Noble-gas compounds

for the noble gases the nominal valencies of n, iv, vi, and viii are possible. 

Xe has twelve outer electrons in XeF4, eight associated with Xe-F bonds in a square-planar arrangement and four as lone pairs in axial positions Xe-F = 1.94 0.01 A.  

This section reports recent studies of molecules which do not fit neatly into any of the previous categories. All the molecules concerned are inorganic ones, and most of the elements concerned are non-metallic ones from the top of Groups V—Vn, plus hydrogen. 

As with F3NO (above) the tetrahedron-edge distances are almost equal in these cases steric effects must be of considerable importance in determining the geometry. 

The other noble-gas elements form compounds much less readily than xenon. For many years, only one binary krypton compound, Krp2, was known with certainty, and 

Analyze We must predict the geometrical stracture given only the molecular formula. 

Plan We must first write the Lewis stmcture for the molecule. We then count the number of electron pairs (domains) around the Xe atom and use that number and the number of bonds to predict the geometry. 

Comment The experimentally determined structure agrees with this prediction. 

Compounds containing the XeFs ion have been characterized. Describe the electron-domain geometry and molecular geometry of this ion. 

Solve There are 36 valence-shell electrons (8 from xenon and 7 fiom each fluorine). If we make four single Xe—F bonds, each fluorine has its octet satisfied. Xe then has 12 electrons in its valence shell, so we expect an octahedral arrangement of six electron pairs. Two of these are nonbonded pairs. Because nonbonded pairs require more volume than bonded pairs (Section 9.2), it is reasonable to expect these nonbonded pairs to be opposite each other. The expected stmcture is square planar, as shown in FIGURE 22.7. 

Describe the electron-domain geometry and molecular geometry of Xep2. 

Bartlett then investigated the crystal structure. It proved to be very similar to that of KSbFs, an ionic compound written K [SbF6] . This suggested that PtF602 also contained a complex ion of the type MFg , and therefore that the positive ion was Reaction 7.5 now becomes  

---

Stable noble-gas compounds have no industrial uses as of this writing but are frequently utilized in laboratories as fluorinating and oxidizing agents. Xenon difluoride and xenon tetrafluoride are relatively mild oxidative fluorinating agents and have been used for the preparation of phosphoms, sulfur, tellurium. 

D. T. Hawkins, W. E. Falconer and N. Bartlett, Noble Gas Compounds, A Bibliography 1962-1976. Plenum Press, New York, 1978. 

Until about 40 years ago, these elements were referred to as "inert gases" they were believed to be entirely unreactive toward other substances. In 1962 Neil Bartlett, a 29-year-old chemist at the University of British Columbia, shook up the world of chemistry by preparing the first noble-gas compound. In the course of his research on platinum-fluorine compounds, he isolated a reddish solid that he showed to be 02+(PtFB-). Bartlett realized that the ionization energy of Xe (1170 kJ/mol) is virtually identical to that of the 02 molecule (1165 kJ/mol). This encouraged him to attempt to make the analogous compound XePtF6. His success opened up a new era in noble-gas chemistry. 

Valence shell electron pair repulsion theory, 1,32-39 effective bond length ratios, 1.34 halogenium species, 3, 312 noble gas compounds, 3,312 repulsion energy coefficient, 1, 33 Valency... 

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. 

Since the discovery of the first noble gas compound, Xe PtF (Bartlett, 1962), a number of compounds of krypton, xenon, and radon have been prepared. Xenon has been shown to have a very rich chemistry, encompassing simple fluorides, XeF2> XeF, and XeF oxides, XeO and XeO oxyf luorides, XeOF2> XeOF, and Xe02 2 perxenates perchlorates fluorosulfates and many adducts with Lewis acids and bases (Bartlett and Sladky, 1973). Krypton compounds are less stable than xenon compounds, hence only about a dozen have been prepared KrF and derivatives of KrF2> such as KrF+SbF, KrF+VF, and KrF+Ta2F11. The chemistry of radon has been studied by radioactive tracer methods, since there are no stable isotopes of this element, and it has been deduced that radon also forms a difluoride and several complex salts. In this paper, some of the methods of preparation and properties of radon compounds are described. For further information concerning the chemistry, the reader is referred to a recent review (Stein, 1983). 

Fields, P. R., Stein, L., and Zirin, M. H., Radon Fluoride Further Tracer Experiments with Radon, in Noble-Gas Compounds (H. H. Hyman, ed), pp. 113-119, University of Chicago Press, Chicago, IL (1963). 

Stein, L., Noble Gas Compounds New Methods for Collecting Radon and Xenon, Chemistry 47, No. 9, 15-20 (1974). 

Claassen, H. H. (1966). The Noble Gases. D. C. Heath, Boston. A very useful introduction to the chemistry of noble gas compounds and structure determination. 

The result here is quite satisfactory because XeF4 does in fact exhibit square planar geometry. It is worth noting, however, that a square planar shape for XeF4 is also predicted by VSEPR theory. Despite the fact that the molecular orbital method has made some inroads as of late, VSEPR is still the best approach available for rationalizing the molecular geometries of noble gas compounds. 

H. H. Hyman (ed.), Noble-Gas Compounds (Chicago, IL, University of Chicago Press, 1963). 

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. 

O O In the early 1960s, Neil Bartlett, at the University of British Columbia, was the first person to synthesize compounds of the noble gas xenon. A number of noble gas compounds (such as XeF2, XeF4, XeFe, and XeOs) have since been synthesized. Consider the reaction of xenon difluoride with fluorine gas to produce xenon tetrafluoride. 

Noble Gas Compounds, Univ of Chicago Press (1963) 5) Anon, FLACS (Florida Chem Soc)... 

Accordingly, he mixed xenon and PtFe and obtained the orange-yellow solid of xenon hexafluoroplatinate—the first noble gas compound. (Although in fact, the compound turned out not to have the structure that Bartlett predicted because at room temperature the XePtFe reacts with another molecule of PtFe to give a product containing [XeF]"[PtFe]- and [PtFslF)... 

The polyhalide ions may conveniently be classified into two groups (Xf-type ions belong to both groups) (I) those I hat are isoelectron ic with noble gas compounds and... 

---