Group 18 — Noble Gases

Members and Occurrence

Group 18 consists of helium ($\ce{He}$), neon ($\ce{Ne}$), argon ($\ce{Ar}$), krypton ($\ce{Kr}$), xenon ($\ce{Xe}$) and radon ($\ce{Rn}$) — together with the synthetic element oganesson ($\ce{Og}$, $Z=118$), which has been produced only in trace amounts and whose chemistry is mostly predicted. All the naturally occurring members are gases and chemically unreactive, forming very few compounds; hence the family name noble gases.

With the exception of radon, all occur in the atmosphere. Their total abundance in dry air is roughly 1% by volume, of which argon is by far the major constituent. Helium — and sometimes neon — is also found trapped in minerals of radioactive origin such as pitchblende, monazite and cleveite, while the chief commercial source of helium is natural gas. Xenon and radon are the rarest members. Radon is not mined at all; it is generated continuously as a decay product of radium:

$$\ce{^{226}_{88}Ra -> ^{222}_{86}Rn + ^{4}_{2}He}$$

Electronic Configuration

Every noble gas closes a shell. The general valence configuration is $\ce{ns^2 np^6}$ — the stable octet — with the single exception of helium, whose first shell is full at just two electrons, $\ce{1s^2}$. It is precisely this closed-shell architecture that accounts for almost every property of the group, from inertness to the extreme ionisation enthalpies.

Why the Noble Gases Are Inert

The near-total chemical reluctance of Group 18 rests on two linked facts, both flowing from the closed valence shell. First, all noble gases except helium have completely filled $\ce{ns^2 np^6}$ valence orbitals (helium is full at $\ce{1s^2}$), so there is no electronic incentive to bond. Second, this stability shows up quantitatively in two energetic terms.

Down the group the ionisation enthalpy falls steadily — from $2372\ \text{kJ mol}^{-1}$ for helium to about $1037\ \text{kJ mol}^{-1}$ for radon — as the valence electrons sit farther from the nucleus. This declining ionisation enthalpy is exactly why xenon, not the lighter members, is the one that yields compounds.

Physical Properties and Trends

All noble gases are monoatomic, colourless, odourless and tasteless, and are only sparingly soluble in water. Because the only force acting between their atoms is the weak London dispersion force, they have very low melting and boiling points. Atomic radius increases down the group with rising atomic number, and the strengthening dispersion forces in the heavier, more polarisable atoms push the boiling points upward from helium to radon.

Helium holds two records worth memorising. It has the lowest boiling point of any known substance, 4.2 K, and it diffuses readily through ordinary laboratory materials such as rubber, glass and plastics — a property exploited in leak detection and cryogenics.

Two trends move in opposite directions and are easy to confuse: ionisation enthalpy falls down the group while boiling point and atomic radius rise. The same closed-shell stability that gives the largest ionisation enthalpy at the top of each period also explains why these gases were the last family to be coaxed into chemistry at all.

Xenon Fluorides — Preparation

For decades every attempt to make a noble-gas compound failed. The breakthrough came in March 1962, when Neil Bartlett noticed that the first ionisation enthalpy of molecular oxygen ($1175\ \text{kJ mol}^{-1}$) was almost identical to that of xenon ($1170\ \text{kJ mol}^{-1}$). Having already made the red salt $\ce{O2^+[PtF6]^-}$, he reasoned xenon should behave likewise — and by mixing $\ce{PtF6}$ with xenon he obtained a red compound formulated as $\ce{Xe^+[PtF6]^-}$. The floodgates opened, and a series of xenon compounds with the most electronegative elements (fluorine and oxygen) followed.

Xenon forms three binary fluorides by direct combination with fluorine; which one results is dictated by the temperature and the $\ce{Xe}:\ce{F2}$ ratio:

$$\ce{Xe(g) + F2(g) ->[673\ K,\ 1\ bar][Xe\ in\ excess] XeF2(s)}$$

$$\ce{Xe(g) + 2F2(g) ->[873\ K,\ 7\ bar][1:5\ ratio] XeF4(s)}$$

$$\ce{Xe(g) + 3F2(g) ->[573\ K,\ 60{-}70\ bar][1:20\ ratio] XeF6(s)}$$

$\ce{XeF6}$ can also be made by the low-temperature reaction of $\ce{XeF4}$ with dioxygen difluoride:

$$\ce{XeF4 + O2F2 ->[143\ K] XeF6 + O2}$$

All three fluorides are colourless crystalline solids that sublime readily at 298 K and are powerful fluorinating agents. They are so reactive that even traces of water hydrolyse them. They also behave amphoterically toward fluoride: with strong fluoride-ion acceptors (Lewis acids) they yield cationic species, and with fluoride-ion donors $\ce{XeF6}$ yields fluoroanions.

$$\ce{XeF2 + PF5 -> [XeF]^+[PF6]^-} \qquad \ce{XeF4 + SbF5 -> [XeF3]^+[SbF6]^-}$$

$$\ce{XeF6 + MF -> M^+[XeF7]^-}\quad(\ce{M = Na, K, Rb, Cs})$$

VSEPR Structures of Xenon Compounds

The shapes of the xenon fluorides follow directly from VSEPR once the lone pairs on the central xenon are counted. $\ce{XeF2}$ has two bond pairs and three lone pairs (five electron pairs, $sp^3d$); the lone pairs take the equatorial positions, leaving the two fluorines axial — hence a linear molecule. $\ce{XeF4}$ has four bond pairs and two lone pairs (six electron pairs, $sp^3d^2$) with the lone pairs trans to each other, giving a square planar shape.

$\ce{XeF6}$ is the awkward one. It has six bond pairs and one lone pair — seven electron pairs in all. Rather than a regular octahedron, the extra lone pair distorts the geometry, and experimentally in the gas phase $\ce{XeF6}$ is a distorted octahedron.

Hydrolysis and Xenon–Oxygen Compounds

Because the fluorides are such strong fluorinating agents, their reactions with water are central to the xenon–oxygen chemistry. $\ce{XeF2}$ hydrolyses slowly to give the free element, hydrogen fluoride and oxygen — a clean way to recover xenon:

$$\ce{2XeF2(s) + 2H2O(l) -> 2Xe(g) + 4HF(aq) + O2(g)}$$

$\ce{XeF4}$ and $\ce{XeF6}$ react with water far more vigorously. Complete hydrolysis builds the colourless, explosive solid xenon trioxide, $\ce{XeO3}$:

$$\ce{6XeF4 + 12H2O -> 4Xe + 2XeO3 + 24HF + 3O2}$$

$$\ce{XeF6 + 3H2O -> XeO3 + 6HF}$$

Crucially, the partial hydrolysis of $\ce{XeF6}$ stops short of the trioxide and yields oxyfluorides — the volatile liquid $\ce{XeOF4}$ and $\ce{XeO2F2}$:

$$\ce{XeF6 + H2O -> XeOF4 + 2HF} \qquad \ce{XeF6 + 2H2O -> XeO2F2 + 4HF}$$

$\ce{XeO3}$ has a pyramidal molecular structure (three bond pairs, one lone pair — analogous to $\ce{NH3}$), while $\ce{XeOF4}$ is a colourless volatile liquid with a square pyramidal shape. These two shapes, alongside the linear/square-planar/distorted-octahedral fluorides, are the exact set NEET likes to test in match-the-column form.

Uses of the Noble Gases

Although the heavy members are chemically interesting, the light noble gases are the ones with broad practical value. Their applications follow straight from their physical properties — inertness, low density, low blood solubility and characteristic discharge colours.