In March of 1962, Bartlett concocted a simple experiment to test his hypothesis. It was this development that led Bartlett to theorize that if PtF 6 could oxidize oxygen, then it might also be able to achieve the "impossible" task of oxidizing xenon, whose ionization potential (energy required to remove an electron) was very similar to that of oxygen. Even though PtF 6 was first prepared some years earlier by researchers at Argonne National Laboratory, its oxidizing power had not been recognized until Bartlett's research. But Bartlett believed that in this case, the PtF 6 component was a more powerful oxidizing agent than even oxygen and was extracting electrons from oxygen, leaving oxygen with a net positive charge. Oxygen normally pulls electrons from other atoms and is thus called an oxidizing agent or oxidant. What was most unusual about this compound was that it contained oxygen in the form of positively charged ions, although oxygen usually has a net negative charge. After much research, they eventually found that the known gaseous fluoride, platinum hexafluoride (PtF 6), was able to oxidize oxygen and produce the red solid, which he and Lohmann had identified as O 2 +PtF 6. With the assistance of his graduate student Derek Lohmann, he vigorously pursued the identity of the red solid. Some years earlier, while experimenting with fluorine and platinum, he had accidentally produced a deep-red solid whose exact chemical composition remained a mystery. In 1961 Neil Bartlett was teaching chemistry at the University of British Columbia in Vancouver, Canada. They predicted that highly reactive atoms such as fluorine might form compounds with xenon, the heaviest of the noble elements and whose electrons, they observed, were not as tightly bound as those of the lighter gases.Įarly Research Yields a Mysterious Compound Among those scientists were Walter Kossel in 1916 and Nobel-prize winning chemist Linus Pauling in 1933. This led chemists to think of them as totally unreactive.Ī few chemists questioned the absolute inertness of the noble gases. These electron arrangements are especially stable, leaving the noble gases without a tendency to gain or loose electrons. For helium, that limit is just two electrons. For all except helium, the maximum capacity of the outer electron shell of the noble gas atom is eight electrons. Their inertness became a basic tenet of chemistry, published in textbooks and taught in classrooms throughout the world.Ĭonventional scientific wisdom held that the noble gas elements could not form compounds because their electronic structure was extremely stable. Helium, neon, argon, krypton, xenon, and radon (all gases at room temperature) were viewed as the "loners" of the Periodic Table. Scientists had always believed that noble gases, also known as inert or rare gases, were chemically unable to react.
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