Editing Azide (section)

== Reactions ==
Azide salts can decompose with release of [[nitrogen gas]]. The decomposition temperatures of the [[alkali metal]] azides are: [[Sodium azide|{{chem2|NaN3}}]] (275&nbsp;°C), [[Potassium azide|{{chem2|KN3}}]] (355&nbsp;°C), [[Rubidium azide|{{chem2|RbN3}}]] (395&nbsp;°C), and [[Caesium azide|{{chem2|CsN3}}]] (390&nbsp;°C). This method is used to produce ultrapure alkali metals:<ref>{{cite book |first=E. |last=Dönges |chapter=Alkali Metals |title=Handbook of Preparative Inorganic Chemistry |edition=2nd |editor-first=G. |editor-last=Brauer |publisher=Academic Press |year=1963 |location=NY |volume=1 |page=475}}</ref>
:{{chem2|2 MN3}} {{overset|heat|→}} {{chem2|2 M + 3 N2}}

[[Protonation]] of azide salts gives toxic [[hydrazoic acid]] in the presence of [[strong acid]]s:

:{{chem2|H+ + N3− → HN3}}

Azide as a [[ligand]] forms numerous [[transition metal azide complex]]es. Some such compounds are [[Explosive|shock sensitive]].

Many [[inorganic]] [[covalent]] azides (e.g., [[fluorine azide]], [[chlorine azide]], [[bromine azide]], [[iodine azide]], [[silicon tetraazide]]) have been described.<ref>{{cite journal |author1=I. C. Tornieporth-Oetting |author2=T. M. Klapötke |name-list-style=amp | title = Covalent Inorganic Azides | journal = [[Angewandte Chemie International Edition in English]] | volume = 34 | issue = 5 | pages = 511–520 | doi = 10.1002/anie.199505111 | year = 1995}}</ref>

The azide anion behaves as a [[nucleophile]]; it undergoes [[nucleophilic substitution]] for both [[aliphatic]] and [[aromatic]] systems. It reacts with [[epoxide]]s, causing a ring-opening; it undergoes [[Michael-like]] [[conjugate addition]] to 1,4-[[Saturated and unsaturated compounds|unsaturated]] [[carbonyl compounds]].<ref name = braese/>

Azides can be used as precursors of the [[metal nitrido complex]]es by being induced to release [[nitrogen|{{chem2|N2}}]], generating a [[metal complex]] in unusual [[oxidation states]] (see ''[[high-valent iron]]'').

=== Redox behaviour and trend to disproportionation ===
{{Main articles|Disproportionation|Comproportionation|Frost diagram}}
[[File:Nitrogen frost diagramm.png|thumb|300px|[[Frost diagram]] for nitrogen species at [[pH]] = 0]]
Azides have an ambivalent [[redox]] behavior: they are both [[Redox|oxidizing]] and [[Redox|reducing]], as they are easily subject to [[disproportionation]], as illustrated by the [[Frost diagram]] of nitrogen. This diagram shows the significant energetic instability of the [[hydrazoic acid]] {{Chem2|HN3}} (or the azide ion) surrounded by two much more stable species, the [[ammonium]] [[ion]] {{Chem2|NH4+}} on the left and the molecular [[nitrogen]] {{Chem2|N2}} on the right. As seen on the Frost diagram the disproportionation reaction lowers ∆G, the [[Gibbs free energy]] of the system {{Nowrap|(−∆G/F {{=}} zE}}, where F is the [[Faraday constant]], z the number of [[electron]]s exchanged in the redox reaction, and E the [[standard electrode potential]]). By minimizing the energy in the system, the disproportionation reaction increases its [[Thermodynamics|thermodynamical]] stability.

=== Destruction by oxidation by nitrite ===
Azides decompose with nitrite compounds such as [[sodium nitrite]]. Each elementary [[redox]] reaction is also a [[comproportionation]] reaction because two different N-species ({{Chem2|N3- and NO2-}}) converge to a same one (respectively {{Chem2|N2, N2O and NO}}) and is favored when the solution is acidified. This is a method of destroying residual azides, prior to disposal.<ref>{{cite book | title = Prudent practices in the laboratory: handling and disposal of chemicals | year = 1995 | publisher = [[National Academy Press]] | location = Washington, D.C. | isbn = 0-309-05229-7 | url = http://books.nap.edu/openbook.php?record_id=4911&page=165 | author = Committee on Prudent Practices for Handling, Storage, and Disposal of Chemicals in Laboratories, Board on Chemical Sciences and Technology, Commission on Physical Sciences, Mathematics, and Applications, National Research Council}}</ref> In the process, nitrogen gas ({{chem2|N2}}) and nitrogen oxides ({{chem2|N2O}} and NO) are formed:

:{{chem2|3 N3- + NO2- + 2 H2O → 5 N2 + 4 OH-}}

:{{chem2|2 N3- + 4 NO2- + 3 H2O → 5 N2O + 6 OH-}}

:{{chem2|N3- + 7 NO2- + 4 H2O → 10 NO + 8 OH-}}

Azide (<em>−⅓</em>) (the [[Reducing agent|reductant]], [[electron donor]]) is [[Redox|oxidized]] in {{Chem2|N2}} (0), [[nitrous oxide]] ({{Chem2|N2O}}) (+1), or [[nitric oxide]] (NO) (+2) while [[nitrite]] (+3) (the [[Oxidizing agent|oxidant]], [[electron acceptor]]) is simultaneously [[Redox|reduced]] to the same corresponding species in each elementary redox reaction considered here above. The respective stability of the reaction products of these three [[comproportionation]] redox reactions is in the following order: {{Chem2|N2 > N2O > NO}}, as can be verified in the Frost diagram for nitrogen.