Lithium

Lithium
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the chemical element named Lithium. For other uses, see Lithium (disambiguation).
3 helium ← lithium → beryllium
H
↑
Li
↓
Na
Periodic Table - Extended Periodic Table


General
Name, Symbol, Number lithium, Li, 3
Chemical series alkali metals
Group, Period, Block 1, 2, s
Appearance silvery white/grey

Standard atomic weight 6.941(2) g·mol−1
Electron configuration 1s2 2s1
Electrons per shell 2, 1
Physical properties
Phase solid
Density (near r.t.) 0.534 g·cm−3
Liquid density at m.p. 0.512 g·cm−3
Melting point 453.69 K
(180.54 °C, 356.97 °F)
Boiling point 1615 K
(1342 °C, 2448 °F)
Critical point (extrapolated)
3223 K, 67 MPa
Heat of fusion 3.00 kJ·mol−1
Heat of vaporization 147.1 kJ·mol−1
Heat capacity (25 °C) 24.860 J·mol−1·K−1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k
at T/K 797 885 995 1144 1337 1610

Atomic properties
Crystal structure cubic body centered
Oxidation states 1
(strongly basic oxide)
Electronegativity 0.98 (Pauling scale)
Ionization energies 1st: 520.2 kJ/mol
2nd: 7298.1 kJ/mol
3rd: 11815.0 kJ/mol
Atomic radius 145 pm
Atomic radius (calc.) 167 pm
Covalent radius 134 pm
Van der Waals radius 182 pm
Miscellaneous
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 92.8 nΩ·m
Thermal conductivity (300 K) 84.8 W·m−1·K−1
Thermal expansion (25 °C) 46 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 6000 m/s
Young's modulus 4.9 GPa
Shear modulus 4.2 GPa
Bulk modulus 11 GPa
Mohs hardness 0.6
CAS registry number 7439-93-2
Selected isotopes
Main article: Isotopes of lithium iso NA half-life DM DE (MeV) DP
6Li 7.5% Li is stable with 3 neutrons
7Li 92.5% Li is stable with 4 neutrons
6Li content may be as low as 3.75% in
natural samples. 7Li would therefore
have a content of up to 96.25%.

References
This box: view • talk • edit
Lithium (IPA: [ˈlɪθiəm]) (from Greek λιθoς (lithos), "stone") 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. Lithium is the 33rd most abundant element on Earth,[1] but due to its high reactivity only appears there naturally in the form of compounds. It corrodes quickly in moist air, forming a black tarnish. On a commercial scale, lithium metal is produced electrolytically from a mixture of lithium chloride and potassium chloride and typically stored under the cover of oil to prevent reactions with air.

Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from natural brines. Trace amounts of lithium are present in the oceans and in some organisms, though it 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 nuclear reaction, and lithium deuteride serves as the fusion fuel in staged thermonuclear weapons.

Contents [hide]
1 Characteristics
1.1 Physical
1.2 Chemical
2 Occurrence
3 Production
4 Isotopes
5 History
6 Applications
6.1 Medical Use
6.2 Other uses
7 Regulation
8 Precautions
9 See also
10 References
11 External links



[edit] Characteristics
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. However, lithium's electronegativity is higher than any other alkali metal, as it has the smallest atomic radius. Thus, many of the properties attributable to alkali metals' weakly-held valence electron are most diminished in lithium; among alkali metals, it is the hardest and least reactive and possesses the highest melting and boiling points.


[edit] Physical
Lithium is soft enough to be cut with a knife, though this is significantly more difficult than cutting sodium. The fresh metal has a silvery-white color, rapidly tarnishing black in air. 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 lowest 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.


[edit] Chemical
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 when exposed to water and water vapours in oxygen. 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.


[edit] Occurrence

Lithium pellets (covered in white lithium hydroxide)On Earth, lithium is widely distributed, but because of its reactivity does not occur in its free form. In keeping with the origin of its name, lithium forms a minor part of almost all igneous rocks and is also found in many natural brines. Lithium is the 33rd most abundant element, contained particularly in the minerals spodumene, lepidolite, petalite, and amblygonite. On average, Earth's crust contains 65 ppm (.0007%) lithium.[1]

In humans lithium compounds have not been found to play a natural biological role; large amounts are slightly toxic. Lithium appears to be an essential trace element for goats, and possibly rats, suggesting a role in humans by analogy. However, the essentiality of ultratrace mineral in humans is far more difficult to determine, due to the difficulty and ethical issues involved with the experiments, which involve total isolation from the environment, and unpalatable semi-synthetic foods.


[edit] 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).[2]

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.[3]

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.[4]

See also Lithium minerals.


[edit] 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.


[edit] History
Petalite, which has lithium in it, was discovered by the Brazilian scientist José Bonifácio de Andrade e Silva in the late 1700s on a trip to Sweden. Lithium was discovered by Johan August Arfwedson in 1817. Arfwedson found the new element within the minerals spodumene and lepidolite in a petalite ore (LiAl(Si2O5)2) that he was analyzing during a routine investigation of some minerals from a mine on the island Utö in Sweden. In 1818 Christian Gmelin was the first to observe that lithium salts give a bright red color in flame. Both men tried and failed to isolate the element from its salts.

The element was not isolated until William Thomas Brande and Sir Humphry Davy later used electrolysis on lithium oxide in 1818. Robert Bunsen and Matiessen isolated larger quantities of the metal by electrolysis of lithium chloride in 1855. Commercial production of lithium metal was achieved in 1923 by the German company Metallgesellschaft through using electrolysis of molten lithium chloride and potassium chloride. It was apparently given the name "lithium" (Greek λιθoς (lithos), meaning "stone") because it was discovered from a mineral while other common alkali metals were first discovered from plant tissue.


[edit] 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.


[edit] Medical Use
Main article: Lithium pharmacology
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 interacts with the normal function of sodium ions to produce numerous changes in the neurotransmitter activity of the brain. Therapeutically useful amounts of lithium are only slightly lower than toxic amounts, so the blood levels of lithium must be carefully monitored during treatment.

Common side effects include muscle tremors, twitching, ataxia, nephrogenic diabetes insipidus (polyuria and polydipsia) and seizures. Most of the side-effects are a result caused by the increased elimination of potassium.


[edit] 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 per cent of mobile phones.[5]
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 apparently no longer plays a role in modern nuclear weapons, having been replaced entirely for the purpose by elemental tritium, which is lighter and easier to handle than lithium salts. [citation needed]
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.[6]
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 + 3T. 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.
Lithium metal is used as a reducing agent in some types of methamphetamine production, particularly in illegal amateur “meth labs.”
Lithium hydroxide is 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.

[edit] 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.

Lithium is a component for thermonuclear weapons (so called "hydrogen bombs") and applications of lithium for this purpose in the nuclear weapons industry is pursued in developing nuclear powers like India, and presumably others.


[edit] 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.


[edit] See also
Lithium compounds

[edit] References
Stwertka, Albert (2002). A Guide to the Elements. New York, NY: Oxford University Press. ISBN 978-0-19-515027-8.
Newton, David E. (1994). The Chemical Elements. New York, NY: Franklin Watts. ISBN 978-0-531-12501-4.
^ a b Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements : A Reference Guide. Westport, Conn.: Greenwood Press, 48. ISBN 0-313-33438-2.
^ A PDF file on lithium prices from the U.S. Geological Survey. Retrieved on September 15, 2005.
^ Los Alamos National Laboratory – Lithium. Retrieved on September 15, 2005.
^ www.roskill.com/reports/prePublication/prepublithium. Retrieved on 2007-06-08.
^ Two ways to play the lithium boom MoneyWeek, 08-01-2007
^ 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.

[edit] External links
Look up lithium in
Wiktionary, the free dictionary.Wikimedia Commons has media related to:
LithiumUSGS: Lithium Statistics and Information
WebElements.com – Lithium
It's Elemental – Lithium
Safety information on Lithium
Information on Lithium and Bipolar Disorder


[hide] v • d • e Alkali Metals
Lithium
Li
Atomic Number: 3
Atomic Weight: 6.941
Melting Point: 453.69
Boiling Point: 1615
Electronegativity: 0.98
|

Sodium
Na
Atomic Number: 11
Atomic Weight: 22.990
Melting Point: 370.87
Boiling Point: 1156
Electronegativity: 0.96
|

Potassium
K
Atomic Number: 19
Atomic Weight: 39.098
Melting Point: 336.58
Boiling Point: 1032
Electronegativity: 0.82
|

Rubidium
Rb
Atomic Number: 37
Atomic Weight: 85.468
Melting Point: 312.46
Boiling Point: 961
Electronegativity: 0.82
|

Caesium
Cs
Atomic Number: 55
Atomic Weight: 132.905
Melting Point: 301.59
Boiling Point: 944
Electronegativity: 0.79
|

Francium
Fr
Atomic Number: 87
Atomic Weight: (223)
Melting Point: ?295
Boiling Point: ?950
Electronegativity: 0.7




Retrieved from "http://en.wikipedia.org/wiki/Lithium"
Categories: All articles with unsourced statements | Articles with unsourced statements since February 2007 | Chemical elements | Alkali metals | Lithium