Gallium compounds

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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. [1] There are also compounds of gallium with negative oxidation states, ranging from -5 to -1, most of these compounds being magnesium gallides (MgxGay).

Contents

Aqueous chemistry

Gallium nitrate nonahydrate Gallium nitrate nonanhydrate.jpg
Gallium nitrate nonahydrate

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+
. [2] : 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). [3] : 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. [4] [2] : 1033  [5] Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts. [3] : 141  Although earlier work suggested Ga(OH)3−
6 as another possible gallate anion, [6] it was not found in later work. [5]

Oxides and chalcogenides

A gallium trioxide crystal Gallium(III) oxide crystal.jpg
A gallium trioxide crystal

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. [4] Reducing Ga
2O
3 with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, Ga
2O. [3] : 285 Ga
2O is a very strong reducing agent, capable of reducing H
2SO
4 to H
2S. [3] : 207  It disproportionates at 800 °C back to gallium and Ga
2O
3. [7]

Gallium(III) sulfide, Ga
2S
3, has 3 possible crystal modifications. [7] : 104  It can be made by the reaction of gallium with hydrogen sulfide (H
2S) at 950 °C. [3] : 162  Alternatively, Ga(OH)
3 can be used at 747 °C: [8]

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. [7] : 104–105  The mercury salt, HgGa
2S
4, can be used as a phosphor. [9]

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. [7] : 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. [7] : 104 

Nitrides and pnictides

Crystal-GaN.jpg
GaP-wafer.jpg
Gallium Arsenide (GaAs) 2" wafer.jpg
Gallium nitride (left), gallium phosphide (middle) 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. [2] : 1034  GaP, GaAs, and GaSb can be synthesized by the direct reaction of gallium with elemental phosphorus, arsenic, or antimony. [7] : 99  They exhibit higher electrical conductivity than GaN. [7] : 101  GaP can also be synthesized by reacting Ga
2O with phosphorus at low temperatures. [10]

Gallium forms ternary nitrides; for example: [7] : 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. [7] : 101 

Halides

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

Gallium trichloride is formed by the reaction of gallium metal with chlorine gas. [4] 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. [3] : 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. [3] : 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. [2] : 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: [4] [2] : 1036  [11]

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: [2] : 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, [2] : 1031  unlike α-AlH
3 in which aluminium has a coordination number of 6. [2] : 1008 

Gallane is unstable above −10 °C, decomposing to elemental gallium and hydrogen. [12]

Organogallium compounds

Organogallium compounds are of similar reactivity to organoindium compounds, less reactive than organoaluminium compounds, but more reactive than organothallium compounds. [13] 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. [14] 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. [13]

Gallium trichloride is a common starting reagent for the formation of organogallium compounds, such as in carbogallation reactions. [15] 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. [13]

See also

Related Research Articles

<span class="mw-page-title-main">Gallium</span> Chemical element, symbol Ga and atomic number 31

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, gallium is in group 13 of the periodic table and is similar to the other metals of the group.

Boron trifluoride is the inorganic compound with the formula BF3. This pungent, colourless, and toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.

<span class="mw-page-title-main">Triphenylphosphine</span> Chemical compound

Triphenylphosphine (IUPAC name: triphenylphosphane) is a common organophosphorus compound with the formula P(C6H5)3 and often abbreviated to PPh3 or Ph3P. It is widely used in the synthesis of organic and organometallic compounds. PPh3 exists as relatively air stable, colorless crystals at room temperature. It dissolves in non-polar organic solvents such as benzene and diethyl ether.

<span class="mw-page-title-main">Gold(III) chloride</span> Chemical compound

Gold(III) chloride, traditionally called auric chloride, is an inorganic compound of gold and chlorine with the molecular formula Au2Cl6. The "III" in the name indicates that the gold has an oxidation state of +3, typical for many gold compounds. It has two forms, the monohydrate (AuCl3·H2O) and the anhydrous form, which are both hygroscopic and light-sensitive solids. This compound is a dimer of AuCl3. This compound has a few uses, such as an oxidizing agent and for catalyzing various organic reactions.

<span class="mw-page-title-main">Calcium sulfide</span> Chemical compound of formula CaS

Calcium sulfide is the chemical compound with the formula CaS. This white material crystallizes in cubes like rock salt. CaS has been studied as a component in a process that would recycle gypsum, a product of flue-gas desulfurization. Like many salts containing sulfide ions, CaS typically has an odour of H2S, which results from small amount of this gas formed by hydrolysis of the salt.

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<span class="mw-page-title-main">Gallium(III) bromide</span> Chemical compound

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<span class="mw-page-title-main">Selenium compounds</span> Chemical compounds containing selenium

Selenium compounds commonly exist in the oxidation states −2, +2, +4, and +6.

Iodine can form compounds using multiple oxidation states. Iodine is quite reactive, but it is much less reactive than the other halogens. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide, iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.

<span class="mw-page-title-main">Gallium trichloride</span> Chemical compound

Gallium trichloride is the chemical compound with the formula GaCl3. Solid gallium trichloride exists as a dimer with the formula Ga2Cl6. It is colourless and soluble in virtually all solvents, even alkanes, which is truly unusual for a metal halide. It is the main precursor to most derivatives of gallium and a reagent in organic synthesis.

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<span class="mw-page-title-main">Organogallium chemistry</span> Chemistry of Organogallium compounds

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Organoplatinum chemistry is the chemistry of organometallic compounds containing a carbon to platinum chemical bond, and the study of platinum as a catalyst in organic reactions. Organoplatinum compounds exist in oxidation state 0 to IV, with oxidation state II most abundant. The general order in bond strength is Pt-C (sp) > Pt-O > Pt-N > Pt-C (sp3). Organoplatinum and organopalladium chemistry are similar, but organoplatinum compounds are more stable and therefore less useful as catalysts.

<span class="mw-page-title-main">Lead compounds</span> Type of compound

Compounds of lead exist with lead in two main oxidation states: +2 and +4. The former is more common. Inorganic lead(IV) compounds are typically strong oxidants or exist only in highly acidic solutions.

<span class="mw-page-title-main">Polonium dioxide</span> Chemical compound

Polonium dioxide (also known as polonium(IV) oxide) is a chemical compound with the formula PoO2. It is one of three oxides of polonium, the other two being polonium monoxide (PoO) and polonium trioxide (PoO3). It is a pale yellow crystalline solid at room temperature. Under lowered pressure (such as a vacuum), it decomposes into elemental polonium and oxygen at 500 °C. It is the most stable oxide of polonium and is an interchalcogen.

<span class="mw-page-title-main">Aluminium compounds</span>

Aluminium (or aluminum) combines characteristics of pre- and post-transition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a post-transition metal, with longer-than-expected interatomic distances. Furthermore, as Al3+ is a small and highly charged cation, it is strongly polarizing and aluminium compounds tend towards covalency; this behaviour is similar to that of beryllium (Be2+), an example of a diagonal relationship. However, unlike all other post-transition metals, the underlying core under aluminium's valence shell is that of the preceding noble gas, whereas for gallium and indium it is that of the preceding noble gas plus a filled d-subshell, and for thallium and nihonium it is that of the preceding noble gas plus filled d- and f-subshells. Hence, aluminium does not suffer the effects of incomplete shielding of valence electrons by inner electrons from the nucleus that its heavier congeners do. Aluminium's electropositive behavior, high affinity for oxygen, and highly negative standard electrode potential are all more similar to those of scandium, yttrium, lanthanum, and actinium, which have ds2 configurations of three valence electrons outside a noble gas core: aluminium is the most electropositive metal in its group. Aluminium also bears minor similarities to the metalloid boron in the same group; AlX3 compounds are valence isoelectronic to BX3 compounds (they have the same valence electronic structure), and both behave as Lewis acids and readily form adducts. Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class.

<span class="mw-page-title-main">Europium compounds</span> Chemical compounds with at least one europium atom

Europium compounds are compounds formed by the lanthanide metal europium (Eu). In these compounds, europium generally exhibits the +3 oxidation state, such as EuCl3, Eu(NO3)3 and Eu(CH3COO)3. Compounds with europium in the +2 oxidation state are also known. The +2 ion of europium is the most stable divalent ion of lanthanide metals in aqueous solution. Lipophilic europium complexes often feature acetylacetonate-like ligands, e.g., Eufod.

Rhenium compounds are compounds formed by the transition metal rhenium (Re). Rhenium can form in many oxidation states, and compounds are known for every oxidation state from -3 to +7 except -2, although the oxidation states +7, +6, +4, and +2 are the most common. Rhenium is most available commercially as salts of perrhenate, including sodium and ammonium perrhenates. These are white, water-soluble compounds. Tetrathioperrhenate anion [ReS4]− is possible.

Neptunium compounds are compounds containg the element neptunium (Np). Neptunium has five ionic oxidation states ranging from +3 to +7 when forming chemical compounds, which can be simultaneously observed in solutions. It is the heaviest actinide that can lose all its valence electrons in a stable compound. The most stable state in solution is +5, but the valence +4 is preferred in solid neptunium compounds. Neptunium metal is very reactive. Ions of neptunium are prone to hydrolysis and formation of coordination compounds.

Americium compounds are compounds containing the element americium (Am). These compounds can form in the +2, +3, and +4, although the +3 oxidation state is the most common. The +5, +6 and +7 oxidation states have also been reported.

References

  1. ↑ Greenwood and Earnshaw, p. 240
  2. 1 2 3 4 5 6 7 8 Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN   978-0-12-352651-9.
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    Ga
    -NMR Spectroscopic Study". J Solution Chem. 37 (10): 1411–1418. doi:10.1007/s10953-008-9314-y. S2CID   95723025.
  6. ↑ 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.
  7. 1 2 3 4 5 6 7 8 9 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.
  8. ↑ Madelung, Otfried (2004). Semiconductors: data handbook (3rd ed.). Birkhäuser. pp. 276–277. ISBN   978-3-540-40488-0.
  9. ↑ Krausbauer, L.; Nitsche, R.; Wild, P. (1965). "Mercury gallium sulfide, HgGa
    2    S
    4    , a new phosphor". Physica. 31 (1): 113–121. Bibcode:1965Phy....31..113K. doi:10.1016/0031-8914(65)90110-2.
  10. ↑ Michelle Davidson (2006). Inorganic Chemistry. Lotus Press. p. 90. ISBN   978-81-89093-39-6.
  11. ↑ Arora, Amit (2005). Text Book Of Inorganic Chemistry. Discovery Publishing House. pp. 389–399. ISBN   978-81-8356-013-9.
  12. ↑ 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.
  13. 1 2 3 Greenwoood and Earnshaw, pp. 262–5
  14. ↑ 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)
  15. ↑ Amemiya, Ryo (2005). "GaCl3 in Organic Synthesis". European Journal of Organic Chemistry. 2005 (24): 5145–5150. doi:10.1002/ejoc.200500512.
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