Hydrogen Peroxide 3 Vs 35

[Jen Ondrey]

Theboiling 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.[7]

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.[13] Crystals of H2O2 are tetragonal with the space group D4

4 or P41212.[14]

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.[15]

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.

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.[16] Sea water contains 0.5 to 14 μg/L of hydrogen peroxide, and freshwater contains 1 to 30 μg/L.[17] 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.[18]

An improved version of Thnard'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.[24]

The bleaching effect of peroxides and their salts on natural dyes had been known since Thnard'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 Weienstein, 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.[17]

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.[31][32]

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.[35][36] 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.

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).[39]

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:

Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acid derivatives,[42] 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.

Peroxisomes are organelles found in virtually all eukaryotic cells.[45] They are involved in the catabolism of very long chain fatty acids, branched chain fatty acids, D-amino acids, polyamines, and biosynthesis of plasmalogens, ether phospholipids, which are found in mammalian brains and lungs.[46] They produce hydrogen peroxide in s process catalyzed by flavin adenine dinucleotide (FAD):[47]

Hydrogen peroxide arises by the degradation of adenosine monophosphate, which yields hypoxanthine. Hypoxanthine is then oxidatively catabolized first to xanthine and then to uric acid, and the reaction is catalyzed by the enzyme xanthine oxidase:[48]

This reaction is important in liver and kidney cells, where the peroxisomes neutralize various toxic substances that enter the blood. Some of the ethanol humans drink is oxidized to acetaldehyde in this way.[49] In addition, when excess H2O2 accumulates in the cell, catalase converts it to H2O through this reaction:

The fenton reaction explains the toxicity of hydrogen peroxides because the hydroxyl radicals rapidly and irreversibly oxidize all organic compounds, including proteins, membrane lipids, and DNA.[50] Hydrogen peroxide is a significant source of oxidative DNA damage in living cells. DNA damage includes formation of 8-Oxo-2'-deoxyguanosine among many other altered bases, as well as strand breaks, inter-strand crosslinks, and deoxyribose damage.[51] By interacting with Cl hydrogen peroxide also lead to chlorinated DNA bases.[51] Hydroxyl radicals readily damage vital cellular components, especially those of the mitochondria.[52][53][54] The compound is a major factor implicated in the free-radical theory of aging, based on its ready conversion into a hydroxyl radical.

The bombardier beetle combine hydroquinone and hydrogen peroxide, leading to a violent exothermic chemical reaction to produce boiling, foul-smelling liquid that partially becomes a gas (flash evaporation) and is expelled through an outlet valve with a loud popping sound.[55][56][57]

About 60% of the world's production of hydrogen peroxide is used for pulp- and paper-bleaching.[30] The second major industrial application is the manufacture of sodium percarbonate and sodium perborate, which are used as mild bleaches in laundry detergents. A representative conversion is:

Sodium percarbonate, which is an adduct of sodium carbonate and hydrogen peroxide, is the active ingredient in such laundry products as OxiClean and Tide laundry detergent. When dissolved in water, it releases hydrogen peroxide and sodium carbonate.[24] By themselves these bleaching agents are only effective at wash temperatures of 60 C (140 F) or above and so, often are used in conjunction with bleach activators, which facilitate cleaning at lower temperatures.

It is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example.[61] Peroxy acids, such as peracetic acid and meta-chloroperoxybenzoic acid also are produced using hydrogen peroxide. Hydrogen peroxide has been used for creating organic peroxide-based explosives, such as acetone peroxide. It is used as an initiator in polymerizations. Hydrogen peroxide reacts with certain di-esters, such as phenyl oxalate ester (cyalume), to produce chemiluminescence; this application is most commonly encountered in the form of glow sticks.

Hydrogen peroxide may be used for the sterilization of various surfaces,[66] including surgical tools,[67] and may be deployed as a vapour (VHP) for room sterilization.[68] H2O2 demonstrates broad-spectrum efficacy against viruses, bacteria, yeasts, and bacterial spores.[69][70] In general, greater activity is seen against Gram-positive than Gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms may increase tolerance in the presence of lower concentrations.[71] Lower levels of concentration (3%) will work against most spores; higher concentrations (7 to 30%) and longer contact times will improve sporicidal activity.[70][72]

Hydrogen peroxide is seen as an environmentally safe alternative to chlorine-based bleaches, as it degrades to form oxygen and water and it is generally recognized as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA).[73]

In bipropellant applications, H2O2 is decomposed to oxidize a burning fuel. Specific impulses as high as 350 s (3.5 kNs/kg) can be achieved, depending on the fuel. Peroxide used as an oxidizer gives a somewhat lower Isp than liquid oxygen but is dense, storable, and non-cryogenic and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle. It may also be used for regenerative cooling of rocket engines. Peroxide was used very successfully as an oxidizer in World War II German rocket motors (e.g., T-Stoff, containing oxyquinoline stabilizer, for both the Walter HWK 109-500 Starthilfe RATO externally podded monopropellant booster system and the Walter HWK 109-509 rocket motor series used for the Me 163B), most often used with C-Stoff in a self-igniting hypergolic combination, and for the low-cost British Black Knight and Black Arrow launchers. Presently, HTP is used on ILR-33 AMBER[76] and Nucleus[77] suborbital rockets.

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