Nitrogen trichloride
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| Names | |
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| Other names
Trichloramine
Agene Nitrogen(III) chloride Trichloroazane Trichlorine nitride | |
| Identifiers | |
3D model (
JSmol ) |
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| ChEBI | |
| ChemSpider | |
ECHA InfoCard
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100.030.029 |
| EC Number |
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| 1840 | |
PubChem CID
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RTECS number
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| UNII | |
CompTox Dashboard (EPA)
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| Properties | |
| NCl3 | |
| Molar mass | 120.36 g·mol−1 |
| 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)
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| trigonal pyramidal | |
| 0.6 D | |
| Thermochemistry | |
Std enthalpy of (ΔfH⦵298)formation |
232 kJ/mol |
| Hazards | |
| NFPA 704 (fire diamond) | |
| 93 °C (199 °F; 366 K) | |
| Related compounds | |
Other anions
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Nitrogen trifluoride Nitrogen tribromide Nitrogen triiodide |
Other cations
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Phosphorus trichloride Arsenic trichloride |
Related chloramines
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Monochloramine Dichloramine |
Related compounds
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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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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]
Preparation and structure
The compound is prepared by treatment of
Intermediates in this conversion include monochloramine and dichloramine, NH2Cl and NHCl2, respectively.
Like ammonia, NCl3 is a
Reactions and uses
The chemistry of NCl3 has been well explored.
explodes to give 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
See also
References
- ^ "Chlorine Chemistry - Chlorine Compound of the Month: Chloramines: Understanding "Pool Smell"". American Chemistry Council. Retrieved 17 December 2019.
- ISBN 978-0-12-352651-9.
- ISBN 978-0-08-037941-8.
- .
- ISBN 978-0-471-29207-4.
- ^ "Health Hazard Evaluation Report: Investigation of Employee Symptoms at an Indoor Water Park" (PDF). NIOSH ENews. 6 (4). August 2008. HETA 2007-0163-3062.
- Annales de Chimie et de Physique. 86 (6): 37–43.
- 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 and Radiochemistry. Vol. 19. pp. 1–63. )
- Kovacic, P.; Lowery, M. K.; Field, K. W. (1970). "Chemistry of N-Bromamines and N-Chloramines". Chemical Reviews. 70 (6): 639–665. .
- 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. .
- 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. .
- 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. .
External links
