It was the evening of 23rd March 1962. Neil Bartlett was in his lab at the University of British Colombia in Vancouver. Everyone else had gone home, including his student. It had taken longer than expected to set up the apparatus for his experiment but now he was ready. A glass flask was filled with a red gas. It was platinum hexafluoride (PtF6), a dangerous, reactive substance. A sealed tube connected it to another container of the colourless inert gas, xenon. Bartlett broke the seal between the two gases. The gases mixed and immediately a yellow solid formed. It was late when Bartlett got home to his irate wife but he was elated. He had shown that the inert noble gas, xenon, was not so inert or noble after all.
Neil Bartlett’s experiment was the first to show that some of the inert gases, variously labeled Group 0, 8 or 18 depending on the Periodic Table you look at, could react with other elements. In 1962 most chemists were of the opinion that the inert gases could not combine with other elements because of their electron arrangements.
The story goes back to the 1890s when William Ramsay and Morris Travers isolated helium, neon, argon, krypton and xenon from air. The elements were all gases that apparently did not react with anything and so they were given a column in the Periodic Table all to themselves. In 1913 Nils Bohr proposed his theory of the structure of the atom with electrons arranged in shells around the nucleus. He worked out that only a particular number of electrons could fit in each shell – 2 in the first, 8 in the second, 18 in the third and so on. The theory explained the arrangement of elements in the Periodic table. Except for helium all the inert gases have 8 electrons in their outer shells.
Three years later Gilbert Lewis in the USA and Walter Kossel in Germany made similar suggestions. They thought that when atoms combine they either transfer electrons to form ions or share electrons to form covalent bonds so that they can get the same electron arrangement as the inert gases. The “octet rule” as suggested by Kossel stated that elements and compounds are stable when their atoms have 8 electrons in the outer shells.
This idea worked very well and explained the structure and properties of compounds such as sodium chloride (NaCl) and carbon dioxide (CO2). It worked so well that chemists came to believe that the inert gases could not form compounds because by losing, gaining or sharing electrons they would break the octet rule.
Neil Bartlett was one of those scientists not bound by rules. He had become absorbed by chemistry as a schoolboy in Newcastle (England) when he made crystals of copper compounds. After graduating with a PhD from Durham in 1958 he took up his position in Vancouver. He had worked with fluorine compounds and found that platinum hexafluoride was a very powerful oxidising agent. It was so good at taking electrons from other atoms that it even made oxygen molecules form a positive ion in a compound. Bartlett noticed that although xenon was supposed to be chemically inert it took about the same amount of energy to knock electrons off xenon atoms as oxygen molecules. That was why he was messing around with xenon and PtF6 on the March evening in 1962.
Bartlett published his work and immediately other chemists repeated his experiment and tried other reactions. Soon there were many compounds of xenon, krypton and radon (the bottom member of the group). Recently even argon has been forced to make a compound although only at very low temperatures. Only helium and neon are remaining noble and aloof from joining with other elements.
Neil Bartlett died in 2008 by which time his lone experiment had become an active field of chemistry. The inert gas compounds are not just the curiosities of chemists – they have uses. As the compounds are unstable and reactive they make some difficult reactions take place. For instance xenon difluoride (XeF2) is used to make the cancer drug 5-fluorouracil. Compounds of krypton and xenon are used in high power lasers. The compounds could replace toxic heavy metal compounds used as catalysts in many manufacturing processes.
1 Find out what happened to Neil Bartlett between 1962 and his death in 2008.
2 Find out the electron arrangement of the ions in sodium chloride and say why they obey the octet rule.
3 Find out about the properties and uses of some inert gas compounds such as xenon hexafluoride (XeF6) and xenon trioxide (XeO3).
4 (A level) Like many compounds platinum hexfluoride (PtF6) does not obey the octet rule as the platinum atom has six covalent bonds. Bartlett found that his compound, XePtF6 was actually made up of ions (XeF)+ and (PtF6)-.
Draw dot and cross diagrams for the ions in xenon hexafluoroplatinate (XePtF6) and state whether they obey the octet rule.
5 In 1962 most chemists accepted the octet rule ruled out compounds of the inert gases. Explain why Bartlett’s conclusion that xenon could form compounds was quickly accepted.
6 In what ways is chemical research probably different today compared to Bartlett’s lonely evening in 1962.
Neil Bartlett, “Forty Years of Fluorine Chemistry” in Fluorine Chemistry at the Millennium, ed. R.E. Banks; (Amsterdam: Elsevier Science, 2000), p. 28.