Zinc cyanide

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Zinc cyanide
Zinc cyanide
3D model (JSmol)
ECHA InfoCard 100.008.331
RTECS number ZH1575000
Molar mass 117.444 g/mol
Appearance white powder
Density 1.852 g/cm3, solid
Melting point 800 °C (1,470 °F; 1,070 K) (decomposes)
0.00005 g/100 mL (20 °C)
Solubility attacked by alkalies, KCN, ammonia
−46.0·10−6 cm3/mol
not listed
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
100 mg/kg, rat (intraperitoneal)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Zinc cyanide is the inorganic compound with the formula Zn(CN)2. It is a white solid that is used mainly for electroplating zinc but also has more specialized applications for the synthesis of organic compounds.


In Zn(CN)2, zinc adopts the tetrahedral coordination environment, all linked by bridging cyanide ligands. The structure consists of two "interpenetrating" structures (blue and red in the picture above). Such motifs are sometimes called "expanded diamondoid" structures. Some forms of SiO2 adopt a similar structure, wherein the tetrahedral Si centres are linked by oxides. The cyanide group shows head to tail disorder with any zinc atom having between one and four carbon neighbours, and the remaining being nitrogen atoms.[2] It shows one of the largest negative coefficients of thermal expansion (exceeding the previous record holder, zirconium tungstate).

Chemical properties[edit]

Typical for an inorganic polymer, Zn(CN)2 is insoluble in most solvents. The solid dissolves in, or more precisely, is degraded by, aqueous solutions of basic ligands such as hydroxide, ammonia, and additional cyanide to give anionic complexes.


Zn(CN)2 is fairly easy to make by combining aqueous solutions of cyanide and zinc ions, for example via the double replacement reaction between KCN and ZnSO4:[3]

ZnSO4 + 2 KCN → Zn(CN)2 + K2SO4

For commercial applications, some effort is made to avoid halide impurities by using acetate salts of zinc:[3][4]

Zn(CH3COO)2 + HCN → Zn(CN)2 + 2 CH3COOH

Zinc cyanide is also produced as a byproduct of certain gold extraction methods. Procedures to isolate gold from aqueous gold cyanide sometimes call for the addition of zinc:

2 [Au(CN)2] + Zn → 2 Au + Zn(CN)2 + 2 CN



The main application of Zn(CN)2 is for electroplating of zinc from aqueous solutions containing additional cyanide.[4]

Organic synthesis[edit]

Zn(CN)2 is used to introduce the formyl group in to aromatic compounds in the Gatterman reaction where it serves a convenient, safer, and non-gaseous alternative to HCN.[5] Because the reaction uses HCl, Zn(CN)2 also supplies the reaction with ZnCl2 in-situ, a Lewis acid catalyst. Examples of Zn(CN)2 being used in this way include the synthesis of 2-Hydroxy-1-nafthaldehyde and Mesitaldehyde.[6]

Zn(CN)2 is also employed as a catalyst for the cyanosilylation of aldehydes and ketones.[7]


  1. ^ http://cameochemicals.noaa.gov/chemical/4808
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  3. ^ a b Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry Vol. 2, 2nd Ed. Newyork: Academic Press. p. 1087. ISBN 9780323161299. 
  4. ^ a b Ernst Gail, Stephen Gos, Rupprecht Kulzer, Jürgen Lorösch, Andreas Rubo and Manfred Sauer "Cyano Compounds, Inorganic" Ullmann's Encyclopedia of Industrial Chemistry Wiley-VCH, Weinheim, 2004. doi:10.1002/14356007.a08_159.pub2
  5. ^ Adams, Roger (1957). Organic Reactions, Volume 9. New York: John Wiley & Sons, Inc. pp. 53–54. ISBN 9780471007265. Retrieved 18 July 2014. 
  6. ^ Adams R., Levine I. (1923). "Simplification of the Gattermann Synthesis of Hydroxy Aldehydes". J. Am. Chem. Soc. 45 (10): 2373–77. doi:10.1021/ja01663a020.  Fuson R. C., Horning E. C., Rowland S. P., Ward M. L. (1955). "Mesitaldehyde". Organic Syntheses. doi:10.15227/orgsyn.023.0057. ; Collective Volume, 3, p. 549 
  7. ^ Rasmussen J. K., Heilmann S. M. (1990). "In situ Cyanosilylation of Carbonyl Compounds: O-Trimethylsilyl-4-Methoxymandelonitrile". Organic Syntheses. doi:10.15227/orgsyn.062.0196. ; Collective Volume, 7, p. 521