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Hydrogen peroxide

Hydrogen Peroxide776
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Hydrogen peroxide is a chemical compound with the formula H22O.

In its pure form, it is a very pale blue liquid, slightly more viscous than water. It is used as an oxidizer, bleaching agent, and antiseptic. Concentrated hydrogen peroxide, or "high-test peroxide", is a reactive oxygen species and has been used as a propellant in rocketry. Its chemistry is dominated by the nature of its unstable peroxide bond.

Hydrogen peroxide is the simplest peroxide (a compound with an oxygen–oxygen single bond). It is unstable and slowly decomposes in the presence of light. Because of its instability, hydrogen peroxide is typically stored with a stabilizer in a weakly acidic solution in a dark-colored bottle.

Hydrogen peroxide is found in biological systems including the human body. Enzymes that use or decompose hydrogen peroxide are classified as peroxidases.

Alexander von Humboldt reported one of the first synthetic peroxides, barium peroxide, in 1799 as a by-product of his attempts to decompose air.

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.

Today, the term "oxygenated water" may appear on retail packaging referring to mixtures containing either water and hydrogen peroxide or water and dissolved oxygen. This could cause personal injury if the difference is not properly understood by the user.

An improved version of Thénard's process used hydrochloric acid, followed by the 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.

Thénard and Joseph Louis Gay-Lussac synthesized sodium peroxide in 1811. The bleaching effect of peroxides and their salts on natural dyes became known around that time, 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.

Pure hydrogen peroxide was long believed to be unstable, as all early attempts to separate it from the water, which is present during synthesis, 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 Wolfenstein, 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 ammonium persulfate, which was itself obtained by the electrolysis of a solution of ammonium bisulfate (NH4HSO4) in sulfuric acid.


Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which the German chemical company BASF developed 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.

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 regenerate the anthraquinone.

Most commercial processes achieve oxidation by bubbling compressed air through a solution of the anthrahydroquinone, with the hydrogen peroxide then extracted from the solution and the anthraquinone recycled back for successive cycles of hydrogenation and oxidation.

The economics of the process depends heavily on effective recycling of the extraction solvents, the hydrogenation catalyst, and the expensive quinone.

Other sources

Small, but detectable, amounts of hydrogen peroxide can be formed by several methods. Small amounts are formed by electrolysis of dilute acid around the cathode where hydrogen evolves if oxygen is bubbled around it.

It is also produced by exposing water to ultraviolet rays from a mercury lamp, or an electric arc while confining it in a UV transparent vessel (e.g. quartz).

It is detectable in ice water after burning a hydrogen gas stream aimed towards it and is also detectable on floating ice. Rapidly cooling humid air blown through an approximately 2,000 °C spark gap results in detectable amounts.

A commercially viable process to produce hydrogen peroxide directly from the environment has been of interest for many years. Efficient direct synthesis is difficult to achieve, as the reaction of hydrogen with oxygen thermodynamically favors the production of water.

Systems for direct synthesis have been developed, most of which employ finely dispersed metal catalysts similar to those used for hydrogenation of organic substrates. None of these has yet reached a point where they can be used for industrial-scale synthesis.


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 H2O2 was around 1.9 million tonnes and grew to 2.2 million in 2006, most of which was at a concentration of 70% or less. In that year, bulk 30% H2O2 sold for around 0.54 USD/kg, equivalent to US$1.50/kg (US$0.68/lb) on a "100% basis".

Hydrogen peroxide occurs in surface water, groundwater, and the atmosphere. It forms upon illumination or natural catalytic action by substances contained in water. Seawater contains 0.5 to 14 ?g/L of hydrogen peroxide, freshwater 1 to 30 ?g/L and air 0.1 to 1 parts per billion.


Hydrogen peroxide is thermodynamically unstable and decomposes to form water and oxygen with a ?Ho of –2884.5 kJ/kg and a ?S of 70.5 J/(mol·K):

The rate of decomposition increases with rising in temperature, concentration, and pH, with cool, dilute, acidic solutions showing the best stability. Decomposition is catalyzed by various compounds, including most transition metals and their compounds (e.g. manganese dioxide (MnO2), silver, and platinum).

Certain metal ions, such as Fe2+ or Ti3+, can cause the decomposition to take a different path, with free radicals such as the hydroxyl radical (HO·) and hydroperoxyl (HOO·) being formed. Non-metallic catalysts include potassium iodide, which reacts particularly rapidly and forms the basis of the elephant toothpaste demonstration.

Hydrogen peroxide can also be decomposed biologically by the 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.



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.

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.

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.

Production of organic compunds

It is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example. It is used in polymerizations, as a flour bleaching agent, and as a treatment for acne. Peroxy acids, such as peracetic acid and meta-chloroperoxybenzoic acid also are produced using hydrogen peroxide.

Hydrogen peroxide has been used for creating organic peroxide-based explosives, such as acetone peroxide.


Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. In advanced oxidation processing, the Fenton reaction gives the highly reactive hydroxyl radical (·OH).

This degrades organic compounds, including those that are ordinarily robust, 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 odor.

Hydrogen peroxide may be used for the sterilization of various surfaces, including surgical tools, and maybe deployed as a vapor (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 may increase tolerance in the presence of lower concentrations.

Lower levels of concentration (3%) will work against most spores; higher concentrations (7 to 30%) and longer contact times will improve 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).

Hydrogen peroxide may be used to treat acne, although benzoyl peroxide is a more common treatment.

Removal of bloodstains

Hydrogen peroxide reacts with blood as a bleaching agent, and so if a bloodstain is fresh, or not too old, liberal application of hydrogen peroxide, if necessary in more than a single application, will bleach the stain fully out. After about two minutes of the application, all residue should be firmly blotted out. Repeat if necessary.

Niche uses

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".

Hydrogen peroxide is also used for tooth whitening. It may be found in most whitening toothpaste. Hydrogen peroxide has shown positive results involving teeth lightness and chroma shade parameters.

It works by oxidizing colored pigments onto the enamel where the shade of the tooth may become lighter. Hydrogen peroxide may be mixed with baking soda and salt to make homemade toothpaste.


High-concentration H2O2 is referred to as "high-test peroxide" (HTTP). 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. 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.

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, H2O2 is decomposed to burn 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, non-cryogenic, and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle.

It may also be used for regenerative cooling of rocket engines. Peroxide was used very successfully as an oxidizer in World War II German rocket motors (e.g. T-Stoff, containing oxyquinoline stabilizer, for both the Walter HWK 109-500 Starthilfe RATO externally podded monopropellant booster system, and 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 an oxidizer or propellant.

Operator error in the use of hydrogen-peroxide torpedoes was named as a possible cause for the sinking 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.

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 solutions in watering solutions. It's 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.


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 cool water and warm water finfish, for control of external columnaris disease in warm-water finfish, and for control of Gyrodactylus spp. in freshwater-reared salmonids.

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.


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 an H22O, 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 aluminum 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, may pose several risks, the main one being that it forms explosive mixtures upon contact with organic compounds.

Highly concentrated hydrogen peroxide is unstable and may cause a boiling liquid expanding vapor explosion (BLEVE) of the remaining liquid. Consequently, a distillation of hydrogen peroxide at normal pressures is highly dangerous.

It is also corrosive, especially when concentrated, but even domestic-strength solutions may 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 (ten times the volume of a 3% solution), leading to internal bloating. Inhaling over 10% can cause severe pulmonary irritation.

With a significant vapor pressure (1.2 kPa at 50 °C), the 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 also has 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 vapors, continuous monitors for hydrogen peroxide should be used. Information on the hazards of hydrogen peroxide is available from OSHA and from the ATSDR.

Adverse effects on wounds

Historically hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics. Now it is thought to inhibit healing and to induce scarring because it destroys newly formed skin cells.

One study found that only very low concentrations (0.03% solution, this is a dilution of typical 3% Peroxide by 100 times) may induce healing, and only if not applied repeatedly. A 0.5% solution was found to impede healing. Surgical use can lead to gas embolism formation.

Despite this, it is still used for wound treatment in many countries, and, in the United States, is prevalent as a major first aid antiseptic.

Dermal exposure to dilute solutions of hydrogen peroxide causes whitening or bleaching of the skin due to microembolism caused by oxygen bubbles in the capillaries.

Use in alternative medicine

Practitioners of alternative medicine have advocated the use of hydrogen peroxide for various conditions, including emphysema, influenza, AIDS, and in particular cancer. There is no evidence of effectiveness and in some cases, it has proved fatal.

The practice calls for the daily consumption of hydrogen peroxide, either orally or by injection, and is based on two precepts. First, that hydrogen peroxide is produced naturally 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, therefore, is believed to kill disease by mimicking the immune response in addition to increasing levels of oxygen within the body. This makes the practice similar to other oxygen-based therapies, such as ozone therapy and hyperbaric oxygen therapy.

Both the effectiveness and safety of hydrogen peroxide therapy are 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.

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 tumor, 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.

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.

During the Second World War, doctors in German concentration camps experimented with the use of hydrogen peroxide injections in the killing of human subjects.

In April 1992, an explosion occurred at the hydrogen peroxide plant at Jarrie in France, due to technical failure of the computerized control system and resulting in one fatality and wide destruction of the plant.

Several people received minor injuries after a hydrogen peroxide spill onboard a flight between the U.S. cities of 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 a propellant for the torpedo, seeped through its container, damaged either by rust or in the loading procedure back on land where an incident involving one of the torpedoes accidentally touching ground went unreported. 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.