{{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, {{chem2|CCl4}}, {{chem2|CS2}}, {{chem2|PCl3}} | MeltingPtC = −40 | BoilingPtC = 71 | Viscosity = }} |Section3={{Chembox Structure | MolShape = trigonal pyramidal | CrystalStruct = orthorhombic (below −40 °C) | Dipole = 0.6 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 formula {{chem2|NCl3}}. 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.<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, 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 1.42 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 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.
<chem>NCl3 + 3H2O -> NH3 + 3HCl + 3/2O2</chem>
Concentrated samples of NCl<sub>3</sub> can explode to give N<sub>2</sub> and chlorine gas.{{cn|date=January 2025}}
<chem>2NCl3 -> N2 + 3Cl2</chem>
In the presence of aluminium trichloride, NCl<sub>3</sub> reacts with some branched hydrocarbons to produce, after a hydrolysis step, amines.<ref name=OS/><ref>{{Cite journal |last=Kovacic |first=Peter |last2=Goralski |first2=Christian T. |last3=Hiller |first3=John J. |last4=Levisky |first4=Joseph A. |last5=Lange |first5=Richard M. |date=March 20, 1965 |title=Amination of Toluene with Trichloramine-Lewis Acid Catalyst |url=https://pubs.acs.org/doi/abs/10.1021/ja01084a021 |journal=Journal of the American Chemical Society |language=en |volume=87 |issue=6 |pages=1262–1266 |doi=10.1021/ja01084a021 |issn=0002-7863|url-access=subscription }}</ref>
==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 = 37–43 |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 °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