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Overview Aluminium chloride (AlCl3) is a compound of aluminium and chlorine. The anhydrous material has a very interesting structure; despite being the halide of a highly electropositive metal, its bonding is principally covalent. This is seen in the fact that it has a low melting and boiling point (it sublimes at 178 °C), and it conducts electricity poorly in the liquid state[1], unlike ionic halides such as sodium chloride. It exists in the solid state as a six-coordinate layer lattice. This melts to a four-coordinate dimer, Al2Cl6 which can vapourise, but at higher temperatures this dissociates into a simple AlCl3 species analogous to BF3.
Aluminum chloride is highly deliquescent, and it can explode in contact with water because of the high hydration. It partially hydrolyses with H2O, forming some hydrogen chloride and/or hydrochloric acid. Aqueous solutions of AlCl3 are fully ionized and conducts electricity well. Such solutions are found to be acidic, indicating that partial hydrolysis of the Al3+ ion is occurring. This can be described (simplified) as:
AlCl3 is probably the most commonly used Lewis acid and also one of the most powerful. It finds widespread application in the chemical industry as a catalyst for Friedel_Crafts reactions, both acylations and alkylations. It also finds use in polymerization and isomerization reactions of hydrocarbons.
Aluminium also forms a lower chloride, aluminum(I) chloride (AlCl), but this is very unstable and only known in the vapour phase[1].
Chemical Properties Aluminium chloride is a powerful Lewis acid, capable of forming stable Lewis acid-base adducts with even weak Lewis bases such as benzophenone or mesitylene[3]. Not surprisingly it forms AlCl4_ in the presence of chloride ion.
In water, partial hydrolysis forms HCl gas or H3O+, as described in the overview above. Aqueous solutions behave similarly to other aluminium salts containing hydrated Al3+ ions- for example giving a gelatinous precipitate of aluminium hydroxide upon reaction with the correct quantity of aqueous sodium hydroxide:
AlCl3(aq) + 3 NaOH(aq) → Al(OH)3(s) + 3 NaCl(aq)
Preparation Aluminium chloride is manufactured by the exothermic reaction of the elements[1], aluminium and chlorine. It is commercially available in large quantities.
Uses The Friedel-Crafts reaction[3] is probably the major use for aluminium chloride, for example in the preparation of anthraquinone (for the dyestuffs industry) from benzene and phosgene.[1] In the general Friedel-Crafts reaction[3] an acyl chloride or alkyl halide reacts with an aromatic system as shown:
With benzene derivatives, the major product is the para isomer. The alkylation reaction has many associated problems (see Friedel-Crafts), so it is less widely used than the acylation reaction. For both reactions the aluminum chloride (and other materials and the equipment) must be moderately dry, although a trace is moisture is necessary for the reaction to proceed. A general problem with the Friedel-Crafts reaction is that the aluminium chloride "catalyst" needs to be present in full stoichiometric quantities in order for the reaction to go to completion, because it complexes strongly with the products (see chemical properties above). This makes it very difficult to recycle, so it must be destroyed after use, generating a large amount of corrosive waste. For this reason chemists are examining the use of more environmentally benign catalysts such as ytterbium(III) trifluoromethanesulfonate or dysprosium(III) trifluoromethanesulfonate, which can be [[recycling|recycled].
Aluminium chloride can also be used to introduce aldehyde groups onto aromatic rings, for example via the Gatterman-Koch reaction which uses carbon monoxide, hydrogen chloride and a copper(I) chloride co_catalyst)[4]:
Aluminium chloride finds a wide variety of other applications in organic chemistry.[5] For example, it can catalyse the "ene reaction", such as the addition of 3-buten-2-one (methyl vinyl ketone) to carvone[6]:
AlCl3 is also widely used for polymerization and isomerization reactions of hydrocarbons. Important examples[1] include the manufacture of ethylbenzene (used to make styrene and thus polystyrene), and also production of dodecylbenzene (used for making detergents).
Precautions Avoid bringing anhydrous AlCl3 in contact with water or bases, or an explosive reaction may result. Gloves and safety goggles should be worn, along with a face shield for larger amounts. The material should be handled in a fume cupboard or chemical hood. When handled in moist air, AlCl3 rapidly absorbs moisture to become a highly acidic and sticky "goo", and it rapidly attacks many materials such as stainless steel and rubber.
Suppliers/Manufacturers Fisher: https://www1.fishersci.com/index.jsp VWR: http://www.vwr.com/index.htm Alfa: http://www.alfa.com/alf/index.htm Aldrich: http://www.sigmaaldrich.com
References - N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, UK, 1984.
- Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- G. A. Olah (ed.), Friedel-Crafts and Related Reactions, Vol. 1, Interscience, New York, 1963.
- L. G. Wade, Organic Chemistry, 5th edition, Prentice Hall, Upper Saddle River, New Jersey, USA, 2003.
- P. Galatsis, in: Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (H. J. Reich, J. H. Rigby, eds.), pp12-15, Wiley, New York, 1999.
- B. B. Snider, Accounts of Chemical Research 13, 426 (1980).
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