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Plutonium compounds
Chemical compounds containing the element plutonium From Wikipedia, the free encyclopedia
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Plutonium compounds are compounds containing the element plutonium (Pu). At room temperature, pure plutonium is silvery in color but gains a tarnish when oxidized.[1] The element displays four common ionic oxidation states in aqueous solution and one rare one:[2]
- Pu(III), as Pu3+ (blue lavender)
- Pu(IV), as Pu4+ (yellow brown)
- Pu(V), as PuO+
2 (light pink)[note 1] - Pu(VI), as PuO2+
2 (pink orange) - Pu(VII), as PuO3−
5 (green)-the heptavalent ion is rare.
The color shown by plutonium solutions depends on both the oxidation state and the nature of the acid anion.[4] It is the acid anion that influences the degree of complexing—how atoms connect to a central atom—of the plutonium species. Additionally, the formal +2 oxidation state of plutonium is known in the complex [K(2.2.2-cryptand)] [PuIICp″3], Cp″ = C5H3(SiMe3)2.[5]
A +8 oxidation state is possible as well in the volatile tetroxide PuO
4.[6] Though it readily decomposes via a reduction mechanism similar to FeO
4, PuO
4 can be stabilized in alkaline solutions and chloroform.[7][6]
Metallic plutonium is produced by reacting plutonium tetrafluoride with barium, calcium or lithium at 1200 °C.[8] Metallic plutonium is attacked by acids, oxygen, and steam but not by alkalis and dissolves easily in concentrated hydrochloric, hydroiodic and perchloric acids.[9] Molten metal must be kept in a vacuum or an inert atmosphere to avoid reaction with air.[9] At 135 °C the metal will ignite in air and will explode if placed in carbon tetrachloride.[10]
Plutonium is a reactive metal. In moist air or moist argon, the metal oxidizes rapidly, producing a mixture of oxides and hydrides.[11] If the metal is exposed long enough to a limited amount of water vapor, a powdery surface coating of PuO2 is formed.[11] Also formed is plutonium hydride but an excess of water vapor forms only PuO2.[9]
Plutonium shows enormous, and reversible, reaction rates with pure hydrogen, forming plutonium hydride.[12] It also reacts readily with oxygen, forming PuO and PuO2 as well as intermediate oxides; plutonium oxide fills 40% more volume than plutonium metal. The metal reacts with the halogens, see below. The following oxyhalides are observed: PuOF, PuOCl, PuOBr, and PuOI. It will react with carbon to form PuC, nitrogen to form PuN, and silicon to form PuSi2.[2][10]
The organometallic chemistry of plutonium complexes is typical for organoactinide species; a characteristic example of an organoplutonium compound is plutonocene.[13][14] Computational chemistry methods indicate an enhanced covalent character in the plutonium-ligand bonding.[12][14]
Powders of plutonium, its hydrides and certain oxides like Pu2O3 are pyrophoric, meaning they can ignite spontaneously at ambient temperature and are therefore handled in an inert, dry atmosphere of nitrogen or argon. Bulk plutonium ignites only when heated above 400 °C. Pu2O3 spontaneously heats up and transforms into PuO2, which is stable in dry air, but reacts with water vapor when heated.[15]
Crucibles used to contain plutonium need to be able to withstand its strongly reducing properties. Refractory metals such as tantalum and tungsten along with the more stable oxides, borides, carbides, nitrides and silicides can tolerate this. Melting in an electric arc furnace can be used to produce small ingots of the metal without the need for a crucible.[9]
Cerium is used as a chemical simulant of plutonium for development of containment, extraction, and other technologies.[16]
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Halides
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Plutonium reacts readily with halogens, forming halides of composition PuX3 (X=F, Cl, Br, I) or PuX4 (X=F).[2][10] The higher plutonium fluorides PuF5 and PuF6 and the higher plutonium chloride PuCl4 are also known, though PuF5 and PuCl4 are only known in the gas phase. Several of these compounds form hydrates, with compositions PuF3·0.40H2O, PuF3·0.75H2O, PuF4·2.5H2O, PuCl3·3H2O, PuCl3·6H2O, and PuBr3·6H2O.[17]
Fluorides
The compounds PuF3, PuF4, and PuF6 are well-characterized. The plutonium(III) fluoride hydrate PuF3·0.75H2O can be prepared from adding hydrofluoric acid to a plutonium(III) solution in nitric acid, from which it precipitates.[18] The hydrate can then be converted to the anhydrous form by heating in hydrofluoric acid stream.[17] However, this method for PuF3 production is not favored as it can lead to problems with filtration after the product is precipitated. Alternatively, it can be produced by reacting plutonium compounds, such as PuO2 or Pu2(C2O4)3 with gaseous fluorinating agents, such as HF, F2, NF3, ClF3, or CCl2F2. This method also produces other plutonium fluorides, but these can be reduced to the trifluoride by action of reducing agents such as hydrogen gas.[19] Due to the toxicity of these fluorinating agents, other methods have also been proposed, such as the reaction of plutonium dioxide with ammonium bifluoride.[20] PuF3 is used to prepare plutonium metal, forming the metal on action by calcium and iodine.[21]
As with PuF3, a plutonium(IV) fluoride (PuF4) hydrate, PuF4·2.5H2O, can be precipitated by hydrofluoric acid in aqueous solution. Surprisingly, however, decomposition of this compound in vacuum yields PuF3 instead of PuF4. Like with PuF3, though, it can be converted to PuF4 by heating in a stream of HF. It is most commonly produced by reacting plutonium dioxide with hydrofluoric acid at high temperatures with oxygen gas. Several other materials can be used in place of plutonium dioxide, like plutonium(IV) nitrate or plutonium peroxide.[17] Anhydrous PuF4 adopts a polymeric structure, isostructural with uranium tetrafluoride. In addition to the anhydrous form, two distinct hydrates of PuF4 are known: the aforementioned PuF4·2.5H2O, and PuF4·xH2O (0.5 ≤ x ≤ 2.0). The hydrates are also isostructural to the corresponding uranium compounds. PuF4·xH2O (0.5 ≤ x ≤ 2.0) adopts a cubic structure, while PuF4·2.5H2O adopts an orthorhombic structure.[22] PuF4 has been found to slowly age over time; a 30 year old sample of PuF4 was found to contain 88% anhydrous PuF4, 8% PuF4·1.6H2O, and 4% PuO2.[23] It is an important intermediate in plutonium metal production.[22]
Upon reaction with fluorine gas at high temperatures, PuF4 is oxidized to plutonium hexafluoride, PuF6. It is also readily produced upon reaction with powerful fluorinating agents, like krypton difluoride and dioxygen difluoride.[17] Unlike the related compound uranium hexafluoride, PuF6 is unstable and highly reactive.[24] PuF6 solid is prone to radiolysis from alpha particles produced in the radioactive decay of plutonium, forming PuF4 and fluorine gas; however, gas phase reaches a stable equilibrium.[25][26] It undergoes slow hydrolysis to plutonyl fluoride PuO2F2. In hydrofluoric acid solution, PuO2F2 forms a hydrate, PuO2F2·H2O, and a solid incorporating HF, PuO2F2·HF·4H2O.[17]
Plutonium also forms a few fluorides intermediate between PuF4 and PuF6, namely plutonium(V) fluoride PuF5, as well as the mixed-valent Pu4F17. Pu4F17 is a brick-red solid formed during the production of PuF6. PuF5 is a weakly-characterized compound only known in the gas phase.[17]
Chlorides
The only stable binary plutonium chloride in the solid phase is plutonium(III) chloride, PuCl3. It can be produced in several ways. For medium-scale reactions (between 1 and 10 grams), the best method of PuCl3 production is the action of hydrochloric acid on plutonium(III) oxalate.[17] Plutonium(III) oxalate can also be converted to plutonium(III) chloride by action of hexachloropropene.[27] Analytically pure samples of PuCl3 can be created by reacting plutonium(IV) oxide prepared by calcination of plutonium(IV) oxalate with phosgene or carbon tetrachloride at elevated temperatures (>500 °C). PuCl3 has been determined to adopt the uranium(III) chloride-type structure. In addition to anhydrous PuCl3, multiple hydrates are also known, with compositions PuCl3·3H2O and PuCl3·6H2O.[17][28] PuCl3·6H2O, which can be synthesized by evaporating HCl solutions,[17] adopts the GdCl3·6H2O-type structure,[note 2] consisting of PuCl2(H2O)+6 cations which are linked by chloride anions.[30] It melts in its own waters of crystallization at 94 °C.[31] PuCl3·3H2O is isostructural with NdCl3·3H2O.[28]
While it does not exist as a solid, dissociating into plutonium(III) chloride and chlorine gas, binary plutonium tetrachloride is known in the gas phase,[29] and it forms several stable adducts, such as with dimethoxyethane[32] or diphenylsulfoxide.[33] Plutonium tetrachloride gas is formed from the reaction of plutonium(III) chloride with chlorine gas at high temperatures, enhancing the volatility of the product.[17]: 1094 The dimethoxyethane adduct, PuCl4(DME)2 (DME=dimethoxyethane), can be prepared by evaporating a plutonium(IV) solution in HCl, adding the residue to a solution of DME, and then adding (CH3)3SiCl.[32] This adduct can be used to prepare other compounds. For example, when it is dissolved in tetrahydrofuran, it forms the mixed-valence complex [PuIIICl2(thf)5]+[PuIVCl5(thf)]− (THF=tetrahydrofuran).[34]
Despite solid PuCl4 being unknown, several stable solid plutonium(IV) chlorides derived from PuCl4, such as the hexachloroplutonates, are well-characterized.[17] The hexachloroplutonate ion (PuCl2−6) is the dominant plutonium species in concentrated hydrochloric acid, and several compounds are known containing this ion. Dicaesium hexachloroplutonate (Cs2PuCl6) can be precipitated from concentrated HCl solution by the addition of caesium chloride.[35] Cs2PuCl6 can be used to prepare other plutonium compounds, such as the cyclopentadienide complex (η5-C5H5)3PuCl. Other hexachloroplutonates are also known, such as K2PuCl6, Rb2PuCl6, (N(CH3)4)2PuCl6, and (N(C2H5)4)2PuCl6.[17]
Bromides
The only stable compound in the plutonium-bromine system is PuBr3. It is formed via several methods, the two best ones being direct synthesis from plutonium and bromine, and reaction of plutonium hydride with hydrobromic acid. It is also prepared by the hydrobromination of various plutonium compounds, like Pu2(C2O4)3, PuCl3, and PuBr3·6H2O, vacuum decompositon of PuBr3·6H2O, or reaction of Pu(OH)4 with bromine. Its structure, termed the PuBr3 structure, consists of PuBr8 polyhedra, where plutonium has a coordination geometry of bicapped trigonal prismatic, which link together to form infinite chains.[17] A hydrate of plutonium(III) bromide, PuBr3·6H2O, is also known, and is formed from plutonium bromide solutions containing water. Its structure consists of the complex [PuBr2(H2O)6]+, and is isostructural with GdCl3·6H2O.[36][37] It also forms a complex with tetrahydrofuran, PuBr3(THF)4 (THF=tetrahydrofuran), which can be formed by reacting plutonium and bromine in THF solution, though the hexahydrate is formed if water is present in the THF solution.[36] The THF complex is used in the synthesis of other plutonium compounds.[38]
A few plutonium bromide complexes are known containing plutonium(IV). Plutonium tetrabromide forms stable complexes with hexamethylphosphoramide (PuBr4·2[(CH3)2N]3PO) triphenylphosphine oxide (PuBr4·2(C6H5)3PO),[39] and tricyclohexylphosphine oxide (PuBr4·2(C6H11)3PO),[40] as well as an unstable complex with acetonitrile (PuBr4·4CH3CN), which decomposes to a plutonium(III) compound.[39] The plutonium(IV) complex ion PuBr2−6 is also known, found in the tetraethylammonium salt [(C2H5)4N]2PuBr6.[41][17]
Iodides
Plutonium(III) iodide can be prepared by reacting plutonium metal with either hydroiodic acid or mercury(II) iodide. It is extremely moisture-sensitive, and when even traces of water are present in its formation conditions, plutonium oxyiodide, PuOI, is formed instead. It adopts the PuBr3-type structure. Several plutonium triiodide complexes are known, such as PuI3(THF)4 (THF=tetrahydrofuran), PuI3(DMSO)4 (DMSO=dimethylsulfoxide), and PuI3(py)4 (py=pyridine). These complexes can be formed by reacting plutonium metal with iodine in tetrahydrofuran, dimethylsulfoxide, or pyridine solutions, respectively. PuI3 is a poorly characterized solid.[17]
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See also
Notes
- The PuO+
2 ion is unstable in solution and will disproportionate into Pu4+ and PuO2+
2; the Pu4+ will then oxidize the remaining PuO+
2 to PuO2+
2, being reduced in turn to Pu3+. Thus, aqueous solutions of PuO+
2 tend over time towards a mixture of Pu3+ and PuO2+
2. UO+
2 is unstable for the same reason.[3]
References
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