|Jmol-3D images||Image 1|
|Molar mass||52.03 g mol−1|
|Density||950 mg mL−1 (at −21 °C)|
|Melting point||−28 °C (−18 °F; 245 K)|
|Boiling point||−21.1 °C; −6.1 °F; 252.0 K|
|Solubility in water||45 g/100 mL (at 20 °C)|
|Solubility||soluble in ethanol, ethyl ether|
|kH||1.9 μmol Pa−1 kg−1|
|Refractive index (nD)||1.327 (18 °C)|
|241.57 J K−1 mol−1|
|Std enthalpy of
|309.07 kJ mol−1|
|Std enthalpy of
|−1.0978–−1.0942 MJ mol−1|
|EU classification||F+ T N|
|R-phrases||R12, R23, R50/53|
|S-phrases||(S1/2), S16, S33, S45, S63|
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
Cyanogen is the chemical compound with the formula (CN)2. It is a colorless, toxic gas with a pungent odor. The molecule is a pseudohalogen. Cyanogen molecules consist of two CN groups – analogous to diatomic halogen molecules, such as Cl2, but far less oxidizing. The two cyano groups are bonded together at their carbon atoms: N≡C−C≡N, although other isomers have been detected. Certain derivatives of cyanogen are also called "cyanogen" even though they contain only one CN group. For example cyanogen bromide has the formula NCBr.
- H2NC(O)C(O)NH2 → NCCN + 2 H2O
although oxamide is manufactured from cyanogen by hydrolysis:
- NCCN + 2 H2O → H2NC(O)C(O)NH2
Cyanogen is typically generated from cyanide compounds. One laboratory method entails thermal decomposition of mercuric cyanide:
- 2 Hg(CN)2 → (CN)2 + Hg2(CN)2
Alternatively, one can combine solutions of copper(II) salts (such as copper(II) sulfate) with cyanides, an unstable copper(II) cyanide is formed which rapidly decomposes into copper(I) cyanide and cyanogen.
- 2 CuSO4 + 4 KCN → (CN)2 + 2 CuCN + 2 K2SO4
Industrially, it is created by the oxidation of hydrogen cyanide, usually using chlorine over an activated silicon dioxide catalyst or nitrogen dioxide over a copper salt. It is also formed when nitrogen and acetylene are reacted by an electrical spark or discharge.
Paracyanogen can be best prepared by heating mercuric cyanide. It can also be prepared by heating silver cyanide, silver cyanate, cyanogen iodide or cyanuric iodide. It can also be prepared by the polymerization of cyanogen at 300-500 °C in the presence of trace impurities. Paracyanogen can also be converted back to cyanogen by heating to 800 °C. Based on experimental evidence, the structure of this polymeric material is thought to be rather irregular, with most of the carbon atoms being of sp2 type and localized domains of pi conjugation.
Cyanogen was first synthesized in 1815 by Joseph Louis Gay-Lussac, who determined its empirical formula and named it. Gay-Lussac coined the word "cyanogène" from the Greek words κυανός (kyanos, blue) and γεννάω (gennao, I create), because cyanide was first isolated by the Swedish chemist Carl Wilhelm Scheele from the pigment "Prussian blue". By the 1850s, cyanogen soap was used by photographers to remove silver stains from their hands.  It attained importance with the growth of the fertilizer industry in the late nineteenth century and is still an important intermediate in the production of many fertilizers. It is also used as a stabilizer in the production of nitrocellulose.
In 1910 spectroscopic analysis of Halley's Comet found cyanogen in its tail. This led to public fear that the Earth would be poisoned. Because of the extremely diffuse nature of a comet's tail, there was no effect when the planet passed through it.
Like other inorganic cyanides, cyanogen is very toxic, as it readily undergoes reduction to cyanide, which poisons the cytochrome c oxidase complex, thus interrupting the mitochondrial electron transfer chain. Cyanogen gas is an irritant to the eyes and respiratory system. Inhalation can lead to headache, dizziness, rapid pulse, nausea, vomiting, loss of consciousness, convulsions, and death, depending on exposure.
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- Hartman, W. W.; Dreger, E. E. (1931), "Cyanogen Bromide", Org. Synth. 11: 30; Coll. Vol. 2: 150
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- Breneman, A. A. (1889). Showing the Progress and Development of Processes for the manufacture of Cyanogen and its Derivates. "The Fixation of Atmospheric Nitrogen". Journal of the American Chemical Society 11 (1): 2–27. doi:10.1021/ja02126a001.
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- Maya, Leon (1993). "Paracyanogen Reexamined". Polymer Science: Part A 31: 2595–2600.
- Gay-Lussac, J. L. (1815). "Recherches sur l'acide prussique". Annales de Chimie 95: 136–231.
- Crookes, William, ed. (1859). Photographic News: A Weekly Record of the Process of the Photography. pp. 11–11.
- Comet's Poisonous Tail
- Halley's Comet 100 years Ago
- Muir, G. D., ed. (1971). Hazards in the Chemical Laboratory. London: The Royal Institute of Chemistry.
- Thomas, N.; Gaydon, A. G.; Brewer, L. (1952). "Cyanogen Flames and the Dissociation Energy of N2". The Journal of Chemical Physics 20 (3): 369–374. Bibcode:1952JChPh..20..369T. doi:10.1063/1.1700426.
- J. B. Conway , R. H. Wilson Jr., A. V. Grosse (1953). "THE TEMPERATURE OF THE CYANOGEN-OXYGEN FLAME". Journal of the American Chemical Society 75 (2): 499. doi:10.1021/ja01098a517.
|Wikimedia Commons has media related to cyanogen.|
- National Pollutant Inventory - Cyanide compounds fact sheet
- CDC - NIOSH Pocket Guide to Chemical Hazards