Aluminium chloride (AlCl3), also known as aluminium trichloride, is the main compound of aluminium and chlorine. It is white, but samples are often contaminated with iron(III) chloride, giving it a yellow color. The solid has a low melting and boiling point. It is mainly produced and consumed in the production of aluminium metal, but large amounts are also used in other areas of the chemical industry. The compound is often cited as a Lewis acid. It is an example of an inorganic compound that reversibly changes from a polymer to a monomer at mild temperature.
3D model (JSmol)
CompTox Dashboard (EPA)
|Molar mass||133.341 g/mol (anhydrous)|
241.432 g/mol (hexahydrate)
|Appearance||white or pale yellow solid,|
|Density||2.48 g/cm3 (anhydrous)|
2.398 g/cm3 (hexahydrate)
|Melting point|| 192.6 °C (378.7 °F; 465.8 K) |
100 °C (212 °F; 373 K)
|Boiling point||180 °C (356 °F; 453 K) (sublimes)|
|439 g/l (0 °C)|
449 g/l (10 °C)
458 g/l (20 °C)
466 g/l (30 °C)
473 g/l (40 °C)
481 g/l (60 °C)
486 g/l (80 °C)
490 g/l (100 °C)
|Solubility||soluble in hydrogen chloride, ethanol, chloroform, carbon tetrachloride |
slightly soluble in benzene
|Vapor pressure||133.3 Pa (99 °C)|
13.3 kPa (151 °C)
|Viscosity||0.35 cP (197 °C)|
0.26 cP (237 °C)
|C12/m1, No. 12|
a = 0.591 nm, b = 0.591 nm, c = 1.752 nm
Lattice volume (V)
Formula units (Z)
Heat capacity (C)
Std enthalpy of
Gibbs free energy (ΔfG˚)
|Safety data sheet||See: data page|
|GHS Signal word||Danger|
|P280, P310, P305+351+338|
|NFPA 704 (fire diamond)|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
380 mg/kg, rat (oral)
3311 mg/kg, rat (oral)
|NIOSH (US health exposure limits):|
IDLH (Immediate danger)
Related Lewis acids
|Supplementary data page|
|Refractive index (n),|
Dielectric constant (εr), etc.
|UV, IR, NMR, MS|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Anhydrous aluminium trichlorideEdit
AlCl3 is probably the most commonly used Lewis acid and also one of the most powerful. It finds application in the chemical industry as a catalyst for Friedel–Crafts reactions, both acylations and alkylations. Important products are detergents and ethylbenzene. It also finds use in polymerization and isomerization reactions of hydrocarbons.
The Friedel–Crafts reaction is the major use for aluminium chloride, for example in the preparation of anthraquinone (for the dyestuffs industry) from benzene and phosgene. In the general Friedel–Crafts reaction, an acyl chloride or alkyl halide reacts with an aromatic system as shown:
The alkylation reaction is more widely used than the acylation reaction, although its practice is more technically demanding because the reaction is more sluggish. For both reactions, the aluminium chloride, as well as other materials and the equipment, should be dry, although a trace of moisture is necessary for the reaction to proceed. A general problem with the Friedel–Crafts reaction is that the aluminium chloride catalyst sometimes is required in full stoichiometric quantities, because it complexes strongly with the products. This complication sometimes generates a large amount of corrosive waste. For these and similar reasons, more recyclable or environmentally benign catalysts have been sought. Thus, the use of aluminium chloride in some applications is being displaced by zeolites.
Aluminium chloride can also be used to introduce aldehyde groups onto aromatic rings, for example via the Gattermann-Koch reaction which uses carbon monoxide, hydrogen chloride and a copper(I) chloride co-catalyst.
Aluminium chloride finds a wide variety of other applications in organic chemistry. For example, it can catalyse the "ene reaction", such as the addition of 3-buten-2-one (methyl vinyl ketone) to carvone:
AlCl3 is also widely used for polymerization and isomerization reactions of hydrocarbons. Important examples include the manufacture of ethylbenzene, which used to make styrene and thus polystyrene, and also production of dodecylbenzene, which is used for making detergents.
Aluminium chloride combined with aluminium in the presence of an arene can be used to synthesize bis(arene) metal complexes, e.g. bis(benzene)chromium, from certain metal halides via the so-called Fischer-Hafner synthesis.
Hydrated aluminium chloridesEdit
The dihydrate has few applications, but aluminium chlorohydrate is a common component in antiperspirants at low concentrations. Hyperhidrosis sufferers need a much higher concentration (12% or higher), sold under such brand names as Driclor.
AlCl3 adopts three different structures, depending on the temperature and the state (solid, liquid, gas). Solid AlCl3 is a sheet-like layered cubic close packed layers. In this framework, the Al centres exhibit octahedral coordination geometry. When aluminium trichloride is in its melted state, it exists as the dimer Al2Cl6, with tetracoordinate aluminium. This change in structure is related to the lower density of the liquid phase (1.78 g/cm3) versus solid aluminium trichloride (2.48 g/cm3). Al2Cl6 dimers are also found in the vapour phase. At higher temperatures, the Al2Cl6 dimers dissociate into trigonal planar AlCl3, which is structurally analogous to BF3. The melt conducts electricity poorly, unlike more-ionic halides such as sodium chloride.
The hexahydrate consists of octahedral [Al(H2O)6]3+ centers and chloride counterions. Hydrogen bonds link the cation and anions. The hydrated form of aluminium chloride has an octahedral molecular geometry, with the central aluminum ion surrounded by six water ligand molecules. This means that the hydrated form cannot act as a Lewis acid since it cannot accept electron pairs, and thus this cannot be used as a catalyst in Friedel-Crafts alkylation of aromatic compounds.
Anhydrous aluminium chloride is a powerful Lewis acid, capable of forming Lewis acid-base adducts with even weak Lewis bases such as benzophenone and mesitylene. It forms tetrachloroaluminate (AlCl4−) in the presence of chloride ions.
Reactions with waterEdit
Aluminium chloride is hygroscopic, having a very pronounced affinity for water. It fumes in moist air and hisses when mixed with liquid water as the Cl− ions are displaced with H2O molecules in the lattice to form the hexahydrate [Al(H2O)6]Cl3 (also white to yellowish in color). The anhydrous phase cannot be regained on heating as HCl is lost leaving aluminium hydroxide or alumina (aluminium oxide):
- Al(H2O)6Cl3 → Al(OH)3 + 3 HCl + 3 H2O
On strong heating (~400 °C), aluminium oxide is formed from the aluminium hydroxide:
- 2 Al(OH)3 → Al2O3 + 3 H2O
Aqueous solutions of AlCl3 are ionic and thus conduct electricity well. Such solutions are found to be acidic, indicative of partial hydrolysis of the Al3+ ion. The reactions can be described (simplified) as
- [Al(H2O)6]3+(aq) ⇌ [Al(OH)(H2O)5]2+(aq) + H+(aq)
- AlCl3 + 3 NaOH → [Al(OH)3] + 3 NaCl
- 2 Al + 3 Cl2 → 2 AlCl3
- 2 Al + 6 HCl → 2 AlCl3 + 3 H2
- 2 Al + 3 CuCl2 → 2 AlCl3 + 3 Cu
In the US in 1993, approximately 21,000 tons were produced, not counting the amounts consumed in the production of aluminium.
Hydrated aluminium trichloride is prepared by dissolving aluminium oxides in hydrochloric acid. Metallic aluminum also readily dissolves in hydrochloric acid ─ releasing hydrogen gas and generating considerable heat. Heating this solid does not produce anhydrous aluminium trichloride, the hexahydrate decomposes to aluminium hydroxide when heated:
- Al(H2O)6Cl3 → Al(OH)3 + 3 HCl + 3 H2O
Symmetry and dipole momentEdit
Aluminium chloride belongs to the point group D3h in its monomeric form and D2h in its dimeric form. Both forms of aluminium chloride, however, do not possess a dipole moment because the bond dipole moments cancel each other out.
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