Editing Bromine (section)

==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>