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Gallium compounds
Gallium compounds are compounds containing the element gallium. These compounds are 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. There are also compounds of gallium with negative oxidation states, ranging from −5 to −1, most of these compounds being magnesium gallides (MgxGay).
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+
. 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).
Alkaline hydroxide solutions dissolve gallium, forming gallate salts (not to be confused with identically named gallic acid salts) containing the Ga(OH)−
4 anion. Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts. Although earlier work suggested Ga(OH)3−
6 as another possible gallate anion, it was not found in later work.
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. Reducing Ga
2O
3 with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, Ga
2O. Ga
2O is a very strong reducing agent, capable of reducing H
2SO
4 to H
2S. It disproportionates at 800 °C back to gallium and Ga
2O
3.
Gallium(III) sulfide, Ga
2S
3, has 3 possible crystal modifications. It can be made by the reaction of gallium with hydrogen sulfide (H
2S) at 950 °C. Alternatively, Ga(OH)
3 can be used at 747 °C:
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. The mercury salt, HgGa
2S
4, 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, Ga
2Se
3 and Ga
2Te
3, have the zincblende structure. They are all semiconductors but are easily hydrolysed and have limited utility.
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Gallium compounds
Gallium compounds are compounds containing the element gallium. These compounds are 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. There are also compounds of gallium with negative oxidation states, ranging from −5 to −1, most of these compounds being magnesium gallides (MgxGay).
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+
. 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).
Alkaline hydroxide solutions dissolve gallium, forming gallate salts (not to be confused with identically named gallic acid salts) containing the Ga(OH)−
4 anion. Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts. Although earlier work suggested Ga(OH)3−
6 as another possible gallate anion, it was not found in later work.
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. Reducing Ga
2O
3 with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, Ga
2O. Ga
2O is a very strong reducing agent, capable of reducing H
2SO
4 to H
2S. It disproportionates at 800 °C back to gallium and Ga
2O
3.
Gallium(III) sulfide, Ga
2S
3, has 3 possible crystal modifications. It can be made by the reaction of gallium with hydrogen sulfide (H
2S) at 950 °C. Alternatively, Ga(OH)
3 can be used at 747 °C:
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. The mercury salt, HgGa
2S
4, 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, Ga
2Se
3 and Ga
2Te
3, have the zincblende structure. They are all semiconductors but are easily hydrolysed and have limited utility.