Gallium
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| General | |||||||||||||||||||
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| Name, Symbol, Number | gallium, Ga, 31 | ||||||||||||||||||
| Chemical series | poor metals | ||||||||||||||||||
| Group, Period, Block | 13, 4, p | ||||||||||||||||||
| Appearance | silvery white | ||||||||||||||||||
| Standard atomic weight | 69.723(1) g·mol−1 | ||||||||||||||||||
| Electron configuration | [Ar] 3d10 4s2 4p1 | ||||||||||||||||||
| Electrons per shell | 2, 8, 18, 3 | ||||||||||||||||||
| Physical properties | |||||||||||||||||||
| Phase | solid | ||||||||||||||||||
| Density (near r.t.) | 5.91 g·cm−3 | ||||||||||||||||||
| Liquid density at m.p. | 6.095 g·cm−3 | ||||||||||||||||||
| Melting point | 302.9146 K (29.7646 °C, 85.5763 °F) | ||||||||||||||||||
| Boiling point | 2477 K (2204 °C, 3999 °F) | ||||||||||||||||||
| Heat of fusion | 5.59 kJ·mol−1 | ||||||||||||||||||
| Heat of vaporization | 254 kJ·mol−1 | ||||||||||||||||||
| Heat capacity | (25 °C) 25.86 J·mol−1·K−1 | ||||||||||||||||||
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| Atomic properties | |||||||||||||||||||
| Crystal structure | orthorhombic | ||||||||||||||||||
| Oxidation states | 3, 1 (amphoteric oxide) | ||||||||||||||||||
| Electronegativity | 1.81 (scale Pauling) | ||||||||||||||||||
| Ionization energies (more) | 1st: 578.8 kJ·mol−1 | ||||||||||||||||||
| 2nd: 1979.3 kJ·mol−1 | |||||||||||||||||||
| 3rd: 2963 kJ·mol−1 | |||||||||||||||||||
| Atomic radius | 130 pm | ||||||||||||||||||
| Atomic radius (calc.) | 136 pm | ||||||||||||||||||
| Covalent radius | 126 pm | ||||||||||||||||||
| Van der Waals radius | 187 pm | ||||||||||||||||||
| Miscellaneous | |||||||||||||||||||
| Magnetic ordering | no data | ||||||||||||||||||
| Thermal conductivity | (300 K) 40.6 W·m−1·K−1 | ||||||||||||||||||
| Speed of sound (thin rod) | (20 °C) 2740 m/s | ||||||||||||||||||
| Mohs hardness | 1.5 | ||||||||||||||||||
| Brinell hardness | 60 MPa | ||||||||||||||||||
| CAS registry number | 7440-55-3 | ||||||||||||||||||
| Selected isotopes | |||||||||||||||||||
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Gallium (IPA: /ˈgaliəm/) is a chemical element that has the symbol Ga and atomic number 31. A soft silvery metallic poor metal, gallium is a brittle solid at low temperatures but liquefies slightly above room temperature and will melt in the hand. It occurs in trace amounts in bauxite and zinc ores. An important application is in the compounds gallium nitride and gallium arsenide, used as a semiconductor, most notably in light-emitting diodes (LEDs).
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[edit] Notable characteristics
Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium metal has a brilliant silvery color and its solid metal fractures conchoidally like glass. Gallium metal expands by 3.1 percent when it solidifies, and therefore storage in either glass or metal containers is avoided, due to the possibility of container rupture with freezing. Gallium shares the higher-density liquid state with only a few materials like germanium, bismuth, antimony and water.
Gallium also attacks most other metals by diffusing into their metal lattice. Gallium for example diffuses into the grain boundaries of Al/Zn alloys[1] or steel.[2], making them very brittle. Gallium metal easily alloys with many metals,[citation needed] and was used in small quantities in the core of the first atomic bomb to help stabilize the plutonium crystal structure.[citation needed]
The melting point temperature of 30°C allows the metal to be melted in one's hand. This metal has a strong tendency to supercool below its melting point/freezing point, thus necessitating seeding in order to solidify. Gallium is one of the metals (with caesium, rubidium, francium and mercury) which are liquid at or near normal room temperature, and can therefore be used in metal-in-glass high-temperature thermometers. It is also notable for having one of the largest liquid ranges for a metal, and (unlike mercury) for having a low vapor pressure at high temperatures. Unlike mercury, liquid gallium metal wets glass and skin, making it mechanically more difficult to handle (even though it is substantially less toxic and requires far fewer precautions). For this reason as well as the metal contamination problem and freezing-expansion problems noted above, samples of gallium metal are usually supplied in polyethylene packets within other containers.
Gallium does not crystallize in any of the simple crystal structures. The stable phase under normal conditions is orthorhombic with 8 atoms in the conventional unit cell. Each atom has only one nearest neighbor (at a distance of 244 pm) and six other neighbors within additional 39 pm. Many stable and metastable phases are found as function of temperature and pressure.
The bonding between the nearest neighbors is found to be of covalent character, hence Ga2 dimers are seen as the fundamental building blocks of the crystal. The compound with arsenic, gallium arsenide is a semiconductor commonly used in light-emitting diodes.
High-purity gallium is attacked slowly by mineral acids.
[edit] History
Gallium (Latin Gallia meaning Gaul (essentially modern France); also gallus, meaning "rooster") was discovered spectroscopically by Lecoq de Boisbaudran in 1875 by its characteristic spectrum (two violet lines) in an examination of a zinc blende from the Pyrenees. Before its discovery, most of its properties had been predicted and described by Dmitri Mendeleev (who called the hypothetical element eka-aluminium) on the basis of its position in his periodic table. Later, in 1875, Boisbaudran obtained the free metal through the electrolysis of its hydroxide in KOH solution. He named the element "gallia" after his native land of France. It was later claimed that, in one of those multilingual puns so beloved of men of science of the early 19th century, he also named it after himself, as 'Lecoq' = the rooster, and Latin for rooster is "gallus"; however, he denied this in an 1877 article.
[edit] Occurrence
Gallium does not exist in free form in nature, nor do any high-gallium minerals exist to serve as a primary source of extraction of the element or its compounds. Gallium is found and extracted as a trace component in bauxite, coal, diaspore, germanite, and sphalerite. The United States Geological Survey (USGC) estimates gallium reserves based on 50 ppm by weight concentration in known reserves of bauxite and zinc ores. Some flue dusts from burning coal have been shown to contain small quantities of gallium, typically less than 1 % by weight.[3][4][5][6]
Most gallium is extracted from the crude aluminium hydroxide solution of the Bayer process for producing alumina and aluminum. A mercury cell electrolysis and hydrolysis of the amalgam with sodium hydroxide leads to sodium gallate. Electrolysis then gives gallium metal. For semiconductor use, further purification is carried out using zone melting, or else single crystal extraction from a melt (Czochralski process). Purities of 99.9999% are routinely achieved and commercially widely available.
As of 2006, the current price for 1 kg gallium of 99.9999% purity seems to be at about US$ 400.[citation needed]
[edit] Applications
Semiconductor and electronic industry. The semiconductor applications are the main reason for the low-cost commercial availability of the extremely high-purity (99.9999+%) metal:
- As a component of the semiconductor Gallium arsenide, the most common application for gallium is analog integrated circuits, with the second largest use being optoelectronic devices (mostly laser diodes and light-emitting diodes.)
- Gallium is used widely as a dopant to dope semiconductors and produce solid-state devices like transistors.
- Gallium is the rarest component of new photovoltaic compounds (such as copper indium gallium selenium sulphide or Cu(In,Ga)(Se,S)2, recently announced by South African researchers) for use in solar panels as an alternative to crystalline silicon, which is currently in short supply.
As a wetting, and alloy improvement agent:
- Because gallium wets glass or porcelain, gallium can be used to create brilliant mirrors.
- Gallium readily alloys with most metals, and has been used as a component in low-melting alloys. The plutonium used in nuclear weapon pits is machined by alloying with gallium to stabilize the allotropes of plutonium.
- Gallium added in quantities up to 2% in common solders can aid wetting and flow characteristics.
As part of an energy storage mechanism:
- When gallium is alloyed with aluminium it can be used to break the bond between hydrogen and oxygen in water. A reaction occurs when water is added to the alloy which produces hydrogen and aluminium oxide. This could potentially provide a solid hydrogen source for transportation purposes, which would be more convenient than a pressurized hydrogen tank.[7] Resmelting the resultant aluminum oxide and gallium mixture to metallic aluminum and gallium and reforming these into electrodes would constitute most of the energy input into the system, while electricity produced by a hydrogen fuel cell could constitute an energy output.[8]The thermodynamic efficiency of the aluminum smelting process is said to be approximately 50 percent. Therefore, at most no more than half the energy that goes into smelting aluminum could be recovered by a fuel cell.
For liquid alloys:
- It has been suggested that a liquid gallium-tin alloy could be used to cool computer chips in place of water. As it conducts heat approximately 65 times better than water it can make a comparable coolant. [1] Gallium has a specific heat capacity of 0.37 J/(g·K). Gallium's high specific gravity of 5.91 gives it a volumetric heat capacity of 0.37 x 5.91 = 2.187 J/cm³, meaning that a volume of gallium will heat by 2.187/4.184 = 0.5227 as much as an equal volume of water in a cooling device. However given water's benign handling characteristics and plentiful abundance in most developed countries, gallium alloys are only really likely to see use in specialized applications such as cooling supercomputers.
- Gallium is used in some high temperature thermometers.
Biomedical applications:
- A low temperature liquid eutectic alloy of gallium, indium, and tin, is widely available in medical thermometers (fever thermometers), replacing problematic mercury. This alloy, with the trade name Galinstan (with the "-stan" referring to the tin), has a freezing point of −20°C.
- Gallium salts such as gallium citrate and gallium nitrate are used as radiopharmaceutical agents in nuclear medicine imaging. (The form or salt is not important, since it is the free dissolved gallium ion Ga3+ which is active). For these applications, a radioactive isotope such as 67Ga is used. The body handles Ga3+ in many ways as though it were iron, and thus it is bound (and concentrates) in areas of inflammation, such as infection, and also areas of rapid cell division. This allows such sites to be imaged by nuclear scan techniques. See gallium scan. This use has largely been replaced by fluorodeoxyglucose (FDG) for positron emission tomography, "PET" scan.
- Gallium nitrate, both oral and topical, is finding use in treating arthritis.[9]
- Much research is being devoted to gallium alloys as substitutes for mercury dental amalgams, but these compounds have yet to see wide acceptance.
- Research is being conducted to determine whether gallium can be used to fight bacterial infections in people with cystic fibrosis. Gallium is similar in size to iron, an essential nutrient for respiration. When gallium is mistakenly picked up by bacteria such as Pseudomonas, the bacteria's ability to respire is interfered with and the bacteria die. The mechanism behind this is that iron is redox active, which allows for the transfer of electrons during respiration, but gallium is redox inactive. [10][11]
Miscellaneous:
- Magnesium gallate containing impurities (such as Mn2+), is beginning to be used in ultraviolet-activated phosphor powder.
- Neutrino detection. Possibly the largest amount of pure gallium ever collected in a single spot was the GALLEX neutrino detector operated in the early 1990's in an Italian mountain tunnel. The detector contained 12.2 tons of watered gallium-71. Solar neutrinos caused a few atoms of Ga-71 to become radioactive Ge-71, which were detected. The solar neutrino flux deduced was found to have a deficit of 40% from theory. This was not explained until better solar neutrino detectors and theories were constructed (see SNO).[2]
[edit] Precautions
While not considered toxic, the data about gallium is inconclusive. Some sources suggest that it may cause dermatitis from prolonged exposure; other tests have not caused a positive reaction. Like most metals, finely divided gallium loses its luster. Powdered gallium appears gray. When gallium is handled with bare hands, the extremely fine dispersion of liquid gallium droplets which results from wetting skin with the metal may appear as a gray skin stain.