9.6: Discovery of Noble Gas Compounds

In 1962 at the University of British Columbia, Neil Bartlett was working with the powerful oxidizer PtF₆ and, because of an accidental leak in his vacuum line, noticed the compound’s reaction with O₂ to generate a solid with formula "PtF₆O₂." The formula suggested Pt in the +10 oxidation state, which was clearly unreasonable because PtF₆ was known to be a more powerful oxidizer than either molecular fluorine (F₂) or molecular oxygen (O₂). Bartlett noticed that the X-ray powder diffraction pattern of the compound was similar to that of Cs⁺AsF₆^-, a salt with the CsCl structure in which octahedral AsF₆^- ions occupy the chloride ion sites. This led Bartlett to propose a formulation of O₂⁺PtF₆^- for his new compound.^[6] Magnetic susceptibility data subsequently confirmed the presence of the paramagnetic O₂⁺ cation, which (see Chapter 2) has a bond order of 2.5. This formulation implies that PtF₆ was a strong enough oxidizing agent to oxidize molecular oxygen.

But just how strong an oxidizer is PtF₆? Its electron affinity could be estimated by using a Born-Haber cycle, filling in the lattice energy of O₂⁺PtF₆^- by means of Kapustinskii's formula:

The electron affinity (EA) for PtF₆ can be calculated as EA = -159 - 1167 + 571 = -751 kJ/mol. To put it in perspective, this is 417 kJ/mol more exothermic than the electron affinity of atomic fluorine (334 kJ). PtF₆ was by far the strongest oxidizer that had ever been made!

Crystals of xenon difluoride (XeF₂), which is made by photolysis of a gas mixture of Xe and F₂. XeF₂ is used as a selective etchant for SiO₂ in certain microelectronic fabrication processes. Prior to Neil Bartlett's discovery of the first xenon compound in 1962, it was believed that elements in group VIII (He, Ne, Ar, Kr, Xe) could not form chemical compounds. Since then chemists have discovered a rich family of noble gas compounds containing Xe, Kr, or Ar.

Bartlett recognized that Xe has ionization energy of +1170 kJ, which is very close to the ionization energy of O₂. Since Xe⁺ should be about the same size as O₂⁺, the lattice energy should be about the same with Xe⁺ in the cation site of the O₂⁺PtF₆^- structure. Since all of the other terms in the Born-Haber cycle for the reaction of Xe with PtF₆ are the same, Bartlett concluded that Xe⁺PtF₆^-, like O₂⁺PtF₆^-, should be a stable compound. He purchased a lecture bottle of xenon gas and reacted the two compounds, producing an orange solid.^[7] While the product initially formed in the reaction may in fact be Xe⁺PtF₆^-, the Xe⁺ free radical is a powerful Lewis acid and reacts further with excess PtF₆. The ultimate product of the reaction is formulated [XeF⁺][Pt₂F₁₁^-], a salt which contains Xe in the +2 oxidation state and Pt in the +5 oxidation state. This was an important discovery because it shattered the dogmatic notion, which derived from the octet rule, that elements in group VIII could not form bonds with other elements. The name of this group was changed from the "inert gases" to the "noble gases." Subsequently, many compounds of Xe and a few of Kr and even Ar (which is much harder to oxidize) were synthesized and characterized.