Review pubs.acs.org/CR
Molecular Hexafluorides Konrad Seppelt* Institut fuer Chemie and Biochemie, Freie Universitaet, Fabeckstr. 34-36, 14195 Berlin, Germany With the sole exception of XeF6, all other molecular hexafluorides are in certain physical aspects very much alike but in other physical aspects very different. Their chemical behavior varies extremely, from SF6 as one of the most chemically inert gases to extremely reactive compounds like PtF6, XeF6, and PuF6. Within each group (main group hexafluorides, transition metal hexafluorides, and actinide hexafluorides), the chemical behavior follows certain trends. This is particularly true for the nine transition metal hexafluorides that, for example, can be used as one-electron CONTENTS oxidants with predictable oxidation force from rather mild to extreme. In summary, there is hardly another class of 1. Introduction A compounds that are so numerous, so seemingly alike but that 2. Main Group Hexafluorides A differ so much in other aspects. 2.1. Sulfur Hexafluoride and Selenium Hexafluoride 2.2. Tellurium Hexafluoride 2.3. Polonium Hexafluoride 2.4. Xenon Hexafluoride 2.5. Radon Hexafluoride 3. Transition Metal Hexafluorides 3.1. Molybdenum and Tungsten Hexafluoride 3.2. Technetium, Ruthenium, and Rhodium Hexafluoride 3.3. Rhenium-, Osmium-, Iridium-, and Platinum Hexafluoride 3.4. Unknown Transition Metal Hexafluorides 4. Actinide Hexafluorides 4.1. Uranium, Neptunium, and Plutonium Hexafluoride 4.2. Americium and Curium Hexafluoride 5. Concluding Remarks Author Information Corresponding Author Notes Biography References
A B B B C C C
2. MAIN GROUP HEXAFLUORIDES SF6, SeF6, TeF6, and XeF6 are very well-documented, while little is known about PoF6 and RnF6. All are best made by reaction of the elements with elemental fluorine. 2.1. Sulfur Hexafluoride and Selenium Hexafluoride
More than 20 000 publications exist about SF6. It was first detected by Moissan in 1900.1 It is an extreme unreactive, even physiologically inert gas. It is the model compound for octahedral symmetry. It is produced worldwide by fluorination of sulfur with elemental fluorine in large quantities; the amount produced can only be compared to that of UF6. Many of the technical uses that are all based on its inertness are now abandoned for ecological reasons, but it remains the only gas that is used in equipment for handling high voltage and strong current electricity (switches, transformers, etc.). The search for its replacement for this special purpose remains unsuccessful today. Its chemical inertness makes it so far almost useless as a starting material for chemical synthesis. Reactions often occur only above 400 °C, e.g., with H2O, Al2O3, and countless other compounds, and end in complete destruction of the molecule. Usually the chemical inertness is ascribed to the shielding of the central sulfur atom by the arrangement of the six fluorine atoms around it. This steric argument is only valid for preventing any kind of “SN2”-type reactions, namely, reactions with a nucleophile followed by expulsion of one F−. But this steric argument cannot hold for a “SN1”-type reaction, where an F− ion would be abstracted first under formation of an intermediate SF5+ cation that then can react with a nucleophile. The inertness against even the strongest Lewis acids is a result of the bonding situation. If only one of the six fluorine atoms is replaced by an even larger methyl group as in CH3−SF5, then the cation CH3−SF4+ can be generated with SbF5.4 Also SClF5
D E G G G H H H H H H H
1. INTRODUCTION There exist 16 well-defined molecular hexafluorides. Only a small number of yet unknown or ill-defined ones remain to be isolated. The first one to be generated has certainly been SF6,1 which up to date remains by far the most important one for technological purposes. This large number of compounds is in contrast to the few molecular hexachlorides (MoCl6, WCl6, ReCl 6 , and UCl 6 , of which two have only recently unambiguously been identified 2) and the only known hexabromide, WBr6. There exist also a small number of mixed molecular hexahalides, SClF5, SBrF5, SeClF5, TeClF5, TeBrF5, and WClF5. Only the first two have found preparative use, as is presented in another paper in this issue.3 © XXXX American Chemical Society
Special Issue: 2015 Fluorine Chemistry Received: March 30, 2014
A
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Figure 1. Two crystal structures of XeF6, showing tetrameric units. Left: The tetrameric unit in the two high-temperature modifications can be described as a trimeric unit of (XeF5+F−)3 with a loosely connected XeF6 unit. Xe−F bond lengths in the monomer unit with close to C2v symmetry range from 183 to 206 pm. Broken lines represent Xe···F distances of 228−245 pm, and dotted lines represent distances of 261−268 pm. Right: At lower temperature the tetrameric (right) units appear with bridging distances of 227−240 pm, which can be described as (XeF5+F−)4. Here the nonbridging Xe−F bond lengths range from 184 to 190 pm.
antiprismatic TeF82−.17 Fluorine atoms can be replaced by oxygen and nitrogen containing ligands: as in HOTeF5,18 (HO) 2 TeF 4 , 18 CH 3 OTeF 5 , 19 (CH 3 O) 2 TeF 4 , 19 H 2 NTeF 5 , 20 (CH3)2NTeF5,21 [(CH3)2N]2TeF4,21 [(CH3)3SiN]2TeF4,20 and others. Of these, HOTeF5 is an important product, although it is generally made from Te(OH)6 and FSO3H. It serves as a source for the highly electronegative −OTeF5 ligand, which has been combined with almost every element in the periodic table from hydrogen over to krypton and xenon up to uranium. The reaction between TeF6 and (CH3)3N is an example of a redox process; it results in (CH3)3N−CH2− N(CH3)2+TeF5−.22
is known to undergo nucleophilic displacement reactions of the fluorine atom opposite to the chlorine atom.5 (C3H7)2SF4 is already very sensitive toward hydrolysis under formation of (C3H7)2SO2.6 The only case where a chemical reaction has been reported under mild conditions and without complete destruction of the molecule is the reduction with (NC)2C C(CN)2 yielding the square pyramidal SF5− anion,7which can be made more readily directly from SF4 and F−. SF6 has been used as fluorination source for metallocenes.8 Reactions at ambient temperature occur under complete destruction with rhodium and nickel complexes.9,10 SF6 has been the focus of many theoretical considerations over the years. Now it is accepted that d-orbital participation in the bonding (“sp3d2 hybridization”) is negligible in the bonding of SF6 and related compounds. Most recent computations give very precise numbers on the energetics and structural data for SF6 and related compounds.11,12 The interpretations of these data state that the 3-center-4-electron type bonding is an incomplete description,12,13 because this would imply a rather weak S−F bond. An increased valence structure model allowing only two 3-center-4-electron bonds has been proposed, in addition to four 2-center-2-electron bonds.13,14 Nevertheless, it remains a qualitative model. In many aspects SeF6 is very similar to SF6. It is also quite stable against nucleophilic attack. The selenium atom is only a little larger than the sulfur atom, which is a consequence of the transition metal contraction. There is again no indication of the formation of either SeF5+ or SeF7− ions. It has even been proposed as a substitute for SF6 in electric switches for extinguishing light arcs.15 However, SeF6 causes selenium poisoning, which is indicative of an enzymatic reduction, in contrast to SF6. The main difference to SF6 is indeed its oxidative capacity; it is easily reduced under ambient conditions, e.g., by I−/H2O. Even at −78 °C a reaction occurs with C6H5Li/ether,16 which could be a convenient source for organic selenium compounds. The higher oxidation force is certainly another consequence of the transition metal contraction, here the stabilization of the s-shell as compared to the p-shell.
2.3. Polonium Hexafluoride
Considering the chemistry of tellurium hexafluoride, it would be important to know more about PoF6. It could be unstable or even nonexistent due to the stabilization of the 6s shell by the lanthanide contraction and the strong relativistic effect. So far it could not be made from readily available 210Po (139-day halflife), most likely due to radioactive self-decay.23 However, a fraction of a mg of 208Po (2.9-year half-life) has been treated with elemental fluorine, and it has been shown that there is a transport of 208Po activity, indicating a volatile compound that realistically can only be PoF6.24 With 209Po (103-year half-life) more chemistry could be done, but similar to 208Po it can only be generated by a p,n-cyclotron reaction from 209Bi in trace amounts. 2.4. Xenon Hexafluoride
Xenon hexafluoride is in this context a unique case. In contrast to all other hexafluorides, it is not monomeric in the solid and liquid or dissolved state, and the molecule does not adopt the octahedral structure. There is already vast literature sources about its molecular structure. An important experimental fact is the early electron diffraction measurement. It clearly showed that the molecule is not octahedral.25 Most likely it has the shape of a slightly C3v distorted octahedron, of course in a dynamic way. When it appears in crystal structures as a molecule with weak secondary contacts to neighboring fluorine atoms, e.g., in two modification of XeF6 or in NO2+XeF7−, it has a C2v distorted octahedral structure;26,27 see Figure 1. The C2v distorted structure can be considered a transition state between C3v distorted states.25 All this points to a “steric activity of the non-bonding electron pair”, but to a much weaker degree than
2.2. Tellurium Hexafluoride
Systematic chemistry starts with TeF6. F− can be added once and twice, forming pentagonal bipyramidal TeF7− and square B
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in less highly coordinated systems. Many theoretical and computational studies have been done over the years to reproduce these experimental findings. Even with the most sophisticated methods, it remains difficult to reproduce the experiments by computations.28 In the solid state, eight modifications have been observed, but not all are fully characterized.27 Often it appears as a tetramer that could be formulated by (XeF5+F−)4. However, hexamers and trimers−monomers are also observed; see Figure 1. More modifications are likely to be found at high pressure. This behavior in the solid state is the most markedly different from all other known molecular hexafluorides. In solutions of inert solvents at low temperatures, it is also tetrameric, as shown by 19F and 129Xe NMR.29 XeF6 is an extremely reactive species. With F− ion acceptors it forms the square pyramidal XeF5+,30 and with F− ion donors it forms the capped octahedral XeF7−26 and the square antiprismatic XeF82−31 ions. The series XeF5+, XeF6, XeF7−, XeF82− is a prominent example of the decreasing steric activity of the nonbonding electron pair with increased crowding. XeF6 is of course an extreme fluorination agent, but this has not been systematically tested. The use of XeF6 is limited by its hydrolysis to extremely explosive XeO3. This reaction occurs stepwise; XeOF4 and XeO2F2 are intermediates, although the latter is better made by reaction from XeOF4 and CsNO3.32 The only other nucleophilic displacement reaction on XeF6 seems to be the formation of unstable Xe(OTeF5)6.33
Table 1. Some Physical Properties of the Hexafluorides, Except XeF6; The Volume Per Molecule Data Refer to the Orthorhombic Phase, Except SF6 and SeF6 SF6
SeF6
TeF6
−50.8 −63.8 156.1(1)43 91.0(75 K)41 347 MoF6 TcF6
−34.5 −45.9 169.0(10)43 97.841 264 RuF6
−37.6 −38.9 181.0(15K)40 102.440 197 RhF6
mp [°C] bp [°C]
17.438 35.0
37.445 55.345
X−Fav [pm]39 vol/molecule [106 pm3]38 ν5(T2u) [cm−1]44
181.3 99.5
181.2 98.3
54.046 ∼70 extrapol48 181.6 97.0
70.047 ∼70 extrapol48 182.1 97.0
116
145
186
mp [°C] subl. [deg] X−Fav [pm] vol/molecule [106 pm3] ν5(T2u) [cm−1]44
mp [°C]38 bp [°C]49 X−Fav [pm]39 vol/molecule [106 pm3]38 ν5(T2u) [cm−1]44 mp [°C] bp [°C] X−F vol/molecule [106 pm3]38 ν5(T2u)44
2.5. Radon Hexafluoride
It seemed that the noble gas radon forms only one fluoride that is nonvolatile. RnF2 has been suggested for its composition.34 It is also possible that several radon fluorides exist that behave similarly in terms of nonvolatility. If RnF6 were to exist, it is likely the highest fluoride possible, due to the increase of relativistic spin-orbit coupling.35 (The still heavier [118]F6 is expected to be unbound, since [118]F4 might have a tetrahedral structure with two inert electron pairs.36) Also because of the influence of relativistics, molecular RnF6 should be regular octahedral.34 Another outcome of the relativistics is that the Rn−F bonds become quite polar, so that highly fluorinebridged structures might exist in the solid, which would result in the nonvolatility of the potential Rn-fluorides.37 Therefore, it could be explained that in HF and halogen fluoride solutions radon fluoride(s) behave in a cationic fashion.34
WF6
ReF6
1.9 17.5 182.3 101.8 127
18.5 33.7 182.4 100.2 147 UF6
69.2 55.7 (subl.) 199.6 115.5 142
192 IrF6
PtF6
33.2 44.4 45.9 53.0 182.5 183.5 99.1 99.6 205 206 NpF6
61.3 69.1 184.8 98.6 211 PuF6
OsF6
54.450 55.250 198.1 115.8 164
50.850 62.250 197.1 118.0 173
orthorhombic (space group Pnma, Z = 4) and ordered.38,39 The phase transitions of the transition metal hexafluorides all occur around 0 °C.38 (The main group hexafluorides have a similar transition at much lower temperatures. TeF6 crystallizes also in Pnma, whereas the crystal symmetries of SeF6 and SF6 are not conclusively determined.40,41 The actinide hexafluorides are ordered in the orthorhombic phase already at room temperature.) Within the precision of the X-ray data39 and the electron diffraction data for the gases,42 the molecules are octahedral. Of course, small deviations, especially dynamic deviations, cannot be detected by these methods. The metal−fluorine bond lengths of MoF6, TcF6, and RuF6 and WF6, ReF6, and OsF6 are the same within experimental error, while there is a small increase for RhF6, and more significantly for IrF6 and PtF6, which can be considered a consequence of the increasing delectron population from d0 to d4. However, at the same time, the volume per molecule in the solid decreases, which is another indication of increasing F···F intermolecular forces; see Table 1. This means lesser partial charge going from left to right in each row, which in turn is a consequence of higher electron affinity of the central metal atom going from left to right. The charge transfer bands that cause the color of the species to change from colorless to dark red point to the same fact.
3. TRANSITION METAL HEXAFLUORIDES Nine transition metal hexafluorides exist that are all wellcharacterized. With few exceptions they are best prepared by direct fluorination of the metals. Two of them, MoF6 and WF6, are commercially available. In Table 1 some physical data are summarized that show the surprising similarity of these compounds. All of them have a narrow liquid range at normal pressure, so they tend to sublimate easily, as is usually the case for spherically shaped molecules. Surprisingly the melting points and boiling points of the second-row hexafluorides are higher than those of the third-row hexafluorides. This can only be explained by an increase of intermolecular F···F interactions. Within a row the melting and boiling points also increase. As a result WF6 is counterintuitively the compound with the highest vapor pressure, and RhF6 is the one with the lowest. All crystallize in the same way, namely, in a high-temperature cubic, disordered way and a low-temperature phase that is
3.1. Molybdenum and Tungsten Hexafluoride
Molybdenum and tungsten hexafluoride have been known for more than 80 years and are the only transition metal hexafluorides that are commercially available. Both compounds are made from the elements. Compared to all other transition metal hexafluorides, they are the least reactive. These arguments explain that both physical and chemical properties are better documented than for all other transition metal C
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Table 2. Experimental and Recently Calculated (Italics) Electron Affinities of the Transition Metal Hexafluorides (eV) MoF6 3.8125 3.6127
TcF6 4.1891 4.1139
WF6 130
3.36 3.5131 4.49132 4.9129 5.14128
RuF6
5.8991 5.6439
6.47126 7.5125 ReF6
113
3.16
133
