View source for Lithium

{{About|the chemical element}}
{{pp-semi-vandalism|small=yes}}

{{Infobox lithium}}
'''Lithium''' is a soft, silver-white [[metal]] that belongs to the [[alkali metal]] group of [[chemical elements]]. It is represented by the [[Chemical symbol|symbol]] '''Li''', and it has the [[atomic number]] 3. Under [[standard conditions for temperature and pressure|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, lithium metal is typically stored in a [[petroleum]] derivative, such as [[mineral oil]] (in which it floats), or petroleum jelly in which it is held below the surface by the viscosity of the medium. When cut open, lithium exhibits a metallic [[luster (mineralogy)|luster]], but contact with moist air [[corrosion|corrodes]] the surface quickly to a dull silvery gray, then black, tarnish. 

The nuclei of lithium are relatively fragile: the two stable lithium isotopes found in nature have lower binding energies per nucleon than any other stable compound nuclides, save [[deuterium]], and [[helium]]-3 (<sup>3</sup>He).<ref name=bind>[[: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.</ref> Though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.<ref>Numerical data from: {{cite journal|doi=10.1086/375492 |first=Katharina |last=Lodders |year=2003 |title=Solar System Abundances and Condensation Temperatures of the Elements |journal=The Astrophysical Journal |volume=591|pages=1220–1247}} Graphed at [[:File:SolarSystemAbundances.jpg]]</ref>

Because of its high [[reactivity (chemistry)|reactivity]], lithium only appears naturally in the form of [[chemical compound|compounds]]. Lithium occurs in a number of [[pegmatite|pegmatitic]] [[mineral]]s, but is also commonly obtained from [[brine]]s and [[clay]]s. On a commercial scale, lithium metal is isolated [[electrolysis|electrolytically]] from a mixture of [[lithium chloride]] and [[potassium chloride]].

Trace amounts of lithium are present in the [[ocean]]s and in some organisms, though the element serves no apparent vital biological function in humans. The lithium [[ion]] Li<sup>+</sup> administered as any of several lithium [[salt (chemistry)|salts]] has proved to be useful as a [[mood stabilizer|mood stabilizing]] drug due to neurological [[lithium pharmacology|effects]] of the ion in the human body. Lithium and its compounds have several industrial applications, including heat-resistant [[glass]] and [[ceramic]]s, high strength-to-weight [[alloy]]s used in [[aircraft]], and [[lithium battery|lithium batteries]]. Lithium also has important links to [[nuclear physics]]. The [[Nuclear transmutation|transmutation]] of lithium atoms to [[tritium]] was the first man-made form of a [[nuclear fusion|nuclear fusion reaction]], and [[lithium deuteride]] serves as a [[nuclear fusion|fusion]] fuel in [[Teller-Ulam design|staged thermonuclear weapon]]s.

==Characteristics==
===Physical===
[[Image:LithiumPelletsUSGOV.jpg|thumb|left|Lithium pellets (covered in white lithium hydroxide)]]
{{See also|Alkali metal}}
Like the other [[alkali metal]]s, lithium has a single [[valence electron]] that is easily given up to form a [[cation]].<ref name=krebs/> Because of this, it is a good conductor of both [[heat]] and [[electricity]] and highly reactive, though it is the least reactive of the alkali metals due to the proximity of its valence electron to its [[atomic nucleus|nucleus]].<ref name=krebs/>

Lithium is soft enough to be cut with a knife, and it is the lightest of the metals of the periodic table. When cut, it possesses a silvery-white color that quickly changes to gray due to oxidation.<ref name=krebs/> It also has a low density (approximately 0.534 g/cm<sup>3</sup>), giving sticks of the metal a similar heft to dowels of wood, such as pine. It floats on water, with which it reacts easily but with noticeably less energy than other alkali metals do. The reaction forms [[hydrogen]] gas and [[lithium hydroxide]] in aqueous solution.<ref name=krebs/> 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.<ref name=emsley/>

Lithium possesses a low [[coefficient of thermal expansion]] and the highest [[specific heat capacity]] of any solid element. Lithium is [[superconductive]] below 400 [[microkelvin|μK]] at standard pressure<ref>{{cite journal|doi=10.1038/nature05820|year=2007|author=Tuoriniemi, J; Juntunen-Nurmilaukas, K; Uusvuori, J; Pentti, E; Salmela, A; Sebedash, A|title=Superconductivity in lithium below 0.4 millikelvin at ambient pressure.|volume=447|issue=7141|pages=187–9|pmid=17495921|journal=Nature}}</ref> and at higher temperatures (more than 9 kelvins) at very high pressures (over 200,000 atmospheres)<ref>{{cite journal|doi=10.1126/science.1078535|year=2002|author=Struzhkin, Vv; Eremets, Mi; Gan, W; Mao, Hk; Hemley, Rj|title=Superconductivity in dense lithium.|volume=298|issue=5596|pages=1213–5|pmid=12386338|journal=Science}}</ref> At cryogenic temperatures, lithium, like sodium, undergoes [[diffusionless transformations|diffusionless phase change transformations]]. At 4.2K it has a [[rhombohedral crystal system]] (with a nine-layer repeat spacing)<ref name="overhauser">{{ cite journal|first = A. W.|last = Overhauser|title=Crystal Structure of Lithium at 4.2 K|doi=10.1103/PhysRevLett.53.64|volume=53|pages=64–65|year=1984|journal = Physical Review Letters}}</ref>; 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.

===Chemical===
In moist air, lithium metal rapidly tarnishes to form a black coating of [[lithium hydroxide]] (LiOH and LiOH·H<sub>2</sub>O), [[lithium nitride]] (Li<sub>3</sub>N) and [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>, the result of a secondary reaction between LiOH and [[carbon dioxide|CO<sub>2</sub>]]).<ref name="kamienski" />

When placed over a flame, lithium gives 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 vapours.<ref>{{cite web |url=http://www.pitt.edu/~alw11/InterestInfo/Articles/Bunsen%20and%20Kirchoff.pdf|title=Chemical Analysis By Observation of Spectra|accessdate=2009-11-05 |last=Kirchoff |first=Gustav|coauthors=Bunsen, Robert |publisher=University of Pittsburgh}}</ref>

Lithium metal 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 [[Fire_extinguisher#Types_of_extinguishing_agents|Types of extinguishing agents]]). Lithium is the only metal which reacts with [[nitrogen]] under [[normal conditions]].

===Lithium compounds===
{{Expand section|date=July 2009}}
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 N<sub>2</sub>, the formation of an [[oxide]] when burnt in O<sub>2</sub>, [[salt (chemistry)|salts]] with similar [[solubility|solubilities]], and thermal instability of the [[carbonate]]s and nitrides.<ref name="kamienski">{{ cite book|first = Conrad W. last = Kamienski, McDonald, Daniel P.; Stark, Marshall W.; Papcun, John R.|chapter =Lithium and lithium compounds|title =Kirk-Othmer Encyclopedia of Chemical Technology|publisher = John Wiley & Sons, Inc.| year = 2004|doi =10.1002/0471238961.1209200811011309.a01.pub2}}</ref>

===Isotopes===
{{Main|Isotopes of lithium}}
Naturally occurring lithium is composed of two stable [[isotope]]s, <sup>6</sup>Li and <sup>7</sup>Li, the latter being the more abundant (92.5% [[natural abundance]]).<ref name=krebs/><ref name=emsley/><ref name=isotopesproject>{{cite web |url=http://ie.lbl.gov/education/parent/Li_iso.htm |title=Isotopes of Lithium|accessdate=2008-04-21|publisher=Berkeley National Laboratory, The Isotopes Project}}</ref> 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]]. Seven [[radioisotope]]s have been characterized, the most stable being <sup>8</sup>Li with a [[half-life]] of 838 [[millisecond|ms]] and <sup>9</sup>Li with a half-life of 178.3&nbsp;ms. All of the remaining [[radioactive]] isotopes have half-lives that are shorter than 8.6&nbsp;ms. The shortest-lived isotope of lithium is <sup>4</sup>Li, which decays through [[proton emission]] and has a half-life of 7.58043 × 10<sup>−23</sup> s.

<sup>7</sup>Li is one of the [[primordial elements]] (or, more properly, primordial isotopes) produced in [[Big Bang nucleosynthesis]]. A small amount of both <sup>6</sup>Li and <sup>7</sup>Li are produced in stars, but are thought to be burned as fast as it is produced.<ref>{{cite journal|title=Lithium Isotopic Abundances in Metal-poor Halo Stars |year=2006|journal=The Astrophysical Journal|doi = 10.1086/503538|volume=644|page=229|author=Asplund, M. ''et al.''}}</ref> Additional small amounts of lithium of both <sup>6</sup>Li and <sup>7</sup>Li may be generated from solar wind, cosmic rays, and early solar system <sup>7</sup>[[Beryllium|Be]] and <sup>10</sup>Be radioactive decay.<ref>{{cite journal |url=http://sims.ess.ucla.edu/PDF/Chaussidon_et_al_Geochim%20Cosmochim_2006a.pdf |doi=10.1016/j.gca.2005.08.016 |first1=M. |last1=Chaussidon |first2=F. |last2=Robert |first3=K.D. |last3=McKeegan |journal=Geochimica et Cosmochimica Acta |volume=70 |number=1|year=2006 |pages=224–245 |title=Li and B isotopic variations in an Allende CAI: Evidence for the in situ decay of short-lived <sup>10</sup>Be and for the possible presence of the short−lived nuclide <sup>7</sup>Be in the early solar system|format=free download pdf}}</ref> <sup>7</sup>Li can also be generated in [[carbon stars]].<ref>{{cite journal |title=Episodic lithium production by extra-mixing in red giants |bibcode=2000A&A...358L..49D |first1=P. A. |last1=Denissenkov |first2=A. |last2=Weiss |journal=Astronomy and Astrophysics |volume=358 |pages=L49–L52 |year=2000}}</ref>

Lithium isotopes fractionate substantially during a wide variety of natural processes,<ref>{{cite journal |year=2004 |first1=H.M. |last1=Seitz |first2=G.P. |last2=Brey |first3=Y. |last3=Lahaye |first4=S. |last4=Durali |first5=S.|last5=Weyer |title=Lithium isotopic signatures of peridotite xenoliths and isotopic fractionation at high temperature between olivine and pyroxenes |journal=Chemical Geology |volume=212 |number=1-2|doi=10.1016/j.chemgeo.2004.08.009|pages=163–177}}</ref> including mineral formation (chemical precipitation), [[metabolism]], and [[ion exchange]]. Lithium ions substitute for [[magnesium]] and [[iron]] in octahedral sites in [[clay]] minerals, where <sup>6</sup>Li is preferred to <sup>7</sup>Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic <sup>11</sup>Li is known to exhibit a [[nuclear halo]].
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==History and etymology==
[[Petalite]] (LiAlSi<sub>4</sub>O<sub>10</sub>, which is lithium aluminium silicate) was first discovered in 1800 by the [[Brazil]]ian [[chemist]] [[José Bonifácio de Andrada e Silva]], who discovered this mineral in a mine on the island of [[Utö, Sweden]].<ref name=mindat>{{cite web |url=http://www.mindat.org/min-3171.html |title=Petalite Mineral Information |accessdate=10 August 2009}}</ref><ref name=webelementshistory>{{cite web |url=http://www.webelements.com/lithium/history.html |title=Lithium:Historical information |accessdate=10 August 2009}}</ref><ref name=discovery>{{cite book |title=Discovery of the Elements |last=Weeks |first=Mary |year=2003 |page=124 |publisher=Kessinger Publishing |location=Whitefish, Montana, United States |isbn=0766138720 |url=http://books.google.com/books?id=SJIk9BPdNWcC&source=gbs_navlinks_s |accessdate=10 August 2009}}</ref> However, it was not until 1817 that [[Johan August Arfwedson]], then working in the laboratory of the chemist [[Jöns Jakob Berzelius]], [[discovery of the chemical elements|detected]] the presence of a new element while analyzing petalite ore.<ref name=berzelius>{{cite web |url=http://www.chemeddl.org/collections/ptl/ptl/chemists/bios/arfwedson.html |title=Johan August Arfwedson |accessdate=10 August 2009 |work= Periodic Table Live!}}</ref><ref name=uwis>{{cite web |url=http://web.archive.org/web/20080605152857/http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm |title=Johan Arfwedson |accessdate=10 August 2009}}</ref><ref name=vanderkrogt>{{cite web|publisher = Elementymology & Elements Multidict|title = Lithium| first = Peter|last =van der Krogt|url = http://web.archive.org/web/20080522144114/http://www.vanderkrogt.net/elements/elem/li.html| accessdate = 2008-09-18}}</ref> This element formed compounds similar to those of [[sodium]] and [[potassium]], though its [[lithium carbonate|carbonate]] and [[lithium hydroxide|hydroxide]] were less [[solubility|soluble in water]] and more [[Base (chemistry)|alkaline]].<ref name=compounds>{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/compounds.html |title=Compounds of the Group 1 Elements |accessdate=10 August 2009 |last=Clark |first=Jim |year=2005}}</ref> Berzelius gave the alkaline material the name "lithos", from the [[Greek language|Greek]] word ''λιθoς'' (transliterated as ''lithos'', meaning "stone"), to reflect its discovery in a solid mineral, as opposed to sodium and potassium, which had been discovered in plant tissues. The name of this element was later standardized as "lithium".<ref name=krebs/><ref name=webelementshistory/><ref name=vanderkrogt/> Arfwedson later showed that this same element was present in the minerals [[spodumene]] and [[lepidolite]].<ref name=webelementshistory/> In 1818, [[Christian Gmelin]] was the first man to observe that lithium salts give a bright red color in flame.<ref name=webelementshistory/> However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.<ref name=webelementshistory/><ref name=vanderkrogt/><ref name="eote">{{cite book|year = 2004|title = Encyclopedia of the Elements: Technical Data - History - Processing - Applications|publisher = Wiley|isbn = 978-3527306664|pages = 287–300|author = Per Enghag}}</ref> This element, lithium, was not isolated until 1821, when [[William Thomas Brande]] isolated the element by performing [[electrolysis]] on [[lithium oxide]], a process that had previously been employed by the chemist [[Sir Humphry Davy]] to isolate the alkali metals potassium and sodium.<ref name=emsley>{{cite book |last=Emsley |first=John |title=Nature's Building Blocks |publisher=Oxford University Press |location=Oxford |year=2001 |isbn=0198503415}}</ref><ref name="eote" /><ref>{{cite web|url = http://www.diracdelta.co.uk/science/source/t/i/timeline/source.html|title = Timeline science and engineering|publisher = DiracDelta Science & Engineering Encyclopedia| accessdate = 2008-09-18}}</ref> Brande also described some pure salts of lithium, such as the chloride, and he performed an estimate of its atomic weight. In 1855, larger quantities of lithium were produced through the electrolysis of [[lithium chloride]] by [[Robert Bunsen]] and [[Augustus Matthiessen]].<ref name=webelementshistory/> The discovery of this procedure henceforth led to commercial production of lithium metal, beginning in 1923, by the German company [[Metallgesellschaft AG]], which performed an electrolysis of a liquid mixture of lithium chloride and [[potassium chloride]].<ref name=webelementshistory/><ref>{{cite web| url = http://www.echeat.com/essay.php?t=29195|title = Analysis of the Element Lithium|first = Thomas|last = Green|date =2006-06-11| publisher = echeat}}</ref>

The production and use of lithium underwent several drastic changes in history. The first major application of lithium became high temperature [[petroleum|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 when in the beginning of the cold war the need for the production of [[Nuclear weapon design|nuclear fusion weapons]] arose and the dominant fusion material [[tritium]] had to be made by irradiating lithium-6. 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 tons of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75% .<ref name="USGSCR1994">{{cite web | url = http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450494.pdf| title = Commodity Report 1994: Lithium}}</ref> <!--With only 7.5% of lithium-6 this makes ca 2200 tons of  lithium-6.-->

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]].<ref name="ciuz2003">{{cite journal | doi = 10.1002/ciuz.200300264 | title = Lithium und seine Verbindungen - Industrielle, medizinische und wissenschaftliche Bedeutung | year = 2003 | last1 = Deberitz | first1 = JüRgen | first2 = Gernot | journal = Chemie in unserer Zeit | volume = 37 | pages = 258 | last2 = Boche}}</ref><ref name="ciuz2003">{{cite journal | doi = 10.1002/ciuz.19850190505 | title = Lithium - wie es nicht im Lehrbuch steht | year = 1985 | last1 = Bauer | first1 = Richard | journal = Chemie in unserer Zeit | volume = 19 | pages = 167 }}</ref> 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.<ref name="USGSCR1994"/> Then, in the mid 1990's several companies started to extract lithium from [[brine]]; this method proved to be less expensive 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.<ref name="USGSCR1994">{{cite web | url = http://minerals.usgs.gov/minerals/pubs/commodity/lithium/myb1-2007-lithi.pdf | title = Minerals Yearbook 2007 : Lithium}}</ref> New companies have expanded brine extraction efforts to meet the rising demand.<ref name="IMR">{{cite book | first = Jessica Elzea | last = Kogel | title = Industrial minerals & rocks: commodities, markets, and uses | isbn = 9780873352338 | pages = 599 | url = http://books.google.com/books?id=zNicdkuulE4C&pg=PA600&lpg=PAPA599 | chapter = Lithium | year = 2006 | publisher = Society for Mining, Metallurgy, and Exploration | location = Littleton, Colo.}})</ref><!--http://books.google.de/books?id=8erDL_DnsgAC&pg=PA339-->

==Occurrence==
<div style="float:right; margin:5px;">
{|
|[[Image:Relative abundance of elements.png|thumb|400px|Lithium is about as common as [[chlorine]] in the Earth's upper continental [[Crust (geology)|crust]], on a per-atom basis.]]
|-
|
{|class="wikitable sortable"
|+Lithium mine production (2008) and reserves in metric tonnes<ref name="minerals.usgs.gov">U.S. Geological Survey, 2009, [http://minerals.usgs.gov/minerals/pubs/mcs/ Mineral commodity summaries] 2009: U.S. Geological Survey, 195 p. Page 95.</ref>
|-
! Country
! Production
! Reserves
! Reserve base
|-
|{{flag|Argentina}}
| 3,200
| Not available
| Not available
|-
|{{flag|Australia}}
| 6,900
| 170,000
| 220,000
|-
|{{flag|Bolivia}}
| 0
| 0
| 5,400,000
|-
|{{flag|Brazil}}
| 180
| 190,000
| 910,000
|-
|{{flag|Canada}}
| 710
| 180,000
| 360,000
|-
|{{flag|Chile}}
| 12,000
| 3,000,000
| 3,000,000
|-
|{{flag|People's Republic of China}}
| 3,500
| 540,000
| 1,100,000
|-
|{{flag|Portugal}}
| 570
| Not available
| Not available
|-
|{{flag|United States}}
| Withheld
| 38,000
| 410,000
|-
|{{flag|Zimbabwe}}
| 300
| 23,000
| 27,000
|-
| '''World total'''
| '''27,400'''
| '''4,100,000'''
| '''11,000,000'''
|}
|}
</div>

===Astronomical occurrence===
{{Main|nucleosynthesis}}
According to modern cosmological theory, both stable isotopes of lithium&mdash;<sup>6</sup>Li and <sup>7</sup>Li&mdash;were among the 3 elements [[nucleosynthesis|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.<ref name=cld>{{cite news|url=http://www.universetoday.com/2006/08/16/why-old-stars-seem-to-lack-lithium/ |title=Why Old Stars Seem to Lack Lithium|date= 16th Aug 2006|author= Fraser Cain}}</ref> 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.<ref name=emsley/>

Though it was one of the 3 first elements to be synthesized in the Big Bang, lithium, as well as [[beryllium]] and [[boron]] are markedly less abundant than the elements with either lower or higher atomic number. This is due to the low temperature necessary to destroy lithium, and a lack of common processes to produce it.<ref name=wesleyan>{{cite web |url=http://web.archive.org/web/20060901133923/http://www.astro.wesleyan.edu/~bill/courses/astr231/wes_only/element_abundances.pdf |title=Element Abundances |accessdate=2009-11-17}}</ref>

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.<ref name=emsley/><ref>{{cite web |url=http://www.universetoday.com/guide-to-space/stars/brown-dwarf |title=Brown Dwarf |accessdate=2009-11-17 |last=Cain |first=Fraser |publisher=Universe Today}}</ref><ref>{{cite web |url=http://www-int.stsci.edu/~inr/ldwarf3.html |title=L Dwarf Classification |accessdate=2009-11-17}}</ref> 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&mdash;neutron stars or black holes&mdash;whose gravity evidently pulls heavier lithium to the surface of a hydrogen-helium star, causing more lithium to be observed.<ref name=emsley/>

===Occurrence on Earth===
{{See also|:category:Lithium minerals|l1=Lithium minerals}}
Lithium is widely distributed on Earth but does not naturally occur in elemental form due to its high reactivity.<ref name=krebs>{{cite book|last = Krebs|first = Robert E.|year = 2006|title = The History and Use of Our Earth's Chemical Elements: A Reference Guide|publisher = Greenwood Press|location = Westport, Conn.|isbn = 0-313-33438-2}}</ref> Estimates for [[crust (geology)|crustal]] content range from 20 to 70 ppm by weight.<ref name="kamienski" /> In keeping with its name, lithium forms a minor part of [[igneous rock]]s, with the largest concentrations in [[granite]]s. Granitic [[pegmatite]]s also provide the greatest abundance of lithium-containing minerals, with [[spodumene]] and [[petalite]] being the most commercially viable sources.<ref name="kamienski" /> A newer source for lithium is [[hectorite]] clay, the only active development of which is through the Western Lithium Corporation in the United States.<ref>{{cite journal|author= Moores, S.|title= Between a rock and a salt lake|journal= Industrial Minerals|date= June 2007|page=58|volume=477}}</ref>

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."<ref>''Handbook of Lithium and Natural Calcium'', Donald Garrett, Academic Press, 2004, cited in ''[http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf The Trouble with Lithium 2]''</ref> At 20&nbsp;mg lithium per kg of Earth's crust <ref>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)]]</ref>, lithium is the 25th most abundant element. [[Nickel]] and [[lead]] have the about the same abundance.

The largest reserve base of lithium is in the [[Salar de Uyuni]] area of [[Bolivia]], which has 5.4 million tons. According to the [[US Geological Survey]], the production and reserves of lithium in metric tons are as follows<ref name="minerals.usgs.gov"/><ref>[http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf Lithium_Microscope]</ref>:

Contrary to the USGS data in the table, other estimates put Chile's reserve base at 7,520,000 metric tons of lithium, and Argentina's at 6,000,000 metric tons.<ref>Clarke, G.M. and Harben, P.W., "Lithium Availability Wall Map". Published June 2009. Referenced at [http://www.lithiumalliance.org/about-lithium/lithium-sources/85-broad-based-lithium-reserves International Lithium Alliance]</ref>

Seawater contains an estimated 230 billion tons of lithium, though at a low concentration of 0.1 to 0.2 ppm.<ref>{{cite web |url=http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html |title=Lithium Occurrence|accessdate=2009-03-13|publisher=Institute of Ocean Energy, Saga University, Japan}}</ref>

==Production==
[[File:Lithium mine, Salar del Hombre Muerto, Argentina.jpg|thumb|Lithium mine, Salar del Hombre Muerto, [[Argentina]]. The brine in this [[salt flat]] is rich in lithium, and the mine concentrates the brine by pumping it into [[Salt evaporation pond|solar evaporation ponds]]. 2009 image from NASA’s [[EO-1]] satellite]]
[[File:Uyuni landsat.JPG|thumb|[[Salar de Uyuni|Uyuni]] salt flat in Bolivia]]
Since the end of [[World War II]] lithium metal production has greatly increased. The metal is separated from other elements in [[igneous rock|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 [[electrolysis|electrolytically]] from a mixture of fused [[lithium chloride]] and [[potassium chloride]]. In 1998 it was about {{nowrap|95 [[United States dollar|US$]] / kg}} (or 43 [[US dollar|US$]]/[[Pound (mass)|pound]]).<ref name="ober">{{cite web |url=http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450798.pdf |title=Lithium|accessdate = 2007-08-19|last=Ober |first=Joyce A |format=pdf |pages = 77–78| publisher=[[United States Geological Survey]]}}</ref>

Deposits of lithium are found in [[South America]] throughout the [[Andes]] mountain chain. [[Chile]] is the leading lithium metal producer, followed by [[Argentina]]. Both countries recover the lithium from brine pools. In the [[United States]] lithium is recovered from brine pools in [[Nevada]].<ref name=CRC>{{cite book| author = Hammond, C. R. |title = The Elements, in Handbook of Chemistry and Physics 81st edition| publisher =CRC press| year = 2000| isbn = 0849304814}}</ref> 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.<ref name="romero">{{cite news|author= Simon Romero|title= In Bolivia, a Tight Grip on the Next Big Resource |url=http://www.nytimes.com/2009/02/03/world/americas/03lithium.html?ref=world|publisher=New York Times |date=February 2, 2009}}</ref> According to the [[US Geological Survey]], Bolivia's [[Uyuni]] Desert has 5.4 million tons of lithium, which can be used to make [[lithium battery|batteries]] for [[hybrid vehicle|hybrid]] and [[electric vehicle]]s.<ref name="romero" /><ref>{{cite web|publisher=USGS|url=http://minerals.usgs.gov/minerals/pubs/mcs/2009/mcs2009.pdf |title=USGS Mineral Commodities Summaries 2009}}</ref>

China may emerge as a significant producer of brine-source [[lithium carbonate]] around 2010. There is potential production of up to 55,000 tons per year if projects in [[Qinghai]] province and [[Tibet]] proceed.<ref name="Legers">{{cite web|url=http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf |format=PDF|title=The Trouble With Lithium 2 |accessdate = 2008-07-07 |date=May 28, 2008|publisher=Meridian International Research}}</ref>

The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tons of the known global lithium reserve base.<ref name="Tahil">{{cite web|url=http://www.meridian-int-res.com/Projects/Lithium_Problem_2.pdf |format=PDF|title=The Trouble with Lithium|accessdate = 2008-07-07|year=2007|publisher=Meridian International Research}}</ref>  This is enough for approximately 4 billion [[electric car]]s.<ref>{{cite web |url=http://www.sequence-omega.net/2009/05/15/lithium-carbonate-supplies-abound/ |title=Lithium Carbonate Supplies Abound! |author=Anthony |date=May 15, 2009 |work=Sequence Omega |accessdate=May 18, 2010 }}</ref>

In 1976 a [[United States National Research Council|National Research Council]] Panel estimated lithium resources at 10.6 million tons for the Western World.<ref>{{cite book|author=Evans, R.K. |year=1978|title=Lithium Reserves and Resources, Energy, Vol 3 |publisher=Pergamon Press}}</ref> With the inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total had nearly tripled by 2008.<ref>{{cite news|author=Evans, R.K. |year=2008|title=An Abundance of Lithium|url= http://www.worldlithium.com/Abstract.html|accessdate=2009-07-07}}</ref><ref>{{cite news|author=Evans, R.K. |year=2008|title=An Abundance of Lithium Part 2|url= http://www.worldlithium.com/AN_ABUNDANCE_OF_LITHIUM_-_Part_2.html|accessdate=2009-07-07}}</ref>

==Applications==
Because of its [[specific heat]] capacity, the highest of all [[solid]]s, lithium is often used in [[coolant]]s for [[heat transfer]] applications.

In the later years of the 20th century lithium became important as an [[anode]] material. Used in [[lithium-ion battery|lithium-ion batteries]] because of its high [[electrochemical potential]], a typical cell can generate approximately 3 [[volt]]s, compared with 2.1 volts for [[lead-acid battery|lead/acid]] or 1.5 volts for [[zinc-carbon cell]]s. Because of its low [[atomic mass]], it also has a high charge- and power-to-weight ratio.

Lithium is also used in the pharmaceutical and fine-chemical industry in the manufacture of [[organolithium reagent]]s, which are used both as strong bases and as reagents for the formation of [[carbon-carbon bond]]s. Organolithiums are also used in [[polymer]] synthesis as [[catalyst]]s/[[radical initiator|initiators]]<ref>{{cite web|url=http://www.sriconsulting.com/CEH/Public/Reports/681.7000/ |title=Organometallics}}</ref> in [[Anionic addition polymerization|anionic polymerization]] of [[functional group|unfunctionalised]] [[olefin]]s.<ref>{{cite journal|title=Polymerization of 1,2-dimethylenecyclobutane by organolithium initiators|journal= Russian Chemical Bulletin| volume =37|year=2005|doi=10.1007/BF00962487|pages=1782–1784|author=Yurkovetskii, A. V.|first2=V. L.|first3=K. L.|last2=Kofman|last3=Makovetskii}}</ref><ref>{{cite journal|doi=10.1021/ma00159a001|title=Functionalization of polymeric organolithium compounds. Amination of poly(styryl)lithium|year=1986|author=Quirk, Roderic P.|journal=Macromolecules|volume=19|pages=1291|first2=Pao Luo|last2=Cheng}}</ref><ref>{{cite book|title=Advances in organometallic chemistry|author= Stone, F. G. A.; West, Robert| publisher= Academic Press|year= 1980|isbn= 0120311186|page=55|url= http://books.google.com/books?id=_gai4kRfcMUC&printsec=frontcover#PPA55,M1}}</ref>

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 percent lithium-6. Large amounts of lithium-6 have been produced by [[isotope separation]] for use in [[nuclear weapon]]s. Lithium-7 gained interest for use in [[nuclear reactor]] [[coolant]]s.

===Medical use===
{{Main|Lithium pharmacology}}
Lithium salts were used during the 19th century to treat [[gout]]. Lithium salts such as [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>), [[lithium citrate]], and [[lithium orotate]] are [[mood stabilizer]]s. They are used in the treatment of [[bipolar disorder]] since, unlike most other mood altering drugs, they counteract both [[mania]] and [[depression (mood)|depression]]. 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.<ref>Pomerantz JM. [http://dbt.consultantlive.com/display/article/1145628/1544855 Hyperparathyroidism Resulting From Lithium Treatment Remains Underrecognized]. ''Drug Benefit Trends.'' 2010;22:62-63</ref> Lithium can also be used to augment [[antidepressant]]s. 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 headache]]s.<ref>{{cite journal|journal = J. R. Soc Med.|year = 1981|volume = 74|issue = 6|pages = 432–436|title = Lithium in migraine and cluster headache: a review.|first = R. C.|last = Peatfield|pmid = 7252959|pmc = 1438791}}</ref>

The active principle in these salts is the lithium ion Li<sup>+</sup>. Although this ion has a smaller diameter than either Na<sup>+</sup> or K<sup>+</sup>, in a watery environment like the [[cytoplasmic fluid]], Li<sup>+</sup> binds to the oxygen atoms of water, making it effectively larger than either Na<sup>+</sup> or K<sup>+</sup> ions. How Li<sup>+</sup> works in the [[central nervous system]] is still a matter of debate. Li<sup>+</sup> elevates brain levels of [[tryptophan]], [[5-HT]] (serotonin), and [[5-HIAA]] (a serotonin metabolite). Serotonin is related to mood stability. Li<sup>+</sup> also reduces [[catecholamine]] activity in the brain (associated with brain activation and mania), by enhancing [[reuptake]] and reducing release. [[therapeutic index|Therapeutically useful amounts]] of lithium (~ 0.6 to 1.2&nbsp;mmol/l) are only slightly lower than toxic amounts (>1.5&nbsp;mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity.

Common side effects of lithium treatment include muscle [[tremor]]s, twitching, [[ataxia]]<ref>{{cite journal|doi = 10.1136/pgmj.55.648.701|title = Lithium neurotoxicity.|year = 1979|author = Newman, P. K.|journal = Postgraduate Medical Journal|volume = 55|pages = 701|pmid = 537955|first2 = M|issue = 648|pmc = 2425735|last2 = Saunders}}</ref> and [[hypothyroidism]]. Long term use is linked to [[hyperparathyroidism]]<ref>{{cite journal|doi=10.1007/BF00549602|title=Chronic lithium intake and hyperparathyroidism|year=1984|author=Prasad, A.|journal=European Journal of Clinical Pharmacology|volume=27|pages=499|pmid=6519159|issue=4}}</ref>, [[hypercalcemia]] (bone loss), [[hypertension]], damage of [[renal tubule|tubuli]] in the [[kidney]], [[nephrogenic diabetes insipidus]] ([[polyuria]] and [[polydipsia]]) and/or glomerular damage - even to the point of [[uremia]]<ref>{{cite journal | url = http://ndt.oxfordjournals.org/cgi/content/abstract/9/9/1250 | title = Kidney damage in long-term lithium patients: A cross-sectional study of patients with 15 years or more on lithium | journal = Nephrol Dial Transplant | year = 1994 | volume = 9 | pages =1250–1254 | pmid = 7816284 | issue = 9 | author = Bendz, H. | author2 = A | author3 = B | author4 = M | author5 = S}}</ref>,  [[seizure]]s<ref>{{cite journal|journal = The Journal of the American Board of Family Practice|volume = 12|issue = 1|pages = 43–47|title = Lithium-induced nephrogenic diabetes insipidus|first = K. A.|last = Stone|url = http://www.jabfm.org/cgi/content/abstract/12/1/43|pmid = 10050642|year = 1999}}</ref> and weight gain.<ref>{{cite news|url=http://www.psycheducation.org/hormones/Insulin/weightgain.htm|title=Weight Gain and Bipolar Disorder Treatment|date= November 2007|publisher= PsychEducation.org}}</ref> Some of the side-effects are a result of the increased elimination of [[potassium]].

There appears to be an increased risk of [[Ebstein's anomaly|Ebstein (cardiac) Anomaly]] in infants born to women taking lithium during the first trimester of [[pregnancy]].

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<ref>{{cite news|url=http://www.upi.com/Science_News/2009/05/01/Lithium-in-drinking-water-may-boost-mood/UPI-66841241235675/|title=Lithium in drinking water may boost mood|date=May 1, 2009 at 11:41 PM|work=Science News|publisher=United Press International|accessdate=2009-05-02}}</ref><ref>{{cite news|url=http://www.telegraph.co.uk/health/healthnews/5251365/Natural-levels-of-lithium-in-drinking-water-help-reduce-suicides.html|title=Natural levels of lithium in drinking water help reduce suicides |last=Alleyne|first=Richard |date=10:01AM BST 01 May 2009|work=Health: Health News|publisher=Telegraph|accessdate=2009-05-02}}</ref><ref>{{cite news|url=http://www.shortnews.com/start.cfm?id=78524|title=Scientists Find Correlation Between Lithium in Drinking Water and Reduced Suicide Rates|date=05/02/2009 03:41 PM |publisher=shortnews.com|accessdate=2009-05-02}}</ref><ref>{{cite journal|year=2009|title=Lithium levels in drinking water and risk of suicide|journal=The British Journal of Psychiatry |publisher=The Royal College of Psychiatrists|pages=194: 464–465| doi=10.1192/bjp.bp.108.055798|url=http://bjp.rcpsych.org/cgi/content/abstract/194/5/464|pmid=19407280|author=Ohgami, H.|volume=194|first2=T|first3=I|first4=N|first5=N|issue=5|last2=Terao|last3=Shiotsuki|last4=Ishii|last5=Iwata}}</ref> but did not address whether lithium in drinking water causes the negative side effects associated with higher doses of the element.<ref>{{cite web|url=http://news.bbc.co.uk/2/hi/health/8025454.stm|title=Lithium in water 'curbs suicide' |date=09:22 GMT, Friday, 1 May 2009 10:22 UK|work=Health:Medical Notes|publisher=BBC|accessdate=2009-05-02}}</ref>

===Other uses===
[[Image:FlammenfärbungLi.png|thumb|upright|The red lithium flame leads to lithium's use in flares and [[pyrotechnics]]]]
*Electrical and electronic uses:
**[[Lithium batteries]] are [[disposable]] ([[primary cell|primary]]) [[Battery (electricity)|batteries]] with lithium metal or lithium compounds as an [[anode]]. Lithium batteries are not to be confused with [[lithium-ion battery|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 phone]]s and [[optical modulator]]s, for such components as [[crystal oscillator|resonant crystals]]. Lithium applications are used in more than 60% of mobile phones.<ref>{{cite web|url = http://nl.computers.toshiba-europe.com/Contents/Toshiba_nl/NL/WHITEPAPER/files/TISBWhitepapertech.pdf|title = You’ve got the power: the evolution of batteries and the future of fuel cells|publisher = Toshiba|format = PDF|accessdate = 2009-05-17}}</ref>
*Chemical uses:
**[[Lithium chloride]] and [[lithium bromide]] are extremely [[hygroscopic]] and are used as [[desiccant]]s.
**Lithium metal is used in the preparation of [[organolithium]] compounds.<!-- and as an [[alloy]]ing agent used to synthesize [[organic compound]]s. ?? is this as a reductant - sounds familiar but can't think of an example -->
*General engineering:
**[[Lithium stearate]] is a common all-purpose, high-temperature [[lithium-based grease|lubricant]].
**When used as a [[flux (metallurgy)|flux]] for [[welding]] or [[soldering]], lithium promotes the fusing of metals during and eliminates the forming of [[oxide]]s by absorbing impurities. Its fusing quality is also important as a flux for producing [[ceramic]]s, [[Vitreous enamel|enamels]] and [[glass]].
*[[Alloy]]s of the metal with [[aluminium]], [[cadmium]], [[copper]] and [[manganese]] are used to make high-performance [[aircraft]] parts (see also [[Al-Li|Lithium-aluminium alloys]]).
*Optics:
**Lithium is sometimes used in focal lenses, including spectacles and the glass for the 200-inch (5.08&nbsp;m) [[telescope]] at [[Mt. Palomar]].{{Citation needed|date=May 2009|reason=see Talk}}
**The high non-linearity of lithium niobate also makes it useful in [[nonlinear optics|non-linear optics applications]].
**[[Lithium fluoride]], artificially grown as [[crystal]], is clear and transparent and often used in specialist optics for [[infrared|IR]], [[ultraviolet|UV]] and VUV ([[vacuum UV]]) applications. It has the lowest [[refractive index]] and the farthest transmission range in the deep UV of all common materials.
*Rocketry:
**Metallic lithium and its complex [[hydride]]s, such a [[lithium aluminium hydride|Li[AlH<sub>4</sub>]]], are used as high energy additives to [[rocket propellant]]s<sup>[3]</sup>.
**[[Lithium peroxide]], [[lithium nitrate]], [[lithium chlorate]] and [[lithium perchlorate]] are used as [[oxidizer]]s in [[rocket propellant]]s, and also in [[oxygen candle]]s that supply [[submarine]]s and [[space capsule]]s with [[oxygen]].<ref>{{cite journal|author = Ernst-Christian, K.|title = Special Materials in Pyrotechnics: III. Application of Lithium and its Compounds in Energetic Systems|year = 2004|journal = [[Propellants, Explosives, Pyrotechnics]]|volume = 29|issue = 2|pages = 67–80|doi = 10.1002/prep.200400032}}</ref>
*Nuclear applications:
**[[Lithium deuteride]] was the [[nuclear fusion|fusion fuel]] of choice in early versions of the [[Nuclear weapon|hydrogen bomb]]. When bombarded by [[neutron]]s, both <sup>6</sup>Li and <sup>7</sup>Li produce [[tritium]]—this reaction, which was not fully understood when [[Teller-Ulam design|hydrogen bombs]] were first tested, was responsible for the runaway yield of the [[Castle Bravo]] [[nuclear test]]. Tritium fuses with [[deuterium]] in a [[Nuclear fusion|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.
**[[Lithium fluoride]] (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred fluoride salt mixture (LiF-[[beryllium fluoride|BeF<sub>2</sub>]]) used in [[molten salt reactor|liquid-fluoride nuclear reactors]]. Lithium fluoride is exceptionally chemically stable and LiF/BeF<sub>2</sub> mixtures have low melting points and the best neutronic properties of fluoride salt combinations appropriate for reactor use.{{Clarify|date=April 2009}}<!-- what neutronic properties are good --> Finely divided lithium fluoride powder has been used for [[Thermoluminescent Dosimeter|thermoluminescent radiation dosimetry]] (TLD): When a sample of such is exposed to radiation, it accumulates [[crystal defect]]s which, when heated, resolve via a release of bluish light whose intensity is proportional to the [[absorbed dose]], thus allowing this to be quantified (see Eugene Tochilin, et al.).
**In conceptualized nuclear [[fusion power]] plants, lithium will be used to produce tritium in [[Magnetic confinement fusion|magnetically confined reactors]] using [[deuterium]] and [[tritium]] as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting [[Plasma (physics)|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. <sup>6</sup>Li + n → <sup>4</sup>He + <sup>3</sup>T. Various means of doing this will be tested at the [[ITER]] reactor being built at [[Cadarache]], France.
**Lithium is used as a source for [[alpha particle]]s, or [[helium]] nuclei. When <sup>7</sup>Li is bombarded by accelerated [[proton]]s <sup>8</sup>[[beryllium|Be]] is formed, which undergoes [[spontaneous fission]] to form two alpha particles. This was the first man-made [[nuclear reaction]], produced by [[John Douglas Cockcroft|Cockroft]] and [[Ernest Walton|Walton]] in 1929.
*Other uses:
**[[Lithium hydroxide]] (LiOH) is an important compound of lithium obtained from [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>). It is a strong [[base (chemitry)|base]], and when heated with a [[fat]] it produces a [[lithium soap]]. Lithium soap has the ability to [[thickening agent|thicken]] [[oil]]s, and it is used to manufacture [[grease (lubricant)|lubricating greases]].
**Lithium hydroxide and [[lithium peroxide]] are used in confined areas, such as aboard [[spacecraft]] and [[submarine]]s, 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 (Li<sub>2</sub>O<sub>2</sub>) in presence of moisture not only absorbs carbon dioxide to form lithium carbonate, but also releases oxygen. For example 2 Li<sub>2</sub>O<sub>2</sub> + 2 CO<sub>2</sub> → 2 Li<sub>2</sub>CO<sub>3</sub> + O<sub>2</sub>.
**Lithium compounds are used as [[pyrotechnic colorant]]s and oxidizers in red [[fireworks]] and [[Flare (pyrotechnic)|flares]].
**The [[Mark 50 torpedo|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]].<ref>{{cite journal|title=Stored Chemical Energy Propulsion System for Underwater Applications| author=Hughes, T.G.; Smith, R.B. and Kiely, D.H. |journal= Journal of Energy|year= 1983|volume=7|issue=2 |pages=128–133|doi=10.2514/3.62644}}</ref>

==Precautions==
[[Image:Limetal.JPG|thumb|left|Lithium ingots with a thin layer of black oxide tarnish]]
Lithium metal is [[corrosive]] and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) initially [[irritation|irritate]] the [[nose]] and [[throat]], while higher exposure can cause a buildup of fluid in the [[lung]]s, 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]].<ref>{{cite book|url = http://books.google.com/books?id=Oo3xAmmMlEwC&pg=PA244|pages = 244–246|isbn = 9780849325236|author = Furr, A. K.|year = 2000|publisher = CRC Press|location = Boca Raton|title = CRC handbook of laboratory safety}}</ref>

<div style="float: right;">
{| class="wikitable"
! style="background:#FF9900" | NFPA 704
|-
| align="left"|{{NFPA 704 | Health = 3 | Flammability = 2 | Reactivity = 2 | Other = <s>W</s>}}
|-
| width="80pt"|Fire diamond for lithium metal
|}
</div>

===Regulation===
Some jurisdictions limit the sale of [[lithium battery|lithium batteries]], which are the most readily available source of lithium metal for ordinary consumers. Lithium can be used to reduce [[pseudoephedrine]] and [[ephedrine]] to [[methamphetamine]] in the [[Birch reduction]] method, which employs solutions of alkali metals dissolved in anhydrous [[ammonia]].

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 circuit|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.<ref>{{cite book|isbn = 9780306447587|pages = 15–16|url = http://books.google.com/books?id=i7U-0IB8tjMC&pg=PA15|author = Samuel C. Levy and Per Bro.|year = 1994|publisher = Plenum Press|location = New York|title = Battery hazards and accident prevention}}</ref>

==See also==
{{Wikipedia-Books|Lithium}}
*[[:Category:Lithium compounds|Lithium compounds]]
*[[Lithium-based grease]]
*[[Dilithium]]
*[[Lithium-ion battery]]

==References==
{{Reflist|colwidth=30em}}

==External links==
{{Wiktionary|lithium}}
{{Commons|Lithium}}
*[http://www.lithiumalliance.org/ International Lithium Alliance]
*[http://minerals.usgs.gov/minerals/pubs/commodity/lithium/ USGS: Lithium Statistics and Information]
*[http://trugroup.com/whitepapers/TRU-Lithium-Outlook-2020.pdf Lithium Supply & Markets 2009 IM Conference 2009 Sustainable lithium supplies through 2020 in the face of sustainable market growth]
*[http://www.webelements.com/lithium/ WebElements.com&nbsp;– Lithium]
*[http://education.jlab.org/itselemental/ele003.html It's Elemental&nbsp;– Lithium]
*[http://www.mcis.soton.ac.uk/Site_Files/pdf/nuclear_history/Working_Paper_No_5.pdf University of Southampton, Mountbatten Centre for International Studies, Nuclear History Working Paper No5.]

{{Compact periodic table}}
{{Lithium compounds}}
{{alkalimetals}}

[[Category:Lithium| ]]
[[Category:Alkali metals]]
[[Category:Chemical elements]]
[[Category:Reducing agents]]

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