Hydrogen peroxide is a chemical compound with the formula H2O2. In its pure form, it is a very pale blue[5]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.[6]
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.[7]
Structure and dimensions of H2O2 in the solid (crystalline) phase
Hydrogen peroxide (H2O2) is a nonplanar molecule with (twisted) C2symmetry; this was first shown by Paul-Antoine Giguère in 1950 using infrared spectroscopy.[9][10] 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.[11] 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.[12]
The molecular structures of gaseous and crystallineH2O2 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 groupD4 4 or P41212.[14]
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.[15]
Phase diagram of H2O2 and water: Area above blue line is liquid. Dotted lines separate solid–liquid phases from solid–solid phases.
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.
Properties of H2O2 and its analogues Values marked * are extrapolated
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.[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]
The amount of hydrogen peroxide in biological systems can be assayed using a fluorometric assay.[19]
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.[20] 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.[21][22][23]
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.[24]
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.[17]
Early attempts failed to produce neat hydrogen peroxide. Anhydrous hydrogen peroxide was first obtained by vacuum distillation.[25]
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.[26]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.[27] 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.[28][29]
Production
Catalytic cycle for the anthraquinone process to produce hydrogen peroxide: an anthraquinone (right) is reduced using hydrogen to produce the corresponding anthrahydroquinone (left). This is oxidized using oxygen to produce hydrogen peroxide and recover anthraquinone.
In 1994, world production of H2O2 was around 1.9 million tonnes and grew to 2.2 million in 2006,[30] 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".[31][clarification needed]
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 palladiumcatalyst. 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]
The net reaction for the anthraquinone-catalyzed process is:[31]
H2 + O2 → H2O2
The economics of the process depend heavily on effective recycling of the extraction solvents, the hydrogenation catalyst and the expensive quinone.
ISO tank container for hydrogen peroxide transportationA tank car designed for transporting hydrogen peroxide by rail
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.
Reactions
Acid-base
Hydrogen peroxide is about 1000 times stronger as an acid than water.[37]
H2O2 ⇌ H+ + HO−2 (pK = 11.65)
Disproportionation
Hydrogen peroxide disproportionates to form water and oxygen with a ΔHo of –2884.5 kJ/kg[38] 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).[39]
Oxidation reactions
The redox properties of hydrogen peroxide depend on pH. In acidic solutions, H2O2 is a powerful oxidizer.
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:
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.
Degradation of hypoxanthine through xanthine to uric acid to form hydrogen peroxide.
The degradation of guanosine monophosphate yields xanthine as an intermediate product which is then converted in the same way to uric acid with the formation of hydrogen peroxide.[48]
thus eliminating the poisonous hydrogen peroxide in the process.
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 reaction of Fe2+ and hydrogen peroxide is the basis of the Fenton reaction, which generates hydroxyl radicals, which are of significance in biology:
Fe(II) + H2O2 → Fe(III)OH + HO·
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 leads 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.
Eggs of sea urchin, shortly after fertilization by a sperm, produce hydrogen peroxide. It is then converted to hydroxyl radicals (HO•), which initiate radical polymerization, which surrounds the eggs with a protective layer of polymer.
As a proposed signaling molecule, hydrogen peroxide may regulate a wide variety of biological processes.[58][59] At least one study has tried to link hydrogen peroxide production to cancer.[60]
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.
The reaction with borax leads to sodium perborate, a bleach used in laundry detergents:
Na2B4O7 + 4 H2O2 + 2 NaOH → 2 Na2B2O4(OH)4 + H2O
Sewage treatment
Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. In advanced oxidation processing, the Fenton reaction[62][63] gives the highly reactive hydroxyl radical (•OH). This degrades organic compounds, including those that are ordinarily robust, such as aromatic or halogenated compounds.[64] It can also oxidize sulfur-based compounds present in the waste; which is beneficial as it generally reduces their odour.[65]
Disinfectant
Hydrogen peroxide may be used for the sterilization of various surfaces,[66] including surgical instruments,[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]
High-concentration H2O2 is referred to as "high-test peroxide" (HTP). It can be used as either a monopropellant (not mixed with fuel) or the oxidizer component of a bipropellant rocket. Use as a monopropellant takes advantage of the decomposition of 70–98% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber, where a catalyst, usually a silver or platinum screen, triggers decomposition, producing steam at over 600 °C (1,100 °F), which is expelled through a nozzle, generating thrust. H2O2 monopropellant produces a maximal specific impulse (Isp) of 161 s (1.6 kN·s/kg). Peroxide was the first major monopropellant adopted for use in rocket applications. Hydrazine eventually replaced hydrogen peroxide monopropellant thruster applications primarily because of a 25% increase in the vacuum specific impulse.[74] Hydrazine (toxic) and hydrogen peroxide (less toxic [ACGIH TLV 0.01 and 1 ppm respectively]) are the only two monopropellants (other than cold gases) to have been widely adopted and utilized for propulsion and power applications.[citation needed] The Bell Rocket Belt, reaction control systems for X-1, X-15, Centaur, Mercury, Little Joe, as well as the turbo-pump gas generators for X-1, X-15, Jupiter, Redstone and Viking used hydrogen peroxide as a monopropellant.[75] The RD-107 engines (used from 1957 to present) in the R-7 series of rockets decompose hydrogen peroxide to power the turbopumps.
In bipropellant applications, H2O2 is decomposed to oxidize a burning fuel. Specific impulses as high as 350 s (3.5 kN·s/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 109-500StarthilfeRATO 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.
In the 1940s and 1950s, the Hellmuth Walter KG–conceived turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant. Operator error in the use of hydrogen peroxide torpedoes was named as possible causes for the sinking of HMS Sidon and the Russian submarine Kursk.[78] SAAB Underwater Systems is manufacturing the Torpedo 2000. This torpedo, used by the Swedish Navy, is powered by a piston engine propelled by HTP as an oxidizer and kerosene as a fuel in a bipropellant system.[79][80]
Household use
Contact lenses soaking in a 3% hydrogen peroxide-based solution. The case includes a catalytic disc which neutralises the hydrogen peroxide over time.
Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.
Hair bleaching
Diluted H2O2 (between 1.9% and 12%) mixed with aqueous ammonia has been used to bleach human hair. The chemical's bleaching property lends its name to the phrase "peroxide blonde".[81]
Hydrogen peroxide is also used for tooth whitening. It may be found in most whitening toothpastes. Hydrogen peroxide has shown positive results involving teeth lightness and chroma shade parameters.[82] It works by oxidizing colored pigments onto the enamel where the shade of the tooth may become lighter.[further explanation needed] Hydrogen peroxide may be mixed with baking soda and salt to make a homemade toothpaste.[83]
Removal of blood stains
Hydrogen peroxide reacts with blood as a bleaching agent, and so if a blood stain is fresh, or not too old, liberal application of hydrogen peroxide, if necessary in more than single application, will bleach the stain fully out. After about two minutes of the application, the blood should be firmly blotted out.[84][85]
Acne treatment
Hydrogen peroxide may be used to treat acne,[86] although benzoyl peroxide is a more common treatment.
Oral cleaning agent
The use of dilute hydrogen peroxide as an oral cleansing agent has been reviewed academically to determine its usefulness in treating gingivitis and plaque. Although there is a positive effect when compared with a placebo, it was concluded that chlorhexidine is a much more effective treatment.[87]
Some horticulturists and users of hydroponics advocate the use of weak hydrogen peroxide solution in watering solutions. Its spontaneous decomposition releases oxygen that enhances a plant's root development and helps to treat root rot (cellular root death due to lack of oxygen) and a variety of other pests.[88][89]
For general watering concentrations, around 0.1% is in use. This can be increased up to one percent for antifungal actions.[90] Tests show that plant foliage can safely tolerate concentrations up to 3%.[91]
Fishkeeping
Hydrogen peroxide is used in aquaculture for controlling mortality caused by various microbes. In 2019, the U.S. FDA approved it for control of Saprolegniasis in all coldwater finfish and all fingerling and adult coolwater and warmwater finfish, for control of external columnaris disease in warm-water finfish, and for control of Gyrodactylus spp. in freshwater-reared salmonids.[92] Laboratory tests conducted by fish culturists have demonstrated that common household hydrogen peroxide may be used safely to provide oxygen for small fish. The hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide.
Removing yellowing from aged plastics
Hydrogen peroxide may be used in combination with a UV-light source to remove yellowing from white or light grey acrylonitrile butadiene styrene (ABS) plastics to partially or fully restore the original color. In the retrocomputing scene, this process is commonly referred to as retrobright.
Safety
Skin shortly after exposure to 35% H2O2
Regulations vary, but low concentrations, such as 5%, are widely available and legal to buy for medical use. Most over-the-counter peroxide solutions are not suitable for ingestion. Higher concentrations may be considered hazardous and typically are accompanied by a safety data sheet (SDS). In high concentrations, hydrogen peroxide is an aggressive oxidizer and will corrode many materials, including human skin. In the presence of a reducing agent, high concentrations of H2O2 will react violently.[93]
While concentrations up to 35% produce only "white" oxygen bubbles in the skin (and some biting pain) that disappear with the blood within 30–45 minutes, concentrations of 98% dissolve paper. However, concentrations as low as 3% can be dangerous for the eye because of oxygen evolution within the eye.[94]
High-concentration hydrogen peroxide streams, typically above 40%, should be considered hazardous due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds (45 kg), or approximately 10 US gallons (38 L), of concentrated hydrogen peroxide.
A commercial bottle of H2O2
Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminium alloys may also be suitable).[95] As it breaks down quickly when exposed to light, it should be stored in an opaque container, and pharmaceutical formulations typically come in brown bottles that block light.[96]
Hydrogen peroxide, either in pure or diluted form, may pose several risks, the main one being that it forms explosive mixtures upon contact with organic compounds.[97]Distillation of hydrogen peroxide at normal pressures is highly dangerous. It is corrosive, especially when concentrated, but even domestic-strength solutions may cause irritation to the eyes, mucous membranes, and skin.[98] Swallowing hydrogen peroxide solutions is particularly dangerous, as decomposition in the stomach releases large quantities of gas (ten times the volume of a 3% solution), leading to internal bloating. Inhaling over 10% can cause severe pulmonary irritation.[99]
With a significant vapour pressure (1.2 kPa at 50 °C),[100] hydrogen peroxide vapour is potentially hazardous. According to U.S. NIOSH, the immediately dangerous to life and health (IDLH) limit is only 75 ppm.[101] The U.S. Occupational Safety and Health Administration (OSHA) has established a permissible exposure limit of 1.0 ppm calculated as an 8-hour time-weighted average (29 CFR 1910.1000, Table Z-1).[97] Hydrogen peroxide has been classified by the American Conference of Governmental Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans".[102] For workplaces where there is a risk of exposure to the hazardous concentrations of the vapours, continuous monitors for hydrogen peroxide should be used. Information on the hazards of hydrogen peroxide is available from OSHA[97] and from the ATSDR.[103]
Wound healing
Historically, hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics.[104]
There is conflicting evidence on hydrogen peroxide's effect on wound healing. Some research finds benefit, while other research find delays and healing inhibition.[105] Its use for home treatment of wounds is generally not recommended.[106]
1.5–3% hydrogen peroxide is used as a disinfectant in dentistry, especially in endodotic treatments together with hypochlorite and chlorhexidine and 1–1.5% is also useful for treatment of inflammation of third molars (wisdom teeth).[107]
Both the effectiveness and safety of hydrogen peroxide therapy is scientifically questionable. Hydrogen peroxide is produced by the immune system, but in a carefully controlled manner. Cells called phagocytes engulf pathogens and then use hydrogen peroxide to destroy them. The peroxide is toxic to both the cell and the pathogen and so is kept within a special compartment, called a phagosome. Free hydrogen peroxide will damage any tissue it encounters via oxidative stress, a process that also has been proposed as a cause of cancer.[113]
Claims that hydrogen peroxide therapy increases cellular levels of oxygen have not been supported. The quantities administered would be expected to provide very little additional oxygen compared to that available from normal respiration. It is also difficult to raise the level of oxygen around cancer cells within a tumour, as the blood supply tends to be poor, a situation known as tumor hypoxia.
Large oral doses of hydrogen peroxide at a 3% concentration may cause irritation and blistering to the mouth, throat, and abdomen as well as abdominal pain, vomiting, and diarrhea.[109] Ingestion of hydrogen peroxide at concentrations of 35% or higher has been implicated as the cause of numerous gas embolism events resulting in hospitalisation. In these cases, hyperbaric oxygen therapy was used to treat the embolisms.[114]
Intravenous injection of hydrogen peroxide has been linked to several deaths.[115][111][112]
The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective, or useful cancer treatment."[110] Furthermore, the therapy is not approved by the U.S. FDA.
Historical incidents
On 16 July 1934, in Kummersdorf, Germany, a propellant tank containing an experimental monopropellant mixture consisting of hydrogen peroxide and ethanol exploded during a test, killing three people.[116]
In June 1955, Royal Navy submarine HMS Sidon sank after leaking high-test peroxide in a torpedo caused it to explode in its tube, killing twelve crew members; a member of the rescue party also succumbed.
In April 1992, an explosion occurred at the hydrogen peroxide plant at Jarrie in France, due to technical failure of the computerised control system and resulting in one fatality and wide destruction of the plant.[118]
The Russian submarine K-141 Kursk sailed to perform an exercise of firing dummy torpedoes at the Pyotr Velikiy, a Kirov-class battlecruiser. On 12 August 2000, at 11:28 local time (07:28 UTC), there was an explosion while preparing to fire the torpedoes. The only credible report to date is that this was due to the failure and explosion of one of the Kursk's hydrogen peroxide-fueled torpedoes. It is believed that HTP, a form of highly concentrated hydrogen peroxide used as propellant for the torpedo, seeped through its container, damaged either by rust or in the loading procedure on land where an incident involving one of the torpedoes accidentally touching ground went unreported. The vessel was lost with all hands.[120]
On 15 August 2010, a spill of about 30 US gallons (110 L) of cleaning fluid occurred on the 54th floor of 1515 Broadway, in Times Square, New York City. The spill, which a spokesperson for the New York City Fire Department said was of hydrogen peroxide, shut down Broadway between West 42nd and West 48th streets as fire engines responded to the hazmat situation. There were no reported injuries.[121]
See also
FOX reagent, used to measure levels of hydrogen peroxide in biological systems
^Easton MF, Mitchell AG, Wynne-Jones WF (1952). "The behaviour of mixtures of hydrogen peroxide and water. Part 1.—Determination of the densities of mixtures of hydrogen peroxide and water". Transactions of the Faraday Society. 48: 796–801. doi:10.1039/TF9524800796. S2CID96669623.
^Brauer G, ed. (1963). Handbook of preparative inorganic chemistry. Vol. 1. Translation editing by Reed F. (2nd ed.). New York: Academic Press. p. 140. ISBN978-0-12-126601-1.
^Chernyshov IY, Vener MV, Prikhodchenko PV, Medvedev AG, Lev O, Churakov AV (4 January 2017). "Peroxosolvates: Formation Criteria, H2O2 Hydrogen Bonding, and Isomorphism with the Corresponding Hydrates". Crystal Growth & Design. 17 (1): 214–220. doi:10.1021/acs.cgd.6b01449. ISSN1528-7483.
^Flohé L (December 2020). "Looking Back at the Early Stages of Redox Biology". Antioxidants. 9 (12): 1254. doi:10.3390/antiox9121254. PMC7763103. PMID33317108. I checked Humboldt's pertinent publication carefully, but was unable to find an unambiguous proof of this assumption; the description of the starting materials ("Alaun-Erden" or "schwere Erden") were just too unprecise to understand what kind of chemical experiments he performed.
^Giguère PA. "Hydrogen peroxide". Access Science. McGraw-Hill Education. doi:10.1036/1097-8542.329200. Archived from the original on 30 November 2018. Retrieved 28 November 2018. Hydrogen peroxide was discovered in 1818 by the French chemist Louis-Jacques Thenard, who named it eau oxygénée (oxygenated water).
^ abJones CW, Clark JH (1999). Applications of Hydrogen Peroxide and Derivatives. Royal Society of Chemistry. ISBN978-0-85404-536-5.
In his 1922 textbook, Joseph Mellor considered three hypothetical molecular structures for hydrogen peroxide, admitting (p. 952): "... the constitution of this compound has not been yet established by unequivocal experiments". See: Joseph William Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, vol. 1 (London, England: Longmans, Green and Co., 1922), p. 952–956.Archived 3 September 2016 at the Wayback Machine
W. C. Schumb, C. N. Satterfield, and R. L. Wentworth (1 December 1953) "Report no. 43: Hydrogen peroxide, Part two"Archived 26 February 2015 at the Wayback Machine, Office of Naval Research, Contract No. N5ori-07819 On p. 178, the authors present six hypothetical models (including cis-trans isomers) for hydrogen peroxide's molecular structure. On p. 184, the present structure is considered almost certainly correct—although a small doubt remained. (Note: The report by Schumb et al. was reprinted as: W. C. Schumb, C. N. Satterfield, and R. L. Wentworth, Hydrogen Peroxide (New York, New York: Reinhold Publishing Corp. (American Chemical Society Monograph), 1955).)
^Penney WG, Sutherland GB (1934). "The theory of the structure of hydrogen peroxide and hydrazine". Journal of Chemical Physics. 2 (8): 492–8. Bibcode:1934JChPh...2..492P. doi:10.1063/1.1749518.
^Penney WG, Sutherland GB (1934). "A note on the structure of H2O2 and H4N2 with particular reference to electric moments and free rotation". Transactions of the Faraday Society. 30: 898–902. doi:10.1039/tf934300898b.
^Ravikumar KS, Kesavan V, Crousse B, Bonnet-Delpon D, Bégué JP (2003). "Mild and Selective Oxidation of Sulfur Compounds in Trifluoroethanol: Diphenyldisulfide and Methyl phenyl Sulfoxide". Org. Synth. 80: 184. doi:10.15227/orgsyn.080.0184.
^Xu WL, Li YZ, Zhang QS, Zhu HS (2004). "A Selective, Convenient, and Efficient Conversion of Sulfides to Sulfoxides". Synthesis (2): 227–232. doi:10.1055/s-2004-44387.
^Giorgio M, Trinei M, Migliaccio E, Pelicci PG (September 2007). "Hydrogen peroxide: a metabolic by-product or a common mediator of ageing signals?". Nature Reviews. Molecular Cell Biology. 8 (9): 722–8. doi:10.1038/nrm2240. PMID17700625. S2CID6407526.
^Gonzalez D, Bejarano I, Barriga C, Rodriguez AB, Pariente JA (2010). "Oxidative Stress-Induced Caspases are Regulated in Human Myeloid HL-60 Cells by Calcium Signal". Current Signal Transduction Therapy. 5 (2): 181–6. doi:10.2174/157436210791112172.
^López-Lázaro M (July 2007). "Dual role of hydrogen peroxide in cancer: possible relevance to cancer chemoprevention and therapy". Cancer Letters. 252 (1): 1–8. doi:10.1016/j.canlet.2006.10.029. PMID17150302.
^Tarr MA, ed. (2003). Chemical degradation methods for wastes and pollutants environmental and industrial applications. New York: M. Dekker. p. 165. ISBN978-0-203-91255-3.
^Pignatello JJ, Oliveros E, MacKay A (January 2006). "Advanced Oxidation Processes for Organic Contaminant Destruction Based on the Fenton Reaction and Related Chemistry". Critical Reviews in Environmental Science and Technology. 36 (1): 1–84. Bibcode:2006CREST..36....1P. doi:10.1080/10643380500326564. S2CID93052585.
^Pera-Titus M, Garcıa-Molina V, Baños MA, Giménez J, Esplugas S (February 2004). "Degradation of chlorophenols by means of advanced oxidation processes: a general review". Applied Catalysis B: Environmental. 47 (4): 219–256. doi:10.1016/j.apcatb.2003.09.010.
^Falagas ME, Thomaidis PC, Kotsantis IK, Sgouros K, Samonis G, Karageorgopoulos DE (July 2011). "Airborne hydrogen peroxide for disinfection of the hospital environment and infection control: a systematic review". The Journal of Hospital Infection. 78 (3): 171–7. doi:10.1016/j.jhin.2010.12.006. PMID21392848.
^Scott R (November 1997). "Homing Instincts". Jane's Navy Steam Generated by Catalytic Decomposition of 80–90% Hydrogen Peroxide Was Used for Driving the Turbopump Turbines of the V-2 Rockets, the X-15 Rocketplanes, the Early Centaur RL-10 Engines and is Still Used on Soyuz for That Purpose Today. International. Archived from the original on 17 July 2011. Retrieved 12 May 2007.
^Sulieman M, Addy M, MacDonald E, Rees JS (May 2004). "The effect of hydrogen peroxide concentration on the outcome of tooth whitening: an in vitro study". Journal of Dentistry. 32 (4): 295–9. doi:10.1016/j.jdent.2004.01.003. PMID15053912.
^Capizzi R, Landi F, Milani M, Amerio P (August 2004). "Skin tolerability and efficacy of combination therapy with hydrogen peroxide stabilized cream and adapalene gel in comparison with benzoyl peroxide cream and adapalene gel in common acne. A randomized, investigator-masked, controlled trial". The British Journal of Dermatology. 151 (2): 481–4. doi:10.1111/j.1365-2133.2004.06067.x. PMID15327558. S2CID2611939.
^Bhattarai SP, Su N, Midmore DJ (2005). Oxygation Unlocks Yield Potentials of Crops in Oxygen-Limited Soil Environments. Advances in Agronomy. Vol. 88. pp. 313–377. doi:10.1016/S0065-2113(05)88008-3. ISBN978-0-12-000786-8.
^French LK, Horowitz BZ, McKeown NJ (July 2010). "Hydrogen peroxide ingestion associated with portal venous gas and treatment with hyperbaric oxygen: a case series and review of the literature". Clinical Toxicology. 48 (6): 533–8. doi:10.3109/15563650.2010.492526. PMID20575671. S2CID25148041.
^"Accident No: DCA-99-MZ-001"(PDF). U.S. National Transportation Safety Board. Archived(PDF) from the original on 3 November 2015. Retrieved 30 October 2015.
Drabowicz, J.; Kiełbasinski, P.; Mikołajczyk, M. (1994). "3. Synthesis of Sulphoxides". The Syntheses of Sulphones, Sulphoxides and Cyclic Sulphides. Wiley. pp. 109–254, See pp. 112–6. doi:10.1002/9780470666357.ch3. ISBN978-0-471-93970-2. OCLC521033898.
Greenwood NN, Earnshaw A (1997). Chemistry of the Elements (2nd ed.). Oxford UK: Butterworth-Heinemann. ISBN1-59124-291-6. OCLC49708420. A great description of properties & chemistry of H2O2.
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