Contents
Hydrogen peroxide
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Hydrogen peroxide
Identifiers
CAS number 7722-84-1
Y
PubChem 784
ChemSpider 763
EC number 231-765-0
UN number 2015 (>60% soln.)
2014 (20–60% soln.)
2984 (8–20% soln.)
RTECS number MX0900000 (>90% soln.)
MX0887000 (>30% soln.)
Properties
Molecular formula H2O2
Molar mass 34.0147 g/mol
Appearance Very light blue color; colorless in solution
Density 1.463 g/cm3
Melting point
-0.43 °C, 273 K, 31 °F
Boiling point
150.2 °C, 423 K, 302 °F
Solubility in water Miscible
Solubility soluble in ether
Acidity (pKa) 11.62 [1]
Refractive index (nD) 1.34
Viscosity 1.245 cP (20 °C)
Dipole moment 2.26 D
Thermochemistry
Std enthalpy of
formation ΔfH
Specific heat capacity, C 1.267 J/g K (gas)
2.619 J/g K (liquid)
Hazards
MSDS ICSC 0164 (>60% soln.)
EU Index 008-003-00-9
EU classification Oxidant (O)
Corrosive (C)
Harmful (Xn)
R-phrases R5, R8, R20/22, R35
S-phrases (S1/2), S17, S26, S28, S36/37/39, S45
NFPA 704
Flash point Non-flammable
LD50 1518 mg/kg
Related compounds
Related compounds Water
Ozone
Hydrazine
Hydrogen disulfide
Y (what is this?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references
Hydrogen peroxide (H2O2) is a very pale blue liquid, slightly more viscous than water, that appears colorless in dilute solution. It is a weak acid, has strong oxidizing properties, and is a powerful bleaching agent. It is used as a disinfectant, antiseptic, oxidizer, and in rocketry as a propellant.[2] The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species.
Hydrogen peroxide is naturally produced in organisms as a byproduct of oxygen metabolism. Nearly all living things (specifically, all obligate and facultative aerobes) possess enzymes known as peroxidases, which harmlessly and catalytically decompose low concentrations of hydrogen peroxide to water and oxygen.
Structure and properties
As with all molecules, the physical properties of hydrogen peroxide are the result of its molecular mass, structure, and distribution of atoms within the molecule.
The preferred molecular structure of any molecule is the configuration that has the lowest internal stress. For hydrogen peroxide, there are two basic structural forms (conformers) available for the molecule. Whereas flat shape of the anti conformer would minimize steric repulsions, the 90° torsion angle of the syn conformer would optimize mixing between the filled p-type orbital of the oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond.[3]
The resulting anticlinal "skewed" shape is a compromise between the two conformers.
Despite the fact that the O-O bond is a single bond, the molecule has a remarkably high barrier to complete rotation of 29.45 kJ/mol (compared with 12.5 kJ/mol for the rotational barrier of ethane). The increased barrier is attributed to repulsion between one lone pair and other lone pairs. The bond angles are affected by hydrogen bonding, which is relevant to the structural difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H2O2.
Analogues of hydrogen peroxide include the chemically identical deuterium peroxide and malodorous hydrogen disulfide.[4] Hydrogen disulfide has a boiling point of only 70.7°C despite having a higher molecular weight, indicating that hydrogen bonding increases the boiling point of hydrogen peroxide.[5]
Aqueous hydrogen peroxide solutions have specific properties that are different from those of the pure chemical due to hydrogen bonding between water and hydrogen peroxide molecules. Specifically, hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression. While pure water melts and freezes at approximately 273K, and pure hydrogen peroxide just 0.4K below that, a 50% (by volume) solution melts and freezes at 221 K.[6]
History
Hydrogen peroxide was first isolated in 1818 by Louis Jacques Thénard by reacting barium peroxide with nitric acid.[7] An improved version of this process used hydrochloric acid, followed by sulfuric acid to precipitate the barium sulfate byproduct. Thénard's process was used from the end of the 19th century until the middle of the 20th century.[8] Modern production methods are discussed below.
For a long time, pure hydrogen peroxide was believed to be unstable, because attempts to separate the hydrogen peroxide from the water, which is present during synthesis, failed. This was because traces of solids and heavy metal ions led to a catalytic decomposition or explosions of the hydrogen peroxide. One hundred percent pure hydrogen peroxide was first obtained through vacuum distillation by Richard Wolffenstein in 1894.[9] At the end of 19th century, Petre Melikishvili and his pupil L. Pizarjevski showed that of the many proposed formulas of hydrogen peroxide, the correct one was H-O-O-H.
The use of H2O2 sterilization in biological safety cabinets and barrier isolators is a popular alternative to ethylene oxide (EtO) as a safer, more efficient decontamination method. H2O2 has long been widely used in the pharmaceutical industry. In aerospace research, H2O2 is used to sterilize satellites.
The FDA has recently granted 510(k) clearance to use H2O2 in individual medical device manufacturing applications. EtO criteria outlined in ANSI/AAMI/ISO 14937 may be used as a validation guideline. Sanyo was the first manufacturer to use the H2O2 process in situ in a cell culture incubator, which is a faster and more efficient cell culture sterilization process.
Manufacture
Formerly inorganic processes were used, employing the electrolysis of an aqueous solution of sulfuric acid or acidic ammonium bisulfate (NH4HSO4), followed by hydrolysis of the peroxodisulfate ((SO4)2)2− that is formed.
Today, hydrogen peroxide is manufactured almost exclusively by the autoxidation of a 2-alkyl anthrahydroquinone (or 2-alkyl-9,10-dihydroxyanthracene) to the corresponding 2-alkyl anthraquinone. Major producers commonly use either the 2-ethyl or the 2-amyl derivative. The cyclic reaction depicted below shows the 2-ethyl derivative, where 2-ethyl-9,10-dihydroxyanthracene (C16H14O2) is oxidized to the corresponding 2-ethylanthraquinone (C16H12O2) and hydrogen peroxide. Most commercial processes achieve this by bubbling compressed air through a solution of the anthracene, whereby the oxygen present in the air reacts with the labile hydrogen atoms (of the hydroxy group), giving hydrogen peroxide and regenerating the anthraquinone. Hydrogen peroxide is then extracted out and the anthraquinone derivative is reduced back to the dihydroxy (anthracene) compound using hydrogen gas in the presence of a metal catalyst. The cycle then repeats itself.[10][11]