Lithium

Lithium (pronounced /ˈlɪθiəm/) is a chemical element with the symbol Li and atomic number 3. It is a soft alkali metal with a silver-white color. Under standard conditions, it is the lightest metal and the least dense solid element. Like all alkali metals, lithium is highly reactive, corroding quickly in moist air to form a black tarnish. For this reason, lithium metal is typically stored under the cover of oil.

According to theory, Lithium (mostly ⁷Li) was one of the few elements synthesized in the Big Bang, although its quantity has vastly decreased. The reasons for its disappearance and the processes by which new lithium is created continue to be important matters of study in astronomy. Lithium is the 33rd most abundant element on Earth,^[1] but due to its high reactivity 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 biological function in humans. Nevertheless, the neurological effect of the lithium ion Li⁺ makes some lithium salts useful as a class of mood stabilizing drugs. Lithium and its compounds have several other commercial 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 splitting of lithium atoms was the first man-made form of nuclear reaction, and lithium deuteride serves as the fusion fuel in staged thermonuclear weapons.

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History and etymology

Petalite (lithium aluminum silicate) was first described in 1800 by the Brazilian scientist José Bonifácio de Andrade e Silva, who discovered the mineral in a Swedish iron mine on the island of Utö. However, it was not until 1817 that Johan August Arfwedson, then a trainee in the laboratory of Jöns Jakob Berzelius, discovered the presence of a new element while analyzing petalite ore. The element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less water soluble and had a larger capacity to neutralize acid. Berzelius gave the alkaline material the name "lithion", from the Greek λιθoς (lithos, "stone"), to reflect its discovery in a mineral, as opposed to sodium and potassium which had been discovered in plant tissue; its name would later be standardized as "lithium". Arfwedson later showed that this same element was present in the mineral ores spodumene and lepidolite. In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color in flame. However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.^[2][3][4]

The element was not isolated until 1821, when William Thomas Brande performed electrolysis on lithium oxide, a process which had previously been employed by Sir Humphry Davy to isolate potassium and sodium.^[3][5]Brande also described pure salts of lithium, such as the chloride, and performed an estimate of its atomic weight. In 1855, Robert Bunsen and Augustus Matthiessen produced large quantities of the metal by electrolysis of lithium chloride. Commercial production of lithium metal began in 1923 by the German company Metallgesellschaft AG through the electrolysis of a molten mixture of lithium chloride and potassium chloride.^[2][6]

Properties

Like other alkali metals, lithium has a single valence electron which it will readily lose to form a cation, indicated by the element's low electronegativity. As a result, lithium is easily deformed, highly reactive, and has lower melting and boiling points than most metals. These and many other properties attributable to alkali metals' weakly-held valence electron are most distinguished in lithium, as it possesses the smallest atomic radius and thus the highest electronegativity of the alkali group. In addition, lithium has a diagonal relationship with magnesium, an element of similar atomic and ionic radius. Chemical resemblances between the two metals include the formation of a nitride in N₂, the formation of an oxide when burnt in O₂, salts with similar solubilities, and thermally-unstable carbonates and nitrides.^[7]

Lithium is soft enough to be cut with a knife, though this is more difficult than cutting sodium. The fresh metal has a silvery-white color which only remains untarnished in dry air.^[7] Lithium has about half the density of water, giving solid sticks of lithium metal the odd heft of a light-to-medium wood like pine. The metal floats highly in hydrocarbons; in the laboratory, jars of lithium are typically composed of black-coated sticks held down in hydrocarbon mechanically by the jar's lid and other sticks.

Lithium is greatly heat-resistant, possessing a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium has also been found to be superconductive below 400 μK. This finding paves the way for further study of superconductivity, as lithium's atomic lattice is the simplest of all metals.

Chemistry

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).^[7]

When placed over a flame, lithium gives off a striking crimson color, but when it burns strongly, the flame becomes a brilliant white. Lithium will ignite and burn in oxygen when exposed to water or water vapours. It is the only metal that reacts with nitrogen at room temperature.

Lithium metal is flammable and potentially explosive when exposed to air and especially water, though it is far less dangerous than other alkali metals in this regard. The lithium-water reaction at normal temperatures is brisk but not violent. Lithium fires are difficult to extinguish, requiring special chemicals designed to smother them (see sodium for details).

See also: Lithium compounds

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). 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 7.58043x10^-23 s.

⁷Li is one of the primordial elements or, more properly, primordial isotopes, produced in Big Bang nucleosynthesis (a small amount of ⁶Li is also produced in stars). Lithium isotopes fractionate substantially during a wide variety of natural processes, including mineral formation (chemical precipitation), metabolism, and ion exchange. Lithium ion substitutes for magnesium and iron in octahedral sites in clay minerals, where 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.

Natural occurrence

See also Lithium minerals.

Lithium is widely distributed on Earth and is the 33rd most abundant element;^[1] however, it does not naturally occur in elemental form due to its high reactivity. Estimates for crustal content range from 20 to 70 ppm by weight.^[7] 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 mineral sources for the element.^[7]

Applications

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

It is an important ingredient in cathode materials, used in rechargeable and single-use batteries because of its high electrochemical potential, light weight, and high current density.

Large quantities of lithium are also used in the manufacture of organolithium reagents, especially n-butyllithium which has many uses in fine chemical and polymer synthesis.

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 can also be used to augment other antidepressant drugs. It is also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.

The active principle in these salts is the lithium ion Li⁺, which having a smaller diameter, can easily displace K⁺ and Na⁺ and even Ca⁺², in spite of its greater charge, occupying their sites in several critical neuronal enzymes and neurotransmitter receptors. Although Li⁺ cannot displace Mg²⁺ and Zn²⁺, because of these ions small size and greater charge (higher charge density, hence stronger bonding), when Mg⁺² or Zn⁺² are present in low concentrations, and Li⁺ is present in high concentrations, the latter can occupy sites normally occupied by Mg⁺² or Zn⁺² in various enzymes. Therapeutically useful amounts of lithium (0.6 to 1.2 mmol/l) are only slightly lower than toxic amounts (>1.5 mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity. Therefore, in theory, coadministration of 400 IU vitamin D, 1 g magnesium citrate (not the insoluble oxide or carbonate), 15 mg Zn (as gluconate or piccolinate, not the insoluble oxide) and 1 pill of vitamin B complex a day, should potentiate the effect of Li,^{[citation needed]} in some cases allowing for the reduction of the therapeutic range to 0.5 to 0.9 mmol/l, of the daily dose of lithium carbonate and of the risk of toxicity.

Common side effects include muscle tremors, twitching, ataxia, hyperparathyroidism (bone loss, hypercalcemia, hypertension, etc,), kidney damage, nephrogenic diabetes insipidus (polyuria and polydipsia) and seizures. Many of the side-effects are a result caused by the increased elimination of potassium.

Other uses

• Lithium chloride and lithium bromide are extremely hygroscopic and frequently used as desiccants.
• Lithium stearate is a common all-purpose high-temperature lubricant.
• Lithium is an alloying agent used to synthesize organic compounds.
• Lithium is used as a flux to promote the fusing of metals during welding and soldering. It also eliminates the forming of oxides during welding by absorbing impurities. This fusing quality is also important as a flux for producing ceramics, enamels, and glass.
• Lithium is sometimes used in glasses and ceramics including the glass for the 200-inch (5.08 m) telescope at Mt. Palomar.
• Alloys of the metal with aluminium, cadmium, copper, and manganese are used to make high performance aircraft parts.
• Lithium niobate is used extensively in telecommunication products, such as mobile phones and optical modulators, for such components as resonant crystals. Lithium products are currently used in more than 60 percent of mobile phones.^[8]
• The high non-linearity of lithium niobate also makes a good choice for non-linear optics applications.
• 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.
• Metallic lithium and its complex hydrides such as e.g. Li[AlH₄] are considered as high energy additives to rocket propellants^[3].
• Lithium peroxide, lithium nitrate, lithium chlorate and lithium perchlorate are used and thought of as oxidizers in both rocket propellants and oxygen candles to supply submarines and space capsules with oxygen.^[9]
• Lithium will be used to produce tritium in magnetically confined nuclear fusion reactors using deuterium and tritium as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium. ⁶Li + n → ⁴He + ³H. Various means of doing this will be tested at the ITER reactor being built at Cadarache, France.
• Lithium is used as a source for alpha particles, or helium nuclei. When ⁷Li is bombarded by accelerated protons, ⁸Be is formed, which undergoes spontaneous fission to form two alpha particles. This was the first man-made nuclear reaction, produced by Cockroft and Walton in 1929.
• Lithium hydroxide (LiOH) is an important compound of lithium obtained from lithium carbonate (Li₂CO₃). It is a strong base, and when heated with a fat, it produces a lithium soap. Lithium soap has the ability to thicken oils and so is used commercially to manufacture lubricating greases. It is also an efficient and lightweight purifier of air. In confined areas, such as aboard spacecraft and submarines, the concentration of carbon dioxide can approach unhealthy or toxic levels. Lithium hydroxide absorbs the carbon dioxide from the air by reacting with it to form lithium carbonate. Any alkali hydroxide will absorb CO₂, but lithium hydroxide is preferred, especially in spacecraft applications, because of the low formula weight conferred by the lithium. Even better materials for this purpose include lithium peroxide (Li₂O₂) that, in presence of moisture, not only absorb carbon dioxide to form lithium carbonate, but also release oxygen. E.g. 2 Li₂O₂ + 2 CO₂ → 2 Li₂CO₃ + O₂.
• Lithium metal is used as a reducing agent in some types of methamphetamine production, particularly in illegal amateur “meth labs.”

Production

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

The metal is produced electrolytically from a mixture of fused lithium and potassium chloride. In 1998 it was about US$ 43 per pound ($95 per kg).^[10]

Chile is currently the leading lithium metal producer in the world, with Argentina next. Both countries recover the lithium from brine pools. In the United States lithium is similarly recovered from brine pools in Nevada.^[11]

China may emerge as a significant producer of brine-based lithium carbonate towards the end of this decade. Potential capacity of up to 45,000 tonnes per year could come on-stream if projects in Qinghai province and Tibet proceed.^{[citation needed]}

Precautions

Lithium metal, due to its alkaline tarnish, is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) can irritate the nose and throat; higher exposure to lithium can cause a build-up of fluid in the lungs, leading to pulmonary edema. The metal itself is usually less a handling hazard than the caustic hydroxide produced when it is in contact with moisture. Lithium should be stored in a non-reactive compound such as naphtha or a hydrocarbon.

Regulation

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. However, the effectiveness of such restrictions in controlling illegal production of methamphetamine remains indeterminate and controversial.

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. However, most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents.

References

1. ^ ^{a b} Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements : A Reference Guide. Westport, Conn.: Greenwood Press, 47-50. ISBN 0-313-33438-2. 
2. ^ ^{a b} Winter, Mark J. Chemistry : Periodic Table : lithium : historical information. Web Elements. Retrieved on August 19, 2007.
3. ^ ^{a b} (2004) Encyclopedia of the Elements: Technical Data - History - Processing - Applications. Wiley, 287-300. ISBN 978-3527306664. 
4. ^ http://www.vanderkrogt.net/elements/elem/li.html
5. ^ http://www.diracdelta.co.uk/science/source/t/i/timeline/source.html
6. ^ http://www.echeat.com/essay.php?t=29195
7. ^ ^{a b c d e} Kamienski et al. "Lithium and lithium compounds". Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons, Inc. Published online 2004. doi:10.1002/0471238961.1209200811011309.a01.pub2
8. ^ Spring, Martin. "Two ways to play the lithium boom", MoneyWeek, 2007-01-08. Retrieved on 2007-08-19. 
9. ^ K. Ernst-Christian (2004). "Special Materials in Pyrotechnics: III. Application of Lithium and its Compounds in Energetic Systems". Propellants, Explosives, Pyrotechnics 29 (2): 67-80. doi:10.1002/prep.200400032.
10. ^ Ober, Joyce A. Lithium (pdf) 77-78. United States Geological Survey. Retrieved on August 19, 2007.
11. ^ Lithium. Los Alamos National Laboratory (December 15, 2003). Retrieved on August 19, 2007.

See also

• Lithium compounds