Xenon hexafluoride
| File:Xenonhexafluorid.svg | |
| Identifiers | |
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3D model (JSmol)
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PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |
| XeF6 | |
| Molar mass | 245.28 g mol-1 |
| Density | 3.56 g cm-3 |
| Melting point | 49.25 °C |
| Boiling point | 75.6 °C |
| reacts with water | |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Xenon hexafluoride is a noble gas compound with the formula XeF6 and the highest of the three binary fluorides of xenon, the other two being XeF2 and XeF4. All are exergonic and stable at normal temperatures. XeF6 is the strongest fluorinating agent of the series. At room temperature, it is a colorless solid that readily sublimes into intensely yellow vapors.
Preparation
Xenon hexafluoride can be prepared by long-term heating of XeF2 at about 300°C and pressure 6 MPa.
With NiF
2 as catalyst, however, this reaction can proceed at 120°C even in xenon-fluorine molar ratios as low as 1:5.[1]
Structure
The structure of XeF6 required several years to establish in contrast to the cases of XeF
2 and XeF
4. In the gas phase the compound is monomeric. VSEPR theory predicts that due to the presence of six fluoride ligands and one lone pair of electrons the structure lacks perfect octahedral symmetry, and indeed electron diffraction combined with high-level calculations indicate that the compound's point group is C3v. The calculated energy for the point group Oh is only insignificantly higher, indicating that the minimum on the energy surface is very shallow. Konrad Seppelt, an authority on noble gas and fluorine chemistry, says, "the structure is best described in terms of a mobile electron pair that moves over the faces and edges of the octahedron and thus distorts it in a dynamic manner.".[2]
129Xe and 19F NMR spectroscopy indicates that in solution the compound assumes a tetrameric structure: four equivalent xenon atoms are arranged in a tetrahedron surrounded by a fluctuating array of 24 fluorine atoms that interchange positions in a "cogwheel mechanism".
XeF
6 crystallizes in 6 possible modifications,[3] including one that contains XeF+
5 ions with bridging F−
ions.[4]
Reactions
Hydrolysis
Xenon hexafluoride hydrolyzes stepwise, ultimately affording xenon trioxide:[5]
- XeF6 + H2O → XeOF4 + 2 HF
- XeOF4 + H2O → XeO2F2 + 2 HF
- XeO2F2 + H2O → XeO3 + 2 HF
XeF6 serves as a Lewis acid, binding one and two fluoride anions:
- XeF6 + F− → XeF−
7 - XeF−
7 + F− → XeF2−
8
Octafluoroxenates
Salts of the octafluoroxenate(VI) anion (XeF2−
8) are very stable, decomposing only above 400 °C.[6][7][8] This anion has been shown to have square antiprismatic geometry, based on single-crystal X-ray counter analysis of its nitrosonium salt, (NO)
2XeF
8.[9] The sodium and potassium salts are formed directly from sodium fluoride and potassium fluoride:[8]
- 2 NaF + XeF
6 → Na
2XeF
8 - 2 KF + XeF
6 → K
2XeF
8
These are thermally less stable than the caesium and rubidium salts, which are synthesized by first forming the heptafluoroxenate salts:
- CsF + XeF
6 → CsXeF
7 - RbF + XeF
6 → RbXeF
7
which are then pyrolysed at 50°C and 20°C, respectively, to form the yellow[10] octafluoroxenate salts:[8][7][6]
- 2 CsXeF
7 → Cs
2XeF
8 + XeF
6 - 2 RbXeF
7 → Rb
2XeF
8 + XeF
6
These salts are hydrolysed by water, yielding various products containing xenon and oxygen.[8]
The two other binary fluorides of xenon do not form such stable adducts with fluoride.
With fluoride acceptors
XeF
6 reacts with strong fluoride acceptors such as RuF
5[4] and BrF
3·AuF
3[11] to form the XeF+
5 cation:
- XeF
6 + RuF
5 → XeF+
5RuF−
6
- XeF
6 + BrF
3·AuF
3 → XeF+
5AuF−
4 + BrF
3
References
- ^ Melita Tramšek; Boris Žemva (2006). "Synthesis, Properties and Chemistry of Xenon(II) Fluoride" (PDF). Acta Chim. Slov. 53 (2): 105–116.
{{cite journal}}: Unknown parameter|day=ignored (help); Unknown parameter|month=ignored (help) - ^ Seppelt, Konrad (1979). "Recent Developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research. 12 (6): 211–216. doi:10.1021/ar50138a004.
{{cite journal}}: Unknown parameter|month=ignored (help) - ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1016/j.jfluchem.2006.04.014, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1016/j.jfluchem.2006.04.014instead. - ^ a b James E. House (2008). Inorganic Chemistry. Academic Press. p. 569. ISBN 0123567866.
- ^ Appelman, E. H. (1964). "Hydrolysis of Xenon Hexafluoride and the Aqueous Solution Chemistry of Xenon". Journal of the American Chemical Society. 86 (11): 2141–2148. doi:10.1021/ja01065a009.
{{cite journal}}: Unknown parameter|coauthors=ignored (|author=suggested) (help); Unknown parameter|month=ignored (help) - ^ a b Holleman, A. F.; Wiberg,, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
{{cite book}}: CS1 maint: extra punctuation (link) - ^ a b Riedel, Erwin; Janiak, Christoph (2007). Anorganische Chemie (7th ed.). Walter de Gruyter. p. 393. ISBN 3110189038.
- ^ a b c d Chandra, Sulekh (2004). Comprehensive Inorganic Chemistry. New Age International. p. 308. ISBN 8122415121.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1126/science.173.4003.1238, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1126/science.173.4003.1238instead. - ^ "Xenon". Encyclopaedia Britannica. Encyclopaedia Britannica Inc. 1995.
- ^ Cotton (2007). Advanced Inorganic Chemistry (6th ed.). Wiley-India. p. 591. ISBN 8126513381.