This is Part 2 of a 4 part series about Scientific Discovery. Read Part 3: The Making of a Miracle Drug.

The noble gases are a family of elements consisting of helium, neon, argon, krypton, xenon, and radon. They have long been considered the prudish snobs of the periodic table, as they seemed to refuse to react with other elements. This is in contrast to carbon, which is widely known to be the whore of the periodic table, hooking up with countless other elements from lithium to nitrogen and sometimes even going so far as to perversely bond with tin, a member of its own periodic family. Enter Neil Bartlett who dared to ask: what if the noble gases aren’t as noble as everyone thinks?

Bartlett spent the better part of March 23, 1962 in his lab setting up an apparatus consisting of two glass chambers connected by glass tubing and separated by a seal. One chamber contained a red gas called platinum hexafluoride and the other contained colorless xenon, one of the noble gases. By 7 p.m. Bartlett had finished setting up the experiment and was alone in the lab. He broke the seal separating the two gases, allowing them to mix. Instantly a light orange powder exploded out of the contact point between the gases, choking the glass tubing. Amazingly, it looked like xenon, formerly known as an ‘inert’ gas, had reacted with platinum hexafluoride. With one experiment Bartlett had overturned almost 75 years of conventional wisdom about the reactivity of the noble gases. At this point he was the only person on Earth who knew that noble gases could react with other compounds. He rushed out of his lab looking for someone to share his discovery with, but the building was completely empty.

But how did Bartlett decide to react xenon with platinum hexafluoride? Earlier he mixed metallic platinum and fluorine gas and accidentally isolated an unknown dark red solid. Rather than tossing the mysterious solid in the trash and moving on, Bartlett decided to figure out what it was. He eventually found that the platinum and fluorine had reacted to give platinum hexafluoride (a known reaction at the time), which had then reacted with oxygen in the air to give the red solid. Based on the chemical structure of the red material, Bartlett deduced that the platinum hexafluoride had oxidized oxygen, meaning it had stripped electrons from molecules of oxygen in the air. This was unusual. Oxygen is a notorious electron hog, and, as its name suggests oxygen usually oxidizes other stuff, other stuff doesn’t oxidize oxygen. For example, think of rust forming – oxygen in the air strips electrons from atoms of iron to make rust. But the opposite had happened here, suggesting that platinum hexafluoride was some sort of super-oxidizer, capable of out-oxidizing oxygen.

Bartlett began to wonder what else he could oxidize with platinum hexafluoride. At around this time he reviewed a table of ionization potentials to use in teaching his undergraduate general chemistry class. Ionization potentials are a measure of how much energy is needed to rip an electron off of a given atom. In other words, it’s a good indication of the energy needed to oxidize something. Bartlett noticed that oxygen and xenon have about the same theoretical ionization potential. At this point you can probably see where this train of thought is leading and so did Bartlett: if he could use platinum hexafluoride to oxidize oxygen, why not use it to oxidize xenon since it takes about the same amount of energy to oxidize both? So despite learning his entire life that noble gases like xenon will refuse to react with anything no matter how hard you try to convince them, he decided to set up a reaction of xenon with platinum hexafluoride.

Structure of Xenon Hexafluoride from J. Chem. Phys. Vol. 102 No. 8, 1995

Read About the Structure of Xenon Hexafluoride in the Journal of Chemical Physics

So there it was – the noble gases aren’t all that noble after all. And once everyone realized this, the flood gates opened and chemists started to make all kinds of different noble gas compounds: xenon difluoride, xenon tetrafluoride, and a variety of xenon oxyfluorides (xenon bound to various numbers of oxygen and fluorine atoms) and xenon oxides (xenon bound to various numbers of oxygen atoms) not to mention radon fluoride and krypton fluoride. It even turned out that xenon difluoride could be made simply by mixing xenon and fluorine gases and exposing the mixture to sunlight. Chemists had tried to react xenon and fluorine 50 years before Bartlett’s discovery but no one had thought to put the mixture in direct sunlight.

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