Redox 66313 225831731 2008-07-15T16:33:36Z DMacks 712163 Reverted edits by [[Special:Contributions/120.28.139.26|120.28.139.26]] ([[User talk:120.28.139.26|talk]]) to last version by BotKung [[Image:redox reaction.png|thumb|300px|right|Illustration of a redox reaction]] '''Redox''' (shorthand for '''reduction/oxidation''' reaction) describes all [[chemical reactions]] in which atoms have their [[oxidation number]] ([[oxidation state]]) changed. This can be either a simple redox process such as the oxidation of [[carbon]] to yield [[carbon dioxide]], or the reduction of [[carbon]] by [[hydrogen]] to yield [[methane]] (CH<sub>4</sub>), or it can be a complex process such as the oxidation of [[sugar]] in the human body through a series of very complex [[electron transfer]] processes. The term ''redox'' comes from the two concepts of '''red'''uction and '''ox'''idation. It can be explained in simple terms: * '''''Oxidation''''' describes the '''''loss''''' of [[electron]]s by a [[molecule]], [[atom]] or [[ion]] * '''''Reduction''''' describes the '''''gain''''' of [[electron]]s by a [[molecule]], [[atom]] or [[ion]] <!-- do not add mnemonic devices or other memory tricks without discussing on talk page --> Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to ''a change in oxidation number''—the actual transfer of electrons may never occur. Thus, oxidation is better defined as an ''increase in oxidation number'', and reduction as a ''decrease in oxidation number''. In practice, the transfer of electrons will always cause a change in oxidation number, but there are many reactions which are classed as "redox" even though no electron transfer occurs (such as those involving [[covalent]] bonds). Non-redox reactions, which do not involve changes in [[formal charge]], are known as [[Metathesis reaction (chemistry)|metathesis]] reactions. [[Image:Redox Halves.png|thumb|400px|The two parts of a redox reaction]] [[Image:Rust03102006.JPG|thumb|250px|Rusting iron]] [[Image:Large bonfire.jpg|thumb|250px|A bonfire. Combustion consists of redox reactions involving [[Radical (chemistry)|free radicals]].]] ==Oxidizing and reducing agents== Substances that have the ability to '''oxidize''' other substances are said to be '''oxidative''' and are known as [[oxidizing agent]]s, '''oxidants''' or '''oxidizers'''. Put another way, the oxidant removes electrons from another substance, and is thus reduced itself. And because it "accepts" electrons it is also called an [[electron acceptor]]. Oxidants are usually chemical substances with elements in high oxidation numbers (e.g., [[hydrogen peroxide|H<sub>2</sub>O<sub>2</sub>]], [[permanganate|MnO<sub>4</sub><sup>&minus;</sup>]], [[chromium trioxide|CrO<sub>3</sub>]], Cr<sub>2</sub>O<sub>7</sub><sup>2&minus;</sup>, [[Osmium(VIII) oxide|OsO<sub>4</sub>]]) or highly [[electronegativity|electronegative]] substances that can gain one or two extra electrons by oxidizing a substance ([[Oxygen|O]], [[Fluorine|F]], [[Chlorine|Cl]], [[Bromine|Br]]). Substances that have the ability to '''reduce''' other substances are said to be '''reductive''' and are known as [[reducing agent]]s, '''reductants''', or '''reducers'''. Put in another way, the reductant transfers electrons to another substance, and is thus oxidized itself. And because it "donates" electrons it is also called an [[electron donor]]. Reductants in chemistry are very diverse. ''Metal reduction''—electropositive elemental [[metal]]s can be used (Li, Na, Mg, Fe, Zn, Al). These metals donate or ''give away'' electrons readily. Other kinds of reductants are ''hydride transfer reagents'' (NaBH<sub>4</sub>, LiAlH<sub>4</sub>), these reagents are widely used in [[organic chemistry]]<ref>{{cite book|last=Hudlický|first=Miloš|title=Reductions in Organic Chemistry|publisher=[[American Chemical Society]]|date=1996|location=Washington, D.C.|pages=429|id=ISBN 0-8412-3344-6}}</ref><ref>{{cite book|last=Hudlický|first=Miloš|title=Oxidations in Organic Chemistry|publisher=[[American Chemical Society]]|date=1990|location=Washington, D.C.|pages=456|id=ISBN 0-8412-1780-7}}</ref>, primarily in the reduction of [[carbonyl]] compounds to [[alcohols]]. Another useful method is reductions involving hydrogen gas (H<sub>2</sub>) with a [[palladium]], [[platinum]], or [[nickel]] [[catalyst]]. These ''catalytic reductions'' are primarily used in the reduction of carbon-carbon double or triple bonds. The chemical way to look at redox processes is that the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or ''reducing agent'' loses electrons and is oxidized and the oxidant or ''oxidizing agent'' gains electrons and is reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a '''redox pair'''. ==Examples of redox reactions== A good example is the reaction between [[hydrogen]] and [[fluorine]]: :<math> \mathrm{H}_{2} + \mathrm{F}_{2} \longrightarrow 2\mathrm {HF}</math> We can write this overall reaction as two half-reactions: the oxidation reaction :<math> \mathrm{H}_{2} \longrightarrow 2\mathrm{H}^{+} + 2e^-</math> and the reduction reaction: :<math> \mathrm{F}_{2} + 2e^- \longrightarrow 2\mathrm{F}^{-}</math> Analysing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above). Elements, even in molecular form, always have an oxidation number of zero. In the first half reaction, hydrogen is oxidized from an oxidation number of zero to an oxidation number of +1. In the second half reaction, fluorine is reduced from an oxidation number of zero to an oxidation number of &minus;1. When adding the reactions together the electrons cancel: :<math>\frac{\begin{array}{rcl} \mathrm{H}_{2} & \longrightarrow & 2\mathrm{H}^{+} + 2e^{-}\\ \mathrm{F}_{2} + 2e^{-} & \longrightarrow & 2\mathrm{F}^{-} \end{array}}{\begin{array}{rcl} \mathrm{H}_{2} + \mathrm{F}_{2} & \longrightarrow & 2\mathrm{H}^{+} + 2\mathrm{F}^{-} \end{array}}</math> And the ions combine to form [[hydrofluoric acid|hydrogen fluoride]]: :<math>\mathrm{H}_{2} + \mathrm{F}_{2}\, \ \longrightarrow \ 2\mathrm{H}^{+} + 2\mathrm{F}^{-}\ \longrightarrow \ 2\mathrm{HF}</math> ===Displacement reactions=== Redox occurs in [[single displacement reactions]] or [[substitution reactions]]. The redox component of this type of reaction is the change of oxidation state (charge) on certain atoms, not the actual exchange of atoms in the compounds. For example, in the reaction between iron and copper(II) sulphate solution: :<math> \mathrm{Fe} + \mathrm{CuSO}_{4} \longrightarrow \mathrm{FeSO}_{4} + \mathrm{Cu}</math> The ionic equation for this reaction is: :<math>\mathrm{Fe} + \mathrm{Cu}^{2+} \longrightarrow \mathrm{Fe}^{2+} + \mathrm{Cu}</math> As two half-equations, it is seen that the iron is oxidized: :<math>\mathrm{Fe} \longrightarrow \mathrm{Fe}^{2+} + 2{e}^{-}</math> And the copper is reduced: :<math>\mathrm{Cu}^{2+} + 2{e}^{-} \longrightarrow \mathrm{Cu}</math> ===Other examples=== * iron(II) oxidizes to iron(III): :Fe<sup>2+</sup> → Fe<sup>3+</sup> + e<sup>&minus;</sup> * [[hydrogen peroxide]] reduces to [[hydroxide]] in the presence of an acid: :H<sub>2</sub>O<sub>2</sub> + 2 e<sup>&minus;</sup> → 2 OH<sup>&minus;</sup> overall equation for the above: :2Fe<sup>2+</sup> + H<sub>2</sub>O<sub>2</sub> + 2H<sup>+</sup> → 2Fe<sup>3+</sup> + 2H<sub>2</sub>O * [[denitrification]], [[nitrate]] reduces to [[nitrogen]] in the presence of an acid: :2NO<sub>3</sub><sup>&minus;</sup> + 10e<sup>&minus;</sup> + 12 H<sup>+</sup> → N<sub>2</sub> + 6H<sub>2</sub>O * iron oxidizes to iron(III) oxide and oxygen is reduced forming iron(III) oxide (commonly known as [[rust]]ing, which is similar to [[tarnish]]ing): :4Fe + 3O<sub>2</sub> → 2 Fe<sub>2</sub>O<sub>3</sub> * [[Combustion]] of [[hydrocarbon]]s, e.g. in an [[internal combustion engine]], produces [[water]], [[carbon dioxide]], some partially oxidized forms such as [[carbon monoxide]] and heat [[energy]]. Complete oxidation of materials containing [[carbon]] produces carbon dioxide. * In [[organic chemistry]], stepwise oxidation of a hydrocarbon produces water and, successively, an [[alcohol]], an [[aldehyde]] or a [[ketone]], [[carboxylic acid]], and then a [[peroxide]]. ==Redox reactions in industry== The primary process of reducing ore to produce metals is discussed in the article on [[Smelting]]. Oxidation is used in a wide variety of industries such as in the production of cleaning products. Redox reactions are the foundation of [[electrochemical cell]]s. ==Redox reactions in biology== <div class="thumb tright" style="background-color: #f9f9f9; border: 1px solid #CCCCCC; margin:0.5em;"> {|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;" |[[Image:Ascorbic acid structure.png|150px|ascorbic acid]] |} {|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;" |[[Image:Dehydroascorbic acid.png|150px|dehydroascorbic acid]] |} <div style="border: none; width:150px;"><div class="thumbcaption"><small>Top: [[ascorbic acid]] ([[reducing agent|reduced form]] of [[Vitamin C]])<br />Bottom: [[dehydroascorbic acid]] ([[oxidizing agent|oxidized form]] of [[Vitamin C]])</small></div></div></div> Many important [[biology|biological]] processes involve redox reactions. [[Cellular respiration]], for instance, is the oxidation of [[glucose]] (C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>) to [[carbon dioxide|CO<sub>2</sub>]] and the reduction of [[oxygen]] to [[water]]. The summary equation for cell respiration is: :C<sub>6</sub>H<sub>12</sub>O<sub>6</sub> + 6 O<sub>2</sub> → 6 CO<sub>2</sub> + 6 H<sub>2</sub>O :The process of cell respiration also depends heavily on the reduction of [[NAD+|NAD<sup>+</sup>]] to NADH and the reverse reaction (the oxidation of NADH to NAD<sup>+</sup>). [[Photosynthesis]] is essentially the reverse of the redox reaction in cell respiration: : 6 CO<sub>2</sub> + 6 H<sub>2</sub>O + [[photon|light energy]] → C<sub>6</sub>H<sub>12</sub>O<sub>6</sub> + 6 O<sub>2</sub> Biological energy is frequently stored and released by means of redox reactions. [[Photosynthesis]] involves the reduction of [[carbon dioxide]] into [[sugar]]s and the oxidation of [[water (molecule)|water]] into molecular [[oxygen]]. The reverse reaction, [[Cellular respiration|respiration]], oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce [[nicotinamide adenine dinucleotide]] (NAD<sup>+</sup>), which then contributes to the creation of a [[proton gradient]], which drives the synthesis of [[adenosine triphosphate]] (ATP) and is maintained by the reduction of oxygen. In animal cells, [[mitochondria]] perform similar functions. See ''[[Membrane potential]]'' article. The term '''redox state''' is often used to describe the balance of [[Nicotinamide adenine dinucleotide|NAD<sup>+</sup>/NADH]] and [[Nicotinamide adenine dinucleotide phosphate|NADP<sup>+</sup>/NADPH]] in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., [[lactic acid|lactate]] and [[pyruvate]], [[beta-hydroxybutyrate]] and [[acetoacetate]]) whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as [[Hypoxia (medical)|hypoxia]], [[Shock (medical)|shock]], and [[sepsis]]. [[Redox signaling]] involves the control of cellular processes by redox processes. ===Redox cycling=== A wide variety of [[aromaticity|aromatic compounds]] are [[enzyme|enzymatically]] reduced to form [[Radical (chemistry)|free radicals]] that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of [[flavoenzyme]]s and their [[coenzyme]]s. Once formed, these anion free radicals reduce molecular oxygen to [[superoxide]] and regenerate the unchanged parent compound. The net reaction is the oxidation of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form superoxide. This catalytic behavior has been described as futile cycle or redox cycling. Examples of redox cycling-inducing molecules are the [[herbicide]] [[paraquat]] and other [[viologen]]s and [[quinone]]s such as [[menadione]]. {{PDFlink|<ref>{{cite web|url = http://www.bioscience.org/2000/v5/d/gutier/gutier.pdf |title = gutier.doc |accessdate = 2008-06-30}}</ref>|2.76&nbsp;[[Mebibyte|MiB]]<!-- application/pdf, 2900046 bytes -->}} ==Balancing redox reactions== Describing the overall electrochemical reaction for a redox process requires a ''balancing'' of the component [[half reaction]]s for oxidation and reduction. For reactions in aqueous solution, this general involves adding [[Proton|H<sup>+</sup>]] , [[Hydroxide|OH<sup>-</sup>]] ion, [[Water (molecule)|H<sub>2</sub>O]] and electrons to compensate the oxidation changes. ===Acid medium=== In acid medium [[Proton|H<sup>+</sup>]] ions and water are added to half reactions to balance the overall reaction. For example, when [[manganese]] (II) reacts with [[sodium bismuthate]]. :<math>\mbox{Reaction unbalanced: }\mbox{Mn}^{2+}(aq) + \mbox{NaBiO}_3(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{MnO}_4^{-}(aq)\,</math> :<math>\mbox{Oxidation: }\mbox{4H}_2\mbox{O}(l)+\mbox{Mn}^{2+}(aq)\rightarrow\mbox{MnO}_4^{-}(aq) + \mbox{8H}^{+}(aq)+\mbox{5e}^{-}\,</math> :<math>\mbox{Reduction: }\mbox{2e}^{-}+ \mbox{6H}^{+}(aq) + \mbox{BiO}_3^{-}(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{3H}_2\mbox{O}(l)\,</math> The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (i.e. multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa). Addition gives: :<math>\mbox{8H}_2\mbox{O}(l)+\mbox{2Mn}^{2+}(aq)\rightarrow\mbox{2MnO}_4^{-}(aq) + \mbox{16H}^{+}(aq)+\mbox{10e}^{-}\,</math> :<math>\mbox{10e}^{-}+ \mbox{30H}^{+}(aq) + \mbox{5BiO}_3^{-}(s)\rightarrow\mbox{5Bi}^{3+}(aq) + \mbox{15H}_2\mbox{O}(l)\,</math> Reaction balanced: :<math>\mbox{14H}^{+}(aq) + \mbox{2Mn}^{2+}(aq)+ \mbox{5NaBiO}_3(s)\rightarrow\mbox{7H}_2\mbox{O}(l) + \mbox{2MnO}_4^{-}(aq)+\mbox{5Bi}^{3+}(aq)+\mbox{5Na}^{+}(aq)\,</math> Similarly for a [[propane]] [[fuel cell]] under acidic conditions: :<math>\mbox{Reaction unbalanced: }\mbox{C}_{3}\mbox{H}_{8}+\mbox{O}_{2}\rightarrow\mbox{CO}_{2}+\mbox{H}_{2}\mbox{O}\,</math> :<math>\mbox{Reduction: }\mbox{4H}^{+} + \mbox{O}_{2}+ \mbox{4e}^{-}\rightarrow\mbox{2H}_{2}\mbox{O}\,</math> :<math>\mbox{Oxidation: }\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20e}^{-}+\mbox{20H}^{+}\,</math> Balancing the number of electrons involved gives: :<math>\mbox{20H}^{+}+\mbox{5O}_{2}+\mbox{20e}^{-}\rightarrow\mbox{10H}_{2}\mbox{O}\,</math> :<math>\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20e}^{-}+\mbox{20H}^{+}\,</math> Equation balanced: :<math>\mbox{C}_{3}\mbox{H}_{8}+\mbox{5O}_{2}\rightarrow\mbox{3CO}_{2}+\mbox{4H}_{2}\mbox{O}\,</math> ===Basic medium=== In basic medium [[Hydroxide|OH<sup>-</sup>]] ions and [[Water (molecule)|water]] are added to half reactions to balance the overall reaction. For example, in the reaction between [[potassium permanganate]] and [[sodium sulfite]]: :<math>\mbox{Reaction unbalanced: }\mbox{KMnO}_{4}+\mbox{Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{MnO}_{2}+\mbox{Na}_{2}\mbox{SO}_{4}+\mbox{KOH}\,</math> :<math>\mbox{Reduction: }\mbox{3e}^{-}+\mbox{2H}_{2}\mbox{O}+\mbox{MnO}_{4}^{-}\rightarrow\mbox{MnO}_{2}+\mbox{4OH}^{-}\,</math> :<math>\mbox{Oxidation: }\mbox{2OH}^{-}+\mbox{SO}^{2-}_{3}\rightarrow\mbox{SO}^{2-}_{4}+\mbox{H}_{2}\mbox{O}+\mbox{2e}^{-}\,</math> Balancing the number of electrons in the two half-cell reactions gives: :<math>\mbox{6e}^{-}+\mbox{4H}_{2}\mbox{O}+\mbox{2MnO}_{4}^{-}\rightarrow\mbox{2MnO}_{2}+\mbox{8OH}^{-}\,</math> :<math>\mbox{6OH}^{-}+\mbox{3SO}^{2-}_{3}\rightarrow\mbox{3SO}^{2-}_{4}+\mbox{3H}_{2}\mbox{O}+\mbox{6e}^{-}\,</math> Equation balanced: :<math>\mbox{2KMnO}_{4}+\mbox{3Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{2MnO}_{2}+\mbox{3Na}_{2}\mbox{SO}_{4}+\mbox{2KOH}\,</math> ==See also== {{wikibooks|General Chemistry/Redox Reactions}} * [[Organic reduction]] * [[Hydrogenation]] * [[Bessemer process]] * [[Bioremediation]] * [[Calvin cycle]] * [[Citric acid cycle]] * [[Electrochemical cell]] * [[Electrochemistry]] * [[Galvanic cell]] * [[Membrane potential]] * [[Oxidative addition|Oxidative addition and reductive elimination]] * [[Reducing agent]] * [[Thermic reaction]] * [[Partial oxidation]] * [[Reduction potential]] ==References== {{reflist}} ==External links== * [http://www.mathcorner.hostwq.net/pages/chemlet/chemlet.html Chemical Equation Balancer], balances chemical equations, including Reduction-Oxidation (redox) reactions and reactions with several distinct solutions. Also teaches how to balance an equation using matrices and linear algebra. * [http://www.shodor.org/UNChem/advanced/redox/redoxcalc.html Redox reactions calculator] * [http://www.chemguide.co.uk/inorganic/redox/definitions.html#top Redox reactions at Chemguide] * [http://www.webqc.org/balance.php Online redox reaction equation balancer, balances equations of any half-cell and full reactions] [[Category:Soil chemistry]] [[Category:Chemical reactions]] [[ar: أكسدة]] [[bs:Redoks reakcija]] [[bg:Окислително-редукционни процеси]] [[ca:Reacció d'oxidació-reducció]] [[cs:Redoxní reakce]] [[da:Redoxreaktion]] [[de:Redoxreaktion]] [[et:Redoksreaktsioon]] [[es:Reducción-oxidación]] [[eo:Redoksa reakcio]] [[fr:Réaction d'oxydo-réduction]] [[ko:산화·환원 반응]] [[hr:Redoks]] [[io:Redoxo]] [[id:Reduksi]] [[it:Ossido-riduzione]] [[he:חמצון חיזור]] [[nl:Redoxreactie]] [[ja:酸化還元反応]] [[pl:Reakcja redoks]] [[pt:Reação redox]] [[ru:Окислительно-восстановительные реакции]] [[sk:Oxidačno-redukčná reakcia]] [[sr:Оксидација]] [[fi:Hapetus-pelkistysreaktio]] [[sv:Redox]] [[th:ปฏิกิริยารีดอกซ์]] [[zh:氧化还原反应]]