Editing Redox (section)

==Examples of redox reactions==
{{Unsourced section|date=December 2023}}[[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:

:{{chem2|H2 -> 2 [[Hydrogen ion|H(+)]] + 2 [[Electron|e(-)]]}}

The reduction reaction converts [[fluorine]] to the fluoride anion:

:{{chem2|F2 + 2 e(-) -> 2 [[Fluoride|F(-)]]}}

The half-reactions are combined so that the electrons cancel:
:{|
|align=right|{{chem|H|2}}
|→
|align=left|2 H<sup>+</sup> + 2 e<sup>−</sup>
|-
|align=right|{{chem|F|2}} + 2 e<sup>−</sup>
|→
|align=left|2 F<sup>−</sup>
|-
|colspan=3|<hr />
|-
|align=right|H<sub>2</sub> + F<sub>2</sub>
|→
|align=left|2 H<sup>+</sup> + 2 F<sup>−</sup>
|}

The protons and fluoride combine to form [[hydrofluoric acid|hydrogen fluoride]] in a non-redox reaction:
:2 H<sup>+</sup> + 2 F<sup>−</sup> → 2 HF
The overall reaction is:

:{{chem2|H2 + F2 -> 2 HF}}

===Metal displacement===

[[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 ZnSO<sub>4</sub> solution connected with a wire and a porous disk to a copper electrode in a CuSO<sub>4</sub> solution.]]

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)}}

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.

The ionic equation for this reaction is:

:{{chem2|Zn + Cu(2+) -> Zn(2+) + Cu}}

As two [[half-reaction]]s, it is seen that the zinc is oxidized:

:{{chem2|Zn -> Zn(2+) + 2 e(-)}}

And the copper is reduced:

:{{chem2|Cu(2+) + 2 e(-) -> Cu}}

===Other examples===
* The reduction of [[nitrate]] to [[nitrogen]] in the presence of an acid ([[denitrification]]):
::{{chem2|2 NO3(-) + 10 e(-) + 12 H(+) -> N2 + 6 H2O}}
* 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.
* 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]].

===Corrosion and rusting===
[[File:Rust screw.jpg|thumb|right|Oxides, such as [[iron(III) oxide]] or [[rust]], which consists of hydrated [[iron(III) oxide]]s Fe<sub>2</sub>O<sub>3</sub>·''n''H<sub>2</sub>O and [[iron(III) oxide-hydroxide]] (FeO(OH), Fe(OH)<sub>3</sub>), form when oxygen combines with other elements.]]
[[File:PyOx.JPG|thumb|Iron rusting in [[pyrite]] cubes]]
* 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:
::{{chem2|4 Fe + 3 O2 -> 2 Fe2O3}}
* The oxidation of iron(II) to iron(III) by [[hydrogen peroxide]] in the presence of an [[acid]]:
::{{chem2|Fe(2+) -> Fe(3+) + e(-)}}
::{{chem2|H2O2 + 2 e(-) -> 2 OH(-)}}
:Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel:
::{{chem2|2 Fe(2+) + H2O2 + 2 H(+) -> 2 Fe(3+) + 2 H2O}}

===Disproportionation===
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).
:{{chem2|S2O3(2-) + 2 H(+) -> S + SO2 + H2O}}
Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.<ref name=Petrucci2017/>{{rp|176}}