Lithium

For other uses, see the chemical element.

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] 6.94±0.06 (abridged)
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	1603 K ​(1330 °C, ​2426 °F)
Density (at 20° C)	0.5334 g/cm3
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)
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
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
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)
Thermal expansion	46.56×10−6/K (at 20 °C)
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)
Young's modulus	4.9 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
Discovery	Johan August Arfwedson (1817)
First isolation	William Thomas Brande (1821)
Isotopes of lithium
Main isotopes 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.

Lithium (Template:PronEng) 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. When cut open, lithium exhibits a metallic lustre, but contact with oxygen quickly turns it back to a dull silvery grey color. Lithium is also highly flammable.

According to theory, lithium was one of the very few elements synthesized in the Big Bang, although its quantity has vastly decreased since that time. The reasons for its disappearance, and the processes by which new lithium is created, continue to be active matters of study in astronomy. Though very light, lithium is nevertheless less common in the universe than any of the first 20 elements.

Due to its high reactivity it only appears naturally on Earth 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 a nuclear reaction, and lithium deuteride serves as the fusion fuel in staged thermonuclear weapons.

History and etymology

Petalite (lithium aluminium silicate) was first discovered in 1800 by the Portuguese 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 "lithos", 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. 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. 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.

Properties

Lithium ingots with a thin layer of black oxide tarnish

Lithium pellets (covered in white lithium hydroxide)

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 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.

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. Lithium has about half the density of water, giving solid sticks of lithium metal the odd heft of a light-to-medium wood such as pine. The metal floats 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 possesses 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.

At cryogenic temperatures, lithium, like sodium, undergoes martensitic transformations when cooled below liquid nitrogen temperatures (77 K). This is a multicrystalline state composed of Face-centered cubic, Body-centered cubic and 9R hexagonal crystals. At liquid helium temperatures (4 K) 9R Hex is the most prevalent structure. The proportion of the different crystalline states is temperature-dependent. On cryogenic cooling lithium at atmospheric pressure, the crystalline state which first predominates is fcc, followed by bcc, followed by 9R hex at the coldest temperatures. On heating solid lithium from deep cryogenic temperatures, the property known as heat of reversion will cause a crystalline state transition from 9R hex to bcc, which absorbs heat and causes cooling. In warming at cryogenic temperatures, there is thus a region of negative specific heat in lithium crystals, due to this state change.

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).

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).

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

Main article: nucleosynthesis

See also: Lithium minerals

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

The stable isotopes lithium-6 and lithium-7 were created in the Big Bang, but the amounts are unclear. There is general agreement that they were larger than the cosmos contains today. Because of the method by which elements are built up by fusion in stars, there is a general trend in the cosmos that the lighter elements are more common. However, lithium (element number 3) is tied with krypton as the 32nd/33rd most abundant element in the cosmos (see Cosmochemical Periodic Table of the Elements in the Solar System), being less common than any element before scandium (element 21). It is not until atomic number 36 (krypton) and beyond, that chemical elements are found to be universally less common in the cosmos than lithium. The reasons have to do with the failure of any good mechanisms to synthesize lithium in the fusion reactions between nuclides in supernovae. Due to the absence of any nuclide with five nucleons which is quasi-stable, nuclei of lithium-5 produced from helium and a proton has no time to fuse with a second proton or neutron to form a six nucleon isotope which might decay to lithium-6, even under extreme conditions of bombardment. Also, the product of helium-helium fusion (berylium-8) is also immediately unstable toward disintigration to helium again, and is thus also not available for formation of lithium. This leaves new formation of the stable isotopes lithium 6 and 7 to rare cosmic ray spallation on carbon or other elements in cosmic dust. Meanwhile, existing Li-6 and Li-7 is destroyed in many nuclear reactions in supernovae, resulting in net removal of lithium from the cosmos.

Lithium is widely distributed on Earth, 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. 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. A newer source for lithium is hectorite clay, the only active development of which is through Western Lithium Corp in the USA.

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." The most important deposit of lithium is in the Salar de Uyuni area of Bolivia, which holds half of the world's reserves. The lithium reserves are estimated at 30.000 tonnes in 2015.

Seawater contains an estimated 230 billion tons of lithium, though at a low concentration of 0.1 to 0.2 ppm.

Major applications of the metal

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 anode materials, used in Lithium-ion 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⁺. Although this ion has a smaller diameter than either Na⁺ or K⁺, in a watery environment like the cytoplasmic fluid, Li⁺ binds to the hydrogen atoms of water making it effectively larger than either Na⁺ or K⁺ ions. How Li⁺ works in the CNS is still a matter of debate. Li⁺ elevates brain levels of tryptophan, 5-HT (serotonin), and 5-HIAA (a serotonin metabolite). The serotonin system is related to stability of mood. Li⁺ also reduces catecholamine activity in the brain (associated with brain activation and mania), by enhancing reuptake and reducing release. 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.

Common side effects of lithium treatment include muscle tremors, twitching, ataxia, hyperparathyroidism as first described by Ashoka Prasad [1], bone loss, hypercalcemia, hypertension, etc.), kidney damage, nephrogenic diabetes insipidus (polyuria and polydipsia) and seizures. Some of the side-effects are a result of the increased elimination of potassium.

Pregnancy - teratogenic properties: Ebstein (cardiac) Anomaly - There appears to be an increased risk of this abnormality in infants of women taking lithium during the first trimester of pregnancy

Other uses

Red lithium flame

• Lithium batteries are disposable (primary) batteries that have lithium metal or lithium compounds as an anode. Lithium batteries are not to be confused with lithium-ion batteries which are high energy-density rechargeable batteries.
• 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-aluminium alloys are used in aerospace applications, such as the external tank of the Space Shuttle, and is planned for the Orion spacecraft.
• 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.
• 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.
• Lithium fluoride (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred 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 and the best neutronic properties of fluoride salt combinations appropriate for reactor use.
• 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.”
• Lithium can be used to make red fireworks
• 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, which generates enormous quantities of heat, in turn used to generate steam from seawater. The steam propels the torpedo in a closed Rankine cycle.

Production and world supply

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.

There are widespread hopes of using lithium ion batteries in electric vehicles, but one 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'".

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).

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.

Nearly half the world's reserves are located in Bolivia, which in 2009 is negotiating with Japanese and French firms to begin production. According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tons of lithium, which can be used to make batteries for hybrid and electric vehicles. This is the largest amount of lithium in any country when compared Chile's 3 million tons of lithium and the United States's 410,000 tons. Icy relations between Bolivia's President Evo Morales and the United States could delay negotiations, and the indigenous people who live on this land could also slow down the process by demanding that they reap their share of the profits.

China may emerge as a significant producer of brine-based lithium carbonate around 2010. Potential capacity of up to 55,000 tonnes per year could come on-stream if projects in Qinghai province and Tibet proceed.

The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tonnes of the known global lithium reserve base.

In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tonnes for the Western World. With the inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total has nearly tripled by 2008.

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 a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium should be stored in a non-reactive compound such as naphtha.

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 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. However, internal shorts have been known to develop due to manufacturing defects or some abuse conditions that can lead to spontaneous thermal runaway .

See also

• Lithium compounds
• Dilithium

porn and lots of it on atoliet

External links

• USGS: Lithium Statistics and Information
• WebElements.com – Lithium
• It's Elemental – Lithium
• University of Southampton, Mountbatten Centre for International Studies, Nuclear History Working Paper No5.