Lithium

Lithium is a Group 1 (IA) element containing just a single valence electron (1s²2s¹). Group 1 elements are called "alkali metals". Lithium is a solid only about half as dense as water and lithium metal is the least dense metal. A freshly cut chunk of lithium is silvery, but tarnishes in a minute or so in air to give a grey surface. Its chemistry is dominated by its tendency to lose an electron to form Li⁺. It is the first element within the second period.

Lithium is mixed (alloyed) with aluminium and magnesium for light-weight alloys, and is also used in batteries, some greases, some glasses, and in medicine.

Lithium does not occur as the free metal in nature because of its high reactivity. Deposits are known all aroun the world. It is a minor component of nearly all igneous rocks and is a component of many natural brines

Lithium: historical information

The mineral petalite (which contains lithium) was discovered by the Brazilian scientist José Bonifácio de Andrada e Silva towards the end of the 18th century while visiting Sweden. Lithium was discovered by Johan August Arfvedson in 1817 during an analysis of petalite ore, an ore now recognised to be LiAl(Si₂O₅)₂, taken from the Swedish island of Utö. Arfvedson subsequently discovered lithium in the minerals spodumene and lepidolite. C.G. Gmelin observed in 1818 that lithium salts colour flames bright red. Neither Gmelin nor Arfvedson were able to isolate the element itself from lithium salts, for example in attempted reductions by heating the oxide with iron or carbon.

The first isolation of elemental lithium was achieved later by W.T. Brande and Sir Humphrey Davy by the electrolysis of lithium oxide. In 1855, Bunsen and Mattiessen isolated larger quantities of the metal by electrolysis of lithium chloride.

In 1923 the first commercial production of lithium metal was achieved by Metallgesellschaft AG in Germany using the electrolysis of a molten mixture of lithium chloride and potassium chloride, exploiting a suggestion made by Guntz in 1893.

Lithium around us

Lithium seems to have no biological role, but does have an effect on the body if swallowed. Sometimes, lithium-based compounds such as lithium carbonate (Li₂CO₃) are used as drugs to treat manic-depressive disorders. The dose is around 0.5 g - 2 g daily.

Lithium does not occur as the free metal in nature because of its high reactivity. Deposits are known all aroun the world. It is a minor component of nearly all igneous rocks and is a component of many natural brines (see below). Large deposits are located in California and Nevada (both in the USA) in several rock forms, particularly spodumene. The four main lithium minerals are spodumene, lepidolite, petalite, and amblygonite.

• spodumene: LiAlSi₂O₆. This is the most important and abundant of the lithium ores. Deposits are located in North America, Brazil, USSR, Spain, parts of Africa, and Argentina. One method of extraction involves converting α-spodumene (the naturally occurring form) to the β-form (a less dense material) by heating to 1100°C, mixing with sulphuric acid, H₂SO₄, and heating to 250°C. This is followed by extracting into water to give a lithium sulphate solution, Li₂SO₄, suitable for further processing.
• lepidolite: K₂Li₃Al₄Si₇O₂₁(OH,F)₃. Deposits are located in Canada and parts of Africa. The mineral sometimes contains caesium and rubidium. Lithium can be extracted by similar methods to those of spodumene
• petalite: LiAlSi₄O₁₀. Deposits are located in parts of Africa and Sweden.
• amblygonite: LiAl(F,OH)PO)₄. Amblygonite occurs in only minor deposits

Lithium is also recovered from lakes such as Searles Lake (California, USA) and Clayton Valley (Nevada, USA). Lithium is extracted from the brine by solar evaporation, precipitation of Group 2 elements if necessary, and precipiation of lithium carbonate by addition of sodium carbonate to the hot brine.

Isolation

 lithium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative lithium ion Li⁺.

The ore spodumene, LiAl(SiO₃)₂, is the most important commercial ore containing lithium. The α form is first converted into the softer β form by heating to around 1100°C. This is mixed carefully with hot sulphuric acid and extracted into water to form lithium sulphate, Li₂SO₄, solution. The sulphate is washed with sodium carbonate, Na₂CO₃, to form a precipitate of the relatively insoluble lithium carbonate, Li₂CO₃.

Li₂SO₄ + Na₂CO₃ → Na₂SO₄ + Li₂CO₃ (solid)

Reaction of lithium carbonate with HCl then provides lithium chloride, LiCl.

Li₂CO₃ + 2HCl → 2LiCl + CO₂ +H₂O

Lithium chloride has a high melting point (> 600°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of LiCl (55%) and KCl (45%) melts at about 430°C and so much less energy and so expense is required for the electrolysis.

cathode: Li⁺(l) + e^- → Li (l)

anode: Cl^-(l) → ¹/₂Cl₂ (g) + e^-

Lithium isotopes

Li-7 is used to control the pH level of the coolant in the primary water circuit of pressurized water reactors. Li-7 is also used for the production of the medical research radioisotope Be-7. Li-6 is used in thermonuclear weapons and the export and use of Li-6 is therefore strictly controlled. Li-6 can also be used for the production of the radioisotope H-3, which is used in biochemistry research.