Editing Cyanide (section)

==Reactions==
===Protonation===
Cyanide is basic. The p''K''<sub>a</sub> of hydrogen cyanide is 9.21. Thus, addition of [[acids]] stronger than hydrogen cyanide to solutions of cyanide salts releases [[hydrogen cyanide]].

===Hydrolysis===
Cyanide is unstable in water, but the reaction is slow until about 170&nbsp;°C. It undergoes [[hydrolysis]] to give [[ammonia]] and [[formate]], which are far less toxic than cyanide:<ref name=Ullmann/>
:{{chem2|CN- + 2 H2O → [[formate|HCO2-]] + [[ammonia|NH3]]}}
[[Cyanide hydrolase]] is an [[enzyme]] that [[catalyzes]] this reaction.

===Alkylation===
Because of the cyanide anion's high [[nucleophile|nucleophilicity]], cyano groups are readily introduced into organic molecules by displacement of a [[halide]] group (e.g., the [[chloride]] on [[methyl chloride]]). In general, organic cyanides are called nitriles. In organic synthesis, cyanide is a C-1 [[synthon]]; i.e., it can be used to lengthen a carbon chain by one, while retaining the ability to be [[wiktionary:functionalize|functionalized]].<ref>{{Ullmann|doi=10.1002/14356007.a17_363|title=Nitriles|year=2000|last1=Pollak|first1=Peter|last2=Romeder|first2=Gérard|last3=Hagedorn|first3=Ferdinand|last4=Gelbke|first4=Heinz-Peter|isbn=3-527-30673-0}}</ref>
:{{chem2|RX + CN- → RCN + X-}}

===Redox===
The cyanide ion is a [[reducing agent|reductant]] and is [[oxidation|oxidized]] by strong [[oxidizing agent]]s such as molecular [[chlorine]] ({{chem2|Cl2}}), [[hypochlorite]] ({{chem2|ClO-}}), and [[hydrogen peroxide]] ({{chem2|H2O2}}). These oxidizers are used to destroy cyanides in [[effluent]]s from [[gold mining]].<ref name="Young_1995">Young, C. A., & Jordan, T. S. (1995, May). Cyanide remediation: current and past technologies. In: Proceedings of the 10th Annual Conference on Hazardous Waste Research (pp. 104–129). Kansas State University: Manhattan, KS. https://engg.ksu.edu/HSRC/95Proceed/young.pdf</ref><ref name="SRK">{{Cite web |title=Cyanide Destruction {{!}} SRK Consulting |author=Dmitry Yermakov |work=srk.com |date= |access-date=2 March 2021 |url= https://www.srk.com/en/publications/cyanide-destruction |language=English}}</ref><ref name="Botz">Botz Michael M. Overview of cyanide treatment methods. Elbow Creek Engineering, Inc. http://www.botz.com/MEMCyanideTreatment.pdf</ref>

===Metal complexation===
The cyanide anion reacts with [[transition metals]] to form [[Cyanometalate|M-CN bonds]]. This reaction is the basis of cyanide's toxicity.<ref>Sharpe, A. G. The Chemistry of Cyano Complexes of the Transition Metals; Academic Press: London, 1976{{page needed|date=July 2015}}</ref> The high affinities of metals for this [[anion]] can be attributed to its negative charge, compactness, and ability to engage in π-bonding.

Among the most important cyanide coordination compounds are the [[potassium ferrocyanide]] and the pigment [[Prussian blue]], which are both essentially nontoxic due to the tight binding of the cyanides to a central iron atom.<ref name=Holl>{{cite book |author1=Holleman, A. F. |author2=Wiberg, E. | title = Inorganic Chemistry | publisher = Academic Press | location = San Diego | year = 2001 | isbn = 978-0-12-352651-9}}</ref>
Prussian blue was first accidentally made around 1706, by heating substances containing iron and carbon and nitrogen, and other cyanides made subsequently (and named after it). Among its many uses, Prussian blue gives the blue color to [[blueprints]], [[bluing (fabric)|bluing]], and [[cyanotype]]s.