Nitrogen trichloride, also known as trichloramine, is the chemical compound with the formulaNCl3. This yellow, oily, and explosive liquid is most commonly encountered as a product of chemical reactions between ammonia-derivatives and chlorine (for example, in swimming pools). Alongside monochloramine and dichloramine, trichloramine is responsible for the distinctive 'chlorine smell' associated with swimming pools, where the compound is readily formed as a product from hypochlorous acid reacting with ammonia and other nitrogenous substances in the water, such as urea from urine.[1]
The compound is generated by treatment of ammonium chloride with calcium hypochlorite. When prepared in an aqueous-dichloromethane mixture, the trichloramine is extracted into the nonaqueous phase.[2] Intermediates in hloramine]] and dichloramine, NH2Cl and NHCl2, respectively.
Nitrogen trichloride, trademarked as Agene, was at one time used to bleach flour,[3] but this practice was banned in the United States in 1949 due to safety concerns.
Structure and properties
Like ammonia, NCl3 is a pyramidal molecule. The N-Cl distances are 1.76Å, and the Cl-N-Cl angles are 107°.[4]
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.
Reactions and uses
The chemistry of NCl3 has been well explored.[5] 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.
In the presence of aluminium trichloride, NCl3 react with some branched hydrocarbons to produce, after a hydrolysis step, amines.[2]
Safety
Nitrogen trichloride can irritate mucous membranes—it is a lachrymatory agent, but has never been used as such.[6][7] The compound (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 several fingers and an eye in two explosions.[8] 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.[9]
Jander, J. (1976). "Recent Chemistry and Structure Investigation of Nitrogen Triiodide, Tribromide, Trichloride, and Related Compounds". Advances in Inorganic Chemistry. Advances in Inorganic Chemistry and Radiochemistry. 19: 1–63. doi:10.1016/S0065-2792(08)60070-9. ISBN9780120236190.
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. Bibcode:1974JMoSp..50...82C. 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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