# Hydrogen peroxide

Hydrogen peroxide

Ball stick model of the hydrogen peroxide molecule
Names
IUPAC name
Hydrogen peroxide
Other names
Dioxidane
Oxidanyl
Perhydroxic acid
0-hydroxyol
Dihydrogen dioxide
Oxygenated water
Peroxaan
Identifiers
3D model (
JSmol
)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard
100.028.878
EC Number
• 231-765-0
IUPHAR/BPS
KEGG
RTECS number
• MX0900000 (>90% soln.)
MX0887000 (>30% soln.)
UNII
UN number 2015 (>60% soln.)
2014 (20–60% soln.)
2984 (8–20% soln.)
• InChI=1S/H2O2/c1-2/h1-2H
Key: MHAJPDPJQMAIIY-UHFFFAOYSA-N
• InChI=1/H2O2/c1-2/h1-2H
Key: MHAJPDPJQMAIIY-UHFFFAOYAL
• OO
Properties
H2O2
Molar mass 34.0147 g/mol
Appearance Very light blue liquid
Odor slightly sharp
Density 1.11 g/cm3 (20 °C, 30% (w/w) solution)[1]
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)
Miscible
Solubility soluble in ether, alcohol
insoluble in petroleum ether
log P -0.43[2]
Vapor pressure 5 mmHg (30 °C)[3]
Acidity (pKa) 11.75
−17.7·10−6 cm3/mol
1.4061
Viscosity 1.245 cP (20 °C)
2.26 D
Thermochemistry
1.267 J/(g·K) (gas)
2.619 J/(g·K) (liquid)
Std enthalpy of
formation
fH298)
−187.80 kJ/mol
Pharmacology
A01AB02 (WHO) D08AX01 (WHO), D11AX25 (WHO), S02AA06 (WHO)
Hazards
GHS labelling:
Danger
H271, H302, H314, H332, H335, H412
P280, P305+P351+P338, P310
NFPA 704 (fire diamond)
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
1518 mg/kg[citation needed]
2000 mg/kg (oral, mouse)[4]
1418 ppm (rat, 4 hr)[4]
227 ppm (mouse)[4]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 1 ppm (1.4 mg/m3)[3]
REL (Recommended)
TWA 1 ppm (1.4 mg/m3)[3]
IDLH
(Immediate danger)
75 ppm[3]
Safety data sheet (SDS) ICSC 0164 (>60% soln.)
Related compounds
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Hydrogen peroxide is a

oxidizer, bleaching agent, and antiseptic, usually as a dilute solution (3%–6% by weight) in water for consumer use, and in higher concentrations for industrial use. Concentrated hydrogen peroxide, or "high-test peroxide", decomposes explosively when heated and has been used as a propellant in rocketry.[6]

Hydrogen peroxide is a

peroxidases
.

## Properties

The boiling point of H2O2 has been extrapolated as being 150.2 °C (302.4 °F), approximately 50 °C (90 °F) higher than water. In practice, hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled at lower temperatures under reduced pressure.[7]

### Structure

Structure and dimensions of H2O2 in the solid (crystalline) phase

Hydrogen peroxide (H2O2) is a nonplanar molecule with (twisted) C2

repulsion between the lone pairs of the adjacent oxygen atoms and dipolar effects between the two O–H bonds. For comparison, the rotational barrier for ethane
is 1040 cm−1 (12.4 kJ/mol).

The approximately 100°

ribonucleic acids and therefore an origin of homochirality in an RNA world.[11]

The molecular structures of gaseous and

tetragonal with the space group D4
4
or P41212.[13]

### Aqueous solutions

In

eutectic mixture, exhibiting freezing-point depression down as low as -56 °C; pure water has a freezing point of 0 °C and pure hydrogen peroxide of -0.43 °C. The boiling point of the same mixtures is also depressed in relation with the mean of both boiling points (125.1 °C). It occurs at 114 °C. This boiling point is 14 °C greater than that of pure water and 36.2 °C less than that of pure hydrogen peroxide.[14]

• Density of aqueous solution of H2O2
H2O2 (
w/w
)
Density
(g/cm3)
Temp.
(°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% 1.45 20
• ### Comparison with analogues

Hydrogen peroxide has several structural analogues with HmX−XHn bonding arrangements (water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, S, N, P). Its melting point is also fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising significantly more readily, indicative of particularly strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs.

Properties of H2O2 and its analogues
Values marked * are extrapolated
Name Formula Molar mass
(g/mol)
Melting
point (°C)
Boiling
point (°C)
Water HOH 18.02 0.00 99.98
Hydrogen peroxide HOOH 34.01 −0.43 150.2*
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*

## Discovery

Alexander von Humboldt is sometimes said to have been the first to report the first synthetic peroxide, barium peroxide, in 1799 as a by-product of his attempts to decompose air, although this is disputed due to von Humboldt's ambiguous wording.[15] 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.[16][17][18] 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.[19]

An improved version of Thénard's process used hydrochloric acid, followed by addition of sulfuric acid to precipitate the barium sulfate byproduct. This process was used from the end of the 19th century until the middle of the 20th century.[20]

The bleaching effect of peroxides and their salts on

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.[21]

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.[23] H2O=O seemed to be just as possible as the modern structure, and as late as in the middle of the 20th century at least half a dozen hypothetical isomeric variants of two main options seemed to be consistent with the available evidence.[24] 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.[25][26]

Previously, hydrogen peroxide was prepared industrially by hydrolysis of ammonium persulfate:

${\displaystyle {\ce {[NH4]2S2O8 + 2 H2O -> 2 [NH4]HSO4 + H2O2}}}$

which was itself obtained by the electrolysis of a solution of ammonium bisulfate ([NH4]HSO4) in sulfuric acid:[27]

${\displaystyle {\ce {2 [NH4]HSO4 -> [NH4]2S2O8 + H2}}}$

## Production

Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was originally developed by BASF in 1939. It begins with the reduction of an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) to the corresponding anthrahydroquinone, typically by hydrogenation on a 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 regenerating the anthraquinone. Most commercial processes achieve oxidation by bubbling compressed air through a solution of the anthrahydroquinone, with the hydrogen peroxide then extracted from the solution and the anthraquinone recycled back for successive cycles of hydrogenation and oxidation.[28][29]

The net reaction for the anthraquinone-catalyzed process is :[28]

H2 + O2 → H2O2

The economics of the process depend 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

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.[30]

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 favours 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.[31][32] One economic obstacle has been that direct processes give a dilute solution uneconomic for transportation. None of these has yet reached a point where it can be used for industrial-scale synthesis.

### Availability

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,

USD/kg, equivalent to US\$1.50/kg (US\$0.68/lb) on a "100% basis"[clarification needed].[28]

### Natural occurrence

Hydrogen peroxide occurs in surface water, in groundwater, and in the

orders of magnitude depending in conditions such as season, altitude, daylight and water vapor content. In rural nighttime air it is less than 0.014 μg/m3, and in moderate photochemical smog it is 14 to 42 μg/m3.[34]

## Reactions

### Decomposition

Hydrogen peroxide decomposes to form water and oxygen with a

kJ/kg[35] and a ΔS
of 70.5 J/(mol·K):

${\displaystyle {\ce {2 H2O2 -> 2 H2O + O2}}}$

The rate of decomposition increases with rise in temperature, concentration, and

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.

### Redox reactions

The redox properties of hydrogen peroxide depend on pH as acidic conditions exacerbate the power of oxidizing agents and basic conditions exacerbate the power of reducing agents. As hydrogen peroxide exhibits ambivalent redox properties, being simultaneously an oxidizer or a reductant, its redox behavior immediately depends on pH.

In acidic solutions, H2O2 is a powerful

oxidizer, stronger than chlorine, chlorine dioxide, and potassium permanganate. When used for cleaning laboratory glassware, a solution of hydrogen peroxide and sulfuric acid is referred to as Piranha solution
.

H2O2 is a source of hydroxyl radicals (OH), which are highly reactive.

Oxidizing
reagent
Reduced
product
Oxidation
potential

(V)
F2 HF 3.0
O3 O2 2.1
H2O2 H2O 1.8
KMnO4 MnO2 1.7
ClO2 HClO 1.5
Cl2 Cl 1.4

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

${\displaystyle {\ce {2 Fe^2+_{(aq)}{}+ H2O2 + 2 H+_{(aq)}-> 2 Fe^3+_{(aq)}{}+ 2 H2O_{(l)}}}}$

and sulfite (SO2−3) is oxidized to sulfate (SO2−4). However, potassium permanganate is reduced to Mn2+ by acidic H2O2.

${\displaystyle {\ce {2 MnO4- + 5 H2O2 + 6 H+ -> 2 Mn^2+ + 8 H2O + 5 O2}}}$[37]

Under

MnO2
).

In basic solutions, hydrogen peroxide is a strong reductant and can reduce a variety of inorganic ions. When H2O2 acts as a reducing agent, oxygen gas is also produced. For example, hydrogen peroxide will reduce sodium hypochlorite and potassium permanganate, which is a convenient method for preparing oxygen in the laboratory:

${\displaystyle {\ce {NaOCl + H2O2 -> O2 + NaCl + H2O}}}$
${\displaystyle {\ce {2 KMnO4 + 3 H2O2 -> 2 MnO2 + 2 KOH + 2 H2O + 3 O2}}}$

### Organic reactions

Hydrogen peroxide is frequently used as an

thioethers to sulfoxides:[38][39]

${\displaystyle {\ce {Ph-S-CH3 + H2O2 -> Ph-S(O)-CH3 + H2O}}}$

Alkaline hydrogen peroxide is used for

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 (CrO3 and H2SO4) forms a blue peroxide CrO(O2)2.

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:

${\displaystyle {\ce {Na2B4O7 + 4 H2O2 + 2 NaOH -> 2 Na2B2O4(OH)4 + H2O}}}$

H2O2 converts carboxylic acids (RCO2H) into peroxy acids (RC(O)O2H), which are themselves used as oxidizing agents. Hydrogen peroxide reacts with acetone to form acetone peroxide and with ozone to form trioxidane. Hydrogen peroxide forms stable adducts with urea (Hydrogen peroxide - urea), sodium carbonate (sodium percarbonate) and other compounds.[41] An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H2O2 in some reactions.

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 manganese ions are reoxidized to make the reaction catalytic. This forms the basis for common monopropellant rockets.

## Biological function

Hydrogen peroxide is formed in humans and other 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.[42] The class of biological enzymes called superoxide dismutase (SOD) 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.[43]

${\displaystyle {\ce {2 O2- + 2 H+ -> O2 + H2O2}}}$
${\displaystyle {\ce {2 H2O2 -> O2 + 2 H2O}}}$

ether phospholipids critical for the normal function of mammalian brains and lungs.[45] Upon oxidation, they produce hydrogen peroxide in the following process catalyzed by flavin adenine dinucleotide (FAD):[46]

${\displaystyle {\ce {R-CH2-CH2-CO-SCoA + O2 ->[{\ce {FAD}}] R-CH=CH-CO-SCoA + H2O2}}}$

Catalase, another peroxisomal enzyme, uses this H2O2 to oxidize other substrates, including phenols, formic acid, formaldehyde, and alcohol, by means of a peroxidation reaction:

${\displaystyle {\ce {H2O2 + R'H2 -> R' + 2 H2O}}}$

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.[47] In addition, when excess H2O2 accumulates in the cell, catalase converts it to H2O through this reaction:

${\displaystyle {\ce {H2O2 ->[{\ce {CAT}}] {1/2O2}+ H2O}}}$

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:[48]

H2O, O2
H2O2
H2O, O2
H2O2
Degradation of hypoxanthine through xanthine to uric acid to form hydrogen peroxide.

The degradation of guanosine monophosphate yields xanthine as an intermediate product which is then converted in the same way to uric acid with the formation of hydrogen peroxide.[48]

Eggs of sea urchin, shortly after fertilization by a sperm, produce hydrogen peroxide. It is then quickly dissociated to HO• radicals. The radicals serve as initiator of radical polymerization, which surrounds the eggs with a protective layer of polymer.[49]

The

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 popping sound.[50][51][52]

Hydrogen peroxide is a

Hydrogen peroxide has roles as a signalling molecule in the regulation of a wide variety of biological processes.

mitochondria.[56][57][58] At least one study has also tried to link hydrogen peroxide production to cancer.[59] These studies have frequently been quoted in fraudulent treatment claims.[citation needed
]

## Uses

### Bleaching

About 60% of the world's production of hydrogen peroxide is used for

Tide laundry detergent. When dissolved in water, it releases hydrogen peroxide and sodium carbonate,[20] 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. It has also been used as a flour bleaching agent and a tooth whitening agent.

### Production of organic compounds

It is used in the production of various

organic peroxide-based explosives, such as acetone peroxide. It is used as an initiator in polymerizations
.

### Sewage treatment

Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. In

halogenated compounds.[63] It can also oxidize sulfur-based compounds present in the waste; which is beneficial as it generally reduces their odour.[64]

### Disinfectant

Hydrogen peroxide may be used for the sterilization of various surfaces,

Gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms may increase tolerance in the presence of lower concentrations.[70] Lower levels of concentration (3%) will work against most spores; higher concentrations (7 to 30%) and longer contact times will improve sporicidal activity.[69][71]

Hydrogen peroxide is seen as an environmentally safe alternative to

### Propellant

Rocket-belt hydrogen-peroxide propulsion system used in a jet pack

High-concentration H2O2 is referred to as "high-test peroxide" (HTP). It can be used either as a

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.[74]

As a bipropellant, H2O2 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, 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.

hypergolic combination, and for the low-cost British Black Knight and Black Arrow launchers. Presently, HTP is used on ILR-33 AMBER[75] and Nucleus[76]
suborbital rockets.

In the 1940s and 1950s, the Hellmuth Walter KG–conceived turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant. Operator error in the use of hydrogen-peroxide torpedoes was named as possible causes for the sinking of HMS Sidon and the Russian submarine Kursk.[77] 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.[78][79]

### Household use

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

Hair bleaching

Diluted H2O2 (between 1.9% and 12%) mixed with

Hydrogen peroxide is also used for tooth whitening. It may be found in most whitening toothpastes. Hydrogen peroxide has shown positive results involving teeth lightness and chroma shade parameters.[81] It works by oxidizing colored pigments onto the enamel where the shade of the tooth may become lighter.[further explanation needed] Hydrogen peroxide may be mixed with baking soda and salt to make a homemade toothpaste.[82]

Removal of blood stains

Hydrogen peroxide reacts with blood as a bleaching agent, and so if a blood stain is fresh, or not too old, liberal application of hydrogen peroxide, if necessary in more than single application, will bleach the stain fully out. After about two minutes of the application, the blood should be firmly blotted out.[83][84]

Acne treatment

Hydrogen peroxide may be used to treat

acne,[85] although benzoyl peroxide
is a more common treatment.

### Niche uses

cyalume
, as found in a glow stick
Glow sticks

Hydrogen peroxide reacts with certain di-

phenyl oxalate ester (cyalume), to produce chemiluminescence; this application is most commonly encountered in the form of glow sticks
.

Horticulture

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.[86][87]

For general watering concentrations around 0.1% is in use and this can be increased up to one percent for anti-fungal actions.[88] Tests show that plant foliage can safely tolerate concentrations up to 3%.[89]

Fishkeeping

Hydrogen peroxide is used in aquaculture for controlling mortality caused by various microbes. In 2019, the U.S. FDA approved it for control of Saprolegniasis in all coldwater finfish and all fingerling and adult coolwater and warmwater finfish, for control of external columnaris disease in warm-water finfish, and for control of Gyrodactylus spp. in freshwater-reared salmonids.[90] Laboratory tests conducted by fish culturists have demonstrated that common household hydrogen peroxide may be used safely to provide oxygen for small fish. The hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide.

Removing yellowing from aged plastics

Hydrogen peroxide may be used in combination with a UV-light source to remove yellowing from white or light grey

retr0bright
.

## Safety

Regulations vary, but low concentrations, such as 5%, are widely available and legal to buy for medical use. Most over-the-counter peroxide solutions are not suitable for ingestion. Higher concentrations may be considered hazardous and typically are accompanied by a safety data sheet (SDS). In high concentrations, hydrogen peroxide is an aggressive oxidizer and will corrode many materials, including human skin. In the presence of a reducing agent, high concentrations of H2O2 will react violently.[91] While concentrations up to 35% produce only "white" oxygen bubbles in the skin (and some biting pain) that disappear with the blood within 30-45 minutes, concentrations of 98% dissolve paper. However concentrations as low as 3% can be dangerous for the eye because of oxygen evolution within the eye.[92]

High-concentration hydrogen peroxide streams, typically above 40%, should be considered hazardous due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds (45 kg), or approximately 10 US gallons (38 L), of concentrated hydrogen peroxide.

Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminium alloys may also be suitable).[93] 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.[94]

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.[95] 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.[96] 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.[97]

With a significant vapour pressure (1.2 kPa at 50 °C),

immediately dangerous to life and health (IDLH) limit is only 75 ppm.[99] 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).[95] 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".[100] 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[95] and from the ATSDR.[101]

### Wound healing

There is conflicting evidence on hydrogen peroxide's effect on wound healing. Some research finds benefit, while other research find delays and healing inhibition.[103] Its use for home treatment of wounds is generally contraindicated.[104] 1.5–3% Hydrogen peroxide is used as a desinfectant in dentistry, especially in endodotic treatments together with hypochlorite and chlorhexidin and 1–1.5% is also useful for treatment of inflammation of third molars (wisdom teeth).[105]

### Use in alternative medicine

Practitioners of

AIDS, and in particular cancer.[106] There is no evidence of effectiveness and in some cases it has proved fatal.[107][108][109][110][111]

Both the effectiveness and safety of hydrogen peroxide therapy is scientifically questionable. Hydrogen peroxide is produced by the immune system, but in a carefully controlled manner. Cells called phagocytes engulf pathogens and then use hydrogen peroxide to destroy them. The peroxide is toxic to both the cell and the pathogen and so is kept within a special compartment, called a phagosome. Free hydrogen peroxide will damage any tissue it encounters via oxidative stress, a process that also has been proposed as a cause of cancer.[112] Claims that hydrogen peroxide therapy increases cellular levels of oxygen have not been supported. The quantities administered would be expected to provide very little additional oxygen compared to that available from normal respiration. It is also difficult to raise the level of oxygen around cancer cells within a tumour, as the blood supply tends to be poor, a situation known as tumor hypoxia.

Large oral doses of hydrogen peroxide at a 3% concentration may cause irritation and blistering to the mouth, throat, and abdomen as well as abdominal pain, vomiting, and diarrhea.[107] Ingestion of hydrogen peroxide at concentrations of 35% or higher has been implicated as the cause of numerous gas embolism events resulting in hospitalisation. In these cases, hyperbaric oxygen therapy was used to treat the embolisms.[113]

Intravenous injection of hydrogen peroxide has been linked to several deaths.[109][110][111] The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective, or useful cancer treatment."[108] Furthermore, the therapy is not approved by the U.S. FDA.

## References

Notes

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5. .
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7. .
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9. .
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11. .
12. .
13. .
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15. PMID 33317108
. I checked Humboldt's pertinent publication carefully, but was unable to find an unambiguous proof of this assumption; the description of the starting materials ("Alaun-Erden" or "schwere Erden") were just too unprecise to understand what kind of chemical experiments he performed.
16. .
17. ^ Thénard LJ (1818). "Observations sur des nouvelles combinaisons entre l'oxigène et divers acides". Annales de chimie et de physique. 2nd series. 8: 306–312. Archived from the original on 3 September 2016. Retrieved 9 February 2016.
18. from the original on 30 November 2018. Retrieved 28 November 2018. Hydrogen peroxide was discovered in 1818 by the French chemist Louis-Jacques Thenard, who named it eau oxygénée (oxygenated water).
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