Chemistry:Aluminium chloride

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Short description: Chemical compound
Aluminium chloride
Aluminium(III) chloride
Aluminium trichloride hexahydrate, pure (top), and contaminated with iron(III) chloride (bottom)
Aluminium trichloride dimer
Aluminium trichloride unit cell
Names
IUPAC name
Aluminium chloride
Other names
Aluminium(III) chloride
Aluminium trichloride
Trichloroaluminum
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 231-208-1
1876
RTECS number
  • BD0530000
UNII
Properties
AlCl
3
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
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]
−704.2 kJ/mol[4]
−628.8 kJ/mol[4]
Pharmacology
1=ATC code }} D10AX01 (WHO)
Hazards
GHS pictograms GHS05: Corrosive
GHS Signal word Danger
HH314Script error: No such module "Preview warning".Category:GHS errors
PP260Script error: No such module "Preview warning".Category:GHS errors, P280, P301+330+331, P303+361+353, P310, P305+351+338+310
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 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 codeNFPA 704 four-colored diamond
0
3
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
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is ☑Y☒N ?)
Infobox references
Tracking categories (test):

Aluminium chloride, also known as aluminium trichloride, is an inorganic compound with the formula AlCl
3. It forms a hexahydrate with the formula [Al(H
2O)
6]Cl
3, 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 metal, 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

Anhydrous

Aluminium-trichloride-3D-structures.png

AlCl
3 adopts three structures, depending on the temperature and the state (solid, liquid, gas). Solid AlCl
3 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 dimer Al
2Cl
6, 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). Al
2Cl
6 dimers are also found in the vapour phase. At higher temperatures, the Al
2Cl
6 dimers dissociate into trigonal planar AlCl
3 monomer, which is structurally analogous to BF
3. The melt conducts electricity poorly,[9] unlike more ionic halides 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(H
2O)
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.

Uses

Alkylation and acylation of arenes

AlCl
3 is a common Lewis-acid catalyst for Friedel-Crafts reactions, both acylations and alkylations.[11] Important products are detergents and ethylbenzene. These types of reactions are the major use for aluminium chloride, for example, in the preparation of anthraquinone (used in the dyestuffs industry) from benzene and phosgene.[9] In the general Friedel-Crafts reaction, an acyl chloride or alkyl halide reacts with an aromatic system as shown:[11]

Benzene Friedel-Crafts alkylation-diagram.svg

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]

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.[16]

AlCl3 formylation.gif

Other applications in organic and organometallic synthesis

Aluminium chloride finds a wide variety of other applications in organic chemistry.[17] For example, it can catalyse the ene reaction, such as the addition of 3-buten-2-one (methyl vinyl ketone) to carvone:[18]

AlCl3 ene rxn.gif

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. bis(benzene)chromium, from certain metal halides via the Fischer–Hafner synthesis. Dichlorophenylphosphine is prepared by reaction of benzene and phosphorus trichloride catalyzed by aluminium chloride.[21]

Reactions

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.[11] It forms tetrachloroaluminate ([AlCl
4]−
) 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(H
2O)
6]Cl
3. The anhydrous phase cannot be regained on heating the hexahydrate. Instead HCl is lost leaving aluminium hydroxide or alumina (aluminium oxide):

[Al(H
2O)
6]Cl
3 → Al(OH)
3 + 3 HCl + 3 H
2O

Like metal aquo complexes, aqueous AlCl
3 is acidic owing to the ionization of the aquo ligands:

[Al(H
2O)
6]3+ ⇌ [Al(OH)(H
2O)
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:

AlCl
3 + 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 Cl
2 → 2 AlCl
3
2 Al + 6 HCl → 2 AlCl
3 + 3 H
2

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

2 Al + 3 CuCl
2 → 2 AlCl
3 + 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(H
2O)
6]Cl
3 → Al(OH)
3 + 3 HCl + 3 H
2O

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.[22] A more complex, basic and hydrated aluminium chloride mineral is cadwaladerite.[23][22]

Safety

Anhydrous AlCl
3 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.[24]

See also

References

  1. ↑ 1.0 1.1 1.2 1.3 Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.45. ISBN 1439855110. 
  2. ↑ 2.0 2.1 Aluminium chloride . Chemister.ru (2007-03-19). Retrieved on 2017-03-17.
  3. ↑ Cite error: Invalid <ref> tag; no text was provided for refs named str
  4. ↑ 4.0 4.1 4.2 4.3 Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 5.5. ISBN 1439855110. 
  5. ↑ 5.0 5.1 5.2 NIOSH Pocket Guide to Chemical Hazards. "#0024". National Institute for Occupational Safety and Health (NIOSH). https://www.cdc.gov/niosh/npg/npgd0024.html. 
  6. ↑ Sigma-Aldrich Co., Aluminium chloride.
  7. ↑ 7.0 7.1 7.2 Helmboldt, Otto; Keith Hudson, L.; Misra, Chanakya; Wefers, Karl; Heck, Wolfgang; Stark, Hans; Danner, Max; Rösch, Norbert (2007). "Ullmann's Encyclopedia of Industrial Chemistry". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a01_527.pub2. 
  8. ↑ In contrast, AlBr
    3     has a more molecular structure, with the Al3+ centers occupying adjacent tetrahedral holes of the close-packed framework of Br−
     ions. Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford Press, Oxford, United Kingdom. ISBN:0198553706.
  9. ↑ 9.0 9.1 9.2 9.3 Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. ISBN 978-0-08-022057-4. https://books.google.com/books?id=OezvAAAAMAAJ&q=0-08-022057-6&dq=0-08-022057-6&source=bl&ots=m4tIRxdwSk&sig=XQTTjw5EN9n5z62JB3d0vaUEn0Y&hl=en&sa=X&ei=UoAWUN7-EM6ziQfyxIDoCQ&ved=0CD8Q6AEwBA. 
  10. ↑ Andress, K.R.; 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. 
  11. ↑ 11.0 11.1 11.2 Olah, G. A., ed (1963). Friedel-Crafts and Related Reactions. 1. New York City: Interscience. 
  12. ↑ Nenitzescu, Costin D.; Cantuniari, Ion P. (1933). "Durch Aluminiumchlorid Katalysierte Reaktion, VI. Mitteil.: Die Umlagerung des Cyclohexans in Metyl-cyclopentan" (in en). Berichte der Deutschen Chemischen Gesellschaft (A and B Series) 66 (8): 1097–1100. doi:10.1002/cber.19330660817. ISSN 1099-0682. https://chemistry-europe.onlinelibrary.wiley.com/doi/abs/10.1002/cber.19330660817. 
  13. ↑ Jonathan T. Reeves; Zhulin Tan; Daniel R. Fandrick; Jinhua J. Song; Nathan K. Yee; Chris H. Senanayake (2012). "Synthesis of Trifluoromethyl Ketones from Carboxylic Acids: 4-(3,4-Dibromophenyl)-1,1,1-trifluoro-4-methylpentan-2-one". Organic Syntheses 89: 210. doi:10.15227/orgsyn.089.0210. 
  14. ↑ Kamil Paruch; Libor Vyklicky; Thomas J. Katz (2003). "Preparation of 9,10-Dimethoxyphenanthrene and 3,6-Diacetyl-9,10-Dimethoxyphenanthrene". Organic Syntheses 80: 227. doi:10.15227/orgsyn.080.0227. 
  15. ↑ Alexander J. Seed; Vaishali Sonpatki; Mark R. Herbert (2002). "3-(4-Bromobenzoyl)propanoic Acid". Organic Syntheses 79: 204. doi:10.15227/orgsyn.079.0204. 
  16. ↑ Wade, L. G. (2003) Organic Chemistry, 5th edition, Prentice Hall, Upper Saddle River, New Jersey, United States. ISBN:013033832X.
  17. ↑ Galatsis, P. (1999) Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents, H. J. Reich, J. H. Rigby (eds.) John Wiley & Sons , New York City . pp. 12–15. ISBN:978-0-471-97925-8.
  18. ↑ Snider, B. B. (1980). "Lewis-acid catalyzed ene reactions". Acc. Chem. Res. 13 (11): 426. doi:10.1021/ar50155a007. 
  19. ↑ Reuben D. Rieke; Stephen E. Bales; Phillip M. Hudnall; Timothy P. Burns; Graham S. Poindexter (1979). "Highly Reactive Magnesium for the Preparation of Grignard Reagents: 1-Norbornanecarboxylic Acid". Organic Syntheses 59: 85. doi:10.15227/orgsyn.059.0085. 
  20. ↑ Sami A. Shama; Carl C. Wamser (1983). "Hexamethyl Dewar Benzene". Organic Syntheses 61: 62. doi:10.15227/orgsyn.061.0062. 
  21. ↑ B. Buchner; L. B. Lockhart Jr. (1951). "Phenyldichlorophosphine". Organic Syntheses 31: 88. doi:10.15227/orgsyn.031.0088. 
  22. ↑ 22.0 22.1 "List of Minerals". International Mineralogical Association. March 21, 2011. https://www.ima-mineralogy.org/Minlist.htm. 
  23. ↑ "Cadwaladerite". https://www.mindat.org/min-845.html. 
  24. ↑ Aluminium Chloride. solvaychemicals.us

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