Hydrogen peroxide

Hydrogen peroxide is a chemical compound with the formula. In its pure form, it is a very pale blue liquid; However at lower concentrations, it appears colorless due to the faintness of the blue coloration. The molecule hydrogen peroxide is asymmetrical and highly polarized. Its strong tendency to form hydrogen bond networks results in greater viscosity compared to water. It is used as an oxidizer, bleaching agent, and antiseptic, usually as a dilute solution in water for consumer use and in higher concentrations for industrial use. Concentrated hydrogen peroxide, or "high-test peroxide", decomposes explosively when heated and has been used as both a monopropellant and an oxidizer in rocketry.
Hydrogen peroxide is a reactive oxygen species and the simplest peroxide, a compound having an oxygen–oxygen single bond. It decomposes slowly into water and elemental oxygen when exposed to light, and rapidly in the presence of organic or reactive compounds. It is typically stored with a stabilizer in a weakly acidic solution in an opaque bottle. Hydrogen peroxide is found in biological systems including the human body. Enzymes that use or decompose hydrogen peroxide are classified as peroxidases.

Properties

The boiling point of has been extrapolated as being, approximately higher than water. In practice, hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled at lower temperatures under reduced pressure.
Hydrogen peroxide forms stable adducts with urea, sodium carbonate and other compounds. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for in some reactions.

Structure

Hydrogen peroxide is a nonplanar molecule with C₂ symmetry; this was first shown by Paul-Antoine Giguère in 1950 using infrared spectroscopy. Although the O−O bond is a single bond, the molecule has a relatively high rotational barrier of 386 cm⁻¹ for rotation between enantiomers via the trans configuration, and 2460 cm⁻¹ via the cis configuration. These barriers are proposed to be due to repulsion between the lone pairs of the adjacent oxygen atoms and dipolar effects between the two O–H bonds. For comparison, the rotational barrier for ethane is 1040 cm⁻¹.
The approximately 100° dihedral angle between the two O–H bonds makes the molecule chiral. It is the smallest and simplest molecule to exhibit enantiomerism. It has been proposed that the enantiospecific interactions of one rather than the other may have led to amplification of one enantiomeric form of ribonucleic acids and therefore an origin of homochirality in an RNA world.
The molecular structures of gaseous and crystalline are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state. Crystals of are tetragonal with the space group D or P4₁2₁2.

Aqueous solutions

In aqueous solutions, hydrogen peroxide forms a eutectic mixture, exhibiting freezing-point depression down as low as −56 °C; pure water has a freezing point of 0 °C and pure hydrogen peroxide of −0.43 °C. The boiling point of the same mixtures is also depressed in relation with the mean of both boiling points. It occurs at 114 °C. This boiling point is 14 °C greater than that of pure water and 36.2 °C less than that of pure hydrogen peroxide.

	Density	Temp.
3%	1.0095	15
27%	1.10	20
35%	1.13	20
50%	1.20	20
70%	1.29	20
75%	1.33	20
96%	1.42	20
98%	1.43	20
100%	1.45	20

Hydrogen peroxide is most commonly available as a solution in water. For consumers, it is usually available from pharmacies at 3 and 6 wt% concentrations. The concentrations are sometimes described in terms of the volume of oxygen gas generated; one milliliter of a 20-volume solution generates twenty milliliters of oxygen gas when completely decomposed. For laboratory use, 30 wt% solutions are most common. Commercial grades from 70% to 98% are also available, but due to the potential of solutions of more than 68% hydrogen peroxide to be converted entirely to steam and oxygen these grades are potentially far more hazardous and require special care in dedicated storage areas. Buyers must typically allow inspection by commercial manufacturers.

Comparison with analogues

Hydrogen peroxide has several structural analogues with bonding arrangements. It has the highest boiling point of this series. Its melting point is also fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising significantly more readily, indicative of particularly strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs.

Name	Formula	Molar mass	Melting point	Boiling point
Water	HOH	18.02	0.00	99.98
Hydrogen peroxide	HOOH	34.01	−0.43	150.2*
Hydrogen disulfide	HSSH	66.15	−89.6	70.7
Hydrazine	H₂NNH₂	32.05	2	114
Hydroxylamine	NH₂OH	33.03	33	58*
Diphosphane	H₂PPH₂	65.98	−99	63.5*

Natural occurrence

Hydrogen peroxide is produced by various biological processes mediated by enzymes.
Hydrogen peroxide has been detected in surface water, in groundwater, and in the atmosphere. It can also form when water is exposed to UV light. Sea water contains 0.5 to 14 μg/L of hydrogen peroxide, and freshwater contains 1 to 30 μg/L. Concentrations in air are about 0.4 to 4 μg/m³, varying over several orders of magnitude depending in conditions such as season, altitude, daylight and water vapor content. In rural nighttime air it is less than 0.014 μg/m³, and in moderate photochemical smog it is 14 to 42 μg/m³.
The amount of hydrogen peroxide in biological systems can be assayed using a fluorometric assay.

Discovery

is sometimes said to have been the first to report the first synthetic peroxide, barium peroxide, in 1799 as a by-product of his attempts to decompose air, although this is disputed due to von Humboldt's ambiguous wording. Nineteen years later Louis Jacques Thénard recognized that this compound could be used for the preparation of a previously unknown compound, which he described as eau oxygénée — subsequently known as hydrogen peroxide.
An improved version of Thénard's process used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate byproduct. This process was used from the end of the 19th century until the middle of the 20th century.
The bleaching effect of peroxides and their salts on natural dyes had been known since Thénard's experiments in the 1820s, but early attempts of industrial production of peroxides failed. The first plant producing hydrogen peroxide was built in 1873 in Berlin. The discovery of the synthesis of hydrogen peroxide by electrolysis with sulfuric acid introduced the more efficient electrochemical method. It was first commercialized in 1908 in Weißenstein, Carinthia, Austria. The anthraquinone process, which is still used, was developed during the 1930s by the German chemical manufacturer IG Farben in Ludwigshafen. The increased demand and improvements in the synthesis methods resulted in the rise of the annual production of hydrogen peroxide from 35,000 tonnes in 1950, to over 100,000 tonnes in 1960, to 300,000 tonnes by 1970; by 1998 it reached 2.7 million tonnes.
Early attempts failed to produce neat hydrogen peroxide. Anhydrous hydrogen peroxide was first obtained by vacuum distillation.
Determination of the molecular structure of hydrogen peroxide proved to be very difficult. In 1892, the Italian physical chemist Giacomo Carrara determined its molecular mass by freezing-point depression, which confirmed that its molecular formula is. seemed to be just as possible as the modern structure, and as late as in the middle of the 20th century at least half a dozen hypothetical isomeric variants of two main options seemed to be consistent with the available evidence. In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide that was very similar to the presently accepted one.

Production

In 1994, world production of was around 1.9 million tonnes and grew to 2.2 million in 2006, most of which was at a concentration of 70% or less. In that year, bulk 30% sold for around 0.54 USD/kg, equivalent to US$1.50/kg on a 100% purity basis.
Hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was originally developed by BASF in 1939. It begins with the reduction of an anthraquinone to the corresponding anthrahydroquinone, typically by hydrogenation on a palladium catalyst. In the presence of oxygen, the anthrahydroquinone then undergoes autoxidation: the labile hydrogen atoms of the hydroxy groups transfer to the oxygen molecule, to give hydrogen peroxide and regenerating the anthraquinone. Most commercial processes achieve oxidation by passing compressed air through a solution of the anthrahydroquinone, with the hydrogen peroxide then extracted from the solution and the anthraquinone recycled back for successive cycles of hydrogenation and oxidation.
The net reaction for the anthraquinone-catalyzed process is:
The economics of the process depend heavily on effective recycling of the extraction solvents, the hydrogenation catalyst and the expensive quinone.