The dis­cov­ery of the most sta­ble el­e­ments in the pe­ri­odic ta­ble led to the gar­ish lights that de­fine big cities. JOEL F. HOOPER ex­plains.

Cosmos - - Cosmos Science Club - JOEL HOOPER is a se­nior re­search fel­low at Monash Univer­sity, in Mel­bourne, Australia. IMAGES 01 Hugo Es­cor­cia / Getty Images 02 Mi­ragec / Getty Images

SINCE THE END of the 19th cen­tury hu­mans have been in­fat­u­ated with the vivid glow of neon. It has be­come a sign of both glitz and kitsch, lead­ing us to the near­est open mo­tel, light­ing up strips of land­marked boule­vards and sym­bol­is­ing those lost to the plea­sures of the night. It is a gas that cre­ates a sense of mys­tery, and one that also holds an ex­ten­sive his­tory.

To un­der­stand what gives neon light its colour, let’s first con­sider the fam­ily that neon gas be­longs to: the no­ble gases. This group of six el­e­ments – the other five are helium, ar­gon, kryp­ton, xenon and radon – have prop­er­ties that make them uniquely sta­ble, seem­ingly par­tic­i­pat­ing in no chem­i­cal in­ter­ac­tions.

This lack of in­ter­ac­tion with other el­e­ments to form com­pounds is the rea­son these six gases are called ‘no­ble’ – it was con­sid­ered a sign of no­bil­ity not to re­act when pro­voked, and mem­bers of the no­bil­ity were ex­pected to be aloof.

Un­der­stand­ing the sta­bil­ity of these gases led to the dis­cov­ery of the struc­ture of the atom, and con­se­quently to how it could be ma­nip­u­lated to cre­ate the red-or­ange glow we know as neon.

It be­gan in 1902, when Dmitri Men­deleev added the no­ble gases to his pe­ri­odic ta­ble. He ar­ranged the el­e­ments in rows and col­umns ac­cord­ing to their atomic weight, re­al­is­ing that cer­tain el­e­ments ap­peared in re­peat­ing (or pe­ri­odic) pat­terns due to their prop­er­ties. The no­ble gases are sit­u­ated on the far right-hand side of the pe­ri­odic ta­ble.

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 the ex­pla­na­tion. Based on work by physi­cist Ernest Rutherford, who had pro­posed that elec­trons or­bit the nu­cleus of an atom much like plan­ets or­bit a star, Bohr pro­posed that elec­trons did in­deed or­bit around the nu­cleus of the atom, but only at cer­tain dis­tances and nowhere in be­tween. He ti­tled these dis­tances “or­bitals” or “shells”.

Bohr’s model was a hit. It was a key step in the de­vel­op­ment of quan­tum me­chan­ics, and in de­ter­min­ing the sta­bil­ity of the no­ble gases.

Us­ing Bohr’s the­ory, a physi­cist named Gilbert Lewis sug­gested that each elec­tron shell was most sta­ble when it contained eight elec­trons, with the ex­cep­tion of the very first shell, which could only ac­com­mo­date two.

Lewis fur­ther ex­plained that most atoms would go to great lengths to stay in this sta­ble state by bor­row­ing, do­nat­ing or shar­ing elec­trons with oth­ers to make sure their outer shell was com­plete. This is now recog­nised as chem­i­cal bond­ing, which al­lows for the for­ma­tion of com­pounds.

This is why the no­ble gases are so sta­ble and very re­sis­tant to form­ing bonds. They have a com­plete outer shell of eight elec­trons (ex­cept for helium, which has just two), so they have no need to share, bor­row or do­nate.

That isn’t to say they can’t un­dergo chem­i­cal bond­ing: xenon and kryp­ton will re­act with flu­o­rine gas, an ex­tremely pow­er­ful ox­i­dant, to form com­pounds such as xenon di­flu­o­ride (XEF ), and helium can re­act with sodium un­der ex­treme pres­sure to form Na He.

The sta­bil­ity of the no­ble gases leads to 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. Be­cause helium doesn’t burn, and is lighter than air, it is also great for fill­ing bal­loons.

Most im­por­tantly, no­ble gases pro­vide us with our neon lights!

One thing to re­mem­ber, though, is that not all so-called neon lights use neon gas. De­pend­ing on its colour, a neon light might use an­other no­ble gas.

The process to cre­ate the light is sim­i­lar to how other lamps work. A high-volt­age elec­tri­cal dis­charge is passed through a tube of low-pres­sure gas. The elec­tri­cal dis­charge can ex­cite elec­trons in the 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 packet of light, known as a pho­ton.

The wave­length of this light cor­re­sponds to its colour, and will de­pend on the dif­fer­ence in energy be­tween the higher and lower shells. Each el­e­ment gives off a dif­fer­ent char­ac­ter­is­tic wave­length of light, and thus colour, when its elec­trons are ex­cited, as each el­e­ment has a dif­fer­ent num­ber of shells its elec­trons can jump up to.

This is why neon lights glow with an elec­tri­fy­ing red-or­ange colour, while ar­gon lamps are laven­der blue and xenon lamps light blue-green. It all comes down to the move­ment of elec­trons be­tween each of the el­e­ment’s shells in or­der to en­sure its outer shell is com­plete.

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.

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