hydrogen peroxide(redirected from Hydrogen dioxide)
Also found in: Dictionary, Thesaurus, Medical, Wikipedia.
hydrogen peroxide,chemical compound, H2O2, a colorless, syrupy liquid that is a strong oxidizing agent and, in water solution, a weak acid. It is miscible with cold water and is soluble in alcohol and ether. Although pure hydrogen peroxide is fairly stable, it decomposes into water and oxygen when heated above about 80°C;; it also decomposes in the presence of numerous catalysts, e.g., most metals, acids, or oxidizable organic materials. A small amount of stabilizer, usually acetanilide, is often added to it. Hydrogen peroxide has many uses. It is available for household use as a 3% (by weight) water solution; it is used as a mild bleaching agent and medicinally as an antiseptic. The 3% solution is sometimes called ten volume strength, since one volume of it releases ten volumes of oxygen when it decomposes. Hydrogen peroxide is available for commercial use in several concentrations. Highly concentrated solutions were first used in World War II by the military, e.g., in fuels for rockets and torpedoes. It is used as a bleaching agent for textiles, e.g., wool and silk, and in paper manufacture. It is also used in chemical manufacture. Hydrogen peroxide is prepared commercially by oxidation of alkylhydroanthraquinones and by electrolysis of ammonium bisulfate. It can also be prepared by reaction of barium peroxide with sulfuric acid and is prepared (with acetone) by oxidation of isopropanol. Hydrogen peroxide was discovered (1818) by L. J. Thenard.
H2O2, the simplest and most important peroxide. It is a transparent, colorless, and odorless liquid with a metallic taste, a boiling point of 150.2°C, and a density of 1.47 g/cm3 at 0°C. Hydrogen peroxide has a melting point of — 0.43°C and can be easily overcooled without solidification. Hydrogen peroxide is completely miscible with water and forms the crystalline hydrate H2O2 • 2H2O. Like water, it is a good solvent for many salts, with which it forms crystalline peroxyhydrates. Hydrogen peroxide was discovered in 1818 by L. J. Thénard.
Very pure hydrogen peroxide is stable, but in the presence of heavy metals and their ions, it decomposes into H2O and O2. Especially efficient catalysts for this decomposition are salts and complexes of iron, copper, and manganese as well as the enzyme catalase. The decomposition of hydrogen peroxide is an exothermic process that can be explosive. Depending on the conditions, hydrogen peroxide can either be a reducing agent or, more commonly, an oxidizing agent. As an oxidizing agent, it liberates, for example, iodine from iodide salts:
2K1 +H2O2 + H2SO4 = I2 + K2SO4 + 2H2O
As a reducing agent, it converts Mn(VII) to Mn(II):
2KMnO4 + 5H2O2 + 3H2SO4 = K2SO4 + 2MnSO4 + 5O2 + 8H2O
These reactions are used for quantitative analysis of hydrogen peroxide in solution.
The mechanism of the oxidation of various substances by hydrogen peroxide is complex. All these oxidations involve formation of the active intermediates HO2 and OH, which are stronger oxidizers than hydrogen peroxide itself. For example, hydrogen peroxide reacts with ferrous ions in solution:
Fe2+ + H2O2 = Fe3+ + OH + OH–
Fenton’s reagent, which is widely used as an oxidizer of organic compounds, is a mixture of solutions of H2O2 and Fe(II) salts.
In the laboratory, hydrogen peroxide is prepared by treating metal peroxides, usually BaO2 or Na2O2, with cold, dilute acids. Industrial preparation involves electrolysis of sulfuric acid and subsequent hydrolysis of the resulting persulfuric acid, H2S2O8:
2H2SO4 → H2S2O8 + 2H+ + 2e–
H2S2O8 + 2H2O = 2H2SO4 + H2O2
Hydrogen peroxide can also be industrially prepared by autoxi-dation of derivatives of anthraquinone and by oxidation of iso-propyl alcohol.
In nature, hydrogen peroxide is formed as an intermediate or side product in the oxidation of many coumpounds by atmospheric oxygen. Traces are found in the various forms of natural atmospheric precipitation. Hydrogen peroxide is formed in plant and animal cells, but its concentration is very low because hydrogen peroxide is decomposed rapidly by the action of the enzymes catalase and peroxidase and because it is rapidly taken up to oxidize organic compounds.
Highly concentrated hydrogen peroxide, upon decomposition over an oxide catalyst, gives a gaseous water-oxygen mixture, which can be heated to 700°C for use as a jet fuel. In the chemical industry, hydrogen peroxide is used as an oxidizing agent, as a starting material for obtaining many peroxides, as an initiator of polymerization, and as a bleach for silk, wool, feathers, and furs.
Among the wastes from chemical production, hydrogen peroxide has acquired special significance as a clean oxidizing agent, one that does not form toxic products and pollute the environment. The production rate of hydrogen peroxide in concentrations from 90 to 98 percent is increasing steadily: this highly concentrated form is stored in aluminum tanks, and sodium pyrophosphate, Na4P2O7, is usually used as a stabilizer. Hydrogen peroxide is not toxic, but its concentrated solutions cause burns upon contact with the skin, mucous membranes, and respiratory tract.
In medicine hydrogen peroxide is used as a disinfecting, deodorizing antiseptic. A 3-percent solution of hydrogen peroxide is used for rinsing and washing in stomatitis, tonsillitis, and gynecological diseases; sometimes it is also used for stopping nosebleeds. Superoxol is used when more concentrated solutions are required. Solutions and salves that contain hydrogen peroxide are also used as depigmenting agents.
REFERENCESchumb, W., C. Satterfield, and R. Wentworth. Perekis’ vodoroda. Moscow, 1958. (Translated from English.)
A. P. PURMAL’