(H2O2), a colourless liquid usually produced as aqueous solutions of various strengths, used principally for bleaching cotton and other textiles and wood pulp, in the manufacture of other chemicals, as a rocket propellant, and for cosmetic and medicinal purposes. Solutions containing more than about 8 percent hydrogen peroxide are corrosive to the skin.
First recognized as a chemical compound in 1818, hydrogen peroxide is the simplest member of the class of s. Of the several processes of manufacture, the principal ones involve reactions of oxygen from the air with certain organic compounds, especially anthraquinone or isopropyl alcohol. Major commercial grades are aqueous solutions containing 35, 50, 70, or 90 percent hydrogen peroxide and small amounts of stabilizers (often tin salts and phosphates) to suppress decomposition.
Hydrogen peroxide decomposes into water and oxygen upon heating or in the presence of numerous substances, particularly salts of such metals as iron, copper, manganese, nickel, or chromium. It combines with many compounds to form crystalline solids useful as mild oxidizing agents; the best-known of these is sodium perborate (NaBO2H2O23H2O or NaBO34H2O). With certain organic compounds, hydrogen peroxide reacts to form hydroperoxides or peroxides, several of which are used to initiate polymerization reactions. In most of its reactions, hydrogen peroxide oxidizes other substances, although it is itself oxidized by a few compounds, such as potassium permanganate.
Pure hydrogen peroxide freezes at -0.43? C (+31.3? F) and boils at 150.2? C (302? F); it is denser than water and is soluble in it in all proportions.
The most important covalent peroxide is hydrogen peroxide, H2O2. When pure, this syrupy, viscous liquid has a pale blue colour, although it appears almost colourless. Many of its physical properties resemble those of water. It has a larger liquid range than water, melting at -0.43? C and boiling at 150.2? C, and it has a higher density (1.44 grams per cubic centimetre at 25? C) than water. The dielectric constant of pure H2O2 is, like that of water, quite high–70.7 at 25? C compared with a value of 78.4 for water at 25? C. However, adding water, which is miscible in all proportions, causes the dielectric constant to increase to a maximum value of 121 at about 35 percent H2O2 and 65 percent H2O. World production of H2O2 is well over one-half million tons per year, making it a major industrial chemical. Most industrial hydrogen peroxide is prepared by a well-conceived process introduced originally by I.G. Farbenindustrie of Germany that uses only hydrogen and oxygen as raw materials. The process involves oxidation of 2-ethylanthraquinol to 2-ethylanthraquinone by passage of air through a solution of the quinol in an organic solvent. The hydrogen peroxide that is produced is extracted into water. The quinone is then reduced back to the quinol by hydrogen in the presence of palladium metal on an inert support. The process is thus a cyclic one. It can be shown by an examination of reduction potentials that aqueous solutions of hydrogen peroxide or the pure liquid should spontaneously decompose to water and oxygen.
2H2O + O2 —> 2H2O2
In the absence of catalysts, minimal decomposition occurs. In the presence of even trace amounts of many metal ions or metal surfaces, however, explosive decomposition can occur. Traces of alkali metal ions dissolved from glass can cause this decomposition, and, for this reason, pure H2O2 (or a concentrated solution) is normally stored in wax-coated or plastic bottles. Hydrogen peroxide is a strong oxidizing agent in either acidic or basic solutions and will also act as a reducing agent toward very strong oxidizing agents, such as the permanganate ion, MnO4-. The largest industrial use of hydrogen peroxide is as a bleach for such materials as textiles, paper pulp, and leather. It is used in dilute solution as a mild antiseptic and disinfectant and is employed in the production of organic stabilizers, polymerization initiators, curing agents, and pharmaceuticals.