Editing Nitrogen trichloride

{{Short description|Chemical compound}}
{{chembox
| Verifiedfields = changed
| Watchedfields = changed
| verifiedrevid = 477001603
| ImageFile1 = NCl3 dimensions.svg
| ImageName1 = Structural formula of nitrogen trichloride
| ImageFile2 = Nitrogen-trichloride-3D-vdW.png
| ImageName2 = Space-filling model of nitrogen trichloride
| ImageCaption2 = {{legend|blue|[[Nitrogen]], N}}{{legend|lime|[[Chlorine]], Cl}}
| ImageFile3 = Nitrogen trichloride.JPG
| ImageSize3 = 
| ImageName3 = Nitrogen trichloride
| OtherNames = Trichloramine<br />Agene<br />Nitrogen(III) chloride<br />Trichloroazane<br />Trichlorine nitride
|Section1={{Chembox Identifiers
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 55361
| InChI = 1/Cl3N/c1-4(2)3
| InChIKey = QEHKBHWEUPXBCW-UHFFFAOYAZ
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 37382
| SMILES = ClN(Cl)Cl
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/Cl3N/c1-4(2)3
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = QEHKBHWEUPXBCW-UHFFFAOYSA-N
| CASNo_Ref = {{cascite|correct|CAS}}
| CASNo = 10025-85-1
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = VA681HRW8W
| PubChem = 61437
| RTECS = QW974000
| EINECS = 233-045-1
| Gmelin = 1840
 }}
|Section2={{Chembox Properties
| Formula = {{chem2|NCl3}}
| N=1|Cl=3
| Appearance = yellow oily liquid
| Odor = [[chlorine]]-like
| Density = 1.653 g/mL
| Solubility = immiscible<br> slowly decomposes
| SolubleOther = soluble in [[benzene]], [[chloroform]], [[carbon tetrachloride|{{chem2|CCl4}}]], [[carbon disulfide|{{chem2|CS2}}]], [[phosphorus trichloride|{{chem2|PCl3}}]]
| MeltingPtC = −40
| BoilingPtC = 71
| Viscosity =
 }}
|Section3={{Chembox Structure
| MolShape = trigonal pyramidal
| CrystalStruct = [[orthorhombic]] (below −40&nbsp;°C)
| Dipole = 0.6 [[Debye|D]]
 }}
|Section4={{Chembox Thermochemistry
| DeltaHf = 232 kJ/mol
 }}
|Section7={{Chembox Hazards
| NFPA-H = 2
| NFPA-F = 0
| NFPA-R = 3
| NFPA-S = Ox
| AutoignitionPtC = 93
| AutoignitionPt_notes = 
 }}
|Section8={{Chembox Related
| OtherAnions = [[Nitrogen trifluoride]]<br />[[Nitrogen tribromide]]<br />[[Nitrogen triiodide]]
| OtherCations = [[Phosphorus trichloride]]<br />[[Arsenic trichloride]]
| OtherFunction_label = [[chloramines]]
| OtherFunction = [[Monochloramine]]<br />[[Dichloramine]]
| OtherCompounds = [[Nitrosyl chloride]]}}
}}
'''Nitrogen trichloride''', also known as '''trichloramine''', is the [[chemical compound]] with the [[chemical formula|formula]] {{chem2|NCl3}}. This yellow, oily, and explosive liquid is most commonly encountered as a product of [[chemical reaction]]s between [[ammonia]]-derivatives and [[chlorine]] (for example, in [[swimming pool]]s). 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]].<ref>{{cite web |date=July 2006 |title=Chloramines: Understanding "Pool Smell"  |url=https://chlorine.americanchemistry.com/Science-Center/Chlorine-Compound-of-the-Month-Library/Chloramines-Understanding-Pool-Smell/ |website=[[American Chemistry Council]] |url-status=dead |archive-url=https://web.archive.org/web/20191217170314/https://chlorine.americanchemistry.com/Science-Center/Chlorine-Compound-of-the-Month-Library/Chloramines-Understanding-Pool-Smell/ |archive-date=17 December 2019 |access-date=17 December 2019}}</ref>

==Preparation and occurrence==
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.<ref name=OS>{{cite journal |doi=10.15227/orgsyn.048.0004 |title=1-Amino-1-Methylcyclohexane |journal=Organic Syntheses |date=1968 |volume=48 |page=4|first1=Peter|last1=Kovacic|first2=Sohan S.|last2=Chaudhary }}</ref> Intermediates in this conversion include [[monochloramine]] and [[dichloramine]], {{chem2|NH2Cl}} and {{chem2|NHCl2}}, respectively.

Nitrogen trichloride, trademarked as [[Agene process|Agene]], was at one time used to bleach [[flour]],<ref>{{Cite journal | doi = 10.1002/jsfa.2740060906| title = Some effects of oxygen on the mixing of bread doughs| journal = Journal of the Science of Food and Agriculture| volume = 6| issue = 9| pages = 501–511| year = 1955| last1 = Hawthorn| first1 = J.| last2 = Todd| first2 = J. P.| bibcode = 1955JSFA....6..501H}}</ref> but this practice was banned in the United States in 1949 due to safety concerns.

==Structure and properties==
Like ammonia, {{chem2|NCl3}} is a [[pyramidal molecule]]. The N-Cl distances are 1.76&nbsp;Å, and the Cl-N-Cl angles are 107°.<ref>{{cite book |author1=Holleman, A. F. |author2=Wiberg, E. | title = Inorganic Chemistry | publisher = Academic Press | location = San Diego | year = 2001 | isbn = 978-0-12-352651-9}}</ref>

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 {{chem2|NCl3}} has been well explored.<ref>{{Greenwood&Earnshaw2nd}}</ref> It is moderately [[polar molecule|polar]] with a [[Molecular dipole moment|dipole moment]] of 0.6 D. The nitrogen center is basic but much less so than ammonia. It is [[Hydrolysis|hydrolyzed]] by hot water to release [[ammonia]] and [[hypochlorous acid]].
:{{chem2|NCl3 + 3 H2O -> NH3 + 3 HOCl}}
Concentrated samples of NCl<sub>3</sub> can explode to give [[Nitrogen|N<sub>2</sub>]] and [[Chlorine|chlorine gas]].{{cn|date=January 2025}}
:{{chem2|2 NCl3 -> N2 + 3 Cl2}}

In the presence of [[aluminium trichloride]], NCl<sub>3</sub> react with some branch hydrocarbon to produce, after a hydrolysis step, [[amine]]s.<ref name=OS/>

==Safety==
Nitrogen trichloride can irritate mucous {{nowrap|membranes{{hsp}}{{mdash}}{{hsp}}}}it is a [[lachrymatory agent]], but has never been used as such.<ref>{{cite book | author = White, G. C. | title = The Handbook of Chlorination and Alternative Disinfectants | edition = 4th | publisher = Wiley | year = 1999 | isbn = 978-0-471-29207-4 | page = 322}}</ref><ref>{{cite journal | id = HETA 2007-0163-3062 | title = Health Hazard Evaluation Report: Investigation of Employee Symptoms at an Indoor Water Park |date=August 2008 | journal = NIOSH ENews | volume = 6 | issue = 4 | url = https://www.cdc.gov/niosh/hhe/reports/pdfs/2007-0163-3062.pdf}}</ref> 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.<ref>{{cite journal | author = Thénard J. L. |author2=Berthollet C. L. |author-link2=Claude Louis Berthollet | title = Report on the work of Pierre Louis Dulong | journal = [[Annales de Chimie et de Physique]] | year = 1813 | volume = 86 | issue = 6 | pages = 37&ndash;43 |author-link=Louis Jacques Thénard}}</ref> In 1813, an {{chem2|NCl3}} explosion blinded Sir [[Humphry Davy]] temporarily, inducing him to hire [[Michael Faraday]] as a co-worker. They were both injured in another {{chem2|NCl3}} explosion shortly thereafter.<ref name="Thomas1991">{{cite book|author=Thomas, J.M.|title=Michael Faraday and The Royal Institution: The Genius of Man and Place (PBK)|url=https://books.google.com/books?id=GN70U1tTe_EC&pg=PA17|year= 1991|publisher=CRC Press|isbn=978-0-7503-0145-9|page=17}}</ref>

==See also==
*[[List of food contamination incidents]]
*[[Nitrogen tribromide]]
*[[Nitrogen triiodide]]

==References==
{{Reflist}}

==Further reading==
*{{Cite journal | author = Jander, J. | title = Recent Chemistry and Structure Investigation of Nitrogen Triiodide, Tribromide, Trichloride, and Related Compounds | journal = Advances in Inorganic Chemistry | year = 1976 | volume = 19 | pages = 1–63 | doi = 10.1016/S0065-2792(08)60070-9 | series = Advances in Inorganic Chemistry and Radiochemistry | isbn = 9780120236190}}
*{{cite journal |author1=Kovacic, P. |author2=Lowery, M. K. |author3=Field, K. W. | title = Chemistry of N-Bromamines and N-Chloramines | journal = Chemical Reviews | year = 1970 | volume = 70 | issue = 6 | pages = 639–665 | doi = 10.1021/cr60268a002}}
*{{cite journal |author1=Hartl, H. |author2=Schöner, J. |author3=Jander, J. |author4=Schulz, H. | title = Die Struktur des Festen Stickstofftrichlorids (−125&nbsp;°C) | journal = Zeitschrift für Anorganische und Allgemeine Chemie | year = 1975 | volume = 413 | issue = 1 | pages = 61–71 | doi = 10.1002/zaac.19754130108}}
*{{cite journal |author1=Cazzoli, G. |author2=Favero, P. G. |author3=Dal Borgo, A. | title = Molecular Structure, Nuclear Quadrupole Coupling Constant and Dipole Moment of Nitrogen Trichloride from Microwave Spectroscopy | journal = Journal of Molecular Spectroscopy | year = 1974 | volume = 50 | issue = 1–3 | pages = 82–89 | doi = 10.1016/0022-2852(74)90219-7 |bibcode=1974JMoSp..50...82C}}
*{{cite journal |author1=Bayersdorfer, L. |author2=Engelhardt, U. |author3=Fischer, J. |author4=Höhne, K. |author5=Jander, J. | title = Untersuchungen an Stickstoff–Chlor-Verbindungen. V. Infrarot- und RAMAN-Spektren von Stickstofftrichlorid | journal = Zeitschrift für Anorganische und Allgemeine Chemie | year = 1969 | volume = 366 | issue = 3–4 | pages = 169–179 | doi = 10.1002/zaac.19693660308}}

==External links==
{{Commons category|Nitrogen trichloride}}
* [https://www.osha.gov/chemicaldata/502 OSHA - Nitrogen trichloride]
* [https://web.archive.org/web/20130619115217/http://www.trianglealumni.org/trichloramine/Trichloramine-References.pdf Nitrogen Trichloride - Health References]

{{nitrogen compounds}}
{{Chlorine compounds}}
{{Nitrides}}
{{Chlorides}}

[[Category:Inorganic amines]]
[[Category:Nitrogen halides]]
[[Category:Inorganic chlorine compounds]]
[[Category:Inorganic nitrogen compounds]]
[[Category:Explosive chemicals]]
[[Category:Nitrogen(III) compounds]]
[[Category:Liquid explosives]]