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2.3 Lithium compounds
3 History and etymology
6.1 Medical use
6.2 Other uses
Lithium (pronounced /ˈlɪθiəm/, LITH-ee-əm) 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 three. 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 petroleum. When cut open, lithium exhibits a metallic luster, but contact with oxygen quickly turns it back to a dull silvery gray color. Lithium in its elemental state is highly flammable.
According to theory, lithium was one of the few elements synthesized in the Big Bang. Since its current estimated abundance in the universe is vastly less than that predicted by physical theories, the processes by which new lithium is created and destroyed, and the true value of its abundance, continue to be active matters of study in astronomy. 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). Though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.
Due to its high reactivity it 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, though 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.
Figure. 0. Silvery white (seen here in oil)
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 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.
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. It also has a low density (approximately 0.534 g/cm3) and thus will float on water, with which it reacts easily. This reaction is energetic, forming hydrogen gas and lithium hydroxide in aqueous solution. Due to its reactivity with water, lithium is usually stored in mineral oil or kerosene.
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 and at higher temperatures (more than 9 kelvin) at very high pressures (over 200,000 atmospheres) 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); 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.
Figure. 1. Lithium pellets (covered in white lithium hydroxide)
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).
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.
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).
2.3 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.
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. 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 both 6Li and 7Li are produced in stars, but are thought to be burned as fast as it is produced. 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. 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.
3. History and etymology
Petalite (LiAlSi4O10, which is lithium aluminium silicate) was first discovered in 1800 by the Brazilian chemist José Bonifácio de Andrade e Silva, who discovered this mineral 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 "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". Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite. In 1818, Christian Gmelin was the first man 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. 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 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 he performed an