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= Gallium =
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Introduction
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Gallium is a chemical element; it has symbol Ga and atomic number 31.
Discovered by the French chemist Paul-Émile Lecoq de Boisbaudran in
1875,
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, 29.7646 °C, 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. A melting point of -19 °C,
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 rather 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, and will melt in a person's hands
at normal human body temperature of 37.0 °C.
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. It 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
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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's volume expands by 3.10% when it
changes from a liquid to a solid so care must be taken when storing it
in containers that may rupture when it changes state. Gallium shares
the higher-density liquid state with a short list of other materials
that includes water, silicon, germanium, bismuth, and plutonium.
Gallium forms alloys with most metals. It readily diffuses into cracks
or grain boundaries of some metals such as aluminium, aluminium-zinc
alloys and steel, 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). 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.
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. 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, 2676 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. Unlike mercury, liquid
gallium metal wets glass and skin, along with most other materials
(with the exceptions of quartz, graphite, gallium(III) oxide and
PTFE), 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. 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
!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. Many stable and metastable
phases are found as function of temperature and pressure.
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. 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. 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.
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.
Isotopes
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Gallium has 30 known isotopes, ranging in mass number from 60 to 89.
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, while isotopes heavier than gallium-71
decay through beta minus decay (electron emission), possibly with
delayed neutron emission, to isotopes of germanium. 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. Gallium-67 and gallium-68 (half-life 67.7 min) are both used
in nuclear medicine.
Chemical properties
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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.
Aqueous chemistry
===================
Strong acids dissolve gallium, forming gallium(III) salts such as
gallium(III) nitrate (gallium nitrate). Aqueous solutions of
gallium(III) salts contain the hydrated gallium ion, . Gallium(III)
hydroxide, , may be precipitated from gallium(III) solutions by adding
ammonia. Dehydrating at 100 °C produces gallium oxide hydroxide,
GaO(OH).
Alkaline hydroxide solutions dissolve gallium, forming 'gallate' salts
(not to be confused with identically named gallic acid salts)
containing the anion. Gallium hydroxide, which is amphoteric, also
dissolves in alkali to form gallate salts. Although earlier work
suggested as another possible gallate anion, it was not found in
later work.
Oxides and chalcogenides
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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, . Reducing with elemental gallium in vacuum
at 500 °C to 700 °C yields the dark brown gallium(I) oxide, . is a
very strong reducing agent, capable of reducing sulfuric acid to
hydrogen sulfide. It disproportionates at 800 °C back to gallium and .
Gallium(III) sulfide, , has 3 possible crystal modifications. It can
be made by the reaction of gallium with hydrogen sulfide () at 950 °C.
Alternatively, can be used at 747 °C:
:2 + 3 → + 6
Reacting a mixture of alkali metal carbonates and with leads to the
formation of 'thiogallates' containing the anion. Strong acids
decompose these salts, releasing in the process. The mercury salt, ,
can be used as a phosphor.
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.
The other binary chalcogenides, and , have the zincblende structure.
They are all semiconductors but are easily hydrolysed and have limited
utility.
Nitrides and pnictides
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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. GaP, GaAs, and GaSb
can be synthesized by the direct reaction of gallium with elemental
phosphorus, arsenic, or antimony. They exhibit higher electrical
conductivity than GaN. GaP can also be synthesized by reacting with
phosphorus at low temperatures.
Gallium forms ternary nitrides; for example:
: + →
Similar compounds with phosphorus and arsenic are possible: and .
These compounds are easily hydrolyzed by dilute acids and water.
Halides
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Gallium(III) oxide reacts with fluorinating agents such as HF or
fluorine to form gallium(III) fluoride, . It is an ionic compound
strongly insoluble in water. However, it dissolves in hydrofluoric
acid, in which it forms an adduct with water, . Attempting to
dehydrate this adduct forms . The adduct reacts with ammonia to form ,
which can then be heated to form anhydrous .
Gallium trichloride is formed by the reaction of gallium metal with
chlorine gas. Unlike the trifluoride, gallium(III) chloride exists as
dimeric molecules, , with a melting point of 78 °C. Equivalent
compounds are formed with bromine and iodine, gallium(III) bromide and
gallium(III) iodide.
Like the other group 13 trihalides, gallium(III) halides are Lewis
acids, reacting as halide acceptors with alkali metal halides to form
salts containing anions, where X is a halogen. They also react with
alkyl halides to form carbocations and .
When heated to a high temperature, gallium(III) halides react with
elemental gallium to form the respective gallium(I) halides. For
example, reacts with Ga to form :
:2 Ga + 3 GaCl (g)
At lower temperatures, the equilibrium shifts toward the left and GaCl
disproportionates back to elemental gallium and . GaCl can also be
produced by reacting Ga with HCl at 950 °C; the product can be
condensed as a red solid.
Gallium(I) compounds can be stabilized by forming adducts with Lewis
acids. For example:
:GaCl + →
The so-called "gallium(II) halides", , are actually adducts of
gallium(I) halides with the respective gallium(III) halides, having
the structure . For example:
:GaCl + →
Hydrides
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Like aluminium, gallium also forms a hydride, , known as 'gallane',
which may be produced by reacting lithium gallanate () with
gallium(III) chloride at −30 °C:
:3 + → 3 LiCl + 4
In the presence of dimethyl ether as solvent, polymerizes to . If no
solvent is used, the dimer ('digallane') is formed as a gas. Its
structure is similar to diborane, having two hydrogen atoms bridging
the two gallium centers, unlike α-aluminium hydride in which aluminium
has a coordination number of 6.
Gallane is unstable above −10 °C, decomposing to elemental gallium and
hydrogen.
Organogallium compounds
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Organogallium compounds are of similar reactivity to organoindium
compounds, less reactive than organoaluminium compounds, but more
reactive than organothallium compounds. 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. 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.
Gallium trichloride is a common starting reagent for the formation of
organogallium compounds, such as in carbogallation reactions. 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.
History
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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.
: Comparison between Mendeleev's 1871 predictions and the known
properties of gallium
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.
Gallium was discovered using spectroscopy by French chemist Paul-Émile
Lecoq de Boisbaudran in 1875 from its characteristic spectrum (two
violet lines) in a sample of sphalerite. Later that year, Lecoq
obtained the free metal by electrolysis of the hydroxide in potassium
hydroxide solution.
He named the element "gallia", from Latin meaning 'Gaul', a name for
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: is French for 'the
rooster', and the Latin word for 'rooster' is . In an 1877 article,
Lecoq denied this conjecture.
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.
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. 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 to
reduce the melting point of alloys.
First blue gallium nitride LED were developed in 1971-1973, but they
were feeble. Only in the early 1990s Shuji Nakamura managed to combine
GaN with indium gallium nitride and develop the modern blue LED, now
making the basis of ubiquitous white LEDs, which Nichia commercialized
in 1993. He and two other Japanese scientists received a Nobel in
Physics in 2014 for this work.
Global gallium production slowly grew from several tens of t/year in
the 1970s til ca. 2010, when it passed 100 t/yr and rapidly
accelerated, by 2024 reaching about 450 t/yr.
Occurrence
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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. The abundance in the Earth's crust is
approximately 16.9 ppm. It is the 34th most abundant element in the
crust. 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, 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.
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. Some coal flue dusts contain small quantities of gallium,
typically less than 1% by weight. 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
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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. 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. 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.
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). Reserves and resources are not relevant for by-products, since
they 'cannot' be extracted independently from the main-products.
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. These figures are significantly
greater than current production (375 t in 2016). 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.
In 2017, the world's production of low-grade gallium was tons--a
decrease 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.
China produced tons of low-grade gallium in 2016 and tons in 2017.
It also accounted for more than half of global LED production. As of
July 2023, China accounted for between 80% and 95% of its production.
Applications
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Semiconductor applications dominate the commercial demand for gallium,
accounting for 98% of the total. The next major application is for
gadolinium gallium garnets. As of 2022, 44% of world use went to light
fixtures and 36% to integrated circuits, with smaller shares equal to
~7% going to photovoltaics and magnets each.
Semiconductors
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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.
Worldwide, gallium arsenide makes up 95% of the annual global gallium
consumption. 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 employ up to 10 times the amount of
GaAs in older models.
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. 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.
Other major applications 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.
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. Gallium is also a component in photovoltaic compounds (such as
copper indium gallium selenium sulfide ) used in solar panels as a
cost-efficient alternative to crystalline silicon.
Galinstan and other alloys
============================
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, in Latin), has a low melting point of −19 °C
(−2.2 °F). 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. 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.
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.
Due to their high surface tension and deformability, gallium-based
liquid metals can be used to create actuators by controlling the
surface tension. Researchers have demonstrated the potentials of using
liquid metal actuators as artificial muscle in robotic actuation.
The plutonium used in nuclear weapon pits is stabilized in the δ phase
and made machinable by alloying with gallium.
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. 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. 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.
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.
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.
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.
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.
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. 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).
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'.
Lubricants: Gallium serves as an additive in glide wax for skis and
other low-friction surface materials.
Flexible electronics: Materials scientists speculate that the
properties of gallium could make it suitable for the development of
flexible and wearable devices.
Hydrogen generation: Gallium disrupts the protective oxide layer on
aluminium, allowing water to react with the aluminium in AlGa to
produce hydrogen gas.
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.
Gallium in the ocean
======================================================================
Advances in trace element testing have allowed scientists to discover
traces of dissolved gallium in the Atlantic and Pacific Oceans. In
recent years, dissolved gallium concentrations have presented in the
Beaufort Sea. These reports reflect the possible profiles of the
Pacific and Atlantic Ocean waters. For the Pacific Oceans, typical
dissolved gallium concentrations are between 4 and 6 pmol/kg at depths
<~150 m. In comparison, for Atlantic waters 25-28 pmol/kg at depths
>~350 m.
Gallium has entered oceans mainly through aeolian input, but having
gallium in our oceans can be used to resolve aluminium distribution in
the oceans. 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. Gallium has a
similar dissolved profile similar to that of aluminium, due to this
gallium can be used as a tracer for aluminium. Gallium can also be
used as a tracer of aeolian inputs of iron. Gallium is used as a
tracer for iron in the northwest Pacific, south and central Atlantic
Oceans. 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.
Precautions
======================================================================
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Metallic gallium is not toxic. However, several gallium compounds are
toxic.
Gallium halide complexes can be toxic. 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.
Inhaled Ga2O3 particles are probably toxic.
External links
======================================================================
* [
http://www.periodicvideos.com/videos/031.htm Gallium] at 'The
Periodic Table of Videos' (University of Nottingham)
* Safety data sheet at
[
http://www.acialloys.com/wp-content/uploads/msds/ga.html
acialloys.com]
* [
http://france-gallium.com/photos-gallium.php High-resolution
photographs of molten gallium, gallium crystals and gallium ingots
under Creative Commons licence]
* [
https://www.lenntech.com/periodic/elements/ga.htm Textbook
information regarding gallium]
*
[
https://minerals.usgs.gov/minerals/pubs/commodity/gallium/index.html
Environmental effects of gallium]
*
[
https://www.usgs.gov/centers/national-minerals-information-center/gallium-statistics-and-information
Gallium Statistics and Information]
* [
https://permanent.fdlp.gov/gpo42065/fs2013-3006.pdf Gallium: A
Smart Metal] United States Geological Survey
*
[
http://arquivo.pt/wayback/20160515021612/http://jcp.aip.org/resource/1/jcpsa6/v26/i4/p784_s1?isAuthorized=no
Thermal conductivity]
* [
http://www.impmc.jussieu.fr/%7Eayrinhac/documents/Ga_data.pdf
Physical and thermodynamical properties of liquid gallium] (doc pdf)
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Gallium
