Aluminium chloride, also known as aluminium trichloride, is an inorganic compound with the formula AlCl3. It forms a hexahydrate with the formula [Al(H2O)6]Cl3, containing six water molecules of hydration. Both the anhydrous form and the hexahydrate are colourless crystals, but samples are often contaminated with iron(III) chloride, giving them a yellow colour.
The anhydrous form is important commercially. It has a low melting and boiling point. It is mainly produced and consumed in the production of aluminium, but large amounts are also used in other areas of the chemical industry.[7] 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.
Structure
Illustration of structures of aluminium chloride
Anhydrous
AlCl3 adopts three structures, depending on the temperature and the state (solid, liquid, gas). Solid AlCl3 has a sheet-like layered structure with cubic close-packed chloride ions. In this framework, the Al centres exhibit octahedral coordination geometry.[8]Yttrium(III) chloride adopts the same structure, as do a range of other compounds. When aluminium trichloride is in its melted state, it exists as the dimerAl2Cl6, 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 planarAlCl3monomer, which is structurally analogous to BF3. The melt conductselectricity poorly,[9] unlike more ionichalides such as sodium chloride.
Aluminium chloride monomer belongs to the point group D3h in its monomeric form and D2h in its dimeric form.
Hexahydrate
The hexahydrate consists of octahedral[Al(H2O)6]3+cation centers and chloride anions (Cl−) as counterions. Hydrogen bonds link the cation and anions.[10]
The hydrated form of aluminium chloride has an octahedral molecular geometry, with the central aluminium ion surrounded by six water ligand molecules. Being coordinatively saturated, the hydrate is of little value as a catalyst in Friedel-Crafts alkylation and related reactions.
The alkylation reaction is more widely used than the acylation reaction, although its practice is more technically demanding. 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.[12] Detailed procedures are available for alkylation[13] and acylation[14][15] of arenes.
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, the use of aluminium chloride has often been displaced by zeolites.[7]
Anhydrous 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− ligands are displaced with H2O molecules to form the hexahydrate [Al(H2O)6]Cl3. The anhydrous phase cannot be regained on heating the hexahydrate. Instead HCl is lost leaving aluminium hydroxide or alumina (aluminium oxide):
Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride at temperatures between 650 and 750 °C (1,202 and 1,382 °F).[9]
In the US in 1993, approximately 21,000 tons were produced, not counting the amounts consumed in the production of aluminium.[7]
Hydrated aluminium trichloride is prepared by dissolving aluminium oxides in hydrochloric acid. Metallic aluminium 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)6]Cl3 → Al(OH)3 + 3 HCl + 3 H2O
Aluminium also forms a lower chloride, aluminium(I) chloride (AlCl), but this is very unstable and only known in the vapour phase.[9]
Natural occurrence
Anhydrous aluminium chloride is not found as a mineral. The hexahydrate, however, is known as the rare mineral chloraluminite.[25] A more complex, basic and hydrated aluminium chloride mineral is cadwaladerite.[26][25]
Safety
Anhydrous AlCl3 reacts vigorously with bases, so suitable precautions are required.
It can cause irritation to the eyes, skin, and the respiratory system if inhaled or on contact.[27]
^Wells AF (1984). Structural Inorganic Chemistry. Oxford, United Kingdom.: Oxford Press. ISBN0198553706. In contrast, AlBr3 has a more molecular structure, with the Al3+ centers occupying adjacent tetrahedral holes of the close-packed framework of Br− ions.
^Andress KR, Carpenter C (1934). "Kristallhydrate II. Die Struktur von Chromchlorid- und Aluminiumchloridhexahydrat". Zeitschrift für Kristallographie – Crystalline Materials. 87. doi:10.1524/zkri.1934.87.1.446. S2CID263857074.
^ abcOlah GA, ed. (1963). Friedel-Crafts and Related Reactions. Vol. 1. New York City: Interscience.
^Nenitzescu CD, Cantuniari IP (1933). "Durch Aluminiumchlorid Katalysierte Reaktion, VI. Mitteil.: Die Umlagerung des Cyclohexans in Metyl-cyclopentan". Berichte der Deutschen Chemischen Gesellschaft (A and B Series). 66 (8): 1097–1100. doi:10.1002/cber.19330660817. ISSN1099-0682.
^Galatsis P (1999). "Aluminum Chloride". In Reich HJ, Rigby JH (eds.). Acidic and Basic Reagents. Handbook of Reagents for Organic Synthesis. New York City: Wiley. pp. 12–15. ISBN978-0-471-97925-8.
^McConaghy JR, Fosselman D (June 2018). "Hyperhidrosis: Management Options". American Family Physician. 97 (11): 729–734. PMID30215934.
^Nawrocki S, Cha J (September 2019). "The etiology, diagnosis, and management of hyperhidrosis: A comprehensive review: Therapeutic options". Journal of the American Academy of Dermatology. 81 (3): 669–680. doi:10.1016/j.jaad.2018.11.066. PMID30710603.
