Ahexafluoride is a chemical compound with the general formula QX_nF₆, QX_nF₆^m−, or QX_nF₆^m+. Many molecules fit this formula. An important hexafluoride ishexafluorosilicic acid (H₂SiF₆), which is a byproduct of the mining ofphosphate rock. In thenuclear industry,uranium hexafluoride (UF₆) is an important intermediate in the purification of this element.

Cationic hexafluorides exist but are rarer than neutral or anionic hexafluorides. Examples are the hexafluorochlorine (ClF₆⁺), and hexafluorobromine (BrF₆⁺)cations.^[1]

Many elements form anionic hexafluorides. Members of commercial interest arehexafluorophosphate (PF₆⁻) andhexafluorosilicate (SiF₆²⁻).

Many transition metals form hexafluoride anions. Often the monoanions are generated by reduction of the neutral hexafluorides. For example,PtF₆⁻ arises by reduction of PtF₆ by O₂. Because of its highly basic nature and its resistance to oxidation, the fluoride ligand stabilizes some metals in otherwise rare high oxidation states, such ashexafluorocuprate(IV),CuF2−6 andhexafluoronickelate(IV),NiF2−6.

Seventeen elements are known to form binary hexafluorides.^[2] Nine of these elements aretransition metals, three areactinides, four arechalcogens, and one is anoble gas. Most hexafluorides aremolecular compounds with lowmelting andboiling points. Four hexafluorides (S, Se, Te, and W) are gases at room temperature (25 °C) and a pressure of 1atm, two are liquids (Re, Mo), and the others are volatile solids. Thegroup 6,chalcogen, andnoble gas hexafluorides are colourless, but the other hexafluorides have colours ranging from white, through yellow, orange, red, brown, and grey, to black.

The molecular geometry of binary hexafluorides is generallyoctahedral, although some derivatives are distorted from O_hsymmetry. For the main group hexafluorides, distortion is pronounced for the 14-electron noble gas derivatives. Distortions in gaseousXeF₆ are caused by its non-bondinglone pair, according toVSEPR theory. In the solid state, it adopts a complex structure involving tetramers and hexamers. According toquantum chemical calculations, ReF₆ and RuF₆ should have tetragonally distorted structures (where the two bonds along one axis are longer or shorter than the other four), but this has not been verified experimentally.^[3]

Polonium hexafluoride is known, but not well-studied. It could not be made from²¹⁰Po, but using the longer-lived isotope²⁰⁸Po and reacting it with fluorine found a volatile product that is almost certainly PoF₆.^[2] The quoted boiling point in the table below is a prediction.

The hexafluorides have a wide range of chemical reactivity.Sulfur hexafluoride is nearly inert and non-toxic due tosteric hindrance (the six fluorine atoms are arranged so tightly around the sulfur atom that it is extremely difficult to attack the bonds between the fluorine and sulfur atoms). It has several applications due to its stability, dielectric properties, and high density.Selenium hexafluoride is nearly as unreactive as SF₆, buttellurium hexafluoride is not very stable and can behydrolyzed by water within 1 day. Also, both selenium hexafluoride and tellurium hexafluoride are toxic, while sulfur hexafluoride is non-toxic. In contrast, metal hexafluorides are corrosive, readily hydrolyzed, and may react violently with water. Some of them can be used asfluorinating agents. The metal hexafluorides have a highelectron affinity, which makes them strong oxidizing agents.^[8]Platinum hexafluoride in particular is notable for its ability to oxidize thedioxygen molecule, O₂, to formdioxygenyl hexafluoroplatinate, and for being the first compound that was observed to react with xenon (seexenon hexafluoroplatinate).

Some metal hexafluorides find applications due to their volatility.Uranium hexafluoride is used in theuranium enrichment process to produce fuel fornuclear reactors.Fluoride volatility can also be exploited fornuclear fuel reprocessing.Tungsten hexafluoride is used in the production ofsemiconductors through the process ofchemical vapor deposition.^[9]

Radon hexafluoride (RnF
₆), the heavier homologue ofxenon hexafluoride, has been studied theoretically,^[10] but its synthesis has not yet been confirmed. Higher fluorides ofradon may have been observed in experiments where unknown radon-containing products distilled together withxenon hexafluoride, and perhaps in the production of radon trioxide: these may have been RnF₄, RnF₆, or both.^[11] It is likely that the difficulty in identifying higher fluorides of radon stems from radon being kinetically hindered from being oxidised beyond the divalent state. This is due to the strong ionicity ofRnF₂ and the high positive charge on Rn in RnF⁺. Spatial separation of RnF₂ molecules may be necessary to clearly identify higher fluorides of radon, of which RnF₄ is expected to be more stable than RnF₆ due tospin–orbit splitting of the 6p shell of radon (Rn^IV would have a closed-shell 6s²
6p²
_1/2 configuration).^[12] The ionicity of the Rn–F bond may also result in a strongly fluorine-bridged structure in the solid, so that radon fluorides may not be volatile.^[2] Continuing the trend, the heavieroganesson hexafluoride should be unbound.^[2]

Krypton hexafluoride (KrF
₆) has been predicted to be stable, but has not been synthesised due to the extreme difficulty of oxidisingkrypton beyond Kr(II).^[13] The synthesis ofamericium hexafluoride (AmF
₆) by thefluorination ofamericium(IV) fluoride (AmF
₄) was attempted in 1990,^[14] but was unsuccessful; there have also been possible thermochromatographic identifications of it andcurium hexafluoride (CmF₆), but it is debated if these are conclusive.^[2]Palladium hexafluoride (PdF
₆), the lighter homologue ofplatinum hexafluoride, has been calculated to be stable,^[15] but has not yet been produced; the possibility ofsilver (AgF₆) andgold hexafluorides (AuF₆) has also been discussed.^[2]Chromium hexafluoride (CrF
₆), the lighter homologue ofmolybdenum hexafluoride andtungsten hexafluoride, was reported, but has been shown to be a mistaken identification of the knownpentafluoride (CrF
₅).^[16]

• Galkin, N. P.; Tumanov, Yu. N. (1971)."Reactivity and Thermal Stability of Hexafluorides".Russian Chemical Reviews.40 (2):154–164.Bibcode:1971RuCRv..40..154G.doi:10.1070/RC1971v040n02ABEH001902.S2CID 250901336. Archived fromthe original on 2015-11-30. Retrieved2012-05-12.

• Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001).Inorganic chemistry. Academic Press.ISBN 978-0-12-352651-9. Retrieved3 March 2011.

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