Lithium
From Freepedia
- This article is about the chemical element Lithium. For other uses, see Lithium (disambiguation).
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| Name, Symbol, Number | lithium, Li, 3 | ||||||||||||||||||||||||
| Chemical series | alkali metals | ||||||||||||||||||||||||
| Group, Period, Block | 1, 2, s | ||||||||||||||||||||||||
| Appearance | silvery white/gray Image:Li,3.jpg | ||||||||||||||||||||||||
| Atomic mass | 6.941(2) g/mol | ||||||||||||||||||||||||
| Electron configuration | 1s2 2s1 | ||||||||||||||||||||||||
| Electrons per shell | 2, 1 | ||||||||||||||||||||||||
| Physical properties | |||||||||||||||||||||||||
| Phase | solid | ||||||||||||||||||||||||
| Density (near r.t.) | 0.534 g/cm³ | ||||||||||||||||||||||||
| Liquid density at m.p. | 0.512 g/cm³ | ||||||||||||||||||||||||
| Melting point | 453.69 K (180.54 °C, 356.97 °F) | ||||||||||||||||||||||||
| Boiling point | 1615 K (1342 °C, 2448 °F) | ||||||||||||||||||||||||
| Heat of fusion | 3.00 kJ/mol | ||||||||||||||||||||||||
| Heat of vaporization | 147.1 kJ/mol | ||||||||||||||||||||||||
| Heat capacity | (25 °C) 24.860 J/(mol·K) | ||||||||||||||||||||||||
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| 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 | nonmagnetic | ||||||||||||||||||||||||
| Electrical resistivity | (20 °C) 92.8 nΩ·m | ||||||||||||||||||||||||
| Thermal conductivity | (300 K) 84.8 W/(m·K) | ||||||||||||||||||||||||
| Thermal expansion | (25 °C) 46 µm/(m·K) | ||||||||||||||||||||||||
| 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 | ||||||||||||||||||||||||
| Notable isotopes | |||||||||||||||||||||||||
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| References | |||||||||||||||||||||||||
Lithium is the chemical element with symbol Li and atomic number 3. In the periodic table, it is located in group 1, among the alkali metals. Lithium in its pure form is a soft, silver white metal, that tarnishes and oxidizes very rapidly in air and water. It is the lightest solid element and is primarily used in heat transfer alloys, in batteries and serves as a component in some drugs known as mood stabilizers.
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Basic features
Lithium is the lightest in weight of all metals and has a density that is only half that of water. Oddly, lithium also exhibits properties of the alkali Earth metals in Group 2. Lithium is a soft, silvery metal, so soft that it can be cut with a sharp knife. Like all alkali metals, lithium possesses a single valence electron, and will readily lose this electron to become a positive ion so that it will not have a full shell, or set of electrons at an energy level. Because of this, lithium reacts easily in water and does not occur freely in nature. Nevertheless it is still less reactive than the chemically similar sodium.
When placed over a flame, this metal 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 oxygen and water. It is the only metal that reacts with nitrogen at room temperature. Lithium has a high specific heat capacity, 3582 J/(kg·K), and a great temperature range in its liquid form, which makes it a useful chemical.
Lithium in its pure form is highly flammable and slightly explosive when exposed to air and especially water. Lithium fires are difficult to extinguish, requiring special chemicals designed to extinguish them. Lithium metal is also corrosive and requires special handling to avoid skin contact. Lithium should be stored in a non-reactive compound such as naphtha or a hydrocarbon. Lithium compounds play no natural biological role and are considered to be slightly toxic. When used as a drug, blood concentrations of Li+ must be carefully monitored.
Applications
Because of its large specific heat (the largest of any solid), lithium is used in heat transfer applications. It is also an important battery anode material due to its high electrochemical potential. In addition to being lighter than the standard dry cell, these batteries produce a higher voltage (3 volts versus 1.5 volts). Other uses:
- Lithium salts such as lithium carbonate (Li2CO3), lithium citrate, and lithium orotate are mood stabilizers 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. Useful amounts of Lithium for this use are only slightly lower than toxic amounts, so the blood levels of Lithium have to be carefully monitored during such a treatment.
- 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.
- Lithium hydroxide is employed to extract carbon dioxide from the air in spacecraft and submarines. Any alkali hydroxide will absorb CO2, but lithium hydroxide is preferred because of its low molecular weight.
- Alloys of the metal with aluminium, cadmium, copper, and manganese are used to make high performance aircraft parts.
- Lithium niobate is used extensively in the telecoms market, such as mobile phones and optical modulators.
- The high non-linearity of lithium niobate also makes a good choice for non-linear applications.
- Lithium deuteride (deuterium is an isotope of hydrogen) is the fusion fuel of choice in the so-called hydrogen bomb. When bombarded by neutrons, both lithium-6 and lithium-7 produce tritium. Tritium fuses with deuterium in a fusion reaction that is relatively easy to achieve.
- Lithium is used as a source for alpha particles, or helium nuclei. When lithium-7 nuclei are bombarded by accelerated protons, some of the lithium nuclei are broken into four protons and four neutrons, which, in turn, 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 hydroxide is an efficient and lightweight purifier of air. In confined areas, such as aboard spacecraft, 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.
History
Petalite, which contains lithium, was first discovered by the Brazilian scientist José Bonifácio de Andrada e Silva toward the end of the 1700s on a trip to Sweden. Lithium was discovered by Johann Arfvedson in 1817. Arfvedson found the new element within the minerals spodumene and lepidolite in a petalite ore, LiAl(Si2O5)2, 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, however.
The element was not isolated until William Thomas Brande and Sir Humphrey Davy later used electrolysis on lithium oxide in 1818. 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 AG through using electrolysis of molten lithium chloride and potassium chloride. It was apparently given the name "lithium" (Greek λιθοσ (lithos), meaning "stone") because it was discovered from a mineral while other common alkali metals were first discovered from plant tissue.
Occurrence
Lithium is widely distributed but does not occur in nature in its free form. Because of its reactivity, it is always found bound with one or more other elements or compounds. It forms a minor part of almost all igneous rocks and is also found in many natural brines. Lithium is the thirty-first most abundant element, contained in trace amounts in the minerals spodumene, lepidolite, and amblygonite. The Earth's crust contains 65 parts per million (ppm) of lithium.
Since the end of World War II, lithium production has greatly increased. The metal is separated from other elements in igneous rocks, and is also extracted from the water of mineral springs. Lepidolite, spodumene, petalite, and amblygonite are the more important minerals containing it.
In the United States lithium is recovered from brine pools in Nevada.[1] Today, most commercial lithium is recovered from brine sources in Chile. The metal, which is silvery in appearance like sodium, potassium and other members of the alkali metal series, is produced electrolytically from a mixture of fused lithium and potassium chloride. There is little market for lithium in its pure metal form and price information is scarce. In 1998 it was about US$ 43 per pound ($95 per kg). [2] Chile is currently the leading pure metal lithium producer in the world.
Isolation (* follow):
cathode: <math>\mbox{Li}^{+}\mbox{*} + \mbox{e}^{-} \to \mbox{Li*}</math>
anode: <math>\mbox{Cl}^{-}\mbox{*} \to \frac{1}{2}\mbox{Cl}_2 (\mbox{gas}) + e^-</math>
Isotopes
Naturally occurring lithium is composed of 2 stable isotopes Li-6 and Li-7 with Li-7 being the most abundant (92.5% natural abundance). Seven radioisotopes have been characterized with the most stable being Li-8 with a half-life of 838 ms and Li-9 with a half-life of 178.3 ms. All of the remaining radioactive isotopes have half-lifes that are less 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.
Lithium-7 is one of the primordial elements (produced in Big Bang nucleosynthesis). Lithium isotopes fractionate substantially during a wide variety of natural processes, including mineral formation (chemical precipitation), metabolism, ion exchange (Li substitutes for magnesium and iron in octahedral sites in clay minerals, where Li-6 is preferential over Li-7), hyperfiltration, and rock alteration.
References
- ^ "Los Alamos National Laboratory – Lithium." Accessed September 15, 2005.
- ^ "A PDF file on lithium prices from the U.S. Geological Survey." Accessed September 15, 2005.
- Book references
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- {{{Author|}}}{{|{{{3}}}}}}|show1| (1998)}}{{{{{Year|}}}}}}|show1|.}} {{|{{{3}}}}}}|show1|[{{{URL}}}}} The History and Use of Our Earth's Chemical Elements : A Reference Guide{{|{{{3}}}}}}|show1|]}}{{|{{{3}}}}}}|show1|, {{{Pages}}}}}{{|{{{3}}}}}}|Show1|, Greenwood Press, Westport, Conn.}}. {{{ID|}}}
- {{{Author|}}}{{|{{{3}}}}}}|show1| (1994)}}{{{{{Year|}}}}}}|show1|.}} {{|{{{3}}}}}}|show1|[{{{URL}}}}} The Chemical Elements{{|{{{3}}}}}}|show1|]}}{{|{{{3}}}}}}|show1|, {{{Pages}}}}}{{|{{{3}}}}}}|Show1|, Franklin Watts, New York, NY}}. {{{ID|}}}
