xenon fluorides (XeF2, XeF4, XeF6)

For sixty years the textbooks were sure of one thing: noble gases do not form compounds. Then in 1962 Neil Bartlett, having seen oxygen lose an electron to the fierce oxidiser PtF6, reasoned that xenon (with almost the same ionisation energy as O2) might do the same — and made the first real noble-gas compound. The flood that followed centred on the xenon fluorides, the cleanest proof that 'inert' was the wrong word.

Three stand out, each a study in VSEPR. XeF2 is made by simply heating xenon with fluorine; xenon keeps three lone pairs and bonds two fluorines, giving five electron domains and a linear molecule (lone pairs equatorial). XeF4, formed with excess fluorine under pressure, has four bonds and two lone pairs — six domains — and the lone pairs go trans to each other, leaving a flat square-planar molecule. XeF6 has six bonds and one lone pair, seven domains, and is not a regular octahedron but a distorted shape, because the stereochemically active lone pair pushes the fluorines out of symmetry. The bonding is highly polar, electrons drawn toward the very electronegative fluorine, and is best described with three-centre, four-electron molecular-orbital pictures rather than the old idea of xenon 'using its d orbitals'.

These compounds matter as a landmark in chemical understanding — they showed that a full octet is a strong preference, not an absolute prohibition, and they reopened a whole branch of chemistry. Practically, XeF2 is a clean, mild fluorinating agent used in the lab and in semiconductor etching. The honest caveats: only fluorine (and oxygen) are electronegative enough to bond xenon under accessible conditions; the molecules are powerful oxidisers and react with water; and the lighter noble gases do not behave this way — krypton gives only the marginal KrF2, while argon, neon, and helium give essentially nothing stable.