lithium [Gr.,=stone], metallic chemical element; symbol Li; at. no. 3; at. wt. 6.941; m.p. about 180.54°C;; b.p. about 1,342°C;; sp. gr. .534 at 20°C;; 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 hydroxide is used to increase the capacity of alkaline storage cells. 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.
Lithium is a chemical element with the symbol Li and atomic number 3. It is a soft alkali metal with a silver-white color. Under standard conditions, it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive, corroding quickly in moist air to form a black tarnish. For this reason, lithium metal is typically stored under the cover of oil. When cut open, lithium exhibits a metallic lustre, but contact with oxygen quickly returns it back to a dull silvery grey color. Lithium is also highly flammable.

According to theory, lithium (mostly 7Li) was one of the few elements synthesized in the Big Bang, although its quantity has vastly decreased. The reasons for its disappearance and the processes by which new lithium is created continue to be important matters of study in astronomy. Lithium is tied with krypton as 32nd or 33rd most abundant element in the cosmos (see Cosmochemical Periodic Table of the Elements in the Solar System), being less common than any element before Rubidium (element 37) except for scandium, gallium, arsenic, and bromine, yet more common than any element beyond krypton (element 36).

Due to its high reactivity it only appears naturally on Earth in the form of compounds. Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from brines and clays; on a commercial scale, lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

Trace amounts of lithium are present in the oceans and in some organisms, though the element serves no apparent biological function in humans. Nevertheless, the neurological effect of the lithium ion Li+ makes some lithium salts useful as a class of mood stabilizing drugs. Lithium and its compounds have several other commercial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, and lithium batteries. Lithium also has important links to nuclear physics: the splitting of lithium atoms was the first man-made form of a nuclear reaction, and lithium deuteride serves as the fusion fuel in staged thermonuclear weapons.

History and etymology

Petalite (lithium aluminium silicate) was first described in 1800 by the Brazilian (then Portuguese) scientist José Bonifácio de Andrade e Silva, who discovered the mineral in a Swedish iron mine on the island of Utö. However, it was not until 1817 that Johan August Arfwedson, then a trainee in the laboratory of Jöns Jakob Berzelius, discovered the presence of a new element while analyzing petalite ore. The element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less water soluble and had a larger capacity to neutralize acid. Berzelius gave the alkaline material the name "lithos", from the Greek λιθoς (lithos, "stone"), to reflect its discovery in a mineral, as opposed to sodium and potassium which had been discovered in plant tissue; its name would later be standardized as "lithium". Arfwedson later showed that this same element was present in the mineral ores spodumene and lepidolite. In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color in flame. However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts. The element was not isolated until 1821, when William Thomas Brande performed electrolysis on lithium oxide, a process which had previously been employed by Sir Humphry Davy to isolate potassium and sodium. Brande also described pure salts of lithium, such as the chloride, and performed an estimate of its atomic weight. In 1855, Robert Bunsen and Augustus Matthiessen produced large quantities of the metal by electrolysis of lithium chloride. Commercial production of lithium metal began in 1923 by the German company Metallgesellschaft AG through the electrolysis of a molten mixture of lithium chloride and potassium chloride.


Like other alkali metals, lithium has a single valence electron which it will readily lose to form a cation, indicated by the element's low electronegativity. As a result, lithium is easily deformed, highly reactive, and has lower melting and boiling points than most metals. These and many other properties attributable to alkali metals' weakly held valence electron are most distinguished in lithium, as it possesses the smallest atomic radius and thus the highest electronegativity of the alkali group.

In addition, lithium has a diagonal relationship with magnesium, an element of similar atomic and ionic radius. Chemical resemblances between the two metals include the formation of a nitride by reaction with N2, the formation of an oxide when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.

Lithium is soft enough to be cut with a knife, though this is more difficult than cutting sodium. The fresh metal has a silvery-white color which only remains untarnished in dry air. Lithium has about half the density of water, giving solid sticks of lithium metal the odd heft of a light-to-medium wood such as pine. The metal floats highly in hydrocarbons; in the laboratory, jars of lithium are typically composed of black-coated sticks held down in hydrocarbon mechanically by the jar's lid and other sticks.

Lithium is greatly heat-resistant, possessing a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium has also been found to be superconductive below 400 μK. This finding paves the way for further study of superconductivity, as lithium's atomic lattice.

At cryogenic temperatures, lithium, like sodium, undergoes martensitic transformations when cooled below liquid nitrogen temperatures (77 oK). This is a multicrystalline state composed of FCC (Face Centered Cubic), BCC (Body Centered Cubic) and R9 Hex. At liquid helium temperatures (4 oK) 9R Hex is the most prevalent crystal. The proportion of the different crystalline states is temperature-dependent. On cryogenic cooling lithium at atmospheric pressure, the crystalline state which first predominates is FCC, followed by BCC, followed by R9 Hex at the coldest temperatures. On heating solid lithium from deep cryogenic temperatures, the property known as heat of reversion will cause a crystalline state transition from R9 Hex to BCC, which absorbs heat and causes cooling. In warming at cryogenic peratures, there is thus a region of negative specific heat in lithium crystals, due to this state change.


In moist air, lithium metal rapidly tarnishes to form a black coating of lithium hydroxide (LiOH and LiOH·H2O), lithium nitride (Li3N) and lithium carbonate (Li2CO3, the result of a secondary reaction between LiOH and CO2).

When placed over a flame, lithium gives off a striking crimson color, but when it burns strongly, the flame becomes a brilliant white. Lithium will ignite and burn in oxygen when exposed to water or water vapours. It is the only metal that reacts with nitrogen at room temperature.

Lithium metal is flammable and potentially explosive when exposed to air and especially water, though it is far less dangerous than other alkali metals in this regard. The lithium-water reaction at normal temperatures is brisk but not violent. Lithium fires are difficult to extinguish, requiring special chemicals designed to smother them (see sodium for details).


Naturally occurring lithium is composed of two stable isotopes 6Li and 7Li, the latter being the more abundant (92.5% natural abundance). Seven radioisotopes have been characterized, the most stable being 8Li with a half-life of 838 ms and 9Li with a half-life of 178.3 ms. All of the remaining radioactive isotopes have half-lives that are shorter than 8.6 ms. The shortest-lived isotope of lithium is 4Li which decays through proton emission and has a half-life of 7.58043x10-23 s.

7Li is one of the primordial elements or, more properly, primordial isotopes, produced in Big Bang nucleosynthesis (a small amount of 6Li is also produced in stars). Lithium isotopes fractionate substantially during a wide variety of natural processes, including mineral formation (chemical precipitation), metabolism, and ion exchange. Lithium ion substitutes for magnesium and iron in octahedral sites in clay minerals, where 6Li is preferred to 7Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic 11Li is known to exhibit a nuclear halo.

Natural occurrence

Lithium is widely distributed on Earth, however, it does not naturally occur in elemental form due to its high reactivity. Estimates for crustal content range from 20 to 70 ppm by weight. In keeping with its name, lithium forms a minor part of igneous rocks, with the largest concentrations in granites. Granitic pegmatites also provide the greatest abundance of lithium-containing minerals, with spodumene and petalite being the most commercially-viable mineral sources for the element.

According to the Handbook of Lithium and Natural Calcium, "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively a few of them are of actual or potential commercial value. Many are very small, others are too low in grade.


Because of its specific heat capacity, the highest of all solids, lithium is often used in heat transfer applications.

It is an important ingredient in anode materials, used in rechargeable and primary batteries because of its high electrochemical potential, light weight, and high current density.

Large quantities of lithium are also used in the manufacture of organolithium reagents, especially n-butyllithium which has many uses in fine chemical and polymer synthesis.

Medical use

Lithium salts were used during the 19th century to treat gout. Lithium salts such as lithium carbonate (Li2CO3), lithium citrate, and lithium orotate are mood stabilizers. They are used in the treatment of bipolar disorder, since unlike most other mood altering drugs, they counteract both mania and depression. Lithium can also be used to augment other antidepressant drugs. It is also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.

The active principle in these salts is the lithium ion Li+, which having a smaller diameter, can easily displace K+ and Na+ and even Ca2+, in spite of its greater charge, occupying their sites in several critical neuronal enzymes and neurotransmitter receptors. Although Li+ cannot displace Mg2+ and Zn2+, because of these ions' small size and greater charge (higher charge density, hence stronger bonding), when Mg2+ or Zn2+ are present in low concentrations, and Li+ is present in high concentrations, the latter can occupy sites normally occupied by Mg2+ or Zn2+ in various enzymes. Therapeutically useful amounts of lithium (~ 0.6 to 1.2 mmol/l) are only slightly lower than toxic amounts (>1.5 mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity.

Common side effects of lithium treatment include muscle tremors, twitching, ataxia, hyperparathyroidism, bone loss, hypercalcemia, hypertension, etc.), kidney damage, nephrogenic diabetes insipidus (polyuria and polydipsia) and seizures. Many of the side-effects are a result caused by the increased elimination of potassium.

Pregnancy - teratogenic properties: Ebstein (cardiac) Anomaly - There appears to be an increased risk of this abnormality in infants of women taking lithium during the first trimester of pregnancy

Other uses

  • Lithium batteries are disposable (primary) batteries that have lithium metal or lithium compounds as an anode. Lithium batteries are not to be confused with lithium-ion batteries which are high energy-density rechargeable batteries
  • Lithium chloride and lithium bromide are extremely hygroscopic and frequently used as desiccants.
  • Lithium stearate is a common all-purpose high-temperature lubricant.
  • Lithium is an alloying agent used to synthesize organic compounds.
  • Lithium is used as a flux to promote the fusing of metals during welding and soldering. It also eliminates the forming of oxides during welding by absorbing impurities. This fusing quality is also important as a flux for producing ceramics, enamels, and glass.
  • Lithium is sometimes used in glasses and ceramics including the glass for the 200-inch (5.08 m) telescope at Mt. Palomar.
  • Alloys of the metal with aluminium, cadmium, copper and manganese are used to make high performance aircraft parts.
  • Lithium-aluminium alloys are used in aerospace applications, such as the external tank of the Space Shuttle, and is planned for the Orion spacecraft.
  • Lithium niobate is used extensively in telecommunication products, such as mobile phones and optical modulators, for such components as resonant crystals. Lithium products are currently used in more than 60 percent of mobile phones.
  • The high non-linearity of lithium niobate also makes a good choice for non-linear optics applications.
  • Lithium deuteride was the fusion fuel of choice in early versions of the hydrogen bomb. When bombarded by neutrons, both 6Li and 7Li produce tritium—this reaction, which was not fully understood when hydrogen bombs were first tested, was responsible for the runaway yield of the Castle Bravo nuclear test. Tritium fuses with deuterium in a fusion reaction that is relatively easy to achieve. Although details remain secret, lithium-6 deuteride still apparently plays a role in modern nuclear weapons, as a fusion material.
  • Metallic lithium and its complex hydrides such as e.g. Li[AlH4] are considered as high energy additives to rocket propellants[3].
  • Lithium peroxide, lithium nitrate, lithium chlorate and lithium perchlorate are used and thought of as oxidizers in both rocket propellants and oxygen candles to supply submarines and space capsules with oxygen.
  • Lithium fluoride (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred fluoride salt mixture (LiF-BeF2) used in liquid-fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF/BeF2 mixtures have low melting points and the best neutronic properties of fluoride salt combinations appropriate for reactor use.
  • Lithium will be used to produce tritium in magnetically confined nuclear fusion reactors using deuterium and tritium as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium. 6Li + n → 4He + 3H. Various means of doing this will be tested at the ITER reactor being built at Cadarache, France.
  • Lithium is used as a source for alpha particles, or helium nuclei. When 7Li is bombarded by accelerated protons, 8Be is formed, which undergoes spontaneous fission to form two alpha particles. This was the first man-made nuclear reaction, produced by Cockroft and Walton in 1929.
  • Lithium hydroxide (LiOH) is an important compound of lithium obtained from lithium carbonate (Li2CO3). It is a strong base, and when heated with a fat, it produces a lithium soap. Lithium soap has the ability to thicken oils and so is used commercially to manufacture lubricating greases.
  • It is also an efficient and lightweight purifier of air. In confined areas, such as aboard spacecraft and submarines, the concentration of carbon dioxide can approach unhealthy or toxic levels. Lithium hydroxide absorbs the carbon dioxide from the air by reacting with it to form lithium carbonate. Any alkali hydroxide will absorb CO2, but lithium hydroxide is preferred, especially in spacecraft applications, because of the low formula weight conferred by the lithium. Even better materials for this purpose include lithium peroxide (Li2O2) that, in presence of moisture, not only absorb carbon dioxide to form lithium carbonate, but also release oxygen. E.g. 2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.
  • Lithium metal is used as a reducing agent in some types of methamphetamine production, particularly in illegal amateur “meth labs.”
  • Lithium can be used to make red fireworks


Since the end of World War II, lithium metal production has greatly increased. The metal is separated from other elements in igneous mineral such as those above, and is also extracted from the water of mineral springs.

There are wide-spread hopes of using lithium ion batteries in electric vehicles, but one study concluded that "realistically achievable lithium carbonate production will be sufficient for only a small fraction of future PHEV and EV global market requirements", that "demand from the portable electronics sector will absorb much of the planned production increases in the next decade", and that "mass production of lithium carbonate is not environmentally sound, it will cause irreparable ecological damage to ecosystems that should be protected and that LiIon propulsion is incompatible with the notion of the 'Green Car'".

The metal is produced electrolytically from a mixture of fused lithium and potassium chloride. In 1998 it was about US$ 43 per pound ($95 per kg).

Chile is currently the leading lithium metal producer in the world, with Argentina next. Both countries recover the lithium from brine pools. In the United States lithium is similarly recovered from brine pools in Nevada.

China may emerge as a significant producer of brine-based lithium carbonate around 2010. Potential capacity of up to 55,000 tonnes per year could come on-stream if projects in Qinghai province and Tibet proceed.

The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tonnes of the known global lithium reserve base.

In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tonnes for the Western World. The inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total has nearly tripled by 2008.


Lithium metal, due to its alkaline tarnish, is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) can irritate the nose and throat; higher exposure to lithium can cause a build-up of fluid in the lungs, leading to pulmonary edema. The metal itself is usually a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium should be stored in a non-reactive compound such as naphtha or a hydrocarbon.


Some jurisdictions limit the sale of lithium batteries, which are the most readily available source of lithium metal for ordinary consumers. Lithium can be used to reduce pseudoephedrine and ephedrine to methamphetamine in the Birch reduction method, which employs solutions of alkali metals dissolved in anhydrous ammonia. However, the effectiveness of such restrictions in controlling illegal production of methamphetamine remains indeterminate and controversial.

Carriage and shipment of some kinds of lithium batteries may be prohibited aboard certain types of transportation (particularly aircraft), because of the ability of most types of lithium batteries to fully discharge very rapidly when short-circuited, leading to overheating and possible explosion. However, most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents.

See also


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