Chemical elements
    Physical Properties
    Chemical Properties
    Hydrogen peroxide
      Physical Properties
      Chemical Properties
      Catalytic Decomposition
      Self Reduction
      Oxidation Processes
      Detection and Estimation

Oxidation Processes with hydrogen peroxide

Hydrogen peroxide possesses exceptional activity as an oxidising agent. Nascent hydrogen is oxidised to water, so that, on electrolysis, hydrogen peroxide solutions undergo reduction at the cathode whilst at the anode also decomposition occurs, the nascent oxygen appearing to have a similar effect to permanganic or hypo- chlorous acid

H2O2 + 2H = 2H2O;
H2O2 + O = H2O + O2.

Hydrogen peroxide readily oxidises alkali nitrites to nitrates in acid solution. In alkaline or neutral solution oxidation does not take place.

Silicic acid gel, when evaporated with a slight excess of 30 per cent, peroxide, yields a highly active amorphous residue which continuously evolves ozonised oxygen. It liberates chlorine from hydrochloric acid, iodine from potassium iodide, decolorises permanganate, and evolves ozone with concentrated sulphuric acid. Possibly a persilicic acid is formed. Thiosulphates are at first oxidised to tetrathionates, the solution becoming alkaline:

2Na2S2O3 + H2O2 = Na2S4O6 + 2NaOH.

The reaction soon becomes more complex, a fact that will account for various published discrepancies.

Concentrated sulphuric acid is oxidised by hydrogen peroxide giving permonosulphuric acid, also called Caro's acid,

H2SO4 + H2O2H2SO5 + H2O;

sulphurous acid yields sulphuric acid and hydrogen sulphide undergoes slow conversion into free sulphur and even into sulphuric acid, the formation of the latter being easily demonstrated by heating together hydrogen sulphide, barium chloride, and hydrogen peroxide in aqueous solution. Hydrogen selenide is oxidised more readily, with formation of selenium. The metallic sulphides become converted into sulphates,

PbS + 4H2O2 = PbSO4 + 4H2O,

and for this reason hydrogen peroxide is frequently applied in the restoration of old paintings in which the white-lead basis of the paint has become blackened by the action of atmospheric hydrogen sulphide.

In alkaline solution cobalt sulphide is oxidised to cobaltic hydroxide, manganese sulphide yields the hydroxide and a deposit of sulphur, whilst zinc sulphide is converted into soluble zincates.

Dilute solutions of hydrogen peroxide (6 per cent.) oxidise yellow phosphorus on warming, phosphorous and phosphoric acids resulting. Amorphous phosphorus is violently attacked by 8 per cent, peroxide, hydrogen phosphide being evolved, phosphorous and phosphoric acids remaining in solution.

Metallic potassium and sodium are explosively converted into the hydroxides when brought into contact with concentrated solutions of hydrogen peroxide; many of the heavier metals such as zinc and iron, and especially aluminium, are readily changed into their respective hydroxides, whilst chromium, arsenic, and molybdenum are oxidised respectively to chromic, arsenic, and molybdic acids. Colloidal tellurium yields telluric acid with very dilute solutions of peroxide; the crystalline modification reacts slowly with 60 per cent, peroxide at 100° C.

Ordinary lead oxide becomes oxidised to the dioxide, and manganese oxide also to its dioxide; many other oxides and hydroxides undergo similar oxidation, the products frequently being unstable peroxidic substances.

In many of these oxidation processes a considerable proportion of the hydrogen peroxide undergoes concurrent decomposition with liberation of gaseous oxidation. Thus an acidified solution of potassium iodide gives a slow formation of iodine, the change being representable as

2HI + H2O2 = 2H2O + I2.

This oxidation, and indeed many others with the same oxidising agent, are greatly accelerated by the presence of certain inorganic substances, particularly iron salts, and especially when these are in the ferrous condition. The addition of a very small quantity of ferrous sulphate to a dilute solution of potassium iodide containing hydrogen peroxide, acetic acid, and starch, reduces in a remarkable manner the time necessary for the production of the well-known blue coloration. Copper salts are less active, but a mixture of copper sulphate and ferrous sulphate is a much better accelerator than would be expected, the copper sulphate appearing disproportionately to augment the activity of the ferrous salt. Complex organic catalysts have also been discovered.

In the presence of hydrochloric acid or hydrobromic acid the oxidation of hydriodic acid may proceed further, the iodine being 'converted into iodic acid possibly by way of iodine trichloride or bromide. However, periodic acid is reduced to iodic acid by hydrogen peroxide and in dilute solution partial reduction even to iodine may occur.

Iodides in general, and also to a less degree bromides and chlorides, even in small quantities, increase the rate of decomposition of hydrogen peroxide in neutral or alkaline solution. This catalytic effect has already been mentioned. Some metals will dissolve in cold, and sometimes in diluted acid solutions in the presence of hydrogen peroxide, even if almost insoluble in them under ordinary conditions. Thus glacial acetic acid containing hydrogen peroxide will attack bismuth, copper, lead, mercury, and silver in the cold; and dilute sulphuric acid charged with peroxide effects the solution of bismuth, copper, mercury, nickel, and silver.

Many per-acids or salts of such acids have been prepared by the action of hydrogen peroxide on the corresponding derivative of the normal acids; thus pernitrates, perborates, percarbonates, permolybdates, pervanadates, pertitanates, and others have been rendered accessible by the strong oxidising character of this substance. Details of such compounds will be found under the heading of the respective parent elements. In this connection it is interesting to note that hydrogen peroxide actually reduces many of these per-acids as well as leading to their formation. Thus an ethereal solution of perchromic acid is gradually decomposed by hydrogen peroxide, as are also permono-phosphoric and pervanadic acids.

With organic substances pure hydrogen peroxide is a powerful and valuable oxidising agent; it will oxidise acetyl chloride to peracetic acid, CH3.CO3H, and acetyl peroxide (CH3.CO)2O2, volatile unstable liquids; another general method for the preparation of such organic per-acids is by the interaction of hydrogen peroxide with the acid itself in the presence of sulphuric or nitric acid as catalyst.

CH3.CO2H + H2O2 = CH3.CO3H + H2O.

The organic per-acids of lower molecular weight are generally pungent, unstable - even explosive - liquids; those of higher molecular weight, such as per benzoic acid, are crystalline compounds and rather more stable.

For many organic oxidation processes, a solution of ordinary 30 per cent, aqueous hydrogen peroxide in acetic acid is used, the latter solvent occasionally being essential. In such cases it is probable that the oxidation is really effected by peracetic acid.

With such a solution of hydrogen peroxide, organic sulphides can be easily and conveniently oxidised to the corresponding sulphoxides and sulphones.


and azo-compounds to azoxy-compounds

R2N2R2N2O, where R and R' represent organic radicles.

Ketones and aldehydes containing respectively the characteristic atomic groups and ,\treact readily with aqueous hydrogen peroxide especially in the presence of a little hydrochloric acid; acetone, CH3.CO.CH3, for example, gives an explosive crystalline acetone peroxide of the molecular formula (C3H6O2)2, and benzaldehyde, C6H5.CHO, yields a fairly stable substance (C6H5CHO2)2. Carbon monoxide appears to be unaffected by hydrogen peroxide.

In the presence of a small quantity of ferrous sulphate an aqueous solution of hydrogen peroxide forms a valuable reagent for the oxidation of polyhydroxy compounds such as glycerol, glycol, and mannitol, a terminal hydroxyl group being invariably converted into an aldehydic one the reaction, for glycerol being representable as follows:


Tartaric acid with the same reagent undergoes oxidation to dihydroxymaleic acid, CO2H.C(OH):C(OH).CO2H, which readily undergoes further oxidation to dihydroxy-tartaric acid, CO2H.C(OH)2.C(OH)2.CO2H, which possesses especial interest on account of the very sparing solubility of its sodium salt. For this reason the acid is recommended by Fenton as a reagent for the quantitative estimation of sodium. The hexose sugars, glucose, fructose, etc., are oxidised by hydrogen peroxide containing the same catalyst (Fenton's reagent) with production of the corresponding osones.

It is interesting to note that, whilst solutions of aniline green and magenta are not bleached by dilute hydrogen peroxide solution in the dark, yet upon exposure to the light of a quartz-mercury lamp the colours readily fade. It would appear, therefore, that under the influence of the light, the peroxide becomes increasingly active.

The bleaching of litmus and of indigo solution (in the latter case with the aid of a little ferrous sulphate) is evidence of the oxidising power of hydrogen peroxide, but probably one of the most striking oxidations effected by hydrogen peroxide is that of benzene to phenol and further into quinol, pyrogallol, quinone, and other products.

A remarkable property of hydrogen peroxide, which may be mentioned here, although in the result the effect is not an oxidising one, is the power of causing organic cyanides to unite with the elements of water, undergoing hydrolysis to the corresponding amides. In acid solution cyanogen gives rise to oxamide, and phenyl cyanide (benzo- nitrile) to benzamide

CN.CN + 2H2O = CO(NH2).CO(NH2);
C6H5.CN + H2O = C6H5.CO.(NH2).

© Copyright 2008-2012 by