cyanogen

Cyanogen
IUPAC name	Ethanedinitrile
Other names	Cyanogen Carbon nitride Dicyan Dicyanogen Nitriloacetonitrile Oxalic acid dinitrile Oxalonitrile Oxalyl cyanide
Identifiers
CAS number	460-19-5 Y
EC number	207-306-5
UN number	1026
RTECS number	GT1925000
Properties
Molecular formula	C2N2
Molar mass	52.04 g/mol
Density	0.95 g/cm3 (liquid, −21 °C)
Melting point	−28 °C
Boiling point	−21 °C
Solubility in water	450 ml/100 ml (20 °C)
Hazards
MSDS	ICSC 1390
EU Index	608-011-00-8
EU classification	Flammable (F) Toxic (T) Dangerous for the environment (N)
R-phrases	R11, R23, R50/53
S-phrases	(S1/2), S23, S45, S60, S61
NFPA 704	4 4 2
Flash point	Flammable gas
Explosive limits	6.6–42.6%
Related compounds
Related compounds	Cyanogen fluoride Cyanogen chloride Cyanogen bromide
Y (what is this?)  (verify) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

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.^[1] Certain derivatives of cyanogen are also called “cyanogen” even though they contain only one CN group. For example cyanogen bromide has the formula NCBr.^[2]

Cyanogen is the anhydride of oxamide:

H₂NC(O)C(O)NH₂ → NCCN + 2 H₂O

Preparation

Cyanogen is typically generated from cyanide compounds. One laboratory method entails thermal decomposition of mercuric cyanide:

2 Hg(CN)₂ → (CN)₂ + 2 HgCN

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.^[3]

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.^[4]

Paracyanogen

Paracyanogen is produced by polymerization of cyanogen through pyrolysis of heavy metal cyanides. ^[5]

History

Cyanogen has a long history and was most likely generated first by Carl Scheele around 1782 while he was studying hydrogen cyanide.^[6][7] The first confirmed synthesis was reported 1802, when it was used to make what is now known as cyanogen chloride. 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.

Safety

Like other inorganic cyanides, cyanogen is very toxic, as it undergoes reduction to cyanide, which binds more strongly than oxygen to 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.^[8]

Cyanogen produces the second hottest known natural flame (after carbon subnitride) with a temperature of over 4525°C (8180°F) when it burns in oxygen.^[9]

See also

• Pseudohalogen

References

1. ^ Ringer, Ashley (Jan 15), "Low-lying singlet excited states of isocyanogen", International Journal of Quantum Chemistry (Wiley) 106 (6): 1137-1140 
2. ^ Hartman W. W.; Dreger, E. E. "Cyanogen Bromide" Organic Syntheses, Collected Volume 2, p.150 (1943).http://www.orgsyn.org/orgsyn/pdfs/CV2P0150.pdf
3. ^ T. K. Brotherton, J. W. Lynn (1959). "The Synthesis And Chemistry Of Cyanogen". Chemical Reviews 59 (5): 841–883. doi:10.1021/cr50029a003. 
4. ^ A. A. Breneman (1959). "Showing the Progress and Development of Processes for the manufacture of Cyanogen and its Derivates (in: THE FIXATION OF ATMOSPHERIC NITROGEN". Journal of the American Chemical Society 11 (1): 2–28. doi:10.1021/ja02126a001. 
5. ^ <a href="http://www.thefreedictionary.com/Paracyanogen">Paracyanogen</a>
6. ^ H. Bauer (1980). "Die ersten organisch-chemischen Synthesen". Naturwissenschaften 67 (1): 1–6. doi:10.1007/BF00424496. 
7. ^ J. Gay-Lussac (1815). "?". Ann. Chim. Et phys. 96: 175. 
8. ^ Muir, GD (ed.) 1971, Hazards in the Chemical Laboratory, The Royal Institute of Chemistry, London.
9. ^ Thomas, N.; Gaydon, A. G.; Brewer, L. (March 1952), "Cyanogen Flames and the Dissociation Energy of N2", The Journal of Chemical Physics 20 (3): 369–374, doi:10.1063/1.1700426, http://scitation.aip.org/getabs/servlet/GetabsServlet?prog=normal&id=JCPSA6000020000003000369000001&idtype=cvips&gifs=yes 

External links

Wikimedia Commons has media related to: cyanogen

• National Pollutant Inventory - Cyanide compounds fact sheet
• PhysOrg.com