Nitrogen trichloride


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). 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]

Nitrogen trichloride
Structural formula of nitrogen trichloride
Space-filling model of nitrogen trichloride
Nitrogen trichloride
Other names
Nitrogen(III) chloride
Trichlorine nitride
  • 10025-85-1 checkY
3D model (JSmol)
  • Interactive image
  • CHEBI:37382 checkY
  • 55361 checkY
ECHA InfoCard 100.030.029 Edit this at Wikidata
EC Number
  • 233-045-1
  • 61437
RTECS number
  • QW974000
  • VA681HRW8W checkY
  • DTXSID4074938 Edit this at Wikidata
  • InChI=1S/Cl3N/c1-4(2)3 checkY
  • InChI=1/Cl3N/c1-4(2)3
  • ClN(Cl)Cl
Appearance yellow oily liquid
Odor chlorine-like
Density 1.653 g/mL
Melting point −40 °C (−40 °F; 233 K)
Boiling point 71 °C (160 °F; 344 K)
slowly decomposes
Solubility soluble in benzene, chloroform, CCl4, CS2, PCl3
orthorhombic (below −40 °C)
trigonal pyramidal
0.6 D
232 kJ/mol
NFPA 704 (fire diamond)
93 °C (199 °F; 366 K)
Related compounds
Other anions
Nitrogen trifluoride
Nitrogen tribromide
Nitrogen triiodide
Other cations
Phosphorus trichloride
Arsenic trichloride
Related chloramines
Related compounds
Nitrosyl chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Preparation and structureEdit

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

Intermediates in this conversion include monochloramine 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°.[2]

Reactions and usesEdit

The chemistry of NCl3 has been well explored.[3] 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.

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 at one time used to bleach flour,[4] but this practice was banned in the United States in 1949 due to safety concerns.


Nitrogen trichloride can irritate mucous membranes—it is a lachrymatory agent, but has never been used as such.[5][6] 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.[7] 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.[8]

See alsoEdit


  1. ^ "Chlorine Chemistry - Chlorine Compound of the Month: Chloramines: Understanding "Pool Smell"". American Chemistry Council. Retrieved 17 December 2019.
  2. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  4. ^ Hawthorn, J.; Todd, J. P. (1955). "Some effects of oxygen on the mixing of bread doughs". Journal of the Science of Food and Agriculture. 6 (9): 501–511. doi:10.1002/jsfa.2740060906.
  5. ^ White, G. C. (1999). The Handbook of Chlorination and Alternative Disinfectants (4th ed.). Wiley. p. 322. ISBN 978-0-471-29207-4.
  6. ^ "Health Hazard Evaluation Report: Investigation of Employee Symptoms at an Indoor Water Park" (PDF). NIOSH ENews. 6 (4). August 2008. HETA 2007-0163-3062.
  7. ^ Thénard J. L.; Berthollet C. L. (1813). "Report on the work of Pierre Louis Dulong". Annales de Chimie et de Physique. 86 (6): 37–43.
  8. ^ Thomas, J.M. (1991). Michael Faraday and The Royal Institution: The Genius of Man and Place (PBK). CRC Press. p. 17. ISBN 978-0-7503-0145-9.

Further readingEdit

  • 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. Vol. 19. pp. 1–63. doi:10.1016/S0065-2792(08)60070-9. ISBN 9780120236190.
  • 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.

External linksEdit

  • OSHA - Nitrogen trichloride
  • Nitrogen Trichloride - Health References