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{{Infobox lithium}}
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==Overview==
Lithium is a [[chemical element]] with [[Chemical symbol|symbol]] Li and [[atomic number]] 3.  It is a soft, silver-white [[metal]] belonging to the [[alkali metal]] [[Group (periodic table)|group]] of chemical elements.  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, it is typically stored in [[mineral oil]].  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.  Because of its high [[reactivity (chemistry)|reactivity]], lithium never occurs freely in nature, and instead, only appears in [[chemical compound|compounds]], which are usually [[ionic compound|ionic]].  Lithium occurs in a number of [[pegmatite|pegmatitic]] minerals, but due to its solubility as an ion is present in ocean water and is commonly obtained from [[brine]]s and [[clay]]s.  On a commercial scale, lithium is isolated [[electrolysis|electrolytically]] from a mixture of [[lithium chloride]] and [[potassium chloride]].


The [[Atomic nucleus|nuclei]] of lithium verge on instability, since the two stable lithium [[isotope]]s found in nature have among the lowest [[binding energies]] per [[nucleon]] of all stable [[nuclide]]s.  Because of its relative nuclear 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.<ref name="Lodders2003">Numerical data from: {{cite doi | 10.1086/375492 }} Graphed at [[:File:SolarSystemAbundances.jpg]]</ref> For related reasons, lithium has important links to [[nuclear physics]].  The [[Nuclear transmutation|transmutation]] of lithium atoms to [[helium]] in 1932 was the first fully man-made [[nuclear reaction]], and [[lithium deuteride|lithium-6 deuteride]] serves as a [[nuclear fusion|fusion]] fuel in [[Teller-Ulam design|staged thermonuclear weapons]].<ref>[http://www.fas.org/nuke/intro/nuke/design.htm Nuclear Weapon Design].  Federation of American Scientists (1998-10-21).  fas.org</ref>


'''Lithium''' ({{pronEng|ˈlɪθiəm}}) is a [[chemical element]] with the symbol '''Li''' and [[atomic number]] 3.  It is a soft [[alkali metal]] with a silver-white color. Under [[standard conditions 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, [[corrosion|corroding]] quickly in moist [[air]] to form a black tarnish. For this reason, lithium metal is typically stored under the cover of [[oil]].
Lithium and its compounds have several industrial applications, including heat-resistant glass and [[ceramic]]s, high strength-to-weight [[alloy]]s used in aircraft, [[lithium battery|lithium batteries]] and [[lithium-ion battery|lithium-ion batteries]]. These uses consume more than half of lithium production.


According to theory, Lithium (mostly <sup>7</sup>Li) was one of the few elements synthesized in the [[Big Bang]], although its quantity has vastly decreased. The reasons for its disappearance and the processes by which new lithium is created continue to be important matters of study in [[astronomy]]. Lithium is the 33rd most abundant element on [[Earth]],<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 | pages = 47-50}}</ref> but due to its high [[reactivity]] 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 all organisms. The element serves no apparent vital biological function, since animals and plants survive in good health without it.  Non-vital functions have not been ruled out. 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 in the treatment of [[bipolar disorder]], due to neurological [[lithium (medication)|effects]] of the ion in the human body.


Trace amounts of lithium are present in the [[ocean]]s and in some organisms, though the element serves no apparent biological function in humans. Nevertheless, the neurological effect of the lithium ion Li<sup>+</sup> makes some lithium [[salt (chemistry)|salt]]s useful as a class of [[mood stabilizer|mood stabilizing]] drugs. Lithium and its compounds have several other commercial 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 fission|splitting]] of lithium atoms was the first man-made form of [[nuclear reaction]], and [[lithium deuteride]] serves as the [[nuclear fusion|fusion]] fuel in [[Teller-Ulam design|staged thermonuclear weapon]]s.
===Atomic and Physical===
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 heat and electricity as well as a highly reactive element, though the least reactive of the alkali metals. Lithium's low reactivity compared to other alkali metals is due to the proximity of its valence electron to its [[atomic nucleus|nucleus]] (the remaining two electrons are in lithium's [[s-orbital|1s orbital]] and are much lower in energy, and therefore they do not participate in chemical bonds).<ref name=krebs/>


== History and etymology==
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.<ref name=krebs/> While it has one of the [[List of elements by melting point|lowest melting points]] among all metals (180 °C), it has the highest melting and boiling points of the alkali metals.<ref>{{RubberBible86th}}</ref>
[[Petalite]] (lithium aluminum silicate) was first described in 1800 by the Brazilian scientist [[José Bonifácio de Andrade e Silva]], who discovered the mineral in a [[Sweden|Swedish]] [[iron mine]] on the island of [[Utö, Sweden|Utö]]. However, it was not until 1817 that [[Johan August Arfwedson]], then a trainee in the laboratory of [[Jöns Jakob Berzelius]], [[discovery of the chemical elements|discovered]] the presence of a new element while analyzing petalite ore. The element formed compounds similar to those of [[sodium]] and [[potassium]], though its [[carbonate]] and [[hydroxide]] were less [[solubility|water soluble]] and had a larger capacity to neutralize acid. Berzelius gave the alkaline material the name "lithion", from the [[Greek language|Greek]] ''λιθoς'' (''lithos'', "stone"), to reflect its discovery in a mineral, as opposed to sodium and potassium which had been discovered in [[plant]] tissue; its name would later be standardized as "lithium". Arfwedson later showed that this same element was present in the mineral ores [[spodumene]] and [[lepidolite]]. In [[1818]], [[Christian Gmelin]] was the first to observe that lithium salts give a bright red color in flame. However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.<ref name=we-hist>{{cite web |last= Winter |first= Mark J|url=http://www.webelements.com/webelements/elements/text/Li/hist.html |title=Chemistry : Periodic Table : lithium : historical information |accessmonthday=August 19 |accessyear=2007 | publisher=Web Elements}}</ref><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}}</ref><ref>http://www.vanderkrogt.net/elements/elem/li.html</ref>


The element was not isolated until 1821, when [[William Thomas Brande]] performed [[electrolysis]] on [[lithium oxide]], a process which had previously been employed by [[Sir Humphry Davy]] to isolate potassium and sodium.<ref name=eote/><ref>http://www.diracdelta.co.uk/science/source/t/i/timeline/source.html </ref>Brande also described pure salts of lithium, such as the chloride, and performed an estimate of its atomic weight. In 1855, [[Robert Bunsen]] and [[Augustus Matthiessen]] produced large quantities of the metal by electrolysis of [[lithium chloride]]. Commercial production of lithium metal began in [[1923]] by the German company [[Metallgesellschaft AG]] through the electrolysis of a molten mixture of lithium chloride and [[potassium chloride]].<ref name=we-hist/><ref> http://www.echeat.com/essay.php?t=29195 </ref>
Lithium has a very low density of 0.534 g/cm<sup>3</sup>, comparable with that of pine wood.  It is the least dense of all elements that are solids at room temperature, the next lightest solid element (potassium, at 0.862 g/cm<sup>3</sup>) being more than 60% denser.  Furthermore, apart from [[helium]] and [[hydrogen]], it is less dense than any liquid element, being only 2/3 as dense as [[liquid nitrogen]] (0.808 g/cm<sup>3</sup>).<ref group=note>Densities for all the gaseous elements can be obtained at Airliquide.com</ref><ref>{{cite web|url=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 |title=Nitrogen, N2, Physical properties, safety, MSDS, enthalpy, material compatibility, gas liquid equilibrium, density, viscosity, inflammability, transport properties |publisher=Encyclopedia.airliquide.com |accessdate=2010-09-29}}</ref> Lithium  can float on the lightest hydrocarbon oils and is one of only three metals that can float on water, the other two being [[sodium]] and [[potassium]].


== Properties ==
Lithium's [[coefficient of thermal expansion]] is twice that of [[aluminium]] and almost four times that of [[iron]].<ref>{{cite web|url=http://www.engineeringtoolbox.com/linear-expansion-coefficients-d_95.html|title = Coefficients of Linear Expansion|publisher = Engineering Toolbox}}</ref> It has 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|bibcode = 2007Natur.447..187T }}</ref> and at higher temperatures (more than 9 K) at very high pressures (>20 GPa)<ref>{{Cite journal|doi=10.1126/science.1078535|year=2002|author=Struzhkin, V.  V.; Eremets, M.  I.; Gan, W; Mao, H.  K.; Hemley, R.  J.|title=Superconductivity in dense lithium|volume=298|issue=5596|pages=1213–5|pmid=12386338|journal=Science|bibcode = 2002Sci...298.1213S }}</ref> At temperatures below 70 K, lithium, like sodium, undergoes [[diffusionless transformations|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.<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|bibcode=1984PhRvL..53...64O}}</ref> Multiple allotropic forms have been reported for lithium at high pressures.<ref>{{cite journal|last1=Schwarz|first1=Ulrich|title=Metallic high-pressure modifications of main group elements|journal=Zeitschrift für Kristallographie|volume=219|page=376|year=2004|doi=10.1524/zkri.219.6.376.34637|issue=6–2004|bibcode = 2004ZK....219..376S }}</ref>
[[Image:LithiumPelletsUSGOV.jpg|thumb|left|Lithium pellets (covered in white lithium hydroxide)]]


Like other [[alkali metal]]s, lithium has a single [[valence electron]] which it will readily lose to form a [[cation]], indicated by the element's low [[electronegativity]]. As a result, lithium is easily deformed, highly reactive, and has lower [[melting point|melting]] and [[boiling point]]s than most metals. These and many other properties attributable to alkali metals' weakly-held valence electron are most distinguished in lithium, as it possesses the smallest [[atomic radius]] and thus the highest electronegativity of the alkali group. In addition, lithium has a [[diagonal relationship]] with [[magnesium]], an element of similar atomic and [[ionic radius]]. Chemical resemblances between the two metals include the formation of a [[nitride]] in N<sub>2</sub>, the formation of an [[oxide]] when burnt in O<sub>2</sub>, [[salt (chemistry)|salts]] with similar [[solubility|solubilities]], and thermally-unstable [[carbonate]]s and nitrides.<ref name=kamienski>Kamienski et al. "Lithium and lithium compounds". ''Kirk-Othmer Encyclopedia of Chemical Technology''. John Wiley & Sons, Inc. Published online '''2004'''. {{doi|10.1002/0471238961.1209200811011309.a01.pub2}}</ref>
Lithium has a specific heat capacity of 3.58 kilojoules per kilogram-Kelvin, the highest of all solids.<ref name= CRC/><ref>http://hilltop.bradley.edu/~spost/THERMO/solidcp.pdf</ref> Because of this, lithium metal is often used in [[coolant]]s for [[heat transfer]] applications.<ref name=CRC/>


Lithium is soft enough to be cut with a knife, though this is more difficult than cutting sodium. The fresh metal has a silvery-white color which only remains untarnished in dry air.<ref name=kamienski/> Lithium has about half the density of water, giving solid sticks of lithium metal the odd heft of a light-to-medium wood like [[pine]]. The metal floats highly in [[hydrocarbon]]s; in the laboratory, jars of lithium are typically composed of black-coated sticks held down in hydrocarbon mechanically by the jar's lid and other sticks.
===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.<ref name=krebs/> Because of its reactivity with water, lithium is usually stored under cover of a hydrocarbon, often [[petroleum jelly]]. Though the heavier alkali metals can be stored in more dense substances, such as [[mineral oil]], lithium is not dense enough to be fully submerged in these liquids.<ref name=emsley/> In moist air, lithium 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" />


Lithium is greatly heat-resistant, possessing a low [[coefficient of thermal expansion]] and the highest [[specific heat capacity]] of any solid element. Lithium has also been found to be [[superconductive]] below 400 [[microkelvin|μK]]. This finding paves the way for further study of superconductivity, as lithium's [[atomic lattice]] is the simplest of all metals.
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.<ref>{{cite journal|doi=10.1039/QJ8611300270|title=XXIV.—On chemical analysis by spectrum-observations|year=1861|journal=Quarterly Journal of the Chemical Society of London|volume=13|issue=3|page=270 }}</ref> 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 nonviolent, as the hydrogen produced will not ignite on its own.  As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically the Class D type (see [[Fire extinguisher#Types of extinguishing agents|Types of extinguishing agents]]). Lithium is the only metal which reacts with [[nitrogen]] under [[Standard conditions for temperature and pressure|normal conditions]].<ref>{{cite book|page=47|url=http://books.google.com/books?id=yb9xTj72vNAC&pg=PA47|title=The history and use of our earth's chemical elements: a reference guide|author=Krebs, Robert E.|publisher=Greenwood Publishing Group|year=2006|isbn=0-313-33438-2}}</ref><ref>{{Cite journal|author1=Institute, American Geological|author2=Union, American Geophysical|author3=Society, Geochemical|title=Geochemistry international|volume =31|issue=1–4|page=115|date=1 January 1994|url=http://books.google.com/books?id=77McAQAAIAAJ}}</ref>


== Chemistry ==
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 [[lithium oxide|oxide]] ({{chem|Li|2|O}}) and peroxide ({{chem|Li|2|O|2}}) 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><ref name = "Greenwood">{{Greenwood&Earnshaw1st|pages=97–99}}</ref> The metal reacts with hydrogen gas at high temperatures to produce [[lithium hydride]] (LiH).<ref>{{cite web
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/>
|url=http://www.lyon.edu/webdata/users/fbeckford/CHM%20120/Lecture%20Notes/Chapter-14.ppt
|title=University of Lyon course online (powerpoint) slideshow
|accessdate=2008-07-27
|author=Beckford, Floyd
|quote=definitions:Slides 8–10 (Chapter 14) |archiveurl =http://web.archive.org/web/20051104025202/http://www.lyon.edu/webdata/users/fbeckford/CHM+120/Lecture+Notes/Chapter-14.ppt |archivedate = 4 November 2005}}</ref>


When placed over a flame, lithium gives off a striking [[crimson]] color, but when it burns strongly, the flame becomes a brilliant white. Lithium will ignite and burn in oxygen when exposed to water or water vapours. It is the only metal that reacts with nitrogen at room temperature.  
Other known [[binary compound]]s include the [[halide]]s ([[lithium fluoride|LiF]], [[lithium chloride|LiCl]], [[lithium bromide|LiBr]], [[Lithium iodide|LiI]]), and the [[sulfide]] ([[lithium sulfide|{{chem|Li|2|S}}]]), the [[superoxide]] ([[Lithium superoxide|{{chem|LiO|2}}]]), [[carbide]] ([[Lithium carbide|{{chem|Li|2|C|2}}]]). Many other inorganic compounds are known, where lithium combines with [[anion]]s to form various salts: [[Lithium borate|borates]], [[Lithium amide|amides]], [[Lithium carbonate|carbonate]], [[Lithium nitrate|nitrate]], or [[borohydride]] ([[Lithium borohydride|{{chem|LiBH|4}}]]). Multiple [[organolithium reagent]]s are known where there is a direct [[covalent bond|bond]] between [[carbon]] and lithium atoms effectively creating a [[carbanion]].  These are extremely powerful [[base (chemistry)|bases]] and [[carbon nucleophile|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.<ref>{{Cite book|pages=3–40|url=http://books.google.com/books?id=z76sVepirh4C&pg=PA16|author=Sapse, Anne-Marie and von R.  Schleyer, Paul |title=Lithium chemistry: a theoretical and experimental overview|publisher=Wiley-IEEE|year=1995|isbn=0-471-54930-4}}</ref> LiHe, a very weakly interacting van der Waals compound, has been detected at very low temperatures.<ref>{{Cite journal|author1=Bretislav Friedrich|title=APS Physics|volume =6|issue=|page=42|date=8 April 2013|url=http://physics.aps.org/articles/v6/42}}</ref>


Lithium metal is flammable and potentially explosive when exposed to air and especially water, though it is far less dangerous than other alkali metals in this regard. The lithium-water reaction at normal temperatures is brisk but not violent. Lithium fires are difficult to extinguish, requiring special chemicals designed to smother them (see [[sodium]] for details).
===Isotopes===
{{seealso|Lithium compounds}}
{{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]].  The two lithium nuclei have lower binding energies per nucleon than any other stable nuclides other than [[deuterium]] and [[helium-3]].  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.<ref name="Lodders2003" />


== Isotopes ==
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&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 7.6 × 10<sup>−23</sup> s.<ref name=nuclidetable>{{cite web|url=http://www.nndc.bnl.gov/chart/reCenter.jsp?z=104&n=158|title=Interactive Chart of Nuclides|publisher=Brookhaven National Laboratory|author=Sonzogni, Alejandro|location=National Nuclear Data Center|accessdate=2008-06-06}}</ref>
{{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]]). 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 7.58043x10<sup>-23</sup> s.
<sup>7</sup>Li is one of the [[primordial elements]] (or, more properly, primordial [[nuclide]]s) 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 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.|bibcode=2006ApJ...644..229A|arxiv = astro-ph/0510636|display-authors=1|last2=Lambert|first2=David L.|last3=Nissen|first3=Poul Erik|last4=Primas|first4=Francesca|last5=Smith|first5=Verne V.  }}</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 hitting heavier atoms, and from 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 |issue=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|bibcode=2006GeCoA..70..224C}}</ref> While lithium is created in stars during the [[Stellar nucleosynthesis]], it is further [[Lithium burning|burnt]]. <sup>7</sup>Li can also be generated in [[carbon star]]s.<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|arxiv = astro-ph/0005356 }}</ref>


<sup>7</sup>Li is one of the [[primordial elements]] or, more properly, primordial isotopes, produced in [[Big Bang nucleosynthesis]] (a small amount of <sup>6</sup>Li is also produced in stars). Lithium isotopes fractionate substantially during a wide variety of natural processes, including [[mineral formation]] (chemical precipitation), [[metabolism]], and [[ion exchange]]. Lithium ion substitutes for [[magnesium]] and [[iron]] in octahedral sites in [[clay]] minerals, where <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.
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 |issue=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]].  The process known as [[atomic vapor laser isotope separation|laser isotope separation]] can be used to separate lithium isotopes.<ref>{{Cite book| page=330| url=http://www.opticsjournal.com/tla.htm| title = Tunable Laser Applications| author = [[F.  J.  Duarte|Duarte, F.  J]]| publisher= CRC Press| year = 2009| isbn=1-4200-6009-0}}</ref>


The exotic <sup>11</sup>Li is known to exhibit a [[nuclear halo]].
Nuclear weapons manufacture and other nuclear physics uses are a major source of artificial lithium fractionation, with the light isotope <sup>6</sup>Li being retained by industry and military stockpiles to such an extent as to slightly but measurably change the <sup>6</sup>Li to <sup>7</sup>Li ratios even in natural sources, such as rivers.  This has led to unusual uncertainty in the standardized [[atomic weight]] of lithium, since this quantity depends on the natural abundance ratios of these naturally-occurring stable lithium isotopes, as they are available in commercial lithium mineral sources.<ref name = "Coplen2002">{{cite journal|doi=10.1351/pac200274101987|title=Isotope-abundance variations of selected elements (IUPAC Technical Report)|year=2002|last1=Coplen|first1=T.B.|last2=Bohlke|first2=J.K.|last3=De Bievre|first3=P.|last4=Ding|first4=T.|last5=Holden|first5=N.E.|last6=Hopple|first6=J.A.|last7=Krouse|first7=H.R.|last8=Lamberty|first8=A.|last9=Peiser|first9=H.S.|last10=N.N.|journal=Pure and Applied Chemistry|volume=74|issue=10|page=1987|display-authors=9}}</ref>


== Natural occurrence ==
===Astronomical===
''See also [[:category:Lithium minerals|Lithium minerals]].''
{{Main|Nucleosynthesis}}
According to modern cosmological theory, lithium—as both of its stable isotopes lithium-6 and lithium-7—was among the 3 elements [[nucleosynthesis|synthesized]] in the [[Big Bang]].<ref>{{cite journal | bibcode= 1985ARA&A..23..319B] | title= Big bang nucleosynthesis – Theories and observations | author1= Boesgaard | first1= A.  M.  | last2= Steigman | first2= G.  | volume= 23 | year= 1985 | page= 319 | journal= IN: Annual review of astronomy and astrophysics.  Volume 23 (A86-14507 04–90).  Palo Alto | doi= 10.1146/annurev.aa.23.090185.001535}}</ref> Though the amount of lithium generated in [[Big Bang nucleosynthesis]] is dependent upon the number of [[photon]]s 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/476/why-old-stars-seem-to-lack-lithium/ |title=Why Old Stars Seem to Lack Lithium|date= 2006-08-16|author= Cain, Fraser }}</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/>


Lithium is widely distributed on Earth and is the 33rd most abundant element;<ref name=krebs/> however, it does not naturally occur in elemental form due to its high reactivity. 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]] rocks, 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 mineral sources for the element.<ref name=kamienski/>
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 of the low temperature necessary to destroy lithium, and a lack of common processes to produce it.<ref name=wesleyan>{{cite web|url=http://www.astro.wesleyan.edu/~bill/courses/astr231/wes_only/element_abundances.pdf |archiveurl=http://web.archive.org/web/20060901133923/http://www.astro.wesleyan.edu/~bill/courses/astr231/wes_only/element_abundances.pdf |archivedate=1 September 2006 |title=Element Abundances |accessdate=2009-11-17}}</ref>


== Applications ==
Lithium is also found in [[brown dwarf]] substellar objects and certain anomalous orange stars.  Because lithium is present in cooler, less-massive brown dwarfs, 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/24593/brown-dwarf/|archiveurl=http://web.archive.org/web/20110225032434/http://www.universetoday.com/24593/brown-dwarf/|archivedate=2011-02-25|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=2013-03-06 | first =Neill | last = Reid | date = 2002-03-10}}</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—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.<ref name=emsley/>
Because of its [[specific heat]] capacity, the highest of all [[solid]]s, lithium is often used in heat transfer applications.


It is an important ingredient in [[cathode]] materials, used in rechargeable and single-use [[Lithium battery|batteries]] because of its high [[electrochemical potential]], light weight, and high current density.
===Terrestrial===
Although lithium is widely distributed on Earth, it 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> 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),<ref>{{cite web|url=http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html |archiveurl=http://web.archive.org/web/20090502142924/http://www.ioes.saga-u.ac.jp/ioes-study/li/lithium/occurence.html |archivedate=2009-05-02 |title=Lithium Occurrence|accessdate=2009-03-13|publisher=Institute of Ocean Energy, Saga University, Japan}}</ref><ref name=enc/> or 25 [[micromolar]];<ref>{{cite web|url=http://www.springerlink.com/content/y621101m3567jku1/ |title=Extraction of metals from sea water|year=1984|publisher=Springer Berlin Heidelberg}}</ref>
higher concentrations approaching 7 ppm are found near [[hydrothermal vents]].<ref name=enc/>


Large quantities of lithium are also used in the manufacture of [[organolithium reagent]]s, especially [[n-butyllithium|''n''-butyllithium]] which has many uses in fine chemical and [[polymer]] synthesis.
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" /> Another significant mineral of lithium is [[lepidolite]].<ref>{{cite book|title=Shriver & Atkins' Inorganic Chemistry|edition=5|publisher=W.  H.  Freeman and Company|place= New York|year= 2010|page=296|isbn=0199236178|author=Atkins, Peter }}</ref> 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> 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.


=== Medical use ===
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 few of them are of actual or potential commercial value.  Many are very small, others are too low in grade."<ref>Garrett, Donald (2004) ''Handbook of Lithium and Natural Calcium'', Academic Press, cited in ''[http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf The Trouble with Lithium 2]'', Meridian International Research (2008)</ref>
{{main|Lithium pharmacology}}


Lithium salts were used during the 19th century to treat [[gout]].
One of the largest ''reserve base''<ref group=note name=res>[http://minerals.usgs.gov/minerals/pubs/mcs/2011/mcsapp2011.pdf Apendixes]. By USGS definitions, reserve base "may encompass those parts of the resources that have a reasonable potential for becoming economically available within planning horizons beyond those that assume proven technology and current economics.  The reserve base includes those resources that are currently economic (reserves), marginally economic (marginal reserves), and some of those that are currently subeconomic (subeconomic resources)."</ref> of lithium is in the [[Salar de Uyuni]] area of [[Bolivia]], which has 5.4 million tonnes.  [[US Geological Survey]], estimates that in 2010 Chile had the largest reserves by far (7.5 million tonnes)<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> and the highest annual production (8,800 tonnes). Other major suppliers include Australia, Argentina and China.
Lithium salts such as [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>), [[lithium citrate]], and [[lithium orotate]] are mood stabilizers. They are used in the treatment of [[bipolar disorder]], since unlike most other mood altering drugs, they counteract both [[mania]] and [[depression (mood)|depression]]. Lithium can also be used to augment other [[antidepressant]] drugs. It is also sometimes prescribed as a preventive treatment for [[migraine]] disease and [[cluster headache]]s.


The active principle in these salts is the lithium ion Li<sup>+</sup>, which having a smaller diameter, can easily displace K<sup>+</sup> and Na<sup>+</sup> and even Ca<sup>+2</sup>, in spite of its greater charge, occupying their sites in several critical neuronal enzymes and neurotransmitter receptors. Although Li<sup>+</sup> cannot displace Mg<sup>2+</sup> and Zn<sup>2+</sup>, because of these ions small size and greater charge (higher charge density, hence stronger bonding), when Mg<sup>+2</sup> or Zn<sup>+2</sup> are present in low concentrations, and Li<sup>+</sup> is present in high concentrations, the latter can occupy sites normally occupied by Mg<sup>+2</sup> or Zn<sup>+2</sup> in various enzymes. Therapeutically useful amounts of lithium (0.6 to 1.2 mmol/l) are only slightly lower than toxic amounts (>1.5 mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity. Therefore, in theory, coadministration of 400 IU vitamin D, 1 g magnesium citrate (not the insoluble oxide or carbonate), 15 mg Zn (as gluconate or piccolinate, not the insoluble oxide) and 1 pill of [[vitamin B]] complex a day, should potentiate the effect of Li, <!--"in theory" is used here very loosely indeed! WHAT theory??--> in some cases allowing for the reduction of the therapeutic range to 0.5 to 0.9 mmol/l, of the daily dose of lithium carbonate and of the risk of toxicity.
In June 2010, the [[New York Times]] reported that American geologists were conducting ground surveys on [[Dry lake|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 as those of Bolivia, which now has the world's largest known lithium reserves."<ref>{{cite news|url=http://www.nytimes.com/2010/06/14/world/asia/14minerals.html?pagewanted=1&hp|title=U.S.  Identifies Vast Riches of Minerals in Afghanistan|accessdate=2010-06-13|work=The New York Times|first=James|last=Risen|date=2010-06-13}}</ref> 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."<ref>{{cite news| url=http://business.timesonline.co.uk/tol/business/industry_sectors/natural_resources/article7149696.ece|location=London|work=The Times |title=Taleban zones mineral riches may rival Saudi Arabia says Pentagon|first1=Jeremy|last1=Page|first2=Michael|last2=Evans|date=2010-06-15}}</ref>


Common side effects include muscle tremors, twitching, [[ataxia]], hyperparathyroidism (bone loss, hypercalcemia, hypertension, etc,), kidney damage, [[nephrogenic diabetes insipidus]] (polyuria and polydipsia) and seizures. Many of the side-effects are a result caused by the increased elimination of potassium.
===Biological===
Lithium is found in trace amount in numerous plants, plankton, and invertebrates, at concentrations of 69 to 5,760 parts per billion (ppb).  In vertebrates the concentration is slightly lower, and nearly all vertebrate tissue and body fluids have been found to contain lithium ranging from 21 to 763 ppb.<ref name=enc/> Marine organisms tend to bioaccumulate lithium more than terrestrial ones.<ref>{{cite journal|last1=Chassard-Bouchaud |first1=C| last2=Galle|first2=P|last3=Escaig|first3=F|last4=Miyawaki|first4=M|title=Bioaccumulation of lithium by marine organisms in European, American, and Asian coastal zones: microanalytic study using secondary ion emission|journal=Comptes rendus de l'Academie des sciences.  Serie III, Sciences de la vie|volume=299|issue=18|pages=719–24|year=1984|pmid=6440674}}</ref> It is not known whether lithium has a physiological role in any of these organisms,<ref name=enc>{{cite web|url=http://www.enclabs.com/lithium.html|accessdate=2010-10-15|title=Some Facts about Lithium|publisher=ENC Labs}}</ref> but nutritional studies in mammals have indicated its importance to health, leading to a suggestion that it be classed as an essential trace element with an RDA of 1&nbsp;mg/day.<ref>{{Cite journal|pmid=11838882|year=2002|last1=Schrauzer|first1=GN|title=Lithium: Occurrence, dietary intakes, nutritional essentiality|volume=21|issue=1|pages=14–21|journal=Journal of the American College of Nutrition|doi=10.1080/07315724.2002.10719188}}</ref> Observational studies in Japan, reported in 2011, suggested that naturally occurring lithium in drinking water may increase human lifespan.<ref>{{cite journal|doi=10.1007/s00394-011-0171-x|title=Low-dose lithium uptake promotes longevity in humans and metazoans|journal=European Journal of Nutrition|year=2011|last1=Zarse|first1=Kim|last2=Terao|first2=Takeshi|last3=Tian|first3=Jing|last4=Iwata|first4=Noboru|last5=Ishii|first5=Nobuyoshi|last6=Ristow|first6=Michael|volume=50|issue=5|pages=387–9|pmid=21301855|pmc=3151375}}</ref>


=== Other uses ===
==History of discovery and use==
* [[Lithium chloride]] and [[lithium bromide]] are extremely [[hygroscopic]] and frequently used as [[desiccant]]s.
[[Petalite]] (LiAlSi<sub>4</sub>O<sub>10</sub>) was discovered in 1800 by the [[Brazil]]ian chemist and statesman [[José Bonifácio de Andrada e Silva]] 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=2009-08-10|work=Mindat.org}}</ref><ref name=webelementshistory>{{cite web|url=http://www.webelements.com/lithium/history.html|title=Lithium:Historical information |accessdate=2009-08-10}}</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=0-7661-3872-0 |url=http://books.google.com/?id=SJIk9BPdNWcC|accessdate=2009-08-10}}</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 |archiveurl=http://web.archive.org/web/20101007084500/http://www.chemeddl.org/collections/ptl/ptl/chemists/bios/arfwedson.html |archivedate=2010-10-07 |title=Johan August Arfwedson |accessdate=2009-08-10 |work= Periodic Table Live!}}</ref><ref name=uwis>{{cite web|url=http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm |archiveurl=http://web.archive.org/web/20080605152857/http://genchem.chem.wisc.edu/lab/PTL/PTL/BIOS/arfwdson.htm |archivedate=5 June 2008 |title=Johan Arfwedson |accessdate=2009-08-10}}</ref><ref name=vanderkrogt>{{cite web|publisher = Elementymology & Elements Multidict|title = Lithium| first = Peter|last =van der Krogt|url =http://elements.vanderkrogt.net/element.php?sym=Li|accessdate = 2010-10-05}}</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=2009-08-10 |last=Clark |first=Jim |year=2005}}</ref> Berzelius gave the alkaline material the name "''lithion''/''lithina''", from the [[Greek language|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 "''lithium''".<ref name=krebs/><ref name=webelementshistory/><ref name=vanderkrogt/>
* [[stearic acid|Lithium stearate]] is a common all-purpose high-temperature [[lubricant]].
* Lithium is an [[alloy]]ing agent used to synthesize [[organic compound]]s.
* Lithium is used as a [[flux (metallurgy)|flux]] to promote the fusing of metals during [[welding]] and [[soldering]]. It also eliminates the forming of oxides during welding by absorbing impurities. This fusing quality is also important as a flux for producing [[ceramic]]s, [[Vitreous enamel|enamels]], and [[glass]].
* Lithium is sometimes used in glasses and ceramics including the glass for the 200-inch (5.08&nbsp;m) [[telescope]] at [[Mt. Palomar]].
* [[Alloy]]s of the metal with [[aluminium]], [[cadmium]], [[copper]], and [[manganese]] are used to make high performance [[aircraft]] parts.
* [[Lithium niobate]] is used extensively in telecommunication products, such as [[mobile phone]]s and [[optical modulator]]s, for such components as resonant crystals. Lithium products are currently used in more than 60 percent of mobile phones.<ref>{{cite news |author=Spring, Martin |title=Two ways to play the lithium boom |url=http://www.moneyweek.com/file/32991/two-ways-to-play-the-lithium-boom.html |publisher=[[MoneyWeek]] |date=2007-01-08 |accessdate=2007-08-19}}</ref>
* The high non-linearity of lithium niobate also makes a good choice for [[nonlinear optics|non-linear optics 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]]&mdash;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 [[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.
* Metallic lithium and its complex [[hydride]]s such as e.g. Li[AlH<sub>4</sub>] are considered as high energy additives to [[rocket propellant]]s<sup>[3]</sup>.
* [[Lithium peroxide]], [[lithium nitrate]], [[lithium chlorate]] and [[lithium perchlorate]] are used and thought of as oxidizers in both rocket propellants and [[oxygen candle]]s to supply submarines and space capsules with oxygen.<ref>{{cite journal | author = K. Ernst-Christian | 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>  
* Lithium will be used to produce tritium in magnetically confined nuclear fusion reactors using deuterium and tritium as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium. <sup>6</sup>Li + n → <sup>4</sup>He + <sup>3</sup>H. 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 Cockroft and Walton in [[1929]].
* [[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, and when heated with a fat, it produces a lithium soap. Lithium soap has the ability to thicken oils and so is used commercially to manufacture lubricating greases. It is also an efficient and lightweight purifier of air. In confined areas, such as aboard [[spacecraft]] and [[submarine]]s, the concentration of carbon dioxide can approach unhealthy or toxic levels. Lithium hydroxide absorbs the carbon dioxide from the air by reacting with it to form lithium carbonate. Any alkali hydroxide will absorb CO<sub>2</sub>, but lithium hydroxide is preferred, especially in spacecraft applications, because of the low formula weight conferred by the lithium. Even better materials for this purpose include lithium peroxide (Li<sub>2</sub>O<sub>2</sub>) that, in presence of moisture, not only absorb carbon dioxide to form lithium carbonate, but also release oxygen. E.g. 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 metal is used as a [[reducing agent]] in some types of [[methamphetamine]] production, particularly in illegal amateur “meth labs.”


== Production ==
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 to observe that lithium salts give a bright red color to flame.<ref name=webelementshistory/> However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.<ref name=webelementshistory/><ref name=vanderkrogt/><ref name="eote">{{Cite book|year = 2004|title = Encyclopedia of the Elements: Technical Data&nbsp;– History&nbsp;– Processing&nbsp;– Applications|publisher = Wiley|isbn = 978-3-527-30666-4|pages = 287–300|author = Enghag, Per }}</ref> 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.<ref name=emsley>{{Cite book|last=Emsley |first=John |title=Nature's Building Blocks |publisher=Oxford University Press |location=Oxford|year=2001 |isbn=0-19-850341-5}}</ref><ref name="eote" /><ref>{{cite journal| publisher=Royal Institution of Great Britain|journal=The Quarterly Journal of Science and the Arts|volume=5| title=The Quarterly journal of science and the arts|year=1818|page=338|format=PDF|accessdate=2010-10-05|url=http://books.google.com/?id=D_4WAAAAYAAJ| author=Various authors}}</ref><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, 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).<ref>{{cite book|url=http://books.google.com/?id=zkIAAAAAYAAJ|first1=William Thomas|last1=Brande|first2=William James|last2=MacNeven|title=A manual of chemistry|year=1821|page=191|accessdate=2010-10-08|publisher=Long}}</ref> 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, 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/free-essay/Analysis-of-the-Element-Lithium-29195.aspx|title = Analysis of the Element Lithium|first = Thomas|last = Green|date =2006-06-11| publisher = echeat}}</ref><ref>{{cite book | url = http://books.google.de/books?id=Ua2SVcUBHZgC&pg=PA99 | page=99 | title = Handbook of Lithium and Natural Calcium Chloride | isbn = 9780080472904 | author1 = Garrett | first1 = Donald E | date = 2004-04-05}}</ref>
Since the end of [[World War II]], lithium metal production has greatly increased. The metal is separated from other elements in igneous mineral such as those above, and is also extracted from the water of [[mineral springs]].


The metal is produced [[electrolysis|electrolytically]] from a mixture of fused lithium and [[potassium chloride]]. In [[1998]] it was about [[US dollar|US$]] 43 per [[Pound (mass)|pound]] ($95 per [[kilogram|kg]]).<ref name=ober>{{cite web |url=http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450798.pdf |title=Lithium |accessmonthday=August 19 |accessyear=2007 |last=Ober |first=Joyce A |format=pdf |pages = 77-78| publisher=[[United States Geological Survey]]}}</ref>
The production and use of lithium underwent several drastic changes in history. The first major application of lithium was in high-temperature [[lithium grease]]s for aircraft engines or similar applications in [[World War II]] and shortly after.  This use was supported by the fact that lithium-based [[soap]]s have a higher melting point than other alkali soaps, and are less corrosive than calcium based soaps. The small market for lithium soaps and the lubricating greases based upon them was supported by several small mining operations mostly in the United States.


[[Chile]] is currently the leading lithium metal producer in the world, with [[Argentina]] next. Both countries recover the lithium from brine pools. In the [[United States]] lithium is similarly recovered from brine pools in [[Nevada]].<ref name=lanl>{{cite web |url=http://periodic.lanl.gov/elements/3.html |title=Lithium |accessmonthday= August 19 |accessyear= 2007|date= December 15, 2003|publisher= [[Los Alamos National Laboratory]]}}</ref>
The demand for lithium increased dramatically during the [[Cold War]] with the production of [[Nuclear weapon design|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%, which was enough to affect the measured [[atomic weight]] of lithium in many standardized chemicals, and even the atomic weight of lithium in some "natural sources" of lithium ion which had been "contaminated" by lithium salts discharged from isotope separation facilities, which had found its way into ground water.<ref name = "Coplen2002"/><ref name="USGSCR1994">{{cite web| url = http://minerals.usgs.gov/minerals/pubs/commodity/lithium/450494.pdf| title = Commodity Report 1994: Lithium|publisher = United States Geological Survey|accessdate = 2010-11-03|year = 1994|first = Joyce A.|last = Ober}}</ref> <!--With only 7.5% of lithium-6 this makes ca.  2,200 tonnes of lithium-6.-->


China may emerge as a significant producer of brine-based lithium carbonate towards the end of this decade. Potential capacity of up to 45,000 tonnes per year could come on-stream if projects in Qinghai province and Tibet proceed.
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&nbsp;– Industrielle, medizinische und wissenschaftliche Bedeutung |year = 2003|last1 = Deberitz|first1 = Jürgen|first2 = Gernot|journal = Chemie in unserer Zeit|volume = 37|page = 258|last2 = Boche| issue = 4}}</ref><ref name="ciuz2003">{{Cite journal| doi = 10.1002/ciuz.19850190505|title = Lithium&nbsp;– wie es nicht im Lehrbuch steht|year = 1985|last1 = Bauer|first1 = Richard |journal = Chemie in unserer Zeit|volume = 19|page = 167| issue = 5}}</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"/> 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 21<sup>st</sup> century.


== Precautions ==
The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.<ref name="USGSYB1994">{{cite web| url = http://minerals.usgs.gov/minerals/pubs/commodity/lithium/myb1-2007-lithi.pdf|title = Minerals Yearbook 2007 : Lithium| publisher = United States Geological Survey|accessdate = 2010-11-03|year = 1994|first = Joyce A.|last = Ober}}</ref> With the surge of lithium demand in batteries in the 2000s, 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 = 978-0-87335-233-8|page = 599|url = http://books.google.com/?id=zNicdkuulE4C&pg=PA600&lpg=PAPA599|chapter = Lithium|year = 2006|publisher = Society for Mining, Metallurgy, and Exploration|location = Littleton, Colo.}}</ref><ref>{{Cite book|url=http://books.google.com/books?id=8erDL_DnsgAC&pg=PA339 |title=Encyclopedia of Chemical Processing and Design: Volume 28&nbsp;– Lactic Acid to Magnesium Supply-Demand Relationships|publisher=M.  Dekker|author=McKetta, John J.|date=18 July 2007 |accessdate=2010-09-29|isbn=978-0-8247-2478-8}}</ref>
Lithium metal, due to its alkaline tarnish, is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) can irritate the nose and throat; higher exposure to lithium can cause a build-up of fluid in the lungs, leading to [[pulmonary edema]]. The metal itself is usually less a handling hazard than the caustic hydroxide produced when it is in contact with moisture. Lithium should be stored in a non-reactive compound such as [[naphtha]] or a hydrocarbon.


=== Regulation ===
==Production==
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. However, the effectiveness of such restrictions in controlling illegal production of methamphetamine remains indeterminate and controversial.
Since the end of [[World War II]] lithium 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 55% [[lithium chloride]] and 45% [[potassium chloride]] at about 450<sup>o</sup> C.<ref>{{Greenwood&Earnshaw2nd|page=73}}</ref> In 1998 it was about {{nowrap|95 US$ / kg}} (or 43 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>


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]]. However, most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents.
Worldwide identified reserves of lithium in 2008 were estimated by the [[US Geological Survey]] as 13 million [[tonne]]s.  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]].<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 = 0-8493-0481-4}}</ref> However, half the world's known reserves are located in [[Bolivia]], a nation sitting along the central eastern slope of the Andes.
In 2009 Bolivia was negotiating with Japanese, French, and Korean firms to begin extraction.<ref name="romero">{{Cite news|author= Romero, Simon |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|work=New York Times |date=2009-02-02}}</ref> According to the [[US Geological Survey]], Bolivia's [[Uyuni]] Desert has 5.4 million tonnes of lithium.<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> A newly discovered deposit in [[Wyoming]]'s [[Rock Springs Uplift]] is estimated at 228,000 tons.  Additional deposits in the same formation were extrapolated to be as much as 18 million tons.<ref>{{cite web|author=Money Game Contributors |url=http://www.businessinsider.com/new-wyoming-lithium-deposit-could-meet-all-us-demand-2013-4 |title=New Wyoming Lithium Deposit |publisher=Business Insider |date=2013-04-26 |accessdate=2013-05-01}}</ref>


== References ==
After an industry wide pricing reduction for lithium carbonate after the Great Financial Crisis, where major suppliers such as [[Sociedad Química y Minera]] (SQM) dropped pricing by 20%<ref>{{cite web|url=http://www.prnewswire.com/news-releases/sqm-announces-new-lithium-prices-62933122.html |title=SQM Announces New Lithium Prices – SANTIAGO, Chile, Sept.  30 /PRNewswire-FirstCall/ |publisher=Prnewswire.com |date=2009-09-30 |accessdate=2013-05-01}}</ref> in light of incoming lithium resource developers and to further defend their market position, pricing in 2012 scaled up due to increased lithium demand.  A 2012 Business Week article outlined the existing oligopoly in the lithium space, "SQM, controlled by billionaire Julio Ponce, is the second-largest, followed by Rockwood, which is backed by Henry Kravis’s KKR & Co., and Philadelphia-based FMC." Global consumption may jump to 300,000 metric tons a year by 2020 from about 150,000 tons in 2012, as demand for lithium batteries has been growing at about 25 percent a year, outpacing the 4 percent to 5 percent overall gain in lithium<ref>{{cite web|last=Riseborough |first=Jesse |url=http://www.businessweek.com/news/2012-06-19/ipad-boom-strains-lithium-supplies-after-prices-triple |title=IPad Boom Strains Lithium Supplies After Prices Triple |publisher=Businessweek |date= |accessdate=2013-05-01}}</ref>
{{reflist|2}}


== External links ==
A potential source is [[Geothermal electricity|geothermal wells]].  Geothermal fluids carry leachates to the surface;<ref name="bourcier">Parker, Ann.  [https://www.llnl.gov/str/JanFeb05/Bourcier.html Mining Geothermal Resources].  Lawrence Livermore National Laboratory</ref> recovery of lithium has been demonstrated in the field.<ref name="Simbol">Patel, P.  (2011-11-16) [http://www.technologyreview.com/news/426131/startup-to-capture-lithium-from-geothermal-plants/   Startup to Capture Lithium from Geothermal Plants]. technologyreview.com</ref>  As the lithium is separated by simple filtration techniques, the process and environmental costs are primarily that of the already-operating geothermal well; relative environmental impacts may thus be positive.<ref name="NYT">Wald, M. (2011-09-28) [http://www.nytimes.com/2011/09/28/business/energy-environment/simbol-materials-plans-to-extract-lithium-from-geothermal-plants.html?_r=1 Start-Up in California Plans to Capture Lithium, and Market Share].  The New York Times</ref>
{{Wiktionary|lithium}}
{{Commons|Lithium}}
* [http://minerals.usgs.gov/minerals/pubs/commodity/lithium/ USGS: Lithium Statistics and Information]
* [http://www.webelements.com/webelements/elements/text/Li/index.html WebElements.com &ndash; Lithium]
* [http://education.jlab.org/itselemental/ele003.html It's Elemental &ndash; Lithium]
* [http://www.state.nj.us/health/eoh/rtkweb/1119.pdf Safety information on Lithium]
* [http://www.bipolar-lives.com/lithium.html Information on Lithium and Bipolar Disorder]


== See also ==
There are differing opinions about the potential growth of lithium production.  According to a 2011 study conducted at [[Lawrence Berkeley National Laboratory]] and the [[University of California Berkeley]], the currently estimated reserve base of lithium should not be a limiting factor for large-scale battery production for electric vehicles, as the study estimated that on the order of 1 billion 40 [[kWh]] Li-based batteries could be built with current reserves.<ref>{{cite web|url=http://www.greencarcongress.com/2011/06/albertus-20110617.html|title=Study finds resource constraints should not be a limiting factor for large-scale EV battery production|publisher=[[Green Car Congress]]|date=2011-06-17|accessdate=2011-06-17}}</ref> Another 2011 study by researchers from the [[University of Michigan]] and [[Ford Motor Company]] found that there are sufficient lithium resources to support global demand until 2100, including the lithium required for the potential widespread use of [[hybrid electric vehicle|hybrid electric]], [[plug-in hybrid|plug-in hybrid electric]] and [[battery electric vehicle]]s.  The study estimated global lithium reserves at 39 million tons, and total demand for lithium during the 90-year period analyzed at 12–20 million tons, depending on the scenarios regarding economic growth and recycling rates.<ref>{{cite web|url=http://www.greencarcongress.com/2011/08/lithium-20110803.html|title=University of Michigan and Ford researchers see plentiful lithium resources for electric vehicles|publisher=[[Green Car Congress]]|date=2011-08-03|accessdate=2011-08-11}}</ref>
* [[:Category:Lithium compounds|Lithium compounds]]


{{alkalimetals}}
However, another study concluded that "realistically achievable lithium carbonate production will be sufficient for only a small fraction of future [[PHEV]] and [[electric vehicle|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'.


[[Category:Chemical elements]]
===Ceramics and glass===
[[Category:Alkali metals]]
Lithium oxide is a widely used flux for processing [[silica]], reducing the [[melting point]] and [[viscosity]] of the material and leading to [[ceramic glaze|glazes]] of improved physical properties including low coefficients for thermal expansion.<ref>http://www.fmclithium.com/Portals/FMCLithiumFineChemicals/Content/Docs/Worldwide%20Demand%20by%20Sector.pdf</ref> Lithium oxides are a component of ovenware.  Worldwide, this is the single largest use for lithium compounds.  [[Lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>) is generally used in this application: upon heating it converts to the oxide.<ref>{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/compounds.html |title=Some Compounds of the Group 1 Elements |last1=Clark |first1=Jim |year=2005 |website=chemguide.co.uk |accessdate=2013-08-08}}</ref>
[[Category:Lithium|*]]
 
===Electrical and electronics===
In the later years of the 20th century, owing to its high [[electrochemical potential]], lithium became an important component of the electrolyte and of one of the electrodes in batteries.  A typical [[lithium-ion battery]] 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 batteries]] are [[disposable]] ([[primary cell|primary]]) [[Battery (electricity)|batteries]] with lithium or its compounds as an [[anode]].<ref>http://www.batteryreview.org/disposable-batteries.html</ref><ref>http://www.emc2.cornell.edu/content/view/battery-anodes.html</ref> 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]].
 
===Lubricating greases===
The third most common use of lithium is in greases.  Lithium hydroxide is a strong [[base (chemistry)|base]] and, when heated with a fat, produces a soap made of lithium [[stearate]].  Lithium soap has the ability to [[thickening agent|thicken]] oils, and it is used to manufacture all-purpose, high-temperature [[grease (lubricant)|lubricating greases]].<ref name=CRC/><ref>{{Cite book|url=http://books.google.com/books?id=J_AkNu-Y1wQC&pg=PA559|page=559|title=Fuels and lubricants handbook: technology, properties, performance, and testing, Volume 1|author=Totten, George E.; Westbrook, Steven R.  and Shah, Rajesh J.|publisher=ASTM International|year=2003|isbn=0-8031-2096-6}}</ref><ref>{{cite book|pages=150–152|url=http://books.google.com/books?id=3FkMrP4Hlw0C&pg=PA152|title=Significance of tests for petroleum products|author=Rand, Salvatore J.  |publisher=ASTM International|year= 2003|isbn=0-8031-2097-4}}</ref>
 
===Metallurgy===
When used as a [[flux (metallurgy)|flux]] for [[welding]] or [[soldering]], metallic lithium promotes the fusing of metals during the process 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]]).<ref>{{cite book|author1=Davis, Joseph R.  ASM International.  Handbook Committee|title=Aluminum and aluminum alloys|url=http://books.google.com/books?id=Lskj5k3PSIcC&pg=PA121|accessdate=2011-05-16|year=1993|publisher=ASM International|isbn=978-0-87170-496-2|pages=121–}}</ref>
 
====Pyrotechnics====
Lithium compounds are used as [[pyrotechnic colorant]]s and oxidizers in red [[fireworks]] and [[flare]]s.<ref name=CRC/><ref>Wiberg, Egon; Wiberg, Nils and Holleman, Arnold Frederick [http://books.google.com/books?id=Mtth5g59dEIC&pg=PA1089 Inorganic chemistry], Academic Press (2001) ISBN 0-12-352651-5, p.  1089</ref>
 
====Air purification====
[[Lithium chloride]] and [[lithium bromide]] are [[hygroscopic]] and are used as [[desiccant]]s for gas streams.<ref name=CRC/>  Lithium hydroxide and [[lithium peroxide]] are the salts most used in confined areas, such as aboard [[spacecraft]] and [[submarine]]s, for carbon dioxide removal and air purification.  Lithium hydroxide absorbs [[carbon dioxide]] from the air by forming 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 reacts with carbon dioxide to form lithium carbonate, but also releases oxygen.<ref>{{cite book|chapter=Air Quality Systems for Related Enclosed Spaces: Spacecraft Air|author=Mulloth, L.M.  and Finn, J.E.|title=The Handbook of Environmental Chemistry|year=2005|volume=4H|pages=383–404|doi=10.1007/b107253}}</ref><ref>{{cite web|url=http://www.dtic.mil/cgi-bin/GetTRDoc?Location=U2&doc=GetTRDoc.pdf&AD=AD0619497|title=Application of lithium chemicals for air regeneration of manned spacecraft|publisher=Lithium Corporation of America & Aeropspace Medical Research Laboratories|year=1965}}</ref> The reaction is as follows:
: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>.
Some of the aforementioned compounds, as well as [[lithium perchlorate]], are used in [[Chemical oxygen generator#Oxygen candle|oxygen candles]] that supply [[submarine]]s with [[oxygen]].  These can also include small amounts of [[boron]], [[magnesium]], [[aluminum]], [[silicon]], [[titanium]], [[manganese]], and [[iron]].<ref>http://pubs.acs.org/doi/abs/10.1021/i360012a016</ref>
 
====Optics====
[[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 one of the lowest [[refractive index]]es and the farthest transmission range in the deep UV of most common materials.<ref>{{Cite book|url=http://books.google.com/books?id=CQ5uKN_MN2gC&pg=PA149|page=149|title=Building Electro-Optical Systems: Making It All Work|author=Hobbs, Philip C.  D.|publisher=John Wiley and Sons|year=2009|isbn=0-470-40229-6}}</ref> 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.<ref>{{Cite book|publisher=World Scientific|url=http://books.google.com/books?id=FY7s7pPSPtgC&pg=PA819|title=Point Defects in Lithium Fluoride Films Induced by Gamma Irradiation|page=819|journal=Proceedings of the 7th International Conference on Advanced Technology & Particle Physics: (ICATPP-7): Villa Olmo, Como, Italy|year=2002|volume=2001|isbn=981-238-180-5}}</ref> Lithium fluoride is sometimes used in focal lenses of [[telescope]]s.<ref name=CRC/><ref>{{Cite journal|last1=Sinton|first1=William M.|title=Infrared Spectroscopy of Planets and Stars|journal=Applied Optics|volume=1|page=105|year=1962|doi=10.1364/AO.1.000105|bibcode = 1962ApOpt...1..105S|issue=2 }}</ref>
 
The high non-linearity of [[lithium niobate]] also makes it useful in [[nonlinear optics|non-linear optics applications]].  It is used extensively in telecommunication products such as mobile phones 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>
 
====Organic and polymer chemistry====
[[Organolithium compound]]s are widely used in the production of polymer and fine-chemicals.  In the polymer industry, which is the dominant consumer of these reagents, alkyl lithium compounds are [[catalyst]]s/[[radical initiator|initiators]].<ref>{{cite web|url=http://chemical.ihs.com/CEH/Public/Reports/681.7000/ |title=Organometallics|work=IHS Chemicals|date=February 2012}}</ref> in [[Anionic addition polymerization|anionic polymerization]] of [[functional group|unfunctionalized]] [[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|issue=9}}</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|page=1291|first2=Pao Luo|last2=Cheng|bibcode = 1986MaMol..19.1291Q|issue=5 }}</ref><ref>{{Cite book|title=Advances in organometallic chemistry|author= Stone, F.  G.  A.; West, Robert| publisher= Academic Press|year= 1980|isbn= 0-12-031118-6|page=55|url=http://books.google.com/?id=_gai4kRfcMUC&printsec=frontcover}}</ref> For the production of fine chemicals, organolithium compounds function as strong bases and as reagents for the formation of [[carbon-carbon bond]]s.  Organolithium compounds are prepared from lithium metal and alkyl halides.<ref>{{Cite book|url=http://books.google.com/books?id=_SJ2upYN6DwC&pg=PA192|page=192|title=Synthetic approaches in organic chemistry|author=Bansal, Raj K.  |year=1996|isbn=0-7637-0665-5}}</ref>
 
Many other lithium compounds are used as reagents to prepare organic compounds.  Some popular compounds include [[lithium aluminium hydride]] (LiAlH<sub>4</sub>), lithium triethylborohydride (LiBH(C<sub>2</sub>H<sub>5</sub>)<sub>3</sub>).
 
====Military applications====
Metallic lithium and its complex [[hydride]]s, such as [[lithium aluminium hydride|Li[AlH<sub>4</sub>]]], are used as high energy additives to [[rocket propellant]]s.<ref name=emsley/> Lithium aluminum hydride can also be used by itself as a [[solid fuel]].<ref>http://media.armadilloaerospace.com/misc/LiAl-Hydride.pdf</ref>
 
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 heat which is used to generate [[steam]].  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>
 
[[Lithium hydride]] containing lithium-6 is used in [[hydrogen bomb]]s.  In the bomb, it is placed around the core of an [[atomic bomb]].<ref>{{cite book|last=Emsley|first=John|title=Nature's Building Blocks|year=2011}}</ref>
 
===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 weapon]]s.<ref>{{cite book|pages=59–60|url=http://books.google.com/books?id=0oa1vikB3KwC&pg=PA60|title=Nuclear Wastelands: A Global Guide to Nuclear Weapons Production and Its Health and Environmental Effects|author=Makhijani, Arjun and Yih, Katherine |publisher=MIT Press|year= 2000|isbn=0-262-63204-7}}</ref> Lithium-7 gained interest for use in [[nuclear reactor]] [[coolant]]s.<ref>{{cite book|url=http://books.google.com/books?id=iRI7Cx2D4e4C&pg=PA278|page=278|title=Nuclear wastes: technologies for separations and transmutation|publisher=National Academies Press|year=1996|isbn=0-309-05226-2|author=National Research Council (U.S.).  Committee on Separations Technology and Transmutation Systems}}</ref>
 
[[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.<ref>{{Cite book|url=http://books.google.com/books?id=yTIOAAAAQAAJ&pg=PA39|page=39|title=How nuclear weapons spread: nuclear-weapon proliferation in the 1990s|author=Barnaby, Frank|publisher=Routledge|year=1993|isbn=0-415-07674-9}}</ref>
 
[[Lithium fluoride]], when highly enriched in the lithium-7 isotope, forms the basic constituent of the 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.  In addition, <sup>7</sup>Li, Be, and F are among the few [[nuclide]]s with low enough [[neutron cross-section|thermal neutron capture cross-sections]] not to poison the fission reactions inside a nuclear fission reactor.<ref group=note>Beryllium and fluorine occur only as one isotope, <sup>9</sup>Be and <sup>19</sup>F respectively.  These two, together with <sup>7</sup>Li, as well as [[deuterium|<sup>2</sup>H]], <sup>11</sup>B, <sup>15</sup>N, <sup>209</sup>Bi, and the stable isotopes of C, and O, are the only nuclides with low enough thermal neutron capture cross sections aside from [[actinide]]s to serve as major constituents of a molten salt breeder reactor fuel.</ref><ref>{{cite journal|last1=Baesjr|first1=C|title=The chemistry and thermodynamics of molten salt reactor fuels|journal=Journal of Nuclear Materials|volume=51|page=149|year=1974|doi=10.1016/0022-3115(74)90124-X|bibcode = 1974JNuM...51..149B }}</ref>
 
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.  Naturally occurring tritium is extremely rare, and must be synthetically produced by surrounding the reacting [[Plasma (physics)|plasma]] with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will fission the lithium to produce more tritium:
:<sup>6</sup>Li + n → <sup>4</sup>He + <sup>3</sup>T.
 
Lithium is also 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 fission to form two alpha particles.  This feat, called "splitting the atom" at the time, was the first fully man-made [[nuclear reaction]].  It was produced by [[John Douglas Cockcroft|Cockroft]] and [[Ernest Walton|Walton]] in 1932.<ref>{{Cite book|url=http://books.google.com/books?id=XyOBx2R2CxEC&pg=PA139|page=139|title=Nobel Prize Winners in Physics|author= Agarwal, Arun|publisher=APH Publishing|year=2008|isbn=81-7648-743-0}}</ref><ref>[http://www-outreach.phy.cam.ac.uk/camphy/cockcroftwalton/cockcroftwalton9_1.htm "'Splitting the Atom': Cockcroft and Walton, 1932: 9.  Rays or Particles?"] Department of Physics,University of Cambridge</ref> (Nuclear reactions and human-directed [[nuclear transmutation]] had been accomplished as early as 1917, but by using natural radioactive bombardment with [[alpha particle]]s).
 
===Medicine===
{{Main|Lithium (medication)}}
In the treatment of [[bipolar disorder]], lithium compounds are used.<ref name = "kean">{{cite book|last=Kean|first=Sam|title=The Disappearing Spoon|year=2011}}</ref>  Lithium salts may also be helpful for related diagnoses, such as [[schizoaffective disorder]] and cyclic [[major depression]].  The active principle in these salts is the lithium ion Li<sup>+</sup>.<ref name = "kean"/> There have been suggestions of increased risk of developing [[Ebstein's anomaly|Ebstein's cardiac anomaly]] in infants born to women taking lithium during the first trimester of pregnancy.<ref name="pmid18982835">{{cite journal |author=Yacobi S, Ornoy A |title=Is lithium a real teratogen? What can we conclude from the prospective versus retrospective studies? A review |journal=Isr J Psychiatry Relat Sci |volume=45 |issue=2 |pages=95–106 |year=2008 |pmid=18982835}}</ref>
 
==Precautions==
<div style="float: right;">
{| class="wikitable"
|-
! style="background:#f90;"|NFPA 704
|-
| style="text-align:left;"|{{NFPA 704|Health = 3|Inflammability = 3|Reactivity = 2|Other = <s>W</s>}}
|-
| style="width:80pt;"|The [[fire diamond]] hazard sign for lithium metal
|}</div>
Lithium is [[corrosive]] and requires special handling to avoid skin contact.  Breathing lithium dust or lithium compounds (which are often alkaline) initially [[irritation|irritate]] the [[human nose|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 [[Corrosive substance|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/?id=Oo3xAmmMlEwC&pg=PA244|pages = 244–246|isbn = 978-0-8493-2523-6|author = Furr, A.  K.|year = 2000|publisher = CRC Press|location = Boca Raton|title = CRC handbook of laboratory safety}}</ref>
 
===Regulation===
Some jurisdictions limit the sale of [[lithium battery|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]].<ref>{{cite web|url=http://www.illinoisattorneygeneral.gov/methnet/understandingmeth/basics.html |title=Illinois Attorney General&nbsp;– Basic Understanding Of Meth |publisher=Illinoisattorneygeneral.gov |accessdate=2010-10-06}}</ref><ref>{{cite journal|url=http://www.ncjolt.org/sites/default/files/7_nc_jl_tech_421.pdf|format=PDF|accessdate=2010-10-05|title=Methamphetamine remediation research act of 2005: Just what the doctor ordered for cleaning up methfields—or sugar pill placebo?| journal=North Carolina Journal of Law & Technology|year=2006|volume=7|first=Aaron R.|last= Harmon}}</ref>
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 = 978-0-306-44758-7|pages = 15–16|url = http://books.google.com/?id=i7U-0IB8tjMC&pg=PA15|author = Bro, Per and Levy, Samuel C.|year = 1994|publisher = Plenum Press|location = New York|title = Battery hazards and accident prevention}}</ref><ref>{{cite web|url=http://www.tsa.gov/travelers/airtravel/assistant/batteries.shtm |archiveurl=http://web.archive.org/web/20120104141539/http://www.tsa.gov/travelers/airtravel/assistant/batteries.shtm |archivedate=2012-01-04 |title=TSA: Safe Travel with Batteries and Devices |publisher=Tsa.gov |date=2008-01-01}}</ref>
 
==Notes==
<references group=note/>


{{Link FA|de}}
==References==
{{Reflist|2}}


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Editor-In-Chief: C. Michael Gibson, M.S., M.D. [2]

Overview

Lithium is a chemical element with symbol Li and atomic number 3. It is a soft, silver-white metal belonging to the alkali metal group of chemical elements. 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 freely in nature, and instead, only appears in compounds, which are usually ionic. Lithium occurs in a number of pegmatitic minerals, but due to its solubility as an ion is present in ocean water and is 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 verge on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative nuclear 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 helium in 1932 was the first fully man-made nuclear reaction, and lithium-6 deuteride serves as a fusion fuel in staged thermonuclear weapons.[2]

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.

Trace amounts of lithium are present in all organisms. The element serves no apparent vital biological function, since animals and plants survive in good health without it. Non-vital 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 in the treatment of bipolar disorder, due to neurological effects of the ion in the human body.

Atomic and Physical

Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[3] 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 alkali metals. Lithium's low reactivity compared to other alkali metals is due to the proximity of its valence electron to its nucleus (the remaining two electrons are in lithium's 1s orbital and are much lower in energy, and therefore they do not participate in chemical bonds).[3]

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.[3] While it has one of the lowest melting points among all metals (180 °C), it has the highest melting and boiling points of the alkali metals.[4]

Lithium has a very low density of 0.534 g/cm3, comparable with that of pine wood. It is the least dense of all elements that are solids at room temperature, the next lightest solid element (potassium, at 0.862 g/cm3) being more than 60% denser. Furthermore, apart from helium and hydrogen, it is less dense than any liquid element, being only 2/3 as dense as liquid nitrogen (0.808 g/cm3).[note 1][5] Lithium can float on the lightest hydrocarbon oils and is one of only three metals that can float on water, the other two being sodium and potassium.

Lithium's coefficient of thermal expansion is twice that of aluminium and almost four times that of iron.[6] It has the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[7] and at higher temperatures (more than 9 K) at very high pressures (>20 GPa)[8] 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.[9] Multiple allotropic forms have been reported for lithium at high pressures.[10]

Lithium has a specific heat capacity of 3.58 kilojoules per kilogram-Kelvin, the highest of all solids.[11][12] Because of this, lithium metal is often used in coolants for heat transfer applications.[11]

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.[3] Because of its reactivity with water, lithium is usually stored under cover of a hydrocarbon, often petroleum jelly. Though the heavier alkali metals can be stored in more dense substances, such as mineral oil, lithium is not dense enough to be fully submerged in these liquids.[13] 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).[14]

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.[15] 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 nonviolent, as the hydrogen produced will not ignite on its own. As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically the Class D type (see Types of extinguishing agents). Lithium is the only metal which reacts with nitrogen under normal conditions.[16][17]

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 (Li
2O) and peroxide (Li
2O
2) when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.[14][18] The metal reacts with hydrogen gas at high temperatures to produce lithium hydride (LiH).[19]

Other known binary compounds include the halides (LiF, LiCl, LiBr, LiI), and the sulfide (Li
2S), the superoxide (LiO
2), carbide (Li
2C
2). Many other inorganic compounds are known, where lithium combines with anions to form various salts: borates, amides, carbonate, nitrate, or borohydride (LiBH
4). Multiple organolithium reagents are known where there is a direct bond between carbon and lithium atoms effectively creating a carbanion. These 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.[20] LiHe, a very weakly interacting van der Waals compound, has been detected at very low temperatures.[21]

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).[3][13][22] 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 nuclides other than deuterium and helium-3. As a result of this, though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.[1]

Seven radioisotopes have been characterized, the most stable being 8Li with a half-life of 838 ms and 9Li with a half-life of 178 ms. All of the remaining radioactive isotopes have half-lives that are shorter than 8.6 ms. The shortest-lived isotope of lithium is 4Li, which decays through proton emission and has a half-life of 7.6 × 10−23 s.[23]

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.[24] 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.[25] While lithium is created in stars during the Stellar nucleosynthesis, it is further burnt. 7Li can also be generated in carbon stars.[26]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[27] 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.[28]

Nuclear weapons manufacture and other nuclear physics uses are a major source of artificial lithium fractionation, with the light isotope 6Li being retained by industry and military stockpiles to such an extent as to slightly but measurably change the 6Li to 7Li ratios even in natural sources, such as rivers. This has led to unusual uncertainty in the standardized atomic weight of lithium, since this quantity depends on the natural abundance ratios of these naturally-occurring stable lithium isotopes, as they are available in commercial lithium mineral sources.[29]

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.[30] 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.[31] 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.[13]

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 of the low temperature necessary to destroy lithium, and a lack of common processes to produce it.[32]

Lithium is also found in brown dwarf substellar objects and certain anomalous orange stars. Because lithium is present in cooler, less-massive brown dwarfs, 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.[13][33][34] 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.[13]

Terrestrial

Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity.[3] 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),[35][36] or 25 micromolar;[37] higher concentrations approaching 7 ppm are found near hydrothermal vents.[36]

Estimates for crustal content range from 20 to 70 ppm by weight.[14] 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.[14] Another significant mineral of lithium is lepidolite.[38] A newer source for lithium is hectorite clay, the only active development of which is through the Western Lithium Corporation in the United States.[39] At 20 mg lithium per kg of Earth's crust,[40] lithium is the 25th most abundant element.

According to the Handbook of Lithium and Natural Calcium, "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively few of them are of actual or potential commercial value. Many are very small, others are too low in grade."[41]

One of the largest reserve base[note 2] of lithium is in the Salar de Uyuni area of Bolivia, which has 5.4 million tonnes. US Geological Survey, estimates that in 2010 Chile had the largest reserves by far (7.5 million tonnes)[42] and the highest annual production (8,800 tonnes). Other major suppliers include Australia, Argentina and China.

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 as those of Bolivia, which now has the world's largest known lithium reserves."[43] 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."[44]

Biological

Lithium is found in trace amount in numerous plants, plankton, and invertebrates, at concentrations of 69 to 5,760 parts per billion (ppb). In vertebrates the concentration is slightly lower, and nearly all vertebrate tissue and body fluids have been found to contain lithium ranging from 21 to 763 ppb.[36] Marine organisms tend to bioaccumulate lithium more than terrestrial ones.[45] It is not known whether lithium has a physiological role in any of these organisms,[36] but nutritional studies in mammals have indicated its importance to health, leading to a suggestion that it be classed as an essential trace element with an RDA of 1 mg/day.[46] Observational studies in Japan, reported in 2011, suggested that naturally occurring lithium in drinking water may increase human lifespan.[47]

History of discovery and use

Petalite (LiAlSi4O10) was discovered in 1800 by the Brazilian chemist and statesman José Bonifácio de Andrada e Silva in a mine on the island of Utö, Sweden.[48][49][50] 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.[51][52][53] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[54] 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 "lithium".[3][49][53]

Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[49] In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame.[49] However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.[49][53][55] 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.[13][55][56][57] 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).[58] In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[49] 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.[49][59][60]

The production and use of lithium underwent several drastic changes in history. The first major application of lithium was in high-temperature lithium greases for aircraft engines or similar applications in World War II and shortly after. This use was supported by the fact that lithium-based soaps have a higher melting point than other alkali soaps, and are less corrosive than calcium based soaps. The small market for lithium soaps and the lubricating greases based upon them 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%, which was enough to affect the measured atomic weight of lithium in many standardized chemicals, and even the atomic weight of lithium in some "natural sources" of lithium ion which had been "contaminated" by lithium salts discharged from isotope separation facilities, which had found its way into ground water.[29][61]

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.[62][62] 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.[61] 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 21st century.

The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.[63] With the surge of lithium demand in batteries in the 2000s, new companies have expanded brine extraction efforts to meet the rising demand.[64][65]

Production

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 55% lithium chloride and 45% potassium chloride at about 450o C.[66] In 1998 it was about 95 US$ / kg (or 43 US$/pound).[67]

Worldwide identified reserves of lithium in 2008 were estimated by the US Geological Survey as 13 million tonnes. 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.[11] However, half the world's known reserves are located in Bolivia, a nation sitting along the central eastern slope of the Andes. In 2009 Bolivia was negotiating with Japanese, French, and Korean firms to begin extraction.[68] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tonnes of lithium.[68][69] A newly discovered deposit in Wyoming's Rock Springs Uplift is estimated at 228,000 tons. Additional deposits in the same formation were extrapolated to be as much as 18 million tons.[70]

After an industry wide pricing reduction for lithium carbonate after the Great Financial Crisis, where major suppliers such as Sociedad Química y Minera (SQM) dropped pricing by 20%[71] in light of incoming lithium resource developers and to further defend their market position, pricing in 2012 scaled up due to increased lithium demand. A 2012 Business Week article outlined the existing oligopoly in the lithium space, "SQM, controlled by billionaire Julio Ponce, is the second-largest, followed by Rockwood, which is backed by Henry Kravis’s KKR & Co., and Philadelphia-based FMC." Global consumption may jump to 300,000 metric tons a year by 2020 from about 150,000 tons in 2012, as demand for lithium batteries has been growing at about 25 percent a year, outpacing the 4 percent to 5 percent overall gain in lithium[72]

A potential source is geothermal wells. Geothermal fluids carry leachates to the surface;[73] recovery of lithium has been demonstrated in the field.[74] As the lithium is separated by simple filtration techniques, the process and environmental costs are primarily that of the already-operating geothermal well; relative environmental impacts may thus be positive.[75]

There are differing opinions about the potential growth of lithium production. According to a 2011 study conducted at Lawrence Berkeley National Laboratory and the University of California Berkeley, the currently estimated reserve base of lithium should not be a limiting factor for large-scale battery production for electric vehicles, as the study estimated that on the order of 1 billion 40 kWh Li-based batteries could be built with current reserves.[76] Another 2011 study by researchers from the University of Michigan and Ford Motor Company found that there are sufficient lithium resources to support global demand until 2100, including the lithium required for the potential widespread use of hybrid electric, plug-in hybrid electric and battery electric vehicles. The study estimated global lithium reserves at 39 million tons, and total demand for lithium during the 90-year period analyzed at 12–20 million tons, depending on the scenarios regarding economic growth and recycling rates.[77]

However, another 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'.

Ceramics and glass

Lithium oxide is a widely used flux for processing silica, reducing the melting point and viscosity of the material and leading to glazes of improved physical properties including low coefficients for thermal expansion.[78] Lithium oxides are a component of ovenware. Worldwide, this is the single largest use for lithium compounds. Lithium carbonate (Li2CO3) is generally used in this application: upon heating it converts to the oxide.[79]

Electrical and electronics

In the later years of the 20th century, owing to its high electrochemical potential, lithium became an important component of the electrolyte and of one of the electrodes in batteries. A typical lithium-ion battery 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.[80][81] 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.

Lubricating greases

The third most common use of lithium is in greases. Lithium hydroxide is a strong base and, when heated with a fat, 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.[11][82][83]

Metallurgy

When used as a flux for welding or soldering, metallic lithium promotes the fusing of metals during the process 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).[84]

Pyrotechnics

Lithium compounds are used as pyrotechnic colorants and oxidizers in red fireworks and flares.[11][85]

Air purification

Lithium chloride and lithium bromide are hygroscopic and are used as desiccants for gas streams.[11] Lithium hydroxide and lithium peroxide are the salts most used in confined areas, such as aboard spacecraft and submarines, for carbon dioxide removal and air purification. Lithium hydroxide absorbs carbon dioxide from the air by forming lithium carbonate, and is preferred over other alkaline hydroxides for its low weight.

Lithium peroxide (Li2O2) in presence of moisture not only reacts with carbon dioxide to form lithium carbonate, but also releases oxygen.[86][87] The reaction is as follows:

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.

Some of the aforementioned compounds, as well as lithium perchlorate, are used in oxygen candles that supply submarines with oxygen. These can also include small amounts of boron, magnesium, aluminum, silicon, titanium, manganese, and iron.[88]

Optics

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.[89] 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.[90] Lithium fluoride is sometimes used in focal lenses of telescopes.[11][91]

The high non-linearity of lithium niobate also makes it useful in non-linear optics applications. It 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.[92]

Organic and polymer chemistry

Organolithium compounds are widely used in the production of polymer and fine-chemicals. In the polymer industry, which is the dominant consumer of these reagents, alkyl lithium compounds are catalysts/initiators.[93] in anionic polymerization of unfunctionalized olefins.[94][95][96] For the production of fine chemicals, organolithium compounds function as strong bases and as reagents for the formation of carbon-carbon bonds. Organolithium compounds are prepared from lithium metal and alkyl halides.[97]

Many other lithium compounds are used as reagents to prepare organic compounds. Some popular compounds include lithium aluminium hydride (LiAlH4), lithium triethylborohydride (LiBH(C2H5)3).

Military applications

Metallic lithium and its complex hydrides, such as Li[AlH4], are used as high energy additives to rocket propellants.[13] Lithium aluminum hydride can also be used by itself as a solid fuel.[98]

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 heat which is used to generate steam. The steam propels the torpedo in a closed Rankine cycle.[99]

Lithium hydride containing lithium-6 is used in hydrogen bombs. In the bomb, it is placed around the core of an atomic bomb.[100]

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.[101] Lithium-7 gained interest for use in nuclear reactor coolants.[102]

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.[103]

Lithium fluoride, when highly enriched in the lithium-7 isotope, forms the basic constituent of the fluoride salt mixture LiF-BeF2 used in liquid fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF-BeF2 mixtures have low melting points. In addition, 7Li, Be, and F are among the few nuclides with low enough thermal neutron capture cross-sections not to poison the fission reactions inside a nuclear fission reactor.[note 3][104]

In conceptualized nuclear fusion power plants, lithium will be used to produce tritium in magnetically confined reactors using deuterium and tritium as the fuel. Naturally occurring tritium is extremely rare, and must be synthetically produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will fission the lithium to produce more tritium:

6Li + n → 4He + 3T.

Lithium is also used as a source for alpha particles, or helium nuclei. When 7Li is bombarded by accelerated protons 8Be is formed, which undergoes fission to form two alpha particles. This feat, called "splitting the atom" at the time, was the first fully man-made nuclear reaction. It was produced by Cockroft and Walton in 1932.[105][106] (Nuclear reactions and human-directed nuclear transmutation had been accomplished as early as 1917, but by using natural radioactive bombardment with alpha particles).

Medicine

Main article: Lithium (medication)

In the treatment of bipolar disorder, lithium compounds are used.[107] Lithium salts may also be helpful for related diagnoses, such as schizoaffective disorder and cyclic major depression. The active principle in these salts is the lithium ion Li+.[107] 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.[108]

Precautions

NFPA 704
Template:NFPA 704
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.[109]

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.[110][111] 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.[112][113]

Notes

  1. ↑ Densities for all the gaseous elements can be obtained at Airliquide.com
  2. ↑ Apendixes. By USGS definitions, reserve base "may encompass those parts of the resources that have a reasonable potential for becoming economically available within planning horizons beyond those that assume proven technology and current economics. The reserve base includes those resources that are currently economic (reserves), marginally economic (marginal reserves), and some of those that are currently subeconomic (subeconomic resources)."
  3. ↑ 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 a molten salt breeder reactor fuel.

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