Lithium

Lithium, 3Li

Freshly cut sample of lithium, with minimal oxides
Lithium
Pronunciation	/ˈlɪθiəm/ ⓘ ​(LITH-ee-əm)
Appearance	silvery-white
Standard atomic weight Ar°(Li)
	[6.938, 6.997][1] 6.94±0.06 (abridged)[2]
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
Group	group 1: hydrogen and alkali metals
Period	period 2
Block	s-block
Electron configuration	[He] 2s1
Electrons per shell	2, 1
Physical properties
Phase at STP	solid
Melting point	453.65 K ​(180.50 °C, ​356.90 °F)
Boiling point	1617 K ​(1344 °C, ​2451 °F)
Density (at 20° C)	0.5334 g/cm3[3]
when liquid (at m.p.)	0.512 g/cm3
Critical point	3220 K, 67 MPa (extrapolated)
Heat of fusion	3.00 kJ/mol
Heat of vaporization	136 kJ/mol
Molar heat capacity	24.860 J/(mol·K)
Specific heat capacity	3582.133 J/(kg·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 states	common: +1 −1[4]
Electronegativity	Pauling scale: 0.98
Ionization energies	1st: 520.2 kJ/mol 2nd: 7298.1 kJ/mol 3rd: 11815.0 kJ/mol
Atomic radius	empirical: 152 pm calculated: 167 pm
Covalent radius	128±7 pm
Van der Waals radius	182 pm
Spectral lines of lithium
Other properties
Natural occurrence	primordial
Crystal structure	​body-centered cubic (bcc) (cI2)
Lattice constant	a = 350.93 pm (at 20 °C)[3]
Thermal expansion	46.56×10−6/K (at 20 °C)[3]
Thermal conductivity	84.8 W/(m⋅K)
Electrical resistivity	92.8 nΩ⋅m (at 20 °C)
Magnetic ordering	paramagnetic
Molar magnetic susceptibility	+14.2×10−6 cm3/mol (298 K)[6]
Young's modulus	4.9[dubious – discuss] GPa
Shear modulus	4.2 GPa
Bulk modulus	11 GPa
Speed of sound thin rod	6000 m/s (at 20 °C)
Mohs hardness	0.6
Brinell hardness	5 MPa
CAS Number	7439-93-2
History
Naming	from the Greek word λιθoς, stone
Discovery	Johan August Arfwedson (1817)
First isolation	William Thomas Brande (1821)
Isotopes of lithium v e
Main isotopes[7] Decay Isotope 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 view talk edit | references

Lithium (from Ancient Greek: λίθος, líthos, 'stone') is a chemical element; it has symbol Li and atomic number 3. It is a soft, silvery-white alkali metal. Under standard conditions, it is the least dense metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable, and must be stored in vacuum, inert atmosphere, or inert liquid such as purified kerosene[8] or mineral oil. It exhibits a metallic luster when pure, but quickly corrodes in air to a dull silvery gray, then black tarnish. It does not occur freely in nature, but occurs mainly as pegmatitic minerals, which were once the main source of lithium. Due to its solubility as an ion, it is present in ocean water and is commonly obtained from brines. Lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

The nucleus of the lithium atom verges 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 its nuclei are very light: it is an exception to the trend that heavier nuclei are less common.[9] For related reasons, lithium has important uses in nuclear physics. The transmutation of lithium atoms to helium in 1932 was the first fully human-made nuclear reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.[10]

Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, lithium grease lubricants, flux additives for iron, steel and aluminium production, lithium metal batteries, and lithium-ion batteries. Batteries alone consume more than three-quarters of lithium production.[11]

Lithium is present in biological systems in trace amounts. Lithium-based drugs are useful as a mood stabilizer and antidepressant in the treatment of mental illness such as bipolar disorder.

Lithium ingots with a thin layer of black nitride tarnish

The alkali metals are also called the lithium family, after its leading element. Like the other alkali metals (which are sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr)), lithium has a single valence electron that, in the presence of solvents, is easily released to form Li+.[12] Because of this, lithium is a good conductor of heat and electricity as well as being chemically reactive, though it is the least reactive of the alkali metals. Lithium's lower reactivity is due to the proximity of its valence electron to its nucleus (the remaining two electrons are in the 1s orbital, much lower in energy, and do not participate in chemical bonds).[12] Molten lithium is significantly more reactive than its solid form.[13][14]

Lithium metal is soft enough to be cut with a knife. It is silvery-white. In air it oxidizes to lithium oxide.[12] Its melting point of 180.50 °C (453.65 K; 356.90 °F)[15] and its boiling point of 1,342 °C (1,615 K; 2,448 °F)[15] are each the highest of all the alkali metals.

Lithium has the lowest density (0.534 g/cm3) of all metals under standard conditions.[16] 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 floating in oil

Lithium's coefficient of thermal expansion is twice that of aluminium and almost four times that of iron.[17] Lithium is superconductive below 400 μK at standard pressure[18] and at higher temperatures (more than 9 K) at very high pressures (>20 GPa).[19] 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 prevalent.[20] Multiple allotropic forms have been identified for lithium at high pressures.[21]

Lithium has a mass specific heat capacity of 3.58 kilojoules per kilogram-kelvin, the highest of all solids.[22][23] Because of this, lithium metal is often used in coolants for heat transfer applications.[22]

Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (95.15% natural abundance).[24][7] Both natural isotopes have anomalously low nuclear binding energy per nucleon (compared to the neighboring elements on the periodic table, helium and beryllium); lithium is the only low numbered element that can produce net energy through nuclear fission. The two lithium nuclei have lower binding energies per nucleon than any other stable nuclides other than hydrogen-1, 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.[9] 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] The 6Li isotope is one of only five stable nuclides to have both an odd number of protons and an odd number of neutrons, the other four stable odd-odd nuclides being hydrogen-2, boron-10, nitrogen-14, and tantalum-180m.[27]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[28] 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 neutron halo, with 2 neutrons orbiting around its nucleus of 3 protons and 6 neutrons. The process known as laser isotope separation can be used to separate lithium isotopes, in particular 7Li from 6Li.[29]

Nuclear weapons manufacture and other nuclear physics applications are a major source of artificial lithium fractionation, with the light isotope 6Li being retained by industry and military stockpiles to such an extent that it has caused slight but measurable change in the 6Li to 7Li ratios 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.[30]

Both stable isotopes of lithium can be laser cooled and were used to produce the first quantum degenerate Bose–Fermi mixture.[31]

Nova Centauri 2013, the place lithium from a stellar nova was detected[32]

Lithium-7 was created in nuclear reactions very early in the history of the universe, in a process called Big Bang nucleosynthesis. The relative amount of lithium was small compared to isotopes of hydrogen and helium. The only other element created was beryllium-7 which decayed into lithium-7. Measured lithium abundance does not match the nucleosynthesis model well, an issue called "cosmological lithium problem;" whether this is due to issues in the difficult experimental measurements or in aspects of the theory is not known. Measurements of primordial lithium abundance are complicated because lithium-7 can be created and destroyed in cosmic ray, classical nova, and supernova reactions, and inside stars. The only known mechanism to produce lithium-6 is via cosmic-ray interactions.[33]

Lithium is also found in brown dwarf substellar objects. 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 them.[34][35] Certain red giant stars can also contain a high concentration of lithium. These stars found to may orbit massive objects—neutron stars or black holes—whose gravity pulls lithium where it can be observed.[36][37]

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

Although lithium is widely distributed on Earth, it does not naturally occur in elemental form due to its high reactivity.[12] 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),[38][39] or 25 micromolar;[40] higher concentrations approaching 7 ppm are found near hydrothermal vents.[39]

Estimates for the Earth's crustal content range from 20 to 70 ppm by weight.[41][42] 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.[41] Another significant mineral of lithium is lepidolite which is now an obsolete name for a series formed by polylithionite and trilithionite.[43][44] Another source for lithium is hectorite clay, the only active development of which is through the Western Lithium Corporation in the United States.[45] At 20 mg lithium per kg of Earth's crust,[46] lithium is the 31st most abundant element.[47]

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."[48]

Chile is estimated (2020) to have the largest reserves by far (9.2 million tonnes),[49] and Australia the highest annual production (40,000 tonnes).[49] One of the largest reserve bases[note 1] of lithium is in the Salar de Uyuni area of Bolivia, which has 5.4 million tonnes. Other major suppliers include Argentina and China.[50][51] As of 2015, the Czech Geological Survey considered the entire Ore Mountains in the Czech Republic as lithium province. Five deposits are registered, one near Cínovec [cs] is considered as a potentially economical deposit, with 160 000 tonnes of lithium.[52] In December 2019, Finnish mining company Keliber Oy reported its Rapasaari lithium deposit has estimated proven and probable ore reserves of 5.280 million tonnes.[53]

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.[54] These estimates are "based principally on old data, which was gathered mainly by the Soviets during their occupation of Afghanistan from 1979–1989".[55] The Department of Defense estimated the lithium reserves in Afghanistan to amount to the ones in Bolivia and dubbed it as a potential "Saudi-Arabia of lithium".[56] In Cornwall, England, the presence of brine rich in lithium was well known due to the region's historic mining industry, and private investors have conducted tests to investigate potential lithium extraction in this area.[57][58]

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 contain lithium ranging from 21 to 763 ppb.[39] Marine organisms tend to bioaccumulate lithium more than terrestrial organisms.[59] Whether lithium has a physiological role in any of these organisms is unknown.[39] Lithium concentrations in human tissue averages about 24 ppb (4 ppb in blood, and 1.3 ppm in bone).[60]

Lithium is easily absorbed by plants[60] and lithium concentration in plant tissue is typically around 1 ppm.[61] Some plant families bioaccumulate more lithium than others.[61] Dry weight lithium concentrations for members of the family Solanaceae (which includes potatoes and tomatoes), for instance, can be as high as 30 ppm while this can be as low as 0.05 ppb for corn grains.[60] Studies of lithium concentrations in mineral-rich soil give ranges between around 0.1 and 50−100 ppm, with some concentrations as high as 100−400 ppm, although it is unlikely that all of it is available for uptake by plants.[61] Lithium accumulation does not appear to affect the essential nutrient composition of plants.[61] Tolerance to lithium varies by plant species and typically parallels sodium tolerance; maize and Rhodes grass, for example, are highly tolerant to lithium injury while avocado and soybean are very sensitive.[61] Similarly, lithium at concentrations of 5 ppm reduces seed germination in some species (e.g. Asian rice and chickpea) but not in others (e.g. barley and wheat).[61]

Many of lithium's major biological effects can be explained by its competition with other ions.[62] The monovalent lithium ion Li+
competes with other ions such as sodium (immediately below lithium on the periodic table), which like lithium is also a monovalent alkali metal. Lithium also competes with bivalent magnesium ions, whose ionic radius (86 pm) is approximately that of the lithium ion[62] (90 pm). Mechanisms that transport sodium across cellular membranes also transport lithium. For instance, sodium channels (both voltage-gated and epithelial) are particularly major pathways of entry for lithium.[62] Lithium ions can also permeate through ligand-gated ion channels as well as cross both nuclear and mitochondrial membranes.[62] Like sodium, lithium can enter and partially block (although not permeate) potassium channels and calcium channels.[62]

The biological effects of lithium are many and varied but its mechanisms of action are only partially understood.[63] For instance, studies of lithium-treated patients with bipolar disorder show that, among many other effects, lithium partially reverses telomere shortening in these patients and also increases mitochondrial function, although how lithium produces these pharmacological effects is not understood.[63][64]

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.[65][66][67][68] 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.[69][70][71][72] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and less alkaline.[73] 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 new element "lithium".[12][67][72]

Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[74][67] In 1818, Christian Gmelin was the first to observe that lithium salts give a bright red color to flame.[67][75] However, both Arfwedson and Gmelin tried and failed to isolate the pure element from its salts.[67][72][76] 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.[36][76][77][78][79] 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).[80] In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[67][81] The discovery of this procedure led to commercial production of lithium in 1923 by the German company Metallgesellschaft AG, which performed an electrolysis of a liquid mixture of lithium chloride and potassium chloride.[67][82][83]

Australian psychiatrist John Cade is credited with reintroducing the use of lithium to treat mania in 1949. Mogens Schou of Denmark continued Cade's research, starting in the 1950s.[84] Throughout the mid-20th century, lithium's mood stabilizing applicability for mania and depression took off in Europe and the United States. As lithium cannot be patented, it was developed as a medicine by academia rather than drug companies.[85]

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 and 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 demand for lithium soaps and lubricating greases was supported by several small mining operations, mostly in the US.

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 US became the prime producer of lithium 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.[30][86]

Lithium is used to decrease the melting temperature of glass and to improve the melting behavior of aluminium oxide in the Hall-Héroult process.[87][88] 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.[86] In the mid-1990s, several companies started to isolate lithium from brine which proved to be a less expensive option than underground or open-pit mining. Most of the mines closed or shifted their focus to other materials because 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 beginning of the 21st century.

The development of lithium-ion batteries increased the demand for lithium and became the dominant use in 2007.[89] With the surge of lithium demand in batteries in the 2000s, new companies have expanded brine isolation efforts to meet the rising demand.[90][91]

Lithium reacts with water easily, but with noticeably less vigor than other alkali metals. The reaction forms hydrogen gas and lithium hydroxide.[12] When placed over a flame, lithium compounds give off a striking crimson color, but when the metal burns strongly, the flame becomes a brilliant silver. Lithium will ignite and burn in oxygen when exposed to water or water vapor. 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).[41] Lithium is one of the few metals that react with nitrogen gas.[92][93]

Because of its reactivity with water, and especially nitrogen, lithium metal is usually stored in a hydrocarbon sealant, often petroleum jelly. Although the heavier alkali metals can be stored under mineral oil, lithium is not dense enough to fully submerge itself in these liquids.[36]

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.[41][94] The metal reacts with hydrogen gas at high temperatures to produce lithium hydride (LiH).[95]

Lithium forms a variety of binary and ternary materials by direct reaction with the main group elements. These Zintl phases, although highly covalent, can be viewed as salts of polyatomic anions such as Si44-, P73-, and Te52-. With graphite, lithium forms a variety of intercalation compounds.[94]

It dissolves in ammonia (and amines) to give [Li(NH3)4]+ and the solvated electron.[94]

Lithium forms salt-like derivatives with all halides and pseudohalides. Some examples include the halides LiF, LiCl, LiBr, LiI, as well as the pseudohalides and related anions. Lithium carbonate has been described as the most important compound of lithium.[94] This white solid is the principal product of beneficiation of lithium ores. It is a precursor to other salts including ceramics and materials for lithium batteries.

The compounds LiBH
4 and LiAlH
4 are useful reagents. These salts and many other lithium salts exhibit distinctively high solubility in ethers, in contrast with salts of heavier alkali metals.

In aqueous solution, the coordination complex [Li(H2O)4]+ predominates for many lithium salts. Related complexes are known with amines and ethers.

Hexameric structure of the n-butyllithium fragment in a crystal

Organolithium compounds are numerous and useful. They are defined by the presence of a bond between carbon and lithium. They serve as metal-stabilized carbanions, although their solution and solid-state structures are more complex than this simplistic view.[96] Thus, these are extremely powerful bases and nucleophiles. They have also been applied in asymmetric synthesis in the pharmaceutical industry. For laboratory organic synthesis, many organolithium reagents are commercially available in solution form. These reagents are highly reactive, and are sometimes pyrophoric.

Like its inorganic compounds, almost all organic compounds of lithium formally follow the duet rule (e.g., BuLi, MeLi). However, in the absence of coordinating solvents or ligands, organolithium compounds form dimeric, tetrameric, and hexameric clusters (e.g., BuLi is actually [BuLi]6 and MeLi is actually [MeLi]4) which feature multi-center bonding and increase the coordination number around lithium. These clusters are broken down into smaller or monomeric units in the presence of solvents like dimethoxyethane (DME) or ligands like tetramethylethylenediamine (TMEDA).[97] As an exception to the duet rule, a two-coordinate lithate complex with four electrons around lithium, [Li(thf)4]+[((Me3Si)3C)2Li]–, has been characterized crystallographically.[98]