|Jmol 3D model
||MX0900000 (>90% soln.)
MX0887000 (>30% soln.)
||2015 (>60% soln.)
2014 (20–60% soln.)
2984 (8–20% soln.)
||Very light blue color; colorless in solution
||1.11 g/cm3 (20 °C, 30% (w/w) solution )
1.450 g/cm3 (20 °C, pure)
||−0.43 °C (31.23 °F; 272.72 K)
||150.2 °C (302.4 °F; 423.3 K) (decomposes)
||soluble in ether, alcohol
insoluble in petroleum ether
||5 mmHg (30 °C)
||1.245 cP (20 °C)
||1.267 J/(g·K) (gas)
2.619 J/(g·K) (liquid)
||A01AB02 (WHO) D08AX01 (WHO), S02AA06 (WHO)
|Safety data sheet
||ICSC 0164 (>60% soln.)
||R5, R8, R20/22, R35
||(S1/2), S17, S26, S28, S36/37/39, S45
|Lethal dose or concentration (LD, LC):
2000 mg/kg (oral, mouse)
||1418 ppm (rat, 4 hr)
||227 ppm (mouse)
|US health exposure limits (NIOSH):
||TWA 1 ppm (1.4 mg/m3)
||TWA 1 ppm (1.4 mg/m3)
Except where otherwise noted, data are given for materials in their standard state
(at 25 °C [77 °F], 100 kPa).
| verify (what is ?)
Hydrogen peroxide is a chemical compound with the formula H
2. In its pure form, it is a colourless liquid, slightly more viscous than water; however, for safety reasons it is normally used as an aqueous solution. Hydrogen peroxide is the simplest peroxide (a compound with an oxygen–oxygen single bond) and finds use as a strong oxidizer, bleaching agent and disinfectant. Concentrated hydrogen peroxide, or "high-test peroxide", is a reactive oxygen species and has been used as a propellant in rocketry.
Hydrogen peroxide is often described as being "water but with one more oxygen atom", a description that can give the incorrect impression of significant chemical similarity between the two compounds. While they have a similar melting point and appearance, pure hydrogen peroxide will explode if heated to boiling, will cause serious contact burns to the skin and can set materials alight on contact. For these reasons it is usually handled as a dilute solution (household grades are typically 3–6% in the U.S. and somewhat higher in Europe). Its chemistry is dominated by the nature of its unstable peroxide bond.
Structure and properties
The boiling point of H
2 has been extrapolated as being 150.2 °C, approximately 50 °C 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.
In aqueous solutions hydrogen peroxide differs from the pure material due to the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression; pure water has a melting point of 0 °C and pure hydrogen peroxide of −0.43 °C, but a 50% (by volume) solution of the two freezes at −51 °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.
and water: Area above blue line is liquid. Dotted lines separate solid+liquid phases from solid+solid phases.
Density of aqueous solution of H2O2
Hydrogen peroxide (H
2) is a nonplanar molecule with (twisted) C2 symmetry. Although the O−O bond is a single bond, the molecule has a relatively high rotational barrier of 2460 cm−1 (29.45 kJ/mol); for comparison, the rotational barrier for ethane is 12.5 kJ/mol. The increased barrier is ascribed to repulsion between the lone pairs of the adjacent oxygen atoms and results in hydrogen peroxide displaying atropisomerism.
The molecular structures of gaseous and crystalline H
2 are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state. Crystals of H
2 are tetragonal with the space group D4
Structure and dimensions of H2O2 in the gas phase
Structure and dimensions of H2O2 in the solid (crystalline) phase
Comparison with analogues
Hydrogen peroxide has several structural analogues with Hm−X−X−Hn bonding arrangements (water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, N, S). 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. All of these analogues are thermodynamically unstable. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs.
Hydrogen peroxide was first described in 1818 by Louis Jacques Thénard, who produced it by treating barium peroxide with nitric acid. An improved version of this process used hydrochloric acid, followed by addition of 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.
Pure hydrogen peroxide was long believed to be unstable, as early attempts to separate it from the water, which is present during synthesis, all failed. This instability was due to traces of impurities (transition-metal salts), which catalyze the decomposition of the hydrogen peroxide. Pure hydrogen peroxide was first obtained in 1894 — almost 80 years after its discovery — by Richard Wolffenstein, who produced it 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. At least half a dozen hypothetical molecular structures 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.
Previously, hydrogen peroxide was prepared industrially by hydrolysis of the ammonium peroxydisulfate, which was itself obtained by the electrolysis of a solution of ammonium bisulfate (NH
4) in sulfuric acid:
8 + 2 H
2O → H
2 + 2 (NH
Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was formalized in 1936 and patented 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; the anthrahydroquinone then undergoes to autoxidation to regenerate the starting anthraquinone, with hydrogen peroxide being produced as a by-product. Most commercial processes achieve oxidation by bubbling compressed air through a solution of the derivatized 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, 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.
The simplified overall equation for the process is deceptively simple:
2 + O
2 → H
The economics of the process depend heavily on effective recycling of the quinone (which is expensive) and extraction solvents, and of the hydrogenation catalyst.
A process to produce hydrogen peroxide directly from the elements has been of interest for many years. Direct synthesis is difficult to achieve, as the reaction of hydrogen with oxygen thermodynamically favours production of water. Systems for direct synthesis have been developed, most of which are based around finely dispersed metal catalysts. None of these has yet reached a point where they can be used for industrial-scale synthesis.
ISO tank container for hydrogen peroxide transportation
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.
In 1994, world production of H
2 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% H
2 sold for around 0.54 USD/kg, equivalent to 1.50 USD/kg (0.68 USD/lb) on a "100% basis".
Hydrogen peroxide is thermodynamically unstable and decomposes to form water and oxygen with a ΔH
o of −98.2 kJ/mol and a ΔS of 70.5 J/(mol·K):
- 2 H
2 → 2 H
2O + O
The rate of decomposition increases with rising temperature, concentration and pH, with cool, dilute, acidic solutions showing the best stability. Decomposition is catalysed by various compounds, including most transition metals and their compounds (e.g. manganese dioxide, silver, and platinum). Certain metal ions, such as Fe2+
, can cause the decomposition to take a different path, with free radicals such as (HO·) and (HOO·) being formed. Non-metallic catalysts include potassium iodide, which reacts particularly rapidly and forms the basis of the elephant toothpaste experiment. Hydrogen peroxide can also be decomposed biologically by enzyme catalase. The decomposition of hydrogen peroxide liberates oxygen and heat; this can be dangerous, as spilling high-concentration hydrogen peroxide on a flammable substance can cause an immediate fire.
Hydrogen peroxide exhibits oxidizing and reducing properties, depending on pH.
In acidic solutions, H
2 is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H
2 can be converted into hydroxyl radicals (·OH), which are highly reactive.
In acidic solutions Fe2+
is oxidized to Fe3+
(hydrogen peroxide acting as an oxidizing agent):
- 2 Fe2+
(aq) + H
2 + 2 H+
(aq) → 2 Fe3+
(aq) + 2 H
and sulfite (SO2−
3) is oxidized to sulfate (SO2−
4). However, potassium permanganate is reduced to Mn2+
by acidic H
2. Under alkaline conditions, however, some of these reactions reverse; for example, Mn2+
is oxidized to Mn4+
In basic solution, hydrogen peroxide can reduce a variety of inorganic ions. When it 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 + H
2 → O
2 + NaCl + H
- 2 KMnO
4 + 3 H
2 → 2 MnO
2 + 2 KOH + 2 H
2O + 3 O
Hydrogen peroxide is frequently used as an oxidizing agent. Illustrative is oxidation of thioethers to sulfoxides:
3 + H
2 → Ph−S(O)−CH
3 + H
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 (CrO
3) forms an unstable blue peroxide CrO(O
This kind of reaction is used industrially to produce peroxoanions. For example, reaction with borax leads to sodium perborate, a bleach used in laundry detergents:
7 + 4 H
2 + 2 NaOH → 2 Na
4 + H
2 converts carboxylic acids (RCO2H) into peroxy acids (RC(O)O2H), which are themselves used as oxidizing agents. Hydrogen peroxide reacts with acetone to form acetone peroxide and with ozone to form trioxidane. Reaction with urea produces the adduct, hydrogen peroxide - urea, used for whitening teeth. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H
2 in some reactions.
Hydrogen peroxide is also one of the two chief chemicals in the defence system of the bombardier beetle, reacting with hydroquinone to discourage predators.
A study published in Nature found that hydrogen peroxide plays a role in the immune system. Scientists found that hydrogen peroxide presence inside cells increased after tissues are damaged in zebrafish, which is thought to act as a signal to white blood cells to converge on the site and initiate the healing process. When the genes required to produce hydrogen peroxide were disabled, white blood cells did not accumulate at the site of damage. The experiments were conducted on fish; however, because fish are genetically similar to humans, the same process is speculated to occur in humans. The study in Nature suggested asthma sufferers have higher levels of hydrogen peroxide in their lungs than healthy people, which could explain why asthma sufferers have inappropriate levels of white blood cells in their lungs.
Hydrogen peroxide has important roles as a signalling molecule in the regulation of a wide variety of biological processes. The compound is a major factor implicated in the free-radical theory of aging, based on how readily hydrogen peroxide can decompose into a hydroxyl radical and how superoxide radical byproducts of cellular metabolism can react with ambient water to form hydrogen peroxide. These hydroxyl radicals in turn readily react with and damage vital cellular components, especially those of the mitochondria. At least one study has also tried to link hydrogen peroxide production to cancer. These studies have frequently been quoted in fraudulent treatment claims.
The amount of hydrogen peroxide in biological systems can be assayed using a fluorimetric assay.
About 60% of the world's production of hydrogen peroxide is used for pulp- and paper-bleaching. The second major industrial application is the manufacture of sodium percarbonate and sodium perborate which are used as mild bleaches in laundry detergents.
It is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example. This is used in polymerisations, as a flour bleaching agent and as a treatment for acne. Peroxy acids, such as peracetic acid and meta-chloroperoxybenzoic acid are also typically produced using hydrogen peroxide.
Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. This is achieved by advanced oxidation processes, such as the Fenton reaction, which use it to generate highly reactive hydroxyl radicals (·OH). These are able to destroy organic contaminates which are ordinarily difficult to remove, such as aromatic or halogenated compounds. It can also oxidize sulfur based compounds present in the waste; which is beneficial as it generally reduces their odour.
Skin shortly after exposure to 35% H
Hydrogen peroxide can be used for the sterilization of various surfaces, including surgical tools and may be deployed as a vapour (VHP) for room sterilization. H2O2 demonstrates broad-spectrum efficacy against viruses, bacteria, yeasts, and bacterial spores. In general, greater activity is seen against gram-positive than gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms can increase tolerance in the presence of lower concentrations. Higher concentrations of H2O2 (10 to 30%) and longer contact times are required for sporicidal activity.
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).
Historically hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics. It is now thought to slow healing and lead to scarring because it destroys newly formed skin cells. Only a very low concentration of H2O2 can induce healing, and only if not repeatedly applied. Surgical use can lead to gas embolism formation. Despite this it is still used for wound treatment in many developing countries.
It is absorbed by skin upon contact and creates a local capillary embolism that appears as a temporary whitening of the skin.
2 (between 1.9% and 12%) mixed with ammonium hydroxide is used to bleach human hair. The chemical's bleaching property lends its name to the phrase "peroxide blonde". Hydrogen peroxide is also used for tooth whitening and can be mixed with baking soda and salt to make a home-made toothpaste.
Hydrogen peroxide may be used to treat acne, although benzoyl peroxide is a more common treatment.
Use in alternative medicine
Practitioners of alternative medicine have advocated the use of hydrogen peroxide for the treatment of various conditions, including emphysema, influenza, AIDS and in particular cancer. The practice calls for the daily consumption of hydrogen peroxide, either orally or by injection and is, in general, based around two precepts. First, that hydrogen peroxide is naturally produced by the body to combat infection; and second, that human pathogens (including cancer: See Warburg hypothesis) are anaerobic and cannot survive in oxygen-rich environments. The ingestion or injection of hydrogen peroxide is therefore believed to kill disease by mimicking the immune response in addition to increasing levels of oxygen within the body. This makes it similar to other oxygen-based therapies, such as ozone therapy and hyperbaric oxygen therapy.
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 by 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 which also has been proposed as a cause of cancer. Claims that hydrogen peroxide therapy increase 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 should also be noted that it is 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. Intravenous injection of hydrogen peroxide has been linked to several deaths.
The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment." Furthermore, the therapy is not approved by the U.S. FDA.
Rocket-belt hydrogen-peroxide propulsion system used in a jet pack
2 is referred to as "high-test peroxide" (HTP). It can be used either as a monopropellant (not mixed with fuel) or as 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 (1,112 °F), which is expelled through a nozzle, generating thrust. H
2 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. 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. 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.
As a bipropellant, H
2 is decomposed to burn a fuel as an oxidizer. 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, noncryogenic and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle. It can 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 for 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.
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 sinkings of HMS Sidon and the Russian submarine Kursk. 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.
Hydrogen peroxide has been used for creating organic peroxide-based explosives, such as acetone peroxide, for improvised explosive devices. Hydrogen peroxide itself also acts as an explosive in high concentrations when placed in an absorbent, and has been used in attacks including the 7 July 2005 London bombings. These explosives tend to degrade quickly and hence are not used as commercial or military explosives.
Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.
- Glow sticks
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.
Some horticulturalists 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.
- Fish aeration
Laboratory tests conducted by fish culturists in recent years have demonstrated that common household hydrogen peroxide can 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.
Regulations vary, but low concentrations, such as 6%, 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 are typically accompanied by a Material Safety Data Sheet (MSDS). 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 H
2 will react violently.
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.
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). Because 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.
Hydrogen peroxide, either in pure or diluted form, can pose several risks, the main one being that it forms explosive mixtures upon contact with organic compounds. Highly concentrated hydrogen peroxide itself is unstable and can cause a boiling liquid expanding vapour explosion (BLEVE) of the remaining liquid. Distillation of hydrogen peroxide at normal pressures is thus highly dangerous. It is also corrosive, especially when concentrated, but even domestic-strength solutions can cause irritation to the eyes, mucous membranes and skin. Swallowing hydrogen peroxide solutions is particularly dangerous, as decomposition in the stomach releases large quantities of gas (10 times the volume of a 3% solution), leading to internal bloating. Inhaling over 10% can cause severe pulmonary irritation.
With a significant vapour pressure (1.2 kPa at 50 °C), hydrogen-peroxide vapour is potentially hazardous. According to U.S. NIOSH, the immediately dangerous to life and health (IDLH) limit is only 75 ppm. 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). Hydrogen peroxide has also been classified by the American Conference of Governmental Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans". 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 and from the ATSDR.
- 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.
- During the Second World War, doctors in German concentration camps experimented with the use of hydrogen peroxide injections in the killing of human subjects.
- Several people received minor injuries after a hydrogen-peroxide spill on board of a flight between the U.S. cities Orlando and Memphis on 28 October 1998.
- 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 rust in the torpedo casing. The vessel was lost with all hands. A similar incident was responsible for the loss of HMS Sidon in 1955.
- 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.
- ^ Easton, M. F.; Mitchell, A. G.; Wynne-Jones, W. F. K. (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. doi:10.1039/TF9524800796.
- ^ a b c d "NIOSH Pocket Guide to Chemical Hazards #0335". National Institute for Occupational Safety and Health (NIOSH).
- ^ a b c "Hydrogen peroxide". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
- ^ Hill, C. N. (2001). A Vertical Empire: The History of the UK Rocket and Space Programme, 1950–1971. Imperial College Press. ISBN 978-1-86094-268-6.
- ^ Brauer, Georg, ed. (1963). Handbook of preparative inorganic chemistry 1. Translation editing by Reed F. (2nd ed.). New York, N.Y.: Academic Press. p. 140. ISBN 978-0-12-126601-1.
- ^ "Hydrogen Peroxide Technical Library" (PDF). Retrieved 3 March 2016.
- ^ Hunt, Robert H.; Leacock, Robert A.; Peters, C. Wilbur; Hecht, Karl T. (1965). "Internal-Rotation in Hydrogen Peroxide: The Far-Infrared Spectrum and the Determination of the Hindering Potential" (PDF). The Journal of Chemical Physics 42 (6): 1931. Bibcode:1965JChPh..42.1931H. doi:10.1063/1.1696228.
- ^ Dougherty, Dennis A.; Anslyn, Eric V. (2005). Modern Physical Organic Chemistry. University Science. p. 122. ISBN 1-891389-31-9.
- ^ Abrahams, S. C.; Collin, R. L.; Lipscomb, W. N. (1 January 1951). "The crystal structure of hydrogen peroxide". Acta Crystallographica 4 (1): 15–20. doi:10.1107/S0365110X51000039.
- ^ Thénard, L. J. (1818). "Observations sur des nouvelles combinaisons entre l'oxigène et divers acides". Annales de chimie et de physique. 2nd series 8: 306–312.
- ^ C. W. Jones, J. H. Clark. Applications of Hydrogen Peroxide and Derivatives. Royal Society of Chemistry, 1999.
- ^ Wolffenstein, Richard (October 1894). "Concentration und Destillation von Wasserstoffsuperoxyd". Berichte der deutschen chemischen Gesellschaft (in German) 27 (3): 3307–3312. doi:10.1002/cber.189402703127.
- ^ G. Carrara (1892) "Sul peso molecolare e sul potere rifrangente dell' acqua ossigenata" (On the molecular weight and on the refractive power of oxygenated water [i.e., hydrogen peroxide]), Atti della Reale Accademia dei Lincei, series 5, 1 (2) : 19–24.
Carrara's findings were confirmed by: W. R. Orndorff and John White (1893) "The molecular weight of hydrogen peroxide and of benzoyl peroxide," American Chemical Journal, 15 : 347–356.
- ^ See, for example:
- In 1882, Kingzett proposed as a structure H2O=O. See: Charles Thomas Kingzett (29 September 1882) "On the activity of oxygen and the mode of formation of hydrogen dioxide," The Chemical News, 46 (1192): 141–142.
- 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.
- W. C. Schumb, C. N. Satterfield, and R. L. Wentworth (1 December 1953) "Report no. 43: Hydrogen peroxide, Part two", Office of Naval Research, Contract No. N5ori-07819 On p. 178, the authors present six hypothetical models 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).)
- ^ See:
- W. G. Penney and G. B. B. M. Sutherland (1934) "The theory of the structure of hydrogen peroxide and hydrazine", Journal of Chemical Physics, 2 (8): 492–498.
- W. G. Penney and G. B. B. M. Sutherland (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.
- ^ a b Campos-Martin, Jose M.; Blanco-Brieva, Gema; Fierro, Jose L. G. (2006). "Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process". Angewandte Chemie International Edition 45 (42): 6962–6984. doi:10.1002/anie.200503779. PMID 17039551.
- ^ H. Riedl and G. Pfleiderer, U.S. Patent 2,158,525 (2 October 1936 in USA, and 10 October 1935 in Germany) to I. G. Farbenindustrie, Germany
- ^ Noritaka Mizuno Gabriele Centi, Siglinda Perathoner, Salvatore Abate "Direct Synthesis of Hydrogen Peroxide: Recent Advances" in Modern Heterogeneous Oxidation Catalysis: Design, Reactions and Characterization 2009, Wiley-VCH. doi:10.1002/9783527627547.ch8
- ^ Edwards, Jennifer K.; Solsona, Benjamin; N, Edwin Ntainjua; Carley, Albert F.; Herzing, Andrew A.; Kiely, Christopher J.; Hutchings, Graham J. (20 February 2009). "Switching Off Hydrogen Peroxide Hydrogenation in the Direct Synthesis Process". Science 323 (5917): 1037–1041. Bibcode:2009Sci...323.1037E. doi:10.1126/science.1168980. PMID 19229032.
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