Interpretation of Oxidation-Reduction Reactions in Terms of the Oxygen-Hydrogen Gas Cell.
Interpretation of Oxidation-Reduction Reactions in Terms of Direct Transfers of Electric Charges.
Arsenic Acid as an Oxidizing Agent.
We may recall the fact that a solution of potassium arseniate, to which dilute hydrochloric acid has been added, will remain clear for some time when the mixture is saturated with hydrogen sulphide (exp.). If a considerable excess of concentrated hydrochloric acid is added to this mixture, hydrogen sulphide immediately forms a dense precipitate (exp.) of arsenic pentasulphide—presumably through the union of quinquivalent arsenic-ion with the sulphide-ion: 2 As5+ + 5 S2− ⇄ As2S5 ↓ (see p. [247]). This behavior suggested that arsenic acid, although a moderately strong acid, might nevertheless be somewhat amphoteric, might have slight basic properties, as well as its ordinary acid functions. The relation is expressed in the equations:[565]
3 H+ + AsO43− ⇄ (HO)3AsO
(HO)3AsO + HOH ⇄ As(OH)5 ⥃ As5+ + 5 HO−.
Since oxidations by arsenic acid involve its reduction to arsenious acid, containing trivalent,[566] in place of quinquivalent arsenic, one might well suspect, that the oxidizing component is the quinquivalent arsenic-ion, As5+, the discharge of two of whose positive charges would cause oxidation (e.g. of iodide-ion), exactly as the discharge of positive charges at the positive pole of an electric current causes oxidation (p. [252]): As5+ + 2 I− ⥂ As3+ + I2. [p285]
In a solution of potassium arseniate, we would have only the faintest trace of the ion As5+, since the addition of an alkali to the system, expressed in the above equations, would carry the reversible reactions towards the left. The addition of dilute hydrochloric acid to the system must carry the reactions towards the right and increase the concentration of As5+; the addition of concentrated acid must increase the concentration of As5+ very much more. Even if the concentration of As5+ remained minute, the oxidizing power would be increased proportionally to the ratio of the concentrations in the first and the last solutions. A millionfold increase in concentration, even when we are dealing with very small numbers, would imply a millionfold increase in the activity of the solution. If, then, the oxidizing component of arsenic acid is the quinquivalent ion, As5+, which would tend to discharge two of its positive (oxidizing) charges, arsenic acid should be a much more powerful oxidizing agent in strong acid solution than in alkaline or neutral solutions.
We thus arrive at the conclusion that the addition of hydrochloric acid to a mixture of arseniate and iodide may be effective, in bringing about the reduction of the arseniate and the oxidation of the iodide, primarily because of its action on arsenic acid, perhaps by facilitating its ionization as a base, and that it is not effective through any action on the iodide, for instance by producing free hydroiodic acid, as is often assumed. This conclusion may easily be tested with the aid of the chemometer (see p. [253]): potassium arseniate against potassium iodide gives only the faintest possible current, barely perceptible with the aid of a very sensitive voltmeter.[567] The addition of hydrochloric acid to the beaker containing the potassium iodide does not increase the potential (it rather decreases it somewhat), whereas the addition of the concentrated acid to the potassium arseniate solution produces a most decided increase in the potential[568] (exp.). It is evident, therefore, that the addition of the acid is primarily and directly intended to increase the oxidizing power of the arsenic acid, rather than to increase the reducing power of the iodide. [p286]
The more common methods of expressing oxidation-reduction reactions of this type are illustrated in the following equations:
Na3AsO4 + 2 HI ⇄ Na3AsO3 + H2O + I2