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| [[File:Peroxide group v.2.png|thumb|upright|Types of peroxides, from top to bottom: peroxide ion, [[organic peroxide]], organic hydroperoxide, [[peracid]]. The peroxide group is marked in '''{{Font color|blue|blue}}'''. R, R<sup>1</sup> and R<sup>2</sup> mark hydrocarbon [[moiety (chemistry)|moieties]].]]
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| A '''peroxide''' is a compound containing an [[oxygen]]–oxygen [[chemical bond|single bond]] or the peroxide [[anion]], O<sub>2</sub><sup>2–</sup>.<ref>{{GoldBookRef|title=peroxides|file=P04510}}</ref> The O−O group is called the '''peroxide group''' or '''peroxo group'''. In contrast to [[oxide|oxide ions]], the oxygen [[atoms]] in the peroxide ion have an [[oxidation state]] of −1.<ref>{{GoldBookRef|title=oxidation state|file=O04365}}</ref>
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| The simplest stable peroxide is [[hydrogen peroxide]]. [[Superoxide]]s, [[dioxygenyl]]s, [[ozone]]s and [[ozonide]]s are considered separately. Peroxide compounds can be roughly classified into [[organic chemistry|organic]] and inorganic. Whereas the inorganic peroxides have an ionic, salt-like character, the organic peroxides are dominated by the covalent bonds. The oxygen-oxygen [[chemical bond]] of peroxide is unstable and easily split into reactive [[radical (chemistry)|radical]]s via [[homolysis (chemistry)|homolytic cleavage]]. For this reason, peroxides are found in nature only in small quantities, in [[water]], [[atmosphere]], plants, and animals. Peroxide ion formation has recently been highlighted as one of the main mechanisms by which oxides accommodate excess oxygen in [[ion]]ic [[crystals]] and may have a large impact on a range of industrial applications including solid oxide [[fuel cells]].<ref>
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| {{cite journal|author=Middleburgh S. C., Lagerlof K.P.D. and Grimes R.W.|title= Accommodation of Excess Oxygen in Group II Oxides |doi=10.1111/j.1551-2916.2012.05452.x|year=2013|journal=Journal of the American Ceramic Society|volume=96|pages=308}}</ref>
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| Peroxides have a [[bleach]]ing effect on organic substances and therefore are added to some [[detergent]]s and [[hair coloring|hair colorants]]. Other large-scale applications include medicine and [[chemical industry]], where peroxides are used in various synthesis reactions or occur as intermediate products. With an annual production of over 2 million [[tonne]]s, hydrogen peroxide is the most economically important peroxide. Many peroxides are unstable and hazardous substances; they cannot be stored and therefore are synthesized ''in situ'' and used immediately.
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| ==In biochemistry==
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| [[File:Ascaridol2.png|thumb|150px|[[Ascaridole]]]]
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| Peroxides are usually very reactive and thus occur in nature only in a few forms. These include, in addition to hydrogen peroxide, a few vegetable products such as [[ascaridole]] and a peroxide derivative of [[prostaglandin]]. Hydrogen peroxide occurs in surface water, groundwater and in the [[atmosphere]]. It forms upon illumination or natural [[catalyst|catalytic]] action by substances containing in [[water]]. Sea water contains 0.5 to 14 mg/L of hydrogen peroxide, freshwater 1 to 30 mg/L and air 0.1 to 1 parts per billion.<ref name=offer/>
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| Hydrogen peroxide is formed in human and animal organisms as a short-lived product in biochemical processes and is [[Toxicity|toxic]] to [[Cell (biology)|cells]]. The toxicity is due to oxidation of [[protein]]s, [[membrane lipid]]s and [[DNA]] by the peroxide ions.<ref>Löffler G. and Petrides, P. E. ''Physiologische Chemie''. 4 ed., p. 288, Springer, Berlin 1988, ISBN 3-540-18163-6 (in German)</ref> The class of biological [[enzyme]]s called SOD ([[superoxide dismutase]]) 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.<ref>Löffler G. and Petrides, P. E. ''Physiologische Chemie''. 4 ed., pp. 321–322, Springer, Berlin 1988, ISBN 3-540-18163-6 (in German)</ref>
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| : <math>\mathrm{2\ O_2 ^-\ +\ 2\ H^+\ \xrightarrow {SOD}\ \ H_2O_2 +\ O_2}</math>
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| : <small> Formation of hydrogen peroxide by superoxide dismutase (SOD) </small>
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| [[Peroxisome]]s are [[organelle]]s found in virtually all [[eukaryotic]] cells.<ref name="pmid20124343">{{cite journal|author = Gabaldón T|title = Peroxisome diversity and evolution|journal =Philos Trans R Soc Lond B Biol Sci. |volume = 365|issue = 1541| pages = 765–73|year = 2010|pmid = 20124343|pmc = 2817229|doi = 10.1098/rstb.2009.0240}}</ref> They are involved in the [[catabolism]] of [[very long chain fatty acid]]s, [[Branched-chain-fatty-acid kinase|branched chain fatty acids]], [[D-amino acids]], [[polyamine]]s, and biosynthesis of [[plasmalogens]], etherphospholipids critical for the normal function of mammalian brains and lungs.<ref name="pmid16756494">{{cite journal|author = Wanders RJ, Waterham HR|title = Biochemistry of mammalian peroxisomes revisited|journal = Annu. Rev. Biochem.|volume = 75|pages = 295–332|year = 2006|pmid = 16756494|doi = 10.1146/annurev.biochem.74.082803.133329}}</ref> Upon oxidation, they produce hydrogen peroxide in the following process:<ref>Nelson, David; Cox, Michael; Lehninger, Albert L. and Cox, Michael M. [http://books.google.com/books?id=wuLQCAOtC4MC&pg=PA663 Lehninger Biochemie], pp. 663–664, Springer, 2001, ISBN 3-540-41813-X (in German)</ref>
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| : <math>\mathrm{R{-}CH_2{-}CH_2{-}CO{-}SCoA\ +\ O_2\ \xrightarrow {FAD}\ \ R{-}CH{=}CH{-}CO{-}SCoA\ +\ H_2O_2}</math>
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| : <small> FAD = [[flavin adenine dinucleotide]] </small>
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| [[Catalase]], another peroxisomal enzyme, uses this H<sub>2</sub>O<sub>2</sub> to oxidize other substrates, including [[phenols]], [[formic acid]], [[formaldehyde]], and [[alcohol]], by means of the peroxidation reaction:
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| :<math>\mathrm{H}_2\mathrm{O}_2 + \mathrm{R'H}_2 \rightarrow \mathrm{R'} + 2\mathrm{H}_2\mathrm{O}</math>, thus eliminating the poisonous hydrogen peroxide in the process.
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| 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.<ref>Riley, Edward P. ''et al.'' (ed.) [http://books.google.com/books?id=TiSL4txuYN0C&pg=PA112 Fetal Alcoholspectrum Disorder Fasd: Management and Policy Perspectives], Wiley-VCH, 2010, ISBN 3-527-32839-4 p. 112</ref> In addition, when excess H<sub>2</sub>O<sub>2</sub> accumulates in the cell, catalase converts it to H<sub>2</sub>O through this reaction:
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| : <math>\mathrm{H_2O_2 \ \xrightarrow {CAT}\ \textstyle\frac12 O_2 + H_2O}</math>
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| Another origin of hydrogen peroxide is the degradation of [[adenosine monophosphate]] which yields [[hypoxanthine]]. Hypoxanthine is then oxidatively [[catabolism|catabolized]] first to [[xanthine]] and then to [[uric acid]], and the reaction is catalyzed by the enzyme [[xanthine oxidase]]:<ref name=lehninger932>Nelson, David; Cox, Michael; Lehninger, Albert L. and Cox, Michael M. [http://books.google.com/books?id=wuLQCAOtC4MC&pg=PA932 Lehninger Biochemie], p. 932, Springer, 2001, ISBN 3-540-41813-X (in German)</ref>
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| [[File:Harnstoffsynthese.png|center|thumb|600px|Degradation of hypoxanthine to uric acid to form hydrogen peroxide. XO = xanthinoxidase]]
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| [[File:Pheropsophus verticalis 01 Pengo.jpg|left|thumb|130px|Australian bombardier beetle]]
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| [[File:Lampyris noctiluca.jpg|thumb|130px|Firefly ''Lampyris noctiluca'']]
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| The degradation of [[guanosine monophosphate]] yields xanthine as an intermediate product which is then converted in the same way to uric acid with the formation of hydrogen peroxide.<ref name=lehninger932/>
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| Eggs of [[sea urchin]], shortly after fertilization by a sperm, produce hydrogen peroxide. It is then quickly dissociated to OH· [[radical (chemistry)|radical]]s. The radicals serve as initiator of [[radical polymerization]], which surrounds the eggs with a protective layer of [[polymer]].<ref>{{cite journal|last1=Kröger|first1=M.|title=History|journal=Chemie in unserer Zeit|volume=23|pages=34|year=1989|doi=10.1002/ciuz.19890230106}}</ref>
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| 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.<ref>{{cite journal|last1=Schildknecht|first1=H.|last2=Holoubek|first2=K.|title=The bombardier beetle and its chemical explosion|journal=Angewandte Chemie|volume=73|pages=1|year=1961|doi=10.1002/ange.19610730102}}</ref><ref name=ncse>{{cite journal |author=Weber CG |title=The Bombadier Beetle Myth Exploded |journal=Creation/Evolution |publisher=[[National Center for Science Education]] |volume=2 |issue=1 |pages=1–5 |date=Winter 1981 |url=http://ncse.com/cej/2/1/bombardier-beetle-myth-exploded}}</ref><ref name=to>{{cite web |author=Isaak, Mark |date=May 30, 2003 |title=Bombardier Beetles and the Argument of Design |work=[[TalkOrigins Archive]] |url=http://www.talkorigins.org/faqs/bombardier.html}}</ref>
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| Hydrogen peroxide is a [[signaling molecule]] of [[Plant pathology|plant defense against pathogens]].<ref>[http://web.archive.org/web/20110723120528/http:/www.helmholtz-muenchen.de/biop/printversionen/pdf/aktuelles/pflanzenschuetzen_no.pdf Wie Pflanzen sich schützen], Helmholtz-Institute of Biochemical Plant Pathology (in German)</ref>
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| In [[firefly]], oxidation of [[Firefly luciferin|luciferins]], which is catalyzed by [[luciferase]]s, yields a peroxy compound [[1,2-dioxetane]]. The dioxetane is unstable and decays spontaneously to [[carbon dioxide]] and excited [[ketone]]s, which release excess energy by emitting light ([[bioluminescence]]).<ref>Aldo Roda [http://books.google.com/books?id=Gq_QDcADZxEC&pg=PA57 Chemiluminescence and Bioluminescence: Past, Present and Future], p. 57, Royal Society of Chemistry, 2010, ISBN 1-84755-812-7</ref>
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| [[File:Luciferin principle.png|thumb|450px|Loss of CO<sub>2</sub> of a dioxetane, giving rise to an excited ketone, which relaxes by emitting light.|center]]
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| {{clear}}
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| ==Hydrogen peroxide==
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| [[File:Riedl-Pfleiderer process.png|thumb|300px|Mechanism of anthraquinone process]]
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| The most widely used synthesis method of [[hydrogen peroxide]] is the [[anthraquinone]] process. There, [[anthraquinone]] is first hydrogenated to anthrahydroquinone. This reduced compound is oxidized with molecular oxygen, regenerating anthraquinone and releasing hydrogen peroxide.<ref name=howi531/> Direct synthesis of hydrogen peroxide from hydrogen and oxygen is inefficient and currently is not practiced industrially.<ref name=howi531/>
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| Many peroxides of mineral acids, such as [[peroxodisulfate]]s<ref>Jander, Blasius, Strähle ''Einführung in das anorganisch-chemische Praktikum''. 14ed., pp. 311–312, Hirzel Verlag, Stuttgart, 1995, ISBN 978-3-7776-0672-9</ref> and percarbonates,<ref>{{cite journal|last1=Riesenfeld|first1=E. H.|last2=Reinhold|first2=B.|title=The existence of real percarbonates and their distinction from carbonates with hydrogen peroxide-crystal|journal=Berichte der deutschen chemischen Gesellschaft|volume=42|pages=4377|year=1909|doi=10.1002/cber.19090420428|issue=4}}</ref> can be obtained by [[anode|anodic]] oxidation of the respective acids. The anode material must be stable to the required high potentials of a few volts and therefore is either platinum or its alloys.<ref>Hamann, Carl H.; Hamnett, A. and Vielstich, Wolf [http://books.google.com/books?id=3YBPeJMKJHIC&pg=PA320 Electrochemistry], Wiley-VCH, 1998, ISBN 3-527-29095-8, p. 320</ref><ref name=j1>{{cite journal|last1=Serrano|first1=K|doi=10.1016/S0013-4686(02)00688-6|title=Electrochemical preparation of peroxodisulfuric acid using boron doped diamond thin film electrodes|year=2002|pages=431|volume=48|journal=Electrochimica Acta |url=http://oatao.univ-toulouse.fr/3010/1/Serrano_3010.pdf|last2=Michaud|first2=P.A.|last3=Comninellis|first3=C.|last4=Savall|first4=A.|issue=4}}</ref>
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| : <math>\mathrm{2\ HSO_4^-\ \xrightarrow \ \ 2\ H^+\ +\ S_2O_8^{2-}\ +\ 2\ e^- \ \ \ \ E^0 = 2.123 V}</math>
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| : <math>\mathrm{2\ SO_4^{2-}\ \xrightarrow \ \ S_2O_8^{2-}\ +\ 2\ e^- \ \ \ \ E^0 = 2.01 V}</math>
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| [[Peroxydisulfuric acid]] was historically used for the production of hydrogen peroxide in a method developed in the early 20th century:<ref name=offer/><ref name=howi531/>
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| : <math>\mathrm{H_2S_2O_8\ +\ 2\ H_2O\longrightarrow\ H_2O_2\ +\ 2\ H_2SO_4}</math>
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| This process requires relatively high concentration of peroxydisulfuric acid as its more dilute solutions evolve oxygen gas instead of peroxide.<ref name=j1/>
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| ==Inorganic peroxides (aside from hydrogen peroxide)==
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| [[File:Sodium-peroxide-unit-cell-3D-balls.png|thumb|[[Unit cell]] of sodium peroxide Na<sub>2</sub>O<sub>2</sub>. The sodium ions are violet and the peroxide ions in red]]
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| Inorganic peroxides are divided into ionic and covalent peroxide. The first class mostly contains the peroxides of the [[alkali metal|alkali]] and [[alkaline earth metal]]s whereas the covalent peroxides are represented by such compounds as hydrogen peroxide and [[peroxymonosulfuric acid]] (H<sub>2</sub>SO<sub>5</sub>). In contrast to the purely ionic character of alkali metal peroxides, peroxides of [[transition metal]]s have a more covalent character.<ref name=volnov/>
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| ===Defect Formation in Fluorite Dioxides===
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| Both CeO<sub>2</sub> and ThO<sub>2</sub> accommodate excess oxygen in their lattice by formaing peroxide molecules.<ref>
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| {{cite journal|author=Middleburgh S. C., Lumpkin G.R. and Grimes R.W.|title= Accommodation of Excess Oxygen in Fluorite Dioxides |url=http://dx.doi.org/10.1016/j.ssi.2013.09.020|year=2013|journal=Solid State Ionics|volume=253|pages=119–122}}</ref> This is in contrast to UO<sub>2</sub>, which takes in excess oxygen by forming charged interstitials. The accommodation of excess oxygen in these systems changes their thermal and diffusion properties. This will have implications in the oxides' use in nuclear fuels and solid oxide fuel cells. Adding oxygen to the CeO<sub>2</sub> system, for example, will increase the diffusivity of the oxygen allowing for more efficient solid oxide fuel cells.
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| ===Peroxide salts===
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| ====Bonding in O<sub>2</sub><sup>2-</sup>====
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| [[File:MO-Peroxid.png|thumb|300px|Molecular orbital diagram of the peroxide ion]]
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| The peroxide ion is composed of two oxygen atoms that are linked by a single bond. The [[molecular orbital diagram]] of the peroxide dianion predicts a doubly occupied antibonding π* orbital and a [[bond order]] of one. The bond length is 149 [[picometer|pm]], which is larger than in the ground state ([[triplet oxygen]]) of the oxygen molecule (<sup>3</sup>O<sub>2</sub>, 121 pm). This translates into the smaller [[Hooke's law|force constant]] of the bond (2.8 [[Newton (unit)|N]]/cm vs. 11.4 N/cm for <sup>3</sup>O<sub>2</sub>) and the smaller [[frequency]] of the molecular vibration (770 cm<sup>−1</sup> vs. 1555 cm<sup>−1</sup> for <sup>3</sup>O<sub>2</sub>).<ref name=howi504>Wiberg, Egon; Wiberg, Nils and Holleman, Arnold Frederick [http://books.google.com/books?id=Mtth5g59dEIC&pg=PA476 ''Inorganic Chemistry''], Academic Press, 2001, ISBN 0-12-352651-5, pp. 475 ff</ref>
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| The peroxide ion can be compared with other molecular oxygen ions [[superoxide]] O<sub>2</sub><sup>−</sup> and [[ozonide]] O<sub>3</sub><sup>−</sup>, but contrary to them, the peroxide is not a radical and not [[Paramagnetism|paramagnetic]].<ref name=howi504/> Owing to the weak bonding between the oxygen atoms, peroxide easily undergoes [[homolysis (chemistry)|homolytic cleavage]] yielding two highly reactive radicals. This cleavage is accelerated by temperature, illumination or [[chemical reaction]]s.
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| : <math>\mathrm{R{-}O{-}O{-}R\longrightarrow\ 2\ R{-}O\cdot}</math>
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| ==== Preparation of peroxide salts ====
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| Most alkali metal peroxides can be synthesized directly by oxygenation of the elements. [[Lithium peroxide]] is form upon treating [[lithium hydroxide]] with hydrogen peroxide:<ref name=volnov>Vol'nov, I. I. ''Peroxides, superoxides and ozonides of alkali and alkaline earth metals'', pp. 21–51, Plenum Press, New York, 1966, no ISBN</ref>
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| : 2 LiOH + H<sub>2</sub>O<sub>2</sub> → Li<sub>2</sub>O<sub>2</sub> + 2 H<sub>2</sub>O
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| Historically, barium peroxide is prepared by oxygenation of [[barium oxide]] at elevated temperature and pressure.<ref name=howi531>Wiberg, Egon; Wiberg, Nils and Holleman, Arnold Frederick [http://books.google.com/books?id=Mtth5g59dEIC&pg=PA471 ''Inorganic Chemistry''], Academic Press, 2001, ISBN 0-12-352651-5, pp. 471–502</ref>
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| : 2 BaO + O<sub>2</sub> → 2 BaO<sub>2</sub>
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| Barium peroxide was once used to produce pure oxygen from air. This process relies on the temperature-dependent chemical balance between barium oxide and peroxide: the reaction of barium oxide with air at 500 °C results in barium peroxide, which upon heating to above 700 °C in oxygen decomposes back to barium oxide releasing pure oxygen.<ref name=howi531/>
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| :<math>\mathrm{2\ BaO\ +\ air\ \xrightarrow {500 ^\circ C}\ \ 2\ BaO_2\ \xrightarrow {700 ^\circ C}\ 2\ BaO\ +\ O_2\ (pure) }</math>
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| ====Properties of peroxide salts====
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| Few reactions are generally formulated for peroxide salt. In excess of dilute acids or water they release hydrogen peroxide.<ref name=volnov/>
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| : <math>\mathrm{Na_2O_2\ +\ 2\ H_2O\ \xrightarrow {H_3O^+}\ \ 2\ NaOH\ +\ H_2O_2}</math>
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| Upon heating, the reaction with water leads to the release of oxygen instead<ref name=volnov/>
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| : <math>\mathrm{2\ Na_2O_2\ + 2\ H_2O\longrightarrow\ 4\ NaOH\ +\ O_2}</math>
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| : <math>\mathrm{2\ Na_2O_2\ \xrightarrow {\Delta T}\ \ 2\ Na_2O\ +\ O_2}</math>
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| The peroxide anion is a stronger nucleophile than hydroxide and displaces hydroxyl from oxyanions e.g. forming perborates and percarbonates. [[Sodium perborate]] and [[sodium percarbonate]] are important consumer and industrial bleaching agents; they stabilize hydrogen peroxide and limit side reactions (e.g. reduction and decomposition note below). The peroxide anion forms an adduct with [[urea]], [[hydrogen peroxide - urea]].
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| 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.
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| ===As a ligand in coordination chemistry===
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| Peroxide functions as a [[bidentate ligand]] in a variety of [[coordination complex]]. Some complexes have only peroxide ligands, e.g., [[chromium(VI) oxide peroxide]] (Cr(O<sub>2</sub>)<sub>4</sub><sup>2-</sup>). Similarly, [[molybdate]] reacts in alkaline media with peroxide to form red peroxomolybdate {Mo(O<sub>2</sub>)<sub>4</sub>}<sup>2–</sup>.<ref>{{cite book|author=Eagleson, Mary |title=Concise encyclopedia chemistry|url=http://books.google.com/books?id=Owuv-c9L_IMC&pg=PA660|year=1994|publisher=Walter de Gruyter|isbn=978-3-11-011451-5|pages=660–}}</ref> The reaction of hydrogen peroxide with aqueous titanium(IV) gives a brightly colored peroxy complex that is a useful test for [[titanium]] as well as hydrogen peroxide. Many [[transition metal dioxygen complex]]es are best described as adducts of peroxide.<ref>Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4</ref>
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| ===Applications===
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| Many inorganic peroxides are used for [[bleach]]ing [[textiles]] and [[paper]] and as a bleaching additive to detergents and cleaning products.<ref name=howi531/> The increasing environmental concerns resulted in the preference of peroxides over chlorine-based compounds and a sharp increase in the peroxide production.<ref name=offer/><ref>Ullmann's Encyclopedia of Industrial Chemistry, Vol A 19, 5 ed., pp. 177–197, VCH, Weinheim, 1991, ISBN 3-527-20138-6</ref> The past use of [[Sodium perborate|perborate]]s as additives to detergents and cleaning products<ref name = "Brotherton">Brotherton, B.J. "Boron: Inorganic Chemistry", in ''Encyclopedia of Inorganic Chemistry'' (1994) Ed. R. Bruce King, John Wiley & Sons ISBN 0-471-93620-0</ref> has been largely replaced by [[Sodium percarbonate|percarbonates]] in order to decrease the emission of boron to the environment. Sodium percarbonate is used in such products as [[OxiClean]] and [[Tide (detergent)|Tide laundry detergent]]. When dissolved in water, it releases hydrogen peroxide and [[soda ash]] (sodium carbonate):<ref name = "Jones">{{cite book|author = Jones, Craig W. |title = Applications of hydrogen peroxide and its derivatives|year = 1999|publisher = [[Royal Society of Chemistry]]|isbn = 0-85404-536-8}}</ref>
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| :2 Na<sub>2</sub>CO<sub>3</sub>·1.5H<sub>2</sub>O<sub>2</sub> → 2 Na<sub>2</sub>CO<sub>3</sub> + 3 H<sub>2</sub>O<sub>2</sub>
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| The use of peroxide compounds in detergents is often reflected in their trade names; for example, [[Persil]] is a combination of the words ''per''borate and ''sil''icate.
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| Some peroxide salts release oxygen upon reaction with carbon dioxide. This reaction is used in generation of oxygen from exhaled carbon dioxide on [[submarine]]s and spaceships. Sodium or lithium peroxides are preferred in space applications because of their lower [[molar mass]] and therefore higher oxygen yield per unit weight.<ref name=howi531/>
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| : <math>\mathrm{2\ Na_2O_2\ +\ 2\ CO_2\longrightarrow\ 2\ Na_2CO_3\ +\ O_2}</math>
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| Alkali metal peroxides can be used for the synthesis of organic peroxides. One example is the conversion of [[benzoyl chloride]] with sodium peroxide to di[[benzoyl peroxide]].<ref>{{cite journal|last1=Gambarjan|first1=Stephan|title=Diphenylamine and Acylperoxyde|journal=Berichte der deutschen chemischen Gesellschaft|volume=42|pages=4003|year=1909|doi=10.1002/cber.190904203164|issue=3}}</ref>
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| [[File:Peroxodibenzoyl.png|center|450px|Synthesis of dibenzoyl]]
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| ==Organic peroxides==
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| [[Organic peroxide]]s can be divided into two major classes, [[peroxy acid]]s and [[organic hydroperoxide]]s. The first class is derived from the [[carboxylic acid]] and the second from [[ether]]s or [[alcohol]]s.
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| ===Preparation===
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| Most peroxy acids can be obtained by the reaction of hydrogen peroxide with the corresponding [[carboxylic acid]]:
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| : <math>\mathrm{R{-}COOH\ +\ H_2O_2\longrightarrow\ R{-}COOOH\ +\ H_2O}</math>
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| : <small> R = organic group </small>
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| Another route employs [[acyl halide]]s instead of the carboxylic acid. This method is used primarily with aromatic compounds in [[Base (chemistry)|basic]] in order to neutralize the resulting [[hydrogen chloride]].
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| : <math>\mathrm{R{-}COCl\ +\ H_2O_2\longrightarrow\ R{-}COOOH\ +\ HCl}</math>
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| Aromatic [[aldehydes]] can be auto-oxidized into peroxycarboxylic acids:
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| : <math>\mathrm{Ar{-}CHO\ +\ O_2\longrightarrow\ Ar{-}COOOH}</math>
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| : <small> Ar = [[aryl]] </small>
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| The products, however, react with the initial aldehyde forming the carboxylic acid:
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| : <math>\mathrm{Ar{-}COOOH\ +\ Ar{-}CHO\longrightarrow\ 2\ Ar{-}COOH}</math>
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| Several synthesis routes are known for [[Aliphatic compound|aliphatic]] peroxides, such as the reaction of dialkylsulfates with alkaline hydrogen peroxide solution.<ref>{{cite journal|last1=Medwedew|first1=S. S.|last2=Alexejewa|first2=E. N.|title=Organic peroxides II. Of the reaction between benzoyl hydroperoxide or benzoyl peroxide and triphenylmethyl|journal=Berichte der deutschen chemischen Gesellschaft (A and B Series)|volume=65|pages=137|year=1932|doi=10.1002/cber.19320650204|issue=2}}</ref><ref>Wiley, Richard Haven "Preparation of diaikyl peroxides" {{US patent|2357298}} Issue date: 1942</ref> In this method, the alkyl sulfate donates the alkyl group and the sulfate ion forms the [[leaving group]].
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| : <math>\mathrm{R_2SO_4\ +\ H_2O_2\longrightarrow\ R{-}O{-}O{-}R\ +\ H_2SO_4}</math>
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| This method can also yield cyclic peroxides.<ref>{{cite journal|last1=Criegee|first1=Rudolf|last2=Müller|first2=Gerhard|title=1.2-Dioxan|journal=Chemische Berichte|volume=89|pages=238|year=1956|doi=10.1002/cber.19560890209|issue=2}}</ref> The four-membered [[dioxetane]]s can be obtained by 2+2 [[cycloaddition]] of oxygen to [[alkene]]s.<ref>Heinz G. O. Becker ''Organikum'', Wiley-VCH, 2001, ISBN 3-527-29985-8, p. 323</ref>
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| [[File:Hock-Phenol.png|thumb|250px|Synthesis of cumene hydroperoxide]]
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| he selective synthesis of hydroperoxides can be carried out by free-radical oxidation of alkanes with oxygen. Here the active site formed by a [[radical initiator]] reacts with oxygen to form a hydroperoxyl. The addition of oxygen results in a more active radical which can further extract hydrogen atoms and release the hydroperoxide, leaving a new radical.<ref name=org206>Heinz G. O. Becker ''Organikum'', Wiley-VCH, 2001, ISBN 3-527-29985-8 pp. 206–207</ref> This process is used industrially for the synthesis of [[phenol]] from [[benzene]] and is called the [[Cumene process]] or Hock process for its [[cumene]] and [[cumene hydroperoxide]] intermediates.<ref>Brückner, R. ''Reaktionsmechanismen: organische Reaktionen, Stereochemie, moderne Synthesemethoden'', pp. 41–42, Spektrum Akademischer Verlag, Munich, 2004, ISBN 3-8274-1579-9 (in German)</ref>
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| This auto-oxidation reaction can be used with common [[solvent]]s from the group of [[ether]]s, such as [[diethyl ether]], [[diisopropyl ether]], [[tetrahydrofuran]] or [[1,4-dioxane]]. It yields a volatile hydroperoxide ether that upon heating can result in a serious explosion.<ref name=org206/>
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| [[File:Schenk-En-Reaktion.png|thumb|250px|Synthesis of hydroperoxides of alkene and singlet oxygen in an [[ene reaction]]]]
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| Peroxides are formed by living organisms through [[ene reaction]]s or [[Diels–Alder reaction]]s between [[alkene]]s and oxygen. Unsaturated fatty acids can serve as the olefinic [[Substrate (chemistry)|substrates]] for the ene reaction and unsaturated [[amino acids]] like [[histidine]] can be the reactant for the Diels-Alder cyclization.<ref>Autor, Anne Pomeroy ''Pathology of Oxygen'', pp. 25–26, Academic Press, New York, 1982, ISBN 0-12-068620-1</ref> [[Rancidification]] (decomposition) of fats is partly caused by the formation of peroxides.
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| ===Properties and applications of organic peroxides===
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| Peroxycarboxylic acids are generally weaker [[acid]]s than the parent carboxylic acids. Like most peroxides, they are strong oxidants and tend to explode at high concentrations and higher temperatures.<ref name=org206/>
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| Organic peracids are used in the synthesis of [[epoxy|epoxies]] via the [[Epoxide|Prilezhaev reaction]].<ref>Beyer, Hans; Walter, Wolfgang and Francke, Wittko ''Lehrbuch der organischen Chemie'', 23 Ed., pp. 76–77, Hirzel Verlag, Stuttgart-Leipzig, 1998, ISBN 3-7776-0808-4</ref> Another important application is the synthesis of [[lactone]]s of cyclic [[ketone]]s in the [[Baeyer–Villiger oxidation]] process.<ref>Vollhardt, Kurt Peter C. and Schore, Neil Eric ''Organische Chemie'', 3rd ed., pp. 818–819, Wiley-VCH, 2000, ISBN 3-527-29819-3</ref> In both cases, electron-poor peroxycarboxylic acids are especially efficient, such as [[meta-chloroperoxybenzoic acid|''meta''-chloroperoxybenzoic acid]] (mCPBA).
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| {{double image|center|Prileschaew-Reaktion.png|400|Baeyer-Villiger oxidation mechanism.png|400|Mechanism of the Prilezhaev reaction|Mechanism of the Baeyer–Villiger oxidation}}
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| [[tert-butyl hydroperoxide|''Tert''-butyl hydroperoxide]] is a common [[oxidant]] in the [[Sharpless epoxidation]], which is used for the [[Enantiomer|stereoselective]] synthesis of epoxides. [[Karl Barry Sharpless]] was awarded the 2001 Nobel prize in Chemistry for this reaction.<ref>Bülle, Jan and Hüttermann, Aloys ''Das Basiswissen der organischen Chemie'', pp. 308–309, Wiley-VCH Weinheim, 2000, ISBN 3-527-30847-4</ref>
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| Peracetic acid is a popular [[disinfectant]] in the medical field and food industry.<ref>Block, S. S. [http://books.google.com/books?id=3f-kPJ17_TYC&pg=PA979 Disinfection, Sterilization and Preservation], Lippincott, Williams & Wilkins, Philadelphia, 2000, ISBN 0-683-30740-1, p. 979</ref> Various peroxide solutions are commercially produced for the cleaning and disinfection of [[contact lens]]es.
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| [[File:Caprolactam-2D-skeletal.png|thumb|100px|Caprolactam]]
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| [[Dibenzoyl peroxide]] is used as a [[radical initiator]]. Its weak peroxide bond cleaves easily yielding reactive benzoxy radicals, which assist [[polymerization]] leading to plastics like [[polyethylene]].<ref>Frankenburg, W. G. [http://books.google.com/books?id=2fLKz-PXvzwC&pg=PA24 Advances in Catalysis and Related Subjects], Vol. 2, pp. 24–26, 1948, Academic Press, ISBN 0-12-007801-5</ref><ref>Wicks, Z. W.; Jones, F. N.; Pappas, S. P. and Wicks, D. A. [http://books.google.com/books?id=KEU0yixqUTYC&pg=PA17 Organic Coatings], 3rd edition, pp. 17–26, 2007, Wiley, New York, ISBN 0-471-69806-7</ref> One of the synthesis methods of the commercially important plastic [[caprolactam]]—the precursor to [[Nylon 6]] (polycaprolactam)—is a Baeyer-Villiger rearrangement of [[cyclohexanone]] with [[peracetic acid]]. This yields [[caprolactone]], which is then converted to caprolactam by reacting it with [[ammonia]].<ref>Arpe, H.–J. [http://books.google.com/books?id=36kHHvzx6M8C&pg=PA284 ''Industrielle Organische Chemie: Bedeutende Vor- und Zwischenprodukte''], pp. 284–285, Wiley-VCH Weinheim, 2007, ISBN 3-527-31540-3</ref>
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| Industrial resins based on acrylic and/or [[methacrylic acid]] esters are invariably produced by radical polymerization with organic peroxides at elevated temperatures.<ref>Thomas Brock, Michael Groteklaes, Peter Mischke [http://books.google.com/books?id=xz2veR6uko0C&pg=PA67 Lehrbuch der Lacktechnologie], Vincentz Network GmbH & Co KG, 2000, ISBN 3-87870-569-7 p. 67</ref> The polymerization rate is adjusted by suitable choice of temperature and type of peroxide.<ref>[http://www.pergan.com/Downloads/Organische_Peroxide_fuer_die_Polymerisation.pdf Organische Peroxide für die Polymerisation]. pergan.com (in German)</ref>
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|
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| Some peroxides are [[drug]]s, whose action is based on the formation of radicals at desired locations in the organism. For example, [[artemisinin]] and its derivatives, such as such [[artesunate]], possess the most rapid action of all current drugs against [[Plasmodium falciparum|falciparum]] [[malaria]].<ref>{{cite journal |author=White NJ |title=Assessment of the pharmacodynamic properties of antimalarial drugs in vivo |journal=Antimicrob. Agents Chemother. |volume=41 |issue=7 |pages=1413–22 |year=1997|pmid=9210658 |pmc=163932}}</ref> Artesunate is also efficient in reducing egg production in ''[[Schistosoma haematobium]]'' infection.<ref>{{cite journal|author=Boulangier D, Dieng Y, Cisse B, ''et al.''|title=Antischistosomal efficacy of artesunate combination therapies administered as curative treatments for malaria attacks|year=2007|journal=Trans R Soc Trop Med Hyg|volume=101|issue=2|pages=113–16|doi=10.1016/j.trstmh.2006.03.003|pmid=16765398}}</ref>
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| [[File:Acetone Peroxide Synthesis V.1.svg|thumb|300px|Synthesis of [[Acetone peroxide|TATP]] from [[acetone]] and H<sub>2</sub>O<sub>2</sub>]]
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| Many organic peroxides can initiate explosive polymerization in materials with unsaturated chemical bonds, and specifically [[Acetone peroxide|triacetone triperoxide]] (TATP) and [[hexamethylene triperoxide diamine]] (HMTD) are powerful explosives. TATP is an inexpensive compound and is relatively easy to make. Whereas most other potent explosives, such as [[trinitrotoluene]] (TNT) or [[RDX]] (the major component of [[C-4 (explosive)|C4]] mixtures), contain nitrogen, which is relatively easy to trace by sniffing techniques, TATP is nitrogen free and therefore is very difficult to detect by conventional screening methods. For this reason, it is an explosive favored by terrorists.<ref name=m/> TATP and HMTD were used in several executed or planned terrorist acts of the early 2000s, most notably in the [[2001 shoe bomb plot]]<ref name="CNN 12-28-01">{{cite news|url=http://archives.cnn.com/2001/US/12/28/inv.reid/|title=Judge denies bail to accused shoe bomber|date=December 28, 2001|publisher=CNN}}</ref><ref name=det>{{cite web|url=http://www.opensourcesinfo.org/journal/2005/7/25/terrorist-use-of-tatp-explosive.html |title=Terrorist Use Of TATP Explosive |publisher=Opensourcesinfo.org |date=2005-07-25 |accessdate=2010-01-18}}</ref> and the [[7 July 2005 London bombings|2005 London Underground bombings]].<ref>[http://www.guardian.co.uk/uk/2006/may/07/theobserver.uknews "The real story of 7/7"], ''[[The Observer]]'', May 7, 2006</ref><ref>[http://www.redorbit.com/news/general/197067/london_bombers_used_everyday_materialsus_police/index.html [[London]] bombers used everyday materials—U.S. police], Reuters, 4 August 2005</ref> Several detection devices have been designed since those events. One, for example, releases a chemical mixture which changes color when interacting with traces of TATP.<ref name=det/>
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| ==Laboratory identification==
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| [[File:Jodprobe.jpg|thumb|[[Iodine test|Iodine-starch test]]. Note the blackening (left) of initially yellowish (right) starch.]]
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| Several analytical methods are used for qualitative and quantitative determination of peroxides.<ref>{{cite journal|last1=Légrádi|first1=L.|last2=Légrádi|first2=J.|title=Detection of peroxides, hydroperoxides and peracids|journal=Mikrochimica Acta|volume=58|pages=119|year=1970|doi=10.1007/BF01218105}}</ref> A simple qualitative detection of peroxides is carried out with the [[iodine test|iodine-starch reaction]].<ref>{{cite journal|last1=Lea|first1=C. H.|title=The Effect of Light on the Oxidation of Fats|journal=Proceedings of the Royal Society B: Biological Sciences|volume=108|pages=175|year=1931|doi=10.1098/rspb.1931.0030|issue=756}}</ref> Here peroxides, hydroperoxides or peracids oxidize the added [[potassium iodide]] into [[iodine]], which reacts with [[starch]] producing a deep-blue color. Commercial paper indicators using this reaction are available. This method is also suitable for quantitative evaluation, but it can not distinguish between different types of peroxide compounds. Discoloration of various [[indigo dye]]s in presence of peroxides is used instead for this purpose.<ref>Veibel, S. ''Analytik organischer Verbindungen'', Akademie-Verlag, Berlin, 1960, p. 262</ref> For example, the loss of blue color in leuco-[[methylene blue]] is selective for hydrogen peroxide.<ref>{{cite journal|last1=Eiss|first1=M. I.|last2=Giesecke|first2=Paul|title=Colorimetric Determination of Organic Peroxides|journal=Analytical Chemistry|volume=31|pages=1558|year=1959|doi=10.1021/ac60153a038|issue=9}}</ref>
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| Quantitative analysis of hydroperoxides is performed using potentiometric [[titration]] with [[lithium aluminium hydride]].<ref>{{cite journal|last1=Higuchi|first1=T.|last2=Zuck|first2=Donald Anton|title=Behaviors of Several Compounds as Indicators in Lithium Aluminum Hydride Titration of Functional Groups|journal=Journal of the American Chemical Society|volume=73|pages=2676|year=1951|doi=10.1021/ja01150a073|issue=6}}</ref> Another way to evaluate the content of peracids and peroxides is the volumetric titration with [[alkoxide]]s such as [[sodium ethoxide]].<ref>{{cite journal|last1=Martin|first1=A. J.|title=Potentiometric titration of hydroperoxide and peracid in Anhydrous Ethylenediamine|journal=Analytical Chemistry|volume=29|pages=79|year=1957|doi=10.1021/ac60121a022}}</ref>
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| == Safety ==
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| [[File:Hydrogen peroxide 35 percent on skin.jpg|thumb|upright|Skin shortly after exposure to 35% H<sub>2</sub>O<sub>2</sub>]]
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| Organic peroxides can accidentally initiate explosive polymerization in materials with unsaturated chemical bonds. Most notably, [[Acetone peroxide|TATP]] and [[Hexamethylene triperoxide diamine|HMTD]] are [[Explosive material|high explosives]], and TATP, because of its high susceptibility to accidental detonation by shock, friction, or sparks, has earned the nickname "Mother of Satan" among certain [[Islamic terrorism|Islamic militant groups]].<ref name=m>{{cite web
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| |author= Genuth, Iddo and Fresco-Cohen, Lucille
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| |title = TATP: Countering the Mother of Satan
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| |publisher = The Future of Things
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| |date = 6 November 2006
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| |url = http://thefutureofthings.com/articles/35/tatp-countering-the-mother-of-satan.html
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| |quote = The tremendous devastative force of TATP, together with the relative ease of making it, as well as the difficulty in detecting it, made TATP one of the weapons of choice for terrorists
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| |accessdate = 24 September 2009}}</ref> TATP can accidentally form as by-products in many commonly used reactions. These reactions range from synthesis of [[MDMA]], where TATP is formed via [[safrole|isosafrole]] oxidation in acetone, to industrial production of [[phenol]], where the second product of the [[cumene process]], [[acetone]], is partially oxidized to peroxide on the second reaction step. Accidental preparation of organic peroxides can occur by mixing ketone solvents (most commonly acetone) with waste materials containing hydrogen peroxide or other oxidizers and leaving the mixture standing for several hours. In addition, many liquid [[ether]]s in the presence of [[air]], light and metals (which act as catalysts) slowly – over a period of months – form highly unstable ether peroxides such as [[diethyl ether peroxide]]. Therefore, ethers are often stored over potassium hydroxide, which not only destroys peroxides but also acts as a powerful [[desiccation|desiccant]].
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| Peroxides are also strong oxidizers and easily react with skin, cotton and wood pulp.<ref>Heinz G. O. Becker ''Organikum'', Wiley-VCH, 2001, ISBN 3-527-29985-8 pp. 741–762</ref> For safety reasons, peroxidic compounds are stored in a cool, opaque container, as heating and illumination accelerate their [[chemical reaction]]s. Small amounts of peroxides, which emerge from storage or reaction vessels are neutralized using reducing agents such as [[iron(II) sulfate]]. The safety measures in industrial plants producing large amounts of peroxides include the following. The equipment is located within reinforced concrete structures with foil windows, which would relieve pressure and not shatter in case of explosion. The products are bottled in small containers and are moved to a cold place promptly after the synthesis. The containers are made of non-reactive materials such as stainless steel, some aluminium alloys or dark glass.<ref>[http://www.ozoneservices.com/articles/004.htm Ozonelab Peroxide compatibility]</ref>
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| ==History==
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| One of the first synthetic peroxides, [[barium peroxide]], was synthesized by [[Alexander von Humboldt]] in 1799 as a by-product of his attempts to decompose air. 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 ''oxidized water'' – now known as hydrogen peroxide.<ref>{{cite journal|title=Der tropfbar flüssige Sauerstoff, oder das oxygenierte Wasser|language=German|author=Gilbert, L. W. |url=http://books.google.com/books?id=xwYAAAAAMAAJ&pg=PA3 |journal=Annals of Physics |year=1820|page=3|volume= 65–66}}</ref> [[Sodium peroxide]] was synthesized in 1811 by Thénard and [[Joseph Louis Gay-Lussac]]. The bleaching effect of peroxides and their salts on [[natural dye]]s became known around that time, but early attempts of industrial production of peroxides failed, and the first plant producing hydrogen peroxide was built only in 1873 in [[Berlin]]. The discovery of the synthesis of hydrogen peroxide by [[electrolysis]] with [[sulfuric acid]] had brought the more efficient electrochemical method. It was first implemented into industry in 1908 in [[Weißenstein]], [[Carinthia (state)|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, and by 1998, it reached 2.7 million tonnes.<ref name=offer>[[Heribert Offermanns|Offermanns, Heribert]]; Dittrich, Gunther; Steiner, Norbert (2000). "Wasserstoffperoxid in Umweltschutz und Synthese". Chemie in unserer Zeit 34 (3): pp.150 {{doi|10.1002/1521-3781(200006)34:3<150::AID-CIUZ150>3.0.CO;2-A|issue=3}}</ref>
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| ==See also==
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| {{colbegin}}
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| *[[Peroxidases]]
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| *[[Peroxide fusion]]
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| *[[Ozone]]
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| *[[Ozonide]], O<sub>3</sub><sup>–</sup>
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| *[[Superoxide]], O<sub>2</sub><sup>–</sup>
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| *[[Dioxygenyl]], O<sub>2</sub><sup>+</sup>
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| {{colend}}
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| ==References==
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| {{reflist|35em}}
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| {{Functional Groups}}
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| [[Category:Anions]]
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| [[Category:Peroxides| ]]
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| [[Category:Organic compounds]]
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| [[Category:Functional groups]]
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| {{Link GA|de}}
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