|UN number||2015 (>60% soln.)
2014 (20–60% soln.)
2984 (8–20% soln.)
|RTECS number||MX0900000 (>90% soln.)
MX0887000 (>30% soln.)
|Molar mass||34.0147 g/mol|
|Appearance||Very light blue color; colorless in solution|
-0.43 °C, 273 K, 31 °F
150.2 °C, 423 K, 302 °F
|Solubility in water||Miscible|
|Solubility||soluble in ether|
|Acidity (pKa)||11.62 |
|Refractive index (nD)||1.34|
|Viscosity||1.245 cP (20 °C)|
|Dipole moment||2.26 D|
|Std enthalpy of
|Specific heat capacity, C||1.267 J/g K (gas)
2.619 J/g K (liquid)
|MSDS||ICSC 0164 (>60% soln.)|
|EU classification||Oxidant (O)
|R-phrases||, , ,|
|S-phrases||, , , , ,|
(what is this?) |
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Hydrogen peroxide (H2O2) is a very pale blue liquid, slightly more viscous than water, that appears colorless in dilute solution. It is a weak base, has strong oxidizing properties, and is a powerful bleaching agent. It is used as a disinfectant, antiseptic, oxidizer, and in rocketry as a propellant. The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species.
Hydrogen peroxide is naturally produced in organisms as a by-product of oxygen metabolism. Nearly all living things (specifically, all obligate and facultative aerobes) possess enzymes known as peroxidases, which harmlessly and catalytically decompose low concentrations of hydrogen peroxide to water and oxygen.
As with all molecules, the physical properties of hydrogen peroxide are the result of its molecular mass, structure, and distribution of atoms within the molecule.
The preferred molecular structure of any molecule is the configuration that has the lowest internal stress. For hydrogen peroxide, there are two basic structural forms (conformers) available for the molecule. Whereas flat shape of the anti conformer would minimize steric repulsions, the 90° torsion angle of the syn conformer would optimize mixing between the filled p-type orbital of the oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond.
The resulting anticlinal "skewed" shape is a compromise between the two conformers.
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). The increased barrier is attributed to repulsion between one lone pair and other lone pairs. The bond angles are affected by hydrogen bonding, which is relevant to the structural difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H2O2.
Analogues of hydrogen peroxide include the chemically identical deuterium peroxide and malodorous hydrogen disulfide. Hydrogen disulfide has a boiling point of only 70.7°C despite having a higher molecular weight, indicating that hydrogen bonding increases the boiling point of hydrogen peroxide.
Aqueous hydrogen peroxide solutions have specific properties that are different from those of the pure chemical due to hydrogen bonding between water and hydrogen peroxide molecules. Specifically, hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression. While pure water melts and freezes at approximately 273K, and pure hydrogen peroxide just 0.4K below that, a 50% (by volume) solution melts and freezes at 221 K.
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 sulfate byproduct. Thénard's process was used from the end of the 19th century until the middle of the 20th century. Modern production methods are discussed below.
For a long time, pure hydrogen peroxide was believed to be unstable, because attempts to separate the hydrogen peroxide from the water, which is present during synthesis, failed. This was because traces of solids and heavy metal ions led to a catalytic decomposition or explosions of the hydrogen peroxide. One hundred percent pure hydrogen peroxide was first obtained through vacuum distillation by Richard Wolffenstein in 1894. At the end of 19th century, Petre Melikishvili and his pupil L. Pizarjevski showed that of the many proposed formulas of hydrogen peroxide, the correct one was H-O-O-H.
The use of H2O2 sterilization in biological safety cabinets and barrier isolators is a popular alternative to ethylene oxide (EtO) as a safer, more efficient decontamination method. H2O2 has long been widely used in the pharmaceutical industry. In aerospace research, H2O2 is used to sterilize satellites.
The FDA has recently granted 510(k) clearance to use H2O2 in individual medical device manufacturing applications. EtO criteria outlined in ANSI/AAMI/ISO 14937 may be used as a validation guideline. Sanyo was the first manufacturer to use the H2O2 process in situ in a cell culture incubator, which is a faster and more efficient cell culture sterilization process.
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 peroxodisulfate ((SO4)2)2− that is formed.
Today, hydrogen peroxide is manufactured almost exclusively by the autoxidation of a 2-alkyl anthrahydroquinone (or 2-alkyl-9,10-dihydroxyanthracene) to the corresponding 2-alkyl anthraquinone. Major producers commonly use either the 2-ethyl or the 2-amyl derivative. The cyclic reaction depicted below shows the 2-ethyl derivative, where 2-ethyl-9,10-dihydroxyanthracene (C16H14O2) is oxidized to the corresponding 2-ethylanthraquinone (C16H12O2) and hydrogen peroxide. Most commercial processes achieve this by bubbling compressed air through a solution of the anthracene, whereby the oxygen present in the air reacts with the labile hydrogen atoms (of the hydroxy group), giving hydrogen peroxide and regenerating the anthraquinone. Hydrogen peroxide is then extracted out 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.
In 1994, world production of H2O2 was around 1.9 million tonnes and grew to 2.2 million in 2006, most of which was at a concentration of 70% or less. In that year bulk 30% H2O2 sold for around US $0.54 per kg, equivalent to US $1.50 per kg (US $0.68 per lb) on a "100% basis".
A new, so-called "high-productivity/high-yield" process, based on an optimized distribution of isomers of 2-amyl anthraquinone, has been developed by Solvay. In July 2008, this process allowed the construction of a "mega-scale" single-train plant in Zandvliet (Belgium). The plant has an annual production capacity more than twice that of the world's next-largest single-train plant. An even-larger plant is scheduled to come onstream at Map Ta Phut (Thailand) in 2011. It can be imagined that this leads to reduction in the cost of production due to economies of scale.
A process to produce hydrogen peroxide directly from the elements has been of interest to producers for many years. The problem with the direct synthesis process is that, in terms of thermodynamics, the reaction of hydrogen with oxygen favors production of water. It had been recognized for some time that a finely dispersed catalyst is beneficial in promoting selectivity to hydrogen peroxide, but, while selectivity was improved, it was still not sufficiently high to permit commercial development of the process. However, an apparent breakthrough was made in the early 2000s by researchers at Headwaters Technology. The breakthrough revolves around development of a minute (nanometer-size) phase-controlled noble metal crystal particles on carbon support. This led, in a joint venture with Evonik Industries, to the construction of a pilot plant in Germany in late 2005. The pilot plant trials to test the commercial feasibility of the process are, it is presumed, ongoing, since little has been revealed about the results or progress of the operation. It is claimed that there are reductions in investment cost because the process is simpler and involves less equipment; however, the process is also more corrosive and unproven. It should be noted that this process results in low concentrations of hydrogen peroxide (about 5–10 wt% versus about 40 wt% through the anthraquione process).
In 2009, another catalyst development was announced by workers at Cardiff University. This development also relates to the direct synthesis, but, in this case, using gold–palladium nanoparticles. Under normal circumstances, the direct synthesis must be carried out in an acid medium to prevent immediate decomposition of the hydrogen peroxide once it is formed. While hydrogen peroxide has a tendency to decompose on its own (which is why, even after production, it is often necessary to add stabilisers to the commercial product when it is to be transported or stored for long periods), the nature of the catalyst can cause this decomposition to accelerate rapidly. It is claimed that the use of this gold-palladium catalyst reduces this decomposition and, as a consequence, little to no acid is required. The process is in a very early stage of development and currently results in very low concentrations of hydrogen peroxide being formed (less than about 1–2 wt%). Nonetheless, it is envisaged by the inventors that the process will lead to an inexpensive, efficient, and environmentally friendly process.
A novel electrochemical process for the production of alkaline hydrogen peroxide has been developed by Dow. The process employs a monopolar cell to achieve an electrolytic reduction of oxygen in a dilute sodium hydroxide solution.
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 (see decomposition); 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 >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 submit to inspection by the small number of commerical manufacturers.
This process is thermodynamically favorable. It has a ΔH
o of −98.2 kJ·mol−1 and a ΔG o of −119.2 kJ·mol−1 and a ΔS of 70.5 J·mol−1·K−1. 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 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.
The liberation of oxygen and energy in the decomposition has dangerous side-effects. Spilling high concentrations of hydrogen peroxide on a flammable substance can cause an immediate fire, which is further fueled by the oxygen released by the decomposing hydrogen peroxide. High test peroxide, or HTP (also called high-strength peroxide) must be stored in a suitable, vented container to prevent the buildup of oxygen gas, which would otherwise lead to the eventual rupture of the container.
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.
A common concentration for hydrogen peroxide is 20-volume, which means that, when 1 volume of hydrogen peroxide is decomposed, it produces 20 volumes of oxygen. A 20-volume concentration of hydrogen peroxide is equivalent to 1.667 mol/dm3 (Molar solution) or about 6%.
Hydrogen peroxide available at drug stores is three-percent solution. In such small concentrations, it is less stable, and decomposes faster. It is usually stabilized with acetanilide, a substance that has toxic side-effects in significant amounts.
H2O2 is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H2O2 can be converted into hydroxyl radicals (.OH) with reactivity second only to fluorine.
|Oxidant||Oxidation potential, V|
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 acidic solutions Fe2+ is oxidized to Fe3+ (hydrogen peroxide acting as an oxidizing agent),
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; for example, Mn2+ is oxidized to Mn4+ (as MnO2).
Hydrogen peroxide is frequently used as an oxidizing agent in organic chemistry. One application is for the oxidation of thioethers to sulfoxides. For example, methyl phenyl sulfide was oxidized to methyl phenyl sulfoxide in 99% yield in methanol in 18 hours (or 20 minutes using a TiCl3 catalyst):
For example, on addition to an aqueous solution of chromic acid (CrO3) or acidic solutions of dichromate salts, it will form an unstable blue peroxide CrO(O2)2. In aqueous solution it rapidly decomposes to form oxygen gas and chromium salts.
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, also known as trioxidane. 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.
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 an 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. It is used in the production of certain organic peroxides such as dibenzoyl peroxide, used in polymerisations 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 peroxy acid. Peracetic acid and meta-chloroperoxybenzoic acid (commonly abbreviated mCPBA) are prepared from acetic acid and meta-chlorobenzoic acid, respectively. The latter is commonly reacted with alkenes to give the corresponding epoxide.
In the PCB manufacturing process, hydrogen peroxide mixed with sulfuric acid was used as the microetch chemical for copper surface roughening preparation.
A combination of a powdered precious metal-based catalyst, hydrogen peroxide, methanol and water can produce superheated steam in one to two seconds, releasing only CO2 and high-temperature steam for a variety of purposes.
Recently, there has been increased use of vaporized hydrogen peroxide in the validation and bio-decontamination of half-suit and glove-port isolators in pharmaceutical production.
Nuclear pressurized water reactors (PWRs) use hydrogen peroxide during the plant shutdown to force the oxidation and dissolution of activated corrosion products deposited on the fuel. The corrosion products are then removed with the cleanup systems before the reactor is disassembled.
Hydrogen peroxide is also used in the oil and gas exploration industry to oxidize rock matrix in preparation for micro-fossil analysis.
A method of producing propylene oxide from hydrogen peroxide has been developed. The process is claimed to be environmentally friendly, since the only significant byproduct is water. It is also claimed the process has significantly lower investment and operating costs. Two of these "HPPO" (hydrogen peroxide to propylene oxide) plants came onstream in 2008: One of them located in Belgium is a Solvay, Dow-BASF joint venture, and the other in Korea is a EvonikHeadwaters, SK Chemicals joint venture. A caprolactam application for hydrogen peroxide has been commercialized. Potential routes to phenol and epichlorohydrin utilizing hydrogen peroxide have been postulated.
A study published in Nature found that hydrogen peroxide plays a role in the immune system. Scientists found that hydrogen peroxide is released after tissues are damaged in zebra fish, which is thought to act as a signal to white blood cells to converge on the site and initiate the healing process. When the genes required to produce hydrogen peroxide were disabled, white blood cells did not accumulate at the site of damage. The experiments were conducted on fish; however, because fish are genetically similar to humans, the same process is speculated to occur in humans. The study in Nature suggested Asthma sufferers have higher levels of hydrogen peroxide in their lungs than healthy people, which could be explained by asthma sufferers having inappropriate levels of white blood cells in their lungs.
H2O2 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, which is expelled through a nozzle, generating thrust. H2O2 monopropellant produces a maximum specific impulse (Isp) of 161 s (1.6 kN·s/kg), which makes it a low-performance monopropellant. Peroxide generates much less thrust than hydrazine, but is not toxic. The Bell Rocket Belt used hydrogen peroxide monopropellant.
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, 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 in World-War-II German rockets (e.g. T-Stoff for the Me-163), and for the low-cost British Black Knight and Black Arrow launchers.
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 require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant, but this was dangerous and has been discontinued by most navies. 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. 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.
While rarely used now as a monopropellant for large engines, small hydrogen peroxide attitude control thrusters are still in use on some satellites. They are easy to throttle, and safer to fuel and handle before launch than hydrazine thrusters. However, hydrazine is more often used in spacecraft because of its higher specific impulse and lower rate of decomposition.
Hydrogen peroxide has been used as an antiseptic and anti-bacterial agent for many years due to its oxidizing effect. While its use has decreased in recent years with the popularity of readily available over the counter products, it is still used by many hospitals, doctors and dentists.
Regulations vary, but low concentrations, such as 3%, are widely available and legal to buy for medical use. Higher concentrations may be considered hazardous and are typically accompanied by a 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 H2O2 will react violently.
High-concentration hydrogen peroxide streams, typically above 40%, should be considered a D001 hazardous waste, due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds, or approximately ten gallons, 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). 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 filter out light.
Hydrogen peroxide, either in pure or diluted form, can pose several risks:
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[[File:|thumb|A picture to show simply how the atoms may fill space. The white is hydrogen and the red is oxygen.]] Hydrogen peroxide is a chemical compound. Its molecular formula is H2O2. It is used as a cleaner, and as hair bleach. In a concentration of 3% (meaning that there are 3 of hydrogen peroxide for 97 of water), it can be used to treat wounds. When mixed with water, hydrogen peroxide can add oxygen gas (O2) to or break down a molecules with charges, but without carbon. 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.
It is also classified as water like substance with another oxygen atom in it.[needs proof]