The bonding in noble gas compounds

The outer electronic configuration of xenon is 5s 5p (7.4) it has no unpaired electrons to pair with an unpaired electron on fluorine. 

One solution is to assume that, when bonding to fluorine, the xenon atom promotes an outer electron to some empty, higher energy orbital. The one usually chosen is one of the five 5d orbitals, which normally become occupied only in the next Period in the Periodic Table. Promotion of, say, a 5s electron gives the configuration shown in 7.5. 

Now two two-electron bonds can be formed from unpaired electrons in the 5 s and 5d orbitals. 

when we apply VSEPR theory to xenon fluorides, we commit ourselves to both Xe—F single bonds, and to using higher energy orbitals, such as 5d, to form those bonds. But this use of higher-energy orbitals is controversial, and alternative treatments have been proposed. 

One alternative uses molecular orbital theory. Suppose we take the F—Xe—F axis in XeF2 to be the z axis. Then we can avoid the use of higher-energy orbitals by forming molecular orbitals from a 2p orbital on each of the fluorines, and the 5p, orbital of xenon. 

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There are currently two approaches to the problem of bonding in noble gas compounds. Neither is completely satisfactory, but between the two they account adequately for the properties of these compounds. The first might be termed a valence bond approach. It would treat the xenon fluorides by means of expanded valence shells through promotion uf electrons lo the SJ orbitals ... 

AU acceptable theories should account for the following facts only the heavier, more readily ionizable noble gases form compounds and only the most electronegative atoms or groups are satisfactory hgands for the noble gases. Two theories of bonding in noble gas compounds are discussed here see also Molecular Orbital Theory and Valence Shell Electron Pair Repulsion Model). 

It is now well known that despite their name the noble gases (at least Rn, Xe, and Kr) do, in fact, participate in interactions normally considered chemical, notably with F but also with other elements, and the Xe-F bond strength is a substantial 30kcal/mole. Noble gas chemistry is accordingly a subject of considerable theoretical interest. Nevertheless, it is extremely unlikely that conditions resulting in the formation of noble gas compounds would be encountered outside the laboratory, so noble gas chemistry will not be important in geochemistry and will not be discussed here. Treatments of noble gas chemistry are presented by Hyman (1963), Classen (1966), Dean (1985), and Pyykko (1997). 

Special consideration will be given to a description of the electronic stmcture and the nature of bonding in noble gas (Ng) containing compounds. Because it is the peculiar electronic structure of the Ng atoms which cause their resistance to forming a chemical bond, we review briefly some of the atomic properties of Ng... 

The molecular structures of the five homoleptic noble gas compounds that are stable enough to be studied in the gas phase are shown in Fig. 19.1. If the noble gas atoms are assumed to form single electron pair bonds to the fluorine atoms, and if the Xe-O bonds in Xe04 are assumed to be double (four electron bonds), then the K and Xe atoms in the trifluorides are surroimded by 10 electrons in the valence shell, the Xe atom in XeP4 by 12, the Xe atom in XeP6 by 14, and in Xe04 by 16 electrons. AU these compounds may thus be described as containing hypervalent noble gas atoms. 

Stable noble gas compounds are restricted to those of xenon. Most of these compounds involve bonds between xenon and the most electronegative elements, fluorine and oxygen. More exotic compounds containing Xe—S, Xe—H, and Xe—C bonds can be formed under carefully controlled conditions, for example in solid matrices at liquid nitrogen temperature. The three Lewis structures below are examples of these compounds in which the xenon atom has a steric munber of 5 and trigonal bipyramidal electron group geometry. 

The method has been confined to main-group compounds presumably because of irregularities expected with unsymmetrical charge distributions in transition metal ions. The noble gas compounds remain outside the scope of the method because of the way in which electronegativity is defined (atom compactness relative to interpolated noble atom compactness). The main weakness of the method when applied to fluorides is in the somewhat arbitrary choice of fluorine bond energies. 

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