Lithium: the essentials
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 Read more »
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.
|Location||ppb by weight||ppb by atoms||Links|
|Human||30 ppb by weight||27 atoms relative to C = 1000000|
Physical properties Read more »
Heat properties Read more »
- Melting point: 453.69 [180.54 °C (356.97 °F)] K
- Boiling point: 1615 [1342 °C (2448 °F)] K
- Enthalpy of fusion: 3.0 kJ mol-1
Crystal structure Read more »
The solid state structure of lithium is: bcc (body-centred cubic).
Lithium: orbital properties Read more »
Lithium atoms have 3 electrons and the shell structure is 2.1. The ground state electronic configuration of neutral Lithium is [He].2s1 and the term symbol of Lithium is 2S1/2.
- Pauling electronegativity: 0.98 (Pauling units)
- First ionisation energy: 520.2 kJ mol‑1
- Second ionisation energy: 7298.1 kJ mol‑1
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 Read more »
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.
|6Li||6.015 122 3(5)||[7.59 (4)]||1||0.8220467|
|7Li||7.016 004 0(5)||[92.41 (4)]||3/2||3.256424|