Gallium

Gallium is a chemical element; it has symbol Ga and atomic number 31. Discovered by the French chemist Paul-Émile Lecoq de Boisbaudran in Paris, France, 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,, 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, well below the freezing point of water, is claimed for the alloy galinstan, 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 compounds in trace amounts in zinc ores and in bauxite. Elemental gallium is a liquid at temperatures greater than, and will melt in a person's hands at normal human body temperature of.
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 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

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, is just above room temperature, and is approximately the same as the average summer daytime temperatures in Earth's mid-latitudes. This melting point is one of the formal temperature reference points in the International Temperature Scale of 1990 established by the International Bureau of Weights and Measures. The triple point of gallium, 302.9166 K, is used by the US National Institute of Standards and Technology 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 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 nor highly toxic 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 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, 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.
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. 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 3d¹⁰ core. This phenomenon recurs with mercury with its "pseudo-noble-gas" 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, producing a significant difference between the linear 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.

Property	a	b	c
α	16	11	31
ρ	543	174	81
ρ	480	154	71.6
ρ	101	30.8	14.3
ρ	13.8	6.8	1.6

Isotopes

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. Isotopes lighter than gallium-69 usually decay through beta plus decay or electron capture to isotopes of zinc, while isotopes heavier than gallium-71 decay through beta minus decay, possibly with delayed neutron emission, to isotopes of germanium. Gallium-70 can decay both ways, to zinc-70 or to germanium-70.
Gallium-67 and gallium-68 are both used for imaging in nuclear medicine.

Chemical properties

Gallium is found primarily in the +3 oxidation state. The +1 oxidation state is also found in some compounds, although it is less common than it is for gallium's heavier congeners indium and thallium. For example, the very stable GaCl₂ contains both gallium and gallium and can be formulated as Ga^IGa^IIICl₄; in contrast, the monochloride is unstable above 0 °C, disproportionating into elemental gallium and gallium chloride. Compounds containing Ga–Ga bonds are true gallium compounds, such as GaS and the dioxane complex Ga₂Cl₄₂.

Aqueous chemistry

Strong acids dissolve gallium, forming gallium salts such as gallium nitrate|. Aqueous solutions of gallium salts contain the hydrated gallium ion,. Gallium hydroxide,, may be precipitated from gallium solutions by adding ammonia. Dehydrating at 100 °C produces gallium oxide hydroxide, GaO.
Alkaline hydroxide solutions dissolve gallium, forming gallate 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

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 oxide,. Reducing with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium 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 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:
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 sulfide and the green gallium 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.