This article is about the chemical element. For other uses, see Lithium (disambiguation). helium ← lithium → beryllium H ↑ Li ↓ Na 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 6.941g·mol−1 Electron configuration 1s2 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 Lithium ( /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, usually ionic ones. 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 1 Characteristics 1.1 Atomic and physical 1.2 Chemistry and compounds 1.3 Isotopes 2 Occurrence 2.1 Astronomical 2.2 Terrestrial 3 History 4 Production 5 Applications 5.1 Electrical and electronics 5.2 Medicinal 5.3 Chemical and industrial 5.4 Nuclear 5.5 Other uses 6 Precautions 6.1 Regulation 7 See also 8 Notes 9 References 10 External links // Characteristics Main article: Alkali metal Atomic and physical 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 14 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 pressure6 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 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 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.1314 Hexameric structure of the LiBu fragment in a crystal 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.1115 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 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).21018 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.20 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.21 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.22 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.23 While lithium is created in stars during the Stellar nucleosynthesis, it is further burnt. 7Li can also be generated in carbon stars.24 Lithium isotopes fractionate substantially during a wide variety of natural processes,25 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.26 Occurrence Lithium is about as common as chlorine in the Earth's upper continental crust, on a per-atom basis. 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.27 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.28 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.102930 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 Terrestrial Lithium mine production (2009) and reserves in tonnes31 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),3233 or 25 micromolar;34 higher concentrations approaching 7 ppm are found near hydrothermal vents.33 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.35 At 20 mg lithium per kg of Earth's crust,36 lithium is the 25th most abundant element. Nickel and lead have about the same abundance. 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.33 Marine organisms tend to bioaccumulate lithium more than terrestrial ones.37 It is not known whether lithium has a physiological role in any of these organisms.33 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."38 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.3139 Other estimates put Argentina's reserve base (7.52 million tonnes) above that of Chile (6 million).40 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." 41 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."42 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.434445 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.464748 This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.49 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".24448

Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.44 In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame.44 However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.444850 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.10505152 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).53 In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.44 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.4454 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%.55 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.5656 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.55 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 2007.57 With the surge of lithium demand in batteries in to 2000s, new companies have expanded brine extraction efforts to meet the rising demand.5859 Production 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).60 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'".61 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.62 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.63 According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tonnes of lithium, which can be used to make batteries for hybrid and electric vehicles.6364 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.61 Worldwide reserves of lithium are estimated to be 23 million tonnes.65 Using the battery efficiency figure of 400 g of lithium per kWh,66 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 67). Applications Usage of lithium in the USA in 200968 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.69 Because of its specific heat capacity, the highest of all solids, lithium is often used in coolants for heat transfer applications.62 Medicinal Main article: Lithium pharmacology 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 depression and mania (though more effective for the latter). Lithium continues to be the gold standard for the treatment of bipolar disorder. It is also helpful for related diagnoses, such as schizoaffective disorder and cyclic major depression. In addition to watching out for the well-known complications of lithium treatment—hypothyroidism and decreased renal function—health care providers should be aware of hyperparathyroidism.70 Lithium can also be used to augment antidepressants. Because of Lithium's nephrogenic diabetes insipidus effects, it can be used to help treat the syndrome of inappropriate antidiuretic hormone hypersecretion (SIADH). It was also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.71 The active principle in these salts is the lithium ion Li+. Although this ion has a smaller diameter than either Na+ or K+, in a watery environment like the cytoplasmic fluid, Li+ binds to the oxygen atoms of water, making it effectively larger than either Na+ or K+ ions. How Li+ works in the central nervous system is still a matter of debate. Li+ elevates brain levels of tryptophan, 5-HT (serotonin), and 5-HIAA (a serotonin metabolite). Serotonin is related to mood stability. Li+ also reduces catecholamine activity in the brain (associated with brain activation and mania), by enhancing reuptake and reducing release. Therapeutically useful amounts of lithium (1.0 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.72 Common side effects of lithium treatment include muscle tremors, twitching, ataxia, and hypothyroidism.73 Long term use is linked to hyperparathyroidism,74 hypercalcemia (bone loss), hypertension, damage of tubuli in the kidney, nephrogenic diabetes insipidus (polyuria and polydipsia) and/or glomerular damage – even to the point of uremia,75 seizures76 and weight gain.77 According to a study in 2009 at Oita University in Japan and published in the British Journal of Psychiatry, communities whose water contained larger amounts of lithium had significantly lower suicide rates,78798081 but did not address whether lithium in drinking water causes the negative side effects associated with higher doses of the element.82 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/initiators83 in anionic polymerization of unfunctionalized olefins.848586 Lithium is used in the preparation of organolithium compounds, which are in turn very reactive and are the basis of many synthetic applications.87 Lithium chloride and lithium bromide are extremely hygroscopic and are used as desiccants.62 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.628889 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.6290 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.91 Lithium-7 gained interest for use in nuclear reactor coolants.92 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.93 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 294 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.

Various means of doing this will be tested at the ITER reactor being built at Cadarache, France.95 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 1929.96 Other uses 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.97 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.98 Lithium fluoride is sometimes used in focal lenses of telescopes.6299 The high non-linearity of lithium niobate also makes it useful in non-linear optics applications. 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.100 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.101 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.102103 For example: 2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2. Precautions NFPA 704 0 3 2 W 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.104 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.105 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.106107 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.108109 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 Notes ^ Densities for all the gaseous elements can be obtained at Airliquide.com ^ 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. References ^ Numerical data from: Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591: 1220–1247. doi:10.1086/375492.  Graphed at File:SolarSystemAbundances.jpg ^ a b c d e f g h Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements: A Reference Guide. Westport, Conn.: Greenwood Press. ISBN 0-313-33438-2.  ^ Lide, D. R., ed. (2005), CRC Handbook of Chemistry and Physics (86th ed.), Boca Raton (FL): CRC Press, ISBN 0-8493-0486-5  ^ "Nitrogen, N2, Physical properties, safety, MSDS, enthalpy, material compatibility, gas liquid equilibrium, density, viscosity, flammability, transport properties". Encyclopedia.airliquide.com. http://encyclopedia.airliquide.com/Encyclopedia.asp?LanguageID=11&CountryID=19&Formula=&GasID=5&UNNumber=&EquivGasID=32&VolLiquideBox=&MasseLiquideBox=&VolGasBox=&MasseGasBox=&RD20=29&RD9=8&RD6=64&RD4=2&RD3=22&RD8=27&RD2=20&RD18=41&RD7=18&RD13=71&RD16=35&RD12=31&RD19=34&RD24=62&RD25=77&RD26=78&RD28=81&RD29=82. Retrieved 2010-09-29.  ^ "Coefficients of Linear Expansion". Engineering Toolbox. http://www.engineeringtoolbox.com/linear-expansion-coefficients-d_95.html.  ^ Tuoriniemi, J; Juntunen-Nurmilaukas, K; Uusvuori, J; Pentti, E; Salmela, A; Sebedash, A (2007). "Superconductivity in lithium below 0.4 millikelvin at ambient pressure.". Nature 447 (7141): 187–9. doi:10.1038/nature05820. PMID 17495921.  ^ Struzhkin, V. V.; Eremets, M. I.; Gan, W; Mao, H. K.; Hemley, R. J. (2002). "Superconductivity in dense lithium". Science 298 (5596): 1213–5. doi:10.1126/science.1078535. PMID 12386338.  ^ Overhauser, A. W. (1984). "Crystal Structure of Lithium at 4.2 K". Physical Review Letters 53: 64–65. doi:10.1103/PhysRevLett.53.64.  ^ Schwarz, Ulrich (2004). "Metallic high-pressure modifications of main group elements". Zeitschrift für Kristallographie 219: 376. doi:10.1524/zkri.219.6.376.34637.  ^ a b c d e f g Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. ISBN 0198503415.  ^ a b c d "Lithium and lithium compounds". Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons, Inc.. 2004. doi:10.1002/0471238961.1209200811011309.a01.pub2.  ^ Kirchoff, Gustav; Bunsen, Robert. "Chemical Analysis By Observation of Spectra". University of Pittsburgh. http://www.pitt.edu/~alw11/InterestInfo/Articles/Bunsen%20and%20Kirchoff.pdf. Retrieved 2009-11-05.  ^ Krebs, Robert E. (2006). The history and use of our earth's chemical elements: a reference guide. Greenwood Publishing Group. p. 47. ISBN 0313334382. http://books.google.com/books?id=yb9xTj72vNAC&pg=PA47.  ^ Institute, American Geological; Union, American Geophysical; Society, Geochemical (1994-01-01). Geochemistry international. 31. p. 115. http://books.google.com/?id=77McAQAAIAAJ&q=Lithium+is+%22the+only+metal%22+which+reacts+with+nitrogen&dq=Lithium+is+%22the+only+metal%22+which+reacts+with+nitrogen.  ^ Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, pp. 97–99, ISBN 0-08-022057-6  ^ Beckford, Floyd. "University of Lyon course online (powerpoint) slideshow". Archived from the original on November 4, 2005. http://web.archive.org/web/20051104025202/http://www.lyon.edu/webdata/users/fbeckford/CHM+120/Lecture+Notes/Chapter-14.ppt. Retrieved 2008-07-27. "definitions:Slides 8–10 (Chapter 14)"  ^ Sapse, Anne-Marie and von R. Schleyer, Paul (1995). Lithium chemistry: a theoretical and experimental overview. Wiley-IEEE. pp. 3–40. ISBN 0471549304. http://books.google.com/books?id=z76sVepirh4C&pg=PA16.  ^ "Isotopes of Lithium". Berkeley National Laboratory, The Isotopes Project. http://ie.lbl.gov/education/parent/Li_iso.htm. Retrieved 2008-04-21.  ^ File:Binding energy curve - common isotopes.svg shows binding energies of stable nuclides graphically; the source of the data-set is given in the figure background. ^ Numerical data from: Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591: 1220–1247. doi:10.1086/375492.  Graphed atFile:SolarSystemAbundances.jpg ^ Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. http://www.nndc.bnl.gov/chart/reCenter.jsp?z=104&n=158. Retrieved 2008-06-06.  ^ Asplund, M. et al. (2006). "Lithium Isotopic Abundances in Metal-poor Halo Stars". The Astrophysical Journal 644: 229. doi:10.1086/503538.  ^ Chaussidon, M.; Robert, F.; McKeegan, K.D. (2006). "Li and B isotopic variations in an Allende CAI: Evidence for the in situ decay of short-lived 10Be and for the possible presence of the short−lived nuclide 7Be in the early solar system" (free download pdf). Geochimica et Cosmochimica Acta 70 (1): 224–245. doi:10.1016/j.gca.2005.08.016. http://sims.ess.ucla.edu/PDF/Chaussidon_et_al_Geochim%20Cosmochim_2006a.pdf.  ^ Denissenkov, P. A.; Weiss, A. (2000). "Episodic lithium production by extra-mixing in red giants". Astronomy and Astrophysics 358: L49–L52. Bibcode: 2000A&A...358L..49D.  ^ Seitz, H.M.; Brey, G.P.; Lahaye, Y.; Durali, S.; Weyer, S. (2004). "Lithium isotopic signatures of peridotite xenoliths and isotopic fractionation at high temperature between olivine and pyroxenes". Chemical Geology 212 (1–2): 163–177. doi:10.1016/j.chemgeo.2004.08.009.  ^ Duarte, F. J (2009). Tunable Laser Applications. CRC Press. p. 330. ISBN 1420060090. http://www.opticsjournal.com/tla.htm.  ^ Fraser Cain (16th Aug 2006). "Why Old Stars Seem to Lack Lithium". http://www.universetoday.com/2006/08/16/why-old-stars-seem-to-lack-lithium/.  ^ "Element Abundances". Archived from the original on 2006-09-01. http://web.archive.org/web/20060901133923/http://www.astro.wesleyan.edu/~bill/courses/astr231/wes_only/element_abundances.pdf. Retrieved 2009-11-17.  ^ Cain, Fraser. "Brown Dwarf". Universe Today. http://www.universetoday.com/24593/brown-dwarf/. Retrieved 2009-11-17.  ^ "L Dwarf Classification". http://www-int.stsci.edu/~inr/ldwarf3.html. Retrieved 2009-11-17.  ^ a b U.S. Geological Survey, 2009, commodity summaries 2009: U.S. Geological Survey ^ "Lithium Occurrence". Institute of Ocean Energy, Saga University, Japan. http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html. Retrieved 2009-03-13.  ^ a b c d "Some Facts about Lithium". ENC Labs. http://www.enclabs.com/lithium.html. Retrieved 2010-10-15.  ^ "Extraction of metals from sea water". Springer Berlin Heidelberg. 1984. http://www.springerlink.com/content/y621101m3567jku1/.  ^ Moores, S. (June 2007). "Between a rock and a salt lake". Industrial Minerals 477: 58.  ^ Taylor, S. R.; McLennan, S. M.; The continental crust: Its composition and evolution, Blackwell Sci. Publ., Oxford, 330 pp. (1985). Cited in Abundances of the elements (data page) ^ Chassard-Bouchaud, C; Galle, P; Escaig, F; Miyawaki, M (1984). "Bioaccumulation of lithium by marine organisms in European, American, and Asian coastal zones: microanalytic study using secondary ion emission". Comptes rendus de l'Academie des sciences. Serie III, Sciences de la vie 299 (18): 719–24. PMID 6440674.  ^ Handbook of Lithium and Natural Calcium, Donald Garrett, Academic Press, 2004, cited in The Trouble with Lithium 2 ^ "Front Matter" (PDF). http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf. Retrieved 2010-09-29.  ^ Clarke, G.M. and Harben, P.W., "Lithium Availability Wall Map". Published June 2009. Referenced at International Lithium Alliance ^ Risen, James (2010-06-13). "U.S. Identifies Vast Riches of Minerals in Afghanistan". The New York Times. http://www.nytimes.com/2010/06/14/world/asia/14minerals.html?pagewanted=1&hp. Retrieved June 13, 2010.  ^ Page, Jeremy; Evans, Michael (2010-06-15). "Taleban zones mineral riches may rival Saudi Arabia says Pentagon". The Times (London). http://business.timesonline.co.uk/tol/business/industry_sectors/natural_resources/article7149696.ece.  ^ "Petalite Mineral Information". http://www.mindat.org/min-3171.html. Retrieved 10 August 2009.  ^ a b c d e f g "Lithium:Historical information". http://www.webelements.com/lithium/history.html. Retrieved 10 August 2009.  ^ Weeks, Mary (2003). Discovery of the Elements. Whitefish, Montana, United States: Kessinger Publishing. p. 124. ISBN 0766138720. http://books.google.com/?id=SJIk9BPdNWcC. Retrieved 10 August 2009.  ^ "Johan August Arfwedson". Periodic Table Live!. http://www.chemeddl.org/collections/ptl/ptl/chemists/bios/arfwedson.html. Retrieved 10 August 2009.  ^ "Johan Arfwedson". Archived from the original on 2008-06-05. http://web.archive.org/web/20080605152857/http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm. Retrieved 10 August 2009.  ^ a b c van der Krogt, Peter. "Lithium". Elementymology & Elements Multidict. http://elements.vanderkrogt.net/element.php?sym=Li. Retrieved 2010-10-05.  ^ Clark, Jim (2005). "Compounds of the Group 1 Elements". http://www.chemguide.co.uk/inorganic/group1/compounds.html. Retrieved 10 August 2009.  ^ a b Per Enghag (2004). Encyclopedia of the Elements: Technical Data – History – Processing – Applications. Wiley. pp. 287–300. ISBN 978-3527306664.  ^ "The Quarterly journal of science and the arts" (PDF). The Quarterly Journal of Science and the Arts (Royal Institution of Great Britain) 5: 338. 1818. http://books.google.com/?id=D_4WAAAAYAAJ. Retrieved 2010-10-05.  ^ "Timeline science and engineering". DiracDelta Science & Engineering Encyclopedia. http://www.diracdelta.co.uk/science/source/t/i/timeline/source.html. Retrieved 2008-09-18.  ^ Brande, William Thomas; MacNeven, William James (1821). A manual of chemistry. p. 191. http://books.google.com/?id=zkIAAAAAYAAJ. Retrieved 2010-10-08.  ^ Green, Thomas (2006-06-11). "Analysis of the Element Lithium". echeat. http://www.echeat.com/essay.php?t=29195.  ^ a b Ober, Joyce A. (1994). "Commodity Report 1994: Lithium". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450494.pdf. Retrieved 2010-11-03.  ^ a b Deberitz, JüRgen; Boche, Gernot (2003). "Lithium und seine Verbindungen – Industrielle, medizinische und wissenschaftliche Bedeutung". Chemie in unserer Zeit 37: 258. doi:10.1002/ciuz.200300264.  ^ Ober, Joyce A. (1994). "Minerals Yearbook 2007 : Lithium". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/lithium/myb1-2007-lithi.pdf. Retrieved 2010-11-03.  ^ Kogel, Jessica Elzea (2006). "Lithium". Industrial minerals & rocks: commodities, markets, and uses. Littleton, Colo.: Society for Mining, Metallurgy, and Exploration. p. 599. ISBN 9780873352338. http://books.google.com/?id=zNicdkuulE4C&pg=PA600&lpg=PAPA599. ) ^ McKetta, John J. (2007-07-18). Encyclopedia of Chemical Processing and Design: Volume 28 – Lactic Acid to Magnesium Supply-Demand Relationships. M. Dekker. ISBN 9780824724788. http://books.google.com/books?id=8erDL_DnsgAC&pg=PA339. Retrieved 2010-09-29.  ^ Ober, Joyce A. "Lithium" (PDF). United States Geological Survey. pp. 77–78. http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450798.pdf. Retrieved 2007-08-19.  ^ a b "The Trouble With Lithium 2" (PDF). Meridian International Research. May 28, 2008. http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf. Retrieved 2008-07-07.  ^ a b c d e f Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN 0849304814.  ^ a b Simon Romero (February 2, 2009). "In Bolivia, a Tight Grip on the Next Big Resource". New York Times. http://www.nytimes.com/2009/02/03/world/americas/03lithium.html?ref=world.  ^ "USGS Mineral Commodities Summaries 2009". USGS. http://minerals.usgs.gov/minerals/pubs/mcs/2009/mcs2009.pdf.  ^ Lithium, USGS Mineral Commodity Summaries, 2010 ^ "How Much Lithium Per Battery" (PDF). http://www.meridian-int-res.com/Projects/How_Much_Lithium_Per_Battery.pdf.  ^ "NISSAN : 100% Electric Zero-emission Nissan LEAF Debuts in Japan". http://www.nissan-global.com/EN/NEWS/2010/_STORY/101203-01-e.html.  ^ USGS (2010). "Lithium" (PDF). http://minerals.usgs.gov/minerals/pubs/commodity/lithium/mcs-2010-lithi.pdf. Retrieved 2010-11-01.  ^ "You’ve got the power: the evolution of batteries and the future of fuel cells" (PDF). Toshiba. http://nl.computers.toshiba-europe.com/Contents/Toshiba_nl/NL/WHITEPAPER/files/TISBWhitepapertech.pdf. Retrieved 2009-05-17.  ^ Pomerantz JM. Hyperparathyroidism Resulting From Lithium Treatment Remains Underrecognized. Drug Benefit Trends. 2010;22:62–63 ^ Peatfield, R. C. (1981). "Lithium in migraine and cluster headache: a review.". J. R. Soc Med. 74 (6): 432–436. PMID 7252959.  ^ (2005) Lithium (LI), Beckman Coulter. (Report). Retrieved 2010-10-05. ^ Newman, P. K.; Saunders, M (1979). "Lithium neurotoxicity.". Postgraduate Medical Journal 55 (648): 701. doi:10.1136/pgmj.55.648.701. PMID 537955.  ^ Prasad, A. (1984). "Chronic lithium intake and hyperparathyroidism". European Journal of Clinical Pharmacology 27 (4): 499. doi:10.1007/BF00549602. PMID 6519159.  ^ Bendz, H.; A; B; M; S (1994). "Kidney damage in long-term lithium patients: A cross-sectional study of patients with 15 years or more on lithium". Nephrol Dial Transplant 9 (9): 1250–1254. PMID 7816284. http://ndt.oxfordjournals.org/cgi/content/abstract/9/9/1250.  ^ Stone, K. A. (1999). "Lithium-induced nephrogenic diabetes insipidus". The Journal of the American Board of Family Practice 12 (1): 43–47. PMID 10050642. http://www.jabfm.org/cgi/content/abstract/12/1/43.  ^ "Weight Gain and Bipolar Disorder Treatment". PsychEducation.org. November 2007. http://www.psycheducation.org/hormones/Insulin/weightgain.htm.  ^ "Lithium in drinking water may boost mood". Science News (United Press International). May 1, 2009 at 11:41 PM. http://www.upi.com/Science_News/2009/05/01/Lithium-in-drinking-water-may-boost-mood/UPI-66841241235675/. Retrieved 2009-05-02.  ^ Alleyne, Richard (10:01AM BST 01 May 2009). "Natural levels of lithium in drinking water help reduce suicides". Health: Health News (London: Telegraph). http://www.telegraph.co.uk/health/healthnews/5251365/Natural-levels-of-lithium-in-drinking-water-help-reduce-suicides.html. Retrieved 2009-05-02.  ^ "Scientists Find Correlation Between Lithium in Drinking Water and Reduced Suicide Rates". shortnews.com. 05/02/2009 03:41 PM. http://www.shortnews.com/start.cfm?id=78524. Retrieved 2009-05-02.  ^ Ohgami, H.; Terao, T; Shiotsuki, I; Ishii, N; Iwata, N (2009). "Lithium levels in drinking water and risk of suicide". The British Journal of Psychiatry (The Royal College of Psychiatrists) 194 (5): 194: 464–465. doi:10.1192/bjp.bp.108.055798. PMID 19407280. http://bjp.rcpsych.org/cgi/content/abstract/194/5/464.  ^ "Lithium in water 'curbs suicide'". Health:Medical Notes (BBC). 09:22 GMT, Friday, 1 May 2009 10:22 UK. http://news.bbc.co.uk/2/hi/health/8025454.stm. Retrieved 2009-05-02.  ^ "Organometallics". http://www.sriconsulting.com/CEH/Public/Reports/681.7000/.  ^ Yurkovetskii, A. V.; Kofman, V. L.; Makovetskii, K. L. (2005). "Polymerization of 1,2-dimethylenecyclobutane by organolithium initiators". Russian Chemical Bulletin 37: 1782–1784. doi:10.1007/BF00962487.  ^ Quirk, Roderic P.; Cheng, Pao Luo (1986). "Functionalization of polymeric organolithium compounds. Amination of poly(styryl)lithium". Macromolecules 19: 1291. doi:10.1021/ma00159a001.  ^ Stone, F. G. A.; West, Robert (1980). Advances in organometallic chemistry. Academic Press. p. 55. ISBN 0120311186. http://books.google.com/?id=_gai4kRfcMUC&printsec=frontcover#PPA55,M1.  ^ Bansal, Raj K. (1996). Synthetic approaches in organic chemistry. p. 192. ISBN 0763706655. http://books.google.com/books?id=_SJ2upYN6DwC&pg=PA192.  ^ Totten, George E.; Westbrook, Steven R. and Shah, Rajesh J. (2003). Fuels and lubricants handbook: technology, properties, performance, and testing, Volume 1. ASTM International. p. 559. ISBN 0803120966. http://books.google.com/books?id=J_AkNu-Y1wQC&pg=PA559.  ^ Rand, Salvatore J. (2003). Significance of tests for petroleum products. ASTM International. pp. 150–152. ISBN 0803120974. http://books.google.com/books?id=3FkMrP4Hlw0C&pg=PA152.  ^ Wiberg, Egon; Wiberg, Nils and Holleman, Arnold Frederick Inorganic chemistry, Academic Press (2001) ISBN 0123526515, p. 1089 ^ Makhijani, Arjun and Yih, Katherine (2000). Nuclear Wastelands: A Global Guide to Nuclear Weapons Production and Its Health and Environmental Effects. MIT Press. pp. 59–60. ISBN 0262632047. http://books.google.com/books?id=0oa1vikB3KwC&pg=PA60.  ^ National Research Council (U.S.). Committee on Separations Technology and Transmutation Systems (1996). Nuclear wastes: technologies for separations and transmutation. National Academies Press. p. 278. ISBN 0309052262. http://books.google.com/books?id=iRI7Cx2D4e4C&pg=PA278.  ^ Barnaby, Frank (1993). How nuclear weapons spread: nuclear-weapon proliferation in the 1990s. Routledge. p. 39. ISBN 0415076749. http://books.google.com/books?id=yTIOAAAAQAAJ&pg=PA39.  ^ Baesjr, C (1974). "The chemistry and thermodynamics of molten salt reactor fuels☆". Journal of Nuclear Materials 51: 149. doi:10.1016/0022-3115(74)90124-X.  ^ "Tritium Breeding". ITER. http://www.iter.org/mach/tritiumbreeding. Retrieved September 2010.  ^ Agarwal, Arun (2008). Nobel Prize Winners in Physics. APH Publishing. p. 139. ISBN 8176487430. http://books.google.com/books?id=XyOBx2R2CxEC&pg=PA139.  ^ Hobbs, Philip C. D. (2009). Building Electro-Optical Systems: Making It All Work. John Wiley and Sons. p. 149. ISBN 0470402296. http://books.google.com/books?id=CQ5uKN_MN2gC&pg=PA149.  ^ Point Defects in Lithium Fluoride Films Induced by Gamma Irradiation. 2001. World Scientific. 2002. 819. ISBN 9812381805. http://books.google.com/books?id=FY7s7pPSPtgC&pg=PA819.  ^ Sinton, William M. (1962). "Infrared Spectroscopy of Planets and Stars". Applied Optics 1: 105. doi:10.1364/AO.1.000105.  ^ Ernst-Christian, K. (2004). "Special Materials in Pyrotechnics: III. Application of Lithium and its Compounds in Energetic Systems". Propellants, Explosives, Pyrotechnics 29 (2): 67–80. doi:10.1002/prep.200400032.  ^ Hughes, T.G.; Smith, R.B. and Kiely, D.H. (1983). "Stored Chemical Energy Propulsion System for Underwater Applications". Journal of Energy 7 (2): 128–133. doi:10.2514/3.62644.  ^ Mulloth, L.M. and Finn, J.E. (2005). "Air Quality Systems for Related Enclosed Spaces: Spacecraft Air". The Handbook of Environmental Chemistry. 4H. pp. 383–404. doi:10.1007/b107253.  ^ "Application of lithium chemicals for air regeneration of manned spacecraft". Lithium Corporation of America & Aeropspace Medical Research Laboratories. 1965. http://www.dtic.mil/cgi-bin/GetTRDoc?Location=U2&doc=GetTRDoc.pdf&AD=AD0619497.  ^ Furr, A. K. (2000). CRC handbook of laboratory safety. Boca Raton: CRC Press. pp. 244–246. ISBN 9780849325236. http://books.google.com/?id=Oo3xAmmMlEwC&pg=PA244.  ^ Yacobi S, Ornoy A (2008). "Is lithium a real teratogen? What can we conclude from the prospective versus retrospective studies? A review". Isr J Psychiatry Relat Sci 45 (2): 95–106. PMID 18982835. http://www.psychiatry.org.il/presentation_sender.asp?info_id=53760.  ^ "Illinois Attorney General – Basic Understanding Of Meth". Illinoisattorneygeneral.gov. http://www.illinoisattorneygeneral.gov/methnet/understandingmeth/basics.html. Retrieved 2010-10-06.  ^ Harmon, Aaron R. (2006). "Methamphetamine remediation research act of 2005: Just what the doctor ordered for cleaning up methfields—or sugar pill placebo?" (PDF). North Carolina Journal of Law & Technology 7. http://jolt.unc.edu/sites/default/files/7_nc_jl_tech_421.pdf. Retrieved 2010-10-05.  ^ Samuel C. Levy and Per Bro. (1994). Battery hazards and accident prevention. New York: Plenum Press. pp. 15–16. ISBN 9780306447587. http://books.google.com/?id=i7U-0IB8tjMC&pg=PA15.  ^ "TSA: Safe Travel with Batteries and Devices". Tsa.gov. 2008-01-01. http://www.tsa.gov/travelers/airtravel/assistant/batteries.shtm. Retrieved 2010-10-06.  External links Look up lithium in Wiktionary, the free dictionary. Wikimedia Commons has media related to: Lithium The Periodic Table of Videos video of Lithium at YouTube International Lithium Alliance USGS: Lithium Statistics and Information Lithium Supply & Markets 2009 IM Conference 2009 Sustainable lithium supplies through 2020 in the face of sustainable market growth WebElements.com – Lithium It's Elemental – Lithium University of Southampton, Mountbatten Centre for International Studies, Nuclear History Working Paper No5. v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties Large version  v · d · e  Lithium compounds Inorganic LiAlH4 · LiAlO2 · LiBF4 · LiBH4 · LiBO2 · LiB3O5 · LiBr · LiCl · LiClO · LiClO3 · LiClO4 · LiCoO2 · LiF · LiH · LiI · LiNH2 · LiNO3 · LiNbO3 · LiOH · LiO2 · LiPF6 · LiTaO3 · Li2B4O7 · Li2CO3 · Li2C2 · Li2MoO4 · Li2O · Li2O2 · Li2S · Li2SO4 · Li3N Organic acetate · citrate · diisopropylamide · orotate · succinate Minerals

Amblygonite · Elbaite · Eucryptite · Jadarite · Lepidolite · Lithiophilite · Petalite · Pezzottaite · Saliotite · Spodumene · Sugilite · Tourmaline · Zabuyelite · Zinnwaldite v · d · eAlkali metals     Lithium Li Atomic Number: 3 Atomic Weight: 6.941 Melting Point: 453.69 K Boiling Point: 1615 K Specific mass: 0.534 g/cm3 Electronegativity: 0.98 Sodium Na Atomic Number: 11 Atomic Weight: 22.990 Melting Point: 370.87 K Boiling Point: 1156 K Specific mass: 0.97 g/cm3 Electronegativity: 0.96 Potassium K Atomic Number: 19 Atomic Weight: 39.098 Melting Point: 336.58 K Boiling Point: 1032 K Specific mass: 0.86 g/cm3 Electronegativity: 0.82 Rubidium Rb Atomic Number: 37 Atomic Weight: 85.468 Melting Point: 312.46 K Boiling Point: 961 K Specific mass: 1.53 g/cm3 Electronegativity: 0.82 Caesium Cs Atomic Number: 55 Atomic Weight: 132.905 Melting Point: 301.59 K Boiling Point: 944 K Specific mass: 1.93 g/cm3 Electronegativity: 0.79 Francium Fr Atomic Number: 87 Atomic Weight: (223) Melting Point: 295(?) K Boiling Point: 950(?) K Specific mass: ? g/cm3 Electronegativity: 0.7

Amblygonite · Elbaite · Eucryptite · Jadarite · Lepidolite · Lithiophilite · Petalite · Pezzottaite · Saliotite · Spodumene · Sugilite · Tourmaline · Zabuyelite · Zinnwaldite v · d · eAlkali metals     Lithium Li Atomic Number: 3 Atomic Weight: 6.941 Melting Point: 453.69 K Boiling Point: 1615 K Specific mass: 0.534 g/cm3 Electronegativity: 0.98 Sodium Na Atomic Number: 11 Atomic Weight: 22.990 Melting Point: 370.87 K Boiling Point: 1156 K Specific mass: 0.97 g/cm3 Electronegativity: 0.96 Potassium K Atomic Number: 19 Atomic Weight: 39.098 Melting Point: 336.58 K Boiling Point: 1032 K Specific mass: 0.86 g/cm3 Electronegativity: 0.82 Rubidium Rb Atomic Number: 37 Atomic Weight: 85.468 Melting Point: 312.46 K Boiling Point: 961 K Specific mass: 1.53 g/cm3 Electronegativity: 0.82 Caesium Cs Atomic Number: 55 Atomic Weight: 132.905 Melting Point: 301.59 K Boiling Point: 944 K Specific mass: 1.93 g/cm3 Electronegativity: 0.79 Francium Fr Atomic Number: 87 Atomic Weight: (223) Melting Point: 295(?) K Boiling Point: 950(?) K Specific mass: ? g/cm3 Electronegativity: 0.7

Amblygonite · Elbaite · Eucryptite · Jadarite · Lepidolite · Lithiophilite · Petalite · Pezzottaite · Saliotite · Spodumene · Sugilite · Tourmaline · Zabuyelite · Zinnwaldite v · d · eAlkali metals     Lithium Li Atomic Number: 3 Atomic Weight: 6.941 Melting Point: 453.69 K Boiling Point: 1615 K Specific mass: 0.534 g/cm3 Electronegativity: 0.98 Sodium Na Atomic Number: 11 Atomic Weight: 22.990 Melting Point: 370.87 K Boiling Point: 1156 K Specific mass: 0.97 g/cm3 Electronegativity: 0.96 Potassium K Atomic Number: 19 Atomic Weight: 39.098 Melting Point: 336.58 K Boiling Point: 1032 K Specific mass: 0.86 g/cm3 Electronegativity: 0.82 Rubidium Rb Atomic Number: 37 Atomic Weight: 85.468 Melting Point: 312.46 K Boiling Point: 961 K Specific mass: 1.53 g/cm3 Electronegativity: 0.82 Caesium Cs Atomic Number: 55 Atomic Weight: 132.905 Melting Point: 301.59 K Boiling Point: 944 K Specific mass: 1.93 g/cm3 Electronegativity: 0.79 Francium Fr Atomic Number: 87 Atomic Weight: (223) Melting Point: 295(?) K Boiling Point: 950(?) K Specific mass: ? g/cm3 Electronegativity: 0.7

Amblygonite · Elbaite · Eucryptite · Jadarite · Lepidolite · Lithiophilite · Petalite · Pezzottaite · Saliotite · Spodumene · Sugilite · Tourmaline · Zabuyelite · Zinnwaldite v · d · eAlkali metals     Lithium Li Atomic Number: 3 Atomic Weight: 6.941 Melting Point: 453.69 K Boiling Point: 1615 K Specific mass: 0.534 g/cm3 Electronegativity: 0.98 Sodium Na Atomic Number: 11 Atomic Weight: 22.990 Melting Point: 370.87 K Boiling Point: 1156 K Specific mass: 0.97 g/cm3 Electronegativity: 0.96 Potassium K Atomic Number: 19 Atomic Weight: 39.098 Melting Point: 336.58 K Boiling Point: 1032 K Specific mass: 0.86 g/cm3 Electronegativity: 0.82 Rubidium Rb Atomic Number: 37 Atomic Weight: 85.468 Melting Point: 312.46 K Boiling Point: 961 K Specific mass: 1.53 g/cm3 Electronegativity: 0.82 Caesium Cs Atomic Number: 55 Atomic Weight: 132.905 Melting Point: 301.59 K Boiling Point: 944 K Specific mass: 1.93 g/cm3 Electronegativity: 0.79 Francium Fr Atomic Number: 87 Atomic Weight: (223) Melting Point: 295(?) K Boiling Point: 950(?) K Specific mass: ? g/cm3 Electronegativity: 0.7