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{{Infobox lithium}} |
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'''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]]. When cut open, lithium exhibits a metallic [[lustre (mineralogy)|lustre]], but contact with oxygen quickly returns it back to a dull silvery grey color. Lithium is also highly flammable. |
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According to theory, lithium (mostly <sup>7</sup>Li) was one of the few elements [[Big Bang nucleosynthesis|synthesized]] in the [[Big Bang]], although its quantity has vastly decreased. The reasons for its disappearance and the processes by which new lithium is created continue to be important matters of study in [[astronomy]]. Lithium is tied with [[krypton]] as 32nd or 33rd most abundant element in the [[cosmos]] (see [[Cosmochemical Periodic Table of the Elements in the Solar System]]), being less common than any element before [[Rubidium]] (element 37) except for [[scandium]], [[gallium]], [[arsenic]], and [[bromine]], yet more common than any element beyond krypton (element 36). |
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'''Hey Ya'll Gwa Nay'''''''''Bold text'''''''''Bold text'''[[[[Link title]] |
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Due to its high [[reactivity]] it only appears naturally on Earth 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]]. |
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== <math>Headline text</math> == |
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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 a [[nuclear reaction]], and [[lithium deuteride]] serves as the [[nuclear fusion|fusion]] fuel in [[Teller-Ulam design|staged thermonuclear weapon]]s. |
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== History and etymology== |
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[[Petalite]] (lithium aluminium silicate) was first described in 1800 by the [[Brazil]]ian ([[Brazilian Declaration of Independence|then]] [[Portugal|Portuguese]]) 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 "lithos", 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 | accessdate = 2007-08-19| 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>{{cite web | publisher = Elementymology & Elements Multidict | title = Lithium| first = Peter | last =van der Krogt | url = http://www.vanderkrogt.net/elements/elem/li.html| accessdate = 2008-09-18}}</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>{{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 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>{{cite web| url = http://www.echeat.com/essay.php?t=29195 | title = Analysis of the Element Lithium | first = Thomas | last = Green | date 2006-06-11| publisher = echeat}}</ref> |
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== Properties == |
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[[Image:Limetal.JPG|thumb|left|Lithium ingots with a thin layer of black oxide tarnish]] |
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[[Image:LithiumPelletsUSGOV.jpg|thumb|left|Lithium pellets (covered in white lithium hydroxide)]] |
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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. |
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In addition, lithium has a [[diagonal relationship]] with [[magnesium]], an element of similar atomic and [[ionic radius]]. Chemical resemblances between the two metals include the formation of a [[nitride]] by reaction with N<sub>2</sub>, the formation of an [[oxide]] when burnt in O<sub>2</sub>, [[salt (chemistry)|salts]] with similar [[solubility|solubilities]], and thermal instability of the [[carbonate]]s and nitrides.<ref name=kamienski>{{ cite book | first = Conrad W. last = Kamienski, McDonald, Daniel P.; Stark, Marshall W.; Papcun, John R. | chapter =Lithium and lithium compounds | title =Kirk-Othmer Encyclopedia of Chemical Technology | publisher = John Wiley & Sons, Inc.| year = 2004 | doi =10.1002/0471238961.1209200811011309.a01.pub2}}</ref> |
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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 such as [[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. |
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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]]. |
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At cryogenic temperatures, lithium, like sodium, undergoes martensitic transformations when cooled below liquid nitrogen temperatures (77 <sup>o</sup>K). This is a multicrystalline state composed of FCC (Face Centered Cubic), BCC (Body Centered Cubic) and R9 Hex. At liquid helium temperatures (4 <sup>o</sup>K) 9R Hex is the most prevalent crystal. The proportion of the different crystalline states is temperature-dependent. On cryogenic cooling lithium at atmospheric pressure, the crystalline state which first predominates is FCC, followed by BCC, followed by R9 Hex at the coldest temperatures. On heating solid lithium from deep cryogenic temperatures, the property known as ''heat of reversion'' will cause a crystalline state transition from R9 Hex to BCC, which absorbs heat and causes cooling. In warming at cryogenic peratures, there is thus a region of negative specific heat in lithium crystals, due to this state change. <ref>REFERENCE see paper by DOUGLAS L. MARTIN circa 1956.</ref> |
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== Chemistry == |
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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/> |
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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. |
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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). |
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== Isotopes == |
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{{main|Isotopes of lithium}} |
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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>{{cite web |url=http://ie.lbl.gov/education/parent/Li_iso.htm |title=Isotopes of Lithium|accessdate=2008-04-21 |author= |date= |work= |publisher=Berkley Lab, The Isotopes Project}}</ref> 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 ms. All of the remaining [[radioactive]] isotopes have half-lives that are shorter than 8.6 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. |
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<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).<ref>{{cite web |url=http://www.journals.uchicago.edu/doi/abs/10.1086/503538 |title=Lithium Isotopic Abundances in Metal-poor Halo Stars |accessdate=2008-04-21 |author= |date=June 10, 2006 |work= |publisher=The Astrophysical Journal}}</ref> 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. The exotic <sup>11</sup>Li is known to exhibit a [[nuclear halo]]. |
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== Natural occurrence == |
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[[Image:Relative abundance of elements.png|thumb|right|Lithium is about as common as [[chlorine]] in the earth's upper continental [[Crust (geology)|crust]], on a per-atom basis.]] |
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{{see also|:category:Lithium minerals|l1=Lithium minerals}} |
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Lithium is widely distributed 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> 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/> |
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According to the ''Handbook of Lithium and Natural Calcium'', "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively a few of them are of actual or potential commercial value. Many are very small, others are too low in grade."<ref>''Handbook of Lithium and Natural Calcium'', [[Donald Garrett]], [[Academic Press]], 2004, cited in ''[http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf The Trouble with Lithium 2]''</ref> |
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== Applications == |
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Because of its [[specific heat]] capacity, the highest of all [[solid]]s, lithium is often used in heat transfer applications. |
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It is an important ingredient in anode materials, used in [[Lithium-ion battery|rechargeable]] and [[Lithium battery|primary]] [[Battery (electricity)|batteries]] because of its high [[electrochemical potential]], light weight, and high current density. |
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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. |
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=== Medical use === |
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{{main|Lithium pharmacology}} |
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Lithium salts were used during the 19th century to treat [[gout]]. Lithium salts such as [[lithium carbonate]] (Li<sub>2</sub>CO<sub>3</sub>), [[lithium citrate]], and [[lithium orotate]] are mood 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.{{Fact|date=September 2008}} |
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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. |
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Common side effects of lithium treatment include muscle [[tremor]]s, twitching, [[ataxia]], [[hyperparathyroidism]], bone loss, [[hypercalcemia]], [[hypertension]], etc.), kidney damage, [[nephrogenic diabetes insipidus]] (polyuria and polydipsia) and [[seizure]]s. Many of the side-effects are a result caused by the increased elimination of potassium. |
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Pregnancy - teratogenic properties: [[Ebstein's anomaly|Ebstein (cardiac) Anomaly]] - There appears to be an increased risk of this abnormality in infants of women taking lithium during the first trimester of pregnancy |
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=== Other uses === |
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* [[Lithium batteries]] are [[disposable]] ([[primary cell|primary]]) [[Battery (electricity)|batteries]] that have lithium metal or lithium compounds as an [[anode]]. Lithium batteries are not to be confused with [[lithium-ion battery|lithium-ion batteries]] which are high energy-density rechargeable batteries |
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* [[Lithium chloride]] and [[lithium bromide]] are extremely [[hygroscopic]] and frequently used as [[desiccant]]s. |
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* [[stearic acid|Lithium stearate]] is a common all-purpose high-temperature [[lubricant]]. |
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* Lithium is an [[alloy]]ing agent used to synthesize [[organic compound]]s. |
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* 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]]. |
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* Lithium is sometimes used in glasses and ceramics including the glass for the 200-inch (5.08 m) [[telescope]] at [[Mt. Palomar]]. |
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* [[Alloy]]s of the metal with [[aluminium]], [[cadmium]], [[copper]] and [[manganese]] are used to make high performance [[aircraft]] parts. |
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* [[Al-Li|Lithium-aluminium alloys]] are used in [[aerospace]] applications, such as the [[Space Shuttle external tank|external tank]] of the [[Space Shuttle]], and is planned for the [[Orion (spacecraft)|Orion spacecraft]]. |
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* [[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> |
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* The high non-linearity of lithium niobate also makes a good choice for [[nonlinear optics|non-linear optics applications]]. |
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* [[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 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. |
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* 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>. |
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* [[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> |
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* Lithium fluoride (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred fluoride salt mixture (LiF-BeF2) used in liquid-fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF/BeF2 mixtures have low melting points and the best neutronic properties of fluoride salt combinations appropriate for reactor use. |
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* 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. |
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* 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. |
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* [[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. |
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* 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>. |
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* Lithium metal is used as a [[reducing agent]] in some types of [[methamphetamine]] production, particularly in illegal amateur “meth labs.” |
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* Lithium can be used to make red fireworks |
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== Production == |
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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]]. |
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There are wide-spread hopes of using [[lithium ion batteries]] in [[electric vehicles]], but one study concluded that "realistically achievable lithium carbonate production will be sufficient for only a small fraction of future [[PHEV]] and [[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'".<ref name=Legers/> |
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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 | accessdate = 2007-08-19|last=Ober |first=Joyce A |format=pdf |pages = 77-78| publisher=[[United States Geological Survey]]}}</ref> |
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[[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 | accessdate = 2007-08-19|date= December 15, 2003|publisher= [[Los Alamos National Laboratory]]}}</ref> |
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China may emerge as a significant producer of brine-based lithium carbonate around 2010. Potential capacity of up to 55,000 tonnes per year could come on-stream if projects in Qinghai province and Tibet proceed.<ref name=Legers>{{cite web|url=http://www.meridian-int-res.com/Projects/Lithium_Microscope.pdf |title=The Trouble With Lithium 2 |accessdate = 2008-07-07 |date=May 28, 2008|publisher=[[Meridian International Research]]}}</ref> |
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The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tonnes of the known global lithium reserve base.<ref name=Tahil>{{cite web|url=http://www.meridian-int-res.com/Projects/Lithium_Problem_2.pdf |title=The Trouble with Lithium | accessdate = 2008-07-07|date=January 2007|publisher=[[Meridian International Research]]}}</ref> |
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In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tonnes for the Western World.<ref>Evans, R.K. (1978) "Lithium Reserves and Resources" Energy, Vol 3 No.3</ref> The inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total has nearly tripled by 2008.<ref>Evans, R.K. (2008) "An Abundance of Lithium" http://www.worldlithium.com/Abstract.html </ref><ref>Evans, R.K. (2008) "An Abundance of Lithium Part 2" http://www.worldlithium.com/AN_ABUNDANCE_OF_LITHIUM_-_Part_2.html</ref> |
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== Precautions == |
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Lithium metal, due to its alkaline tarnish, is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) can irritate the nose and throat; higher exposure to lithium can cause a build-up of fluid in the lungs, leading to [[pulmonary edema]]. The metal itself is usually a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium should be stored in a non-reactive compound such as [[naphtha]] or a hydrocarbon.{{Facts|date=February 2008}} |
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=== Regulation === |
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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.{{Facts|date=February 2008}} |
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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.{{Facts|date=February 2008}} |
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==See also== |
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*[[:Category:Lithium compounds|Lithium compounds]] |
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*[[Dilithium]] |
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== References == |
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{{reflist|2}} |
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== External links == |
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{{Wiktionary|lithium}} |
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{{Commons|Lithium}} |
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*[http://minerals.usgs.gov/minerals/pubs/commodity/lithium/ USGS: Lithium Statistics and Information] |
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*[http://www.webelements.com/lithium/ WebElements.com – Lithium] |
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*[http://education.jlab.org/itselemental/ele003.html It's Elemental – Lithium] |
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*[http://www.bipolar-lives.com/lithium.html Information on Lithium and Bipolar Disorder] |
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*[http://www.mcis.soton.ac.uk/Site_Files/pdf/nuclear_history/Working_Paper_No_5.pdf University of Southampton, Mountbatten Centre for International Studies, Nuclear History Working Paper No5.] |
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{{alkalimetals}} |
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{{Compact periodic table}} |
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[[Category:Chemical elements]] |
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[[Category:Alkali metals]] |
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[[Category:Lithium|*]] |
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[[cs:Lithium]] |
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[[co:Litiu]] |
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[[cy:Lithiwm]] |
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[[da:Lithium]] |
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[[de:Lithium]] |
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[[et:Liitium]] |
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[[el:Λίθιο]] |
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[[es:Litio]] |
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[[eo:Litio]] |
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[[eu:Litio]] |
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[[fa:لیتیم]] |
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[[fo:Lithium]] |
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[[fr:Lithium]] |
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[[fur:Liti]] |
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[[ga:Litiam]] |
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[[gv:Litçhey]] |
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[[gl:Litio]] |
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[[ko:리튬]] |
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[[haw:Likiuma]] |
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[[hy:Լիթիում]] |
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[[hi:लिथियम]] |
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[[hr:Litij]] |
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[[io:Litio]] |
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[[id:Litium]] |
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[[is:Litín]] |
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[[it:Litio]] |
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[[he:ליתיום]] |
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[[jv:Litium]] |
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[[kn:ಲಿಥಿಯಮ್]] |
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[[ka:ლითიუმი]] |
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[[sw:Lithi]] |
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[[ht:Lityòm]] |
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[[ku:Lîtyûm]] |
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[[la:Lithium]] |
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[[lv:Litijs]] |
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[[lb:Lithium]] |
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[[lt:Litis]] |
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[[jbo:roksodna]] |
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[[hu:Lítium]] |
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[[mk:Литиум]] |
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[[ml:ലിഥിയം]] |
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[[ms:Litium]] |
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[[nl:Lithium]] |
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[[ja:リチウム]] |
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[[no:Litium]] |
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[[nn:Litium]] |
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[[oc:Liti]] |
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[[uz:Litiy]] |
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[[pa:ਲਿਥੀਅਮ]] |
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[[nds:Lithium]] |
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[[pl:Lit (pierwiastek)]] |
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[[pt:Lítio]] |
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[[ksh:Lithium]] |
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[[ro:Litiu]] |
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[[qu:Lityu]] |
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[[ru:Литий]] |
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[[sa:लिथियम]] |
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[[sq:Litiumi]] |
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[[simple:Lithium]] |
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[[sk:Lítium]] |
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[[sl:Litij]] |
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[[sr:Литијум]] |
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[[sh:Litij]] |
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[[su:Litium]] |
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[[fi:Litium]] |
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[[sv:Litium]] |
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[[ta:லித்தியம்]] |
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[[th:ลิเทียม]] |
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[[vi:Liti]] |
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[[tr:Lityum]] |
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[[uk:Літій]] |
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[[wa:Litiom]] |
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[[yi:ליטיום]] |
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[[zh-yue:鋰]] |
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[[zh:锂]] |
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