(4) If one of the ions of acetic acid is suppressed by the addition of some agent, equilibrium is again destroyed and the resulting change is always in the direction of reëstablishing a condition of equilibrium on the basis of the law of equilibrium. Instances of this common case, such as the neutralization of acetic or any other acid by a base, or the driving out of a weak acid, from its salts, by a strong acid, or of a weak base by a strong base, considered from the point of view of the equilibrium law, should be worked out by the student.[221] [p115]
(5) The displacement of an acid (or base) in salts by another acid (or base) is subject always to the law that equilibrium is reached, when all the equilibrium constants are satisfied in the system. Very frequently, the application of the law will lead, apparently, to the incorrect conclusion that the stronger acid (or base) will always displace, more or less completely, the weaker[222]—an inference, out of which grew, indeed, the characterization of acids and bases as strong and weak. Yet, when the laws of equilibrium, as the result of the peculiar values of constants, demand that, on the contrary, a strong acid or base should be displaced by a weak, or even a feeble one, we find that the change in this direction occurs with equal ease. Numerous instances will be given where a weak acid (or a weak base) does this to a certain extent (Chap. X), and others where precipitation of salts facilitates the action of weaker acids greatly by the introduction of new, physical constants. The following case of the liberation of hydrochloric acid by the exceedingly weak acid, hydrocyanic acid, is important because it shows a reversal of the common action without the formation of any precipitate, and especially because it brings out most strikingly the relations between ionization and chemical activity in a case of special importance to analytical chemists.
The Exceptional Ionization of Mercuric Cyanide and Its Consequences.
Exp. Into the parallel tubes of the conductivity apparatus (p. [77]) equivalent quantities[223] of solutions of mercuric cyanide[224] Hg(CN)2, mercuric chloride [p116] HgCl2, and barium chloride BaCl2, are introduced. The barium chloride represents an ordinary salt of the same type as the mercury salts, and the current passing through its solution makes the little lamp glow. The electrodes in the mercuric chloride solution must be brought quite close together before sufficient current will pass through the solution to bring its little lamp to redness. In the case of the cyanide we can, at most, get a faint, dull glow by bringing the electrodes together as closely as we can, without allowing them to touch each other and short-circuit the current.
It is evident that these mercury salts are not as readily ionizable as are ordinary salts. This difference, as may be anticipated, shows itself also in the chemical behavior of their solutions and makes necessary special precautions on the part of the analyst in examining mercury compounds. For instance, whereas mercuric oxide is readily precipitated by the addition of sodium hydroxide to solutions of the nitrate and even of the moderately ionized chloride, one fails to get a precipitate of oxide from the cyanide solution (exp.), and if we relied on this test, we should overlook the mercury entirely. That traces of the mercuric-ion are present, as indicated by the minimal conductivity of the cyanide, is confirmed by the fact that from the cyanide solution the sulphide, which is much less soluble than is the oxide, may be precipitated by the addition of ammonium sulphide (exp.). That the sulphide is in fact less soluble[225] than the oxide is shown by the conversion of the latter into the former by the action of ammonium sulphide (exp.).
As a further result of the abnormally slight ionization of these mercury salts, the analyst, unless he is on his guard, may also have difficulty in discovering the presence of their negative ions. Thus, while sodium chloride readily gives hydrogen chloride when treated with concentrated sulphuric acid, mercuric chloride, although it is a soluble salt, does not (exp.), and reliance on this test alone might lead to a gross error.[226] In the case of the cyanide, it is correspondingly difficult to recognize the [p117] cyanide-ion (see the laboratory experiment, Part III). The ordinary tests fail to show its presence until the mercury has been removed from the solution by precipitation as a sulphide. Mercuric cyanide being a deadly poison which analysts are liable to meet and have met with in criminal cases, it is clear that a knowledge of these facts is vital to analytical accuracy.
Now, the very slight ionization of mercuric cyanide enables us to realize, in the following experiment, the case where an exceedingly weak acid without the formation of any precipitate involving physical constants, may displace a much stronger acid from its salts. Hydrocyanic acid is one of the weakest acids (table, p. [104]), the constant for the ratio [H+] × [CN−] / [HNC] being 0.7E−9. It is so weak an acid that the addition of a dilute solution to methyl orange will not redden the indicator, but will have only a barely perceptible effect on it (exp.). Mercuric chloride solutions also are almost neutral to methyl orange (exp.) (very slight decomposition of the salt by water makes the solution very slightly acid, not enough to produce more than an orange color with methyl orange). Now, mercuric chloride, while it is not very easily ionized, is, we found, very much more readily ionized than is mercuric cyanide. The consequence is that when we add hydrocyanic acid to a mercuric chloride solution, the equilibrium between mercuric chloride and its ions and between hydrocyanic acid and its ions will be decidedly displaced, the mercuric-ion combining with the cyanide-ion to form the scarcely ionizable mercuric cyanide. As a result, more and more of the molecular mercuric chloride and hydrocyanic acid will be ionized; and since the other ions, the chloride and the hydrogen ions, form a readily ionizable electrolyte, hydrochloric acid, these ions (H+ and Cl−) will accumulate in the solution and we shall have sufficient ionized hydrochloric acid liberated to make the solution decidedly acid.
Exp. When the two solutions described above are mixed, a strongly acid solution, colored a bright pink, results.
In the following equations the [dark arrows] indicate the direction in which the action goes when the solutions are mixed:
HgCl2 ⥂ Hg2+ + 2 Cl−
2 HCN ⥂ 2 CN− + 2 H+
2 CN− + Hg2+ ⥂ Hg(CN)2
2 Cl− + 2 H+ ⇄ 2 HCl