Gallium

2007 Schools Wikipedia Selection. Related subjects: Chemical elements

31 zinc ← gallium → germanium Al ↑ Ga ↓ In Periodic Table - Extended Periodic Table
General
Name, Symbol, Number	gallium, Ga, 31
Chemical series	poor metals
Group, Period, Block	13, 4, p
Appearance	silvery white
Atomic mass	69.723 (1) g/mol
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
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 1310 1448 1620 1838 2125 2518
Atomic properties
Crystal structure	orthorhombic
Oxidation states	3 ( amphoteric oxide)
Electronegativity	1.81 (Pauling scale)
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
Main article: Isotopes of gallium iso NA half-life DM DE ( MeV) DP 69Ga 60.11% Ga is stable with 38 neutrons 71Ga 39.89% Ga is stable with 40 neutrons
References

Gallium ( IPA: /ˈgaliəm/) is a chemical element that has the symbol Ga and atomic number 31. A rare, 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 compound gallium arsenide, used as a semiconductor, most notably in light-emitting diodes (LEDs).

Notable characteristics

Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium metal has a brilliant silvery colour 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 water and bismuth.

Gallium also attacks most other metals by diffusing into their metal lattice — another reason why it is important to keep gallium away from metal containers such as steel or aluminium. Gallium metal easily alloys with many metals, and was used in small quantities in the core of the first atomic bomb to help stabilize the plutonium crystal structure.

The melting point temperature of 30 °C allows the metal to be melted in the hand. This metal has a strong tendency to supercool below its melting point, thus necessitating seeding in order to solidify. Gallium is one of the metals (with caesium, 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 neighbour (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.

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 and, in one of those multilingual puns so beloved of men of science of the early 19th century, after himself, as 'Lecoq' = the rooster, and Latin for rooster is "gallus".

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 rather found and extracted as a trace component in bauxite, coal, diaspore, germanite, and sphalerite. The USGS 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 as much as 1.5 percent gallium.

Most gallium is extracted from the crude aluminium hydroxide solution of the Bayer process for producing alumina and aluminium. 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.

Applications

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.)

Other uses include:

• Since gallium wets glass or porcelain, gallium can be used to create brilliant mirrors.
• Used widely as a dopant to dope semiconductors and produce solid-state devices like transistors.
• 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. Much research is being devoted to gallium alloys as substitutes for mercury dental amalgams, but such compounds have yet to see wide acceptance.
• Gallium added in quantities up to 2% in common solders can aid wetting and flow characteristics.
• Gallium is used in some high temperature thermometers.
• A 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, has a freezing point of −20 °C.
• Magnesium gallate containing impurities (such as Mn⁺²), is beginning to be used in ultraviolet-activated phosphor powder.
• 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. 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 4.184/2.187 = 1.9 times more than 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 specialised applications such as cooling supercomputers.
• 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 Ga⁺³ which is active). For these applications, a radioactive isotope such as ⁶⁷Ga is used. The body handles Ga⁺³ 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 FDG PET scan.
• 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.

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.