HOW NOBLE GAS CAN FORMED COMPOUND ?

By :

 Putu Eka Surya Putra (0713031001)

Chemistry Eduacation department, Faculty of Mathematics and Sciences, 

Ganesha University of Education

     Xenon is now known to react with highly electronegative species. For example, xenon will react  with fluorine to form XeF6. In 1962, N. Bartlett showed that Xe will react with PtF6 to form a red solid, which we now know is a xenoplatinum salt but he thought was XePtF6. At about the same time, R. Hoppe at al. carried out the same reaction, and also formed XeF2 by a different reaction. Xe has the ability hybridize its filled s and p orbitals with d orbitals to form hybrid orbitals which can bond to F, but of course it "violates" the Octet Rule. Of course, P and S also regularly violate this rule, so this is no big deal. It's a rule, not a law of nature...

     As to the specific value you cited, I have not been able to dig that up. Pauling's scale was developed before Xe's chemistry was known and so Pauling electronegativity tables do not have an entry for Xe typically (although I guess it is possible that he proposed a value for Xe in the 70s or 80s. That would not be surprising to me at all). Kr is also now known to undergo some reactions.

     There are several electronegativity scales in addition to Pauling's. The two best known are the Mullekin and Allred-Rochow scales,and there is a fourth one which is called the "spectroscopic electronegativity." The Allred-Rochow scale agrees pretty well trend-wise with the Pauling scale, and gives an electronegativity value of 2.4. This same scale also predicts electronegativity values for the other noble gases:

He=5.50, Ne=4.84, Ar=3.20, Kr=2.94, Xe=2.40.

   

     It was initially 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 (except helium, which has no p orbital at all), 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, although subsequent predictions for XeF₈ indicated that it would 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 the lighter ones. Hence, the outermost electrons experience a shielding effect from the inner electrons that makes them more easily ionized, 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.