Gallium

Gallium is a chemical element with symbol Ga and atomic number 31. It is in group 13 of the periodic table, and thus has similarities to the other metals of the group, aluminium, indium, and thallium. Gallium does not occur as a free element in nature, but as gallium(III) compounds in trace amounts in zinc ores and in bauxite.[6] Elemental gallium is a soft, silvery blue metal at standard temperature and pressure, a brittle solid at low temperatures, and a liquid at temperatures greater than 29.76 °C (85.57 °F) (above room temperature, but below the normal human body temperature of 98.6 °F (37.0 °C), hence, the metal will melt in a person's hands).

The melting point of gallium is used as a temperature reference point. Gallium alloys are used in thermometers as a non-toxic and environmentally friendly alternative to mercury, and can withstand higher temperatures than mercury. The alloy galinstan (70% gallium, 21.5% indium, and 10% tin) has an even lower melting point of −19 °C (−2 °F), well below the freezing point of water.

Since its discovery in 1875, gallium has been used to make alloys with low melting points. It is also used in semiconductors as a dopant in semiconductor substrates.

Gallium is predominantly used in electronics. Gallium arsenide, the primary chemical compound of gallium in electronics, is used in microwave circuits, high-speed switching circuits, and infrared circuits. Semiconducting gallium nitride and indium gallium nitride produce blue and violet light-emitting diodes (LEDs) and diode lasers. Gallium is also used in the production of artificial gadolinium gallium garnet for jewelry.

Gallium has no known natural role in biology. Gallium(III) behaves in a similar manner to ferric salts in biological systems and has been used in some medical applications, including pharmaceuticals and radiopharmaceuticals.

Gallium,  31Ga
Gallium crystals
Gallium
Pronunciation/ˈɡæliəm/ ​(GAL-ee-əm)
Appearancesilvery blue
Standard atomic weight Ar, std(Ga)69.723(1)[1]
Gallium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Al
↑
Ga
↓
In
zinc ← gallium → germanium
Atomic number (Z)31
Groupgroup 13 (boron group)
Periodperiod 4
Blockp-block
Element category  post-transition metal
Electron configuration[Ar] 3d10 4s2 4p1
Electrons per shell
2, 8, 18, 3
Physical properties
Phase at STPsolid
Melting point302.9146 K ​(29.7646 °C, ​85.5763 °F)
Boiling point2673 K ​(2400 °C, ​4352 °F)[2]
Density (near r.t.)5.91 g/cm3
when liquid (at m.p.)6.095 g/cm3
Heat of fusion5.59 kJ/mol
Heat of vaporization256 kJ/mol[2]
Molar heat capacity25.86 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 1310 1448 1620 1838 2125 2518
Atomic properties
Oxidation states−5, −4, −2, −1, +1, +2, +3[3] (an amphoteric oxide)
ElectronegativityPauling scale: 1.81
Ionization energies
  • 1st: 578.8 kJ/mol
  • 2nd: 1979.3 kJ/mol
  • 3rd: 2963 kJ/mol
  • (more)
Atomic radiusempirical: 135 pm
Covalent radius122±3 pm
Van der Waals radius187 pm
Color lines in a spectral range
Spectral lines of gallium
Other properties
Natural occurrenceprimordial
Crystal structure ​orthorhombic
Orthorhombic crystal structure for gallium
Speed of sound thin rod2740 m/s (at 20 °C)
Thermal expansion18 µm/(m·K) (at 25 °C)
Thermal conductivity40.6 W/(m·K)
Electrical resistivity270 nΩ·m (at 20 °C)
Magnetic orderingdiamagnetic
Magnetic susceptibility−21.6·10−6 cm3/mol (at 290 K)[4]
Young's modulus9.8 GPa
Poisson ratio0.47
Mohs hardness1.5
Brinell hardness56.8–68.7 MPa
CAS Number7440-55-3
History
Namingafter Gallia (Latin for: France), homeland of the discoverer
PredictionDmitri Mendeleev (1871)
Discovery and first isolationLecoq de Boisbaudran (1875)
Main isotopes of gallium
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
66Ga syn 9.5 h β+ 66Zn
67Ga syn 3.3 d ε 67Zn
68Ga syn 1.2 h β+ 68Zn
69Ga 60.11% stable
70Ga syn 21 min β− 70Ge
ε 70Zn
71Ga 39.89% stable
72Ga syn 14.1 h β− 72Ge
73Ga syn 4.9 h β− 73Ge

Physical properties

Gallium kristallisiert
Crystallization of gallium from the melt

Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium metal has a silvery color and its solid metal fractures conchoidally like glass. Gallium liquid expands by 3.10% when it solidifies; therefore, it should not be stored in glass or metal containers because the container may rupture when the gallium changes state. Gallium shares the higher-density liquid state with a short list of other materials that includes water, silicon, germanium, antimony, bismuth, and plutonium.[7]

Gallium attacks most other metals by diffusing into the metal lattice. For example, it diffuses into the grain boundaries of aluminium-zinc alloys[8] and steel,[9] making them very brittle. Gallium easily alloys with many metals, and is used in small quantities in the plutonium-gallium alloy in the plutonium cores of nuclear bombs to stabilize the plutonium crystal structure.[10]

The melting point of gallium, at 302.9146 K (29.7646 °C, 85.5763 °F), is just above room temperature, and is approximately the same as the average summer daytime temperatures in Earth's mid-latitudes. This melting point (mp) is one of the formal temperature reference points in the International Temperature Scale of 1990 (ITS-90) established by the International Bureau of Weights and Measures (BIPM).[11][12][13] The triple point of gallium, 302.9166 K (29.7666 °C, 85.5799 °F), is used by the US National Institute of Standards and Technology (NIST) in preference to the melting point.[14]

The melting point of gallium allows it to melt in the human hand, and then refreeze if removed. The liquid metal has a strong tendency to supercool below its melting point/freezing point: Ga nanoparticles can be kept in the liquid state below 90 K.[15] Seeding with a crystal helps to initiate freezing. Gallium is one of the four non-radioactive metals (with caesium, rubidium, and mercury) that are known to be liquid at, or near, normal room temperature. Of the four, gallium is the only one that is neither highly reactive (rubidium and caesium) nor highly toxic (mercury) 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 for having (unlike mercury) a low vapor pressure at high temperatures. Gallium's boiling point, 2673 K, is more than eight times higher than its melting point on the absolute scale, the greatest ratio between melting point and boiling point of any element.[16] Unlike mercury, liquid gallium metal wets glass and skin, along with most other materials (with the exceptions of quartz, graphite, and Teflon), making it mechanically more difficult to handle even though it is substantially less toxic and requires far fewer precautions. Gallium painted onto glass is a brilliant mirror.[17] For this reason as well as the metal contamination and freezing-expansion problems, samples of gallium metal are usually supplied in polyethylene packets within other containers.

Properties of gallium for different crystal axes[18]
Property a b c
α (~25 °C, µm/m) 16 11 31
ρ (29.7 °C, nΩ·m) 543 174 81
ρ (0 °C, nΩ·m) 480 154 71.6
ρ (77 K, nΩ·m) 101 30.8 14.3
ρ (4.2 K, pΩ·m) 13.8 6.8 1.6

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. Within a unit cell, each atom has only one nearest neighbor (at a distance of 244 pm). The remaining six unit cell neighbors are spaced 27, 30 and 39 pm farther away, and they are grouped in pairs with the same distance.[19] Many stable and metastable phases are found as function of temperature and pressure.[20]

The bonding between the two nearest neighbors is covalent; hence Ga2 dimers are seen as the fundamental building blocks of the crystal. This explains the low melting point relative to the neighbor elements, aluminium and indium. This structure is strikingly similar to that of iodine and forms because of interactions between the single 4p electrons of gallium atoms, further away from the nucleus than the 4s electrons and the [Ar]3d10 core. This phenomenon recurs with mercury with its "pseudo-noble-gas" [Xe]4f145d106s2 electron configuration, which is liquid at room temperature.[21] The 3d10 electrons do not shield the outer electrons very well from the nucleus and hence the first ionisation energy of gallium is greater than that of aluminium.[7]

The physical properties of gallium are highly anisotropic, i.e. have different values along the three major crystallographical axes a, b, and c (see table), producing a significant difference between the linear (α) and volume thermal expansion coefficients. The properties of gallium are strongly temperature-dependent, particularly near the melting point. For example, the coefficient of thermal expansion increases by several hundred percent upon melting.[18]

Isotopes

Gallium has 31 known isotopes, ranging in mass number from 56 to 86. Only two isotopes are stable and occur naturally, gallium-69 and gallium-71. Gallium-69 is more abundant: it makes up about 60.1% of natural gallium, while gallium-71 makes up the remaining 39.9%. All the other isotopes are radioactive, with gallium-67 being the longest-lived (half-life 3.261 days). Isotopes lighter than gallium-69 usually decay through beta plus decay (positron emission) or electron capture to isotopes of zinc, although the lightest few (with mass numbers 56 through 59) decay through prompt proton emission. Isotopes heavier than gallium-71 decay through beta minus decay (electron emission), possibly with delayed neutron emission, to isotopes of germanium, while gallium-70 can decay through both beta minus decay and electron capture. Gallium-67 is unique among the light isotopes in having only electron capture as a decay mode, as its decay energy is not sufficient to allow positron emission.[22] Gallium-67 and gallium-68 (half-life 67.7 min) are both used in nuclear medicine.

Chemical properties

Gallium is found primarily in the +3 oxidation state. The +1 oxidation state is also found in some compounds, although it is less common than it is for gallium's heavier congeners indium and thallium. For example, the very stable GaCl2 contains both gallium(I) and gallium(III) and can be formulated as GaIGaIIICl4; in contrast, the monochloride is unstable above 0 °C, disproportionating into elemental gallium and gallium(III) chloride. Compounds containing Ga–Ga bonds are true gallium(II) compounds, such as GaS (which can be formulated as Ga24+(S2−)2) and the dioxan complex Ga2Cl4(C4H8O2)2.[23]

Aqueous chemistry

Strong acids dissolve gallium, forming gallium(III) salts such as Ga2(SO4)3 (gallium sulfate) and Ga(NO
3)
3 (gallium nitrate). Aqueous solutions of gallium(III) salts contain the hydrated gallium ion, [Ga(H
2O)
6]3+
.[24]:1033 Gallium(III) hydroxide, Ga(OH)
3, may be precipitated from gallium(III) solutions by adding ammonia. Dehydrating Ga(OH)
3 at 100 °C produces gallium oxide hydroxide, GaO(OH).[25]:140–141

Alkaline hydroxide solutions dissolve gallium, forming gallate salts (not to be confused with identically-named gallic acid salts) containing the Ga(OH)−
4 anion.[26][24]:1033[27] Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts.[25]:141 Although earlier work suggested Ga(OH)3−
6 as another possible gallate anion,[28] it was not found in later work.[27]

Oxides and chalcogenides

Gallium reacts with the chalcogens only at relatively high temperatures. At room temperature, gallium metal is not reactive with air and water because it forms a passive, protective oxide layer. At higher temperatures, however, it reacts with atmospheric oxygen to form gallium(III) oxide, Ga
2O
3.[26] Reducing Ga
2O
3 with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, Ga
2O.[25]:285 Ga
2O is a very strong reducing agent, capable of reducing H
2SO
4 to H
2S.[25]:207 It disproportionates at 800 °C back to gallium and Ga
2O
3.[29]

Gallium(III) sulfide, Ga
2S
3, has 3 possible crystal modifications.[29]:104 It can be made by the reaction of gallium with hydrogen sulfide (H
2S) at 950 °C.[25]:162 Alternatively, Ga(OH)
3 can be used at 747 °C:[30]

2 Ga(OH)
3 + 3 H
2S → Ga
2S
3 + 6 H
2O

Reacting a mixture of alkali metal carbonates and Ga
2O
3 with H
2S leads to the formation of thiogallates containing the [Ga
2S
4]2−
 anion. Strong acids decompose these salts, releasing H
2S in the process.[29]:104–105 The mercury salt, HgGa
2S
4, can be used as a phosphor.[31]

Gallium also forms sulfides in lower oxidation states, such as gallium(II) sulfide and the green gallium(I) sulfide, the latter of which is produced from the former by heating to 1000 °C under a stream of nitrogen.[29]:94

The other binary chalcogenides, Ga
2Se
3 and Ga
2Te
3, have the zincblende structure. They are all semiconductors but are easily hydrolysed and have limited utility.[29]:104

Nitrides and pnictides

Gallium nitride (left) and gallium arsenide (right) wafers

Crystal-GaN
Gallium Arsenide (GaAs) 2" wafer

Gallium reacts with ammonia at 1050 °C to form gallium nitride, GaN. Gallium also forms binary compounds with phosphorus, arsenic, and antimony: gallium phosphide (GaP), gallium arsenide (GaAs), and gallium antimonide (GaSb). These compounds have the same structure as ZnS, and have important semiconducting properties.[24]:1034 GaP, GaAs, and GaSb can be synthesized by the direct reaction of gallium with elemental phosphorus, arsenic, or antimony.[29]:99 They exhibit higher electrical conductivity than GaN.[29]:101 GaP can also be synthesized by reacting Ga
2O with phosphorus at low temperatures.[32]

Gallium forms ternary nitrides; for example:[29]:99

Li
3Ga + N
2 → Li
3GaN
2

Similar compounds with phosphorus and arsenic are possible: Li
3GaP
2 and Li
3GaAs
2. These compounds are easily hydrolyzed by dilute acids and water.[29]:101

Halides

Gallium(III) oxide reacts with fluorinating agents such as HF or F
2 to form gallium(III) fluoride, GaF
3. It is an ionic compound strongly insoluble in water. However, it dissolves in hydrofluoric acid, in which it forms an adduct with water, GaF
3·3H
2O. Attempting to dehydrate this adduct forms GaF
2OH·nH
2O. The adduct reacts with ammonia to form GaF
3·3NH
3, which can then be heated to form anhydrous GaF
3.[25]:128–129

Gallium trichloride is formed by the reaction of gallium metal with chlorine gas.[26] Unlike the trifluoride, gallium(III) chloride exists as dimeric molecules, Ga
2Cl
6, with a melting point of 78 °C. Eqivalent compounds are formed with bromine and iodine, Ga
2Br
6 and Ga
2I
6.[25]:133

Like the other group 13 trihalides, gallium(III) halides are Lewis acids, reacting as halide acceptors with alkali metal halides to form salts containing GaX−
4 anions, where X is a halogen. They also react with alkyl halides to form carbocations and GaX−
4.[25]:136–137

When heated to a high temperature, gallium(III) halides react with elemental gallium to form the respective gallium(I) halides. For example, GaCl
3 reacts with Ga to form GaCl:

2 Ga + GaCl
3 ⇌ 3 GaCl (g)

At lower temperatures, the equilibrium shifts toward the left and GaCl disproportionates back to elemental gallium and GaCl
3. GaCl can also be produced by reacting Ga with HCl at 950 °C; the product can be condensed as a red solid.[24]:1036

Gallium(I) compounds can be stabilized by forming adducts with Lewis acids. For example:

GaCl + AlCl
3 → Ga+
[AlCl
4]−

The so-called "gallium(II) halides", GaX
2, are actually adducts of gallium(I) halides with the respective gallium(III) halides, having the structure Ga+
[GaX
4]−
. For example:[26][24]:1036[33]

GaCl + GaCl
3 → Ga+
[GaCl
4]−

Hydrides

Like aluminium, gallium also forms a hydride, GaH
3, known as gallane, which may be produced by reacting lithium gallanate (LiGaH
4) with gallium(III) chloride at −30 °C:[24]:1031

3 LiGaH
4 + GaCl
3 → 3 LiCl + 4 GaH
3

In the presence of dimethyl ether as solvent, GaH
3 polymerizes to (GaH
3)
n. If no solvent is used, the dimer Ga
2H
6 (digallane) is formed as a gas. Its structure is similar to diborane, having two hydrogen atoms bridging the two gallium centers,[24]:1031 unlike α-AlH
3 in which aluminium has a coordination number of 6.[24]:1008

Gallane is unstable above −10 °C, decomposing to elemental gallium and hydrogen.[34]

Organogallium compounds

Organogallium compounds are of similar reactivity to organoindium compounds, less reactive than organoaluminium compounds, but more reactive than organothallium compounds.[35] Alkylgalliums are monomeric. Lewis acidity decreases in the order Al > Ga > In and as a result organogallium compounds do not form bridged dimers as organoaluminum compounds do. Organogallium compounds are also less reactive than organoaluminum compounds. They do form stable peroxides.[36] These alkylgalliums are liquids at room temperature, having low melting points, and are quite mobile and flammable. Triphenylgallium is monomeric in solution, but its crystals form chain structures due to weak intermolecluar Ga···C interactions.[35]

Gallium trichloride is a common starting reagent for the formation of organogallium compounds, such as in carbogallation reactions.[37] Gallium trichloride reacts with lithium cyclopentadienide in diethyl ether to form the trigonal planar gallium cyclopentadienyl complex GaCp3. Gallium(I) forms complexes with arene ligands such as hexamethylbenzene. Because this ligand is quite bulky, the structure of the [Ga(η6-C6Me6)]+ is that of a half-sandwich. Less bulky ligands such as mesitylene allow two ligands to be attached to the central gallium atom in a bent sandwich structure. Benzene is even less bulky and allows the formation of dimers: an example is [Ga(η6-C6H6)2] [GaCl4]·3C6H6.[35]

History

Small gallium droplets fusing together

In 1871, the existence of gallium was first predicted by Russian chemist Dmitri Mendeleev, who named it "eka-aluminium" from its position in his periodic table. He also predicted several properties of eka-aluminium that correspond closely to the real properties of gallium, such as its density, melting point, oxide character and bonding in chloride.[38]

Comparison between Mendeleev's 1871 predictions and the known properties of gallium[39]
Property Mendeleev's predictions Actual properties
Atomic weight ~68 69.723
Density 5.9 g/cm3 5.904 g/cm3
Melting point Low 29.767 °C
Formula of oxide M2O3 Ga2O3
Density of oxide 5.5 g/cm3 5.88 g/cm3
Nature of hydroxide amphoteric amphoteric

Mendeleev further predicted that eka-aluminium would be discovered by means of the spectroscope, and that metallic eka-aluminium would dissolve slowly in both acids and alkalis and would not react with air. He also predicted that M2O3 would dissolve in acids to give MX3 salts, that eka-aluminium salts would form basic salts, that eka-aluminium sulfate should form alums, and that anhydrous MCl3 should have a greater volatility than ZnCl2: all of these predictions turned out to be true.[39]

Gallium was discovered using spectroscopy by French chemist Paul Emile Lecoq de Boisbaudran in 1875 from its characteristic spectrum (two violet lines) in a sample of sphalerite.[40] Later that year, Lecoq obtained the free metal by electrolysis of the hydroxide in potassium hydroxide solution. He named the element "gallia", from Latin Gallia meaning Gaul, after his native land of France. It was later claimed that, in one of those multilingual puns so beloved by men of science in the 19th century, he had also named gallium after himself: "Le coq" is French for "the rooster" and the Latin word for "rooster" is "gallus". In an 1877 article, Lecoq denied this conjecture.[41] Originally, de Boisbaudran determined the density of gallium as 4.7 g/cm3, the only property that failed to match Mendeleev's predictions; Mendeleev then wrote to him and suggested that he should remeasure the density, and de Boisbaudran then obtained the correct value of 5.9 g/cm3, that Mendeleev had predicted almost exactly.[39]

From its discovery in 1875 until the era of semiconductors, the primary uses of gallium were high-temperature thermometrics and metal alloys with unusual properties of stability or ease of melting (some such being liquid at room temperature). The development of gallium arsenide as a direct band gap semiconductor in the 1960s ushered in the most important stage in the applications of gallium.[17]

Occurrence

Gallium does not exist as a free element in the Earth's crust, and the few high-content minerals, such as gallite (CuGaS2), are too rare to serve as a primary source.[42] The abundance in the Earth's crust is approximately 16.9 ppm.[43] This is comparable to the crustal abundances of lead, cobalt and niobium. Yet unlike these elements, gallium does not form its own ore deposits with concentrations of > 0.1 wt.% in ore. Rather it occurs at trace concentrations similar to the crustal value in zinc ores,[42][44] and at somewhat higher values (~ 50 ppm) in aluminium ores, from both of which it is extracted as a by-product. This lack of independent deposits is due to gallium's geochemical behaviour, showing no strong enrichment in the processes relevant to the formation of most ore deposits.[42]

The United States Geological Survey (USGS) estimates that more than 1 million tons of gallium is contained in known reserves of bauxite and zinc ores.[45][46] Some coal flue dusts contain small quantities of gallium, typically less than 1% by weight.[47][48][49][50] However, these amounts are not extractable without mining of the host materials (see below). Thus, the availability of gallium is fundamentally determined by the rate at which bauxite, zinc ores (and coal) are extracted.

Production and availability

6N Gallium sealed in vacuum ampoule
99.9999% (6N) gallium sealed in vacuum ampoule

Gallium is produced exclusively as a by-product during the processing of the ores of other metals. Its main source material is bauxite, the chief ore of aluminium, but minor amounts are also extracted from sulfidic zinc ores (sphalerite being the main host mineral). In the past, certain coals were an important source.

During the processing of bauxite to alumina in the Bayer process, gallium accumulates in the sodium hydroxide liquor. From this it can be extracted by a variety of methods. The most recent is the use of ion-exchange resin.[6] Achievable extraction efficiencies critically depend on the original concentration in the feed bauxite. At a typical feed concentration of 50 ppm, about 15% of the contained gallium is extractable.[6] The remainder reports to the red mud and aluminium hydroxide streams. Gallium is removed from the ion-exchange resin in solution. Electrolysis then gives gallium metal. For semiconductor use, it is further purified with zone melting or single-crystal extraction from a melt (Czochralski process). Purities of 99.9999% are routinely achieved and commercially available.[51]

Bauxite Jamaica 1984
Bauxite mine in Jamaica (1984)

Its by-product status means that gallium production is constrained by the amount of bauxite, sulfidic zinc ores (and coal) extracted per year. Therefore, its availability needs to be discussed in terms of supply potential. The supply potential of a by-product is defined as that amount which is economically extractable from its host materials per year under current market conditions (i.e. technology and price).[52] Reserves and resources are not relevant for by-products, since they cannot be extracted independently from the main-products.[53] Recent estimates put the supply potential of gallium at a minimum of 2,100 t/yr from bauxite, 85 t/yr from sulfidic zinc ores, and potentially 590 t/yr from coal.[6] These figures are significantly greater than current production (375 t in 2016).[54] Thus, major future increases in the by-product production of gallium will be possible without significant increases in production costs or price. The average price in for low-grade gallium was $120 per kilogram in 2016 and $135-140 per kilogram in 2017.[55]

In 2017, the world's production of low-grade gallium was ca. 315 tons — an increase of 15% from 2016. China, Japan, South Korea, Russia, and Ukraine were the leading producers, while Germany ceased primary production of gallium in 2016. The yield of high-purity gallium was ca. 180 tons, mostly originating from China, Japan, Slovakia, UK and U.S. The 2017 world annual production capacity was estimated at 730 tons for low-grade and 320 tons for refined gallium.[55]

China produced ca. 250 tons of low-grade gallium in 2016 and ca. 300 tons in 2017. It also accounted for more than half of global LED production.[55]

Applications

Semiconductor applications dominate the commercial demand for gallium, accounting for 98% of the total. The next major application is for gadolinium gallium garnets.[56]

Semiconductors

Blue LED and Reflection
Gallium-based blue LEDs

Extremely high-purity (>99.9999%) gallium is commercially available to serve the semiconductor industry. Gallium arsenide (GaAs) and gallium nitride (GaN) used in electronic components represented about 98% of the gallium consumption in the United States in 2007. About 66% of semiconductor gallium is used in the U.S. in integrated circuits (mostly gallium arsenide), such as the manufacture of ultra-high-speed logic chips and MESFETs for low-noise microwave preamplifiers in cell phones. About 20% of this gallium is used in optoelectronics.[45]

Worldwide, gallium arsenide makes up 95% of the annual global gallium consumption.[51] It amounted $7.5 billion in 2016, with 53% originating from cell phones, 27% from wireless communications, and the rest from automotive, consumer, fiber-optic, and military applications. The recent increase in GaAs consumption is mostly related to the emergence of 3G and 4G smartphones, which use 10 times more GaAs than older models.[55]

Gallium arsenide and gallium nitride can also be found in a variety of optoelectronic devices which had a market share of $15.3 billion in 2015 and $18.5 billion in 2016.[55] Aluminium gallium arsenide (AlGaAs) is used in high-power infrared laser diodes. The semiconductors gallium nitride and indium gallium nitride are used in blue and violet optoelectronic devices, mostly laser diodes and light-emitting diodes. For example, gallium nitride 405 nm diode lasers are used as a violet light source for higher-density Blu-ray Disc compact data disc drives.[57]

Other major application of gallium nitride are cable television transmission, commercial wireless infrastructure, power electronics, and satellites. The GaN radio frequency device market alone was was estimated at $370 million in 2016 and $420 million in 2016.[55]

Multijunction photovoltaic cells, developed for satellite power applications, are made by molecular-beam epitaxy or metalorganic vapour-phase epitaxy of thin films of gallium arsenide, indium gallium phosphide, or indium gallium arsenide. The Mars Exploration Rovers and several satellites use triple-junction gallium arsenide on germanium cells.[58] Gallium is also a component in photovoltaic compounds (such as copper indium gallium selenium sulfide Cu(In,Ga)(Se,S)2) used in solar panels as a cost-efficient alternative to crystalline silicon.[59]

Galinstan and other alloys

Galinstan on glass
Galinstan from a broken thermometer, easily wetting a piece of glass
Gallium alloy 3D prints (26519727708)
Owing to their low melting points, gallium and its alloys can be shaped into various 3D forms using 3D printing and additive manufacturing

Gallium readily alloys with most metals, and is used as an ingredient in low-melting alloys. The nearly eutectic alloy of gallium, indium, and tin is a room temperature liquid used in medical thermometers. This alloy, with the trade-name Galinstan (with the "-stan" referring to the tin, stannum in Latin), has a low freezing point of −19 °C (−2.2 °F).[60] It has been suggested that this family of alloys could also be used to cool computer chips in place of water.[61] Gallium alloys have been evaluated as substitutes for mercury dental amalgams, but these materials have yet to see wide acceptance.

Because gallium wets glass or porcelain, gallium can be used to create brilliant mirrors. When the wetting action of gallium-alloys is not desired (as in Galinstan glass thermometers), the glass must be protected with a transparent layer of gallium(III) oxide.[62]

The plutonium used in nuclear weapon pits is stabilized in the δ phase and made machinable by alloying with gallium.[63]

Biomedical applications

Although gallium has no natural function in biology, gallium ions interact with processes in the body in a manner similar to iron(III). Because these processes include inflammation, a marker for many disease states, several gallium salts are used (or are in development) as pharmaceuticals and radiopharmaceuticals in medicine. Interest in the anticancer properties of gallium emerged when it was discovered that 67Ga(III) citrate injected in tumor-bearing animals localized to sites of tumor. Clinical trials have shown gallium nitrate to have antineoplastic activity against non-Hodgkin’s lymphoma and urothelial cancers. A new generation of gallium-ligand complexes such as tris(8-quinolinolato)gallium(III) (KP46) and gallium maltolate has emerged.[64] Gallium nitrate (brand name Ganite) has been used as an intravenous pharmaceutical to treat hypercalcemia associated with tumor metastasis to bones. Gallium is thought to interfere with osteoclast function, and the therapy may be effective when other treatments have failed.[65] Gallium maltolate, an oral, highly absorbable form of gallium(III) ion, is an anti-proliferative to pathologically proliferating cells, particularly cancer cells and some bacteria that accept it in place of ferric iron (Fe3+). Researchers are conducting clinical and preclinical trials on this compound as a potential treatment for a number of cancers, infectious diseases, and inflammatory diseases.[66]

When gallium ions are mistakenly taken up in place of iron(III) by bacteria such as Pseudomonas, the ions interfere with respiration, and the bacteria die. This happens because iron is redox-active, allowing the transfer of electrons during respiration, while gallium is redox-inactive.[67][68]

A complex amine-phenol Ga(III) compound MR045 is selectively toxic to parasites resistant to chloroquine, a common drug against malaria. Both the Ga(III) complex and chloroquine act by inhibiting crystallization of hemozoin, a disposal product formed from the digestion of blood by the parasites.[69][70]

Radiogallium salts

Gallium-67 salts such as gallium citrate and gallium nitrate are used as radiopharmaceutical agents in the nuclear medicine imaging known as gallium scan. The radioactive isotope 67Ga is used, and the compound or salt of gallium is unimportant. The body handles Ga3+ in many ways as though it were Fe3+, and the ion is bound (and concentrates) in areas of inflammation, such as infection, and in areas of rapid cell division. This allows such sites to be imaged by nuclear scan techniques.[71]

Gallium-68, a positron emitter with a half-life of 68 min, is now used as a diagnostic radionuclide in PET-CT when linked to pharmaceutical preparations such as DOTATOC, a somatostatin analogue used for neuroendocrine tumors investigation, and DOTA-TATE, a newer one, used for neuroendocrine metastasis and lung neuroendocrine cancer, such as certain types of microcytoma. Gallium-68's preparation as a pharmaceutical is chemical, and the radionuclide is extracted by elution from germanium-68, a synthetic radioisotope of germanium, in gallium-68 generators.[72]

Other uses

Gallium is used for neutrino detection. Possibly the largest amount of pure gallium ever collected in a single spot is the Gallium-Germanium Neutrino Telescope used by the SAGE experiment at the Baksan Neutrino Observatory in Russia. This detector contains 55–57 tonnes (~9 cubic metres) of liquid gallium.[73] Another experiment was the GALLEX neutrino detector operated in the early 1990s in an Italian mountain tunnel. The detector contained 12.2 tons of watered gallium-71. Solar neutrinos caused a few atoms of 71Ga to become radioactive 71Ge, which were detected. This experiment showed that the solar neutrino flux is 40% less than theory predicted. This deficit was not explained until better solar neutrino detectors and theories were constructed (see SNO).[74]

Gallium is also used as a liquid metal ion source for a focused ion beam. For example, a focused gallium-ion beam was used to create the world's smallest book, Teeny Ted from Turnip Town.[75] Another use of gallium is as an additive in glide wax for skis, and other low-friction surface materials.[76]

A well-known practical joke among chemists is to fashion gallium spoons and use them to serve tea to unsuspecting guests, since gallium has a similar appearance to its lighter homolog aluminium. The spoons then melt in the hot tea.[77]

Gallium
Hazards
GHS pictograms The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)
GHS signal word Danger
H290, H318
P280, P305, P351, P338, P310[78]
NFPA 704

Precautions

Metallic gallium is not toxic. However, exposure to gallium halide complexes can result in acute toxicity.[80] The Ga3+ ion of soluble gallium salts tends to form the insoluble hydroxide when injected in large doses; precipitation of this hydroxide resulted in renal toxicity in animals. In lower doses, soluble gallium is tolerated well and does not accumulate as a poison, instead being excreted mostly through urine. Excretion of gallium occurs in two phases: the first phase has a biological half-life of 1 hour, while the second has a biological half-life of 25 hours.[71]

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Bibliography

External links

Boron group

The boron group are the chemical elements in group 13 of the periodic table, comprising boron (B), aluminium (Al), gallium (Ga), indium (In), thallium (Tl), and perhaps also the chemically uncharacterized nihonium (Nh). The elements in the boron group are characterized by having three electrons in their outer energy levels (valence layers). These elements have also been referred to as the triels.Boron is classified as a metalloid while the rest, with the possible exception of nihonium, are considered post-transition metals. Boron occurs sparsely, probably because bombardment by the subatomic particles produced from natural radioactivity disrupts its nuclei. Aluminium occurs widely on earth, and indeed is the third most abundant element in the Earth's crust (8.3%). Gallium is found in the earth with an abundance of 13 ppm. Indium is the 61st most abundant element in the earth's crust, and thallium is found in moderate amounts throughout the planet. Nihonium is never found in nature and therefore is termed a synthetic element.

Several group 13 elements have biological roles in the ecosystem. Boron is a trace element in humans and is essential for some plants. Lack of boron can lead to stunted plant growth, while an excess can also cause harm by inhibiting growth. Aluminium has neither a biological role nor significant toxicity and is considered safe. Indium and gallium can stimulate metabolism; gallium is credited with the ability to bind itself to iron proteins. Thallium is highly toxic, interfering with the function of numerous vital enzymes, and has seen use as a pesticide.

Copper indium gallium selenide

Copper indium gallium (di)selenide (CIGS) is a I-III-VI2 semiconductor material composed of copper, indium, gallium, and selenium. The material is a solid solution of copper indium selenide (often abbreviated "CIS") and copper gallium selenide. It has a chemical formula of CuIn(1-x)Ga(x)Se2 where the value of x can vary from 0 (pure copper indium selenide) to 1 (pure copper gallium selenide). CIGS is a tetrahedrally bonded semiconductor, with the chalcopyrite crystal structure, and a bandgap varying continuously with x from about 1.0 eV (for copper indium selenide) to about 1.7 eV (for copper gallium selenide).

Digallane

Digallane (systematically named digallane(6) and di-μ-hydrido-bis(dihydridogallium)) is an inorganic compound with the chemical formula GaH2(H)2GaH2 (also written [{GaH2(μ-H)}2] or [Ga2H6]). It is the dimer of the monomeric compound gallane. The eventual preparation of the pure compound, reported in 1989,

was hailed as a "tour de force." Digallane had been reported as early as 1941 by Wiberg; however, this claim could not be verified by later work by Greenwood and others.

Galium

Galium is a large genus of annual and perennial herbaceous plants in the family Rubiaceae, occurring in the temperate zones of both the Northern and Southern Hemispheres. Some species are informally known as bedstraw.There are over 600 species of Galium, with estimates of 629 to 650 as of 2013. The field madder, Sherardia arvensis, is a close relative and may be confused with a tiny bedstraw. Asperula is also a closely related genus; some species of Galium (such as woodruff, G. odoratum) are occasionally placed therein.

Gallium(I) oxide

Gallium(I) oxide or gallium suboxide is an inorganic compound with the formula Ga2O.

Gallium(II) sulfide

Gallium(II) sulfide, GaS, is a chemical compound of gallium and sulfur. The normal form of gallium(II) sulfide as made from the elements has a hexagonal layer structure containing Ga24+ units which have a Ga-Ga distance of 248pm. This layer structure is similar to GaTe, GaSe and InSe. An unusual metastable form, with a distorted wurtzite structure has been reported as being produced using MOCVD. The metal organic precursors were di-tert-butyl gallium dithiocarbamates, for example GatBu2(S2CNMe2) and this was deposited onto GaAs. The structure of the GaS produced in this way is presumably Ga2+ S2−.Single layers of gallium sulfide are dynamically stable two-dimensional semiconductors, in which the valence band has an inverted Mexican-hat shape, leading to a Lifshitz transition as the hole-doping is increased.

Gallium(III) bromide

Gallium(III) bromide (GaBr3) is a chemical compound, and one of four Gallium trihalides.

Gallium(III) fluoride

Gallium(III) fluoride (GaF3) is a chemical compound. It is a white solid that melts under pressure above 1000 °C but sublimes around 950 °C. It has the FeF3 structure where the gallium atoms are 6-coordinate. GaF3 can be prepared by reacting F2 or HF with Ga2O3 or by thermal decomposition of (NH4)3GaF6. GaF3 is virtually insoluble in water. Solutions of GaF3 in HF can be evaporated to form the trihydrate, GaF3·3H2O, which on heating gives a hydrated form of GaF2(OH). Gallium(III) fluoride reacts with mineral acids to form hydrofluoric acid.

Gallium(III) oxide

Gallium(III) trioxide is an inorganic compound with the formula Ga2O3. It exists as several polymorphs, all of which are white, water-insoluble solids. Although no commercial applications exist, Ga2O3 is an intermediate in the purification of gallium, which is consumed almost exclusively as gallium arsenide.

Gallium arsenide

Gallium arsenide (GaAs) is a compound of the elements gallium and arsenic. It is a III-V direct bandgap semiconductor with a zinc blende crystal structure.

Gallium arsenide is used in the manufacture of devices such as microwave frequency integrated circuits, monolithic microwave integrated circuits, infrared light-emitting diodes, laser diodes, solar cells and optical windows.GaAs is often used as a substrate material for the epitaxial growth of other III-V semiconductors including indium gallium arsenide, aluminum gallium arsenide and others.

Gallium nitride

Gallium nitride (GaN) is a binary III/V direct bandgap semiconductor commonly used in light-emitting diodes since the 1990s. The compound is a very hard material that has a Wurtzite crystal structure. Its wide band gap of 3.4 eV affords it special properties for applications in optoelectronic, high-power and high-frequency devices. For example, GaN is the substrate which makes violet (405 nm) laser diodes possible, without use of nonlinear optical frequency-doubling.

Its sensitivity to ionizing radiation is low (like other group III nitrides), making it a suitable material for solar cell arrays for satellites. Military and space applications could also benefit as devices have shown stability in radiation environments.Because GaN transistors can operate at much higher temperatures and work at much higher voltages than gallium arsenide (GaAs) transistors, they make ideal power amplifiers at microwave frequencies. In addition, GaN offers promising characteristics for THz devices.

Gallium scan

A gallium scan (also called "gallium imaging") is a type of nuclear medicine test that uses either a gallium-67 (67Ga) or gallium-68 (68Ga) radiopharmaceutical to obtain images of a specific type of tissue, or disease state of tissue. Gallium salts like gallium citrate and gallium nitrate may be used. The form of salt is not important, since it is the freely dissolved gallium ion Ga3+ which is active. Both 67Ga and 68Ga salts have similar uptake mechanisms. Gallium can also be used in other forms, for example 68Ga-PSMA is used for cancer imaging. The gamma emission of gallium 67 is imaged by a gamma camera, while the positron emission of gallium 68 is imaged by positron emission tomography (PET).

Gallium salts are taken up by tumors, inflammation, and both acute and chronic infection, allowing these pathological processes to be imaged. Gallium is particularly useful in imaging osteomyelitis that involves the spine, and in imaging older and chronic infections that may be the cause of a fever of unknown origin.

Gallium trichloride

Gallium trichloride is the chemical compound with the formula GaCl3. Solid gallium trichloride exists as a dimer with the formula Ga2Cl6. It is colourless and soluble in virtually all solvents, even alkanes, which is truly unusual for a metal halide. It is the main precursor to most derivatives of gallium and a reagent in organic synthesis.As a Lewis acid, GaCl3 is milder than aluminium trichloride. Gallium(III) is easier to reduce than Al(III), so the chemistry of reduced gallium compounds is more extensive than for aluminium. Ga2Cl4 is known whereas the corresponding Al2Cl4 is not. The coordination chemistry of Ga(III) and Fe(III) are similar, and gallium(III) compounds have been used as diamagnetic analogues of ferric compounds.

Indium-111 WBC scan

The indium white blood cell scan, also called "indium leukocyte imaging", "indium-111 scan", or simply "indium scan", is a nuclear medicine procedure in which white blood cells (mostly neutrophils) are removed from the patient, tagged with the radioisotope Indium-111, and then injected intravenously into the patient. The tagged leukocytes subsequently localize to areas of relatively new infection. The study is particularly helpful in differentiating conditions such as osteomyelitis from decubitus ulcers for assessment of route and duration of antibiotic therapy.In imaging of infections, the gallium scan has a sensitivity advantage over the indium white blood cell scan in imaging osteomyelitis (bone infection) of the spine, lung infections and inflammation, and in detecting chronic infections. In part, this is because gallium binds to neutrophil membranes, even after neutrophil death, whereas localization of neutrophils labeled with indium requires them to be in relatively good functional order. However, indium leukocyte imaging is better at localizing acute (i.e., new) infections, where live neutrophils are still rapidly and actively localizing to the infection, for imaging for osteomyelitis that does not involve the spine, and for locating abdominal and pelvic infections.

Both the gallium scan and indium-111 white blood cell imaging may be used to image fever of unknown origin (elevated temperature without an explanation). However, the indium leukocyte scan will localize only to the approximately 25% of such cases which are caused by acute infections, while gallium is more broadly sensitive, localizing to other sources of fever, such as chronic infections and tumors. Gallium may be a better choice for spleen study because gallium does not normally accumulate in the spleen.

Isotopes of gallium

Natural gallium (31Ga) consists of a mixture of two stable isotopes: gallium-69 and gallium-71. The most commercially important radioisotopes are gallium-67 and gallium-68.

Gallium-67 (half-life 3.3 days) is a gamma-emitting isotope (the gamma emitted immediately after electron-capture) used in standard nuclear medical imaging, in procedures usually referred to as gallium scans. It is usually used as the free ion, Ga3+. It is the longest-lived radioisotope of gallium.

The shorter-lived gallium-68 (half-life 68 minutes) is a positron-emitting isotope generated from germanium-68 in gallium-68 generators, for use in a small minority of diagnostic PET scans. For this use, it is usually attached as a tracer to a carrier molecule, which gives the resulting radiopharmaceutical a different tissue-uptake specificity from the ionic Ga-67 radioisotope normally used in standard gallium scans.

List of semiconductor materials

Semiconductor materials are nominally small band gap insulators. The defining property of a semiconductor material is that it can be doped with impurities that alter its electronic properties in a controllable way.Because of their application in the computer and photovoltaic industry—in devices such as transistors, lasers, and solar cells—the search for new semiconductor materials and the improvement of existing materials is an important field of study in materials science.

Most commonly used semiconductor materials are crystalline inorganic solids. These materials are classified according to the periodic table groups of their constituent atoms.

Different semiconductor materials differ in their properties. Thus, in comparison with silicon, compound semiconductors have both advantages and disadvantages. For example, gallium arsenide (GaAs) has six times higher electron mobility than silicon, which allows faster operation; wider band gap, which allows operation of power devices at higher temperatures, and gives lower thermal noise to low power devices at room temperature; its direct band gap gives it more favorable optoelectronic properties than the indirect band gap of silicon; it can be alloyed to ternary and quaternary compositions, with adjustable band gap width, allowing light emission at chosen wavelengths, and allowing e.g. matching to wavelengths with lowest losses in optical fibers. GaAs can be also grown in a semi-insulating form, which is suitable as a lattice-matching insulating substrate for GaAs devices. Conversely, silicon is robust, cheap, and easy to process, whereas GaAs is brittle and expensive, and insulation layers can not be created by just a few an oxide layer; GaAs is therefore used only where silicon is not sufficient.

By alloying multiple compounds, some semiconductor materials are tunable, e.g., in band gap or lattice constant. The result is ternary, quaternary, or even quinary compositions. Ternary compositions allow adjusting the band gap within the range of the involved binary compounds; however, in case of combination of direct and indirect band gap materials there is a ratio where indirect band gap prevails, limiting the range usable for optoelectronics; e.g. AlGaAs LEDs are limited to 660 nm by this. Lattice constants of the compounds also tend to be different, and the lattice mismatch against the substrate, dependent on the mixing ratio, causes defects in amounts dependent on the mismatch magnitude; this influences the ratio of achievable radiative/nonradiative recombinations and determines the luminous efficiency of the device. Quaternary and higher compositions allow adjusting simultaneously the band gap and the lattice constant, allowing increasing radiant efficiency at wider range of wavelengths; for example AlGaInP is used for LEDs. Materials transparent to the generated wavelength of light are advantageous, as this allows more efficient extraction of photons from the bulk of the material. That is, in such transparent materials, light production is not limited to just the surface. Index of refraction is also composition-dependent and influences the extraction efficiency of photons from the material.

Ribonucleotide reductase inhibitor

Ribonucleotide reductase inhibitors are a family of anti-cancer drugs that interfere with the growth of tumor cells by blocking the formation of deoxyribonucleotides (building blocks of DNA).

Examples include:

motexafin gadolinium.

hydroxyurea

fludarabine, cladribine, gemcitabine, tezacitabine, and triapine

gallium maltolate, gallium nitrate

SAGE (Soviet–American Gallium Experiment)

SAGE (Soviet–American Gallium Experiment, or sometimes Russian-American Gallium Experiment) is a collaborative experiment devised by several prominent physicists to measure the solar neutrino flux.

Semiconductor

A semiconductor material has an electrical conductivity value falling between that of a metal, like copper, gold, etc. and an insulator, such as glass. Their resistance decreases as their temperature increases, which is behaviour opposite to that of a metal. Their conducting properties may be altered in useful ways by the deliberate, controlled introduction of impurities ("doping") into the crystal structure. Where two differently-doped regions exist in the same crystal, a semiconductor junction is created. The behavior of charge carriers which include electrons, ions and electron holes at these junctions is the basis of diodes, transistors and all modern electronics. Some examples of semiconductors are silicon, germanium, and gallium arsenide. After silicon, gallium arsenide is the second most common semiconductor used in laser diodes, solar cells, microwave frequency integrated circuits, and others. Silicon is a critical element for fabricating most electronic circuits.

Semiconductor devices can display a range of useful properties such as passing current more easily in one direction than the other, showing variable resistance, and sensitivity to light or heat. Because the electrical properties of a semiconductor material can be modified by doping, or by the application of electrical fields or light, devices made from semiconductors can be used for amplification, switching, and energy conversion.

The conductivity of silicon is increased by adding a small amount of pentavalent (antimony, phosphorus, or arsenic) or trivalent (boron, gallium, indium) atoms (part in 108). This process is known as doping and resulting semiconductors are known as doped or extrinsic semiconductors. Apart from doping, the conductivity of a semiconductor can equally be improved by increasing its temperature. This is contrary to the behaviour of a metal in which conductivity decreases with increase in temperature.

The modern understanding of the properties of a semiconductor relies on quantum physics to explain the movement of charge carriers in a crystal lattice. Doping greatly increases the number of charge carriers within the crystal. When a doped semiconductor contains mostly free holes it is called "p-type", and when it contains mostly free electrons it is known as "n-type". The semiconductor materials used in electronic devices are doped under precise conditions to control the concentration and regions of p- and n-type dopants. A single semiconductor crystal can have many p- and n-type regions; the p–n junctions between these regions are responsible for the useful electronic behavior.

Although some pure elements and many compounds display semiconductor properties, silicon, germanium, and compounds of gallium are the most widely used in electronic devices. Elements near the so-called "metalloid staircase", where the metalloids are located on the periodic table, are usually used as semiconductors.

Some of the properties of semiconductor materials were observed throughout the mid 19th and first decades of the 20th century. The first practical application of semiconductors in electronics was the 1904 development of the cat's-whisker detector, a primitive semiconductor diode used in early radio receivers. Developments in quantum physics in turn allowed the development of the transistor in 1947 and the integrated circuit in 1958.

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