How noble gases give us neon lights

The discovery of the most stable elements in the Periodic Table led to the garish lights that define big cities. Joel F. Hooper explains.

The reluctance of noble gases to form chemical bonds is the key to making neon lights.
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Discovered around the end of the 19th century, the noble gases are the most stable group of the chemical elements.

The six gases: helium, neon, argon, krypton, xenon and radon have a myriad of modern uses. When they were first discovered, their strangely stand-offish properties were a mystery.

Uniquely stable, they seemed to participate in no chemical reactions. But by understanding the stability of the noble gases, physicists discovered the key to chemical bonding itself. Dmitri Mendeleev added the noble gases to his periodic table in 1902, where he arranged the elements in rows and columns according to their atomic weight.

Mendeleev was able to see repeating (or periodic) patterns in their properties. The noble gases appeared regularly in the periodic table, occurring in every eighth position, at least amongst the lighter elements.

Physicists struggled to find a model that would explain this curious observation. What was the significance of the number eight?

In 1912, a young Danish physicist named Niels Bohr came up with a new explanation that forever changed our understanding of the atom.

Having just completed his PhD studies in Denmark, he moved to England to work with a physicist named Ernest Rutherford, who had just proposed that electrons orbit around the dense nucleus of an atom, much like planets orbiting a star.

Bohr recognised that Rutherford's model was on the right track, but it didn't quite fit with experimental data. He instead proposed that electrons do indeed orbit around the positively charged nucleus of the atom, but that they can only sit at certain distances from the nucleus.

Bohr suggested that electrons could occupy these energy “shells” which surrounded the nucleus of an atom like layers of an onion, but that they could not reside between shells. Bohr’s model was a hit, and a key step in the development of quantum mechanics.

Using the Bohr model, a physicist named Gilbert Lewis quickly came up with an explanation for the incredible stability of the noble gases. According to Lewis, each electron shell around the atom is most stable when it contains eight electrons, with the exception of the very first shell, which can only accommodate two.

Atoms will go to great lengths to make sure their outermost shell contains an optimal eight electrons. When these shells are incomplete, atoms will try to fill up their outer shell by borrowing, donating or sharing electrons with others. This gives rise to chemical bonding, the combination of elements to form compounds.

Because the noble gases have a complete outer shell of eight electrons (except for helium, which has just two), they are very resistant to forming bonds.

Chemical bonding is possible: xenon and krypton will react with fluorine gas, an extremely powerful oxidant, to form compounds such as xenon difluoride (XeF2). Helium and neon were thought to have no stable compounds, until a team of scientists led by Artem Oganov from Stony Brook University in New York reported in 2017 that helium could react with sodium under extreme pressure to form Na2He.

Neon, nestled in the periodic table.
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But it’s the stability of the noble gases which drives their practical uses. Argon is used as a shielding gas in welding, while helium is used as a cryogenic coolant in MRI machines and superconductors. And because helium doesn’t burn, it can be used to give buoyancy to blimps, without risking another Hindenburg disaster.

Our most common encounter with the noble gases, however, is probably in gas discharge lamps, such as neon lights.

In these lamps, a high voltage electrical discharge is passed through a tube of low-pressure neon gas. This electrical discharge can excite electrons in the neon atoms, causing them to jump from a low-lying and stable shell to a higher shell.

This, of course, breaks Lewis’ rule of eight electrons in the outer shell, so the excited electron will eventually “relax” back to its preferred shell. As this electron relaxes back to the lower shell, it sheds some energy, in the form of a package, or photon, of light.

The wavelength of this light, corresponding to its colour, will depend on the difference in energy between the higher and lower shells. This means that each element gives off a few characteristic wavelengths of light when it’s electrons are excited, due to the different shells its electrons can jump up to.

This is why neon lights give an electrifying red-orange colour, while argon lamps are lavender blue and xenon lamps can be blue-green. It all comes down to how the excited electrons in each atom are able to find their way back home, and complete their outer shell.

So we can trace both the stability of the noble gases, as well as the bright and lurid colours of the neon light back to the same quirk in quantum mechanics: Gilbert Lewis’ rule of eight.

Joel Hooper is a senior research fellow at Monash University, in Melbourne, Australia.