Cyanogen is the chemical compound with the formula(CN)2. Its structure is N≡C−C≡N. The simplest stable carbon nitride, it is a colorless and highly toxicgas with a pungentodor. The molecule is a pseudohalogen. Cyanogen molecules are linear, and 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, though other isomers have been detected.[6] The name is also used for the CN radical,[7] and hence is used for compounds such as cyanogen bromide (Br−C≡N)[8] (but see also Cyano radical). When burned at increased pressure with oxygen, it is possible to get a blue tinted flame, the temperature of which is about 4,800 °C (8,670 °F) (a higher temperature is possible with ozone). It is as such regarded as the gas with the second highest temperature of burning (after dicyanoacetylene).
though oxamide is manufactured from cyanogen by hydrolysis:[9]
N≡C−C≡N + 2 H2O → H2N−C(=O)−C(=O)−NH2
Preparation
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
Or, 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.[10]
For the two less stable isomers of cyanogen, the order of the atoms differs. Isocyanogen (or cyanogen cyanide) is −C≡N+−C≡N.[12] It has been detected in the interstellar medium.[13]
Paracyanogen is a polymer of cyanogen. It can be best prepared by heating mercury(II) cyanide. It can also be prepared by heating silver cyanide, silver cyanate, cyanogen iodide or cyanuric iodide.[14] It can also be prepared by the polymerization of cyanogen at 300 to 500 °C (572 to 932 °F) in the presence of trace impurities. Paracyanogen can also be converted back to cyanogen by heating to 800 °C (1,470 °F).[9] 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 π conjugation.[15]
History
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, to create), because cyanide was first isolated by Swedish chemist Carl Wilhelm Scheele from the pigment Prussian blue.[16] It attained importance with the growth of the fertilizer industry in the late 19th century and remains an important intermediate in the production of many fertilizers. It is also used as a stabilizer in the production of nitrocellulose.
Cyanogen is commonly found in comets.[17] In 1910 a spectroscopic analysis of Halley's Comet found cyanogen in the comet's tail, which led to public fear that the Earth would be poisoned as it passed through the tail. People in New York wore gas masks, and merchants sold quack "comet pills" claimed to neutralize poisoning.[17] Because of the extremely diffuse nature of the tail, there was no effect when the planet passed through it.[18][19]
Safety
Like other cyanides, cyanogen is very toxic, as it readily undergoes reduction to cyanide, which poisons the cytochrome c oxidase complex, thus interrupting the mitochondrialelectron 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.[20] Lethal dose through inhalation typically ranges from 100 to 150 milligrams (1.5 to 2.3 grains).
Cyanogen produces the second-hottest-known natural flame (after dicyanoacetylene aka carbon subnitride) with a temperature of over 4,525 °C (8,177 °F) when it burns in oxygen.[21][22]
^"oxalonitrile (CHEBI:29308)". Chemical Entities of Biological Interest. UK: European Bioinformatics Institute. 27 October 2006. Main. Retrieved 6 June 2012.
^ abNIOSH Pocket Guide to Chemical Hazards. Department of Health and Human Services, Centers for Disease Control, National Institute for Occupational Safety & Health. September 2007. p. 82.
^Ringer, A. L.; Sherrill, C. D.; King, R. A.; Crawford, T. D. (2008). "Low-lying singlet excited states of isocyanogen". International Journal of Quantum Chemistry. 106 (6): 1137–1140. Bibcode:2008IJQC..108.1137R. doi:10.1002/qua.21586.
^Bickelhaupt, F. Matthias; Nibbering, Nico M. M.; Van Wezenbeek, Egbert M.; Baerends, Evert Jan (1992). "Central Bond in the Three CN.cntdot.dimers NC-CN, CN-CN and CN-NC: Electron Pair Bonding and Pauli Repulsion Effects". The Journal of Physical Chemistry. 96 (12): 4864–4873. doi:10.1021/j100191a027.
^Bircumshaw, L. L.; F. M. Tayler; D. H. Whiffen (1954). "Paracyanogen: its formation and properties. Part I". J. Chem. Soc.: 931–935. doi:10.1039/JR9540000931.
^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.
