Hexafluoride
A hexafluoride is a chemical compound with the general formula QXnF6, QXnF6m−, or QXnF6m+. Many molecules fit this formula. An important hexafluoride is hexafluorosilicic acid (H2SiF6), which is a byproduct of the mining of phosphate rock. In the nuclear industry, uranium hexafluoride (UF6) is an important intermediate in the purification of this element.
Contents
1 Hexafluoride cations
2 Hexafluoride anions
3 Binary hexafluorides
3.1 Binary hexafluorides of the chalcogens
3.2 Binary hexafluorides of the noble gases
3.3 Binary hexafluorides of the transition metals
3.4 Binary hexafluorides of the actinides
3.5 Chemical properties of binary hexafluorides
3.6 Applications of binary hexafluorides
3.7 Predicted binary hexafluorides
4 Literature
5 References
Hexafluoride cations
Cationic hexafluorides exist but are rarer than neutral or anionic hexafluorides. Examples are the hexafluorochlorine (ClF6+), and hexafluorobromine (BrF6+) cations.[1]
Hexafluoride anions
Structure of the hexafluorophosphate anion, PF6−.
Many elements form anionic hexafluorides. Members of commercial interest are hexafluorophosphate (PF6−) and hexafluorosilicate (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 as hexafluorocuprate2− and hexafluoronickelate2−.
Binary hexafluorides
Hexafluoride-forming elements
Octahedral structure of SF6
Seventeen elements are known to form binary hexafluorides.[citation needed] Nine of these elements are transition metals, three are actinides, four are chalcogens, and one is a noble gas. Most hexafluorides are molecular compounds with low melting and boiling points. Four hexafluorides (S, Se, Te, and W) are gases at room temperature (25 °C) and a pressure of 1 atm, two are liquids (Re, Mo), and the others are volatile solids. The group 6, chalcogen, and noble 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 generally octahedral, although some derivatives are distorted from Oh symmetry. For the main group hexafluorides, distortion is pronounced for the 14-electron noble gas derivatives. Distortions in gaseous XeF6 are caused by its non-bonding lone pair, according to VSEPR theory. In the solid state, it adopts a complex structure involving tetramers and hexamers. According to quantum 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.[2]
The status of polonium hexafluoride is unclear: some experimental results suggest that it may have been synthesized, but it was not well characterized. The quoted boiling point in the table below is thus a prediction. Despite this situation, some sources describe it without comment as a known compound.
Binary hexafluorides of the chalcogens
| Compound | Formula | m.p (°C) | b.p. (°C) | subl.p. (°C) | MW | solid ρ (g cm−3) (at m.p.)[3] | Bond distance (pm) | Colour |
|---|---|---|---|---|---|---|---|---|
| Sulfur hexafluoride | SF 6 | −50.8 | −63.8 | 146.06 | 2.51 (−50 °C) | 156.4 | colourless | |
| Selenium hexafluoride | SeF 6 | −34.6 | −46.6 | 192.95 | 3.27 | 167–170 | colourless | |
Tellurium hexafluoride[4] | TeF 6 | −38.9 | −37.6 | 241.59 | 3.76 | 184 | colourless | |
Polonium hexafluoride[5][6] | PoF 6 | ≈ −40? | 3.76 | 322.99 | colourless[6] |
Binary hexafluorides of the noble gases
| Compound | Formula | m.p (°C) | b.p. (°C) | subl.p. (°C) | MW | solid ρ (g cm−3) | Bond (pm) | Colour |
|---|---|---|---|---|---|---|---|---|
| Xenon hexafluoride | XeF 6 | 49.5 | 75.6 | 245.28 | 3.56 | colourless |
Binary hexafluorides of the transition metals
| Compound | Formula | m.p (°C) | b.p. (°C) | subl.p. (°C) | MW | solid ρ (g cm−3) | Bond (pm) | Colour |
|---|---|---|---|---|---|---|---|---|
| Molybdenum hexafluoride | MoF 6 | 17.5 | 34.0 | 209.94 | 3.50 (−140 °C)[2] | 181.7[2] | colourless | |
| Technetium hexafluoride | TcF 6 | 37.4 | 55.3 | (212) | 3.58 (−140 °C)[2] | 181.2[2] | yellow | |
| Ruthenium hexafluoride | RuF 6 | 54 | 215.07 | 3.68 (−140 °C)[2] | 181.8[2] | dark brown | ||
| Rhodium hexafluoride | RhF 6 | ≈ 70 | 216.91 | 3.71 (−140 °C)[2] | 182.4[2] | black | ||
| Tungsten hexafluoride | WF 6 | 2.3 | 17.1 | 297.85 | 4.86 (−140 °C)[2] | 182.6[2] | colourless | |
| Rhenium hexafluoride | ReF 6 | 18.5 | 33.7 | 300.20 | 4.94 (−140 °C)[2] | 182.3[2] | yellow | |
| Osmium hexafluoride | OsF 6 | 33.4 | 47.5 | 304.22 | 5.09 (−140 °C)[2] | 182.9[2] | yellow | |
| Iridium hexafluoride | IrF 6 | 44 | 53.6 | 306.21 | 5.11 (−140 °C)[2] | 183.4[2] | yellow | |
| Platinum hexafluoride | PtF 6 | 61.3 | 69.1 | 309.07 | 5.21 (−140 °C)[2] | 184.8[2] | deep red |
Binary hexafluorides of the actinides
| Compound | Formula | m.p (°C) | b.p. (°C) | subl.p. (°C) | MW | solid ρ (g cm−3) | Bond (pm) | Colour |
|---|---|---|---|---|---|---|---|---|
| Uranium hexafluoride | UF 6 | 64.052 | 56.5 | 351.99 | 5.09 | 199.6 | grey | |
| Neptunium hexafluoride | NpF 6 | 54.4 | 55.18 | (358) | 198.1 | orange | ||
| Plutonium hexafluoride | PuF 6 | 52 | 62 | (356) | 5.08 | 197.1 | brown |
Chemical properties of binary hexafluorides
The hexafluorides have a wide range of chemical reactivity. Sulfur hexafluoride is nearly inert and non-toxic due to steric 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, but tellurium hexafluoride is not very stable and can be hydrolyzed by water within 1 day. Also, both selenium hexafluoride and tellurium hexafluoride are toxic, unlike sulfur hexafluoride (which is non-toxic). In contrast, metal hexafluorides are corrosive, readily hydrolyzed, and may react violently with water. Some of them can be used as fluorinating agents. The metal hexafluorides have a high electron affinity, which makes them strong oxidizing agents.[7]Platinum hexafluoride in particular is notable for its ability to oxidize the dioxygen molecule, O2, to form dioxygenyl hexafluoroplatinate, and for being the first compound that was observed to react with xenon (see xenon hexafluoroplatinate).
Applications of binary hexafluorides
Some metal hexafluorides find applications due to their volatility. Uranium hexafluoride is used in the uranium enrichment process to produce fuel for nuclear reactors. Fluoride volatility can also be exploited for nuclear fuel reprocessing. Tungsten hexafluoride is used in the production of semiconductors through the process of chemical vapor deposition.[8]
Predicted binary hexafluorides
Radon hexafluoride (RnF
6), the heavier homologue of xenon hexafluoride, has been studied theoretically,[9] but has not yet been synthesised. Krypton hexafluoride (KrF
6) has been predicted to be stable, but has not been synthesised due to the extreme difficulty of oxidising krypton beyond Kr(II).[10] The synthesis of americium hexafluoride (AmF
6) by the fluorination of americium(IV) fluoride (AmF
4) was attempted in 1990,[11] but was unsuccessful. Palladium hexafluoride (PdF
6), the lighter homologue of platinum hexafluoride, has been calculated to be stable,[12] but has not yet been produced. Chromium hexafluoride (CrF
6), the lighter homologue of molybdenum hexafluoride and tungsten hexafluoride, was reported, but has been shown to be a mistaken identification of the known pentafluoride (CrF
5).[13]
Literature
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..mw-parser-output cite.citation{font-style:inherit}.mw-parser-output .citation q{quotes:"""""""'""'"}.mw-parser-output .citation .cs1-lock-free a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/6/65/Lock-green.svg/9px-Lock-green.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .citation .cs1-lock-limited a,.mw-parser-output .citation .cs1-lock-registration a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/d/d6/Lock-gray-alt-2.svg/9px-Lock-gray-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .citation .cs1-lock-subscription a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/a/aa/Lock-red-alt-2.svg/9px-Lock-red-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration{color:#555}.mw-parser-output .cs1-subscription span,.mw-parser-output .cs1-registration span{border-bottom:1px dotted;cursor:help}.mw-parser-output .cs1-ws-icon a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/4/4c/Wikisource-logo.svg/12px-Wikisource-logo.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output code.cs1-code{color:inherit;background:inherit;border:inherit;padding:inherit}.mw-parser-output .cs1-hidden-error{display:none;font-size:100%}.mw-parser-output .cs1-visible-error{font-size:100%}.mw-parser-output .cs1-maint{display:none;color:#33aa33;margin-left:0.3em}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration,.mw-parser-output .cs1-format{font-size:95%}.mw-parser-output .cs1-kern-left,.mw-parser-output .cs1-kern-wl-left{padding-left:0.2em}.mw-parser-output .cs1-kern-right,.mw-parser-output .cs1-kern-wl-right{padding-right:0.2em}
References
^ Wiberg, Wiberg & Holleman 2001, p. 436.
^ abcdefghijklmnopqrs Drews, T.; Supeł, J.; Hagenbach, A.; Seppelt, K. (2006). "Solid state molecular structures of transition metal hexafluorides". Inorganic Chemistry. 45 (9): 3782–3788. doi:10.1021/ic052029f. PMID 16634614.
^ Wilhelm Klemm and Paul Henkel "Über einige physikalische Eigenschaften von SF6, SeF6, TeF6 und CF4" Z. anorg. allgem. Chem. 1932, vol. 207, pages 73–86. doi:10.1002/zaac.19322070107
^ "4. Physical Constants of Inorganic Compound". CRC Handbook of Chemistry and Physics (90 ed.). Boca Raton, FL: CRC Press. 2009. pp. 4–95. ISBN 978-1-4200-9084-0.
^ CAS #35473-38-2
^ ab Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils, ed., Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, p. 594, ISBN 0-12-352651-5
^ Bartlett, N. (1968). "The Oxidizing Properties of the Third Transition Series Hexafluorides and Related Compounds". Angewandte Chemie International Edition. 7 (6): 433–439. doi:10.1002/anie.196804331.
^ http://www.timedomaincvd.com/CVD_Fundamentals/films/W_WSi.html
^ Filatov, M.; Cremer, D. (2003). "Bonding in radon hexafluoride: An unusual relativistic problem". Physical Chemistry Chemical Physics. 2003 (5): 1103–1105. Bibcode:2003PCCP....5.1103F. doi:10.1039/b212460m.
^ Dixon, D. A.; Wang, T. H.; Grant, D. J.; Peterson, K. A.; Christe, K. O.; Schrobilgen, G. J. (2007). "Heats of Formation of Krypton Fluorides and Stability Predictions for KrF4 and KrF6 from High Level Electronic Structure Calculations". Inorganic Chemistry. 46 (23): 10016–10021. doi:10.1021/ic701313h. PMID 17941630.
^ Malm, J. G.; Weinstock, B.; Weaver, E. E. (1958). "The Preparation and Properties of NpF6; a Comparison with PuF6". The Journal of Physical Chemistry. 62 (12): 1506–1508. doi:10.1021/j150570a009.
^ Aullón, G.; Alvarez, S. (2007). "On the Existence of Molecular Palladium(VI) Compounds: Palladium Hexafluoride". Inorganic Chemistry. 46 (7): 2700–2703. doi:10.1021/ic0623819. PMID 17326630.
^ Riedel, S.; Kaupp, M. (2009). "The highest oxidation states of the transition metal elements" (PDF). Coordination Chemistry Reviews. 253 (5–6): 606–624. doi:10.1016/j.ccr.2008.07.014.
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