3D model (JSmol)
|UN number||2015 (>60% soln.)|
2014 (20-60% soln.)
2984 (8-20% soln.)
CompTox Dashboard (EPA)
|Molar mass||34.0147 g/mol|
|Appearance||Very light blue color; colorless in solution|
|Density||1.11 g/cm3 (20 °C, 30% (w/w) solution )|
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)|
|Solubility||soluble in ether, alcohol|
insoluble in petroleum ether
|Vapor pressure||5 mmHg (30 °C)|
Refractive index (nD)
|Viscosity||1.245 cP (20 °C)|
Heat capacity (C)
|1.267 J/(g·K) (gas)|
2.619 J/(g·K) (liquid)
Std enthalpy of
|A01AB02 (WHO) D08AX01 (WHO), D11AX25 (WHO), S02AA06 (WHO)|
|Safety data sheet||ICSC 0164 (>60% soln.)|
|GHS signal word||danger|
|H271, H302, H314, H332, H335, H412|
|P280, P305+351+338, P310|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
2000 mg/kg (oral, mouse)
LC50 (median concentration)
|1418 ppm (rat, 4 hr)|
LCLo (lowest published)
|227 ppm (mouse)|
|US health exposure limits (NIOSH):|
|TWA 1 ppm (1.4 mg/m3)|
|TWA 1 ppm (1.4 mg/m3)|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Hydrogen peroxide is a chemical compound with the formula . In its pure form, it is a very pale blue, clear liquid, slightly more viscous than water. Hydrogen peroxide is the simplest peroxide (a compound with an oxygen-oxygen single bond). It is used as an oxidizer, bleaching agent, and antiseptic. Concentrated hydrogen peroxide, or "high-test peroxide", is a reactive oxygen species and has been used as a propellant in rocketry. Its chemistry is dominated by the nature of its unstable peroxide bond.
Hydrogen peroxide is unstable and slowly decomposes in the presence of light. Because of its instability, hydrogen peroxide is typically stored with a stabilizer in a weakly acidic solution. Hydrogen peroxide is found in biological systems including the human body. Enzymes that use or decompose hydrogen peroxide are classified as peroxidases.
The boiling point of has been extrapolated as being 150.2 °C, approximately 50 °C higher than water. In practice, hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled at lower temperatures under reduced pressure.
In aqueous solutions hydrogen peroxide differs from the pure substance due to the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression; 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.
Hydrogen peroxide is a nonplanar molecule as shown by Paul-Antoine Giguère in 1950 using infrared spectroscopy, with (twisted) C2symmetry. Although the O-O bond is a single bond, the molecule has a relatively high rotational barrier of 2460 cm-1 (29.45 kJ/mol); for comparison, the rotational barrier for ethane is 12.5 kJ/mol. The increased barrier is ascribed to repulsion between the lone pairs of the adjacent oxygen atoms and results in hydrogen peroxide displaying atropisomerism.
The molecular structures of gaseous and crystalline are significantly different. This difference is attributed to the effects of hydrogen bonding, which is absent in the gaseous state. Crystals of are tetragonal with the space group D4
Hydrogen peroxide has several structural analogues with Hm-X-X-Hn bonding arrangements (water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, N, S). Its melting point is also fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising significantly more readily, indicative of particularly strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. All of these analogues are thermodynamically unstable. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs.
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 (French: oxygenated water) - subsequently known as hydrogen peroxide. (Today this term refers instead to water containing dissolved oxygen (O2).)
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.
Thénard and Joseph Louis Gay-Lussac synthesized sodium peroxide in 1811. The bleaching effect of peroxides and their salts on natural dyes became known around that time, but early attempts of industrial production of peroxides failed, and the first plant producing hydrogen peroxide was built in 1873 in Berlin. The discovery of the synthesis of hydrogen peroxide by electrolysis with sulfuric acid introduced the more efficient electrochemical method. It was first implemented into industry in 1908 in Weißenstein, Carinthia, Austria. The anthraquinone process, which is still used, was developed during the 1930s by the German chemical manufacturer IG Farben in Ludwigshafen. The increased demand and improvements in the synthesis methods resulted in the rise of the annual production of hydrogen peroxide from 35,000 tonnes in 1950, to over 100,000 tonnes in 1960, to 300,000 tonnes by 1970; by 1998 it reached 2.7 million tonnes.
Pure hydrogen peroxide was long believed to be unstable, as early attempts to separate it from the water, which is present during synthesis, all failed. This instability was due to traces of impurities (transition-metal salts), which catalyze the decomposition of the hydrogen peroxide. Pure hydrogen peroxide was first obtained in 1894--almost 80 years after its discovery--by Richard Wolffenstein, who produced it by vacuum distillation.
Determination of the molecular structure of hydrogen peroxide proved to be very difficult. In 1892 the Italian physical chemist Giacomo Carrara (1864-1925) determined its molecular mass by freezing-point depression, which confirmed that its molecular formula is H2O2. At least half a dozen hypothetical molecular structures seemed to be consistent with the available evidence. In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide that was very similar to the presently accepted one.
Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process, which was formalized in 1936 and patented in 1939. It begins with the reduction of an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) to the corresponding anthrahydroquinone, typically by hydrogenation on a palladium catalyst. 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.
The net reaction for the anthraquinone-catalyzed process is simply:
A process to produce hydrogen peroxide directly from the elements has been of interest for many years. Direct synthesis is difficult to achieve, as the reaction of hydrogen with oxygen thermodynamically favours production of water. Systems for direct synthesis have been developed, most of which employ finely dispersed metal catalysts similar to those used for hydrogenation of organic substrates. None of these has yet reached a point where they can be used for industrial-scale synthesis.
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 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% sold for around 0.54 USD/kg, equivalent to US$1.50/kg (US$0.68/lb) on a "100% basis".
Hydrogen peroxide occurs in surface water, groundwater and in the atmosphere. It forms upon illumination or natural catalytic action by substances contained in water. Sea water contains 0.5 to 14 ?g/L of hydrogen peroxide, freshwater 1 to 30 ?g/L and air 0.1 to 1 parts per billion.
The rate of decomposition increases with rise in temperature, concentration, and pH, with cool, dilute, acidic solutions showing the best stability. Decomposition is catalysed by various compounds, including most transition metals and their compounds (e.g. manganese dioxide (MnO2), silver, and platinum). Certain metal ions, such as or , can cause the decomposition to take a different path, with free radicals such as the hydroxyl radical (HO·) and hydroperoxyl (HOO·) being formed. Non-metallic catalysts include potassium iodide, which reacts particularly rapidly and forms the basis of the elephant toothpaste experiment. Hydrogen peroxide can also be decomposed biologically by the enzyme catalase. The decomposition of hydrogen peroxide liberates oxygen and heat; this can be dangerous, as spilling high-concentration hydrogen peroxide on a flammable substance can cause an immediate fire.
Hydrogen peroxide exhibits oxidizing and reducing properties, depending on pH.
In acidic solutions, is one of the most powerful oxidizers known and is stronger than chlorine, chlorine dioxide, and potassium permanganate. When used for removing organic stains from laboratory glassware it is referred to as Piranha solution. Also, through catalysis, can be converted into hydroxyl radicals (·OH), which are highly reactive.
In acidic solutions is oxidized to (hydrogen peroxide acting as an oxidizing agent):
In basic solution, hydrogen peroxide can reduce a variety of inorganic ions. When it acts as a reducing agent, oxygen gas is also produced. For example, hydrogen peroxide will reduce sodium hypochlorite and potassium permanganate, which is a convenient method for preparing oxygen in the laboratory:
Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acid derivatives, and for the oxidation of alkylboranes to alcohols, the second step of hydroboration-oxidation. It is also the principal reagent in the Dakin oxidation process.
It also converts metal oxides into the corresponding peroxides. For example, upon treatment with hydrogen peroxide, chromic acid( + ) forms an unstable blue peroxide CrO(.
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. An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for 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 manganous ions are reoxidized to make the reaction catalytic. This forms the basis for common monopropellant rockets.
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. 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.
Peroxisomes are organelles found in virtually all eukaryotic cells. They are involved in the catabolism of very long chain fatty acids, branched chain fatty acids, D-amino acids, polyamines, and biosynthesis of plasmalogens, etherphospholipids critical for the normal function of mammalian brains and lungs. Upon oxidation, they produce hydrogen peroxide in the following process:
This reaction is important in liver and kidney cells, where the peroxisomes neutralize various toxic substances that enter the blood. Some of the ethanol humans drink is oxidized to acetaldehyde in this way. In addition, when excess H2O2 accumulates in the cell, catalase converts it to H2O through this reaction:
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:
Eggs of sea urchin, shortly after fertilization by a sperm, produce hydrogen peroxide. It is then quickly dissociated to OH· radicals. The radicals serve as initiator of radical polymerization, which surrounds the eggs with a protective layer of polymer.
The bombardier beetle has a device which allows it to shoot corrosive and foul-smelling bubbles at its enemies. The beetle produces and stores hydroquinone and hydrogen peroxide, in two separate reservoirs in the rear tip of its abdomen. When threatened, the beetle contracts muscles that force the two reactants through valved tubes into a mixing chamber containing water and a mixture of catalytic enzymes. When combined, the reactants undergo a violent exothermic chemical reaction, raising the temperature to near the boiling point of water. The boiling, foul-smelling liquid partially becomes a gas (flash evaporation) and is expelled through an outlet valve with a loud popping sound.
Hydrogen peroxide has roles as a signalling molecule in the regulation of a wide variety of biological processes. The compound is a major factor implicated in the free-radical theory of aging, based on how readily hydrogen peroxide can decompose into a hydroxyl radical and how superoxide radical byproducts of cellular metabolism can react with ambient water to form hydrogen peroxide. These hydroxyl radicals in turn readily react with and damage vital cellular components, especially those of the mitochondria. At least one study has also tried to link hydrogen peroxide production to cancer. These studies have frequently been quoted in fraudulent treatment claims.
The second major industrial application is the manufacture of sodium percarbonate and sodium perborate, which are used as mild bleaches in laundry detergents. Sodium percarbonate, which is an adduct of sodium carbonate and hydrogen peroxide, is the active ingredient in such laundry products as OxiClean and Tide laundry detergent. When dissolved in water, it releases hydrogen peroxide and sodium carbonate,
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 is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example. It is used in polymerisations, as a flour bleaching agent, and as a treatment for acne. Peroxy acids, such as peracetic acid and meta-chloroperoxybenzoic acid also are produced using hydrogen peroxide. Hydrogen peroxide has been used for creating organic peroxide-based explosives, such as acetone peroxide.
Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. In advanced oxidation processing, the Fenton reaction gives the highly reactive hydroxyl radical (·OH). This degrades organic compounds, including those that are ordinarily robust, such as aromatic or halogenated compounds. It can also oxidize sulfur based compounds present in the waste; which is beneficial as it generally reduces their odour.
Hydrogen peroxide may be used for the sterilization of various surfaces, including surgical tools, and may be deployed as a vapour (VHP) for room sterilization. H2O2 demonstrates broad-spectrum efficacy against viruses, bacteria, yeasts, and bacterial spores. In general, greater activity is seen against Gram-positive than Gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms may increase tolerance in the presence of lower concentrations. Higher concentrations of H2O2 (10 to 30%) and longer contact times are required for sporicidal activity.
Hydrogen peroxide is seen as an environmentally safe alternative to chlorine-based bleaches, as it degrades to form oxygen and water and it is generally recognized as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA).
Historically hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics. Now it is thought to inhibit healing and to induce scarring, because it destroys newly formed skin cells. One study found that only very low concentrations (0.03% solution, this is a dilution of typical 3% Peroxide by 100 times) may induce healing, and only if not applied repeatedly. A 0.5% solution was found to impede healing. Surgical use can lead to gas embolism formation. Despite this, it is still used for wound treatment in many countries, and, in the United States, is prevalent as a major first aid antiseptic.
Dermal exposure to dilute solutions of hydrogen peroxide causes whitening or bleaching of the skin due to microembolism caused by oxygen bubbles in the capillaries.
Diluted (between 1.9% and 12%) mixed with aqueous ammonia has been used to bleach human hair. The chemical's bleaching property lends its name to the phrase "peroxide blonde". Hydrogen peroxide is also used for tooth whitening. It may be found in most whitening toothpastes. Hydrogen peroxide has shown positive results involving teeth lightness and chroma shade parameters. It works by oxidizing colored pigments onto the enamel where the shade of the tooth indeed, may become lighter. Hydrogen peroxide may be mixed with baking soda and salt to make a home-made toothpaste.
Practitioners of alternative medicine have advocated the use of hydrogen peroxide for various conditions, including emphysema, influenza, AIDS, and cancer, although there is no evidence of effectiveness and in some cases, it may even be fatal.
The practice calls for the daily consumption of hydrogen peroxide, either orally or by injection, and is based on two assumptions. First, that hydrogen peroxide is produced naturally by the body to combat infection; and second, that human pathogens (including cancer: See Warburg hypothesis) are anaerobic and cannot survive in oxygen-rich environments. The ingestion or injection of hydrogen peroxide therefore is believed to kill disease by mimicking the immune response in addition to increasing levels of oxygen within the body. This makes the practice similar to other oxygen-based therapies, such as ozone therapy and hyperbaric oxygen therapy.
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. 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 difficult to raise the level of oxygen around cancer cells within a tumour, as the blood supply tends to be poor, a situation known as tumor hypoxia.
Large oral doses of hydrogen peroxide at a 3% concentration may cause irritation and blistering to the mouth, throat, and abdomen as well as abdominal pain, vomiting, and diarrhea.Intravenous injection of hydrogen peroxide has been linked to several deaths.
The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective, or useful cancer treatment." Furthermore, the therapy is not approved by the U.S. FDA.
High-concentration is referred to as "high-test peroxide" (HTP). It can be used either as a monopropellant (not mixed with fuel) or as the oxidizer component of a bipropellant rocket. Use as a monopropellant takes advantage of the decomposition of 70-98% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber, where a catalyst, usually a silver or platinum screen, triggers decomposition, producing steam at over 600 °C (1,112 °F), which is expelled through a nozzle, generating thrust. monopropellant produces a maximal specific impulse (Isp) of 161 s (1.6 kN·s/kg). Peroxide was the first major monopropellant adopted for use in rocket applications. Hydrazine eventually replaced hydrogen-peroxide monopropellant thruster applications primarily because of a 25% increase in the vacuum specific impulse. Hydrazine (toxic) and hydrogen peroxide (less-toxic [ACGIH TLV 0.01 and 1 ppm respectively]) are the only two monopropellants (other than cold gases) to have been widely adopted and utilized for propulsion and power applications. The Bell Rocket Belt, reaction control systems for X-1, X-15, Centaur, Mercury, Little Joe, as well as the turbo-pump gas generators for X-1, X-15, Jupiter, Redstone and Viking used hydrogen peroxide as a monopropellant.
As a bipropellant, 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 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. T-Stoff, containing oxyquinoline stabilizer, for both the Walter HWK 109-500 Starthilfe RATO externally podded monopropellant booster system, and for the Walter HWK 109-509 rocket motor series used for the Me 163B), most often used with C-Stoff in a self-igniting hypergolic combination, and for the low-cost British Black Knight and Black Arrow launchers.
In the 1940s and 1950s, the Hellmuth Walter KG-conceived turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant. Operator error in the use of hydrogen-peroxide torpedoes was named as possible causes for the sinkings of HMS Sidon and the Russian submarine Kursk. SAAB Underwater Systems is manufacturing the Torpedo 2000. This torpedo, used by the Swedish Navy, is powered by a piston engine propelled by HTP as an oxidizer and kerosene as a fuel in a bipropellant system.
Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.
Some horticulturalists and users of hydroponics advocate the use of weak hydrogen peroxide solution in watering solutions. Its spontaneous decomposition releases oxygen that enhances a plant's root development and helps to treat root rot (cellular root death due to lack of oxygen) and a variety of other pests.
Laboratory tests conducted by fish culturists in recent years have demonstrated that common household hydrogen peroxide 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.
Regulations vary, but low concentrations, such as 6%, are widely available and legal to buy for medical use. Most over-the-counter peroxide solutions are not suitable for ingestion. Higher concentrations may be considered hazardous and 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 will react violently.
High-concentration hydrogen peroxide streams, typically above 40%, should be considered hazardous due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations, if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds (45 kg), or approximately 10 US gallons (38 L), of concentrated hydrogen 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 block light.
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. Highly concentrated hydrogen peroxide is unstable and may cause a boiling liquid expanding vapour explosion (BLEVE) of the remaining liquid. Distillation of hydrogen peroxide at normal pressures is thus, highly dangerous. It is also corrosive, especially when concentrated, but even domestic-strength solutions may 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 (ten times the volume of a 3% solution), leading to internal bloating. Inhaling over 10% can cause severe pulmonary irritation.
With a significant vapour pressure (1.2 kPa at 50 °C), hydrogen-peroxide vapour is potentially hazardous. According to U.S. NIOSH, the immediately dangerous to life and health (IDLH) limit is only 75 ppm. The U.S. Occupational Safety and Health Administration (OSHA) has established a permissible exposure limit of 1.0 ppm calculated as an 8-hour time-weighted average (29 CFR 1910.1000, Table Z-1). Hydrogen peroxide also has been classified by the American Conference of Governmental Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans". For workplaces where there is a risk of exposure to the hazardous concentrations of the vapours, continuous monitors for hydrogen peroxide should be used. Information on the hazards of hydrogen peroxide is available from OSHA  and from the ATSDR.
Hydrogen peroxide was discovered in 1818 by the French chemist Louis-Jacques Thenard, who named it eau oxygénée (oxygenated water).