Lithium floating in oil
Spectral lines of lithium
|Name, symbol||lithium, Li|
|Lithium in the periodic table|
|Standard atomic weight||6.94(1)|
|Element category||alkali metal|
|Group, block||group 1 (alkali metals), s-block|
|Electron configuration||[He] 2s1|
|per shell||2, 1|
|Melting point||453.65 K (180.50 °C, 356.90 °F)|
|Boiling point||1603 K (1330 °C, 2426 °F)|
|Density (near r.t.)||0.534 g·cm−3 (at 0 °C, 101.325 kPa)|
|Liquid density||at m.p.: 0.512 g·cm−3|
|Critical point||3220 K, 67 MPa (extrapolated)|
|Heat of fusion||3.00 kJ·mol−1|
|Heat of vaporization||136 kJ·mol−1|
|Molar heat capacity||24.860 J·mol−1·K−1|
|Oxidation states||+1 (a strongly basic oxide)|
|Electronegativity||0.98 (Pauling scale)|
|Ionization energies||1st: 520.2 kJ·mol−1
2nd: 7298.1 kJ·mol−1
3rd: 11815.0 kJ·mol−1
|Atomic radius||empirical: 152 pm|
|Covalent radius||128±7 pm|
|Van der Waals radius||182 pm|
|Crystal structure||body-centered cubic (bcc)|
|Speed of sound||thin rod: 6000 m·s−1 (at 20 °C)|
|Thermal expansion||46 µm·m−1·K−1 (at 25 °C)|
|Thermal conductivity||84.8 W·m−1·K−1|
|Electrical resistivity||92.8 nΩ·m (at 20°C)|
|Young's modulus||4.9 GPa|
|Shear modulus||4.2 GPa|
|Bulk modulus||11 GPa|
|Discovery||Johan August Arfwedson (1817)|
|First isolation||William Thomas Brande (1821)|
|Most stable isotopes|
|6Li content may be as low as 3.75% in
natural samples. 7Li would therefore
have a content of up to 96.25%.
Lithium (from Greek: λίθος lithos, "stone") is a chemical element with symbol Li and atomic number 3. It is a soft, silver-white metal belonging to the alkali metal group of chemical elements. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable. For this reason, it is typically stored in mineral oil. When cut open, lithium exhibits a metallic luster, but contact with moist air corrodes the surface quickly to a dull silvery gray, then black tarnish. Because of its high reactivity, lithium never occurs freely in nature, and instead, only appears in compounds, which are usually ionic. Lithium occurs in a number of pegmatitic minerals, but due to its solubility as an ion, is present in ocean water and is commonly obtained from brines and clays. On a commercial scale, lithium is isolated electrolytically from a mixture of lithium chloride and potassium chloride.
The nuclei of lithium verge on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative nuclear instability, lithium is less common in the solar system than 25 of the first 32 chemical elements even though the nuclei are very light in atomic weight. For related reasons, lithium has important links to nuclear physics. The transmutation of lithium atoms to helium in 1932 was the first fully man-made nuclear reaction, and lithium-6 deuteride serves as a fusion fuel in staged thermonuclear weapons.
Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, lithium batteries and lithium-ion batteries. These uses consume more than half of lithium production.
Trace amounts of lithium are present in all organisms. The element serves no apparent vital biological function, since animals and plants survive in good health without it. Non-vital functions have not been ruled out. The lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood-stabilizing drug in the treatment of bipolar disorder, due to neurological effects of the ion in the human body.
- 1 Properties
- 2 Occurrence
- 3 History
- 4 Production
- 5 Uses
- 6 Precautions
- 7 See also
- 8 Notes
- 9 References
- 10 External links
Atomic and physical
Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation. 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).
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. 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.
Lithium has a very low density of 0.534 g/cm3, comparable with that of pine wood. It is the least dense of all elements that are solids at room temperature, the next lightest solid element (potassium, at 0.862 g/cm3) being more than 60% denser. Furthermore, apart from helium and hydrogen, it is less dense than any liquid element, being only 2/3 as dense as liquid nitrogen (0.808 g/cm3).[note 1] Lithium can float on the lightest hydrocarbon oils and is one of only three metals that can float on water, the other two being sodium and potassium.
Lithium's coefficient of thermal expansion is twice that of aluminium and almost four times that of iron. It has the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure and at higher temperatures (more than 9 K) at very high pressures (>20 GPa) 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. Multiple allotropic forms have been reported for lithium at high pressures.
Lithium has a specific heat capacity of 3.58 kilojoules per kilogram-Kelvin, the highest of all solids. Because of this, lithium metal is often used in coolants for heat transfer applications.
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. Because of its reactivity with water, lithium is usually stored under cover of a hydrocarbon, often petroleum jelly. Though the heavier alkali metals can be stored in more dense substances, such as mineral oil, lithium is not dense enough to be fully submerged in these liquids. 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).
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. Lithium is flammable, and it is potentially explosive when exposed to air and especially to water, though less so than the other alkali metals. The lithium-water reaction at normal temperatures is brisk but nonviolent, as the hydrogen produced will not ignite on its own. As with all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers (Class D type). Lithium is the only metal which reacts with nitrogen under normal conditions.
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
2) when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides. The metal reacts with hydrogen gas at high temperatures to produce lithium hydride (LiH).
Other known binary compounds include the halides (LiF, LiCl, LiBr, LiI), and the sulfide (Li
2S), the superoxide (LiO
2), carbide (Li
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. LiHe, a very weakly interacting van der Waals compound, has been detected at very low temperatures.
Lithium has also been found to exhibit ferromagnetism in its gaseous form, under certain conditions.
Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (92.5% natural abundance). Both natural isotopes have anomalously low nuclear binding energy per nucleon compared to the next lighter and heavier elements, helium and beryllium, which means that alone among stable light elements, lithium can produce net energy through nuclear fission. The two lithium nuclei have lower binding energies per nucleon than any other stable nuclides other than deuterium and helium-3. As a result of this, though very light in atomic weight, lithium is less common in the Solar System than 25 of the first 32 chemical elements. 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.
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. 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. While lithium is created in stars during the Stellar nucleosynthesis, it is further burnt. 7Li can also be generated in carbon stars.
Lithium isotopes fractionate substantially during a wide variety of natural processes, 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.
Nuclear weapons manufacture and other nuclear physics uses are a major source of artificial lithium fractionation, with the light isotope 6Li being retained by industry and military stockpiles to such an extent as to slightly but measurably change the 6Li to 7Li ratios even in natural sources, such as rivers. This has led to unusual uncertainty in the standardized atomic weight of lithium, since this quantity depends on the natural abundance ratios of these naturally-occurring stable lithium isotopes, as they are available in commercial lithium mineral sources.
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. 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.
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.
Lithium is also found in brown dwarf substellar objects and certain anomalous orange stars. Because lithium is present in cooler, less-massive brown dwarfs, but is destroyed in hotter red dwarf stars, its presence in the stars' spectra can be used in the "lithium test" to differentiate the two, as both are smaller than the Sun. 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.
|People's Republic of China||5,200||3,500,000|
Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity. 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), or 25 micromolar; higher concentrations approaching 7 ppm are found near hydrothermal vents.
Estimates for the Earth's crustal content range from 20 to 70 ppm by weight. 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. Another significant mineral of lithium is lepidolite. A newer source for lithium is hectorite clay, the only active development of which is through the Western Lithium Corporation in the United States. At 20 mg lithium per kg of Earth's crust, lithium is the 25th most abundant element.
According to the Handbook of Lithium and Natural Calcium, "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively few of them are of actual or potential commercial value. Many are very small, others are too low in grade."
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.
In June 2010, the New York Times reported that American geologists were conducting ground surveys on dry salt lakes in western Afghanistan believing that large deposits of lithium are located there. "Pentagon officials said that their initial analysis at one location in Ghazni Province showed the potential for lithium deposits as large as those of Bolivia, which now has the world's largest known lithium reserves." 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."
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. Marine organisms tend to bioaccumulate lithium more than terrestrial ones. It is not known whether lithium has a physiological role in any of these organisms, 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. Observational studies in Japan, reported in 2011, suggested that naturally occurring lithium in drinking water may increase human lifespan.
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. 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. This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline. Berzelius gave the alkaline material the name "lithion/lithina", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium which was known partly for its high abundance in animal blood. He named the metal inside the material "lithium".
Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite. In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame. However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts. 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. 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). In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen. 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.
The production and use of lithium underwent several drastic changes in history. The first major application of lithium was in high-temperature lithium greases for aircraft engines or similar applications in World War II and shortly after. This use was supported by the fact that lithium-based soaps have a higher melting point than other alkali soaps, and are less corrosive than calcium based soaps. The small market for lithium soaps and the lubricating greases based upon them was supported by several small mining operations mostly in the United States.
The demand for lithium increased dramatically during the Cold War with the production of nuclear fusion weapons. Both lithium-6 and lithium-7 produce tritium when irradiated by neutrons, and are thus useful for the production of tritium by itself, as well as a form of solid fusion fuel used inside hydrogen bombs in the form of lithium deuteride. The United States became the prime producer of lithium in the period between the late 1950s and the mid-1980s. At the end, the stockpile of lithium was roughly 42,000 tonnes of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75%, which was enough to affect the measured atomic weight of lithium in many standardized chemicals, and even the atomic weight of lithium in some "natural sources" of lithium ion which had been "contaminated" by lithium salts discharged from isotope separation facilities, which had found its way into ground water.
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. 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. But in the mid-1990s, several companies started to extract lithium from brine which proved to be a less expensive method than underground or even open-pit mining. Most of the mines closed or shifted their focus to other materials as only the ore from zoned pegmatites could be mined for a competitive price. For example, the US mines near Kings Mountain, North Carolina closed before the turn of the 21st century.
The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007. With the surge of lithium demand in batteries in the 2000s, new companies have expanded brine extraction efforts to meet the rising demand.
Since the end of World War II lithium production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits. The metal is produced electrolytically from a mixture of fused 55% lithium chloride and 45% potassium chloride at about 450o C. In 1998 it was about 95 US$ / kg (or 43 US$/pound).
Worldwide identified reserves of lithium in 2008 were estimated by the US Geological Survey as 13 million tonnes. Deposits of lithium are found in South America throughout the Andes mountain chain. Chile is the leading lithium producer, followed by Argentina. Both countries recover the lithium from brine pools. In the United States lithium is recovered from brine pools in Nevada. 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. According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tonnes of lithium. A newly discovered deposit in Wyoming's Rock Springs Uplift is estimated at 228,000 tons. Additional deposits in the same formation were extrapolated to be as much as 18 million tons.
After an industry wide pricing reduction for lithium carbonate after the Great Financial Crisis, where major suppliers such as Sociedad Química y Minera (SQM) dropped pricing by 20% in light of incoming lithium resource developers and to further defend their market position, pricing in 2012 scaled up due to increased lithium demand. A 2012 Business Week article outlined the existing oligopoly in the lithium space, "SQM, controlled by billionaire Julio Ponce, is the second-largest, followed by Rockwood, which is backed by Henry Kravis’s KKR & Co., and Philadelphia-based FMC." Global consumption may jump to 300,000 metric tons a year by 2020 from about 150,000 tons in 2012, as demand for lithium batteries has been growing at about 25 percent a year, outpacing the 4 percent to 5 percent overall gain in lithium
A potential source is geothermal wells. Geothermal fluids carry leachates to the surface; recovery of lithium has been demonstrated in the field. As the lithium is separated by simple filtration techniques, the process and environmental costs are primarily that of the already-operating geothermal well; relative environmental impacts may thus be positive.
There are differing opinions about the potential growth of lithium production. A 2008 study concluded that "realistically achievable lithium carbonate production will be sufficient for only a small fraction of future PHEV and EV global market requirements", that "demand from the portable electronics sector will absorb much of the planned production increases in the next decade", and that "mass production of lithium carbonate is not environmentally sound, it will cause irreparable ecological damage to ecosystems that should be protected and that LiIon propulsion is incompatible with the notion of the 'Green Car'".
However, 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. 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.
On June 9, 2014, the Financialist publication, produced by the Credit Suisse company, stated that demand for lithium is growing at more than 12 percent a year; according to Credit Suisse, this rate exceeds projected availability by 25 percent. The publication compared the 2014 lithium situation with oil, whereby "higher oil prices spurred investment in expensive deepwater and oil sands production techniques"; that is, the price of lithium will continue to rise until more expensive production methods that can boost total output receive the attention of investors.
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. Lithium carbonate (Li2CO3) is generally used in this application: upon heating it converts to the oxide.
Electrical and electronics
In the later years of the 20th century, owing to its high electrode potential, lithium became an important component of the electrolyte and of one of the electrodes in batteries. Because of its low atomic mass, it has a high charge- and power-to-weight ratio. A typical lithium-ion battery can generate approximately 3 volts per cell, compared with 2.1 volts for lead-acid or 1.5 volts for zinc-carbon cells. Lithium-ion batteries, which are rechargeable and have a high energy density, should not be confused with lithium batteries, which are disposable (primary) batteries with lithium or its compounds as the anode. Other rechargeable batteries that use lithium include the lithium-ion polymer battery, lithium iron phosphate battery, and the nanowire battery.
The third most common use of lithium is in greases. Lithium hydroxide is a strong base and, when heated with a fat, produces a soap made of lithium stearate. Lithium soap has the ability to thicken oils, and it is used to manufacture all-purpose, high-temperature lubricating greases.
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).
Other chemical and industrial uses
Lithium chloride and lithium bromide are hygroscopic and are used as desiccants for gas streams. Lithium hydroxide and lithium peroxide are the salts most used in confined areas, such as aboard spacecraft and submarines, for carbon dioxide removal and air purification. Lithium hydroxide absorbs carbon dioxide from the air by forming lithium carbonate, and is preferred over other alkaline hydroxides for its low weight.
- 2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.
Some of the aforementioned compounds, as well as lithium perchlorate, are used in oxygen candles that supply submarines with oxygen. These can also include small amounts of boron, magnesium, aluminum, silicon, titanium, manganese, and iron.
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. 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. Lithium fluoride is sometimes used in focal lenses of telescopes.
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.
Organic and polymer chemistry
Organolithium compounds are widely used in the production of polymer and fine-chemicals. In the polymer industry, which is the dominant consumer of these reagents, alkyl lithium compounds are catalysts/initiators. in anionic polymerization of unfunctionalized olefins. For the production of fine chemicals, organolithium compounds function as strong bases and as reagents for the formation of carbon-carbon bonds. Organolithium compounds are prepared from lithium metal and alkyl halides.
Many other lithium compounds are used as reagents to prepare organic compounds. Some popular compounds include lithium aluminium hydride (LiAlH4), lithium triethylborohydride (LiBH(C2H5)3). n-Butyllithium (C4H9Li) and tert-butyl lithium (C4H9Li) are commonly used as extremely strong bases called superbase.
The Mark 50 Torpedo stored chemical energy propulsion system (SCEPS) uses a small tank of sulfur hexafluoride gas which is sprayed over a block of solid lithium. The reaction generates heat which is used to generate steam. The steam propels the torpedo in a closed Rankine cycle.
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. Lithium-7 gained interest for use in nuclear reactor coolants.
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.
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]
In conceptualized nuclear fusion power plants, lithium will be used to produce tritium in magnetically confined reactors using deuterium and tritium as the fuel. Naturally occurring tritium is extremely rare, and must be synthetically produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will fission the lithium to produce more tritium:
- 6Li + n → 4He + 3T.
Lithium is also used as a source for alpha particles, or helium nuclei. When 7Li is bombarded by accelerated protons 8Be is formed, which undergoes fission to form two alpha particles. This feat, called "splitting the atom" at the time, was the first fully man-made nuclear reaction. It was produced by Cockroft and Walton in 1932. (Nuclear reactions and human-directed nuclear transmutation had been accomplished as early as 1917, but by using natural radioactive bombardment with alpha particles).
In 2013, the US Government Accountability Office said the Lithium-7 critical to the operation of 65 out of 100 American nuclear reactors “places their ability to continue to provide electricity at some risk”. The problem stems from the decay of US nuclear infrastructure. The US shut down most of its machinery in 1963, given a huge surplus, mostly consumed during the twentieth century. The report said it would take five years and $10 million to $12 million.
Reactors that use lithium-7 heat water under high pressure and transfer heat through heat exchangers that are prone to corrosion. The reactors use lithium to counteract the corrosive effects of boric acid, which is added to the water to absorb excess neutrons.
Lithium is useful in the treatment of bipolar disorder. Lithium salts may also be helpful for related diagnoses, such as schizoaffective disorder and cyclic major depression. The active part of these salts is the lithium ion Li+. They may increase the risk of developing Ebstein's cardiac anomaly in infants born to women who take lithium during the first trimester of pregnancy.
|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.
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. 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.
- Lithium (medication)
- Lithium compounds
- Lithium-based grease
- Lithium-air battery
- Lithium-ion battery
- Densities for all the gaseous elements can be obtained at Airliquide.com
- 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)."
- 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.
- Numerical data from: Lodders, Katharina (July 10, 2003). "Solar System Abundances and Condensation Temperatures of the Elements" (PDF). The Astrophysical Journal (The American Astronomical Society) 591 (2): 1220–1247. Bibcode:2003ApJ...591.1220L. doi:10.1086/375492. Graphed at File:SolarSystemAbundances.jpg
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