Aluminium chloride

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Aluminium chloride
Aluminium chloride
Aluminium trichloride dimer
Identifiers
CAS number 7446-70-0 (anhydrous) YesY, 10124-27-3 (hydrate), 7784-13-6 (hexahydrate)
PubChem 24012
ChemSpider 22445 YesY
UNII LIF1N9568Y YesY
ChEBI CHEBI:30114 YesY
RTECS number BD0530000
ATC code D10AX01
Jmol-3D images Image 1
Image 2
Properties
Molecular formula AlCl3
Molar mass 133.34 g/mol (anhydrous)
241.43 g/mol (hexahydrate)
Appearance white or pale yellow solid,
hygroscopic
Density 2.48 g/cm3 (anhydrous)
1.3 g/cm3 (hexahydrate)
Melting point 192.4 °C (378.3 °F; 465.5 K)
(anhydrous)
100 °C (212 °F; 373 K)
(hexahydrate)
180 °C (356 °F; 453 K)
(sublimes)
Boiling point 120 °C (248 °F; 393 K) (hexahydrate)
Solubility in water 43.9 g/100 ml (0 °C)
44.9 g/100 ml (10 °C)
45.8 g/100 ml (20 °C)
46.6 g/100 ml (30 °C)
47.3 g/100 ml (40 °C)
48.1 g/100 ml (60 °C)
48.6 g/100 ml (80 °C)
49 g/100 ml (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)[1]
Viscosity 0.35 cP (197 °C)
0.26 cP (237 °C)[1]
Structure
Crystal structure Monoclinic, mS16
Space group C12/m1, No. 12
Coordination
geometry
Octahedral (solid)
Tetrahedral (liquid)
Molecular shape Trigonal planar
(monomeric vapour)
Thermochemistry
Specific
heat capacity
C
91 J/mol·K[1]
Std molar
entropy
So298
111 J/mol·K[2]
Std enthalpy of
formation
ΔfHo298
−704.2 kJ/mol[1][2]
Gibbs free energy ΔG -628.6 kJ/mol[1]
Hazards
MSDS External MSDS
GHS pictograms The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[3]
GHS signal word Danger
GHS hazard statements H314[3]
GHS precautionary statements P280, P310, P305+351+338[3]
EU classification Corrosive C
R-phrases R34
S-phrases (S1/2), S7/8, S28, S45
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazards (white): no codeNFPA 704 four-colored diamond
LD50 anhydrous:
380 mg/kg, rat (oral)
hexahydrate:
3311 mg/kg, rat (oral)
Related compounds
Other anions Aluminium fluoride
Aluminium bromide
Aluminium iodide
Other cations Boron trichloride
Gallium trichloride
Indium(III) chloride
Magnesium chloride
Related Lewis acids Iron(III) chloride
Boron trifluoride
Supplementary data page
Structure and
properties
n, εr, etc.
Thermodynamic
data
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Aluminium chloride (AlCl3) is the main compound of aluminium and chlorine. It is white, but samples are often contaminated with iron trichloride, giving it a yellow colour. 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 chemical industry. The compound is often cited as a Lewis acid. It is an example of an inorganic compound that "cracks" at mild temperature, reversibly changing from a polymer to a monomer.

Structure[edit]

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.[4] In the melt, aluminium trichloride 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) vs 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,[5] unlike more ionic halides such as sodium chloride.

Reactions[edit]

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.[6] It forms tetrachloroaluminate AlCl4 in the presence of chloride ions.

Aluminium chloride reacts with calcium and magnesium hydrides in tetrahydrofuran forming tetrahydroaluminates.

Reactions with water[edit]

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 AlCl3·6H2O (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), the aluminium oxide is formed from the aluminium hydroxide via:

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+ [Al(OH)(H2O)5]2+ + H+

Aqueous solutions behave similarly to other aluminium salts containing hydrated Al3+ ions, giving a gelatinous precipitate of aluminium hydroxide upon reaction with dilute sodium hydroxide:

AlCl3 + 3 NaOH → Al(OH)3 + 3 NaCl

Synthesis[edit]

Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride at temperatures between 650 to 750 °C.[5]

2 Al + 3 Cl2 → 2 AlCl3
2 Al + 6 HCl → 2 AlCl3 + 3 H2

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 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 oxide when heated to 300 °C:[7]

2 AlCl3 + 3 H2O → Al2O3 + 6 HCl

Aluminium also forms a lower chloride, aluminium(I) chloride (AlCl), but this is very unstable and only known in the vapour phase.[5]

Uses[edit]

Anhydrous aluminium trichloride[edit]

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[6] is the major use for aluminium chloride, for example in the preparation of anthraquinone (for the dyestuffs industry) from benzene and phosgene.[5] In the general Friedel–Crafts reaction, an acyl chloride or alkyl halide reacts with an aromatic system as shown:[6]

Benzene Friedel-Crafts alkylation-diagram.svg

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.[citation needed] 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 trichloride 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.[8]

AlCl3 formylation.gif

Aluminium chloride finds a wide variety of other applications in organic chemistry.[9] For example, it can catalyse the "ene reaction", such as the addition of 3-buten-2-one (methyl vinyl ketone) to carvone:[10]

AlCl3 ene rxn.gif

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.[5]

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 chlorides[edit]

The hexahydrate has few applications, but aluminium chlorohydrate is a common component in antiperspirants at low concentrations.[7] Hyperhidrosis sufferers need a much higher concentration (12% or higher), sold under such brand names as Drysol, DryDerm, sunsola, Maxim, Odaban, CertainDri, B+Drier, Chlorhydrol, Anhydrol Forte and Driclor.

Safety[edit]

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.[11]

Aluminum chloride has been established as a neurotoxin.[12][13][14][15]

References[edit]

  1. ^ a b c d e http://chemister.ru/Database/properties-en.php?dbid=1&id=353
  2. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 0-618-94690-X. 
  3. ^ a b c Sigma-Aldrich Co., Aluminum chloride. Retrieved on 2014-05-05.
  4. ^ In contrast, AlBr3 has a more molecular structure, with the Al3+ centers occupying adjacent tetrahedral holes of the close-packed framework of Br ions.A. F. Wells, Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom, 1984.
  5. ^ a b c d e N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, Pergamon Press, Oxford, United Kingdom, 1984.
  6. ^ a b c G. A. Olah (ed.), Friedel-Crafts and Related Reactions, Vol. 1, Interscience, New York, 1963.
  7. ^ a b c Otto Helmboldt, L. Keith Hudson, Chanakya Misra, Karl Wefers, Wolfgang Heck, Hans Stark, Max Danner, Norbert Rösch "Aluminum Compounds, Inorganic" in Ullmann's Encyclopedia of Industrial Chemistry 2007, Wiley-VCH, Weinheim.doi:10.1002/14356007.a01_527.pub2
  8. ^ L. G. Wade, Organic Chemistry, 5th edition, Prentice Hall, Upper Saddle River, New Jersey, United States, 2003.
  9. ^ P. Galatsis, in: Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, (H. J. Reich, J. H. Rigby, eds.), pp. 12–15, Wiley, New York, 1999.
  10. ^ B. B. Snider (1980). "Lewis-acid catalyzed ene reactions". Acc. Chem. Res. 13 (11): 426. doi:10.1021/ar50155a007. 
  11. ^ http://www.solvaychemicals.us/static/wma/pdf/5/1/1/8/ALCL.pdf
  12. ^ He BP, Strong MJ (January 2000). "A morphological analysis of the motor neuron degeneration and microglial reaction in acute and chronic in vivo aluminum chloride neurotoxicity". J. Chem. Neuroanat. 17 (4): 207–15. doi:10.1016/S0891-0618(99)00038-1. PMID 10697247. 
  13. ^ Zubenko GS, Hanin I (October 1989). "Cholinergic and noradrenergic toxicity of intraventricular aluminum chloride in the rat hippocampus". Brain Res. 498 (2): 381–4. doi:10.1016/0006-8993(89)91121-9. PMID 2790490. 
  14. ^ Peng JH, Xu ZC, Xu ZX, et al. (August 1992). "Aluminum-induced acute cholinergic neurotoxicity in rat". Mol. Chem. Neuropathol. 17 (1): 79–89. doi:10.1007/BF03159983. PMID 1388451. 
  15. ^ Banks, W.A.; Kastin, A.J. (1989). "Aluminum-induced neurotoxicity: alterations in membrane function at the blood–brain barrier". Neurosci Biobehav Rev 13 (1): 47–53. doi:10.1016/S0149-7634(89)80051-X. PMID 2671833. 

External links[edit]