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Lithium isotopes fractionate substantially during a wide variety of natural processes,<ref>{{cite journal |year=2004 |first1=H.M. |last1=Seitz |first2=G.P. |last2=Brey |first3=Y. |last3=Lahaye |first4=S. |last4=Durali |first5=S.|last5=Weyer |title=Lithium isotopic signatures of peridotite xenoliths and isotopic fractionation at high temperature between olivine and pyroxenes |journal=Chemical Geology |volume=212 |number=1-2|doi=10.1016/j.chemgeo.2004.08.009|pages=163–177}}</ref> including mineral formation (chemical precipitation), [[metabolism]], and [[ion exchange]]. Lithium ions substitute for [[magnesium]] and [[iron]] in octahedral sites in [[clay]] minerals, where <sup>6</sup>Li is preferred to <sup>7</sup>Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic <sup>11</sup>Li is known to exhibit a [[nuclear halo]].
Lithium isotopes fractionate substantially during a wide variety of natural processes,<ref>{{cite journal |year=2004 |first1=H.M. |last1=Seitz |first2=G.P. |last2=Brey |first3=Y. |last3=Lahaye |first4=S. |last4=Durali |first5=S.|last5=Weyer |title=Lithium isotopic signatures of peridotite xenoliths and isotopic fractionation at high temperature between olivine and pyroxenes |journal=Chemical Geology |volume=212 |number=1-2|doi=10.1016/j.chemgeo.2004.08.009|pages=163–177}}</ref> including mineral formation (chemical precipitation), [[metabolism]], and [[ion exchange]]. Lithium ions substitute for [[magnesium]] and [[iron]] in octahedral sites in [[clay]] minerals, where <sup>6</sup>Li is preferred to <sup>7</sup>Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic <sup>11</sup>Li is known to exhibit a [[nuclear halo]].
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==History and etymology==
==History and etymology==

Revision as of 15:21, 27 May 2010

This article is about the chemical element. For other uses, see Lithium (disambiguation).

Lithium, 3Li
Freshly cut sample of lithium, with minimal oxides
Lithium
Pronunciation/ˈlɪθiəm/ ​(LITH-ee-əm)
Appearancesilvery-white
Standard atomic weight Ar°(Li)
Lithium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
H
↑
Li
↓
Na
helium ← lithium → beryllium
Atomic number (Z)3
Groupgroup 1: hydrogen and alkali metals
Periodperiod 2
Block  s-block
Electron configuration[He] 2s1
Electrons per shell2, 1
Physical properties
Phase at STPsolid
Melting point453.65 K ​(180.50 °C, ​356.90 °F)
Boiling point1617 K ​(1344 °C, ​2451 °F)
Density (at 20° C)0.5334 g/cm3[3]
when liquid (at m.p.)0.512 g/cm3
Critical point3220 K, 67 MPa (extrapolated)
Heat of fusion3.00 kJ/mol
Heat of vaporization136 kJ/mol
Molar heat capacity24.860 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 797 885 995 1144 1337 1610
Atomic properties
Oxidation statescommon: +1
−1[4]
ElectronegativityPauling scale: 0.98
Ionization energies
  • 1st: 520.2 kJ/mol
  • 2nd: 7298.1 kJ/mol
  • 3rd: 11815.0 kJ/mol
Atomic radiusempirical: 152 pm
Covalent radius128±7 pm
Van der Waals radius182 pm
Color lines in a spectral range
Spectral lines of lithium
Other properties
Natural occurrenceprimordial
Crystal structure ​body-centered cubic (bcc) (cI2)
Lattice constant
Body-centered cubic crystal structure for lithium
a = 350.93 pm (at 20 °C)[3]
Thermal expansion46.56×10−6/K (at 20 °C)[3]
Thermal conductivity84.8 W/(m⋅K)
Electrical resistivity92.8 nΩ⋅m (at 20 °C)
Magnetic orderingparamagnetic
Molar magnetic susceptibility+14.2×10−6 cm3/mol (298 K)[5]
Young's modulus4.9 GPa
Shear modulus4.2 GPa
Bulk modulus11 GPa
Speed of sound thin rod6000 m/s (at 20 °C)
Mohs hardness0.6
Brinell hardness5 MPa
CAS Number7439-93-2
History
DiscoveryJohan August Arfwedson (1817)
First isolationWilliam Thomas Brande (1821)
Isotopes of lithium
Main isotopes[6] Decay
abun­dance half-life (t1/2) mode pro­duct
6Li [1.9%, 7.8%] stable
7Li [92.2%, 98.1%] stable
Significant variation occurs in commercial samples because of the wide distribution of samples depleted in 6Li.
 Category: Lithium
| references

Lithium is a soft, silver-white metal that belongs to the alkali metal group of chemical elements. It is represented by the symbol Li, and it has the atomic number 3. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable. For this reason, lithium metal is typically stored in a petroleum derivative, such as mineral oil (in which it floats), or petroleum jelly in which it is held below the surface by the viscosity of the medium. When cut open, lithium exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a dull silvery gray, then black, tarnish.

The nuclei of lithium are relatively fragile: the two stable lithium isotopes found in nature have lower binding energies per nucleon than any other stable compound nuclides, save deuterium, and helium-3 (3He).[7] Though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.[8]

Because of its high reactivity, lithium only appears naturally in the form of compounds. Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from brines and clays. On a commercial scale, lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

Trace amounts of lithium are present in the oceans and in some organisms, though the element serves no apparent vital biological function in humans. The lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood stabilizing drug due to neurological effects of the ion in the human body. Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, and lithium batteries. Lithium also has important links to nuclear physics. The transmutation of lithium atoms to tritium was the first man-made form of a nuclear fusion reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.

Characteristics

Physical

Lithium pellets (covered in white lithium hydroxide)
See also: Alkali metal

Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[9] Because of this, it is a good conductor of both heat and electricity and highly reactive, though it is the least reactive of the alkali metals due to the proximity of its valence electron to its nucleus.[9]

Lithium is soft enough to be cut with a knife, and it is the lightest of the metals of the periodic table. When cut, it possesses a silvery-white color that quickly changes to gray due to oxidation.[9] It also has a low density (approximately 0.534 g/cm3), giving sticks of the metal a similar heft to dowels of wood, such as pine. It floats on water, with which it reacts easily but with noticeably less energy than other alkali metals do. The reaction forms hydrogen gas and lithium hydroxide in aqueous solution.[9] Because of its reactivity with water, lithium is usually stored under cover of a viscous hydrocarbon, often petroleum jelly; though the heavier alkali metals can be stored in less dense substances, such as mineral oil, lithium is not dense enough to be fully submerged in these liquids.[10]

Lithium possesses a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[11] and at higher temperatures (more than 9 kelvins) at very high pressures (over 200,000 atmospheres)[12] At cryogenic temperatures, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2K it has a rhombohedral crystal system (with a nine-layer repeat spacing)[13]; at higher temperatures it transforms to face-centered cubic and then body-centered cubic. At liquid-helium temperatures (4 K) the rhombohedral structure is the most prevalent.

Chemical

In moist air, lithium metal rapidly tarnishes to form a black coating of lithium hydroxide (LiOH and LiOH·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 gives off a striking crimson color, but when it burns strongly the flame becomes a brilliant silver. Lithium will ignite and burn in oxygen when exposed to water or water vapours.[15]

Lithium metal is flammable, and it is potentially explosive when exposed to air and especially to water, though less so than the other alkali metals. The lithium-water reaction at normal temperatures is brisk but not violent, though the hydrogen produced can ignite. As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically Class D type (see Types of extinguishing agents). Lithium is the only metal which reacts with nitrogen under normal conditions.

Lithium compounds

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 when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.[14]

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).[9][10][16] Both natural isotopes have anomalously low nuclear binding energy per nucleon compared to the next lighter and heavier elements, helium and beryllium, which means that alone among stable light elements, lithium can produce net energy through nuclear fission. Seven radioisotopes have been characterized, the most stable being 8Li with a half-life of 838 ms and 9Li 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 4Li, which decays through proton emission and has a half-life of 7.58043 × 10−23 s.

7Li is one of the primordial elements (or, more properly, primordial isotopes) 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 it is produced.[17] Additional small amounts of lithium of both 6Li and 7Li may be generated from solar wind, cosmic rays, and early solar system 7Be and 10Be radioactive decay.[18] 7Li can also be generated in carbon stars.[19]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[20] 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.

History and etymology

Petalite (LiAlSi4O10, which is lithium aluminium silicate) was first discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada e Silva, who discovered this mineral in a mine on the island of Utö, Sweden.[21][22][23] 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.[24][25][26] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[27] Berzelius gave the alkaline material the name "lithos", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to sodium and potassium, which had been discovered in plant tissues. The name of this element was later standardized as "lithium".[9][22][26] Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[22] In 1818, Christian Gmelin was the first man to observe that lithium salts give a bright red color in flame.[22] However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.[22][26][28] This element, lithium, was not isolated until 1821, when William Thomas Brande isolated the element by performing electrolysis on lithium oxide, a process that had previously been employed by the chemist Sir Humphry Davy to isolate the alkali metals potassium and sodium.[10][28][29] Brande also described some pure salts of lithium, such as the chloride, and he performed an estimate of its atomic weight. In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[22] The discovery of this procedure henceforth led to commercial production of lithium metal, beginning in 1923, by the German company Metallgesellschaft AG, which performed an electrolysis of a liquid mixture of lithium chloride and potassium chloride.[22][30]

The production and use of lithium underwent several drastic changes in history. The first major application of lithium became high temperature grease for aircraft engines or similar applications in World War II and shortly after. This small market was supported by several small mining operations mostly in the United States. The demand for lithium increased dramatically when in the beginning of the cold war the need for the production of nuclear fusion weapons arose and the dominant fusion material tritium had to be made by irradiating lithium-6. The United States became the prime producer of lithium in the period between the late 1950s and the mid 1980s. At the end the stockpile of lithium was roughly 42,000 tons of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75% .[31]

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.[32][32] 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.[31] Then, in the mid 1990's several companies started to extract lithium from brine; this method proved to be less expensive than underground or even open pit mining. Most of the mines closed or shifted their focus to other materials as only the ore from zoned pegmatites could be mined for a competitive price. For example, the US mines near Kings Mountain, North Carolina closed before the turn of the century. The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.[31] New companies have expanded brine extraction efforts to meet the rising demand.[33]

Occurrence

Lithium is about as common as chlorine in the Earth's upper continental crust, on a per-atom basis.
Lithium mine production (2008) and reserves in metric tonnes[34]
Country Production Reserves Reserve base
 Argentina 3,200 Not available Not available
 Australia 6,900 170,000 220,000
 Bolivia 0 0 5,400,000
 Brazil 180 190,000 910,000
 Canada 710 180,000 360,000
 Chile 12,000 3,000,000 3,000,000
 People's Republic of China 3,500 540,000 1,100,000
 Portugal 570 Not available Not available
 United States Withheld 38,000 410,000
 Zimbabwe 300 23,000 27,000
World total 27,400 4,100,000 11,000,000

Astronomical occurrence

Main article: nucleosynthesis

According to modern cosmological theory, both stable isotopes of lithium—6Li and 7Li—were among the 3 elements synthesized in the Big Bang. Though the amount of lithium generated in Big Bang nucleosynthesis is dependent upon the number of photons per baryon, for accepted values the lithium abundance can be calculated, and there is a "cosmological lithium discrepancy" in the universe: older stars seem to have less lithium than they should, and some younger stars have far more. The lack of lithium in older stars is apparently caused by the "mixing" of lithium into the interior of stars, where it is destroyed.[35] Furthermore, lithium is produced in younger stars. Though it transmutes into two atoms of helium due to collision with a proton at temperatures above 2.4 million degrees Celsius (most stars easily attain this temperature in their interiors), lithium is more abundant than predicted in later-generation stars, for causes not yet completely understood.[10]

Though it was one of the 3 first elements to be synthesized in the Big Bang, lithium, as well as beryllium and boron are markedly less abundant than the elements with either lower or higher atomic number. This is due to the low temperature necessary to destroy lithium, and a lack of common processes to produce it.[36]

Lithium is also found in brown dwarf stars and certain anomalous orange stars. Because lithium is present in cooler, less-massive brown dwarf stars, but is destroyed in hotter red dwarf stars, its presence in the stars' spectra can be used in the "lithium test" to differentiate the two, as both are smaller than the Sun.[10][37][38] Certain orange stars can also contain a high concentration of lithium. Those orange stars found to have a higher than usual concentration of lithium (such as Centaurus X-4) orbit massive objects—neutron stars or black holes—whose gravity evidently pulls heavier lithium to the surface of a hydrogen-helium star, causing more lithium to be observed.[10]

Occurrence on Earth

See also: Lithium minerals

Lithium is widely distributed on Earth but does not naturally occur in elemental form due to its high reactivity.[9] 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] 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]

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."[40] At 20 mg lithium per kg of Earth's crust [41], lithium is the 25th most abundant element. Nickel and lead have the about the same abundance.

The largest reserve base of lithium is in the Salar de Uyuni area of Bolivia, which has 5.4 million tons. According to the US Geological Survey, the production and reserves of lithium in metric tons are as follows[34][42]:

Contrary to the USGS data in the table, other estimates put Chile's reserve base at 7,520,000 metric tons of lithium, and Argentina's at 6,000,000 metric tons.[43]

Seawater contains an estimated 230 billion tons of lithium, though at a low concentration of 0.1 to 0.2 ppm.[44]

Production

Lithium mine, Salar del Hombre Muerto, Argentina. The brine in this salt flat is rich in lithium, and the mine concentrates the brine by pumping it into solar evaporation ponds. 2009 image from NASA’s EO-1 satellite
Uyuni salt flat in Bolivia

Since the end of World War II lithium metal production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits.

The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride. In 1998 it was about 95 US$ / kg (or 43 US$/pound).[45]

Deposits of lithium are found in South America throughout the Andes mountain chain. Chile is the leading lithium metal producer, followed by Argentina. Both countries recover the lithium from brine pools. In the United States lithium is recovered from brine pools in Nevada.[46] Nearly half the world's known reserves are located in Bolivia, a nation sitting along the central eastern slope of the Andes. In 2009 Bolivia is negotiating with Japanese, French, and Korean firms to begin extraction.[47] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tons of lithium, which can be used to make batteries for hybrid and electric vehicles.[47][48]

China may emerge as a significant producer of brine-source lithium carbonate around 2010. There is potential production of up to 55,000 tons per year if projects in Qinghai province and Tibet proceed.[49]

The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tons of the known global lithium reserve base.[50] This is enough for approximately 4 billion electric cars.[51]

In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tons for the Western World.[52] With the inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total had nearly tripled by 2008.[53][54]

Applications

Because of its specific heat capacity, the highest of all solids, lithium is often used in coolants for heat transfer applications.

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

Lithium is also used in the pharmaceutical and fine-chemical industry in the manufacture of organolithium reagents, which are used both as strong bases and as reagents for the formation of carbon-carbon bonds. Organolithiums are also used in polymer synthesis as catalysts/initiators[55] in anionic polymerization of unfunctionalised olefins.[56][57][58]

Lithium-6 is valued as a source material for tritium production and as a neutron absorber in nuclear fusion. Natural lithium contains about 7.5 percent lithium-6. Large amounts of lithium-6 have been produced by isotope separation for use in nuclear weapons. Lithium-7 gained interest for use in nuclear reactor coolants.

Medical use

Main article: Lithium pharmacology

Lithium salts were used during the 19th century to treat gout. Lithium salts such as lithium carbonate (Li2CO3), lithium citrate, and lithium orotate are mood stabilizers. They are used in the treatment of bipolar disorder since, unlike most other mood altering drugs, they counteract both mania and depression. Lithium continues to be the gold standard for the treatment of bipolar disorder. It is also helpful for related diagnoses, such as schizoaffective disorder and cyclic major depression. In addition to watching out for the well-known complications of lithium treatment—hypothyroidism and decreased renal function—health care providers should be aware of hyperparathyroidism.[59] Lithium can also be used to augment antidepressants. Because of Lithium's nephrogenic diabetes insipidus effects, it can be used to help treat the syndrome of inappropriate antidiuretic hormone hypersecretion (SIADH). It was also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.[60]

The active principle in these salts is the lithium ion Li+. Although this ion has a smaller diameter than either Na+ or K+, in a watery environment like the cytoplasmic fluid, Li+ binds to the oxygen atoms of water, making it effectively larger than either Na+ or K+ ions. How Li+ works in the central nervous system is still a matter of debate. Li+ elevates brain levels of tryptophan, 5-HT (serotonin), and 5-HIAA (a serotonin metabolite). Serotonin is related to mood stability. Li+ also reduces catecholamine activity in the brain (associated with brain activation and mania), by enhancing reuptake and reducing release. Therapeutically useful amounts of lithium (~ 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.

Common side effects of lithium treatment include muscle tremors, twitching, ataxia[61] and hypothyroidism. Long term use is linked to hyperparathyroidism[62], hypercalcemia (bone loss), hypertension, damage of tubuli in the kidney, nephrogenic diabetes insipidus (polyuria and polydipsia) and/or glomerular damage - even to the point of uremia[63], seizures[64] and weight gain.[65] Some of the side-effects are a result of the increased elimination of potassium.

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

According to a study in 2009 at Oita University in Japan and published in the British Journal of Psychiatry, communities whose water contained larger amounts of lithium had significantly lower suicide rates[66][67][68][69] but did not address whether lithium in drinking water causes the negative side effects associated with higher doses of the element.[70]

Other uses

The red lithium flame leads to lithium's use in flares and pyrotechnics

Precautions

Lithium ingots with a thin layer of black oxide tarnish

Lithium metal is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) initially 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.[74]

NFPA 704
NFPA 704
safety square
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 2: Must be moderately heated or exposed to relatively high ambient temperature before ignition can occur. Flash point between 38 and 93 °C (100 and 200 °F). E.g. diesel fuelInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
3
2
2
Fire diamond for lithium metal

Regulation

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

Carriage and shipment of some kinds of lithium batteries may be prohibited aboard certain types of transportation (particularly aircraft) because of the ability of most types of lithium batteries to fully discharge very rapidly when short-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.[75]

See also

Template:Wikipedia-Books

References

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  2. ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ a b c Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
  4. ^ Li(–1) has been observed in the gas phase; see R. H. Sloane; H. M. Love (1947). "Surface Formation of Lithium Negative Ions". Nature. 159: 302–303. doi:10.1038/159302a0.
  5. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
  6. ^ Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
  7. ^ File:Binding energy curve - common isotopes.svg shows binding energies of stable nuclides graphically; the source of the data-set is given in the figure background.
  8. ^ Numerical data from: Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal. 591: 1220–1247. doi:10.1086/375492. Graphed at File:SolarSystemAbundances.jpg
  9. ^ a b c d e f g Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements: A Reference Guide. Westport, Conn.: Greenwood Press. ISBN 0-313-33438-2.
  10. ^ a b c d e f Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. ISBN 0198503415.
  11. ^ Tuoriniemi, J; Juntunen-Nurmilaukas, K; Uusvuori, J; Pentti, E; Salmela, A; Sebedash, A (2007). "Superconductivity in lithium below 0.4 millikelvin at ambient pressure". Nature. 447 (7141): 187–9. doi:10.1038/nature05820. PMID 17495921.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  12. ^ Struzhkin, Vv; Eremets, Mi; Gan, W; Mao, Hk; Hemley, Rj (2002). "Superconductivity in dense lithium". Science. 298 (5596): 1213–5. doi:10.1126/science.1078535. PMID 12386338.{{cite journal}}: CS1 maint: multiple names: authors list (link)
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