# Redox: From Wikipedia, the free encyclopedia

<p class="callout success">Source: [https://en.wikipedia.org/wiki/Redox](https://en.wikipedia.org/wiki/Redox)</p>

<div class="vector-body" id="bkmrk-from-wikipedia%2C-the-"><div class="vector-body"><div class="noprint" id="bkmrk-from-wikipedia%2C-the--0">From Wikipedia, the free encyclopedia</div><div id="bkmrk--0"></div><div id="bkmrk--1"></div><div id="bkmrk--2"></div>[Jump to navigation](https://en.wikipedia.org/wiki/Redox#mw-head)[Jump to search](https://en.wikipedia.org/wiki/Redox#searchInput)<div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="hatnote navigation-not-searchable" role="note">For other uses, see [Redox (disambiguation)](https://en.wikipedia.org/wiki/Redox_(disambiguation) "Redox (disambiguation)").</div><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/a/a8/NaF.gif/350px-NaF.gif)](https://en.wikipedia.org/wiki/File:NaF.gif)<div class="thumbcaption"><div class="magnify">  
</div>[Sodium](https://en.wikipedia.org/wiki/Sodium "Sodium") and [fluorine](https://en.wikipedia.org/wiki/Fluorine "Fluorine") bonding ionically to form [sodium fluoride](https://en.wikipedia.org/wiki/Sodium_fluoride "Sodium fluoride"). Sodium loses its outer [electron](https://en.wikipedia.org/wiki/Electron "Electron") to give it a stable [electron configuration](https://en.wikipedia.org/wiki/Electron_configuration "Electron configuration"), and this electron enters the fluorine atom [exothermically](https://en.wikipedia.org/wiki/Exothermic "Exothermic"). The oppositely charged ions are then attracted to each other. The sodium is oxidized; and the fluorine is reduced.</div></div></div><div class="thumb tright"><div class="thumbinner"><span class="mw-tmh-player video"><span aria-label="47 seconds" class="mw-tmh-duration mw-tmh-label">0:47</span></span><div class="thumbcaption"><div class="magnify">  
</div>Demonstration of the reaction between a strong oxidizing and a reducing agent. When a few drops of [glycerol](https://en.wikipedia.org/wiki/Glycerol "Glycerol") (mild reducing agent) are added to powdered [potassium permanganate](https://en.wikipedia.org/wiki/Potassium_permanganate "Potassium permanganate") (strong oxidizing agent), a violent redox reaction accompanied by self-ignition starts.</div></div></div><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/1/19/Redox_example.svg/170px-Redox_example.svg.png)](https://en.wikipedia.org/wiki/File:Redox_example.svg)<div class="thumbcaption"><div class="magnify">  
</div>Example of a <a class="mw-selflink selflink">reduction–oxidation</a> reaction between sodium and chlorine, with the <span class="nowrap">*OIL RIG*</span> mnemonic<sup class="reference" id="bkmrk-%5B1%5D">[\[1\]](https://en.wikipedia.org/wiki/Redox#cite_note-1)</sup></div></div></div></div></div></div></div>**Redox** (**reduction–oxidation**, <span class="rt-commentedText nowrap"><span class="IPA nopopups noexcerpt" lang="en-fonipa">[/<span title="/ˈ/: primary stress follows">ˈ</span><span title="'r' in 'rye'">r</span><span title="/ɛ/: 'e' in 'dress'">ɛ</span><span title="'d' in 'dye'">d</span><span title="/ɒ/: 'o' in 'body'">ɒ</span><span title="'k' in 'kind'">k</span><span title="'s' in 'sigh'">s</span>/](https://en.wikipedia.org/wiki/Help:IPA/English "Help:IPA/English")</span></span> [*RED-oks*](https://en.wikipedia.org/wiki/Help:Pronunciation_respelling_key "Help:Pronunciation respelling key"), <span class="rt-commentedText nowrap"><span class="IPA nopopups noexcerpt" lang="en-fonipa">[/<span title="/ˈ/: primary stress follows">ˈ</span><span title="'r' in 'rye'">r</span><span title="/iː/: 'ee' in 'fleece'">iː</span><span title="'d' in 'dye'">d</span><span title="/ɒ/: 'o' in 'body'">ɒ</span><span title="'k' in 'kind'">k</span><span title="'s' in 'sigh'">s</span>/](https://en.wikipedia.org/wiki/Help:IPA/English "Help:IPA/English")</span></span> [*REE-doks*](https://en.wikipedia.org/wiki/Help:Pronunciation_respelling_key "Help:Pronunciation respelling key")<sup class="reference" id="bkmrk-%5B2%5D">[\[2\]](https://en.wikipedia.org/wiki/Redox#cite_note-2)</sup>) is a type of [chemical reaction](https://en.wikipedia.org/wiki/Chemical_reaction "Chemical reaction") in which the [oxidation states](https://en.wikipedia.org/wiki/Oxidation_state "Oxidation state") of [substrate](https://en.wikipedia.org/wiki/Substrate_(chemistry) "Substrate (chemistry)") change.<sup class="reference" id="bkmrk-%5B3%5D">[\[3\]](https://en.wikipedia.org/wiki/Redox#cite_note-3)</sup>

<div class="vector-body" id="bkmrk-oxidation%C2%A0is-the%C2%A0los"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output">- **Oxidation** is the *loss* of electrons or an *increase* in the oxidation state of a chemical or atoms within it.
- **Reduction** is the *gain* of electrons or a *decrease* in the oxidation state of a chemical or atoms within it.

</div></div></div></div>There are two classes of redox reactions:

<div class="vector-body" id="bkmrk-electron-transfer%C2%A0wh"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output">- *electron-transfer* where only one (usually) electron flows from the reducing agent to the oxidant. This type of redox reaction is often discussed in term of redox couples and electrode potentials.
- *atom transfer*, where an atom transfers from one substrate to another. For example, in the [rusting](https://en.wikipedia.org/wiki/Rusting "Rusting") of iron, the oxidation state of iron atoms increases as it converts to an oxide and simultaneously the oxidation state of oxygen decreases as it accepts electrons released by iron. Although oxidation reactions are commonly associated with the formation of oxides, other chemical species can serve the same function.<sup class="reference" id="bkmrk-%5B4%5D">[\[4\]](https://en.wikipedia.org/wiki/Redox#cite_note-gale-4)</sup> In [hydrogenation](https://en.wikipedia.org/wiki/Hydrogenation "Hydrogenation"), [C=C](https://en.wikipedia.org/wiki/Alkene "Alkene") (and other) bonds are reduced by transfer of hydrogen atoms.

<div aria-labelledby="mw-toc-heading" class="toc" role="navigation"> [ ] <div class="toctitle" dir="ltr" lang="en">  
</div></div></div></div></div></div>## Contents

<div class="vector-body" id="bkmrk-1terminology-1.1oxid"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div aria-labelledby="mw-toc-heading" class="toc" id="bkmrk-1terminology-1.1oxid-0" role="navigation"><div class="toctitle" dir="ltr" lang="en"><label class="toctogglelabel" for="toctogglecheckbox"></label>  
</div>- [<span class="tocnumber">1</span><span class="toctext">Terminology</span>](https://en.wikipedia.org/wiki/Redox#Terminology)
    - [<span class="tocnumber">1.1</span><span class="toctext">Oxidants</span>](https://en.wikipedia.org/wiki/Redox#Oxidants)
    - [<span class="tocnumber">1.2</span><span class="toctext">Reducers</span>](https://en.wikipedia.org/wiki/Redox#Reducers)
    - [<span class="tocnumber">1.3</span><span class="toctext">Electronation and deelectronation</span>](https://en.wikipedia.org/wiki/Redox#Electronation_and_deelectronation)
- [<span class="tocnumber">2</span><span class="toctext">Rates, mechanisms, and energies</span>](https://en.wikipedia.org/wiki/Redox#Rates,_mechanisms,_and_energies)
- [<span class="tocnumber">3</span><span class="toctext">Standard electrode potentials (reduction potentials)</span>](https://en.wikipedia.org/wiki/Redox#Standard_electrode_potentials_(reduction_potentials))
- [<span class="tocnumber">4</span><span class="toctext">Examples of redox reactions</span>](https://en.wikipedia.org/wiki/Redox#Examples_of_redox_reactions)
    - [<span class="tocnumber">4.1</span><span class="toctext">Metal displacement</span>](https://en.wikipedia.org/wiki/Redox#Metal_displacement)
    - [<span class="tocnumber">4.2</span><span class="toctext">Other examples</span>](https://en.wikipedia.org/wiki/Redox#Other_examples)
    - [<span class="tocnumber">4.3</span><span class="toctext">Corrosion and rusting</span>](https://en.wikipedia.org/wiki/Redox#Corrosion_and_rusting)
    - [<span class="tocnumber">4.4</span><span class="toctext">Disproportionation</span>](https://en.wikipedia.org/wiki/Redox#Disproportionation)
- [<span class="tocnumber">5</span><span class="toctext">Redox reactions in industry</span>](https://en.wikipedia.org/wiki/Redox#Redox_reactions_in_industry)
- [<span class="tocnumber">6</span><span class="toctext">Redox reactions in biology</span>](https://en.wikipedia.org/wiki/Redox#Redox_reactions_in_biology)
    - [<span class="tocnumber">6.1</span><span class="toctext">Redox cycling</span>](https://en.wikipedia.org/wiki/Redox#Redox_cycling)
- [<span class="tocnumber">7</span><span class="toctext">Redox reactions in geology</span>](https://en.wikipedia.org/wiki/Redox#Redox_reactions_in_geology)
- [<span class="tocnumber">8</span><span class="toctext">Redox reactions in soils</span>](https://en.wikipedia.org/wiki/Redox#Redox_reactions_in_soils)
- [<span class="tocnumber">9</span><span class="toctext">Mnemonics</span>](https://en.wikipedia.org/wiki/Redox#Mnemonics)
- [<span class="tocnumber">10</span><span class="toctext">See also</span>](https://en.wikipedia.org/wiki/Redox#See_also)
- [<span class="tocnumber">11</span><span class="toctext">References</span>](https://en.wikipedia.org/wiki/Redox#References)
- [<span class="tocnumber">12</span><span class="toctext">Further reading</span>](https://en.wikipedia.org/wiki/Redox#Further_reading)
- [<span class="tocnumber">13</span><span class="toctext">External links</span>](https://en.wikipedia.org/wiki/Redox#External_links)

</div></div></div></div></div>## <span class="mw-headline" id="bkmrk-terminology">Terminology</span>

"Redox" is a [combination](https://en.wikipedia.org/wiki/Portmanteau "Portmanteau") of the words "reduction" and "oxidation". The term "redox" was first used in 1928.<sup class="reference" id="bkmrk-%5B5%5D">[\[5\]](https://en.wikipedia.org/wiki/Redox#cite_note-5)</sup> The processes of oxidation and reduction occur simultaneously and cannot occur independently.<sup class="reference" id="bkmrk-%5B4%5D-0">[\[4\]](https://en.wikipedia.org/wiki/Redox#cite_note-gale-4)</sup> 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,<sup class="reference" id="bkmrk-%5B6%5D">[\[6\]](https://en.wikipedia.org/wiki/Redox#cite_note-6)</sup> e.g., <span class="chemf nowrap">[Fe<sup>2+</sup>  ](https://en.wikipedia.org/wiki/Iron "Iron")</span>/ <span class="chemf nowrap">[Fe<sup>3+</sup>  ](https://en.wikipedia.org/wiki/Iron "Iron")</span>.The oxidation alone and the reduction alone are each called a *[half-reaction](https://en.wikipedia.org/wiki/Half-reaction "Half-reaction")* because two half-reactions always occur together to form a whole reaction.

### <span class="mw-headline" id="bkmrk-oxidants">Oxidants</span>

*Oxidation* originally implied reaction with oxygen to form an oxide. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving the loss of electrons. 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](https://en.wikipedia.org/wiki/Oxidizing_agent "Oxidizing agent"), oxidants, or oxidizers. The oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an [electron acceptor](https://en.wikipedia.org/wiki/Electron_acceptor "Electron acceptor"). Oxidants are usually chemical substances with elements in high oxidation states (e.g., <span class="chemf nowrap">[H  
<sub>2</sub>O  
<sub>2</sub>](https://en.wikipedia.org/wiki/Hydrogen_peroxide "Hydrogen peroxide")</span>, <span class="chemf nowrap">[MnO<sup>−</sup>  
<sub>4</sub>](https://en.wikipedia.org/wiki/Permanganate "Permanganate")</span>, <span class="chemf nowrap">[CrO  
<sub>3</sub>](https://en.wikipedia.org/wiki/Chromium_trioxide "Chromium trioxide")</span>, <span class="chemf nowrap">[Cr  
<sub>2</sub>O<sup>2−</sup>  
<sub>7</sub>](https://en.wikipedia.org/wiki/Dichromate "Dichromate")</span>, <span class="chemf nowrap">[OsO  
<sub>4</sub>](https://en.wikipedia.org/wiki/Osmium(VIII)_oxide "Osmium(VIII) oxide")</span>), or else highly [electronegative](https://en.wikipedia.org/wiki/Electronegativity "Electronegativity") elements ([O<sub>2</sub>](https://en.wikipedia.org/wiki/Oxygen "Oxygen"), [F<sub>2</sub>](https://en.wikipedia.org/wiki/Fluorine "Fluorine"), [Cl<sub>2</sub>](https://en.wikipedia.org/wiki/Chlorine "Chlorine"), [Br<sub>2</sub>](https://en.wikipedia.org/wiki/Bromine "Bromine")) that can gain extra electrons by oxidizing another substance.<sup class="noprint Inline-Template Template-Fact">\[*[<span title="This claim needs references to reliable sources. (January 2021)">citation needed</span>](https://en.wikipedia.org/wiki/Wikipedia:Citation_needed "Wikipedia:Citation needed")*\]</sup>

Oxidizers are oxidants but the term is mainly reserved for sources of oxygen, particularly in the context of explosions. [Nitric acid](https://en.wikipedia.org/wiki/Nitric_acid "Nitric acid") is an oxidizer.

<div class="vector-body" id="bkmrk-the%C2%A0international%C2%A0pi"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/e/e5/GHS-pictogram-rondflam.svg/170px-GHS-pictogram-rondflam.svg.png)](https://en.wikipedia.org/wiki/File:GHS-pictogram-rondflam.svg)<div class="thumbcaption"><div class="magnify">  
</div>The [international](https://en.wikipedia.org/wiki/Globally_Harmonized_System_of_Classification_and_Labeling_of_Chemicals "Globally Harmonized System of Classification and Labeling of Chemicals") [pictogram](https://en.wikipedia.org/wiki/GHS_hazard_pictograms "GHS hazard pictograms") for oxidizing chemicals</div></div></div><div class="hatnote navigation-not-searchable" role="note">Main article: [Oxidizing agent](https://en.wikipedia.org/wiki/Oxidizing_agent "Oxidizing agent")</div></div></div></div></div>[Oxygen](https://en.wikipedia.org/wiki/Oxygen "Oxygen") is the quintessential oxidizer.

### <span class="mw-headline" id="bkmrk-reducers">Reducers</span>

<div class="vector-body" id="bkmrk-main-article%3A%C2%A0reduci"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="hatnote navigation-not-searchable" role="note">Main article: [Reducing agent](https://en.wikipedia.org/wiki/Reducing_agent "Reducing agent")</div></div></div></div></div>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](https://en.wikipedia.org/wiki/Reducing_agent "Reducing agent"), reductants, or reducers. The reductant (reducing agent) transfers electrons to another substance and is thus itself oxidized. And, because it donates electrons, the reducing agent is also called an [electron donor](https://en.wikipedia.org/wiki/Electron_donor "Electron donor"). Electron donors can also form [charge transfer complexes](https://en.wikipedia.org/wiki/Charge_transfer_complex "Charge transfer complex") with electron acceptors. The word *reduction* originally referred to the loss in weight upon heating a metallic [ore](https://en.wikipedia.org/wiki/Ore "Ore") such as a [metal oxide](https://en.wikipedia.org/wiki/Metal_oxide "Metal oxide") to extract the metal. In other words, ore was "reduced" to metal. [Antoine Lavoisier](https://en.wikipedia.org/wiki/Antoine_Lavoisier "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](https://en.wikipedia.org/wiki/Chemical_species "Chemical species") which transfer the equivalent of one [electron](https://en.wikipedia.org/wiki/Electron "Electron") in <a class="mw-selflink selflink">redox</a> reactions. The term is common in biochemistry.<sup class="reference" id="bkmrk-%5B7%5D">[\[7\]](https://en.wikipedia.org/wiki/Redox#cite_note-7)</sup> A reducing equivalent can be an electron, a hydrogen atom, as a [hydride ion](https://en.wikipedia.org/wiki/Hydrogen_anion "Hydrogen anion").<sup class="reference" id="bkmrk-%5B8%5D">[\[8\]](https://en.wikipedia.org/wiki/Redox#cite_note-:1-8)</sup>

Reductants in chemistry are very diverse. [Electropositive](https://en.wikipedia.org/wiki/Electropositive "Electropositive") elemental [metals](https://en.wikipedia.org/wiki/Metal "Metal"), such as [lithium](https://en.wikipedia.org/wiki/Lithium "Lithium"), [sodium](https://en.wikipedia.org/wiki/Sodium "Sodium"), [magnesium](https://en.wikipedia.org/wiki/Magnesium "Magnesium"), [iron](https://en.wikipedia.org/wiki/Iron "Iron"), [zinc](https://en.wikipedia.org/wiki/Zinc "Zinc"), and [aluminium](https://en.wikipedia.org/wiki/Aluminium "Aluminium"), are good reducing agents. These metals donate or *give away* electrons relatively readily. They transfer electrons.

*Hydride transfer reagents*, such as [NaBH<sub>4</sub>](https://en.wikipedia.org/wiki/Sodium_borohydride "Sodium borohydride") and [LiAlH<sub>4</sub>](https://en.wikipedia.org/wiki/Lithium_aluminium_hydride "Lithium aluminium hydride"), reduce by atom transfer: they transfer the equivalent of hydride or H<sup>-</sup>. These reagents widely used in \[the reduction of [carbonyl](https://en.wikipedia.org/wiki/Carbonyl "Carbonyl") compounds to [alcohols](https://en.wikipedia.org/wiki/Alcohols "Alcohols").<sup class="reference" id="bkmrk-%5B9%5D">[\[9\]](https://en.wikipedia.org/wiki/Redox#cite_note-9)</sup><sup class="reference" id="bkmrk-%5B10%5D">[\[10\]](https://en.wikipedia.org/wiki/Redox#cite_note-10)</sup> A related method of reduction involves the use of hydrogen gas (H<sub>2</sub>) as sources of H atoms.

### <span class="mw-headline" id="bkmrk-electronation-and-de-0">Electronation and deelectronation</span>

The electrochemist [John Bockris](https://en.wikipedia.org/wiki/John_Bockris "John Bockris") proposed the words *electronation* and *deelectronation* to describe reduction and oxidation processes, respectively, when they occur at [electrodes](https://en.wikipedia.org/wiki/Electrode "Electrode").<sup class="reference" id="bkmrk-%5B11%5D">[\[11\]](https://en.wikipedia.org/wiki/Redox#cite_note-11)</sup> These words are analogous to [protonation](https://en.wikipedia.org/wiki/Protonation "Protonation") and [deprotonation](https://en.wikipedia.org/wiki/Deprotonation "Deprotonation").<sup class="reference" id="bkmrk-%5B12%5D">[\[12\]](https://en.wikipedia.org/wiki/Redox#cite_note-12)</sup> They have not been widely adopted by chemists worldwide, although [IUPAC](https://en.wikipedia.org/wiki/IUPAC "IUPAC") has recognized the term electronation.<sup class="reference" id="bkmrk-%5B13%5D">[\[13\]](https://en.wikipedia.org/wiki/Redox#cite_note-13)</sup>

## <span id="bkmrk--4"></span><span class="mw-headline" id="bkmrk-rates%2C-mechanisms%2C-a-0">Rates, mechanisms, and energies</span>

Redox reactions can occur slowly, as in the formation of [rust](https://en.wikipedia.org/wiki/Rust "Rust"), or rapidly, as in the case of burning fuel. Electron transfer reactions are generally fast, occurring within the time of mixing.

The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred. Such reactions can also be quite complex, i.e. involve many steps. The mechanisms of electron-transfer reactions occur by two distinct pathways, [inner sphere electron transfer](https://en.wikipedia.org/wiki/Inner_sphere_electron_transfer "Inner sphere electron transfer") and [outer sphere electron transfer](https://en.wikipedia.org/wiki/Outer_sphere_electron_transfer "Outer sphere electron transfer").

Analysis of bond energies and ionization energies in water allow calculation of the thermodynamic aspects of redox reactions.

## <span id="bkmrk--5"></span><span class="mw-headline" id="bkmrk-standard-electrode-p-0">Standard electrode potentials (reduction potentials)</span>

Each half-reaction has a *standard [electrode potential](https://en.wikipedia.org/wiki/Electrode_potential "Electrode potential")* (*E*<sup>o</sup>  
<sub>cell</sub>), which is equal to the potential difference or [voltage](https://en.wikipedia.org/wiki/Voltage "Voltage") at equilibrium under [standard conditions](https://en.wikipedia.org/wiki/Standard_state "Standard state") of an [electrochemical cell](https://en.wikipedia.org/wiki/Electrochemical_cell "Electrochemical cell") in which the [cathode](https://en.wikipedia.org/wiki/Cathode "Cathode") reaction is the [half-reaction](https://en.wikipedia.org/wiki/Half-reaction "Half-reaction") considered, and the [anode](https://en.wikipedia.org/wiki/Anode "Anode") is a [standard hydrogen electrode](https://en.wikipedia.org/wiki/Standard_hydrogen_electrode "Standard hydrogen electrode") where hydrogen is oxidized:

<div class="vector-body" id="bkmrk-1%E2%81%842%C2%A0h2%C2%A0%E2%86%92-h%2B%C2%A0%2B-e%E2%88%92."><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><span class="frac" role="math"><span class="num">1</span>⁄<span class="den">2</span></span> H<sub>2</sub> → H<sup>+</sup> + e<sup>−</sup>.</dd></dl></div></div></div></div>The electrode potential of each half-reaction is also known as its *reduction potential* *E*<sup>o</sup>  
<sub>red</sub>, 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> → <span class="frac" role="math"><span class="num">1</span>⁄<span class="den">2</span></span> 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., −0.763 V for Zn<sup>2+</sup>).<sup class="reference" id="bkmrk-%5B14%5D">[\[14\]](https://en.wikipedia.org/wiki/Redox#cite_note-14)</sup>

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

<div class="vector-body" id="bkmrk-eocell%C2%A0%3D%C2%A0eocathode%C2%A0%E2%80%93"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>*E*<sup>o</sup>  
<sub>cell</sub> = *E*<sup>o</sup>  
<sub>cathode</sub> – *E*<sup>o</sup>  
<sub>anode</sub></dd></dl></div></div></div></div>However, the potential of the reaction at the anode is sometimes expressed as an *oxidation potential*:

<div class="vector-body" id="bkmrk-eoox%C2%A0%3D%C2%A0%E2%80%93eored."><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>*E*<sup>o</sup>  
<sub>ox</sub> = –*E*<sup>o</sup>  
<sub>red</sub>.</dd></dl></div></div></div></div>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

<div class="vector-body" id="bkmrk-eocell%C2%A0%3D%C2%A0eored%28catho"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>*E*<sup>o</sup>  
<sub>cell</sub> = *E*<sup>o</sup>  
<sub>red(cathode)</sub> + *E*<sup>o</sup>  
<sub>ox(anode)</sub></dd></dl></div></div></div></div>## <span class="mw-headline" id="bkmrk-examples-of-redox-re-0">Examples of redox reactions</span>

<div class="vector-body" id="bkmrk-illustration-of-a-re"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/e/ee/Redox_reaction.png/300px-Redox_reaction.png)](https://en.wikipedia.org/wiki/File:Redox_reaction.png)<div class="thumbcaption"><div class="magnify">  
</div>Illustration of a redox reaction</div></div></div></div></div></div></div>In the reaction between [hydrogen](https://en.wikipedia.org/wiki/Hydrogen "Hydrogen") and [fluorine](https://en.wikipedia.org/wiki/Fluorine "Fluorine"), hydrogen is being oxidized and fluorine is being reduced:

<div class="vector-body" id="bkmrk-h2%C2%A0%2B%C2%A0f2%C2%A0%E2%86%92-2%C2%A0hf"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><span class="chemf nowrap">H  
<sub>2</sub></span> + <span class="chemf nowrap">F  
<sub>2</sub></span> → 2 HF</dd></dl></div></div></div></div>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-reactions](https://en.wikipedia.org/wiki/Half-reaction "Half-reaction"). The oxidation reaction converts hydrogen to protons:

<div class="vector-body" id="bkmrk-h2%C2%A0%E2%86%92-2%C2%A0h%2B%C2%A0%2B-2%C2%A0e%E2%88%92"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><span class="chemf nowrap">H  
<sub>2</sub></span> → 2 [H<sup>+</sup>](https://en.wikipedia.org/wiki/Hydrogen_ion "Hydrogen ion") + 2 [e<sup>−</sup>](https://en.wikipedia.org/wiki/Electron "Electron")</dd></dl></div></div></div></div>The reduction reaction converts fluorine to the fluoride anion:

<div class="vector-body" id="bkmrk-f2%C2%A0%2B-2%C2%A0e%E2%88%92%C2%A0%E2%86%92-2%C2%A0f%E2%88%92"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><span class="chemf nowrap">F  
<sub>2</sub></span> + 2 e<sup>−</sup> → 2 [F<sup>−</sup>](https://en.wikipedia.org/wiki/Fluoride "Fluoride")</dd></dl></div></div></div></div>The half reactions are combined so that the electrons cancel:

<div class="vector-body" id="bkmrk-h2-%E2%86%92-2%C2%A0h%2B%C2%A0%2B-2%C2%A0e%E2%88%92-f2%C2%A0"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><table><tbody><tr><td align="right"><span class="chemf nowrap">H  
<sub>2</sub></span></td><td>→</td><td align="left">2 H<sup>+</sup> + 2 e<sup>−</sup></td></tr><tr><td align="right"><span class="chemf nowrap">F  
<sub>2</sub></span> + 2 e<sup>−</sup></td><td>→</td><td align="left">2 F<sup>−</sup></td></tr><tr><td colspan="3">---

</td></tr><tr><td align="right">H<sub>2</sub> + F<sub>2</sub></td><td>→</td><td align="left">2 H<sup>+</sup> + 2 F<sup>−</sup></td></tr></tbody></table>

</dd></dl></div></div></div></div>The protons and fluoride combine to form [hydrogen fluoride](https://en.wikipedia.org/wiki/Hydrofluoric_acid "Hydrofluoric acid") in a non-redox reaction:

<div class="vector-body" id="bkmrk-2%C2%A0h%2B%C2%A0%2B-2%C2%A0f%E2%88%92%C2%A0%E2%86%92-2%C2%A0hf"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>2 H<sup>+</sup> + 2 F<sup>−</sup> → 2 HF</dd></dl></div></div></div></div>The overall reaction is:

<div class="vector-body" id="bkmrk-h2%C2%A0%2B%C2%A0f2%C2%A0%E2%86%92-2%C2%A0hf-0"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd><span class="chemf nowrap">H  
<sub>2</sub></span> + <span class="chemf nowrap">F  
<sub>2</sub></span> → 2 HF</dd></dl></div></div></div></div>### <span class="mw-headline" id="bkmrk-metal-displacement">Metal displacement</span>

<div class="vector-body" id="bkmrk-a-redox-reaction-is-"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/6/6d/Galvanic_cell_with_no_cation_flow.svg/350px-Galvanic_cell_with_no_cation_flow.svg.png)](https://en.wikipedia.org/wiki/File:Galvanic_cell_with_no_cation_flow.svg)<div class="thumbcaption"><div class="magnify">  
</div>A redox reaction is the force behind an [electrochemical cell](https://en.wikipedia.org/wiki/Electrochemical_cell "Electrochemical cell") like the [Galvanic cell](https://en.wikipedia.org/wiki/Galvanic_cell "Galvanic cell") pictured. The battery is made out of a zinc electrode in a ZnSO<sub>4</sub> solution connected with a wire and a porous disk to a copper electrode in a CuSO<sub>4</sub> solution.</div></div></div></div></div></div></div>In this type of reaction, a metal atom in a compound (or in a solution) is replaced by an atom of another metal. For example, [copper](https://en.wikipedia.org/wiki/Copper "Copper") is deposited when [zinc](https://en.wikipedia.org/wiki/Zinc "Zinc") metal is placed in a [copper(II) sulfate](https://en.wikipedia.org/wiki/Copper(II)_sulfate "Copper(II) sulfate") solution:

Zn(s)+ CuSO<sub>4</sub>(aq) → ZnSO<sub>4</sub>(aq) + Cu(s)

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.

The ionic equation for this reaction is:

<div class="vector-body" id="bkmrk-zn-%2B-cu2%2B%C2%A0%E2%86%92-zn2%2B%C2%A0%2B-c"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>Zn + Cu<sup>2+</sup> → Zn<sup>2+</sup> + Cu</dd></dl></div></div></div></div>As two [half-reactions](https://en.wikipedia.org/wiki/Half-reaction "Half-reaction"), it is seen that the zinc is oxidized:

<div class="vector-body" id="bkmrk-zn-%E2%86%92-zn2%2B%C2%A0%2B-2%C2%A0e%E2%88%92"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>Zn → Zn<sup>2+</sup> + 2 e<sup>−</sup></dd></dl></div></div></div></div>And the copper is reduced:

<div class="vector-body" id="bkmrk-cu2%2B%C2%A0%2B-2%C2%A0e%E2%88%92%C2%A0%E2%86%92-cu"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>Cu<sup>2+</sup> + 2 e<sup>−</sup> → Cu</dd></dl></div></div></div></div>### <span class="mw-headline" id="bkmrk-other-examples">Other examples</span>

<div class="vector-body" id="bkmrk-the-reduction-of%C2%A0nit"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output">- The reduction of [nitrate](https://en.wikipedia.org/wiki/Nitrate "Nitrate") to [nitrogen](https://en.wikipedia.org/wiki/Nitrogen "Nitrogen") in the presence of an acid ([denitrification](https://en.wikipedia.org/wiki/Denitrification "Denitrification")): <div class="mwe-math-element"><div class="mwe-math-mathml-display mwe-math-mathml-a11y">{\displaystyle {\ce {2NO3- + 10e- + 12H+ -&gt; N2 + 6 H2O}}}</div>![{\displaystyle {\ce {2NO3- + 10e- + 12H+ -> N2 + 6 H2O}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/2a8176f7eac06a0e9951f4032d48e2b64c4946a1)</div>
- The [combustion](https://en.wikipedia.org/wiki/Combustion "Combustion") of [hydrocarbons](https://en.wikipedia.org/wiki/Hydrocarbon "Hydrocarbon"), such as in an [internal combustion engine](https://en.wikipedia.org/wiki/Internal_combustion_engine "Internal combustion engine"), produces [water](https://en.wikipedia.org/wiki/Water "Water"), [carbon dioxide](https://en.wikipedia.org/wiki/Carbon_dioxide "Carbon dioxide"), some partially oxidized forms such as [carbon monoxide](https://en.wikipedia.org/wiki/Carbon_monoxide "Carbon monoxide"), and heat [energy](https://en.wikipedia.org/wiki/Energy "Energy"). Complete oxidation of materials containing [carbon](https://en.wikipedia.org/wiki/Carbon "Carbon") produces carbon dioxide.
- The stepwise oxidation of a hydrocarbon by oxygen, in [organic chemistry](https://en.wikipedia.org/wiki/Organic_chemistry "Organic chemistry"), produces water and, successively: an [alcohol](https://en.wikipedia.org/wiki/Alcohol_(chemistry) "Alcohol (chemistry)"), an [aldehyde](https://en.wikipedia.org/wiki/Aldehyde "Aldehyde") or a [ketone](https://en.wikipedia.org/wiki/Ketone "Ketone"), a [carboxylic acid](https://en.wikipedia.org/wiki/Carboxylic_acid "Carboxylic acid"), and then a [peroxide](https://en.wikipedia.org/wiki/Peroxide "Peroxide").

</div></div></div></div>### <span class="mw-headline" id="bkmrk-corrosion-and-rustin-0">Corrosion and rusting</span>

<div class="vector-body" id="bkmrk-oxides%2C-such-as%C2%A0iron"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/4/4d/Rust_screw.jpg/220px-Rust_screw.jpg)](https://en.wikipedia.org/wiki/File:Rust_screw.jpg)<div class="thumbcaption"><div class="magnify">  
</div>Oxides, such as [iron(III) oxide](https://en.wikipedia.org/wiki/Iron(III)_oxide "Iron(III) oxide") or [rust](https://en.wikipedia.org/wiki/Rust "Rust"), which consists of hydrated [iron(III) oxides](https://en.wikipedia.org/wiki/Iron(III)_oxide "Iron(III) oxide") Fe<sub>2</sub>O<sub>3</sub>·*n*H<sub>2</sub>O and [iron(III) oxide-hydroxide](https://en.wikipedia.org/wiki/Iron(III)_oxide-hydroxide "Iron(III) oxide-hydroxide") (FeO(OH), Fe(OH)<sub>3</sub>), form when oxygen combines with other elements</div></div></div><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/e/e1/PyOx.JPG/220px-PyOx.JPG)](https://en.wikipedia.org/wiki/File:PyOx.JPG)<div class="thumbcaption"><div class="magnify">  
</div>Iron rusting in [pyrite](https://en.wikipedia.org/wiki/Pyrite "Pyrite") cubes</div></div></div>- The term [corrosion](https://en.wikipedia.org/wiki/Corrosion "Corrosion") refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. [Rusting](https://en.wikipedia.org/wiki/Rust "Rust"), the formation of [iron oxides](https://en.wikipedia.org/wiki/Iron_oxide "Iron oxide"), is a well-known example of electrochemical corrosion; it forms as a result of the oxidation of [iron](https://en.wikipedia.org/wiki/Iron "Iron") metal. Common rust often refers to [iron(III) oxide](https://en.wikipedia.org/wiki/Iron(III)_oxide "Iron(III) oxide"), formed in the following chemical reaction: <div class="mwe-math-element"><div class="mwe-math-mathml-display mwe-math-mathml-a11y">{\displaystyle {\ce {4Fe + 3O2 -&gt; 2Fe2O3}}}</div>![{\displaystyle {\ce {4Fe + 3O2 -> 2Fe2O3}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/18fc9027e7b3b3f86e1b0a4b873f60c698242fa3)</div>
- The oxidation of iron(II) to iron(III) by [hydrogen peroxide](https://en.wikipedia.org/wiki/Hydrogen_peroxide "Hydrogen peroxide") in the presence of an acid: <div class="mwe-math-element"><div class="mwe-math-mathml-display mwe-math-mathml-a11y">{\displaystyle {\ce {Fe^{2+}-&gt;{Fe^{3+}}+e-}}}</div>![{\displaystyle {\ce {Fe^{2+}->{Fe^{3+}}+e-}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/95b386fd8edc42e4111d8e24afada2fc42e63592)</div><div class="mwe-math-element"><div class="mwe-math-mathml-display mwe-math-mathml-a11y">{\displaystyle {\ce {H2O2 + 2e- -&gt; 2OH-}}}</div>![{\displaystyle {\ce {H2O2 + 2e- -> 2OH-}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/cde47282ad6b63737b024398931d022da3574816)</div>Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel: <div class="mwe-math-element"><div class="mwe-math-mathml-display mwe-math-mathml-a11y">{\displaystyle {\ce {{2Fe^{2+}}+{H2O2}+2H+-&gt;{2Fe^{3+}}+2H2O}}}</div>![{\displaystyle {\ce {{2Fe^{2+}}+{H2O2}+2H+->{2Fe^{3+}}+2H2O}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/939e479583693c6d8891be76ef00dab384d14775)</div>

</div></div></div></div>### <span class="mw-headline" id="bkmrk-disproportionation">Disproportionation</span>

A [disproportionation](https://en.wikipedia.org/wiki/Disproportionation "Disproportionation") reaction is one in which a single substance is both oxidized and reduced. For example, [thiosulfate](https://en.wikipedia.org/wiki/Thiosulfate "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](https://en.wikipedia.org/wiki/Sulfur_dioxide "Sulfur dioxide") (oxidation state +4).

<div class="vector-body" id="bkmrk-s2o32-%C2%A0%2B-2-h%2B%C2%A0%E2%86%92-s-%2B-"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>S<sub>2</sub>O<sub>3</sub><sup>2-</sup> + 2 H<sup>+</sup> → S + SO<sub>2</sub> + H<sub>2</sub>O</dd></dl></div></div></div></div>Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.<sup class="reference" id="bkmrk-%5B15%5D">[\[15\]](https://en.wikipedia.org/wiki/Redox#cite_note-15)</sup>

## <span class="mw-headline" id="bkmrk-redox-reactions-in-i-0">Redox reactions in industry</span>

[Cathodic protection](https://en.wikipedia.org/wiki/Cathodic_protection "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](https://en.wikipedia.org/wiki/Sacrificial_anode "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 [galvanized](https://en.wikipedia.org/wiki/Galvanization "Galvanization") steel, in which a sacrificial coating of zinc on steel parts protects them from rust.<sup class="noprint Inline-Template Template-Fact">\[*[<span title="This claim needs references to reliable sources. (January 2021)">citation needed</span>](https://en.wikipedia.org/wiki/Wikipedia:Citation_needed "Wikipedia:Citation needed")*\]</sup>

Oxidation is used in a wide variety of industries such as in the production of [cleaning products](https://en.wikipedia.org/wiki/Category:Cleaning_products "Category:Cleaning products") and oxidizing [ammonia](https://en.wikipedia.org/wiki/Ammonia "Ammonia") to produce [nitric acid](https://en.wikipedia.org/wiki/Nitric_acid "Nitric acid").

Redox reactions are the foundation of [electrochemical cells](https://en.wikipedia.org/wiki/Electrochemical_cell "Electrochemical cell"), which can generate electrical energy or support [electrosynthesis](https://en.wikipedia.org/wiki/Electrosynthesis "Electrosynthesis"). Metal [ores](https://en.wikipedia.org/wiki/Ore "Ore") often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by [smelting](https://en.wikipedia.org/wiki/Smelting "Smelting") at high temperature in the presence of a reducing agent. The process of [electroplating](https://en.wikipedia.org/wiki/Electroplating "Electroplating") uses redox reactions to coat objects with a thin layer of a material, as in [chrome-plated](https://en.wikipedia.org/wiki/Chrome_plating "Chrome plating") [automotive](https://en.wikipedia.org/wiki/Automotive "Automotive") parts, [silver plating](https://en.wikipedia.org/wiki/Plating#Silver_plating "Plating") [cutlery](https://en.wikipedia.org/wiki/Cutlery "Cutlery"), [galvanization](https://en.wikipedia.org/wiki/Galvanization "Galvanization") and [gold-plated](https://en.wikipedia.org/wiki/Gold-plated "Gold-plated") [jewelry](https://en.wikipedia.org/wiki/Jewelry "Jewelry").<sup class="noprint Inline-Template Template-Fact">\[*[<span title="This claim needs references to reliable sources. (January 2021)">citation needed</span>](https://en.wikipedia.org/wiki/Wikipedia:Citation_needed "Wikipedia:Citation needed")*\]</sup>

## <span class="mw-headline" id="bkmrk-redox-reactions-in-b-0">Redox reactions in biology</span>

<div class="vector-body" id="bkmrk-top%3A%C2%A0ascorbic-acid%C2%A0%28"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tleft"><table border="0" cellpadding="2" cellspacing="0" width="150px"><tbody><tr><td>[![ascorbic acid](https://upload.wikimedia.org/wikipedia/commons/thumb/5/57/Ascorbic_acid_structure.svg/150px-Ascorbic_acid_structure.svg.png)](https://en.wikipedia.org/wiki/File:Ascorbic_acid_structure.svg "ascorbic acid")</td></tr></tbody></table>

<table border="0" cellpadding="2" cellspacing="0" width="150px"><tbody><tr><td>[![dehydroascorbic acid](https://upload.wikimedia.org/wikipedia/commons/thumb/3/3a/Dehydroascorbic_acid_2.svg/150px-Dehydroascorbic_acid_2.svg.png)](https://en.wikipedia.org/wiki/File:Dehydroascorbic_acid_2.svg "dehydroascorbic acid")</td></tr></tbody></table>

<div><div class="thumbcaption"><small>Top: [ascorbic acid](https://en.wikipedia.org/wiki/Ascorbic_acid "Ascorbic acid") ([reduced form](https://en.wikipedia.org/wiki/Reducing_agent "Reducing agent") of [Vitamin C](https://en.wikipedia.org/wiki/Vitamin_C "Vitamin C"))  
Bottom: [dehydroascorbic acid](https://en.wikipedia.org/wiki/Dehydroascorbic_acid "Dehydroascorbic acid") ([oxidized form](https://en.wikipedia.org/wiki/Oxidizing_agent "Oxidizing agent") of [Vitamin C](https://en.wikipedia.org/wiki/Vitamin_C "Vitamin C"))</small></div></div></div><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/5/51/Extremely_overripe_banana.jpg/170px-Extremely_overripe_banana.jpg)](https://en.wikipedia.org/wiki/File:Extremely_overripe_banana.jpg)<div class="thumbcaption"><div class="magnify">  
</div>[Enzymatic browning](https://en.wikipedia.org/wiki/Food_browning#Enzymatic_browning "Food browning") is an example of a redox reaction that takes place in most fruits and vegetables.</div></div></div></div></div></div></div>Many important [biological](https://en.wikipedia.org/wiki/Biology "Biology") processes involve redox reactions. Before some of these processes can begin iron must be assimilated from the environment.<sup class="reference" id="bkmrk-%5B16%5D">[\[16\]](https://en.wikipedia.org/wiki/Redox#cite_note-16)</sup>

[Cellular respiration](https://en.wikipedia.org/wiki/Cellular_respiration "Cellular respiration"), for instance, is the oxidation of [glucose](https://en.wikipedia.org/wiki/Glucose "Glucose") (C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>) to [CO<sub>2</sub>](https://en.wikipedia.org/wiki/Carbon_dioxide "Carbon dioxide") and the reduction of [oxygen](https://en.wikipedia.org/wiki/Oxygen "Oxygen") to [water](https://en.wikipedia.org/wiki/Water "Water"). The summary equation for cell respiration is:

<div class="vector-body" id="bkmrk-c6h12o6%C2%A0%2B-6%C2%A0o2%C2%A0%E2%86%92-6%C2%A0c"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>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</dd></dl></div></div></div></div>The process of cell respiration also depends heavily on the reduction of [NAD<sup>+</sup>](https://en.wikipedia.org/wiki/Nicotinamide_adenine_dinucleotide "Nicotinamide adenine dinucleotide") to NADH and the reverse reaction (the oxidation of NADH to NAD<sup>+</sup>). [Photosynthesis](https://en.wikipedia.org/wiki/Photosynthesis "Photosynthesis") and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cell respiration:

<div class="vector-body" id="bkmrk-6%C2%A0co2%C2%A0%2B-6%C2%A0h2o-%2B%C2%A0ligh"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>6 CO<sub>2</sub> + 6 H<sub>2</sub>O + [light energy](https://en.wikipedia.org/wiki/Photon "Photon") → C<sub>6</sub>H<sub>12</sub>O<sub>6</sub> + 6 O<sub>2</sub></dd></dl></div></div></div></div>Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of [carbon dioxide](https://en.wikipedia.org/wiki/Carbon_dioxide "Carbon dioxide") into [sugars](https://en.wikipedia.org/wiki/Sugar "Sugar") and the oxidation of [water](https://en.wikipedia.org/wiki/Water_(molecule) "Water (molecule)") 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](https://en.wikipedia.org/wiki/Nicotinamide_adenine_dinucleotide "Nicotinamide adenine dinucleotide") (NAD<sup>+</sup>) to NADH, which then contributes to the creation of a [proton gradient](https://en.wikipedia.org/wiki/Proton_gradient "Proton gradient"), which drives the synthesis of [adenosine triphosphate](https://en.wikipedia.org/wiki/Adenosine_triphosphate "Adenosine triphosphate") (ATP) and is maintained by the reduction of oxygen. In animal cells, [mitochondria](https://en.wikipedia.org/wiki/Mitochondria "Mitochondria") perform similar functions. See the *[Membrane potential](https://en.wikipedia.org/wiki/Membrane_potential "Membrane potential")* article.

[Free radical](https://en.wikipedia.org/wiki/Free_radical "Free radical") reactions are redox reactions that occur as a part of [homeostasis](https://en.wikipedia.org/wiki/Homeostasis "Homeostasis") and killing microorganisms, 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](https://en.wikipedia.org/wiki/Antioxidant "Antioxidant"). Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.

The term *redox state* is often used to describe the balance of [GSH/GSSG](https://en.wikipedia.org/wiki/Glutathione "Glutathione"), NAD<sup>+</sup>/NADH and [NADP<sup>+</sup>/NADPH](https://en.wikipedia.org/wiki/Nicotinamide_adenine_dinucleotide_phosphate "Nicotinamide adenine dinucleotide phosphate") 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., [lactate](https://en.wikipedia.org/wiki/Lactic_acid "Lactic acid") and [pyruvate](https://en.wikipedia.org/wiki/Pyruvate "Pyruvate"), [beta-hydroxybutyrate](https://en.wikipedia.org/wiki/Beta-hydroxybutyrate "Beta-hydroxybutyrate"), and [acetoacetate](https://en.wikipedia.org/wiki/Acetoacetate "Acetoacetate")), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as [hypoxia](https://en.wikipedia.org/wiki/Hypoxia_(medical) "Hypoxia (medical)"), [shock](https://en.wikipedia.org/wiki/Shock_(circulatory) "Shock (circulatory)"), and [sepsis](https://en.wikipedia.org/wiki/Sepsis "Sepsis"). Redox mechanism also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the [CoRR hypothesis](https://en.wikipedia.org/wiki/CoRR_hypothesis "CoRR hypothesis") for the function of DNA in mitochondria and chloroplasts.

### <span class="mw-headline" id="bkmrk-redox-cycling">Redox cycling</span>

Wide varieties of [aromatic compounds](https://en.wikipedia.org/wiki/Aromaticity "Aromaticity") are [enzymatically](https://en.wikipedia.org/wiki/Enzyme "Enzyme") reduced to form [free radicals](https://en.wikipedia.org/wiki/Radical_(chemistry) "Radical (chemistry)") that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of flavoenzymes and their [coenzymes](https://en.wikipedia.org/wiki/Coenzyme "Coenzyme"). Once formed, these anion free radicals reduce molecular oxygen to [superoxide](https://en.wikipedia.org/wiki/Superoxide "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](https://en.wikipedia.org/wiki/Futile_cycle "Futile cycle") or redox cycling.

## <span class="mw-headline" id="bkmrk-redox-reactions-in-g-0">Redox reactions in geology</span>

<div class="vector-body" id="bkmrk-blast-furnaces-of%C2%A0t%C5%99"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="thumb tright"><div class="thumbinner">[![](https://upload.wikimedia.org/wikipedia/commons/thumb/f/f4/VysokePece1.jpg/250px-VysokePece1.jpg)](https://en.wikipedia.org/wiki/File:VysokePece1.jpg)<div class="thumbcaption"><div class="magnify">  
</div>Blast furnaces of [Třinec Iron and Steel Works](https://en.wikipedia.org/wiki/T%C5%99inec_Iron_and_Steel_Works "Třinec Iron and Steel Works"), Czech Republic</div></div></div></div></div></div></div>Minerals are generally oxidized derivatives of metals. Iron is mined as its [magnetite](https://en.wikipedia.org/wiki/Magnetite "Magnetite") (Fe<sub>3</sub>O<sub>4</sub>). Titanium is mined as its dioxide, usually in the form of [rutile](https://en.wikipedia.org/wiki/Rutile "Rutile") (TiO<sub>2</sub>). 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 furnaces](https://en.wikipedia.org/wiki/Blast_furnace "Blast furnace") 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:<sup class="reference" id="bkmrk-%5B17%5D">[\[17\]](https://en.wikipedia.org/wiki/Redox#cite_note-17)</sup>

<div class="vector-body" id="bkmrk-fe2o3%C2%A0%2B-3co-%E2%86%92-2fe-%2B-"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><dl><dd>Fe<sub>2</sub>O<sub>3</sub> + 3CO → 2Fe + 3CO<sub>2</sub></dd></dl></div></div></div></div>## <span class="mw-headline" id="bkmrk-redox-reactions-in-s-0">Redox reactions in soils</span><span class="mw-editsection"><span class="mw-editsection-bracket">\[</span>[edit](https://en.wikipedia.org/w/index.php?title=Redox&action=edit&section=16 "Edit section: Redox reactions in soils")<span class="mw-editsection-bracket">\]</span></span>

Electron transfer reactions are central to myriad processes and properties in soils, and electron "activity", 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.<sup class="reference" id="bkmrk-%5B18%5D">[\[18\]](https://en.wikipedia.org/wiki/Redox#cite_note-18)</sup> 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.<sup class="reference" id="bkmrk-%5B19%5D">[\[19\]](https://en.wikipedia.org/wiki/Redox#cite_note-19)</sup><sup class="reference" id="bkmrk-%5B20%5D">[\[20\]](https://en.wikipedia.org/wiki/Redox#cite_note-20)</sup>

## <span class="mw-headline" id="bkmrk-mnemonics">Mnemonics</span><span class="mw-editsection"><span class="mw-editsection-bracket">\[</span>[edit](https://en.wikipedia.org/w/index.php?title=Redox&action=edit&section=17 "Edit section: Mnemonics")<span class="mw-editsection-bracket">\]</span></span>

<div class="vector-body" id="bkmrk-main-article%3A%C2%A0list-o"><div class="vector-body"><div class="mw-body-content mw-content-ltr" dir="ltr" lang="en"><div class="mw-parser-output"><div class="hatnote navigation-not-searchable" role="note">Main article: [List of chemistry mnemonics](https://en.wikipedia.org/wiki/List_of_chemistry_mnemonics "List of chemistry mnemonics")</div></div></div></div></div>The key terms involved in redox can be confusing.<sup class="reference" id="bkmrk-%5B21%5D">[\[21\]](https://en.wikipedia.org/wiki/Redox#cite_note-Robertson-21)</sup><sup class="reference" id="bkmrk-%5B22%5D">[\[22\]](https://en.wikipedia.org/wiki/Redox#cite_note-Chemistry_ConceptsAndApplications-22)</sup> 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.<sup class="reference" id="bkmrk-%5B23%5D">[\[23\]](https://en.wikipedia.org/wiki/Redox#cite_note-Rodgers-23)</sup> These [mnemonics](https://en.wikipedia.org/wiki/Mnemonics "Mnemonics") are commonly used by students to help memorise the terminology:<sup class="reference" id="bkmrk-%5B24%5D">[\[24\]](https://en.wikipedia.org/wiki/Redox#cite_note-Zumhahl-24)</sup>

<div class="vector-body" id="bkmrk-%22oil-rig%22-%E2%80%94%C2%A0oxidatio"><div class="mw-body-content mw-content-ltr" dir="ltr" id="bkmrk-%22oil-rig%22-%E2%80%94%C2%A0oxidatio-0" lang="en"><div class="mw-parser-output">- "[OIL RIG](https://en.wikipedia.org/wiki/Oil_rig "Oil rig")" — **o**xidation **i**s **l**oss of electrons, **r**eduction **i**s **g**ain of electrons<sup class="reference" id="bkmrk-%5B21%5D-0">[\[21\]](https://en.wikipedia.org/wiki/Redox#cite_note-Robertson-21)</sup><sup class="reference" id="bkmrk-%5B22%5D-0">[\[22\]](https://en.wikipedia.org/wiki/Redox#cite_note-Chemistry_ConceptsAndApplications-22)</sup><sup class="reference" id="bkmrk-%5B23%5D-0">[\[23\]](https://en.wikipedia.org/wiki/Redox#cite_note-Rodgers-23)</sup><sup class="reference" id="bkmrk-%5B24%5D-0">[\[24\]](https://en.wikipedia.org/wiki/Redox#cite_note-Zumhahl-24)</sup>
- "LEO the lion says GER \[grr\]" — **l**oss of **e**lectrons is **o**xidation, **g**ain of **e**lectrons is **r**eduction<sup class="reference" id="bkmrk-%5B21%5D-1">[\[21\]](https://en.wikipedia.org/wiki/Redox#cite_note-Robertson-21)</sup><sup class="reference" id="bkmrk-%5B22%5D-1">[\[22\]](https://en.wikipedia.org/wiki/Redox#cite_note-Chemistry_ConceptsAndApplications-22)</sup><sup class="reference" id="bkmrk-%5B23%5D-1">[\[23\]](https://en.wikipedia.org/wiki/Redox#cite_note-Rodgers-23)</sup><sup class="reference" id="bkmrk-%5B24%5D-1">[\[24\]](https://en.wikipedia.org/wiki/Redox#cite_note-Zumhahl-24)</sup>
- "LEORA says GEROA" — the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent).<sup class="reference" id="bkmrk-%5B23%5D-2">[\[23\]](https://en.wikipedia.org/wiki/Redox#cite_note-Rodgers-23)</sup>
- "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat") — reduction occurs at the cathode and the anode is for oxidation
- "RED CAT gains what AN OX loses" – reduction at the cathode gains (electrons) what anode oxidation loses (electrons)
- "PANIC" – Positive Anode and Negative is Cathode. This applies to [electrolytic cells](https://en.wikipedia.org/wiki/Electrolytic_cell "Electrolytic cell") 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 or voltaic cells](https://en.wikipedia.org/wiki/Galvanic_cell "Galvanic cell"), such as [fuel cells](https://en.wikipedia.org/wiki/Fuel_cell "Fuel cell"), produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.

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