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Hydrogen peroxide is a chemical compound with the formula H
. It is the simplest peroxide (a compound with an oxygen-oxygen single bond) and in its pure form is a colorless liquid, slightly more viscous than water. For safety reasons it is normally encountered as an aqueous solution, also colorless. Hydrogen peroxide is a strong oxidizer and is used as a 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.[1]

Organisms naturally produce trace quantities of hydrogen peroxide, most notably by a respiratory burst as part of the immune response.


  • Structure and properties 1
    • Properties 1.1
      • In aqueous solutions 1.1.1
    • Structure 1.2
    • Comparison with analogues 1.3
  • Discovery 2
  • Manufacture 3
    • Availability 3.1
  • Reactions 4
    • Decomposition 4.1
    • Redox reactions 4.2
    • Organic reactions 4.3
    • Precursor to other peroxide compounds 4.4
  • Biological function 5
  • Applications 6
    • Industrial 6.1
    • Medical 6.2
      • Disinfectant 6.2.1
      • Cosmetic applications 6.2.2
      • Use in alternative medicine 6.2.3
    • Propellant 6.3
      • Explosives 6.3.1
    • Other uses 6.4
  • Safety 7
    • Historical incidents 7.1
  • See also 8
  • References 9
  • External links 10

Structure and properties


The boiling point of H
has been extrapolated as being 150.2 °C, approximately 50 degrees higher than water; however, in practice hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled under reduced pressure via a variety of techniques.[2]

In aqueous solutions

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 mixture is also depressed in relation with the median of both boiling points (125.1°C). It occurs at 114°C. This boiling point is 14° greater than that of pure water and 36.2° less than that of pure hydrogen peroxide.[3]

Phase diagram of H
and water: Area above blue line is liquid. Dotted lines separate solid+liquid phases from solid+solid phases.
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Hydrogen peroxide

i/ATC_code_D08" id="whe_lnki_15" title="ATC code D08">D08, S02
Jmol-3D images Image 1
Molecular formula H2O2
Molar mass 34.0147 g/mol
Appearance Very light blue color; colorless in solution
Odor slightly sharp
Density 1.135 g/cm3 (20 °C, 30-percent)
1.450 g/cm3 (20 °C, pure)
Melting point −0.43 °C (31.23 °F; 272.72 K)
Boiling point 150.2 °C (302.4 °F; 423.3 K) (decomposes)
Solubility in water Miscible
Solubility soluble in ether, alcohol
insoluble in petroleum ether
Acidity (pKa) 11.75
Refractive index (nD) 1.4061
Viscosity 1.245 cP (20 °C)
Dipole moment 2.26 D
heat capacity
1.267 J/g K (gas)
2.619 J/g K (liquid)
Std enthalpy of
-187.80 kJ/mol
MSDS ICSC 0164 (>60% soln.)
EU Index 008-003-00-9
EU classification Oxidant (O)
Corrosive (C)
Harmful (Xn)
R-phrases R5, R8, R20/22, R35
S-phrases (S1/2), S17, S26, S28, S36/37/39, S45
NFPA 704
Flash point Non-flammable
LD50 1518 mg/kg
Related compounds
Related compounds Water
Hydrogen disulfide
Dioxygen difluoride
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 YesY   YesY/N?)
Density of aqueous solution of H2O2
H2O2 (v/v) Density (g/cm3) Temperature (°C)
3% 1.0095 15
27% 1.10 20
35% 1.13 20
50% 1.20 20
70% 1.29 20
75% 1.33 20
96% 1.42 20
98% 1.43 20
100% 1.450 20


Hydrogen peroxide (H
), is a nonplanar molecule possessing (twisted) C2 symmetry. Although the O−O bond is a single bond, the molecule has a relatively high barrier to rotation of 2460 cm-1 (29.45 kJ/mol);[4] 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
are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state.[5] Crystals of H
are tetragonal with the space group D_4^4 P4_12_1.[6]

O-O bond length = 147.4 pm O-H bond length = 95.0 pm
Structure and dimensions of H2O2 in the gas phase
O-O bond length = 145.8 pm O-H bond length = 98.8 pm
Structure and dimensions of H2O2 in the solid (crystalline) phase
Properties of H2O2 and its analogues
values marked * are extrapolated
Name Formula Molar mass (g mol−1) Mpt (°C) Bpt (°C)
Hydrogen peroxide HOOH 34.01 −0.43 150.2*
Water HOH 18.02 0.00 99.98
Hydrogen disulfide HSSH 66.15 −89.6 70.7
Hydrazine H2NNH2 32.05 2 114
Hydroxylamine NH2OH 33.03 33 58*
Diphosphane H2PPH2 65.98 −99 63.5*

Comparison with analogues

Hydrogen peroxide has a number of structural analogues with Hm-E-E-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 do not exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide as well. 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.[7] 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.[8]

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. However, this instability was due to traces of impurities (transition metals 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 via vacuum distillation.[9]

Determination of the molecular structure of hydrogen peroxide proved to be very difficult. It was only in 1892 that the Italian physical chemist Giacomo Carrara (1864–1925) determined its molecular weight by freezing point depression, which confirmed that its molecular formula is H2O2.[10] At least half a dozen hypothetical molecular structures seemed to be consistent with the available evidence.[11] In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide which was very similar to the presently accepted one and which subsequent evidence cumulatively proved to be correct.[12]


Previously, hydrogen peroxide has been prepared industrially by hydrolysis of the ammonium peroxydisulfate, which was itself obtained via the electrolysis of a solution of ammonium bisulfate (NH
) in sulfuric acid.

(NH4)2S2O8 + 2 H2O → H2O2 + 2 (NH4)HSO4

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 via 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.[13][14]

Hydrogen peroxide production with the Riedl-Pfleiderer process

The simplified overall equation for the process is deceptively simple:[13]

+ O

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, in terms of thermodynamics, the reaction of hydrogen with oxygen favours production of water. Systems for direct synthesis have been developed; most of which are based around finely dispersed metal catalysts.[15][16] However none of these have 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 >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
was around 1.9 million tonnes and grew to 2.2 million in 2006,[17] most of which was at a concentration of 70% or less. In that year bulk 30% H
sold for around US $0.54 per kg, equivalent to US $1.50 per kg (US $0.68 per lb) on a "100% basis".[18][19]



Hydrogen peroxide is thermodynamically unstable and decomposes to form water and oxygen with a ΔHo of −98.2 kJ·mol−1 and a ΔS of 70.5 J·mol−1·K−1.

2 H
→ 2 H
+ 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).[20] Certain metal ions, such as Fe2+
or Ti3+
, 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 concentrations of hydrogen peroxide on an inflammable substance can cause an immediate fire.

Redox reactions

Hydrogen peroxide exhibits oxidizing and reducing properties, depending on pH.

In acidic solutions, H
is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H
can be converted into hydroxyl radicals (OH), which are highly reactive.

Oxidant/Reduced product Oxidation potential, V
Fluorine/Hydrogen fluoride 3.0
Ozone/Oxygen 2.1
Hydrogen peroxide/Water 1.8
Potassium permanganate/Manganese dioxide 1.7
Chlorine dioxide/HClO 1.5
Chlorine/Chloride 1.4

In acidic solutions Fe2+
is oxidized to Fe3+
(hydrogen peroxide acting as an oxidizing agent),

2 Fe2+
(aq) + H
+ 2 H+
(aq) → 2 Fe3+
(aq) + 2 H

and sulfite (SO2−
) is oxidized to sulfate (SO2−
). However, potassium permanganate is reduced to Mn2+
by acidic H
. Under alkaline conditions, however, some of these reactions reverse; for example, Mn2+
is oxidized to Mn4+
(as MnO

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
+ NaCl + H
2 KMnO
+ 3 H
→ 2 MnO
+ 2 KOH + 2 H
+ 3 O

Organic reactions

Hydrogen peroxide is frequently used as an oxidizing agent. Illustrative is oxidation of thioethers to sulfoxides.[21][22]

+ H
→ Ph−S(O)−CH
+ 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
) form 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:

+ 4 H
+ 2 NaOH → 2 Na
+ H

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 it interacts with ozone to form hydrogen trioxide, also known as 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
in some reactions.

Biological function

Hydrogen peroxide is also one of the two chief chemicals in the defense 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 the hydrogen peroxide presence inside cells increased after tissues are damaged in zebra fish, 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.[23][24]

Hydrogen peroxide has important roles as a signaling molecule in the regulation of a wide variety of biological processes.[25] 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.[26] These hydroxyl radicals in turn readily react with and damage vital cellular components, especially those of the mitochondria.[27] At least one study has also tried to link hydrogen peroxide production to cancer.[28] 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.[29]



About 60% of the world's production of hydrogen peroxide is used for pulp- and paper-bleaching.[17] 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 dibenzoyl peroxide being a high volume example. It 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 aromatic or halogenated compounds. It can also oxidize sulphur based compounds present in the waste; which is beneficial as it generally reduces their odour.[31]



Skin shortly after exposure to 35% H

Hydrogen peroxide is seen as an environmentally safe alternative to chlorine-based bleaches. It can be used for the disinfection of various surfaces[32] and is generally recognized as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA).[33] However studies have found it to be ineffective in certain cases and hospitals and other medical institutions are now being advised to use chlorine-based bleaches for disinfection.[34]

Historically, hydrogen peroxide was commonly 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.[35] Only a very low concentration of H2O2 can induce healing, and only if not repeatedly applied.[36] Surgical use can lead to gas embolism formation.[37]

It is absorbed by skin upon contact and creates a local capillary embolism that appears as a temporary whitening of the skin.[38]

Cosmetic applications

Diluted H
(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".[39] Hydrogen peroxide is also used for tooth whitening and can be mixed with baking soda and salt to make a home-made toothpaste.[40]

Hydrogen peroxide may be used to treat acne,[41] 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.[42] The practise calls for the daily consumption of hydrogen peroxide, either orally or by injection and is, in general, based around 2 precepts. Firstly that hydrogen peroxide is naturally produced by the body to combat infection. Secondly, 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 disputed by mainstream scientists. 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.[43] 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.

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.[44] Intravenous injection of hydrogen peroxide has been linked to several deaths.[45][46][47]

The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment"[48] The therapy is not approved by the U.S. FDA.


Rocket Belt hydrogen peroxide propulsion system used in a jet pack

High concentration H
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
monopropellant produces a maximum specific impulse (Isp) of 161 s (1.6 kN·s/kg), which makes it a low-performance monopropellant. Peroxide generates much less thrust than hydrazine. The Bell Rocket Belt used hydrogen peroxide monopropellant.

As a bipropellant H
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 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 Walter 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, but this was dangerous and has been discontinued by most navies. Hydrogen peroxide leaks were blamed for the sinkings of HMS Sidon and the Russian submarine Kursk. It was discovered, for example, by the Japanese Navy in torpedo trials, that the concentration of H
in right-angle bends in HTP pipework can often lead to explosions in submarines and torpedoes. 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.[49][50]

While rarely used now as a monopropellant for large engines, small hydrogen peroxide attitude control thrusters are still in use on some satellites. They are easy to throttle, and safer to fuel and handle before launch than hydrazine thrusters. However, hydrazine is more often used in spacecraft because of its higher specific impulse and lower rate of decomposition.


Hydrogen peroxide has been used for creating acetone peroxide, for improvised explosive devices, including the 7 July 2005 London bombings.[51] These explosives tend to degrade quickly and hence are not used as commercial or military explosives.

Other uses

Chemiluminescence of cyalume, as found in a glow stick

Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.

Glow sticks

Hydrogen peroxide reacts with certain esters, such as cyalume and phenyl oxalate ester, 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.[52][53][54]

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.[55][56]


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
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 , of concentrated hydrogen peroxide.

Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances.[57] 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).[58] 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 filter out light.[59]

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.[60] Highly concentrated hydrogen peroxide itself is unstable, and can then cause a boiling liquid expanding vapor 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.[61] 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 bleeding. Inhaling over 10% can cause severe pulmonary irritation.[62]

With a significant vapor pressure (1.2 kPa at 50 °C[CRC Handbook of Chemistry and Physics, 76th Ed, 1995–1996]), hydrogen peroxide vapor is potentially hazardous. According to U.S. NIOSH, the Immediately Dangerous to Life and Health (IDLH) limit is only 75 ppm.[63] The U.S. Occupational Safety and Health Administration (OSHA) has established a permissible exposure limit of 1.0 ppm calculated as an eight hour time weighted average (29 CFR 1910.1000, Table Z-1)[60] and 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."[64] Information on the hazards of hydrogen peroxide is available from OSHA[60] and from the ATSDR.[65]

Historical incidents

  • On 16 July 1934, in Kummersdorf, Germany, a rocket engine using hydrogen peroxide exploded, killing three people. As a result of this incident, Wernher von Braun decided not to use hydrogen peroxide as an oxidizer in the rockets he developed afterward.
  • Several people received minor injuries after a hydrogen peroxide spill on board a flight between the U.S. cities Orlando and Memphis on 28 October 1998.[67]
  • The Russian submarine K-141 Kursk sailed out to sea 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.[1]

See also



  1. ^ Hill, C. N. (2001). A Vertical Empire: The History of the UK Rocket and Space Programme, 1950–1971. Imperial College Press.  
  2. ^ Riley, edited by Georg Brauer ; translated by Scripta Technica, Inc. Translation editor Reed F. (1963). Handbook of preparative inorganic chemistry. Volume 1 (2nd ed. ed.). New York, N.Y.: Academic Press. p. 140.  
  3. ^ 60% hydrogen peroxide msds 50% H2O2 MSDS
  4. ^ 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". The Journal of Chemical Physics 42 (6): 1931.  
  5. ^ Dougherty, Dennis A.; Eric V. Anslyn (2005). Modern Physical Organic Chemistry. University Science. p. 122.  
  6. ^ Abrahams, S. C.; Collin, R. L.; Lipscomb, W. N. (1 January 1951). "The crystal structure of hydrogen peroxide". Acta Crystallographica 4 (1): 15–20.  
  7. ^ L. J. Thénard (1818). "Observations sur des nouvelles combinaisons entre l’oxigène et divers acides".  
  8. ^ C. W. Jones, J. H. Clark. Applications of Hydrogen Peroxide and Derivatives. Royal Society of Chemistry, 1999.
  9. ^ Wolffenstein, Richard (October 1894). "Concentration und Destillation von Wasserstoffsuperoxyd". Berichte der deutschen chemischen Gesellschaft (in German) 27 (3): 3307–3312.  
  10. ^ 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.
  11. ^ See, for example:
    • In 1882, Kingzett proposed as a structure H2O=O. See: Charles Thomas Kingzett (September 29, 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), pages 952-956.
    • W. C. Schumb, C.N. Satterfield, and R.L. Wentworth (December 1, 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).)
  12. ^ 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.
  13. ^ a b Jose M. Campos-Martin, Gema Blanco-Brieva, Jose L. G. Fierro (2006). "Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process". Angewandte Chemie International Edition 45 (42): 6962–6984.  
  14. ^ 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
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  • J. Drabowicz, et al. (1994). G. Capozzi, et al., ed. The Syntheses of Sulphones, Sulphoxides and Cyclic Sulphides. Chichester UK: John Wiley & Sons. pp. 112–6.  
  • N.N. Greenwood, A. Earnshaw (1997). Chemistry of the Elements (2nd ed.). Oxford UK: Butterworth-Heinemann.  A great description of properties & chemistry of H
  • J. March (1992). Advanced Organic Chemistry (4th ed.). New York: Wiley. p. 723. 

External links

  • Hydrogen Peroxide at The Periodic Table of Videos (University of Nottingham)
  • Material Safety Data Sheet
  • ATSDR Agency for Toxic Substances and Disease Registry FAQ
  • International Chemical Safety Card 0164
  • NIOSH Pocket Guide to Chemical Hazards
  • Process flow sheet of Hydrogen Peroxide Production by anthrahydroquinone autoxidation
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