Aluminium compounds

Source: Wikipedia, the free encyclopedia.
Sample of aluminium sulfate hexadecahydrate, Al2(SO4)3·16H2O.

Lewis acids and readily form adducts.[4] Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class.[5]

Reactions of aluminium

Aluminium reacts with most nonmetals upon heating, forming compounds such as

intermetallic compounds involving metals from every group on the periodic table. Aluminium has a high chemical affinity to oxygen, which renders it suitable for use as a reducing agent in the thermite reaction. A fine powder of aluminium reacts explosively on contact with liquid oxygen; under normal conditions, however, aluminium forms a thin oxide layer that protects the metal from further corrosion by oxygen, water, or dilute acid, a process termed passivation.[2][6] This layer is destroyed by contact with mercury due to amalgamation or with salts of some electropositive metals.[2] As such, the strongest aluminium alloys are less corrosion-resistant due to galvanic reactions with alloyed copper,[7] and aluminium's corrosion resistance is greatly reduced by aqueous salts, particularly in the presence of dissimilar metals.[1] In addition, although the reaction of aluminium with water at temperatures below 280 °C is of interest for the production of hydrogen, commercial application of this fact has challenges in circumventing the passivating oxide layer, which inhibits the reaction, and in storing the energy required to regenerate the aluminium.[8]

Primarily because it is corroded by dissolved

chlorides, such as common sodium chloride, household plumbing is never made from aluminium.[9] However, because of its general resistance to corrosion, aluminium is one of the few metals that retains silvery reflectance in finely powdered form, making it an important component of silver-colored paints. Aluminium mirror finish has the highest reflectance of any metal in the 200–400 nm (UV) and the 3,000–10,000 nm (far IR) regions; in the 400–700 nm visible range it is slightly outperformed by tin and silver and in the 700–3000 nm (near IR) by silver, gold, and copper.[10]

In hot concentrated

aluminates—protective passivation under these conditions is negligible.[9] The reaction with aqueous alkali is often written:[2]

Al + NaOH + H2O → NaAlO2 + 3/2 H2

although the aluminium species in solution is probably instead the hydrated tetrahydroxoaluminate anion, [Al(OH)4]− or [Al(H2O)2(OH)4]−.[2]

Oxidizing acids do not effectively attack high-purity aluminium because the oxide layer forms and protects the metal; aqua regia will nevertheless dissolve aluminium. This allows aluminium to be used to store reagents such as nitric acid, concentrated sulfuric acid, and some organic acids.[11]

Inorganic compounds

The vast majority of compounds, including all aluminium-containing minerals and all commercially significant aluminium compounds, feature aluminium in the oxidation state 3+. The coordination number of such compounds varies, but generally Al3+ is either six- or four-coordinate. Almost all compounds of aluminium(III) are colorless.[2]

Aluminium hydrolysis as a function of pH. Coordinated water molecules are omitted. (Data from Baes and Mesmer)[12]

In aqueous solution, Al3+ exists as the hexaaqua cation [Al(H2O)6]3+, which has an approximate

pKa of 10−5.[13] Such solutions are acidic as this cation can act as a proton donor, progressively hydrolysing to [Al(H2O)5(OH)]2+, [Al(H2O)4(OH)2]+, and so on. As pH increases these mononuclear species begin to aggregate together by the formation of hydroxide bridges,[2] forming many oligomeric ions, such as the Keggin ion [Al13O4(OH)24(H2O)12]7+.[13] The process ends with precipitation of aluminium hydroxide, Al(OH)3. This is useful for clarification of water, as the precipitate nucleates on suspended particles in the water, hence removing them. Increasing the pH even further leads to the hydroxide dissolving again as aluminate, [Al(H2O)2(OH)4]−, is formed. Aluminium hydroxide forms both salts and aluminates and dissolves in acid and alkali, as well as on fusion with acidic and basic oxides:[2]

Al2O3 + 3 SiO2 fuse→  Al2(SiO3)3
Al2O3 + CaO fuse→  Ca(AlO2)2

This behaviour of Al(OH)3 is termed amphoterism, and is characteristic of weakly basic cations that form insoluble hydroxides and whose hydrated species can also donate their protons. Further examples include Be2+, Zn2+, Ga3+, Sn2+, and Pb2+; indeed, gallium in the same group is slightly more acidic than aluminium. One effect of this is that aluminium salts with weak acids are hydrolysed in water to the aquated hydroxide and the corresponding nonmetal hydride: aluminium sulfide yields hydrogen sulfide, aluminium nitride yields ammonia, and aluminium carbide yields methane. Aluminium cyanide, acetate, and carbonate exist in aqueous solution but are unstable as such; only incomplete hydrolysis takes place for salts with strong acids, such as the halides, nitrate, and sulfate. For similar reasons, anhydrous aluminium salts cannot be made by heating their "hydrates": hydrated aluminium chloride is in fact not AlCl3·6H2O but [Al(H2O)6]Cl3, and the Al–O bonds are so strong that heating is not sufficient to break them and form Al–Cl bonds instead:[2]

2[Al(H2O)6]Cl3 heat→  Al2O3 + 6 HCl + 9 H2O

All four

heat of formation. Each aluminium atom is surrounded by six fluorine atoms in a distorted octahedral arrangement, with each fluorine atom being shared between the corners of two octahedra in a structure related to but distorted from that of ReO3. Such {AlF6} units also exist in complex fluorides such as cryolite, Na3AlF6, but should not be considered as [AlF6]3− complex anions as the Al–F bonds are not significantly different in type from the other M–F bonds.[14] Such differences in coordination between the fluorides and heavier halides are not unusual, occurring in SnIV and BiIII as well for example; even bigger differences occur between CO2 and SiO2.[14] AlF3 melts at 1,290 °C (2,354 °F) and is made by reaction of aluminium oxide with hydrogen fluoride gas at 700 °C (1,292 °F).[14]

Mechanism of the Friedel–Crafts acylation, using AlCl3 as a catalyst

With heavier halides, the coordination numbers are lower. The other trihalides are

Lewis acidic nature makes them useful as catalysts for the Friedel–Crafts reactions. Aluminium trichloride has major industrial uses involving this reaction, such as in the manufacture of anthraquinones and styrene; it is also often used as the precursor for many other aluminium compounds and as a reagent for converting nonmetal fluorides into the corresponding chlorides (a transhalogenation reaction).[14]

AlCl3 + 3 LiZ → 3 LiCl + AlZ3 (Z = R, NR2, N=CR2)
AlCl3 + 4 LiZ → 3 LiCl + LiAlZ4 (Z = R, NR2, N=CR2, H)
BF3 + AlCl3 → AlF3 + BCl3

Aluminium forms one stable oxide with the

polymorphs). Many other intermediate and related structures are also known.[13] Most are produced from ores by a variety of wet processes using acid and base. Heating the hydroxides leads to formation of corundum. These materials are of central importance to the production of aluminium and are themselves extremely useful. Some mixed oxide phases are also very useful, such as spinel (MgAl2O4), Na-β-alumina (NaAl11O17), and tricalcium aluminate (Ca3Al2O6, an important mineral phase in Portland cement).[13]

The only stable

zinc blende structure. All four can be made by high-temperature (and possibly high-pressure) direct reaction of their component elements.[17]

Rarer oxidation states

Although the great majority of aluminium compounds feature Al3+ centers, compounds with lower oxidation states are known and are sometimes of significance as precursors to the Al3+ species.

Aluminium(I)

Main article: Aluminium(I) compounds

AlF, AlCl, AlBr, and AlI exist in the gaseous phase when the respective trihalide is heated with aluminium, and at cryogenic temperatures. Their instability in the condensed phase is due to their ready disproportionation to aluminium and the respective trihalide: the reverse reaction is favored at high temperature (although even then they are still short-lived), explaining why AlF3 is more volatile when heated in the presence of aluminium, as is aluminium when heated in the presence of AlCl3.[14]

A stable derivative of aluminium monoiodide is the cyclic adduct formed with triethylamine, Al4I4(NEt3)4. Also of theoretical interest but only of fleeting existence are Al2O and Al2S. Al2O is made by heating the normal oxide, Al2O3, with silicon at 1,800 °C (3,272 °F) in a vacuum. Such materials quickly disproportionate to the starting materials.[18]

Aluminium(II)

Very simple Al(II) compounds are invoked or observed in the reactions of Al metal with oxidants. For example,

aluminium monoxide, AlO, has been detected in the gas phase after explosion[19] and in stellar absorption spectra.[20] More thoroughly investigated are compounds of the formula R4Al2 which contain an Al–Al bond and where R is a large organic ligand.[21]

Organoaluminium compounds and related hydrides

Main article:
Organoaluminium compound
Structure of trimethylaluminium, a compound that features five-coordinate carbon.

A variety of compounds of empirical formula AlR3 and AlR1.5Cl1.5 exist.

heterocyclic and cluster organoaluminium compounds involving Al–N bonds.[23]

The industrially most important aluminium hydride is lithium aluminium hydride (LiAlH4), which is used in as a reducing agent in organic chemistry. It can be produced from lithium hydride and aluminium trichloride:[25]

4 LiH + AlCl3 → LiAlH4 + 3 LiCl

The simplest hydride, aluminium hydride or alane, is not as important. It is a polymer with the formula (AlH3)n, in contrast to the corresponding boron hydride that is a dimer with the formula (BH3)2.[25]

References

  1. ^ a b c Greenwood and Earnshaw, pp. 222–4
  2. ^ a b c d e f g h i Greenwood and Earnshaw, pp. 224–7
  3. ^ Greenwood and Earnshaw, pp. 112–3
  4. ^ King, p. 241
  5. ^ King, pp. 235–6
  6. ISBN 978-0-08-044495-6. Archived
    from the original on 21 May 2016.
  7. ISBN 978-0-340-63207-9
    .
  8. ^ "Reaction of Aluminum with Water to Produce Hydrogen" (PDF). U.S. Department of Energy. 1 January 2008. Archived from the original (PDF) on 14 September 2012.
  9. ^
    ISBN 978-0-8031-2610-7. Archived
    from the original on 24 April 2016.
  10. ISBN 978-0-7503-0688-1
    .
  11. ISBN 978-3-527-30673-2
    .
  12. ISBN 978-0-89874-892-5
    .
  13. ^ a b c d e f Greenwood and Earnshaw, pp. 242–52
  14. ^ a b c d e f Greenwood and Earnshaw, pp. 233–7
  15. ISBN 978-1-136-37393-0
    .
  16. ^ Roscoe, Henry Enfield; Schorlemmer, Carl (1913). A treatise on chemistry. Macmillan. p. 718. Aluminium forms one stable oxide, known by its mineral name corundum.
  17. ^ a b Greenwood and Earnshaw, pp. 252–7
  18. doi:10.1002/anie.199601291
    .
  19. S2CID 4163250
    .
  20. doi:10.1086/147348
    .
  21. ISBN 978-0-12-031151-4
    .
  22. ISBN 978-3-527-29390-2
    .
  23. ^ a b Greenwood and Earnshaw, pp. 257–67
  24. doi:10.1016/S0022-328X(00)86043-X
  25. ^ a b Greenwood and Earnshaw, pp. 227–32

Bibliography

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