Noble gas compound
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Noble gas compounds are chemical compounds that include an element from column 18 of the periodic table, the noble gases.
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[edit] History and background
Until the 20th century it was believed that the noble gases could not form compounds due to their full valence shell of electrons that rendered them very chemically stable and unreactive.
All noble gases have full s and p outer electron shells (i.e. 8 outer shell electrons, except helium, which has 2, but is nonetheless stable), and so do not form chemical compounds easily. Because of their high ionization energy and almost zero electron affinity, they were not expected to be reactive at all.
In 1933, however, Linus Pauling predicted that the heavier noble gases would be able to form compounds with fluorine and oxygen. Specifically, he predicted the existence of krypton hexafluoride and xenon hexafluoride (XeF6), speculated that XeF8 might exist as an unstable compound, and suggested that xenic acid would form perxenate salts.[1] These predictions proved quite accurate, except that XeF8 is now predicted to be not only thermodynamically unstable, but kinematically unstable,[2] and as of 2006 has not been made.
The heavier noble gases have more electron shells than those near the top. Hence, the outermost electrons experience a shielding effect from the inner electrons that makes it easier to ionize them since they are less strongly attracted to the positively-charged nucleus. This results in an ionization energy low enough to form stable compounds with the most electronegative elements, fluorine and oxygen.
[edit] Pre-1962 compounds
Prior to 1962, the only isolated compounds of noble gases were clathrates (including clathrate hydrates). Other compounds such as coordination compounds were observed only by spectroscopic means.[1]
Hydrates are formed by compressing the noble gases in water. It is believed that the water molecule (strong dipole) induces a weak dipole in the noble gas atoms, resulting in dipole-dipole interaction. Heavier atoms are more influenced than smaller ones, hence Xe·6H2O is the most stable hydrate. The existence of these compounds has, however, been disputed in recent years.[citation needed]
Clathrates (also known as cage compounds) are compounds of noble gases in which they are trapped within cavities of crystal lattices of certain organic and inorganic substances. The essential condition for their formation is that the guest (noble gas) atoms should be of appropriate size to fit in the cavities of the host crystal lattice. For instance, Ar, Kr and Xe can form clathrates with β-quinol, but not He or Ne.
Clathrates have been used for separation of He and Ne from Ar, Kr and Xe, and also for the transportation of Ar, Kr and Xe. In addition,85Kr clathrate provides a safe source of beta particles, while 133Xe clathrate provides a useful source of gamma rays.
Coordination compounds such as Ar·BF3 were postulated to exist at low temperatures, but have never been confirmed. Also, compounds such as WHe2 and HgHe2 were reported to have been formed by electron bombardment, but recent research has shown that He is probably adsorbed on the surface of the metal, hence these compounds cannot be called true chemical compounds.
[edit] True noble gas compounds
In 1962, Neil Bartlett noticed that the highly oxidising compound platinum hexafluoride ionised O2 to O2+ (dioxygenyl). As the ionisation energy of this, at 1165 kJmol-1 is close to the ionisation energy of 1170 kJmol-1 required for Xe to Xe+ transition, (And that of radon is lower) he tried the reaction of Xe with PtF6, which yielded what was at the time believed to be xenon hexafluoroplatinate, a crystalline solid with the formula Xe+[PtF6]-, later shown to be more complex, containing both XeFPtF6 and XeFPt2F11This was the first real compound of any noble gas.[1]
In recent years, several compounds of noble gases, particularly xenon, have been prepared. Among these are the xenon fluorides (xenon difluoride (XeF2), xenon tetrafluoride (XeF4), xenon hexafluoride (XeF6), oxyfluorides(XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4) and oxides (xenon trioxide (XeO3), xenon tetroxide (XeO4)). Xenon difluoride can be produced by the simple exposure of Xe and F2 gasses to sunlight; while the mixing of the two gasses had been tried over 50 years before in an attempt to produce a reaction, nobody had thought to simply expose the mixture to sunlight.
Radon has reacted with fluorine to form radon fluoride (RnF2), which glows with a yellow light in the solid state. Krypton is able to react with fluorine to form krypton fluoride KrF2, and short-lived excimers of Xe2 and noble gas halides such as xenon chloride (XeCl2) are used in excimer lasers. The discovery of argon fluoride (ArF2) was announced in 2003. In theory Neon fluoride may be preparable, but this looks doubtful.
Recently xenon has been shown to produce a wide variety of compounds of the type XeOxY2 Where x is 1,2 or 3 and Y is any electronegative group, such as CF3 or OTeF5. The range of compounds is impressive, running into the thousands and involving bonds between xenon and oxygen, nitrogen, carbon, and even gold, as well as perxenic acid, several halides, and complex ions; a range of compounds seen in the neighbouring element Iodine. The compound Xe2Sb2F11 contains a Xe-Xe bond, the longest element-element bond known. (308.71pm)
[edit] Applications
Most applications of noble gas compounds are either as oxidising agents or as a means to store noble gases in a dense form. Xenic acid is a valuable oxidising agent because it has no potential for introducing impurities: the xenon is simply liberated as a gas. It is rivalled only by ozone in this respect.[1] The perxenates are even more powerful oxidising agents, and the xenon fluorides are good fluorinating agents.
Radioactive isotopes of krypton and xenon are difficult to store and dispose, and compounds of these elements may be more easily handled than the gaseous forms.[1]
[edit] References
- ^ a b c d e Holloway, John H. (1968). Noble-Gas Chemistry. London: Methuen.
- ^ Seppelt, Konrad (June 1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research 12: 211–216.
