric acid (sp. gr. 1.12), for which, approximately, [H^{+}] = 0.25. Then (equation (IV), p. 201) [S^{2−}] = (1.1E−23) / (0.25)^2 = 1.8E−22, and [Pb^{2+}] × [S^{2−}] = 1E−4 × 1.8E−22 = 1.8E−26, which is a much lower value than that given by Knox, and which still is not claimed to represent the limit of insolubility. Experiments, made in this laboratory, confirm this result and show further, that lead-ion in a concentration of 1E−5 is precipitated in the presence of 0.25 molar hydrochloric acid ([H^{+}] = 0.22). Then [S^{2−}] = 2.3E−22 and [Pb^{2+}] × [S^{2−}] = 2.3E−22 × 10^{−5} = 2E−27, which does not yet express the limit of insolubility.

[419] The fractions are not prepared in the lecture, but the first fraction is kept suspended in part of the solution of the two sulphates and may be kept so for years. The last fraction is kept in a separate container.

[420] A large excess of acid is liable to interfere with the precipitation of the ‹last traces› of cupric sulphide and is avoided in exact work.

[421] Noyes and Bray use, approximately, [H^{+}] = 0.25 [‹J. Am. Chem. Soc.›, «29», 137 (1907)]. Tests in this laboratory showed that 1 milligram of cadmium-ion, or of lead-ion, in 100 c.c., is readily precipitated by hydrogen sulphide in the presence of 0.25 molar hydrochloric acid, ([H^{+}] = 0.22).

[422] Kahlbaum's "Krystallviolett," [(CH_{3})_{2}NC_{6}H_{4}]_{2}C : C_{6}H_{4}N(CH_{3})_{2}Cl, is referred to.

[423] An indelible ink pencil (violet) may, in most cases, be used in place of the solution. The details for the application of the indicator are given in the instructions for laboratory practice, Lab. Manual, pp. 31, 102, 103.

[424] See Blyth, ‹Poisons, etc.›, p. 608 (1895), in regard to the detection of traces of lead.

[p216]