Catalytic Decomposition of hydrogen peroxide

Hydrogen peroxide can be regarded as stable only in cold dilute aqueous solutions free from even traces of alkali, compounds of the heavy metals, and suspended solid matter, and when protected from exposure to bright illumination. Thus in solution in tap water decomposition occurs fifty times as rapidly as in conductivity water. Minute quantities of alkalis markedly accelerate the rate of decomposition, possibly by the intermediate formation of unstable "salts " of the peroxide, the activity of the catalyst being dependent upon concentration of the hydroxyl ion which it yields. Exposure to ultra-violet radiation also induces the rapid decomposition of hydrogen peroxide, the reaction in this case being unimoleeular, and also differing from the spontaneous or purely thermal decomposition of the solutions in being accelerated both by alkalis and acids.

On account of the extreme sensitiveness of hydrogen peroxide to external influences, it is almost impossible to consider the decomposition process apart from the effect of catalysts because the minute traces of foreign matter inevitably present in a " pure " substance and the walls of the containing vessel are capable in this case of producing a disproportionate effect on the stability of the compound. Even water exerts a distinct Catalytic Decomposition of hydrogen peroxide effect. Ordinary pure hydrogen peroxide and its solutions, as is to be expected, decompose more rapidly when the temperature is raised, but, as has already been stated, the decomposition is sufficiently gradual to permit distillation under reduced pressure. The thermal decomposition is a bimolecular process, and in the case of the pure compound the heat generated during the chemical change tends to make this become explosive. With pure hydrogen peroxide the oxygen liberated measured at 14° and 760° mm. is 475 times the volume of the original hydrogen peroxide; with solutions of the substance the volume of oxygen (at N.T.P.) producible by the decomposition of unit volume of the solution supplies a convenient method of expressing the concentration; solutions are therefore frequently sold not as of a certain percentage concentration but as " ten volume," " twenty volume," etc.:

2H₂O₂ = 2H₂O + O₂.

A ten-volume strength corresponds approximately to a 3 per cent, solution, and a twenty-volume strength to a 6 per cent. Perhydrol is practically a 30 per cent, solution corresponding to a one hundred-volume strength.

Many other substances than those mentioned above possess the power of markedly influencing the rate of decomposition of hydrogen peroxide, especially in feebly alkaline solution. Carbon, silver, gold, the platinum metals, and many other substances in a fine state of division are exceedingly active, but if the metal exposes only a smooth polished surface, the result may be relatively inappreciable. In colloidal solution the activity of the noble metals is still greater, and the effect of one part of colloidal platinum in more than 100,000,000 times its weight of water exerts a distinct Catalytic Decomposition of hydrogen peroxide action. Whether the effect of such metals is due to a mere surface action or to the continuous formation and decomposition of an intermediate unstable oxide is uncertain, but the former view appears more probable, because all finely divided substances exhibit a similar although often feebler effect; silica powder, for example, has a very considerable accelerating influence on the decomposition. It is a very remarkable fact that these colloidal substances lose their power in the presence of almost equally small quantities of such substances as mercuric chloride, phosphorus and arsenic hydrides, hydrocyanic acid, and hydrogen sulphide, which are therefore described as "poisons" to the catalyst. Some organic substances such as alcohol act similarly. The decomposition of hydrogen peroxide is usually regarded as a reaction of the first order 8 or monomolecular, whether in acid or in neutral solution. The reaction is convenient to study, since its rate can be followed by titration with permanganate, or volumetrically, by measuring the volume of evolved oxygen.

In alkaline solution the activity of the colloidal platinum increases to a maximum with increase of alkalinity, and then decreases. In this respect it behaves in an analogous manner to certain inorganic ferments. Exposure to Rontgen rays retards the reaction. Colloidal rhodium, palladium, iridium, silver, and gold behave in an analogous manner to platinum. By increasing the pressure of oxygen in contact with the decomposing solution from one to 200 atmospheres the rate of decomposition is not appreciably affected.

Other metals, the commonest being lead, bismuth, and manganese, in powder form exert a more moderate effect on the decomposition. Mercury would also fall into this class of moderate accelerators, but the catalytic action in this case is remarkable in being periodic or rhythmic. When the concentration of hydrogen ion is reduced to an almost negligible quantity by the addition of a little sodium acetate solution, a clean mercury surface in contact with hydrogen peroxide solution of approximately 10 per cent, concentration, at periodic intervals of about one second, becomes coated with a bronze film which suddenly disappears with a burst of oxygen from the contact layer of the two liquids; the substance of the film, which is alternately formed and decomposed, is probably an unstable oxide, possibly mercurous peroxide.

Copper, nickel, cobalt, and cadmium have only a feeble effect on the rate of decomposition of the substance.

Many compounds, especially various metallic oxides, also induce very rapid decomposition of hydrogen peroxide without themselves being permanently changed. In addition to the solutions of the alkali hydroxides already mentioned, manganese dioxide, cobalt oxide, and lead oxide (massicot) are remarkably active, and as might be expected a colloidal solution of manganese dioxide is also able to exert powerful catalytic influence. The effect in such cases may be partly a surface effect, but is also probably due in part to the intermediate formation and decomposition of unstable highly oxidised derivatives.

The oxides of iron, bismuth, copper, cerium, and magnesium are capable of exerting an appreciable influence on the rate of decomposition, but much depends on the physical condition of the solid, freshly precipitated iron oxide, for example, being more effective than the ignited substance; aluminium hydroxide is rather exceptional in behaving as a "negative catalyst" and retarding the decomposition.

Amongst other inorganic catalysts are to be included the iodides (also bromides and chlorides but less active) and chromates or dichromates. The agent in the former case appears to be the iodide ion, the mechanism of the reaction probably involving oxidation to hypoiodite which then reacts with more hydrogen peroxide with formation of iodide, and water and free oxygen (compare the reaction with hypochlorites).

H₂O₂ + I' = H₂O + IO'; H₂O₂ + b>IO' = H₂O + I' + O₂.

The action of iodides is catalytic only in neutral solution; in alkaline solution oxidation occurs with formation of free iodine. Whilst chromates and dichromates accelerate the decomposition of hydrogen peroxide, the action is not purely catalytic because some of the chromate or dichromate undergoes permanent reduction, so that the change falls more correctly into the next section.

With iodine the following reactions are believed to occur:

(i) 5H₂O₂ + I₂ = 2HIO₃ + 4H₂O,
(ii) 5H₂O₂ + 2HIO₃ = 5O₂ + I₂ + 6H₂O,
(iii) 2H₂O₂ = 2H₂O + O₂,

the last-named reaction being catalysed by the iodic acid - iodine couple. In ammonium hydroxide solution the reaction takes place in accordance with the equation

2NH₃ + I₂ + H₂O₂ = 2NH₄I + O₂.

Manganous sulphate and ferric salts in general accelerate the decomposition of hydrogen peroxide. The sulphate is less active than the chloride or nitrate. With dilute solutions of the salts the effect is proportional to the concentration of the peroxide and that of the iron ions, whilst in the presence of acids it is inversely proportional to the hydrogen-ion concentration. The temperature coefficient of the reaction is 3.25 for ten degrees.

If a few drops of potassium ferrocyanide solution are added to dilute hydrogen peroxide (1 per cent.), and kept in the dark, decomposition of the latter is exceedingly slow. On placing in direct sunlight for a few moments, however, brisk evolution of oxygen takes place and continues, even after removal from the light. The effect is not due to rise of temperature, but, presumably, to some catalyst generated under the influence of the light.

Even carbon, in the form of charcoal, catalytically decomposes hydrogen peroxide, its activity being apparently connected with its absorptive power for gases.

Certain complex organic substances are known to catalyse the decomposition of hydrogen peroxide: of these "catalases" the best known is the "haemase" present in blood, in the presence of which the decomposition process is greatly accelerated and approximates to a unimolecular reaction. It is a remarkable circumstance that many of the "poisons" which destroy the catalytic power of the colloidal noble metals have a similar effect on the power of haemase, but the list of poisons is not quite the same for the two types of catalyst. In view of the marked influence of these poisons or negative catalysts it is possible that the various preservatives mentioned earlier are effective in a similar manner, namely, by checking the activity of traces of positive catalysts such as alkali.

It is interesting to note that metallic salts may catalytically decompose hydrogen peroxide in organic solvents such as amyl acetate and quinoline. In the latter solvent, if not more than 2 per cent, of water is present, the velocity of decomposition in the presence of manganese acetate corresponds to that required for a bimolecular reaction. But if the quinoline is saturated with water, "the reaction is monomolecular. On the other hand, not a few organic substances tend to stabilise hydrogen peroxide solutions. Amongst these are oxalic, uric and benzoic acids, acetanilide.

Concentrated solutions of sodium chloride preserve the peroxide, provided a catalyst such as sodium hydroxide is excluded. Dilute sulphuric acid is very effective, even 0.00066 gram of the acid per litre exerting a marked retarding effect upon the rate of decomposition of 30 per cent, peroxide solution. There would appear to be no simple relationship between the retardation effect and the concentration of the sulphuric acid. A yellow bottle is preferable to a white or blue one.

Although not strictly a case of catalysis, the effect of radium radiation on the rate of decomposition of hydrogen peroxide is most conveniently mentioned here; the penetrating rays are the most effective.