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Hexafluoride
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<h2 id="hexafluoride-cations">Hexafluoride cations</h2> <p>Cationic hexafluorides exist but are rarer than neutral or anionic hexafluorides. Examples are the hexafluorochlorine (ClF6+), and hexafluorobromine (BrF6+) <a href="/facts/Ion/3bePpDna">cations</a>.<a class="footnote-ref" id="fnref:1" href="#fn:1"><sup>1</sup></a> </p> <h2 id="hexafluoride-anions">Hexafluoride anions</h2> <p>Many elements form anionic hexafluorides. Members of commercial interest are <a href="/facts/Hexafluorophosphate/mXYXC2Bf">hexafluorophosphate</a> (PF6−) and <a href="/facts/Hexafluorosilicate/J35wRTRh">hexafluorosilicate</a> (SiF62−). </p><p>Many transition metals form hexafluoride anions. Often the monoanions are generated by reduction of the neutral hexafluorides. For example, <a href="/facts/Dioxygenyl_hexafluoroplatinate/AUcQD8RL">PtF6−</a> 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 <a href="/facts/Caesium_hexafluorocuprate(IV)/GLTWxj0c">hexafluorocuprate(IV)</a>, CuF2−6 and <a href="/facts/Hexafluoronickelate(IV)/o48Hpm0l">hexafluoronickelate(IV)</a>, NiF2−6. </p> <h2 id="binary-hexafluorides">Binary hexafluorides</h2> <p>Seventeen elements are known to form binary hexafluorides.<a class="footnote-ref" id="fnref:2" href="#fn:2"><sup>2</sup></a> Nine of these elements are <a href="/facts/Transition_metal/0slGWoLQ">transition metals</a>, three are <a href="/facts/Actinide/kHC4lxJV">actinides</a>, four are <a href="/facts/Chalcogen/nXLGmnqL">chalcogens</a>, and one is a <a href="/facts/Noble_gas/j0yHMEAE">noble gas</a>. Most hexafluorides are <a href="/facts/Molecule/N9ljSfFq">molecular</a> compounds with low <a href="/facts/Melting_point/MY0w0h0k">melting</a> and <a href="/facts/Boiling_point/E2BK9QoS">boiling points</a>. Four hexafluorides (S, Se, Te, and W) are gases at room temperature (25 °C) and a pressure of 1 <a href="/facts/Atmosphere_(unit)/aTbYALII">atm</a>, two are liquids (Re, Mo), and the others are volatile solids. The <a href="/facts/Group_6_element/s8cA1Rwd">group 6</a>, <a href="/facts/Chalcogen/nXLGmnqL">chalcogen</a>, and <a href="/facts/Noble_gas/j0yHMEAE">noble gas</a> hexafluorides are colourless, but the other hexafluorides have colours ranging from white, through yellow, orange, red, brown, and grey, to black. </p><p>The molecular geometry of binary hexafluorides is generally <a href="/facts/Octahedral_molecular_geometry/bNXLVvks">octahedral</a>, although some derivatives are distorted from Oh <a href="/facts/Symmetry_group/CNqeseMp">symmetry</a>. For the main group hexafluorides, distortion is pronounced for the 14-electron noble gas derivatives. Distortions in gaseous <a href="/facts/Xenon_hexafluoride/QU16jh1g">XeF6</a> are caused by its non-bonding <a href="/facts/Lone_pair/Q6kEAC4g">lone pair</a>, according to <a href="/facts/VSEPR_theory/WqsRgVRV">VSEPR theory</a>. In the solid state, it adopts a complex structure involving tetramers and hexamers. According to <a href="/facts/Quantum_chemical/H9AXV0lx">quantum chemical</a> 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.<a class="footnote-ref" id="fnref:3" href="#fn:3"><sup>3</sup></a> </p><p><a href="/facts/Polonium_hexafluoride/4hQcWfvK">Polonium hexafluoride</a> is known, but not well-studied. It could not be made from 210Po, but using the longer-lived isotope 208Po and reacting it with fluorine found a volatile product that is almost certainly PoF6.<a class="footnote-ref" id="fnref:4" href="#fn:4"><sup>4</sup></a> The quoted boiling point in the table below is a prediction. </p> <h3>Binary hexafluorides of the chalcogens</h3> <table><tbody><tr><th>Compound</th><th><a href="/facts/Chemical_formula/hghUtthl">Formula</a></th><th><a href="/facts/Melting_point/MY0w0h0k">m.p</a> (°C)</th><th><a href="/facts/Boiling_point/E2BK9QoS">b.p.</a> (°C)</th><th><a href="/facts/Sublimation_point/KDoKokMn">subl.p.</a> (°C)</th><th><a href="/facts/Molar_mass/m2laJErl">MW</a></th><th>solid <a href="/facts/Density/92p6hh9I">ρ</a> (g cm−3) (at m.p.)<a class="footnote-ref" id="fnref:5" href="#fn:5"><sup>5</sup></a></th><th>Bond distance (<a href="/facts/Picometer/Cb54mVBX">pm</a>)</th><th>Colour</th></tr><tr><td><a href="/facts/Sulfur_hexafluoride/jeee0PJ8">Sulfur hexafluoride</a></td><td>SF6</td><td>−50.8</td><td></td><td>−63.8</td><td>146.06</td><td>2.51 (−50 °C)</td><td>156.4</td><td>colourless</td></tr><tr><td><a href="/facts/Selenium_hexafluoride/aBFQIy2g">Selenium hexafluoride</a></td><td>SeF6</td><td>−34.6</td><td></td><td>−46.6</td><td>192.95</td><td>3.27</td><td>167–170</td><td>colourless</td></tr><tr><td><a href="/facts/Tellurium_hexafluoride/aEKw0fYb">Tellurium hexafluoride</a><a class="footnote-ref" id="fnref:6" href="#fn:6"><sup>6</sup></a></td><td>TeF6</td><td>−38.9</td><td>−37.6</td><td></td><td>241.59</td><td>3.76</td><td>184</td><td>colourless</td></tr><tr><td><a href="/facts/Polonium_hexafluoride/4hQcWfvK">Polonium hexafluoride</a><a class="footnote-ref" id="fnref:7" href="#fn:7"><sup>7</sup></a><a class="footnote-ref" id="fnref:8" href="#fn:8"><sup>8</sup></a></td><td>PoF6</td><td></td><td>≈ −40?</td><td></td><td>322.99</td><td></td><td></td><td>colourless<a class="footnote-ref" id="fnref:9" href="#fn:9"><sup>9</sup></a></td></tr></tbody></table> <h3>Binary hexafluorides of the noble gases</h3> <table><tbody><tr><th>Compound</th><th><a href="/facts/Chemical_formula/hghUtthl">Formula</a></th><th><a href="/facts/Melting_point/MY0w0h0k">m.p</a> (°C)</th><th><a href="/facts/Boiling_point/E2BK9QoS">b.p.</a> (°C)</th><th><a href="/facts/Sublimation_point/KDoKokMn">subl.p.</a> (°C)</th><th><a href="/facts/Molar_mass/m2laJErl">MW</a></th><th>solid <a href="/facts/Density/92p6hh9I">ρ</a> (g cm−3)</th><th>Bond (<a href="/facts/Picometer/Cb54mVBX">pm</a>)</th><th>Colour</th></tr><tr><td><a href="/facts/Xenon_hexafluoride/QU16jh1g">Xenon hexafluoride</a></td><td>XeF6</td><td>49.5</td><td>75.6</td><td></td><td>245.28</td><td>3.56</td><td></td><td>colourless (solid)yellow (gas)</td></tr></tbody></table> <h3>Binary hexafluorides of the transition metals</h3> <table><tbody><tr><th>Compound</th><th><a href="/facts/Chemical_formula/hghUtthl">Formula</a></th><th><a href="/facts/Melting_point/MY0w0h0k">m.p</a> (°C)</th><th><a href="/facts/Boiling_point/E2BK9QoS">b.p.</a> (°C)</th><th><a href="/facts/Sublimation_point/KDoKokMn">subl.p.</a> (°C)</th><th><a href="/facts/Molar_mass/m2laJErl">MW</a></th><th>solid <a href="/facts/Density/92p6hh9I">ρ</a> (g cm−3)</th><th>Bond (<a href="/facts/Picometer/Cb54mVBX">pm</a>)</th><th>Colour</th></tr><tr><td><a href="/facts/Molybdenum_hexafluoride/99vwmHAB">Molybdenum hexafluoride</a></td><td>MoF6</td><td>17.5</td><td>34.0</td><td></td><td>209.94</td><td>3.50 (−140 °C)<a class="footnote-ref" id="fnref:10" href="#fn:10"><sup>10</sup></a></td><td>181.7<a class="footnote-ref" id="fnref:11" href="#fn:11"><sup>11</sup></a></td><td>colourless</td></tr><tr><td><a href="/facts/Technetium_hexafluoride/VpxJ5c7X">Technetium hexafluoride</a></td><td>TcF6</td><td>37.4</td><td>55.3</td><td></td><td>(212)</td><td>3.58 (−140 °C)<a class="footnote-ref" id="fnref:12" href="#fn:12"><sup>12</sup></a></td><td>181.2<a class="footnote-ref" id="fnref:13" href="#fn:13"><sup>13</sup></a></td><td>yellow</td></tr><tr><td><a href="/facts/Ruthenium_hexafluoride/Fy2uTTcx">Ruthenium hexafluoride</a></td><td>RuF6</td><td>54</td><td></td><td></td><td>215.07</td><td>3.68 (−140 °C)<a class="footnote-ref" id="fnref:14" href="#fn:14"><sup>14</sup></a></td><td>181.8<a class="footnote-ref" id="fnref:15" href="#fn:15"><sup>15</sup></a></td><td>dark brown</td></tr><tr><td><a href="/facts/Rhodium_hexafluoride/HGrRWsFQ">Rhodium hexafluoride</a></td><td>RhF6</td><td>≈ 70</td><td></td><td></td><td>216.91</td><td>3.71 (−140 °C)<a class="footnote-ref" id="fnref:16" href="#fn:16"><sup>16</sup></a></td><td>182.4<a class="footnote-ref" id="fnref:17" href="#fn:17"><sup>17</sup></a></td><td>black</td></tr><tr><td><a href="/facts/Tungsten_hexafluoride/Hf0S5IVa">Tungsten hexafluoride</a></td><td>WF6</td><td>2.3</td><td>17.1</td><td></td><td>297.85</td><td>4.86 (−140 °C)<a class="footnote-ref" id="fnref:18" href="#fn:18"><sup>18</sup></a></td><td>182.6<a class="footnote-ref" id="fnref:19" href="#fn:19"><sup>19</sup></a></td><td>colourless</td></tr><tr><td><a href="/facts/Rhenium_hexafluoride/yE3MJb6B">Rhenium hexafluoride</a></td><td>ReF6</td><td>18.5</td><td>33.7</td><td></td><td>300.20</td><td>4.94 (−140 °C)<a class="footnote-ref" id="fnref:20" href="#fn:20"><sup>20</sup></a></td><td>182.3<a class="footnote-ref" id="fnref:21" href="#fn:21"><sup>21</sup></a></td><td>yellow</td></tr><tr><td><a href="/facts/Osmium_hexafluoride/H29lBtak">Osmium hexafluoride</a></td><td>OsF6</td><td>33.4</td><td>47.5</td><td></td><td>304.22</td><td>5.09 (−140 °C)<a class="footnote-ref" id="fnref:22" href="#fn:22"><sup>22</sup></a></td><td>182.9<a class="footnote-ref" id="fnref:23" href="#fn:23"><sup>23</sup></a></td><td>yellow</td></tr><tr><td><a href="/facts/Iridium_hexafluoride/Buqb1nnD">Iridium hexafluoride</a></td><td>IrF6</td><td>44</td><td>53.6</td><td></td><td>306.21</td><td>5.11 (−140 °C)<a class="footnote-ref" id="fnref:24" href="#fn:24"><sup>24</sup></a></td><td>183.4<a class="footnote-ref" id="fnref:25" href="#fn:25"><sup>25</sup></a></td><td>yellow</td></tr><tr><td><a href="/facts/Platinum_hexafluoride/Nqqdd4Xc">Platinum hexafluoride</a></td><td>PtF6</td><td>61.3</td><td>69.1</td><td></td><td>309.07</td><td>5.21 (−140 °C)<a class="footnote-ref" id="fnref:26" href="#fn:26"><sup>26</sup></a></td><td>184.8<a class="footnote-ref" id="fnref:27" href="#fn:27"><sup>27</sup></a></td><td>deep red</td></tr></tbody></table> <h3>Binary hexafluorides of the actinides</h3> <table><tbody><tr><th>Compound</th><th><a href="/facts/Chemical_formula/hghUtthl">Formula</a></th><th><a href="/facts/Melting_point/MY0w0h0k">m.p</a> (°C)</th><th><a href="/facts/Boiling_point/E2BK9QoS">b.p.</a> (°C)</th><th><a href="/facts/Sublimation_point/KDoKokMn">subl.p.</a> (°C)</th><th><a href="/facts/Molar_mass/m2laJErl">MW</a></th><th>solid <a href="/facts/Density/92p6hh9I">ρ</a> (g cm−3)</th><th>Bond (<a href="/facts/Picometer/Cb54mVBX">pm</a>)</th><th>Colour</th></tr><tr><td><a href="/facts/Uranium_hexafluoride/pV5BCMkJ">Uranium hexafluoride</a></td><td>UF6</td><td>64.052</td><td></td><td>56.5</td><td>351.99</td><td>5.09</td><td>199.6</td><td>colorless</td></tr><tr><td><a href="/facts/Neptunium_hexafluoride/lwBW2oxH">Neptunium hexafluoride</a></td><td>NpF6</td><td>54.4</td><td>55.18</td><td></td><td>(351)</td><td></td><td>198.1</td><td>orange</td></tr><tr><td><a href="/facts/Plutonium_hexafluoride/0WyT3R20">Plutonium hexafluoride</a></td><td>PuF6</td><td>52</td><td>62</td><td></td><td>(358)</td><td>5.08</td><td>197.1</td><td>brown</td></tr></tbody></table> <h3>Chemical properties of binary hexafluorides</h3> <p>The hexafluorides have a wide range of chemical reactivity. <a href="/facts/Sulfur_hexafluoride/jeee0PJ8">Sulfur hexafluoride</a> is nearly inert and non-toxic due to <a href="/facts/Steric_effects/TnX5c8HG">steric hindrance</a> (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. <a href="/facts/Selenium_hexafluoride/aBFQIy2g">Selenium hexafluoride</a> is nearly as unreactive as SF6, but <a href="/facts/Tellurium_hexafluoride/aEKw0fYb">tellurium hexafluoride</a> is not very stable and can be <a href="/facts/Hydrolysis/JTvwIHPs">hydrolyzed</a> 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 as <a href="/facts/Fluorinating_agent/vYrGVWcJ">fluorinating agents</a>. The metal hexafluorides have a high <a href="/facts/Electron_affinity/1mf4A396">electron affinity</a>, which makes them strong oxidizing agents.<a class="footnote-ref" id="fnref:28" href="#fn:28"><sup>28</sup></a> <a href="/facts/Platinum_hexafluoride/Nqqdd4Xc">Platinum hexafluoride</a> in particular is notable for its ability to oxidize the <a href="/facts/Dioxygen/UVNYnKUQ">dioxygen</a> molecule, O2, to form <a href="/facts/Dioxygenyl_hexafluoroplatinate/AUcQD8RL">dioxygenyl hexafluoroplatinate</a>, and for being the first compound that was observed to react with xenon (see <a href="/facts/Xenon_hexafluoroplatinate/RtRHxpjp">xenon hexafluoroplatinate</a>). </p> <h3>Applications of binary hexafluorides</h3> <p>Some metal hexafluorides find applications due to their volatility. <a href="/facts/Uranium_hexafluoride/pV5BCMkJ">Uranium hexafluoride</a> is used in the <a href="/facts/Uranium_enrichment/vSb2Tcsb">uranium enrichment</a> process to produce fuel for <a href="/facts/Nuclear_reactor/8xqjIbge">nuclear reactors</a>. <a href="/facts/Fluoride_volatility/Ckc3ltaq">Fluoride volatility</a> can also be exploited for <a href="/facts/Nuclear_fuel_reprocessing/aTMUoYaW">nuclear fuel reprocessing</a>. <a href="/facts/Tungsten_hexafluoride/Hf0S5IVa">Tungsten hexafluoride</a> is used in the production of <a href="/facts/Semiconductors/zFf3mGTE">semiconductors</a> through the process of <a href="/facts/Chemical_vapor_deposition/7khLxeU5">chemical vapor deposition</a>.<a class="footnote-ref" id="fnref:29" href="#fn:29"><sup>29</sup></a> </p> <h3>Predicted binary hexafluorides</h3> <h4>Radon hexafluoride</h4> <p><a href="/facts/Radon_hexafluoride/8bp389W2">Radon hexafluoride</a> (RnF6), the heavier homologue of <a href="/facts/Xenon_hexafluoride/QU16jh1g">xenon hexafluoride</a>, has been studied theoretically,<a class="footnote-ref" id="fnref:30" href="#fn:30"><sup>30</sup></a> but its synthesis has not yet been confirmed. Higher fluorides of <a href="/facts/Radon/0V9arWKl">radon</a> may have been observed in experiments where unknown radon-containing products distilled together with <a href="/facts/Xenon_hexafluoride/QU16jh1g">xenon hexafluoride</a>, and perhaps in the production of radon trioxide: these may have been RnF4, RnF6, or both.<a class="footnote-ref" id="fnref:31" href="#fn:31"><sup>31</sup></a> 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 of <a href="/facts/Radon_difluoride/MwhWcqrW">RnF2</a> 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 to <a href="/facts/Spin%E2%80%93orbit_interaction/Pv6maL1S">spin–orbit</a> splitting of the 6p shell of radon (RnIV would have a closed-shell 6s26p21/2 configuration).<a class="footnote-ref" id="fnref:32" href="#fn:32"><sup>32</sup></a> 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.<a class="footnote-ref" id="fnref:33" href="#fn:33"><sup>33</sup></a> Continuing the trend, the heavier <a href="/facts/Oganesson/zs5MQDjV">oganesson</a> hexafluoride should be unbound.<a class="footnote-ref" id="fnref:34" href="#fn:34"><sup>34</sup></a> </p> <h4>Others</h4> <p><a href="/facts/Krypton_hexafluoride/sw1lWkRv">Krypton hexafluoride</a> (KrF6) has been predicted to be stable, but has not been synthesised due to the extreme difficulty of oxidising <a href="/facts/Krypton/llkNlCke">krypton</a> beyond Kr(II).<a class="footnote-ref" id="fnref:35" href="#fn:35"><sup>35</sup></a> The synthesis of <a href="/facts/Americium_hexafluoride/Ddgq5eRg">americium hexafluoride</a> (AmF6) by the <a href="/facts/Halogenation/vYrGVWcJ">fluorination</a> of <a href="/facts/Americium(IV)_fluoride/AMHCqcDT">americium(IV) fluoride</a> (AmF4) was attempted in 1990,<a class="footnote-ref" id="fnref:36" href="#fn:36"><sup>36</sup></a> but was unsuccessful; there have also been possible thermochromatographic identifications of it and <a href="/facts/Curium_hexafluoride/61KPzfjz">curium hexafluoride</a> (CmF6), but it is debated if these are conclusive.<a class="footnote-ref" id="fnref:37" href="#fn:37"><sup>37</sup></a> <a href="/facts/Palladium_hexafluoride/dATsgIff">Palladium hexafluoride</a> (PdF6), the lighter homologue of <a href="/facts/Platinum_hexafluoride/Nqqdd4Xc">platinum hexafluoride</a>, has been calculated to be stable,<a class="footnote-ref" id="fnref:38" href="#fn:38"><sup>38</sup></a> but has not yet been produced; the possibility of silver (AgF6) and <a href="/facts/Gold_hexafluoride/LoGzyWN7">gold hexafluorides</a> (AuF6) has also been discussed.<a class="footnote-ref" id="fnref:39" href="#fn:39"><sup>39</sup></a> <a href="/facts/Chromium_hexafluoride/YtKvuG87">Chromium hexafluoride</a> (CrF6), the lighter homologue of <a href="/facts/Molybdenum_hexafluoride/99vwmHAB">molybdenum hexafluoride</a> and <a href="/facts/Tungsten_hexafluoride/Hf0S5IVa">tungsten hexafluoride</a>, was reported, but has been shown to be a mistaken identification of the known <a href="/facts/Chromium_pentafluoride/UU3regxv">pentafluoride</a> (CrF5).<a class="footnote-ref" id="fnref:40" href="#fn:40"><sup>40</sup></a> </p> <h2 id="literature">Literature</h2> <ul><li>Galkin, N. P.; Tumanov, Yu. N. (1971). <a href="https://web.archive.org/web/20151130091739/http://www.turpion.org/php/paper.phtml?journal_id=rc&paper_id=1902">"Reactivity and Thermal Stability of Hexafluorides"</a>. <i>Russian Chemical Reviews</i>. 40 (2): 154–164. <a href="/facts/Bibcode_(identifier)/9HtdQSGB">Bibcode</a>:<a href="https://ui.adsabs.harvard.edu/abs/1971RuCRv..40..154G">1971RuCRv..40..154G</a>. <a href="/facts/Doi_(identifier)/muM9Etpq">doi</a>:<a href="https://doi.org/10.1070%2FRC1971v040n02ABEH001902">10.1070/RC1971v040n02ABEH001902</a>. <a href="/facts/S2CID_(identifier)/ldJsHa2Y">S2CID</a> <a href="https://api.semanticscholar.org/CorpusID:250901336">250901336</a>. Archived from <a href="http://www.turpion.org/php/paper.phtml?journal_id=rc&paper_id=1902">the original</a> on 2015-11-30. Retrieved 2012-05-12.</li></ul> <h2 id="sources">Sources</h2> <ul><li>Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). <a href="https://books.google.com/books?id=Mtth5g59dEIC"><i>Inorganic chemistry</i></a>. Academic Press. <a href="/facts/ISBN_(identifier)/15AdSPa9">ISBN</a> 978-0-12-352651-9. Retrieved 3 March 2011.</li></ul> Wikimedia Commons has media related to Hexafluorides. <h2 id="references">References</h2> <ol> <li id="fn:1"><p>Wiberg, Wiberg & Holleman 2001, p. 436. - Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN 978-0-12-352651-9. Retrieved 3 March 2011. <a href="https://books.google.com/books?id=Mtth5g59dEIC" target="_blank">https://books.google.com/books?id=Mtth5g59dEIC</a> <a href="#fnref:1" class="footnote-back-ref">↩</a></p></li> <li id="fn:2"><p>Seppelt, Konrad (2015). "Molecular Hexafluorides". Chemical Reviews. 115 (2): 1296–1306. doi:10.1021/cr5001783. 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