Equally interesting are Kahlenberg's observations of interaction between hydrogen chloride in benzene solution and a similar solution of copper oleate. Each solution shows absence of appreciable conductivity, yet, when the solutions are mixed, precipitation of copper chloride occurs instantly. Whether we have here an instantaneous action between non-ionized molecules, as claimed by the observer, or whether the minimal ionization[157] of the hydrogen chloride and copper oleate, the existence of which we have a right to assume, is sufficient to account for this rapid action, both components being in solution and intimately mixed, is a question of the greatest interest. But until quantitative measurements of all the factors involved in chemical actions in benzene solution are obtained, a very difficult, but necessary task, which the discoverer of the action omitted to perform, no definite conclusion whatever can be based on such results, interesting as they are. Water, although it is only minimally ionized (tables, Chapter VI), hydrolyzes salts like potassium cyanide and aluminium chloride almost instantly, and it has been rigorously proved that the resulting condition of equilibrium involves the ions of water[158] (Chapter X). With the possibility that the well-known enormous speeds of action of ions may completely offset the tremendous reduction in concentration of the hydrogen-ion, in a benzene solution of hydrogen chloride as compared with an aqueous solution, further analysis of the relations is imperative.[159] Until such investigations have been carried out, we must consider [p087] it possible that the reactions of hydrogen chloride in benzene solution may be reactions of its non-ionized molecules or reactions of its ions. In view of the undoubted minute concentrations of the latter, as compared with aqueous solutions, and in view of the inertness of non-ionized hydrogen chloride in aqueous solutions as compared with the activity of its ions, a benzene solution of hydrogen chloride should show far less ionic activity than an aqueous solution, and that such is the case is brought out clearly by Kahlenberg's interesting experiments.
Some Applications of the Chemical Activity of Ions to Qualitative Analysis.
Perhaps the most instructive case of this kind, that we can study, is that of iron in ferrous and ferric salts. Exceedingly sensitive tests are known for the ferrous and the ferric ions. Thiocyanates produce an intensely red salt, Fe(SCN)3, when added, for instance, to ferric chloride; potassium ferrocyanide, K4Fe(CN)6, precipitates ferric ferrocyanide, Fe4[Fe(CN)6]3, Prussian blue, from ferric chloride solutions; ammonium hydroxide precipitates quantitatively the insoluble red ferric hydroxide (exps.). With ferrous salts, potassium ferricyanide K3Fe(CN)6 precipitates ferro-ferricyanide Fe3[Fe(CN)6]2, Turnbull's blue; ammonium sulphide precipitates black ferrous sulphide (exps.). Now, in two of the reagents used, potassium ferro- and ferricyanide, iron is present according to the formulæ given. If one should attempt to demonstrate its presence by means of these tests—among the most sensitive and most reliable tests known in analysis—one would fail utterly. Thiocyanates do not produce even the faintest tinge of pink in potassium ferricyanide solution[163]; ammonium hydroxide does not precipitate any ferric hydroxide (exps.). Ammonium sulphide does not precipitate the least trace of a black sulphide from a ferrocyanide solution, and when the latter is mixed with the ferricyanide solution, no trace, either of Prussian or Turnbull's blue, is shown (exps.). The contrast between the behavior of these salts and ferrous and ferric salts is now sharply and definitely interpreted, as being the result of the contrast in their ionization,—the color tests we use are extremely sensitive tests only for the ferric and ferrous ions, Fe3+ and Fe2+, respectively,—but potassium ferrocyanide ionizes into potassium ions and the negative ferrocyanide ions Fe(CN)64−, and shows the actions of ferrous ions as little as chlorate ions ClO3− exhibit the reactions of chloride ions Cl−. Potassium ferricyanide, in turn, gives rise to trivalent, negative ferricyanide ions Fe(CN)63− and not to ferric [p089] ions.[164] If any doubts arise on this point, one can decide the question readily by experiment. When a concentrated solution of potassium ferricyanide is placed in a U-tube under a solution of some colorless electrolyte, such as sodium sulphate, and plates connected with a battery are inserted, there is no difficulty (exp.) in seeing that the yellow ion,[165] containing the iron, moves to the positive pole and not to the negative. The iron is, therefore, as a matter of experiment, part of a negatively charged substance.
That iron is really present in these compounds can be shown most effectively if we destroy the salts:
Exp. Dry, pulverized potassium ferrocyanide is intimately mixed with dry potassium carbonate and the mixture heated in a hard glass test tube. When the whole mass has become red-hot, insuring complete decomposition, the hot (not red-hot) tube is plunged into water; the salts are extracted and particles of metallic iron are left undissolved. The action is
K4Fe(CN)6 + K2CO3 → 6 KCN + FeCO3
FeCO3 → FeO + CO2
and FeO + KCN → KCNO + Fe.
If the iron is dissolved in a little dilute hydrochloric acid and oxidized to the ferric condition, by the addition of a few drops of bromine water, the intensely red solution characteristic of ferric salts may be readily obtained, when a thiocyanate is added to the solution.
Chapter V Footnotes
[114] Loomis. (Cf. Whetham, Theory of Solution, p. 320 (1902).)
[115] Loomis; and E. H. Griffith. (Cf. Whetham, loc. cit.)