Noble Gas Compounds

Core Idea

The heavier noble gases, above all xenon, form stable compounds with the most electronegative elements, fluorine and oxygen. Neil Bartlett showed this in 1962 by reacting xenon with platinum hexafluoride. The xenon fluorides XeF₂, XeF₄ and XeF₆ have shapes that VSEPR predicts well, and their bonding is usually described with three-center four-electron bonds rather than d-orbital use. Reactivity rises down group 18 as ionization energy falls.

How It's Best Learned

Retrace Bartlett's argument first: compare the ionization energy of O₂ with that of xenon and predict the reaction with PtF₆. Then draw Lewis structures and VSEPR shapes for XeF₂, XeF₄, XeF₆ and XeO₃. Finish with the three-center four-electron bond for linear F–Xe–F so the "expanded octet" picture is replaced by a better one.

Common Misconceptions

• Noble gases cannot form compounds because their shells are full. (Xenon and krypton form well-characterized fluorides, and xenon also forms oxides and oxofluorides.)
• Bonding in XeF₂ requires xenon to use its 5d orbitals. (Calculations show d-orbital participation is small; a three-center four-electron bond built from p orbitals explains the bonding.)
• All noble gases are equally unreactive. (Reactivity increases down the group as ionization energy falls; helium and neon form no stable neutral compounds under ordinary conditions.)

Explainer

For decades, chemists treated the noble gases as inert. Their filled valence shells and high ionization energies seemed to rule out bonding. In 1962, Neil Bartlett changed that view. He had found that platinum hexafluoride, PtF₆, is a strong enough oxidizing agent to remove an electron from O₂, forming O₂⁺[PtF₆]⁻. He then noticed that the first ionization energy of xenon is almost equal to that of O₂. When he mixed xenon with PtF₆, an orange-yellow solid formed. Its exact composition proved more complex than first thought, but it was the first noble gas compound.

Soon after, chemists found that xenon reacts directly with fluorine. By varying the ratio of reactants, the temperature and the pressure, they obtained XeF₂, XeF₄ and XeF₆. Their shapes follow VSEPR rules. XeF₂ has three lone pairs and two bonding pairs, giving a linear molecule. XeF₄ has two lone pairs and four bonding pairs, giving a square planar molecule. XeF₆ has one lone pair and six bonding pairs, and its structure is a distorted octahedron. Hydrolysis of XeF₆ gives xenon trioxide, XeO₃, a dangerously explosive solid.

How should we describe the bonding? An older picture used xenon 5d orbitals to "expand the octet." Modern calculations show these orbitals contribute little. A better model is the three-center four-electron bond. In F–Xe–F, one xenon 5p orbital combines with one p orbital on each fluorine to make three molecular orbitals: bonding, nonbonding and antibonding. Four electrons fill the bonding and nonbonding orbitals. The nonbonding orbital sits mainly on the fluorines, so the fluorines carry partial negative charge and xenon carries partial positive charge. This explains why only very electronegative partners form stable bonds to xenon.

The trend across group 18 follows ionization energy. Radon is the most easily ionized, but its intense radioactivity makes its chemistry hard to study. Krypton forms KrF₂ and a few related species. Argon forms no stable compounds under ordinary conditions, although HArF has been detected at very low temperatures in a solid matrix. Helium and neon form no stable neutral compounds under ordinary conditions.

Noble gas chemistry is a useful test of bonding models. The same three-center four-electron idea also helps describe other hypervalent species, such as the triiodide ion, I₃⁻.