|Name, Symbol, Number||gallium, Ga, 31|
|Chemical series||poor metals|
|Group, Period, Block||13, 4, p|
|Atomic mass||69.723(1) g/mol|
|Electron configuration||[Ar] 3d10 4s2 4p1|
|Electrons per shell||2, 8, 18, 3|
|Density (near r.t.)||5.91 g/cm³|
|Liquid density at m.p.||6.095 g/cm³|
|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|
|Heat of vaporization||254 kJ/mol|
|Heat capacity||(25 °C) 25.86 J/(mol·K)|
|Electronegativity||1.81 (Pauling scale)|
|1st: 578.8 kJ/mol|
|2nd: 1979.3 kJ/mol|
|3rd: 2963 kJ/mol|
|Atomic radius||130 pm|
|Atomic radius (calc.)||136 pm|
|Covalent radius||126 pm|
|Van der Waals radius||187 pm|
|Magnetic ordering||no data|
|Thermal conductivity||(300 K) 40.6 W/(m·K)|
|Speed of sound (thin rod)||(20 °C) 2740 m/s|
|Brinell hardness||60 MPa|
|CAS registry number||7440-55-3|
Gallium (chemical symbol Ga, atomic number 31) is a rare, soft, silvery metal. It is a brittle solid at low temperatures, but it liquefies slightly above room temperature and melts in the hand. It is one of only a few materials that expands when freezing (like water), and its liquid form has a higher density than the solid form (like water). Gallium occurs in trace amounts in bauxite (an aluminum ore) and zinc ores.
Gallium is most commonly used in the form of the compound gallium(III) arsenide, which is a semiconductor useful for integrated circuits, light-emitting diodes (LEDs), and laser diodes. The nitride and phosphide of gallium are also valuable semiconductor materials, and gallium itself is used as a dopant in semiconductors. In addition, this metal is a component in low-melting temperature alloys, and its alloy with indium and tin is used in medical thermometers to replace mercury. Also, gallium can wet (coat) glass to create brilliant mirrors.
Gallium does not exist in free form in nature, nor are there any gallium-rich minerals that might serve as primary sources of extraction of the element or its compounds. Rather, gallium is extracted as a trace component from bauxite, coal, diaspore, germanite, and sphalerite. 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 aluminum 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 percent are routinely achieved and widely available commercially.
Before gallium was discovered, the element and many of its properties had been predicted and described by Dmitri Mendeleev, on the basis of its position in the periodic table. Mendeleev called the hypothetical element eka-aluminum.
In 1875, Lecoq de Boisbaudran discovered gallium by the technique known as spectroscopy. When examining a sample of zinc blende from the Pyrenees, he noticed two unique violet lines in its spectrum, indicative of a previously unknown element. Later, he obtained the free metal by the electrolysis of its hydroxide in KOH solution. He named the element "gallia" after his native land of France; also, in one of those multilingual puns so beloved of men of science of the early nineteenth century, he named it after himself—Lecoq means "the rooster" in French, and Latin for rooster is gallus.
In the periodic table, gallium lies in group 13 (former group 3A), between aluminum and indium, and in the same group as thallium. Consequently, its properties resemble those of these three elements. In addition, it is situated in period 4, between zinc and germanium. It is also said to be one of the "poor metals"—elements located between the transition metals and metalloids in the periodic table.
High-purity, metallic gallium has a brilliant, silvery color. By contrast, like most metals, finely divided gallium loses its luster—powdered gallium appears gray. The solid form fractures conchoidally, like glass. When liquid gallium solidifies, it expands by 3.1 percent. Thus, its liquid state has a higher density than the solid state—a property characteristic of only a few materials like water and bismuth. Also, given the property of expansion during solidification, gallium is not stored in either glass or metal containers to prevent the container from rupturing when the element freezes.
Gallium also diffuses into the crystal lattice of most other metals. This is another reason why it is important to keep gallium away from metal containers such as steel or aluminum. Gallium easily alloys with many other metals, and it was used in small quantities in the core of the first atomic bomb to help stabilize the plutonium crystal structure.
Given its melting point of 30°C, the metal readily melts in the hand. Also, the liquid form has a strong tendency to supercool below its melting point, and it needs to be seeded for solidification to begin. Gallium is one of the metals—along with cesium, francium, and mercury)—that is liquid at or near normal room temperature. It 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 wets (coats) glass and skin, making it mechanically more difficult to handle, although it is substantially less toxic and requires far fewer precautions. For this reason, as well as the metal contamination and freezing expansion problems noted above, samples of gallium metal are usually supplied in polyethylene packets within other containers.
Gallium does not crystallize into any of the simple crystal structures. The stable phase under normal conditions is orthorhombic, with eight atoms in the conventional unit cell. Each atom has only one nearest neighbor (at a distance of 244 picometers) and six other neighbors within an additional 39-picometer radius. The bonding between nearest neighbors has covalent character. Also, the element has many stable and metastable phases, depending on the temperature and pressure conditions.
High-purity gallium is attacked slowly by mineral acids.
Many isotopes of gallium are known, ranging from 56Ga to 86Ga. Among them, there are two stable isotopes: 69Ga and 71Ga, at relative abundances estimated at 60.11 percent and 39.89 percent, respectively. The radioisotopes, by contrast, have extremely short half-lives.
Gallium can form a number of compounds. Some of them are mentioned below.
Gallium, its alloys, and its compounds have many applications. Some of them are listed below.
Gallium is not considered toxic, but the data about its effects are inconclusive. Some sources suggest that it may cause dermatitis from prolonged exposure; other tests have not caused a positive reaction. When the element is handled with bare hands, the skin acquires a gray stain from an extremely fine dispersion of liquid gallium droplets.
All links retrieved May 19, 2017.
New World Encyclopedia writers and editors rewrote and completed the Wikipedia article in accordance with New World Encyclopedia standards. This article abides by terms of the Creative Commons CC-by-sa 3.0 License (CC-by-sa), which may be used and disseminated with proper attribution. Credit is due under the terms of this license that can reference both the New World Encyclopedia contributors and the selfless volunteer contributors of the Wikimedia Foundation. To cite this article click here for a list of acceptable citing formats.The history of earlier contributions by wikipedians is accessible to researchers here:
The history of this article since it was imported to New World Encyclopedia: