Editing Bromine (section)

==Properties==
Bromine is the third [[halogen]], being a [[Nonmetal (chemistry)|nonmetal]] in group 17 of the periodic table. Its properties are thus similar to those of [[fluorine]], [[chlorine]], and [[iodine]], and tend to be intermediate between those of chlorine and iodine, the two neighbouring halogens. Bromine has the electron configuration [Ar]4s{{sup|2}}3d{{sup|10}}4p{{sup|5}}, with the seven electrons in the fourth and outermost shell acting as its [[valence electron]]s. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell.<ref name="Greenwood800">Greenwood and Earnshaw, pp. 800–4</ref> Corresponding to [[periodic trend]]s, it is intermediate in [[electronegativity]] between chlorine and iodine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than chlorine and more reactive than iodine. It is also a weaker oxidising agent than chlorine, but a stronger one than iodine. Conversely, the [[bromide]] ion is a weaker reducing agent than iodide, but a stronger one than chloride.<ref name="Greenwood800" /> These similarities led to chlorine, bromine, and iodine together being classified as one of the original triads of [[Johann Wolfgang Döbereiner]], whose work foreshadowed the [[periodic law]] for chemical elements.<ref name="purdue">{{cite web | title = Johann Wolfgang Dobereiner| publisher = Purdue University| url = https://chemed.chem.purdue.edu/genchem/history/dobereiner.html| archive-url = https://web.archive.org/web/20141114215946/https://chemed.chem.purdue.edu/genchem/history/dobereiner.html| url-status = dead| archive-date = 2014-11-14| access-date = 2008-03-08}}</ref><ref>{{cite web | title = A Historic Overview: Mendeleev and the Periodic Table | publisher = NASA | url = https://genesismission.jpl.nasa.gov/educate/scimodule/UnderElem/UnderElem_pdf/HistOverST.pdf | access-date = 2008-03-08 | archive-date = 7 April 2021 | archive-url = https://web.archive.org/web/20210407165616/https://genesismission.jpl.nasa.gov/educate/scimodule/UnderElem/UnderElem_pdf/HistOverST.pdf | url-status = live }}</ref> It is intermediate in [[atomic radius]] between chlorine and iodine, and this leads to many of its atomic properties being similarly intermediate in value between chlorine and iodine, such as first [[ionisation energy]], [[electron affinity]], enthalpy of dissociation of the X{{sub|2}} molecule (X = Cl, Br, I), ionic radius, and X–X bond length.<ref name="Greenwood800" /> The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour.<ref name="Greenwood793">Greenwood and Earnshaw, p. 793–4</ref>

All four stable halogens experience intermolecular [[van der Waals force]]s of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of bromine are intermediate between those of chlorine and iodine. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of bromine are again intermediate between those of chlorine and iodine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure.<ref name="Greenwood800" /> The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid that freezes at −7.2&nbsp;°C and boils at 58.8&nbsp;°C. (Iodine is a shiny black solid.) This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group.<ref name="Greenwood800" /> Specifically, the colour of a halogen, such as bromine, results from the [[atomic electron transition|electron transition]] between the [[HOMO/LUMO|highest occupied]] antibonding ''π{{sub|g}}'' molecular orbital and the lowest vacant antibonding ''σ{{sub|u}}'' molecular orbital.<ref name="Greenwood804">Greenwood and Earnshaw, pp. 804–9</ref> The colour fades at low temperatures so that solid bromine at −195&nbsp;°C is pale yellow.<ref name="Greenwood800" />

Liquid bromine is infrared-transparent.<ref>{{cite web|url=https://labphoto.tumblr.com/post/163685893555/bromination-using-elemental-bromine-did-you-know|title=Bromination using elemental bromine....|date=1 Aug 2017|first=Kristof|last=Heged&uuml;s|publisher=[[Tumblr]]|access-date=12 January 2025|archive-url=https://web.archive.org/web/20171210055012/https://labphoto.tumblr.com/post/163685893555/bromination-using-elemental-bromine-did-you-know|archive-date=10 December 2017|website=Pictures from an Organic Chemistry Laboratory|url-status=live}}</ref>  

Like solid chlorine and iodine, solid bromine crystallises in the [[orthorhombic crystal system]], in a layered arrangement of Br{{sub|2}} molecules. The Br–Br distance is 227&nbsp;pm (close to the gaseous Br–Br distance of 228&nbsp;pm) and the Br···Br distance between molecules is 331&nbsp;pm within a layer and 399&nbsp;pm between layers (compare the van der Waals radius of bromine, 195&nbsp;pm). This structure means that bromine is a very poor conductor of electricity, with a conductivity of around 5&nbsp;×&nbsp;10{{sup|−13}}&nbsp;Ω{{sup|−1}}&nbsp;cm{{sup|−1}} just below the melting point, although this is higher than the essentially undetectable conductivity of chlorine.<ref name="Greenwood800" />

At a pressure of 55&nbsp;[[GPa]] (roughly 540,000 times atmospheric pressure) bromine undergoes an insulator-to-metal transition. At 75&nbsp;GPa it changes to a face-centered orthorhombic structure. At 100&nbsp;GPa it changes to a body centered orthorhombic monatomic form.<ref>{{Cite journal|author = Duan, Defang|title = Ab initio studies of solid bromine under high pressure|journal = Physical Review B|volume = 76|date = 2007-09-26|doi=10.1103/PhysRevB.76.104113|page = 104113|bibcode = 2007PhRvB..76j4113D|issue = 10 |display-authors=etal}}</ref>

===Isotopes===
{{main|Isotopes of bromine}}
Bromine has two stable [[isotope]]s, {{sup|79}}Br and {{sup|81}}Br. These are its only two natural isotopes, with {{sup|79}}Br making up 51% of natural bromine and {{sup|81}}Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used for [[nuclear magnetic resonance]], although {{sup|81}}Br is more favourable. The relatively 1:1 distribution of the two isotopes in nature is helpful in identification of bromine containing compounds using mass spectroscopy. Other bromine isotopes are all radioactive, with [[half-life|half-lives]] too short to occur in nature. Of these, the most important are {{sup|80}}Br (''t''{{sub|1/2}} = 17.7&nbsp;min), {{sup|80m}}Br (''t''{{sub|1/2}} = 4.421&nbsp;h), and {{sup|82}}Br (''t''{{sub|1/2}} = 35.28&nbsp;h), which may be produced from the [[neutron activation]] of natural bromine.<ref name="Greenwood800" /> The most stable bromine radioisotope is {{sup|77}}Br (''t''{{sub|1/2}} = 57.04&nbsp;h). The primary decay mode of isotopes lighter than {{sup|79}}Br is [[electron capture]] to isotopes of [[selenium]]; that of isotopes heavier than {{sup|81}}Br is [[beta decay]] to isotopes of [[krypton]]; and {{sup|80}}Br may decay by either mode to stable {{sup|80}}Se or {{sup|80}}Kr. Br isotopes from <sup>87</sup>Br and heavier undergo beta decay with neutron emission and are of practical importance because they are fission products.<ref name="NUBASE">{{NUBASE 2003}}</ref>