Self Reduction of hydrogen peroxide

There is a numerous class of chemical substances containing oxygen which, when placed in contact with hydrogen peroxide, cause the latter to change into water and oxygen whilst they themselves simultaneously lose oxygen. This Self Reduction result is again to be attributed in many cases to the primary formation of unstable more highly oxidised molecules. Many oxides of the noble metals, e.g. Au₂O₃, PtO₂, and HgO, exhibit this behaviour with neutral or alkaline solutions and emerge in the metallic state. Silver oxide behaves similarly and the action has been carefully investigated; apparently the change is not strictly representable by the equation

Ag₂O + H₂O₂ = 2Ag + H₂O + O₂

as might be expected. More oxygen is evolved than would be generated by the hydrogen peroxide alone, but the quantity is not exactly double. The probable course of the reaction is the formation of a peroxidic derivative of silver which then decomposes into silver and oxygen, the finely divided silver produced, possibly accompanied by some oxygen derivative of silver, catalytically accelerating the independent decomposition of the hydrogen peroxide.

In the majority of such reactions, however, the oxidised compound and the hydrogen peroxide appear to undergo reduction to an equivalent extent. Ozone and hydrogen peroxide react as follows:

H₂O₂ + O₃ = H₂O + 2O₂

this relationship holds only for the process in alkaline solution, in acid solution an excessive quantity of ozone undergoing Self Reduction or decomposition except when a very large excess of hydrogen peroxide is present. This is explained on the assumption that the interaction between ozone and hydrogen peroxide is accompanied by the spontaneous decomposition of ozone, this latter reaction being catalytically accelerated by the peroxide.

Although manganese dioxide behaves merely as a catalyst towards a neutral or alkaline solution of hydrogen peroxide, yet in acid solution reduction occurs to manganous oxide, MnO, or a corresponding salt, the hydrogen peroxide undergoing simultaneous reduction with the liberation of an equal amount of free oxygen. Lead dioxide and hydrogen peroxide in acidic solution likewise undergo mutual reduction. In both cases half the liberated oxygen may be attributed to the dioxide and half to the peroxide, but it is also possible that the mechanism of the reaction may involve the transference of an oxygen atom from the metallic peroxide to the hydrogen peroxide which thereby becomes oxidised to water and oxygen. If the latter view is correct, the liberated oxygen comes entirely from the hydrogen peroxide, but with either explanation the final result is the same. As manganous oxide, MnO, and lead oxide, PbO, are oxidised by hydrogen peroxide in the presence of alkali giving the corresponding dioxides, it is quite possible that the catalytic effect of manganese dioxide and of the less active lead dioxide on an alkaline solution of hydrogen peroxide is due to the repeated oxidation and reduction of the monoxides by the hydrogen peroxide.

Already, in the consideration of methods for the preparation of oxygen, mention has been made of the interaction of hydrogen peroxide with alkaline solutions of ferricyanides and with acidic solutions of hypochlorites and permanganates (hypochlorous and permanganic acids) with formation of the corresponding ferrocyanides, chlorides, and manganese salts. In these cases the reactions occur quantitatively, and can be represented as follows:

2K₃FeC₆N₆ + 2KOH + H₂O₂ = 2K₄FeC₆N₆ + 2H₂O + O₆;
HOCl + H₂O₂ = HCl + H₂O + O₂;
2HMnO₄ + 2H₂SO₄ + 5H₂O₂ = 2MnSO₄ + 8H₂O + 5O₂.

Again, the view is held by some chemists that in the interaction of permanganic acid and hydrogen peroxide the liberated oxygen originates entirely from the latter substance, which is oxidised by the permanganic acid with production of water and free oxygen. In addition to this, the observation that at temperatures below 0° C., interaction, as demonstrated by the decolorisation of the permanganate, will occur without any marked effervescence of oxygen, has led to the suggestion that a higher oxide, possibly hydrogen trioxide, H₂O₃, or hydrogen tetroxide, H₂O₄, is an unstable intermediate product, and that the evolution of oxygen occurs in the Self Reduction of this. This view has been vigorously combated, and it appears highly probable that the lack of effervescence in the cold is due merely to the oxygen remaining in supersaturated solution and perhaps in part as persulphuric acid if sulphuric acid is originally present. Even below 0° C. most of the oxygen may be liberated in the gaseous condition, and the reaction is so definite in its results at the ordinary temperature as to constitute a trustworthy and convenient method for the estimation of the concentration of hydrogen peroxide solutions.

In alkaline media, chromium oxide undergoes oxidation to a chromate. The importance of the medium is again seen in this case, because in acid solution a bichromate or chromic acid yields an unstable blue solution containing a labile perchromie acid which is stated to have the composition HCrO₄ or H₂CrO₅, according to the relative properties of the reagents. The blue solution soon decomposes, giving, in the presence of sulphuric acid, ordinary chromium sulphate and oxygen; the blue compound can be extracted with ether in which a more stable solution is obtained. The blue coloration has for years been made use of as a convenient and delicate test for chromic acid and for hydrogen peroxide.

Potassium persulphate also reacts slowly with hydrogen peroxide solutions, giving potassium hydrogen sulphate and oxygen -

H₂O₂ + K₂S₂O₈ = 2KHSO₄ + O₂.

A suspension of silver chloride in potassium hydroxide solution is rapidly reduced by hydrogen peroxide in accordance with the equation

2AgClH₂O₂ + 2KOH = 2Ag + 2KCl + 2H₂O + O₂.