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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 point1603 K ​(1330 °C, ​2426 °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 (/[invalid input: 'icon']ˈlɪθiəm/ LITH-ee-əm) (from lithos, Greek for stone) is a soft, silver-white metal with symbol Li and atomic number 3. It belongs 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.[7] 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 deuteride serves as a fusion fuel in staged thermonuclear weapons.

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. Nonvital functions have not been ruled out. The lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood-stabilizing drug in the treatment of bipolar disorder, due to neurological effects of the ion in the human body.

Properties

Main article: Alkali metal

Atomic and physical

alt1
alt2
Lithium pellets covered in white lithium hydroxide (left) and ingots with a thin layer of black oxide tarnish (right)

Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[8] 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).[8]

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.[8] 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.[9]

It is the lightest metal in the periodic table, so light that it can float on water and even on oil, and it is one of three metals that can (the other two are sodium and potassium). It has a very low density, of approximately 0.534 g/cm3, which gives sticks of the metal a similar heft to dowels of a medium density wood, such as pine. It floats on water but also reacts with it.[8]

Lithium floating in oil

It is the least dense of all elements that are not gases at room temperature. The next lightest element is over 60% more dense (potassium, at 0.862 g/cm3). Furthermore, aside from helium and hydrogen, it is the least dense element in a solid or liquid state, being only 2/3 as dense as liquid nitrogen (0.808 g/cm3).[note 1][10]

Lithium's coefficient of thermal expansion is twice that of aluminium and almost four times that of iron.[11] It has the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[12] and at higher temperatures (more than 9 K) at very high pressures (>20 GPa)[13] 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.[14] Multiple allotropic forms have been reported for lithium at high pressures.[15]

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.[8] 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.[16] 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).[17]

Hexameric structure of the n-butyllithium fragment in a crystal

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.[18] Lithium is inflammable, 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, 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 Class D type (see Types of extinguishing agents). Lithium is the only metal which reacts with nitrogen under normal conditions.[19][20]

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.[17][21] The metal reacts with hydrogen gas at high temperatures to produce lithium hydride (LiH).[22]

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.[23]

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).[8][16][24] 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.[25] 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.[7] 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.[26]

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.[27] 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.[28] While lithium is created in stars during the Stellar nucleosynthesis, it is further burnt. 7Li can also be generated in carbon stars.[29]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[30] 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.[31]

Occurrence

Lithium is about as common as chlorine in the Earth's upper continental crust, on a per-atom basis.

Astronomical

Main article: Nucleosynthesis

According to modern cosmological theory, lithium—as both of its stable isotopes lithium-6 and lithium-7—was among the 3 elements synthesized in the Big Bang. Though the amount of lithium generated in Big Bang nucleosynthesis is dependent upon the number of photons per baryon, for accepted values the lithium abundance can be calculated, and there is a "cosmological lithium discrepancy" in the Universe: older stars seem to have less lithium than they should, and some younger stars have far more. The lack of lithium in older stars is apparently caused by the "mixing" of lithium into the interior of stars, where it is destroyed.[32] 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.[16]

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.[33]

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.[16][34][35] 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.[16]

Terrestrial

Lithium mine production (2011) and reserves in tonnes[36]
Country Production Reserves[note 2]
 Argentina 3,200 850,000
 Australia 9,260 970,000
 Brazil 160 64,000
 Canada (2010) 480 180,000
 Chile 12,600 7,500,000
 People's Republic of China 5,200 3,500,000
 Portugal 820 10,000
 Zimbabwe 470 23,000
World total 34,000 13,000,000
See also: Lithium minerals

Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity.[8] 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),[37][38] or 25 micromolar;[39] higher concentrations approaching 7 ppm are found near hydrothermal vents.[38]

Estimates for crustal content range from 20 to 70 ppm by weight.[17] 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.[17] Another significant mineral of lithium is lepidolite.[40] A newer source for lithium is hectorite clay, the only active development of which is through the Western Lithium Corporation in the United States.[41] At 20 mg lithium per kg of Earth's crust,[42] lithium is the 25th most abundant element. Nickel and lead have about the same abundance.

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."[43]

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) and the highest annual production (8,800 tonnes). Other major suppliers include Australia, Argentina and China.[36][44] Other estimates put Chile's reserve base (7,520 million tonnes) above that of Argentina (6 million).[45]

In June 2010, the New York Times reported that American geologists were conducting ground surveys on dry salt lakes in western Afghanistan believing that large deposits of lithium are located there. "Pentagon officials said that their initial analysis at one location in Ghazni Province showed the potential for lithium deposits as large of those of Bolivia, which now has the world's largest known lithium reserves."[46] 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."[47]

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.[38] Marine organisms tend to bioaccumulate lithium more than terrestrial ones.[48] It is not known whether lithium has a physiological role in any of these organisms,[38] 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.[citation needed] Observational studies in Japan, reported in 2011, suggested that naturally occurring lithium in drinking water may increase human lifespan.[49]

History

Johan August Arfwedson is credited with the discovery of lithium in 1817

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.[50][51][52] 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.[53][54][55] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[56] Berzelius gave the alkaline material the name "lithion/lithina", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium which was known partly for its high abundance in animal blood. He named the metal inside the material as "lithium".[8][51][55]

Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[51] In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame.[51] However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.[51][55][57] 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.[16][57][58][59] 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).[60] In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[51] 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.[51][61]

The production and use of lithium underwent several drastic changes in history. The first major application of lithium became high temperature grease for aircraft engines or similar applications in World War II and shortly after. This small market was supported by several small mining operations mostly in the United States. The demand for lithium increased dramatically during the Cold War with the production of nuclear fusion weapons. Both lithium-6 and lithium-7 produce tritium when irradiated by neutrons, and are thus useful for the production of tritium by itself, as well as a form of solid fusion fuel used inside hydrogen bombs in the form of lithium deuteride. The United States became the prime producer of lithium in the period between the late 1950s and the mid 1980s. At the end the stockpile of lithium was roughly 42,000 tonnes of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75%.[62]

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.[63][63] 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.[62] But in the mid-1990s, several companies started to extract lithium from brine which proved to be a less expensive method than underground or even open pit mining. Most of the mines closed or shifted their focus to other materials as only the ore from zoned pegmatites could be mined for a competitive price. For example, the US mines near Kings Mountain, North Carolina closed before the turn of the century. The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.[64] With the surge of lithium demand in batteries in to 2000s, new companies have expanded brine extraction efforts to meet the rising demand.[65][66]

Production

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Satellite images of the Salar del Hombre Muerto, Argentina (left), and Uyuni, Bolivia (right), salt flats are rich in lithium. The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in the left image).

Since the end of World War II lithium production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride. In 1998 it was about 95 US$ / kg (or 43 US$/pound).[67]

Worldwide identified reserves of lithium in 2008 were estimated by the US Geological Survey as 13 million tonnes.[36] 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.[68] 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.[69] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tonnes of lithium.[69][70]

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.[71] 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.[72]

Applications

Usage of lithium in the USA in 2010[73]
  Ceramics and glass (29%)
  Batteries (27%)
  Lubricating greases (12%)
  Continuous casting (5%)
  Air treatment (4%)
  Polymers (3%)
  Primary aluminum production (2%)
  Pharmaceuticals (2%)
  Other (16%)

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. Lithium oxides are a component of ovenware. Worldwide, this is the single largest use for lithium compounds (see chart).

Electrical and electronics

In the later years of the 20th century lithium became important as an anode material. Used in lithium-ion batteries because of its high electrochemical potential, a typical cell can generate approximately 3 volts, compared with 2.1 volts for lead/acid or 1.5 volts for zinc-carbon cells. Because of its low atomic mass, it also has a high charge- and power-to-weight ratio. Lithium batteries are disposable (primary) batteries with lithium or its compounds as an anode. Lithium batteries are not to be confused with lithium-ion batteries, which are high energy-density rechargeable batteries. Other rechargeable batteries include the lithium-ion polymer battery, lithium iron phosphate battery, and the nanowire battery. New technologies are constantly being announced.

Lubricating greases

The third most common use of lithium is in greases. Lithium hydroxide is a strong base, and when heated with a fat it produces a soap made of lithium stearate. Lithium soap has the ability to thicken oils, and it is used to manufacture all-purpose, high-temperature lubricating greases.[68][74][75]

Other chemical and industrial uses

Lithium use in flares and pyrotechnics is due to its rose-red flame

Inorganic lithium salts

Lithium chloride and lithium bromide are extremely hygroscopic and are used as desiccants.[68] Lithium hydroxide (LiOH) is an important compound of lithium obtained from lithium carbonate (Li2CO3).

Metallic lithium and its complex hydrides, such as Li[AlH4], are used as high energy additives to rocket propellants.[16]

Air purification

Lithium peroxide, lithium nitrate, lithium chlorate and lithium perchlorate are used as oxidizers in rocket propellants, and also in oxygen candles that supply submarines with oxygen.[76]

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 reacting with it to form lithium carbonate, and is preferred over other alkaline hydroxides for its low weight. Lithium peroxide (Li2O2) in presence of moisture not only absorbs carbon dioxide to form lithium carbonate, but also releases oxygen.[77][78] For example:

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.

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.[79] 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.[80] Lithium fluoride is sometimes used in focal lenses of telescopes.[68][81]

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.[82]

Elemental lithium and reagents prepared from it

Because of its specific heat capacity, the highest of all solids, lithium metal is often used in coolants for heat transfer applications.[68]

The launch of a torpedo using lithium as fuel

The Mark 50 Torpedo stored chemical energy propulsion system (SCEPS) uses a small tank of sulfur hexafluoride gas which is sprayed over a block of solid lithium. The reaction generates enormous heat which is used to generate steam from seawater. The steam propels the torpedo in a closed Rankine cycle.[83]

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] Lithium compounds are also used as pyrotechnic colorants and oxidizers in red fireworks and flares.[68][85]

Lithium metal is also used in the pharmaceutical and fine-chemical industry in the manufacture of organolithium reagents, which are used both as strong bases and as reagents for the formation of carbon-carbon bonds. Organolithium compounds are also used in polymer synthesis as catalysts/initiators[86] in anionic polymerization of unfunctionalized olefins.[87][88][89] Lithium is used in the preparation of organolithium compounds, which are in turn very reactive and are the basis of many synthetic applications.[90]

Nuclear

Lithium-6 is valued as a source material for tritium production and as a neutron absorber in nuclear fusion. Natural lithium contains about 7.5% lithium-6 from which large amounts of lithium-6 have been produced by isotope separation for use in nuclear weapons.[91] Lithium-7 gained interest for use in nuclear reactor coolants.[92]

Lithium deuteride was used as fuel in the Castle Bravo nuclear device.

Lithium deuteride was the fusion fuel of choice in early versions of the hydrogen bomb. When bombarded by neutrons, both 6Li and 7Li produce tritium — this reaction, which was not fully understood when hydrogen bombs were first tested, was responsible for the runaway yield of the Castle Bravo nuclear test. Tritium fuses with deuterium in a fusion reaction that is relatively easy to achieve. Although details remain secret, lithium-6 deuteride still apparently plays a role in modern nuclear weapons, as a fusion material.[93]

Lithium fluoride, 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][94]

In conceptualized nuclear fusion power plants, lithium will be used to produce tritium in magnetically confined reactors using deuterium and tritium as the fuel. Tritium does not occur naturally[citation needed] and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium:

6Li + n → 4He + 3T.

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.[95][96] (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 pharmacology

In the treatment of bipolar disorder, lithium compounds continue to be the standard against which newer medications are measured. 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+, although detailed mechanisms are debated.

Precautions

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 (red): no hazard codeInstability 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
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.[97]

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.[98]

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.[99][100] 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.[101][102]

See also

Notes

  1. ^ Densities for all the gaseous elements can be obtained at Airliquide.com
  2. ^ a b 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.

References

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