My watch list  

Noble gas compound

Noble gas compounds are chemical compounds that include an element from Group 18 of the periodic table, the noble gases.


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][2] These predictions proved quite accurate, except that XeF8 is now predicted to be not only thermodynamically unstable, but kinematically unstable,[3] 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.

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.[2]


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 He and Ne cannot fit because they are too small.

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

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.


Hydrates are formed by compressing the noble gases in water. It is believed that the water molecule, a 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]

True noble gas compounds

In 1962, Neil Bartlett noticed that the highly oxidising compound platinum hexafluoride ionised O2 to O2+. As the ionisation energy of O2 to O2+ (1165 kJ mol–1) is nearly equal to the ionisation energy of Xe to Xe+ (1170 kJ mol–1), he tried the reaction of Xe with PtF6. This yielded a crystalline product xenon hexafluoroplatinate, whose formula was propsed to be Xe+[PtF6]. [2][4] It was later shown[citation needed] that the compound is actually more complex, containing both XeFPtF6 and XeFPt2F11. This was the first real compound of any noble gas.

Later in 1962 the first simple (two-element) noble gas compound (xenon tetraflouride) was synthesized by Howard Claassen by subjecting xenon and fluorine to a high temperature.[5]

In recent years, several compounds of noble gases, particularly xenon, have been prepared. Among these are the xenon fluorides (XeF2, XeF4, XeF6), oxyfluorides (XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4) and oxides (XeO3 and XeO4). Xenon difluoride can be produced by the simple exposure of Xe and F2 gases to sunlight; while the mixing of the two gases 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 RnF2, which glows with a yellow light in the solid state. Krypton is able to react with fluorine to form KrF2, and short-lived excimers of Xe2 and noble gas halides such as XeCl2 are used in excimer lasers. The discovery of ArF2 was announced in 2003 but so far is not confirmed. No conventional compounds of He or Ne are known.

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, N(SO2F)2 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 (308.71 pm).


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.[2] 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.[2]


  1. ^ Linus Pauling (June 1933). "The Formulas of Antimonic Acid and the Antimonates". J. Am. Chem. Soc. 55, (5): 1895 - 1900. doi:10.1021/ja01332a016.
  2. ^ a b c d e Holloway, John H. (1968). Noble-Gas Chemistry. London: Methuen. 
  3. ^ Seppelt, Konrad (June 1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research 12: 211–216.
  4. ^ Bartlett, N. (1962). "Xenon hexafluoroplatinate Xe+[PtF6]". Proceedings of the Chemical Society of London (6): 218. doi:10.1039/PS9620000197.
  5. ^ Claassen, H. H.; Selig, H.; Malm, J. G. (1962). "Xenon Tetrafluoride". J. Am. Chem. Soc. 84 (18): 3593. doi:10.1021/ja00877a042.


This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Noble_gas_compound". A list of authors is available in Wikipedia.
Your browser is not current. Microsoft Internet Explorer 6.0 does not support some functions on Chemie.DE