 Aluminium trichloride hexahydrate, pure (top), and contaminated withiron(III) chloride (bottom) | |||
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| Names | |||
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| IUPAC name Aluminium chloride | |||
| Other names Aluminium(III) chloride Aluminium trichloride Trichloroaluminum | |||
| Identifiers | |||
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3D model (JSmol) |
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| ChEBI | |||
| ChemSpider |
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| ECHA InfoCard | 100.028.371 | ||
| EC Number |
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| 1876 | |||
| RTECS number |
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| Properties | |||
| AlCl3 | |||
| Molar mass |
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| Appearance | Colourless crystals,hygroscopic | ||
| Density |
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| Melting point | |||
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| Solubility |
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| Vapor pressure |
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| Viscosity |
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| Structure | |||
| Monoclinic,mS16 | |||
| C12/m1, No. 12[3] | |||
a = 0.591 nm,b = 0.591 nm,c = 1.752 nm[3] | |||
Lattice volume (V) | 0.52996 nm3 | ||
Formula units (Z) | 6 | ||
| Octahedral (solid) Tetrahedral (liquid) | |||
| Trigonal planar (monomeric vapour) | |||
| Thermochemistry | |||
| 91.1 J/(mol·K) (solid)[4] 125.5 J/(mol K) (liquid) 82.46 J/(mol K) (gas)[5] | |||
Std molar entropy(S⦵298) | 109.3 J/(mol·K) (solid)[4] 172.91 J/(mol·K) (liquid) 314.44 J/(mol·K) (gas)[5] | ||
Std enthalpy of formation(ΔfH⦵298) | −704.2 kJ/mol (solid)[4] −674.80 kJ/mol (liquid) -584.59 kJ/mol (gas)[5] | ||
Gibbs free energy(ΔfG⦵) | −628.8 kJ/mol[4] | ||
| Pharmacology | |||
| D10AX01 (WHO) | |||
| Hazards | |||
| GHS labelling:[7] | |||
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| Danger | |||
| H314 | |||
| P260,P280,P301+P330+P331,P303+P361+P353,P305+P351+P338+P310,P310 | |||
| NFPA 704 (fire diamond) | |||
| Lethal dose or concentration (LD, LC): | |||
LD50 (median dose) | 380 mg/kg, rat (oral, anhydrous) 3311 mg/kg, rat (oral, hexahydrate) | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible) | None[6] | ||
REL (Recommended) | 2 mg/m3[6] | ||
IDLH (Immediate danger) | N.D.[6] | ||
| Related compounds | |||
Otheranions | |||
Othercations | |||
RelatedLewis acids | |||
Related compounds | |||
| Supplementary data page | |||
| Aluminium chloride (data page) | |||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Aluminium chloride, also known asaluminium trichloride, is aninorganic compound with the formulaAlCl3. It forms ahexahydrate with the formula[Al(H2O)6]Cl3, containing sixwater molecules of hydration. Both the anhydrous form and the hexahydrate are colourless crystals, but samples are often contaminated withiron(III) chloride, giving them a yellow colour.
The anhydrous form is commercially important. 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.[8] The compound is often cited as aLewis acid. It is aninorganic compound thatreversibly changes from apolymer to amonomer at mild temperature.
AlCl3 adopts three structures, depending on thetemperature and thestate (solid, liquid, gas). SolidAlCl3 has a sheet-like layered structure with cubic close-packed chloride ions. In this framework, the Al centres exhibitoctahedral coordination geometry.[9]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 thedimer (Al2Cl6point group D2h), withtetracoordinate 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 thevapour phase. At higher temperatures, theAl2Cl6 dimersdissociate intotrigonal planarAlCl3monomer (point group D3h), which is structurally analogous toBF3. The meltconductselectricity poorly,[10] unlike moreionichalides such assodium chloride.
The hexahydrate consists ofoctahedral[Al(H2O)6]3+cation centers and chlorideanions (Cl−) ascounterions.Hydrogen bonds link the cation and anions.[11]The hydrated form of aluminium chloride has an octahedral molecular geometry, with the central aluminium ion surrounded by sixwater ligand molecules. Being coordinatively saturated, the hydrate is of little value as acatalyst inFriedel-Crafts alkylation and related reactions.
AlCl3 is a common Lewis-acidcatalyst forFriedel-Crafts reactions, both acylations and alkylations.[12] These types of reactions are the major use for aluminium chloride, for example, in the preparation ofanthraquinone (used in thedyestuffs industry) frombenzene andphosgene.[10] In the general Friedel-Crafts reaction, anacyl chloride oralkyl halide reacts with anaromatic system as shown:[12]
Thealkylation reaction is more widely used than theacylation 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.[13] Detailed procedures are available for alkylation[14] and acylation[15][16] of arenes.
A general problem with the Friedel-Crafts reaction is that the aluminium chloride catalyst sometimes is required in fullstoichiometric quantities, because itcomplexes strongly with the products. This complication always generates a large amount ofcorrosive waste. For these and similar reasons, the use of aluminium chloride has rarely been displaced byzeolites.[8]
Aluminium chloride can also be used to introducealdehyde groups ontoaromatic rings, for example via theGattermann-Koch reaction which usescarbon monoxide,hydrogen chloride and acopper(I) chlorideco-catalyst.[17]
Aluminium chloride finds a wide variety of other applications inorganic chemistry.[18] For example, it can catalyse theene reaction, such as the addition of3-buten-2-one (methyl vinyl ketone) tocarvone:[19]
It is used to induce a variety of hydrocarbon couplings and rearrangements.[20][21]
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 theFischer–Hafner synthesis.Dichlorophenylphosphine is prepared by reaction ofbenzene andphosphorus trichloride catalyzed by aluminium chloride.[22]
| Clinical data | |
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| AHFS/Drugs.com | Monograph |
| License data | |
| Routes of administration | Topical |
| ATC code |
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| Identifiers | |
| CompTox Dashboard(EPA) | |
| ECHA InfoCard | 100.028.371 |
| Data page | |
| Aluminium chloride (data page) | |
Topical aluminum chloride hexahydrate is used for the treatment ofhyperhidrosis (excessive sweating).[23][24][25]
Anhydrous aluminium chloride is a powerfulLewis acid, capable of forming Lewis acid-baseadducts with even weakLewis bases such asbenzophenone andmesitylene.[12] It formstetrachloroaluminate ([AlCl4]−) in the presence ofchloride ions.
Aluminium chloride reacts withcalcium andmagnesium hydrides intetrahydrofuran forming tetrahydroaluminates.[citation needed]
Anhydrous aluminium chloride ishygroscopic, having a very pronounced affinity for water. It fumes in moist air and hisses when mixed with liquid water as theCl− ligands are displaced withH2O 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):
Likemetal aquo complexes,aqueousAlCl3 is acidic owing to the ionization of theaquo ligands:
Aqueous solutions behave similarly to otheraluminiumsalts containing hydratedAl3+ions, giving a gelatinousprecipitate ofaluminium hydroxide upon reaction with dilutesodium hydroxide:
Aluminium chloride is manufactured on a large scale by theexothermic reaction ofaluminium metal with chlorine orhydrogen chloride at temperatures between 650 and 750 °C (1,202 and 1,382 °F).[10]
Aluminium chloride may be formed via asingle displacement reaction betweencopper(II) chloride and aluminium.
In the US in 1993, approximately 21,000 tons were produced, not counting the amounts consumed in the production of aluminium.[8]
Hydrated aluminium trichloride is prepared by dissolving aluminium oxides inhydrochloric 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 toaluminium hydroxide when heated:
Aluminium also forms a lowerchloride,aluminium(I) chloride (AlCl), but this is very unstable and only known in the vapour phase.[10]
Anhydrous aluminium chloride is not found as a mineral. The hexahydrate, however, is known as the rare mineral chloraluminite.[26] A more complex, basic and hydrated aluminium chloride mineral iscadwaladerite.[27][26]
Aluminium chlorides were known in the 18th century as muriate of alumina, marine alum, argillaceous marine salt,[28] muriated clay.[29] It was first chemically studied in the 1830s.[30]
AnhydrousAlCl3 is strongly corrosive and releases hydrochloric acid in contact with water.[8]
In contrast,AlBr3 has a more molecular structure, with theAl3+ centers occupying adjacent tetrahedral holes of the close-packed framework ofBr− ions.
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