Bromine compounds

Bromine compounds are compounds containing the element bromine (Br). These compounds usually form the -1, +1, +3 and +5 oxidation states. 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 potentials of the X₂/X⁻ couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 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.^[1]

Hydrogen bromide

The simplest compound of bromine is

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 gases, with the exception of xenon in the very unstable XeBr₂; 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.)^[4]

Bromination of metals with Br₂ tends to yield lower oxidation states than chlorination with Cl₂ 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,

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:^[4]

3 WBr₅ + Al thermal gradient→475 °C → 240 °C 3 WBr₄ + AlBr₃

EuBr₃ + 1/2 H₂ → EuBr₂ + HBr

2 TaBr₄ 500 °C→  TaBr₃ + TaBr₅

Most of the bromides of the pre-transition metals (groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states) are mostly ionic, while nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine.^[4]

Bromine halides

The halogens form many binary,

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.^[5] 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.^[4] Bromine monofluoride in ethanol readily leads to the monobromination of the aromatic compounds PhX (para-bromination occurs for X = Me, Bu^t, OMe, Br; meta-bromination occurs for the deactivating X = –CO₂Et, –CHO, –NO₂); this is due to heterolytic fission of the Br–F bond, leading to rapid electrophilic bromination by Br⁺.^[4]

At room temperature, bromine trifluoride (BrF₃) 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 KBrF₄ and BrF₂SbF₆ remain reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to form BrF⁺
₂ and BrF⁻
₄ and thus conducts electricity.^[6]

Bromine pentafluoride (BrF₅) 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 °C, and on a small scale by the fluorination of potassium bromide at 25 °C. It also reacts violently with water and is a very strong fluorinating agent, although chlorine trifluoride is still stronger.^[7]

Polybromine compounds

Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such as peroxydisulfuryl fluoride (S₂O₆F₂) can oxidise it to form the cherry-red Br⁺
₂ cation. A few other bromine cations are known, namely the brown Br⁺
₃ and dark brown Br⁺
₅.^[8] The tribromide anion, Br⁻
₃, has also been characterised; it is analogous to triiodide.^[5] Neutral Br
₂ molecules are incorporated into an adduct with the anionic copper-bromine complex Cu₂Br₆^2-.^[9]

Bromine oxides and oxoacids

So-called "bromine dioxide", a pale yellow crystalline solid, may be better formulated as bromine perbromate, BrOBrO₃. It is thermally unstable above −40 °C, violently decomposing to its elements at 0 °C. Dibromine trioxide, syn-BrOBrO₂, is also known; it is the anhydride of hypobromous acid and bromic acid. It is an orange crystalline solid which decomposes above −40 °C; if heated too rapidly, it explodes around 0 °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.^[12]

The four

Hypobromous acid is unstable to disproportionation. The hypobromite ions thus formed disproportionate readily to give bromide and bromate:^[10]

3 BrO− ⇌ 2 Br− + BrO− 3

K = 1015

Bromous acids and

Like the other carbon–halogen bonds, the C–Br bond is a common functional group that forms part of core

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. The principal reactions for organobromides include dehydrobromination, Grignard reactions, reductive coupling, and nucleophilic substitution.^[15]

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⁺, a potent electrophile. The enzyme

An old qualitative test for the presence of the alkene functional group is that alkenes turn brown aqueous bromine solutions colourless, forming a bromohydrin with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilic bromonium intermediate. This is an example of a halogen addition reaction.^[18]

See also

• Category:Bromine compounds
• Organobromine compounds

References