Chemical Properties

Why the Noble Gases Resist Reacting

Noble gases are, as a rule, the least chemically reactive elements in the periodic table. Two structural reasons explain this inertness:

• (i) Every member except helium (which is $1s^2$) carries a completely filled $ns^2np^6$ valence-shell configuration — an especially stable electronic arrangement.
• (ii) They combine very high ionisation enthalpy with strongly positive electron gain enthalpy — meaning it costs a lot of energy either to pull an electron away from them or to force one onto them.

The First Noble-Gas Compound

For decades after their discovery, chemists tried and failed to make the noble gases react with anything. That changed in March 1962, when Neil Bartlett, working at the University of British Columbia, first prepared a red compound formulated as $O_2^+PtF_6^-$. Bartlett then noticed something striking: the first ionisation enthalpy of molecular oxygen (1175 kJ mol⁻¹) is almost identical to that of xenon (1170 kJ mol⁻¹). Reasoning that if $PtF_6$ could pull an electron off $O_2$, it should be able to do the same to Xe, he mixed $PtF_6$ directly with xenon gas — and successfully isolated a second red compound, $Xe^+PtF_6^-$. This single result opened the door: chemists soon went on to synthesise a range of xenon compounds, mostly with the two most electronegative elements, fluorine and oxygen.

Beyond xenon, the noble-gas chemistry that has been built up is much more limited. Only krypton difluoride, $KrF_2$, has been studied in any detail among krypton compounds. Radon compounds such as $RnF_2$ have only ever been identified by radiotracer technique, never actually isolated as bulk samples (because every isotope of radon is radioactive and short-lived). No genuine compounds of argon, neon or helium are known at all.

Xenon–Fluorine Compounds

Xenon forms three binary fluorides — $XeF_2$, $XeF_4$ and $XeF_6$ — each obtained by direct combination of the elements under a specific set of conditions:

$XeF_6$ can also be made a second way, by reacting $XeF_4$ with dioxygen difluoride at 143 K:

All three fluorides are colourless crystalline solids that sublime readily at 298 K. They behave as powerful fluorinating agents, and every one of them is hydrolysed rapidly by even trace amounts of water. $XeF_2$, for instance, hydrolyses to give elemental xenon back, along with HF and oxygen:

Shapes of the xenon fluorides (VSEPR): The structural diagrams referenced here show five xenon-based molecules built up from a central orange Xe atom, with lone pairs drawn as shaded lobes carrying a pair of dots, and bonded atoms shown as smaller spheres joined to Xe by solid (near) or dashed (far) bonds to suggest depth.

• $XeF_2$ has two bonding pairs and three lone pairs, giving a linear shape — the diagram shows the two F atoms directly opposite each other through Xe, with the three lone pairs splayed around the equatorial plane.
• $XeF_4$ has four bonding pairs and two lone pairs, giving a square planar shape — the four F atoms sit at the corners of a square around Xe, with the two lone pairs positioned above and below that plane.
• $XeF_6$ has six bonding pairs and one lone pair (seven electron pairs in total), and as a result adopts a distorted octahedral shape rather than a perfect octahedron — the lone pair pushes into the electron arrangement, an effect confirmed experimentally in the gas phase, with one lone-pair lobe shown pointing up and the six F atoms arranged around Xe in a near-octahedral pattern.

Xenon fluorides also react as Lewis acids or bases towards other fluorides. With a fluoride-ion acceptor, they form a cationic species:

With a fluoride-ion donor (an alkali metal fluoride $MF$, where $M$ = Na, K, Rb or Cs), $XeF_6$ instead forms a fluoroanion:

Xenon–Oxygen Compounds

Hydrolysing $XeF_4$ or $XeF_6$ with water produces the trioxide, $XeO_3$:

If instead $XeF_6$ is hydrolysed only partially, the products are oxyfluorides rather than the fully oxidised trioxide:

$XeO_3$ is a colourless, explosive solid. With three bonding pairs and one lone pair on xenon, it adopts a pyramidal shape — the referenced diagram shows the three green O atoms forming a tripod-like base beneath Xe, with the single lone pair pointing up at the apex opposite them.

$XeOF_4$ is a colourless, volatile liquid. Here xenon carries five bonding pairs (one Xe=O and four Xe–F) plus one lone pair, giving a square pyramidal shape — the diagram shows the O atom at the apex above Xe, the four F atoms forming the square base, and the lone pair pointing down opposite the apical oxygen.

Uses of the Noble Gases …