Array ( [0] => {{short description|Chemical reaction in which oxidation states of atoms are changed}} [1] => {{other uses}} [2] => {{Use mdy dates|date=November 2020}}{{Use American English|date=November 2020}} [3] => [4] => [[File:NaF.gif|upright=1.6|thumb|right|[[Sodium]] "gives" one outer electron to [[fluorine]], bonding them to form [[sodium fluoride]]. The sodium atom is oxidized, and the fluorine is reduced.]] [5] => [[File:16. Реакција меѓу силно оксидационо и редукционо средство.webm|thumb|right|upright=1.25|When a few drops of [[glycerol]] (mild reducing agent) are added to powdered [[potassium permanganate]] (strong oxidizing agent), a violent redox reaction accompanied by self-ignition starts.]] [6] => {{redox_example.svg}} [7] => [8] => '''Redox''' ({{IPAc-en|ˈ|r|ɛ|d|ɒ|k|s}} {{respell|RED|oks}}, {{IPAc-en|ˈ|r|iː|d|ɒ|k|s}} {{respell|REE|doks}}, '''reduction–oxidation'''{{Cite web|url=https://en.oxforddictionaries.com/definition/redox|title=redox – definition of redox in English {{!}} Oxford Dictionaries|website=Oxford Dictionaries {{!}} English|access-date=2017-05-15|archive-url=https://web.archive.org/web/20171001031248/https://en.oxforddictionaries.com/definition/redox|archive-date=2017-10-01|url-status=dead}} or '''oxidation–reduction'''{{cite book |last1=Petrucci |first1=Ralph H. |last2=Harwood |first2=William S. |last3=Herring |first3=F. Geoffrey |title=General Chemistry |date=2002 |publisher=Prentice-Hall |isbn=0-13-014329-4 |edition=8th}}{{rp|150}}) is a type of [[chemical reaction]] in which the [[oxidation state]]s of a [[reactant]] change.{{cite web|url=http://www.wiley.com/college/boyer/0470003790/reviews/redox/redox.htm|title=Redox Reactions|publisher=wiley.com|archive-url=https://web.archive.org/web/20120530081215/http://www.wiley.com/college/boyer/0470003790/reviews/redox/redox.htm|archive-date=2012-05-30|url-status=live|access-date=2012-05-09}} Oxidation is the loss of [[electron]]s or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state. [9] => [10] => There are two classes of redox reactions: [11] => * [[Electron transfer|Electron-transfer]] – Only one (usually) electron flows from the atom, ion or molecule being oxidized to the atom, ion, or molecule that is reduced. This type of redox reaction is often discussed in terms of redox couples and electrode potentials. [12] => * [[Atom transfer]] – An atom transfers from one substrate to another. For example, in the [[rusting]] of [[iron]], the oxidation state of iron atoms increases as the iron converts to an [[oxide]], and simultaneously the oxidation state of oxygen decreases as it accepts electrons released by the iron. Although oxidation reactions are commonly associated with the formation of oxides, other chemical species can serve the same function.{{cite encyclopedia | last = Haustein | first = Catherine Hinga | title = Oxidation-reduction reaction |encyclopedia=The Gale Encyclopedia of Science |edition=5th |editor1= K. Lee Lerner |editor2=Brenda Wilmoth Lerner | publisher = Gale Group | location = Farmington Hills, MI | year = 2014 | url = https://link.gale.com/apps/doc/CV2644031629/SCIC?u=dc_demo&sid=SCIC&xid=baa9fde9 }} In [[hydrogenation]], bonds like [[alkene|C=C]] are reduced by [[Transfer hydrogenation|transfer of hydrogen atoms]]. [13] => [14] => ==Terminology== [15] => "Redox" is a [[portmanteau]] of the words "reduction" and "oxidation". The term "redox" was first used in 1928.{{OEtymD|redox}} [16] => [17] => The processes of oxidation and reduction occur simultaneously and cannot occur independently. In redox processes, 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 is involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form,{{cite journal |journal= Pure and Applied Chemistry |volume= 92 |issue= 4 |title= Terminology of electrochemical methods of analysis (IUPAC Recommendations 2019) |first1= José M. |last1= Pingarrón |first2= Ján |last2= Labuda |first3= Jiří |last3= Barek |first4= Christopher M. A. |last4= Brett |first5= Maria Filomena |last5= Camões |first6= Miroslav |last6= Fojta |first7= D. Brynn |last7= Hibbert |year= 2020 |pages= 641–694 |doi= 10.1515/pac-2018-0109 |doi-access= free }} e.g., {{chem|link=Iron|Fe|2+}}/ {{chem|link=Iron|Fe|3+}}.The oxidation alone and the reduction alone are each called a ''[[half-reaction]]'' because two half-reactions always occur together to form a whole reaction. [18] => [19] => ===Oxidants=== [20] => {{Main|Oxidizing agent}} [21] => Oxidation originally implied a reaction with oxygen to form an oxide. Later, the term was expanded to encompass [[Chemical substance|substance]]s that accomplished chemical reactions similar to those of oxygen. Ultimately, the meaning was generalized to include all processes involving the loss of electrons or the increase in the oxidation state of a chemical species.{{cite book |last1=Petrucci |first1=Ralph H. |last2=Harwood |first2=William S. |last3=Herring |first3=F. Geoffrey|title=General Chemistry: Principles and Modern applications |date=2017 |location=Toronto |isbn=978-0-13-293128-1 |edition=11th |publisher=Pearson }}{{rp|A49}} Substances that have the ability to oxidize other substances (cause them to lose electrons) are said to be oxidative or oxidizing, and are known as [[oxidizing agent]]s, oxidants, or oxidizers. The oxidant removes electrons from another substance, and is thus itself reduced.{{rp|A50}} Because it "accepts" electrons, the oxidizing agent is also called an [[electron acceptor]]. Oxidants are usually chemical substances with elements in high oxidation states{{rp|159}} (e.g., {{chem|link=dinitrogen tetroxide|N|2|O|4}}, {{chem|link=permanganate|MnO|4|-}}, {{chem|link=chromium trioxide|CrO|3}}, {{chem|link=dichromate|Cr|2|O|7|2-}}, {{chem|link=Osmium(VIII) oxide|OsO|4}}), or else highly [[electronegativity|electronegative]] elements (e.g. [[Oxygen|O2]], [[Fluorine|F2]], [[Chlorine|Cl2]], [[Bromine|Br2]], [[Iodine|I2]]) that can gain extra electrons by oxidizing another substance.{{rp|909}} [22] => [23] => Oxidizers are oxidants, but the term is mainly reserved for sources of oxygen, particularly in the context of explosions. [[Nitric acid]] is a strong oxidizer.{{cite web |title=Nitric Acid Fact Sheet |url=https://essr.umd.edu/sites/default/files/2021-10/NitricAcidFactSheet.pdf |website=Department of Environmental Safety, Sustainability & Risk |publisher=University of Maryland |access-date=12 February 2024}} [24] => [25] => [[File:GHS-pictogram-rondflam.svg|thumb|upright|The [[Globally Harmonized System of Classification and Labeling of Chemicals|international]] [[GHS hazard pictograms|pictogram]] for oxidizing chemicals]] [26] => [27] => ===Reductants=== [28] => {{Main|Reducing agent}} [29] => Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known as [[reducing agent]]s, reductants, or reducers. The reductant transfers electrons to another substance and is thus itself oxidized.{{rp|159}} Because it donates electrons, the reducing agent is also called an [[electron donor]]. Electron donors can also form [[charge transfer complex]]es with electron acceptors. The word reduction originally referred to the loss in weight upon heating a metallic [[ore]] such as a [[metal oxide]] to extract the metal. In other words, ore was "reduced" to metal.{{cite book |last1=Whitten |first1=Kenneth W. |last2=Gailey |first2=Kenneth D. |last3=Davis |first3=Raymond E. |title=General Chemistry |date=1992 |publisher=Saunders College Publishin |isbn=0-03-072373-6 |page=147 |edition=4th}} [[Antoine Lavoisier]] demonstrated that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving a gain of electrons. Reducing equivalent refers to [[chemical species]] which transfer the equivalent of one [[electron]] in redox reactions. The term is common in [[biochemistry]].{{Cite book| vauthors = Jain JL | title = Fundamentals of Biochemistry | publisher = S. Chand | year = 2004 | isbn = 81-219-2453-7 }} A reducing equivalent can be an electron or a hydrogen atom as a [[Hydrogen anion|hydride ion]].{{Cite book|title=Lehninger Principles of Biochemistry | first1 = Albert L | last1 = Lehninger | first2 = David L | last2 = Nelson | first3 = Michael M | last3 = Cox | name-list-style = vanc |isbn=9781464126116|edition=Seventh|location=New York, NY|oclc=986827885|date = 2017-01-01}} [30] => [31] => Reductants in chemistry are very diverse. [[Electropositive]] elemental [[metal]]s, such as [[lithium]], [[sodium]], [[magnesium]], [[iron]], [[zinc]], and [[aluminium]], are good reducing agents. These metals donate electrons relatively readily.{{Citation needed|date=December 2023}} [32] => [33] => [[Hydride transfer reagents]], such as [[sodium borohydride|NaBH4]] and [[lithium aluminium hydride|LiAlH4]], reduce by atom transfer: they transfer the equivalent of hydride or H−. These reagents are widely used in the reduction of [[carbonyl]] compounds to [[alcohols]].{{cite book|last=Hudlický|first=Miloš|title=Reductions in Organic Chemistry |publisher=American Chemical Society |year=1996|location=Washington, D.C.|pages=429|isbn=978-0-8412-3344-7}}{{cite book|last=Hudlický|first=Miloš|title=Oxidations in Organic Chemistry|publisher=American Chemical Society|year=1990|location=Washington, D.C.|pages=[https://archive.org/details/oxidationsinorga00hudl/page/456 456]|isbn=978-0-8412-1780-5|url-access=registration|url=https://archive.org/details/oxidationsinorga00hudl/page/456}} A related method of reduction involves the use of hydrogen gas (H2) as sources of H atoms.{{rp|288}} [34] => [35] => ===Electronation and deelectronation=== [36] => The [[Electrochemistry|electrochemist]] [[John Bockris]] proposed the words electronation and deelectronation to describe reduction and oxidation processes, respectively, when they occur at [[electrode]]s.{{Cite book|last1=Bockris |first1=John O'M. |last2=Reddy |first2=Amulya K. N. |title=Modern Electrochemistry |publisher=Plenum Press|year= 1970|pages=352–3}} These words are analogous to [[protonation]] and [[deprotonation]].{{cite book |last1=Bockris |first1=John O'M. |last2=Reddy |first2=Amulya K.N. |volume=1 |title=Modern Electrochemistry |orig-year=1970 |publisher=Springer Science & Business Media |year=2013 |isbn=9781461574675 |url=https://books.google.com/books?id=0xzlBwAAQBAJ&pg=PA494 |page=494 |access-date=29 March 2020 |quote=The homogeneous proton-transfer reactions described are similar to homogeneous electron-transfer reactions in that the overall electron-transfer reaction can be decomposed into one electronation reaction and one deelectronation reaction.}} They have not been widely adopted by chemists worldwide,{{Citation needed|date=December 2023}} although [[IUPAC]] has recognized the terms electronationIUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Compiled by A. D. McNaught and A. Wilkinson. Blackwell Scientific Publications, Oxford (1997). Online version (2019-) created by S. J. Chalk. {{ISBN|0-9678550-9-8}}. https://goldbook.iupac.org/terms/view/R05222 and de-electronation.IUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Compiled by A. D. McNaught and A. Wilkinson. Blackwell Scientific Publications, Oxford (1997). Online version (2019-) created by S. J. Chalk. {{ISBN|0-9678550-9-8}}. https://goldbook.iupac.org/terms/view/O04362 [37] => [38] => ==Rates, mechanisms, and energies== [39] => {{Expand section|date=April 2023}} [40] => Redox reactions can occur slowly, as in the formation of [[rust]], or rapidly, as in the case of burning [[fuel]]. Electron transfer reactions are generally fast, occurring within the time of mixing.{{cn|date=February 2024}} [41] => [42] => The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred. Such reactions can also be quite complex, involving many steps. The mechanisms of electron-transfer reactions occur by two distinct pathways, [[inner sphere electron transfer]] and [[outer sphere electron transfer]].{{cn|date=February 2024}} [43] => [44] => Analysis of bond energies and [[Ionization energy|ionization energies]] in water allow calculation of the thermodynamic aspects of redox reactions.{{cn|date=February 2024}} [45] => [46] => ==Standard electrode potentials (reduction potentials)== [47] => {{Refimprove|section|date=December 2023}} [48] => Each half-reaction has a standard [[electrode potential]] (''E''{{su|p=o|b=cell}}), which is equal to the potential difference or [[voltage]] at equilibrium under [[standard state|standard conditions]] of an [[electrochemical cell]] in which the [[cathode]] reaction is the [[half-reaction]] considered, and the [[anode]] is a [[standard hydrogen electrode]] where hydrogen is oxidized: [49] => :{{1/2}}H2 → H+ + e− [50] => [51] => The electrode potential of each half-reaction is also known as its reduction potential (''E''{{su|p=o|b=red}}), or potential when the half-reaction takes place at a cathode. The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H+ + e− → {{1/2}}H2 by definition, positive for oxidizing agents stronger than H+ (e.g., +2.866 V for F2) and negative for oxidizing agents that are weaker than H+ (e.g., −0.763V for Zn2+).{{rp|873}} [52] => [53] => For a redox reaction that takes place in a cell, the potential difference is: [54] => :''E''{{su|p=o|b=cell}} = ''E''{{su|p=o|b=cathode}} – ''E''{{su|p=o|b=anode}} [55] => [56] => However, the potential of the reaction at the anode is sometimes expressed as an ''oxidation potential'': [57] => :''E''{{su|p=o|b=ox}} = –''E''{{su|p=o|b=red}} [58] => The oxidation potential is a measure of the tendency of the reducing agent to be oxidized but does not represent the physical potential at an electrode. With this notation, the cell voltage equation is written with a plus sign [59] => :''E''{{su|p=o|b=cell}} = ''E''{{su|p=o|b=red(cathode)}} + ''E''{{su|p=o|b=ox(anode)}} [60] => [61] => ==Examples of redox reactions== [62] => {{Unsourced section|date=December 2023}}[[File:Redox reaction.png|thumb|upright=1.35|right|Illustration of a redox reaction]] [63] => In the reaction between [[hydrogen]] and [[fluorine]], hydrogen is being oxidized and fluorine is being reduced: [64] => [65] => :{{chem2|H2 + F2 -> 2 HF}} [66] => [67] => This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the F-F bond. This reaction can be analyzed as two [[half-reaction]]s. The oxidation reaction converts hydrogen to [[proton]]s: [68] => [69] => :{{chem2|H2 -> 2 [[Hydrogen ion|H(+)]] + 2 [[Electron|e(-)]]}} [70] => [71] => The reduction reaction converts [[fluorine]] to the fluoride anion: [72] => [73] => :{{chem2|F2 + 2 e(-) -> 2 [[Fluoride|F(-)]]}} [74] => [75] => The half reactions are combined so that the electrons cancel: [76] => :{| [77] => |align=right|{{chem|H|2}} [78] => |→ [79] => |align=left|2 H+ + 2 e− [80] => |- [81] => |align=right|{{chem|F|2}} + 2 e− [82] => |→ [83] => |align=left|2 F− [84] => |- [85] => |colspan=3|
[86] => |- [87] => |align=right|H2 + F2 [88] => |→ [89] => |align=left|2 H+ + 2 F− [90] => |} [91] => [92] => The protons and fluoride combine to form [[hydrofluoric acid|hydrogen fluoride]] in a non-redox reaction: [93] => :2 H+ + 2 F− → 2 HF [94] => The overall reaction is: [95] => [96] => :{{chem2|H2 + F2 -> 2 HF}} [97] => [98] => ===Metal displacement=== [99] => [100] => [[File:Galvanic cell with no cation flow.svg|thumb|upright=1.6|A redox reaction is the force behind an [[electrochemical cell]] like the [[Galvanic cell]] pictured. The battery is made out of a zinc electrode in a ZnSO4 solution connected with a wire and a porous disk to a copper electrode in a CuSO4 solution.]] [101] => [102] => In this type of reaction, a [[metal]] atom in a compound or solution is replaced by an atom of another metal. For example, [[copper]] is deposited when [[zinc]] metal is placed in a [[copper(II) sulfate]] solution: [103] => [104] => :{{chem2|Zn (s) + CuSO4 (aq) -> ZnSO4 (aq) + Cu (s)}} [105] => [106] => In the above reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc. [107] => [108] => The ionic equation for this reaction is: [109] => [110] => :{{chem2|Zn + Cu(2+) -> Zn(2+) + Cu}} [111] => [112] => As two [[half-reaction]]s, it is seen that the zinc is oxidized: [113] => [114] => :{{chem2|Zn -> Zn(2+) + 2 e(-)}} [115] => [116] => And the copper is reduced: [117] => [118] => :{{chem2|Cu(2+) + 2 e(-) -> Cu}} [119] => [120] => ===Other examples=== [121] => * The reduction of [[nitrate]] to [[nitrogen]] in the presence of an acid ([[denitrification]]): [122] => ::{{chem2|2 NO3(-) + 10 e(-) + 12 H(+) -> N2 + 6 H2O}} [123] => * The [[combustion]] of [[hydrocarbon]]s, such as 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. [124] => * The stepwise oxidation of a hydrocarbon by oxygen, in [[organic chemistry]], produces water and, successively: an [[Alcohol (chemistry)|alcohol]], an [[aldehyde]] or a [[ketone]], a [[carboxylic acid]], and then a [[peroxide]]. [125] => [126] => ===Corrosion and rusting=== [127] => [[File:Rust screw.jpg|thumb|right|Oxides, such as [[iron(III) oxide]] or [[rust]], which consists of hydrated [[iron(III) oxide]]s Fe2O3·''n''H2O and [[iron(III) oxide-hydroxide]] (FeO(OH), Fe(OH)3), form when oxygen combines with other elements.]] [128] => [[File:PyOx.JPG|thumb|Iron rusting in [[pyrite]] cubes]] [129] => * The term [[corrosion]] refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. [[Rust]]ing, the formation of [[iron oxide]]s, is a well-known example of electrochemical corrosion: it forms as a result of the oxidation of [[iron]] metal. Common rust often refers to [[iron(III) oxide]], formed in the following chemical reaction: [130] => ::{{chem2|4 Fe + 3 O2 -> 2 Fe2O3}} [131] => * The oxidation of iron(II) to iron(III) by [[hydrogen peroxide]] in the presence of an [[acid]]: [132] => ::{{chem2|Fe(2+) -> Fe(3+) + e(-)}} [133] => ::{{chem2|H2O2 + 2 e(-) -> 2 OH(-)}} [134] => :Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel: [135] => ::{{chem2|2 Fe(2+) + H2O2 + 2 H(+) -> 2 Fe(3+) + 2 H2O}} [136] => [137] => ===Disproportionation=== [138] => A [[disproportionation]] reaction is one in which a single substance is both oxidized and reduced. For example, [[thiosulfate]] ion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur (oxidation state 0) and [[sulfur dioxide]] (oxidation state +4). [139] => :{{chem2|S2O3(2-) + 2 H(+) -> S + SO2 + H2O}} [140] => Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.{{rp|176}} [141] => [142] => ==Redox reactions in industry== [143] => [[Cathodic protection]] is a technique used to control the corrosion of a metal surface by making it the [[cathode]] of an electrochemical cell. A simple method of protection connects protected metal to a more easily corroded "[[sacrificial anode]]" to act as the [[anode]]. The sacrificial metal, instead of the protected metal, then corrodes. A common application of cathodic protection is in [[galvanization|galvanized]] steel, in which a sacrificial coating of zinc on steel parts protects them from rust.{{cn|date=February 2024}} [144] => [145] => Oxidation is used in a wide variety of industries such as in the production of [[:Category:Cleaning products|cleaning products]] and oxidizing [[ammonia]] to produce [[nitric acid]].{{cn|date=February 2024}} [146] => [147] => Redox reactions are the foundation of [[electrochemical cell]]s, which can generate electrical energy or support [[electrosynthesis]]. Metal [[ore]]s often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by [[smelting]] at high temperature in the presence of a reducing agent. The process of [[electroplating]] uses redox reactions to coat objects with a thin layer of a material, as in [[Chrome plating|chrome-plated]] [[automotive]] parts, [[Plating#Silver plating|silver plating]] [[cutlery]], [[galvanization]] and [[gold-plated]] [[jewelry]].{{cn|date=February 2024}} [148] => [149] => ==Redox reactions in biology== [150] => {{Unsourced section|date=December 2023}}
[151] => {|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;" [152] => |[[File:Ascorbic acid structure.svg|150px|ascorbic acid]] [153] => |} [154] => {|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;" [155] => |[[File:Dehydroascorbic acid 2.svg|150px|dehydroascorbic acid]] [156] => |} [157] =>
Top: [[ascorbic acid]] ([[reducing agent|reduced form]] of [[Vitamin C]])
Bottom: [[dehydroascorbic acid]] ([[oxidizing agent|oxidized form]] of [[Vitamin C]])
[158] => [[File:Extremely overripe banana.jpg|thumb|upright|[[Food browning#Enzymatic browning|Enzymatic browning]] is an example of a redox reaction that takes place in most fruits and vegetables.]] [159] => Many important [[biology|biological]] processes involve redox reactions. Before some of these processes can begin iron must be [[Assimilation (biology)|assimilated]] from the environment.{{Cite book|chapter-url=https://www.degruyter.com/document/doi/10.1515/9783110589771-005|doi = 10.1515/9783110589771-005|chapter = Titles of Volumes 1–44 in the Metal Ions in Biological Systems Series|title = Metals, Microbes, and Minerals - the Biogeochemical Side of Life|year = 2021|pages = xxiii-xxiv|publisher = De Gruyter|isbn = 9783110588903|s2cid = 242013948}} [160] => [161] => [[Cellular respiration]], for instance, is the oxidation of [[glucose]] (C6H12O6) to [[carbon dioxide|CO2]] and the reduction of [[oxygen]] to [[water]]. The summary equation for cell respiration is: [162] => [163] => :{{chem2|C6H12O6 + 6 O2 -> 6 CO2 + 6 H2O + Energy}} [164] => [165] => The process of cell respiration also depends heavily on the reduction of [[Nicotinamide adenine dinucleotide|NAD+]] to NADH and the reverse reaction (the oxidation of NADH to NAD+). [[Photosynthesis]] and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration: [166] => [167] => :{{chem2|6 CO2 + 6 H2O + [[photon|light energy]] -> C6H12O6 + 6 O2}} [168] => [169] => [[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, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce [[nicotinamide adenine dinucleotide]] (NAD+) to NADH, 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. [170] => [171] => {{See also|Membrane potential}} [172] => [173] => [[Free radical]] reactions are redox reactions that occur as a part of [[homeostasis]] and killing [[microorganism]]s, where an electron detaches from a molecule and then reattaches almost instantaneously. Free radicals are a part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or an [[antioxidant]]. [174] => [175] => The term redox state is often used to describe the balance of [[Glutathione|GSH/GSSG]], NAD+/NADH and [[Nicotinamide adenine dinucleotide phosphate|NADP+/NADPH]] in a biological system such as a [[Cell (biology)|cell]] or [[Organ (biology)|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. Redox mechanisms also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the [[CoRR hypothesis]] for the function of [[DNA]] in [[Mitochondrion|mitochondria]] and [[chloroplast]]s. [176] => [177] => ===Redox cycling=== [178] => Wide varieties 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 a [[futile cycle]] or redox cycling. [179] => [180] => ==Redox reactions in geology== [181] => [[File:VysokePece1.jpg|thumb|250px|right|Blast furnaces of [[Třinec Iron and Steel Works]], Czech Republic]] [182] => Minerals are generally oxidized derivatives of metals. Iron is mined as its [[magnetite]] (Fe3O4). Titanium is mined as its dioxide, usually in the form of [[rutile]] (TiO2). To obtain the corresponding metals, these oxides must be reduced, which is often achieved by heating these oxides with carbon or carbon monoxide as reducing agents. [[Blast furnace]]s are the reactors where iron oxides and coke (a form of carbon) are combined to produce molten iron.The main chemical reaction producing the molten iron is:{{Ullmann |doi=10.1002/14356007.a14_461.pub2|title=Iron|year=2006|last1=Oeters|first1=Franz|last2=Ottow|first2=Manfred|last3=Meiler|first3=Heinrich|last4=Lüngen|first4=Hans Bodo|last5=Koltermann|first5=Manfred|last6=Buhr|first6=Andreas|last7=Yagi|first7=Jun-Ichiro|last8=Formanek|first8=Lothar|last9=Rose|first9=Fritz|last10=Flickenschild|first10=Jürgen|last11=Hauk|first11=Rolf|last12=Steffen|first12=Rolf|last13=Skroch|first13=Reiner|last14=Mayer-Schwinning|first14=Gernot|last15=Bünnagel|first15=Heinz-Lothar|last16=Hoff|first16=Hans-Georg}} [183] => :{{chem2|Fe2O3 + 3 CO -> 2 Fe + 3 CO2}} [184] => [185] => == Redox reactions in soils == [186] => [[Electron transfer]] reactions are central to myriad processes and properties in soils, and [[redox potential]], quantified as Eh (platinum electrode potential ([[voltage]]) relative to the standard hydrogen electrode) or pe (analogous to pH as -log electron activity), is a master variable, along with pH, that controls and is governed by chemical reactions and biological processes. Early theoretical research with applications to flooded soils and [[paddy rice]] production was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.{{cite journal |last=Ponnamperuma |first=Felix Nelson |date=1992 |title=The chemistry of submerged soils |journal=Advances in Agronomy |volume=24 |pages=29–96 |doi=10.1016/S0065-2113(08)60633-1 |isbn=9780120007240 |url=https://scholar.google.fr/scholar_url?url=https://www.researchgate.net/profile/Clinton-Rissmann/post/What-are-the-likely-chemical-changes-in-a-soil-as-it-becomes-increasingly-anaerobic/attachment/59d649f079197b80779a455a/AS%253A473170680520704%25401489824089664/download/Chemistry%2Bof%2BSubmerged%2Bsoils%2Bby%2BPonnamperuma.pdf&hl=fr&sa=X&ei=znv9ZISfHcKsmgG8yaLgBQ&scisig=AFWwaeakfkMSaifWn6g52i2Tultp&oi=scholarr |access-date=10 September 2023}} Later work built on this foundation, and expanded it for understanding redox reactions related to heavy metal oxidation state changes, [[pedogenesis]] and morphology, organic compound degradation and formation, [[free radical]] chemistry, [[wetland]] delineation, [[soil remediation]], and various methodological approaches for characterizing the redox status of soils.{{cite journal |last1=Bartlett |first1=Richmond J. |last2=James |first2=Bruce R. |date=1991 |title=Redox chemistry of soils |journal=Advances in Agronomy |volume=39 |pages=151–208}}{{cite book|last1=James |first1=Bruce R. |title=Handbook of soil sciences: properties and processes |edition=second |last2=Brose |first2=Dominic A. |publisher=[[CRC Press]] |year=2012 |isbn=978-1-4398-0305-9 |editor-last1=Huang |editor-first1=Pan Ming |location=Boca Raton, Florida |pages=14-1 -- 14-24 |chapter=Oxidation-reduction phenomena |editor-last2=Li |editor-first2=Yuncong |editor-last3=Sumner |editor-first3=Malcolm E.}} [187] => [188] => [240] => [241] => == Mnemonics == [242] => [243] => {{Main|List of chemistry mnemonics}} [244] => The key terms involved in redox can be confusing.{{cite book [245] => |last1= Robertson |first1= William [246] => |title= More Chemistry Basics [247] => |url= https://books.google.com/books?id=hIzuarlXtH4C&pg=PA82 [248] => |publisher= National Science Teachers Association [249] => |year= 2010 [250] => |page= 82 [251] => |isbn=978-1-936137-74-9}}{{cite book [252] => |last1= Phillips |first1= John [253] => |last2= Strozak |first2= Victor [254] => |last3= Wistrom |first3= Cheryl [255] => |title= Chemistry: Concepts and Applications [256] => |publisher= Glencoe McGraw-Hill [257] => |year= 2000 [258] => |page= 558 [259] => |isbn= 978-0-02-828210-7}} For example, a reagent that is oxidized loses electrons; however, that reagent is referred to as the reducing agent. Likewise, a reagent that is reduced gains electrons and is referred to as the oxidizing agent.{{cite book [260] => |last1= Rodgers |first1= Glen [261] => |title= Descriptive Inorganic, Coordination, and Solid-State Chemistry [262] => |url= https://books.google.com/books?id=g_ybia0hGw8C&pg=PA330 [263] => |publisher= Brooks/Cole, Cengage Learning [264] => |year= 2012 [265] => |page= 330 [266] => |isbn=978-0-8400-6846-0}} These [[mnemonics]] are commonly used by students to help memorise the terminology:{{cite book [267] => |last1= Zumdahl |first1= Steven [268] => |last2= Zumdahl |first2= Susan [269] => |title= Chemistry [270] => |url= https://books.google.com/books?id=IdhYqXy37KIC&pg=PA160 [271] => |publisher= Houghton Mifflin [272] => |year= 2009 [273] => |page= 160 [274] => |isbn=978-0-547-05405-6}} [275] => * "[[Oil rig|OIL RIG]]" — '''o'''xidation '''i'''s '''l'''oss of electrons, '''r'''eduction '''i'''s '''g'''ain of electrons [276] => * "LEO the lion says GER [grr]" — '''l'''oss of '''e'''lectrons is '''o'''xidation, '''g'''ain of '''e'''lectrons is '''r'''eduction [277] => * "LEORA says GEROA" — the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent). [278] => * "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat") — reduction occurs at the cathode and the anode is for oxidation [279] => * "RED CAT gains what AN OX loses" – reduction at the cathode gains (electrons) what anode oxidation loses (electrons) [280] => * "PANIC" – Positive Anode and Negative is Cathode. This applies to [[electrolytic cell]]s which release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. These [[galvanic cell|galvanic or voltaic cell]]s, such as [[fuel cell]]s, produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode. [281] => [282] => ==See also== [283] => {{div col|colwidth=20em}} [284] => * [[Anaerobic respiration]] [285] => * [[Bessemer process]] [286] => * [[Bioremediation]] [287] => * [[Calvin cycle]] [288] => * [[Chemical equation]] [289] => * [[Chemical looping combustion]] [290] => * [[Citric acid cycle]] [291] => * [[Electrochemical series]] [292] => * [[Electrochemistry]] [293] => * [[Electrolysis]] [294] => * [[Electron equivalent]] [295] => * [[Electron transport chain]] [296] => * [[Electrosynthesis]] [297] => * [[Galvanic cell]] [298] => * [[Hydrogenation]] [299] => * [[Membrane potential]] [300] => * [[Microbial fuel cell]] [301] => * [[Murburn concept]] [302] => * [[Nucleophilic abstraction]] [303] => * [[Organic redox reaction]] [304] => * [[Oxidative addition|Oxidative addition and reductive elimination]] [305] => * [[Oxidative phosphorylation]] [306] => * [[Partial oxidation]] [307] => * [[Pro-oxidant]] [308] => * [[Redox gradient]] [309] => * [[Redox potential]] [310] => * [[Reducing agent]] [311] => * [[Reducing atmosphere]] [312] => * [[Reduction potential]] [313] => * [[Exothermic|Thermic reaction]] [314] => * [[Transmetalation]] [315] => * [[Sulfur cycle]] [316] => {{div col end}} [317] => [318] => ==References== [319] => {{reflist}} [320] => [321] => ==Further reading== [322] => * {{cite book |editor1-last=Schüring |editor1-first=J. |editor2-last=Schulz |editor2-first=H. D. |editor3-last=Fischer |editor3-first=W. R. |editor4-last=Böttcher |editor4-first=J. |editor5-last=Duijnisveld |editor5-first=W. H. |title=Redox: Fundamentals, Processes and Applications |date=1999 |publisher=Springer-Verlag |location=Heidelberg |isbn=978-3-540-66528-1 |pages=246 |hdl=10013/epic.31694.d001 }} [323] => * {{Cite book|doi=10.1021/bk-2011-1071|isbn=978-0-8412-2652-4|title=Aquatic Redox Chemistry|series=ACS Symposium Series|year=2011|editor1-last=Tratnyek|editor1-first=Paul G.|editor2-last=Grundl|editor2-first=Timothy J.|editor3-last=Haderlein|editor3-first=Stefan B.|volume=1071 }} [324] => [325] => ==External links== [326] => {{wikiquote}} [327] => {{Commons category|Redox reactions}} [328] => {{Reaction mechanisms}} [329] => {{Authority control}} [330] => [331] => [[Category:Soil chemistry]] [332] => [[Category:Chemical reactions]] [333] => [[Category:Articles containing video clips]] [334] => [[Category:Redox]] [335] => [[Category:Reaction mechanisms]] [] => )
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Redox

Redox is a term derived from the words "reduction" and "oxidation," which are two fundamental chemical processes that occur together in reactions involving the transfer of electrons between molecules. It refers to a type of chemical reaction in which one molecule is reduced, or gains electrons, while another molecule is oxidized, or loses electrons.

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It refers to a type of chemical reaction in which one molecule is reduced, or gains electrons, while another molecule is oxidized, or loses electrons. The Wikipedia page on redox provides a comprehensive overview of this important concept in chemistry. It explains the basic principles of redox reactions, including the involvement of half-reactions, oxidation numbers, and the role of electron transfer. The page also delves into the various applications of redox reactions, such as in electrochemistry, energy production, and biological systems. The history of redox is explored, highlighting its development from early observations of substances like oxygen and hydrogen, to the formulation of the modern understanding of electron transfer. The page also discusses the different types of redox reactions, including combination reactions, decomposition reactions, and displacement reactions. Furthermore, the article provides examples of redox reactions in everyday life, such as rusting, combustion, and photosynthesis. It also explains how redox reactions are utilized in various industries, such as pharmaceuticals, environmental remediation, and metallurgy. The page on redox includes additional sections dedicated to the measurement and balancing of redox reactions, as well as the concept of redox potential and its importance in predicting the spontaneity of a reaction. The role of redox reactions in both chemical and biological systems is explored, emphasizing their significance in processes such as cellular respiration and oxidation of molecules for energy production. In conclusion, the Wikipedia page on redox provides a detailed and comprehensive overview of this fundamental concept in chemistry. It covers the history, principles, applications, and various types of redox reactions, making it a valuable resource for students, researchers, and anyone interested in understanding the role of electron transfer in chemical processes.

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