Redox: Difference between revisions

Latest revision as of 04:32, 15 November 2025

Template:Short description Script error: No such module "other uses". Template:Use mdy dates Template:Use American English

File:NaF.gif

Sodium "gives" one outer electron to fluorine, bonding them to form sodium fluoride. The sodium atom is oxidized, and fluorine is reduced.

File:16. Реакција меѓу силно оксидационо и редукционо средство.webm

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.

Template:Redox example.svg

Redox (Template:IPAc-en Template:Respell, Template:IPAc-en Template:Respell, reduction–oxidation^[1] or oxidation–reduction^[2]Template:Rp) is a type of chemical reaction in which the oxidation states of the reactants change.^[3] Oxidation is the loss of electrons or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state. The oxidation and reduction processes occur simultaneously in the chemical reaction.

There are two classes of redox reactions:

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.
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 forming oxides, other chemical species can serve the same function.^[4] In hydrogenation, bonds like C=C are reduced by transfer of hydrogen atoms.

Terminology

"Redox" is a portmanteau of "reduction" and "oxidation." The term was first used in a 1928 article by Leonor Michaelis and Louis B. Flexner.^[5]^[6]

Oxidation is a process in which a substance loses electrons. Reduction is a process in which a substance gains electrons.

The processes of oxidation and reduction occur simultaneously and cannot occur independently.^[4] 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,^[7] e.g., Template:Chem/ Template:Chem.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.^[4]

In electrochemical reactions the oxidation and reduction processes do occur simultaneously but are separated in space.

Oxidants

Script error: No such module "Labelled list hatnote".

Oxidation originally implied a reaction with oxygen to form an oxide. Later, the term was expanded to encompass substances 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.^[8]Template:Rp 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 agents, oxidants, or oxidizers. The oxidant removes electrons from another substance, and is thus itself reduced.^[8]Template:Rp Because it "accepts" electrons, the oxidizing agent is also called an electron acceptor. Oxidants are usually chemical substances with elements in high oxidation states^[2]Template:Rp (e.g., Template:Chem, Template:Chem, Template:Chem, Template:Chem, Template:Chem), or else highly electronegative elements (e.g. O₂, F₂, Cl₂, Br₂, I₂) that can gain extra electrons by oxidizing another substance.^[2]Template:Rp

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.^[9]

File:GHS-pictogram-rondflam.svg

The international pictogram for oxidizing chemicals

Reductants

Script error: No such module "Labelled list hatnote". 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 agents, reductants, or reducers. The reductant transfers electrons to another substance and is thus itself oxidized.^[2]Template:Rp Because it donates electrons, the reducing agent is also called an electron donor. Electron donors can also form charge transfer complexes 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.^[10] 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.^[10] Reducing equivalent refers to chemical species which transfer the equivalent of one electron in redox reactions. The term is common in biochemistry.^[11] A reducing equivalent can be an electron or a hydrogen atom as a hydride ion.^[12]

Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate electrons relatively readily.^[13]

Hydride transfer reagents, such as NaBH₄ and 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.^[14]^[15] A related method of reduction involves the use of hydrogen gas (H₂) as sources of H atoms.^[2]Template:Rp

Electronation and de-electronation

The electrochemist John Bockris proposed the words electronation and de-electronation to describe reduction and oxidation processes, respectively, when they occur at electrodes.^[16] These words are analogous to protonation and deprotonation.^[17] IUPAC has recognized the terms electronation^[18] and de-electronation.^[19]

Rates, mechanisms, and energies

Script error: No such module "Unsubst". 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.^[20]

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^[21] and outer sphere electron transfer.^[22]

Analysis of bond energies and ionization energies in water allows calculation of the thermodynamic aspects of redox reactions.^[23]

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential (ETemplate:Su), which is equal to the potential difference or voltage at equilibrium under 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:^[24]

Template:1/2 H₂ → H⁺ + e⁻

The electrode potential of each half-reaction is also known as its reduction potential (ETemplate:Su), 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⁻ → Template:1/2H₂ 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⁺).^[8]Template:Rp

For a redox reaction that takes place in a cell, the potential difference is:

ETemplate:Su = ETemplate:Su − ETemplate:Su

However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:

ETemplate:Su = −ETemplate:Su

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

ETemplate:Su = ETemplate:Su + ETemplate:Su

Examples of redox reactions

File:Redox reaction.png

Illustration of a redox reaction

In the reaction between hydrogen and fluorine, hydrogen is being oxidized and fluorine is being reduced:

@@ Line 14: / Line 14: @@
 ==Terminology==
-"Redox" is a [[portmanteau]] of the words "REDuction" and "OXidation." The term "redox" was first used in 1928.<ref>{{OEtymD|redox}}</ref>
+"Redox" is a [[portmanteau]] of "reduction" and "oxidation." The term was first used in a 1928 article by [[Leonor Michaelis]] and [[Louis B. Flexner]].<ref>{{OEtymD|redox}}</ref><ref name="i381">{{cite journal | last=Michaelis | first=L. | last2=Flexner | first2=Louis B. | title=Oxidation-reduction systems of biological significance | journal=Journal of Biological Chemistry | volume=79 | issue=2 | date=1928 | doi=10.1016/S0021-9258(20)79958-8 | doi-access=free | pages=689–722}}</ref>
 Oxidation is a process in which a substance loses electrons. Reduction is a process in which a substance gains electrons.
@@ Line 25: / Line 25: @@
 {{Main|Oxidizing agent}}
-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.<ref name="Petrucci2017">{{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 }}</ref>{{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.<ref name="Petrucci2017" />{{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<ref name="Petrucci2002" />{{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|O<sub>2</sub>]], [[Fluorine|F<sub>2</sub>]], [[Chlorine|Cl<sub>2</sub>]], [[Bromine|Br<sub>2</sub>]], [[Iodine|I<sub>2</sub>]]) that can gain extra electrons by oxidizing another substance.<ref name="Petrucci2002" />{{rp|909}}
+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.<ref name="Petrucci2017">{{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 }}</ref>{{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 agents, oxidants, or oxidizers. The oxidant removes electrons from another substance, and is thus itself reduced.<ref name="Petrucci2017" />{{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<ref name="Petrucci2002" />{{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|O<sub>2</sub>]], [[Fluorine|F<sub>2</sub>]], [[Chlorine|Cl<sub>2</sub>]], [[Bromine|Br<sub>2</sub>]], [[Iodine|I<sub>2</sub>]]) that can gain extra electrons by oxidizing another substance.<ref name="Petrucci2002" />{{rp|909}}
 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.<ref>{{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}}</ref>
@@ Line 33: / Line 33: @@
 ===<span class="anchor" id="reducing equivalent"></span>Reductants===
 {{Main|Reducing agent}}
-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.<ref name=Petrucci2002/>{{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.<ref name=Whitten>{{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}}</ref> [[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.<ref name=Whitten/> Reducing equivalent refers to [[chemical species]] which transfer the equivalent of one [[electron]] in redox reactions. The term is common in [[biochemistry]].<ref>{{Cite book| vauthors = Jain JL | title = Fundamentals of Biochemistry | publisher = S. Chand  | year = 2004 | isbn = 81-219-2453-7 }}</ref> A reducing equivalent can be an electron or a hydrogen atom as a [[Hydrogen anion|hydride ion]].<ref name="Lehninger-2017">{{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}}</ref>
+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 agents, reductants, or reducers. The reductant transfers electrons to another substance and is thus itself oxidized.<ref name=Petrucci2002/>{{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.<ref name=Whitten>{{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}}</ref> [[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.<ref name=Whitten/> Reducing equivalent refers to [[chemical species]] which transfer the equivalent of one [[electron]] in redox reactions. The term is common in [[biochemistry]].<ref>{{Cite book| vauthors = Jain JL | title = Fundamentals of Biochemistry | publisher = S. Chand  | year = 2004 | isbn = 81-219-2453-7 }}</ref> A reducing equivalent can be an electron or a hydrogen atom as a [[Hydrogen anion|hydride ion]].<ref name="Lehninger-2017">{{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}}</ref>
 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.<ref>{{cite web | url=https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch19/oxred_3.php#top | title=Oxidizing and Reducing Agents }}</ref>
 [[Hydride transfer reagents]], such as [[sodium borohydride|NaBH<sub>4</sub>]] and [[lithium aluminium hydride|LiAlH<sub>4</sub>]], reduce by atom transfer: they transfer the equivalent of hydride or H<sup>−</sup>. These reagents are widely used in the reduction of [[carbonyl]] compounds to [[alcohols]].<ref>{{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}}</ref><ref>{{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}}</ref>  A related method of reduction involves the use of hydrogen gas (H<sub>2</sub>) as sources of H atoms.<ref name=Petrucci2002/>{{rp|288}}
+===Electronation and de-electronation===
+The [[Electrochemistry|electrochemist]] [[John Bockris]] proposed the words electronation and de-electronation to describe reduction and oxidation processes, respectively, when they occur at [[electrode]]s.<ref>{{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}}</ref> These words are analogous to [[protonation]] and [[deprotonation]].<ref>{{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.}}</ref> [[IUPAC]] has recognized the terms electronation<ref>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/R05222</ref> and de-electronation.<ref>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</ref>
 ==Rates, mechanisms, and energies==
@@ Line 49: / Line 52: @@
 ==Standard electrode potentials (reduction potentials)==
 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:<ref>{{Cite book |title=Chemistry: the central science |date=2015 |publisher=Pearson |isbn=978-0-321-91041-7 |editor-last=Brown |editor-first=Theodore L. |edition=13 |location=Boston, Mass. |pages=Chapter 4}}</ref>
-:{{1/2}}H<sub>2</sub> → H<sup>+</sup> + e<sup>−</sup>
+:{{1/2}} H<sub>2</sub> → H<sup>+</sup> + e<sup>−</sup>
 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<sup>+</sup> + e<sup>−</sup> → {{1/2}}H<sub>2</sub> by definition, positive for oxidizing agents stronger than H<sup>+</sup> (e.g., +2.866 V for F<sub>2</sub>) and negative for oxidizing agents that are weaker than H<sup>+</sup> (e.g., −0.763V for Zn<sup>2+</sup>).<ref name="Petrucci2017" />{{rp|873}}
@@ Line 62: / Line 65: @@
 ==Examples of redox reactions==
-{{Unsourced section|date=December 2023}}[[File:Redox reaction.png|thumb|upright=1.35|right|Illustration of a redox reaction]]
+[[File:Redox reaction.png|thumb|upright=1.35|right|Illustration of a redox reaction]]
 In the reaction between [[hydrogen]] and [[fluorine]], hydrogen is being oxidized and fluorine is being reduced:
 :{{chem2|H2 + F2 -> 2 HF}}
-This spontaneous reaction releases a large amount of energy (542 kJ per 2 g of hydrogen) because two H-F bonds are much stronger than one H-H bond and one F-F bond. This reaction can be analyzed as two [[half-reaction]]s. The oxidation reaction converts hydrogen to [[proton]]s:
+This spontaneous reaction releases a large amount of energy (542 kJ per 2 g of hydrogen) because two H-F bonds are much stronger than one H-H bond and one F-F bond.<ref>{{Cite web |date=May 9, 2016 |title=15.11: Bond Enthalpies and Exothermic or Endothermic Reactions |url=https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_(Moore_et_al.)/15%3A_Thermodynamics-_Atoms_Molecules_and_Energy/15.11%3A_Bond_Enthalpies_and_Exothermic_or_Endothermic_Reactions |access-date=2025-11-15 |website=Chemistry LibreTexts |language=en}}</ref> This reaction can be analyzed as two [[half-reaction]]s. The oxidation reaction converts hydrogen to [[proton]]s:
 :{{chem2|H2 -> 2 [[Hydrogen ion|H(+)]] + 2 [[Electron|e(-)]]}}
@@ Line 104: / Line 107: @@
 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:
-:{{chem2|Zn (s) + CuSO4 (aq) -> ZnSO4 (aq) + Cu (s)}}
+:{{chem2|Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s)}}
 In the above reaction, zinc metal displaces the copper(II) ion from the copper sulfate solution, thus liberating free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc.
@@ Line 150: / Line 153: @@
 ==Redox reactions in biology==
-{{Unsourced section|date=December 2023}}<div class="thumb tleft skin-invert-image" style="; border: 1px solid #CCCCCC; margin:0.5em;">
+{{Unsourced section|date=December 2023}}
-{|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;"
+[[File:Extremely overripe banana.jpg|thumb|125px|upright|[[Food browning#Enzymatic browning|Enzymatic browning]] is an example of a redox reaction that takes place in most fruits and vegetables.]]
-|[[File:Ascorbic acid structure.svg|150px|ascorbic acid]]
+{{multiple image
-|}
+ | direction = vertical
-{|border="0" width=150px border="0" cellpadding="2" cellspacing="0" style="font-size: 85%; border: 1px solid #CCCCCC; margin: 0.3em;"
+ | width = 125px
-|[[File:Dehydroascorbic acid 2.svg|150px|dehydroascorbic acid]]
+ | image1 = Ascorbic acid structure.svg
-|}
+ | alt1 =
-<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>
+ | caption1 = [[Ascorbic acid]] ([[reducing agent|reduced form]] of [[vitamin C]])
-[[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.]]
+ | image2 = Dehydroascorbic acid 2.svg
+ | alt2 =
+ | caption2 = [[Dehydroascorbic acid]] ([[oxidizing agent|oxidized form]] of vitamin C)
+}}
 Many essential [[biology|biological]] processes involve redox reactions. Before some of these processes can begin, iron must be [[Assimilation (biology)|assimilated]] from the environment.<ref>{{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}}</ref>
-[[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 cellular respiration is:
+[[Cellular respiration|Aerobic cellular respiration]], for instance, is the oxidation of substrates [in this case: [[glucose]] (C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>)] and the reduction of [[oxygen]] to [[water]]. The summary equation for aerobic respiration is:
-:{{chem2|C6H12O6 + 6 O2 -> 6 CO2 + 6 H2O + Energy}}
+:{{chem2|C6H12O6 + 6 O2 -> 6 CO2 + 6 H2O + Energy}}{{Citation needed|date=September 2025|reason=First of all, at no point does glycolysis produce CO2; all CO2 comes from the Citric Acid Cycle and pyruvate oxidation. Second of all, reduction of O2 produces 12 H2O, not 6.}}
 The process of cellular respiration also depends heavily on the reduction of [[Nicotinamide adenine dinucleotide|NAD<sup>+</sup>]] to NADH and the reverse reaction (the oxidation of NADH to NAD<sup>+</sup>). [[Photosynthesis]] and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cellular respiration:
@@ Line 243: / Line 249: @@
 <!-- All sources cited below name some or all of the mnemonics listed. There is a discussion on the talk page about the inclusion of mnemonics. -->
 {{Main|List of chemistry mnemonics}}
 The key terms involved in redox can be confusing.<ref name=Robertson>{{cite book
 |last1= Robertson |first1= William

Template:Chem	→	2 H⁺ + 2 e⁻
Template:Chem + 2 e⁻	→	2 F⁻

H₂ + F₂	→	2 H⁺ + 2 F⁻

Redox: Difference between revisions

Latest revision as of 04:32, 15 November 2025

Terminology

Oxidants

Reductants

Electronation and de-electronation

Rates, mechanisms, and energies

Standard electrode potentials (reduction potentials)

Examples of redox reactions

Metal displacement

Other examples

Corrosion and rusting

Disproportionation

Redox reactions in industry

Redox reactions in biology

Redox cycling

Redox reactions in geology

Redox reactions in soils

Mnemonics

See also

References

Further reading

External links

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