Hydrogen peroxide

Hydrogen peroxide
General
Systematic name	Dihydrogen dioxide
Other names	Hydrogen peroxide hydrogen dioxide
Molecular formula	H₂O₂
Molar mass	34.0146 g/mol.
Appearance	Colourless liquid.
CAS number
Properties
Density and phase	1.4 g/cm³, liquid
Solubility in water	Miscible.
Melting point	-11 °C (262.15 K)
Boiling point	150.2 °C (423.35 K)
Acidity (pK_a)	11.65
Viscosity	1.245 cP at 20 °C
Structure
Molecular shape	?
Dipole moment	2.26 D
Hazards
MSDS	External MSDS
Main hazards	Oxidant, corrosive.
NFPA 704
Flash point	Non-flammable.
R/S statement	R: , , , , S: , , , , , , , ,
RTECS number	MX0900000
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Other anions	?
Other cations	Sodium peroxide
Related compounds	Water ozone hydrazine
Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Chemical infobox

Hydrogen peroxide (H₂O₂) is a clear liquid, slightly more viscous than water. It has strong oxidizing properties and is therefore a powerful bleaching agent that has found use as a disinfectant, as an oxidizer, and in rocketry (particularly in high concentrations as high-test peroxide (HTP) as a monopropellant), and in bipropellant systems.

History

Hydrogen peroxide was first isolated in 1818 by Louis Jacques Thénard by reacting barium peroxide with nitric acid. An improved version of this process used hydrochloric acid, followed by sulfuric acid to precipitate the barium chloride that is obtained as a byproduct. Thenard's process was used from the end of the 19th century until the middle of the 20th century. Craig W Jones, J H Clark. Applications of Hydrogen Peroxide and Deriatives. Royal Society of Chemistry, 1999. For more modern manufacture methods, see below.

Uses

Domestic uses

It is commonly used (in very low concentrations, typically around 5%) to bleach human hair, hence the phrases peroxide blonde and bottle blonde. It burns the skin upon contact in sufficient concentration. In lower concentrations (3%), it is used medically for cleaning wounds and removing dead tissue. The Food and Drug Administration has approved 3% hydrogen peroxide ("Food-Grade," or without added chemical stabilisers) for use as a mouthwash. Commercial peroxide solutions (most H₂O₂ bought over the counter from pharmacies) are not suitable for ingestion as they contain additional harmful chemicals.

The Food and Drug Administration (FDA) has classified hydrogen peroxide as a Low Regulatory Priority (LRP) drug for use in controlling fungus on fish and fish eggs. Hydrogen peroxide has been experimentally proven to be effective against Amyloodinium sp., a marine fish ectoparasite. It was used at a dosage of 25ppm for 30 minutes to treat Pacific Threadfin, Polydactylus sexfilis infected with Amyloodinium ocellatum. (Montgomery et al., 1999b). Some species of fish tolerate the treatment well, but others are highly sensitive to the chemical (Noga, 2000). Results may also vary between juvenile and adult fish.

Hydrogen peroxide is effective against other ectoparasites, such as Ambiphrya and Gyrodactylus spp. (Rach et al. 2000). Sodium percarbonate is a compound that releases hydrogen peroxide when dissolved in water. Sodium percarbonate was demonstrated to kill the freshwater ectoparasite, Ichthyophthirius multifiliis, at the theront stage or free-swimming, infective stage (Buchmann, et al., 2002). It is currently used in Denmark with rainbow trout Oncorhynchus mykiss at a concentration of 50-100 mg/L is twice a week without any apparent ill effects on the fish.

If more than 50% of the theronts died the concentration of hydrogen peroxide was recorded as effective (Buchmann, et al., 2002). A dosage of 12.5mg/L at a temperature of 12°C kills theronts within 3 hours. This same dosage was not effective against the tomocysts stage of Ichthyophthirius multifiliis. However, dosages of 12.5 mg/L for 180 min and 62.5 mg/L for 90 min were effective against theronts (Buchmann, et al., 2002).

The life cycle of Ichthyophthirius multifiliis is temperature dependent, so the warmer the water temperature the shorter the duration of the parasites life cycle. At 12°C, the medium time frame for tomocsyts to hatch is 9 days and the attached parasitic stage has a duration of 10-12 days. This means that at 12°C treatment should continue daily for a minimum of three weeks. Caution should be taken when using hydrogen peroxide as accidental spills may have an immediate adverse effect on human skin.

Studies to test the effectiveness and safety of hydrogen peroxide for the treatment of other ectoparasites such as Cryptocaryon irritans are fully warranted (Montgomery-Brock, D. et al., 2000). However, this treatment is considered to be highly experimental, therefore it cannot be recommended. The side effects and survival rate when using hydrogen peroxide may not prove to be acceptable. Protective clothing and safety glasses should be worn when using a dose of 35% and higher. The water temperature should be carefully monitored when treating with hydrogen peroxide, because it becomes more toxic as the temperature rises. At this point, the safety, effectiveness, correct dosage and duration of treatment for this experimental method have not been established.

Some gardeners and hydroponics implementers have professed the value of hydrogen peroxide in their watering solutions. They claim its spontaneous decomposition releases oxygen to the plant that can enhance root development and also help treat root rot, which is cellular root death due to lack of oxygen. 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. Citation Reference Hydrogen-peroxide releases oxygen by decomposition when it is added to water. Hydrogen peroxide has been gaining in popularity for the treatment of hydrogen sulfide and iron. Catalytic carbon and redox media perform well with hydrogen peroxide pretreatment. Generally 90% of the reaction between hydrogen peroxide and hydrogen sulfide takes place within 10 to 15 minutes, with the balance reacting in an additional 20 to 30 minutes. The sulfur in hydrogen sulfide (H2S) is in the -2 state. In a neutral solution, hydrogen peroxide will oxidize hydrogen sulfide to elemental sulfur via the following reaction:

8 H₂S(g) + 8 H₂O₂(aq) = S₈(s) + 16 H₂O(l)

No acid is produced. The reaction takes a while, so if you're bubbling the hydrogen sulfide through the peroxide solution, you need to recycle the gas stream through the peroxide solution. Metal ions catalyze the reaction. To be more specific for doses of chemical feed levels for oxidation of iron, manganese and hydrogen sulfide in domestic water supplies, here are some figures:

Iron: For each ppm Fe feed = 0.3 - 0.5 ppm, 20 minutes

Manganese: For each ppm Mn feed = 0.8 - 1.0 ppm, 20 minutes

Hydrogen Sulfide: For each ppm H2S feed = 1.0 - 1.5 ppm, 30 minutes

(all above figures are for minimum retention time). When more than one constituent is to be oxidized (i.e. iron & H2S) add the above values to determine the total ppm feed needed to oxidize two or more.

H2O2 is one of the most powerful oxidizers known -- stronger than chlorine, chlorine dioxide, and potassium permanganate. And through catalysis, H2O2 can be converted into hydroxyl radicals (.OH) with reactivity second only to fluorine.

Oxidant	Oxidation Potential, V
Fluorine	3.0
Hydroxyl radical	2.8
Ozone	2.1
Hydrogen peroxide	1.8
Potassium permanganate	1.7
Chlorine dioxide	1.5
Chlorine	1.4

Hydrogen peroxide is strong oxidizer effective in controlling sulfide and organic related odors in wastewater collection and treatment systems. It is typically applied to a wastewater system most frequently where there is a retention time of less than five hours and at least 30 minutes prior to the point where the hydrogen sulfide is released. Hydrogen peroxide will oxidize the hydrogen sulfide present and in addition promote bio-oxidation of organic odors. Hydrogen peroxide decomposes to oxygen and water adding dissolved oxygen to the system thereby reducing Biological Oxygen Demand (BOD).

Commercial peroxide, as bought at the drugstore in a 2.5%-3% solution, can be used to remove bloodstains from carpets and clothing. If a few tablespoons of peroxide are poured onto the stain, they will bubble up in the area of the blood. After a few minutes the excess liquid can be wiped up with a cloth or paper towel and the stain will be gone. Care should be taken, however, as hydrogen peroxide will bleach or discolor many fabrics.

Hydrogen peroxide is used in glow sticks as an oxidising agent. It reacts with phenyl oxalate ester to form an unstable CO₂ dimer which in turn causes an added dye to reach an excited state, the latter relaxing to release photons of light.

Storage

Household hydrogen peroxide solutions are commonly found in concentrations of 3% solution stored in dark containers, since hydrogen peroxide decomposes in light.

Possession of any concentration over 30% w/w requires a special permit (in the U.S.) due to the explosive reactivity of concentrated hydrogen peroxide. The best way to store hydrogen peroxide is in a dark refrigerator.

Industrial applications

About 50% of the world's production of hydrogen peroxide in 1994 was used for pulp- and paper-bleaching. Other bleaching applications are becoming more important as hydrogen peroxide is seen as a more environmentally-benign alternative to chlorine based bleaches.

Other major industrial applications for hydrogen peroxide include the manufacture of sodium percarbonate and sodium perborate, used as mild bleaches in laundry detergents. In addition it is used in the production of certain organic peroxides such as dibenzoyl peroxide, used as free radical initiators in polymerisation and other chemical processes. Hydrogen peroxide is also used in the production of epoxides such as propylene oxide. Reaction with carboxylic acids produces a corresponding "per-acid"; for industrial use peracetic acid is prepared in this way from acetic acid. MCPBA, used extensively in the laboratory, is likewise prepared from meta-chlorobenzoic acid.

Use as propellant

The use of H₂O₂ as a propellant takes advantage of the decomposition of 70-98+% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber where usually a metal (esp. silver or platinum) catalyst triggers decomposition, and the oxygen/steam that is produced is hot enough (>600 °C) either to be used directly or to combust a fuel. As a monopropellant (not mixed with fuel), it produces a maximum specific impulse (I_sp) of 161 s (1.6 kN·s/kg), which makes it a low-performance monopropellant. The famous Bell Rocket Belt used hydrogen peroxide monopropellant. When 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 was used very successfully as an oxidizer for the low-cost British launchers, Black Knight and Black Arrow.

Compared to hydrazine, peroxide is less toxic, but it is also much less powerful. Peroxide gives a slightly lower I_sp than liquid oxygen, but is dense, storable, noncryogenic and can be more easily used to drive gas turbines to give high pressures. It also can be used for regenerative cooling of rocket engines.

In the 1940s and 1950s, the Walter turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and maintenance-demanding compared to the conventional diesel-electric power system. Some torpedoes used hydrogen peroxide as oxidizer or propellant, but this use has been discontinued by most navies for safety reasons. Hydrogen peroxide leaks were blamed for the sinkings of HMS Sidon and the Russian submarine Kursk. It was discovered, for example, by the Japanese Navy in torpedo trials, that the concentration of H₂O₂ in right-angle bends in HTP pipework can often lead to explosions in submarines and torpedoes.

While its application as a monopropellant for large engines has waned, small thrusters for attitude control that run on hydrogen peroxide are still in use on some satellites, and provide benefits on the spacecraft, making it easier to throttle and safer loading and handling of fuel before launch (as compared to hydrazine monopropellant). However, hydrazine is a more popular monopropellent in spacecraft because of its higher specific impulse and lower rate of decomposition.

Therapeutic use

Hydrogen peroxide has been used as an antiseptic and anti-bacterial agent for many years. While its use has decreased in recent years due to the popularity of better-smelling and more readily-available over the counter products, it is still used by many hospitals, doctors and dentists in sterilising, cleaning and treating everything from floors to Root canal procedures.

Thirty-five percent food-grade hydrogen peroxide has been marketed under names such as Oxywater and H₂O₂, with claims of medicinal or therapeutic value as "hydrogen peroxide therapy." Advocates of the product claim that it can be diluted and used for "hyper-oxygenation therapy" to treat AIDS, cancer, and many other conditions; some also claim that information about these beneficial uses of peroxide have been suppressed by the medical and scientific communities.

Recently, alternative medical practitioners have advocated administering doses of hydrogen peroxide intravenously in extremely low (less than one percent) concentrations for hydrogen peroxide therapy — a controversial alternative medical treatment for cancer. However, according to the American Cancer Society, "there is no scientific evidence that hydrogen peroxide is a safe, effective or useful cancer treatment." They advise cancer patients to "remain in the care of qualified doctors who use proven methods of treatment and approved clinical trials of promising new treatments." CA Cancer J Clin. 1993 Jan-Feb;43(1):47-56. "Questionable methods of cancer management: hydrogen peroxide and other 'hyperoxygenation' therapies." PMID 8422605 Internal use of hydrogen peroxide has a history of causing fatal blood disorders, and its recent use as a theraputic treatment has been linked to several deaths.CBS News, Jan 12 2005, "A Perscription for Death?" http://www.cbsnews.com/stories/2005/01/12/60II/main666489.shtml Snopes.com, "Hydrogen Peroxide." http://www.snopes.com/medical/healthyself/peroxide.asp

Another therapeutic use of hydrogen peroxide is to cure colds and flus. Some alternative medicine practitioners recommend inserting a few drops into each ear at the first sign of a cold. According to their claims, for most people, this will eliminate the cold virus at first attempt within a few hours, with stated 80% success rate. However the medical evidence of its efficacy is lacking.

Hydrogen Peroxide is GRAS (Generally Recognised As Safe) as an antimicrobial agent, an oxidizing agent and more by the US Food and Drug Administration Hydrogen Peroxide can also be used as a toothpaste when mixed with correct quantities of Baking Soda and Salt [http://www.fda.gov/bbs/topics/CONSUMER/CON00065.html.

Physical properties

Hydrogen peroxide adopts a "skewed" shape, due to repulsion between the lone pairs on the oxygen atoms. Despite the fact that the O-O bond is a single bond, the molecule has a remarkably high barrier to complete rotation of 29.45 kJ/mol (compared with 12.5 kJ/mol for the rotational barrier of ethane); this is also caused by the lone-pair repulsion. The bond angles are affected greatly by hydrogen bonding, which explains the difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H₂O₂.

Chemical properties

Hydrogen peroxide can decompose spontaneously into water and oxygen. It usually acts as an oxidizing agent, but there are many reactions where it acts as a reducing agent, releasing oxygen as a by-product. It also readily forms both inorganic and organic peroxides.

Decomposition

Hydrogen peroxide often decomposes exothermically into water and oxygen gas spontaneously:

2 H₂O₂ → 2 H₂O + O₂ + Energy

This process is very favorable; it has a ΔH^o of −98.2 kJ/mol and a ΔG^o of −119.2 kJ/mol and a ΔS of 70.5 J/mol K. The rate of decomposition is dependent on the temperature and concentration of the peroxide, as well as the pH and the presence of impurities and stabilizers. Hydrogen peroxide is incompatible with many substances that catalyse its decomposition, including most of the transition metals and their compounds. Common catalysts include manganese dioxide, potassium permanganate, and silver. The same reaction is catalysed by the enzyme catalase, found in the liver, whose main function in the body is the removal of toxic byproducts of metabolism and the reduction of oxidative stress. The decomposition occurs more rapidly in alkali, so acid is often added as a stabilizer.

Spilling high concentration peroxide on a flammable substance can cause an immediate fire fueled by the oxygen released by the decomposing hydrogen peroxide. High-strength peroxide (also called high-test peroxide, or HTP) must be stored in a vented container to prevent the buildup of oxygen gas, which would otherwise lead to the eventual rupture of the container. Any container must be made of a compatible material such as PTFE, polyethylene, stainless steel or aluminium and undergo a cleaning process (passivation) to remove all contamination prior to the introduction of peroxide. (Note that while compatible at room temperature, polyethylene can explode with peroxide in a fire.)

In the presence of certain catalysts, such as Fe²⁺ or Ti³⁺, the decomposition may take a different path, with free radicals such as HO· (hydroxyl) and HOO· being formed. A combination of H₂O₂ and Fe²⁺ is known as Fenton's reagent.

Redox reactions

In aqueous solution, hydrogen peroxide can oxidize or reduce a variety of inorganic ions. When it acts as an reducing agent, oxygen gas is also produced. In acid solution Fe²⁺ is oxidized to Fe³⁺,

2 Fe²⁺(aq) + H₂O₂ + 2 H⁺(aq) → 2 Fe³⁺(aq) + 2H₂O(l)

and sulfite (SO₃²⁻) is oxidized to sulfate (SO₄²⁻). However, potassium permanganate is reduced to Mn²⁺ by acidic H₂O₂. Under alkaline conditions, however, some of these reactions reverse; Mn²⁺ is oxidized to Mn⁴⁺ (as MnO₂), yet Fe³⁺ is reduced to Fe²⁺.

2 Fe³⁺ + H₂O₂ + 2 OH⁻ → 2 Fe²⁺ + 2 H₂O + O₂

Hydrogen peroxide is frequently used as an oxidising agent in organic chemistry. One application is for the oxidation of thioethers to sulfoxides. For example, methyl phenyl sulfide was oxidised to methyl phenyl sulfoxide in 99% yield in methanol in 18 hours (or 20 minutes using a TiCl₃ catalyst):

Ph-S-CH₃ + H₂O₂ → Ph-S(O)-CH₃ + H₂O

Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acids, and also for oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation.

Formation of peroxide compounds

Hydrogen peroxide is a weak acid, and it can form hydroperoxide or peroxide salts or derivatives of many metals. For example, with aqueous solutions of chromic acid (CrO₃), it can form an unstable blue peroxide CrO(O₂)₂. It can also produce peroxoanions by reaction with anions; for example, reaction with borax leads to sodium perborate, a bleach used in laundry detergents:

Na₂B₄O₇ + 4 H₂O₂ + 2 NaOH → 2 Na₂B₂O₄(OH)₄ + H₂O

H₂O₂ converts carboxylic acids (RCOOH) into peroxy acids (RCOOOH), which are themselves used as oxidizing agents. Hydrogen peroxide reacts with acetone to form acetone peroxide, and it interacts with ozone to form hydrogen trioxide. Reaction with urea produces carbamide peroxide, used for whitening teeth. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H₂O₂ in some reactions.

Hydrogen peroxide reacts with ozone to form trioxidane.

Basicity

Hydrogen peroxide is a much weaker base than water, but it can still form adducts with very strong acids. The superacid HF/SbF₅ forms unstable compounds containing the ^*⁺ ion.

Manufacture

Hydrogen peroxide is manufactured today almost exclusively by the autoxidation of 2-ethyl-9,10-dihydroxyanthracene to 2-ethylanthraquinone and hydrogen peroxide using oxygen from the air. The anthraquinone derivative is then extracted out and reduced back to the dihydroxy compound using hydrogen gas in the presence of a metal catalyst. The overall equation for the process is deceptively simple:

H₂ + O₂ → H₂O₂

However the economics of the process depend on effective recycling of the quinone and extraction solvents, and of the hydrogenation catalyst.

Formerly inorganic processes were used, employing the electrolysis of an aqueous solution of sulfuric acid or acidic ammonium bisulfate (NH₄HSO₄), followed by hydrolysis of the peroxydisulfate ((SO₄)₂)²⁻ which is formed.

In 1994, world production of H₂O₂ was around 1.9 million tonnes, most of which was at a concentration of 70% or less. In that year bulk 30% H₂O₂ sold for around US $0.54 per kg, equivalent to US $1.50 per kg (US $0.68 per lb) on a "100% basis".

Concentration

Hydrogen peroxide works best as a propellant in extremely high concentrations. However, there are very few suppliers of high-purity hydrogen peroxide, and they are averse to selling to any but the largest institutions. As a result, amateurs wishing to use this for rocket fuel usually have to purchase 70% or lower-purity (most of the remaining 30% is water, and sometimes there are traces of stabilizing materials, such as tin), and increase its concentration themselves. Many try distillation, but this is extremely dangerous with hydrogen peroxide; peroxide vapour can detonate at a temperature of about 70 °C (158 °F). A safer approach is sparging, possibly followed by fractional freezing, but, even when using this method, contaminants may still often cause explosions.

In the 1950s, high-test peroxide was more readily available, but because of safety concerns bulk manufacturers have since switched over to handling lower concentrations of H₂O₂ whenever possible. Some amateur groups have expressed interest in manufacturing their own peroxide, for their use and for sale in small quantities to others.

Hazards

Hydrogen peroxide vapour can detonate above 70°C (158°F), so it is critical to keep solutions and vapour cool. Distillation of hydrogen peroxide at normal pressures is highly dangerous. Hydrogen peroxide vapours can form sensitive contact explosives with hydrocarbons such as greases. Hazardous reactions ranging from ignition to explosion have been reported with alcohols, ketones, carboxylic acids (particularly acetic acid), amines and phosphorus.

Hydrogen peroxide, if spilled on clothing (or other flammable materials), will preferentially evaporate water until the concentration reaches sufficient strength, then clothing will spontaneously ignite. Leather generally contains metal ions from the tanning process and will often catch fire quite quickly.

Concentrated hydrogen peroxide (>50%) is corrosive, and 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 bleeding. Severe pulmonary irritation by inhalation over 10%.

The IARC lists hydrogen peroxide in Group 3: not classifiable as to its carcinogenicity to humans. One study on mice found an increase in cancers of the digestive system following ingestion of hydrogen peroxide, but other animal studies have proven inconclusive. Hydrogen peroxide is produced as a byproduct of oxygen metabolism, and virtually all organisms possess enzymes known as peroxidases, which catalyse the decomposition of hydrogen peroxide to water and oxygen (see Decomposition above).

A leak of high-test peroxide (85%-98% hydrogen peroxide) from a torpedo caused an explosion that sealed the fate of the Russian submarine Kursk. Also an earlier, very similar accident on HMS Sidon (P259) claimed 13 lives.

References

J. Drabowicz et al., in The Syntheses of Sulphones, Sulphoxides and Cyclic Sulphides, p112-116, G. Capozzi et al., eds., John Wiley & Sons, Chichester, UK, 1994. ISBN 0-471-93970-6.
N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997. A great description of properties & chemistry of H₂O₂.
J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
W. T. Hess, Hydrogen Peroxide, in Kirk-Othmer Encyclopedia of Chemical Technology, 4th edition, Wiley, New York, Vol. 13, 961-995 (1995).

External links

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