Aluminium bromide


Aluminium bromide
Ball and stick model of dimeric aluminium bromide
Aluminium bromide from 1967 In large ampoules
Preferred IUPAC name
Aluminium bromide
Other names
Aluminic bromide

Aluminium(III) bromide

Aluminium tribromide
  • 7727-15-3 checkY
  • 7784-11-4 hexahydrate checkY
3D model (JSmol)
  • Interactive image
  • dimer: Interactive image
  • 22818 checkY
  • 9040513 hemi(acetyl bromide) checkY
  • 9499890 ethanethiol checkY
ECHA InfoCard 100.028.891 Edit this at Wikidata
EC Number
  • 231-779-7
  • 24409
  • 11062022 tertikis(tetrachloromethane)
  • 10865226 hemi(acetyl bromide)
  • 11324936 ethanethiol
  • 6093832 tris(pyridine)
RTECS number
  • BD0350000
  • IY20HBK5LK checkY
UN number 1725
  • DTXSID2064783 Edit this at Wikidata
  • InChI=1S/Al.3BrH/h;3*1H/q+3;;;/p-3 checkY
  • InChI=1/Al.3BrH/h;3*1H/q+3;;;/p-3
  • Br[Al](Br)Br
  • dimer: Br[Al-]1(Br)[Br+][Al-]([Br+]1)(Br)Br
AlBr3·6H2O (hexahydrate)
Molar mass 266.694 g/mol (anhydrous)
374.785 g/mol (hexahydrate)[1]
Appearance white to pale yellow powder[1]
Odor pungent
Density 3.2 g/cm3 (anhydrous)
2.54 g/cm3 (hexahydrate)[1]
Melting point 97.5 °C (anhydrous)
93 °C (hexahydrate)[1]
Boiling point 255 (anhydrous)[1]
very soluble, partially hydrolyses indicated by a fuming solution and an optional appearance of white precipitate
Solubility slightly soluble in methanol, diethyl ether, acetone
Monoclinic, mP16 (anhydrous)
P21/c, No. 14
a = 0.7512 nm, b = 0.7091 nm, c = 1.0289 nm
α = 90°, β = 96.44°, γ = 90°
100.6 J/(mol·K)
180.2 J/(mol·K)
-572.5 kJ/mol
GHS pictograms GHS05: CorrosiveGHS07: Exclamation mark
GHS Signal word Danger
H302, H314
P260, P264, P270, P280, P301+P312, P301+P330+P331, P303+P361+P353, P304+P340, P305+P351+P338, P310, P321, P330, P363, P405, P501
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
1598 mg/kg (oral, rat)
Related compounds
Other anions
aluminium trichloride
aluminium triiodide
Other cations
boron tribromide
Related compounds
iron(III) bromide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Aluminium bromide is any chemical compound with the empirical formula AlBrx. Aluminium tribromide is the most common form of aluminium bromide.[3] It is a colorless, sublimable hygroscopic solid; hence old samples tend to be hydrated, mostly as aluminium tribromide hexahydrate (AlBr3·6H2O).


The dimeric form of aluminium tribromide (Al2Br6) predominates in the solid state, in solutions in noncoordinating solvents (e.g. CS2), in the melt, and in the gas phase. Only at high temperatures do these dimers break up into monomers:

Al2Br6 → 2 AlBr3 ΔH°diss = 59 kJ/mol

The species aluminium monobromide forms from the reaction of HBr with Al metal at high temperature. It disproportionates near room temperature:

6/n "[AlBr]n" → Al2Br6 + 4 Al

This reaction is reversed at temperatures higher than 1000 °C. Aluminium monobromide has been crystallographically characterized in the form the tetrameric adduct Al4Br4(NEt3)4 (Et = C2H5). This species is electronically related to cyclobutane. Theory suggest that the diatomic aluminium monobromide condenses to a dimer and then a tetrahedral cluster Al4Br4, akin to the analogous boron compound.[4]

Al2Br6 consists of two AlBr4 tetrahedra that share a common edge. The molecular symmetry is D2h.

The monomer AlBr3, observed only in the vapor, can be described as trigonal planar, D3h point group. The atomic hybridization of aluminium is often described as sp2. The Br-Al-Br bond angles are 120 °.


Experiment showing synthesis of aluminum bromide from the elements.

By far the most common form of aluminium bromide is Al2Br6. This species exists as hygroscopic colorless solid at standard conditions. Typical impure samples are yellowish or even red-brown due to the presence of iron-containing impurities. It is prepared by the reaction of HBr with Al:

2 Al + 6 HBr → Al2Br6 + 3 H2

Alternatively, the direct bromination occurs also:

2 Al + 3 Br2 → Al2Br6


A demonstration of the reaction of the exothermic reaction of the strong Lewis acid (Al2Br6) and strong Lewis base (H2O).

Al2Br6 dissociates readily to give the strong Lewis acid, AlBr3. Regarding the tendency of Al2Br6 to dimerize, it is common for heavier main group halides to exist as aggregates larger than implied by their empirical formulae. Lighter main group halides such as boron tribromide do not show this tendency, in part due to the smaller size of the central atom.

Consistent with its Lewis acidic character, Al2Br6 is hydrolyzed by water with evolution of HBr and formation of Al-OH-Br species. Similarly, it also reacts quickly with alcohols and carboxylic acids, although less vigorously than with water. With simple Lewis bases (L), Al2Br6 forms adducts, such as AlBr3L.

Aluminium tribromide reacts with carbon tetrachloride at 100 °C to form carbon tetrabromide:

4 AlBr3 + 3 CCl4 → 4 AlCl3 + 3 CBr4

and with phosgene yields carbonyl bromide and aluminium chlorobromide:[citation needed]

AlBr3 + COCl2 → COBr2 + AlCl2Br

Al2Br6 is used as a catalyst for the Friedel-Crafts alkylation reaction.[3] Related Lewis acid-promoted reactions include as epoxide ring openings and decomplexation of dienes from iron carbonyls. It is a stronger Lewis acid than the more common Al2Cl6.


Aluminium tribromide is a highly reactive material.[5]


  1. ^ a b c d e f Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.45. ISBN 1-4398-5511-0.
  2. ^ Troyanov, Sergey I.; Krahl, Thoralf; Kemnitz, Erhard (2004). "Crystal structures of GaX3(X= Cl, Br, I) and AlI3". Zeitschrift für Kristallographie. 219 (2–2004): 88–92. doi:10.1524/zkri. S2CID 101603507.
  3. ^ a b Paquette, Leo A. (2001). Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X. ISBN 0471936235.
  4. ^ Dohmeier, Carsten; Loos, Dagmar; Schnöckel, Hansgeorg (1996). "Aluminum(I) and Gallium(I) Compounds: Syntheses, Structures, and Reactions". Angewandte Chemie International Edition in English. 35 (2): 129. doi:10.1002/anie.199601291.
  5. ^ Renfew, Malcom M. (1991). "Hazardous laboratory chemicals: Disposal guide (Armour, M.A.)". Journal of Chemical Education. 68 (9): A232. Bibcode:1991JChEd..68Q.232R. doi:10.1021/ed068pA232.2.