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Hydrogen peroxide

2007 Schools Wikipedia Selection. Related subjects: Chemical compounds

                          Hydrogen peroxide
                          Hydrogen peroxide

   Hydrogen peroxide Hydrogen peroxide
                               General
   Systematic name     Dihydrogen dioxide
   Other names         Hydrogen peroxide
                       hydrogen dioxide
   Molecular formula   H[2]O[2]
   Molar mass          34.0147 g/mol.
   Appearance          Very pale blue colour; colorless in solution.
   CAS number          [7722-84-1]
                              Properties
   Density and phase   1.4 g/cm^3, 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 c P at 20 °C
                              Structure
   Molecular shape      ?
   Dipole moment       2.26 D
                               Hazards
   MSDS                30% hydrogen peroxide msds
                       60% hydrogen peroxide msds
   Main hazards        Oxidant, corrosive.
   NFPA 704

                       0
                       3
                       1
                       OX
   Flash point         Non-flammable.
   R/S statement       R: R5, R8, R20, R22,R35
                       S: S1, S2, S17, S26,S28,
                       S36, S37, S39, S45
   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)
   Infobox disclaimer and references

   Hydrogen peroxide (H[2]O[2]) is a very pale blue liquid which appears
   colourless in a dilute solution, 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 byproduct. Thenard's process was used
   from the end of the 19th century until the middle of the 20th century.
   Modern manufacture methods are discussed below.

Uses

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 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 "per-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.

Domestic uses

   Diluted H[2]O[2] (around 5%) is used to bleach human hair, hence the
   phrases peroxide blonde and bottle blonde. It can absorb into skin upon
   contact and create a local skin capillary embolism which appears as a
   temporary whitening of the skin. It whitens skeletons that are to be
   put on display. 3% H[2]O[2] is used medically for cleaning wounds,
   removing dead tissue, or as an oral debriding agent. Most
   over-the-counter peroxide solutions are not, however, suitable for
   ingestion.

   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. See ectoparasite.

   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 exposed to catalysts.

   Hydrogen peroxide is increasingly popular 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 H[2]S) is in the -2 state. In a neutral
   solution, hydrogen peroxide will oxidize hydrogen sulfide to elemental
   sulfur via the following reaction: 8 H[2]S(g) + 8 H[2]O[2](aq) →
   S[8](s) + 16 H[2]O(l)

   The reaction is slow but may be catalyzed by metal ions. 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 & H[2]S) add the above values to determine
   the total ppm feed needed to oxidize two or more.

   Hydrogen peroxide is a 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 odours. 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 CO[2] dimer which
   in turn causes an added dye to reach an excited state, the latter
   relaxing to release photons of light.

Storage

   Small quantities of many different concentrations and grades can be
   legally stored and used with few regulations.

   Hydrogen peroxide should be stored in a container made from a material
   that doesn't react with the chemical. Numerous materials and processes
   are available, and these vary based on the concentration and grade
   (purity) of the hydrogen peroxide. In general, it is an oxidizer and
   should be stored away from fuel sources and sources of catalytic
   contamination. Because oxygen is formed during the natural
   decomposition of the peroxide, the resulting increase in pressure can
   cause a glass container to break. Therfore, H[2]O[2] should be stored
   in vented plastic containers.

Use as propellant

   H[2]O[2] 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, and the hot (>600 °C)
   oxygen/steam produced is used directly for thrust. H[2]O[2]
   monopropellant produces a maximum specific impulse (I[sp]) of 161 s
   (1.6 kN·s/kg), which makes it a low-performance monopropellant.
   Compared to hydrazine, peroxide is less toxic, but it is also much less
   powerful. The famous Bell Rocket Belt used hydrogen peroxide
   monopropellant.

   As a bipropellant, H[2]O[2] 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 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. Peroxide was used very successfully as an oxidizer for early
   World-War-II era German rockets, and for the low-cost British
   launchers, Black Knight and Black Arrow.

   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 H[2]O[2] 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.

   Recently, H2O2/ propylene as an approach to inexpensive Single Stage To
   Orbit has been proposed; this involves a main fuel tank containing
   propylene, with a bladder floating in it containing the H2O2. This
   combination offers 15% superior ISP to O2/RP4 (a kerosene used as
   rocket propellant), avoiding the need for turbines, cryogenic storage
   or hardware, and greatly reduced cost for the construction of the
   booster; the potential of this and other alternate systems is discussed
   in some detail at Dunn Engineering which is offered as a citation.

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.

   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." Internal use of hydrogen
   peroxide has a history of causing fatal blood disorders, and its recent
   use as a therapeutic treatment has been linked to several deaths.^,

   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. Like
   benzoyl peroxide, hydrogen peroxide is also sometimes used in the
   treatment of acne.

   Hydrogen peroxide is also used as an emetic in veterinary practice.

Physical properties

   Structure of hydrogen peroxide

   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). The increased barrier is also attributed to
   lone-pair lone-pair repulsion. 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 H[2]O[2].

Chemical properties

   H[2]O[2] is one of the most powerful oxidizers known -- stronger than
   chlorine, chlorine dioxide, and potassium permanganate. And through
   catalysis, H[2]O[2] 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 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 (disproportionates) exothermically
   into water and oxygen gas spontaneously:

          2 H[2]O[2] → 2 H[2]O + O[2] + Energy

   This process is very favorable; it has a ΔH^[DEL: o :DEL] of −98.2 kJ/
   mol and a ΔG^[DEL: o :DEL] 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 Fe^2+ or Ti^3+, the
   decomposition may take a different path, with free radicals such as HO·
   ( hydroxyl) and HOO· being formed. A combination of H[2]O[2] and Fe^2+
   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 Fe^2+ is oxidized to Fe^3+,

          2 Fe^2+(aq) + H[2]O[2] + 2 H^+(aq) → 2 Fe^3+(aq) + 2H[2]O(l)

   and sulfite (SO[3]^2−) is oxidized to sulfate (SO[4]^2−). However,
   potassium permanganate is reduced to Mn^2+ by acidic H[2]O[2]. Under
   alkaline conditions, however, some of these reactions reverse; Mn^2+ is
   oxidized to Mn^4+ (as MnO[2]), yet Fe^3+ is reduced to Fe^2+.

          2 Fe^3+ + H[2]O[2] + 2 OH^− → 2 Fe^2+ + 2 H[2]O + O[2]

   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 TiCl[3] catalyst):

          Ph-S-CH[3] + H[2]O[2] → Ph-S(O)-CH[3] + H[2]O

   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 (CrO[3]), it can form an unstable blue
   peroxide CrO(O[2])[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:

          Na[2]B[4]O[7] + 4 H[2]O[2] + 2 NaOH → 2 Na[2]B[2]O[4](OH)[4] +
          H[2]O

   H[2]O[2] 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 H[2]O[2] in some reactions.

   Hydrogen peroxide reacts with ozone to form trioxidane.

Alkalinity

   Hydrogen peroxide is a much weaker base than water, but it can still
   form adducts with very strong acids. The superacid HF/ SbF[5] forms
   unstable compounds containing the [H[3]O[2]]^+ 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:

   H[2] + O[2] → H[2]O[2]

   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
   (NH[4]HSO[4]), followed by hydrolysis of the peroxydisulfate
   ((SO[4])[2])^2− which is formed.

   In 1994, world production of H[2]O[2] was around 1.9 million tonnes,
   most of which was at a concentration of 70% or less. In that year bulk
   30% H[2]O[2] 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-- roughly over 70%. Although any concentration of
   peroxide will generate some hot gas (oxygen plus some steam), at
   concentrations above approximately 67%, the heat of decomposing
   hydrogen peroxide becomes large enough to completely vaporize all the
   liquid at standard temperature. This represents a safety turning point,
   since decomposition of any concentration above this amount is capable
   of transforming the liquid entirely to heated gas (the higher the
   concentration, the hotter the resulting gas), and this hot steam/oxygen
   mixture can then be used to generate maximal thrust, power, or work.

   Normal propellant grade concentrations therefore vary from 70 to 98%,
   with common grades of 70, 85, 90, and 98%. Many of these grades and
   variations are described in detail in the United States propellant
   specification number MIL-P-16005 Revision F, which is currently
   available. The available suppliers of high concentration propellant
   grade hydrogen peroxide are generally one of the large commercial
   companies which make other grades of hydrogen peroxide; including
   Solvay Interox, FMC, and Degussa. Other companies which have made
   propellant grade hydrogen peroxide in the recent past include Air
   Liquide and DuPont. Note that DuPont recently sold its hydrogen
   peroxide manufacturing business to Degussa.

   Propellant grade hydrogen peroxide is available to qualified buyers.
   Typically this chemical is only sold to commercial companies or
   government institutions which have the ability to properly handle and
   utilize the material.

   Non-professionals have purchased 70% or lower concentration hydrogen
   peroxide (the remaining 30% is water with traces of impurities and
   stabilizing materials, such as tin salts, phosphates, nitrates, and
   other chemical additives), and increased its concentration themselves -
   a potentially extremely dangerous practice that should not be
   encouraged. Many amateurs try distillation, but this is extremely
   dangerous with hydrogen peroxide; peroxide vapor can ignite or detonate
   depending on specific combinations of temperature and pressure. In
   general any boiling mass of high concentration hydrogen peroxide at
   ambient pressure will produce vapor phase hydrogen peroxide which can
   detonate. This hazard is mitigated, but not entirely eliminated with
   vacuum distillation. Vacuum distillation of propellant grade hydrogen
   peroxide is still hazardous and is best done by qualified laboratories
   or companies. Other approaches for concentrating hydrogen peroxide are
   sparging and fractional crystallization.

   High concentration hydrogen peroxide is readily available in 70, 90,
   and 98% concentrations in sizes of 1 gallon, 30 gallon, and bulk tanker
   truck volumes. Propellant grade hydrogen peroxide is being used on
   current military systems and is in numerous defense and aerospace
   research and development programs. Many privately funded rocket
   companies are using hydrogen peroxide, notably Blue Origin. Some
   amateur groups have expressed interest in manufacturing their own
   peroxide, for their use and for sale in small quantities to others. The
   production of hydrogen peroxide by amateurs is potentially dangerous to
   both the producers of the chemical, persons in the vicinity of the
   chemical, and users of the chemical.

Hazards

   Hydrogen peroxide, either in pure or diluted form, can pose several
   risks:
     * Above roughly 70% concentrations, hydrogen peroxide can give off
       vapor that can detonate above 70 °C (158 °F) at normal atmospheric
       pressure. This can then BLEVE the remaining liquid. Distillation of
       hydrogen peroxide at normal pressures is thus highly dangerous, and
       must be avoided.

     * Hydrogen peroxide vapors 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.
       The saying is 'peroxides kill chemists'.

     * 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 almost immediately.

     * 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%.

   Hydrogen peroxide is naturally produced as a byproduct of oxygen
   metabolism, and virtually all organisms possess enzymes known as
   peroxidases, which apparently harmlessly catalytically decomposes low
   concentrations of hydrogen peroxide to water and oxygen (see
   Decomposition above).

   In one incident, several people were injured after a hydrogen peroxide
   spill on board Northwest Airlines Flight 957 because they mistook it
   for water.

   For more information on the risks of working with this chemical,
   consult an MSDS. It reacts to make water and oxygen gas.
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