Editing Nitrite (section)

==Reactions==
=== Acid-base properties ===
Nitrite is the conjugate base of the weak acid [[nitrous acid]]:
:HNO<sub>2</sub> {{eqm}} H<sup>+</sup> + {{chem|NO|2|-}};{{spaces|5}} [[acid dissociation constant|p''K''<sub>a</sub>]] ≈ 3.3 at 18&nbsp;°C<ref name=scdb>[http://www.acadsoft.co.uk/scdbase/scdbase.htm IUPAC SC-Database] {{Webarchive|url=https://web.archive.org/web/20170619235720/http://www.acadsoft.co.uk/scdbase/scdbase.htm |date=19 June 2017 }} A comprehensive database of published data on equilibrium constants of metal complexes and ligands</ref>
Nitrous acid is also highly unstable, tending to [[disproportionate]]:
:3&nbsp;HNO<sub>2</sub> (aq) {{eqm}} H<sub>3</sub>O<sup>+</sup> + {{chem|NO|3|-}} + 2&nbsp;NO
This reaction is slow at 0&nbsp;°C.<ref name=p461/> Addition of acid to a solution of a nitrite in the presence of a [[reducing agent]], such as iron(II), is a way to make [[nitric oxide]] (NO) in the laboratory.

=== Oxidation and reduction ===
The formal [[oxidation state]] of the nitrogen atom in nitrite is +3. This means that it can be either oxidized to oxidation states +4 and +5, or reduced to oxidation states as low as −3. Standard [[reduction potential]]s for reactions directly involving nitrous acid are shown in the table below:<ref>{{Greenwood&Earnshaw2nd|page=431}}</ref> 
:{|class="wikitable"
|-
!Half-reaction||''E''<sup>0</sup> ([[Volt|V]])
|-
| {{chem|NO|3|-}} + 3 H<sup>+</sup> + 2 e<sup>−</sup> {{eqm}} HNO<sub>2</sub> + H<sub>2</sub>O|| +0.94
|-
| 2 HNO<sub>2</sub> + 4 H<sup>+</sup> + 4 e<sup>−</sup> {{eqm}} H<sub>2</sub>N<sub>2</sub>O<sub>2</sub> + 2 H<sub>2</sub>O||+0.86
|-
| N<sub>2</sub>O<sub>4</sub> + 2 H<sup>+</sup> + 2 e<sup>−</sup> {{eqm}} 2 HNO<sub>2</sub>||+1.065
|-
| 2 HNO<sub>2</sub>+ 4 H<sup>+</sup> + 4 e<sup>−</sup> {{eqm}} N<sub>2</sub>O + 3 H<sub>2</sub>O||+1.29
|}
The data can be extended to include products in lower oxidation states. For example:
:H<sub>2</sub>N<sub>2</sub>O<sub>2</sub> + 2 H<sup>+</sup> + 2 ''e''<sup>−</sup> {{eqm}} N<sub>2</sub> + 2 H<sub>2</sub>O;{{spaces|5}} ''E''<sup>0</sup> = +2.65&nbsp;V

Oxidation reactions usually result in the formation of the [[nitrate]] ion, with nitrogen in oxidation state +5. For example, oxidation with [[permanganate]] ion can be used for quantitative analysis of nitrite (by [[titration]]):
:5 {{chem|NO|2|-}} + 2 {{chem|MnO|4|-}} + 6 H<sup>+</sup> → 5 {{chem|NO|3|-}} + 2 Mn<sup>2+</sup> + 3 H<sub>2</sub>O

The product of reduction reactions with nitrite ion are varied, depending on the [[reducing agent]] used and its strength. With [[sulfur dioxide]], the products are NO and N<sub>2</sub>O; with tin(II) (Sn<sup>2+</sup>) the product is [[hyponitrous acid]] (H<sub>2</sub>N<sub>2</sub>O<sub>2</sub>); reduction all the way to ammonia (NH<sub>3</sub>) occurs with [[hydrogen sulfide]]. With the [[hydrazine|hydrazinium]] cation ({{chem|N|2|H|5|+}}) the product of nitrite reduction is [[hydrazoic acid]] (HN<sub>3</sub>), an unstable and explosive compound:

:HNO<sub>2</sub> + {{chem|N|2|H|5|+}} → HN<sub>3</sub> + H<sub>2</sub>O + H<sub>3</sub>O<sup>+</sup>
which can also further react with nitrite:
:HNO<sub>2</sub> + HN<sub>3</sub> → N<sub>2</sub>O + N<sub>2</sub> + H<sub>2</sub>O

This reaction is unusual in that it involves compounds with nitrogen in four different oxidation states.<ref name=p461>{{Greenwood&Earnshaw2nd|pages=461–464}}</ref>

===Analysis of nitrite===
{{See also|Nitrite test}}
Nitrite is detected and analyzed by the [[Griess test|Griess Reaction]], involving the formation of a deep red-colored [[azo dye]] upon treatment of a {{chem|NO|2|-}}-containing sample with [[sulfanilic acid]] and naphthyl-1-amine in the presence of acid.<ref>{{Cite journal |last=Ivanov |first=V. M. |date=2004-10-01 |title=The 125th Anniversary of the Griess Reagent |url=https://doi.org/10.1023/B:JANC.0000043920.77446.d7 |journal=Journal of Analytical Chemistry |language=en |volume=59 |issue=10 |pages=1002–1005 |doi=10.1023/B:JANC.0000043920.77446.d7 |s2cid=98768756 |issn=1608-3199|url-access=subscription }}</ref>

=== Coordination complexes ===
{{main|Transition metal nitrite complex}}
Nitrite is an [[ambidentate ligand]] and can form a wide variety of [[coordination complex]]es by binding to metal ions in several ways.<ref name=p461/> Two examples are the red nitrito complex [Co(NH<sub>3</sub>)<sub>5</sub>(ONO)]<sup>2+</sup> is [[metastable]], [[isomer]]izing to the yellow [[Nitropentaamminecobalt(III) chloride|nitro complex [Co(NH<sub>3</sub>)<sub>5</sub>(NO<sub>2</sub>)]<sup>2+</sup>]]. Nitrite is processed by several enzymes, all of which utilize coordination complexes.