Hydrogen peroxide is a chemical compound with the formula H
2. In its pure form, it is a pale blue, clear liquid, slightly more
viscous than water.
Hydrogen peroxide is the simplest peroxide (a
compound with an oxygen–oxygen single bond). 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 unstable and slowly decomposes in the presence of
base or a catalyst. Because of its instability, hydrogen peroxide is
typically stored with a stabilizer in a weakly acidic solution.
Hydrogen peroxide is found in biological systems including the human
body. Enzymes that use or decompose hydrogen peroxide are classified
1.1 Aqueous solutions
1.3 Comparison with analogues
4.3 Organic reactions
4.4 Precursor to other peroxide compounds
5 Biological function
6.3 Production of organic compounds
6.4.1 Cosmetic applications
6.4.2 Use in alternative medicine
6.6 Other uses
7.1 Historical incidents
8 See also
10 External links
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
In aqueous solutions hydrogen peroxide differs from the pure material
due to the effects of hydrogen bonding between water and hydrogen
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. 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.
Phase diagram of H
2 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
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
Properties of H2O2 and its analogues
values marked * are extrapolated
Molar mass (g/mol)
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.
Alexander von Humboldt
Alexander von Humboldt synthesized one of the first synthetic
peroxides, barium peroxide, in 1799 as a by-product of his attempts to
Nineteen years later
Louis Jacques Thénard
Louis Jacques Thénard recognized that this
compound could be used for the preparation of a previously unknown
compound, which he described as oxidized water – subsequently known
as hydrogen peroxide. 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
Joseph Louis Gay-Lussac
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, and 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 implemented into
industry 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 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 autoxidation to
regenerate the starting anthraquinone, with hydrogen peroxide 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 groups), 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 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
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
A tank car designed for transporting hydrogen peroxide by rail
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
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
2 sold for around 0.54 USD/kg, equivalent to 1.50 USD/kg (0.68 USD/lb)
on a "100% basis".
Hydrogen peroxide occurs in surface water, groundwater and in the
atmosphere. It forms upon illumination or natural catalytic action by
substances contained in water. Sea water contains 0.5 to 14 μg/L
of hydrogen peroxide, freshwater 1 to 30 μg/L and air 0.1 to 1
parts per billion.
Hydrogen peroxide is thermodynamically unstable and decomposes to form
water and oxygen with a ΔHo of −98.2 kJ/mol and a ΔS of
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+
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 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.
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
2 can be converted into hydroxyl radicals (·OH), which are highly
Oxidation potential, V
potassium permanganate/manganese dioxide
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+ (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
2 → O
2 + NaCl + H
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
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 + H
4) 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.
reacts with acetone to form acetone peroxide and with ozone to form
Hydrogen peroxide forms stable adducts with urea (hydrogen
peroxide - urea), sodium carbonate (sodium percarbonate) and other
compounds. An acid-base adduct with triphenylphosphine oxide is a
useful "carrier" for H
2 in some reactions.
The peroxide anion is a stronger nucleophile than hydroxide and
displaces hydroxyl from oxyanions e.g. forming perborates and
Sodium perborate and sodium percarbonate are important
consumer and industrial bleaching agents; they stabilize hydrogen
peroxide and limit side reactions (e.g. reduction and decomposition
note below). The peroxide anion forms an adduct with urea, hydrogen
Hydrogen peroxide is both an oxidizing agent and reducing agent. The
oxidation of hydrogen peroxide by sodium hypochlorite yields singlet
oxygen. The net reaction of a ferric ion with hydrogen peroxide is a
ferrous ion and oxygen. This proceeds via single electron oxidation
and hydroxyl radicals. This is used in some organic chemistry
oxidations, e.g. in the Fenton's reagent. Only catalytic quantities of
iron ion is needed since peroxide also oxidizes ferrous to ferric ion.
The net reaction of hydrogen peroxide and permanganate or manganese
dioxide is manganous ion; however, until the peroxide is spent some
manganous ions are reoxidized to make the reaction catalytic. This
forms the basis for common monopropellant rockets.
Hydrogen peroxide is formed in human and animals as a short-lived
product in biochemical processes and is toxic to cells. The toxicity
is due to oxidation of proteins, membrane lipids and
DNA by the
peroxide ions. The class of biological enzymes called SOD
(superoxide dismutase) is developed in nearly all living cells as an
important antioxidant agent. They promote the disproportionation of
superoxide into oxygen and hydrogen peroxide, which is then rapidly
decomposed by the enzyme catalase to oxygen and water.
displaystyle ce 2O2^ - +2H+->[ atop ce SOD ] H2O2 +O2
Formation of hydrogen peroxide by superoxide dismutase (SOD)
Peroxisomes are organelles found in virtually all eukaryotic
cells. They are involved in the catabolism of very long chain
fatty acids, branched chain fatty acids, D-amino acids, polyamines,
and biosynthesis of plasmalogens, etherphospholipids critical for the
normal function of mammalian brains and lungs. Upon oxidation,
they produce hydrogen peroxide in the following process:
displaystyle ce R-CH2-CH2-CO-SCoA +O2->[ atop ce FAD ]
FAD = flavin adenine dinucleotide
Catalase, another peroxisomal enzyme, uses this H2O2 to oxidize other
substrates, including phenols, formic acid, formaldehyde, and alcohol,
by means of the peroxidation reaction:
displaystyle ce H2O2 + R'H2 -> R' + 2H2O
, 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. In addition, when excess H2O2 accumulates in the cell,
catalase converts it to H2O through this reaction:
displaystyle ce H2O2->[ atop ce CAT ] 1/2O2 +H2O
Another origin of hydrogen peroxide is the degradation of adenosine
monophosphate which yields hypoxanthine.
Hypoxanthine is then
oxidatively catabolized first to xanthine and then to uric acid, and
the reaction is catalyzed by the enzyme xanthine oxidase:
Degradation of hypoxanthine through xanthine to uric acid to form
Australian bombardier beetle
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.
Eggs of sea urchin, shortly after fertilization by a sperm, produce
hydrogen peroxide. It is then quickly dissociated to OH· radicals.
The radicals serve as initiator of radical polymerization, which
surrounds the eggs with a protective layer of polymer.
The bombardier beetle has a device which allows it to shoot corrosive
and foul-smelling bubbles at its enemies. The beetle produces and
stores hydroquinone and hydrogen peroxide, in two separate reservoirs
in the rear tip of its abdomen. When threatened, the beetle contracts
muscles that force the two reactants through valved tubes into a
mixing chamber containing water and a mixture of catalytic enzymes.
When combined, the reactants undergo a violent exothermic chemical
reaction, raising the temperature to near the boiling point of water.
The boiling, foul-smelling liquid partially becomes a gas (flash
evaporation) and is expelled through an outlet valve with a loud
Hydrogen peroxide is a signaling molecule of plant defense against
Hydrogen peroxide has 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. Sodium percarbonate, which is an adduct of sodium
carbonate and hydrogen peroxide, is the active ingredient in such
OxiClean and Tide laundry detergent. When dissolved in
water, it releases hydrogen peroxide and sodium carbonate:
2 Na2CO3·3 H2O2 → 2 Na2CO3 + 3 H2O2
Production of organic compounds
It is used in the production of various organic peroxides with
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 produced using hydrogen
Hydrogen peroxide has been used for creating organic
peroxide-based explosives, such as acetone peroxide.
Skin shortly after exposure to 35% H
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 odour.
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-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
Food and Drug Administration
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 inhibit healing and to induce
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.
Dermal exposure to dilute solutions of hydrogen peroxide cause
whitening or bleaching of the skin due to microembolism caused by
oxygen bubbles in the capillaries.
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. It can be found in most whitening toothpastes.
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 can indeed become
Hydrogen peroxide 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 various conditions, including emphysema,
AIDS and cancer, although there is no evidence of
effectiveness and in some cases it may even be
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
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 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.
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.
Further information: High-test peroxide
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
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
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
World War II
World War II German rocket motors (e.g. T-Stoff, containing
oxyquinoline stabilizer, for both the
Walter HWK 109-500
Walter HWK 109-500 Starthilfe
RATO externally podded monopropellant booster system, and for the
Walter HWK 109-509
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
Chemiluminescence of cyalume, as found in a glow stick
Hydrogen peroxide has various domestic uses, primarily as a cleaning
and disinfecting agent.
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.
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
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
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.
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
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 a flight between the U.S. cities of Orlando and Memphis on 28
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 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
FOX reagent, used to measure levels of hydrogen peroxide in biological
^ 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.
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Institute for Occupational Safety and Health (NIOSH).
^ a b c "
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T. (1965). "Internal-Rotation in
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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.
Peroxide (New York, New York: Reinhold Publishing
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text: authors list (link)
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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.
^ Soyuz using hydrogen peroxide propellant (
^ "Ways to use
Peroxide in the Garden". Using Hydrogen
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^ Bhattarai SP, Su N, Midmore DJ; Su; Midmore (2005). "Oxygation
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Hydrogen Peroxide". Retrieved 3 March
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Process flow sheet of
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Peroxide Handbook by Rocketdyne
Antiseptics and disinfectants (D08)
Biguanides and amidines
Phenol and derivatives
Quaternary ammonium compounds
Propanol (propyl alcohol)
Isopropanol (isopropyl alcohol)
Ethanol (ethyl alcohol)#
‡Withdrawn from market
§Never to phase III
Stomatological preparations (A01)
Infection and antiseptics
Drugs used for diseases of the ear (S02)
Aluminium triacetate (Burow's solution)
Analgesics and anesthetics
Human hair color
Brown (varieties: Chestnut • Auburn)
Red (varieties: Auburn • Titian)
Hair dye stripping
Disappearing blonde gene
Melanocortin 1 receptor
Molecules detected in outer space
Magnesium monohydride cation
Hydrogen cyanide (HCN)
Hydrogen isocyanide (HNC)
Protonated molecular hydrogen
Protonated carbon dioxide
Protonated hydrogen cyanide
Buckminsterfullerene (C60 fullerene, buckyball)
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Atomic and molecular astrophysics
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Earliest known life forms
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Molecules in stars
Nexus for Exoplanet System Science
PAH world hypothesis
Polycyclic aromatic hydrocarbon
Polycyclic aromatic hydrocarbon (PAH)
RNA world hypothesis
TRP channel modulators
Sanshool (ginger, Sichuan and melegueta peppers)
Allyl isothiocyanate (mustard, radish, horseradish, wasabi)
CR gas (dibenzoxazepine; DBO)
CS gas (2-chlorobenzal malononitrile)
Farnesyl thiosalicylic acid
Ligustilide (celery, Angelica acutiloba)
Linalool (Sichuan pepper, thyme)
Methyl salicylate (wintergreen)
Oleocanthal (olive oil)
Paclitaxel (Pacific yew)
Polygodial (Dorrigo pepper)
Shogaols (ginger, Sichuan and melegueta peppers)
Thiopropanal S-oxide (onion)
Umbellulone (Umbellularia californica)
Adhyperforin (St John's wort)
Hyperforin (St John's wort)
Cooling Agent 10
Rutamarin (Ruta graveolens)
Steviol glycosides (e.g., stevioside) (Stevia rebaudiana)
Sweet tastants (e.g., glucose, fructose, sucrose; indirectly)
Rutamarin (Ruta graveolens)
Triptolide (Tripterygium wilfordii)
Sanshool (ginger, Sichuan and melegueta peppers)
Bisandrographolide (Andrographis paniculata)
Camphor (camphor laurel, rosemary, camphorweed, African blue basil,
Capsaicin (chili pepper)
Carvacrol (oregano, thyme, pepperwort, wild bergamot, others)
Dihydrocapsaicin (chili pepper)
Eugenol (basil, clove)
Evodiamine (Euodia ruticarpa)
Homocapsaicin (chili pepper)
Homodihydrocapsaicin (chili pepper)
Low pH (acidic conditions)
Nonivamide (PAVA) (PAVA spray)
Nordihydrocapsaicin (chili pepper)
Paclitaxel (Pacific yew)
Phorbol esters (e.g., 4α-PDD)
Piperine (black pepper, long pepper)
Polygodial (Dorrigo pepper)
Rutamarin (Ruta graveolens)
Resiniferatoxin (RTX) (Euphorbia resinifera/pooissonii)
Shogaols (ginger, Sichuan and melegueta peppers)
Thymol (thyme, oregano)
Tinyatoxin (Euphorbia resinifera/pooissonii)
Cannabigerolic acid (cannabis)
See also: Receptor/signaling modulators • Ion channel modulators