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Data Zone
| Classification: | Gallium is an ‘other metal’ |
| Color: | silvery-blue |
| Atomic weight: | 69.723 |
| State: | solid |
| Melting point: | 29.76 oC, 302.91 K |
| Boiling point: | 2200 oC, 2473 K |
| Electrons: | 31 |
| Protons: | 31 |
| Neutrons in most abundant isotope: | 40 |
| Electron shells: | 2,8,18,3 |
| Electron configuration: | [Ar] 3d10 4s2 4p1 |
| Density @ 20oC: | 5.907 g/cm3 |
Discovery of Gallium
Before the discovery of gallium its existence and main properties were predicted by Russian chemist Dmitri Mendeleev. He named the hypothetical element eka-aluminum as he predicted the element would sit below aluminum on the periodic table.
Gallium was discovered by French chemist Paul E. Lecoq de Boisbaudran through a spectroscope in 1875 in Paris.
Its now characteristic spectrum (two violet lines) identified it as a new element.
De Boisbaudran extracted gallium in the first instance from a zinc blende ore from the Pyrenees and obtained initially only 0.65 grams from 430 kilograms of ore. He isolated gallium by electrolysis of its hydroxide in potassium hydroxide solution.
The origin of the name comes from the Latin word ‘Gallia’, meaning France.
Appearance and Characteristic
Harmful effects:
Gallium is considered to be non-toxic.
Characteristics:
Gallium is a silvery, glass-like, soft metal. It sits close to the non-metals in the periodic table and its metallic properties aren’t as obviously metallic as most other metals. Solid gallium is brittle and is a poorer electrical conductor than lead.
Gallium is a silvery, glass-like, soft metal. It sits close to the non-metals in the periodic table and its metallic properties aren’t as obviously metallic as most other metals. Solid gallium is brittle and is a poorer electrical conductor than lead.
The solid metal fractures conchoidally. (Conchoidally means like a shell – the fractured surfaces are curved like a sea shell.)
Gallium has the second largest liquid range of any element and is one of the few metals that is liquid near room temperature (m.pt. 29.76 oC, 85.6 oF ), melting in the hand.
Bromine is the only non-metallic element that is liquid at or around room-temperature.
Gallium liquid clings to or wets glass and similar surfaces.
Gallium also has the unusual property that (like water) it expands as it freezes.
Uses of Gallium
Low melting gallium alloys are used in some medical thermometers as non-toxic substitutes for mercury.
Gallium arsenide is used in semiconductor production mainly for laser diodes, light-emitting diodes and solar panels. It is also used to create brilliant mirrors.
Abundance and Isotopes
Abundance earth’s crust: 19 parts per million by weight, 5.5 parts per million by moles
Abundance solar system: 40 parts per billion by weight, 0.6 parts per billion by moles
Cost, pure: $220 per 100g
Cost, bulk: $ per 100g
Source: Gallium does not exist free in nature and there are no minerals with any substantial gallium content. Commercially, most gallium is extracted as a byproduct of aluminum and zinc production. Gallium is also extracted from the flue dusts of coal.
Isotopes: Gallium has 24 isotopes whose half-lives are known, with mass numbers 61 to 84. Of these, two are stable: 69Ga and 71Ga with natural abundances of 60.1% and 39.9% respectively.
References
1. Photo by Foobar, GNU FD.
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NEW WORLD ENCYCLOPEDIA
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.
Occurrence and isolation
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 hydrolysisof 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.
History
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.
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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 | ||||||||||||||||||
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| 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/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) | ||||||||||||||||||
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| Atomic properties | |||||||||||||||||||
| Crystal structure | orthorhombic | ||||||||||||||||||
| Oxidation states | 3 (amphoteric oxide) | ||||||||||||||||||
| Electronegativity | 1.81 (Pauling scale) | ||||||||||||||||||
| Ionization energies (more) | 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 | ||||||||||||||||||
| Miscellaneous | |||||||||||||||||||
| Magnetic ordering | no data | ||||||||||||||||||
| Thermal conductivity | (300 K) 40.6 W/(m·K) | ||||||||||||||||||
| Speed of sound (thin rod) | (20 °C) 2740 m/s | ||||||||||||||||||
| Mohs hardness | 1.5 | ||||||||||||||||||
| Brinell hardness | 60 MPa | ||||||||||||||||||
| CAS registry number | 7440-55-3 | ||||||||||||||||||
| Notable isotopes | |||||||||||||||||||
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Notable characteristics
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.
Isotopes
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.
Compounds
Gallium can form a number of compounds. Some of them are mentioned below.
- Gallium(III) arsenide (GaAs): It is an important semiconductor, used for such devices as microwave-frequency integrated circuits (Monolithic Microwave Integrated Circuits, or MMICs), infrared light-emitting diodes (LEDs), laser diodes, and solar cells. Some of its electronic properties are superior to those of silicon. For instance, GaAs devices can function at higher frequencies (above 250 gigahertz), generating less noise, and can be operated at higher power levels. Also, they have a direct band gap, so they can be used to emit light.
- Gallium(III) hydroxide (Ga(OH)3): This is the normal mineral form of gallium in the Earth's crust. It does not occur as a discrete mineral, but gallium replaces aluminum in ores such as bauxite. Gallium hydroxide isamphoteric, that is, it can behave as an acid as well as a base. In strongly acidic conditions, the ion Ga3+ is formed; in strongly basic conditions, Ga(OH)4-, is formed.
- Gallium(III) nitride (GaN): This hard, mechanically stable material is a binary semiconductor with a wide, direct band gap. It is used in optoelectronic devices such as high-brightness, blue LEDs and blue laser diodes. Its sensitivity to ionizing radiation is low, making it a suitable material for solar cell arrays for satellites. It is being investigated for use in high-frequency, high-power transistors that can operate at high temperatures.
- Gallium(III) phosphide (GaP): This solid, crystalline material has the appearance of pale orange pieces. It is odorless and insoluble in water, and it melts at 1,480°C. It is a semiconductor with an indirect band gap (2.26 electronvolt). Sulfur or tellurium may be added as dopants to turn gallium phosphide into an n-type semiconductor; or zinc may be added as a dopant to prepare a p-type semiconductor. GaP is used for the manufacture of low- and standard-brightness red, orange, and green LEDs.
Applications
Gallium, its alloys, and its compounds have many applications. Some of them are listed below.
- The most common applications of gallium are in the form of the semiconductor gallium(III) arsenide. This compound is used mainly for analog integrated circuits, and also for optoelectronic devices such as LEDs and laser diodes.
- Gallium is widely used as a dopant in semiconductors, to produce solid-state devices such as transistors.
- Given that gallium can wet glass or porcelain, it 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.
- When added in quantities up to 2 percent in common solders, gallium can aid wetting and flow characteristics.
- Gallium is used in some high-temperature thermometers.
- An alloy of gallium, indium, and tin (trade name Galinstan) is widely available in medical thermometers (fever thermometers), replacing problematic mercury. This alloy 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.
- Gallium salts, such as gallium citrate or gallium nitrate containing a radioactive isotope of gallium, have been used in nuclear medicine imaging. This use, however, has largely been replaced by FDG PET scans.
- Much research is being devoted to gallium alloys as substitutes for mercury dental amalgams, but such compounds have yet to gain wide acceptance.
- Gallium is the rarest component of new photovoltaic compounds (such as copper indium gallium selenium sulphide or Cu(In,Ga)(Se,S)2, announced by South African researchers) for use in solar panels as an alternative to crystalline silicon, which is currently in short supply.
- 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]
Precautions
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.
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