Editing Bromine

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{{Infobox bromine|engvar=en-GB}}
[[File:Bróm.jpg|thumb|Flask of bromine showing liquid and vapor form above]]

'''Bromine''' is a [[chemical element]]; it has [[chemical symbol|symbol]] '''Br''' and [[atomic number]] 35. It is a volatile red-brown [[liquid]] at room temperature that evaporates readily to form a similarly coloured vapour. Its properties are intermediate between those of [[chlorine]] and [[iodine]]. Isolated independently by two chemists, [[Carl Jacob Löwig]] (in 1825) and [[Antoine Jérôme Balard]] (in 1826), its name was derived {{ety|grc|''βρῶμος'' (bromos)|stench}}, referring to its sharp and pungent smell.

Elemental bromine is very reactive and thus does not occur as a [[free element]] in nature. Instead, it can be isolated from colourless soluble crystalline mineral halide [[Ionic salt|salts]] analogous to [[table salt]], a property it shares with the other [[halogen]]s. While it is rather rare in the Earth's crust, the high solubility of the [[bromide]] ion (Br{{sup|−}}) has caused its [[Bromine cycle|accumulation in the oceans]]. Commercially the element is easily extracted from brine [[evaporation pond]]s, mostly in the [[United States]] and [[Israel]]. The mass of bromine in the oceans is about one three-hundredth that of chlorine.

At [[standard conditions for temperature and pressure]] it is a liquid; the only other element that is liquid under these conditions is [[Mercury (element)|mercury]]. At high temperatures, [[organobromine compound]]s readily dissociate to yield free bromine atoms, a process that stops [[free radical]] chemical [[chain reaction]]s. This effect makes organobromine compounds useful as [[fire retardant]]s, and more than half the bromine produced worldwide each year is put to this purpose. The same property causes ultraviolet [[sunlight]] to dissociate volatile organobromine compounds in the [[atmosphere]] to yield free bromine atoms, causing [[ozone depletion]]. As a result, many organobromine compounds—such as the [[pesticide]] [[Bromomethane|methyl bromide]]—are no longer used. Bromine compounds are still used in [[well drilling fluids]], in [[photographic film]], and as an intermediate in the manufacture of [[organic compound|organic]] chemicals.

Large amounts of bromide salts are toxic from the action of soluble bromide ions, causing [[bromism]]. However, bromine is beneficial for human [[eosinophil]]s,<ref name="pmid2538427" /> and is an essential trace element for [[collagen]] development in all animals.<ref name="pmid24906154" /> Hundreds of known organobromine compounds are generated by terrestrial and marine plants and animals, and some serve important biological roles.<ref name="Gribble99" /> As a [[pharmaceutical]], the simple bromide ion (Br{{sup|−}}) has inhibitory effects on the central nervous system, and bromide [[Salt (chemistry)|salts]] were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses as [[antiepileptic]]s.

==History==
[[File:Antoine Jérôme Balard 1870s.jpg|thumb|upright|left|[[Antoine Jerome Balard|Antoine Balard]], one of the discoverers of bromine]]

Bromine was discovered independently by two chemists, [[Carl Jacob Löwig]]<ref name="L1" /> and [[Antoine Jérôme Balard|Antoine Balard]],<ref name="Bal1826" /><ref name="Balard" /> in 1825 and 1826, respectively.<ref>{{Cite journal|title = The discovery of the elements: XVII. The halogen family|last = Weeks|first = Mary Elvira|author-link=Mary Elvira Weeks|journal = Journal of Chemical Education|date = 1932|volume = 9|page = 1915|doi = 10.1021/ed009p1915|bibcode=1932JChEd...9.1915W|issue = 11}}</ref>

Löwig isolated bromine from a mineral water spring from his hometown [[Bad Kreuznach]] in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with [[diethyl ether]]. After evaporation of the ether, a brown liquid remained. With this liquid as a sample of his work he applied for a position in the laboratory of [[Leopold Gmelin]] in [[Heidelberg]]. The publication of the results was delayed and Balard published his results first.<ref name="Löwig" />

Balard found bromine chemicals in the ash of [[seaweed]] from the [[salt marsh]]es of [[Montpellier]]. The seaweed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance were intermediate between those of chlorine and iodine; thus he tried to prove that the substance was [[iodine monochloride]] (ICl), but after failing to do so he was sure that he had found a new element and named it muride, derived from the [[Latin]] word {{lang|la|muria}} ("brine").<ref name="Balard" /><ref name="OEtymD" /><ref>{{L&S|muria|ref}}</ref>

After the French chemists [[Louis Nicolas Vauquelin]], [[Louis Jacques Thénard]], and [[Joseph-Louis Gay-Lussac]] approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the [[Académie des Sciences]] and published in ''Annales de Chimie et Physique''.<ref name="Bal1826" /> In his publication, Balard stated that he changed the name from ''muride'' to ''brôme'' on the proposal of M. Anglada. The name ''brôme'' (bromine) derives from the [[Greek language|Greek]] {{lang|grc|βρῶμος}} ({{transliteration|grc| brômos}}, "stench").<ref name="Bal1826" /><ref name="Bal1826b" /><ref name="OEtymD">{{OEtymD|bromine}}</ref><ref>{{LSJ|brw{{=}}mos2|βρῶμος|ref}}.</ref> Other sources claim that the French chemist and physicist [[Joseph-Louis Gay-Lussac]] suggested the name ''brôme'' for the characteristic smell of the vapors.<ref name="b1" /><ref name="Wisniak" /> Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in [[Stassfurt]] enabled its production as a by-product of [[potash]].<ref name="Greenwood790">Greenwood and Earnshaw, p. 790</ref>

<!-- Some also suggest that it may have been discovered by [[Bernard Courtois]], the man who discovered iodine.<ref>{{cite web
|url = https://www.vanderkrogt.net/elements/elem/br.html
|title = Bromine
|publisher = vanderkrogt.net }}</ref>-->
Apart from some minor medical applications, the first commercial use was the [[daguerreotype]]. In 1840, bromine was discovered to have some advantages over the previously used iodine vapor to create the light sensitive [[silver halide]] layer in daguerreotypy.<ref>{{Cite book|title = The Daguerreotype: Nineteenth-century Technology and Modern Science|first = M. Susan|last = Barger|author2=White, William Blaine|publisher = JHU Press|date = 2000|isbn = 978-0-8018-6458-2|chapter = Technological Practice of Daguerreotypy| pages =31–35}}</ref>

By 1864, a 25% solution of liquid bromine in .75 molar aqueous potassium bromide<ref>{{cite web|  url=https://chestofbooks.com/health/materia-medica-drugs/Materia-Medica-Therapeutics-Inorganic-Substances/Hospital-Gangrene-Erysipelas-Bromine-Bromum-Treatment.html| title=Hospital Gangrene - Erysipelas. Bromine (Bromum) Treatment}} The formula commonly employed was: "bromine, 1 oz.; bromide of potassium, 160 gr.; water, 4 oz."</ref> was widely used<ref>{{cite web| url=https://www.civilwarmed.org/hospital-gangrene-in-the-civil-war/| title=Hospital Gangrene in the Civil War | date=19 July 2023 }}</ref> to treat [[gangrene]] during the American Civil War, before the publications of [[Joseph Lister]] and [[Pasteur]].<ref>{{cite journal| title=Treatment of War Wounds: A Historical Review| author1= Manring, M. M.| author2= Hawk, Alan| author3= Calhoun, Jason H.| author4= Anderson, Romney C.| journal=Clinical Orthopaedics and Related Research| date=2009| volume=467| issue=8| pages=2168–2191| doi=10.1007/s11999-009-0738-5| pmid=19219516| pmc=2706344}} In 1863, the Union medical officer Middleton Goldsmith (1818–1887), stationed in Louisville, KY, reported the results of a treatment protocol that called for débridement of all necrotic tissue and application of a mixture of bromine, bromide of potassium, and water applied to dressings.</ref> 

[[Potassium bromide]] and [[sodium bromide]] were used as [[anticonvulsant]]s and [[sedative]]s in the late 19th and early 20th centuries, but were gradually superseded by [[chloral hydrate]] and then by the [[barbiturate]]s.<ref>{{Cite book|title = A History of Psychiatry: From the Era of the Asylum to the Age of Prozac|first = Edward|last = Shorter| publisher = John Wiley and Sons|date = 1997|isbn = 978-0-471-24531-5|page =200}}</ref> In the early years of the [[First World War]], bromine compounds such as [[xylyl bromide]] were used as [[poison gas]].<ref name="Borden_chemwarfare">{{cite book | title = Medical Aspects of Chemical Warfare | year = 2008 | publisher = [[Borden Institute]] | author1 = Corey J Hilmas |author2=Jeffery K Smart |author3=Benjamin A Hill | chapter = Chapter 2: History of Chemical Warfare (pdf) | pages = 12–14 | chapter-url = https://www.bordeninstitute.army.mil/published_volumes/chemwarfare/CHAP2_Pg_09-76.pdf | access-date = 20 November 2016 | archive-date = 26 August 2012 | archive-url = https://web.archive.org/web/20120826100423/https://www.bordeninstitute.army.mil/published_volumes/chemwarfare/CHAP2_Pg_09-76.pdf | url-status = dead }}</ref>

==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>

==Chemistry and compounds==
{{Main|Bromine compounds}}
{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ style="margin-bottom: 5px;" | Halogen bond energies (kJ/mol)<ref name="Greenwood804" />
|-
! X
! XX
! HX
! BX{{sub|3}}
! AlX{{sub|3}}
! CX{{sub|4}}
|-
! F
| 159
| 574
| 645
| 582
| 456
|-
! Cl
|243
|428
|444
|427
|327
|-
! Br
|193
|363
|368
|360
|272
|-
! I
|151
|294
|272
|285
|239
|}
Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the [[standard electrode potential]]s of the X{{sub|2}}/X{{sup|−}} couples (F, +2.866&nbsp;V; Cl, +1.395&nbsp;V; Br, +1.087&nbsp;V; I, +0.615&nbsp;V; At, approximately +0.3&nbsp;V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.<ref name="Greenwood804" />

===Hydrogen bromide===
The simplest compound of bromine is [[hydrogen bromide]], HBr. It is mainly used in the production of inorganic [[bromide]]s and [[alkyl bromide]]s, and as a catalyst for many reactions in organic chemistry. Industrially, it is mainly produced by the reaction of [[hydrogen]] gas with bromine gas at 200–400&nbsp;°C with a [[platinum]] catalyst. However, reduction of bromine with [[red phosphorus]] is a more practical way to produce hydrogen bromide in the laboratory:<ref name="Greenwood809">Greenwood and Earnshaw, pp. 809–12</ref>
: 2 P + 6 H{{sub|2}}O + 3 Br{{sub|2}} → 6 HBr + 2 H{{sub|3}}PO{{sub|3}}
: H{{sub|3}}PO{{sub|3}} + H{{sub|2}}O + Br{{sub|2}} → 2 HBr + H{{sub|3}}PO{{sub|4}}

At room temperature, hydrogen bromide is a colourless gas, like all the hydrogen halides apart from [[hydrogen fluoride]], since hydrogen cannot form strong [[hydrogen bond]]s to the large and only mildly electronegative bromine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen bromide at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised.<ref name="Greenwood809" /> Aqueous hydrogen bromide is known as [[hydrobromic acid]], which is a strong acid (p''K''{{sub|a}} = −9) because the hydrogen bonds to bromine are too weak to inhibit dissociation. The HBr/H{{sub|2}}O system also involves many hydrates HBr·''n''H{{sub|2}}O for ''n'' = 1, 2, 3, 4, and 6, which are essentially salts of bromine [[anion]]s and [[hydronium]] [[cation]]s. Hydrobromic acid forms an [[azeotrope]] with boiling point 124.3&nbsp;°C at 47.63&nbsp;g HBr per 100&nbsp;g solution; thus hydrobromic acid cannot be concentrated beyond this point by distillation.<ref name="Greenwood812">Greenwood and Earnshaw, pp. 812–6</ref>

Unlike [[hydrogen fluoride]], anhydrous liquid hydrogen bromide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its [[dielectric constant]] is low and it does not dissociate appreciably into H{{sub|2}}Br{{sup|+}} and {{chem|HBr|2|-}} ions – the latter, in any case, are much less stable than the [[bifluoride]] ions ({{chem|HF|2|-}}) due to the very weak hydrogen bonding between hydrogen and bromine, though its salts with very large and weakly polarising cations such as [[caesium|Cs{{sup|+}}]] and [[quaternary ammonium cation|{{chem|NR|4|+}}]] (R = [[methyl group|Me]], [[ethyl group|Et]], [[butyl group|Bu{{sup|''n''}}]]) may still be isolated. Anhydrous hydrogen bromide is a poor solvent, only able to dissolve small molecular compounds such as [[nitrosyl chloride]] and [[phenol]], or salts with very low [[lattice energy|lattice energies]] such as tetraalkylammonium halides.<ref name="Greenwood812" />

===Other binary bromides===
[[File:Bromid stříbrný.PNG|thumb|right|[[Silver bromide]] (AgBr)]]
Nearly all elements in the periodic table form binary bromides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the [[noble gas]]es, with the exception of [[xenon]] in the very unstable [[Xenon dibromide|XeBr{{sub|2}}]]); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond [[bismuth]]); and having an electronegativity higher than bromine's ([[oxygen]], [[nitrogen]], [[fluorine]], and [[chlorine]]), so that the resultant binary compounds are formally not bromides but rather oxides, nitrides, fluorides, or chlorides of bromine. (Nonetheless, [[nitrogen tribromide]] is named as a bromide as it is analogous to the other nitrogen trihalides.)<ref name="Greenwood821">Greenwood and Earnshaw, pp. 821–4</ref>

Bromination of metals with Br{{sub|2}} tends to yield lower oxidation states than chlorination with Cl{{sub|2}} when a variety of oxidation states is available. Bromides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrobromic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen bromide gas. These methods work best when the bromide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative bromination of the element with bromine or hydrogen bromide, high-temperature bromination of a metal oxide or other halide by bromine, a volatile metal bromide, [[carbon tetrabromide]], or an organic bromide. For example, [[niobium(V) oxide]] reacts with carbon tetrabromide at 370&nbsp;°C to form [[niobium(V) bromide]].<ref name="Greenwood821" /> Another method is halogen exchange in the presence of excess "halogenating reagent", for example:<ref name="Greenwood821" />
:FeCl{{sub|3}} + BBr{{sub|3}} (excess) → FeBr{{sub|3}} + BCl{{sub|3}}
When a lower bromide is wanted, either a higher halide may be reduced using hydrogen or a metal as a reducing agent, or thermal decomposition or [[disproportionation]] may be used, as follows:<ref name="Greenwood821" />
: 3 WBr{{sub|5}} + Al {{overunderset|→|thermal gradient|475&nbsp;°C → 240&nbsp;°C}} 3 WBr{{sub|4}} + AlBr{{sub|3}}
: EuBr{{sub|3}} + {{sfrac|1|2}} H{{sub|2}} → EuBr{{sub|2}} + HBr
: 2 TaBr{{sub|4}} {{overunderset|→|500&nbsp;°C|&nbsp;}} TaBr{{sub|3}} + TaBr{{sub|5}}

Most metal bromides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Both ionic and covalent bromides are known for metals in oxidation state +3 (e.g. [[scandium bromide]] is mostly ionic, but [[aluminium bromide]] is not). [[Silver bromide]] is very insoluble in water and is thus often used as a qualitative test for bromine.<ref name="Greenwood821" />

===Bromine halides===
The halogens form many binary, [[diamagnetic]] [[interhalogen]] compounds with stoichiometries XY, XY{{sub|3}}, XY{{sub|5}}, and XY{{sub|7}} (where X is heavier than Y), and bromine is no exception. Bromine forms a monofluoride and monochloride, as well as a trifluoride and pentafluoride. Some cationic and anionic derivatives are also characterised, such as {{chem|BrF|2|-}}, {{chem|BrCl|2|-}}, {{chem|BrF|2|+}}, {{chem|BrF|4|+}}, and {{chem|BrF|6|+}}. Apart from these, some [[pseudohalogen|pseudohalides]] are also known, such as [[cyanogen bromide]] (BrCN), bromine [[thiocyanate]] (BrSCN), and bromine [[azide]] (BrN{{sub|3}}).<ref name="Greenwood824">Greenwood and Earnshaw, pp. 824–8</ref>

The pale-brown [[bromine monofluoride]] (BrF) is unstable at room temperature, disproportionating quickly and irreversibly into bromine, bromine trifluoride, and bromine pentafluoride. It thus cannot be obtained pure. It may be synthesised by the direct reaction of the elements, or by the comproportionation of bromine and bromine trifluoride at high temperatures.<ref name="Greenwood824" /> [[Bromine monochloride]] (BrCl), a red-brown gas, quite readily dissociates reversibly into bromine and chlorine at room temperature and thus also cannot be obtained pure, though it can be made by the reversible direct reaction of its elements in the gas phase or in [[carbon tetrachloride]].<ref name="Greenwood821" /> Bromine monofluoride in [[ethanol]] readily leads to the monobromination of the [[aromaticity|aromatic]] compounds PhX (''para''-bromination occurs for X = Me, Bu{{sup|''t''}}, OMe, Br; ''meta''-bromination occurs for the deactivating X = –CO{{sub|2}}Et, –CHO, –NO{{sub|2}}); this is due to heterolytic fission of the Br–F bond, leading to rapid electrophilic bromination by Br{{sup|+}}.<ref name="Greenwood821" />

At room temperature, [[bromine trifluoride]] (BrF{{sub|3}}) is a straw-coloured liquid. It may be formed by directly fluorinating bromine at room temperature and is purified through distillation. It reacts violently with water and explodes on contact with flammable materials, but is a less powerful fluorinating reagent than [[chlorine trifluoride]]. It reacts vigorously with [[boron]], [[carbon]], [[silicon]], [[arsenic]], [[antimony]], iodine, and [[sulfur]] to give fluorides, and will also convert most metals and many metal compounds to fluorides; as such, it is used to oxidise [[uranium]] to [[uranium hexafluoride]] in the nuclear power industry. Refractory oxides tend to be only partially fluorinated, but here the derivatives KBrF{{sub|4}} and BrF{{sub|2}}SbF{{sub|6}} remain reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to form {{chem|BrF|2|+}} and {{chem|BrF|4|-}} and thus conducts electricity.<ref name="Greenwood828">Greenwood and Earnshaw, pp. 828–31</ref>

[[Bromine pentafluoride]] (BrF{{sub|5}}) was first synthesised in 1930. It is produced on a large scale by direct reaction of bromine with excess fluorine at temperatures higher than 150&nbsp;°C, and on a small scale by the fluorination of [[potassium bromide]] at 25&nbsp;°C. It also reacts violently with water and is a very strong fluorinating agent, although chlorine trifluoride is still stronger.<ref name="Greenwood832">Greenwood and Earnshaw, pp. 832–5</ref>

===Polybromine compounds===
Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such as [[peroxydisulfuryl fluoride]] (S{{sub|2}}O{{sub|6}}F{{sub|2}}) can oxidise it to form the cherry-red {{chem|Br|2|+}} cation. A few other bromine cations are known, namely the brown {{chem|Br|3|+}} and dark brown {{chem|Br|5|+}}.<ref name="Greenwood842">Greenwood and Earnshaw, pp. 842–4</ref> The tribromide anion, {{chem|Br|3|-}}, has also been characterised; it is analogous to [[triiodide]].<ref name="Greenwood824" />

===Bromine oxides and oxoacids===
{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ Standard reduction potentials for aqueous Br species<ref name="Greenwood853" />
! {{nowrap|E°(couple)}}!!{{nowrap|''a''(H{{sup|+}}) {{=}} 1}}<br>(acid)!!{{nowrap|E°(couple)}}!!{{nowrap|''a''(OH{{sup|−}}) {{=}} 1}}<br>(base)
|-
|Br{{sub|2}}/Br{{sup|−}}||+1.052|||Br{{sub|2}}/Br{{sup|−}}||+1.065
|-
|HOBr/Br{{sup|−}}||+1.341||BrO{{sup|−}}/Br{{sup|−}}||+0.760
|-
|{{chem|BrO|3|-}}/Br{{sup|−}}||+1.399||{{chem|BrO|3|-}}/Br{{sup|−}}||+0.584
|-
|HOBr/Br{{sub|2}}||+1.604||BrO{{sup|−}}/Br{{sub|2}}||+0.455
|-
|{{chem|BrO|3|-}}/Br{{sub|2}}||+1.478||{{chem|BrO|3|-}}/Br{{sub|2}}||+0.485
|-
|{{chem|BrO|3|-}}/HOBr||+1.447||{{chem|BrO|3|-}}/BrO{{sup|−}}||+0.492
|-
|{{chem|BrO|4|-}}/{{chem|BrO|3|-}}||+1.853||{{chem|BrO|4|-}}/{{chem|BrO|3|-}}||+1.025
|}
[[Bromine oxide]]s are not as well-characterised as [[chlorine oxide]]s or [[iodine oxide]]s, as they are all fairly unstable: it was once thought that they could not exist at all. [[Dibromine monoxide]] is a dark-brown solid which, while reasonably stable at −60&nbsp;°C, decomposes at its melting point of −17.5&nbsp;°C; it is useful in [[bromination]] reactions<ref name="handin">{{Citation
 | last1 = Perry
 | first1 = Dale L.
 | last2 = Phillips
 | first2 = Sidney L.
 | year = 1995
 | title = Handbook of Inorganic Compounds
 | publisher = CRC Press
 | isbn = 978-0-8493-8671-8
 | pages = 74
 | url = https://books.google.com/books?id=0fT4wfhF1AsC&q=%22Bromine+dioxide%22&pg=PA74
 | access-date = 25 August 2015
 | archive-date = 25 July 2021
 | archive-url = https://web.archive.org/web/20210725075132/https://books.google.com/books?id=0fT4wfhF1AsC&q=%22Bromine+dioxide%22&pg=PA74
 | url-status = live
 }}</ref> and may be made from the low-temperature decomposition of [[bromine dioxide]] in a vacuum. It oxidises iodine to [[iodine pentoxide]] and [[benzene]] to [[1,4-benzoquinone]]; in alkaline solutions, it gives the [[hypobromite]] anion.<ref name="Greenwood850">Greenwood and Earnshaw, pp. 850–1</ref>

So-called "[[bromine dioxide]]", a pale yellow crystalline solid, may be better formulated as bromine [[perbromate]], BrOBrO{{sub|3}}. It is thermally unstable above −40&nbsp;°C, violently decomposing to its elements at 0&nbsp;°C. [[Dibromine trioxide]], ''syn''-BrOBrO{{sub|2}}, is also known; it is the anhydride of [[hypobromous acid]] and [[bromic acid]]. It is an orange crystalline solid which decomposes above −40&nbsp;°C; if heated too rapidly, it explodes around 0&nbsp;°C. A few other unstable radical oxides are also known, as are some poorly characterised oxides, such as [[dibromine pentoxide]], [[tribromine octoxide]], and bromine trioxide.<ref name="Greenwood850" />

The four [[oxoacid]]s, [[hypobromous acid]] (HOBr), [[bromous acid]] (HOBrO), [[bromic acid]] (HOBrO{{sub|2}}), and [[perbromic acid]] (HOBrO{{sub|3}}), are better studied due to their greater stability, though they are only so in aqueous solution. When bromine dissolves in aqueous solution, the following reactions occur:<ref name="Greenwood853">Greenwood and Earnshaw, pp. 853–9</ref>
:{|
|-
| Br{{sub|2}} + H{{sub|2}}O || {{eqm}} HOBr + H{{sup|+}} + Br{{sup|−}} || ''K''{{sub|ac}} = 7.2 × 10{{sup|−9}} mol{{sup|2}} l{{sup|−2}}
|-
| Br{{sub|2}} + 2 OH{{sup|−}} || {{eqm}} OBr{{sup|−}} + H{{sub|2}}O + Br{{sup|−}} || ''K''{{sub|alk}} = 2 × 10{{sup|8}} mol{{sup|−1}} l
|}

Hypobromous acid is unstable to disproportionation. The [[hypobromite]] ions thus formed disproportionate readily to give bromide and bromate:<ref name="Greenwood853" />
:{|
|-
| 3 BrO{{sup|−}} {{eqm}} 2 Br{{sup|−}} + {{chem|BrO|3|-}} || ''K'' = 10{{sup|15}}
|}

Bromous acids and [[bromite]]s are very unstable, although the [[strontium]] and [[barium]] bromites are known.<ref name="Greenwood862">Greenwood and Earnshaw, pp. 862–5</ref> More important are the [[bromate]]s, which are prepared on a small scale by oxidation of bromide by aqueous [[hypochlorite]], and are strong oxidising agents. Unlike chlorates, which very slowly disproportionate to chloride and perchlorate, the bromate anion is stable to disproportionation in both acidic and aqueous solutions. Bromic acid is a strong acid. Bromides and bromates may comproportionate to bromine as follows:<ref name="Greenwood862" />
:{{chem|BrO|3|-}} + 5 Br{{sup|−}} + 6 H{{sup|+}} → 3 Br{{sub|2}} + 3 H{{sub|2}}O

There were many failed attempts to obtain perbromates and perbromic acid, leading to some rationalisations as to why they should not exist, until 1968 when the anion was first synthesised from the radioactive [[beta decay]] of unstable {{chem|83|Se|O|4|2-}}. Today, perbromates are produced by the oxidation of alkaline bromate solutions by fluorine gas. Excess bromate and fluoride are precipitated as [[silver bromate]] and [[calcium fluoride]], and the perbromic acid solution may be purified. The perbromate ion is fairly inert at room temperature but is thermodynamically extremely oxidising, with extremely strong oxidising agents needed to produce it, such as fluorine or [[xenon difluoride]]. The Br–O bond in {{chem|BrO|4|-}} is fairly weak, which corresponds to the general reluctance of the 4p elements [[arsenic]], [[selenium]], and bromine to attain their group oxidation state, as they come after the [[scandide contraction]] characterised by the poor shielding afforded by the radial-nodeless 3d orbitals.<ref name="Greenwood871">Greenwood and Earnshaw, pp. 871–2</ref>

===Organobromine compounds===
{{main|Organobromine compound}}
[[File:N-Bromosuccinimide.svg|thumb|upright|Structure of [[N-Bromosuccinimide|''N''-bromosuccinimide]], a common brominating reagent in organic chemistry]]
Like the other carbon–halogen bonds, the C–Br bond is a common functional group that forms part of core [[organic chemistry]]. Formally, compounds with this functional group may be considered organic derivatives of the bromide anion. Due to the difference of electronegativity between bromine (2.96) and carbon (2.55), the carbon atom in a C–Br bond is electron-deficient and thus [[electrophilic]]. The reactivity of organobromine compounds resembles but is intermediate between the reactivity of [[organochlorine compound|organochlorine]] and [[organoiodine compound]]s. For many applications, organobromides represent a compromise of reactivity and cost.<ref name="KO" />

Organobromides are typically produced by additive or substitutive bromination of other organic precursors. Bromine itself can be used, but due to its toxicity and volatility, safer brominating reagents are normally used, such as [[N-Bromosuccinimide|''N''-bromosuccinimide]]. The principal reactions for organobromides include [[dehydrohalogenation|dehydrobromination]], [[Grignard reaction]]s, [[Wurtz reaction|reductive coupling]], and [[nucleophilic substitution]].<ref name="KO">Ioffe, David and Kampf, Arieh (2002) "Bromine, Organic Compounds" in ''Kirk-Othmer Encyclopedia of Chemical Technology''. John Wiley & Sons. {{doi| 10.1002/0471238961.0218151325150606.a01}}.</ref>

Organobromides are the most common organohalides in nature, even though the concentration of bromide is only 0.3% of that for chloride in sea water, because of the easy oxidation of bromide to the equivalent of Br{{sup|+}}, a potent electrophile. The enzyme [[bromoperoxidase]] catalyzes this reaction.<ref>{{cite journal|doi=10.1021/ja047925p|pmid=15548002|title=Vanadium Bromoperoxidase-Catalyzed Biosynthesis of Halogenated Marine Natural Products|journal=Journal of the American Chemical Society|volume=126|issue=46|pages=15060–6|year=2004|last1=Carter-Franklin|first1=Jayme N.|last2=Butler|first2=Alison|bibcode=2004JAChS.12615060C }}</ref> The oceans are estimated to release 1–2 million tons of [[bromoform]] and 56,000 tons of [[bromomethane]] annually.<ref name="Gribble99" />

[[Image:Alkene-bromine-addition-2D-skeletal.png|upright=1.8|thumb|Bromine addition to alkene reaction mechanism]]
An old qualitative test for the presence of the [[alkene]] functional group is that alkenes turn brown aqueous bromine solutions colourless, forming a [[halohydrin|bromohydrin]] with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilic [[halonium ion|bromonium]] intermediate. This is an example of a [[halogen addition reaction]].<ref name="Clayden">{{cite book | last1 = Clayden | first1 = Jonathan | author-link1 = Jonathan Clayden | last2 = Greeves | first2 = Nick | last3 = Warren | first3 = Stuart | author-link3 = Stuart Warren | title = Organic Chemistry | edition = 2nd | publisher = Oxford University Press | date = 2012 | isbn = 978-0-19-927029-3 |pages=427–9}}</ref>

==Occurrence and production==
[[File:STS028-96-65.jpg|thumb|View of salt evaporation pans on the Dead Sea, where [[Jordan]] (right) and Israel (left) produce salt and bromine]]
Bromine is significantly less abundant in the crust than fluorine or chlorine, comprising only 2.5&nbsp;[[parts per million]] of the Earth's crustal rocks, and then only as bromide salts. It is significantly more abundant in the oceans, resulting from long-term [[Leaching (chemical science)|leaching]]. There, it makes up 65&nbsp;parts per million, corresponding to a ratio of about one bromine atom for every 660 chlorine atoms. Salt lakes and brine wells may have higher bromine concentrations: for example, the [[Dead Sea]] contains 0.4% bromide ions.<ref name="Greenwood795">Greenwood and Earnshaw, pp. 795–6</ref> It is from these sources that bromine extraction is mostly economically feasible.<ref name="o1">{{Cite journal|title = Minerals From Sea Salt|first = John A.|last = Tallmadge|author2 = Butt, John B.|author3 = Solomon Herman J.|journal = Ind. Eng. Chem.|date = 1964|volume = 56 |pages = 44–65| doi = 10.1021/ie50655a008|issue = 7}}</ref><ref name="o2">{{Cite journal|journal = Clinics in Dermatology|volume = 14 |date = 1996|pages = 659–664|doi = 10.1016/S0738-081X(96)00101-0|pmid = 8960809 |title = Climatotherapy at the Dead Sea in Jordan|first = Oumeish Youssef|last = Oumeish|issue = 6}}</ref><ref name="o3">{{Cite journal|journal = Hydrological Processes|volume = 14|issue = 1|pages = 145–154| date = 2008|title = The water balance of the Dead Sea: an integrated approach|first = Radwan A.|last = Al-Weshah|doi = 10.1002/(SICI)1099-1085(200001)14:1<145::AID-HYP916>3.0.CO;2-N|bibcode = 2000HyPr...14..145A }}</ref> Bromine is the tenth most abundant element in seawater.<ref>{{Cite web |title=Get the Facts About the Element Bromine |url=https://www.thoughtco.com/bromine-element-facts-606510 |access-date=2024-08-10 |website=ThoughtCo |language=en}}</ref>

The main sources of bromine production are [[Israel]] and [[Jordan]].<ref>{{cite web |title=Major countries in worldwide bromine production from 2018 to 2022 |url=https://www.statista.com/statistics/264926/world-bromine-production/ |publisher=Statista |access-date=26 April 2023}}</ref> The element is liberated by halogen exchange, using chlorine gas to oxidise Br{{sup|−}} to Br{{sub|2}}. This is then removed with a blast of steam or air, and is then condensed and purified.<ref>{{cite web |title=Process operations at Octel Amlwch |url=https://www.octelamlwch.co.uk/operations/ |website=Octel Bromine Works |access-date=27 July 2021 |archive-date=27 July 2021 |archive-url=https://web.archive.org/web/20210727211804/http://www.octelamlwch.co.uk/operations/ |url-status=live }}</ref> Today, bromine is transported in large-capacity metal drums or lead-lined tanks that can hold hundreds of kilograms or even tonnes of bromine. The bromine industry is about one-hundredth the size of the chlorine industry. Laboratory production is unnecessary because bromine is commercially available and has a long shelf life.<ref name="Greenwood798">Greenwood and Earnshaw, pp. 798–9</ref>

==Applications==
A wide variety of organobromine compounds are used in [[Industry (manufacturing)|industry]]. Some are prepared from bromine and others are prepared from [[hydrogen bromide]], which is obtained by burning [[hydrogen]] in bromine.<ref name="Ullmann">{{Cite book|author=Mills, Jack F. |title=Bromine: in Ullmann's Encyclopedia of Chemical Technology|publisher= Wiley-VCH Verlag|location= Weinheim|date= 2002|doi=10.1002/14356007.a04_391|chapter=Bromine|isbn=978-3527306732}}</ref>

===Flame retardants===
[[File:Tetrabromobisphenol A.svg|thumb|left|Tetrabromobisphenol A]]

[[Brominated flame retardant]]s represent a commodity of growing importance, and make up the largest commercial use of bromine. When the brominated material burns, the flame retardant produces [[hydrobromic acid]] which interferes in the radical [[chain reaction]] of the [[oxidation]] reaction of the fire. The mechanism is that the highly reactive hydrogen radicals, oxygen radicals, and [[hydroxyl radical]]s react with hydrobromic acid to form less reactive bromine radicals (i.e., free bromine atoms). Bromine atoms may also react directly with other radicals to help terminate the free radical chain-reactions that characterise combustion.<ref>{{Cite journal|journal = Journal of Fire Sciences|volume = 14|pages = 426–442| date = 1996|doi = 10.1177/073490419601400602|title = Mechanisms for Flame Retardancy and Smoke suppression – A Review|first =Joseph|last = Green|issue = 6 |s2cid = 95145090}}</ref><ref>{{Cite journal|journal = Polymer Degradation and Stability|volume = 77|date = 2002|pages = 325–331| doi = 10.1016/S0141-3910(02)00067-8|title = Fire retardant mechanism of aliphatic bromine compounds in polystyrene and polypropylene|first = Jelle|last = Kaspersma|author2 = Doumena, Cindy|author3 = Munrob Sheilaand|author4 = Prinsa, Anne-Marie|issue = 2}}</ref>

To make brominated polymers and plastics, bromine-containing compounds can be incorporated into the polymer during [[polymerisation]]. One method is to include a relatively small amount of brominated monomer during the polymerisation process. For example, [[vinyl bromide]] can be used in the production of [[polyethylene]], [[polyvinyl chloride]] or [[polypropylene]]. Specific highly brominated molecules can also be added that participate in the polymerisation process. For example, [[tetrabromobisphenol A]] can be added to [[polyester]]s or epoxy resins, where it becomes part of the polymer. Epoxies used in [[printed circuit board]]s are normally made from such flame retardant [[resin]]s, indicated by the FR in the abbreviation of the products ([[FR-4]] and [[FR-2]]). In some cases, the bromine-containing compound may be added after polymerisation. For example, [[decabromodiphenyl ether]] can be added to the final polymers.<ref>{{Cite journal|journal = Journal of Fire Sciences|volume = 22|pages = 25–40| date = 2004|doi = 10.1177/0734904104038107|title = A Review of Current Flame Retardant Systems for Epoxy Resins|first = Edward D.|last = Weil|author2=Levchik, Sergei|s2cid = 95746728}}</ref>

A number of gaseous or highly volatile brominated [[halomethane]] compounds are non-toxic and make superior fire suppressant agents by this same mechanism, and are particularly effective in enclosed spaces such as submarines, airplanes, and spacecraft. However, they are expensive and their production and use has been greatly curtailed due to their effect as ozone-depleting agents. They are no longer used in routine fire extinguishers, but retain niche uses in aerospace and military automatic fire suppression applications. They include [[bromochloromethane]] (Halon 1011, CH{{sub|2}}BrCl), [[bromochlorodifluoromethane]] (Halon 1211, CBrClF{{sub|2}}), and [[bromotrifluoromethane]] (Halon 1301, CBrF{{sub|3}}).<ref name="UllmannF">Siegemund, Günter; Schwertfeger, Werner; Feiring, Andrew; Smart, Bruce; Behr, Fred; Vogel, Herward; McKusick, Blaine (2002) "Fluorine Compounds, Organic" Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. {{doi|10.1002/14356007.a11_349}}</ref>

===Other uses===
[[File:Bromo-Seltzer_Tower_MD2.jpg|thumb|upright|Baltimore's [[Emerson Bromo-Seltzer Tower]], originally part of the headquarters of Emerson Drug Company, which made [[Bromo-Seltzer]]]]
[[Silver bromide]] is used, either alone or in combination with [[silver chloride]] and [[silver iodide]], as the light sensitive constituent of [[photographic emulsion]]s.<ref name="Greenwood798" />

[[1,2-Dibromoethane|Ethylene bromide]] was an [[Gasoline additive|additive in gasolines]] containing lead anti-[[engine knocking]] agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine use in 1966 in the US. This application has declined since the 1970s due to environmental regulations (see below).<ref>{{Cite journal|journal = Environment International|volume = 29|date = 2003| doi = 10.1016/S0160-4120(03)00121-1|title = An overview of commercially used brominated flame retardants, their applications, their use patterns in different countries/regions and possible modes of release|first = Mehran|last = Alaeea|author2 = Ariasb, Pedro|author3 = Sjödinc, Andreas|author4 = Bergman, Åke|issue = 6|pmid=12850087|pages = 683–9| bibcode=2003EnInt..29..683A }}</ref>

[[Brominated vegetable oil]] (BVO), a complex mixture of plant-derived triglycerides that have been reacted to contain atoms of the element bromine bonded to the molecules, is used primarily to help emulsify citrus-flavored soft drinks, preventing them from separating during distribution.

Poisonous [[bromomethane]] was widely used as [[pesticide]] to [[Fumigation|fumigate]] soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used.<ref name="USGSYB2007" /> These volatile organobromine compounds are all now regulated as [[ozone depletion]] agents. The [[Montreal Protocol|Montreal Protocol on Substances that Deplete the Ozone Layer]] scheduled the phase out for the [[ozone depleting]] chemical by 2005, and organobromide pesticides are no longer used (in housing fumigation they have been replaced by such compounds as [[sulfuryl fluoride]], which contain neither the chlorine or bromine organics which harm ozone). Before the Montreal protocol in 1991 (for example) an estimated 35,000 tonnes of the chemical were used to control [[nematode]]s, [[fungi]], [[weed]]s and other soil-borne diseases.<ref>{{cite web|title = Alternatives to Methyl Bromide for the Control of Soil-Borne Diseases and Pests in California|first = Belinda|last = Messenger|author2 = Braun, Adolf|date = 2000|publisher = Pest Management Analysis and Planning Program|url = https://www.cdpr.ca.gov/docs/emon/methbrom/alt-anal/sept2000.pdf|access-date = 2008-11-17|archive-url = https://web.archive.org/web/20100421041145/https://www.cdpr.ca.gov/docs/emon/methbrom/alt-anal/sept2000.pdf|archive-date = 2010-04-21|url-status = dead}}</ref><ref>{{Cite journal|title = Economics of the "Critical Use" of Methyl bromide under the Montreal Protocol|doi = 10.1093/cep/byi028|journal =Contemporary Economic Policy|volume = 23|pages = 376–393|date = 2008|first = Stephen J.|last = Decanio|author2=Norman, Catherine S.|issue = 3}}</ref>

In pharmacology, inorganic [[bromide]] compounds, especially [[potassium bromide]], were frequently used as general sedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used as anticonvulsants in both veterinary and human medicine, although the latter use varies from country to country. For example, the U.S. [[Food and Drug Administration]] (FDA) does not approve bromide for the treatment of any disease, and [[sodium bromide]] was removed from over-the-counter sedative products like [[Bromo-Seltzer]], in 1975.<ref>{{cite book|author=Adams, Samuel Hopkins |title=The Great American fraud|url=https://archive.org/details/greatamericanfr03adamgoog|access-date=2011-06-25 |date=1905|publisher=Press of the American Medical Association}}</ref> Commercially available organobromine pharmaceuticals include the vasodilator [[nicergoline]], the sedative [[brotizolam]], the anticancer agent [[pipobroman]], and the antiseptic [[merbromin]]. Otherwise, organobromine compounds are rarely pharmaceutically useful, in contrast to the situation for [[organofluorine chemistry|organofluorine compounds]]. Several drugs are produced as the bromide (or equivalents, hydrobromide) salts, but in such cases bromide serves as an innocuous counterion of no biological significance.<ref name="KO" />

Other uses of organobromine compounds include high-density drilling fluids, dyes (such as [[Tyrian purple]] and the indicator [[bromothymol blue]]), and pharmaceuticals. Bromine itself, as well as some of its compounds, are used in water treatment, and is the precursor of a variety of inorganic compounds with an enormous number of applications (e.g. [[silver bromide]] for photography).<ref name="Greenwood798" /> [[Zinc–bromine battery|Zinc–bromine batteries]] are hybrid [[Flow battery|flow batteries]] used for stationary electrical power backup and storage; from household scale to industrial scale.

Bromine is used in cooling towers (in place of chlorine) for controlling bacteria, algae, fungi, and [[zebra mussel]]s.<ref>Buecker, Brad (1998-01-07) [https://www.power-eng.com/emissions/choose-the-right-cooling-tower-chemicals/#gref Choose the Right Cooling Tower Chemicals]. ''Power Engineering''. {{Webarchive|url=https://web.archive.org/web/20210810000356/https://www.power-eng.com/emissions/choose-the-right-cooling-tower-chemicals/#gref|date=10 August 2021}} Choose the Right Cooling Tower Chemicals | Power Engineering |1998</ref>

Because it has similar antiseptic qualities to chlorine, bromine can be used in the same manner as chlorine as a disinfectant or antimicrobial in applications such as swimming pools. Bromine came into this use in the United States during [[World War II]] due to a predicted shortage of chlorine.<ref>{{Cite journal |last1=VanderVelde |first1=T. L. |last2=Mallmann |first2=W. L. |last3=Moore |first3=A. V. |date=1948 |title=A Comparative Study Of Chlorine And Bromine For Swimming Pool Disinfection |url=https://www.jstor.org/stable/26324332 |journal=The Sanitarian |volume=11 |issue=2 |pages=47–52 |jstor=26324332 |pmid=18891247 |issn=0096-560X}}</ref> However, bromine is usually not used outside for these applications due to it being relatively more expensive than chlorine and the absence of a stabilizer to protect it from the sun. For indoor pools, it can be a good option as it is effective at a wider pH range. It is also more stable in a heated pool or hot tub.<ref>{{Cite book |last=Blanchard |first=Kristine |title=Pool Care For Dummies |date=2023 |publisher=John Wiley and Sons |isbn=978-1-394-16611-4 |edition=1st |location=Indianapolis |pages=259–260}}</ref>

==Biological role and toxicity==
{{main|Vanadium bromoperoxidase|Eosinophil peroxidase|Bromism}}

A 2014 study suggests that bromine (in the form of bromide ion) is a necessary cofactor in the biosynthesis of [[collagen IV]], making the element [[essential element|essential]] to [[basement membrane]] architecture and tissue development in animals.<ref name="pmid24906154">{{Cite journal| display-authors = 5| author = McCall AS| author2 = Cummings CF| author3 = Bhave G| author4 = Vanacore R| author5 = Page-McCaw A| author6 = Hudson BG|title = Bromine Is an Essential Trace Element for Assembly of Collagen IV Scaffolds in Tissue Development and Architecture|journal = Cell|volume = 157|issue = 6|pages = 1380–92|date = 2014|pmid = 24906154|doi=10.1016/j.cell.2014.05.009| pmc=4144415}}</ref> Nevertheless, no clear deprivation symptoms or syndromes have been documented in mammals.<ref name="Nielsen2000">{{cite book|last1=Nielsen|first1=Forrest H.|title=Clinical Nutrition of the Essential Trace Elements and Minerals |chapter=Possibly Essential Trace Elements|year=2000|pages= 11–36|chapter-url=https://archive.org/details/clinicalnutritio0000unse_l0q5/page/11|doi=10.1007/978-1-59259-040-7_2|isbn=978-1-61737-090-8}}</ref> In other biological functions, bromine may be non-essential but still beneficial when it takes the place of chlorine. For example, in the presence of hydrogen peroxide, [[H2O2|H{{sub|2}}O{{sub|2}}]], formed by the [[eosinophil]], and either chloride, iodide, thiocyanate, or bromide ions, [[eosinophil peroxidase]] provides a potent mechanism by which eosinophils kill multicellular [[parasitism|parasites]] (such as the nematode worms involved in [[filariasis]]) and some [[bacteria]] (such as [[tuberculosis]] bacteria). Eosinophil peroxidase is a [[haloperoxidase]] that preferentially uses bromide over chloride for this purpose, generating [[hypobromite]] ([[hypobromous acid]]), although the use of chloride is possible.<ref name="pmid2538427">{{Cite journal| author = Mayeno AN| author2 = Curran AJ| author3 = Roberts RL| author4 = Foote CS|title = Eosinophils preferentially use bromide to generate halogenating agents|journal = J. Biol. Chem.|volume = 264|issue = 10|pages = 5660–8|date = 1989| doi = 10.1016/S0021-9258(18)83599-2|pmid = 2538427|doi-access = free}}</ref>

[[File:2-Octyl 4-bromo-3-oxobutanoate.svg|thumb|right|upright=1.2|Octan-2-yl 4-bromo-3-oxobutanoate, an organobromine compound found in mammalian cerebrospinal fluid]]

α-Haloesters are generally thought of as highly reactive and consequently toxic intermediates in organic synthesis. Nevertheless, mammals, including humans, cats, and rats, appear to biosynthesize traces of an α-bromoester, 2-octyl 4-bromo-3-oxobutanoate, which is found in their [[cerebrospinal fluid]] and appears to play a yet unclarified role in inducing REM sleep.<ref name="Gribble99" /> Neutrophil myeloperoxidase can use H{{sub|2}}O{{sub|2}} and Br{{sup|−}} to brominate deoxycytidine, which could result in DNA mutations.<ref name="pmid11096071">{{Cite journal| author = Henderson JP| author2 = Byun J| author3 = Williams MV| author4 = Mueller DM|title = Production of brominating intermediates by myeloperoxidase.|journal = J. Biol. Chem.|volume = 276|issue = 11|pages = 7867–75|date = 2001|pmid =11096071| doi = 10.1074/jbc.M005379200|doi-access = free}}</ref> Marine organisms are the main source of organobromine compounds, and it is in these organisms that bromine is more firmly shown to be essential. More than 1600 such organobromine compounds were identified by 1999. The most abundant is [[methyl bromide]] (CH{{sub|3}}Br), of which an estimated 56,000 tonnes is produced by marine algae each year.<ref name="Gribble99">{{Cite journal|title = The diversity of naturally occurring organobromine compounds|author = Gribble, Gordon W. |journal = Chemical Society Reviews| volume = 28|pages = 335–346|date = 1999|doi = 10.1039/a900201d|issue = 5}}</ref> The essential oil of the Hawaiian alga ''[[Asparagopsis taxiformis]]'' consists of 80% [[bromoform]].<ref>{{Cite journal|title = Volatile halogen compounds in the alga Asparagopsis taxiformis (Rhodophyta)|first = B. Jay|last = Burreson|author2= Moore, Richard E.|author3= Roller, Peter P.|journal = [[Journal of Agricultural and Food Chemistry]]|date = 1976|volume = 24|pages = 856–861|doi = 10.1021/jf60206a040|issue = 4| bibcode=1976JAFC...24..856B }}</ref> Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme, [[vanadium bromoperoxidase]].<ref>{{Cite journal|journal = Natural Product Reports|date = 2004|volume = 21|issue = 1|pmid = 15039842|doi = 10.1039/b302337k|title = The role of vanadium bromoperoxidase in the biosynthesis of halogenated marine natural products|first = Alison|last = Butler|author2=Carter-Franklin, Jayme N.|s2cid = 19115256|pages = 180–8}}</ref>

The bromide anion is not very toxic: a normal daily intake is 2 to 8&nbsp;milligrams.<ref name="Nielsen2000" /> However, high levels of bromide chronically impair the membrane of neurons, which progressively impairs neuronal transmission, leading to toxicity, known as [[bromism]]. Bromide has an [[elimination half-life]] of 9 to 12 days, which can lead to excessive accumulation. Doses of 0.5 to 1&nbsp;gram per day of bromide can lead to bromism. Historically, the therapeutic dose of bromide is about 3 to 5&nbsp;grams of bromide, thus explaining why chronic toxicity (bromism) was once so common. While significant and sometimes serious disturbances occur to neurologic, psychiatric, dermatological, and gastrointestinal functions, death from bromism is rare.<ref name="pdo2003">{{cite book |last1=Olson |first1=Kent R. |title=Poisoning & drug overdose |url=https://books.google.com/books?id=vuec3nTovyUC |edition=4th |date=1 November 2003 |publisher=Appleton & Lange |isbn=978-0-8385-8172-8 |pages=140–141 |access-date=5 November 2016 |archive-date=24 December 2016 |archive-url=https://web.archive.org/web/20161224042040/https://books.google.com/books?id=vuec3nTovyUC |url-status=live }}</ref> Bromism is caused by a neurotoxic effect on the brain which results in [[somnolence]], [[psychosis]], [[seizures]] and [[delirium]].<ref>{{cite book |last1=Galanter |first1=Marc |last2=Kleber |first2=Herbert D. |title=The American Psychiatric Publishing Textbook of Substance Abuse Treatment |url=https://books.google.com/books?id=6wdJgejlQzYC |edition=4th |date=1 July 2008 |publisher=American Psychiatric Publishing Inc |location=United States of America |isbn=978-1-58562-276-4 |page=217 |access-date=5 November 2016 |archive-date=27 April 2021 |archive-url=https://web.archive.org/web/20210427154856/https://books.google.com/books?id=6wdJgejlQzYC |url-status=live }}</ref>

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Elemental bromine (Br{{sub|2}}) is toxic and causes [[chemical burns]] on human flesh. Inhaling bromine gas results in similar irritation of the respiratory tract, causing coughing, choking, shortness of breath, and death if inhaled in large enough amounts. Chronic exposure may lead to frequent bronchial infections and a general deterioration of health. As a strong oxidising agent, bromine is incompatible with most organic and inorganic compounds.<ref name="msds">{{cite web |url=https://www.sciencelab.com/msds.php?msdsId=9927659 |title=Material Safety Data Sheet: Bromine MSDS |author=Science Lab.com |website=sciencelab.com |access-date=27 October 2016 |archive-url=https://web.archive.org/web/20171115230355/https://www.sciencelab.com/msds.php?msdsId=9927659 |archive-date=15 November 2017 |url-status=dead }}</ref> Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames.<ref name="Greenwood798" /> The [[Occupational Safety and Health Administration]] (OSHA) of the [[United States]] has set a [[permissible exposure limit]] (PEL) for bromine at a time-weighted average (TWA) of 0.1&nbsp;ppm. The [[National Institute for Occupational Safety and Health]] (NIOSH) has set a [[recommended exposure limit]] (REL) of TWA 0.1&nbsp;ppm and a short-term limit of 0.3&nbsp;ppm. The exposure to bromine [[immediately dangerous to life and health]] (IDLH) is 3&nbsp;ppm.<ref>{{PGCH|0064}}</ref> Bromine is classified as an [[List of extremely hazardous substances|extremely hazardous substance]] in the United States as defined in Section 302 of the U.S. [[Emergency Planning and Community Right-to-Know Act]] (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.<ref name="gov-right-know">{{Cite journal  | title = 40 C.F.R.: Appendix A to Part 355—The List of Extremely Hazardous Substances and Their Threshold Planning Quantities | url = https://edocket.access.gpo.gov/cfr_2008/julqtr/pdf/40cfr355AppA.pdf | edition = 1 July 2008 | access-date = 29 October 2011 | archive-url = https://web.archive.org/web/20120225051612/https://edocket.access.gpo.gov/cfr_2008/julqtr/pdf/40cfr355AppA.pdf | archive-date = 25 February 2012 | url-status = dead |journal =[[Federal Register]] | publisher = [[United States Government Publishing Office|Government Printing Office]] }}</ref>

== References==
{{reflist|refs=
<ref name="b1">On page 341 of his article, A. J. Balard (1826) "''Mémoire sur une substance particulière contenue dans l'eau de la mer''" [Memoir on a peculiar substance contained in sea water], ''Annales de Chimie et de Physique'', 2nd series, vol. 32, [https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA337 pp. 337–381] {{Webarchive|url=https://web.archive.org/web/20160505131749/https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA337 |date=5 May 2016 }}, Balard states that Mr. Anglada persuaded him to name his new element ''brôme''. However, on page 382 of the same journal – "''Rapport sur la Mémoire de M. Balard relatif à une nouvelle Substance''" [Report on a memoir by Mr. Balard regarding a new substance], ''Annales de Chimie et de Physique'', series 2, vol. 32, [https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA382 pp. 382–384.] {{Webarchive|url=https://web.archive.org/web/20160511032613/https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA382 |date=11 May 2016 }} – a committee of the French Academy of Sciences claimed that they had renamed the new element ''brôme''.</ref>

<ref name="Balard">{{cite journal|volume = 28|journal = Annals of Philosophy|date = 1826|pages = 381–387 and 411–426|first = Antoine|last = Balard|title = Memoir on a peculiar Substance contained in Sea Water|url = https://books.google.com/books?id=A-M4AAAAMAAJ|access-date = 5 June 2020|archive-date = 17 July 2021|archive-url = https://web.archive.org/web/20210717094507/https://books.google.com/books?id=A-M4AAAAMAAJ|url-status = live}}</ref>

<ref name="Bal1826">{{cite journal|first = A. J.|last = Balard|date = 1826|title = ''Mémoire sur une substance particulière contenue dans l'eau de la mer''|trans-title = Memoir on a peculiar substance contained in sea water|language = fr|journal = Annales de Chimie et de Physique|series = 2nd series|volume = 32|url = https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA337|pages = 337–381|access-date = 5 January 2016|archive-date = 5 May 2016|archive-url = https://web.archive.org/web/20160505131749/https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA337|url-status = live}}</ref>

<ref name="Bal1826b">{{cite journal|first = L. N.|last = Vauquelin|author2 = Thenard, L.J.|author3 = Gay-Lussac, J.L.|date = 1826|title = ''Rapport sur la Mémoire de M. Balard relatif à une nouvelle Substance''|trans-title = Report on a memoir by Mr. Balard regarding a new substance|language = fr|journal = Annales de Chimie et de Physique|series = 2nd series|volume = 32|url = https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA382|pages = 382–384|access-date = 5 January 2016|archive-date = 11 May 2016|archive-url = https://web.archive.org/web/20160511032613/https://books.google.com/books?id=vBIAAAAAMAAJ&pg=PA382|url-status = live}}</ref>

<!--<ref name=emsley>{{cite book|title=Nature's Building Blocks: An A-Z Guide to the Elements|last=Emsley|first=John|publisher=Oxford University Press|date=2001|location=Oxford, England, UK|isbn=0-19-850340-7|chapter=Bromine|pages=69–73}}</ref>-->

<ref name="L1">
{{cite book|first = Carl Jacob|last = Löwig|title = Das Brom und seine chemischen Verhältnisse |trans-title=Bromine and its chemical relationships|language = de|publisher = Carl Winter|place = Heidelberg|date = 1829|url=https://archive.org/details/bub_gb_UGFQAAAAcAAJ}}</ref>

<!-- <ref name=L2>{{cite journal|first = Carl|last = Löwig|year = 1827|title = Über Brombereitung und eine auffallende Zersetzung des Aethers durch Chlor (On the preparation of bromine and a striking decomposition of ether by chlorine)|journal = Magazine für Pharmacie|volume = 21|url = https://books.google.com/books?id=bO43AAAAMAAJ&pg=PA31|pages = 31–36}}</ref><ref name=L3>{{cite journal|first = Carl|last = Löwig|year =1828|title = Über einige Bromverbindungen und über Bromdarstellung" (On some bromine compounds and on the production of bromine)|journal = [[Annalen der Physik und Chemie]]|volume = 14|url = https://books.google.com/books?id=vG0EAAAAYAAJ&pg=PA485|pages = 485–499|doi = 10.1002/andp.18280901113|issue = 11|bibcode = 1828AnP....90..485L }}</ref><ref name=L4>{{cite journal|first = Carl|title = Ueber einige Bromverbindungen und über Bromdarstellung|last = Löwig|journal = Annalen der Physik|volume = 90|issue = 11|pages = 485–499|year = 1828|doi = 10.1002/andp.18280901113|bibcode = 1828AnP....90..485L }}</ref> -->

<ref name="Löwig">{{Cite journal|title = Nekrolog: Carl Löwig|author-link = Hans Heinrich Landolt|first = Hans Heinrich|last = Landolt|journal = [[Berichte der deutschen chemischen Gesellschaft]]|volume = 23|pages = 905–909|date = 1890|url = https://gallica.bnf.fr/ark:/12148/bpt6k907222/f920.chemindefer|doi = 10.1002/cber.18900230395|issue = 3|access-date = 24 February 2022|archive-date = 9 February 2022|archive-url = https://web.archive.org/web/20220209200258/https://gallica.bnf.fr/ark:/12148/bpt6k907222/f920.chemindefer|url-status = live}}</ref>

<ref name="o1">{{Cite journal|title = Minerals From Sea Salt|first = John A.|last = Tallmadge|author2 = Butt, John B.|author3 = Solomon Herman J.|journal = Ind. Eng. Chem.|date = 1964|volume = 56 |pages = 44–65| doi = 10.1021/ie50655a008|issue = 7}}</ref>

<ref name="o2">{{Cite journal|journal = Clinics in Dermatology|volume = 14 |date = 1996|pages = 659–664|doi = 10.1016/S0738-081X(96)00101-0|pmid = 8960809 |title = Climatotherapy at the Dead Sea in Jordan|first = Oumeish Youssef|last = Oumeish|issue = 6}}</ref>

<ref name="o3">{{Cite journal|journal = Hydrological Processes|volume = 14|issue = 1|pages = 145–154| date = 2008|title = The water balance of the Dead Sea: an integrated approach|first = Radwan A.|last = Al-Weshah|doi = 10.1002/(SICI)1099-1085(200001)14:1<145::AID-HYP916>3.0.CO;2-N|bibcode = 2000HyPr...14..145A }}</ref>

<!--<ref name="USGSCR2007">{{cite web|url = https://minerals.usgs.gov/minerals/pubs/commodity/bromine/mcs-2008-bromi.pdf|title = Commodity Report 2007: Bromine|first = Phyllis A.|last=Lyday| publisher = United States Geological Survey|access-date = 2008-09-03}}</ref>-->

<ref name="USGSYB2007">{{cite web|date=May 2010|url = https://minerals.usgs.gov/minerals/pubs/commodity/bromine/myb1-2006-bromi.pdf|title = Mineral Yearbook 2007: Bromine|first = Phyllis A.|last = Lyday|publisher = United States Geological Survey|access-date = 2008-09-03|archive-date = 19 October 2017|archive-url = https://web.archive.org/web/20171019221335/https://minerals.usgs.gov/minerals/pubs/commodity/bromine/myb1-2006-bromi.pdf|url-status = live}}</ref>

<ref name="Wisniak">{{Cite journal|first = Jaime|last = Wisniak|url = https://revista.cnic.edu.cu/revistaCQ/sites/default/files/articulos/CQ-2004-1-035-040.pdf|title = Antoine-Jerôme Balard. The discoverer of bromine|journal = Revista CENIC Ciencias Químicas|volume = 35|issue = 1|pages = 35–40|date = 2004|access-date = 24 February 2022|archive-date = 25 March 2016|archive-url = https://web.archive.org/web/20160325093029/http://revista.cnic.edu.cu/revistaCQ/sites/default/files/articulos/CQ-2004-1-035-040.pdf|url-status = live}}</ref>}}

== General and cited references ==
* {{Greenwood&Earnshaw2nd}}

{{Bromine compounds}}
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[[Category:Bromine| ]]
[[Category:Chemical elements]]
[[Category:Diatomic nonmetals]]
[[Category:Gases with color]]
[[Category:Halogens]]
[[Category:Oxidizing agents]]
[[Category:Reactive nonmetals]]