Hydrogen Peroxide Formation

Of the various methods by which hydrogen peroxide has been artificially produced the oldest is that involving the interaction of barium peroxide and a mineral acid. As already mentioned, peroxides (superoxides) in general exhibit this reaction.

In 1882 it was observed by M. Traube that during the electrolysis of dilute sulphuric acid, hydrogen peroxide may be found both at the anode and the cathode.1 Its formation at the anode is to be ascribed to the decomposition of persulphuric acid, H₂S₂O₈, which is produced in some quantity. At the cathode the presence of the hydrogen peroxide is due to the reduction of dissolved oxygen, and it has been shown more recently that in solution in dilute sulphuric acid under a pressure of 100 atmospheres oxygen can be so successfully reduced to hydrogen peroxide as to give a yield of more than 80 per cent, of that calculated from the current used, and solutions of 2.3 per cent, concentration can be readily obtained. That the process is merely one of chemical reduction is confirmed by the fact that the dissolved oxygen can be reduced to hydrogen peroxide also by treatment with hydrogen in the presence of the metal palladium.

In the electrolysis of solutions of alkali hydroxides under suitable conditions, hydrogen peroxide is produced at the anode by the combination of the discharged hydroxyl ions.

In contradistinction to ozone, hydrogen peroxide is not produced during the electrolysis of solutions of fluorides.

As has been mentioned earlier, the slow atmospheric oxidation of substances at the ordinary temperature (autoxidation) frequently gives rise to peroxidic substances of powerful oxidising properties; these substances are unstable and in their decomposition give rise to ozone, and, if water is present, also to hydrogen peroxide. The slow oxidation of phosphorus in the presence of water is stated to be accompanied by the formation of a little hydrogen peroxide. In such cases the mechanism of the change probably involves the formation of an additive compound of the unstable primary peroxide compound with water, which subsequently decomposes with formation of hydrogen peroxide.

The corrosion of metals is another case of slow oxidation, and in the presence of water appreciable quantities of hydrogen peroxide may be produced; the amalgams of the metals frequently give better results than the metals themselves. If the hydrogen peroxide is removed from the solution as rapidly as it is produced (the addition of barium hydroxide, for example, will remove the hydrogen peroxide as a precipitate of barium peroxide), the yield of hydrogen peroxide may become almost quantitative, and in the case of zinc the final result of the reaction may be represented:

Zn + O₂ + 2H₂O = Zn(OH)₂ + H₂O₂.

Other metals, e.g. magnesium, cadmium, and lead, can be made to yield similar results, as also does palladium hydride when allowed to oxidise in the presence of water containing a little sulphuric acid. It is possible, however, that, at least in the ease of lead, the hydrogen peroxide owes its formation to the reduction of dissolved oxygen by nascent hydrogen produced during the corrosion.

Many organic substances, such as alcohols, ethers, acetone, and especially unsaturated compounds like turpentine, are capable of slow oxidation, the action being favoured by exposure to sunlight; in the presence of moisture, hydrogen peroxide is frequently to be found amongst the products of the chemical change. The disinfectant power of the " Sanitas " preparations, the basis of which is obtained by the atmospheric oxidation of wet turpentine oil, is largely due to hydrogen peroxide.

Hydrogen and oxygen, and even steam and oxygen, can be made to combine at low temperatures under the influence of the silent electric discharge, and the process may be regarded as a slow oxidation of hydrogen comparable with the preceding. Also the silent electric discharge generally favours the production of unstable substances. For these two reasons, therefore, the formation of hydrogen peroxide might be expected, and the expectation is justified by experiment. By working with a well-cooled gaseous mixture, it is possible to obtain a considerable yield of hydrogen peroxide, whilst at -80° C. the yield is almost quantitative and the product pure.

Although radium salts will decompose hydrogen peroxide, they likewise form this substance when their rays act upon water. Kernbaum concluded that the β rays are the most effective agents, and suggested that the reaction takes places as follows:

2H₂O = H₂O₂ + H₂.

Hydrogen peroxide occurs frequently amongst the products of gaseous combustion in the presence of moisture, and of the combination of hydrogen and oxygen by flame or explosion. On account of the instability of hydrogen peroxide, it is advisable to cool the products as rapidly as possible, for example, by allowing the flame to impinge on the surface of cold water or ice, when the condensed liquid will exhibit the reactions of hydrogen peroxide.

Whether the production of hydrogen peroxide in the last case is a direct process or is due to the further interaction of water vapour and oxygen at high temperature is rather uncertain. Certainly at very high temperatures oxygen and water-vapour will combine, with formation of some hydrogen peroxide which can be detected after rapid cooling. Indeed, the combination of water-vapour and oxygen may be effected even at 130° C. under the influence of a silent electric discharge, and at very high temperatures even water-vapour alone, without the addition of an excess of oxygen, will undergo slight conversion into hydrogen peroxide. It is therefore possible that the slight formation of hydrogen peroxide in the combustion of hydrogen, of moist carbon monoxide, or cyanogen, may be due to a purely thermal influence on the water, or water and oxygen, present. The observation that traces of hydrogen peroxide are produced when an arc discharge is formed, using very dilute sulphuric acid as a cathode, may have a similar explanation.