Gallium

This article is about the chemical element. For other uses, see Gallium (disambiguation).

Gallium is a chemical element with the symbol Ga and atomic number 31. Discovered by the French chemist Paul-Émile Lecoq de Boisbaudran in 1875,^[6] gallium is in group 13 of the periodic table and is similar to the other metals of the group (aluminium, indium, and thallium).

Gallium, ₃₁Ga

Gallium
Pronunciation	/ˈɡæliəm/ ​(GAL-ee-əm)
Appearance	silvery blue
Standard atomic weight Ar°(Ga)
	69.723±0.001 69.723±0.001 (abridged)^[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
Group	group 13 (boron group)
Period	period 4
Block	p-block
Electron configuration	[Ar] 3d^10 4s^2 4p^1
Electrons per shell	2, 8, 18, 3
Physical properties
Phase at STP	solid
Melting point	302.9146 K ​(29.7646 °C, ​85.5763 °F)
Boiling point	2673 K ​(2400 °C, ​4352 °F)^[2]
Density (near r.t.)	5.91 g/cm^3
when liquid (at m.p.)	6.095 g/cm^3
Heat of fusion	5.59 kJ/mol
Heat of vaporization	256 kJ/mol^[2]
Molar heat capacity	25.86 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, −3,^[3] −2, −1, 0, +1, +2, +3^[4] (an amphoteric oxide)
Electronegativity	Pauling scale: 1.81
Ionization energies	1st: 578.8 kJ/mol 2nd: 1979.3 kJ/mol 3rd: 2963 kJ/mol (more)
Atomic radius	empirical: 135 pm
Covalent radius	122±3 pm
Van der Waals radius	187 pm
Spectral lines of gallium
Other properties
Natural occurrence	primordial
Crystal structure	​orthorhombic
Speed of sound thin rod	2740 m/s (at 20 °C)
Thermal expansion	18 µm/(m⋅K) (at 25 °C)
Thermal conductivity	40.6 W/(m⋅K)
Electrical resistivity	270 nΩ⋅m (at 20 °C)
Magnetic ordering	diamagnetic
Molar magnetic susceptibility	−21.6×10^−6 cm^3/mol (at 290 K)^[5]
Young's modulus	9.8 GPa
Poisson ratio	0.47
Mohs hardness	1.5
Brinell hardness	56.8–68.7 MPa
CAS Number	7440-55-3
History
Naming	after Gallia (Latin for: France), homeland of the discoverer
Prediction	Dmitri Mendeleev (1871)
Discovery and first isolation	Lecoq de Boisbaudran (1875)
Isotopes of gallium v e
Main isotopes Decay abun­dance half-life (t1/2) 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
Category: Gallium view talk edit | references

Elemental gallium is a soft, silvery metal at standard temperature and pressure. In its liquid state, it becomes silvery white. If enough force is applied, solid gallium may fracture conchoidally. Since its discovery in 1875, gallium has widely been used to make alloys with low melting points. It is also used in semiconductors, as a dopant in semiconductor substrates.

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. An even lower melting point of −19 °C (−2 °F), well below the freezing point of water, is claimed for the alloy galinstan (62–⁠95% gallium, 5–⁠22% indium, and 0–⁠16% tin by weight), but that may be the freezing point with the effect of supercooling.

Gallium does not occur as a free element in nature, but as gallium(III) compounds in trace amounts in zinc ores (such as sphalerite) and in bauxite. Elemental gallium is a liquid at temperatures greater than 29.76 °C (85.57 °F), and will melt in a person's hands at normal human body temperature of 37.0 °C (98.6 °F).

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 and diode lasers. Gallium is also used in the production of artificial gadolinium gallium garnet for jewelry. Gallium is considered a technology-critical element by the United States National Library of Medicine and Frontiers Media.^[7][8]

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.

Physical properties

 

Crystallization of gallium from the melt

Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium is a silvery blue metal that 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, bismuth, and plutonium.^[9]

Gallium forms alloys with most metals. It readily diffuses into cracks or grain boundaries of some metals such as aluminium, aluminium–zinc alloys^[10] and steel,^[11] causing extreme loss of strength and ductility called liquid metal embrittlement.

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).^[12][13][14] 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.^[15]

The melting point of gallium allows it to melt in the human hand, and then solidify 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.^[16] 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 (as are rubidium and caesium) nor highly toxic (as is 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 nearly nine times higher than its melting point on the absolute scale, the greatest ratio between melting point and boiling point of any element.^[17] Unlike mercury, liquid gallium metal wets glass and skin, along with most other materials (with the exceptions of quartz, graphite, gallium(III) oxide^[18] and PTFE),^[19] making it mechanically more difficult to handle even though it is substantially less toxic and requires far fewer precautions than mercury. Gallium painted onto glass is a brilliant mirror.^[19] 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^[20]

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.^[21] Many stable and metastable phases are found as function of temperature and pressure.^[22]

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 may form 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.^[23] 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.^[9] Ga₂ dimers do not persist in the liquid state and liquid gallium exhibits a complex low-coordinated structure in which each gallium atom is surrounded by 10 others, rather than 11–12 neighbors typical of most liquid metals.^[24][25]

The physical properties of gallium are highly anisotropic, i.e. have different values along the three major crystallographic 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.^[20]

Isotopes

Main article: Isotopes of gallium

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 (mass numbers 56–59) 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.^[26] Gallium-67 and gallium-68 (half-life 67.7 min) are both used in nuclear medicine.

Chemical properties

Main article: Gallium compounds

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 Ga^IGa^IIICl₄; 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.^[27]

Aqueous chemistry

Strong acids dissolve gallium, forming gallium(III) salts such as Ga(NO
₃)
₃ (gallium nitrate). Aqueous solutions of gallium(III) salts contain the hydrated gallium ion, [Ga(H
₂O)
₆]³⁺
.^[28]: 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).^{[29]: 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.^{[30][28]: 1033 [31]} Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts.^[29]: 141 Although earlier work suggested Ga(OH)³⁻
₆ as another possible gallate anion,^[32] it was not found in later work.^[31]

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.^[30] Reducing Ga
₂O
₃ with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, Ga
₂O.^[29]: 285 Ga
₂O is a very strong reducing agent, capable of reducing H
₂SO
₄ to H
₂S.^[29]: 207 It disproportionates at 800 °C back to gallium and Ga
₂O
₃.^[33]

Gallium(III) sulfide, Ga
₂S
₃, has 3 possible crystal modifications.^[33]: 104 It can be made by the reaction of gallium with hydrogen sulfide (H
₂S) at 950 °C.^[29]: 162 Alternatively, Ga(OH)
₃ can be used at 747 °C:^[34]

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.^{[33]: 104–105} The mercury salt, HgGa
₂S
₄, can be used as a phosphor.^[35]

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.^[33]: 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.^[33]: 104

Nitrides and pnictides

 

 

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

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.^[28]: 1034 GaP, GaAs, and GaSb can be synthesized by the direct reaction of gallium with elemental phosphorus, arsenic, or antimony.^[33]: 99 They exhibit higher electrical conductivity than GaN.^[33]: 101 GaP can also be synthesized by reacting Ga
₂O with phosphorus at low temperatures.^[36]

Gallium forms ternary nitrides; for example:^[33]: 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.^[33]: 101

Halides

See also: Gallium 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.^{[29]: 128–129}

Gallium trichloride is formed by the reaction of gallium metal with chlorine gas.^[30] 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.^[29]: 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.^{[29]: 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.^[28]: 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:^{[30][28]: 1036 [37]}

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:^[28]: 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,^[28]: 1031 unlike α-AlH
₃ in which aluminium has a coordination number of 6.^[28]: 1008

Gallane is unstable above −10 °C, decomposing to elemental gallium and hydrogen.^[38]

Organogallium compounds

Main article: Organogallium chemistry

Organogallium compounds are of similar reactivity to organoindium compounds, less reactive than organoaluminium compounds, but more reactive than organothallium compounds.^[39] Alkylgalliums are monomeric. Lewis acidity decreases in the order Al > Ga > In and as a result organogallium compounds do not form bridged dimers as organoaluminium compounds do. Organogallium compounds are also less reactive than organoaluminium compounds. They do form stable peroxides.^[40] 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.^[39]

Gallium trichloride is a common starting reagent for the formation of organogallium compounds, such as in carbogallation reactions.^[41] 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.^[39]

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.^[42]

Comparison between Mendeleev's 1871 predictions and the known properties of gallium^[43]

Property	Mendeleev's predictions	Actual properties
Atomic weight	~68	69.723
Density	5.9 g/cm^3	5.904 g/cm^3
Melting point	Low	29.767 °C
Formula of oxide	M2O3	Ga2O3
Density of oxide	5.5 g/cm^3	5.88 g/cm^3
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.^[43]

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.^[44] Later that year, Lecoq obtained the free metal by electrolysis of the hydroxide in potassium hydroxide solution.^[45]

He named the element "gallia", from Latin Gallia meaning Gaul, after his native land of France. It was later claimed that, in a multilingual pun of a kind favoured 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.^[45]

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 exactly.^[43]

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 bandgap semiconductor in the 1960s ushered in the most important stage in the applications of gallium.^[19] In the late 1960s, the electronics industry started using gallium on a commercial scale to fabricate light emitting diodes, photovoltaics and semiconductors, while the metals industry used it^[46] to reduce the melting point of alloys.^[47]

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.^[48] The abundance in the Earth's crust is approximately 16.9 ppm.^[49] 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,^[48][50] 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.^[48]

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.^[51][52] Some coal flue dusts contain small quantities of gallium, typically less than 1% by weight.^{[53][54][55][56]} 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

 

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).^[57][58] 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.^[57] 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.^[57] 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.^[59]

 

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).^[60] Reserves and resources are not relevant for by-products, since they cannot be extracted independently from the main-products.^[61] 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.^[57] These figures are significantly greater than current production (375 t in 2016).^[62] 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 for low-grade gallium was $120 per kilogram in 2016 and $135–140 per kilogram in 2017.^[63]

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.^[63]

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.^[63]

Applications

Semiconductor applications dominate the commercial demand for gallium, accounting for 98% of the total. The next major application is for gadolinium gallium garnets.^[64]

Semiconductors

 

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.^[51]

Worldwide, gallium arsenide makes up 95% of the annual global gallium consumption.^[59] It amounted to $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.^[63]

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.^[63] 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.^[65]

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 estimated at $370 million in 2016 and $420 million in 2016.^[63]

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.^[66] 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.^[67]

Galinstan and other alloys

 

Galinstan easily wetting a piece of ordinary glass

 

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 melting point of −19 °C (−2.2 °F).^[68] It has been suggested that this family of alloys could also be used to cool computer chips in place of water, and is often used as a replacement for thermal paste in high-performance computing.^[69][70] Gallium alloys have been evaluated as substitutes for mercury dental amalgams, but these materials have yet to see wide acceptance. Liquid alloys containing mostly gallium and indium have been found to precipitate gaseous CO₂ into solid carbon and are being researched as potential methodologies for carbon capture and possibly carbon removal.^[71][72]

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.^[73]

The plutonium used in nuclear weapon pits is stabilized in the δ phase and made machinable by alloying with gallium.^[74][75]

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.^[76] 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.^[77] 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.^[78]

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.^[79][80]

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.^[81][82]

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.^[83]

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.^[84]

Other uses

Neutrino detection: Gallium is used for neutrino detection. Possibly the largest amount of pure gallium ever collected in a single location 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.^[85] 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 (solar neutrino problem) was not explained until better solar neutrino detectors and theories were constructed (see SNO).^[86]

Ion source: 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.^[87]

Lubricants: Gallium serves as an additive in glide wax for skis and other low-friction surface materials.^[88]

Flexible electronics: Materials scientists speculate that the properties of gallium could make it suitable for the development of flexible and wearable devices.^[89][90]

Hydrogen generation: Gallium disrupts the protective oxide layer on aluminium, allowing water to react with the aluminium in AlGa to produce hydrogen gas.^[91]

Humor: 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.^[92]

Gallium in the ocean

Advances in trace element testing have allowed scientists to discover traces of dissolved gallium in the Atlantic and Pacific Oceans ^[93] In recent years, dissolved gallium concentrations have presented in the Beaufort Sea.^[93][94] These reports reflect the possible profiles of the Pacific and Atlantic Ocean waters.^[94] For the Pacific Oceans, typical dissolved gallium concentrations are between 4–6 pmol/kg at depths <~150 m. In comparison, for Atlantic waters 25–28 pmol/kg at depths >~350 m.^[94]

Gallium has entered oceans mainly through aeolian input, but having gallium in our oceans can be used to resolve aluminium distribution in the oceans.^[95] The reason for this is that gallium is geochemically similar to aluminium, just less reactive. Gallium also has a slightly larger surface water residence time than aluminium.^[95] Gallium has a similar dissolved profile similar to that of aluminium, due to this gallium can be used as a tracer for aluminium.^[95] Gallium can also be used as a tracer of aeolian inputs of iron.^[96] Gallium is used as a tracer for iron in the northwest Pacific, south and central Atlantic Oceans.^[96] For example, in the northwest Pacific, low gallium surface waters, in the subpolar region suggest that there is low dust input, which can subsequently explain the following high-nutrient, low-chlorophyll environmental behavior.^[96]

Precautions

Gallium

Hazards
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H290, H318
Precautionary statements	P280, P305, P310, P338, P351^[97]
NFPA 704 (fire diamond)	^[98]   1 0 0

Metallic gallium is not toxic. However, exposure to gallium halide complexes can result in acute toxicity.^[99] The Ga³⁺ ion of soluble gallium salts tends to form the insoluble hydroxide when injected in large doses; precipitation of this hydroxide resulted in nephrotoxicity 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.^[83]

References

1. "Standard Atomic Weights: Gallium". CIAAW. 1987.
2. Zhang Y; Evans JRG; Zhang S (2011). "Corrected Values for Boiling Points and Enthalpies of Vaporization of Elements in Handbooks". J. Chem. Eng. Data. 56 (2): 328–337. doi:10.1021/je1011086.
3. Ga(−3) has been observed in LaGa, see Dürr, Ines; Bauer, Britta; Röhr, Caroline (2011). "Lanthan-Triel/Tetrel-ide La(Al,Ga)_x(Si,Ge)_1-x. Experimentelle und theoretische Studien zur Stabilität intermetallischer 1:1-Phasen" (PDF). Z. Naturforsch. (in German). 66b: 1107–1121.
4. Hofmann, Patrick (1997). Colture. Ein Programm zur interaktiven Visualisierung von Festkörperstrukturen sowie Synthese, Struktur und Eigenschaften von binären und ternären Alkali- und Erdalkalimetallgalliden (PDF) (Thesis) (in German). PhD Thesis, ETH Zurich. p. 72. doi:10.3929/ethz-a-001859893. hdl:20.500.11850/143357. ISBN 978-3728125972.
5. Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
6. Scerri, Eric (2020). The Periodic Table: Its Story and Its Significance. Oxford University Press. p. 149. ISBN 978-0-19-091436-3.
7. Cobelo-García, A.; Filella, M.; Croot, P.; Frazzoli, C.; Du Laing, G.; Ospina-Alvarez, N.; Rauch, S.; Salaun, P.; Schäfer, J.; Zimmermann, S. (2015). "COST action TD1407: network on technology-critical elements (NOTICE)—from environmental processes to human health threats". Environmental Science and Pollution Research International. 22 (19): 15188–15194. doi:10.1007/s11356-015-5221-0. ISSN 0944-1344. PMC 4592495. PMID 26286804.
8. Romero-Freire, Ana; Santos-Echeandía, Juan; Neira, Patricia; Cobelo-García, Antonio (2019). "Less-Studied Technology-Critical Elements (Nb, Ta, Ga, In, Ge, Te) in the Marine Environment: Review on Their Concentrations in Water and Organisms". Frontiers in Marine Science. 6. doi:10.3389/fmars.2019.00532. ISSN 2296-7745.
9. Greenwood and Earnshaw, p. 222
10. Tsai, W. L; Hwu, Y.; Chen, C. H.; Chang, L. W.; Je, J. H.; Lin, H. M.; Margaritondo, G. (2003). "Grain boundary imaging, gallium diffusion and the fracture behavior of Al–Zn Alloy – An in situ study". Nuclear Instruments and Methods in Physics Research Section B. 199: 457–463. Bibcode:2003NIMPB.199..457T. doi:10.1016/S0168-583X(02)01533-1.
11. Vigilante, G. N.; Trolano, E.; Mossey, C. (June 1999). "Liquid Metal Embrittlement of ASTM A723 Gun Steel by Indium and Gallium". Defense Technical Information Center. Retrieved 2009-07-07.
12. Preston–Thomas, H. (1990). "The International Temperature Scale of 1990 (ITS-90)" (PDF). Metrologia. 27 (1): 3–10. Bibcode:1990Metro..27....3P. doi:10.1088/0026-1394/27/1/002. S2CID 250785635. (PDF) from the original on 2007-06-18.
13. "ITS-90 documents at Bureau International de Poids et Mesures".
14. Magnum, B. W.; Furukawa, G. T. (August 1990). (PDF). National Institute of Standards and Technology. NIST TN 1265. Archived from the original (PDF) on 2003-07-04.
15. Strouse, Gregory F. (1999). "NIST realization of the gallium triple point". Proc. TEMPMEKO. 1999 (1): 147–152. Retrieved 2016-10-30.
16. Parravicini, G. B.; Stella, A.; Ghigna, P.; Spinolo, G.; Migliori, A.; d'Acapito, F.; Kofman, R. (2006). "Extreme undercooling (down to 90K) of liquid metal nanoparticles". Applied Physics Letters. 89 (3): 033123. Bibcode:2006ApPhL..89c3123P. doi:10.1063/1.2221395.
17. Greenwood and Earnshaw, p. 224
18. Chen, Ziyu; Lee, Jeong-Bong (2019). "Gallium Oxide Coated Flat Surface as Non-Wetting Surface for Actuation of Liquid Metal Droplets". 2019 IEEE 32nd International Conference on Micro Electro Mechanical Systems (MEMS). pp. 1–4. doi:10.1109/memsys.2019.8870886. ISBN 978-1-7281-1610-5.
19. Greenwood and Earnshaw, p. 221
20. Rosebury, Fred (1992). Handbook of Electron Tube and Vacuum Techniques. Springer. p. 26. ISBN 978-1-56396-121-2.
21. Bernascino, M.; et al. (1995). "Ab initio calculations of structural and electronic properties of gallium solid-state phases". Phys. Rev. B. 52 (14): 9988–9998. Bibcode:1995PhRvB..52.9988B. doi:10.1103/PhysRevB.52.9988. PMID 9980044.
22. "Phase Diagrams of the Elements", David A. Young, UCRL-51902 "Prepared for the U.S. Energy Research & Development Administration under contract No. W-7405-Eng-48". (1975)
23. Greenwood and Earnshaw, p. 223
24. Yagafarov, O. F.; Katayama, Y.; Brazhkin, V. V.; Lyapin, A. G.; Saitoh, H. (November 7, 2012). "Energy dispersive x-ray diffraction and reverse Monte Carlo structural study of liquid gallium under pressure". Physical Review B. 86 (17): 174103. Bibcode:2012PhRvB..86q4103Y. doi:10.1103/PhysRevB.86.174103 – via APS.
25. Drewitt, James W. E.; Turci, Francesco; Heinen, Benedict J.; Macleod, Simon G.; Qin, Fei; Kleppe, Annette K.; Lord, Oliver T. (April 9, 2020). "Structural Ordering in Liquid Gallium under Extreme Conditions". Physical Review Letters. 124 (14): 145501. Bibcode:2020PhRvL.124n5501D. doi:10.1103/PhysRevLett.124.145501. PMID 32338984. S2CID 216177238.
26. Audi, Georges; Bersillon, Olivier; Blachot, Jean; Wapstra, Aaldert Hendrik (2003), "The NUBASE evaluation of nuclear and decay properties", Nuclear Physics A, 729: 3–128, Bibcode:2003NuPhA.729....3A, doi:10.1016/j.nuclphysa.2003.11.001
27. Greenwood and Earnshaw, p. 240
28. Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN 978-0-12-352651-9.
29. Downs, Anthony John (1993). Chemistry of aluminium, gallium, indium, and thallium. Springer. ISBN 978-0-7514-0103-5.
30. Eagleson, Mary, ed. (1994). Concise encyclopedia chemistry. Walter de Gruyter. p. 438. ISBN 978-3-11-011451-5.
31. Sipos, P. L.; Megyes, T. N.; Berkesi, O. (2008). "The Structure of Gallium in Strongly Alkaline, Highly Concentrated Gallate Solutions—a Raman and ⁷¹
    Ga
    -NMR Spectroscopic Study". J Solution Chem. 37 (10): 1411–1418. doi:10.1007/s10953-008-9314-y. S2CID 95723025.
32. Hampson, N. A. (1971). Harold Reginald Thirsk (ed.). Electrochemistry—Volume 3: Specialist periodical report. Great Britain: Royal Society of Chemistry. p. 71. ISBN 978-0-85186-027-5.
33. Greenwood, N. N. (1962). Harry Julius Emeléus; Alan G. Sharpe (eds.). Advances in inorganic chemistry and radiochemistry. Vol. 5. Academic Press. pp. 94–95. ISBN 978-0-12-023605-3.
34. Madelung, Otfried (2004). Semiconductors: data handbook (3rd ed.). Birkhäuser. pp. 276–277. ISBN 978-3-540-40488-0.
35. Krausbauer, L.; Nitsche, R.; Wild, P. (1965). "Mercury gallium sulfide, HgGa
    ₂S
    ₄, a new phosphor". Physica. 31 (1): 113–121. Bibcode:1965Phy....31..113K. doi:10.1016/0031-8914(65)90110-2.
36. Michelle Davidson (2006). Inorganic Chemistry. Lotus Press. p. 90. ISBN 978-81-89093-39-6.
37. Arora, Amit (2005). Text Book Of Inorganic Chemistry. Discovery Publishing House. pp. 389–399. ISBN 978-81-8356-013-9.
38. Downs, Anthony J.; Pulham, Colin R. (1994). Sykes, A. G. (ed.). Advances in Inorganic Chemistry. Vol. 41. Academic Press. pp. 198–199. ISBN 978-0-12-023641-1.
39. Greenwoood and Earnshaw, pp. 262–5
40. Uhl, W. and Halvagar, M. R.; et al. (2009). "Reducing Ga-H and Ga-C Bonds in Close Proximity to Oxidizing Peroxo Groups: Conflicting Properties in Single Molecules". Chemistry: A European Journal. 15 (42): 11298–11306. doi:10.1002/chem.200900746. PMID 19780106.{{cite journal}}: CS1 maint: multiple names: authors list (link)
41. Amemiya, Ryo (2005). "GaCl₃ in Organic Synthesis". European Journal of Organic Chemistry. 2005 (24): 5145–5150. doi:10.1002/ejoc.200500512.
42. Ball, Philip (2002). The Ingredients: A Guided Tour of the Elements. Oxford University Press. p. 105. ISBN 978-0-19-284100-1.
43. Greenwood and Earnshaw, p. 217.
44. Lecoq de Boisbaudran, Paul Émile (1875). "Caractères chimiques et spectroscopiques d'un nouveau métal, le gallium, découvert dans une blende de la mine de Pierrefitte, vallée d'Argelès (Pyrénées)". Comptes Rendus Hebdomadaires des Séances de l'Académie des Sciences. 81: 493–495.
45. Weeks, Mary Elvira (1932). "The discovery of the elements. XIII. Some elements predicted by Mendeleeff". Journal of Chemical Education. 9 (9): 1605–1619. Bibcode:1932JChEd...9.1605W. doi:10.1021/ed009p1605.
46. Petkof, Benjamin (1978). "Gallium" (PDF). GPO. USGS Minerals Yearbook. (PDF) from the original on 2021-06-02.
47. "An Overview of Gallium". AZoNetwork. 18 December 2001.
48. Frenzel, Max (2016). "The distribution of gallium, germanium and indium in conventional and non-conventional resources – Implications for global availability (PDF Download Available)". ResearchGate. doi:10.13140/rg.2.2.20956.18564. Retrieved 2017-06-02.
49. Burton, J. D.; Culkin, F.; Riley, J. P. (2007). "The abundances of gallium and germanium in terrestrial materials". Geochimica et Cosmochimica Acta. 16 (1): 151–180. Bibcode:1959GeCoA..16..151B. doi:10.1016/0016-7037(59)90052-3.
50. Frenzel, Max; Hirsch, Tamino; Gutzmer, Jens (July 2016). "Gallium, germanium, indium, and other trace and minor elements in sphalerite as a function of deposit type — A meta-analysis". Ore Geology Reviews. 76: 52–78. doi:10.1016/j.oregeorev.2015.12.017.
51. Kramer, Deborah A. "Mineral Commodity Summary 2006: Gallium" (PDF). United States Geological Survey. (PDF) from the original on 2008-05-14. Retrieved 2008-11-20.
52. Kramer, Deborah A. "Mineral Yearbook 2006: Gallium" (PDF). United States Geological Survey. (PDF) from the original on 2008-05-09. Retrieved 2008-11-20.
53. Xiao-quan, Shan; Wen, Wang & Bei, Wen (1992). "Determination of gallium in coal and coal fly ash by electrothermal atomic absorption spectrometry using slurry sampling and nickel chemical modification". Journal of Analytical Atomic Spectrometry. 7 (5): 761. doi:10.1039/JA9920700761.
54. "Gallium in West Virginia Coals". West Virginia Geological and Economic Survey. 2002-03-02. from the original on 11 March 2002.
55. Font, O; Querol, Xavier; Juan, Roberto; Casado, Raquel; Ruiz, Carmen R.; López-Soler, Ángel; Coca, Pilar; Peña, Francisco García (2007). "Recovery of gallium and vanadium from gasification fly ash". Journal of Hazardous Materials. 139 (3): 413–23. doi:10.1016/j.jhazmat.2006.02.041. PMID 16600480.
56. Headlee, A. J. W. & Hunter, Richard G. (1953). "Elements in Coal Ash and Their Industrial Significance". Industrial and Engineering Chemistry. 45 (3): 548–551. doi:10.1021/ie50519a028.
57. Frenzel, Max; Ketris, Marina P.; Seifert, Thomas; Gutzmer, Jens (March 2016). "On the current and future availability of gallium". Resources Policy. 47: 38–50. doi:10.1016/j.resourpol.2015.11.005.
58. Frenzel, Max; Hirsch, Tamino; Gutzmer, Jens (2016). "Gallium, germanium, indium, and other trace and minor elements in sphalerite as a function of deposit type — A meta-analysis". Ore Geology Reviews. 76: 52–78. doi:10.1016/j.oregeorev.2015.12.017. ISSN 0169-1368.
59. Moskalyk, R. R. (2003). "Gallium: the backbone of the electronics industry". Minerals Engineering. 16 (10): 921–929. doi:10.1016/j.mineng.2003.08.003.
60. Frenzel, M; Tolosana-Delgado, R; Gutzmer, J (2015). "Assessing the supply potential of high-tech metals – A general method". Resources Policy. 46: 45–58. doi:10.1016/j.resourpol.2015.08.002.
61. Frenzel, Max; Mikolajczak, Claire; Reuter, Markus A.; Gutzmer, Jens (June 2017). "Quantifying the relative availability of high-tech by-product metals – The cases of gallium, germanium and indium". Resources Policy. 52: 327–335. doi:10.1016/j.resourpol.2017.04.008.
62. Gallium – In: USGS Mineral Commodity Summaries (PDF). United States Geological Survey. 2017. (PDF) from the original on 2017-04-27.
63. Galium. USGS (2018)
64. Greber, J. F. (2012) "Gallium and Gallium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, doi:10.1002/14356007.a12_163.
65. Coleman, James J.; Jagadish, Chennupati; Catrina Bryce, A. (2012-05-02). Advances in Semiconductor Lasers. pp. 150–151. ISBN 978-0-12-391066-0.
66. Crisp, D.; Pathare, A.; Ewell, R. C. (2004). "The performance of gallium arsenide/germanium solar cells at the Martian surface". Acta Astronautica. 54 (2): 83–101. Bibcode:2004AcAau..54...83C. doi:10.1016/S0094-5765(02)00287-4.
67. Alberts, V.; Titus J.; Birkmire R. W. (2003). "Material and device properties of single-phase Cu(In,Ga)(Se,S)₂ alloys prepared by selenization/sulfurization of metallic alloys". Thin Solid Films. 451–452: 207–211. Bibcode:2004TSF...451..207A. doi:10.1016/j.tsf.2003.10.092.
68. Surmann, P; Zeyat, H (Nov 2005). "Voltammetric analysis using a self-renewable non-mercury electrode". Analytical and Bioanalytical Chemistry. 383 (6): 1009–13. doi:10.1007/s00216-005-0069-7. ISSN 1618-2642. PMID 16228199. S2CID 22732411.
69. Knight, Will (2005-05-05). . Archived from the original on 2007-02-11. Retrieved 2008-11-20.
70. Martin, Yves. "High Performance Liquid Metal Thermal Interface for Large Volume Production" (PDF).
71. "Technology solidifies carbon dioxide - ASME". www.asme.org. Retrieved 2022-09-05.
72. "New way to turn carbon dioxide into coal could 'rewind the emissions clock'". www.science.org. Retrieved 2022-09-05.
73. United States. Office of Naval Research. Committee on the Basic Properties of Liquid Metals, U.S. Atomic Energy Commission (1954). Liquid-metals handbook. U.S. Govt. Print. Off. p. 128.
74. Sublette, Cary (2001-09-09). "Section 6.2.2.1". Nuclear Weapons FAQ. Retrieved 2008-01-24.
75. Besmann, Theodore M. (2005). "Thermochemical Behavior of Gallium in Weapons-Material-Derived Mixed-Oxide Light Water Reactor (LWR) Fuel". Journal of the American Ceramic Society. 81 (12): 3071–3076. doi:10.1111/j.1151-2916.1998.tb02740.x.
76. Chitambar, Christopher R. (2018). "Chapter 10. Gallium Complexes as Anticancer drugs". In Sigel, Astrid; Sigel, Helmut; Freisinger, Eva; Sigel, Roland K. O. (eds.). Metallo-Drugs: Development and Action of Anticancer Agents. Metal Ions in Life Sciences. Vol. 18. Berlin: de Gruyter GmbH. pp. 281–301. doi:10.1515/9783110470734-016. ISBN 9783110470734. PMID 29394029.
77. . Archived from the original on 2009-06-08. Retrieved 2009-07-07.
78. Bernstein, L. R.; Tanner, T.; Godfrey, C. & Noll, B. (2000). "Chemistry and Pharmacokinetics of Gallium Maltolate, a Compound With High Oral Gallium Bioavailability". Metal-Based Drugs. 7 (1): 33–47. doi:10.1155/MBD.2000.33. PMC 2365198. PMID 18475921.
79. "A Trojan-horse strategy selected to fight bacteria". INFOniac.com. 2007-03-16. Retrieved 2008-11-20.
80. Smith, Michael (2007-03-16). "Gallium May Have Antibiotic-Like Properties". MedPage Today. Retrieved 2008-11-20.
81. Goldberg D. E.; Sharma V.; Oksman A.; Gluzman I. Y.; Wellems T. E.; Piwnica-Worms D. (1997). "Probing the chloroquine resistance locus of Plasmodium falciparum with a novel class of multidentate metal(III) coordination complexes". J. Biol. Chem. 272 (10): 6567–72. doi:10.1074/jbc.272.10.6567. PMID 9045684. S2CID 3408513.
82. Biot, Christophe; Dive, Daniel (2010). "Bioorganometallic Chemistry and Malaria". Medicinal Organometallic Chemistry. Topics in Organometallic Chemistry. Vol. 32. p. 155. doi:10.1007/978-3-642-13185-1_7. ISBN 978-3-642-13184-4. S2CID 85940061.
83. Nordberg, Gunnar F.; Fowler, Bruce A.; Nordberg, Monica (7 August 2014). Handbook on the Toxicology of Metals (4th ed.). Academic Press. pp. 788–90. ISBN 978-0-12-397339-9.
84. Banerjee, Sangeeta Ray; Pomper, Martin G. (June 2013). "Clinical Applications of Gallium-68". Appl. Radiat. Isot. 76: 2–13. doi:10.1016/j.apradiso.2013.01.039. PMC 3664132. PMID 23522791.
85. . 2001-10-19. Archived from the original on 2010-07-05. Retrieved 2009-06-24.
86. "Neutrino Detectors Experiments: GALLEX". 1999-06-26. Retrieved 2008-11-20.
87. "Nano lab produces world's smallest book" 2015-10-13 at the Wayback Machine. Simon Fraser University. 11 April 2007. Retrieved 31 January 2013.
88. US 5069803, Sugimura, Kentaro; Hasimoto, Shoji & Ono, Takayuki, "Use of a synthetic resin composition containing gallium particles in the glide surfacing material of skis and other applications", issued 1995 
89. Kleiner, Kurt (3 May 2022). "Gallium: The liquid metal that could transform soft electronics". Knowable Magazine. doi:10.1146/knowable-050322-2. Retrieved 31 May 2022.
90. Tang, Shi-Yang; Tabor, Christopher; Kalantar-Zadeh, Kourosh; Dickey, Michael D. (26 July 2021). "Gallium Liquid Metal: The Devil's Elixir". Annual Review of Materials Research. 51 (1): 381–408. Bibcode:2021AnRMS..51..381T. doi:10.1146/annurev-matsci-080819-125403. ISSN 1531-7331. S2CID 236566966. Retrieved 31 May 2022.
91. Amberchan, Gabriella; et al. (2022-02-14). "Aluminum Nanoparticles from a Ga–Al Composite for Water Splitting and Hydrogen Generation". ACS Applied Nano Materials. 5 (2): 2636–2643. doi:10.1021/acsanm.1c04331. ISSN 2574-0970.
92. Kean, Sam (2010). The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table of the Elements. Boston: Little, Brown and Company. ISBN 978-0-316-05164-4.
93. Orians, K. J.; Bruland, K. W. (April 1988). "Dissolved Gallium in the Open Ocean". Nature. 332 (21): 717–19. Bibcode:1988Natur.332..717O. doi:10.1038/332717a0. S2CID 4323435.
94. McAlister, Jason A.; Orians, Kristin J. (20 December 2015). "Dissolved Gallium in the Beaufort Sea of the Western Arctic Ocean: A GEOTRACES cruise in the International Polar Year". Marine Chemistry. 177 (Part 1): 101–109. doi:10.1016/j.marchem.2015.05.007. Retrieved 29 August 2021 – via ScienceDirect.
95. Shiller, A. M. (June 1998). "Dissolved Gallium in the Atlantic Ocean". Marine Chemistry. 61 (1): 87–99. doi:10.1016/S0304-4203(98)00009-7.
96. Shiller, A. M.; Bairamadgi, G. R. (August 2006). "Dissolved Gallium in the northwest Pacific and the south and central Atlantic Oceans: Implications for aeolian Fe input and reconsideration of Profiles". Geochemistry, Geophysics, Geosystems. 7 (8): n/a. Bibcode:2006GGG.....7.8M09S. doi:10.1029/2005GC001118. S2CID 129738391.
97. "Gallium 203319". Sigma Aldrich.
98. "MSDS – 203319". Sigma Aldrich.
99. Ivanoff, C. S.; Ivanoff, A. E.; Hottel, T. L. (February 2012). "Gallium poisoning: a rare case report". Food Chem. Toxicol. 50 (2): 212–5. doi:10.1016/j.fct.2011.10.041. PMID 22024274.

Bibliography

• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.

External links

• Gallium at The Periodic Table of Videos (University of Nottingham)
• Safety data sheet at acialloys.com
• High-resolution photographs of molten gallium, gallium crystals and gallium ingots under Creative Commons licence
• – textbook information regarding gallium
• Environmental effects of gallium
• [httpd://minerals.usgs.gov/minerals/pubs/commodity/gallium/460798.pdf Price development of gallium 1959–1998]
• Gallium: A Smart Metal United States Geological Survey
• Thermal conductivity
• Physical and thermodynamical properties of liquid gallium (doc pdf)