Overview of the Periodic Table
Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12 | Boron (Group 13) | Group 13 - the Triels
Carbon (Group 14) | Group 14 - The Tetrels | Nitrogen (Group 15) | Group 15 - The Pentels
Oxygen (Group 16) | Group 16 - The Chalcogens | Group 17 - The Halogens | Group 18 - The Noble Gases
An Overview of the Transition Metals | Group 3 | Group 4 | Group 5 | Group 6 | Group 7 | Group 8 | Group 9 | Group 10 | Group 11
The Lanthanides | The Actinides

The remaining elements of group 13, aluminium, gallium, indium and thallium have markedly different physical and chemical properties in comparison to boron. Aluminium is the most abundant metal in the Earth’s crust and occurs in igneous rocks such as felspars and micas. It also occurs in many clays and in minerals such as cryolite (Na₃AlF₆ and spinel (MgAl₂O₄) but the most important mineral is bauxite which is essentially a hydrated oxide/hydroxide (Al₂O₃·n H₂O) from which the metal is extracted. Many gemstones are also impure forms of the oxide, Al₂O₃, containing small amount of transition metal ions that give them their colour, for example, ruby (Cr) and sapphire (Co)

Gallium, indium and thallium occur in much smaller quantities in the Earth’s crust. Gallium can be found in small amounts in bauxite but generally the other metals are found in low abundance in the sulphide ores of other metals such as zinc, iron and copper. Most gallium is produced as a by-product of the aluminium industry whilst indium and thallium are extracted from the dusts produced through the roasting of the sulphide ores of metals like zinc and through sulphide roasting for sulphuric acid manufacture.

Preparation, Properties and Uses.

Aluminium is produced on an industrial scale from the mineral bauxite. The mineral is purified and then dissolves in an aqueous solution of sodium hydroxide which produces the hydroxide Al(OH)₄^-. Insoluble impurites, mostly iron salts, are removed and then Al(OH)₃·3 H₂O is precipitated on cooling. After the water molecules of crystallisation have been removed by heating it is dissolved in molten cryolite at 940-980 °C and electrolysed using large carbon electrodes. Before electrolysis techniques became available for the extraction of aluminium, it was one of the most expensive metals in the world and was even exhibited besides the crown jewels at the Parid exhibition in 1855. After the introduction of electrolytic methods, the price fell by over a thousand fold over a century. Aluminium is a vastly important material affecting our daily lives. It has wide application in the construction industry, it is used to build cars and aircraft, it is used to manufacture power lines and is used for kitchen foil and in drink cans. It is a strong, hard metal and is quite electropositive and reactive. It avoids corrosion however as it quickly forms a hard surface film of oxide. It will dissolve in aqueous NaOH and will formed compounds with most non-metal elements with heating.

Gallium is obtained as a by-product of the aluminium industry. Solutions of aluminium salts are electrolysed and consequently enriched in gallium. Further electrolysis of the extract yields gallium metal. Gallium is an important element in the electronics industry and has applications in semiconductors. One of the most importan of these materials in gallium arsenide, GaAs, which can emit light and is used in diode lasers and LEDs. It is also used in low temperature solders and the fluorescent compound MgGa₂O₄ is used in photocopiers.

Indium is also extracted by electrolytic process and like gallium is also an important element in the electronics industry. Many indium compounds find applications in semiconducting materials for transistors, thermistors and photo active devices. Thallium is extracted from the fine dusts from the sulphide ores of other elements by solvation in warm dilute acids. It is then separated from impurities and purified by electrolysis and deposition. Thallium and all it’s compounds are extremely toxic. The element has no major uses however the thallium(I) compound Tl₂SO₄ was once used as rodent poison. It has since been banned in most countries due to it’s odourless and tasteless nature making the risk of accidental poisoning unacceptably high.

Table 1: Physical properties of the group 13 element. Boron is included for comparison.

These latter group 13 elements differ from boron firstly in that they are low melting soft metals. This softness is related to the elements having few valence electrons for metallic bonding. They are also much more reactive at lower temperatures and have ionisation energies low enough for the formation of cations. The chemistry of aluminium is dominated by the 3+ oxidation state. Compounds in the +1 oxidation state are also known for Ga, In and Tl. Indeed, the +3 oxidation state become less stable relative to the +1 state down the group. In aqueous solutions, thallium exist almost exclusively as Th^I.

When we look at the elements of group 13, we noticed some interesting patterns in their physical properties. On going down a given group of the periodic table we expect the ionisation energies to fall as outer electrons are accommodated in successively higher shells and are held less tightly. This trend can be seen going from boron to aluminium, however, the ionisation energies for the next element down, gallium, are higher. The rise in ionisation energy from Al to Ga is due to the different arrangements of the inner electrons for the two. Inner electrons somewhat shield the outermost electrons from the positive charge of the nucleus. This isn’t perfect however, and so the outermost electrons feel what is called an effective nuclear charge. When we build up our electrons for aluminium, the preceeding element is magnesium where we have filled the 3s orbital. One row down underneath magnesium as have calcium but to get to gallium we have the first row of the d block. Because we have an extra ten elements and ten electrons which do not effectively shield the nuclear charge, the outer electrons of gallium feel a greater effective nuclear charge than those for aluminium and so are harder to remove. This also has an effect on the atomic radius. Since the outer electron of gallium are more tightly held than those of aluminium, the atoms are also smaller. This effect is called the d block contraction. On going to indium we observe the expected reduction in ionisation energy and increase in atomic size. When we reach thallium however we again see a rise in ionisation energy and see that the atomic radius is only marginally larger than for indium. This is similar effect to what we saw for gallium but is due to the filling of the 4f orbitals of the lanthanide elements.

Oxides

The only oxide of aluminium is alumina, Al₂O₃, which comes in a variety of hydrated and anhydrous forms and also occur in mineral. They are all white or transparent. α-Al₂O₃ and γ-Al₂O₃ are the two anhydrous forms and differ in the arrangment of the Al and O atoms. α-Al₂O₃ is a hard substance that is stable at high temperatures and resist hydration by water and reaction with acids. γ-Al₂O₃ on the other hand will readily take up water and dissolve in acidic solutions. Hydrated forms with the formula Al(O)OH exist in a variety of structure in many minerals. The hydroxide Al(OH)₃ doesn’t exist in mineral but can be produced as a precipitate by bubbling carbon dioxide through basic solutions of Al(OH)₄^-. Gallium and indium form oxides, Ga₂O₃ and In₂O₃, which are similar to alumina and also form similar hydroxides. The oxides of thallium are quite different however. The thallium(III) oxide, Tl₂O₃, is a brown-black powder that decomposes to the black tahllium(I) compound, Tl₂O on heating to 100 °C.

Halides

Almuinium, gallium and indium all form trifluorides, MF₃ which are ionic in nature and have high melting points (~ 1000 °C). The chlorides, bromides and iodides of these metals are covalent and much more volatile having much lower melting points. They exist as dimeric molecules with the formula M₂X₆ using two halide atoms to bridge the metals which have tetrahedral configuration, see figure 1. They are also soluble in nonpolar organic solvents.

Figure 1: Structure of In₂I₆, typical of the trihalides of group 13.

I I I
\ / \ / 
In In
/ \ / \ 
I I I

The trihalides of thallium are much less stable than those of the rest of the group. Although TlF₃ is stable, the chloride, Th^IIICl₃ decomposes through loss of Cl₂ at about 40 °C to give the monohalide Th^ICl. The tribromide also decomposes losing Br₂ but give a compound with the formula [Th^IIIBr₄]^-[Th^I]⁺. The triiodide behave differently again. This has the formula ThI₃ but the thallium is Th^I rather than Th^III. In this compound the iodides exist as the linear I₃^- molecular ion which has the structure [ I-I-I ]^-. This is because thallium(III) is too oxidising to form the expected triiodide and will oxidise two I^- ions to form iodine, I₂. The third iodide ion then combines with the iodine molecule to give the observed molecular anion. Strangely however, the +3 oxidation states can be achieved on addition of I^- to this compound to produce [Th^IIII₄]^-.

Much like the trihalides of boron, the trihalides of the metals of group 13 are also Lewis acids and form adducts with electron donor molecules and halide ions. This properties is employed in many reactions such as the Friedel-Crafts reactions. In these reactions hydrogen atoms on organic aromatic molecules such as benzene are replaced by an alkyl or acyl organic group. This is achieved by generating an organic cation, by removing a hailde using an aluminium compound such as AlCl₃, which then replaces H⁺.

RCl + AlCl₃→ R⁺ + AlCl₄^-

R⁺ + C₆H₆→ C₆H₅R + H⁺

Gallium and indium will also form some metal(I) monohalides but these are most stable for thallium. They are produced by the addition of halide ion salts to acidic solutions of soluble thallium(I) compounds. The chloride and bromide are light sensitive compounds much like silver bromide.

Aqueous Chemistry

The metals of group 13 all form octahedral hexaaqua ions, [M(H₂O)₆]³⁺ which occur in hydrated salts such as halides, sulphates, nitrates and perchlorates and in solution these are acidic liberating H⁺

[M(H₂O)₆]³⁺→ [M(H₂O)₅(OH)]²⁺ + H⁺