18.12 Occurrence, Preparation, and Properties of the Noble Gases

Occurrence, Isolation, and Applications of Noble Gases

Noble gases occupy Group 18 of the periodic table and stand out because of their near-total chemical inertness. Understanding why they behave this way, and the few exceptions to that rule, ties together several periodic trends you've already studied.

Noble Gases in Everyday Life

Six elements make up the noble gases: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). All of them exist in Earth's atmosphere, though most are present only in trace amounts:

• Argon is by far the most abundant at roughly 1% of the atmosphere.
• Neon sits at about 0.0018%, helium at ~0.0005%, krypton at ~0.00011%, and xenon at a tiny ~0.000009%.

Helium is also the second most abundant element in the universe (after hydrogen) because it's continuously produced by nuclear fusion inside stars.

Isolation Methods

Noble gases are obtained through two main routes:

1. Fractional distillation of liquid air — Air is cooled until it liquefies, then slowly warmed. Each noble gas boils off at a different temperature, so they can be collected separately.
2. Extraction from natural gas — Helium often co-occurs with methane in underground deposits. Cryogenic distillation cools the natural gas mixture until everything except helium liquefies, leaving helium gas behind.

Applications

Each noble gas fills a practical niche because of its specific physical properties:

• Helium — Fills balloons and airships (low density gives buoyancy). Serves as a coolant in MRI machines and nuclear reactors because of its extremely low boiling point and high thermal conductivity. Also provides an inert atmosphere for welding and semiconductor manufacturing, preventing unwanted oxidation.
• Neon — Used in illuminated signs, where it emits its characteristic orange-red glow when an electric current passes through it.
• Argon — Fills incandescent light bulbs to keep the tungsten filament from oxidizing, which extends the bulb's lifespan. Also used to create inert atmospheres for welding and for preserving sensitive documents like the Declaration of Independence.
• Krypton and xenon — Found in high-intensity discharge lamps and automotive headlights, where they produce bright white light. Xenon is also used in camera flash lamps and certain medical devices (dental curing lights, dermatology equipment).
• Radon — Was once used in radiotherapy for cancer treatment, but its use has largely been abandoned because radon is radioactive and poses serious health risks with prolonged exposure.

Physical Properties and Chemical Reactivity of Noble Gases

Physical Properties

Noble gases have remarkably low boiling and melting points compared to nearly every other element. The reason comes down to intermolecular forces: because noble gas atoms have complete electron shells, the only attractions between them are London dispersion forces, which are the weakest type of intermolecular force.

Two key trends as you move down Group 18 (He → Rn):

• Boiling and melting points increase. Larger atoms have more electrons, which means stronger London dispersion forces. Helium has the lowest boiling point of any element at 4.2 K (−268.9 °C), while radon's boiling point is 211.5 K (−61.7 °C).
• Ease of liquefaction increases. Helium's dispersion forces are so weak that it's the hardest substance to liquefy. Radon, with the strongest dispersion forces in the group, is the easiest noble gas to liquefy, though its forces are still weak compared to most other elements.

Chemical Reactivity

Noble gases are generally unreactive because their outer electron shells are already full (the stable octet, or in helium's case, a filled 1s shell). There's very little energetic incentive for them to form bonds.

That said, reactivity does increase going down the group. Heavier noble gases have larger atomic radii and greater electron shielding, which makes their outermost electrons easier to pull into a bond.

Xenon compounds with fluorine — Xenon reacts with fluorine under high temperature and pressure to form several stable compounds:

1. Xenon difluoride ($XeF_2$) — linear geometry, colorless solid
2. Xenon tetrafluoride ($XeF_4$) — square planar geometry, colorless solid
3. Xenon hexafluoride ($XeF_6$) — distorted octahedral geometry, colorless solid

Xenon compounds with oxygen — Xenon also forms oxide compounds under extreme conditions:

1. Xenon trioxide ($XeO_3$) — pyramidal geometry, colorless solid (and highly explosive)
2. Xenon tetroxide ($XeO_4$) — tetrahedral geometry, colorless solid

Krypton can form krypton difluoride ($KrF_2$), but this compound is unstable and decomposes readily at room temperature.

The lighter noble gases (He, Ne, Ar) have not been observed to form stable compounds. Their higher ionization energies and smaller atomic sizes make it far too difficult to coax their electrons into bonding.

Periodic Trends in Noble Gases

These trends tie together everything above:

• Electron configuration: All noble gases have a full outer shell ($ns^2np^6$, except helium which is $1s^2$). This is the root of their stability.
• Ionization energy: Noble gases have the highest ionization energies in their respective periods, which is why they resist losing electrons and forming bonds.
• Atomic radius: Increases from He to Rn. Larger radius means weaker hold on outer electrons, which explains why Xe and Kr can form compounds while He, Ne, and Ar cannot.

These three trends work together: as atomic radius grows and ionization energy drops going down the group, both physical properties (higher boiling points from stronger dispersion forces) and chemical behavior (greater willingness to form compounds) shift in a predictable way.