Xenon tetrafluoride bonding

Chlorine pentafluoride and xenon tetrafluoride appear in Figure 9-26. Each has an inner atom with a steric number of 6, but their electron group arrangements include lone pairs. As a result, CIF5 has a square pyramidal shape, whereas XeF4 has a square planar shape. Pictures can help us determine whether or not the bond polarities cancel ... 

Describe the bonding of chlorine trifluoride and xenon tetrafluoride. 

This 3c4e approach may be used to understand the bonding and shape of xenon tetrafluoride,... 

Possible interactions between xenon difluoride and xenon tetrafluoride molecules as a result of bond dipoles. The magnitude of x would be larger than that ofy, because the Xe atom is polarized in two directions (by two fluorine atoms) in the difluoride and in four directions (by four fluorine atoms) in the tetrafluoride. 

Electron Distribution in the Xenon Fluorides and Xenon Oxide Tetrafluoride by ESCA and Evidence for Orbital Independence in the Xenon-Fluorine Bonding... 

T. X. Carroll, R. W. Shaw, T. D. Thomas, C. Kindle and N. Bartlett, Electron Distribution in the Xenon Fluorides and Xenon Oxide Tetrafluoride by ESCA and Evidence for Orbital Independence in the Xenon-Fluorine Bonding, J. Am. Chem. Soc. 96 (1974) 1989-1996. 

When a molecule has four bonded atoms and two lone pairs, however, the lone pairs always lie at opposite vertices to avoid the stronger 90° lone pair-lone pair repulsions. This positioning gives the square planar shape (AX4E2), as in xenon tetrafluoride (XeF4) ... 

This would mean that there are 10 electrons in the valence shell of the Ng atom in xenon difluoride or krypton difluoride and 12 or 14 electrons for xenon tetrafluoride or hexafluoride, and even more for the octafluoroxenate ion, [XeFs] . Since one s and three p orbitals can accommodate only eight electrons, this would require the participation of d orbitals. In fact, the currently favored model uses only s and p atomic orbitals [16]. For example, XeF2 can be constmcted with a three-center-two-electron (3c-2e) bond, like NF5 (Chapter 4, Fig. 4.5) without using d orbitals (Fig. 5.1). Perhaps one should not worry much about which orbitals are involved, because as has been pointed out, bonding is not an observable quantity only bonding distancies and electron density are amenable to observation although... 

Draw a Lewis structure of the noble gas compound xenon tetrafluoride (Xep4) in which all F atoms are bonded to the central Xe atom. 

On the basis of simple electron pair repulsion theory the difluoride should be linear, with the fluorine atoms at the poles of a trigonal bipyramidal arrangement of electron pairs. The tetrafluoride should be a square planar arrangement of fluorine atoms about the xenon atom, there being a non-bonding electron pair both above and below the plane. The hexafluoride, since it possesses a nonbonding electron pair, should show some deviation from octahedral symmetry. The structural evidence has so far been in accord with all of these expectations. Table 10 contains some of the structural data obtained from X-ray and neutron diffraction studies of the di and tetra-... 

In view of the instability of the trioxide it is interesting to speculate on the stability of the oxyfluorides. Xenon oxide tetrafluoride should be an exothermic compound since the atomic heat of formation using average bond energies of 32 kcal. and 17 kcal. respectively for Xe-F and Xe-O is —160 kcal. mole-. The combined heat of atomization of four fluorine atoms (4 x 18.3 kcal.) and one oxygen atom (59.2 kcals.), 132.4 kcal., is less than this. Xenon dioxide difluoride, by such considerations, should be endothermic. It is of interest that xenon oxide tetrafluoride is easily made by controlled hydrolysis of xenon hexafluoride and is a thermally stable entity (m.p. —28) whereas xenon dioxide difluoride has been observed only mass spectrometrically < >. 

Inert xenon actually forms many compounds, especially with highly electronegative fluorine. The AFf° values for xenon difluoride, tetrafluoride, and hexafluoride are —105, —284, and —402 kJ/mol, respectively. Find the average bond energy of the Xe—F bonds in each fluoride. 

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