Editing Bromine (section)

===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>