Xenon Compounds: Fluorides, Oxides and Oxyfluoride

For a very long time the noble gases were believed to be entirely, categorically unreactive --

hence the older name "inert gases" -- until 1962, when Neil Bartlett showed that xenon could

indeed be made to react, forming a compound with platinum hexafluoride. Of all the noble gases,

xenon is the one that forms a genuine, well-characterised chemistry of compounds, for two

structural reasons: it is comparatively large (so its outer electrons are held less tightly

than in the smaller noble gases above it) and it has a comparatively low first ionisation enthalpy for a noble gas (still very high compared with most elements, but the lowest among

the noble gases with any practical, non-radioactive chemistry available for study) -- low enough

that a sufficiently electronegative, aggressively oxidising partner, above all fluorine, can pull

electron density away from xenon and force it into genuine covalent bonds.

Xenon's fluorides are all prepared by direct combination of xenon and fluorine gas, sealed

together (usually in a nickel vessel, since nickel forms its own passivating fluoride layer that

resists further attack) and heated or exposed to sunlight, with the product formed depending on

the Xe:F$_2$ ratio and the reaction conditions used. Xenon difluoride, $\text{XeF}_2$, forms

from a $1:5$ Xe:F$_2$ ratio under exposure to sunlight; it is linear (bond angle

$180^\circ$), with xenon $sp^3d$ hybridised: of the five electron domains around xenon in the

underlying trigonal bipyramid, the three lone pairs occupy the equatorial positions

(minimising their mutual $90^\circ$ repulsions) and the two fluorine atoms are pushed into the

two axial positions, directly opposite one another. Xenon tetrafluoride, $\text{XeF}_4$,

forms from the same $1:5$ ratio but at higher temperature and pressure ($673\ \text{K}$,

$6\text{-}7$ bar); it is square planar, with xenon $sp^3d^2$ hybridised: of the six electron

domains in the underlying octahedron, the two lone pairs occupy axial positions (opposite

each other, again minimising mutual repulsion), leaving the four fluorine atoms arranged in a

perfect square in the equatorial plane. Xenon hexafluoride, $\text{XeF}_6$, forms from a

Xe-rich $1:20$ ratio under still more forcing conditions ($573\ \text{K}$, $60\text{-}70$ bar); it

is distorted octahedral, with xenon $sp^3d^3$ hybridised: here there are seven electron

domains (six fluorine atoms plus one lone pair) that cannot be accommodated in any regular,

symmetric polyhedron, so the single remaining lone pair sits over one triangular face of what

would otherwise be a perfect octahedron, pushing the six Xe-F bonds measurably away from their

ideal $90^\circ$ positions and giving the molecule its characteristic, non-rigid distorted shape.

Controlled hydrolysis of the fluorides gives access to xenon's oxygen compounds. Careful,

complete hydrolysis of $\text{XeF}_6$ with water gives xenon trioxide, $\text{XeO}_3$:

$\text{XeO}_3$ is a dangerously explosive solid whose molecule is pyramidal, xenon $sp^3$

hybridised with one lone pair, directly analogous in shape to $\text{NH}_3$/$\text{PH}_3$ but

with three Xe=O (or resonance-delocalised Xe-O) bonds in place of three Xe-H bonds. Partial,

low-temperature hydrolysis of $\text{XeF}_4$ with a controlled, limited amount of water gives

xenon oxydifluoride, $\text{XeOF}_2$:

$\text{XeOF}_2$ is, like $\text{ClF}_3$/$\text{BrF}_3$, T-shaped: xenon is $sp^3d$

hybridised, with five electron domains around it (one Xe=O, two Xe-F, and two lone pairs …