Aluminium chloride

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Aluminium chloride
Aluminium(III) chloride
Aluminium(III) chloride
Aluminium trichloride hexahydrate, pure (top), and contaminated with iron(III) chloride (bottom)
Aluminium trichloride dimer
Aluminium trichloride dimer
Aluminium trichloride unit cell
Aluminium trichloride unit cell
Names
IUPAC name
Aluminium chloride
Other names
Aluminium(III) chloride
Aluminium trichloride
Trichloroaluminum
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
 100.028.371 Edit this at Wikidata
EC Number
  • 231-208-1
1876
RTECS number
  • BD0530000
UNII
  • InChI=1S/Al.3ClH/h;3*1H/q+3;;;/p-3 checkY
    Key: VSCWAEJMTAWNJL-UHFFFAOYSA-K checkY
  • InChI=1/Al.3ClH/h;3*1H/q+3;;;/p-3
    Key: VSCWAEJMTAWNJL-DFZHHIFOAR
SMILES
  • dimer
    : Cl[Al-]1(Cl)[Cl+][Al-]([Cl+]1)(Cl)Cl
  • hexahydrate: [OH2+][Al-3]([OH2+])([OH2+])([OH2+])([OH2+])[OH2+].[Cl-].[Cl-].[Cl-]
Properties
AlCl3
Molar mass
  • 133.341 g/mol (anhydrous)
  • 241.432 g/mol (hexahydrate)
[1]
Appearance Colourless crystals,
hygroscopic
Density
  • 2.48 g/cm3 (anhydrous)
  • 2.398 g/cm3 (hexahydrate)
[1]
Melting point
  • 180 °C (356 °F; 453 K) (anhydrous, sublimes)[1]
  • 100 °C (212 °F; 373 K) (hexahydrate, decomposes)[1]
  • 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)
[2]
Viscosity
  • 0.35 cP (197 °C)
  • 0.26 cP (237 °C)
[2]
Structure
Monoclinic, mS16
C12/m1, No. 12[3]
a = 0.591 nm, b = 0.591 nm, c = 1.752 nm[3]
0.52996 nm3
6
Octahedral (solid)
Tetrahedral (liquid)
Trigonal planar
(monomeric vapour)
Thermochemistry
91.1 J/(mol·K)[4]
109.3 J/(mol·K)[4]
Std enthalpy of
formation
(ΔfH⦵298)
 −704.2 kJ/mol[4]
−628.8 kJ/mol[4]
Pharmacology
D10AX01 (WHO)
Hazards
GHS labelling:[6]
GHS05: Corrosive
Danger
H314
P260, P280, P301+P330+P331, P303+P361+P353, P305+P351+P338+P310, P310
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
3
0
2
Lethal dose or concentration (LD, LC):
380 mg/kg, rat (oral, anhydrous)
3311 mg/kg, rat (oral, hexahydrate)
NIOSH (US health exposure limits):
PEL (Permissible)
 None[5]
REL (Recommended)
 2 mg/m3[5]
IDLH
(Immediate danger)
 N.D.[5]
Related compounds
Other anions
Other cations
Related
Lewis acids
Supplementary data page
Aluminium chloride (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)
Aluminium chloride
Clinical data
AHFS/Drugs.comMonograph
License data
ECHA InfoCard
100.028.371 Edit this at Wikidata
Data page
Aluminium chloride (data page)

Aluminium chloride, also known as aluminium trichloride, is an

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 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.

Lewis acid. It is 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

ionic halides such as sodium chloride
.

Hexahydrate

The hexahydrate consists of

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
Friedel-Crafts alkylation
and related reactions.

Uses

Alkylation and acylation of arenes

AlCl3 is a common Lewis-acid

aromatic system as shown:[11]

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

corrosive waste. For these and similar reasons, the use of aluminium chloride has often been displaced by zeolites.[7]

Aluminium chloride can also be used to introduce

co-catalyst.[16]

Other applications in organic and organometallic synthesis

Aluminium chloride finds a wide variety of other applications in

3-buten-2-one (methyl vinyl ketone) to carvone:[18]

It is used to induce a variety of hydrocarbon couplings and rearrangements.[19][20]

Aluminium chloride combined with aluminium in the presence of an arene can be used to synthesize bis(arene) metal complexes, e.g.

Fischer–Hafner synthesis. Dichlorophenylphosphine is prepared by reaction of benzene and phosphorus trichloride catalyzed by aluminium chloride.[21]

Medical

Topical aluminum chloride hexahydrate is used for the treatment of hyperhidrosis (excessive sweating).[22][23][24]

Reactions

Anhydrous aluminium chloride is a powerful

Lewis bases such as benzophenone and mesitylene.[11] It forms tetrachloroaluminate ([AlCl4]−) in the presence of chloride
ions.

Aluminium chloride reacts with calcium and magnesium hydrides in tetrahydrofuran forming tetrahydroaluminates.[citation needed]

Reactions with water

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):

[Al(H2O)6]Cl3 → Al(OH)3 + 3 HCl + 3 H2O

Like

aquo ligands
:

[Al(H2O)6]3+ ⇌ [Al(OH)(H2O)5]2+ + H+

Aqueous solutions behave similarly to other

precipitate of aluminium hydroxide upon reaction with dilute sodium hydroxide
:

AlCl3 + 3 NaOH → Al(OH)3 + 3 NaCl

Synthesis

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]

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

Aluminium chloride may be formed via a single displacement reaction between copper(II) chloride and aluminium.

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.[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

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]

History

Aluminium chlorides were known in the 18th century as muriate of alumina, marine alum, argillaceous marine salt,[27] muriated clay.[28] It was first chemically studied in the 1830s.[29]

Safety

Anhydrous AlCl3 is strongly corrosive and releases hydrochloric acid in contact with water.[7]

See also

References

  1. ^
    ISBN 1-4398-5511-0
    .
  2. ^ a b "Properties of substance: Aluminium chloride". Chemister.ru. 2007-03-19. Archived from the original on 2014-05-05. Retrieved 2017-03-17.
  3. ^
    S2CID 100796636
    .
  4. ^
    ISBN 1-4398-5511-0
    .
  5. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0024". National Institute for Occupational Safety and Health (NIOSH).
  6. ^ Sigma-Aldrich Co., Aluminium chloride.
  7. ^
    ISBN 978-3527306732
    .
  8. ISBN 0198553706
    . In contrast, AlBr3 has a more molecular structure, with the Al3+ centers occupying adjacent tetrahedral holes of the close-packed framework of Br− ions.
  9. ^
    ISBN 978-0-08-022057-4
    .
  10. S2CID 263857074
    .
  11. ^ a b c Olah GA, ed. (1963). Friedel-Crafts and Related Reactions. Vol. 1. New York City: Interscience.
  12. ISSN 1099-0682
    .
  13. doi:10.15227/orgsyn.089.0210
    .
  14. doi:10.15227/orgsyn.080.0227
    .
  15. doi:10.15227/orgsyn.079.0204
    .
  16. ISBN 013033832X
    .
  17. ISBN 978-0-471-97925-8
    .
  18. doi:10.1021/ar50155a007
    .
  19. doi:10.15227/orgsyn.059.0085
    .
  20. doi:10.15227/orgsyn.061.0062
    .
  21. doi:10.15227/orgsyn.031.0088
    .
  22. PMID 30215934
    .
  23. PMID 30710603
    .
  24. ^ "Aluminum Chloride (Topical) (Monograph)". American Society of Health System Pharmacists (ASHP). drugs.com.
  25. ^ a b "List of Minerals". www.ima-mineralogy.org. International Mineralogical Association. March 21, 2011.
  26. ^ "Cadwaladerite". www.mindat.org.
  27. ^ de Fourcroy AF (1790). Elements of natural history, and of chemistry: being the second edition of the elementary lectures on those sciences ... enlarged and improved by the author ... Translated into English, with ... notes; and an historical preface by the translator W. Nicholson.
  28. ^ Berthollet CL (1791). Elements of the Art of Dyeing ... Translated ... by William Hamilton. Stephen Couchman; sold by J. Johnson.
  29. ^ Mather WW (1835). "Chloride of Aluminium and its Analysis". In Silliman B (ed.). The American Journal of Science. Vol. 27. Kline Geology Laboratory, Yale University. pp. 241–253 (249).
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