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lithium (lĭthˈēəm) [Gr.,=stone], metallic chemical element; symbol Li; at. no. 3; interval in which at. wt. ranges 6.938–6.997; m.p. about 180.54℃; b.p. about 1,342℃; sp. gr. .534 at 20℃; valence +1. Lithium is a soft, silver-white metal. It is one of the alkali metals in Group 1 of the periodic table. It is the least dense metal. Because it has high specific heat, it has found some use in cooling systems for nuclear reactors; such use is limited because lithium is very corrosive. Lithium metal is prepared by electrolysis of fused lithium chloride. Lithium reacts with water less readily than sodium. It burns in air with a brilliant white flame. Lithium forms many inorganic compounds, among them a hydride (LiH), a nitride (Li3N), an oxide (lithia, Li2O), a hydroxide (LiOH), a carbide (Li2C2), a carbonate (Li2CO3), and a phosphate (Li3PO4). When heated it reacts directly with the halogens to form halides. Lithium aluminum hydride (LiAlH4) is an important reagent in organic chemistry. Lithium also forms numerous organic compounds. One compound of major importance is lithium stearate, produced by cooking tallow (or other animal fat) with lithium hydroxide; lithium stearate is used to transform oil into lithium-base lubricating greases, which have found extensive use in the automotive industry. Lithium carbonate is used in special glasses and ceramic glazes. Lithium chloride and bromide are used as brazing and welding fluxes; they are also used in air conditioning systems because they are very hygroscopic, i.e., they absorb moisture. Lithium is also used in electric storage cells, or batteries; it is used in disposible, typically button-shaped batteries and in rechargeable lithium-ion batteries, which are widely used in portable electronic devices. Lithium compounds are used in the nuclear energy industry, in the preparation of plastics and synthetic rubber, and in the synthesis of vitamin A. Lithium is added in small amounts to magnesium, aluminum, or lead-base alloys; it is also used as a degasifier in iron, steel, and copper refining. In addition, lithium is used to scavenge small amounts of oxygen and nitrogen in electronic vacuum tubes. Trace amounts of lithium and its compounds color a flame bright red; they are used in pyrotechnics. Lithium in the salt form has recently come into use as a medical treatment for bipolar disorder. Lithium is widely distributed in nature; it is found in the soil, in plants, in animals, and in the human body. It is also found in the sun. Lithium may be profitably extracted from ores containing as little as 1% lithium (measured as lithium oxide). Some commercially important minerals are lepidolite, petalite, spodumene, and amblygonite. Lithium is also produced from brines such as those in Searles Lake, Calif., and in the Great Salt Lake, Utah. Lithium was discovered in 1817 by J. A. Arfvedson.
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The following article is from The Great Soviet Encyclopedia (1979). It might be outdated or ideologically biased.



Li, a chemical element in Group I of the Mendeleev periodic system. Atomic number, 3; atomic weight 6.941; an alkali metal. Natural lithium is composed of two stable isotopes, 6Li (7.42 percent) and 7Li (92.58 percent).

Lithium was discovered in the mineral petalite by the Swedish chemist J. A. Arfvedson in 1817; its name is derived from the Greek lithos, “stone.” In 1818 the British chemist H. Davy first obtained metallic lithium.

Occurrence in nature. Lithium is a characteristic element of the earth’s crust (content, 3.2 X 10 –3 percent by mass) and accumulates in the final products of the differentiation of magma (pegmatites). Only a small quantity of lithium occurs in the mantle: hyperbasic rocks contain a total of 5 X 10–5 percent (1.5 X 10–3 percent in basic rock, 2 X 10–3 percent in intermediate rock, and 4 X 10–3 in acidic rock). The similarity of ionic radii for Li+, Fe2+, and Mg2+ facilitates the inclusion of lithium in the lattices of ferromagnesian silicates (pyroxenes and amphiboles). Granitoids contain lithium as an isomorphic admixture in mica. Twenty-eight independent lithium minerals (silicates, phosphates, and so on) are known to exist only in pegmatites and in the biosphere; they are all rare. Lithium migrates relatively poorly through the biosphere and plays a smaller role in living matter than the other alkali metals. Lithium is readily extracted from water by clay; however, its content in the world’s oceans is comparatively low (1.5 X 10–5 percent). Commercial lithium deposits are associated with the occurrence of magmatic rocks (pegmatites and pneumatolites) and with the biosphere (salt lakes).

Physical and chemical properties. Solid lithium is a silver-white metal that rapidly becomes covered with a dark gray tarnish composed of lithium nitride, Li3N, and lithium oxide, Li2O. At normal temperature lithium crystallizes in a cubic body-centered lattice, a = 3.5098 angstroms (Å); atomic radius, 1.57 Å; ionic radius of Li+, 0.68 Å. At temperatures below — 195°C, the lattice is converted into the hexagonal close-packed form. Lithium is the lightest metal (density, 0.534 g/cm3 at 20°C). Melting point, 180.5°C; boiling point, 1317°C; specific heat at 0°-100°C, 3.31 X 103 joules per (kg.°K), or 0.790 cal/(g.deg); thermal coefficient of linear expansion at 0°-100°C, 5.6 X 10–5; specific electric resistance at 20°C, 9.29 X 10–8 ohm · m (9.29 microhms. cm); temperature coefficient of electric resistance at 0°-100°C, 4.50 X 10–3. Lithium is paramagnetic. The metal is extremely ductile and tough and thus is suitable for pressure molding, rolling, and drawing out into wire. Its hardness on the Mohs’ scale is 0.6 (harder than sodium and potassium), and it may be cut easily with a knife. Pressure of fluidity at 15°-20°C, 17 meganewtons per sq m (MN/m2), or 1.7 kilograms-force per sq mm (kgf/mm2); elastic modulus, 5 GN/m2 (500 kgf/mm2); tensile strength, 116 MN/m2 (11.8 kgf/mm2); relative elongation, 50–70 percent. Lithium vapors turn flames a carmine-red color.

The configuration of the outer electron shell of the lithium atom is 2s1; it is univalent in all known compounds. Lithium yields lithium oxide, Li2O (the peroxide Li2O2 may be obtained only by the indirect method), upon interaction with oxygen or upon heating in air (it burns with a blue flame). Lithium reacts less vigorously with water than other alkali metals; it forms lithium hydroxide, LiOH, and hydrogen. Inorganic acids dissolve lithium vigorously (it has the highest electrode potential of all metals—3.02 V).

Lithium combines with halogens (with iodine upon heating) to form halides, the most important of which is lithium chloride. Upon heating with sulfur, lithium produces lithium sulfide, Li2S; with hydrogen, lithium hydride, LiH. Lithium reacts slowly with nitrogen at room temperature and vigorously at 250° C to form lithium nitride, Li3N. Direct reaction of lithium with phosphorus is not achieved, although the phosphides Li3P, LiP, and Li2P2 may be obtained under specific conditions. Heating of lithium with carbon yields lithium carbide, Li2C2; heating with silicon yields lithium suicide. Binary lithium compounds (for example, Li2O, LiH, Li3N, Li2C2, and LiCl), as well as lithium hydroxide, LiOH, are highly reactive; upon heating or melting they destroy many metals as well as porcelain and quartz. Lithium carbonate, Li2CO3, lithium fluoride, LiF, lithium phosphate, Li3PO4, and other lithium compounds are similar in properties and conditions of formation to the corresponding derivatives of magnesium and calcium.

Lithium forms numerous organolithium compounds; this determines its large role in organic synthesis.

Lithium is a component of many alloys. It forms high-concentration solid solutions with several metals (magnesium, zinc, and aluminum) and intermetallic compounds (LiAg, LiHg, LiMg2, LiAl, and so on) with many metals. The intermetallic compounds are often extremely hard and refractory and are not significantly changed by exposure to air; several of them are semiconductors. Studies have been made of more than 30 binary and a number of ternary systems that contain lithium; the corresponding alloys have already come to be used in engineering.

Preparation and use. Lithium compounds are prepared by the hydrometallurgical processing of concentrates (products of the concentration of lithium ore). Spodumene, a basic silicate mineral, is processed by the lime, sulfate, and sulfuric-acid methods.

The lime method involves the decomposition of spodumene by limestone at 1150–1200°C:

Li2O · Al2O3 · 4SiO2 + 8CaCO3

= Li2O · Al2O3 + 4(2CaO · SiO2) + 8CO2

Upon lixiviation of the sinter with water in the presence of excess lime, the lithium aluminate dissociates to form lithium hydroxide:

Li2O· Al2O3 + Ca(OH)2 = 2LiOH + CaO · Al2O3

In the sulfate method, and other aluminosilicates are sintered with potassium sulfate:

Li2O · Al2O3 · 4SiO 2 + K2SO4

= Li2SO4 + K2O · Al2O3 · 4SiO2

The lithium sulfate is dissolved in water, and soda is used to precipitate lithium carbonate from the solution:

Li2 SO4 + Na2CO3 = Li2CO3 ↓ + Na2SO4

The sulfuric-acid method also involves the preparation first of a lithium sulfate solution and then of lithium carbonate; the spodumene is decomposed by the action of sulfuric acid at 250°-300°C (this reaction is applicable only to the β-variety of spodumene):

β-Li2O · Al2O3 · 4SiO2 + H2SO4

= Li2SO4 + H2O · Al2O3 · 4SiO2

This method is used for processing ores not dressed with spodumene, if their Li2O content is not less than 1 percent. Phosphate lithium materials are readily decomposed by acids; however, newer methods involve decomposition by a mixture of gypsum and lime at 950°–1050°C, with subsequent hydrous treatment of the sinter and precipitation of lithium carbonate from the solutions.

Metallic lithium is produced by electrolysis of a fused mixture of equal parts (by weight) of lithium and potassium chlorides at 400°–460°C. The electrolysis cells are lined with magnesite, alun-dum, mullite, talc, graphite, and other materials that are resistant to the fused electrolyte; graphite rods serve as the anode, and iron rods are used for the cathode. Raw metallic lithium contains mechanical inclusions and impurities (potassium, magnesium, calcium, aluminum, silicon, and iron, but primarily sodium). The inclusions are removed by remelting, and the impurities are eliminated by refining at reduced pressure. Metallothermic methods of lithium extraction are currently of prime interest.

The main area of application of lithium is nuclear power engineering. The isotope 6Li is the sole industrial source for the preparation of tritium according to the following reaction:

The thermal neutron capture cross sections (σ) for lithium isotopes differ sharply: for 6Li, 945; for 7Li, 0.033; and for a natural mixture, 67 (in barns). This is important in relation to the industrial use of lithium in the manufacture of control rods for reactor protection systems. Liquid lithium in the form of the isotope Li is used as a heat-transfer agent in uranium reactors. Fused 7LiF is used to dissolve compounds of uranium and thorium in homogeneous reactors. The largest consumer of lithium compounds is the silicate industry, which uses lithium minerals, LiF, Li2CO3, and many specially prepared compounds. In ferrous metallurgy, lithium and its compounds and alloys are widely used in the deoxidation, alloying, and modification of many types of alloys. In nonferrous metallurgy, lithium processing imparts good structural properties and ductility and high strength to alloys. Aluminum alloys with a total lithium content of 0.1 percent (Aeron and Scleron) are widely known; in addition to lightness, they have high strength, ductility, and corrosion resistance and are very promising for aircraft construction. The addition of 0.04 percent lithium to lead-calcium bearing alloys increases their hardness and reduces friction. Lithium compounds are used in the manufacture of plastic lubricants. Based on its importance in modern engineering, lithium is considered to be one of the most valuable rare elements.


Lithium in the organism. Lithium is always a component of living organisms, but its biological role has not yet been adequately explained. It has been established that lithium increases resistance to disease in plants and intensifies the photochemical activity of chloroplasts in tomato leaves and the synthesis of nicotine in tobacco plants. Among marine organisms, the ability to concentrate lithium is most pronounced in red and brown algae; among terrestrial plants, in the families Ranunculaceae (meadow rue and buttercup) and Solanaceae (matrimony vine). In animal organisms, lithium concentration occurs primarily in the liver and lungs.


Pliushchev, V. E., and B. D. Stepin. Khimiia i tekhnologiia soedinenii litiia, rubidiia i tseziia. Moscow, 1970.
Landolt, P., and M. Sittig. “Litii.” In Spravochnik po redkim metallam. Moscow, 1965. (Translated from English.)
The Great Soviet Encyclopedia, 3rd Edition (1970-1979). © 2010 The Gale Group, Inc. All rights reserved.


A chemical element, symbol Li, atomic number 3, atomic weight 6.939; an alkali metal.
McGraw-Hill Dictionary of Scientific & Technical Terms, 6E, Copyright © 2003 by The McGraw-Hill Companies, Inc.


a soft silvery element of the alkali metal series: the lightest known metal, used as an alloy hardener, as a reducing agent, and in batteries. Symbol: Li; atomic no.: 3; atomic wt.: 6.941; valency: 1; relative density: 0.534; melting pt.: 180.6?C; boiling pt.: 1342?C
Collins Discovery Encyclopedia, 1st edition © HarperCollins Publishers 2005