Editing Bromine (section)

===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>