Editing Selenium (section)

===Halogen compounds===
Selenium reacts with [[fluorine]] to form [[selenium hexafluoride]]:

{{block indent|{{chem2|Se8 + 24 F2 -> 8 SeF6}}}}

In comparison with its sulfur counterpart ([[sulfur hexafluoride]]), [[selenium hexafluoride]] (SeF<sub>6</sub>) is more reactive and is a toxic [[pulmonary]] irritant.<ref>{{cite book |last1=Proctor |first1=Nick H. |title=Proctor and Hughes' chemical hazards of the workplace |last2=Hathaway |first2=Gloria J. |publisher=Wiley-IEEE |year=2004 |isbn=978-0-471-26883-3 |editor=Hughes, James P. |edition=5th |page=625}}</ref> [[Selenium tetrafluoride]] is a laboratory-scale [[fluorinating agent]].

The only stable [[chloride]]s are [[selenium tetrachloride]] (SeCl<sub>4</sub>) and [[selenium monochloride]] (Se<sub>2</sub>Cl<sub>2</sub>), which might be better known as selenium(I) chloride and is structurally analogous to [[disulfur dichloride]]. Metastable solutions of [[selenium dichloride]] can be prepared from [[sulfuryl chloride]] and selenium (reaction of the elements generates the [[Selenium tetrachloride|tetrachloride]] instead), and constitute an important reagent in the preparation of selenium compounds (e.g. Se<sub>7</sub>). The corresponding [[bromide]]s are all known, and recapitulate the same stability and structure as the chlorides.<ref>{{cite book
|title=Handbook of chalcogen chemistry: new perspectives in sulfur, selenium and tellurium |author=Xu, Zhengtao |editor=Devillanova, Francesco A. |publisher=Royal Society of Chemistry |year=2007 |isbn=978-0-85404-366-8 |page=460}}</ref>

The [[iodide]]s of selenium are not well known, and for a long time were believed not to exist.<ref name=":1">{{Cite journal |last1=Gopal |first1=Madhuban |last2=Milne |first2=John |date=October 1992 |title=Spectroscopic evidence for selenium iodides in carbon disulfide solution: Se3I2, Se2I2, and SeI2 |url=https://pubs.acs.org/doi/abs/10.1021/ic00048a017 |journal=Inorganic Chemistry |language=en |volume=31 |issue=22 |pages=4530–4533 |doi=10.1021/ic00048a017 |issn=0020-1669}}</ref> There is limited [[Spectroscopy|spectroscopic]] evidence that the lower iodides may form in bi-elemental solutions with nonpolar solvents, such as [[carbon disulfide]]<ref>{{Cite journal |last=McCullough |first=James D. |date=December 1939 |title=Evidence for Existence of a Selenium Iodide |url=https://pubs.acs.org/doi/abs/10.1021/ja01267a052 |journal=Journal of the American Chemical Society |language=en |volume=61 |issue=12 |pages=3401–3402 |doi=10.1021/ja01267a052 |bibcode=1939JAChS..61.3401M |issn=0002-7863}}</ref> and [[carbon tetrachloride]];<ref name=":1" /> but even these appear to [[Photosensitivity|decompose under illumination]].<ref>Rao, M.&nbsp;R.&nbsp;Aswatha&nbsp;Narayana. [https://www.ias.ac.in/public/Volumes/seca/012/04/0410-0415.pdf "Selenium iodide"]. In ''Proceedings of the Indian Academy of Sciences-Section&nbsp;A'', vol.&nbsp;12, pp.&nbsp;410-415. Springer India, 1940.</ref>

Some selenium oxyhalides—[[seleninyl fluoride]] (SeOF<sub>2</sub>) and [[selenium oxychloride]] (SeOCl<sub>2</sub>)—have been used as specialty solvents.<ref name="house2008" />