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.[1]

Nitrogen trichloride
Structural formula of nitrogen trichloride
Space-filling model of nitrogen trichloride
Names
Other names
Trichloramine
Agene
Nitrogen(III) chloride
Trichloroazane
Trichlorine nitride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.029
EC Number
  • 233-045-1
1840
RTECS number
  • QW974000
UNII
Properties
NCl3
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)
immiscible
slowly decomposes
Solubility soluble in benzene, chloroform, CCl4, CS2, PCl3
Structure
orthorhombic (below −40 °C)
trigonal pyramidal
0.6 D
Thermochemistry
Std enthalpy of
formation (ΔfH⦵298)
 232 kJ/mol
Hazards
not listed
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
 Monochloramine
Dichloramine
Related compounds
 Nitrosyl chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

Preparation and structure

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 uses

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.

Safety

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]

gollark: Retractable toast heating pins?
gollark: I still hold that a chemical hypertoaster could be more effective.
gollark: It saddens me when I see something with a screen which lacks a browser I can open osmarks.net on.
gollark: And check the weather.
gollark: So you can monitor live toast status.

See also

References
  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 reading
  • 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. 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.
Salts and covalent derivatives of the nitride ion
NH3
N2H4		 He(N2)11
Li3N Be3N2 BN β-C3N4
g-C3N4
CxNy		 N2 NxOy NF3 Ne
Na3N Mg3N2 AlN Si3N4 PN
P3N5		 SxNy
SN
S4N4		 NCl3 Ar
K3N Ca3N2 ScN TiN VN CrN
Cr2N		 MnxNy FexNy CoN Ni3N CuN Zn3N2 GaN Ge3N4 As Se NBr3 Kr
Rb3N Sr3N2 YN ZrN NbN β-Mo2N Tc Ru Rh PdN Ag3N CdN InN Sn Sb Te NI3 Xe
Cs3N Ba3N2   Hf3N4 TaN WN Re Os Ir Pt Au Hg3N2 TlN Pb BiN Po At Rn
Fr3N Ra3N2   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
La CeN Pr Nd Pm Sm Eu GdN Tb Dy Ho Er Tm Yb Lu
Ac Th Pa UN Np Pu Am Cm Bk Cf Es Fm Md No Lr
This article is issued from Wikipedia. The text is licensed under Creative Commons - Attribution - Sharealike. Additional terms may apply for the media files.