Nitrogen trichloride

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Nitrogen trichloride, also known as trichloramine, is the chemical compound with the formula NCl₃. This yellow, oily, pungent-smelling and explosive liquid is most commonly encountered as a byproduct of chemical reactions between ammonia-derivatives and chlorine (for example, in swimming pools).

Video Nitrogen trichloride

Preparation and structure

The compound is prepared by treatment of ammonium salts, such as ammonium nitrate with chlorine.

Intermediates in this conversion include chloramine and dichloramine, NH₂Cl and NHCl₂, respectively.

Like ammonia, NCl₃ is a pyramidal molecule. The N-Cl distances are 1.76 Å, and the Cl-N-Cl angles are 107°.

Nitrogen trichloride can form in small amounts when public water supplies are disinfected with monochloramine, and in swimming pools by disinfecting chlorine reacting with urea in urine and sweat from bathers. Nitrogen trichloride, trademarked as Agene, was used to artificially bleach and age flour, but was banned in 1949: In humans Agene was found to cause severe and widespread neurological disorders leading to its banning in 1947. Dogs that ate bread made from treated flour suffered epileptic-like fits; the toxic agent was methionine sulfoximine.

Maps Nitrogen trichloride

Reactions

The chemistry of NCl₃ has been well explored. It is moderately polar with a dipole moment of 0.6 D. The nitrogen center is basic but much less so than ammonia. It is hydrolyzed by hot water to release ammonia and hypochlorous acid.

NCl₃ + 3 H₂O -> NH₃ + 3 HOCl

NCl₃ explodes to give N₂ and chlorine gas. This reaction is inhibited for dilute gases.

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Safety

Nitrogen trichloride can irritate mucous membranes--it is a lachrymatory agent, but has never been used as such. The pure substance (rarely encountered) is a dangerous explosive, being sensitive to light, heat, even moderate shock, and organic compounds. Pierre Louis Dulong first prepared it in 1812, and lost two fingers and an eye in two explosions. In 1813, an NCl₃ explosion blinded Sir Humphry Davy temporarily, inducing him to hire Michael Faraday as a co-worker. They were both injured in another NCl₃ explosion shortly thereafter.

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See also

• List of food contamination incidents

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References

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Further reading

• Jander, J. (1976). "Recent Chemistry and Structure Investigation of Nitrogen Triiodide, Tribromide, Trichloride, and Related Compounds". Advances in Inorganic Chemistry. 19: 1-63. doi:10.1016/S0065-2792(08)60070-9. 
• Kovacic, P.; Lowery, M. K.; Field, K. W. (1970). "Chemistry of N-Bromamines and N-Chloramines". Chemical Reviews. 70 (6): 639-665. doi:10.1021/cr60268a002. 
• Hartl, H.; Schöner, J.; Jander, J.; Schulz, H. (1975). "Die Struktur des Festen Stickstofftrichlorids (-125 °C)". Zeitschrift für Anorganische und Allgemeine Chemie. 413 (1): 61-71. doi:10.1002/zaac.19754130108. 
• Cazzoli, G.; Favero, P. G.; Dal Borgo, A. (1974). "Molecular Structure, Nuclear Quadrupole Coupling Constant and Dipole Moment of Nitrogen Trichloride from Microwave Spectroscopy". Journal of Molecular Spectroscopy. 50 (1-3): 82-89. doi:10.1016/0022-2852(74)90219-7. 
• Bayersdorfer, L.; Engelhardt, U.; Fischer, J.; Höhne, K.; Jander, J. (1969). "Untersuchungen an Stickstoff-Chlor-Verbindungen. V. Infrarot- und RAMAN-Spektren von Stickstofftrichlorid". Zeitschrift für Anorganische und Allgemeine Chemie. 366 (3-4): 169-179. doi:10.1002/zaac.19693660308. 

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External links

• OSHA - Nitrogen trichloride
• Nitrogen Trichloride - Health References

Source of article : Wikipedia