Hydrogen peroxide
| Hydrogen peroxide | |
|---|---|
 |
|
| General | |
| Systematic name | hydrogen peroxide |
| Other names | hydrogen
dioxide hydroperoxide |
| Molecular formula | H2O2 |
| Molar mass | 34.01 g/mol |
| pH | 4.5 |
| Appearance | clear liquid |
| CAS number | [7722-84-1] |
| Properties | |
| Density and phase | 1.4 g/cm3, liquid |
| Solubility in water | miscible |
| Melting point | -11 °C (262 K) |
| Boiling point | 141 °C (414 K) |
| Acidity (pKa) | 11.65 |
| Viscosity | 1.245 c P at 20 °C |
| Structure | |
| Molecular shape | skewed |
| Dipole moment | 2.26 D |
| Hazards | |
| MSDS | External MSDS |
| EU classification | Oxidant
(O) Corrosive (C) |
| R-phrases | R5, R8, R20/R22, R35 |
| S-phrases | S1/S2, S17, S26, S28, S36/S37/S39, S45 |
| NFPA 704 |
   |
| Flash point | not flammable |
| RTECS number | MX0900000 |
| Supplementary data page | |
|
Structure
and properties |
n, εr, etc. |
|
Thermodynamic data |
Phase
behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Related compounds | |
| Other cations | sodium peroxide |
| Related peroxides |
methyl
peroxide hydrogen trioxide |
| Related compounds |
water oxygen, ozone hydrazine |
| Except where noted otherwise, data are given
for materials in their standard state (at 25°C, 100 kPa) Infobox disclaimer and references |
|
Hydrogen peroxide (H2O2) is a clear liquid, slightly more viscous than water, that 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.
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.
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.
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.
Storage
Household hydrogen peroxide solutions are commonly found in concentrations of 3% solution. Because hydrogen peroxide decomposes in the presence of light, it should be stored in a cool environment out of direct sunlight. It should also be stored out of reach of children in a well-marked bottle, as ingestion of large amounts can cause gastrointestinal problems including burns, pain and vomiting. Hydrogen peroxide decomposes naturally into oxygen and water. This property means that caution must be taken when opening a bottle that has been stored for a long period of time. Be sure not to open an old bottle near a flame or spark.
Stronger solutions such as a 35% solution have been responsible for at least one death when ingested undiluted.
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 (>600C) either to be used directly or to combust a fuel. As a monopropellant (not mixed with fuel), it produces a maximum specific impulse (Isp) of 161 s (1.6 kN•s/kg), which makes it a low-performance monopropellant. The famous Bell rocket pack 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 much less toxic, but it is also much less powerful. Peroxide gives a slightly lower Isp 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."
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 FDA [2]. Hydrogen Peroxide can also be used as a toothpaste when mixed with correct quantities of Baking Soda and Salt [3].
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; 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
ΔHo of −98.2
kJ/ mol and a
ΔGo of −119.2 kJ/mol. 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 or aluminium (not stainless steel) 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 a 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.
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 (H3O2)+ 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. 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, 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.
- MATERIAL SAFETY DATA SHEET
- ATSDR Agency for Toxic Substances and Disease Registry FAQ