Editing Nitrogen (section)

===Halides and oxohalides===
[[File:Nitrogen trichloride.JPG|thumb|right|[[Nitrogen trichloride]]]]
All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable; examples include NClF<sub>2</sub>, NCl<sub>2</sub>F, NBrF<sub>2</sub>, NF<sub>2</sub>H, [[fluoroamine|NFH<sub>2</sub>]], [[Dichloramine|NCl<sub>2</sub>H]], and [[Monochloramine|NClH<sub>2</sub>]].<ref name="Greenwood438">Greenwood and Earnshaw, pp. 438–42</ref>

[[Nitrogen trifluoride]] (NF<sub>3</sub>, first prepared in 1928) is a colourless and odourless gas that is thermodynamically stable, and most readily produced by the [[electrolysis]] of molten [[ammonium fluoride]] dissolved in anhydrous [[hydrogen fluoride]]. Like [[carbon tetrafluoride]], it is not at all reactive and is stable in water or dilute aqueous acids or alkalis. Only when heated does it act as a fluorinating agent, and it reacts with [[copper]], arsenic, antimony, and bismuth on contact at high temperatures to give [[tetrafluorohydrazine]] (N<sub>2</sub>F<sub>4</sub>). The cations {{chem|NF|4|+}} and {{chem|N|2|F|3|+}} are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such as [[arsenic pentafluoride]]), as is ONF<sub>3</sub>, which has aroused interest due to the short N–O distance implying partial double bonding and the highly polar and long N–F bond. Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room temperature and above to give the radical NF<sub>2</sub>•. [[Fluorine azide]] (FN<sub>3</sub>) is very explosive and thermally unstable. [[Dinitrogen difluoride]] (N<sub>2</sub>F<sub>2</sub>) exists as thermally interconvertible ''cis'' and ''trans'' isomers, and was first found as a product of the thermal decomposition of FN<sub>3</sub>.<ref name="Greenwood438" />

[[Nitrogen trichloride]] (NCl<sub>3</sub>) is a dense, volatile, and explosive liquid whose physical properties are similar to those of [[carbon tetrachloride]], although one difference is that NCl<sub>3</sub> is easily hydrolysed by water while CCl<sub>4</sub> is not. It was first synthesised in 1811 by [[Pierre Louis Dulong]], who lost three fingers and an eye to its explosive tendencies. As a dilute gas it is less dangerous and is thus used industrially to bleach and sterilise flour. [[Nitrogen tribromide]] (NBr<sub>3</sub>), first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100&nbsp;°C. [[Nitrogen triiodide]] (NI<sub>3</sub>) is still more unstable and was only prepared in 1990. Its adduct with ammonia, which was known earlier, is very shock-sensitive: it can be set off by the touch of a feather, shifting air currents, or even [[alpha particle]]s.<ref name="Greenwood438" /><ref>{{ cite journal | author = Bowden, F. P. | title = Initiation of Explosion by Neutrons, α-Particles, and Fission Products | journal = Proceedings of the Royal Society of London A | year = 1958 | volume = 246 | issue = 1245 | pages = 216–19 | doi = 10.1098/rspa.1958.0123 | bibcode = 1958RSPSA.246..216B | s2cid = 137728239 }}</ref> For this reason, small amounts of nitrogen triiodide are sometimes synthesised as a demonstration to high school chemistry students or as an act of "chemical magic".<ref>{{cite book | author1 = Ford, L. A. | author2 = Grundmeier, E. W. | title = Chemical Magic | publisher = Dover | year = 1993 | page = [https://archive.org/details/chemicalmagic00ford_0/page/76 76] | isbn = 978-0-486-67628-9 | url-access = registration | url = https://archive.org/details/chemicalmagic00ford_0/page/76 }}</ref> [[Chlorine azide]] (ClN<sub>3</sub>) and [[bromine azide]] (BrN<sub>3</sub>) are extremely sensitive and explosive.<ref>{{ cite journal |author1=Frierson, W. J. |author2=Kronrad, J. |author3=Browne, A. W. | title = Chlorine Azide, ClN<sub>3</sub>. I | journal = [[Journal of the American Chemical Society]] | year = 1943 | volume = 65 | issue = 9 | pages = 1696–1698 | doi = 10.1021/ja01249a012 |bibcode=1943JAChS..65.1696F }}</ref><ref name="solid">{{cite journal|last=Lyhs|first=Benjamin|author2=Bläser, Dieter|author3=Wölper, Christoph|author4=Schulz, Stephan|author5=Jansen, Georg|title=Solid-State Structure of Bromine Azide|journal=Angewandte Chemie International Edition|date=20 February 2012|volume=51|issue=8|pages=1970–1974|doi=10.1002/anie.201108092|pmid=22250068|url=https://duepublico2.uni-due.de/servlets/MCRFileNodeServlet/duepublico_derivate_00073133/Accepted_Manuscript_Angew_Chem_Int_Ed_2012_51_1970.pdf|access-date=25 August 2021|archive-date=25 August 2021|archive-url=https://web.archive.org/web/20210825213503/https://duepublico2.uni-due.de/servlets/MCRFileNodeServlet/duepublico_derivate_00073133/Accepted_Manuscript_Angew_Chem_Int_Ed_2012_51_1970.pdf|url-status=live}}</ref>

Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO<sub>2</sub>). The first is very reactive gases that can be made by directly halogenating nitrous oxide. [[Nitrosyl fluoride]] (NOF) is colourless and a vigorous fluorinating agent. [[Nitrosyl chloride]] (NOCl) behaves in much the same way and has often been used as an ionising solvent. [[Nitrosyl bromide]] (NOBr) is red. The reactions of the nitryl halides are mostly similar: [[nitryl fluoride]] (FNO<sub>2</sub>) and [[nitryl chloride]] (ClNO<sub>2</sub>) are likewise reactive gases and vigorous halogenating agents.<ref name="Greenwood438" />