Lithium

Lithium
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This article is about the chemical element. For other uses, see Lithium (disambiguation).
Good article
helium ← lithium → beryllium
H
↑
Li
↓
Na
Element 1: Hydrogen (H), Other non-metal

Element 2: Helium (He), Noble gas
Element 3: Lithium (Li), Alkali metal

Element 4: Beryllium (Be), Alkaline earth metal

Element 5: Boron (B), Metalloid

Element 6: Carbon (C), Other non-metal

Element 7: Nitrogen (N), Other non-metal

Element 8: Oxygen (O), Other non-metal

Element 9: Fluorine (F), Halogen

Element 10: Neon (Ne), Noble gas
Element 11: Sodium (Na), Alkali metal

Element 12: Magnesium (Mg), Alkaline earth metal

Element 13: Aluminium (Al), Other metal

Element 14: Silicon (Si), Metalloid

Element 15: Phosphorus (P), Other non-metal

Element 16: Sulfur (S), Other non-metal

Element 17: Chlorine (Cl), Halogen

Element 18: Argon (Ar), Noble gas
Element 19: Potassium (K), Alkali metal

Element 20: Calcium (Ca), Alkaline earth metal

Element 21: Scandium (Sc), Transition metal

Element 22: Titanium (Ti), Transition metal

Element 23: Vanadium (V), Transition metal

Element 24: Chromium (Cr), Transition metal

Element 25: Manganese (Mn), Transition metal

Element 26: Iron (Fe), Transition metal

Element 27: Cobalt (Co), Transition metal

Element 28: Nickel (Ni), Transition metal

Element 29: Copper (Cu), Transition metal

Element 30: Zinc (Zn), Transition metal

Element 31: Gallium (Ga), Other metal

Element 32: Germanium (Ge), Metalloid

Element 33: Arsenic (As), Metalloid

Element 34: Selenium (Se), Other non-metal

Element 35: Bromine (Br), Halogen

Element 36: Krypton (Kr), Noble gas
Element 37: Rubidium (Rb), Alkali metal

Element 38: Strontium (Sr), Alkaline earth metal

Element 39: Yttrium (Y), Transition metal

Element 40: Zirconium (Zr), Transition metal

Element 41: Niobium (Nb), Transition metal

Element 42: Molybdenum (Mo), Transition metal

Element 43: Technetium (Tc), Transition metal

Element 44: Ruthenium (Ru), Transition metal

Element 45: Rhodium (Rh), Transition metal

Element 46: Palladium (Pd), Transition metal

Element 47: Silver (Ag), Transition metal

Element 48: Cadmium (Cd), Transition metal

Element 49: Indium (In), Other metal

Element 50: Tin (Sn), Other metal

Element 51: Antimony (Sb), Metalloid

Element 52: Tellurium (Te), Metalloid

Element 53: Iodine (I), Halogen

Element 54: Xenon (Xe), Noble gas
Element 55: Caesium (Cs), Alkali metal

Element 56: Barium (Ba), Alkaline earth metal

Element 57: Lanthanum (La), Lanthanoid

Element 58: Cerium (Ce), Lanthanoid

Element 59: Praseodymium (Pr), Lanthanoid

Element 60: Neodymium (Nd), Lanthanoid

Element 61: Promethium (Pm), Lanthanoid

Element 62: Samarium (Sm), Lanthanoid

Element 63: Europium (Eu), Lanthanoid

Element 64: Gadolinium (Gd), Lanthanoid

Element 65: Terbium (Tb), Lanthanoid

Element 66: Dysprosium (Dy), Lanthanoid

Element 67: Holmium (Ho), Lanthanoid

Element 68: Erbium (Er), Lanthanoid

Element 69: Thulium (Tm), Lanthanoid

Element 70: Ytterbium (Yb), Lanthanoid

Element 71: Lutetium (Lu), Lanthanoid

Element 72: Hafnium (Hf), Transition metal

Element 73: Tantalum (Ta), Transition metal

Element 74: Tungsten (W), Transition metal

Element 75: Rhenium (Re), Transition metal

Element 76: Osmium (Os), Transition metal

Element 77: Iridium (Ir), Transition metal

Element 78: Platinum (Pt), Transition metal

Element 79: Gold (Au), Transition metal

Element 80: Mercury (Hg), Transition metal

Element 81: Thallium (Tl), Other metal

Element 82: Lead (Pb), Other metal

Element 83: Bismuth (Bi), Other metal

Element 84: Polonium (Po), Metalloid

Element 85: Astatine (At), Halogen

Element 86: Radon (Rn), Noble gas
Element 87: Francium (Fr), Alkali metal

Element 88: Radium (Ra), Alkaline earth metal

Element 89: Actinium (Ac), Actinoid

Element 90: Thorium (Th), Actinoid

Element 91: Protactinium (Pa), Actinoid

Element 92: Uranium (U), Actinoid

Element 93: Neptunium (Np), Actinoid

Element 94: Plutonium (Pu), Actinoid

Element 95: Americium (Am), Actinoid

Element 96: Curium (Cm), Actinoid

Element 97: Berkelium (Bk), Actinoid

Element 98: Californium (Cf), Actinoid

Element 99: Einsteinium (Es), Actinoid

Element 100: Fermium (Fm), Actinoid

Element 101: Mendelevium (Md), Actinoid

Element 102: Nobelium (No), Actinoid

Element 103: Lawrencium (Lr), Actinoid

Element 104: Rutherfordium (Rf), Transition metal

Element 105: Dubnium (Db), Transition metal

Element 106: Seaborgium (Sg), Transition metal

Element 107: Bohrium (Bh), Transition metal

Element 108: Hassium (Hs), Transition metal

Element 109: Meitnerium (Mt)

Element 110: Darmstadtium (Ds)

Element 111: Roentgenium (Rg)

Element 112: Copernicium (Cn), Transition metal

Element 113: Ununtrium (Uut)

Element 114: Ununquadium (Uuq)

Element 115: Ununpentium (Uup)

Element 116: Ununhexium (Uuh)

Element 117: Ununseptium (Uus)

Element 118: Ununoctium (Uuo)
Lithium has a body-centered cubic crystal structure
3Li
Periodic table
Appearance
silver-white (seen here floating in oil)
General properties
Name, symbol, number lithium, Li, 3
Pronunciation /ˈlɪθiəm/ LI-thee-əm
Element category alkali metal
Group, period, block 1, 2, s
Standard atomic weight 6.941
Electron configuration 1s2 2s1 or [He]2s1
Electrons per shell 2, 1 (Image)
Physical properties
Phase solid
Density (near r.t.) 0.534 g·cm−3
Liquid density at m.p. 0.512 g·cm−3
Melting point 453.69 K, 180.54 °C, 356.97 °F
Boiling point 1615 K, 1342 °C, 2448 °F
Critical point (extrapolated)
3223 K, 67 MPa
Heat of fusion 3.00 kJ·mol−1
Heat of vaporization 147.1 kJ·mol−1
Specific heat capacity (25 °C) 24.860 J·mol−1·K−1
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 797 885 995 1144 1337 1610
Atomic properties
Oxidation states +1, -1
(strongly basic oxide)
Electronegativity 0.98 (Pauling scale)
Ionization energies 1st: 520.2 kJ·mol−1
2nd: 7298.1 kJ·mol−1
3rd: 11815.0 kJ·mol−1
Atomic radius 152 pm
Covalent radius 128±7 pm
Van der Waals radius 182 pm
Miscellanea
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 92.8 nΩ·m
Thermal conductivity (300 K) 84.8 W·m−1·K−1
Thermal expansion (25 °C) 46 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 6000 m/s
Young's modulus 4.9 GPa
Shear modulus 4.2 GPa
Bulk modulus 11 GPa
Mohs hardness 0.6
CAS registry number 7439-93-2
Most stable isotopes
Main article: Isotopes of lithium
iso NA half-life DM DE (MeV) DP
6Li 7.5% 6Li is stable with 3 neutrons
7Li 92.5% 7Li is stable with 4 neutrons
6Li content may be as low as 3.75% in
natural samples. 7Li would therefore
have a content of up to 96.25%.
v · d · e · r

Lithium (play /ˈlɪθiəm/, LI-thee-əm) is a soft, silver-white metal that belongs to the alkali metal group of chemical elements. It is represented by the symbol Li, and it has the atomic number 3. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable. For this reason, it is typically stored in mineral oil. When cut open, lithium exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a dull silvery gray, then black, tarnish. Because of its high reactivity, lithium never occurs free in nature, and instead, only appears in compounds, which are usually ionic. Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from brines and clays. On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

The nuclei of lithium are not far from being unstable, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. As a result, they can be used in fission reactions as well as fusion reactions of nuclear devices. Due to its near instability, lithium is less common in the solar system than 25 of the first 32 chemical elements even though the nuclei are very light in atomic weight.[1] For related reasons, lithium has important links to nuclear physics. The transmutation of lithium atoms to tritium was the first man-made form of a nuclear fusion reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.

Trace amounts of lithium are present in the oceans and in all organisms. The element serves no apparent vital biological function, since animal and plants survive in good health without it. Nonvital functions have not been ruled out. The lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood-stabilizing drug due to neurological effects of the ion in the human body. Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, lithium batteries and lithium-ion batteries. These uses consume more than half of lithium production.
Contents
[hide]

* 1 Characteristics
o 1.1 Atomic and physical
o 1.2 Chemistry and compounds
o 1.3 Isotopes
* 2 Occurrence
o 2.1 Astronomical
o 2.2 Terrestrial
o 2.3 Biological
* 3 History
* 4 Production
* 5 Applications
o 5.1 Electrical and electronics
o 5.2 Ceramics
o 5.3 Chemical and industrial
o 5.4 Nuclear
o 5.5 Medicine
o 5.6 Weaponry
* 6 Precautions
o 6.1 Regulation
* 7 See also
* 8 Notes
* 9 References
* 10 External links

[edit] Characteristics
Main article: Alkali metal
[edit] Atomic and physical
alt1
alt2
Lithium pellets covered in white lithium hydroxide (left) and ingots with a thin layer of black oxide tarnish (right)

Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[2] Because of this, it is a good conductor of heat and electricity as well as a highly reactive element, though the least reactive of the even-more highly reactive alkali metals. Lithium's low reactivity compared to other alkali metals is thought to be due to the proximity of its valence electron to its nucleus (the remaining two electrons in lithium's 1s orbital and are much lower in energy, and therefore they do not participate in chemical bonds).[2]

Lithium metal is soft enough to be cut with a knife. When cut, it possesses a silvery-white color that quickly changes to gray due to oxidation.[2] While it has one of the lowest melting points among all metals (180 °C), it has the highest melting point of the alkali metals.[3]

It is the lightest metal in the periodic table. It has a very low density, of approximately 0.534 g/cm3, which gives sticks of the metal a similar heft to dowels of a medium density wood, such as pine. It floats on water but also reacts with it.[2] It is the least dense of all elements that are not gasses at room temperature. The next lightest element is over 60% more dense (potassium, at 0.862 g/cm3). Furthermore, aside from helium and hydrogen, it is the least dense element in a solid or liquid state, being only 2/3 as dense as liquid nitrogen (0.808 g/cm3).[note 1][4]

Lithium's coefficient of thermal expansion is twice that of aluminum and almost four times that of iron.[5] It has the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[6] and at higher temperatures (more than 9 K) at very high pressures (>20 GPa)[7] At temperatures below 70 K, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2 K it has a rhombohedral crystal system (with a nine-layer repeat spacing); at higher temperatures it transforms to face-centered cubic and then body-centered cubic. At liquid-helium temperatures (4 K) the rhombohedral structure is the most prevalent.[8] Multiple allotropic forms have been reported for lithium at high pressures.[9]
[edit] Chemistry and compounds

Lithium reacts with water easily, but with noticeably less energy than other alkali metals do. The reaction forms hydrogen gas and lithium hydroxide in aqueous solution.[2] Because of its reactivity with water, lithium is usually stored under cover of a viscous hydrocarbon, often petroleum jelly. Though the heavier alkali metals can be stored in less dense substances, such as mineral oil, lithium is not dense enough to be fully submerged in these liquids.[10] In moist air, lithium 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).[11]
Hexameric structure of the n-butyllithium fragment in a crystal

When placed over a flame, lithium compounds give off a striking crimson color, but when it burns strongly the flame becomes a brilliant silver. Lithium will ignite and burn in oxygen when exposed to water or water vapors.[12] Lithium is flammable, and it is potentially explosive when exposed to air and especially to water, though less so than the other alkali metals. The lithium-water reaction at normal temperatures is brisk but not violent, though the hydrogen produced can ignite. As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically Class D type (see Types of extinguishing agents). Lithium is the only metal which reacts with nitrogen under normal conditions.[13][14]

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 (Li2O) and peroxide (Li2O2) when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.[11][15] The metal reacts with hydrogen gas at high temperatures to produce lithium hydride (LiH).[16]

Other known binary compounds include the halides (LiF, LiCl, LiBr, LiI), and the sulfide (Li2S), the superoxide (LiO2), carbide (Li2C2). Many other inorganic compounds are known, where lithium combines with anions to form various salts: borates, amides, carbonate, nitrate, or borohydride (LiBH4). Multiple organolithium reagents are known where there is a direct bond between carbon and lithium atoms effectively creating a carbanion that are extremely powerful bases and nucleophiles. In many of these organolithium compounds, the lithium ions tend to aggregate into high-symmetry clusters by themselves, which is relatively common for alkali cations.[17]
[edit] Isotopes
Main article: Isotopes of lithium

Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (92.5% natural abundance).[2][10][18] Both natural isotopes have anomalously low nuclear binding energy per nucleon compared to the next lighter and heavier elements, helium and beryllium, which means that alone among stable light elements, lithium can produce net energy through nuclear fission. The two lithium nuclei have lower binding energies per nucleon than any other stable compound nuclides other than deuterium, and helium-3.[19] As a result of this, though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.[1] 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 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 7.6 × 10−23 s.[20]

7Li is one of the primordial elements (or, more properly, primordial nuclides) produced in Big Bang nucleosynthesis. A small amount of both 6Li and 7Li are produced in stars, but are thought to be burned as fast as produced.[21] Additional small amounts of lithium of both 6Li and 7Li may be generated from solar wind, cosmic rays hitting heavier atoms, and from early solar system 7Be and 10Be radioactive decay.[22] While lithium is created in stars during the Stellar nucleosynthesis, it is further burnt. 7Li can also be generated in carbon stars.[23]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[24] including mineral formation (chemical precipitation), metabolism, and ion exchange. Lithium ions substitute 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. The process known as laser isotope separation can be used to separate lithium isotopes.[25]
[edit] Occurrence
Lithium is about as common as chlorine in the Earth's upper continental crust, on a per-atom basis.
[edit] Astronomical
Main article: Nucleosynthesis

According to modern cosmological theory, lithium—as both of its stable isotopes lithium-6 and lithium-7—was among the 3 elements synthesized in the Big Bang. Though the amount of lithium generated in Big Bang nucleosynthesis is dependent upon the number of photons per baryon, for accepted values the lithium abundance can be calculated, and there is a "cosmological lithium discrepancy" in the Universe: older stars seem to have less lithium than they should, and some younger stars have far more. The lack of lithium in older stars is apparently caused by the "mixing" of lithium into the interior of stars, where it is destroyed.[26] Furthermore, lithium is produced in younger stars. Though it transmutes into two atoms of helium due to collision with a proton at temperatures above 2.4 million degrees Celsius (most stars easily attain this temperature in their interiors), lithium is more abundant than predicted in later-generation stars, for causes not yet completely understood.[10]

Though it was one of the three first elements (together with helium and hydrogen) to be synthesized in the Big Bang, lithium, together with beryllium and boron are markedly less abundant than other nearby elements. This is a result to the low temperature necessary to destroy lithium, and a lack of common processes to produce it.[27]

Lithium is also found in brown dwarf stars and certain anomalous orange stars. Because lithium is present in cooler, less-massive brown dwarf stars, but is destroyed in hotter red dwarf stars, its presence in the stars' spectra can be used in the "lithium test" to differentiate the two, as both are smaller than the Sun.[10][28][29] Certain orange stars can also contain a high concentration of lithium. Those orange stars found to have a higher than usual concentration of lithium (such as Centaurus X-4) orbit massive objects—neutron stars or black holes—whose gravity evidently pulls heavier lithium to the surface of a hydrogen-helium star, causing more lithium to be observed.[10]
[edit] Terrestrial
Lithium mine production (2009) and reserves in tonnes[30] Country↓ Production↓ Reserves↓
Argentina 2,200 800,000
Australia 4,400 580,000
Brazil 110 190,000
Canada 480 180,000
Chile 7,400 7,500,000
People's Republic of China 2,300 540,000
Portugal 490 Not available
United States Withheld 38,000
Zimbabwe 350 23,000
World total 18,000 9,900,000
See also: Lithium minerals

Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity.[2] The total lithium content of seawater is very large and is estimated as 230 billion tonnes, where the element exists at a relatively constant concentration of 0.14 to 0.25 parts per million (ppm),[31][32] or 25 micromolar;[33] higher concentrations approaching 7 ppm are found near hydrothermal vents.[32]

Estimates for crustal content range from 20 to 70 ppm by weight.[11] 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 sources.[11] A newer source for lithium is hectorite clay, the only active development of which is through the Western Lithium Corporation in the United States.[34] At 20 mg lithium per kg of Earth's crust,[35] lithium is the 25th most abundant element. Nickel and lead have about the same abundance.

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."[36]

The largest reserve base of lithium is in the Salar de Uyuni area of Bolivia, which has 5.4 million tonnes. US Geological Survey, estimates that in 2009 Chile had the largest reserves by far (7.5 million tonnes) and the highest annual production (7,400 tonnes). Other major suppliers include Australia, Argentina and China.[30][37] Other estimates put Argentina's reserve base (7.52 million tonnes) above that of Chile (6 million).[38]

In June 2010, the New York Times reported that American geologists were conducting ground surveys on dry salt lakes in western Afghanistan believing that large deposits of lithium are located there. "Pentagon officials said that their initial analysis at one location in Ghazni Province showed the potential for lithium deposits as large of those of Bolivia, which now has the world’s largest known lithium reserves." [39] These estimates are "based principally on old data, which was gathered mainly by the Soviets during their occupation of Afghanistan from 1979–1989" and "Stephen Peters, the head of the USGS’s Afghanistan Minerals Project, said that he was unaware of USGS involvement in any new surveying for minerals in Afghanistan in the past two years. 'We are not aware of any discoveries of lithium,' he said."[40]
[edit] Biological

Lithium is found in trace amount in numerous plants, plankton, and invertebrates, at concentrations of 69 to 5,760 parts per billion (ppb). In vertebrates the concentration is slightly lower, and nearly all vertebrate tissue and body fluids have been found to contain lithium ranging from 21 to 763 ppb.[32] Marine organisms tend to bioaccumulate lithium more than terrestrial ones.[41] It is not known whether lithium has a physiological role in any of these organisms,[32] but nutritional studies in mammals have indicated its importance to health, leading to a suggestion that it be classed as an essential trace element with an RDA of 1 mg/day. [1] Observational studies in Japan, reported in 2011, suggested that naturally occurring lithium in drinking water may increase human lifespan. [2]
[edit] History
Johan August Arfwedson is credited with the discovery of lithium in 1817

Petalite (LiAlSi4O10) was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada e Silva in a mine on the island of Utö, Sweden.[42][43][44] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jakob Berzelius, detected the presence of a new element while analyzing petalite ore.[45][46][47] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[48] Berzelius gave the alkaline material the name "lithion/lithina", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium which was known partly for its high abundance in animal blood. He named the metal inside the material as "lithium".[2][43][47]

Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[43] In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame.[43] However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.[43][47][49] It was not isolated until 1821, when William Thomas Brande obtained it by electrolysis of lithium oxide, a process that had previously been employed by the chemist Sir Humphry Davy to isolate the alkali metals potassium and sodium.[10][49][50][51] Brande also described some pure salts of lithium, such as the chloride, and, estimating that lithia (lithium oxide) contained about 55% metal, estimated the atomic weight of lithium to be around 9.8 g/mol (modern value ~6.94 g/mol).[52] In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[43] The discovery of this procedure henceforth led to commercial production of lithium, beginning in 1923, by the German company Metallgesellschaft AG, which performed an electrolysis of a liquid mixture of lithium chloride and potassium chloride.[43][53]

The production and use of lithium underwent several drastic changes in history. The first major application of lithium became high temperature grease for aircraft engines or similar applications in World War II and shortly after. This small market was supported by several small mining operations mostly in the United States. The demand for lithium increased dramatically during the Cold War with the production of nuclear fusion weapons. Both lithium-6 and lithium-7 produce tritium when irradiated by neutrons, and are thus useful for the production of tritium by itself, as well as a form of solid fusion fuel used inside hydrogen bombs in the form of lithium deuteride. The United States became the prime producer of lithium in the period between the late 1950s and the mid 1980s. At the end the stockpile of lithium was roughly 42,000 tonnes of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75%.[54]

Lithium was used to decrease the melting temperature of glass and to improve the melting behavior of aluminium oxide when using the Hall-Héroult process.[55][55] These two uses dominated the market until the middle of the 1990s. After the end of the nuclear arms race the demand for lithium decreased and the sale of Department of Energy stockpiles on the open market further reduced prices.[54] But in the mid-1990s, several companies started to extract lithium from brine which proved to be a less expensive method than underground or even open pit mining. Most of the mines closed or shifted their focus to other materials as only the ore from zoned pegmatites could be mined for a competitive price. For example, the US mines near Kings Mountain, North Carolina closed before the turn of the century. The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.[56] With the surge of lithium demand in batteries in to 2000s, new companies have expanded brine extraction efforts to meet the rising demand.[57][58]
[edit] Production
alt1
alt2
Satellite images of the Salar del Hombre Muerto, Argentina (left), and Uyuni, Bolivia (right), salt flats are rich in lithium. The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in the left image).

Since the end of World War II lithium production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride. In 1998 it was about 95 US$ / kg (or 43 US$/pound).[59]

There are widespread 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'".[60]

Deposits of lithium are found in South America throughout the Andes mountain chain. Chile is the leading lithium producer, followed by Argentina. Both countries recover the lithium from brine pools. In the United States lithium is recovered from brine pools in Nevada.[61] Nearly half the world's known reserves are located in Bolivia, a nation sitting along the central eastern slope of the Andes. In 2009 Bolivia is negotiating with Japanese, French, and Korean firms to begin extraction.[62] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tonnes of lithium.[62][63] China may emerge as a significant producer of brine-source lithium carbonate around 2010. There is potential production of up to 55,000 tonnes per year if projects in Qinghai province and Tibet proceed.[60]

Worldwide reserves of lithium are estimated to be 23 million tonnes.[64] Using the battery efficiency figure of 400 g of lithium per kWh,[65] this gives a total maximum lithium battery capacity of 52 billion kWh which, assuming it is used exclusively for car batteries, is enough for approximately 2 billion cars with a 24 kWh battery (like a Nissan Leaf [66]).
[edit] Applications
Usage of lithium in the USA in 2010[67]
Ceramics and glass (31%)
Batteries (23%)
Lubricating greases (9%)
Other (15%)
More[show]
Air treatment (6%)
Primary aluminum production (6%)
Continuous casting (4%)
Rubber and thermoplastics (4%)
Pharmaceuticals (2%)
[edit] Electrical and electronics

In the later years of the 20th century lithium became important as an anode material. Used in lithium-ion batteries because of its high electrochemical potential, a typical cell can generate approximately 3 volts, compared with 2.1 volts for lead/acid or 1.5 volts for zinc-carbon cells. Because of its low atomic mass, it also has a high charge- and power-to-weight ratio. Lithium batteries are disposable (primary) batteries with lithium or its compounds as an anode. Lithium batteries are not to be confused with lithium-ion batteries, which are high energy-density rechargeable batteries. Other rechargeable batteries include the lithium-ion polymer battery, lithium iron phosphate battery, and the nanowire battery. New technologies are constantly being announced.

Lithium niobate is used extensively in telecommunication products such as mobile phones and optical modulators, for such components as resonant crystals. Lithium applications are used in more than 60% of mobile phones.[68] Because of its specific heat capacity, the highest of all solids, lithium is often used in coolants for heat transfer applications.[61]
[edit] Ceramics

Lithium oxide is a widely used flux for processing silica, reducing the melting point and viscosity of the material and leading to glazes of improved physical properties including low coefficients for thermal expansion. Lithium oxides are a component of ovenware.
[edit] Chemical and industrial
Lithium use in flares and pyrotechnics is due to its red flame

Lithium is also used in the pharmaceutical and fine-chemical industry in the manufacture of organolithium reagents, which are used both as strong bases and as reagents for the formation of carbon-carbon bonds. Organolithium compounds are also used in polymer synthesis as catalysts/initiators[69] in anionic polymerization of unfunctionalized olefins.[70][71][72] Lithium is used in the preparation of organolithium compounds, which are in turn very reactive and are the basis of many synthetic applications.[73]

Lithium chloride and lithium bromide are extremely hygroscopic and are used as desiccants.[61] 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 soap made of lithium stearate. Lithium soap has the ability to thicken oils, and it is used to manufacture all-purpose, high-temperature lubricating greases.[61][74][75]

When used as a flux for welding or soldering, lithium promotes the fusing of metals during and eliminates the forming of oxides by absorbing impurities. Its fusing quality is also important as a flux for producing ceramics, enamels and glass. Alloys of the metal with aluminium, cadmium, copper and manganese are used to make high-performance aircraft parts (see also Lithium-aluminium alloys). Lithium compounds are also used as pyrotechnic colorants and oxidizers in red fireworks and flares.[61][76]
[edit] Nuclear

Lithium-6 is valued as a source material for tritium production and as a neutron absorber in nuclear fusion. Natural lithium contains about 7.5% lithium-6 from which large amounts of lithium-6 have been produced by isotope separation for use in nuclear weapons.[77] Lithium-7 gained interest for use in nuclear reactor coolants.[78]
Lithium deuteride was used as fuel in the Castle Bravo nuclear device.

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.[79]

Lithium fluoride as highly enriched in the lithium-7 isotope forms the basic constituent of the fluoride salt mixture LiF-BeF2 that used in liquid-fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF-BeF2 mixtures have low melting points. In addition, 7Li, Be, and F are among the few nuclides with low enough thermal neutron capture cross-sections to not poison the fission reactions inside a nuclear fission reactor.[note 2][80]

In conceptualized nuclear fusion power plants, lithium will be used to produce tritium in magnetically confined 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 + 3T.

Lithium is also 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.[81]
[edit] Medicine
Main article: Lithium (medication)

Lithium compounds continue to be the gold standard for the treatment of bipolar disorder and is helpful for related diagnoses, such as schizoaffective disorder and cyclic major depression. The active principle in these salts is the lithium ion Li+, although detailed mechanisms are debated.

Lithium fluoride, artificially grown as crystal, is clear and transparent and often used in specialist optics for IR, UV and VUV (vacuum UV) applications. It has one of the lowest refractive indexes and the farthest transmission range in the deep UV of most common materials.[82] Finely divided lithium fluoride powder has been used for thermoluminescent radiation dosimetry (TLD): when a sample of such is exposed to radiation, it accumulates crystal defects which, when heated, resolve via a release of bluish light whose intensity is proportional to the absorbed dose, thus allowing this to be quantified.[83] Lithium fluoride is sometimes used in focal lenses of telescopes.[61][84] The high non-linearity of lithium niobate also makes it useful in non-linear optics applications. Lithium hydroxide and lithium peroxide are used in confined areas, such as aboard spacecraft and submarines, for air purification. Lithium hydroxide absorbs carbon dioxide from the air by reacting with it to form lithium carbonate, and is preferred over other alkaline hydroxides for its low weight. Lithium peroxide (Li2O2) in presence of moisture not only absorbs carbon dioxide to form lithium carbonate, but also releases oxygen.[85][86] For example:

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.

[edit] Weaponry
The launch of a torpedo using lithium as fuel

Metallic lithium and its complex hydrides, such a Li[AlH4], are used as high energy additives to rocket propellants.[10] Lithium peroxide, lithium nitrate, lithium chlorate and lithium perchlorate are used as oxidizers in rocket propellants, and also in oxygen candles that supply submarines and space capsules with oxygen.[87] The Mark 50 Torpedo stored chemical energy propulsion system (SCEPS) uses a small tank of sulfur hexafluoride gas which is sprayed over a block of solid lithium. The reaction generates enormous heat which is used to generate steam from seawater. The steam propels the torpedo in a closed Rankine cycle.[88]
[edit] Precautions
NFPA 704
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The fire diamond hazard sign for lithium metal

Lithium is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) initially irritate the nose and throat, while higher exposure can cause a buildup of fluid in the lungs, leading to pulmonary edema. The metal itself is a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium is safely stored in non-reactive compounds such as naphtha.[89]

There have been suggestions of increased risk of developing Ebstein's cardiac anomaly in infants born to women taking lithium during the first trimester of pregnancy.[90]
[edit] Regulation

Some jurisdictions limit the sale of lithium batteries, which are the most readily available source of lithium 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.[91][92] 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 in a process called thermal runaway. Most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents. Internal shorts have been known to develop due to manufacturing defects or damage to batteries that can lead to spontaneous thermal runaway.[93][94]
[edit] See also
Book: Lithium
Wikipedia Books are collections of articles that can be downloaded or ordered in print.

* Dilithium
* Lithium compounds
* Lithium-based grease
* Lithium-ion battery

[edit] Notes

1. ^ Densities for all the gaseous elements can be obtained at Airliquide.com
2. ^ Beryllium and fluorine occur only as one isotope, 9Be and 19F respectively. These two, together with 7Li, as well as 2H, 11B, 15N, 209Bi, and the stable isotopes of C, and O, are the only nuclides with low enough thermal neutron capture cross sections aside from actinides to serve as major constituents of an molten salt breeder reactor fuel.

Posted in: Alkali metals