Editing Gallium (section)

==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 GaCl<sub>2</sub> contains both gallium(I) and gallium(III) and can be formulated as Ga<sup>I</sup>Ga<sup>III</sup>Cl<sub>4</sub>; in contrast, the monochloride is unstable above 0&nbsp;°C, [[disproportionating]] into elemental gallium and gallium(III) chloride. Compounds containing Ga–Ga bonds are true gallium(II) compounds, such as [[gallium(II) sulfide|GaS]] (which can be formulated as Ga<sub>2</sub><sup>4+</sup>(S<sup>2−</sup>)<sub>2</sub>) and the [[dioxan]] complex Ga<sub>2</sub>Cl<sub>4</sub>(C<sub>4</sub>H<sub>8</sub>O<sub>2</sub>)<sub>2</sub>.<ref name="GreenwoodEarnshaw2nd"/>{{rp|240}}

===Aqueous chemistry===
Strong acids dissolve gallium, forming gallium(III) salts such as [[gallium(III) nitrate|{{chem|Ga(NO|3|)|3}}]] (gallium nitrate). [[Aqueous]] solutions of gallium(III) salts contain the hydrated gallium ion, {{chem|[Ga(H|2|O)|6|]|3+}}.<ref name="wiberg_holleman">{{cite book |title= Inorganic chemistry |author1= Wiberg, Egon |author2= Wiberg, Nils |author3= Holleman, Arnold Frederick |publisher= Academic Press |date= 2001 |isbn= 978-0-12-352651-9}}</ref>{{rp|1033}} [[Gallium(III) hydroxide]], {{chem|Ga(OH)|3}}, may be precipitated from gallium(III) solutions by adding [[ammonia]]. Dehydrating {{chem|Ga(OH)|3}} at 100&nbsp;°C produces gallium oxide hydroxide, GaO(OH).<ref name="downs">{{cite book |title= Chemistry of aluminium, gallium, indium, and thallium |author= Downs, Anthony John |publisher= Springer |date= 1993 |isbn= 978-0-7514-0103-5}}</ref>{{rp|140–141}}

Alkaline [[hydroxide]] solutions dissolve gallium, forming ''gallate'' salts (not to be confused with identically named [[gallic acid]] salts) containing the {{chem|Ga(OH)|4|-}} anion.<ref name="eagleson" /><ref name="wiberg_holleman" />{{rp|1033}}<ref name="sipos" /> Gallium hydroxide, which is [[amphoteric]], also dissolves in alkali to form gallate salts.<ref name="downs" />{{rp|141}} Although earlier work suggested {{chem|Ga(OH)|6|3-}} as another possible gallate anion,<ref>{{cite book |title= Electrochemistry—Volume 3: Specialist periodical report |author= Hampson, N. A. |editor= Harold Reginald Thirsk |publisher= Royal Society of Chemistry |location= Great Britain |date= 1971 |isbn= 978-0-85186-027-5 |page= 71 |url= https://books.google.com/books?id=vN0Y7KMGqNcC}}</ref> it was not found in later work.<ref name="sipos">{{cite journal |last1= Sipos |first1= P. L. |last2= Megyes |first2= T. N. |last3= Berkesi |year= 2008 |first3= O. |title= The Structure of Gallium in Strongly Alkaline, Highly Concentrated Gallate Solutions—a Raman and {{SimpleNuclide|Ga|71}}-NMR Spectroscopic Study |journal= J Solution Chem |volume= 37 |issue= 10 |pages= 1411–1418 |doi= 10.1007/s10953-008-9314-y |s2cid= 95723025 }}</ref>

===Oxides and chalcogenides===
Gallium reacts with the [[chalcogen]]s only at relatively high temperatures. At room temperature, gallium metal is not reactive with air and water because it forms a [[passivation (chemistry)|passive]], protective [[oxide]] layer. At higher temperatures, however, it reacts with atmospheric [[oxygen]] to form [[gallium(III) oxide]], {{chem|Ga|2|O|3}}.<ref name="eagleson">{{cite book
|title= Concise encyclopedia chemistry
|url= https://archive.org/details/conciseencyclope00eagl
|url-access= registration
|editor= Eagleson, Mary
|publisher= Walter de Gruyter
|date= 1994
|isbn= 978-3-11-011451-5
|page= [https://archive.org/details/conciseencyclope00eagl/page/438 438]
}}</ref> Reducing {{chem|Ga|2|O|3}} with elemental gallium in vacuum at 500&nbsp;°C to 700&nbsp;°C yields the dark brown [[gallium(I) oxide]], {{chem|Ga|2|O}}.<ref name="downs" />{{rp|285}} {{chem|Ga|2|O}} is a very strong [[reducing agent]], capable of reducing [[sulfuric acid|{{chem|H|2|SO|4}}]] to [[hydrogen sulfide|{{chem|H|2|S}}]].<ref name="downs" />{{rp|207}} It disproportionates at 800&nbsp;°C back to gallium and {{chem|Ga|2|O|3}}.<ref name="emeleus_sharpe">{{cite book |title= Advances in inorganic chemistry and radiochemistry |volume= 5 |author= Greenwood, N. N. |editor= Harry Julius Emeléus |editor-link= Harry Julius Emeléus |editor2= Alan G. Sharpe |publisher= Academic Press |date= 1962
|isbn= 978-0-12-023605-3 |pages= 94–95}}</ref>

[[Gallium(III) sulfide]], {{chem|Ga|2|S|3}}, has 3 possible crystal modifications.<ref name="emeleus_sharpe" />{{rp|104}} It can be made by the reaction of gallium with [[hydrogen sulfide]] ({{chem|H|2|S}}) at 950&nbsp;°C.<ref name="downs" />{{rp|162}} Alternatively, {{chem|Ga(OH)|3}} can be used at 747&nbsp;°C:<ref>{{cite book |title= Semiconductors: data handbook |author= Madelung, Otfried |edition= 3rd |publisher= Birkhäuser |date= 2004 |isbn= 978-3-540-40488-0 |pages= 276–277}}</ref>
:2 {{chem|Ga(OH)|3}} + 3 {{chem|H|2|S}} → {{chem|Ga|2|S|3}} + 6 {{chem|H|2|O}}

Reacting a mixture of alkali metal carbonates and {{chem|Ga|2|O|3}} with {{chem|H|2|S}} leads to the formation of ''thiogallates'' containing the {{chem|[Ga|2|S|4|]|2-}} anion. Strong acids decompose these salts, releasing {{chem|H|2|S}} in the process.<ref name="emeleus_sharpe" />{{rp|104–105}} The mercury salt, {{chem|HgGa|2|S|4}}, can be used as a [[phosphor]].<ref>{{cite journal |author= Krausbauer, L. |author2= Nitsche, R. |author3= Wild, P. |year= 1965 |title= Mercury gallium sulfide, {{chem|HgGa|2|S|4}}, a new phosphor |journal= Physica |volume= 31 |issue= 1 |pages= 113–121 |doi= 10.1016/0031-8914(65)90110-2|bibcode= 1965Phy....31..113K}}</ref>

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&nbsp;°C under a stream of nitrogen.<ref name="emeleus_sharpe" />{{rp|94}}

The other binary chalcogenides, {{chem|Ga|2|Se|3}} and {{chem|Ga|2|Te|3}}, have the [[zincblende (crystal structure)|zincblende]] structure. They are all semiconductors but are easily [[hydrolysis|hydrolysed]] and have limited utility.<ref name="emeleus_sharpe" />{{rp|104}}

===Nitrides and pnictides===
{{Multiple image
|width= 160
|image1= Crystal-GaN.jpg
|image2= Gallium Arsenide (GaAs) 2" wafer.jpg
|footer= Gallium nitride (left) and gallium arsenide (right) wafers
}}
Gallium reacts with ammonia at 1050&nbsp;°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 [[zinc sulfide|ZnS]], and have important [[semiconductor|semiconducting]] properties.<ref name="wiberg_holleman" />{{rp|1034}} GaP, GaAs, and GaSb can be synthesized by the direct reaction of gallium with elemental phosphorus, arsenic, or antimony.<ref name="emeleus_sharpe" />{{rp|99}} They exhibit higher electrical conductivity than GaN.<ref name="emeleus_sharpe" />{{rp|101}} GaP can also be synthesized by reacting {{chem|Ga|2|O}} with phosphorus at low temperatures.<ref>{{cite book |title= Inorganic Chemistry |author= Michelle Davidson |publisher= Lotus Press |date= 2006 |isbn= 978-81-89093-39-6 |page= 90}}</ref>

Gallium forms ternary [[nitride]]s; for example:<ref name="emeleus_sharpe" />{{rp|99}}
:{{chem|Li|3|Ga}} + {{chem|N|2}} → {{chem|Li|3|GaN|2}}

Similar compounds with phosphorus and arsenic are possible: {{chem|Li|3|GaP|2}} and {{chem|Li|3|GaAs|2}}. These compounds are easily hydrolyzed by dilute [[acid]]s and water.<ref name="emeleus_sharpe" />{{rp|101}}

===Halides===
{{See also|Gallium halides}}
Gallium(III) oxide reacts with [[Halogenation|fluorinating agents]] such as [[hydrogen fluoride|HF]] or [[fluorine|{{chem|F|2}}]] to form [[gallium(III) fluoride]], {{chem|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, {{chem|GaF|3|·3H|2|O}}. Attempting to dehydrate this adduct forms {{chem|GaF|2|OH·''n''H|2|O}}. The adduct reacts with ammonia to form {{chem|GaF|3|·3NH|3}}, which can then be heated to form anhydrous {{chem|GaF|3}}.<ref name="downs" />{{rp|128–129}}

[[Gallium trichloride]] is formed by the reaction of gallium metal with [[chlorine]] gas.<ref name="eagleson" /> Unlike the trifluoride, gallium(III) chloride exists as dimeric molecules, {{chem|Ga|2|Cl|6}}, with a melting point of 78&nbsp;°C. Equivalent compounds are formed with bromine and iodine, [[gallium(III) bromide|{{chem|Ga|2|Br|6}}]] and [[gallium(III) iodide|{{chem|Ga|2|I|6}}]].<ref name="downs" />{{rp|133}}

Like the other group 13 trihalides, gallium(III) halides are [[Lewis acid]]s, reacting as halide acceptors with alkali metal halides to form salts containing {{chem|GaX|4|-}} anions, where X is a halogen. They also react with [[haloalkane|alkyl halides]] to form [[carbocation]]s and {{chem|GaX|4|-}}.<ref name="downs" />{{rp|136–137}}

When heated to a high temperature, gallium(III) halides react with elemental gallium to form the respective gallium(I) halides. For example, {{chem|GaCl|3}} reacts with Ga to form {{chem|GaCl}}:
:2 Ga + {{chem|GaCl|3}} {{eqm}} 3 GaCl (g)

At lower temperatures, the equilibrium shifts toward the left and GaCl disproportionates back to elemental gallium and {{chem|GaCl|3}}. GaCl can also be produced by reacting Ga with HCl at 950&nbsp;°C; the product can be condensed as a red solid.<ref name="wiberg_holleman" />{{rp|1036}}

Gallium(I) compounds can be stabilized by forming adducts with Lewis acids. For example:
:GaCl + {{chem|AlCl|3}} → {{chem|Ga|+|[AlCl|4|]|-}}

The so-called "gallium(II) halides", {{chem|GaX|2}}, are actually adducts of gallium(I) halides with the respective gallium(III) halides, having the structure {{chem|Ga|+|[GaX|4|]|-}}. For example:<ref name="eagleson" /><ref name="wiberg_holleman" />{{rp|1036}}<ref name="arora">{{cite book
|title= Text Book Of Inorganic Chemistry
|author= Arora, Amit
|publisher= Discovery Publishing House
|date= 2005
|isbn= 978-81-8356-013-9
|pages= 389–399
}}</ref>
:GaCl + {{chem|GaCl|3}} → {{chem|Ga|+|[GaCl|4|]|-}}

===Hydrides===
Like [[aluminium]], gallium also forms a [[hydride]], {{chem|GaH|3}}, known as ''[[gallane]]'', which may be produced by reacting lithium gallanate ({{chem|LiGaH|4}}) with [[gallium(III) chloride]] at −30&nbsp;°C:<ref name="wiberg_holleman" />{{rp|1031}}
:3 {{chem|LiGaH|4}} + {{chem|GaCl|3}} → 3 LiCl + 4 {{chem|GaH|3}}

In the presence of [[dimethyl ether]] as solvent, {{chem|GaH|3}} polymerizes to {{chem|(GaH|3|)|''n''}}. If no solvent is used, the dimer {{chem|Ga|2|H|6}} (''[[digallane]]'') is formed as a gas. Its structure is similar to [[diborane]], having two hydrogen atoms bridging the two gallium centers,<ref name="wiberg_holleman" />{{rp|1031}} unlike α-[[aluminium hydride|{{chem|AlH|3}}]] in which aluminium has a coordination number of 6.<ref name="wiberg_holleman" />{{rp|1008}}

Gallane is unstable above −10&nbsp;°C, decomposing to elemental gallium and [[hydrogen]].<ref name="sykes">{{cite book
|title= Advances in Inorganic Chemistry
|volume= 41
|author1= Downs, Anthony J.
|author2= Pulham, Colin R.
|editor= Sykes, A. G.
|publisher= Academic Press
|date= 1994
|isbn= 978-0-12-023641-1
|pages= 198–199
}}</ref>

===Organogallium compounds===
{{Main|Organogallium chemistry}}
Organogallium compounds are of similar reactivity to [[Organoindium chemistry|organoindium]] compounds, less reactive than [[Organoaluminium chemistry|organoaluminium]] compounds, but more reactive than [[Thallium#Organothallium compounds|organothallium]] compounds.<ref name="GreenwoodEarnshaw2nd"/>{{rp|262-5}} Alkylgalliums are monomeric. [[Lewis acid]]ity 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.<ref>{{cite journal |last1=Uhl |first1=Werner |last2=Reza Halvagar |first2=Mohammad |last3=Claesener |first3=Michael |title=Reducing Ga-H and Ga-C Bonds in Close Proximity to Oxidizing Peroxo Groups: Conflicting Properties in Single Molecules |journal=Chemistry – A European Journal |date=26 October 2009 |volume=15 |issue=42 |pages=11298–11306 |doi=10.1002/chem.200900746 |pmid=19780106 }}</ref> 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.<ref name="GreenwoodEarnshaw2nd"/>{{rp|262-5}}

Gallium trichloride is a common starting reagent for the formation of organogallium compounds, such as in [[carbometalation|carbogallation]] reactions.<ref>{{cite journal |doi= 10.1002/ejoc.200500512 |volume=2005 |issue=24 |title=GaCl<sub>3</sub> in Organic Synthesis |year=2005 |journal=European Journal of Organic Chemistry |pages=5145–5150 |last1= Amemiya |first1= Ryo}}</ref> Gallium trichloride reacts with [[lithium]] cyclopentadienide in [[diethyl ether]] to form the trigonal planar gallium cyclopentadienyl complex GaCp<sub>3</sub>. Gallium(I) forms complexes with [[arene]] [[ligand]]s such as [[hexamethylbenzene]]. Because this ligand is quite bulky, the structure of the [Ga(η<sup>6</sup>-C<sub>6</sub>Me<sub>6</sub>)]<sup>+</sup> is that of a [[half sandwich compound|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(η<sup>6</sup>-C<sub>6</sub>H<sub>6</sub>)<sub>2</sub>] [GaCl<sub>4</sub>]·3C<sub>6</sub>H<sub>6</sub>.<ref name="GreenwoodEarnshaw2nd"/>{{rp|262-5}}