Chemical compound
Hydrogen peroxide is a chemical compound with the formula H2O2. In its pure form, it is a very pale blue liquid that is slightly more viscous than water. It is used as an oxidizer, bleaching agent, and antiseptic, usually as a dilute solution (3%–6% by weight) 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 H2O2 has been extrapolated as being 150.2 °C (302.4 °F), approximately 50 °C (90 °F) 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 (Hydrogen peroxide - urea), sodium carbonate (sodium percarbonate) and other compounds. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H2O2 in some reactions.
Structure
Hydrogen peroxide (H2O2) is a nonplanar molecule with (twisted) 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 (4.62 kJ/mol) for rotation between enantiomers via the trans configuration, and 2460 cm−1 (29.4 kJ/mol) 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 (12.4 kJ/mol).
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 H2O2 are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state. Crystals of H2O2 are tetragonal with the space group D 4 4 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 (125.1 °C). 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.
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 (with the temperature of the steam increasing as the concentration increases above 68%) 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 HmX−XHn bonding arrangements (water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, S, N, P). 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.
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 forms upon illumination of water. 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
Alexander von Humboldt 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 ("oxygenated water") – 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 (1864–1925) determined its molecular mass by freezing-point depression, which confirmed that its molecular formula is H2O2. H2O=O 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 H2O2 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% H2O2 sold for around 0.54 USD/kg, equivalent to US$1.50/kg (US$0.68/lb) on a "100% basis".
Today, 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 (such as 2-ethylanthraquinone or the 2-amyl derivative) 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 bubbling 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:
H2 + O2 → H2O2
The economics of the process depend heavily on effective recycling of the extraction solvents, the hydrogenation catalyst and the expensive quinone.
Historical methods
Hydrogen peroxide was once prepared industrially by hydrolysis of ammonium persulfate:
[NH4]2S2O8 + 2 H2O → 2 [NH4]HSO4 + H2O2
[NH4]2S2O4 was itself obtained by the electrolysis of a solution of ammonium bisulfate ([NH4]HSO4) in sulfuric acid.
Other routes
Small amounts are formed by electrolysis, photochemistry, and electric arc, and related methods.
A commercially viable route for hydrogen peroxide via the reaction of hydrogen with oxygen favours production of water but can be stopped at the peroxide stage. One economic obstacle has been that direct processes give a dilute solution uneconomic for transportation. None of these has yet reached a point where it can be used for industrial-scale synthesis.
Reactions
Hydrogen peroxide is about 1000 times stronger as an acid than water.
H2O2 ⇌ H+ + HO2− pK = 11.65
Disproportionation
Hydrogen peroxide disproportionates to form water and oxygen with a ΔH o of −2884.5 kJ/kg and a ΔS of 70.5 J/(mol·K):
2 H2O2 → 2 H2O + O2
The rate of decomposition increases with rise in temperature, concentration, and pH. H2O2 is unstable under alkaline conditions. Decomposition is catalysed by various redox-active ions or compounds, including most transition metals and their compounds (e.g. manganese dioxide (MnO2), silver, and platinum).
Oxidation reactions
The redox properties of hydrogen peroxide depend on pH. In acidic solutions, H2O2 is a powerful oxidizer.
Sulfite (SO2−3) is oxidized to sulfate (SO2−4).
Reduction reactions
Under alkaline conditions, hydrogen peroxide is a reductant. When H2O2 acts as a reducing agent, oxygen gas is also produced. For example, hydrogen peroxide will reduce sodium hypochlorite and potassium permanganate, which is a convenient method for preparing oxygen in the laboratory:
NaOCl + H2O2 → O2 + NaCl + H2O
2 KMnO4 + 3 H2O2 → 2 MnO2 + 2 KOH + 2 H2O + 3 O2
The oxygen produced from hydrogen peroxide and sodium hypochlorite is in the singlet state.
Although usually a reductant, alkaline hydrogen peroxide converts Mn(II) to the dioxide:
H2O2 + Mn2+ + 2 OH− → MnO2 + 2 H2O
In a related reaction, potassium permanganate is reduced to Mn2+ by acidic H2O2:
2 MnO4− + 5 H2O2 + 6 H+ → 2 Mn2+ + 8 H2O + 5 O2
Organic reactions
Hydrogen peroxide is frequently used as an oxidizing agent. Illustrative is oxidation of thioethers to sulfoxides:
Ph-S-CH3 + H2O2 → Ph-S(O)-CH3 + H2O
Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acid derivatives, and for the oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation. It is also the principal reagent in the Dakin oxidation process.
Precursor to other peroxide compounds
Hydrogen peroxide is a weak acid, forming hydroperoxide or peroxide salts with many metals. It also converts metal oxides into the corresponding peroxides. For example, upon treatment with hydrogen peroxide, chromic acid (CrO3 and H2SO4) forms a blue peroxide CrO(O2)2.
Biochemistry
Production
The aerobic oxidation of glucose in the presence of the enzyme glucose oxidase produces hydrogen peroxide. The conversion affords gluconolactone:
C6H12O7 + O2 → C6H10O7 + H2O2
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