Aluminium — Properties & Uses

Aluminium in Group 13

Aluminium is the second member of Group 13, the boron family, with the outer electronic configuration $\ce{[Ne]3s^2 3p^1}$. It is the most abundant metal and the third most abundant element in the earth's crust (8.3% by mass), occurring chiefly in the minerals bauxite, $\ce{Al2O3.2H2O}$, and cryolite, $\ce{Na3AlF6}$. Where boron is a typical non-metal, aluminium is a genuine, highly electropositive metal — yet it retains a clear chemical kinship with boron, most visibly in the electron deficiency of its trivalent compounds.

The key to aluminium's chemistry is its oxidation state. Although the first member boron is forced into purely covalent bonding by its very high sum of the first three ionisation enthalpies, the corresponding sum falls considerably from boron to aluminium. Aluminium is therefore able to form $\ce{Al^3+}$ ions and exhibits a stable, dominant $+3$ oxidation state. The lower $+1$ state, which becomes progressively important for the heavier congeners gallium, indium and thallium through the inert pair effect, is essentially absent for aluminium.

The presence of vacant $3d$ orbitals in the valence shell of aluminium is decisive. It allows aluminium to expand its covalence beyond four — for example, where boron can form only $\ce{[BF4]^-}$, aluminium can form the octahedral $\ce{[AlF6]^3-}$ ion using $sp^3d^2$ hybridisation. The same vacant orbitals let aluminium chloride in acidified aqueous solution exist as the octahedral hexaaquaaluminium(III) ion, $\ce{[Al(H2O)6]^3+}$.

Physical Properties

Aluminium is a bright silvery-white metal with high tensile strength. It has high electrical and thermal conductivity; on a weight-to-weight basis its electrical conductivity is twice that of copper, a fact that underpins its use in overhead transmission cables. It is soft, ductile and readily worked into pipes, tubes, rods, wires, plates and foils. Aluminium forms useful alloys with copper, manganese, magnesium, silicon and zinc, which extends its mechanical range considerably.

The strongly negative standard electrode potential, $E^{\ominus}(\ce{Al^3+/Al}) = -1.66\ \text{V}$, confirms that aluminium has a high tendency to form $\ce{Al^3+}$(aq) ions and is far more electropositive than its heavier congener thallium, for which the $+3$ state is powerfully oxidising. A noteworthy quirk of the family is that the atomic radius of gallium (135 pm) is actually smaller than that of aluminium (143 pm), because the ten additional $3d$ electrons in gallium screen the nuclear charge only poorly.

Reactivity Towards Air

On thermodynamic grounds aluminium should react readily with both air and water, yet in practice the metal is stable in both. The reason is kinetic protection: a very thin, continuous film of oxide forms on the surface, only about $10^{-4}$ to $10^{-6}$ mm thick, and this film shields the metal from further attack. If the protective covering is destroyed — for instance by amalgamating the surface with mercury — the exposed metal decomposes even cold water, forming $\ce{Al2O3}$ and liberating dihydrogen.

When amorphous aluminium is heated in air it burns to the oxide, and at high temperature it combines with dinitrogen to give the nitride. Aluminium articles are often anodised: electrolysing dilute $\ce{H2SO4}$ with aluminium as the anode builds a much thicker oxide layer (around $10^{-2}$ mm) that can take up pigments, giving a coloured, decorative finish.

The amphoteric nature of $\ce{Al2O3}$ — intermediate between the acidic $\ce{B2O3}$ and the basic oxides of indium and thallium — is itself a direct expression of the metallic-yet-borderline character that defines aluminium's place in the group.

Amphoteric Character: Acids and Alkalies

The single most examined property of aluminium is its amphoterism: it dissolves in both mineral acids and aqueous alkalies, liberating dihydrogen in each case. Boron, by contrast, reacts with neither acids nor alkalies at moderate temperature, so the contrast within Group 13 is sharp.

In dilute hydrochloric acid, aluminium dissolves to give aluminium(III) ions and dihydrogen gas:

$$\ce{2Al(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2(g)}$$ $$\ce{2Al(s) + 6H+(aq) -> 2Al^3+(aq) + 3H2(g)}$$

With aqueous sodium hydroxide, aluminium again liberates dihydrogen, this time forming a soluble aluminate — sodium tetrahydroxoaluminate(III):

$$\ce{2Al(s) + 2NaOH(aq) + 6H2O(l) -> 2Na[Al(OH)4](aq) + 3H2(g)}$$

The aluminate is also written as $\ce{NaAlO2.2H2O}$, sodium aluminate. Because aluminium reacts with both an acid and a base to give salt plus hydrogen, it is amphoteric in the fullest sense. A common laboratory and household demonstration draws on this: a mixture of dilute $\ce{NaOH}$ and aluminium pieces generates dihydrogen vigorously and is used to clear blocked drains.

One further subtlety belongs here. In the trivalent state most aluminium compounds are covalent and are hydrolysed in water; the trichloride on hydrolysis forms tetrahedral $\ce{[Al(OH)4]^-}$-type species ($sp^3$ at Al), while in acidified solution the octahedral $\ce{[Al(H2O)6]^3+}$ ion ($sp^3d^2$, using the $3d$ orbitals) is obtained. This versatility of geometry is again traceable to those vacant $3d$ orbitals.

Aluminium Chloride and the Al2Cl6 Dimer

Anhydrous aluminium chloride is a white solid that sublimes at 453 K. It is made by passing dry hydrogen chloride or chlorine over heated aluminium under anhydrous conditions:

$$\ce{2Al + 6HCl -> Al2Cl6 + 3H2}$$ $$\ce{2Al + 3Cl2 -> Al2Cl6}$$

The crucial structural fact is that aluminium chloride exists as a dimer, $\ce{Al2Cl6}$, at room temperature and dissociates to the monomer $\ce{AlCl3}$ only at high temperature. The reason is electron deficiency. Aluminium has only three valence electrons; when all three form $\ce{Al-Cl}$ covalent bonds, the aluminium atom is left with only six electrons in its valence shell. To complete the octet, each aluminium accepts a lone pair from a chlorine atom of a neighbouring $\ce{AlCl3}$ unit through coordinate (dative) bonds. Two such bridges are formed, giving the dimer, and the arrangement of chlorine atoms about each aluminium is roughly tetrahedral.

When treated with water the dimer is hydrolysed, giving hydrated aluminium ions and chloride ions:

$$\ce{Al2Cl6 + 12H2O -> 2[Al(H2O)6]^3+ + 6Cl^-}$$

Because anhydrous aluminium chloride is electron deficient, it is a strong Lewis acid, and it is this property that makes it the classic catalyst for the Friedel–Crafts reaction. Aluminium chloride is also partially hydrolysed by atmospheric moisture, liberating $\ce{HCl}$ gas, which is why white fumes appear around a bottle of the anhydrous solid.

Alums and Potash Alum

When two salts each capable of independent existence are mixed and the solution is allowed to crystallise, crystals containing both salts are formed; in solution, however, all the ions exist freely. Such substances are called double salts. The classic example is the crystallisation of a solution of potassium sulphate and aluminium sulphate, which deposits transparent octahedral crystals of potash alum.

The solid potash alum contains the ions $\ce{[K(H2O)6]^+}$, $\ce{[Al(H2O)6]^3+}$ and $\ce{SO4^2-}$, and it gives the tests of all its constituent ions in solution — the signature of a double salt. Double sulphates of similar composition and properties are collectively called alums.

The trivalent aluminium cation may be replaced by a trivalent metal ion of similar ionic size, such as $\ce{Ti^3+}$, $\ce{Cr^3+}$, $\ce{Fe^3+}$ or $\ce{Co^3+}$, and the monovalent potassium may be replaced by ammonium, $\ce{NH4+}$. All these alums are isomorphous — they crystallise in the same octahedral form. Potash alum, $\ce{KAl(SO4)2.12H2O}$, is by far the most important; it is used as a mordant in the dyeing industry and in the purification of water.

Uses of Aluminium

Aluminium and its alloys are among the most widely used engineering materials, a consequence of the combination of low density, high tensile strength, high conductivity and the self-protecting oxide film. Because aluminium can be drawn and rolled into pipes, tubes, rods, wires, plates and foils, it finds uses across packaging, utensil-making, construction, and the aeroplane and transportation industries.

A practical caution accompanies the metal's everyday use. Although aluminium is protected by its oxide film, prolonged contact with water can disturb that film, so aluminium utensils should not be kept standing in water overnight. The old NCERT also notes that the use of aluminium and its compounds for domestic purposes has been reduced considerably because of concerns over their toxicity. Even so, in structural and electrical engineering aluminium remains indispensable.