Group 13 of the Periodic Table contains the elements boron, aluminium, gallium, indium and thallium. The 'Header' element of each Group in the Periodic Table often displays properties anomalous to the rest of the Group. Group 13 is no exception in that this is the first group of the Periodic Table to contain a non-metal: ie, boron. The remaining four elements of group 13 - aluminium, gallium, indium and thallium, sometimes known as the 'Poor Metals' - have markedly different physical and chemical properties from boron. For this reason, boron and aluminium appear to be the most dissimilar elements in the Group, the only obvious similarity being their electronic configurations (see box at right which shows, for each element, its atomic mass, chemical symbol, name, atomic number and electron configuration).
10.81 B 26.98 Al 69.72 Ga 114.82 In 204.38 Tl
Occurrence of the elements
Except for the very abundant aluminium, all the elements in this group are rare. Aluminium is the most abundant metal in the Earth, making up about 8% of the Earth's crust and occurring in igneous rocks such as feldspars and micas. It also occurs in many clays and in minerals such as cryolite (Na3AlF6) and spinel (MgAl2O4) but the most important mineral is bauxite1 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 amounts of transition metal ions that give them their colour: for example, ruby (chromium) and sapphire (cobalt).
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 while 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.
Boron is a strange and fascinating element with some unique properties which make it extremely valuable for a wide variety of applications. It is a non-metal of high melting point and very low electrical conductivity.
Being a non-metal, it has a weakly acidic oxide which is soluble in water.
The main reason for the differences in the chemistries of boron and aluminium (and therefore the rest of the Group) is that boron has a small compact atom, while aluminium has a relatively large atom (see Column 5 of Table 1 below). Thus, unlike the rest of Group 13, boron is non-metallic. Its chemistry is actually much more similar to that of carbon or silicon in the neighbouring Group 14. The so-called 'diagonal relationship' to silicon is particularly striking. However, its unusual electrophilicity and other characteristics make boron chemically unique.
There is also a diagonal relationship between aluminium and beryllium in Group 2.
As the chemistry of boron is not typical of this Group, boron will be the subject of a separate Entry.
In historical terms, Man has unwittingly been using compounds of aluminium for thousands of years in the form of aluminosilicate pottery. This material came to be known as 'metal of clay' but, because aluminium is a relatively reactive metal, it was not possible to separate the metal from the other elements to which it was bonded. Indeed, it was the great French chemist Lavoisier who, in 1782, first realised that this clay was predominantly the oxide of a hitherto unknown metal.
Native aluminium metal itself had never ever existed until Sir Humphry Davy probably achieved it in 1809 when he reacted iron with alumina in an electric arc. This produced an iron-aluminium alloy and, for a brief instant aluminium2 would have existed in its free metallic state.
In 1825 HC Oerstedt, a Dane, produced a tiny sample of aluminium in the laboratory by chemical means. Twenty years later the German scientist Frederick Wohler produced aluminium lumps as big as pinheads. By 1854 Sainte-Clair Deville had made improvements in Wohler's method and produced aluminium globules the size of marbles. Sainte-Clair Deville was encouraged by Napoleon III to produce aluminium commercially and at the Paris exhibition in 1855 aluminium bars were exhibited next to the crown jewels.
It was not until 31 years later, however, that an economical means of commercial production was discovered. On 23 February, 1886, working with home-made apparatus in the family woodshed, a 22-year-old American undergraduate named Charles Martin Hall passed an electric current through a solution of cryolite and alumina, thus achieving a separation of aluminium from the oxygen with which it is combined.
At about the same time, a French man, Paul LT Herault, achieved the same result using a similar process. However, being primarily interested in aluminium alloys, he did not immediately recognise its importance.
In 1888, the German chemist Karl Joseph Bayer was issued a German patent for an improved process for making Bayer Aluminium Oxide (alumina), thus laying the foundations for what is now sometimes referred to as the 'Aluminium Age'. The Bayer and Hall-Herault processes enabled economic production of the world's most plentiful and versatile structural element.
Commercial Production of Aluminium
Aluminium is quite electropositive and reactive, coming above zinc in the Reactivity Series of metals. It cannot therefore be extracted from its ore (bauxite) using carbon. It is, in fact, produced on an industrial scale by the electrolysis of bauxite; a process which only became possible in the 19th Century following Faraday's discovery of magnetic induction of electricity, thus allowing generation on a commercial scale. The electrolytic production of aluminium is now known as the Bayer-Hall-Herault process.
In this process, the electrolyte - in this case, bauxite - needs to be molten to allow the passage of the electric current. The melting point of bauxite (aluminium oxide) is very high, in excess of 2000°C. This means large amounts of energy would be needed to melt it, and energy costs money! For this reason some cryolite (aluminium fluoride - a less common ore of aluminium) is mixed in and this reduces the melting point to just below 1000°C. Aluminium metal is produced at the negative electrode (cathode) and oxygen gas is liberated at the anode. Both electrodes are made of graphite, an allotrope of carbon, and, at the high temperature used in the electrolysis, these gradually burn away and have to be continually replaced.
Before electrolytic techniques became available, aluminium was one of the most expensive metals in the world and, as described above, was even exhibited beside the crown jewels at the Paris exhibition in 1855. Napoleon III also had a set of aluminium tableware for the use of his most important guests, while guests of lower rank or status had to make do with gold or silverware. The statue popularly known as Eros in London's Piccadilly Circus, set up in 1893 to honour Lord Shaftesbury, was intended to be a prestigious monument and was therefore cast in the exotic and precious metal, aluminium. However, after the introduction of electrolytic methods, the price of aluminium fell by over a thousandfold over the course of a century.
Properties and Uses
Aluminium is an extremely important metal affecting our daily lives. It is a strong, light metal with a Mohs hardness. of 2.75. It has wide application in the construction industry and, being a light metal, it is used to build cars and aircraft.
Its electrical conductivity is about 60% that of copper, but its lightness and price favour its use in electrical transmission lines.
It also has familiar applications in that it is used for kitchen foil and in drink cans.
Being high in the Reactivity Series, aluminium powder is used in the extraction of high melting point metals from their ores using the thermite reaction.
Despite being relatively high in the Reactivity Series, aluminium avoids corrosion as it quickly forms a hard surface film of oxide, which is impenetrable to moisture. It will, however, dissolve in aqueous sodium hydroxide and will form compounds with most non-metal elements with heating.
The melting point of aluminium is fairly low (660°C).
Aluminium oxide (alumina) is amphoteric and is insoluble in water.
Aluminium forms both ionic and covalent compounds, and thus the oxide, fluoride and sulphate are predominantly ionic while the chloride is predominantly covalent.
Boron and aluminium share some properties in that they each form complexes.
Role of Aluminium in Biology
The ubiquity of aluminium in nature would suggest a biological function yet, until recently, no specific function had been found. However, aluminium compounds are toxic to most plants and animals - in animals they act as neurotoxins. Due to its high reactivity, Al3+ is able to interfere with several biological functions, including enzymatic activities in key metabolic pathways, including Krebs Cycle enzymes such as succinic dehydrogenase.
Aluminium ions (Al3+) in aluminosilicate soils are mobilised by acid rain, and may drain into streams and ponds. This causes a sticky mucus to accumulate in the gills of fish and eventually kills them. Trees and other plants which absorb Al3+ ions will be damaged.
For some years now there has been concern about the possible role of aluminium in a number of neurological disorders such as Alzheimer's Disease. Scientists have observed increased levels of aluminium in the brain tissues of some patients suffering from Alzheimer's Disease, amyotrophic lateral sclerosis and Parkinson's Disease. Although various hypotheses have been put forward to explain this, there is insufficient evidence to say that aluminium is causative.
If aluminium is a cause, then it could enter our bodies in a number of ways. An obvious route is through diet. If food is cooked or stored in aluminium containers then some of the aluminium may well dissolve and be ingested. Acidic foods such as fruit and vegetable juices (especially rhubarb), tomatoes and sauerkraut tend to increase the aluminium content more than other foods. However, it is considered that the amount of aluminium ingested in this way is very small compared to that naturally present in certain foodstuffs or present in additives.
Foods which are naturally high in aluminium include tea leaves and some herbs and spices. However, the intake from these is again low because only small quantities of these foodstuffs are generally consumed.
Food additives which contain aluminium include anti-caking agents calcium aluminium silicate (E556) sodium aluminosilicate (E554) and the baking powder raising agent, sodium aluminium phosphate (E541).
However, the amount of aluminium absorbed from foodstuffs, whether naturally present or from additives, is likely to be low compared to that ingested from pharmaceutical products such as antacids, buffered analgesics and anti-diarrhoeal agents. For example, some indigestion tablets are pure aluminium hydroxide!
Aluminium in the form of aluminium chlorohydrate is also present in antiperspirant formulations.
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.
Aluminium in the form of Al3+ may also enter our bodies through drinking tap water, where there may be residues following water treatment (see below).
Properties and Uses of Gallium
Gallium is one of three elements that naturally occur as a liquid at or close to room temperature, the other two being mercury and caesium. With a melting point of 29.76 °C and a boiling point of 2204 °C, gallium has a liquid range of 2107°C, one of the largest liquid ranges of any metal, so it has found use in high temperature thermometers. As it easily forms alloys with most metals it has been used to create low melting alloys, and is also used in low temperature solders.
Gallium is an important element in the electronics industry and has applications as doping material in semiconductors. One of the most important of these materials in gallium arsenide, GaAs, which can produce laser light directly from electricity and is used in diode lasers and LEDs. Some of us may be old enough to recall the original digital watches, with their bright red or green displays (due to GaAs or gallium nitride, GaN, respectively) which appeared at the touch of a button.
The fluorescent compound MgGa2O4 is used in photocopiers.
Large amounts of gallium trichloride (GaCl3; 90 tons or two to three years' supply) have been gathered to build the Gallium Neutrino Observatory (GNO), an observatory located in Italy built to detect solar neutrinos which are produced inside the sun during the process of nuclear fusion. The reaction:
nu + 7131Ga ⇒ 7132>Ge + e-
has a very low detection rate (less than one interaction per day in 30 tonnes of gallium) thus making gallium unique for this purpose.
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 dissolution in warm dilute acids. It is then separated from impurities and purified by electrolysis and deposition. Thallium and all its compounds are extremely toxic. The element itself has no major uses; however the thallium(I) compound Tl2SO4 was once used as rodenticide and ant killer. Its use for these purposes has since been banned in most countries due to its odourless and tasteless nature making the risk of accidental poisoning unacceptably high.
The electrical conductivity of thallium sulphide changes with exposure to infrared light, and so this compound is used in photocells. Thallium oxide has been used to produce glasses with a high refractive index, and is also used in the manufacture of photocells.
At one time, thallium sulphate was used in medicine as a depilatory agent, and thallium carbonate was used to treat mildew in textiles.
Thallium bromide-iodide crystals have been used as infrared optical materials. Thallium has been used in treating ringworm and other skin infections; however, its use has been limited because of the narrow margin between toxicity and therapeutic benefits.
Table 1: Physical properties of the group 13 elements. Boron is included for comparison.
point / °C
point / °C
|1st, 2nd& 3rdIonisation
energies / kJ mol-1
|Metal radius / pm 3||M3+ ionic
radius / pm
|B||2180||c.3650||801, 2427, 3659||80-90||-|
|Al||660||2467||577, 1816, 2744||143||54|
|Ga||30||2403||579, 1979, 2962||135||62|
|In||156||2080||558, 1820, 2704||167||80|
|Tl||304||1357||589, 1970, 2877||170||89|
The last four 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 exists almost exclusively as ThI.
There are some interesting patterns in the physical properties of the Group 13. On going down any given group of the periodic table one expects 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, but 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 the electronic configuration for aluminium is built up, the preceding element is magnesium where the 3s orbital is full. Calcium is one row down directly underneath magnesium, but to get to gallium the first row of the d block has to be filled. Because there is 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 electrons 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 the expected reduction in ionisation energy and increase in atomic size is observed. When thallium is reached, however, a rise in ionisation energy is observed but the atomic radius is only marginally larger than for indium. This is a similar effect to that observed for gallium but is due to the filling of the 4f orbitals of the lanthanide elements.
Some Important Salts
The only oxide of aluminium is alumina, Al2O3, which comes in a variety of hydrated and anhydrous forms, and also occurs in minerals. They are all white or transparent. α-Al2O3 and γ-Al2O3 are the two anhydrous forms and differ in the arrangement of the Al and O atoms. α-Al2O3 is a hard substance that is stable at high temperatures and resists 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 structures in many minerals. The hydroxide Al(OH)3 doesn't exist in minerals 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 thallium(I) compound, Tl2O on heating to 100 °C.
Aluminium, 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, ThIIICl3 decomposes through loss of Cl2 at about 40 °C to give the monohalide ThICl. The tribromide also decomposes losing Br2 but gives a compound with the formula [ThIIIBr4]-[ThI]+. The triiodide behaves differently again. This has the formula ThI3 but the thallium is ThI rather than ThIII. 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 [ThIIII4]-.
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 property 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 halilde 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.
Aluminium (III) sulphate is sometimes added by water companies to water supplies in order to remove fine particles, colour and bacteria. As aluminium sulphate has an acidic pH, water companies usually adjust the water to a pH of between 7 and 8 and, under these alkaline conditions aluminium hydroxide precipitates out as fine solid particles which can then be removed by means of sand filters. A mishap with this process was responsible for Britain's worst mass water poisoning incident at Camelford, Cornwall in 1988.
Thallium (I) sulphate Tl2SO4 adopts the same structure as potassium sulphate, K2SO4. Thallium (I) sulphate is soluble in water and is highly toxic since the thallium (I) cation is very similar to potassium and sodium cations, which are essential for life. Once the thallium ion enters the cell, many of the processes that transport potassium and sodium are disrupted. Once it has entered the body, thallium sulphate concentrates in the kidneys, liver, brain, and other body tissues. Thallium (I) sulphate has been used as a rodenticide, and was Saddam Hussein's poison of choice for dealing with Iraqi dissidents.
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+
For this reason, solutions of the salts are acidic; indeed, solutions of aluminium salts, for example, are as acidic as vinegar.